Chemical Solution Deposition Of Semiconductor Films (Food Science & Technology Seri)

CHEMICAL SOLUTION DEPOSITION OF SEMICONDUCTOR FILMS

Gary Hodes Weizmann Institute of Science Rehovot, Israel

Marcel Dekker, Inc.

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ISBN: 0-8247-0851-2 This book is printed on acid-free paper. Headquarters Marcel Dekker, Inc. 270 Madison Avenue, New York, NY 10016 tel: 212-696-9000; fax: 212-685-4540 Eastern Hemisphere Distribution Marcel Dekker AG Hutgasse 4, Postfach 812, CH-4001 Basel, Switzerland tel: 41-61-260-6300; fax: 41-61-260-6333 World Wide Web http://www.dekker.com The publisher offers discounts on this book when ordered in bulk quantities. For more information, write to Special Sales/Professional Marketing at the headquarters address above. Copyright © 2003 by Marcel Dekker, Inc. All Rights Reserved. Neither this book nor any part may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopying, microfilming, and recording, or by any information storage and retrieval system, without permission in writing from the publisher. Current printing (last digit): 10 9 8 7 6 5 4 3

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To my parents, for their dedication to my education and their faith in my abilities

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Preface

Chemical solution deposition (CSD); known also as chemical bath deposition (CBD) and simply chemical deposition (CD, the form we will use in this book) was first described in 1869, and it has been used since to deposit films of many different semiconductors. It is probably the simplest method available for this purpose—all that is needed is a vessel to contain the solution (an aqueous solution made up of a few, usually common, chemicals) and the substrate on which deposition is required. Various “complications,” such as some mechanism for stirring and a thermostated bath to maintain a specific and constant temperature, are options that may be useful. In spite of this extreme experimental simplicity, understanding the mechanisms involved in the deposition and the ability to widen the range of deposits obtained—both in composition and the control of numerous other properties—is usually not so simple. Also in spite of its simplicity, it has not been exploited as a technique as much as might be expected. However, CD has experienced somewhat of a renaissance recently, due largely to its overwhelmingly successful use in depositing buffer layers of CdS (and similar materials) in thin-film photovoltaic cells. The deposition of the CdS, as with many other semiconductors that have been deposited by CD, is often recipe oriented; there seem to be almost as many different “recipes” as there are groups. Notwithstanding the wide interest and use of this technique, at the time when the idea to write this book was conceived, there was no recent comprehen-

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sive review or any general summing up of the field of CD. The first general review on the subject is that by Chopra et al. in 1982 [1]. Nine years later, a review by Lokhande [2] was published with an emphasis on describing the deposition of the various semiconductors that had been deposited up to then. A comprehensive and general review was published by Lincot et al. in 1998 [3] (about a year after the writing of this book had commenced) and, in the same year, two more reviews, one a more specific review by Nair et al. [4] on their extensive work in the field connected with solar energy–related issues and the other, by Savadogo [5], describing CD (and electrodeposited) semiconductors used as solar energy materials. In the last two years, two new reviews have appeared, by Mane and Lokhande [6] and most recently by Niesen and DeGuire [7], the latter covering also other solution deposition methods, such as SILAR (Successive Ion Layer Adsorption and Reaction), electroless deposition and liquid phase deposition (which can be considered to be a subset of CD) and emphasizes oxides, although sulphides and selenides are also covered. The driving force for this book is the perceived need for a detailed coverage of a field that has expanded enormously in recent years. While the title of the technique suggests that the book is aimed mainly at chemists, this would be an incorrect impression. Many of those who should find this book useful will be “physicists” or “engineers” who are dealing with thin-film photovoltaic cells. Some of these readers may have only a superficial background in chemistry, and for these, Chapter 1, Fundamentals, which deals with the science (largely chemistry) behind the technique, will be very important background reading. In this chapter, material will be found on such topics relevant to CD as principles of precipitation and solubility product; nucleation; growth; colloids; aggregation and sticking. Even those with a good chemistry background are advised to read this chapter, if only to refresh their knowledge. In one respect, this book is organized somewhat differently than usual. It contains a fairly comprehensive review of CD in the form of Chapter 2, General Review. Most of the material in this review will be expanded on in the relevant chapters, and one might ask why it is included at all. The reason is that most people do not read a book of this type from cover to cover; they read those chapters or parts of chapters they consider relevant to their purposes. In doing this, they are likely to miss matter from other chapters that is also relevant. For this reason, it is strongly recommended that all readers read through the first three chapters. Chapter 1, as already noted, is to acquire or refresh the relevant scientific background. Chapter 2 should give a good overview of what has been done without having to go into too much detail. Apart from this, however, Chapter 2 contains detail not found elsewhere; a short history of CD, details on the effects of substrate nature and variation (if I were to rewrite this book, an additional chapter would be devoted solely to the substrate), some “recipes” for depositing certain films and, very important, descriptions, and, where possible, explanations of the reasons for

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a particular recipe and the expected effects of various changes to those recipes. Also, other methods that are related to CD are treated briefly in this chapter. In order to promote the ease of reading and “flow” in this chapter, references are, for the most part, not given unless they are not clearly provided in the chapter and section relevant to that particular topic elsewhere in the book. The importance of the third “required” chapter, Chapter 3, Mechanisms of Chemical Deposition, should be self-evident. Understanding the possible mechanisms of the deposition is the best defense against what is commonly (and often with considerable justification) thought of as a “recipe-oriented” field. Not that that a reading of this chapter will automatically endow the reader with the ability to know what mechanism is actually operating in every case—unfortunately, often far from it. However, it is hoped that it will help in the movement toward that goal. Having worked (it is hoped) through these “compulsory” chapters, most readers will want more detail on specific aspects that are important to them. How can films of X be made? How are the various experimental parameters expected to affect the properties of this film? Why is CD so useful for photovoltaic cells? How can nanocrystalline films of Y be made with a specific crystal size? Such questions will be answered in subsequent chapters (or, if not answered, at least information will be given to allow the reader to plan experiments in order to find the answers). The next five chapters deal with deposition of specific groups of semiconductors. In Chapter 4, II–VI Semiconductors, all the sulphides, selenides, and (what little there is on) tellurides of cadmium (most of the chapter), zinc (a substantial part), and mercury (a small part). (Oxides are left to a later chapter.) This chapter is, understandably, a large one, due mainly to the large amount of work carried out on CdS and to a lesser extent on CdSe. Chapter 5, PbS and PbSe, provides a separate forum for PbS and PbSe, which provided much of the focus for CD in earlier years. The remaining sulphides and selenides are covered in Chapter 6, Other Sulphides and Selenides. There are many of these compounds, thus, this is a correspondingly large chapter. Chapter 7, Oxides and Other Semiconductors, is devoted mainly to oxides and some hydroxides, as well as to miscellaneous semiconductors that have only been scantily studied (elemental selenium and silver halides). These previous chapters have been limited to binary semiconductors, made up of two elements (with the exception of elemental Se). Chapter 8, Ternary Semiconductors, extends this list to semiconductors composed of three elements, whether two different metals (most of the studies) or two different chalcogens. The final two chapters deal with “applications” (in the scientific as well as commercial sense) of CD films. As already mentioned, photovoltaic cells is the one subject that has given CD a push in the last decade, while photoelectrochemical cells was probably the main driving force for such studies in the decade before that. Chapter 9 deals with Photovoltaic and Photoelectrochemical Properties.

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Finally, the tendency for CD films to be nanocrystalline and often to exhibit quantum-size effects is treated in the final chapter, Chapter 10, Nanocrystallinity and Size Quantization in CD Semiconductor Films. The layout of this book means that there will be some overlap between sections. However, this system should allow those readers interested in one or more specific sections to skip the others, thereby making the book more efficient for the individual reader. An example of this is the use of quantum-size effects to elucidate CD mechanisms. This is treated, with different emphasis, both in Chapter 3 (Mechanisms of Chemical Deposition) and in Chapter 10 (Nanocrystallinity and Size Quantization in CD Semiconductor Films). This book covers a field that, from its title, may appear to be limited. In reality, it is surprisingly multidisciplinary. Inorganic chemistry and film formation are, of course, fields that are evident from the title. Certainly, those concerned with the deposition of semiconductor films for any purpose should find this book useful and informative. However, it will also be valuable to those working in other fields. The considerable section on semiconductor quantum dots, for example, will be of interest to those working with low-dimensional semiconductors. Scientists and engineers working in thin-film solar cells will find a compendium of research on CD buffer layers in these cells. There is much in the book relevant to colloid scientists. Even biologists are not forgotten: the slow formation and deposition of inorganic compounds characteristic of CD has a lot in common with biomineralization, and hopefully this book will be useful to those working in the field of biomimetics. And, of course, the original application of CD, for near- to mid-infrared detectors, will attract those designing or using optoelectronic equipment in this wavelength range. Finally, a word concerning the coverage of the literature. When starting out on this enterprise, the intention was to try and cover the field more or less completely, with the exception of some of the earlier work, mostly on PbS. Of course, some papers might occupy a considerable amount of book space, while, at the other extreme, others might just be mentioned in a table. During the long course of putting this book together, it became increasingly clear, from the “new” literature that appeared (not necessarily chronologically new but just new to the author), even toward the final stages of writing, that an appreciable number of relevant papers would remain unknown to the author. For these, the author expresses regret, not only to the authors of such works, but also to the readers of this book. It is hoped that the expanding literature that is continually appearing on the subject does not render this book out of date too rapidly. From the objective scientific viewpoint, it can be hoped that this process will not be too slow either.

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REFERENCES 1. KL Chopra, RC Kainthla, DK Pandya, AP Thakoor. In: Physics of Thin Films, Vol. 12. Academic Press, New York and London, 1982, pp 167. 2. CD Lokhande. Mater. Chem. Phys. 28:1, 1991. 3. D Lincot, M Froment, H Cachet. In: RC Alkire, DM Kolb, Eds. Adv. Electrochem. Sci. Eng., New York: Wiley-VCH, 1998, Vol. 6, p 165. 4. PK Nair, MTS Nair, VM Garcia, OL Arenas, Y Pena, A Castillo, IT Ayala, O Gomezdaza, A Sanchez, J Campos, H Hu, R Suarez, ME Rincon. Sol. Energy Mater. Sol. Cells 52:313, 1998. 5. O Savadogo. Sol. Energy Mater. Sol. Cells 52:361, 1998. 6. RS Mane, CD Lokhande Mater. Chem. Phys. 65:1, 2000. 7. TP Niesen, MR De Guire J. Electroceram. 6:169, 2001.

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Contents

Dedication Preface 1. Fundamentals 2. General Review 3. Mechanisms of Chemical Deposition 4. II-VI Semiconductors 5. PbS and PbSe 6. Other Sulphides and Selenides 7. Oxides and Other Semiconductors

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8. Ternary Semiconductors 9. Photovoltaic and Photoelectrochemical Properties 10. Nanocrystallinity and Size Quantization in Chemical Deposited Semiconductor Films

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1 Fundamentals

The purpose of this chapter is to give sufficient background in the chemical principles involved in CD. For those with a good background in chemistry, a quick read through this chapter as a refresher course will probably be sufficient. However, the chapter is written, to a large extent, keeping in mind that not all readers will have a good chemistry background. The emphasis is on a qualitative or semiquantitative understanding of the principles involved, sufficient to understand the concepts that arise throughout the book. A deliberate policy has been made not to go too deeply into the details of these fundamentals where it is considered unnecessary; references to further reading will be given for those who wish such additional detail.

1.1 SOLUTION CHEMISTRY 1.1.1 Basic Terminology The pH of a solution is the negative logarithm of the hydrogen ion concentration in the solution: pH
log [H]

(1.1)

(Note for thermodynamic purists: Here and for the rest of this book, concentrations are used in place of activities.) Concentrations are denoted by square brackets; thus [H] means the concentration of hydrogen ions.

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The pH of pure water at 25°C is 7. Most (but not all) CD reactions take place in basic solutions at typical pH values of 9–12. Since hydroxide intermediates are often important in CD, it is worth noting that a pH of 10 is equivalent to a hydroxide ion concentration of 104 M at 25°C (since the ion product of water, [H][OH],
1014 at this temperature). As will be discussed shortly, this ion product is very temperature dependent, and so the OH concentration at any particular pH varies considerably with temperature. The pH of pure (and also not so pure) water is very sensitive to small concentrations of acids and bases. One drop of concentrated sulphuric acid added to a liter of water will change the pH by 4 pH units (from 7 to ca. 3). Solution pH can be stabilized by a buffer (although there may be cases where a stable pH is not desirable); addition of (not too large) quantities of acid or base to a buffered solution will not affect the pH much. Buffers are usually mixtures of weak acids or bases and their salts. A common example in CD is the use of an ammonium salt  (NH 4 X ) to control the pH of an ammonia solution. The equilibrium of ammonia in water is given by  NH3  H2O D NH 4  OH

(1.2)

Since hydroxide ions are formed when ammonia dissolves in water, the pH of an aqueous ammonia solution is alkaline. The value of pH can be calculated from a knowledge of the equilibrium constant, K, of this equilibrium. The equilibrium constant for an equilibrium of general type aA  bB    dD D eE  f F    hH

(1.3)

is given by [E]e[F] ƒ . . . [H]h  K
[A]a[B]b . . . [D]d

(1.4)

K for the ammonia dissolution, Eq. (1.2), is given by  [NH 4 ][OH ] K

1.8  105 [NH3]

(at 25°C)

(1.5)

For example, for a 1 M solution of ammonia ([NH3]
1), since the NH 4 and OH concentrations are equal [from Eq. (1.2)], [OH] can be calculated to be 4.2  10-3 M. Knowing that the ion product of water, [H][OH],
1014 at this temperature allows us to convert [OH] to [H] and thus to find the pH (which is 11.62). If an ammonium salt is added to the ammonia solution, the NH 4 concentration is now dictated by the concentration of ammonium salt added rather than by that existing due to the weak dissociation of ammonia. Thus, assuming the same 1 M ammonia as before, but adding 0.1 M NH 4 (say, as NH4Cl), then, ignoring the few percent correction due to the extra NH 4 arising from the NH3 dissocia-

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tion, [OH] is given by 1.8  10-5 [OH]

1.8  104 M 0.1

(1.6)

and the pH of the solution becomes 10.25. The buffering action of this solution can be understood by considering equilibrium Eq. (1.2). If extra OH is added to the solution, the equilibrium is shifted to the left; i.e., it tends to remove OH. It also removes NH 4 ; but if the concentration of this ion is high to begin with, then this change will not affect the pH greatly. (This is the reason that the combination of ammonia and ammonium ions is a better buffer than ammonia by itself.) A similar argument can be made for addition of acid, through the equilibrium: NH3  H D NH 4

(1.7) NH 4.

In this case, hydrogen ions are consumed in converting ammonia to This buffering action requires nonionized base (or acid) to operate, hence the requirement of a weak base or acid together with its salt As well as this buffering action, addition of ammonium ion also decreases the pH of an ammonia solution as shown above. This is an important effect—more important than the buffering action in many CD processes.

1.1.2 Hydrolysis of Metal Ions Most cations are hydrated in aqueous solutions to a greater or lesser extent: Mx  nH2O D M(H2O)x n

(1.8)

The water is polarized and attracted by the positively charged cation. The greater the positive charge on the cation and the nearer the water can approach the cation, the greater will be this polarization and attraction. Thus small, highly charged (high-valence) cations will in general be more strongly solvated than large, monovalent ones. Continuing the argument, the positive charge on the cation attracts electrons from the oxygen of the water molecules. This, in turn, can result in the transfer of electron density from the OMH bonds to the (now electron-deficient) oxygen, as exemplified here:

(1.9)

This will weaken the OMH bond and may even break it, resulting in formation of a metal hydroxide and a hydrogen ion, the latter which will be hydrated by a

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molecule of the surrounding water: (x1) M(H2O)x  H3O n  H2O D M(H2O)n1OH

(1.10)



The H3O (a hydrated hydrogen ion) is acidic; therefore this equilibrium generates acidity in the solution. The more the cation attracts electron density from the water, therefore, the more acidic is the cation. As with solvation, small, highly charged cations should be more acidic than large cations with a small charge. The charge in particular is a very important factor in determining the degree of acidity of cations. Therefore monovalent cations are generally basic, while trivalent ones are acidic. Tetravalent cations, such as Sn4 and particularly Ti4 are so highly acidic that their simple cations either do not exist in water or do so only under very acidic conditions. The hydrated metal hydroxy complex in Eq. (1.10) is a soluble species. However, if the pH is sufficiently high, the metal hydroxide, which is relatively insoluble for most metals (apart from the alkali group metals) will precipitate. The pH value at which hydroxide precipitation occurs can be related to the acidity of the cation and is approximately equal to the pKa of the cation, where the pKa is minus the logarithm of the equilibrium constant of Eq. (1.10).

1.1.3 Solubility Product A central concept necessary to understanding the mechanisms of CD is that of the solubility product (Ksp). The solubility product gives the solubility of a sparingly soluble ionic salt (this includes salts normally termed “insoluble”). Consider a very sparingly soluble salt (say, CdS) in equilibrium with its saturated aqueous solution: CdS(s) D Cd2  S2

(1.11)

(where subscript s represents the solid phase). The CdS dissolves in water to give a small concentration of Cd and S ions. This concentration is defined by the solubility product, Ksp, the product of the concentrations of the dissolved ions: Ksp
[Cd2][S2]

(1.12)

or more generally, for the dissolution: MaXb D aMn  bX m Ksp
[Mn]a[Xm]b

(1.13) (1.14)

The more soluble is the salt, the greater the ion product and the greater is Ksp. However, Ksp also depends on the number of ions involved. Thus Bi2S3 has a value of Ksp
[Bi3]2[S2]3
10100. The very low value is due, in part, to the relatively large number of atoms in the Bi2S3 molecule and therefore of ions involved in the equilibrium. A list of approximate values of Ksp for some of the

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TABLE 1.1 Values of the Solubility Product (at 25C) for Compounds Relevant to CD Solid Ag2S Ag2Se AgCl AgBr AgI As2S3 Bi(OH)3 Bi2S3 Bi2Se3 Cd(OH)2 CdS

K sp

Solid

K sp

Solid

K sp

Solid

K sp

3  1050 1054 2  1010 8  1013 1016 2  1022 6  1031 10100 10130 2  1014 1028

CdSe CdTe Co(OH)2 CoS CuOH Cu(OH)2 Cu2S CuS CuSe Fe(OH)2 Fe(OH)3

4  1035 1042 5  1015 1021 1  1014 2  1020 1048 5  1036 2  1040 5  1017 3  1039

FeS HgS HgSe In(OH)3 In2S3 Mn(OH)2 MnS Ni(OH)2 NiS NiSe PbCO3

1018 6  1053 4  1059 6  1034 6  1076 5  1013 1013 3  1016 1021 2  1026 1013

Pb(OH)2a PbS PbSe Sn(OH)2 Sn(OH)4 SnS SnS2 SnSe Zn(OH)2 ZnS ZnSe

1015–1020 1028 1037 5  1028 1  1056 1026 6  1057 5  1034 1016 3  1025 1027

There is often a large variation in values from source to source—in some cases, some orders of magnitude. For this reason, only one significant figure (at most) is given before the exponent. A table of solubility products for many sulphides based on a reevaluated value for the second dissociation constant of H2S is given in Ref. 1. The values in that study are typically some orders of magnitude lower than the ones shown here. a This is probably a hydrated lead oxide rather than a simple hydroxide.

semiconductors and related salts encountered in CD is given in Table 1.1. The values for oxides are not so readily available as for sulphides and selenides. However, it must also be kept in mind that deposition of oxides often occurs via a hydroxide or hydrated oxide, and therefore the relevant value is that of this hydroxide or hydrated oxide. Some explanation is required here concerning the S2 ion. In actual fact, in all but highly alkaline solutions (and the solutions used in CD, while mostly alkaline, are not that alkaline), most of the sulphur ion will be in the form of HS rather than S2. This is due to the equilibrium between the two species: HS D S2  H

Ka
1017.3

(1.15)

or alternatively, in terms of hydroxide concentration which is related to the hydrogen ion concentration through the ion product of water: HS  OH D S2  H2O

Ka
103.3

(1.16)

(at room temperature). Thus at a pH of 11 (a common value in CD), which gives a value for [OH] at room temperature of 103 M, the S2 concentration will be [S2]
103.3 [HS][OH]
106.3 [HS]

(1.17) 

Therefore the main sulphur ion in solution will be HS .

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Since Ksp is given in terms of [S2], we can write [S2] in terms of [H] using Eq. (1.15): [S2]
1017.3[HS]/[H]

(1.18)

and derive a solubility product of a sulphide, MxSy, in terms of the dominant HS concentration: Ksp
[M2y/x]x[S2]y
[M2y/x]x(1017.3[HS]/[H])y

(1.19)

A list of such solubility products for metal sulphides, as well as updated conventional ones, has been given by Licht [1]. In this book, we will continue to use the more conventional solubility products, partly because they are more common and partly because the relevant equilibria are less unwieldly to describe. In applying the solubility product concept to CD, it is often useful to consider it in terms of what concentration of ions is required in solution before precipitation occurs. Thus, for CdS, with a Ksp value of ca. 1028 (from Table 1.1), [Cd2][S2]
1028 M.

(1.20)

While the concentration of each ion in this example will be equal when dissolution of the solid is considered, for formation of the solid from the ions, they may be completely different; it is the product of the concentrations that is important. Thus a solution 0.2 M in sulphide ion and 1027 M in Cd2 will (in principle) precipitate CdS (the ion product will be greater than Ksp) while 0.1 M sulphide and the same 1027 M of Cd2 will (only just) not. Ksp can be derived theoretically from the free energies of formation of the species involved in the dissolution equilibria. Thus, for the equilibrium MaXb(s) D aMc(aq)  bXd(aq)

(1.21)

the free energy of the dissolution is given by G0
aG0(aMc(aq))  b G0(Xd(aq))  G0(MaXb(s))

(1.22)

And since G0
RT ln K

(1.23)

G0 ln Ksp
 RT

(1.24)

then

Since Ksp is a thermodynamic quantity, the ion product that should result in precipitation may not necessarily do so for kinetic reasons (hence the term used earlier to qualify precipitation: “in principle”). This would be a case of supersaturation. In practice, however, the solubility product does give a fairly good idea of when precipitation will occur in most cases.

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CD reactions sometimes proceed via a metal hydroxide intermediate; the concentration of OH ions in the solution is particularly important in such cases. Since almost all CD reactions are carried out in aqueous solutions, the pH of the deposition solution will give this concentration. In translating pH into OH concentration, the very temperature-dependent ionization constant of water should be kept in mind, as mentioned previously. The reason for this can be seen from Table 1.2, which gives the OH concentration in water at a pH of 10 (a typical pH value for many CD reactions), calculated from the ionization constant of water, Kw from the relation log Kw
log [H][OH]
log [H]  log [OH]
pH  log [OH]

(1.25)

The OH concentration increases by nearly two orders of magnitude between 0 and 60°C. The OH concentration increases (decreases) by one order of magnitude for every unit increase (decrease) in pH. This means that the formation of a metal hydroxide (whether as a colloid or as a precipitate) in aqueous solution will be strongly dependent on temperature when the product of the free metal ions and OH ions is close to the hydroxide solubility product, although increase in Ksp with temperature may partially offset this effect.

1.1.4 Complexation Most CD reactions are carried out in alkaline solution. To prevent precipitation of metal hydroxides, a complexing agent (often called a ligand, since complexing agents to cations are electron donors) is added. The complexant also reduces the concentration of free metal ions, which helps to prevent rapid bulk precipitation of the desired product. This section gives the basics of the theory of complexation.

TABLE 1.2 Effect of Temperature on OH Concentration in Water Temp. (C)

-log10 Kw

OH concentration at pH
10

0 10 20 30 40 50 60

14.944 14.535 14.167 13.833 13.535 13.262 13.017

1.138  105 2.917  105 6.808  105 1.469  104 2.917  104 5.470  104 9.616  104

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If a KOH solution is added to a solution of a Cd salt, Cd(OH)2 will precipitate immediately. From the Ksp of Cd(OH)2 (2  1014 at room temperature), and assuming a pH of, say, 11 ([OH]
103 at the same temperature), from Eq. (1.14), we can calculate that a Cd concentration above 2  108 M is enough to initiate Cd(OH)2 formation. If ammonium hydroxide (ammonia in water)—a common complexant for Cd in CD—is added to a suspension of Cd(OH)2, the Cd(OH)2 will redissolve, assuming enough ammonia has been added. How much is enough ammonia? This can be calculated from the stability constant of the complex between ammonia and Cd. The equilibrium of this reaction to form the cadmium tetraamine complex is given by Cd2  4NH3 D Cd(NH3) 2 4

(1.26)

and the stability constant of this equilibrium, Ks, by [Cd(NH3)2 4 ] Ks

1.3  107 2 [Cd ][NH3]4

(1.27)

As calculated previously, Cd(OH)2 will precipitate when the free Cd2 concentration is larger than 2  108 M (at a pH of 11 and at room temperature). From Eq. (1.27), for a total Cd concentration of 0.1 M, we can calculate that a free NH3 concentration of 0.79 M will result in a free Cd2 concentration of 2  108 M. Add to this the ammonia tied up by complexation (0.1  4 M), the minimum NH3 concentration required to prevent precipitation of Cd(OH)2 is therefore 1.19 M. At a Cd concentration of 0.01 M (more typical of many depositions), the corresponding concentration of ammonia is ca. 0.5 M. At a deposition temperature of 60°C (CdS deposition is generally carried out at elevated temperatures, usually 60°C), the ion product of water is 13, and therefore the OH concentration at a pH
11 will be 102 M. Calculating the minimum concentration of ammonia required to prevent precipitation of Cd(OH)2 at 60°C and 0.01 M total Cd gives a value of 1.44 M. The value of pH chosen is typical of these solutions. For higher values of pH, and at higher temperatures at the same pH (both of which mean an increased [OH]), more ammonia will be required to prevent precipitation of Cd(OH)2. This calculation ignored the decrease in the stability constant of the complex with increasing temperature (see later) as well as the increase in Ksp that normally occurs with an increase in temperature. These two effects act in opposite directions; for most cases, their combined effect will be much smaller than that of the temperature-dependent ion product of water. Another simplification is the assumption of only one complex species; this simplification is reasonable for most purposes. If a solution contains an excess of one of the ions of a sparingly soluble salt, this will modify the solubility of the sparingly soluble salt according to the common ion effect. As an example of this effect, we might consider the precipitation

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of Zn(OH)2 by hydroxide according to the reaction Zn2  2OH D Zn(OH)2

(1.28)



An excess of OH (the common ion) should shift the reaction to the right, i.e., to more complete precipitation of the Zn(OH)2. This effect is a general one, but the conclusions are not always valid; the example (deliberately) given here is one where it is not valid. The reason is that OH can form a complex with Zn2 2 (Zn(OH)2 from solution and reduc4 —the zincate ion), thus removing free Zn ing the degree of precipitation. For a sufficiently high concentration of OH, which can be calculated from the stability constant of the zinc–hydroxide (zincate) complex, the Zn(OH)2 will completely redissolve. The stability constant of a complex is temperature dependent—increased temperature generally leads to increased dissociation of the complex. Qualitatively, this can be explained by the Le Chatelier principle, which states that if there is a change in a reaction parameter, the reaction will proceed in a direction that opposes that change. Thus an increase in temperature will cause the reaction to go in the direction in which heat is absorbed, which is dissociation of the complex. More quantitatively, the relation between equilibrium constant and temperature is given approximately by H0 ln K
 RT

( a constant)

(1.29)

where H0 is the standard enthalpy change in the process and R is the gas constant. This is the integration of the van’t Hoff equation, hence the constant term. (The derivation of this equation can be found in any elementary physical chemistry textbook and there is no need to repeat it here—the result is what is important for us.) This equation is approximate for a number of reasons. One is that it ignores changes in entropy that often will act in the opposite direction for complex formation. However, the trend is generally correct. The stability constant of a complex does not, according to Eq. (1.27), depend on the concentrations of the species comprising the complex. For very dilute solutions, however, complexes become less stable than expected from their “literature” stability complex. The reason for this lies in the fact that the equilibrium shown in Eq. (1.26) is not strictly correct; a more accurate representation would be 2 Cd(H2O)2 6  4NH3 D Cd(NH3)4  6H2O

(1.30)

(hydrolysis and hydration of ammonia and ammonium ion is ignored, although for an accurate representation, it should be considered—it will not affect the argument). Since ammonia is a much stronger ligand than water (water of hydration can be considered as a ligand), it will exchange all the water as long as the ammonia concentration is not too low. If it is very low, then not all the water will necessarily be exchanged, and a different equilibrium (or mixture of equilibria) with

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a different equilibrium constant will exist. A classic example of this effect is given by the cobalt complex with thiocyanate, SCN:  2 Co(H2O)2 6  4SCN D Co(SCN)4  6H2O

pink

(1.31)

blue 2

The pink color of the hydrated Co ion turns blue when a high concentration of SCN is added; if diluted with water, this solution reverts to pink. This color transition is reversible. An aqueous solution of Co2 is pink, while anhydrous Co(II) salts are typically blue (a fact well-known to chemists and even to schoolchildren who have experimented with invisible inks). The color change to blue on addition of SCN to a noncomplexed Co2 solution is caused by dehydration of the Co2 due to exchange of water with the SCN, a stronger ligand than water. However, SCN is not a very much stronger ligand than water to Co2, and therefore a relatively high concentration is required to exchange all the water. At intermediate concentrations of SCN, mixed aquo-thiocyanato complexes can be formed, which are pink. The stronger the ligand relative to water, the less the concentration required to exchange the water. It is worth noting that a statistical effect (different combinations of the various complexants) may result in mixtures of complexants binding more strongly to a cation than would be expected based on the individual stability constants of the complexants [2]. If a compound containing more than one cation is to be deposited, complexation could be used to offset the difference in Ksp between the individual metal compounds. As an example, consider the deposition of (Cd,Hg)S. From Table 1.1, the value of Ksp for HgS is much lower than for CdS. This means that under the conditions of CD, where the sulphide ion is slowly formed, we would expect only HgS to form (until almost all the Hg was used up). Some Cd might be incorporated into the deposit by adsorption, but the deposit should, according to considerations of solubility product, be predominantly HgS. The concentration of Hg in the deposit can be decreased by choosing a complex (or mixture of complexes) that complexes Hg more strongly than it does Cd. In this case, since Hg forms very strong complexes with many ligands, there is a large choice. Skyllas-Kazacos et al. deposited films of (Cd,Hg)S using a combination of ammonia and cyanide (the latter is a strong complex for both cations, but more so for Hg) [3]. In addition, the Hg concentration was much smaller than that of cyanide, while the Cd concentration was larger. This means that there was enough cyanide to complex the Hg but not enough for the Cd. A further factor that may have allowed codeposition of Cd was the use of the chloride anion, which is a moderately strong complex for Hg2 but only a weak one for Cd2. This combination of factors allowed codeposition of the Cd and Hg as sulphides, but the concentrations of Hg in the films were larger (by a factor of ca. 4) than in the deposition solution. The addition of a large concentration of iodide to the deposition solution would probably

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have been even better, since iodide is an extremely weak complexant for Cd and a very strong one for Hg. It would therefore have removed most of the free Hg2 ions while only changing the Cd2 concentration a relatively small amount. This judicious use of a complexant to allow codeposition of two cations with widely differing values of Ksp is, unfortunately, not always useful. An example is the deposition of (Cd,Zn)S—a material of interest particularly because of its potential use in photovoltaic cells (see Chapter 9). The stability constants of Cd and Zn ions are in most cases very similar for any particular complex (although, of course, they vary greatly from one complex to another). This reflects the similarity of the chemistry of these two ions (and the difference between them and the Hg2 ion). Therefore complexation is very limited as a means to control the concentration of one of these ions relative to the other. The use of complexation to allow codeposition of alloys is well known in electroplating. The best-known example is that of brass (Cu/Zn) plating, where cyanide, which is a stronger complex for Cu than it is for Zn, brings the deposition potentials of the two metals, originally far apart, to almost the same value. There is a direct connection between this effect and the equivalent one for CD. This arises from the fact that, for both CD and electrodeposition of alloys (we include mixed metal compounds in the term alloy), the effect of the complexant is to lower the concentration of free cations. For CD this affects the deposition through the solubility product, while for electrodeposition it affects the deposition potential through the Nernst equation: [Ox] RT E
E 0   ln  (1.32) nF [Red] where the oxidized species, Ox, the cation in this case, is reduced in concentration, resulting in a more negative deposition potential, E, compared to the standard potential, E0. In the case of metal electrodeposition, the reduced species, Red, is the metal that, since it is a solid, can be taken as unity concentration. n is the number of electrons transferred per molecule of reaction (e.g., for Cd deposition from Cd2, n
2) and F is Faraday’s constant (ca. 96,500 coulombs/mole). The shift in potential due to complexation, E, (
E  E 0) can be approximated by E (in mV)
60 log [cation]/n

(1.33)

From Eq. (1.27), we can write log Ks
log [complexed cation]  log [cation]  a log [ligand]

(1.34)

where a is the number of ligand molecules in the complex. Combining Eqs. (1.33) and (1.34) we get log Ks
n E (in mV)/60  log [complexed cation]  a log [ligand]

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(1.35)

For all but very weak complexes, the concentration of complex ion (and often also the free ligand) is normally very much larger than that of the free cation. With this in mind, Eq. (1.35) can often be approximated by log Ks
n E (in mV)/60

(1.36)

This [or, more accurately, Eq (1.35)] allows us to calculate values of the stability constant of a complexant from tables of electrochemical potentials. For example, a shift of 300 mV in potential due to complexation gives an (approximate) value for the stability constant of that complex of 105 (for n
1) or, for the more common case in CD, where divalent cations (n
2) predominate, 1010.

1.2 NUCLEATION AND GROWTH CD can occur either by initial homogeneous nucleation in solution or by heteronucleation on a substrate, depending on the deposition mechanism (see Chapter 3). For this reason, we consider both types of nucleation.

1.2.1 Homogeneous Nucleation According to simple solubility considerations, a precipitate will be formed when the product of the concentrations of anions and cations exceeds the solubility product. From another viewpoint, phase transformation occurs when the free energy of the new phase is lower than that of the initial (metastable) phase. However, there are many examples where the ion product exceeds Ksp, yet no precipitation occurs—the phenomenon of supersaturation. The solubility product also does not provide information on how the particles of the precipitate form—nucleation. Nucleation involves various physical processes, and both thermodynamic and kinetic aspects must be considered. Homogeneous nucleation can occur due to local fluctuations in the solution— whether in concentration, temperature, or other variables. The first stage in growth is collision between individual ions or molecules to form embryos (embryos are nuclei that are intrinsically unstable against redissolution—see later). Embryos grow by collecting individual species that collide with them. While these species may be ions, atoms, or molecules in general, for CD, adsorption of ions on the embryo seems to be the most probable growth mechanism. They may also grow by collisions between embryos; however, unless the embryo concentration is large, this is less likely. These embryos may redissolve in the solution before they have a chance to grow into stable particles (nuclei). Because of the high surface areas, and therefore high surface energies of such small nuclei, they are thermodynamically unstable against redissolution. They may, however, be kinetically stabilized by low temperatures, which increase their lifetime, possibly enough for them to grow to a size where they are thermodynamically stable. This is an important reason why

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smaller particles can be formed at lower temperatures in a precipitation reaction; the subcritical embryos last long enough to grow into stable particles, while at higher temperatures they would redissolve, reducing the density of nuclei. This results in an increase of the particle size, since there is more reactant per nucleus. The critical radius, Rc, is the size where the embryo (nucleus) has a 50:50 chance of either redissolving or growing into a stable nucleus; it is determined by the balance between the surface energy required to form the embryo, Es
4 R2

( is the surface energy per unit area)

(1.37)

and the energy released when a spherical particle is formed, Ev
4 R3 L/3 ( is the density of the solid and L is the heat of solution).

(1.38)

This balance is shown in Fig. 1.1. The typical size of Rc is about 100 molecules— between 1 and 2 nm in diameter. Solvent molecules can adsorb on the embryos and change their surface energy; the critical radius will therefore depend not only on the material of the nucleating phase but also on the solution phase.

1.2.2 Heterogeneous Nucleation In heterogeneous nucleation, subcritical embryos (or even individual ions) can adsorb onto the substrate. The energy required to form an interface between the

FIG. 1.1 Energetics of nucleation. The critical radius, Rc, depends on the balance between surface and volume energies of the growing particle.

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FIG. 1.2 Processes involved in heterogeneous nucleation on a surface.

embryo and the solid substrate will usually be less than that required for homogeneous nucleation, where no such interface exists. Therefore heterogeneous nucleation is energetically preferred over homogeneous nucleation and can occur near equilibrium saturation conditions, compared with the high degree of supersaturation often required for homogeneous nucleation. These subcritical nuclei can grow, either by surface diffusion or by material addition from solution. It should also be noted that nuclei that are subcritical in solution may be supercritical when adsorbed on a substrate. This is a consequence of reduced contact between nucleus and solution as well as stabilization of the adsorbed nucleus. These processes are shown schematically in Fig. 1.2. It was noted earlier that even individual ions may adsorb onto a surface. More specifically, depending on the surface chemistry of the substrate, individual ions or molecular species may actually be chemisorbed, creating a nucleus for reaction and further growth. Pure homogeneous nucleation is probably less common that might appear from the above discussion. Because of the greater ease of nucleation on a solid phase than homogeneously, any solid matter in the solution will act as a nucleation center. It is difficult to prepare solutions without some solid phase (usually dust particles)—careful filtering is necessary to attain such particle-free solutions. That this is so can be seen from the simple test of shining a laser beam (preferably a green or blue laser, since scattering is greatly enhanced compared with a red one) through an visibly “clear” solution; the resulting scattering by dust particles is almost always evident.

1.2.3 Crystal Growth Once (stable) nuclei have formed, there are several ways in which they can increase in size. One is a continuation of the process of embryo growth discussed earlier: adsorption of ionic species from the solution onto the nucleus. Crystal

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growth of this type can be considered a self-assembling process. Thus for CdS, either Cd2 or S2 will adsorb (as discussed later, since a crystal, and in particular a polar one, is made up of different faces, the adsorption properties of each may be different, and therefore both types of ions may adsorb on any one crystal). The next growth step will then be adsorption of the oppositely charged ion to give an additional CdS molecule. This process can continue until either all the ions of any one type are used up or growth is blocked, e.g., by aggregation or by blocking of the crystal surface by a foreign adsorbed species. Also, growth may continue but in a different geometric orientation, giving rise to twinning, polycrystallinity, etc. Another mechanism for crystal growth is known as Ostwald ripening. If a small nucleus or embryo is close to a larger crystal, the ions formed by (partial) dissolution of the smaller, less stable crystal can be incorporated into the larger crystal. As the smaller crystal becomes even smaller, its dissolution will become ever more favorable and eventually it will disappear. The result is that the larger crystals grow at the expense of the smaller ones. If the concentration of particles is sufficiently high, then the probability of collisions between these particles becomes high. This can result in either aggregation or coalescence. When two particles approach each other, the van der Waals force of attraction (see section 1.3.1) between them will often cause them to stick together. This can continue until a large particle (large in relation to the original particle size) comprising the individual particles has formed (Fig. 1.3A). This is the process of aggregation, and the resulting large particle is called an aggregate. (In colloidal chemistry, the alternative terms of flocculation and floc are often encountered.) The properties of the aggregate may be similar to those of the individual particles in some ways (such as X-ray diffraction

FIG. 1.3 Aggregation (A) and coalescence (B) of individual particles.

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peak broadening, quantum size effects) and very different in others (e.g., light scattering, sedimentation). In an aggregate, there are grain boundaries between individual crystallites. However, in some cases, particularly if the temperature is high enough to allow appreciable diffusion of the crystal atoms, surface diffusion may occur where two (or more) particles have aggregated, resulting in the formation of a neck. This is termed coalescence. Coalescence may continue until one large particle is formed from the original two or more particles (Fig. 1.3B).

1.2.4 Particle Size Distribution If nucleation occurs in a very short time, whereas growth occurs separately, often over a much longer time but without further nucleation, then the size distribution is likely to be narrow, since all the original nuclei should be of similar size and grow at the same rate. The opposite case, where nucleation and growth occur simultaneously, usually results in a wide size distribution. Homogeneous nucleation normally requires a supersaturated solution, while growth can occur close to the saturation concentration. Therefore rapid nucleation can occur if supersaturation is rapidly reached. This nucleation lowers the concentration of reactants below that needed to cause further nucleation. If one of the reactants is supplied at a low concentration after nucleation has occurred (such as by in situ homogeneous formation in the solution), then growth can occur without further nucleation, resulting in a narrow size distribution. In CD, where the reaction is slow, it might be expected that nucleation and growth will always occur together, resulting in a relatively wide size distribution. This is indeed expected for heterogeneous nucleation on a substrate. However, for mechanisms where homogeneous nucleation of an intermediate phase occurs rapidly in the solution but conversion to the final compound is a slow process, nucleation and growth can still be separated.

1.3 FORCES BETWEEN PARTICLES AND BETWEEN PARTICLES AND SURFACES Once nanoparticles have been formed, whether in an early state of growth or in a more or less final size, their fate depends on the forces between the individual particles and between particles and solid surfaces in the solution. While particles initially approach each other by transport in solution due to Brownian motion, convection, or sedimentation, when close enough, interparticle forces will determine their final state. If the dominant forces are repulsive, the particles will remain separate in colloidal form. If attractive, they will aggregate and eventually precipitate. In addition, they may adsorb onto a solid surface (the substrate or the walls of the vessel in which the reaction is carried out). For CD, both attractive particle–sur-

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face and particle–particle forces are required for film formation. In all the stages of the CD process, except for a very few studies, that of adhesion of the film to the substrate is probably the least understood; it is rarely even considered in mechanistic studies. Why do some particles stick to a particular substrate and others not? Also (and easier to understand, at least intuitively), why do the particles stick to each other in building up a film? In order to understand these sticking phenomena so crucial for the CD process, we consider the various forces involved—repulsive as well as attractive— involving the particles. We discuss first the more obvious forces and then some less obvious ones that nonetheless may be important in some cases. Since the dominant force in CD is the van der Waals attraction, we will begin with this interaction.

1.3.1 van der Waals Forces The main interaction that determines whether and to what extent particles will adhere to each other and also (if there is no specific chemical interaction) to a substrate is, in most cases, the van der Waals interaction. The van der Waals force of attraction is a universal interaction that operates between all particles, whether atoms, molecules, clusters, charged, or noncharged. The attraction is due to an induced dipole–induced dipole interaction between particles. The dipoles arise from fluctuations in electron density around the ion cores, resulting in transient changes in the charge density distribution. This transient dipole in one particle induces an equal and opposite dipole in the other one, resulting in an attraction. It may be thought that all the transient dipoles in the randomly orientated particles would cancel each other and average to zero. This is indeed the case for an ensemble of particles. However, the correlation between the dipole in one particle and the induced dipole in another at any time is not zero—the correlations, and therefore the attractive interactions, do not average out to zero. Such charge fluctutations are a universal property of matter and occur even in a completely nonpolar material. If there is a permanent dipole in (some of) the interacting particles, these dipoles will also contribute to the van der Waals interaction. For purely nonpolar particles, the interaction is known as the London, or dispersion, energy. The van der Waals interaction between atoms or molecules, E, varies as the inverse sixth power of the distance, d, between them: 1 E  d6

(1.39)

(The minus sign signifies an attractive interaction.) For macroscopic (this includes microscopic and nanoscopic) bodies, this interaction is much less short-range, and the distance dependence varies both with the geometry of the interacting bodies and with the distance of separation. For macroscopic bodies, it is usually assumed

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TABLE 1.3 Interaction Energies and Forces of Attraction Between Two Bodies with Different Geometries Geometry

Energy

Force

16Ar 31r 32  9d 6

32Ar13r 23  3d 7

16Ar6  9d 6

32Ar 6  3d 7

Ar1r2  6(r1  r2)d

Ar 1r 2 2 6(r 1  r 2)d

Ar  12d

Ar 2 12d

2Ar3  9d3

2Ar3  3d 4

Ar  6d

Ar 2 6d

A is the Hamaker constant, r (r1, r 2) the radius of the spherical particles, and d the distance between surfaces of the two bodies. Note that larger particles will interact more strongly (more adherent films?).

that the interactions between all the different bodies are additive. Table 1.3 shows values of this interaction for various geometries of two interacting bodies relevant to CD. These can be divided into particle–particle attraction (formation and growth of aggregates) and particle–plane surface (i.e., the substrate) interaction. Two different distance scales are shown, depending on whether the separation is considerably larger or smaller than the radius of the particle. Clearly there will also be intermediate separations, where the separation and radius are comparable, with intermediate dependence on the separation. Cases where two interacting spheres are identical or of different size are also shown. Both are relevant for CD, where initial aggregation will occur between elementary colloids of approximately the same dimensions but further interaction can occur between two particles of very different sizes. The Hamaker coefficient, A, is a measure of the interaction and is dependent on the material of the particle as well as on the surrounding medium. Heavy atoms, which are generally more polarizable (i.e., the electron distribution can be more

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easily perturbed), generally have a larger value of A than lighter atoms and therefore a greater attractive interaction. A has units of energy and values that vary typically from several times the thermal energy, kT, to several tens of kT in air or a vacuum and typically an order of magnitude less in liquid media. At large distances between particles, correlation between fluctuations in one particle and the induced dipole in another breaks down. This occurs when the time taken for the interaction (acting at the speed of light) is comparable to the characteristic scale of the electron fluctuations, viz. the plasma frequency. The plasma frequency ranges typically from 10 eV down to 2 eV (closer to the former for many dielectrics and to the latter for metals), which translates into a length scale of between 600 and 100 nm. At this distance scale, the (at this point, very weak) van der Waals forces are termed retarded forces, because of the appreciable time required for the transmitting dipole electromagnetic field to reach the receiving species. At this distance, the attraction falls off approximately as the inverse seventh power of the distance. It is probable that diffusional and convective motion and electrostatic interactions will dominate at such distances and the van der Waals interaction will be negligible in most cases. These various relationships between force and particle separation imply that the attractive force between particles will become infinite when they touch. In reality, other short-range forces will modify this relationship when r is very small, in particular the repulsion from overlap of atomic orbitals. The van der Waals attraction will then be balanced by this overlap repulsion. At these short distances (a few tenths of a nanometer), the van der Waals attraction will be strong enough to hold the particles fairly strongly together. This balance between van der Waals forces of attraction and overlap repulsion forces is shown schematically in Fig. 1.4, where the very steep repulsive interaction at atomic distances is due to the overlap repulsion. Hydration forces (see section 1.3.3) may also result in repulsion between surfaces at somewhat greater separations. Particle adhesion occurs when the distance between bodies is that of an atomic spacing. From Table 1.3, the force between a sphere of radius r and a flat surface at close approach is Ar  6d 2

(1.40)

At contact, d is the atomic spacing. For a solid where the van der Waals forces dominate, the work needed to separate two unit areas from contact to infinity is given by A  12 z2

(1.41)

where z is the atomic spacing [
d in Eq. (1.40) at contact]. This is the energy required to produce two new surfaces, i.e., 2, where  is the surface energy of the

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FIG. 1.4 Resultant interaction energy between two particles with van der Waals attractive interactions and electron overlap repulsion interactions.

solid. Therefore, A 
 24 z2

(1.42)

Substituting Eq. (1.42) into Eq. (1.40) gives the force of adhesion at contact in terms of the surface energy: F
4 r

(1.43)

This relation is clearly very simplified, being based on a number of approximations, such as the validity of the use of the Hamaker constant at such close distances and the particle and surface being of the same material. Also, the relationship between surface force and van der Waals forces does not hold for many solids, in particular for metals where metallic bonding is important. Nonetheless, if taken as an indication of the forces holding particles to each other and to surfaces, it does give a feel for these forces.

1.3.2 Electrostatic Forces For solid particles dispersed in a liquid medium, there exists, in most cases, a layer of charge separation at the phase boundary—the electrical double layer. A num-

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ber of processes can cause this double layer. One of the most common is adsorption of charged species at the solid–liquid interface. For example, in a colloidal sol of CdS prepared by precipitation of a Cd salt with sulphide ions, Cd2 or S2 (HS), depending on which is in excess, will adsorb at the CdS surface. Another mechanism for formation of surface-charged species is surface dissociation. A common example of this is the case of metal oxides in water; the water may dissociate at the oxide surface as follows:   D MOH  OH

D

MMO  H2O

MOMOH  H

(1.44)

leaving a positively or negatively charged surface. Such a reaction will, obviously, be very pH dependent. Yet another possibility for formation (or change) in a double layer is by accumulation of one charge type. This may occur by doping the solid with an ion of valence different from that of the solid (e.g., In3 in CdS) or by illumination with super-bandgap illumination. In the first case, the (in this case) electron donor (In3) is immobile, while the donated electron is mobile; if the electron is transferred to the liquid, then the solid will become positively charged. The same occurs for an illuminated solid where electron/hole formation occurs, if one of the charges is preferentially injected into the liquid. Even in the absence of charge injection into the liquid, localized (e.g., trapped) charges will affect the double layer if the countercharge is (relatively) delocalized over the particle. The double layer close to a (near) surface-localized donor will be different than that for the rest of the particle. Such an effect is probably not important in large particles, where such fluctuations can be evened out. For very small particles, however, where only a single “dopant” may exist, this effect may be appreciable. Even for a pure and perfectly stoichiometric particle, the double layer need not be homogeneous around the particle. For example, a CdS particle will consist of different crystal faces. Most notably, the opposite polar faces, consisting of only Cd or of S atoms, can be expected to possess different double layers. Incidentally, these polar faces might be expected to attract other polar faces of opposite polarity and repel those of the same polarity. Such an effect would lead to some form of self-assembly. However, any effect of this nature will be much smaller (if it exists at all) in solution compared to vacuum or air, due to adsorption of ions from solution onto the polar faces, which will tend to neutralize this effect. The charge at the surface of a solid (including any adsorbed species) will be balanced by a countercharge in the electrolyte; the double layer as a whole is electrically neutral. The countercharges remain in the vicinity of the surface-adsorbed charge but, due to thermal motion, do not accumulate at the surface but move in a more or less diffuse cloud surrounding the particle. The extent to which this layer

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of countercharge extends into the solution—how diffuse it is—depends on the concentration of charged species in the solution. This gives the solution screening length. The screening length is analogous to the space charge layer width in a semiconductor; in the same way, the screening length, , is a function of the square root of the charge (ionic) concentration.



2e2z 2n 
 0kT



1/2

(1.45)

where z is the ionic charge, n is the ionic concentration, 0 is the permittivity of free space, and  is the dielectric constant of the material. (The dielectric constant is normally taken as constant. It should be pointed out, however, that for nanoparticles of several nanometers or less, the value of  decreases with particle size, assuming the particles are in a medium of smaller dielectric constant than the particles themselves, a reasonable assumption for our purposes. This effect is treated in some detail by Lannoo et al. [4].) The thickness of this diffuse layer (also known as the Gouy layer) is the inverse of ; i.e., the potential drop across the diffuse layer, d, decays to d/e (where e in this case is the natural exponent, 2.718) over a distance 1. Figure 1.5 shows a schematic representation of the double layer at a planar solid–liquid interface. The potential drop across the Helmholz layer is shown as linear (in the presence of specific adsorption, it will not be completely linear), followed by a tailing-off of the potential into the diffuse layer. For concentrated solutions (0.1 M) the diffuse layer is typically a nanometer or less, while for dilute solutions it may be tens or even hundreds of nanometers.

FIG. 1.5 Schematic diagram of the electrical double layer at a solid–liquid interface.

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The crucial importance of the double layer when dealing with colloidal particles dispersed in a solution is due to the repulsion of one particle by another. While overall the particles are neutral, because the diffuse layer can extend into the solution, the unbalanced charge in the diffuse layer of one particle experiences a repulsion by that of another particle. Normally, from the Coulomb law of electrostatics, the force between two equal (in charge and in sign) particles is given by q1q2 F
 (1.46) 4 0d where F is the force acting between two charges, q1 and q2, separated by a distance d. Because of the presence of the diffuse layer, however, the repulsion force between two particles is strongly dependent on the screening length, , and is approximately proportional to exp (-d); the force of repulsion between two colloids will decrease exponentially with distance. Overlap of the diffuse layers of approaching particles prevents them from getting close to each other, and the particles form a stable colloidal solution. More relevant for our purposes, for a moderately concentrated electrolyte of the type normally encountered in CD (on the order of 0.1 M or more total concentration), this diffuse layer is around 1 nm or less. The diffuse layer screens the surface charge and allows the like-charged particles to approach each other closely, to the point where the van der Waals forces of attraction dominate, causing aggregation. This is the basis of salting-out of a colloid; addition of a strong electrolyte to the colloid reduces the thickness of the diffuse layer, allowing closer approach of the particles to each other and eventual aggregation and precipitation. The competition between the attractive van der Waals and repulsive electrostatic forces and the importance of this competition in colloid stability is known as the DLVO theory, named for the scientists who developed a theoretical analysis of the overall interaction (Derjaguin, Landau, Verwey, and Overbeek). The resultant interaction is shown in Fig. 1.6. For a dilute electrolyte, there is a relatively large barrier to aggregation where the double-layer repulsion dominates the interaction to the greatest extent and the interacting particles fall into what is known as the secondary minimum. At this point they are kinetically stabilized against aggregation. As the electrolyte concentration increases, the barrier becomes smaller and eventually disappears, resulting in the particles becoming trapped in the primary minimum, i.e., aggregation. It is worth noting that the electrostatic force can be attractive as well as repulsive, depending on the sign of the two charges. For the case where a single colloidal species is present in the solution, it will be repulsive, since all the particles will have the same charge. Two different colloidal species of opposite charge could conceivably be present in the deposition solution, either because the CD process involves conversion of one species into another (e.g., metal hydroxide and metal chalcogenide) or because two or more different cations, or even different valence states of the same cation, are present.

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FIG. 1.6 DLVO interactions showing the energetics of colloidal particles as a competition between electrostatic double-layer repulsion and van der Waals attractions. The primary minimum is due to strong short-range electron overlap repulsion (shown in Figure 1.4 but not shown here).

1.3.3 Entropic and Other Short-Range Forces Apart from double-layer, van der Waals, and electron-overlap interactions, any two bodies in a liquid medium (even if only one, or even neither, is charged) will experience a (usually) repulsive component of force as they approach each other. This is due essentially to entropic considerations. As two bodies approach each other very closely, the species in solution have increasingly less room in which to move; the entropy of these species therefore decreases, producing a repulsive force between the two bodies. In its simplest form, this force is usually considered to exist between two infinite flat plates. In the context of forces between colloidal particles of the type common in CD, this entropic force should be less important, since the solution species between two approaching particles can be relatively easily pushed out of the intervening space, both due to the small size and due to the approximately spherical shape common in these systems. In fact, in this case there may even be a weak attraction—the depletion interaction—due to the smaller density of solution species in the space where the two particles are closest to each other (from which they have been pushed out) and the surrounding solution. The entropic force of repulsion may be important, however, when considering the sticking behavior of larger aggregates at plane surfaces.

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This entropic force is important where adsorption of polymers occurs on colloidal particles. This is due to interaction between polymer chains on the interacting particles: As the particles approach each other to the point where the polymer chains of the two particles interact, there is an decrease in entropy due to confinement of the chains, in an analogous manner to the solution species discussed earlier, with the same result—repulsion. This is the basis of polymeric stabilization of colloids; it is generally undesirable in CD, since adhesion and aggregation are preferred in this case. However, in view of the fact that the presence of such polymers (and other stabilizing adsorbates) may prevent the aggregation needed to build up a CD layer, it is important to be aware of the effect. There are other close-range forces related to entropy changes, including various interactions between solution species and a solid surface, such as solvation (in water, hydration) forces. Hydration forces can occur when hydrated cations are adsorbed at interacting surfaces. As these surfaces approach each other closely, loss of water of hydration is necessary in order to allow closer approach. While these forces can be repulsive, attractive or oscillating, they are most likely to be repulsive under the conditions of CD. Such forces may be very important for CD, which is almost always carried out in the presence of a high ionic concentration. For example they could be a cause of poor adhesion of some CD films. Solvation forces are treated in detail in Israelachvili’s book—see Further Reading at the end of this chapter, “Forces” subsection. If this treatment of forces between particles fails to convince the reader that it is natural for particles to stick together, one can resort to the more intuitive approach. It is well known that inorganic colloids require a stabilizing agent to prevent their sticking together and eventually precipitating. In other words, precautions usually have to be taken to prevent the natural tendency of these particles eventually to stick to each other.

1.4 CHARACTERIZATION TECHNIQUES— SOME CAUTIONS Many techniques have been used to characterize CD films. The purpose of this section is not to review all these techniques, but only to draw attention to some of them that are sometimes misinterpreted. Three common techniques used are transmission electron microscopy together with electron diffraction, powder X-ray diffraction, and optical absorption (or transmission) spectroscopy.

1.4.1 Transmission Electron Microscopy/Electron Diffraction and X-Ray Powder Diffraction Transmission electron microscopy (TEM) is used to image nanocrystal (lateral) size, shape and size distribution. Electron diffraction (ED) provides information

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on the composition of the deposit, crystal phase, and orientation. X-ray diffraction (XRD) also provides similar information on composition (more accurately than ED) and phase as well as crystallite orientation. In the last, ED is superior in many ways, since a much smaller area can be selected (selected area diffraction—SAD) and, in addition, azimuthal lattice alignment between deposit and substrate (epitaxy) can be determined from ED but not from the commonly used powder XRD measurements; powder XRD reveals texturing (one particular crystal face parallel to the substrate for all crystals for perfect texturing) but not orientation (crystal lattices in any direction parallel to the substrate of all crystals aligned in the same way). There are XRD measurements that can distinguish orientation, but while becoming somewhat more common, these are still rather infrequently used, at least in CD studies, compared to the normal “–2” measurement. CD films are often nanocrystalline. One very important use of XRD when dealing with nanocrystals is to estimate crystal dimensions through the Scherrer relationship: 1.3 crystal diameter
 (2) cos 

(1.47)

where  is the X-ray wavelength (0.1541 nm for Cu K radiation, a commonly used source), (2) is the peak full width at half maximum (FWHM) in radians, and  is the peak position.* The shape of the crystal can also modify this relationship, which is valid for a spherical crystal (close to the shape often encountered). As a rough and useful rule of thumb, a peak FWHM of 1° at an angle of 2
25° (a common approximate position) means a crystal size of ca. 12 nm (for Cu K radiation), and the size is inversely proportional to the FWHM. Actually, to be more precise, what is measured is not necessarily crystal size but coherence length, the length over which the periodicity of the crystal is complete. An example of a coherence length smaller than the crystal size is a twinned crystal; XRD measures the size of each individual twin. Other causes for XRD peaks being broader than expected based on crystal size is the presence of strain in the crystals or other defects, such as dislocations, which destroy the long-range lattice order. Separation between crystal size and strain can be made if several different peaks are present, since the angular dependences of the two factors are different (see Ref. 5 for an example of this). Thus, the interpretation of XRD peak broadening should be carried out with care and preferably using complementary TEM measurements. The opposite case, where the peaks are narrower than expected based on crystal size, does not occur; a narrow peak means a (relatively) large coherence length and * Note that in the common /2 measurement, the  in cos  is half of the 2 value. For example, if the peak being measured is at 2
25°, cos  will be cos 12.5°. For small angles, the error in taking cos 2 instead of cos  is not too large (10% at 2
25°, for example). However this error becomes larger as the peak angle increases.

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therefore crystal size. However, even here, interpretation is not always straightforward. For example, the XRD pattern of a deposit of tall cylinders of small cross section will give a peak width characteristic of the height but not of the cross section (the latter will be seen in TEM images). Thus the TEM and XRD sizes will necessarily be different in such a case. Another example is where there is a mixture of large and small crystals. Even if the large crystals constitute a relatively small fraction of the total material, they may in some cases dominate the XRD pattern, since peak heights decrease with decreasing crystal size due to increase of peak width and (ideally) constant peak area for the same quantity of material. If the crystal size becomes very small (a few nanometers), the XRD peaks will be very wide and also relatively weak. There is no shortage of examples in the literature where samples have been classified as “amorphous” or “poorly crystalline” either on the basis of the lack of any XRD peak or because the peak(s) were very broad. When carrying out an XRD measurement on a CD film, in particular, a particularly thin one (some tens of nanometers or less), if no peaks are seen in the measurement, it is advisable to repeat the measurement over a narrow range (where a major peak is expected) and with a very slow scan (e.g., 10°/hr or even slower). If a thin-film attachment is available, this will reduce the likelihood of such misinterpretations. It is useful to remember that except for compounds that are commonly amorphous, CD semiconductor films are rarely truly amorphous.

1.4.2 Optical Absorption Optical absorption spectroscopy is often carried out on CD films to verify that the films have a bandgap expected from the deposited semiconductor. Additionally, since CD films are often nanocrystalline and the most apparent effect of very small crystal size is the increasing bandgap due to size quantization (the effect is visible to the eye if the bandgap is in the visible region of the spectrum), absorption (or transmission) optical spectroscopy is clearly a fast and simple pointer to crystal size, since bandgap–size correlations have been made for a number of semiconductor colloids and films. There are some potential problems that should be taken into account when interpreting such spectra. A spectrophotometer measures transmission (and maybe also reflection) but not absorption. What is measured as absorption is a transmission measurement that is mathematically manipulated to convert it to absorption. Absorption is usually measured as absorbance, A, which by definition is given by A
log10 Io/I

(1.48)

where I is the intensity of the transmitted light and I0 that of the incident light. The transmission, T, is I T
 I0

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(1.49)

The spectrophotometer measures the transmission and, if an absorption measurement is carried out, converts the transmission into absorbance using these equations. This conversion works fine for samples where there is no reflection, either specular or diffuse, as is the case for nonturbid solutions. However, for films there is invariably some reflection, which is often quite large, particularly for films of high dielectric constant (or refractive index) materials, such as PbS and PbSe. Additionally, if the films are not completely transparent, then scattering introduces an extra element of reflection. Therefore, to measure the real absorption of a film, a reflection measurement must also be carried out and correction for this reflection made. The correction will be approximate and depends on the nature of the film itself. However, that most commonly used is T Tcorr
 1R

(1.50)

where Tcorr is the corrected transmission, T is the measured transmission, and R is the reflectance. This correction neglects reflection from the film/substrate interface (assuming the front face of the film faces the illumination source), and it can be calculated that this will give a value of Tcorr that is too small (therefore an absorption that is too high). Use of (1  R)2 in the denominator of Eq. (1.50) includes this reflection but tends to give a value of Tcorr that is too high. For very thin films, where reflection cannot be assumed to originate from a surface (i.e., where the film thickness is not much greater than the depth from which reflection can occur), the calculation is more complicated. Fortunately, reflection from such thin films is also normally low, and therefore the correction is less important. To collect scattered transmission and correct for diffuse reflectance, a spectrophotometer with an integrating sphere should be used. This is important if films are not very transparent. In many cases, the lack of correction for reflection will not affect the shape of the optical spectrum very much, but it will just give an inflated value for absorption. However, there are frequently cases where the shape of the absorption spectrum is also appreciably changed after the reflection correction is carried out. Also, if the primary absorption is a weak one, then correction for reflection, and in particular for scattering, is crucial, since the absorption may be masked by these effects. The absorption coefficient, , of the semiconductor can be derived from the absorption (or transmission) spectrum according to the Beer–Lambert equation applied to solids: I
I0 et

(1.51)

where t is the film thickness. For bulk semiconductors, the relationship between  derived from the absorption spectrum and the semiconductor bandgap, Eg, is given by: (h)n
C(h  Eg)

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(1.52)

where n is 2 for a direct transition and 0.5 for an indirect one and C is a constant. A plot of (h)n vs. h should then give a straight line (over much of the absorption onset region), which extrapolates at a zero value of (h)n to the value for Eg. This calculation is based on the density of states in the valence and conduction bands of the bulk semiconductor. For semiconductors in the quantum size regime, however, the density of states may be quite different than in the bulk. Additionally, a large size distribution, meaning distribution of bandgaps, will smear out the onset. In practice, however, this extrapolation often appears to give a reasonable value of the bandgap. The important thing is to be aware of the limitations of the measurement. Just because a few points on a plot based on Eq. (1.52) give a straight line does not automatically mean that the bandgap can be obtained from the extrapolation of this line; not too infrequently, a weaker absorption onset at longer wavelengths has been ignored, although consideration of the entire spectrum, together with the expected behavior of the material, would lead one to conclude that this weaker absorption determines the bandgap.

FURTHER READING Solution Chemistry Just about any textbook on inorganic chemistry. A particularly useful one, which the author referred to on many occasions while writing this chapter, is: G Wulfsberg. Principles of Descriptive Inorganic Chemistry. Mill Valley, CA: University Science Books, 1991.

For more extensive tables of solubility product than given here, Lange’s Handbook of Chemistry (J.A. Dean, 11th ed. New York: McGraw-Hill, 1999) gives an extensive list. The standard CRC Handbook of Chemistry and Physics (CRC Press) also gives a useful, if less extensive, list. Reference 1 also provides an extensive list of sulphides.

Nucleation and Growth HC Freyhardt, ed. Crystals: Growth, Properties, and Applications. New York: SpringerVerlag, 1983, Vol. 9 (this is more mathematical than the other treatments). HK Henisch. Crystal Growth in Gels. University Park: Pennsylvania State University Press, 1973. AE Nielsen. Kinetics of Precipitation. New York: Pergamon Press, 1964. BN Roy. Crystal Growth from Melts. New York: Wiley, 1992.

Forces F Evans, H Wennerström. The Colloidal Domain: Where Physics, Chemistry, Biology, and Technology Meet. New York: VCH, 1994.

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PC Hiemenz. Principles of colloid and surface chemistry. New York: Marcel Dekker, 1986. J Israelachvili. Intermolecular and Surface Forces. Orlando, FL: Academic Press, 1992. H Ohshima, K Furusawa, eds. Electrical Phenomena at Interfaces: Fundamentals, Measurements, and Applications. New York: Marcel Dekker, 1998.

REFERENCES 1. 2. 3. 4. 5.

S Licht. J. Electrochem. Soc. 135:2971, 1988. P O’Brien, DJ Otway, D Smith-Boyle. Thin Solid Films 361:17, 2000. M Skyllas-Kazacos, JF McCann, R Arruzza. Appl. Surf. Sci. 22/23:1091, 1985. M Lannoo, C Delerue, G Allan. Phys. Rev. Lett. 74:3415, 1995. SB Qadri, JP Yang, EF Skelton, BR Ratna. Appl. Phys. Lett. 70:1020, 1997.

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2 General Review

2.1 A (BRIEF) HISTORY OF CHEMICAL DEPOSITION Chemical deposition (CD) of films is not a new technique. As early as 1835, Liebig reported the first deposition of silver—the silver mirror deposition—using a chemical solution technique [1]. The first reported CD of a compound semiconductor film appears to be formation of “lüsterfarben” (lustrous colors) on various metals from thiosulphate solutions of lead acetate, copper sulphate, and antimony tartrate, giving films of PbS, Cu-S or Sb-S, which possessed “splendid” colors (interference colors resulting from various thicknesses of the deposited films) [2]. More “recent” studies of this general process have invoked an electrochemical mechanism for some thiosulphate depositions, based on the dependence of deposition on either the nature (standard electrochemical potential) of the metal substrate or on a contacting non-noble metal (which can be looked at as an internal electrochemical deposition) [3–5]. However, while it is probable that an electrochemical or mixed electrochemical/chemical mechanism may be applicable on some metal substrates, some of these depositions do appear to be true CD processes. PbS is probably the clearest of these; others were Cu2S, Ag-S, Bi-S, Sb-S. Fe, Ni, and Co all formed apparent sulphide films on Fe substrates, while ammonium molybdate deposited a film from a thiosulphate solution that did not contain S and was probably an oxide. Beutel and Kutzelnigg cover a wide range of depo-

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sitions from thiosulphate solutions—both CD and electrochemical [5]. Only in a few cases were these films characterized other than by color. In 1884, Emerson-Reynolds reported deposition of PbS films by reaction between thiourea (thiocarbamide) and alkaline lead tartrate, where “the metallic sulphide . . . became firmly attached as a specular layer to the sides of the vessel” [6]. A wide range of substrates, apart from that just mentioned (a glass beaker), was successfully used for this deposition; porcelain, ebonite, iron, steel, and brass were specifically mentioned. Even more important, the deposits were very adherent, as quantified by their ability to “withstand considerable friction with a washleather, and under this treatment take a fine polish.” Infrared photoconductivity in CD PbS films was reported nearly a century ago [7,8], and this application has been a central driving force for subsequent investigations in CD lead chalcogenide films. The early literature invariably mentions the pioneering work of Kutscher in Germany, during World War II, in developing CD PbS and PbSe films for infrared detectors. However, the apparent lack, in all these references to Kutscher’s studies, of any published papers might suggest (to the overly suspicious reader) a possible military involvement in these studies. These (and subsequent) studies succeeded to the extent that CD was, and apparently still is, the main technique used in making commercial PbS and PbSe infrared detectors (vacuum evaporation was the only competing technology) [9,10]. For a long time, CD was then essentially limited to PbS and PbSe. It was not until 1961 that deposition of CdS, now the most widely studied material in CD, was explicitly reported [11] (although CdS deposited from a thiosulphate solution which “sticks obstinately to the glass” was already noted in 1912 [11a]). The range of materials deposited by CD was gradually extended, particularly in the 1980s, to include sulphides and selenides of many metals, some oxides, and also many ternary compounds (Tables 2.1 and 2.3 in this chapter list films deposited by CD). Chemical deposition received a major impetus after CdS films, chemically deposited onto CdTe (and, later, onto CuInSe2) films, were shown to give superior photovoltaic (PV) cells compared with the previously evaporated CdS. The first reference to CD CdS used in thin-film PV cells appears to be from Uda et al. [12], although no special importance was attached to the CD technique in that paper. Birkmire et al. showed that CD CdS was as good as evaporated (Cd,Zn)S as the heterojunction partner in CuInSe2-based thin-film cells, giving 10.6% efficiency [13]. Two years later, the efficiency of CuInSe2 cells using CD CdS had increased to 12.8% [14]. In 1991, Chu et al. used CD CdS to make high-efficiency (13.4%) CdTe/CdS thin-film cells, explicitly stressing the beneficial role of the CD CdS [15] and followed this a year later with a 14.5% cell using the CD CdS [16]. Nowadays, Cd is almost universally used to form the CdS layer on both CdTe and CuIn(Ga)Se2 thin-film PV cells.

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Another cause of interest in this technique is due to the fact that the crystals in most as-deposited CD films are very small. Considering the current interest in nanoparticles, CD is an excellent technique to deposit nanocrystalline films. More specifically, if the nanocrystals are small enough, they exhibit size quantization, the most obvious manifestation of which is an increase in the optical bandgap with decrease in crystal size, as was shown for CD CdSe [17] and later for CD PbSe [18,19]. In fact, the changes in optical spectra that occurred in these films as a function of nanocrystal size were exploited to provide information on the different mechanisms of the deposition process [20]. Chemical deposition has also been emphasized as a technique to form solar control coatings. Solar control coatings are envisaged for use on windows in hot climates and possess the (ideal) characteristic of moderate to high visible transmission to provide adequate lighting, together with high infrared (0.7–2.5 m) reflectance to minimize heating by solar energy. CD is a potentially suitable method to prepare these coatings on the large areas of glass that would be needed. Most of the work in this field has been carried out by Nair and Nair in Mexico using various semiconductor films, mainly PbS [21–23] and CuxS [23]. See also this group’s recent review on this work [24]. These coatings are normally yellowish or neutral by transmitted light and various shades of gold, blue, or purple by reflected light.

2.2 WHAT IS CHEMICAL DEPOSITION? Chemical deposition refers to the deposition of films on a solid substrate from a reaction occurring in a solution (almost always aqueous). Using the prototypical CdS as an example, a Cd salt in solution can be converted to CdS by adding sulphide ions (e.g., as H2S or Na2S); CdS immediately precipitates (unless the solution is very dilute—a few millimolar or less, in which case CdS often forms as a colloidal sol). Another pathway for CdS formation, one that does not require free sulphide ions, is decomposition of a Cd-thiocomplex (a compound that binds to Cd through a sulphur atom). In CD, the trick (or at least one of them) is to control the rate of these reactions so that they occur slowly enough to allow the CdS either to form gradually on the substrate or to diffuse there and adhere either to the substrate itself (at the early stages of deposition) or to the growing film, rather than aggregate into larger particles in solution and precipitate out. This rate control can be accomplished by generating the sulphide slowly in the deposition solution. The rate of generation of sulphide, and therefore reaction rate, can be controlled through a number of parameters, in particular the concentration of sulphide-forming precursor, solution temperature, and pH. The CdS forms through a number of different possible pathways: simple ionic reaction between Cd2 and sulphide ion; topotactic conversion of Cd(OH)2, which may be present in the deposition solution, to CdS by sulphide; and decomposition of a complex between Cd (whether as a free ion or as a Cd compound,

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e.g., Cd(OH)2) and the sulphide precursor (often thiourea, which, like other chalcogenide precursors, also acts as a complexant for metal ions). Although CD can be carried out in both acidic and alkaline solutions, most CD reactions have been carried out in alkaline solutions. This is necessary for selenide deposition using selenosulphate (see later), which is unstable in acid solution (of course, the chalcogenide precursor must not be too stable under all conditions, otherwise they will not work). Therefore to prevent (at least bulk) precipitation of metal hydroxides in the deposition solution, the metal ion must be complexed. There is a very wide range of possible complexing agents available; the most used are intermediate in complexing strength—not too weak, in order to prevent bulk precipitation of hydroxide, but not too strong, which may prevent deposition of the desired film altogether.

2.3 WHAT MATERIALS CAN BE DEPOSITED BY CD? In principle, CD can be used to deposit any compound that satisfies four basic requirements. The compound can be made by simple precipitation. This generally, although not exclusively, refers to the formation of a stoichiometric compound formed by ionic reaction. The compound should be relatively (and preferably highly) insoluble in the solution used (except in a very few cases, this has been water). The compound should be chemically stable in the solution. If the reaction proceeds via the free anion, then this anion should be relatively slowly generated (to prevent sudden precipitation). If the reaction is of the complex-decomposition type, then decomposition of the metal complex should similarly occur relatively slowly (see Sec. 2.5 for a description of reaction mechanisms). Of course there are other specific factors that need to be taken into account, particularly whether the compound will form an adherent film on the substrate or not. However, the preceding four factors are general requirements.

2.4 FORMATION OF THE (CHALCOGENIDE) ANIONS One of the requirements for CD is either slow release of the anion—in most cases a chalcogenide anion—or slow decomposition of a suitable complex containing a chalcogenide atom. This section is confined to the former and discusses slow anion release. The precursors used up to now, together with some of the reactions leading to anion formation, will be briefly described (more details can be found in Chap. 3).

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2.4.1 Oxide Many oxides have been deposited by CD using many different techniques, some of which are described below. Section 2.9.4.1 provides a wider overview of these methods. Often, what is deposited is a hydroxide or hydrated oxide. In many cases, hydroxide ion is generated slowly. There are a number of methods to do this, the most common being hydrolysis of urea: (NH2)2CBO  2H2O → (NH4)2CO3

(2.1)

where the carbonate formed partially hydrolyses to give OH. Another technique is reduction of nitrate to nitrite by an alkylamineborane:    NO 3  H2O  2e (from alkylamineborane) → NO 2  2OH

(2.2)

Additionally, hydroxide may initially be present, but the reaction may be slowed down e.g. by complexation of the metal ions or by carrying out the reaction at low temperature. The hydroxide reacts with the metal ion to give the metal hydroxide, a hydrated oxide or an oxide, depending on the chemistry of the particular metal(hydr)oxy system. The hydroxides or hydrated oxides can be heated in air or oxygen to form the oxides. Another method used to deposit oxides, particularly those with a higher oxidation state than the starting cation, uses persulphate, S2O 2 8 , a strong oxidizing agent. While the exact mechanism of oxide formation using persulphate is unclear, it appears to involve internal electrochemical reactions; e.g., for PbO2: Pb2  2H2O → PbO2  4H  2e 2  S2O2 8  2e → 2SO 4

(2.3a) (2.3b)

However, it is possible that free radicals, such as · OH, are involved, since persulphate hydrolysis can proceed with formation of H2O2, which is itself sometimes used to deposit oxides:  2 S2O2 8  2H2O → 2SO 4  H2O2  2H

(2.4)

2.4.2 Sulphide Thiourea (SC(NH2)2), the sulphur analogue of urea, is the most commonly used sulphur precursor. There are a number of possible decomposition routes for thiourea in aqueous solution (it is invariably used in alkaline solutions). Probably the most important is SC(NH2)2  OH D HS  CN2H2  H2O

(2.5)

which generates sulphide ion (the cyanamide, CN2H2, can hydrolyze further, but this need not concern us at present). Actually, aqueous solutions of thiourea are

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not very unstable; the presence of a cation that can precipitate an insoluble sulphide is necessary for the decomposition to proceed at a reasonable rate. This implies that reaction (2.5) is actually an equilibrium; the metal ion removes sulphide ion, which drives the equilibrium continually to the right. Thioacetamide (H3C.C(S)NH2) has also been commonly used in CD. It has the advantage, compared with thiourea, that it works in both acid and alkaline solution. A general decomposition reaction for sulphide formation is H3C.C(S)NH2  2H2O → CH3COOH  H2S  NH3

(2.6) 

In alkaline solution, the sulphide will be in the form of sulphide (HS and S2) ions. Thiosulphate (S2O32) was the original sulphur source in early CD processes, and, while less commonly used nowadays, it still has a place in the modern CD literature. It is most commonly used in somewhat acidic solutions, although it has also been employed in alkaline solution. Thiosulphate is unstable in fairly acidic solutions and decomposes to give elemental sulphur, e.g.,   S2O 2 3  H → S  HSO 3

(2.7)

It is often suggested that the thiosulphate, a mild reducing agent, reduces this sulphur to sulphide. There are a number of other potential sulphide-forming reactions, depending on pH (see Sec. 3.2.1). It should be pointed out here that the SMS bond in thiosulphate is easily broken. In view of the strong complexes it forms with some metal cations, the probability of a mechanism whereby the SMS bond of the complex is broken, leading to metal sulphide formation without formation of sulphide ion, should be seriously considered (see Sec. 2.5).

2.4.3 Selenide Selenourea, the selenium analogue of thiourea, which hydrolyzes in the same way to give selenide ions, was once the most common source of Se. It is an unstable compound that requires the presence of a reducing agent—usually Na2SO3. Dimethylselenourea is more stable than selenourea but still difficult to work with. The most common Se precursor used nowadays is sodium selenosulphate (Na2SeSO3), which can be considered as the analogue of thiosulphate, with one S atom substituted by Se. It is much more stable (and cheaper) than selenourea and therefore simplified greatly the deposition of selenides. It can only be used in alkaline solutions (it decomposes at pH values lower than ca. 7 to precipitate elemental red Se). Its alkaline hydrolysis is usually given as   2 SeSO2 3  OH → HSe  SO 4

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(2.8)

although it is very probable that SO 2 4 is the end product of a more complicated reaction. As for most other mechanisms discussed earlier, Eq. (2.8) is probably also an equilibrium between selenide and some sulphur–oxygen intermediate, with the metal cation removing the selenide as for the thiourea decomposition in the presence of metal cations.

2.4.4 Telluride There are only a very few reports of telluride deposition by CD. This is due to a number of factors: the relative instability of the tellurium analogues of the Sand Se-forming precursors; the strong reducing conditions necessary to form telluride ions; and the rapid reaction of telluride ions with oxygen dissolved in the solution. CdTe films have been formed by using hydrazine—a strong reducing agent—in a deposition solution containing TeO2, and very thin films have been observed to form when H2Te was added to a solution of a Cd salt. Recently, the deposition of CdTe has been reported using the tellurium analogue of selenosulphate (tellurosulphate). While Te is apparently only very slightly soluble in sulphite, it is apparently enough to deposit tellurides. For both methods, stoichiometric films are more difficult to obtain than for sulphides or selenides, with elemental Te typically also formed. A very recent study has also described conversion of CD Cd(OH)2 films to CdTe using a solution of Te in hydroxymethane sulphinic acid.

2.4.5 Halides Halides (confined at present to silver halides) can be deposited by hydrolyzing a water-soluble halogeno-alcohol (halohydrin) to slowly form halide ions in the presence of Ag ions: X(CH2)nOH  H2O D X  H  HO(CH2)nOH

(2.9)

The solubility products of most halides are much higher in general than those of chalcogenides. Those of the silver halides are fairly low, which allows these depositions to take place readily.

2.4.6 Other Anions Although CD seems to have been limited to chalcogenides (including oxides and hydroxides) and isolated cases of carbonates, silver halides, and elemental Se, it should be possible to deposit salts of other anions. There are a number of other anions that can be slowly and homogeneously generated. These are discussed in Chapter 3.

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2.5 MECHANISMS OF CHEMICAL DEPOSITION 2.5.1 Introduction The mechanisms of CD processes can be divided into two different processes: formation of the required compound by ionic reactions involving free anions, and decomposition of metal complexes. These two categories can be further divided in two: formation of isolated single molecules that cluster and eventually form a crystal or particle, and mediation of a solid phase, usually the metal hydroxide. We consider first the pathways involving free anions and defer to later those where a metal complex decomposes. A starting point for discussing the mechanisms of CD is to consider a simple precipitation reaction. If H2S is added to an aqueous solution of a Cd salt, yellow CdS precipitates out immediately. H2S precipitates the sulphides of most cations (the alkaline and alkaline earth sulphides are soluble in water); this is the basis of the well-known (at least, in the author’s university days) inorganic analytical scheme. Such a precipitation will not, however, result in a film on a substrate or on the walls of the reaction vessel (actually, it may do so to a very slight extent but this film would be extremely thin). To form a visible film of CdS, conditions must be chosen so that bulk precipitation is prevented or at least slowed down drastically. This is the purpose of the chalcogenide precursors, discussed in the previous section. They slowly generate the chalcogenide, allowing slow formation of the metal chalcogenide (CdS in the present example). The formation of the film, based on the formation of chalcogenide ions, can occur by two fundamentally different processes. We continue to use CdS as the example.

2.5.2 Ion-by-Ion Mechanism The simplest mechanism, often assumed to be the operative one in general, is commonly called the ion-by-ion mechanism, since it occurs by sequential ionic reactions. The basis of this mechanism, illustrated for CdS, is given by Cd2  S2 → CdS

(2.10)

If the ion product [Cd2][S2] exceeds the solubility product, Ksp, of CdS (1028; Table 1.1), then, neglecting kinetic problems of nucleation, CdS will form as a solid phase (see Chap. 1). If the reaction is carried out in alkaline solution (by far the most common case), then a complex is needed to keep the metal ion in solution and to prevent the hydroxide from precipitating out (but see later). Since the decomposition of the chalcogenide precursor can be controlled over a very wide range (by temperature, pH, concentration), the rate of CdS formation can likewise be well controlled. Of course, the CdS should form a film on the substrate and (at least ideally) not precipitate in the solution. This

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aspect of the ion-by-ion mechanisms (and all other mechanisms) is treated in Section 2.6.

2.5.3 Hydroxide Cluster Mechanism It was stated earlier that complexation of the Cd was necessary to prevent Cd(OH)2 precipitation. However, very often (more often than realized), Cd(OH)2 (or metal hydroxides in general) are important reaction intermediates in the CD process. If the complex concentration is not high enough to prevent completely the formation of Cd(OH)2, then a relatively small amount of Cd(OH)2 may be formed, not as a visible precipitate, but as a colloid. Since Cd(OH)2 is colorless and colloids typically do not scatter light, unless they aggregate to a large extent (in which case a suspension is the result), this means that the Cd(OH)2 colloid may not be visible to the eye. The CdS is then formed by reaction of slowly generated S2 ion with the Cd(OH)2: Cd 2  2OH → Cd(OH)2

(2.11)

followed by Cd(OH)2  S2 → CdS  2OH

(2.12)

Reaction (2.12) occurs because Ksp for CdS (1028; Table 1.1) is much smaller than that for Cd(OH)2 (2  1014). Another way of looking at this is that the free energy of formation of CdS is more negative than that of Cd(OH)2. It has also been suggested that the hydroxide cluster can act as a catalyst for thiourea decomposition. In this case, sulphide formation will occur preferentially at the surface of the hydroxide rather than nucleate separately in the solution. Such a course is logical based on the previous discussion of the effect of metal cations on the equilibrium of thiourea decomposition. It should be borne in mind that the mechanism may change in the course of the deposition. As the metal is depleted from solution, the complex:metal ratio will increase and may pass the point where no solid hydroxide phase is present in the solution. In this case, the ion-by-ion process will occur (initially in parallel with the hydroxide mechanism, later maybe exclusively) if the conditions are suitable.

2.5.4 Complex-Decomposition Mechanism The chalcogenide precursors possess many talents. Apart from forming the chalcogenide ions, they also form complexes with metal ions. As noted at the beginning of this section, and ignoring the distinction between ion-by-ion and hydroxide cluster mechanisms treated previously, CD processes can be divided according to two basic mechanisms: participation of free sulphide ions (the

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commonly accepted mechanism in most cases, although this does not necessarily mean that it is always the correct one), and a pathway involving decomposition of a metal/chalcogen-containing complex without formation of free sulphide. Unfortunately, it is usually difficult to distinguish between these two processes, and the former is assumed more out of inertia than because of any clear proof. In the specific case of the CdS deposition using thiourea, a complex-decomposition mechanism has been proposed in a number of different investigations, based on kinetic studies of the film formation process. Here we can revert to intuition and suggest that, in the case of strong complexation between the chalcogen compound and the metal ion (e.g., as occurs between thiosulphate and Hg, Ag, and Cu), it may seem more logical for the fairly weak SMS bond to break than the very strong metal–chalcogen bond. To be fair, these very strongly complexed cations are also those whose chalcogenides have a very low solubility product, and therefore very little free sulphide would be needed to form those metal chalcogenides. As for the free chalcogenide processes, the complex-decomposition mechanism can occur either by an ion-by-ion (or molecule-by-molecule, since free ions need not be involved directly) pathway, e.g., 2  [SO3MSMHgMSMSO3]2  H2O → HgS  SO2 (2.13) 4  2H  S2O3

(the molecule of HgS can interact with other HgS molecules to form clusters and eventually crystals. Of course, they may also redissolve. These aspects are treated in the following section), or by a solid-phase intermediate, e.g., [Cd(OH)2]n  SC(NH2)2 → [Cd(OH)2]nMSMC(NH2)2 [Cd(OH)2]nMSMC(NH2)2 → [Cd(OH)2]n1CdS  CN2H2  2H2O

(2.14) (2.15)

with eventual exchange of all the hydroxide in the Cd(OH)2 to CdS.

2.6 NUCLEATION, ADHESION, AND FILM GROWTH Probably the least-known aspect of the CD process is what determines the nucleation on the substrate and the subsequent film growth. In considering this aspect, we will treat the ion-by-ion and hydroxide cluster mechanisms separately, although there will be many features in common. The principles discussed should be the same for both the free chalcogenide and the complex-decomposition mechanisms.

2.6.1 Ion-by-Ion Growth For nucleation to occur homogeneously in a particle-free solution by the ion-byion process, supersaturation, usually a high degree of supersaturation, is typically required (Chap. 1, section 1.2.1). The presence of a surface (the substrate or the

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walls of the reaction vessel) introduces a degree of heterogeneity that facilitates nucleation. For this reason, depositions that proceed via the ion-by-ion process tend to occur mainly on the substrate or other surfaces, rather than involving a large amount of precipitate typical of the hydroxide mechanism. The surface can be considered a catalyst for the nucleation. As discussed in Chapter 1, the most important force involved in adhesion of the deposit to a substrate in general is the van der Waals force of attraction. In the initial stages of growth, there may be specific chemical interactions between the deposit and substrate. For example, if gold is used as a substrate, S, Se, and many of their compounds interact chemically with the gold to form S(Se)–Au bonds. This would promote good adhesion of the deposit to the gold. There could also be chemical and electrostatic interactions between surfaces of the individual crystals. For example, the positive S(Se) face of polar crystals could bind to the negative metal face of an adjacent crystal if the relative orientations are suitable (in practice, this will probably not occur, since the crystal faces will adsorb solution species as they grow). However, the van der Waals interaction between the crystals in the strongly ionic solution is enough in most cases to ensure adhesion of the crystals to one another. The fact that reasonably adherent films can be grown on apparently unreactive substrates, such as plastics, and even on such an inert and hydrophobic material as Teflon suggests that while such specific interactions between the semiconductor and substrate may improve adhesion to the substrate, they are not essential for film formation. Once nucleation has begun on a substrate (this usually includes the inside walls of the reaction vessel), it generally becomes easier for the film to grow, since deposition usually occurs more readily on the nucleated surface than on the clean surface. The crystals will continue to grow until blocked by some process, such as steric hindrance by nearby crystals or adsorption of surface-active substances from the solution. The former is probably the dominant reason for growth termination in most cases.

2.6.2 Hydroxide Mechanism Nucleation of the chalcogenide is much simpler in this process, since a solid phase—the metal hydroxide (or other solid phase)—is already present and the process proceeds by a substitution reaction on that solid phase. In this case, the initial step in the deposition is adhesion of the hydroxide to the substrate. This hydroxide is then converted into, e.g., CdS, forming a primary deposit of CdS clusters. More Cd(OH)2 and, as the reaction proceeds, CdS and partially converted hydroxide diffuses/convects to the substrate, where it may stick, either to uncovered substrate (in the early stages of deposition) or to already deposited material. This is essentially the same process as aggregation, described in Chap-

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ter 1 in the section on forces, and again is a consequence of van der Waals attraction. Since the initial nucleation of hydroxide occurs homogeneously in the solution, the CdS also is formed homogeneously and therefore usually precipitates out in the solution to a large extent. This precipitation occurs if the isolated crystals aggregate to a sufficient extent to form large flocs. Film formation occurs when high-surface-energy particles (single nanocrystals or small aggregates) reach the substrate (or any other surface) before they precipitate out in the form of large aggregates. This aggregation and homogeneous precipitation can be minimized, in some cases even prevented, by judicious choice of deposition parameters. Thus, while extensive precipitation suggests a hydroxide mechanism (some precipitation can occur in the ion-by-ion process), its absence does not always mean that the ion-by-ion process occurs. An expected difference between ion-by-ion and hydroxide (or any other cluster) mechanisms is that in the latter, since colloids from the solution stick to the substrate surface, the crystal size is not expected to change greatly with film thickness (it may increase to some extent, since the colloids themselves can grow via an ion-by-ion process on the crystals). For ion-by-ion growth, it is likely that crystal growth occurs on nucleii already present on the substrate, and therefore crystal size can increase with increasing deposition. The foregoing description assumes adsorption of colloidal metal hydroxide from solution onto the substrate as the primary nucleation step. However, hydroxides can also adsorb on solid surfaces at pH values below that of bulk hy-

FIG. 2.1 Schematic diagram showing the probable steps involved in the ion-by-ion mechanism. A: Diffusion of Cd and S ions to the substrate. B: Nucleation of the Cd and S ions facilitated by the substrate to form CdS nucleii. C: Growth of the CdS nucleii by adsorption of Cd and S ions from solution and nucleation of new CdS crystals. D: Continued growth of CdS crystals, which adhere to each other through van der Waals forces (possibly also chemical interactions).

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FIG. 2.2 Schematic diagram showing the probable steps involved in the hydroxide mechanism. A: Diffusion of hydroxide colloidal particles to the substrate, where they adhere (B) and react with S ions (either generated homogeneously in solution or catalyzed by the hydroxide surface). This reaction results in exchange of the hydroxide by sulphide, probably starting at the surface of the colloid and proceeding inward (C). This reaction will occur both at the surface-adsorbed colloids and at those dispersed in the solution. Reaction will continue (as long as the supply of sulphide continues) until most of the hydroxide is converted to sulphide (D); eventually the primary particles of CdS will adhere to each other to form an aggregated film (E); usually the nonadsorbed particles will also aggregate and precipitate out of the solution.

droxide precipitation, an effect which has been related to the presence of an electric field at the substrate/solution interface. This will certainly affect the nucleation, since this can now occur only on the substrate and not in solution. It is not so clear whether it will affect further crystal growth or not. The basic features of the ion-by-ion and hydroxide cluster film-forming mechanisms are shown schematically in Figures 2.1 and 2.2, respectively. Film formation involving complex decomposition will proceed in a similar manner (Fig. 2.3 shows this for a molecule-by-molecule deposition).

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FIG. 2.3 Schematic diagram illustrating possible steps in the complex-decomposition mechanism. The complex (CdMSML, where L is a ligand or part of the S-forming species) decomposes to CdS on the substrate (possibly catalyzed by the substrate) and, to a greater or lesser extent, also homogeneously in the solution (A, B). The CdS nuclei formed grow by adsorption and decomposition of more complex species (C) until a film of aggregated crystals is formed (D) in the same manner as for the previous two mechanisms.

2.7 KINETICS OF THE DEPOSITION Due to the different pathways that can occur in the CD process, the kinetics can vary widely from one deposition to another, as reflected by the rather wide range of activation energies (not commonly measured, but measured often enough to draw some conclusions) found. Regarding the time taken for a deposition, some depositions can be completed in a few minutes or less, while others can proceed for days and still be far from termination. This section is only meant to give a general picture; Chapter 3 should be consulted for more specific details and examples. Kinetic studies on the growth of CD films show, in most cases, an induction period at the beginning of the process where no clearly observable growth occurs, an approximately linear growth region, and a termination step where no further growth occurs (see Fig. 2.4). Strangely enough, this type of growth kinetics often occurs regardless of the deposition mechanism. For the ion-by-ion growth, it is very simple to explain. Deposition begins only when the chalcogenide concentration is high enough to allow nucleation to occur—the induction time corresponds to this buildup of chalcogenide concentration. Growth then occurs on these initial nucleii, along with new nucleation—the approximately linear region of growth. As the limiting reactant is used up, growth will start to slow down and eventually stop due to depletion of the reactants. For the cluster mechanism, while growth and termination can be similarly explained, the induction period is less obvious. The hydroxide cluster can start to adsorb on the substrate immediately after immersion of the substrate in the deposition solution, yet experiments have shown that film growth often does not occur for some time. While the reason for this is not clear, it may be connected with the

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fact that the hydroxide particles often do not form a film, beyond some primary adsorption on a surface; only when reaction to form the metal chalcogenide occurs does film formation develop. In this case, we can again invoke the need for a minimum concentration of chalcogenide ion. Some studies of the deposition rate have suggested that the rate-limiting step is a chemical rather than a diffusion process, which supports the formation of the chalcogenide as this limiting step rather than diffusion of cluster species to the substrate. Also, as described in the previous section, metal hydroxide might deposit on the substrate under conditions where it will not form in the solution, but this is likely to be confined to the surface layer. It must be stressed, however, that due to the various possible processes involved in CD, the results of one or even of several studies cannot automatically be extrapolated to all other depositions of the same compound. If reaction is allowed to proceed until the termination stage is reached, the terminal thickness of many CD films is typically several hundred nanometers, although it may reach a micron or more in some cases. This terminal thickness depends to a large extent on the deposition parameters. To take an extreme case, addition of sulphide to a solution of Cd ions will give an immediate precipitate of CdS, but no (or at most an extremely thin) deposit on the walls of the deposition vessel, which may thicken somewhat with time, but will not be visible (which means a terminal thickness less than ca. 20 nm). For a normal CD reaction, if precipitation occurs homogeneously in solution, then that precipitate is lost for film

FIG. 2.4 Typical shape of the curve reprenting time dependence of film thickness during growth.

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deposition, with resulting reduction in terminal thickness. Therefore, the ion-byion process, with its lesser tendency for homogeneous precipitation, will usually result in a larger terminal thickness than the cluster process, for comparable initial reactant concentrations. Of course, film thickness can made as small as desired simply by removing the substrate when the desired thickness has been reached (although very thin films may not be homogeneous but, rather, clusters spread heterogeneously on the substrate). Alternatively, films thicker than the terminal thickness may be obtained by repeated deposition (there will be a limit to the thickness even here, since thick films tend to peel off the substrate).

2.8 SUBSTRATE To a first approximation, films can be deposited by CD on any surface (this is one of the advantages of CD). Of course, there will be certain obvious exceptions, such as substrates that are unstable in the deposition solution (this is rarely a problem, in practice) or “dirty” substrates. In several studies, CdS has been shown to form quite adherent films on Teflon, and this attests to the ability of CD to form films on a wide range of substrates. An important advantage of CD is that the shape of the substrate is usually not important – very irregularly-shaped substrates can be used. To a second approximation, the nature of the substrate is usually important in order to obtain an adherent film; some substrates result in more adherent films than others. Rough substrates are better in this respect (in common with most deposition procedures), probably due to the greater actual surface area of contact per geometric surface area and the possibility of anchoring of the initial deposit in pores of the substrate. Oxides [this includes glass, conducting oxides such as tin oxide and indium tin oxide, and, to a lesser extent, silica (quartz)], in spite of appearing inert, are actually quite reactive in terms of their adsorption properties. This is due to the presence of hydroxyl surface groups, which can form fairly strong hydrogen bonds. Yet there can be very noticeable differences in the adhesion to different glasses and between glass and silica (deposits tend to be less adherent to silica than to glasses in general). An early study on CD PbS films using different glass substrates found large differences in film formation on the different substrates; no (or, at best, only patchy) films were formed on borosilicate glass or on silica, whereas lead flint glasses, followed by zinc crown glass, resulted in the best films [25]. The ability to form good PbS films on these latter glasses was ascribed to the ability of the PbO or ZnO in these glasses to from insoluble sulphides. This would enhance binding to the depositing film. The possibility of ion exchange between metal ions in the glass and those in solution may also play a role in binding the initial CD film. Glass substrates can be sensitized, usually with a solution of SnCl2, which hydrolyzes to give nuclei of tin hydroxide or oxide, on the surface. While in most cases of CD, such sensitization is not used, and not required, there have been re-

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ports of better layers (more adherent and/or homogeneous layers, faster deposition) using such sensitization. For example, it has been shown that, for PbSe deposition from selenosulphate solutions, film formation begins immediately on SnCl2-sensitized glass, instead of after an induction period as on plain glass. Additionally, film formation occurs (at least initially) in the absence of bulk precipitation in solution, in contrast to the parallel film deposition and bulk precipitation that occurs when nonsensitized glass is used [26]. This suggests that a high supersaturation is required for nucleation to occur on untreated glass (as well in solution) and that this is not the case when nuclei are already present on the glass from the sensitization (or from previously deposited PbSe). It can be concluded that such sensitization of glass (and probably many other surfaces) should be considered if satisfactory growth is not obtained without it. Although probably not important for most requirements, it should also be kept in mind that such films will contain a small amount of tin at the film/substrate interface. Metals make good substrates in general, either because chalcogenides tend to adsorb strongly on the noble metals, in particular gold, or the non-noble metals are covered with a (usually hydroxylated in the deposition solution) oxide layer. In addition, if the metal in the deposition solution has a sufficiently negative potential, an internal electrochemical reduction may occur (remember that electrodeposition can often be carried out from CD solutions). This was suggested a long time ago for deposition of various metal sulphides from thiosulphate solutions on certain metal substrates [3,5]. A large variety of CD films have been deposited on different polymer surfaces subjected to various activation treatments [27]. The most effective treatment was immersing the substrate in KMnO4 solution for 24 hr, which formed a brown Mn-O film and subsequent removal of this film with, e.g., conc. HCl. It was suggested that the permanganate introduces carboxylic groups on the originally hydrophobic surface. Films have also been deposited on monolayers—both Langmuir and self-assembled. In many cases, such depositions have studied nucleation of very thin layers of deposit (more accurately, scattered nanocrystals), but deposition of thicker films, such as PbS [28] and Fe(O)OH [29] on self-assembled monolayers and various metal sulphides and selenides on Langmuir–Blodgett films [30] have been studied, in some cases to understand how a well-defined—either chemically and/or geometrically—substrate can control nucleation and growth geometry. Deposition does not occur on some monolayers, or at least considerably less readily than on non-monolayer-covered substrate. An example of such a monolayer is octadecylphosphonic acid; CdS was found not to deposit on this monolayer but did grow on the free areas of mica partially covered with the monolayer [31]. The use of monolayers as substrates has been exploited to pattern deposits. The principle behind this idea is that monolayers with either hydrophobic or hydrophilic end groups can be patterned onto a substrate. Deposition will usually

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occur only (or at least highly preferentially) on the hydrophilic endgroups. Whether this is due to a simple physical interaction between solution (aqueous) species and a hydrophilic surface, due to electrical charge on the solvated endgroup, or to some other specific interaction (or some combination of these effects) is not clear. An example of this patterning is shown here for TiO2 deposition [32]. A long-chain thioacetate-terminated trichlorosilane is self-assembled on an oxidized Si substrate (Fig. 2.5a and b). The thioacetate is somewhat hydrophobic. Exposure of this monolayer through a mask (a fine mesh grid) to UV radiation

FIG. 2.5 Self-assembly of a long-chain thioacetate-terminated trichlorosilane on an oxidized Si substrate (a and b). Exposure of this monolayer through a fine mesh grid mask to UV radiation (c) oxidizes the somewhat hydrophobic-thioacetate to hydrophilic sulphonate endgroups (d). Deposition of TiO2 from an aqueous solution of TiCl4/HCl at 80°C on the UV-exposed (hydrophilic) regions of the substrate (e). (From Ref. 32.)

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FIG. 2.6 SEM micrograph of patterned TiO2 deposit (left side) and Ti elemental mapping of this sample (right side). (From Ref. 32.)

oxidizes the thioacetate to hydrophilic sulphonate groups (c and d). Deposition of TiO2 from an aqueous solution of TiCl4/HCl at 80°C occurred only on the UV-exposed (hydrophilic) regions of the substrate (Fig. 2.6). See Section 4.1.5.3 and Section 5.2.4.2) for other examples of patterned deposition (of CdS and PbS, respectively) using different patterning procedures. For any particular substrate, the adhesion can also depend on the deposition parameters, although little is reported on this aspect (for a good reason—it is not really understood). The surface of glass (of whatever type) is hydroxylated in aqueous solution and the concentration of the various species is clearly pH dependent, as can be seen from Eq. (2.16): J M JSiMOMH

D

J   M JSiMO  H

(2.16)

[Since silica is very acidic, the dissociation to give a positively charged surface (MSiMOH 2 ) will only occur in very acidic solution (pH  2) and is not commonly encountered in CD. Other less acidic oxide surfaces may, however, be positively charged in CD solutions.] This dissociation can affect interaction between glass and various species in the solution as a function of the solution pH. Note that water (also ammonia) may interact, through hydrogen bonding, with the hydroxylated silica. In the end, adhesion as a function of deposition parameters (in contrast to the nature of the substrate) can usually not be reliably predicted, and there is no substitute for experimental experience. Apart from adhesion, the crystallographic properties of the CD film are sometimes dependent on the nature of the substrate (although more often there does not seem to be any dependence of this type). One example is epitaxial deposition on a crystallographically ordered substrate [epitaxial here means a struc-

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tural relationship between the crystal lattice of the deposit and of the substrate— i.e., the crystal axes of the deposit are aligned (but not necessarily parallel to) the crystal axes of the substrate]. Examples of this will be discussed later and in other chapters. Also, even if the deposition is not epitaxial (and only a few cases have been reported where epitaxy occurs), different texturing of the film may occur. Texturing refers to the preferred orientation of the crystals of the deposit perpendicular to the substrate; thus, textured films of hexagonal CdS often have the basal face (the Cd or S face) pointing upward. If orientation occurs also in the plane of the substrate (azimuthal orientation), then the deposit is called oriented. The term orientation is often (erroneously) used where texturing is meant. A monolayer-covered substrate may dictate the crystallographic form of a CD film, depending on the interaction between the monolayer endgroups and the semiconductor. Thus, epitaxial growth of PbS has been accomplished at arachidic acid monolayers where the PbS interatomic spacings along the (111) plane are well matched with the monolayer packing [33]. Texturing has been observed in some cases, even when the monolayers themselves were not well ordered. It will be obvious that the cluster mechanism of deposition is unlikely to lead to an oriented film, since the clusters would have to align themselves with the substrate lattice, either on adsorption or subsequently. Therefore an epitaxial film is highly suggestive of an ion-by-ion growth, which is more likely to be directed by the substrate. One example has been described of CD (PbS) on a poled ferroelectric substrate. The PbS crystal size was larger (ca. 1 m) on the poled substrate than on the unpoled (or a glass) substrate (ca. 0.3 m) [34]. Other changes in the electrical properties of the films were noted. The differences were ascribed to the electric field and charge accumulation at the ferroelectric surface (more details can be found in Sec. 5.2.4.3).

2.9 DEPOSITION OF SPECIFIC SEMICONDUCTORS This section will give an overview of the various semiconductors that have been deposited by CD. The groups of semiconductors will be divided in much the same way as in the rest of the book: II–VI IV–VI Other binary sulphides and selenides Oxides, other binary semiconductors, and elemental Se Ternary compounds A few specific examples of experimental details will be given in this section, with explanation of the importance of the different variables.

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2.9.1 II–VI Compounds This comprises the most studied group and includes CdS (the single most studied compound), CdSe, CdTe, ZnS, ZnSe, HgS, HgSe, and various mixtures of these compounds.

2.9.1.1

CdS

Film Deposition. We begin with a common “recipe” for CdS deposition and follow the deposition process using this recipe. This is described in detail as a way of understanding both the deposition process and which factors can affect the final film. STEP 1: PREPARATION OF STOCK SOLUTIONS. Usually, stock solutions of the solid reactants are prepared in advance, since many depositions are normally carried out. The three reagents required are:

CdSO4: Other water-soluble Cd salts can be used equally well, such as the chloride or acetate; there is no clear evidence in general that the nature of the anion is important, although there are a few studies that reported differences in the deposit depending on the anion used. NH4OH: Since this is required in high concentrations and is a liquid, a stock solution is not required, but it can be used directly from a bottle of concentrated ammonia (concentrated ammonia is ca. 15 M). Thiourea (SC(NH2)2): This solution will slowly precipitate sulphur (seen as a fine white precipitate—white instead of the usual yellow of sulphur possibly due to size quantization of the finely divided precipitate?). However, it can usually be kept over a period of weeks or even months in a stoppered bottle without major adverse effects, although for optimum reproducibility a fresh solution may be preferred. If a prepared stock solution is used, filtration of this solution before use will minimize the presence of sulphur particles in the deposition solution; such particles can act as nucleation centers and accelerate precipitation in the solution. STEP 2: PREPARATION OF THE SUBSTRATE(S). While almost any substrate can be used, we will use glass microscope slides in this example; this is a common substrate and makes it easy to see the CdS film. The microscope slide can be cut to whatever size and shape is convenient. The slide should be cleaned well, since films usually do not adhere well to “dirty” surfaces. Suitable cleaning agents are trichloroethylene or/and sulphochromic acid, and the slide should be well rinsed with pure water. If the slide is clean, water dropped onto it will form a film (hydrophilic surface), while on a “dirty” (hydrophobic) slide the water will form drops. Needless to say, the part of the slide where deposition is to occur should not be touched with the hands after this treatment.

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STEP 3: PREPARATION OF THE DEPOSITION SOLUTION. Solution compositions, even based on exactly the same constituents as stated earlier, vary widely from one literature source to another. We give a typical “average” composition and discuss the effect of variations from this composition. This “average” solution is made up of (concentrations in final solution):

CdSO4 10 mM Aqueous ammonia 1M Thiourea 50 mM Solution pH 11 Deposition temperature 70°C Concentrated ammonium hydroxide is added to a stock solution of CdSO4 (or other Cd salt). Initially, Cd(OH)2 precipitates, but this redissolves in excess ammonia to give the cadmium ammine complex: Cd2  4NH3 D Cd(NH3) 2 4

(2.17)

This solution is heated to 70°C in a thermostated water bath. A stock solution of thiourea is then added to bring the thiourea concentration up to 50 mM. The final solution pH can be adjusted with KOH (more basic) or acetic acid (more acidic). STEP 4: DEPOSITION. The substrate(s) is immersed in the preceding solution and placed in the water bath at 70°C. (Note that CdS usually also forms homogeneously in solution, and this can sediment onto the substrate, where it forms a loose coating and may prevent growth of an adherent film. For this reason, the substrate should be placed vertically in solution; or if placed at an angle or even horizontally, only the underside of the substrate, where sedimentation does not occur, should be used.) The time of deposition is variable (although it should be reproducible from one run to another under the same conditions), and it is difficult to give a particular time. Usually a thin film (tens of nanometers) will form in some minutes, and this will slowly thicken over some tens of minutes to hours to a typical terminal thickness of ca. 200 nm. The simplest way of determining the optimum time (and “optimum” will depend on the application) is to simply look at the film. With a little experience (or with the more quantitative help of a spectrophotometer), the approximate thickness can be estimated from how deep the yellow color is. Thickness can also be estimated from the transmission spectrum by measuring the drop in transmission over the near-bandgap region; this method is less dependent on reflection and scattering losses (as long as they are not too large) than simply measuring absorption at a fixed wavelength. In any case, prior calibration of the spectra with known thicknesses of the film is required. When the desired thickness is attained, the substrate is removed and rinsed with water. This rinsing can also be carried out in an ultrasonic bath, which more effectively removes loose particles.

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VARIATIONS IN THE PROCESS. There are many variables in this (experimentally simple) process; concentration of the various reactants, pH, and temperature are the main ones. Other, less important (in most cases) ones are stirring of solution and illumination of the solution during deposition. We will ignore these last two here. The effect of reactant concentration can be divided into two separate influences. The simplest is obvious: Lower overall concentrations result in a slower rate. This does not necessarily mean a thinner film, however—sometimes the opposite. The reason for this is clear if we return to our introductory discussion on the CD process—rapid precipitation. It is clear that if the reaction is too fast, it will terminate with most of the product precipitating homogeneously in solution rather than depositing on the substrate (which requires time to occur). This results in a very thin film, if any film at all. Similarly, for the less extreme case of a CD reaction that terminates, not within a second, but still in a short time, the final film thickness will be small. At the other extreme, if the reaction is extremely slow, a thick film can be built up, but it may take a very long time for this to occur (weeks, even months). It is therefore evident that there is an optimum rate for the reaction, which can be controlled by a combination of reactant concentrations, temperature, and pH. A separate effect of concentration is the ratio between the metal ion and the complexant concentrations. This ratio determines, often more than the overall reactant concentrations, the reaction rate, since it controls the concentration of free metal ions in solution. It can also determine the reaction pathway. Further discussion of this factor will be left for the next example (CdSe), since it has been treated in more detail for that case. The solution pH influences a number of factors, and it is not always simple to predict its effect. Thus, thiourea decomposition (in alkaline solution) is generally faster at higher pH. The probability of the presence of a solid phase of Cd(OH)2 and its concentration in the solution are both increased at higher pH (higher OH concentration). The pH is determined, in the example given earlier, by the concentration of ammonia. However, it can be adjusted independent of the ammonia concentration. Addition of an ammonium salt, which with ammonia, acts as a buffer, will lower the pH through the following equilibrium:  NH3  H2O D NH 4  OH

(2.18)

Ammonia is alkaline in water because of this equilibrium, which produces hydroxide ions. The addition of an ammonium salt (NH 4 ions) will push the equilibrium back toward the left, i.e., lower hydroxide concentration, therefore lower pH. Increase of pH can be effected by addition of sodium or potassium hydroxide. This explains why some CdS depositions based on the Cd/NH3/thiourea formula include either an ammonium salt or an alkali metal hydroxide. Finally, the effect of temperature on increasing the reaction rate (possibly also the mechanism—see next example) is again obvious, since the thiourea de-

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composition will be faster at higher temperature. Additionally, since the stability constant of the complex is usually smaller at higher temperatures (see Chap. 1), there will be a higher free metal ion concentration, which again translates into faster rate (although, since sulphide generation is usually rate determining, this effect may not be large). Temperature programming during the deposition has also been employed. This can be useful, for example, if the deposition solution changes the substrate in some beneficial way (as occurs in some photovoltaic cell substrates—see Chap. 9). The initial low temperature delays coverage of the substrate and allows more time for this surface treatment to occur. If the temperature is then increased, deposition occurs at a reasonable rate. There are other bath compositions based on different sulphide-generating precursors and/or complexing agents. Thioacetamide and thiosulphate are two of the former, while ethylenediamine is a common example of a complexant that has been used instead of ammonia. The volatility of ammonia, and its gradual loss in an open deposition bath, is circumvented by using a less volatile complexant, such as ethylenediamine. Some Properties of CdS Films. It should be noted that all properties discussed in this chapter refer to as-deposited films (not annealed) unless specifically stated otherwise. In general, annealing increases crystal size and reduces dark conductivity. The latter obviously depends to a large degree on the annealing atmosphere. STRUCTURE. CdS can exist in three different crystal structures: hexagonal (wurtzite), cubic (zincblende)—both tetrahedrally coordinated and cubic (rocksalt), which is sixfold coordinated. Except in a few cases, the rocksalt modification of CdS has been observed only at very high pressures: CD films of this phase have never been reported. The other two phases have been reported to occur in CD films under various conditions. The wurtzite phase is thermodynamically slightly more stable, and invariably forms if the zincblende phase is heated above 300–400°C. The low-temperature CD method therefore can allow the formation of the zincblende phase, and this phase is commonly obtained in CD CdS films. Very often, a mixture of wurtzite and zincblende phases has been reported in the literature. There are many variables that affect the crystal structure, including the nature of the complex, the substrate, and sometimes even stirring. ORIENTATION. In common with with other deposition techniques, a preferential texturing of the film in the (111) (for zincblende) or (0002) (for wurtzite) direction is often reported. However, nontextured (or weakly textured) films are probably more commonly obtained. As for crystal structure, the degree of texturing depends on several factors, an important one being the nature of the substrate. It is likely that highly textured deposits form by the ion-by-ion mechanism; a pure solid-phase cluster mechanism is less likely to result in strong texturing.

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In certain cases, epitaxial growth has been observed on crystallographically well-defined substrates, such as single crystal InP and CuInSe2. This type of growth, even more than textured growth, demands ion-by-ion growth as well as a crystallographically well-defined substrate. Lattice parameters (at least at some defined ratio and angular match) fairly close to that of CdS are also required for a relatively defect-free interface, although there is some flexibility here. For example, CdS with a high density of stacking faults has been epitaxially deposited on GaP (7% mismatch). OPTICAL PROPERTIES. The most commonly reported optical properties are optical transmission, with some studies also on photoluminescence. The importance of the optical transmission for CdS in particular lies in its use in photovoltaic cells, where it acts as a window layer. The CdS should be as transparent as possible to the incoming radiation. The transmission is a function of thickness, bandgap, and film structure (is the film transparent or scattering?). The bandgap in most studies is constant (ca. 2.45 eV at room temperature), although somewhat larger values have been obtained due to size quantization in very small crystallites. Photoluminescence of the films varies greatly, both in intensity and in spectral shape, from one report to another. This is not surprising, since this property is very dependent on the state of the surface of the individual crystals. A red (ca. 1.8 eV) defect emission is usually seen, but green, yellow, and infrared peaks have also been reported. The various wavelengths are related to different defects in/on the crystals; even the green emission is probably due to a shallow defect emission. ELECTRICAL PROPERTIES AND PHOTOCONDUCTIVITY. Electrical resistivity of CD CdS films is commonly studied. Values for this (dark) resistivity vary over many orders of magnitude from one film to another, usually for reasons that are not understood. Values as high as 109 -cm and as low as 15 -cm have been reported for undoped films (doped films have been reported with still lower resistivities). Since the films are often highly resistive, it is not surprising that they exhibit strong photoconduction. Photoconduction occurs due to the formation of free carriers by illumination, and if the free carrier concentration is low to begin with (low conductivity), then the photogenerated carriers will usually dominate the conductivity. This is in contrast to a relatively conducting semiconductor (high doping level), when the extra photogenerated (majority) carriers are only a small perturbation to those present in the dark. Light:dark photoconductivity ratios (sensitivity) as high as 109 have been reported. In many cases, the photocurrent decay time is measured in hours, and this is explained (in a general way) by slow states; the nature of these states is not usually known.

2.9.1.2

CdSe

Film Deposition. Chemical deposition films of CdSe have been relatively widely investigated, largely for photoanodes in photoelectrochemical cells (see

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Sec. 2.12.1.2). Here we will give an example of CdSe deposition, emphasizing the differences between the procedures for CdSe and CdS. The most important single difference in the deposition is, of course, the chalcogenide-forming reagent. A procedure used commonly by the author will be described, where the complexant is nitrilotriacetic acid (NTA), which is related to the more common EDTA (ethylenediaminetetracetic acid) and has the chemical formula N(CH2COOH)3. It is a strong complexant for Cd2 (and many other cations), although less so than EDTA, but stronger than the often-used ammonia. It also has an advantage (in most cases) over ammonia in that it is nonvolatile and is not lost during deposition. STEP 1: PREPARATION OF STOCK SOLUTIONS. The aqueous Cd solution is prepared as for CdS. As for CdS, the nature of the anion does not appear to be very important—the sulphate is commonly used.

NTA. NTA itself is not very soluble in water. Either it is used as the Na or K salt, or the acid can be dissolved together with three equivalents of KOH or NaOH to form a solution of the salt. This salt will be referred to hereafter simply as NTA. Sodium selenosulphate (Na2SeSO3). This compound is made by dissolving elemental Se in an aqueous solution of sodium sulphite: Na2SO3  Se D Na2SeSO3

(2.19)

The prepared solution will slowly deposit Se as a black precipitate over a period of weeks. Also, the solution is considerably less stable than thiourea. If kept out of excessive contact with air, it will be usable for about a month if high reproducibility of the deposition kinetics is not important. However, it is important to be aware of the slow decomposition and loss of reactivity of this reactant. If a freshly made Na2SeSO3 solution is used to deposit CdSe, the reaction will proceed much faster than if an aged solution is used; this fact should be taken into account in preparing the overall deposition solution. The solution is made up from an aqueous solution 0.4 M in Na2SO3 and 0.2 M in Se. This solution typically is stirred at 60°C for a couple of hours (the Se dissolves slowly, the reason a fresh solution is usually not made up every time). If this solution is prepared over a longer time (maybe 6 hours and/or at a higher temperature, even boiling), it will undergo accelerated aging. This results in more reproducible deposition conditions on one hand, but at the cost of reduced reactivity and a shorter lifetime of the solution. STEP

2: PREPARATION OF THE SUBSTRATE(S).

The same as described earlier

for CdS. STEP 3: PREPARATION OF THE DEPOSITION SOLUTION. As for CdS, the conditions can be very variable. The following is a typical deposition solution, but re-

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member that allowances should be made for the varying activity of the selenosulphate. If deposition occurs too rapidly (slowly), parameters should be changed to slow down (speed up) the reaction [e.g., lower (higher) selenosulphate concentration, higher (lower) NTA:Cd ratio, lower (higher) temperature]. The following procedure is followed to obtain a final concentration of 60 mM each Cd2 and Na2SeSO3 and ca. 100 mM NTA (the reason for the “ca.” will be explained shortly). Take the calculated amount of stock Cd solution and add water. Then add the stock NTA solution. If the concentrated Cd and NTA solutions are mixed together without adding water between them, they may form a gel and the solution will then be useless. The pH should then be adjusted to ca. 8 by adding KOH or NaOH. It is important that the pH at this stage be greater than 7, since selenosulphate immediately decomposes to red selenium at a pH lower than 7. This solution is brought to the temperature at which the deposition is to be carried out. The selenosulphate (also at the same temperature) is then added. This will increase the pH (since the selenosulphate is alkaline). KOH is added as necessary to bring the pH to ca. 10 (9.5–10 is a suitable range; it should be neither much lower nor much higher than this). If the pH increases to much above 10 on the addition of selenosulphate, dilute acetic acid can be carefully added to lower the pH to the required range. STEP 4: DEPOSITION. This is the same as for CdS, except that the deposition usually is slower and takes from hours to days, depending on conditions. Orange/red coloration in the solution, corresponding to the start of CdSe formation, should occur some minutes after adding the selenosulphate. If it occurs immediately or almost immediately, the reaction may be too fast and only a thin terminal thickness may be obtained. On the other hand, if coloration has still not started after about 30 min, the deposition will probably be very slow. VARIATIONS IN THE PROCESS. The reader may have noticed that some of the preparation details just given are rather vague. This reflects the varying activity of the selenosulphate solution, both with time and from one batch to another. It is more important to understand the effects of the various parameters and to be able to vary them logically than to follow an exact recipe. These effects will be described in detail in Chapter 3, and it is recommended that anyone wanting to deposit these films read that chapter before carrying out the deposition. Here, a brief explanation of these various factors is given. The deposition temperature obviously will increase the reaction rate (and the deposition rate, although, again, if the reaction is too rapid, only a thin film will be obtained). Another effect of increased temperature is increased crystal size in the film. The crystal size varies (assuming the NTA:Cd ratio is not too high), typically from 4 nm at close to 0°C to ca. 8 nm at 80°C. This change can be seen by a change in color from yellow-orange (low temperature) to red (high temperature) due to size quantization (see Sec. 2.12.2 and Chapter 10).

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If crystal size is not an important issue, then higher temperatures are usually more convenient, since the deposition is faster. Adhesion of the film to the substrate also appears to be better in general at higher temperatures (larger crystal size? See table 1.3). Solution composition is important not only because the reaction rate increases with the concentrations of selenosulphate and/or Cd, but even more so through the ratio between the NTA and Cd concentrations. The higher this ratio, the slower the reaction, since the free Cd2 concentration is less. The optimum value for this ratio will depend, among other factors, on the deposition temperature. At higher temperatures, a higher NTA:Cd ratio is required to prevent too rapid reaction. Thus the 100 mM concentration of NTA (for 60 mM Cd) given earlier is typical for ca. room temperature deposition, but would probably need to be increased to 110 mM or even more at substantially higher temperatures. Even more important, if the NTA:Cd ratio increases above a certain value (which depends on the other solution parameters but which is typically between 1.7 and 2.1), the mechanism of deposition changes from a hydroxide cluster mechanism to ion-by-ion deposition. The latter is considerably slower, does not deposit homogeneously in solution (or much less than the cluster mechanism), and results in larger crystal size (typically 8–20 nm, depending on deposition temperature). Also, in contrast to the cluster mechanism, where deposition is (should be) homogeneous and the films transparent throughout the deposition, the ion-by-ion mechanism usually results in films that appear nonhomogeneous and highly scattering in the early stages of deposition, but become homogeneous and transparent with time (typically several days at lower temperatures, less at higher ones). If all the foregoing gives the reader the impression that CdSe deposition is more of an art than a science, this would not be a gross misunderstanding. Experience is certainly useful here, more so than for CdS deposition, which is more reproducible, probably due to the more stable thiourea (and possibly also because of the shorter deposition time typical for the CdS deposition, based on one of Murphy’s laws—the more time you allow for something to occur, the greater the opportunity that something will go wrong). Some Properties of CdSe Films. STRUCTURE. CdSe forms the same three crystal structures as described earlier for CdS. The main difference between the CD films of the two materials is that, while CdS can be commonly found in both the wurtzite and sphalerite forms, CdSe is more commonly deposited in the cubic zincblende form. Mixtures of the two forms have been reported in some cases, particularly when a visible Cd(OH)2 precipitate is present in the initial deposition solution. As for CdS, CdSe has been epitaxially deposited on single crystal InP. As expected, epitaxy occurred only for the ion-by-ion mechanism, where individual species could either adsorb on or migrate to the ideal lattice position.

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The CdSe crystal size in the hydroxide mechanism is probably determined by the size of the CdOH)2 particles in the solution and on the substrate, while that formed through the ion-by-ion mechanism will depend on the heterogeneous nucleation on the substrate, and is invariably larger. OPTICAL PROPERTIES. CD CdSe films often exhibit size quantization, with a blue shift in their absorption spectrum of as much as 0.6 eV. Photoluminescence of the films (in the absence of water vapor) usually is dominated, as for CdS, by a broad defect emission varying from ca. 1.4 to 1.6 eV. A (close to) band-to-band emission is often also observed, usually (but not always) at a lower intensity than the broad defect emission. In the presence of water vapor, however, the band-to-band emission often dominates. ELECTRICAL PROPERTIES. Resistivity studies on CdSe are much less widespread than on CdS films. The dark conductivity of undoped films is high (108 -cm is typical), and the photocurrent sensitivity is less than for CdS films (even under illumination, the films are normally very resistive).

2.9.1.3

CdTe

Only two different studies on true CdTe deposition by CD appear to exist. In the first, CdTe was deposited from solutions containing TeO2 and hydrazine, the latter presumably slowly reducing the TeO2 to telluride ion. In this study, the main interest was not on the deposition itself but on the further use of the CdTe films, and not much information was given on the films themselves. More recently, CdTe was deposited using what was apparently the Te analogue of selenosulphate, Na2TeSO3. In all these depositions, the films were not very stoichiometric and included considerable amounts of Te. Some structural, optical, and electrical properties of these latter films were given (described in Chap. 4). In a very recent study, CD Cd(OH)2 films were converted to CdTe using a solution of Te in hydroxymethane sulphinic acid, which acted as a telluride source. The conversion of the hydroxide to CdTe was incomplete, but there did not seem to be any free Te in the films. Properties of these films were also described.

2.9.1.4

ZnS and ZnSe

The most important difference between CD of CdS(Se) and ZnS(Se) is related to the difference between the solubility products of the hydroxide and chalcogenides of the two metals. Considering the sulphides, the various values of Ksp are: Cd(OH)2—2  1014; CdS—1028; Zn(OH)2—1.1016; ZnS—3.1025. The differences between the pairs of Ksp are 2  1014 (for Cd) and 3  108 (for Zn). Since hydroxide ions are present at much higher concentrations than sulphide, hydroxide formation, and stability against sulphurization, is much more likely for Zn than for Cd.

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ZnS is commonly deposited by CD. However, in many cases, the ZnS is not stoichiometric and contains hydroxide in one form or another. ZnSe, on the other hand, is more difficult to deposit from selenosulphate baths, in spite of the somewhat lower solubility product of ZnSe (4.1026) compared with ZnS. In fact, a strong reducing agent such as hydrazine is normally required. The hydrazine presumably reduces the selenosulphate to give a high enough selenide concentration to allow ZnSe to form, although it is possible that other factors are also important. (The sulphite, present in excess in selenosulphate, is itself a reducing agent, although much weaker than hydrazine.) Activation energies for ZnS(Se) depositions are generally considerably lower than for CdS and PbS deposition, and this suggests a different deposition mechanism (although, as we have already seen, the deposition mechanism is largely a function of the deposition conditions and not simply the material itself). Also, the crystal size of ZnS(Se) is usually smaller than for CdS or PbS. Both these factors suggest that ZnS and ZnSe form by a pure cluster mechanism. In fact, nowhere in the literature is there evidence for ion-by-ion growth of ZnS or ZnSe.

2.9.1.5

HgS and HgSe

There are only several reports on the deposition of the mercury chalcogenides (some ternaries containing Hg have also been described). HgS films have been deposited from a simple chemical precipitation reaction between mercuric chloride and sodium sulphide. Under suitable conditions, a film is formed along with the precipitate. The thickness of the films were ca. 0.7 m, which is thicker than normal for CD films, even more so considering the rapid nature of the reaction, which normally only leads to a very thin film, if any at all. The films apparently deposit from the colloidal HgS formed on mixing the reactants. HgS films have also been deposited from “more conventional” deposition solutions using thiourea and the tetraiodide complex of mercury—a strong complex—in alkaline solution. Both these and the previous films showed an optical absorption with a gradual absorption onset at 700 nm and a sharp one at 400 nm. The thiosulphate mercury(II) complex in ammoniacal solution has also been used to deposit films of HgS, both in ammoniacal and nonammoniacal solutions. The deposit from the former was predominantly -HgS (cinnabar). As with the previous films, a sharp optical absorption onset at ca. 400 nm was observed, together with a more gradual one extending, in this case, to beyond 800 nm and dependent on film thickness. The nonammoniacal solution gave crystal sizes (and optical bandgaps) that varied with deposition temperature from 3 nm (2.4 eV) at 0°C to 8 nm (1.9 eV) at 85°C and corresponding resistivities between 104 and 103 -cm. HgSe was first deposited from an iodide-complexed solution using selenosemicarbazide. Two other depositions were described, both using selenosulphate. One used an alkaline Hg–formamide complex. The as-deposited films did not show an XRD pattern, suggesting that the material was either amorphous or

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very small nanocrystalline. The latter in particular is supported by the large bandgap measured from the absorption spectrum (1.42 eV compared to semimetal bulk HgSe with a negative bandgap). Electrical conductivity measurements indicated a midgap Fermi-level characteristic of intrinsic semiconducting material. In the other selenosulphate method, HgSe films were deposited from an ammoniacal bath onto polyester substrates. The films were strongly (111) textured. The crystal size was 7.7 nm, resulting in a strong blue shift in the optical absorption spectrum, and a measured bandgap of 2.5 eV (compared to a negative—semimetal— bandgap of bulk HgSe). The sheet resistance of these films (13 k-cm2) was relatively low, considering the small crystal size.

2.9.2 IV–VI Compounds 2.9.2.1

Deposition

As noted at the beginning of this chapter, most of the early studies in CD focused on PbS, followed by PbSe, driven by their photoconducting properties. For optimum use as photoconductors, the deposited films were annealed in an oxygencontaining atmosphere. Most of this section will focus on nonannealed films, and annealed films will be treated only very briefly. More details on the annealed as well as as-deposited films will be given in Chapter 5. Many different bath compositions (different refers to bath constituents rather than simply to different concentrations) have been used to deposit PbS and PbSe, more so than the II–VI materials. Both alkaline and acid baths have been described, although the former are much more common. The chalcogen sources are similar to those used for the II–VI compounds, including thiourea, most commonly used for PbS and selenourea (originally the source of choice) and, more commonly nowadays, selenosulphate for PbSe. The variety of complexing agents used is large; various carboxylic acids (most commonly citrate), triethanolamine, nitrilotriacetate, hydroxide, and even selenosulphate itself have all been used. Addition of thiosulphate to a citrate/selenosulphate bath for PbSe resulted in a deposit that did not show an XRD pattern and was assumed to be amorphous; blue shifts in the optical spectra could be explained either by structural changes or, possibly, by size quantization in crystals too small to be seen in the XRD spectrum. PbS has been deposited from an acidic bath using thiosulphate. Also, very thin films of PbS (maximum absorbance 0.015) have been grown on quartz immersed overnight in a solution of Pb ions together with polyvinyl alcohol (ostensibly to protect against aggregation of colloids) through which H2S has been passed.

2.9.2.2

Film Structure and Morphology

PbS and PbSe are almost always found in the rocksalt (RS) crystal form. All structural investigations on CD films have shown this form, with one exception; PbSe deposited from hydroxide complex at high hydroxide concentrations and at rela-

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tively high temperatures have produced a novel rhombohedral modification with an external hexagonal morphology and crystals typically 1–3 m in size. For PbSe, crystal size varies considerably, depending on deposition parameters, more so than for the II–VI compounds. Crystals as small as 4 nm and larger than a micron can be grown. The main factors, as for the II–VI materials, are temperature and complex:Pb ratio. The use of a hydroxide complex gives the widest range of crystal sizes. PbSe often exhibits a bimodal distribution of crystal sizes. PbS, in contrast, tends to form only relatively large crystals. Chemical deposition films of PbS and PbSe are generally not strongly textured. One report has described (200) textured PbS films on glass if H2O2 is present in the deposition solution.

2.9.2.3

Electrical Properties

The IV–VI films are usually p-type, both as deposited and after annealing in air. One study, where PbS was deposited from a bath containing hydrazine, found the deposit on glass to be n-type temporarily but converted to p-type on air exposure. By depositing the PbS on a trivalent metal coating (such as Al), the n-type conductivity could be stabilized for a longer time. Electrical resistivity of the films, both PbS and PbSe, has often been reported to be of the order of 105 -cm as deposited, with a reduction of about an order of magnitude after annealing in air. However, this can vary considerably from one type of deposition to another. Resistivities greater than 109 -cm have been reported in some cases, which invariably drop to the k-cm range after air annealing. These high-resistivity films are probably those with a very small crystal size (small meaning ca. 10 nm or less). There are many studies on photoconductivity in these films, many of them early ones and focused on annealed films (since air annealing is necessary for optimal photoconductivity). The use of a chemical oxidant (which never seems to be specified) gives much higher photosensitivity for as-deposited films than for films deposited without oxidant, although even here annealing is used to obtain maximum performance. Some studies on photoconductivity in as-deposited PbSe films have shown shifts in photoconductivity spectral response, with onsets shifted to 2.2 m instead of the ca. 4.5 m more typical of annealed films. As with optical absorption studies, these shifts can be attributed to size quantization.

2.9.3 Other Sulphides and Selenides A large range of other metal sulphides and selenides have been deposited by CD. Since these will be individually described in Chapter 6, it will be sufficient here to list all binary sulphides and selenides (along with oxides) in Table 2.1, along with up to three references to each compound.

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TABLE 2.1 Binary Semiconductors Deposited by CD Sulphides

Selenides

Ag2S [35–37] As2S3 [42, 43] Bi2S3 [44–46] CdS [50–52] CoS [58] CuxS [62–64]

HgS [71–73] In2S3 [76–78] MnS [81–83] Mo-S [86,87] NiS [89] PbS [92–94] Sb2S3 [43, 99, 100] SnS [104–106] (also Sn2S3) SnS2 [42, 106, 108] TlS [114] ZnS [117–119]

Ag2Se [38, 39] Bi2Se3 [47–49] CdSe [20, 53, 54] CoSe [59] CuxSe [65–67]

Oxides Ag2O/AgO [40, 41]

CdO [55–57] CoO [60, 61] Cu2O [68] FeO(OH) [29] Fe2O3 [69] Fe3O4 [70]

HgSe [38, 74, 75] In2O3 [79, 80] MnO2 [84, 85] Mo-Se [86-88] NiSe [89] PbSe [53, 95, 96] Sb2Se3 [101–103] SnSe [107]

TlSe [115,116] ZnSe [120–122]

NiO [90, 91] PbO2 [97, 98] Sb2O3 [102] SnO2 [109–111] TiO2 [32, 112, 113] Tl2O3 [84, 98] ZnO [123-125]

Only three references to pure CD deposition of tellurides, all of them for CdTe, have been found [126–128] and therefore no column for tellurides is given here. SiO2, Y2O3, and ZrO2, not given in the table, have also been deposited, as have Ag halides.

2.9.4 Oxides, halides, and elemental Se 2.9.4.1

Oxides

While most studies on CD have been on sulphides and selenides, considerable work has also been carried out on oxide films. The films are most often formed by reaction of hydroxide ions with a metal salt. While it might be expected that the product is a hydroxide rather than an oxide, in many reported cases oxides are directly formed. This is probably due to two factors: Many of the metal ions used (e.g., Pb, Sn, Tl,) do not readily, if at all, form simple hydroxides; others (Ag, Cu, Mn) are very readily oxidized even in aqueous solutions. Ni(II) hydroxide is fairly readily dehydrated, particularly in the presence of the persulphate ion used to deposit the oxides in some cases. Zn(OH)2, Cd(OH)2 and In(OH)3 are reasonably stable; Zn(OH)2 can be easily dehydrated while the other two require annealing to form the oxide.

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Persulphates (also called peroxydisulphates) (S2O 2 8 ) are very strong oxidizing agents and have been used to deposit oxides, in particular those with oxidation states higher than those of the original cation. PbO2, MnO2, and Tl2O3 deposited from solutions of Pb2, Mn2, and Tl are examples of this type of deposition. Sometimes a small concentration of Ag ions is needed in the deposition solution as a catalyst. Ag is a known catalyst for oxidations 2 using S2O 2 8 (S2O 8 oxidizes it to Ag(III), and this is then the active oxidizing agent). Fe-oxides, ZnO and In2O3 have been deposited using dimethylamineborane (DMAB) or trimethylamineborane and Zn or In nitrate. The nitrate anion is important and is believed to be reduced by the DMAB to nitrite and hydroxide. It is not clear why ZnO, and not the hydroxide, is the final product. In the case of In, the deposited films were In(OH)3 and required annealing at ca. 200°C to form the oxide. Homogeneous precipitation using urea (which hydrolyzes to give an alkaline solution) has been used extensively, and in a few cases films of basic salts (sulphates of Al and Sn(IV) and formate of Fe) have been obtained. These are not considered semiconductors in the conventional sense, but do provide examples for extension of the CD method beyond the conventional sulphide-selenide-oxide compounds. Many (hydr)oxides have been deposited by slowing down natural hydrolysis, usually by complexation (e.g., AgO from an alkaline triethanolamine-complexed silver bath). Highly-acidic cations will readily hydrolyse even under acidic conditions. Fluoro-complexes of some of these (e.g., Ti, Si) can be controllably hydrolyzed by addition of boric acid, which reacts with the F, thereby destroying the complex and allowing hydrolysis. A (not-closed) heated solution of ammonia gradually loses ammonia. If a cation is complexed with ammonia, the free-cation concentration will gradually increase as ammonia is lost (a rare example of slow cation release rather than anion release). It will also increase with an increase in temperature, due to the decrease in stability of complexes with increasing temperature. As one example using this principle, thin films of mixed ZnO/Zn(OH)2, which converted to ZnO on heating over 200°C, were deposited from a heated aqueous Zn-ion/ammonia bath. Apart from these methods, there are others that are relevant to this section. Aqueous solutions of permanganate will slowly decompose, forming a brown film [MnO2 or possibly MnO(OH)] on the walls of the vessel in which they are stored. Increase of either acidity or alkalinity of the solution can accelerate this decomposition reaction. As another example of film formation due to slow solution decomposition, the author possesses a glass bottle with a green, highly tenacious film of (presumably) Cr2O3 resulting from years of storage of some (unfortunately unknown) Cr-containing solution.

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TABLE 2.2 Electrical and Optical Properties of CD Oxidesa Resistivity (-cm)

Material

1

Cu2O Fe3O4 Fe2O3

10 2  103 2

In2O3

109 33

In2O3 MnO2

2.102

NiO NiO PbO2 SnO2

105 3  102 103 101

Annealed at 350C As deposited Annealed at 200C

Conductivity type

Absorption onset (nm)

Reference

p

ca. 600 Black ca. 570

68 129 130

n

ca. 400

80

79 84

pH
6.3 pH
8

2  103 SnO2

103

(5% Sb)

104 109 4  104 2  104 to 4  102

TiO2 Tl2O3 ZnO

ZnO ZnO:Al

Special conditions


102 2  104 to 2  102

p

Annealed at 250C Annealed at 400C Annealed at 250C Annealed at 400C

Depending on boron content

n

ca. 700
430 1.7 350

90 131 98 109

n n

310

n

n

600
380

450 340

32 98 132

133

In contrast to the rest of this book, optical properties are given as an approximate absorption edge in the absorption (transmission) spectrum, rather than as a value of bandgap. This gives those who are not familiar with semiconductors a better feel for the appearance of the film. a As deposited, unless stated otherwise in the “Special conditions” column.

Since one common use of oxide films is for transparent, conducting coatings, the resistivities of these films were usually measured. Table 2.2 shows some basic electrical and optical properties of some of these films.

2.9.4.2

Halides

Silver halides, in particular AgI and AgBr, have been deposited by hydrolysis of halogenoalcohols (halohydrins) to free halide ions in a solution of Ag under

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acidic conditions. For AgCl, a better method was found to be simple precipitation by adding a solution of NaCl to one of AgNO3. The concentration of the solutions is important for this latter deposition—ca. 20 mM (much larger concentrations result in rapid aggregation and little film formation, while much lower concentrations give a very slow deposition rate). A visible AgCl film (visible by scattering) is formed very rapidly—in seconds—which is unique in CD, which normally requires much longer to form a visible film. The crystal size of these silver halide films is large compared to most CD films, rarely less than 100 nm and sometimes as large as a micron.

2.9.4.3

Elemental Se

Virtually all the semiconductors deposited by CD are compound semiconductors, the one exception being elemental Se. This has been deposited from solutions of selenosulphate, which rapidly form Se if acidified. By control of the pH, this reaction can be controlled to allow Se deposition to occur. Se films have also been deposited from colloids of Se (prepared by reducing SeO2 solutions) by photodeposition, whereby the light activates the formation of films.

2.9.5 Coprecipitation of Metal Chalcogenides— Ternary Compounds Chemical deposition is not limited to binary compounds. Ternary (and higher) compounds can be deposited by this technique. For the same reason as for the non II–VI and IV–VI compounds in Section 2.9.3, this section will suffice with a table of ternary compounds reported up to now, with two additions. The first is a brief consideration of the principles involved in the deposition of materials containing three or more elements. The second is to identify, in the table, which deposits have been clearly demonstrated to be a true single-phase solid solution rather than a mixture of two or more phases. If, e.g., thiourea is added to a mixture of Cd and Zn ions complexed with ammonia, then, depending on the mechanism and experimental conditions of deposition, the deposit could be CdS, ZnS, a mixture of the two, a single-phase CdxZnxS compound, or some combination of these. Structural characterization, most commonly XRD, together with elemental analysis, can usually reveal the nature of the product. There a number of factors involved in determining just what the product turns out to be. An obvious one is consideration of the solubility of the various products (and intermediates): The lower the value of Ksp of one binary product relative to another, the more likely that product is (in principle) to deposit preferentially. This simple consideration is complicated by a number of other factors. One is the tendency of metal ions to coadsorb on the (usually high-surface-area) pri-

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mary deposit, even if that metal chalcogenide (or whatever compound) is much more soluble than the primary one. Such adsorbed species can become occluded in the growing crystal as subsequent layers are built up. Another possibility, most likely for codeposition of structurally similar materials, is formation of a true solid solution. Even if the relevant binary materials are not structurally similar, a true single-phase ternary deposit may be obtained, although in this case the composition range is likely to be much narrower than in the previous case. For codeposition of two different cations, compound formation may occur, particularly if the two cations have quite different acidic and basic properties. Chapter 8 describes these principles in more detail. Table 2.3 lists ternaries that have been deposited, together with indication of when clear single compounds formation was verified. While solid solution formation is usually the goal of these studies, it should be kept in mind that separate phases, either as a composite or as separate layers, may be required for some purposes. For example, bilayers of CdS/ZnO and CdS/ZnS have been deposited from single solutions. These depositions depend on the preferential deposition of CdS over ZnS and, in the case of the former, the often-encountered greater ease of formation of the oxide (hydroxide) than the sulphide of Zn. TABLE 2.3 Ternary Materials Deposited by CD Material (Cd, Bi)S (Cd, Hg)S (Cd, Pb)S Cd(S, Se) (Cd, Zn)S

References

(Cd, Zn)Se CuBiS2

134 139, 140* 140, 145, 146–150* 152–154*, 155, 156a 158, 159*, 160, 161, 162–164* 166*, 167 169

Cu3BiS3 CuInS2 CuInSe2 Bi2 (S, Se)3 (Pb, Sn)Se (Pb, Bi)S

171* (after anneal) 173, 174*b 176–180* 183* 185* 186*

Material

References

In(S, OH) (Pb, Hg)S Pb(S, Se) (Sb, Bi)2S3 Sn(S, OH)

135, 136*, 137, 138 144–144* 151* 157* (after anneal) 165

(Zn, Cd)O Zn(O, OH), Zn(S, O, OH) Zn(S, OH) Zn(Se, OH) Zn(S, Se) (Pb, Cu)S (Cd, Sn)O

168* 170 170, 172 175 181, 182* 184 (two phase) 133*

Where a predominantly single phase (even over only part of the composition range) was at least reasonably clearly demonstrated or could be inferred from the results, at least at some composition, the relevant reference is followed by an asterix, although sometimes this refers to annealed films. a Probably solid solution, based on Ref. 154. b See relevant section in Chapter 8.

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2.10 NONAQUEOUS DEPOSITIONS 2.10.1 Introduction Although almost all CD processes have been carried out in aqueous solution, there are a few examples in nonaqueous media (specifically, in carboxylic acids). It is worth considering the differences between aqueous and nonaqueous deposition in general first. Solubility is an important criterion. The reagents involved need to be soluble in the solvent (this is, of course, obvious, but sometimes the obvious needs pointing out). This limits the choice of reagents, compared to aqueous solutions. For example, sodium selenosulphate, the most common selenide-producing reagent, is insoluble in most organic solvents at any useful concentration. Selenourea (or one of its derivatives) is more useful in this case. When considering the choice of metal salt, halides and particularly iodides as well as perchlorates tend to be more soluble in organic solvents than most other common anions. Next there is the question of whether the anion-producing reaction, which is normally a hydrolysis, can occur in the absence of water. Alternate reactions may be needed. This is not a problem for depositions which occur by the complex decomposition mechanism. Finally, even if these criteria are satisfied, there remains the question of whether the product will adhere to form a film or just precipitate homogeneously in the solution. This is the most difficult criterion to answer a priori. The hydroxide and/or oxy groups present on many substrates in aqueous solutions are likely to be quite different in a nonaqueous solvent (depending on whether hydroxide groups are present or not). Another factor that could conceivably explain the general lack of film formation in many organic solvents is the lower Hamaker constant of water compared with many other liquids; this means that the interaction between a particle in the solvent and a solid surface will be somewhat more in water than in most other liquids (see Chapter 1, van der Waals forces). From the author’s own experience, although slow precipitation can be readily accomplished from nonaqueous solutions, film formation appears to be the exception rather than the rule. The few examples described in the literature are confined to carboxylic acid solvents (see later). What are the advantages of deposition from nonaqueous solutions? One is the possibility to form films of compounds that are soluble, or not sufficiently insoluble, in water. A (potential) example of this is formation of halides using, e.g., chlorohydrins, which are in general soluble in organic solvents, to generate Cl. This could be expanded to materials that tend to form hydroxy-chalcogenides rather than pure chalcogenides, such as ZnS and In2S3. If water (including water of hydration from the salts) is rigorously excluded, the hydroxy impurity cannot form from solvents that do not contain hydroxy groups.

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2.10.2 Examples of CD from Nonaqueous Solutions Films of various stoichiometries of Sn-S have been deposited from carboxylic acid (acetic, propionic, butyric) solutions of elemental S and SnCl2. Depending on deposition conditions, in particular whether some water was present and how much, as well as the presence of complexing agents, films of approximate composition SnS, Sn2S3, or SnS2 could be formed. Interestingly, various Sn-S films were also formed on the walls of the deposition vessel above the liquid level (by several centimeters); this was attributed to reaction between volatile SnCl4 and H2S, both formed in the deposition bath. Bi2S3 was deposited from glacial acetic acid solutions of Bi(NO3)3 mixed with formaldehyde solutions of Na2S2O3 or thioacetamide. Sb2S3 films were deposited in a similar manner from acetic acid solutions. These depositions are described in more detail in Chapter 6.

2.11 RELATED DEPOSITION TECHNIQUES The CD technique is based on either slow formation of a reactive anion or slow decomposition of a complex compound. However, there are other techniques, not involving these slow steps, that nevertheless are sometimes called chemical deposition or chemical solution deposition or are closely related. These techniques are dealt with very briefly in this section.

2.11.1 Successive Ion-Layer Adsorption and Reaction (SILAR) Process The SILAR process is, as its name suggests, somewhat analogous to molecular beam epitaxy (MBE), although the films obtained are most often not epitaxial. Like MBE, SILAR proceeds via a layer-by-layer buildup of the film, except in solution instead of in a vacuum. In the SILAR process, the substrate is immersed first in a solution containing the metal cation, rinsed, then immersed in a solution containing the desired anion, and again rinsed. This gives (ideally) one monolayer of the deposit. The process is then repeated for as many times as needed to obtain the required thickness. (It is probably not important if the anion or the cation is adsorbed first; after the first cycle, the process should be the same, although it is conceivable that differences in adhesion to the substrate may result, depending on the initial order.) The rinsing steps are important, since without them relatively large reservoirs of one ion would remain on the substrate, and clusters of the semiconductor, rather than a film, would result. In fact, by omitting the rinsing step, films can be built up much more quickly. Thus, by successively immersing a glass slide in fairly strong solutions of Na2S and a Cd salt (say, 0.1M), visible yellow films

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FIG. 2.7 Simplified schematic diagram of an automated SILAR process for CdS.

of CdS appear after only several cycles. However, such films are less likely to be homogeneous, compared with a properly prepared SILAR film. Cu2O, ZnO, and SnS have been deposited in this way (see Ref. 187 and references therein). The true SILAR process is based on the expectation that, after each rinse step, only one monolayer of the previously adsorbed ion will remain, which will favor a layertype growth. The technique is slow and tedious, but automation of the process can be carried out, whereby the substrate is attached to stepping motors that alternately immerse and remove it from a series of beakers. This is shown schematically in Fig. 2.7 for the example of CdS. The substrate is attached to an arm that can be moved both vertically (to immerse/remove it) and horizontally (or in a circle) to position it above different reaction vessels. While the layers are usually polycrystalline, epitaxial layers of both zincblende and wurtzite CdS have been grown on various single-crystal substrates with lattice parameters close to those of CdS, as might be expected for an ion-byion growth. Besides CdS, many other semiconductors have been deposited by the SILAR technique as well as organic conducting polymers, such as polypyrrole and polyaniline. For representative references and to locate some groups working in this field, see Refs. 188–195.

2.11.2 Pyrolysis of Precursor Films This method is based on pyrolysis of a metal chalcogenide–containing precursor. Heating CD hydroxide films to form oxides is a simple example and is common in the deposition of some oxides. This will be treated in Chapter 7, which

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deals with oxides. Somewhat further from CD as understood in this book, the method is also widely used as the well-known sol-gel method for oxide films, in particular SiO2 and TiO2. In the sol-gel method, a more-or-less viscous colloidal gel is deposited on a substrate either by dipping the substrate in the gel and slowly raising it from the solution or by spin-coating. The film, whose thickness can be controlled by the viscosity of the sol and either the rate of removal from solution or the rotational speed of the spin-coater, is then pyrolyzed to form the oxide. This technique is not so easily extended to nonoxides. Sulphur, for example, does not form the cross-linking bonds needed to form the sol-gel as readily as does oxygen. However, a related method has been used, albeit to a very small extent, to form CdS films. It is based on the thermal decomposition (at ca. 300°C) of a Cd–thiourea complex, which is formed as a film by slowly withdrawing the substrate from a methanolic solution of a Cd salt and thiourea [196]. This can certainly be extended to other metal sulphides, using other complexes of sulphur (and also selenium). However, the complex and anion of the metal salt need to be chosen so that all the by-products of the pyrolysis reaction are volatile, otherwise the film will be contaminated with the nonvolatile by-products. For example, using cadmium nitrate and thiourea, all the by-products are volatile: Cd(NO3)2  (NH2)2CS → CdS  CO2  2H2O  2N2O

(2.20)

2.11.3 Exchange Reactions Exchange reactions to convert one material into another by immersion in suitable solutions are well known. Such a reaction is the basis for formation of the oncepopular CdS/Cu2S photovoltaic cell, where a CdS film was immersed in a CuCl solution and part of the film converted to Cu2S. Since the solubility product of Cu2S is lower than that of CdS (or, in electrochemical language, Cu is more noble than Cd), exchange is thermodynamically favorable. Examples of such exchange reactions using CD films are CdS and CdSe to the corresponding silver salts [197] and SnS2 to Ag2S [198].

2.11.4 Intermixing by Annealing Multilayers By depositing two (or more) different layers and annealing them, intermixing of the layers can lead to ternary and multinary compounds, although clear compound formation does not always occur. Thus, annealing (at 150°C, a relatively low temperature) ZnS-CuxS and PbS-CuxS films resulted in extensive interdiffusion of the metallic elements but no XRD confirmation of solid solution formation [199]. On the other hand, Sb2S3-CuS layers converted fully to CuSbS2 at 400°C, which ex-

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hibited p-type conductivity and a direct bandgap of 1.5 eV [200]. By evaporating In onto CD Sb2S3, InSb was formed after annealing at 300°C.

2.11.5 Gas–solution Interface Reaction Films of semiconductors up to a few hundred nanometers in thickness have been formed by the simple reaction between H2S or H2Se and the surface of an aqueous solution of the metal ion. This technique was described long ago for PbS and PbSe [201]. A more comprehensive description of the method, extended to a number of different metal sulphides, has been given [202]. It is stressed that the gas phase is passed over the solution surface and not bubbled through the solution, which would break up the film. In fact, if the gas flow continues for too long (typically more than a few minutes), the film tends to break up and precipitate. Since the substrate for these films is a liquid surface, the films can be (carefully) picked up and transferred to another surface or possibly even be self-supporting in small areas.

2.12 SPECIFIC TOPICS RELEVANT FOR CD FILMS 2.12.1 Solar Cells 2.12.1.1 Solid-State Photovoltaic Cells Probably the most important factor responsible for the renewal of interest in the CD technique is the almost universal use of CD CdS films in thin-film photovoltaic cells based on either Cu(Ga)InSe2 (abbreviated here as CIS, which in-

FIG. 2.8 Schematic diagram of the CIS/CdS and CdTe/CdS photovoltaic cells. The back contact to the CdTe cell—Cu-doped carbon paste – is a commonly used one, but there are several modifications to this contact as well as completely different ones in use.

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cludes Cu(In,Ga)Se2, usually abbreviated in the literature as CIGS, CuInSe2, and CuInS2) or CdTe absorber films. Schematic diagrams of these cells are shown in Figure 2.8. The common component of both is a thin (50–100 nm) film of CD CdS called the window layer [since incident light (ideally) passes through it] or the buffer layer (which points out the lack of understanding of exactly what this layer does). The CIS cell is a substrate (frontwall) cell (light passes through the semiconductors in the direction of the substrate), while the CdTe cell is normally in the superstrate (backwall) configuration (light passes first through the substrate). With only a few exceptions, the CdS is deposited from a standard ammonia/thiourea bath at ca. 70°C, with variations in the concentrations of reactants, the use of temperature programming, and some variation in pH (using an ammonium chloride buffer). It is notable that, in spite of many attempts to substitute CdS by another CD material (driven by the desire for a environmentally friendlier material), CdS remains the best material to date for this purpose, both for CIS and CdTe cells. Other materials deposited by CD include various Zn(OH,S), Zn(OH,Se), and In(OH,S) “compounds” and In(OH)3. The three first materials appear to be incompletely sulphided or selenided hydroxides, and it is not clear whether they are a mixed or a single phase. Also, it is usually unclear whether oxide or hydroxide also occurs [although one XPS study of In(OH,S) has demonstrated the absence of either In(OH)3 or In2S3 in the film]. While some of these buffer layers approach CdS in terms of cell efficiency, they are invariably inferior. Studies have been undertaken in an attempt to understand why CdS appears to be so unique in this role and why CD is the best technique to deposit it. The CD solution clearly plays an important role, not only in depositing the CdS, but also in its effect on the absorber surface. In the case of CuInSe2, the solution removes native oxides (of In, Cu, and Se), removes excess CuxSe, and forms an interface that is Cu-deficient and contains Cd. In fact, Cd has been shown to diffuse a small but appreciable distance (ca. 10 nm) into Cu(Ga)InSe2 films but not into CuInSe2 single crystals. It is not clear whether this is due to differences in composition or in crystallinity. However, the diffusion of Cd into CIS is believed to be related to the presence of a Cu-depleted region at the CIS surface; the CD bath, as already noted, is instrumental in forming such a region. Additionally, it is thought that removal of surface oxygen substituting for Se vacancies by replacement with S may increase band bending by modification of the surface charge. Since the envisionaged application of a CD process in thin-film solar cells is a large-scale one, efforts have been made to optimize the deposition process used, particularly in minimizing the waste Cd-containing solutions. Dilute Cd solutions (ca. 1 mM), a flow system with filtration, and a heated substrate have been employed to this end. The heated substrate means that deposition occurs preferentially on the substrate rather than on the cooler walls of the deposition vessel. Also, ethylenediamine has been used as a complexant rather than the much more volatile ammonia.

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There are a number of studies on other photovoltaic properties of CD films (see Chapter 9). As an example, p-n junctions have been fabricated by depositing PbS on a glass substrate partially coated with a trivalent metal, such as Al. The PbS deposited on the glass is (as usual) p-type, while that deposited on the metal coating was n-type, at least for some time. Photovoltages up to 0.1 V (at room temperature) and 0.28 V (at 90 K) were measured (the bandgap of PbS is ca. 0.4 eV – even lower at low temperatures so only small photovoltages are expected).

2.12.1.2 Photoelectrochemical Cells One of the attractive features of CD is its simplicity (in terms of carrying out the deposition, that is, not always in understanding the deposition itself). The same property of simplicity is often ascribed to photoelectrochemical cells (PECs). Therefore it is not surprising that CD has often been used to fabricate the semiconductor electrodes for PECs. A PEC is the liquid junction analogue of a solid-state Schottky cell. In its simplest configuration, a semiconductor with an ohmic contact is immersed in an electrolyte and connected via a load to a second counterelectrode (often platinum or graphite). Super-bandgap light incident on the semiconductor creates electron/hole pairs that are separated by the electric field (space charge layer) formed by the contact between the semiconductor and the electrolyte. One charge type (holes, for the more common case of n-type semiconductors) is injected into the electrolyte. The holes oxidize some electrolyte species, while the electrons are extracted through the ohmic contact and flow through an external circuit to the counterelectrode, where a reduction occurs. If the same species is both oxidized (at the semiconductor electrode) and reduced (at the counterelectrode), no net change occurs in the electrolyte, and electrical energy is generated in the external circuit. This type of cell is called a regenerative PEC. If different species are electrolyzed, chemical energy can be produced. Most PECs made using CD semiconductor films are of the former type. In nanocrystalline semiconductor films (commonly obtained in CD), the crystal size may be too small to support an appreciable space charge layer. Charges in that case are separated by differing kinetics between electron and hole injection into the electrolyte. The upcoming discussion on nonannealed films treats this in somewhat more detail. (Chapter 9 discusses PECs and their principles of operation more fully.) The majority of PEC studies have been carried out on either CdS or CdSe films, although many other CD semiconductors have been demonstrated to exhibit PEC activity. In most, though not all, cases, these films were annealed for optimum PEC response. Films annealed at temperatures above ca. 300°C usually exhibit a large degree of crystal growth, and therefore such films will be discussed separately from as-deposited films, which, in most cases, are composed of crystals 20 nm in size.

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Annealed Films. For many years, the CdSe photoelectrode—polysulphide electrolyte PEC—was probably the most studied system in PEC research. The bandgap of (bulk—see later) CdSe is ca. 1.7 eV, which is close to the theoretical optimum of 1.5 eV for photovoltaic cells in general; and in relative terms, that system was fairly stable in terms of self-oxidation of the semiconductor film in the electrolyte by the photogenerated holes. Simulated solar conversion efficiencies up to 6.8% on Ti substrates have been reported for annealed CD CdSe films in polysulphide electrolyte based on a low-ammonia-concentration–selenosulphate bath. Several successive depositions were required to build up an optimum final film thickness of 2.5 m (when most of the light was absorbed). The initial deposit was annealed to improve adherence and the final multideposited film was annealed at 550°C in air, followed by etching and zinc ion treatment. By incorporating silicotungstic acid (STA) in the deposition bath, larger particle size (after annealing at 430°C) and much better conversion efficiencies (compared to the STA-free bath) were obtained—11.7% based on tungsten-halogen illumination (solar efficiency is lower). The effect of the improved performance was not clear. It was suggested that the STA improved charge transfer kinetics at the CdSe/electrolyte interface. The larger particle size may also be an important factor. Chemical deposition CD CdS has shown much lower efficiencies in a PEC. This is due to its higher bandgap, which allows only a small fraction of solar radiation to be absorbed in the film. Nonannealed films. Although the conversion efficiencies are much lower than those of annealed films, the PEC properties of as-deposited films show other interesting properties, connected with their nanocrystalline and somewhat porous morphology. As already noted, there is usually no appreciable space charge layer in these nanocrystals. Since the films are porous and electrolyte can reach (most of) the surface of all the nanocrystals, charge generation can be considered to take place at the surface of the individual crystals—there is no need for a field to provide the driving force for charge drift. Electrons and holes are then separated by different kinetics for electron and hole transfer to the electrolyte, which in turn is affected by the relative trapping efficiencies of the charges at the nanocrystal surface. If one charge (say, holes) is removed rapidly by the electrolyte, the electrons can get to the back contact without recombination. In practice, much of the photogenerated charge is lost, probably by injection of both carriers into the electrolyte (indirect recombination). A unique property of these films results from this mechanism. CdSe, as deposited, behaves as an n-type semiconductor (holes are transferred to the electrolyte, while electrons are extracted at the back contact). In this case, the holes are preferentially trapped at the surface, and are more readily injected into solu-

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tion. Mild etching in dilute HCl changes the distribution of trapping states at the surface in such a way that electrons are now more readily trapped and are preferentially injected into the (identical polysulphide) electrolyte, the film behaving as a p-type semiconductor. It is notable that this behavior, typical for a crystal size of 4–5 nm, is not observed when the crystal size is 20 nm. The very high surface area and high density of trapping states appear to be determining here. As-deposited CdS has been studied as a photoelectrode with various dopants (Al, As, Cu) incorporated in the deposition bath. The emphasis in these studies was on PEC efficiencies, which were very low in all cases, although doping, particularly with As, did have a beneficial effect. An interesting variation of photoactivity has been observed during the deposition of several semiconductor films (CdSe, CdS, PbS, Bi2S3). Illumination during deposition increases the deposition rate and, in some cases, increases the crystal size somewhat. While some effects can be ascribed to local heating, increasing deposition rate, the main effect is probably electron-hole generation by super-bandgap light and reduction of the chalcogen species to chalcogenide by the photogenerated electrons at the growing crystal surface, resulting in the formation of more metal chalcogenide. Some other semiconductors have been investigated as photoelectrodes, in all cases giving low but appreciable photoactivity. These include Bi2S3, PbSe, Sb2Se3, SnS2, HgS, and Ag2S. Chemical deposition has also been used to form both the photoelectrode (CdSe) and the storage electrode (Ag2S) in a photoelectrochemical storage cell, where the Ag2S acts as an in situ storage electrode.

2.12.2 Quantum Size Effects in CD Films Films of materials deposited at or near room temperature (and in this respect 100°C is considered to be near room temperature) tend to have a small crystal size. This is not surprising since high temperatures are normally required to impart sufficient mobility to a freshly deposited species in order for recrystallization to occur. This small crystal size, which at one time was almost universally considered to be a disadvantage, is increasingly considered to be an advantage as interest in nanocrystalline and nanoparticle materials grows. The term nanocrystalline usually refers to materials with a crystal size from a nanometer up to hundreds of nanometers (at this upper limit, the term microcrystalline starts to take over). The size of the crystals formed in CD films is often small enough that quantum-size effects become apparent. The terms quantum size effect and size quantization are normally used to describe a material whose energy structure differs from that of the bulk material. As crystal (or, more generally, particle) size decreases, charges (electrons and holes) in the particles are constrained in an increasingly small volume. When the particle size becomes smaller than the Bohr diameter of the charges in the bulk material (between 2 and 20 nm for many ma-

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terials), the quantum-size effect is manifested as an increase in the bandgap, Eg, of the material and separation of the energy bands into discrete levels. The increase in Eg is most commonly seen as a blue shift (i.e., to higher energy) of the optical absorption spectrum (see Chapter 10 for more details on this topic.) This effect is shown in Figure 2.9 for CdSe films deposited from baths containing Cd complexed with NTA (nitrilotriacetate) and Na2SeSO3 as a Se source. The nanocrystal size, measured by both XRD and TEM, varied from ca. 3 nm up to 20 nm with increase in temperature and/or change in mechanism from a cluster mechanism to an ion-by-ion deposition. The optical bandgap shifts from 1.8 eV (for bulk, zincblende CdSe) to ca. 2.4 eV for the smallest nanocrystals (ca. 3 nm). The main difference between the two mechanisms as they relate to crystal size (discussed in Sec. 2.6) is that the cluster mechanism is three dimensional while the ion-by-ion one is mainly two dimensional. Crystal size in the former is limited largely by the amount of reactant per nucleus: The more nuclei, the smaller the final crystal size, since the same concentration of reactants is divided over more nuclei. Temperature affects this by stabilizing (kinetically) smaller nuclei as temperature is lowered, thus increasing the number of nuclei at lower temperature,

FIG. 2.9 Transmission spectra of CD CdSe films deposited at various temperatures from CdSO4/NTA/Na2SeSO3 baths. All samples deposited by hydroxide cluster mechanism except 80°C HC (high complex), which proceeded via the ion-by-ion mechanism. The effective bandgap can be approximated by the wavelength (photon energy) a little shorter (higher) than the absorption onset. A second absorption knee, ca. 0.4 eV to higher energy of the initial onset, seen clearly in the 41°C and 80°C samples, is due to a transition from the spin-orbit valence level to the lowest conduction level and is commonly observed in these films.

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resulting in a smaller final crystal size. For the two-dimensional ion-by-ion mechanism, the final crystal size is probably determined by the number of nuclei on the substrate (initially bare, eventually covered with deposit), which grow until they touch other crystals and cannot grow further in the horizontal (to the substrate) direction. Increased temperature also may allow increased coalescence of the individual crystals into larger ones. Other CD semiconductors have been shown to exhibit size quantization. PbSe shows the effect very clearly, since quantum size effects can be clearly seen in this material, even in crystals up to several tens of nanometers in size (due to the small effective mass of the excited electron-hole pair). Shifts of greater than 1 eV have been demonstrated, from the bulk bandgap of 0.28 eV to 1.5 eV. CdS, deposited from the usual ammonia/thiourea bath, normally gives films that are not quantized, with crystal size larger than 10 nm. Some exceptions do occur, however. Small increases in bandgap have been found when high thiourea concentrations were used. At high values of pH (12), smaller crystal sizes (down to ca. 5 nm) have been obtained with blue spectral shifts corresponding to increases in bandgap of up to a few hundred meV for very thin films. The crystal size increased with increase in film thickness. Using NTA as a complex and working under conditions where the cluster mechanism is operative, 5-nm nanocrystals of CdS exhibiting quantum size effects have also been obtained. This crystal size did not vary much with film thickness. Using a process of electrochemically induced CD of CdS (see Chapter 4), nanocrystalline CdS films were grown using 2mercaptoethanol as a strongly adsorbing growth-termination (capping) agent. By increasing the concentration of mercaptoethanol, crystal size was reduced down to 4.1 nm (and bandgap increased up to 2.69 eV). Both ZnS and ZnSe films have been grown that show moderate increases in bandgap (up to a few hundred meV). The Zn chalcogenides generally exhibit smaller increases in bandgap than the corresponding Cd compounds of the same size, due to their larger effective masses. HgS, deposited from thiosulphate solution, has been described with crystal sizes that depend on deposition temperature, from 3 nm to 8 nm and corresponding variation in apparent bandgaps from 2.4 to 1.9 eV. HgSe has been deposited by CD exhibiting different bandgaps and ones strongly shifted from the bulk value (bulk HgSe is a semimetal with a negative bandgap). Values as high as 2.5 eV for 7.7 nm crystal size have been reported. Another bath composition gave a bandgap of 1.42 eV, although this was not explained through size quantization but because of an amorphous structure. Bi2S3 has been deposited with film thickness–dependent crystal sizes, from 5.2 to 8 nm and corresponding bandgaps, measured from the absorption spectra, from nearly 2.3 down to 1.8 eV. Cu-S films, of various stoichiometries, have also shown small quantum size effects.

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Apart from these reported quantum size effects, there are a number of materials that, while quantum size effects have not been explicitly invoked, have been reported to show blue spectral shifts that may be explained in the same manner. These include Cu1.8S, various CuS compositions (Eg 1.8—2 eV), Sb2S3, and SnS. These systems are all described in Chapter 10. While shifts due to size quantization have most commonly been seen in absorption spectroscopy, other spectroscopies, such as photoelectrochemical photocurrent, photovoltage (using a vibrating Kelvin probe), photoluminescence, and photoconductivity spectroscopies have all shown quantum shifts in various CD films.

2.13 APPLICATIONS (ACTUAL AND POTENTIAL) OF CD FILMS We conclude this chapter with applications of CD semiconductor films, both those that have been realized and potential uses. The most important use of CD films for many years was to make PbS and PbSe films for photoconductive detectors [10]. Such detectors, made by CD, are still in use today, although they are facing competition from photovoltaic III–V detectors. It should be noted that for good photosensitivity, air-annealing of the CD films is carried out, and this annealing treatment is connected with partial oxidation of the PbS and PbSe surfaces. Today, the the most important “application” for CD films is the use of CD CdS as the window (or buffer) layer in thin-film photovoltaic cells [16]. Both CdTe- and CuInSe2-type absorber films use this procedure. Such cells have reached the pilot plant stage, and there appears to be no obvious competitor for the CD CdS at present. The study of CD semiconductors, and in particular CdSe and CdS, for use as photoelectrodes in photoelectrochemical cells is connected with this use, although much farther from likely application. This study was driven, to a large extent, by the simplicity of deposition of the films, a particularly sought-after requirement for this purpose, both from the point of view of applications and because it allowed many groups (usually chemists) who maybe did not have ready access to conventional vacuum deposition systems to prepare the films. The success of CD CdS in photovoltaic cells has driven related research with potential applications in other semiconductor devices. Since the CD process seems to play a role in the favorable properties of the CdS windows by decreasing interface recombination, studies of its passivation properties on other interfaces and surfaces have been carried out, with considerable success. For example, when a very thin film (ca. 6 nm) was deposited between InP and SiO2, the resulting reduction of the interface state density led to improved electrical properties of metal-insulator-semiconductor capacitors and field effect transistors (FETs)

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based on this interface [203,204]. These improvements were correlated with removal of native oxides and protection against oxide regrowth by the CdS, as well as with removal of phosphorus vacancies at the interface. Improvements were also obtained in InAlAs/InGaAs transistors and metal-semiconductor-metal photodetectors by the same treatment [205]. It is clear that these effects are connected not only with the properties of the CD films but also with reaction between the substrate semiconductor surface and the CD solution. Another example of CD films applied toward electronic devices is the fabrication of thin-film field effect transistors by depositing CdS (50 nm) onto oxidized n-Si and annealing at 400°C [206]. Another potential application is for solar control coatings. Thin films of certain sulphides—in particular those of copper and lead—are reasonably transmissive to visible solar radiation while reflecting most of the infrared radiation. If used as a window coating, the heat rejection from the solar radiation will result in reduced heating of the interior by solar radiation, compared to a noncoated glass window [23,24]. Such coatings are commonly in use, particularly in large buildings, and often take on an attractive golden appearance by reflected light. They are applied by vacuum techniques. If the material usage is sufficiently high (as has been demonstrated for CD CdS used for photovoltaic cells), CD is an attractive alternative to deposit these coatings. The high infrared reflectivity of these films implies highly doped (therefore relatively highly electrically conducting) films. However, another type of solar control coating, a solar shield for passive cooling, requires films that have a high transmission in the mid-infrared. Passive cooling occurs when a surface emits more radiation than it absorbs [207]. The (cloudless) atmosphere is relatively transparent to mid-infrared radiation (between ca. 8–13 m—the atmospheric window); therefore, radiation of these wavelengths can be emitted from a surface at ambient temperature through the atmosphere to space, which is a sink at a very low temperature (theoretically 4 K, in practice higher, but still very much lower than the surface of the Earth). This leads to a cooling effect at night and explains why car surfaces and grass become covered with much more dew during a clear night than a cloudy one. However, during the day, solar radiation would swamp this cooling effect. To prevent this, a shield is required that is transparent to the 8–13-m region but that blocks the complete solar spectrum. Low-bandgap semiconductors, such as PbS and PbSe on polyethylene (one of the very few materials, readily available in large areas, that is transparent to the atmospheric window region), are suitable for this purpose [208], but they must be close to intrinsic to prevent free carrier absorption and reflection in the infrared region. Nanocrystalline films (and CD films are often nanocrystalline) are more likely to be intrinsic than large-grain ones: unlike a macroscopic crystal, it is possible to obtain small nanocrystals with zero doping concentration, i.e., truly intrinsic. Other potential applications for CD films have been suggested and studied.

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Related to the solar control coatings just described, solar absorbing coatings can be used, e.g., for water heating. Multilayered stacks of CD PbS and CdS on Nicoated Cu can be configured to minimize reflectivity in the solar region and minimize emittance in the thermal infrared region [209]. The relatively high electrical conductivity of CD CuxS films has been exploited to form an electrical contact to ferroelectric films, as a partially transparent conducting film on plastic and also as a cupric ion sensor [210]. Our brief overview has given uses for CD films and rationales for their study. It should be clear that more uses are likely to develop as the present resurgence in interest increases the pool of knowledge in the field and allows deposition of better-quality films, higher reproducibility, new materials and old ones in different forms than usual.

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3 Mechanisms of Chemical Deposition

In this chapter, we critically discuss the various mechanisms for CD. Since the kinetics of deposition are intimately related to the mechanism, this aspect is also treated here. This chapter is divided into several sections dealing with different aspects of the mechanisms. However, there is a large degree of overlap of the material in these different sections; indeed, it was often not obvious in which section a specific topic should be placed. While the main discussion of any specific point will be limited to the most relevant section, the same point will often be treated, more briefly, in other sections. However, for those readers who are looking for a specific topic in this chapter, it would be prudent not to be confined to only what appears to be the relevant section. For example, elements of the mechanism, probably with different emphases, are dealt with also in the sections on nucleation and on kinetics, while many issues of nucleation are considered in the mechanistic studies. There is one example of a CD process (for deposition of tin sulphides) in which elemental sulphur dissolved in a nonaqueous solvent is used as a source for S. Since this appears to be the only example in the literature for this type of film deposition, it will be discussed in Chapter 6 together with the relevant study on tin sulphides. However, there is no reason to believe that this process may not be applicable to other materials. Metal sulphides (and selenides) are known to form, as precipitates, by reacting certain metal salts with dissolved elemental chalcogen, although visible film formation seems to be limited, up to now, to this one example.

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3.1 THE FOUR BASIC MECHANISMS In spite of the fact that CD has been in use for a long time and that the reactions involved appear to be quite straightforward, the mechanism of the CD process is often unclear. There is a good reason for this: There are several different mechanisms of CD. These can be divided into four fundamentally different types. Using the common thiourea deposition of CdS as an example, these are: The simple ion-by-ion mechanism 2 Cd(NH3) 2  4NH3 4 D Cd

(3.1) (dissociation of complex to free Cd2 ions) (NH2)2CS  2OH → S2  CN2H2  2H2O (3.2) (formation of sulphide ion) (CdS formation by ionic reaction) (3.3) Cd2  S2 → CdS The simple cluster (hydroxide) mechanism nCd2  2nOH D [Cd(OH)2]n

(3.4) (formation of a solid Cd(OH)2 cluster) [Cd(OH)2]n  nS2 [from Reaction (3.2)] → (3.5) nCdS  2nOH (exchange reaction)

The complex decomposition ion-by-ion mechanism: (NH2)2CS  Cd2 D [(NH2)2CSMCd]2 [(NH2)2CSMCd]2  2OH → CdS  CN2H2  2H2O

(3.6) (3.7)*

The complex decomposition cluster mechanism: [Cd(OH)2]n [from Reaction (3.4)]  (NH2)2CS D (3.8) [Cd(OH)2]n1(OH)2CdMSMC(NH2)2 [Cd(OH)2]n1(OH)2CdMSMC(NH2)2 → (3.9) [Cd(OH)2]n1CdS  CN2H2  2H2O which continues until conversion of all the Cd(OH)2 to CdS. The first two mechanisms involve free sulphide ions (or other anions), while the last two are based on breaking of a carbon–chalcogen bond and do not involve formation of free chalcogenide. Most mechanistic studies have assumed the for* The use of the simple thiourea–Cd ligand in Eq. (3.7) is for simplicity. There are a number of different complexes involving Cd and thiourea, in particular, some containing hydroxy groups. As long a a solid phase of Cd(OH)2 is not present, then such hydroxy–thiourea Cd complexes involve an ionby-ion-type of mechanism, as exemplified in Eq. (3.7).

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mation of free anions (not always justifiably). It is also possible that more than one mechanism occurs in parallel or that conditions change during deposition so that the deposition mechanism also changes. Since most mechanistic studies were carried out for CdS deposition from an ammoniacal solution and using thiourea as a sulphide source, and a smaller number for CdSe deposition using selenosulphate as selenide source, this chapter will mostly use these materials as examples. However, it should be stressed that the concepts discussed in this chapter are, for the most part, valid for the deposition of all semiconductors.

3.2 SLOW FORMATION OF THE CHALCOGENIDE IONS The rate-limiting step in CD for the first two mechanisms is almost always formation of the chalcogenide ion. This reaction should be slow; otherwise fast, homogeneous precipitation of the metal chalcogenide will occur with little film formation. (Even rapid precipitation can lead to a film; however, this film will be extremely thin and in most cases not visible.) Almost all the literature on CD is limited to sulfides (mostly), selenides, and oxides (including hydrated oxides and hydroxides). Anion-forming reactions are described in this section.

3.2.1 Sulphides 3.2.1.1

Thiourea

Thiourea (the sulphur analogue of urea) is the most commonly used sulphur precursor. In alkaline solution (in which depositions involving thiourea are carried out), the first step of the hydrolysis gives sulphide ions and cyanamide: SC(NH2)2  OH D HS  CN2H2  H2O

(3.10)

The cyanamide can hydrolyze further, with the overall reaction, if carried to completion, going via urea to ammonium carbonate: H2O 2H2O CN2H2  → CO(NH2)2  → (NH4)2CO3

(3.11)

Cyanamide can also react with ammonia to give guanidine: NH 3 CN2H2  → (NH2)2CBNH

(3.12)

All of these compounds can be (and have been) found as impurities in CD films. However, the important step is the formation of sulphide ion. In neutral and acidic solutions, thiourea can be decomposed to thiocyanate ion [1], which can be useful if the intention is to deposit thiocyanates:  SC(NH2)2 → NH 4  CNS

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(3.13)

It should be kept in mind that many of these decomposition reactions are equilibria. The decomposition of thiourea in the absence of a metal ion will normally be much slower than in the presence of such an ion. The metal ion removes sulphide as metal sulphide—the less soluble the sulphide, the more effective the removal at very low sulphide concentrations. This continuous removal of sulphide shifts the equilibrium to the direction of more sulphide production. The same principle holds for many other anion precursors.

3.2.1.2

Dimethylthiourea

Dimethylthiourea is much less commonly used than the more available thiourea: (CH3)NHC(S)NH(CH3)  H2O → (CH3)NHC(O)NH(CH3)  H2S

3.2.1.3

(3.14)

Thioacetamide

Thioacetamide has been used for a long time as an analytical reagent to precipitate metal sulphides (see Ref. 2 for many relevant references). Thioacetamide can be hydrolyzed over a wide range of pH and is often used for CD in acidic baths. In a strongly acidic bath (pH  ca. 2), H2S is formed: H3C.C(S)NH2  2H2O → CH3COOH  H2S  NH3

(3.15)

It has been shown [3] that this reaction can proceed by two pathways, one in which the carbon–sulphur bond is broken first: H3C.C(S)NH2  H2O → H3C.C(O)NH2  H2S

(3.16)

forming acetamide as an intermediate, or a pathway in which the carbon–nitrogen bond is first broken to give thioacetic acid: H3C.C(S)NH2  H2O → H3C.C(S)OH  NH3

(3.17)

which then is hydrolyzed to H2S and acetic acid. The H2S dissolves in water as hydrosulphide ion: H2S  H2O D HS  H3O

(3.18)

In an alkaline bath, the overall reaction is: H3C.C(S)NH2  2OH → CH3COO  NH3  HS

(3.19)

Hydrolysis in alkaline solution is considerably faster than in acid solutions. At intermediate pH values, particularly in weakly acidic solutions (pH  2), metal sulphide formation using thioacetamide may proceed through decomposition of a metal ion (or solid phase)–thioacetamide complex rather than through intermediate formation of sulphide [4] (see Sec. 3.3.3). Thioacetamide in pure water is fairly stable and does not readily hydrolyze at room temperature.

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3.2.1.4

Thiosulphate

The earliest CD processes were carried out using thiosulphate. Although thiourea (and to some extent thioacetamide) are now more commonly used to deposit sulphides, thiosulphate is still sometimes used. While the reaction pathways listed below are intended to suggest possibilities for reactions involving thiosulphate, it must be noted that the mechanism(s) for these reactions is (are) still not clear. Mechanisms have been proposed in the CD literature, but no convincing proof for any particular one has been forwarded. Thiosulphate depositions are most often carried out in a weakly acidic bath (pH 3). Several reactions are possible in such solutions: 2 S2O 2 3  H2O → H2S  SO 4

(3.20)

with the H2S dissolving to give HS, as in Reaction (3.18):  → S  HSO S2O 2 3 H 3  S2O 2  2H → S  SO  H2O 3 2

(3.21) (3.22)

and in alkaline solution: 2   S2O 2 3  OH → HS  SO 4

(3.23)

It has been suggested that the thiosulphate, a reducing agent, may act as an electron donor and reduce the elemental sulphur formed in Reactions (3.21) and (3.22), forming sulphide ions: S  2e → S2

(3.24)

Because of the strong complexes thiosulphate forms with some metal ions, it is very possible that these metal–thiosulphate complexes undergo a complex-decomposition mechanism (Section 3.3.3). However, one early study on the formation of PbS on boiling Pb2 and thiosulphate in water found that PbS formed more readily when excess thiosulphate was present [5], which suggests that decomposition of thiosulphate to sulphide might be the dominant pathway under the conditions of that study.

3.2.2 Selenides 3.2.2.1

Selenourea

Selenourea (SeC(NH2)2) was apparently first used for CD by Kutzscher’s group in Germany during World War II. It appears that this work was not published; references to it come through other sources [6]. The first published use of selenourea for CD appears to be in 1949 by Milner and Watts [7] to deposit PbSe for photoconductive cells. It was the main reactant used to form selenide films (mainly PbSe) through the 1960s, after which selenosulphate (see later) became the dom-

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inant selenide precursor. Other examples of its use are for ZnSe [8] and HgSe [9]. Selenourea is an unstable compound that requires the presence of a reducing agent—usually Na2SO3—to minimize oxidation to elemental Se. The selenide formation presumably parallels thiourea hydrolysis in Eq. (3.10): SeC(NH2)2  OH → HSe  CN2H2  H2O

3.2.2.2

(3.25)

Dimethylselenourea

Dimethylselenourea (CH3)2NC(NH2)Se (for preparation see Ref. 10) was used to deposit films of PbSe [11]. It is more stable than selenourea, although still not very stable in solution and, like selenourea, was used together with Na2SO3 to prevent oxidation. The optimum pH for this deposition was 9.8; bulk precipitation occurred at a pH of 10.1, while the deposition slowed greatly as the pH decreased even a few tenths of a pH unit. The activity of this solution, as measured by the rate of deposition, was reported to increase with time (hours). This was explained, in view of the measured reduction in the rate of reaction by sulphite (which was assumed to complex with the dimethylselenourea), by reduction of the concentration of sulphite with time (by oxidation to sulphate). This means that, for reproducible results, the age of the dimethylselenourea/sulphite solution should be taken into account or an aged solution (at least 10 hr old) should be used (the change in activity of the solution slows down to a large degree after this time).

3.2.2.3

Selenosemicarbazide

Selenosemicarbazide (H2NN(H)C(Se)NH2) was used by Velykanov et al. to make Cd, Zn, Ag, and Hg selenide films and precipitates from aqueous alkaline solutions [12]. As with the other Se precursors, Na2SO3 was used to stabilize the selenosemicarbazide against rapid decomposition. This reagent was reputed to be more stable than the previous two.

3.2.2.4

Selenosulphate

Selenosulphate (Na2SeSO3), which is the analogue of thiosulphate with the active S atom substituted by Se, was used by Kitaev for deposition of PbSe [13] and CdSe films [14]. Since it is more stable, simpler to prepare, and cheaper than selenourea, it simplified the deposition of selenides and has for a long time, with only a few exceptions, been the precursor invariably used to deposit selenides. It can be prepared by dissolving elemental Se in an aqueous sulphite solution: 2 Se  SO 2 3 D SeSO 3

(3.26)

The Se dissolves slowly (typically an hour or two at 60–70°C). Although more stable than selenourea, it does slowly hydrolyze and becomes less active with age—fairly rapidly for the first few days after preparation and then slowly; it is still usable for weeks after preparation without any special storage demands, al-

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though at increasingly reduced activity. Storage under cold, oxygen-free conditions will slow the aging further. This gradual change in activity should be taken into account in depositions using this material, since it is not always practical to make it fresh for each deposition. The mechanism for hydrolysis of selenosulphate is often given as   2 SeSO 2 3  OH → HSe  SO 4

(3.27)

SO 2 4

However, although is likely to be a final product, the reaction is probably not as simple as this. A suggested first step for the hydrolysis is [15] 2   2SeSO 2 3  H2O → HSe  SeS2O 6  OH

(3.28)

Selenosulphate is unstable in acidic solutions. If the pH of a fairly concentrated selenosulphate solution is reduced below ca. 7, red Se will precipitate out. This property has been used to prepare films of Se by slightly acidifying dilute solutions of selenosulphate [16].

3.2.3 Tellurides Very few examples of telluride deposition have been reported. There are a number of reasons why tellurides are more difficult to deposit by CD than sulphides or selenides. One is the very negative redox potential of the telluride/tellurium couple (E0
1.14 V). This means that a strong reducing solution is necessary. Another, directly related reason is the instability of telluride ion (and possible tellurium precursors). Even dissolved oxygen will rapidly oxidize telluride. The strong reducing solution will need to reduce dissolved oxygen, preventing this reaction. Thus a combination of a suitable Te source and a strong reducing agent with exclusion (or reduction) of oxygen can succeed in forming tellurides. In fact, such conditions have also been found necessary to form films of ZnSe using selenosulfate and are preferable for deposition of ZnS films (see Sec. 3.5 and Chap. 4). CdTe has been deposited by hydrazine reduction of TeO2 [17,18]. The potential of hydrazine oxidation is sufficiently negative to allow formation of telluride: N2H4(aq)  4OH → N2  4H2O  4e Te  2e D Te

2

E0
1.16 V E0
1.14 V

(3.29) (3.30)

especially because only very small concentrations of telluride need to exist at any time, causing a positive shift of the Te reduction through the Nernst equation (Eq. 1.32). Te was considered to be essentially insoluble in Na2SO3 under normal conditions, although preparation of tellurosulphate, the analogue of selenosulphate, under hydrothermal conditions has been reported [19]. A recent study has described CdTe deposition using this reagent [20]. Apparently the solubility, while low, is suf-

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ficient to be useful. At this stage, it is assumed that the tellurosulphate reacts analogously to selenosulphate (Eqs. 3.26–3.28). As with the TeO2 reduction just described, because of the great ease of oxidation of telluride to elemental Te, the films invariably contain elemental Te. (See Sec. 4.3 for details of this deposition.) While not true CD, a novel telluriding agent, Te dissolved in an alkaline solution of hydroxymethanesulphinic acid, has been used to convert Cd(OH)2 films to CdTe [21]. While there is some doubt as to the nature of the active telluriding agent, from the description of the preparation process of this reagent—a change in color from deep purple (characteristic of polytelluride ions, Te 2 x ) to a faint pink (pure Te2 is colorless but will be colored this way if traces of Te 2 x are present) as the preparation proceeds—it does appear to contain free telluride ion. It should be noted that elemental Te can slowly dissolve in concentrated air-free alkaline solutions, with the formation of polytelluride and its characteristic purple color.

3.2.4 Oxides Oxides have been commonly deposited by CD. In many cases, the deposit is a hydroxide or hydrated oxide formed by reaction of the metal ions with slowly generated hydroxide. A variety of precursors has been used to generate the hydroxide.

3.2.4.1

Urea

Urea is slowly hydrolyzed in water to form ammonium carbonate: (NH2)2CBO  2H2O → (NH4)2CO3

(3.31)

Carbonates are alkaline since they dissociate to some extent to form hydroxide ions:   CO 2 3  H2O D OH  HCO 3

(K
1.8  104)

(3.32)

The hydroxide so formed will react with the metal ion to form the metal hydroxide, hydrated oxide, or oxide, depending on the relative stability of the various oxides and hydroxides. (The resulting hydroxides or oxy hydroxides can be heated in air or oxygen to form the oxides.) In addition, insoluble carbonates may also form. The competition between hydroxide and carbonate will depend on their solubility products. Carbonates tend to be more soluble than hydroxides of the same metal ion. On the other hand, the value of K for equilibrium (3.32) (1.8  104) means that the concentration of hydroxide will be ca. four orders of magnitude less than that of carbonate (assuming no other pH-determining species is present). While the urea method has been commonly used in the past to form bulk precipitates of basic salts, including oxides, for analytical purposes, it was noted that transparent, adherent films were typically formed on the walls of the beaker in which the precipitation was carried out [22]. Notably “basic stannic sulfate,” which was probably mainly SnO2 since very little sulphate was found in the ac-

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companying precipitate, was found to adhere so tightly to the walls of the beaker that it was difficult to remove [23].

3.2.4.2 Dimethylamineborane (DMAB) and Trimethylamineborane Dimethylamineborane (DMAB) and trimethylamineborane have been used to deposit Zn [24], In [25] and Fe [25a] oxides using the metal nitrates. The nitrate anion is believed to be reduced by the DMAB to nitrite and hydroxide:   (CH3)2NHBH3  2H2O → BO 2  (CH3)2NH  7H  6e

NO 3



NO 2



 H2O  2e →  2OH x  M  xOH → M(OH)x

(3.33) (3.34) (3.35)

In the case of Zn the oxide forms spontaneously, while for In the hydroxide film is heated to form the oxide. This mechanism is actually a mixed electrochemical process involving partial anodic and cathodic reactions (i.e., charge exchange occurs at an interface) rather than a pure chemical deposition process.

3.2.4.3

Persulphate

Persulphate (also called peroxydisulphate) (S2O 82) is a very strong oxidizing agent and has been used to deposit oxides, often in a higher oxidation state that that of the original metal ion. Films grown using this reagent include -PbO2, NiO, MnO2, and Tl2O3. The two latter required a small concentration of Ag ions in the deposition solution as a catalyst. Ag is a known catalyst for oxidations using S2O 2 8 (it is oxidized by persulphate to Ag(III), which is then the active oxidation agent). It is probable that Pb and Ni also act likewise, while Mn and Tl do not (or much less so). A study of the use of persulphate for the deposition of PbO2 has provided strong evidence that the deposition actually is a mixed electrolytic process (similar to that proposed for dimethylamineborane depositions) rather than a pure chemical deposition [26]. The partial electrolytic reactions were given as Pb2  2H2O → PbO2  4H  2e 2  S2O 2 8  2e → 2SO 4

(3.36a) (3.36b)

Persulphates hydrolyze to form (finally) sulphate and hydrogen peroxide; the probable overall reaction can be given by 2  S2O 2 8  2H2O → 2SO 4  H2O2  2H

(3.37)

This reaction occurs via various radical species, and it is also possible that some oxide depositions occur by these radicals (the formation of which is probably initiated by fission of the S2O 2 8 ion to two sulphate radicals) or directly by the hydrogen peroxide, which itself usually involves a free-radical reaction. Hydrogen peroxide itself has been used directly to form oxide films in a few cases.

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3.2.4.4

Other oxide-forming reactions

Aqueous solutions of acidic metal salts are usually inherently unstable and hydrolyze readily to oxides (the hydroxides of these metals tend to be relatively unstable), in some cases forming films. Such hydrolysis can be more readily controlled by adding boric acid to fluoro-complexes of the metal, e.g.: H2TiF6  2H2O D TiO2  6HF

(3.38a)

followed by removal of the HF by the boric acid:  H3BO3  4HF D BF 4  H3O  2H2O

(3.38b)

Ref. 27, which describes such reactions as liquid phase deposition, gives more details on this method. Finally, many ammoniacal or amine solutions of metal salts will form films of hydroxides or oxides on heating or even long standing at room temperature. Ammonia is very volatile and will gradually be lost (in an open system), resulting in reduced complexation of the metal ions. The pH will also drop, but it is likely that the increase in concentration of free metal ions due to loss of ammonia dominates. Chapter 7 treats the deposition of oxides in detail.

3.2.5 Halides Halide ions can be formed by hydrolysis of alkyl halides, as shown for AgCl precipitation [28]: RMCl  OH → Cl  ROH

(3.39)

Haloalcohols (halohydrins) which are more water soluble than alkyl halides, have been used to generate chloride ions [29], and are therefore more suitable for aqueous CD. The hydrolysis of haloalcohols has been used to deposit films of AgBr and AgI, e.g., ICH2CH2OH  H2O → I  HOCH2CH2OH  H

(3.40)

3.2.6 Other Anions Although they do not appear to have been used (at least deliberately) to form films, there are other slow anion-generating reactions. Although most of the compounds formed in these reactions are not considered semiconductors in the normally accepted sense, they merit at least a mention here.

3.2.6.1

Sulphate

Sulphate is formed by reaction between persulphate and thiosulphate: 2 2 2 S2O 2 8  2S2O 3 → 2SO 4  S4O 6

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(3.41)

This would be limited to cations that form insoluble sulphates but soluble persulphates and thiosulphates (Ba and Sr were demonstrated by Lamer and Dinegar [30]). Thiosulphate, in particular, forms soluble complexes with many cations and therefore should (often) not present a problem in this respect, as long as the metal sulphide is not formed under the conditions of the deposition. In addition, solvents other than water can be used in principle, and therefore it might be possible to deposit sulphates that are soluble in water but insoluble in another solvent.

3.2.6.2

Phosphates

Phosphates have been formed by slow hydrolysis of trialkyl phosphates, hydrogen phosphate ions, or metaphosphoric acid, which liberate phosphate ions (see p. 46 in Ref. 31). This may be of interest as a precursor for the preparation of phosphide semiconductors.

3.2.6.3

Arsenates

Arsenates have been described in one case exploiting the fact that the zirconyl cation forms a water-soluble arsenite but insoluble arsenate. By adding nitric acid to a solution of zirconyl chloride and sodium arsenite, the arsenite was oxidized to arsenate by the nitric acid, precipitating the insoluble zirconyl arsenate [32]. As for phosphates (and probably more readily), arsenates might be reduced to arsenides. It should be stressed that these reactions were used to form precipitates and not films. There is no guarantee that films can be formed using these reactions. However, it is reasonable to expect that, under the right conditions, it may be possible to produce films of these compounds. It is left as an exercise for the curious reader to find these “right” conditions. Details of these and other slow precipitations are given in Ref. 31. There are a number of examples of homogeneous precipitation of hydroxides based on slow cation release, such as destruction of the Fe–EDTA complex with H2O2 (see Ref. 33). In CD, the only well-defined example of this is heating an ammonia complex (e.g. of Cd2). The loss of ammonia by volatilization will gradually increase the concentration of free Cd2 ions.

3.3 MECHANISMS OF CHEMICAL DEPOSITION Having dealt with slow formation of the reacting anion, we now turn to the various mechanisms by which the CD compounds are formed. For the most part, the details of nucleation and film growth are left to the following section. Here we concentrate on the reactions taking place that form the semiconductor material. There are four main mechanisms for the compound formation, as outlined in Sec. 3.1; which one is operative depends on the specific process and reaction parameters. 1. 2.

Simple ion-by-ion mechanism Simple cluster (hydroxide) mechanism

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3. 4.

Complex-decomposition ion-by-ion mechanism Complex-decomposition cluster mechanism.

3.3.1 (Simple) Ion-by-Ion Mechanism The conceptually simplest mechanism, often assumed to be the operative one in general, is commonly called the ion-by-ion mechanism, since it occurs by sequential ionic reactions. The basis of this mechanism, illustrated for CdS, is given by Cd2  S2 D CdS

(adsorbed on the substrate)

(3.42)

If the ion product, [Cd2][S2], exceeds the solubility product, Ksp, of CdS (ca. 1028; Table 1.1), then CdS can form as a solid phase, although a larger ionic product may be required if supersaturation occurs. If the ion product does not exceed Ksp, no solid phase will form, except possibly transiently due to local fluctuations in the solution, and the small solid nucleii will redissolve before growing to a stable size. For that reason, the precipitation process is shown as an equilibrium rather than as a one-way reaction. This reasoning, that no solid phase phase will form if the ion product does not exceed Ksp, may not be true under certain circumstances. CdS formation might occur at the substrate surface under conditions where none can be formed homogeneously. This is an important point; however, since it will be treated in more detail in the following section, on hydroxide formation, it therefore will not be discussed further in this section. Equation (3.42) gives the overall reaction for formation of CdS. However, the process is more complicated than this and comprises a number of reactions and equilibria. The mechanism involves the formation of S2 ions and control of Cd2 concentration. The S2 can be formed by a number of methods, as already described. Here we consider the most common one, the decomposition of thiourea by aqueous alkaline solution: (NH2)2CS  2OH → S2  CN2H2  2H2O

(3.43)

Since the S2 concentration can be made as low as desired simply by controlling the rate of Reaction (3.43) (e.g., using low temperatures and/or relatively low pH), in principle, even at relatively high free-Cd2 concentrations, the deposition rate should be easily controlled. Since an alkaline pH is required to decompose the thiourea to sulphide, a complexing agent is needed to keep Cd2 in solution and to prevent Cd(OH)2 from precipitating. As explained in Chapter 1 [Eqs. (1.26) and (1.27) and calculation following these equations], using ammonia as a complexant for Cd (0.1 M total Cd concentration) and at a pH of 11, a concentration of ammonia
1.19 M will be needed to prevent formation of Cd(OH)2 at room tem-

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perature—more at the higher temperatures more commonly encountered in CdS deposition. (We shall see, in the following section, that although this statement is perfectly accurate, it does not mean that CdS deposition will not occur at lower concentrations of ammonia at the same pH.) When reading the literature, in many (probably most) cases it is not clear whether the deposition proceeds by an ion-by-ion process. The reason is that, unless another mechanism is specifically discussed, it is often assumed that the deposition proceeds via the ion-by-ion mechanism. If the exact deposition parameters are known, which mechanism is operative can, in most cases, be calculated. Two criteria have often been cited in the literature as proof of deposition via the ion-by-ion mechanism. One is epitaxial deposition of the CD film. (Epitaxy refers to growth of one material on another in such a way as to result in coherence between the lattice of the substrate and the deposit. Often—although not necessarily—the lattice of the deposit is aligned in the same direction as that of the substrate.) This is based on the expectation that a cluster mechanism will not result in an epitaxial film; for this to occur, clusters of maybe thousands of atoms would need to be able to rearrange themselves on the substrate. Some examples of epitaxial growth are given in Sections 3.4.2 and 4.1.5.2. The second criterion, sometimes seen in the older literature, is that films formed via the cluster mechanism should be poorly adherent and optically scattering, while those formed via the ion-by-ion mechanism will be adherent and transparent. Unlike the epitaxy, this criterion is faulty; films formed by the cluster mechanism can be highly transparent and strongly adherent, while there are examples of films that proceed via the ion-by-ion mechanism that are not well adherent and are optically scattering. The degree of adhesion of the film and its transparency say little, if anything, about the mechanism of the deposition. Thus, following this reasoning, it was claimed that ion-by-ion deposition occurred under conditions where a visible Cd(OH)2 suspension was present [34]. While the presence of Cd(OH)2 does not exclude the possibility that ion-by-ion deposition occurs, it is unlikely that ion-by-ion deposition would be dominant when a high concentration of Cd(OH)2 is present. More recently, there have been a small number of studies that provide strong evidence for the ion-by-ion mechanism. It must be pointed out, however, that while it is not very difficult to distinguish between an ion-by-ion and cluster mechanism in most cases, it is much more difficult to distinguish between a simple ion-by-ion and a complex-decomposition ion-by-ion mechanism. Therefore most investigations that conclude an ion-by-ion mechanism is operative, while usually assuming the simple ion-by-ion process, do not distinguish between the simple and complex-decomposition pathways. One investigation has shown a clear-cut boundary in the crystal size of films (CdSe, CdS, and, to a lesser extent, PbSe), depending on whether the deposition occurred via an ion-by-ion or a cluster mechanism [15]. The solution conditions

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were constant except for a change in the ratio between complexant and Cd. A small crystal size was obtained under conditions where Cd(OH)2 was proven to exist (as a colloid, not visible to the eye) in the deposition solution, while a larger crystal size was obtained when no Cd(OH)2 was present as a solid phase. Thus the separation between the conditions where the ion-by-ion and cluster mechanisms occurred was clearly shown. This study is dealt with in more detail in the following section. A recent investigation of CdS deposited from a thiourea bath found that the initial deposit (on mica) formed as islands ca. 0.5 nm high and 10–40 nm across; further growth led to an increase in height without changing the lateral dimensions [35]. Such a growth mode supports an ion-by-ion mechanism, since a cluster mechanism, whereby (presumably fairly symmetric) hydroxide clusters adsorb on a substrate, is expected to lead to an equally symmetric initial growth mode. The island size in this study was measured by atomic force microscopy, where the tip size can determine the measured lateral dimensions of particles smaller than the tip. For this reason, scanning tunneling microscopy (or electron microscopy for lateral dimensions) would be more reliable in such measurements. This investigation was carried out using triethanolamine as complexant and under conditions where deposition was slower than usual. As always the case, the results and interpretation cannot be extrapolated to other bath conditions.

3.3.2 Hydroxide (Cluster) Mechanism We noted in the earlier example that Cd(OH)2 precipitation should be avoided and calculated the concentration of ammonia required to prevent this precipitation. In reality, CDs are quite often carried out under conditions where a metal hydroxide (or hydrated oxide) is formed. This might seem to imply that a precipitate of, e.g., Cd(OH)2 is formed at the start of such depositions. In fact, this is (usually) not the case; the Cd(OH)2 is formed, but either as a colloid rather than a precipitate, or as an adsorbed species on the substrate but not in the bulk of the solution. Since Cd(OH)2 is colorless and colloids usually do not scatter light (otherwise it is termed a suspension), this means that the Cd(OH)2 colloid is not apparent to the eye. The CdS is then formed by reaction of S2 ion with the Cd(OH)2: Cd2  2OH → Cd(OH)2

(3.44)

followed by Cd(OH)2  S2 → CdS  2OH

(3.45)

The driving force for Eq. (3.45) is the much lower value of Ksp for CdS (ca. 1028; Table 1.1) than for Cd(OH)2 (2  1014), which reflects the more negative free energy of formation of the former. This means that sulphide will readily substitute for hydroxide in the case of Cd. An idea of the amount of sulphide needed to con-

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vert Cd(OH)2 into CdS can be calculated from the ratio of the two solubility products: [Cd2][S2] Ksp (CdS) 
 Ksp(Cd(OH)2) [Cd2][OH]2 (3.46)

[S2] 1028
5  1015
  2
 [OH ] 2  1014

At a pH of 11 ([OH]
103 M at room temperature), this gives a value for [S2] of 5  1021 M. Thus it requires a very low concentration of S2 indeed to begin to convert Cd(OH)2 into CdS. Even taking into account that most of the sulphide species will be HS rather than S2, at pH 11 the [S2]:[HS] ratio is 106.3 (or 5  107) [Eq. (1.17)], and the required HS concentration will be 5  1021
1014 M  5  107

(3.46a)

still a very low concentration. At the higher temperatures usually used for CdS deposition, the OH concentration will be about an order of magnitude higher (because of the strong temperature dependence of the ionic product of water), and the required S2 concentration will be about two orders of magnitude higher than required at room temperature. The participation of Cd(OH)2 in the deposition of CdS (and other metal chalcogenides) has been demonstrated or suggested on many occasions. Kitaev et al. presented a theoretical thermodynamic treatment of the Cd2/ammonia/ thiourea system to show when Cd(OH)2 should be present as a solid phase in the deposition solution [36]. A graphical representation of this analysis is shown in Figure 3.1. This graph is based on two equilibria: the solubility product of Cd(OH)2 and the stability constant of the ammonia (ammine) complex of Cd. Consider first the former: Cd2  2OH D Cd(OH)2

Ksp
2  1014 2

(3.47) 

Based on this equilibrium, we can express [Cd ] in terms of [OH ] for the case where Cd(OH)2 will just precipitate; at higher pH it will precipitate and at lower pH it will not. Since [OH] can be converted into [H] (and therefore pH) through the ion product of water (see Chap. 1), a graph can be made of pH vs. [Cd2] (or p[Cd2], which, analogously to pH, is equal to minus the logarithm of [Cd2]). This is the hydroxide line in Figure 3.1. Its physical meaning is that above this line, Cd(OH)2 will be present in the solution, while below it there will be no Cd(OH)2. A similar (somewhat more complicated) calculation can be made based on the stability constant of the Cd–tetrammine complex [see Eq. (1.27)] and using the hydrolysis of ammonia (see Refs. 34 and 36 for details of the calculation), which

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FIG. 3.1 Regions of stability for the Cd–ammonia system for 0.1 M total Cd concentration and at room temperature. The hydroxide line separates conditions where Cd(OH)2 will (above the line) or will not (below the line) thermodynamically exist. The complex line gives the concentration of free Cd2 at any pH value (where the pH is determined only by the ammonia concentration and not, e.g., by added alkali). (Adapted from Ref. 34.)

gives the free-Cd2 concentration as a function of pH. This is shown as the complex line in Figure 3.1. (The calculation is made for a total-Cd2 concentration of 0.1 M, but it is only slightly dependent on this concentration; e.g., for a Cd2 concentration of 1 mM, the line will shift ca. 0.2 pH units to lower pH values. Also, the presence of other ammine and hydroxy complexes should strictly speaking be taken into account; however, these considerations will not result in major changes to the overall picture.) Since the free-Cd2 concentration is a function of ammonia concentration, the p[Cd2] can be also identified with a p[NH3] (top axis). Considering both the hydroxide and complex lines together, for the left side of the figure, where the complex line is above the hydroxide line, the concentration of free Cd2 will always be high enough to form Cd(OH)2. Where the complex line is below the hydroxide line, however (the right side of the figure), Cd(OH)2 will not form at pH values above the complex line but below the hydroxide line, since the Cd2 concentration will not be high enough to form

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Cd(OH)2. Extra OH ions (e.g., added KOH) are required to increase the pH and form Cd(OH)2 in this region. This type of analysis (which can be extended to other metal ions and other complexants) provides a basis for choosing conditions for depositions involving hydroxide clusters. An elegant example of this was given by Kitaev et al [36] for CdS deposition in the absence of ammonia (or any other complex, except for thiourea itself, which is a weak complexant for Cd). Extrapolation of the hydroxide line to p[Cd2]
1 (0.1 M free Cd2) gives a pH of 7.7, or, alternatively, a pOH of 6.3; above this value of pOH, no Cd(OH)2 is formed. However, at higher temperature, the ion product of water increases considerably (Table 1.2). Thus, while pOH (
pH for pure water) at room temperature is 7, at 80°C it becomes 6.3. Thus, while no Cd(OH)2 is formed in a 0.1 M solution of a Cd salt (ignoring effects of the Cd salt and thiourea on the solution pH), at 80°C it can thermodynamically form. Kitaev et al. demonstrated experimentally that a solution of 0.1 M cadmium acetate and thiourea produced a film of CdS only at 90°C and above. This analysis, of course, is based on the hydroxide cluster model. Deposition by another mechanism could conceivably occur at lower temperatures given enough time. Rieke and Bentjen made a detailed study of the role of Cd(OH)2 in the deposition of CdS films from the ammoniacal thiourea bath [37]. They first studied Cd(OH)2 formation in the absence of thiourea. Cd(OH)2 formation, as measured visually by laser scattering (which shows up small amounts of turbidity in the solution), began at a pH of ca. 10.4 (the solid phase, once formed, could exist for quite long periods of time down to a pH as low as 10). Surface analysis (XPS) showed that Cd(OH)2 formed on silicon substrates at a pH value as low as 9, even though no Cd(OH)2 formed in the solution under those conditions. The Cd(OH)2, about 1.5 nm average thickness, was apparently stabilized against dissolution by the substrate. In the presence of thiourea, a SEM study of the early stages of deposition showed that large, platelike crystals of CdS formed at a pH where Cd(OH)2 was present in the solution but that this deposit was very nonadherent and could easily be wiped off. However, at lower values of pH, where Cd(OH)2 had previously been shown to occur only on the substrate, strongly adherent CdS spherules were observed. The density (in terms of surface coverage) of this deposit decreased with decreasing pH, and the coverage became low at low values of pH, where surface-adsorbed Cd(OH)2 did not occur. Therefore, only films deposited under conditions where appreciable amounts of surface-adsorbed, but not bulk, Cd(OH)2 occurred were of good quality (adherent and specularly reflecting). In connection with this study, in particular the suggestion that the Cd(OH)2 was stabilized by the substrate against dissolution, it has been shown that Co(OH)2 can form at a solid (SiO2) surface at a pH lower than that necessary to cause bulk precipitation of Co(OH)2 [38]. This was explained by the effect of the electric field at the solid/liquid interface on the dielectric constant of the interface

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region and, through this, on the free energy (and therefore the solubility product) of the relevant hydroxide dissolution. It was noted that this field effect should increase with an increase in the charge of the cation. In terms of CD, this would mean that the difference in the ease of deposition of the hydroxide at a solid surface and homogeneously in solution will increase with the charge on the cation. Furthermore, this argument is not limited to hydroxides but should be valid for all depositions, including chalcogenides (and therefore also ion-by-ion deposition), as was briefly mentioned in the previous section. Betenekov et al. [39] used an isotopic tracer technique to show that, for their range of solution compositions, the initial deposition involved adsorption of Cd(OH)2 on the glass substrate. At the beginning of the reaction, only Cd was observed to form on the substrate and this was interpreted to be due to Cd(OH)2, since any other insoluble Cd compounds that might be formed from the deposition solution (containing CdCl2, NaOH, NH4OH, and thiourea dissolved in water) were expected to contain either S or C.* However, they concluded that the deposition proceeded, not by reaction between Cd(OH)2 and sulphide formed by decomposition of thiourea, but rather by decomposition of a Cd(OH)2–thiourea complex (see Sec. 3.3.3.1). O’Brien and Saeed studied the CdS deposition, only using ethylenediamine instead of ammonia as complexant because of the better-defined coordination properties of the former [40]. In common with the former studies, they also found the presence of Cd(OH)2 necessary for the deposition of good-quality CdS films by comparing the conditions needed to obtain such films with those calculated using the relevant thermodynamic parameters. A short digression at this point is required to define the term good-quality film. The use of the adjective good-quality depends very much on what is required from the film. In the context of CD and thin-film preparation in general, good quality almost always refers to two parameters: good adhesion and specular reflection (smooth). Of course, requirements can be envisaged where these properties, in particular the latter, are not desired, such as if a rough “scattering” film is required. However, keeping this caveat in mind, we will continue to use the term good-quality as it is almost always used in the CD literature. While most mechanistic investigations have been carried out on CdS, other semiconductors, in particular CdSe, have also been studied with regard to the deposition mechanism. Kainthla et al., in their study of the formation of CdSe films from ammoniacal solutions containing sodium selenosulfate, noted that when a visible precipitate of Cd(OH)2 was present in their solutions (obtained by adding * It needs to be repeatedly emphasized that the mechanism deduced in any one investigation is not necessarily valid for different experimental conditions. At the same time, it does appear likely that for this very well-studied system of CdS deposition from ammonia/thiourea baths, Cd(OH)2 is, in many cases, involved in the deposition.

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KOH), a thin terminal thickness (ca. 80 nm) of CdSe was obtained, compared with several hundred nanometers obtained under optimal conditions (no precipitate) [41]. If the pH was too high and Cd(OH)2 precipitation was extensive, no film was observed to form. In fact, an analysis of their results suggests that films are obtained only if Cd(OH)2 is present (not necessarily as a visible precipitate). While they do not specifically consider Cd(OH)2 as a chemical intermediate in their reaction, they do conclude that Cd(OH)2—both adsorbed on the substrate and in the solution—acts as a nucleation center for CdSe formation. Gorer and Hodes carried out a study of CdSe deposition from selenosulphate solutions of Cd complexed with nitrilotriacetate—NTA [15], which is a stronger complex for Cd than is ammonia and with which is easy to obtain conditions where no Cd(OH)2 is present, even at relatively high values of pH. A change in reaction mechanism, from hydroxide cluster to ion-by-ion, was observed by monitoring the optical transmission spectra of the films deposited on glass. The basis for this investigation is the change in semiconductor bandgap, and thus in its spectrum, when the crystal size becomes very small—the quantum size effect (discussed in detail in Chap. 10). The CdSe crystal size (and therefore the bandgap) is not strongly dependent on deposition parameters within a fairly wide range of parameters. However, it was observed that if the ratio between the NTA and Cd concentrations was increased above a certain value—called the critical complex ratio—the crystal size suddenly increased (for fixed values of temperature and pH). This was seen first by a red shift in the optical spectrum (see Chap. 10 for examples of these changes) and was subsequently verified by direct measurements of the CdSe crystal sizes in the films using TEM and XRD. A variety of analytical techniques were then used to verify that Cd(OH)2 was present in the solution when the complex:Cd ratio was below the critical value (Rc) but absent above it. Cd(OH)2 absorbs in the UV range of the spectrum, and spectral monitoring of Se-free solutions showed that it was present only below Rc. Light scattering by a blue laser also confirmed the presence of a heterogeneous phase below Rc but not above it. Similar XPS analyses to those employed by Rieke and Bentjen for CdS showed that Cd adsorbed on the glass substrate, immersed in Se-free solutions, only below Rc. This is seen in Table 3.1: Appreciable amounts of Cd (as Cd(OH)2) were seen only when the pH was sufficiently high and the complex:Cd ratio relatively low.

TABLE 3.1 Cd:Si Ratio Measured by XPS pH Cd:Si

9.0 0.020

10.0 0.024

10.5 0.17

11.0 0.45

11.0 0.005

For glass substrates immersed in Cd2/nitrilotriacetate solutions (no selenosulphate) at different pH values for 5 min at 40C. NTA:Cd ratio
1.63 except for the rightmost column, where it is 2.25.

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Theoretical thermodynamic calculations of the conditions under which Cd(OH)2 should form were also carried out, based on the solubility product of Cd(OH)2, the stability constants of the Cd-NTA system, and the ion product of water at different temperatures. The values of Rc computed from these calculations agreed with those measured experimentally. In agreement with Kainthla et al. and contrary to at least some conditions for thiourea-based CdS deposition, Gorer and Hodes found that adsorption of a Cd hydroxide species on the substrate occurred only under conditions where the solid hydroxide formed also in the solution. This need not necessarily be interpreted as contradictory; it may be due to the different conditions involved. An important factor in the reaction for CdSe deposition via the hydroxide mechanism was an observed gradual increase in pH, ca. 0.8 pH units over the course of the deposition. It is not clear what the cause of this increase is. However, it means that while Cd(OH)2 may not be formed at the start of the deposition, it may form during the deposition. This can then explain the induction time where no apparent reaction takes place initially. For the preceding experiments, the standard conditions were to set the solution pH to 10. Coloration of the solution (indicating formation of CdSe) occurred when the pH reached ca. 10.3 (depending, of course, on the temperature and complex:Cd ratio). If the pH was adjusted to 10.3 at the beginning of the reaction, coloration began almost immediately rather than after a more typical time of several minutes. This increase in pH did not occur for the ion-by-ion mechanism, probably because the excess NTA necessary in this case also acts as a buffer. In any case, an increase in pH is not required for this mechanism. These results for CdSe were extended to the deposition of PbSe (using citrate as complex) and CdS (using NTA and thiourea) and found to be also applicable for these cases [15]. For PbSe, the transition, while clearly evident, was not so sharp, since citrate is a weak complex and a solid phase is always present, but to greatly differing degrees depending on the conditions. The colloidal phase in this case is a hydrated lead oxide, which, in a selenosulphate-free solution, adsorbs strongly onto the substrate (see Sec. 3.4.3 and Fig. 3.5). This hydrated oxide is rapidly selenized to PbSe, and the process involves breaking down of the relatively large oxide crystals during reaction with the selenosulphate and recrystallization of the PbSe product. This can be seen by following the reaction of precipitated hydrated lead oxide with selenosulphate by XRD (Fig. 3.2), where the sharp peaks of the oxide transform into very broad peaks indicative of an almost amorphous structure, which themselves vanish as sharp PbSe peaks appear. A quartz crystal microbalance study of the kinetics of CdSe deposition from the preceding solution (nitrilotriacetate) showed an importance difference in the mode of growth of the CdSe below and above Rc [42]. Below Rc (hydroxide mechanism), a periodicity in the deposition rate was observed, with a period corresponding to a film thickness of ca. 5 nm—i.e., the approximate size of a single

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FIG. 3.2 XRD spectra showing the process of PbSe formation from the reaction of precipitated hydrated lead oxide with Na2SeSO3 solution. (a) Starting material; (b–e) after 1.5, 3, 4.5, and 6 mn reaction, respectively. (Adapted from Ref. 46.)

crystal. Above Rc, the period corresponded to 0.3 nm, a single CdSe monolayer, as expected from an ion-by-ion growth. This interpretation is based on growth by coverage of the surface by one layer—whether a layer of crystallites or a single Cd-Se layer—at a time. The critical ratio concept already described was derived for a fixed concentration of metal ions. As pointed out by O’Brien and McAleese [43], while useful at high metal concentrations, it requires modification when low concentrations are employed, due to the changes in stability of complexes when dilute. They developed a system whereby, on a plots of total metal concentration against total complex concentration, a curve defining a constant value of free-metal concentration was drawn. Figure 3.3 shows several such plots for the Cd–ethylenediamine system studied by them. From these plots, the need for much higher complex:metal ratios for dilute solutions is evident. The complex:metal ratio required for a constant concentration of free metal, shown in the top curve, is reasonably constant for higher concentrations but increases strongly at low concentrations. We postpone detailed explanation of what determines the different crystal sizes in the two mechanisms to Section 3.4. At this point, it is enough to say that the CdSe crystal size in the hydroxide mechanism will be determined mainly by the size of the Cd(OH)2 particles in the solution and on the substrate, while that

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formed in the ion-by-ion mechanism will depend on the heterogeneous nucleation on the substrate. It should be repeated at this stage that this mechanism is dependent on a large difference between the solubility products of the hydroxide and chalcogenide of the required metal. The situation for ZnS, for example, is considerably less favorable. The values of Ksp for the hydroxide and sulphide of Zn are 8  1017 and 3  1025, respectively. The same calculation for Zn as carried out earlier for CdS shows that a S2 concentration of 4  1015 M is required at pH
11 to convert the hydroxide into sulphide; this is a million times more than that required to form CdS. This does not tell us at this stage whether ZnS will form or not, but only that it is less likely to than CdS. We will return to this problem when we discuss the specific deposition of II–VI compounds in Chapter 4. It is quite possible that the mechanism will change in the course of the deposition. As the metal is depleted from solution, the complex:metal ratio will increase

FIG. 3.3 Equivalent solution contour plots for solutions of Cd2 and ethylenediamine [en] at 50°C. The curves represent conditions for a constant concentration of free Cd2
109M. Bottom curve (solid line): [total Cd] against [ethylenediamine]. Top curve (dotted line): the ratio of [ethylenediamine] to [total Cd] against [ethylenediamine]. (Adapted from Ref. 43.)

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and may pass the point where no solid hydroxide phase is present in the solution. In this case, the ion-by-ion process will occur if the conditions are suitable. This is most likely if the initial conditions were close to the border between ion-by-ion and cluster mechanisms. The opposite may also occur: initial deposition by an ion-byion pathway followed by clusters, which either build up gradually in the solution or change their aggregation properties, adhering to the film. This has been shown to occur for CdS deposited from a thiourea bath using a combination of quartz crystal microbalance (which measures the mass of the deposit) and electrochemical impedence spectroscopy (which provides indirect structural information) [44]. Change in aggregation properties of the colloids present during the deposition was suggested as the cause of the change in deposition. It is also possible that both mechanisms occur in parallel. Thus, deposition might begin with nucleation of Cd(OH)2, but growth might occur on this Cd(OH)2 via an ion-by-ion mechanism: Cd(OH)2  S2  Cd2 → Cd(OH)2CdS

(3.48)

either by itself or together with conversion of the Cd(OH)2 to CdS, as in Eq. (3.45). Alternatively, each mechanism may occur independent of each other. An example of where this latter case appears to occur is in the deposition of PbSe [45,46], where, under certain conditions, two spatially separated domains of small crystals (cluster mechanism) and larger crystals (typical of the ion-by-ion mechanism) are obtained. The relative concentrations of these two domains can, of course, be varied by changing the solution composition, but the fact that they occur as separated domains on the substrate suggests that differences in the substrate from one region to the other also plays a role. Another factor that could lead to a change in the details of the mechanism with time is the buildup of reaction products in the solution. For example, the cyanamide formed in the decomposition of thiourea [Eq. (3.10)] can react with Cd ions adsorbed on the surface of the CdS to give the sparingly soluble CdCN2:  (CdS)n Cd 2 ads  CN2H2 D CdSCdCN2(ads)  2H

(3.49)

The CdCN2 would then react with any sulphide present to form the more insoluble CdS. This would be parallel to the more straightforward sulphidization of (CdS)n Cd 2 ads to CdS. Before finishing this section, we present a few words concerning the effect of deposition temperature on the terminal film thickness. Although exceptions have been reported, it is a general observation that a lower deposition temperature results in ultimately thicker films (of course, the deposition rate is slower). There may be a number of reasons for this. One is that higher temperatures favor the hydroxide cluster mechanism due mainly to the higher hydroxide concentration (the strongly temperature-dependent ion product of water) and also to the lower stability of complexes at high temperature (which may be offset by higher solubility of the metal hydroxide). This effect has been shown for CdSe deposition, where a

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higher complex:Cd ratio is needed to prevent Cd(OH)2 formation at higher temperatures [15]. The ion-by-ion mechanism, since it often does not result in homogeneous precipitation, or at least less than in the cluster mechanism, tends to yield films of larger terminal thickness. This will be valid even if the mechanism is a mixture of the two mechanisms. Another effect of low temperature is (in the absence of stirring) reduced mass transport. This can cause two opposing effects: reduction of reactants reaching the substrate and reduction of aggregation (collisions between colloidal particles). The concentration of free chalcogenide ions will also increase essentially exponentially with increasing temperature. However, it is difficult to predict the effect this will have on terminal thickness; rates of homogeneous precipitation and film formation will both increase. It is possible that thinner films will occur for the ion-by-ion mechanism, since there is an increased probability of the occurrence of homogeneous precipitation with increasing chalcogenide ion concentration (the product of chalcogenide concentration and metal ion will increase—also because the free-metal ion concentration will be somewhat higher at higher temperature due to reduced stability of the metal complex).

3.3.3 Complex-Decomposition Mechanism The complex-decomposition mechanism can, as with the free-anion mechanisms, be divided into ion-by-ion and cluster pathways. However, since experimental data relating directly to the complex decomposition mechanism is rather sparse, it will be dealt with in one main section rather than two. The cluster pathway has been more emphasized in these studies and will be dealt with first. It is very important to stress at this point that almost all experiment data described below could be explained in terms of a simple (anion-mediated) mechanism. Equally valid, most of the data described in the previous section could be explained by a complex-decomposition mechanism.

3.3.3.1

Cluster Mechanism

The basis of this mechanism is that a solid phase is formed but, instead of reacting directly with a free anion, it forms an intermediate complex with the “anionforming” reagent. Continuing with CdS deposited from a thiourea bath as our example, this would be given as MCd(OH)2  (NH2)2CS D Cd(OH)2SC(NH2)2

(3.50)

where MCd(OH)2 is one molecule in the solid-phase cluster. This complex, or a similar one containing also ammine ligands, then decomposes to CdS: MCd(OH)2SC(NH2)2 → MCdS  CN2H2  2H2O i.e., the SMC bond of the thiourea breaks, leaving the S bound to Cd.

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(3.51)

Such a mechanism was suggested by Betenekov et al. based on their isotopic tracer technique discussed in the previous section [39]. They suggested that Cd(OH)2 forms initially on the substrate and catalyzes the thioureau decomposition. Of course the catalytic effect of the solid surface could be to decompose thiourea to sulphide ion and not necessarily to catalyze the complex-decomposition mechanism. A similar catalytic effect of PbS on the decomposition of thiourea had been suggested previously by Norr [47]. Kinetic measurements by Rieke and Bentjen suggested that CdS likewise catalyzed thiourea decomposition [37]. OrtegaBorges and Lincot also deduced such a mechanism based on kinetic measurements of the CdS deposition using a quartz crystal microbalance [48]. In this case, the measurements were found to fit best with a complex-decomposition model. Both they and Rieke and Bentjen found optimum deposition to occur under conditions where Cd(OH)2 was formed as a surface species on the substrate but not in the bulk of the solution. Kinetic measurements also led Doña and Herrero to a similar conclusion of a complex-decomposition mechanism, but with the main difference that the initial adsorbed species is not Cd(OH)2 itself but an ammine–hydroxide [49]:  Cd(NH3) 2 4  2OH D [Cd(OH)2(NH3)2]ads  2NH3

(3.52) 

They based this modification on the known adsorbance of OH on glass and on the common occurrence of transition metal mixed water–ammonia complexes with coordination number of 4. Parallel structural studies of the deposited CdS showed textured growth, supporting a mechanism whereby alternate Cd and S species were involved, in an ion-by-ion process. Such a growth suggests adsorption of a molecular hydroxy-ammine species rather than a cluster. In fact, the mechanism of Ortega-Borges and Lincot also does not differentiate between a hydroxide cluster and molecule. Unfortunately, it is nontrivial to distinguish reliably between the complexdecomposition and sulphide-formation mechanisms. For example, in the study of PbS (as a precipitate) formation from thiourea [47], the two main results used to support complex decomposition were: (a) very little sulphide was formed in alkaline solutions of thiourea and (b) addition of PbS powder catalyzed the reaction, seen by the disappearance of the induction time for precipitation and more rapid PbS formation when PbS was added at the start of the reaction. However, these results would also be obtained in a free-anion mechanism, for the following reasons: (a) Thiourea decomposition is an equilibrium reaction [see Eq. (3.10)]. Formation of sulphide will shift the equilibrium back to the left. If a metal ion, which forms a sulphide with a low-solubility product, is present in the solution, however, it will remove even a very low concentration of sulphide continually as formed.

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(b) The presence of a large solid surface in the solution will reduce the induction time, even if the mechanism proceeds through free-sulphide formation, since the initial nucleation step will be facilitated. These factors do not argue against the complex-decomposition mechanism, but they should not be too readily interpreted, in the absence of other evidence, as evidence against the sulphide mechanism. Granted, this is an old study, but it does point up the difficulty in distinguishing between the two mechanisms. Kinetic studies and subsequent fitting of the data from these studies to various models [48,49] appear to be the best way of approaching this problem at present.

3.3.3.2

Ion-by-Ion Mechanism

Consider the complexation of free Cd2 by thiourea to give a Cd–thiourea complex ion: Cd2  (NH2)2CS D [(NH2)2CSMCd]2

(3.53)

This ion could, in principle, hydrolyze by breaking the SMC bond to form CdS: [(NH2)2CSMCd]2  2OH → CdS  CN2H2  2H2O

(3.54)

This would lead to CdS formation in solution. If the Cd2 is adsorbed on the substrate (either directly or indirectly through a hydroxide linkage) or on previously deposited CdS, then the same reaction would occur. If the CdS so formed remained bound to the substrate (it is assumed that CdS generated on previously deposited CdS would remain bound), the result would be film growth by an ion-by-ion, complex-decomposition mechanism. As with the cluster mechanism, it is difficult to distinguish experimentally between the complex-decomposition mechanism and the free-anion pathway. Some studies have involved deposition (or precipitation) from acidic solutions. It is reasonable to assume that no hydroxide is present under these conditions for most metal ions commonly used in CD and that deposition occurs via an ion-by-ion mechanism. Thioacetamide decomposition at intermediate pH values, particularly in weakly acidic solutions (pH  2), has been suggested to occur through a thioacetamide complex rather than through intermediate formation of sulphide [4]. Of course, this process may also occur in parallel with either acid or alkaline hydrolysis of thioacetamide to (ultimately) sulphide at certain pH ranges. It is also possible that this complex-decomposition reaction occurs at both high and low pH values in certain cases. The Cd–thiourea example that has been mainly used up to now in this section is a very weak complex. However, there are examples where the chalcogenide precursor is a strong complexant to the metal and may also be used as the complexant. Depositions based on thiosulphate as a S source are good examples of

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this. Also, thiosulphate depositions are most often carried out under at least somewhat acidic conditions, which means that they are more likely to involve an ionby-ion mechanism. Two old examples are the hydrolysis of silver thiosulphate to give (bulk) silver sulphide [50]: Ag2S2O3  H2O → Ag2S  H2SO4

(3.55)

and the deposition of a number of metal sulphides, including PbS and CuxS by heating thiosulphate solutions of the metal salts [51]. While less common nowadays as a S source than, e.g., thiourea, thiosulphate reactions have been used to deposit films of many different metal sulphides. The mechanism often suggested in these studies is reduction of elemental S, formed by acidic decomposition of thiosulphate, by the thiosulphate itself, forming sulphide ions (see Sec. 3.2.1.4). However, no mechanistic studies of these reactions appear to have been undertaken. While there is no convincing proof in the literature to distinguish which mechanism is operative in these cases (sulphide-ion formation or complex decomposition), chemical intuition leads us to expect the latter where the metal–chalcogen bond is strong (metal ions as Ag and Cu form very strong complexes through the labile S atom of thiosulphate), and it seems reasonable to expect this metal–sulphur bond to remain intact during the reaction rather than formation of sulphide ion to occur. At the same time, those same metals from sulphides with very-low-solubility products (understandably, since both the solubility product and strength of complexation are related in the same way to the strength of the metal–sulphur bond). Therefore, very low concentrations of free sulphide are needed to form the metal sulphide. While thiosulphate is not very commonly used to deposit sulphides nowadays, its Se counterpart, selenosulphate, is the most common reagent for selenide deposition. By analogy with thiosulphate, it might be argued that the mechanism involves formation of selenide either through hydrolysis or through reduction of Se, which forms even more readily in selenosulphate than does S in thiosulphate, or by complexation with metal ion or metal hydroxide and breaking of the SeMS bond (complex-decomposition mechanism). It can only be said that the mechanism of selenide formation using selenosulphate has not been unambiguously determined. An important difference between depositions using selenosulphate and thiosulphate is that the former are carried out in alkaline solution, in contrast to the (mainly) acidic conditions used for thiosulphate depositions. This means that both ion-by-ion and cluster mechanisms can occur using this reagent, as has been shown to be the case for CdSe, CdS, and PbSe [15,46]. Once again, it requires emphasizing that, although selenosulphate depositions are invariably assumed to occur via free selenide ions, this has not been proven, and a complex-decomposition pathway cannot be excluded in selenosulphate depositions in general. Thermodynamic analyses of metal sulphide formation from thiosulphate [52] and thiourea [53] and metal selenide formation from selenosulphate [13,54]

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have been made. These analyses are based on the assumption that free chalcogenide ions are formed.

3.4 NUCLEATION, ADHESION, AND FILM GROWTH Probably the least-known aspect of the CD process is what determines the nucleation on the substrate. Why do adherent films grow under some conditions and poorly adherent films or even no film at all under others, even when slow precipitation occurs in solution? In considering this aspect, the two basic mechanisms— hydroxide and ion-by-ion—may behave very differently, although there are also features in common. When considering nucleation, the anion-mediated mechanisms and the complex-decomposition mechanisms will behave similarly in most cases. Some basic features of nucleation will first be considered, followed by issues specific to each. Film growth can involve continuous nucleation, sticking of colloids from solution and growth of individual crystals in the film. The last in particular is considered in Chapter 10, where it is an important factor in the context of quantum size effects.

3.4.1 General Features of Nucleation and Adhesion The basic science behind nucleation and forces between materials have been treated in Chapter 1. For those interested in this section, it is assumed that this basic science is (more or less, at least) understood. However, the basics treated in Chapter 1, while important to an understanding of film (as opposed to isolated crystal) formation, are not enough by themselves to provide a phenomenological explanation of film formation. We would ideally like to be able to predict in advance, from fundamental principles, whether a particular bath formulation will result in adherent films or not. We cannot! However, if we cannot reliably predict adhesion, we can at least choose conditions so that the probability of adhesion is good. Considering first adsorption of metal ions or neutral species directly on the substrate, there are a number of possible mechanisms for this process. Most simply, there will an equilibrium between metal species in solution and a solid surface leading to dynamic adsorption of the metal. Adsorption of metal ions onto solid surfaces has been extensively studied, to a large extent because of the use of oxide surfaces to adsorb heavy metal ions and remove them from solution (see Ref. 55 for an example and list of other references on this subject). This adsorption may go even farther with ion exchange between the solution metal ions and ions in the substrate (again, glass is a good example of where this may occur). Coulombic attraction between charged species in solution and a surface may play a part in initial adsorption on the surface. Under the high pH values more com-

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mon in CD, most oxide surfaces (including glass) tend to be negatively charged because of the acid–base equilibrium of the oxide [see Eq. (2.16)]. Positive ions (e.g., Cd2, Cd(OH)) will be attracted to this surface by coulombic forces (in practice, this attraction can be reduced by the solvation shell of the metal ion). Specific chemical interactions between primary particles or even reactants and the substrate is another parameter that can aid in prediction of adhesion, at least qualitatively. Chalcogenide ions will chemisorb to many metals, in some cases forming a surface compound with the metal. Au, Ag, and Cu are the best examples. For Au in a solution of sulphide ions, the surface can be considered a gold sulphide entity; for Cu, bulk sulphidization occurs, leading to eventual disintegration of the entire Cu to CuxS. However, since chalcogenide ion requires time to be formed in a CD process, it is more likely that specific (or even nonspecific, as described earlier) adsorption of the metal ions or species is dominant in the initial adsorption process. A study of Cd2 adsorption on SiO2 from ethanol/cyclohexane mixtures followed by CdS formation by reaction with H2S has concluded that the CdS bonds to the silica via SiMOH linkages [56]. While not normally considered, it is possible that the chalcogen presursor is bonded, through the chalcogen atom, to the surface of some substrates, mainly metals. Considering now particle adsorption, the section on forces in Chapter 1 concluded by stressing that it was normal for particles to stick together (the van der Waals attractive forces eventually dominate in CD processes unless a protective surface layer is present—and sometimes even then). Yet this property of “sticking together” and also of “sticking to the substrate” clearly can vary greatly; sufficient particle adhesion to cause aggregation may not be sufficient to form an adherent film (adherent meaning that the particles stick both to the substrate and to each other—both are necessary). However, since colloids do normally aggregate, the main task in CD is to ensure that these aggregated particles (and therefore the primary ones attached to the substrate) do adhere well to the substrate. Why was the phrase and therefore the primary ones attached to the substrate just used? Why not just aggregated particles? It is probably not a very inaccurate generalization to state that adherent films of a reasonable thickness will, in most cases, not form from a solution in which aggregation has already visibly started and in which no new product is being formed before the substrate is immersed. As a clear example of this, if, during CdSe deposition from a solution in which the hydroxide mechanism is operative, the substrate is placed in the deposition solution after precipitation starts, although some deposition will occur, it is usually poor quality (poorly adherent and patchy). Clearly the initial adsorption process is important, not just for the initial deposit but for the entire film. This will not be a surprise to the many who deal with covering a surface—whether by electrodeposition, vacuum coatings such as evaporation or sputtering, or, to take even more common examples, painting or gluing, where the state of the initial surface is very important. However, while a clean substrate is as important for CD as in any coating process, there must be other factors involved. Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

The fact that primary particles will stick where aggregated ones will not can have (at least) two causes. One is that the primary particle may be different than the final aggregate (e.g., in the hydroxide mechanism, the first stage is adsorption of hydroxide particles). The other is explained by surface energy: A single particle has a larger external surface-to-volume ratio than an aggregate of the same particles, meaning a larger surface energy and therefore greater potential to stick to a surface in which it comes in contact. While it should always be remembered that adsorption from solution will certainly reduce this surface energy, perhaps drastically, the difference between single and aggregated particles remains valid. Finally, and this is likely to be important for nucleation in many cases, the effect of the electric field (Helmholz layer) at the substrate/solution interface in promoting formation of a deposit under conditions where none can form in the solution (described in Sec. 3.3.2) should be considered. Whether or not this occurs, and to what extent, can be experimentally measured.

3.4.2 Ion-by-Ion Mechanism Most nucleation studies of CD have treated either the hydroxide or hydroxidecomplex mechanisms (see later) or have not clearly defined which mechanism was, indeed, operative under the conditions of the experiments. Due to the paucity of dependable experimental data, therefore, we consider nucleation and growth by the ion-by-ion mechanism, to a large extent, from a theoretical viewpoint. Figure 2.4 showed a general-form curve of film formation as a function of time. This form is valid in many cases regardless of the mechanism. For the ionby-ion mechanism, an induction period is generally necessary for sufficient chalcogenide ion to build up and form a solid metal chalcogenide phase. It is probable that some metal ions adsorb on the substrate, e.g., by an ion exchange, an electrostatic mechanism, or simple equilibrium (see Sec. 3.4.1). However, while this stage may be important for film initiation, it is not normally considered as a growth stage in the usual sense of the word. Growth can be considered to begin when stable clusters of the deposit begin to form on the substrate. Ideally, deposition occurs only on the substrate and not in solution. This is possible due to the effect of a surface (even one in which no chemical interaction occurs with any constituents of the deposition solution—and such interaction probably does occur to a greater or lesser extent; water, for example, interacts with many different surfaces). It is easier to nucleate on a surface than from a homogeneous solution. The possible effect of the electric field at the substrate/solution interface on promoting nucleation has already been described. Additionally, some surfaces are easier to nucleate on than others. This is the basis of the sensitization of some surfaces, usually glass by a SnCl2 treatment. The SnCl2 hydrolyzes to give nuclei of hydrated tin oxide on the glass surface, and these nuclei then form nucleation centers for growth of the CD film. Such sensitization may reduce (even to zero) the initial induction (nucleation) time of deposition. An example of this is Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

the deposition of CdS on a SnO2/glass (conducting glass) surface that has already been covered with a layer of CdS. Essentially no induction time was found for film growth, compared with a few minutes induction time for deposition, from the same solution and under otherwise identical conditions, on bare conducting glass [57]. This was explained by catalytic decomposition of the thiourea by the predeposited CdS, and hence immediate film growth, compared to the slower decomposition of thiourea in the absence of this catalytic surface. Once nucleation of clusters has begun, growth can occur, since in most cases the depositing material will deposit on itself more readily than on a substrate (since the reactants, almost by definition, chemically bond to the product). Thus CdS will chemisorb sulphide and/or cadmium ions, depending on the absorption properties of the CdS, in particular the crystal face involved. The crystal size of deposits formed via the ion-by-ion process tends to be larger than those of the same material deposited via the hydroxide mechanism (see Sec. 3.4.3). This can be explained by slower nucleation, resulting in fewer nucleii; growth by homogeneous formation of chalcogenide ion tends to favor slower nucleation for the ionby-ion process. These nuclei therefore have more lateral room to grow in the plane of the substrate.* In principle, large crystals (of the order of microns) could be expected in this case. While true in some cases, crystal size was usually much smaller than this. Growth termination by adsorption of various species that are not involved in the growth process or by defect formation are possible reasons for this. One possible measure of ion-by-ion growth is that, in contrast to a cluster mechanism, deposition of fresh material will generally be preferred on already-existing deposit, as already noted. This means that the number of nucleii may not increase greatly (after the very early stages of film formation) and that the film will grow by growth of these initial crystals. This has been reported, for example, in the case of Ag2S deposition from a thiourea/ammonia bath [58]. While the mechanism of this deposition is not certain, the combination of this growth pattern together with the strong complexation of the Ag by thiourea (strong AgMS bond in the thiourea complex) suggests an ion-by-ion complex-decomposition mechanism. Growth of various semiconductors onto certain single-crystal substrates has resulted in epitaxial growth in a number of cases. This epitaxy has been well studied for CdS deposition by Lincot et al. [59–63]. Although the epitaxy requires a certain degree of lattice matching between semiconductor and substrate, chemical interactions between the constituents of the deposition solution and the substrate are important as well (discussed in more detail in Chap. 4). It is a reasonable assumption that epitaxial deposition occurs via an ion-by-ion process. Indeed, it has

* It is a characteristic of the ion-by-ion mechanism for CdSe deposition from NTA solutions that thin films are highly scattering, and this scattering decreases as film thickness increases. It is likely that this scattering is due to voids in the thin films, which are a result of the low density of initial nuclei.

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been observed that epitaxy ceases when CdS colloids begin to appear in the solution, allowing a cluster growth to occur [63]. Termination of growth due to depletion of reactants will eventually occur. The reactants do not have to be completely used up. The concentration of free metal ions will decrease not only due to decrease in total metal concentration, but also because the ratio of the metal ion to complexant (the concentration of the latter usually remains constant throughout the reaction unless ammonia is used in an open vessel) also decreases. Additionally, as the chalcogen source is used up, the chalcogenide-forming (or chalcogen complexation) reaction slows down. Thus, at some point, the rate of deposition will slow down to an unacceptable degree, even though there may be an appreciable fraction of the reactants remaining in the deposition solution. For the ion-by-ion growth, as the deposition slows and there is more time for rearrangement of newly formed material to its most stable configuration, the likelihood of larger and better-formed crystals may increase, as shown in an early study of PbSe deposition [13].

3.4.3 Cluster Mechanism The general shape of the growth of CD films as a function of time is often similar for the cluster mechanism as for the ion-by-ion mechanism (Fig. 2.4). Figure 3.4 shows an actual example of CdSe deposition from a solution (containing nitrilo-

FIG. 3.4 Time dependence of CdSe (nitrilotriacetate bath, hydroxide mechanism, room temperature) film growth measured by quartz crystal microbalance.

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triacetate as complex and under conditions where Cd(OH)2 is present as a colloid) measured by a quartz crystal balance. This deposition was not continued to the inevitable termination point, but the initial stages of the induction period followed by essentially linear growth are clear. Since there is a colloidal phase present in the solution from the very beginning, why is there an induction period at all in this growth? Should these colloids not stick to the substrate and build up a film? The answer, at least in this specific case (and probably in many others), is yes and no; yes, the colloids will stick to the substrate, and no, they need not build up to form a film. There are a number of studies that show that Cd(OH)2 forms on the substrate from the start of the process. An early radiochemical study of CdS deposition showed the initial presence of Cd, free of S, on glass immersed in the (ammoniacal) deposition solution both with and without thiourea present [39]. The amount of Cd (considered to be Cd(OH)2) was constant with time (after an initial short time) until S began to be detected in the deposit, at which point it started to increase. Cd(OH)2 was found on the substrate, using XPS analyses, for CdSe deposition on glass [15] and for CdS deposition on Si using an ammonia bath [37], only using deposition solutions that did not contain the chalcogen presursor. The Cd(OH)2 only formed on the substrate at a high-enough pH (typically 9 and above, although this value will depend greatly on other solution parameters, in particular complex:Cd ratio and temperature). In both cases, the coverage of the substrates by the Cd(OH)2 was usually sparse, as evidenced by the predominant substrate (Si) signals, and never increased to a level where the Si was not predominant. TEM studies of the Cd(OH)2 formation in the CdSe deposition, where selenosulphate was present, confirmed this sparse coverage by Cd(OH)2[15]; CdSe formed only after some time. Similar results were obtained for PbSe deposition; immersion of Au-coated glass in an alkaline solution of (selenosulphate-free) citrate-complexed PbAc2 resulted in adhesion of hydrated lead oxide (there is no stable simple lead hydroxide) to the Au, which was not washed off by water (Fig. 3.5), although in this case, PbSe formed more rapidly than did CdSe when selenosulphate was present [46]. All these results show that Cd(OH)2 colloids do adsorb on a substrate (either under conditions where Cd(OH)2 is present in solution or, according to the studies of Rieke and Bentjen and Ortega-Borges and Lincot [48], even when it is not present in solution but under solution conditions close to solid hydroxide formation). The induction period when “no” deposition is seen in the hydroxide-cluster deposition therefore is understood to mean that a fast and nongrowing Cd(OH)2 adsorption has occurred, which is too fast and/or too little to measure by the experimental methods used to make the kinetic curves, and that only when the hydroxide starts to convert into the chalcogenide, by reaction of the slowly formed chalcogenide ion with the hydroxide, does real film formation proceed. An obvious question at this point is: Why does CdS (CdSe) grow, i.e., continue to deposit and form a thick film, but Cd(OH)2 does not? Rieke and Bentjen

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FIG. 3.5 TEM image of Au film-on-glass immersed for 3 min in a solution of PbAc2 (60 mM) and trisodium citrate (160 mM) at pH 10.8. Electron diffraction showed the black crystals to be hydrated lead oxide.

discussed the adsorption of various Cd and Cd-OH species onto the Si substrate. The Si will be covered with an oxide (therefore it can be treated similarly to quartz and even approximated to glass for the present purposes), and this oxide will be negatively charged at the pH values involved. Adsorption of positive ions from solution (Cd2, Cd(OH)), would be favored, and would eventually neutralize this charge, after which further adsorption of positively charged species would no longer occur. This scenario, however, was considered to conflict with the observation that greater amounts of Cd(OH)2 were found on the (Si) substrate at higher pH values, where the concentrations of the positive species would be greatly reduced. Adsorption of neutral Cd(OH)2 was considered more likely. This adsorption could occur in different ways. Direct adsorption of Cd(OH)2 colloids from solution was one possibility. However, since Cd(OH)2 could form, at least to some extent, on the substrate from solutions where no Cd(OH)2 was present (see earlier), surface-catalyzed adsorption was considered, such as [48].  Cd(NH3) 2 4  2OH  surface site D [Cd(OH)2]ads  4NH3

(3.56)

where the OH may or may not be that originally bound at the surface, or, in the hydroxide–chalcogenide complex-decomposition mechanism, by adsorption of the ammine-hydroxide species in Eq. (3.52). The various observations that, at relatively low pH, the coverage of the substrate by Cd(OH)2 was poor indicate the dynamic equilibrium between adsorbed Cd(OH)2 and the solution; the Cd(OH)2 was in a continual state of dissolution and deposition. Since the concentrations of both free Cd2 and OH was constant (in

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the absence of large-scale depletion of the reactants, the case when no chalcogenide precursor was added), this dynamic equilibrium resulted in a steady (and low) concentration of Cd(OH)2 on the substrate. When chalcogenide ion begins to form and converts the hydroxide to the chalcogenide (the simple hydroxide mechanism), or when the chalcogenide precursor forms a complex with the surface-adsorbed species followed by complex dissociation (the complex-decomposition mechanism), this disturbs the equilibrium, allowing more surface-adsorbed hydroxide to form, resulting in film growth. [Note that films of Cd(OH)2 (and other hydroxides) can be grown by CD under conditions where generation of OH occurs; i.e., the system is not in equilibrium]. This is one likely scenario for the linear part of the film growth. There are other possibilities, and it is likely that more than one mechanism of growth occurs. These other possibilities include direct adsorption of colloids from the solution and parallel ion-by-ion growth on the primary deposit. Direct adsorption of colloids can certainly occur. However, this mechanism is often associated with poorly adherent and optically scattering films. It occurs with these properties if the substrate is placed so that sedimentation of colloids from the solution occurs directly onto the substrate. For this reason, the substrate should be placed either vertical in the solution or, if non vertical, with only the lower side of the deposit (which forms on both sides of the substrate) retained. In principle, direct adhesion of colloids could result in adherent films, but it is more likely that this involves isolated colloids or small aggregates that have greater contact area than a large aggregate. Greater contact area here means that 200 (for example) colloids that have adsorbed onto the surface one by one forming a single aggregate would have a larger contact area to that surface than the same 200 colloid aggregate that adsorbed, as an aggregate, in a single step. There is also the question of whether a Cd(OH)2 colloid would adhere better to a CdS surface than to itself (for the case where hydroxide is present in solution). Since the hydroxide is in a dynamic equilibrium, nonaggregated particles (or particles containing only a small number of crystals) will be more likely to be present for the hydroxide than for the chalcogenide (where the equivalent equilibrium is likely only for very tiny nuclei). Parallel ion-by-ion growth might occur on previously deposited (by the hydroxide mechanism) film. However, at least in the simple hydroxide mechanism, where a solid hydroxide is present in solution, this is not likely to be a major factor in the growth, except under solution conditions close to the transition between hydroxide and ion-by-ion growth. The reason for this is that the chalcogenide is being formed homogeneously throughout the solution. In the case where hydroxide is present in solution, most of this chalcogenide will react with colloidal metal hydroxide in solution (as seen in most cases by precipitated metal chalcogenide), and the concentration at the substrate will be very low—enough to convert hydroxide to chalcogenide perhaps, but less likely to form a new phase.

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For the case where no hydroxide is present in solution but is formed on the substrate while the homogeneously formed chalcogenide will not react with hydroxide in solution, some or even much of it may still not reach the substrate. This is due to the relative lack of stability of the chalcogenide ions, which could be oxidized homogeneously, e.g., by dissolved oxygen. This will be more important for selenide than for the considerably more stable sulphide; but even sulphide is not stable in very low concentrations unless oxygen is rigorously excluded. This potential problem will not exist for the complex-decomposition mechanism, where there is a high concentration of the chalcogenide complex. The crystal size of the deposit obtained by the simple hydroxide mechanism will be essentially that of the Cd(OH)2 clusters, which are converted into, e.g., CdS. Such colloids tend to be very small (typically in the region of 5 nm), and therefore the CdS crystal size should be similar. It may be larger if some form of ion-by-ion growth occurs on these primary crystals. Also, it may be different if the hydroxide is present on the substrate but not in solution, since the size of the hydroxide deposit on the substrate will, at least in part, be affected by different factors than that formed homogeneously in the solution. In general, crystal size in the simple hydroxide mechanism (hydroxide colloid present in solution) is smaller than that formed in an ion-by-ion process. This has been shown for CdS, CdSe, and PbSe, where typical crystal sizes for the hydroxide (ion-by-ion) mechanisms were found to be (in nm): CdS—5 (70); CdSe—5 (15); PbSe—5 (10–1000) [15]. In contrast to the ion-byion mechanism, the crystal size of films formed via a cluster mechanism is not expected to grow greatly (some modest growth can and usually does occur) with increasing film thickness, since film growth occurs by sequential addition of new clusters. Thus, in our observation of CdSe growth via the hydroxide mechanism, the color of the film (a sensitive measure of crystal size due to size quantization; see Chap. 10) changes only a little during the deposition (color refers to spectral position and not to depth of color, which, of course, does increase with film thickness). SEM micrographs in the study of Rieke and Bentjen [37] showed that, although the number density of CdS nucleii/unit area on the substrate was constant with time after deposition started, the size of the nucleii increased linearly with time. Additionally, the size distribution of the nucleii, both in the early and later stages of growth, was quite narrow. These film growth kinetics were identified with burst-type nucleation, well known in homogeneous solution precipitation, where a homogeneous reaction in solution causes sudden nucleation whenever a critical concentration of one of the reactants is reached. This nucleation reduces the concentration of this reactant so that further growth occurs only on existing nucleii (nucleation usually requires a large supersaturation, while growth on an existing nucleus does not). This type of nucleation usually results in a narrow size distribution, as seen here. In a quartz crystal microbalance investigation of CdSe film growth rate from a selenosulphate/ammonia/triethanolamine bath with different Cd:selenosulphate

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ratios, two peaks were observed in the growth rate vs. time plot of all solutions [64]. This was explained by a two-stage growth. From electron microscopic examination of the growing films, the first stage was attributed to instantaneous nucleation and 2-D lateral growth to cover the substrate, while the second stage was due to 3-D nucleation and growth at random sites on the first layer. Aggregation of colloidal particles was invoked as the mechanistic pathway. After all the speculation involved in the foregoing discussion of film growth, the termination step is as simple to explain as for the ion-by-ion mechanism. Growth can occur as long as the concentration of chalcogenide anion is high enough to allow Reaction (3.5), the conversion of the hydroxide to the chalcogenide, to occur [or, for the complex-decomposition mechanism, sufficient chalcogenide precursor as, e.g., in Eq. (3.7)]. It is also possible that depletion of the metal hydroxide occurs first (in which case the mechanism may change to ionby-ion, as described earlier). In either case, termination is simply due to depletion of the reactants. Typically in the cluster mechanism, most of the reactants are lost in homogeneous precipitation.

3.4.4 Complex-Decomposition Mechanism The initial nucleation stage of the complex-decomposition mechanism is probably similar to the simple free-anion mechanism. Either ionic or molecular metal species (ion-by-ion) or Cd(OH)2 (cluster) adsorbs on the substrate. However, instead of conversion of the hydroxide to sulphide by topotactic reaction with sulphide ions, the chalcogenide precursor (in almost all studies of this mechanism, that is thiourea) adsorbs on the Cd(OH)2 surface to form a hydroxide–thiourea complex, which then decomposes to CdS. A possible difference between the simple hydroxide mechanism and the complex decomposition is in the manner of crystal growth. We noted earlier that crystal growth in the simple hydroxide mechanism may occur via an ion-by-ion process but probably not to a large degree, since the concentration of chalcogenide ion would be very low indeed (it would react rapidly with Cd(OH)2 in the solution). However, for the hydroxide-decomposition mechanism, the chalcogenide reactant is not the free chalcogenide ion but the precursor, which is present in much higher concentrations. Therefore once a solid phase capable of catalyzing the chalcogenide precursor has formed, the crystal growth is quite likely to switch over to a predominantly ion-by-ion process, such as in Reaction (3.51). This means that, as for the pure ion-by-ion process, the crystal size might be expected to be larger than for the pure hydroxide mechanism, since the ion-by-ion growth favors, in principle, crystal growth rather than renucleation. Typical crystal sizes for CdS prepared from the ammonia/thiourea bath, which appears under many experimental conditions to proceed via the hydroxide-complex mechanism, are in the region of 10 nm to several tens of nanometers, larger than the approximately 5 nm obtained for CdS and CdSe from the simple hydroxide-cluster mechanism. Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

3.5 KINETICS OF DEPOSITION Kinetic studies have been a popular topic in CD, particularly for CdS. This would suggest that the present section will be a large one to reflect this activity. In fact, the reverse will be the case. This will be a relatively short section that will not try to cover even a moderately large part of the kinetic studies. The reason is that kinetic measurements have been used, to a large extent, to study the mechanisms of deposition, and this has been dealt with already (not the details of the kinetic studies, but the conclusions). Additionally, since a CD process can vary widely in rate—from a few minutes to days and weeks—often depending strongly on small changes in one or another concentration of a particular reactant, the important information to be learned from kinetic studies (apart from mechanistic diagnosis) is how this rate depends on experimental conditions, and this can be done with a few selected examples from the literature. Ortega-Borges and Lincot [48] carried out a detailed kinetic study of CdS deposition from the standard ammonia (ammonium)/thiourea bath using a quartz crystal microbalance to measure film thickness. They measured a deposition rate with fractional values of reaction order [Cd]0.6[Tu]0.8 rate
K  [NH3]3.3[H]1.5

(3.57)

They therefore concluded that several different rate-determining steps were involved in the deposition. Figure 3.6 shows the dependence of the deposition rate on the concentration of the reactants (Cd, thiourea, ammonia, and pH—the last varied through introduction of ammonium ion) (a) as well as an Arrhenius plot of the deposition (b) for the CdS deposition. From the kinetic data, they deduced the hydroxide-complex-decomposition mechanism, given earlier in Eqs. (3.50) and (3.51) and, more specifically, as  Cd(NH3) 2 4  2OH  surface site D Cd(OH)2(ads)  4NH3

(3.58)

This first step represents a reversible adsorption of Cd(OH)2 on the substrate. The reaction order of 1.5 for hydroxide [given by H in Eq. (3.57)] implied the participation of two hydroxide ions in the process. The next step was formation of a complex between the adsorbed Cd(OH)2 and thiourea: Cd(OH)2(ads)  (NH2)2CS → Cd(OH)2SC(NH2)2(ads)

(3.59)

This complex then decomposes to CdS: Cd(OH)2SC(NH2)2(ads) → CdS  CN2H2  2H2O

(3.60)

probably by nucleophilic attack of the thiourea S atom on the Cd(OH)2. The activation energy of the deposition (85 KJ/mole), measured from the Arrhenius plot in Figure 3.6b, is very similar to that of thiourea decomposition, suggesting the slow nature of Reaction (3.60). Earlier measurements of PbSe deposition from a Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

FIG. 3.6 (a) Log–log plots of reactant concentration (X) vs. deposition rate for CdS deposited from an ammonia/thiourea bath. Standard conditions: [Cd]
14 mM; [Tu] (thiourea)
28 mM; [NH3]
1.74 M; T
60°C. Reactants are Cd (total concentration); Tu; NH3; pH (adjusted by adding ammonium ion) that gives hydroxide concentration. (b) Temperature dependence of deposition rate. (Adapted from Ref. 48).

selenourea bath gave an activation energy of 60 KJ/mole [65], with the lower value compared to CdS deposition presumably reflecting the greater instability of selenourea compared to thiourea. Using the same basic system and similar experiments, Doña and Herrero measured reaction orders for the various species comparable to those measured in the study of Ortega-Borges and Lincot, except for that of ammonia, which was 1.8 instead of 3.3 [49]: [Cd]0.9[Tu]1.1[OH]1.7 rate
K  [NH3]1.8

(3.61)

This, together with the known tendency of metal ions to form mixed hydroxy–ammine complexes, suggested to them that two ammonia molecules were involved in the first step and that the adsorbed species in Reaction (3.58) was a hydroxy–ammine species, viz. Cd(OH)2(NH3)2. Decomposition of the hydroxide– ammine–thiourea complex was then assumed to occur by nucleophilic attack of an ammonia species on the SBC bond of the thiourea. Based on this mechanism, a detailed theoretical model has recently been proposed for CdS deposition from the thiourea/ammonia bath [65a]. Prediction of different aspects of the deposition kinetics using this model provided a very good fit with the relevant experimental data. Rieke and Bentjen [37] studied the kinetics of CdS deposition from an ammoniacal thiourea bath using SEM. As discussed earlier, they found that goodCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

quality films formed in a pH range where Cd(OH)2 formed on the substrate (Si, in their experiments) but not in the bulk of the solution. Their kinetic study was made in this pH range (specifically, pH
9.55). No CdS deposited initially, but the rate of formation of CdS increased with time, eventually becoming more or less constant over the time of their experiment. This is characteristic of an autocatalytic reaction, where the initial deposit accelerates the rate of further deposition. O’Brien and Saeed, using ethylenediamine as compexant, higher deposition temperatures, and glass as a substrate, found that the thickness of the CdS film increased linearly with time (after an initial induction period) and also that there was no increase in the size of the nucleii (both in contrast to the previous study) [40]. In spite of the different experimental conditions, the mechanism of the depositions in both studies appears to be essentially the same, i.e., hydroxide-mediated catalysis of thiourea decomposition. It must be kept in mind that the kinetics of CD, as with the deposition mechanism, can be very different from one system to another. Two connected examples of this are given here. In one study of PbSe deposition from a citrate-complexed selenourea solution containing hydrazine, the rate was proportional to the pH and to the selenourea concentrations but independent of the Pb and citrate concentrations [65]. This was explained by a rate-determining step involving decomposition of selenourea at the (catalytic) PbSe surface by hydroxide. It is noteworthy that the Pb concentration was typically an order of magnitude less than that of selenourea. Therefore the independence of the rate on Pb (or citrate, which determines the concentration of free Pb2) concentration, suggests that formation of selenide ion, and not a complex-decomposition mechanism, occurs. The second example is seen in the study of PbSe deposition by Kainthla et al. from selenosulphate solution [41]. In most examples of CD from alkaline solution, the deposition rate increases with increase in pH. This is due to both the greater rate of decomposition of the chalcogenide precursor at higher pH (this decomposition usually involves hydroxide ions) and, in many cases, the greater probability of solid hydroxide formation (as long as this is not excessive). However, for PbSe deposition using citrate as complex for the Pb and selenosulphate as Se precursor, the opposite occurs: The deposition rate decreases with increase in pH. This is due to the specific hydroxy–citrate complex formed: 2 Pb(OH)C6H5O 2  OH  C6H5O 3 7 D Pb 7 

(3.62)

Increase in pH (
increase in [OH ]) shifts the equilibrium to the left, resulting in a lower concentration of free Pb2 ions and thus a slower reaction to give PbSe. This means that, in contrast to the deposition from a selenourea bath described earlier, the rate is dependent on Pb concentration and possibly independent of hydroxide concentration at a constant free-Pb2 concentration. This would then suggest that the opposite mechanism, i.e., a complex decomposition, is effective for the selenosulphate bath. It is stressed that these conclusions on selenide formation Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

or complex-decomposition mechanisms are indications of which mechanism is taking place, but are far from being firm proof of this. For the deposition of ZnS and ZnSe, hydrazine is normally used to form films at a reasonable rate. The role of the hydrazine is not obvious. It is tempting to assume that hydrazine, being a strong reductant, reduces the chalcogen precursor to chalcogenide ion, as was assumed for CdTe deposition (see Sec. 3.2.3). However, this appears to be oversimplified. In a study of the effect of various amines (including hydrazine) on the deposition rate (and composition) of ZnS films deposited from ammonia/thiourea baths, Mokili et al. found a strong dependence of the rate on the type of amine added [66], as shown in Figure. 3.7. While it is difficult to separate the effect of concentration from the different types of amines in this experiment, it is clear that an increase in rate is general on addition of amine (apart from the initial induction time using triethanolamine). Since the amines also act as complexation agents, they would, on this basis, be expected to reduce the deposition rate (by reducing the free Cd2 concentration). The fact that the opposite occurs implies that they must increase the thiourea (or selenosulphate for selenides) decomposition. In this respect, all the amines used have pronounced reducing properties, with redox potentials of 0.46 (triethanolamine), 0.56 (ethanolamine), and 1.16 (hydrazine), which parallels the order of increase in deposition rate seen in Figure 3.7. O’Brien et al. discuss various theories for

FIG. 3.7 Evolution of ZnS film thickness with time. ZnCl2/NH4Cl/thiourea/NH3 bath at 85°C. The effect of various amines on growth. Triethanolamine—0.2 M; ethanolamine— 0.7 M; hydrazine—3 M. (Adapted from Ref. 66).

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the accelerating effect of hydrazine on the CD process [67]. While agreeing that the hydrazine increases thiourea decomposition, the specific details of the effect of hydrazine (and other amines) on the CD process are still not fully understood. The activation energies of the deposition for both ZnS [68] and ZnSe [69], measured from Arrhenius plots, are 21 and 26 KJ/mole, respectively, much smaller than the values for CdS (85 KJ/mole) or PbSe (60 KJ/mole) described earlier. Stirring does not affect the deposition rate for either ZnS or ZnSe, so the deposition is not under diffusion control. In interpreting activation energies for CD processes, it is important to remember that what is measured is the film growth, and this is not necessarily the same as the rate of formation of the metal chalcogenide, much of which is usually formed homogeneously in the solution. The Zn compounds were both probably formed by a cluster mechanism, in contrast to the ion-by-ion complex-decomposition mechanism probable for the CdS. The “activation energy of deposition” for the ZnS(Se) therefore depends to a large extent on the rate of sticking of clusters, although other factors could also be involved, such as a parallel ion-by-ion (whether by complex decomposition or free chalcogenide ions). To interpret such results correctly will require a study of activation energies of different compounds deposited under different and controlled mechanistic pathways and preferably also measuring the total amount of product formed (in solution as well as on the substrate and other surfaces). It is also relevant that the crystal size for ZnS and ZnSe deposits is typically smaller than for CdS or CdSe deposited from an ion-by-ion bath; this supports a cluster mechanism for these depositions. In fact, in contrast to CdS and CdSe, there are no cases in the literature where ZnS or ZnSe have been clearly shown to have been deposited via an ion-by-ion mechanism.

3.6 DEPOSITION FROM ACIDIC BATHS While the majority of CD reactions have been carried out in alkaline baths, there have been a number of sulphide depositions reported in acidic baths. These are all (with one exception, Sn-S from acidic nonaqueous S baths [70]; see Chap. 6) based on two sulphide precursors: thiosulphate (more commonly used in the very early studies, but still sometimes used) and thioacetamide. Beutel and Kutzelnigg [71] described a large selection of colored films deposited on various metals using CD (and also electrodeposition) from metal salt–thiosulphate solutions. No characterization of these films (other than interference colors) was made. However, it is clear that CD of metal sulphides did occur in some of these cases (see Chap. 6). Lokhande described a variety of sulphides (CdS, ZnS, Bi2S3, Sb2S3, As2S3, Cu2S, Ag2S, SnS2, and PdS2) deposited using thiosulphate at a pH of typically ca. 3 [72]. Other studies on deposition from acidic thiosulphate baths are: PbS [73], Ag2S [74], Bi2S3 [75], CuxS [76,77], Sb2S3 [78], SnS2 [79].

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The mechanism of the thiosulphate reaction is not clear. Lokhande, in his studies, has suggested an internal reduction involving the reaction  2 2S2O 2 3 → S4O 6  2e

(3.63)

These electrons reduce elemental S formed in Reaction (3.21) or (3.22) to give sulphide (or hydrosulphide) ions, as in Reaction (3.24), which react with the metal ions. The general mechanism of sulphide generation has been assumed in most studies using thiosulphate, e.g., Nair et al [78] for Sb2S3 deposition and Grozdanov et al. in their study of CuxS deposition from Cu-thiosulphate solutions [76], although in this latter study it is noted that the mechanism may be more complicated in this Cu-S system. As has already been pointed out, in spite of the fact that the free sulphide mechanism is invariably assumed in thiosulphate depositions, there is no evidence up to now against the complex-decomposition mechanism. No thorough mechanistic or kinetic studies have been made on this system. Since the studies on CD using thiosulphate have not attempted to differentiate between these different mechanisms (and since such differentiation may be difficult, this is not surprising), we are left with the conclusion that there is no clear consensus on which mechanism is operative. Also, the mechanism may vary depending on conditions (as for the alkaline baths), and of course a combination of mechanisms may be operative in some cases. Other points to note when considering thiosulphate as a reagent in CD is that thiosulphate is a strong complex for a number of metals and, since it is a fair reducing agent, also may reduce the metal ions (as is known to occur for Cu2 to Cu). Many sulphides have been deposited using thioacetamide in acidic solutions (Chapter 6 describes most of these). For depositions using thioacetamide, as with thiosulphate, there are no detailed mechanistic studies. Both H2S formation and complex decomposition are possible in acid solutions, as discussed in Section 3.2.1.3. Deposition of CdS was accomplished using thioacetamide in acidic solution by exploiting electrolytic proton reduction to increase the pH locally at the cathode (substrate), and the mechanism was believed to be a surface-catalyzed decomposition of a Cd–thioacetamide complex [80]. Because of the acidic conditions, with the exception of very acidic cations, nucleation resulting from a solid colloidal phase is unlikely. For a deposition that occurs through free (hydro)sulphide ion, it is probable that nucleation occurs when the concentration of this sulphide ion is high enough to cause precipitation of the metal sulphide, possibly catalyzed by the substrate surface. Similarly for a complex-decomposition pathway, a high-enough concentration of the final product to permit solid-phase formation will be required. It should be remembered that the concentration of free metal ion will, in most cases, be higher than for alkaline baths, due to the (usually) weaker complexing strengths of the sulphide precur-

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sors. This will allow formation of a solid phase at lower concentrations of sulphide than in more heavily complexed alkaline solutions. There appear to be no cases of selenide deposition from acidic baths. Selenosulphate is not stable under even mildly acidic conditions, and all selenoureabased baths have been alkaline ones. Oxides or hydroxides have been deposited from acid baths, in particular readily hydrolyzable acidic metal ions. These have been discussed in Section 3.2.4.

3.7 EFFECT OF STIRRING There have been only a few studies on the effects of stirring the deposition solution on the deposited film. Overall, stirring affects CD films mainly by preventing deposition of loosely adhering, large aggregates. These loose deposits are readily removed by the stirring action. This is important, since they block the substrate, preventing normal adherent film growth. Such nonadherent deposits can also be prevented without stirring by placing the substrate in the bath at an angle; the deposit on the upper surface, which will usually be a mixture of adherent and loosely adherent material, can be removed (by wiping with a reagent that dissolves the film, often dilute HCl), leaving the film on the lower surface, which does not collect precipitated deposit. Such loosely adhering CdS films in nonstirred solutions have been reported by Kaur et al. [34] and by Doña and Herrero [49]. The latter and also OrtegaBorges and Lincot [48] found that the rate of deposition is affected by stirring only at low stirring rates, and the effect is not large. There is no apparent difference between low and fast stirring rates. This implies that even slow stirring is enough to prevent sticking of large, loosely adhering particles. Apart from deposition rate, the study by Kaur et al. found that stirring could, in some cases, strongly affect the film quality and that the effect of stirring was dependent on the concentration of ammonia (relative to the Cd). For low ammonia concentrations, where a visible Cd(OH)2 phase existed, and for high ammonia concentrations, where it is probable that no solid hydroxide phase occurred, strongly stirred solutions resulted in formation of CdS films with predominantly wurtzite structure, while films deposited from unstirred solutions contained a large amount of zincblende CdS. Precipitate formed homogeneously in the solution was found to be zincblende, and the effect of stirring was to minimize sticking of colloidal particles from solution, thereby reducing the zincblende component in the films (and also resulting in more adherent and specularly reflecting films—adsorption of large aggregated colloids from solution caused light scattering and reduction in specular reflectance). For the intermediate case where there was just enough ammonia to dissolve the visible Cd(OH)2 (but a colloidal Cd(OH)2 probably was present in solution), stirring did not have any major effect—thick, powdery, and zincblende

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films were obtained in both cases. This is unexpected, since these are conditions where good films are often obtained. There is one unusual case where stirring had a very strong effect on the deposition rate: Ag2S deposited from a thiourea bath [81]. The deposition was very slow in the unstirred solution and increased, more rapidly at first, then linearly, up to the maximum stirring speed of ca. 1100 rpm. The energy of activation of this deposition was 20.4 kJ/mole, much less than values typical of reaction-controlled depositions under similar conditions involving, e.g., CdS or PbS, and similar to those found for ZnS and ZnSe. However, the Zn-S(Se) depositions were independent of stirring. It therefore appears that the mechanism of this Ag2S deposition is diffusion controlled and is unlike other mechanisms discussed previously. As pointed out by Ortega-Borges and Lincot, the relative independence of the majority of CD processes on hydrodynamic conditions explains the excellent lateral homogeneity characteristic of this technique, since the deposition depends not so much on mass transport in the solution as on chemical reaction rates.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

WHR Shaw, DG Walker. J. Am. Chem. Soc. 78:5769, 1956. DF Bowersox, DM Smith, EH Swift. Talanta 3:282, 1960. OM Peeters, CJ de Ranter. J. Chem. Soc. Perkin II 1832, 1974. R Williams, PN Yocom, FS Stofko. J. Colloid Interface Sci. 106:388, 1985. WH Perkins, AT King. J. Chem. Soc. 103:301, 1913. WN Arnquist. Proc. IRE 47:1420, 1959. CJ Milner, BN Watts. Nature 163:322, 1949. GA Kitaev, TP Sokolova. Russ. J. Inorg. Chem. 15:167, 1970. GA Kitaev, VM Makova. USSR Patent 377,445:C.A. 79:84806j, 1973. RA Zingaro, FC Bennett Jr, GW Hammar. J. Org. Chem. 18:292, 1953. RA Zingaro, DO Skovlin. J. Electrochem. Soc. 111:42, 1964. AA Velykanov, EK Ostrovskaya, NP Garina, VA Turacova, AA Tchurkan. Ukr. Chim. Zh. 49:764, 1983. GM Fofanov, GA Kitaev. Russ. J. Inorg. Chem. 14:322, 1969. GA Kitaev, TS Terekhova. Russ. J. Inorg. Chem. 15:25, 1970. S Gorer, G Hodes. J. Phys. Chem. 98:5338, 1994. GA Kitaev, GM Fofanov. Zh. Prikl. Khim. 43:1694, 1970. GK Padam, SK Gupta. Appl. Phys. Lett. 53:865, 1988. RW Buckley. In: 11th ECPV Solar Energy Conf: Montreux, Switzerland, 1992, p 962. AV Kalyakina, RI Pelyukpashidi. Tr. Khim. Met. Inst. Akad. Nauk. Kaz. SSR 17:114, 1973. VB Patil, PD More, DS Sutrave, GS Shahane, RN Mulik, LP Deshmukh. Mater. Chem. Phys. 65:282, 2000. M Sotelo-Lerma, RA Zingaro, SJ Castillo. J. Organomet. Chem. 623:81, 2001.

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22. 23. 24. 25. 25a. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61.

L Gordon. Anal. Chem. 24:459, 1952. HH Willard, L Gordon. Anal. Chem. 25:170, 1953. M Izaki, T Omi. J. Electrochem. Soc. 144:L3, 1997. M Izaki, O Shinoura. Electrochem. Solid State Lett. 1:215, 1998. M Izaki. O Shinoura. Adv. Mater. 13:142, 2001. W Mindt. J. Electrochem. Soc. 117:615, 1970. TP Niesen, MR DeGuire. J. Electroceram. 6:169, 2001 DH Klein, L Gordon, TH Walnut. Talanta 3:177, 1959. L Gordon, JI Peterson, BP Burtt. Anal. Chem. 27:1770, 1955. VK LaMer, RH Dinegar. J. Am. Chem. Soc. 73:380, 1951. L Gordon, ML Salutsky, HH Willard. Precipitation from Homogeneous Solutions. New York: Wiley, 1959. JR Gump, GR Sherwood. Anal. Chem. 22:496, 1950. PFS Cartwright. Analyst 92:319, 1967. I Kaur, DK Pandya, KL Chopra. J. Electrochem. Soc. 127:943, 1980. ML Breen, JT Woodward, DK Schwartz, AW Apblett. Chem. Mat. 10:710, 1998. GA Kitaev, AA Uritskaya, SG Mokrushin. Russ. J. Phys. Chem. 39:1101, 1965. PC Rieke, SB Bentjen. Chem. Mat. 5:43, 1993. RO James, TW Healy. J. Colloid Interface Sci. 40:53, 1972. ND Betenekov, VP Medvedev, AS Zhukovskaya, GA Kitaev. Sov. Radiochem. 20:524, 1979. P O’Brien, T Saeed. J. Crystal Growth 158:497, 1996. RC Kainthla, DK Pandya, KL Chopra. J. Electrochem. Soc. 127:277, 1980. H Cachet, M Froment, G Maurin. J. Electroanal. Chem. 406:239, 1996. P O’Brien, J McAleese. J. Mater. Chem. 8:2309, 1998. D Lincot, RO Borges. J. Electrochem. Soc. 139:1880, 1992. S Gorer, A Albu-Yaron, G Hodes. J. Phys. Chem. 99:16442, 1995. S Gorer, A Albu-Yaron, G Hodes. Chem. Mater. 7:1243, 1995. MK Norr. J. Phys. Chem. 65:1278, 1961. R Ortega-Borges, D Lincot. J. Electrochem. Soc. 140:3464, 1993. JM Doña, J Herrero. J. Electrochem. Soc. 144:4081, 1997. EB Andersen. Z. Phys. Chem. B. 32:237, 1936. E Beutel, A Kutzelnigg. Z. Elektrochem. 36:523, 1930. GA Kitaev, AA Uritskaya. Russ. J. Appl. Chem. 72:592, 1999. LE Yatlova, AA Uritskaya, GA Kitaev, TI Dzyuba. Russ. J. Appl. Chem. 66:1534, 1993. GA Kitaev, AZ Khvorenkova. Russ. J. Appl. Chem. 71:1325, 1998. J Liu, SM Howard, KN Han Langmuir. 9:3635, 1993. I Dekany, L Turi, G Galbacs, JH Fendler. J. Colloid Interface Sci. 195:307, 1997. D Lincot, J Vedel. In: 10th ECPV Solar Energy Conf. Lisbon Portugal: 1991, p 931. TP Bol’shchikovsa, GA Kitaev, VI Dvoinin, MV Degtyarev, LM Dvoskina. Izv. Akad. SSSR, Neorg. Mater. 16:387, 1980. D Lincot, R Ortega-Borges, M Froment. Appl. Phys. Lett. 64:569, 1994. M Froment, MC Bernard, R Cortes, B Mokili, D Lincot. J. Electrochem. Soc. 142:2642, 1995. R Cortes, M Froment, B Mokili, D Lincot. Philos. Mag. Letts. 73:209, 1996.

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62. 63. 64. 65. 65a. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81.

D Lincot, B Mokili, R Cortes, M Froment. Microsc. Microanal. Microstruct. 7:217, 1996. MJ Furlong, M Froment, MC Bernard, R Cortés, AN Tiwari, M Krejci, H Zogg, D Lincot. J. Crystal Growth 193:114, 1998. H Cachet, H Essaaidi, M Froment, G Maurin. J. Electroanal. Chem. 396:175, 1995. AB Lundin, GA Kitaev. Inorg. Mater. 1:1905, 1965. M. Kostoglou, N. Andritsos, AJ Karabelas. Ind. Eng. Chem. Res. 39:3272, 2000. B Mokili, M Froment, D Lincot. J. de Phys. IV 5:261, 1995. P O’Brien, DJ Otway, D Smith-Boyle. Thin Solid Films 361:17, 2000. JM Doña, J Herrero. J. Electrochem. Soc. 141:205, 1994. JM Doña, J Herrero. J. Electrochem. Soc. 142:764, 1995. RD Engelken, HE McCloud, C Lee, M Slayton, H Ghoreishi. J. Electrochem. Soc. 134:2696, 1987. E Beutel, A Kutzelnigg. Monats. 58:295, 1931. CD Lokhande. Mater. Chem. Phys. 28:145, 1991. KM Gadave, SA Jodgudri, CD Lokhande. Thin Solid Films 245:7, 1994. SS Dhumure, CD Lokhande. Thin Solid Films 240:1, 1994. CD Lokhande, AU Ubale, PS Patil. Thin Solid Films 302:1, 1997. I Grozdanov, CK Barlingay, SK Dey. Thin Solid Films 250:67, 1994. I Grozdanov, M Najdoski. J. Solid State Chem. 114:469, 1995. MTS Nair, Y Pena, J Campos, VM García, PK Nair, J. Electrochem. Soc. 145:2113, 1998. CD Lokhande. J. Phys. D: Appl. Phys. 23:1703, 1990. K Yamaguchi, T Yoshida, T Sugiura, H Minoura. J. Phys. Chem. B 102:9677, 1998. H Meherzi-Maghraoui, P Cowache, D Lincot, M Dachraoui. J. Chim. Phys. 96:259, 1999.

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4 II–VI Semiconductors

There has been a clear emphasis, in the CD literature, on II–VI semiconductors, mostly CdS, some CdSe, and recently on ZnS. This being the case, the reader may reasonably expect this chapter to be a voluminous one. On the other hand, many of these studies have focused on deposition mechanisms and kinetics (which are dealt with in the previous chapter), with photovoltaic cells, and, to a lesser extent, with quantum size effects, both of which will be dealt with in subsequent chapters. Two detailed descriptions of the experimental procedure (for CdS and CdSe) are given in Chapter 2. This leaves the obvious question: “What’s left?” The present chapter will answer that question. This includes properties of the films not explicitly discussed in other sections, such as crystal structure, optical and electrical properties, as well as variants of the deposition process. Also, more detail will be given on non-Cd chalcogenides. In short, there is indeed much left.

4.1 CdS A point concerning CdS deposition. Many studies have used what is referred here to as the standard deposition bath. This bath is made up of a Cd salt, ammonia (sometimes with an ammonium salt to lower the pH) to complex the Cd and adjust the pH, thiourea, and deposition temperatures usually in the range of 60–90°C. Of course, this bath still allows for large differences in reactant concentrations (e.g., the Cd concentration varies from a low of 1 mM to as much as 100

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mM). If the reader has read the previous chapter, it should be obvious that not only is the concentration of various reactants important, but so is the ratio between the Cd (or other metal ion) and the complexant. With this caveat, we will use the term standard bath to cover all concentrations, unless there is a specific reason to do otherwise. Since the majority of studies on CdS used this standard bath in one form or another, the films discussed in this section can be assumed to have been deposited from such a bath unless otherwise stated. Also for this reason, it is more natural to begin with properties of the films and afterwards to discuss variations in deposition.

4.1.1 Crystallography In several cases where epitaxial growth occurs involving the ion-by-ion mechanism, the crystal structure is dictated by the substrate structure. This is treated separately in Section 4.1.5. Many papers state that one or the other crystal modification is obtained, without giving either diffraction data or/and where the data is ambiguous. The energy difference between the hexagonal (wurtzite) and cubic (zincblende) phases is very small (the former is slightly more stable); hence both are often found together. This commonly leads to the presence of twins and stacking faults in the crystals. The density of stacking faults in films deposited from a standard bath increases with increased thiourea concentration or decreased Cd concentration, and is typically 1011–1012 cm3 [1]. If the cubic phase is annealed at ca. 400°C or above, the hexagonal phase is normally obtained. In view of the lack of a convincing explanation of why one or the other crystal structure is formed, a sampling of reported crystal structures is given in list form, together with differences in preparation from the standard bath (if any) that may give clues to the crystal modification obtained.

4.1.1.1 Hexagonal (Wurtzite) When insufficient NH3 was added and Cd(OH)2 was present as a clearly visible suspension, hexagonal CdS formed on the substrate if the solution was well stirred; i.e., the precipitate in solution, which was cubic, was not allowed to accumulate on the substrate [2]. A high resolution transmission electron microscopy (HTEM) study of the early stages of CdS deposition on a carbon-coated TEM grid showed only hexagonal CdS to be formed, while hexagonal with some cubic CdS was formed by precipitation in the solution [3]. Using CdI2 in a standard bath, hexagonal CdS was obtained. (If CdCl2 was used, the deposit appeared to be zincblende, although it may also have been highly textured wurtzite) [4].) Using CdCl2 (Cd(Ac)2 gave no XRD) hexagonal CdS with moderate texture (0002) was deposited on glass [5].

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With ethylenediamine as complexant and with Cd(OH)2 present in solution, some hexagonal CdS was formed, although cubic CdS might also have been present [6]. A citrate/ammonia bath gave predominantly hexagonal CdS [7]. Some particularly clear examples of predominantly hexagonal formation were from an acid bath using thioacetamide [8], a triethanolamine bath under conditions where ion-by-ion deposition was believed to occur and the deposition rate was slow [9], and a nitrilotriacetate bath, where deposition was also slow but a hydroxide cluster deposition was shown to take place [10]. (Ion-by-ion growth, under conditions similar to the last example, but with a high enough complex concentration to prevent Cd(OH)2 formation, also showed apparently highly textured hexagonal CdS, although in this case the predominant presence of this form was not unambiguous [10].)

4.1.1.2 Cubic (zincblende) If just enough ammonia was added to dissolve Cd(OH)2, the cubic form was obtained [2] (regardless of stirring—see Ref. 2). From a standard solution although the concentrations of reactants used were not given [11]. Mainly cubic obtained from a standard bath on SnO2/glass over a range of conditions (including with and without ammonium buffer and using ethylenediamine instead of ammonia) [12]. From a bath with low Cd concentration (1 mM) and high ammonia concentrations (2 M) suggesting that the conditions were such as to favor ion-byion deposition [13]. Another study with low Cd concentration (2–5 mM), ammonia concentrations ca. 300 times higher than the Cd concentration, and added ammonium ions (which reduces the pH and therefore favors ion-by-ion deposition) likewise found only the cubic phase. Preferential (111) texture of cubic CdS on ITO/glass [14].

4.1.1.3 Mixed Hexagonal/Cubic A mixture of phases was often reported. This is not surprising considering the small energy difference between them. Some examples follow. Standard deposition giving thick films (close to 1 m thick) on glass resulted in films that were ca. 90% cubic and 10% hexagonal [15]. Either not enough NH3 to dissolve the Cd(OH)2 and not stirred or a large excess of NH3 and stirred [2]. Standard deposition on tin oxide/glass [16]. The precipitate formed in solution was predominantly cubic, with some hexagonal [17].

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Detailed analysis of different XRD techniques led to the conclusion that CdS used in CdS/CdTe PV cells was polytype, with essentially random stacking of cubic and hexagonal structures in individual crystals [18]. This study goes a long way to explaining the wide variation in apparent crystal structure. A triethanolamine/ammonia bath gave a mixed deposit [19,20]. Some comments on the role of the anion of the Cd salt are in order. Due to the small energy difference between the two phases, small changes in adsorption of solution species onto the growing crystals may be enough to dictate the final crystal structure. Early studies have shown differences in the crystal structure of CdS precipitates, depending on the anion of the Cd salt. Halides resulted in hexagonal CdS, sulphate gave cubic, while nitrate could give either, depending on temperature and pH [21]. Films deposited from a typical NH3/thiourea bath, using either the iodide or chloride salt of Cd, were studied by both XRD and ED [4]. The films deposited using CdCl2 were highly textured (111) zincblende (with crystal size ca. 15 nm). Those deposited from CdI2 showed sharp hexagonal reflections that were not highly textured. In addition, these sharp peaks rode on broad peaks, which, while not discussed, suggest that most of the film is made up of a much smaller crystal size, which might be cubic. Using acetate or chloride anions, a well-defined peak [(probably (111)] and some much smaller peaks were found for acetate and very weak (but narrow) peaks for chloride [22]. If ammonium salts were added (therefore lowering the pH), well-defined and strongly textured (111) peaks were obtained for both anions. The presence of foreign ions can obviously affect the crystal structure. More noticeable, however, was the effect of these ions on the crystallinity. Adding Cu (by CuCl in the deposition solution) caused a decrease in the intensity of the (zincblende) XRD peaks with increasing Cu concentration [23,24]. However, there was no apparent change in peak width, implying that the crystal size did not change appreciably, since a reduction in crystal size—whether by reduction of the coherence length of a fixed crystal size by defects or by actual change in crystal size—would result in peak broadening. Similar results were obtained for films doped using CuCl2 and subjected to annealing in air at 300°C [25]. Doping with a variety of cations (Cu, Ag, As, In) was in all cases reported to result in loss of the XRD pattern [15].

4.1.2 Crystal Size There are a number of factors that determine crystal size. Probably the two most important are the deposition mechanism (ion-by-ion growth, in general, will result in larger crystal size than the hydroxide mechanism, discussed in detail in Chap. 3) and specific adsorption of anions onto the growing crystal (this can affect both crystal structure and size). In most cases, the CdS crystal size from the standard bath was typically 10–20 nm, although sometimes it could be several times larger than this, particuCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

larly from lower-pH (buffered) solutions. This is larger than the typical size of the accompanying CdS precipitate, which tends to be between 5 and 10 nm. For example, a crystal size of some tens of nanometers was deposited on a carbon-coated TEM grid, but the precipitate in the same solution was 3–6 nm [3]. This was attributed to an ion-by-ion mechanism for the film vs. a hydroxide cluster mechanism in solution. Some large (typically 100 nm but reported up to 1 m) hexagonal-shaped thin crystals of Cd(OH)2 were also formed in the solution precipitate. If the growth begins by a cluster mechanism but ion-by-ion growth (whether by free sulphide or by complex decomposition) occurs in parallel, then an intermediate crystal size is a logical outcome, since the ion-by-ion growth can occur on the small hydroxide-formed clusters, leading to crystal growth. The nature of the anion of the Cd salt was found to affect the crystal size in some cases, although it appears that such effects are not universal but related to other variables in the deposition process. In one report, the use of CdCl2 gave large crystals (probably 100 nm), but with CdAc2 no XRD pattern was observed [5]. From the optical spectrum of the CdS deposited from an acetate bath, a crystal size of ca. 5 nm can be inferred based on size quantization. Another study [22] found fairly narrow XRD peaks (crystal size at least 20 nm) using CdAc2. CdS deposited from a CdI2 solution gave an XRD pattern of sharp peaks (see the previous section) on a broad background [4]. This, together with the blue-shifted optical spectrum, suggests that most of the film is made up of a much smaller crystal size. The crystal size of films deposited using CdCl2 in the same study was ca. 15 nm (estimated from the XRD pattern). Using a nitrilotriacetate (NTA) bath, the deposition mechanism could be more easily controlled than from the ammonia bath; NTA is a much stronger complexant than ammonia, allowing pure ion-by-ion deposition if the NTA:Cd ratio is high enough. The crystal size from such an ion-by-ion deposition was 70 nm (instrument broadening limited), while from a hydroxide-mediated NTA bath it was 5 nm [10]. A crystal size of ca. 10 nm was reported from an acidic thioacetamide bath [8]. The only other acidic bath where crystal size could be extracted was the photodeposition method using a thiosulphate solution, where, from the XRD, a size of 10 nm could be estimated [26] If a comparison of crystal size and structure is made from this and the previous section, a general trend appears suggesting that crystals that grow in the hexagonal modification are, in general (and there are exceptions), larger than those that are cubic.

4.1.3 Optical Properties 4.1.3.1 Transmission/Absorption Spectroscopy Many studies present optical absorption or transmission spectra of the resulting films. (A reminder that a spectrophotometer measures transmission, not abCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

sorbance: The “absorbance” measured by a spectrophotometer is a mathematical manipulation of the transmission and, if reflection is present, will not be accurate without reflection correction. See Sec. 1.4 for details.) The purpose of these spectra is usually to show that the deposits are of high quality (usually interpreted to mean transparent in the subbandgap region; in most cases, scattering is undesirable, although there may be exceptions, depending on the intended application of the films) and are indeed made of the material claimed (as seen from the bandgap value, which can be estimated from these spectra). Scattering is usually caused by optically large (comparable to the wavelength of the light) nonhomogeneous aggregates; this often occurs by sedimentation of colloidal aggregates onto the substrate. However, it can also occur even if no colloidal phase is present in the solution. There have been few studies on control of scattering in CD films. One study reported more aggregates and lower transmission at lower deposition temperatures from a citrate/ammonia bath [27]. However, another study, using a triethanolamine bath, reported more aggregates at higher deposition temperatures, although not in a regular manner, resulting in generally lower transmission at higher deposition temperatures [28]. Therefore, as is generally the case, such specific results should not automatically be applied to all CD CdS films. Application of a magnetic field perpendicular to the substrate caused an increase in transmission of the film [29], although it is not clear whether this is due to increased specular reflectance of the field-free films or decreased scattering of the films deposited with the field. It may be indirectly inferred from the transmission spectra that there is no less (maybe even more) scattering in the films deposited with a field; but Atomic Force Microscope (AFM) morphology studies give the impression that these films are smoother, which would be one (but not a unique) indication of less scattering and greater specular reflectance. Most CD films reported are fairly transparent to very transparent, typically between 60 and 90% transparent in the subbandgap region, although lower values are not infrequently seen. Since the optical spectra are most often not corrected for specular reflection, this reflection will reduce the transmission, but the “quality” of such films can be very high. In fact, a high specular reflectance is indicative of “good-quality” films, since films with a considerable degree of scattering exhibit low specular reflectance. The other type of information that can be extracted from optical spectra via the bandgap is an estimation of crystal size if the semiconductor is in the sizequantized domain. This is due to the blue spectral shift caused by size quantization: The smaller the crystal size, the larger the blue shift and the larger the bandgap. This is discussed in more detail in Chapter 10. Here we note briefly some studies where such shifts have been seen. In most cases, the standard bath gives a crystal size larger than the largest size that will show an appreciable blue shift (for CdS, this value is ca. 6 nm). There are some exceptions, however, with minor modifications of the deposition conditions. The use of iodide [4] and acetate [5] as anions of the Cd salt resulted in blue shifts Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

of 0.1 eV or more. Blue shifts of ca. 0.2 eV were measured in CdS films deposited from citrate/ammonia baths [7,27]. Transmission spectra can be modified by doping. Thus, if Cu (as CuCl) is added to a triethanolamine/ammonia bath for CdS deposition, the effective bandgap measured from the spectra shifted from 2.35 eV (no Cu) to just over 2.0 eV [23,24]. The less steep onset and more pronounced absorption tail of the Cu:CdS, together with the very low value of bandgap, suggests that this large shift is due either to a subbandgap transition arising from Cu impurity states in the CdS or even to absorption in a separate phase of Cu-S, which would probably itself be quantized. In a somewhat similar study, but with films annealed at 300°C in air, a drop in bandgap due to Cu doping from 2.48 eV to 2.38 eV was measured [25]. In contrast, another study reported an increase in bandgap from 2.4 to 2.48 eV upon doping with either Cu, Ag, As, or In (as well as a less steep onset). In principle, doping can affect the measured absorption spectrum in different ways. The most obvious is introduction of levels in the gap (which would result in an apparent lowering of the bandgap). Amorphization (as seen by the loss in the XRD pattern) is commonly found to occur upon doping; the resulting disorder could cause tailing of the states near the band edges. This tailing would normally be seen as a decrease in the bandgap, although an increase has been explained by splitting of the tailed levels from the bands [15]. Increase in bandgap can also occur by filling the lower-lying conduction (valence) band levels with electrons (holes), thereby requiring a larger photon energy to promote an electron from the valence to the conduction band (the Burstein–Moss shift). Antimony doping has been shown to have strong effects on the optical spectra [30]. The bandgap decreased from 2.47 eV (pure CdS) to 1.7 eV (nominally 0.075% Sb) and then increased to 2.86 eV for 0.1% Sb. The strong drop in bandgap (the absorption was strong) for moderate Sb levels suggests an impurity band, while the increase for higher doping levels could, by itself, be explained by the Burstein–Moss shift. However, these explanations would require effective removal of the impurity band at some intermediate Sb level, which would not normally be expected. The crystallinity of the films does not appear (from the XRD spectra) to change appreciably with doping, thus removing amorphicity as a possible explanation for the effects.

4.1.3.2 Photoluminescence Photoluminescence (PL) is light emitted when photogenerated electrons and holes recombine. In that sense, it is the opposite of absorption. However, while an optical absorption spectrum in the great majority of cases shows the valence–conduction band transition (i.e., the bandgap), photoluminescence spectra are much more complex as a rule. The bandgap emission (sometimes called band-to-band emission) may or may not occur (it often does not), and subbandgap transitions of longer wavelength are commonly seen. These transitions are from various surface Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

or bulk states in the bandgap, and they therefore can give information about these states, which are not seen, or at best are seen very weakly, in absorption spectra. In the majority of cases, most of the energy of photogenerated electron–hole pairs is not emitted in radiative transitions, but converted to heat in nonradiative transitions (i.e., the luminescence efficiency is commonly low). It is outside the scope of this book to go into detail on the explanations for the various PL spectra measured in CD CdS films; several examples are given with possible origins for the various spectral peaks. The pH of deposition (adjusted by adding NH4OH, therefore pH increased but free Cd2 decreased) affected the PL spectra of the CdS films deposited from a standard solution [31]. A broad, red luminescence (ca. 1.2–2.0 eV with peak at 1.68 eV) was characteristic of all the spectra, regardless of deposition pH. At pH
11.5, a narrow (0.18 eV half-width) green peak (2.255 eV) appeared, but it did not occur above or below this pH value). This peak, ca. 0.2 eV less than the bandgap, could be either a shallow-donor-to-shallow-acceptor transition or a band-to-fairly shallow interband state transition. Different crystal sizes (some in the quantum size regime) were obtained by varying the film thickness. Three interconnected PL peaks at ca. 1.83, 1.35, and 1.06 eV were obtained (no green emission) [32]. A model of transitions from a deep donor (Cd–O complex) level to various other levels was suggested to explain these peaks. Another deposition, probably from a standard solution (although the details of the deposition were not complete), gave a dominant peak in the green region (2.38 eV) and a broad low-energy shoulder extending to ca. 1.5 eV [33]. Deconvoluting the spectrum revealed, besides the green peak, a small yellow peak (2.25 eV), attributed to a Cd interstitial–Cd vacancy (iCd–VCd) complex and red band (1.80 eV) associated with sulphur vacancies. PL spectra of CdS deposited from two different acidic baths have been reported. From an acid thioacetamide bath, a broad band centered around ca. 1.5 eV was obtained [8]. The most likely cause for this luminescence was suggested to be valence band hole–S vacancy recombination. Films deposited under illumination from a thiosulphate solution exhibited a broad band from ca. 1.46–2.0 eV (peak at ca. 1.66 eV) [26]. The wide range of different PL spectra obtained shows just how much the various films vary from each other and the sensitivity of the (mainly surface) electronic structure of the CdS to the deposition parameters.

4.1.4 Resistivity and Photoconduction of As-Deposited CdS 4.1.4.1 Dark Resistivity Note that only as-deposited films are considered here. As will be evident to anyone familiar with semiconductor processing, annealing of these films can be car-

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ried out to drastically change their electrical and photoconductive properties, depending on the annealing conditions. While a considerable amount of work has been expended on studying the various effects of annealing on CD CdS films, this is mostly outside the scope of this book. (See, however, the end of Sec. 4.1.7, which gives some information on the effects of rapid thermal annealing and subsequent removal of oxygen on the electrical properties. Also, since it is relevant for PbS(Se) photoconductors and for photovoltaic cell use, both of which usually require some annealing, it will be treated somewhat in Chap. 5 and 9.) Two points are worth noting here. One is that, in general, annealing in hydrogen or vacuum invariably reduces the dark resistivity of CD CdS, while reannealing in air or oxygen increases it again. Oxygen can chemisorb on the CdS surface, extracting electrons from the conduction band, and therefore decrease the free-electron concentration (hence increase in resistivity). The second point is that, for photovoltaic cell use, a lower resistivity does not necessarily mean a better cell; in fact the opposite may even be true. The dark resistivity of CD CdS is often, although by no means always, very high. This may be reasonably attributed to the high degree of stoichiometry usually obtained with CD films. This stoichiometry is certainly expected for ion-byion growth, and is probable also for hydroxide-mediated growth (both simple and complex) as long as all the hydroxide has been converted. It is likely that the cases where low resistivity has been reported can be explained by nonstoichiometry. In one study [34], the activation energy of the dark conductivity (measured at and above room temperature) was found to be 1 V, from which it can be inferred that the CdS Fermi level is very close to the bandgap center, meaning that the CdS is highly intrinsic and free of common bulk defects, in particular S vacancies. Weak n-type behavior with very low donor concentration (1012 cm3) and considerably higher deep trap densities (1015–1016 cm3) were found on standard films deposited on quartz using space charge–limited current measurements [35]. This high resistance may be responsible for the fact that the commonly used Ag contacts to the CdS behave in an ohmic manner [36], although Ag is not normally considered to be a good ohmic contact to CdS in general. Two reasons can be given for this. One is that the high resistivity of the CdS means that even an appreciable contact resistance may be negligible. Another factor is that, since the high-resistance CdS is often close to intrinsic (i.e., the Fermi level is close to the bandgap center), which by definition means a higher value of work function, even a high-work-function metal is less likely to form a Schottky (blocking) contact to the CdS. This is a good point to bring up briefly a property of very small crystals (often obtained in CD), which is dealt with in more detail in Chapter 9. The crystal size is, in most cases, much smaller than the size of any space charge layer that would be formed. This means that in an isolated nanocrystal, unless the doping level is very high (usually it is not, as attested to by the high resistivities more of-

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ten obtained), there will not be much built-in electric field. The situation of a film of aggregated nanocrystals is not so obvious, but it is likely that a space charge layer of the normal type will not be obtained. This has important implications in considering contacts both between nanocrystal and metal contact and between nanocrystals. As far as (photo)conducting properties are concerned, terms such as Schottky contact or grain boundary barriers need to be considered with this point in mind. Values for the resistivities of various CD CdS films are given in Table 4.1. The first thing that can be observed is the lack of any obvious correlation between the resistivity and deposition conditions in most cases. Only some conditions are given in the table; the details of the concentrations are not always available; even if they are, it would be oversimplistic to try to compare them based only on concentration. For example, the ratio of Cd salt to complexant is no less (probably

TABLE 4.1 Dark Conductivity and Photoconductivity of CD CdS Films Resistivity (-cm) Bath conditions CdCl2 90C CdSO4 70C CdAc2 or CdCl2 80C CdAc2 Cd(NO3)2 room temperature CdSO4 90C (80–95C) pH 10–11 (80–95C) pH 9–10 CdCl2 CdI2 CdAc2 NH4 85C Triethanolamine 26C Triethanolamine 75C Citrate/ammonia 60C Citrate/ammonia 75C Citrate/ammonia 50C Citrate/ammonia 70C Thioacetamide (pH
8) 40C Temp. of minimum resistance 63C Temp. of minimum resistance 78C CdCl2 85C No magnetic field CdCl2 85C Magnetic field

Dark

Light

Reference

9

700 1000 103–104 10 –1000 Small effect — —

38 39 5 16 40 11 22 22 4 4 41 28, 36 28, 36 27 27 7 7 37

10 108 106–107 104–106 1012 104 340–600 15–150 106–108 108–1010 103–104 109 109 5.107 108 2  108 2  108 104 10 400 2  105 60

102–104 103–105 — 1000 1 2000 5 200 4 — — — — —

42

The first column gives the Cd salt and deposition temperature. If not specifically noted, the bath is a standard ammonia/thiourea bath.

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even more) important than the absolute concentrations themselves. Also, so many parameters are interrelated. In an attempt to say something intelligent about these resistivities, there appears to be some correlation between the pH and resistivity, with low resistivity obtained when the pH is relatively low (only a few experiments have been carried out at relatively low values of pH; also note Ref. 22, which describes an anomalously low resistivity even at “normal” values of pH). The bath described by Ito and Shiraishi [37] is very different from the previous ones, for three reasons: the relatively low pH (
8), the use of thioacetamide instead of thiourea, and the flow system used in this deposition. Very low values of dark resistivity were obtained with this bath and with an unusual temperature dependence (a minimum of 10 cm was found at 63°C, which increased on either side of this temperature value). It was suggested that Cl, from the NH4Cl buffer, acted as a dopant; however, other chloride baths gave much higher resistivities. Some weak correlation between film morphology and resistivity was noted for films, deposited from a pH 9.5 bath: Films deposited from a closed system (no evaporation or loss of ammonia) were more uniform and had somewhat lower (two to three times) resistivity (ca. 1.5  104 -cm) than films deposited from an open bath [43]. The triethanolamine bath gives consistently high resistivities, independent of the deposition conditions, and the citrate/ammonia bath behaves similarly, although with somewhat lower resistivities. It should be noted that chemisorption of oxygen on the CdS is known to increase the resistivity, and some (many?) differences may well be due to different surface chemistries of the CdS crystals. Thus, by definition the complexant can bind to surface Cd and cover the surface in some cases. This can lead to (at least) two opposing effects: “insulation” of one crystal from another and prevention of oxygen chemisorption (this latter need not, however, necessarily lead to lower resistivity but the opposite if the adsorbed complex acts in a similar way to the oxygen). This surface adsorption may explain, e.g., both the high resistivity of the triethanolamine-bath films and their relative independence from other deposition conditions. It is worth noting that most (although not all) CD CdS films have a rather similar crystal size; therefore this factor does not seem to be as important as might have been expected. Films deposited from a standard bath but under application of a magnetic field perpendicular to the substrate exhibited resistivities that decreased strongly as the strength of the magnetic field increased [42]. The resistivity, which for a film deposited without the field was 2  105 -cm, was as low as 60 -cm on application of a 77-mT magnetic field. From the dark resistivity/temperature values, a level with an activation energy of 0.68 eV was found for the field-free films and a shallow level of 0.053 eV and others at
0.17 eV for those deposited in the presence of the field. The shallow level at 0.053 eV was suggested to be due to excess Cd. It was suggested that the magnetic field might affect the rate of arrival of cad-

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mium and sulphide ions at the substrate and that more sulphur vacancies are formed in the presence of the field, hence the lower resistivity (see Sec. 4.1.6.8 for more details on this deposition).

4.1.4.2

Photoconductivity

CD CdS films are usually strongly photoconducting as deposited. This is in contrast to most other CD films, which tend to be only weakly or moderately photoconductive in the as-deposited state [44]. Values of the resistivity under illumination are given in Table 4.1 where available. Most experiments have used illumination intensities comparable to solar irradiation. As is the case for the dark resistivity, the dependence of the sensitivity of the photoconductivity (defined here as the ratio between light and dark conductivity) on the deposition parameters is far from clear-cut. Some observations can be made, however. The first (obvious) one is that for a high sensitivity, the dark resistivity must be high. Apart from this, there does seem to be a general trend (clear-cut in the triethanolamine and citrate baths and seen also by the lack of appreciable photoconductivity in the one low- (room-) temperature-deposited film reported [40]) of an increase in photosensitivity (due to decrease in light resistivity) with increasing deposition temperature. For the standard baths, the sensitivity varies in most cases between 103 and 6 10 . A CdI2 bath resulted in somewhat greater sensitivity (as well as dark resistivity) than a CdCl2 bath [4]. The deposition temperatures of these two baths were different, but it was reported that the film properties were independent of the deposition temperature. Another study found no appreciable difference in either dark or light resistivity between acetate and chloride baths [5]. Interestingly, there was apparently a large difference in crystal size between the two baths (see Sec. 4.1.2 on crystal size), which implies that the crystal size is not an important factor in determining the resistivity or photoconductivity, at least for this bath. The triethanolamine bath showed a distinct trend of the photosensitivity with deposition temperature [28,34,36,45,46]. The photosensitivity was higher (which, since the dark resistivity was temperature independent, means the light resistivity was lower) for higher temperatures, with the major change occurring between 30 and 45°C. The citrate/ammonia bath has much in common with the triethanolamine/ ammonia bath with high light sensitivities, particularly at higher deposition temperature, and little temperature dependence of dark resistivity. This is in spite of the very different bath compositions and concentrations. In particular, the citrate bath contained much lower concentrations of Cd and, as a result, was more highly complexed. One study showed a very strong dependence of dark resistivity on measuring temperature and a much weaker dependence of the resistivity under illumina-

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tion, with the result that the sensitivity was very temperature dependent [15] (see Sec. 4.1.4.3). The decay time of the photoconductivity, , is another important parameter for which, as with the dark resistivity and photoconductivity, there is often no obvious correllation with the deposition parameters. In many cases, the decay is very slow (hours), particularly for the triethanolamine bath, where it can be greater than 10 hr. Decay times of hours have been reported for the standard bath [4,16], although if CdI2 was used instead of CdCl2 in the latter,  dropped to seconds (the deposition temperature was different for the chloride and iodide bath, but reportedly this did not affect the properties). For films deposited from the citrate bath, temperature-dependent decay times from 1 min (deposition at 75°C) to tens of minutes (60°C) were reported in one study [27] and tens of minutes with only a small deposition temperature dependence in the other [7]. Differences in the deposition conditions of these two studies were described earlier. From the triethanolamne bath, much longer decay times were observed at higher deposition temperatures, as shown in Figure 4.1 [34]. Another study, using ammoniumbuffered standard films (possibly ion-by-ion deposition), found decay times of seconds for single films, which increased to several minutes for multiply deposited films [47]. The photosensitivity of these latter films was less than for most others (ca. 30 for the single films and 10 for the multiple ones).

FIG. 4.1 Time dependence of photoconductivity of CdS films deposited from triethanolamine/ammonia/thiourea bath at two different temperatures (26°C and 75°C). The two plots at each temperature differ by the ratio between the Cd and thiourea concentrations: [Thiourea]:[Cd]
0.25 for the upper plots at each temperature and 0.5 for the lower plots. (Adapted from Ref. 34 with permission from Elsevier Science).

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It is important to note that oxygen adsorption has a strong effect on the photoconductivity decay of the triethanolamine films [36]; it is probable that this is not limited only to these films. Oxygen greatly decreases the decay lifetime, as seen by the increase in  for aged films compared with freshly deposited ones. In this same study, the photoconductivity sensitivity and decay time both decreased greatly with increase in measuring temperature (the former due mainly to the activated decrease in dark resistivity with increase in temperature, but also due partly to decrease in resistance under illumination). The possible effects of oxygen on photoconducting parameters are discussed in this reference (general background on photoconductivity is given in Ref. 34). Here it is enough to note that adsorbed oxygen is believed to extract electrons from the CdS conduction band and also to introduce deep trapping centers (interband surface states) that increase the carrier lifetimes, thereby increasing the photoconductivity decay time. Oxygen can also affect electron mobility between grains by modification of grain boundary barriers between crystals. A comparison was made between films deposited from standard baths using either CdCl2 or CdAc2 [48]. While for most measurements these films were annealed at 300°C, and therefore are not compared with the as-deposited films here, thermally stimulated current (TSC) measurements were carried out on asdeposited films. Such TSC provides an indication of the density and energy of trapping centers: The magnitude of the current, obtained by heating the sample and exciting charges out of traps into a band, is an indication of the trap density and the temperature at which the current is generated is a measure of the trap energy. The trap density was much higher and the traps considerably deeper for the acetate-prepared films than for the chloride ones (after annealing, the trap density was higher for the chloride films, seen as a large increase in TSC for the chloride film after heat treatment). We can conclude this section with the insight, gained from this overview of the electrical and photoconductivity properties of these films, that, in spite of the many studies already carreid out, a comprehensive and systematic study of these properties and their correlation with a wide range of deposition parameters is still needed in order to understand what determines these properties. These studies should also include postdeposition treatments—not so much annealing, which has been carried out, but surface treatments (e.g., immersion in triethanolamine), which could show the importance (or lack of it) of the crystal surface condition.

4.1.4.3

Electrical Properties of Doped CdS Films

Doping can be divided into two parts: native doping (e.g., S vacancies) and extrinsic doping by foreign elements. This section deals with the latter, not because it is more important but because there is little in the literature to link native doping with the electrical properties of CD films. It will be enough to note that the few measurements of ND (donor density) carried out tend to give values typically

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around 1016 cm3 (e.g., Ref. 49 for epitaxial, hexagonal CdS on (1¯ 1¯ 1¯) InP from C–V measurements) or even larger. This is somewhat surprising, since, if these films are so highly stoichiometric as expected, lower values might be expected. Two comments here. One is that the doping may originate from the surface, since the surface-to-volume ratio of the CdS crystallites is high. The other, also related to the high surface area, is that errors in interpreting C–V measurements commonly used to derive ND may arise because of the lack of knowledge of the true surface area. We will now consider individual dopants used in CD CdS. Boron. Boron, substituted for Cd, is a donor in CdS. B-doped films were deposited by adding boric acid to a standard deposition bath, with the B:Cd ratio varying from 105 to 102 [16]. The boron was assumed to occur in the form of borate ions (BO 2 3 ). The dark resistivity dropped nearly three orders of magnitude with optimum B content (B:Cd ratio in solution of 0.001), from 2.104 to 30 -cm. At higher B concentrations, the dark resistivity again increased until, at a B:Cd ratio of 0.01, the original resistivity of the undoped CdS was regained, and it did not change with increased B content. However, the resistivity under illumination decreased to ca. 3 -cm, almost independent of the B content. Thus high B content increased the photosensitivity of the CdS, although only by a factor of 2–3. The rate of decay of the photocurrent was greatly reduced by B doping, from ca. one hour for undoped films to as much as several tens of hours for doped ones. This suggests a deep trap resulting from the B, separate from the shallow donor that is responsible for the drop in resistivity. Nitrogen. Nitrogen ions (N with energy of 130 KeV) were implanted into CdS deposited from a triethanolamine/ammonia bath [20]. The resistivity of the as-deposited films was ca. 108 -cm and dropped, depending on ion dose, up to seven orders of magnitude for an ion dose of 1017 ion/cm2. Even more notable, the conductivity type changed from n-type (the normal type for CdS) to p-type, as measured by hot probe. An acceptor level, 0.6 eV above the valence band, was introduced by the ion implantation. Copper. Copper was doped into triethanolamine/ammonia films by adding CuCl in the deposition solution [23,24]. Resistivity dropped from 109 cm (undoped) to a minimum of 0.5 -cm for optimum doping (the Cu content of the CdS was not measured) and the conductivity was p-type. Both higher and lower Cu concentrations in solution gave higher resistivities; it was surmised that too high concentrations of Cu in the deposition solution resulted in rapid precipitation of the Cu as CuxS, depleting the solution of Cu. While the high-conductivity p-type CdS was, not surprisingly, not photoconductive, films with smaller amounts of Cu were photoconductive (although the photoconductivity gain was less than for nondoped CdS, and the response time of the photoconduction, both

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rise and fall, was much shorter than for the undoped film). The Cu-induced states resulted in recombination centers rather than long-lived trap states characteristic of CD CdS films. Another study of Cu-doped CdS films reported different results, although there were differences in the preparation; in particular the films were annealed in air at 300°C, and CuCl2 (rather than CuCl) was used [25]. In this work, dark resistivity did not vary greatly for low Cu concentrations (ca. 107 -cm), but increased by nearly two orders of magnitude for high Cu concentration. Also, the photoconductivity response increased with Cu content. The annealing carried out in this study calls for caution in any comparison with the previous one. A third study (no variation in Cu content) found a decrease in dark resistance (from ca. 108 to ca. 104 -cm upon Cu (also Ag, As, and In) doping [15]. The dark resistivity was very highly temperature dependent, especially for the undoped CdS (ca. 106 -cm at 35°C and 104 -cm at 50°C). Since the resistance under illumination (⬇105 -cm for undoped and ⬇104 -cm for doped) was much less temperature sensitive, the photoconductivity response was very temperature dependent, being more pronounced at lower temperatures. Lithium. Lithium can act as an acceptor in CdS. Shikalgar and Pawar studied electrical [50] and photoconducting [51] properties of Li-doped CdS [standard deposition with 0.1% (by weight of CdS) Li salt added to the deposition solution]. The addition of Li increased the dark resistivity by a factor of 3–4 (resistivities were given as ohms and the exact geometry of the measuring system was not described: however, specific resistivities could be estimated to be of the order of 107 -cm). The room-temperature energy of activation for both doped and undoped films was ca. 22 meV, i.e., shallow donor conduction. Above ca. 60°C, the resistivity dropped much more rapidly as a function of temperature, with an activation energy of ca. 1.2 eV (intrinsic conductivity). In addition, the Li-doped film exhibited an intermediate level at 0.16 eV in the temperature range of ca. 40–65°C, ascribed to a Li acceptor level. The Li-doped films, like the undoped ones, were n-type; CdS is difficult to dope p-type due to self-compensation, and since the resistivity of CD CdS films are normally very high, it is not surprising that acceptor doping does not increase this resistivity very greatly. The light:dark conductivity ratio of these Li-doped films was not explicitly given, although an order of magnitude value of 106 could be inferred from the results. The photocurrent–time behavior for the CdS:Li (the equivalent data for the undoped films were not given) was history dependent. Initially, the photocurrent increased linearly with time (over a maximum measured time of 10 min), but increased more rapidly and exponentially with larger photocurrents after light–dark cycling. In all cases, the decay was multiexponential and slow, typically tens of minutes. These measurements were carried out in vacuum; if air was introduced, the steady-state photocurrent decreased, attributed to oxygen adsorption on the surface of the CdS crystals, resulting in extraction of electrons from the conduc-

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tion band. Interestingly, the peak of the photocurrent of the CdS:Li was blueshifted (480 nm) compared to the undoped CdS (520 nm); no explanation could be given for this effect. Aluminum. Aluminum, as a trivalent ion, should be an n-type dopant for CdS. A small decrease in resistivity (by a factor of 2) to ca. 103 -cm was found when Al2(SO4)3 was added to a standard bath [52]. In another study, Al was added as Al2(SO4)3 to a thiourea bath (85°C) of relatively low pH (9.5) [53]. The resistivity of the undoped CdS was ca. 3  105 -cm and decreased at least an order of magnitude on doping with Al. Codoping with chlorine (as NH4Cl in the solution) decreased the resistivity almost another order of magnitude. In both studies, an excess of Al resulted in an increase of resistivity. This was explained by excess Al3 occupying interstitial positions: however, interstitial Al3 might be expected to increase the n-type doping and therefore decrease the resistivity, and it is not clear why the resistivity should increase. It may be that an insulating Al(OH)3 phase occurs if too much Al is added.

4.1.5 Substrate-Dependent Growth and Epitaxy 4.1.5.1

Introduction

There are a number of studies that report the effects of the substrate on the CdS films. With the exception of epitaxial deposition, which will constitute the main part of this section, it is usually difficult to explain any specific substrate effect. Also, it should be borne in mind that each specific study is confined to one deposition bath and that a substrate effect obtained for one bath need not necessarily be obtained for a different one. Some examples of substrate effects on the film can be given. Strong (0002) or (111) texturing was obtained on glass substrates but much weaker texturing on SnO2 /glass [16]. Much poorer crystallinity (this may also mean smaller crystal size) of the CdS was obtained on Si than on glass or ITO/glass [54]. Using XRD peak shifts and optical absorption spectroscopy, the presence of strain in as-grown CdS on both glass and ITO/glass was inferred [55]. The strain was greater for the films deposited on the ITO, and this was attributed to mismatch strain between the CdS and ITO.

4.1.5.2

Epitaxy

Various investigations into the epitaxial deposition of CdS onto different singlecrystal substrates have been carried out by Lincot et al. On InP, which is closely lattice matched to CdS (0.1% difference), epitaxial deposition (c-axis of hexagonal CdS perpendicular to the substrate) occurs on the (1¯1¯1¯) P polar face of the InP but polycrystalline deposition on the (111) In face [49,56]. This difference was clearly due to differing chemical or electrostatic interaction between the InP faces

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and the constituents of the deposition solution, since the lattice spacing is the same for both polar faces. The degree of epitaxy on the P face was also dependent on the deposition conditions: in particular, higher temperatures resulted in better epitaxy, as might be expected due to the higher mobility of the depositing species on the surface. The epitaxy was maintained up to at least 100 nm film thickness (the maximum thickness studied). On (100) etched (Br-MeOH) InP, cubic, fairly well-oriented CdS was formed, although with many small-angle grain boundaries. In the absence of the etch treatment (using only H2SO4 to remove native oxide, as was also employed following the Br-MeOH treatment), only polycrystalline CdS was deposited, which showed that not only the crystal face, but also the manner in which that face was pretreated, is important [57]. The formation of the cubic phase in contrast to the hexagonal phase formed on(1¯1¯1¯) InP was attributed to the lack of lattice match between the (cubic) (100) face of InP and hexagonal CdS. GaP has a much larger mismatch with CdS (7%) compared with InP. Yet a fair degree of epitaxy was obtained for CD CdS on the (1¯1¯1¯) GaP surface [58]. In this case, a mixture of cubic and hexagonal CdS with a large density of stacking faults, presumed due to strain relaxation arising from the large mismatch, was obtained. Because of the importance of the junction between CD CdS and CuInSe2 (CIS) for thin-film photovoltaic cells (see Chap. 9), as well as the relatively small mismatch between CdS and CIS (0.7%), deposition onto oriented CIS films has also been studied [59,60]. Two different CIS faces were studied—(100) and (112). As with the (100) P face of InP, because there was no lattice match between this face and any hexagonal CdS face, cubic, epitaxial CdS was deposited. On (112) CIS, which matches either (111) cubic or (0001) hexagonal CdS, a mixture of both phases was deposited, with a moderate degree of epitaxy, which improved if the CIS was first subjected to a cyanide treatment (cyanide removes excess CuxSe and various Se species and generally cleans up the surface). The epitaxy also improved with increase in temperature. The transition temperature was quite abrupt: Below 60°C, the films were polycrystalline, while above this temperature they were epitaxial, with increasing perfection as the temperature increased. Since the deposition at lower temperatures was much slower than at higher ones (for a temperature difference of ca. 40°C, the deposition rate increased 30 times), this implied that the increased mobility of the depositing species on the surface was not necessarily the main factor in the temperature dependence of epitaxy. It was suggested that faster and more complete decomposition of reaction intermediates was an important factor in determining the epitaxy. If so, this is presumably true for epitaxial deposition on other substrates. The composition of the deposition solution was important in order to obtain epitaxy on CIS, which, in contrast to the epitaxy on InP, was not obtained using “standard” solutions. Instead, low Cd concentrations (maybe more important, low

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Cd:NH3 ratios) and lower than usual pH (adjusted by adding ammonium ions) were necessary. These are the factors that favor an ion-by-ion mechanism over one involving Cd(OH)2, which may explain the need for this modified solution (most commonly used solutions operate in or near the region where Cd(OH)2 can exist, at least on the substrate). This explanation still leaves open the question of why good epitaxial growth can be obtained on InP under some conditions but not CIS from a “standard” solution. While there is no present answer to this question, considering the sensitivity of the epitaxy to the chemical properties of the surface (such as etch or differences between P and In polar faces of InP), this should not be too surprising. The epitaxy was maintained for CdS thicknesses up to 100 nm, after which the deposit became polycrystalline. This transition coincided with the visual formation of CdS in the solution, which resulted in a switch of the mechanism from an ion-by-ion growth, necessary to obtain epitaxy, to one involving colloidal species. Since, in principle, conditions can be chosen so that only an ion-by-ion growth occurs, it can be expected that much thicker epitaxial films are obtainable from CD on suitable substrates.

4.1.5.3 Deposition on Monolayers: Selective Growth and Patterning If a substrate is not clean, films either do not grow or grow with poor adherence on the “dirty” parts of the substrate. This has been exploited by partially covering the substrate with a monolayer. When a mica substrate was incompletely covered by a monolayer of octadecylphosphonic acid, CdS growth was found to occur preferentially on the mica [9]. This was shown also for CdS deposited on an octadecyltrichlorosilane- (OTS)-coated Si substrate and was used to pattern the CdS deposit by applying the OTS onto the Si using a patterned stamp [61]. Either the OTS could be removed by sonication or, even without removal of the OTS, deposition occurred only on the bare Si if the CdS was not too thick (ca. 50 nm). Edges with a variation of ca. 100 nm could be deposited by this method.

4.1.6 Variations in Preparation 4.1.6.1

Variation in pH

While deposition rate normally increases with increase in pH for the standard bath, using an ammonium salt to lower the solution pH resulted in the opposite behavior; i.e., increased pH led to slower deposition [22]. The pH in these experiments was increased by adding NH4OH. Increased pH (and NH4OH) results in two opposing effects: Thiourea decomposition increases but free [Cd2] decreases. Since the deposition rate for the solution with no added NH 4 increases with increase in pH, the former apparently outweighs the latter. When extra ammonium ion is added, much more ammonium hydroxide is needed to increase the

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pH to the original value, and the decrease in free [Cd2] dominates the reaction kinetics. More generally, addition of an ammonium salt (apart from the hydroxide, which increases the pH) increases the complexation of Cd without increasing the pH (it actually decreases the pH of an NH4OH solution due to buffer action). This results in a decrease in deposition rate due to a lower free-Cd2 concentration and, if the pH decreases, reduction in the thiourea decomposition rate. A slower deposition allows the formation of thicker films, since less CdS will be formed homogeneously in the solution. This has been shown in many studies (e.g., Refs. 22, 40, and 43). In addition to slower thiourea decomposition, lower pH will also decrease the likelihood of Cd(OH)2 formation and will therefore favor (relatively) an ion-by-ion rather than cluster mechanism. Only a few acidic baths have been described (see also sections 4.1.6.6 and 4.1.6.7 below). In one, thiosulphate was used at a pH of between 2 and 4 and a temperature of 85°C [131]. The bandgap was 2.35 eV and the resistivity 104–105 -cm. Thioacetamine has also been used at a pH of 5 [8]. The films from this bath were clearly hexagonal. The rationale for using an acid bath is to prevent the formation of hydroxy species; this is a major problem for ZnS but much less so for CdS.

4.1.6.2

Variation in Complexant

Cyanide, a stronger complexant than others used, has been employed as a complex for CdS deposition [62]. Except for the fact that thicker films could be obtained (ca. 1 m compared to a few hundred nanometers with the standard method), the properties of the films made with cyanide [crystal structure, crystal size, bandgap (measured to be 2.2 eV, an anomalously low value but the same as that of films deposited from a standard bath in the same study)] were the same as those of the standard bath. Solution composition details were not given. Films up to 3 m thick were obtained with the triethanolamine/ammonia bath [19]. It is probable that this larger-than-normal thickness is due to deposition occurring via an ion-by-ion mechanism, due to the additional complexing by the triethanolamine, the somewhat lower pH than usual (10), and, for the 3-m film, the low deposition temperature employed (30°C), factors that reduce free-Cd2 and/or OH concentrations, thereby favoring the ion-by-ion mechanism. This would result in thicker films since no (or at least less) CdS is wasted as a homogeneous precipitate in the solution. Ethylenediamine has been used as a complexant [6]. It is a stronger complexant than ammonia and therefore only needs to be used in low concentration (between two and four times that of Cd).

4.1.6.3

Variation in Thiourea Concentration

Increase in the concentration of thiourea clearly leads to an increase in deposition rate. Additionally, it has been seen that the defect density (measured from TEM micrographs as structural defects such as stacking faults) decreased greatly with

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an increase in thiourea concentration from 1011–1012cm2 for a thiourea concentration of 28 mM to 1010cm2 for a concentration of 100 mM [12].

4.1.6.4

Anion Effects

Some effects of anions have been noted previously. These are often not consistent and, in general, it is difficult to attribute the effect of the anion of the Cd salt to any specific effect with any degree of confidence. Two studies on the effect of the anion on the rate of deposition did find small but significant differences (a factor of 2), which more or less were in agreement [63,64]. The latter found the rate to increase in the order: CdI2, CdSO4, Cd(NO3)2, Cd(CH3COO)2, and CdCl2; this series corresponds approximately with the decrease in (negative) electrode potential and correspondingly to the decreased strength of complexation between Cd2 and the respective anion. Additionally, the rate decreased with increase in concentration of the anion (added as an alkali metal salt). These observations suggest that the effect may be due to mild additional complexation compared to that of ammonia alone.

4.1.6.5

Surfactants

The addition of surfactants to the standard CdS bath resulted in a reduction in the rate of deposition and an increase in the terminal thickness [65]. Surfactants adsorb onto surfaces (both the substrate and colloidal particles in the solution), and therefore it is not surprising that the growth rate is reduced. The adsorption of the surfactant onto CdS colloidal particles also can prevent flocculation and precipitation, thereby increasing the CdS available for deposition (hence, presumably, the increased terminal thickness). At the same time, it is possible that the presence of surfactants (or any strongly adsorbed species) might prevent sticking of the colloidal particles to the substrate and to each other, in the same way as they prevent flocculation, which is exactly the sticking together of the colloids. However, there is no evidence that this actually occurs in this study. Another effect of the surfactant, inferred from the slightly higher than usual bandgap (2.52 eV), is the small size of the CdS crystals; although not measured, it can be assumed to be ca. 5 nm from the bandgap shift, presumably due to size quantization (see Chap. 10). This is not surprising since crystal growth is in competition with adsorption of the surfactant.

4.1.6.6

Electrochemical/Chemical Deposition

Yamaguchi et al. described an interesting extension of the CD process for CdS using a parallel electrochemical step [66]. They termed this process electrochemically induced chemical deposition. It is based on electroreduction of protons in solution, which results in an increase in pH locally at the electrode. They used thioacetamide as a sulphur source. In the acid solutions in which the deposition is carried out (pH between 1.6 and 4.6), no film deposition of CdS occurs (although it does precipitate in the solution) in the absence of the electrochemical proton reduction. In the presence of proton reduction, CdS films were formed. These films

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were characterized by relatively large, hexagonally faceted wurtzite CdS crystals (crystal sizes from a few tens of nanometers up to 300 nm, the larger sizes being formed at lower pH). The CdS precipitated in solution was quite different; it was not faceted and had smaller crystal size (ca. 15 nm). It was suggested that the film growth proceeded by a surface-catalyzed decomposition of a Cd–thioacetamide complex and that the electrochemical proton reduction affected the surface properties (presumably the surface of both substrate and growing CdS) in such a way as to reduce the activation energy needed for the deposition reaction. The growth was an atom-by-atom (or ion-by-ion) process, leading to larger crystal size than normally obtained by the hydroxide-mediated particle growth. By adding a strongly adsorbing species (2-mercaptoethanol) to such a deposition bath, they were able to reduce the crystal size by varying amounts due to surface capping of the growing crystals, preventing further crystal growth but allowing nucleation to proceed. Thus, the film thickness was not strongly affected provided the mercaptoethanol concentration was low (10 mM); above this concentration, film growth was prevented, as would be expected, since adhesion between crystals and substrate or between different crystals would probably be poor in the presence of an adsorbed coating [67]. The resulting nanocrystalline films exhibited quantum size effects (see Chap. 10 for more details).

4.1.6.7

Illumination-Induced Growth

There have been a few experiments related to the effect of illumination of the growth of CdS films. Simple heating of the deposition bath by absorption of the radiation is one obvious factor that can affect the deposition [68]. However, even in this case, other effects occur, since the color of the bath was reported to darken if UV (sunlight) illumination was employed. Based on previous studies of illuminated CdS colloids when elemental Cd was formed, both as a film and in solution [69], as well as the known tendency of ZnS to undergo reduction to metallic Zn under UV illumination, this darkening may be assumed to be caused by elemental Cd. There are several possible mechanisms that may explain such an effect; reduction of the CdS by photogenerated electrons is one possibility. A variant of CD was based on illumination of a solution containing thiosulphate and cadmium ions by UV light [26,70,71]. CdS was deposited only on the illuminated portion of the substrate. Since only light absorbed by thiosulphate (wavelength shorter than 300 nm) was effective, the effect was attributed to photodecomposition of thiosulphate to elemental S and solvated electrons and subsequent reaction with Cd2. 2 S2O 2 3  h D S  SO 3  2 2S2O 2 3  h → S4O 6  2e  2 2 SO 2 3  S2O 3  h → S3O 6  2e 2  Cd  S  2e → CdS

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(4.1) (4.2) (4.3) (4.4)

The optimal pH was ca. 3.5 (at lower values, the film was contaminated with elemental S, which forms spontaneously in the dark), while the deposition rate slowed down at higher values). The CdS was sphalerite, and, from an examination of the XRD, a crystal size of 10 nm could be estimated. The bandgap was 2.42 eV, the literature value for CdS.

4.1.6.8 Deposition Under the Influence of an External Magnetic Field Deposition has been carried out from a standard bath with a magnetic field applied, both parallel and perpendicular to the substrate [29]. Differences were found in the film properties only for a field perpendicular to the substrate. The transmission of the films in the nonabsorbing region was ca. 10% higher (see Sec. 4.1.3.1). The films deposited with the field were up to three times thicker than those deposited using the same conditions but in the absence of a field. The dark resistivity of the field-applied films was much less than that of the field-free ones (see Sec. 4.1.4.1). The cause of these effects is not clear. Similar measurements were carried out using an external electric field [72]. Some differences in morphology and optical properties were measured, depending on the direction of the field with respect to the substrate. It is not clear, from the experimental setup, why the field should influence the deposition, since the field is external and should drop across the air and the glass walls of the reaction vessel.

4.1.6.9

Deposition on/in Porous Silicon

Porous silicon is under extensive study, largely due to its luminescence properties. For electroluminescence, however, some form of contact has to be made with the Si, and this necessitates deposition of another phase inside the pores of the Si in order to contact as much as possible of the internal area of the high-surface-area Si. With this in mind, CdS has been deposited inside the pores of porous silicon via a two-stage method [73]. Cd(OH)2 was deposited from an ammoniacal bath at pH 8, followed by conversion of the Cd(OH)2 to CdS by treatment with thioacetamide at pH 8. This was repeated several times until the pores were essentially filled with CdS. The reason that this two-stage process was needed is that either the Si was unstable at the temperatures and pH values needed to deposit CdS from a thiourea solution, or CdS was formed in solution rather than on the Si surface using thioacetamide.

4.1.6.10

Bath Geometry

One of the disadvantages of the CD process as usually carried out is the large waste of materials (for example, in CdS deposition, most of the Cd—often over 90%—is unused in the film deposition because it deposits homogeneously in solution and/or on the walls of the reaction vessel). Probably more important than

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the material loss is the environmental concern of disposing of this Cd (or other heavy metal) if it is not recycled. The same goes for the ammonia often used as a complexant—much is lost to the atmosphere and, in common with other complexants, much is wasted to tie up the heavy metal ions. Use of low concentrations of metal ion (ca. 1 mM) presents a partial solution to this problem. However, for any industrial process, a continuous-flow system seems the best option. Ito and Shiraishi flowed a solution of thioacetamide and a Cd salt into a 0.5-mm-thick flow space [37]. A detailed flow system has been described for CdS by Boyle et al. [74]. There are several features in this system, shown schematically in Figure 4.2. Probably the most important is the locally heated substrate. Since CD reactions are usually very temperature dependent, by heating only the substrate (in this case, by resistive heating), deposition is limited, to a large extent, to the substrate. This system also uses ethylenediamine instead of ammonia, which greatly decreases loss by evaporation as occurs with ammonia in an open system. Filters are employed in the flow system to remove any colloidal matter formed. Fresh reagents can be added, as required, to the recirculating closed-loop system. In connection with the foregoing studies, a batch process, whereby the solution was filtered after deposition and complete reaction and reused (with the addition of more reactants as required), was also shown to be feasible [75]. While the deposition rate slowed down for successive depositions, this could be compensated for by increasing the concentrations of various reactants from run to run. The photovoltaic parameters of Cu(In,Ga)Se2/CdS solar cells fabricated using this approach was not found to vary from deposition to deposition.

FIG. 4.2 Flow system for continuous deposition using a locally heated substrate. (After Ref. 74 with permission from Elsevier Science.)

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By minimizing the spacing between substrates, reagent utilization can be maximized by increasing the ratio of substrate surface to solution volume [27]. The maximum film thickness was reduced if the substrate spacing was too close— depending on the deposition parameters, and this was related to a “critical thickness” of reagent layer at the substrate surface connected with the presence of colloidal particles in the solution. Typically some 1- to 10-mm spacing was necessary to obtain the maximum thickness (200–400 nm in this case). This study was carried out under conditions where a hydroxide cluster mechanism was operative (the solution was already turbid in the early stages of deposition), and the results cannot be extrapolated to other mechanisms.

4.1.7 Impurities in Chemical Deposition CdS There have been a number of studies involving impurities in the CdS films, with various results. It must be emphasized that if the films are not very well rinsed after preparation (and possibly even if they are), some of the ions involved in the preparation may be present as adsorbed species. The most comprehensive study, involving a range of analytical techniques, sums up the probable impurity situation for films deposited from standard baths [76]. The main impurity, not unexpectedly, is oxygen (ca. 11 atomic %). Evidence was presented to show that this O was probably mainly in the forms of carbonate and adsorbed water. The carbonate could come from two sources: dissolution of atmospheric CO2 and (see Eq. (3.11)) from decomposition of thiourea. Nitrogen (ca. 5 at. %) occurs as carbon–nitrogen bonds, probably mainly cyanamide (NCN2), although other C–N bonded compounds were also believed to be present. If cyanamide is present as the Cd salt, this would tie up 5% of the Cd. The Cd:S ratio was found to be only slightly higher than unity (ca. 1.02), and some of the Cd may be bound to carbonate. Therefore other C–N species are likely also to be present, e.g., cyanide, several of which could adsorb to one Cd or even to a CdS moiety. By reducing the concentration of thiourea in the bath, C–N impurities in the CdS film could be reduced almost to zero [77]. Whatever the nature of the C–N impurity, much of it could be removed simply by dissolution in water at 60°C [78] It should be pointed out that this deposition was carried out for films ca. 50 nm thick; the study was carried out with CdS window layers for solar cells in mind, which are usually thin. It is possible that much longer depositions result in different impurities. Thus the sparingly soluble cadmium carbonate and cyanamide will be converted to CdS if enough sulphide ion is formed with time (or, for the complex-decomposition mechanism, if enough adsorbed thiourea decomposes on the surface of the solid phases). Of course, longer time also means more thiourea decomposition products. Another study found much smaller concentrations of oxygen in the films (4 at.%) [79], and most of this oxygen was attributed to bound water. Although

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the deposition rates were comparable for this study and that in Ref. 76, the deposition solutions and conditions were quite different. A decrease in the O content, measured by XPS, on Ar ion sputtering together with a Cd:S ratio close to unity led Danaher et al. to propose that the O (not quantified) was present as surface sulphate [38]. In this same investigation, SIMS analysis (much more sensitive than XPS, which is limited to ca. 0.1% concentration) found a variety of impurities, including decomposition products of thiourea, CdO, and Cd(OH)2, but these were not seen in XPS, showing that they were present in very low concentration. It is worth noting that analysis of the deposition solution after deposition was complete, and after filtration of the solid precipitate showed the presence of urea and guanidine, but not cyanamide, and that the amounts of these compounds were less than those stoichiometrically expected, suggesting further decomposition of urea to ammonia and carbonate [75] (see Sec. 3.2.1.1). Films deposited from chloride and iodide (otherwise standard) baths were compared [80]. I (ca. 3 at.%) was found in the iodide-deposited films but 1% Cl in the chloride-deposited ones. About 5 at.% O was also found in both films. The excess Cd was believed to occur as Cd–O, Cd–OH, and, additionally in the iodide films, Cd–I or Cd–(I–O) species. In most cases, the Cd:S ratio in these films was slightly greater than unity (usually between 1.02 and 1.1). A ratio of less than unity (0.92) was found for multiple layers (i.e., two or more layers deposited one on the other [47]; for a single layer using the same deposition solution, the ratio was unity). The oxygen concentration varied from 8% (from an iodide bath) to 10–12% (from a chloride or sulphate bath). As in the previous study, the oxygen was believed to be present mainly as Cd(OH)2 or CdO. Even larger concentrations of oxygen were found at low concentrations of ammonia (up to 18%) or at lower deposition temperature. Another (XPS) study, however, found the ratio to be typically 1.3 [81], Some general conclusions can be drawn concerning oxygen in the films, in spite of the large spread reported in different studies, both in amounts and in interpretation of its source. The first thing to note is that, since the crystal size of the CdS is often in the region of 10 nm, around 10% of all the atoms will be located at a crystal surface. Thus adsorption of either oxygen or water will already show a relatively large amount of oxygen. However, unless this oxygen substitutes for sulphur, such adsorption will not change the Cd:S ratio. Other sources of oxygen, such as hydroxide, carbonate, and oxide (the last is less likely) will increase this ratio. As discussed earlier, carbonate can form from either dissolution of atmospheric CO2 or from the decomposition products of thiourea. Cd(OH)2 is more likely to be formed when the pH is high (unbuffered solutions) or the ammonia concentration is low (less complexation, which probably outweights the slightly lower pH). The effect of temperature on Cd(OH)2 formation is complicated. A

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higher temperature means a higher hydroxide concentration at a constant pH and less effective complexation on one hand, but also faster decomposition of thiourea resulting in more efficient conversion of the hydroxide to sulphide on the other. From the results of Nakada et al. [81] discussed earlier, where more oxygen is found at lower temperatures, it appears that the latter effect is dominant. Rapid thermal annealing (RTA) in vacuum of CD CdS films has been shown to remove most of the oxygen that occurs in these films [82]. Typical annealing conditions were: heating rate—100°C/min; maximum temperature— 600°C for 1 min; cooling rate—50°C/min. X-ray photoelectron spectroscopy (XPS) showed that, apart from the immediate surface, oxygen was effectively removed from CdS films (deposited from a thiourea bath). Dark resistivity was drastically reduced after this treatment, from ca. 107 -cm for the as-deposited film to ca. 1–10 -cm after the RTA treatment. As can be expected, the photoconduction sensitivity also decreased drastically, from a light:dark resistivity ratio of ca. 104 to ca. 1.6 after annealing. This decrease in resistivity was attributed to removal of electron traps that originated from the adsorbed oxygen. The authors also suggested that, unlike conventional annealing, which results in a loss of stoichiometry, e.g., by formation of Cd vacancies if annealed in vacuum, the RTA process does not change stoichiometry. Pronounced thickness effects on resistivity were also noted in these RTA CdS films. For example, a 95-nm-thick film showed a resistivity of 15 -cm, which decreased to 0.2 -cm for a thickness of 150 nm (with no pronounced change for even thicker films). Also, storage (in a dessicator, presumably in air) increased the resistance of the thin films, about an order of magnitude for the 95 nm film after 50 days, with continuing increase but, for a 250 nm-thick film, only a small, initial increase (ca. 50%). RTA of the stored films decreased the resistivity to their original value before storage. The effect of storage was attributed to oxygen adsorption.

4.2 CdSe 4.2.1 A Mechanistic Introduction Before going into details of the various aspects of specific CdSe depositions, and although it is not intended to deal with mechanistic aspects here (they have been considered already in Chap. 3), it bears mentioning that, although in contrast to CdS, the complex-decomposition mechanism has not been discussed with respect to CdSe deposition, it is still possible that this mechanism does occur in some, or even many, cases. If there is no evidence specifically in favor of this mechanism in general, there is also none against it. This point is stressed here since, in the literature on CdSe (and selenides in general), it is automatically assumed that the reaction proceeds via free selenide ions.

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4.2.2 Selenosulphate as Se Source CdSe was deposited using a selenosulphate source as far back as 1970 [83]. Mirror films of CdSe were reported to form from an ammoniacal solution of Cd2 only under conditions where Cd(OH)2 was present in the solution and at a pH  11.75. Under the conditions of those experiments, it was reported that this was the minimum pH required to convert Cd(OH)2 to CdSe. At higher values of pH, the rate of conversion to CdSe increased, but so did homogeneous precipitation, with the result that the films were thinner. A detailed study of CdSe deposition was carried out using an ammoniacomplexed solution with selenosulphate [84]. Most of this study was concerned with kinetic measurements, already discussed in Chapter 3. Two different types of solution were considered: a clear solution where there was no visible Cd(OH)2 and one with added KOH to give a visible Cd(OH)2 suspension. The former required heating to at least 45°C for deposition to occur (although it is likely that deposition would occur even at room temperature after enough time). The CdSe was of the zincblende structure. With a visible Cd(OH)2 suspension present, deposition occured at room temperature, but the terminal thickness was only ca. 80 nm. The higher the pH, the lower the terminal thickness, since more of the Cd precipitated in the solution. The CdSe from this bath was a mixture of wurtzite and zincblende structures. The deposition rate and terminal thickness of the films were somewhat dependent on the nature of the substrate, both somewhat larger for Ge and Si than for glass. A modification of this method used lower concentrations of ammonia (0.2 M for a Cd concentration of ca. 50 mM) in a sealed vessel, thus preventing irreproducibilities due to escape of ammonia vapor [85]. Treatment of the Ti and stainless steel substrates by soaking in a suspension of Cd(OH)2 improved the homogeneity of the films. At an ammnonia concentration of 0.3 M, no deposition occurred (at least within the time frame of these depositions—about one hour). A triethanolamine/ammonia bath has been used for CdSe [19]. While this system resulted in thick films for CdS (up to a few microns), CdSe films deposited under the same conditions, only using selenosulphate instead of thiourea, were thinner [although films of 500 nm were obtained at 30°C that did not show signs of satuation (of thickness) after 25 hr—the longest time measured]. Ethylenediamine has also been used as a complex for Cd, with both precipitates and films formed [86]. In this case, the emphasis was on the precipitates, and no characterization was carried out on the films. Nitrilotriacetic acid (NTA) [N(CH2COOH)3] is a complexant for many metal ions (see Sec. 2.9.1.2 for information on this compound). The sodium or potassium salts of NTA have been used to complex Cd for CD of CdSe from selenosulphate solutions [10,87]. The rate of film growth depends on many factors, as discussed in Chapter 3; experimental details for CdSe deposition from this solution are given in Chapter 2. However, growth times are generally longer than for Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

ammonia-based baths—typically from a few hours to a few days for film thicknesses in the range of 100—300 nm. The most notable property of these films is their change in color with deposition conditions—from yellow if deposited at temperatures below 0°C and under conditions where the hydroxide mechanism is operative, to very deep red (thick films appear black by reflected light) for high-temperature, ion-by-ion depositions. This variation in color is a consequence of size quantization, discussed in detail for these films in Chapter 10, and the CdSe (always zincblende) crystal size varies from 3 to 20 nm. This color variation can be translated into a variation of CdSe bandgap from ca. 2.3 eV for the yellow films to the bulk value (for zincblende CdSe) of ca. 1.8 ev. Annealing the films causes crystal growth and therefore loss of the size quantization effects. The major crystal growth, which corresponds to the phase change from zincblende to wurtzite, occurs between 300 and 400°C [88,89]. Since the optical spectra of these films are so sensitive to the crystal size (a change of 10% in crystal size can result in an easily measured spectral shift), measurements of the spectra provide a sensitive technique to investigate the effect of different deposition parameters on crystal size. Thus, while the crystal size is not strongly dependent on the various concentrations of reactants (apart from the NTA:Cd ratio in the region where the mechanism changes), small increases (of the order of 10–20%) in crystal size are observed if the Cd and/or selenosulphate concentration is decreased considerably in the hydroxide cluster mechanism regime [90]. This can be rationalized, in a general way, by the greater likelihood of small, thermodynamically unstable CdSe nuclei growing to a stable configuration if the reactant concentrations are greater (see Section 1.2 for a discussion of nucleation and growth), since the growth rate will be faster. The greater the concentration of nuclei, the smaller the final crystal size for a fixed reactant concentration. This reasoning can also explain the observation that the average crystal size increases somewhat as deposition proceeds and reactant concentration decreases [89,91], although this growth might also occur by deposition of new CdSe on previously deposited crystals; probably both mechanisms are operative to a greater or lesser extent. From broadening of photoluminescence peaks with increasing deposition time, it was inferred that the crystal size distribution also increased with deposition time [91]. Illumination of the solution during deposition by supra-bandgap light affects crystal growth, probably via photoelectrodeposition of CdSe on the growing crystals [90,92]. The solution used for CD can also be used for electrodeposition of CdSe. Light absorbed by the individual CdSe crystals forms electron/hole pairs, and the electrons can reduce the solution at the CdSe surface in the same way as those supplied by an external power supply. Figure 4.3 shows the transmission spectra of CD films deposited in the dark and under illumination for two different deposition temperatures. The red shift in the spectra of the illuminated films indicates a larger crystal size of these size-quantized samples (by ca. 1 nm) compared to the nonilluminated ones. Additionally, the shape of the spectrum changes for Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

FIG. 4.3 Transmission spectra of CdSe films deposited from a selenosulphate/NTA bath in the dark and under illumination (tungsten halogen lamp) at two different temperatures (6°C and 55°C).

the low-temperature illuminated film; the onset becomes less steep. This spectral shape is typical for films electrodeposited from this solution [87]. Since the rate of electrodeposition is essentially temperature independent while that of CD is strongly temperature dependent, the effect of the illumination (through the relative amount of photoelectrodeposited CdSe) will be greater for low-temperature films, seen particularly clearly by the pronounced change in shape of the low-temperature film. The growth does not occur for very weak illumination, suggesting that the photoelectrochemical deposition is not very efficient, and other processes (electron/hole recombination or parasitic electrochemical reactions) dominate. Also, the crystal size saturates for light intensities above a certain level. This was interpreted to mean that one electron/hole pair was sufficient to influence the growth process [92]. This would not be expected for a photoelectrochemical growth process, as described earlier. A more logical explanation for the saturation, particularly in view of the probable low quantum efficiency of the photoelectrochemical CdSe deposition discussed earlier, is that one of the charges is removed more rapidly into the solution than the other. If, e.g., this is an electron, then a hole will be left (probably trapped) on the crystal. Absorption of another photon will form another electron/hole pair, which will then recombine rapidly by Auger recombination (three-body interaction) rather than form more CdSe. While the pH of the deposition solution (based on the cluster mechanism) has been found to increase by as much as 0.8 during the deposition (see Sec. 3.3.2), this increase was found to be considerably greater, up to 2.2 pH units, under illumination [92]. This could provide some clues about the mechanism of the CdSe formation under illumination. A possible pathway that could account for

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this increase in pH (i.e., overall generation of OH) is given as 2   Cd(OH)2  SeSO 2 3  2e → CdSe  SO 3  2OH

(4.5)

This reduction has, of course, to be balanced by the local oxidation reaction. The most likely reaction, oxidation of sulphite, leads to an increase in acidity that would cancel out the rise in pH of Eq. (4.5). Probably other oxidation reactions that do not generate acidity occur that result in a net increase in pH (an example would be photocorrosion of the CdSe by photogenerated holes). Another possibility that could explain the effect of illumination is a change in the electric double layer surrounding the CdSe particles, either adsorbed on the substrate or in the solution, which could lower a potential barrier to adsorption and coalescence, as suggested previously for film formation from Se colloids under illumination [93]. Partial coalescence would reduce the blue spectral shift due to size quantization. However, the spectral shape is not expected to undergo a fundamental change in this case. The photoelectrochemical explanation therefore appears more reasonable. Addition of silicotungstic acid (STA) to a selenosulphate/ammonia/triethanolamine bath resulted in a reduced rate of deposition but a larger final thickness (greater than 1 m could be obtained). This was attributed to adsorption of the STA on the individual CdSe crystals (typically 4–5 nm in size), which impedes aggregation of the invidual crystals [94]. Reduced aggregation will slow down both film growth (which relies on aggregation if the mechanism is one of colloidal growth as appears to be the case here) and precipitation of CdSe in solution, which will result in loss of CdSe for film formation. From the XRD pattern of such CdSe films reported in an earlier study [95], a crystal size of ca. 4 nm could be estimated. The STA may also play a role in limiting the crystal size by capping, although even without the STA, small crystal sizes ca. 5 nm are usually obtained from similar deposition baths.

4.2.3 Selenourea Source Films have been deposited using selenourea and an ammonia-complexed solution at 65°C [96]. Zincblende CdSe was obtained with an optical spectrum corresponding to a bandgap of 1.84 eV (the bulk room-temperature bandgap of zincblende CdSe is ca. 1.8 eV). Analysis of electrical conductivity measurements indicated charge transfer occurred via a variable hopping mechanism through fairly deep states (a level 0.29 eV below the conduction band was found from these measurements).

4.2.4 N,N-Dimethylselenourea Source N,N-Dimethylselenourea was used with ammonia baths and additional citrate or tartrate complexation at pH of 11.3 (citrate bath) or 10.4 (tartrate) to deposit CdSe on glass at room temperature [97]. No XRD of the films was detected. From the

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optical spectra, the bandgaps were high (ca. 2.1 eV, even larger for thinner films) suggesting crystal sizes of ca. 5 nm or less (see Chap. 10). The XRD spectra of such a small crystal size can be missed in normal powder XRD if special care is not taken, mainly slow scanning. Electrical and photoconductive properties of these films are described in Section 4.2.7.2. Another study of these films concentrated on the particle size of the films and is discussed in Chapter 10 [98].

4.2.5 Selenosemicarbazide Source One example is given in the literature where selenosemicarbazide was used to deposit CdSe from solutions containing different complexants [99]. This reagent was apparently more stable than selenourea or N,N-dimethylselenourea, and it is surprising that it does not appear to have been subsequently used. The CdSe films were specular and had a resistivity of 108–109 -cm, which dropped about an order of magnitude on illumination.

4.2.6 Epitaxial Deposition Using a nitrilotriacetate solution and a complex:Cd ratio high enough to prevent Cd(OH)2 formation (ion-by-ion mechanism), epitaxial growth of zincblende CdSe was obtained from a selenosulphate solution on both (111) and (1¯1¯1¯) polar faces of single-crystal InP [100]. (For lower complex:Cd ratios, in the regime of the hydroxide cluster mechanism, the deposits were always polycrystalline, as expected for this mechanism.) The degree of epitaxy improved with increasing temperature and was high at 90°C. Additionally, there was a high density of twins in the deposits obtained at low temperatures, but less in those obtained at 90°C. The addition of silicotungstic acid to the deposition solution destroyed the epitaxial growth, presumably due to blocking of the InP surface (and also the growing CdSe) by the strongly adsorbed silicotungstic acid. This study also reported that films deposited on carbon membranes at temperatures 80°C were of hexagonal (wurtzite) structure, with a high density of planar defects, in contrast to the zincblende obtained from both hydroxide and ionby-ion mechanisms at lower temperatures and to the epitaxial films on InP at all temperatures.

4.2.7 Some Specific Optoelectronic Properties 4.2.7.1

Photoluminescence (see also Sec. 10.2.3)

Photoluminescence of CdSe films deposited from the selenosulphate/nitrilotriacetate bath varied both in intensity and in spectral shape. As an example of the former, a sample on glass, broken from a glass slide when the film thickness was ca. 30 nm, gave a much stronger signal than the original sample left in the deposition

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solution until the thickness reached ca. 100 nm. Such an effect, however, was not reproducible (unpublished results). Regarding spectral shape, both bandgap emission with weak, if any, subbandgap signal [87,91] and deep subbandgap signal with relatively weak bandgap emission [89,101] have been observed with no apparent difference between the samples. (Note that the use of the term bandgap emission refers to emission close to the bandgap energy. The emission may be from very shallow traps.) Although the measurement temperatures for these different experiments were not all the same, temperature does not usually have a very pronounced effect on the ratio between the two peaks, except below 50 K, and cannot explain the observed differences. In humid atmosphere, where water vapor is adsorbed on the CdSe, the predominant emission is close to bandgap [101] (see Sec. 4.2.7.3). Most studies do not state whether measurements were carried out in the open atmosphere (low-temperature measurements, of course, are not). An investigation on films deposited from an N,N-dimethylselenourea/citrate/ammonia bath showed both bandgap and lower-energy (ca. 1.75 eV) peaks [98]. The latter was attributed to larger crystal size (see Chap. 10), but it is likely that the low-energy peak is a subbandgap response arising from surface states. The temperature dependence of the emission spectra provides useful information on the source of the emission. Figure 4.4 shows a series of emission spec-

FIG. 4.4 Photoluminescence spectra taken at different temperatures of a CdSe film deposited from a selenosulphate/NTA bath. Crystal size of CdSe ca. 4 nm. (From ref. 101 with permission from Elsevier Science).

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tra of a (4-nm-crystal-size) CdSe film taken at different temperatures. Each spectrum is made up of a bandgap peak (at high photon energy) and a subbandgap peak at lower energy. The bandgap peak becomes weaker as the measurement temperature increased. This is normal for a band-to-band transition. The subbandgap peak, however, initially increases up to ca. 50 K, and then decreases with further increase in temperature. This behavior (and also incident light intensity dependence of the emission) is typical of donor–acceptor recombination. In the present case, the nanocrystals are essentially intrinsic and are not expected to contain bulk dopants. However, the surface states (see following section), after trapping charge, can behave in much the same way, and the recombination can be explained by recombination of surface-trapped electrons and holes with the emission red-shifted ca. 0.5 eV from the bandgap [101]. In the study by Trojanek et al. [89], this emission is shifted nearly 0.7 eV from the bandgap and the parallel increase in emission energy as a function of crystal size with the theoretical (effective mass approximation) conduction band shift interpreted to mean that the emission occurred from shallow-trapped electrons to deep-trapped holes. Time-resolved measurements of photogenerated (very intense illumination, up to 0.56 GW/cm2) electron/hole recombination on CD (selenosulphate/NTA bath) CdSe of different crystal sizes has shown that the trapping of electrons, probably in surface states, occurs in ca. 0.5 ps, and a combination of (intensity-dependent) Auger recombination and shallow-trapped recombination occurs in a time frame of ca. 50 ps. A much slower (not measured) decay due to deeply trapped charges also occurred [102]. A different time-resolved photoluminescence study on similar films attributed emission to recombination from localized states [103]. In particular, the large difference in luminescence efficiency and lifetime between samples annealed in air and in vacuum evidenced the surface nature of these states. A photoluminescence study of CdSe deposited from a selenourea/ammonia solution onto glass at 80°C and relatively low pH (7–8) was made [31]. An emission peak centered at 1.445 eV (860 nm) was observed with a tail to the low energy side. Such an emission must be due to deep traps, since the shift from the bandgap emission is ca. 0.4 eV, a value close to, or somewhat less than, that for films deposited from selenosulphate solution (see earlier). Annealing in air shifted the emission to higher energies.

4.2.7.2

Photoconductivity

There appears to be a fundamental difference between films deposited from a selenosulphate source and those deposited from a dimethylselenourea source (most of the detailed photoconductivity studies are on the latter). As-deposited films using a dimethylselenourea source had a resistivity of ca. 2  1012 /sq (⬇108 -cm for a film thickness of 0.5 m). The resistivity dropped only a little under illumination, with a maximum decrease of less than an order of magnitude. For

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films deposited at a higher temperature (50°C), the resistivity was up to an order of magnitude lower (ca. 107 -cm), which dropped ca. six times under illumination. Air-annealing increased the photosensitivity greatly (up to ca. 107 at 450°C), in contrast to the case for CdS, where good photosensitivity was obtained for asdeposited films (but always deposited at high temperatures) and was reduced on air-annealing. The dark resistivity of the CdSe films increased by nearly an order of magnitude after annealing, again in contrast to the decrease usually obtained for CdS films. This behavior is also in contrast to CdSe films deposited from selenosulphate/triethanolamine/ammonia solutions, where the relatively low resistivity of the as-deposited film (5103 -cm) dropped to a few ohm-centimeters on annealing (430°C in air) [95], or to another study of selenosulphate/ammonia films with dark resistivity of 107–108 -cm, which dropped to 1–10 -cm after annealing at 280°C in vacuum [104]. On the other hand, in another study on selenosulphate/ammonia films, which reported a dark resistivity of 105–106 -cm as deposited, the dark resistivity initially increased on heating in air (ca. four times at 180°C) then decreased only an order of magnitude on heating at 340°C [105]. The photosensitivity of these latter films was low but appreciable in the as-deposited state (varying from 5 to 50) and increased up to ca. 103 after annealing at 180°C before decreasing again at higher annealing temperatures. While it is clear that CdSe films deposited from dimethylselenourea possess higher photosensitivity after annealing compared to those deposited from a selenosulphate bath, the reasons for the difference are not understood. One possibility that might explain this difference was suggested, based on a comparison of CdSe films deposited from a dimethylselenourea bath with CdS films. It was hypothesized that the difference in the effect of annealing on dark resistivity was due to the formation of conducting CdO due to the oxidation of CdS, which did not occur in CdSe [106]. The ease of oxidation of the Cd chalcogenides is normally CdTe  CdSe  CdS; however, this usually refers to oxidation of the chalcogen, e.g., CdTe to CdTeOx, rather than to CdO. More details on the effect of annealing on the photoconduction properties are given in Refs. 106 and 107. Another characteristic of these films, different from most CD CdS films is the fast photocurrent decay (no more than a few seconds at most) of the films, both as deposited and annealed. Clearly, the trapping centers in these films are also efficient recombination centers, which may, at least in part, explain the low photosensitivity of the as-deposited films. CdSe films prepared by the selenosulphate/citrate process were less photosensitive than those deposited by the dimethylselenourea method, although details were not given [106]. Some studies of these films were made after immersion in solutions of Hg or Cu ions, when ion exchange reactions occurred, converting the surface of the crystallites to (partial) Hg–Se and Cu–Se compounds. As expected, such treatments could affect the photoconductivity of the films greatly.

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4.2.7.3

Surface States on CdSe Films

An important aspect of semiconductor films in general with regard to electronic properties is the effect of intrabandgap states, and particularly surface states, on these properties. Surface states are electronic states in the “forbidden” gap that exist because the perfect periodicity of the semiconductor crystal, on which band theory is based, is broken at the surface. Change of chemistry due to bonding of various adsorbates at the surface is often an important factor in this respect. For CD semiconductor films, which are usually nanocrystalline, the surface-tovolume ratio may be very high (several tens of percent of all the atoms may be situated at the surface for 5 nm crystals), and the effects of such surface states are expected to be particularly high. Some aspects of surface states probed by photoluminescence studies are discussed in the previous section. Surface treatments of CD CdSe films deposited from selenosulphate/NTA solutions have a pronounced effect on various optical, electrical, and optoelectronic properties of the films, due to interaction with or modification of such surface states. Mild etching (dilute HCl) of the films reverses the direction of current flow both in CdSe/polysulphide photoelectrochemical cells [108] and in Kelvin probe surface photovoltage (SPV) measurements in air [109]. These studies are discussed in more detail in Chapter 9, in Section 9.2 on photoelectrochemical cells. At this point, it is sufficient to state that the effect is believed to be due to preferential trapping of either electrons or holes at surface states that are modified by the etching process. The adsorption of water vapor on these CdSe films acts to passivate, at least to a large extent, some of these surface states. In particular, strong subbandgap signals have been observed in SPV measurements of these CdSe films (as deposited) only when measured in a dry ambient; in normal atmospheric humidity, no such signal occurs and only suprabandgap light gives rise to an SPV signal [109]. Parallel results have been observed in photoluminescence measurements, which are particularly sensitive to surface states: The predominant subbandgap emission that occurs in a dry atmosphere changes to a predominantly (near) bandgap emission in a humid atmosphere [101]. The asymmetrical nature of these states, seen in optically detected magnetic resonance (ODMR) spectroscopy, is further evidence for their surface nature; bulk states are expected to be symmetric [110]. It is important to note that these effects are seen only in small-crystal-size nanocrystalline films (the foregoing experiments were carried out on 4- to 5-nm-crystal-size films). No such effects were observed if the crystal size was ca. 20 nm; the surface-to-volume ratio is already much smaller for this size. Current–voltage spectroscopy of individual CdSe quantum dots deposited mostly by electrodeposition, but also by CD, using a conducting AFM (atomic force microscope) tip also showed directly the presence of surface states in a dry atmosphere but not in a humid ambient [111].

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4.3 CdTe There appear to be only two independent reports in the literature on CD CdTe (which is the only telluride reported in the CD literature). In the first one, the CdTe was deposited with the main purpose of carrying out electron spin resonance and morphological studies of the effect of annealing, while another paper, based apparently on the first, described the deposition, with the main purpose of further use in photovoltaic cells. Therefore only limited detail on the actual deposition or properties of the as-deposited films were given. This is surprising in view of the fact that the first report is the first case of a CD telluride. In the original study by Padam and Gupta [112], the deposition solution contained triethanolamine/ammonia-complexed CdAc2 and the Te source was TeO2 with hydrazine hydrate as a reducing agent. The nature of the TeO2 solution was not clear, since TeO2 is only slightly soluble in water; it may have been dissolved in a hydroxide solution, in which it is much more soluble. The deposition was carried out at 90°C. Electron diffraction of the as-deposited films showed both zincblende and wurtzite phases of CdTe (the zincblende phase is the more stable and commonly encountered one). Buckley used the same technique to deposit films for photovoltaic cells [113], only with CdCl2 as Cd source and apparently a lower Cd concentration. Electron diffraction of these films showed a predominantly zincblende structure with some wurtzite phase. The films were p-type with a resistivity of 20  5 -cm. These values, and the subsequent photovoltaic cells, apparently refer to asdeposited films; no reference to annealing was made in this study. The second, more recent investigation described Te dissolved in sodium sulphite as a source of telluride (tellurosulphate) [114]. It was previously believed that Te was insoluble in sulphite solution under normal conditions, although there is one previous reference in the literature to this reagent prepared under pressure at high temperatures [115]. The CdTe deposition described in Ref. 114 indicates that the solubility is sufficient to allow deposition of tellurides. The deposition was carried out in a solution of CdSO4 containing triethanolamine, ammonia, and NaOH. Both deposition rate and film thickness were maximal at 75°C deposition temperature. As with the previous TeO2 deposition, both zincblende and wurtzite (dominant) phases of CdTe were obtained. Elemental analysis showed a small Cd excess. This appears to be in contradiction to the XRD analyses, which showed considerable amounts of Te (also TeO2, particularly after heating at 100°C). Although the films were apparently highly scattering, making bandgap measurement more difficult, the bandgap (direct), measured from the optical spectrum, was ca. 1.4 eV, close to the literature value. The (room-temperature) resistivity was ca. 106 -cm and the conductivity n-type. The carrier density was ca. 1019 cm3. In experiments carried out by us to repeat these two methods, we have been able to deposit CdTe, although stoichiometry control was difficult, with Te oc-

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curring in the hydrazine method and mainly TeO2 impurity in the tellurosulphate one. Probably careful attention to removal of oxygen from the solution and maybe addition of hydrazine to the tellurosulphate-based solution would improve the stoichiometry. It is clear that other telluride films, in particular with metals that form very insoluble chalcogenides, such as Pb, Cu, and Hg, should be accessible using these methods. Although not strictly CD, CdTe films were very recently prepared by treating CD Cd(OH)2 films with a telluriding solution [116]. The Cd(OH)2 films were deposited from an alkaline H2O2 bath containing citrate-complexed Cd. These films were treated with a solution of elemental tellurium dissolved in hydroxymethanesulphinic acid. It appears that this solution contains telluride ions, although it has not yet been well characterized. The Cd(OH)2 was converted (incompletely) into CdTe films. The bandgap was ca. 1.5 eV (approximate due to the highly scattering nature of the films). The dark resistivity was ca. 5  108 -cm, which decreased to ca. 7  107 -cm upon illumination.

4.4 ZnS 4.4.1 Introduction Chemical deposition of ZnS has been the subject of considerable activity, the main reason for which is its hoped-for substitution for CdS in thin-film photovoltaic cells. Since the chemistries of Zn and Cd are similar in many ways, it might be expected that deposition of their chalcogenides is also similar. However, there is a dominant difference in their properties that results in the fact that ZnS is considerably more difficult to deposit by CD than CdS. This difference is manifested by the difference in solubility products between the respective hydroxides and chalcogenides. Considering, for example, the sulphides, the relevant values of Ksp are: Cd(OH)2 2  1014 Zn(OH)2 8  1017

CdS 1028 ZnS 3  1025

The deposition for mechanisms proceeding through hydroxide clusters is dependent on a large difference between the solubility products of the hydroxide and sulphide, since the sulphide exchanges the hydroxide. The situation for ZnS is therefore much less favorable than that for CdS: About a million times higher concentration of sulphide is required to form ZnS than CdS [see Eq. (3.46)], as shown in graphical form in Figure 4.5. More sulphide is also required at higher temperatures because of the strongly increasing ion product of water with increasing temperature, resulting in higher hydroxide concentrations for any particular value of pH. (It is again stressed that in dealing with values of solubility products, there are large variations, sometimes of orders of magnitude, between one source and another. Therefore calculations based on these values are correspondingly impre-

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FIG. 4.5 Steady state concentration of sulphide ion needed to convert the hydroxides of Cd and Zn into the corresponding sulphides at 25°C and 60°C.

cise. However, to take the present case as an example, whether the difference in solubility product ratios between the Zn and Cd hydroxides and sulphides is 106 or an order or so larger or smaller in magnitude does not qualitatively alter the conclusion.) The obvious solution to this problem is to deposit ZnS at a lower pH when the OH concentration will be lower. However, with thiourea, for example, lowering the pH results in slower hydrolysis of the thiourea and therefore a lower sulphide concentration (presumably also a reduced decomposition rate if the thiourea–hydroxide complex-decomposition mechanism is effective). This has been circumvented, as described later, by working with alternative sulphiding agents at low values of pH. Deposition by a pure ion-by-ion mechanism should also solve this problem, since no hydroxide is involved. However, in this case we encounter the problem of the high Ksp of ZnS compared to CdS, which again means that more sulphide is needed. For thiourea, this means a higher pH, which again means that strong complexation is needed to prevent Zn(OH)2 formation, by reducing the free [Cd2]. However, this will also reduce the rate of ZnS deposition. While there are many examples in the literature of cluster deposition of ZnS, there does not even seem to be one unambiguous case of ion-by-ion deposition of this semiconductor. As already implied, most depositions of ZnS have been carried out under conditions where Zn(OH)2 can be formed. From the forgoing general discussion, this means that, even if ZnS is identified (by XRD, for example), there is in most studies no evidence for the absence of some hydroxide species. More recent stud-

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ies have addressed this problem, and there are some papers with “Zn(OH,S)” in the title. This probably usually refers to a mixture of ZnS and unreacted Zn(OH)2 and not to a true ternary (more accurately, quaternary) compound. For that reason, such films are discussed in this chapter rather than in Chapter 8 (Ternary Semiconductors). A recent condensed review on ZnS deposition, with emphasis on the differences in CD of CdS and ZnS, is given in Ref. 117. The majority of studies involved ammonia-complexed baths with thiourea as a source of S. Hydrazine was also used in most cases, although there were several studies where ZnS was obtained without hydrazine. Amines (triethanolamine or ethanolamine) were also used, again both with and without hydrazine (which is itself a type of amine). One effect of the hydrazine is to speed up the deposition. Hydrazine, a strong reducing agent is expected to increase the rate of sulphide formation. However, it seems that this is only part of the picture. A study of the effect of various amines on the rate of ZnS deposition showed that although hydrazine gave the fastest rate, other amines (ethanolamine and triethanolamine) also increased the deposition rate [118] (see Fig. 3.7). The amines all act as complexants; therefore they would reasonably be expected to reduce the deposition rate by reducing the free-Cd2 concentration. Amines have reducing properties (redox potentials of hydrazine, ethanolamine, and triethanolamine are 1.16, 0.56, and 0.46, respectively). It appears that the amines accelerate thiourea decomposition. However, the mechanism of this effect is not yet clearly understood [119]. Kinetic studies have found a value for the activation energy of ZnS film formation from thiourea/ammonia-based baths of 5 kCal/mole (21 kJ/mole), too low to be chemical reaction controlled (unlike CdS, which is thiourea-decomposition controlled) but not diffusion controlled, since stirring the deposition solution has no effect [120,121]. This was explained by a rate-determining step of Zn-ligand dissociation from a hydrazine complex. However, addition of an additional complexant into the solution will only act to reduce the free-Zn2 concentration, not only due to the extra complexing power of the hydrazine, but also because of the statistical factor that complexes comprising more than one ligand will be further stabilized by the fact that different combinations of the ligands can be found in the complexes [119]. There is clearly a fundamental difference in deposition mechanism of these ZnS depositions compared with that normally encountered with CdS. More details on the difference in activation energies of deposition for CdS and ZnS (ZnSe) are discussed in Section 3.5.

4.4.2 Specific Studies of ZnS Deposition Some specific properties of ZnS films grown from thiourea/ammonia/hydrazine baths, sometimes with added compounds, are given next. Doña and Herrero noted that hydrazine was not essential for growth but that it speeded up the growth rate (by a factor of ca. 3) and improved the homogeneity Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

and the specular reflectance (a consequence of better microhomogeneity) [120]. No XRD pattern was observed for these films, indicating either highly disordered and/or very small crystals. However, TEM/ED did show that the sphalerite phase was formed with a crystal size of ⬇10 nm. The stoichiometry measured by EDS was Zn:S
1:1. This implies that either pure ZnS or a hydroxy sulphide (Zn,OH,S) was formed. Optical transmission of the films was a spectral average of ca. 85% beyond the bandgap. The bandgap measured from the spectrum was 3.76 eV (literature for zincblende ZnS
3.6 eV), suggesting size quantization, although the measured crystal (ca. 10 nm) is at least twice the size required for such an effect. Resistivity was nearly 109 -cm and independent of bath composition. Temperature dependence of resisitivity gave an activation energy of 0.95 eV, which was ascribed, based on previous studies, to an acceptor level above the valence band due to Cu impurity. Using various amines added to the ammonia bath (in most cases with added hydrazine), sphalerite ZnS films were obtained with a crystal size of ca. 3 nm [118]. Rutherford Backscattering Spectroscopy (RBS) analyses showed that there was about twice as much Zn in the films as S. (More basic solution and more hydrazine gave more stoichiometric films). Extended X-ray Absorption Fine Structure (EXAFS) and Fourier Transform Infra-red (FTIR) spectroscopy showed that the films did not have Zn-O groups but rather Zn-OH ones [122] and that there is probably a mixture of ZnS and unreacted Zn(OH)2, quite likely as a ZnS shell around a Zn(OH)2 core. Optical spectra gave a bandgap of ca. 3.85 eV, considerably blueshifted from the bulk value of 3.6 eV, as expected from such small crystals. Nucleation studies of ZnO and ZnS on glass and SnO2-glass from ammonia/thiourea baths (sometimes also with hydrazine) were carried out [123]. The deposition conditions, mainly pH, were varied. On glass, both ZnS and ZnO could be deposited, depending on conditions. On SnO2-glass, however, only ZnO was formed (a few percent S could be obtained at high pH). This suggested that a surface-activated mechanism was important for nucleation of ZnS and less so for ZnO. Most depositions were carried out at high temperatures. One exception was a room-temperature deposition from an ammonia/thiourea bath (no hydrazine) [124]. No structural characterization was made, but the optical spectra were consistent with ZnS. The spectra show high transmission and low absorption in the suprabandgap region, along with low specular reflectance. Some samples, which had been deposited for a relatively long time, exhibited strong absorption (and somewhat increased reflectance) in the IR region, which could be ascribed to free carriers. The bandgap, at between 3.7 and 3.8 eV, was slightly higher than the bulk value for sphalerite ZnS. Some room-temperature depositions were also carried out from a thiourea/ammonia/ammonium sulphate/hydrazine bath. Again, high transmission in the suprabandgap region and a bandgap of 3.75 eV were measured. A bath without ammonia, using triethanolamine/hydrazine/thiourea and added NaOH to a pH ca. 12.3, was described [125]. Contrary to most other depoCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

sitions, the ZnS had the wurtzite structure, measured by XRD. Optical spectroscopy gave a bandgap of 3.68 eV Transmission above the bandgap was ca. 30%, implying a large amount of scattering (specular reflection would not be as high as 70%). The resistivity was ca. 106 -cm, and the films were n-type. Thioacetamide has also been used to deposit ZnS. In this case, depositions could be carried out under acid or neutral conditions as well as alkaline ones. Three different studies were carried out using thioacetamide. In one [126], two different baths were investigated—one with and the other without ammonia (the latter was probably slightly acidic and certainly had a considerably lower pH than the ammonia bath). Zincblende ZnS was obtained from both baths. The crystal size (TEM) was 6–8 nm (ammonia bath at 90°C); from the ammonia-free bath, the crystal size was not given, but from the sharpness of the ED pattern, it is probably larger. The bandgap (temperature independent) for films deposited from the ammonia-free bath was 3.6 eV (
literature value). For films deposited from the ammonia bath, values of bandgap up to 4.05 eV (5°C deposition), which depended on deposition temperature (3.8 eV at 90°C; 3.95 eV at 30°C), were measured. Urea was used in the ammonia-free bath to improve adherence, although it is not known why this affected adherence. ZnS was grown from a thioacetamide bath using triethanolamine to complex the Zn2 and buffered by NH3/NH 4 to a pH of 10 [127]. Glass coated with these ZnS films was found to be a good substrate for other CD semiconductors that might otherwise exhibit poor adhesion, and this was the purpose of this study. It was noted from optical transmission that for films no thicker than ca. 0.15 m, scattering was negligible, but it became increasingly marked for thicker films. XPS depth-sputtered analysis of these films indicated that the ZnS–glass interface was not sharp, and it was suggested that Zn diffuses into the glass to some extent, explaining the good adhesion of the films [128]. A more complete characterization of these films was subsequently carried out [129]. No XRD pattern was found for the as-deposited films, implying amorphous or very small crystal size. Even after a short anneal at 500°C, a crystal dimension as low as 13 nm (depending on crystal orientation) was measured by XRD, implying a much smaller size in the as-deposited film. In line with the crystal size, optical absorption showed bandgaps between 3.85 and 3.95 eV, higher than the bulk value, presumably due to size quantization. These films were not photoconductive as deposited but became so on air-annealing with an optimum annealing temperature of 388°C, when the photosensitivity increased to 104. The film resistivity decreased from ca. 5  107 -cm (as deposited) to ca. 104 -cm after air-annealing at 400–500°C. The third study used a solution of ZnCl2 and thioacetamide at a pH of 2.45 [130]. Films were deposited on ITO/glass at 70°C. In spite of the acid conditions, which precluded formation of Zn(OH)2, a precipitate formed in the solution in

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parallel to film formation. The crystal size for both film and precipitate was ca. 3 nm. The energy of activation was measured for both the precipitate and the film formation (usually, it is only measured for film formation) and found to be the same for both, with a value of 9.5 kCal/mole (40 kJ/mole); this value is intermediate between that usually found for CdS deposition, where the rate-determining step is a chemical reaction, and for ZnS deposited from an alkaline thiourea bath [120] (21 kJ/mole; see Sec. 3.5). Whether this difference is due to the expected different mechanisms (ion-by-ion probable for the acidic bath and cluster mechanism for the alkaline one) or to differences in thiourea and thioacetamide decomposition is unknown at present. Thiosulphate has been used in an acid bath (pH 2–4) at 85°C [131]. The films were poor (not uniform, powdery, and nonadherent). The only other characterizations given were the bandgap (3.4 eV) and resistivity (106–107 -cm). Thiosulphate has also been used to deposit ZnS using a photochemical reaction, in the same manner as used previously for CdS (see Sec. 4.1.6.7 for a description of the proposed mechanism). In brief, UV light of wavelength shorter than 300 nm (from a Hg lamp) forms hydrated electrons and elemental S from thiosulphate solution, and this reacts with Zn2 to give ZnS [132]. The ZnS was highly nonstoichiometric with excess Zn. This is not surprising, since ZnS itself is sensitive to UV light, with the formation of elemental Zn due to the strong reducing action of photogenerated electrons. While referring to precipitation in solution and not CD of films, it is of interest to mention that N-allylthiourea was used as a sulphur source in a Zn2/ammonia bath [133]. Pure ZnS (as detected by XRD) was obtained at a pH of ca. 11.0 at 90°C. At lower pH values (and at 80°C), only ZnO was obtained; at higher values, a mixture of ZnO and ZnS was formed. The ZnS was the wurtzite form in all cases. The ZnS crystal size was ca. 5 nm at pH
11.0 and slightly smaller (ca. 4 nm) at higher pH values, which gave a mixed phase of ZnS and ZnO. The ZnO crystal size was much larger (ca. 200 nm). The ZnS fraction increased as the ammonia concentration (at constant pH) increased (lower [Zn2 ]) or as the pH decreased (at constant ammonia concentration) from 13 to 11 (lower [OH]). Such experiments should help in choosing optimum conditions for CD. A variant of CD has been described where ZnS was precipitated as a gel by adding concentrated S2 to a concentrated solution of a Cd salt. The pH was then reduced by HNO3 to a value between 5 and 7, when a semitransparent sol formed. Heating this sol between 100 and 200°C (in an autoclave) resulted in the formation of zincblende ZnS films [134]. If the S2 was not in excess (twice the Cd concentration), some ZnO, together with some wurtzite ZnS, also formed. Addition of a CuCl2 solution to the pH-adjusted sol and heating at 140°C resulted in Cudoped ZnS (particle size 60  10 nm) that showed several photoluminescence bands [135].

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4.5 ZnSe 4.5.1 Introduction In contrast to CdSe, which in most cases was prepared using selenosulphate as a Se source, ZnSe was deposited in the majority of studies using selenourea. One report was given where ZnSe (among other selenides) was deposited from an ammoniacal solution at 20°C [99] using selenosemicarbazide as a Se source. Other than that the ZnSe films were specular in nature, no details on their properties were given. There are two reported studies where selenosulphate was used. In one [136], it was noted that hydrazine was essential to obtain film formation; for ZnS, while it was usually preferable to include hydrazine in the deposition solution, it was apparently not essential and films could be made without it. This difference probably reflects the higher solubility of ZnSe compared to CdSe, the lower solubility of the hydroxide of Zn compared to that of Cd and either the faster formation of selenide ion from selenourea compared to selenosulphate or a different deposition mechanism for the different Se sources. The other study reportedly obtained ZnSe without using hydrazine (triethanolamine and ammonia were used together as complexants) [137]. Unusually high-concentration selenosulphate solution (0.8 M) was used. No structural or analytical characterization was carried out on the films, although the bandgap measured from the optical absorption was in the correct region for ZnSe. Hydrazine was also used in all the studies employing selenourea (but not where N,N-dimethylselenourea was used—see later). It is not clear to what degree hydrazine was essential in these studies. In an early theoretical study of ZnSe deposition from Zn–ammine/selenourea baths, the use of hydrazine is not mentioned [138]. On the other hand, in a patent by the same authors describing ZnSe deposition, the use of hydrazine hydrate was the main issue in the claims [138a]. In most of these studies, ammonia was used as a complexant. The conditions varied somewhat from study to study, but it appears, either explicitly or by resorting to educated guesses, that in all cases the deposition occurred by a hydroxide cluster mechanism. The pH was usually ca. 11.5, and deposition temperatures varied from 50°C to 80°C. The exception was the deposition described in the patent by Kitaev and Sokolova, where only ZnCl2, hydrazine hydrate, and selenourea ( acetic acid to acidify the very alkaline solution to a pH of ca. 9) were used [138a]. It was claimed that deposition could be obtained over a temperature range of 10–70°C. No characterization or properties of these particular films were given.

4.5.2 Depositions Using Selenourea 4.5.2.1

Structure

XRD of the films gave very broad, ill-defined peaks [139,140]. In one study, the precipitated powder from the deposition solution was found to be wurtzite, while

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grazing-angle XRD (more sensitive for thin films) gave a poorly defined, very broad spectrum [141]. In another study [140], while no normal XRD pattern was obtained, grazing-angle XRD showed a mixture of wurtzite and zincblende ZnSe together with a little ZnO. TEM/ED showed zincblende ZnSe in one study [139] and predominantly wurtzite in another [140]. Crystal size was measured in two cases by TEM: 2–2.5 nm [139] and ca. 10 nm [140].

4.5.2.2

Composition

Elemental analyses found such films to be Zn rich (with respect to Se) [139,142]. The films were oxygen rich, and it is probable that, as was often found with ZnS, the films are a mixture of ZnSe and Zn(OH)2, together with a little ZnO in some cases [140]. An XPS depth profile study found the surface of the films to be more stoichiometric (although still Zn rich) than near the substrate [143].

4.5.2.3

Optical Spectroscopy (Bandgap)

Optical spectroscopy was used to measure bandgap (bulk value ca. 2.6 eV). Such values could be correlated with the crystal size, with values of 2.7 eV [140] and 2.9 eV [139,142], the latter due to size quantization in very small crystals (ca. 2.5 nm). One study noted that the film was a transparent white (colorless) turning to orange, presumably due to Se formation on exposure of excess selenourea to air [141]. The white film, when annealed in air at 300°C, showed a broad ZnSe diffraction peak (by 400°C, the film was converted to ZnO). This, together with the very broad peaks (of both the 300°C annealed sample and of the powder precipitated in the solution) suggested that the white color was due to size quantization (white implies a bandgap of 3 eV—bulk ZnSe is pale yellow). The presence of oxide/hydroxide in these films would result in a weakening of the yellow color of pure ZnSe but would not be expected to change the spectral position.

4.5.2.4

Electrical Resistivity

Electrical resistivity was measured in only one case, with a value of 2  108 -cm [140].

4.5.3 Deposition Using N,N-Dimethylselenourea A modification of the general procedure used a substituted selenourea (N,Ndimethylselenourea) instead of selenourea and no hydrazine. Citrate was used, together with ammonia, presumably as a co-complexant, although it may also have functioned as a mild reducing agent, and the deposition was carried out at a relatively low temperature of 50°C [144]. No XRD pattern of the films was found, but the powder precipitated in the solution was zincblende phase with very broad peaks (3 nm coherence length). In contrast to other measurements, where the Zn:Se ratio was greater than 1, the

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Zn:Se ratio of these films was less than 1 (0.85). As with the other films, there were considerable amounts of oxygen (this was measured by XPS and includes normal surface-adsorbed oxygen). A value for the bandgap of 2.63 eV was measured from the optical spectrum. This (very slightly) higher value than the bulk value was explained by the small crystal size. (Size quantization for a crystal size of 3 nm would be expected to give a larger blue shift than this). The films were ntype with a resistivity of ca. 3  107 -cm, within an order of magnitude of the value measured from a selenourea bath [140].

4.5.4 Deposition Using Selenosulphate There is one reported brief study of ZnSe deposition using selenosulphate [136]. Considering the understandable preference for using selenosulphate rather than selenourea for CdSe depositions in most cases (selenosulphate is more stable and simpler to make and to handle), it is surprising that this is not also the case for ZnSe. It is possible that this is simply a case of “inertia”; i.e., most researchers follow essentially the same recipe (although the selenosulphate technique described here predated the other studies). Hydrazine was used (in its absence, no ZnSe was formed). A mixture of both triethanolamine and ammonia, together with NaOH, was used as complexant/pH adjuster; no explanation of the solution composition was given. The deposition was carried out at 100°C (the highest deposition temperature of all these processes). The films were light yellow (characteristic of ZnSe), which changed to a light reddish color. The reddish color might have been due to formation of Se under the conditions of the deposition, although Se would be expected to dissolve in the selenosulphate. X-ray diffraction showed the films (presumably the reddish ones) to be wurtzite ZnSe. A bandgap of 2.62 eV was measured from the optical spectra. The films were n-type with a relatively low resistivity of ca. 104 -cm.

4.5.5 Deposition Using Selenosemicarbazide Selenosemicarbazide was used in one study to make, among other materials, ZnSe films and precipitates from aqueous ammoniacal solutions at 20°C [99]. The films were specular, but no further information was provided.

4.5.6 Miscellaneous Methods A novel CD technique used metallic Al to reduce Se. Elemental Se and metallic Al foil, together with ZnCl2, were dissolved in NaOH solution and heated to 80°C in an autoclave with substrates of Teflon or alumina [145]. (Metal chalcogenide films have been chemically deposited onto Teflon in a number of reports.) It is notable that the films deposited onto alumina were reported to be 1.8 m thick— much thicker than ZnSe CD from other solutions. The films were sphalerite ZnSe

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with a crystal size greater than 10 nm, estimated from XRD peak widths. The Zn:Se ratio was essentially unity, in contrast to other reports, where the films were invariably nonstoichiometric. (At deposition temperatures of 125°C and above, ZnO was also formed.) The mechanism of this process was suggested to be: 3Se  6OH  → 2Se2  SeO 2 3  3H2O 2Al 

SeO 2 3



2Al(OH) 4

 3H2O  2OH →  Se  Se2  Zn(OH) 2 → ZnSe  4OH 4

2

(4.6) (4.7) (4.8)

Since the presence of the Al was essential to obtain films, Reaction (47) or other reactions, possibly involving reduction of Se species with nascent hydrogen (formed by the dissolution of Al in the NaOH solution), were probably the important steps in the deposition. Another method of ZnSe deposition is not true CD but is related and worthy of mention. H2Se, prepared by decomposition of CdSe with HCl, was passed over an aqueous solution of zinc acetate [146,147]. The thin ZnSe film (ca. 50–100 nm thick) formed at the gas–solution interface could be lifted up and placed on any desired substrate. The films deposited at 80°C were found to be zincblende ZnSe by both XRD and ED (no XRD pattern was observed for films deposited at 2°C). The optical bandgap was 2.62 eV, and resistivity was 107 -cm.

4.6 HgS HgS possesses a very low value of Ksp (6  1053) and therefore is expected to be relatively simple to deposit. In fact, apart from some relatively early literature on ternary mercury sulphides with lead [148,149] and cadmium [150], which will be discussed in Chapter 8, only three separate studies on CD HgS have been found. Perakh and Ginsburg [151] deposited HgS films using two different techniques. One was a standard CD method using thiourea and HgCl2 complexed with iodide (iodide is a strong complexant for Hg2) in an alkaline solution. The other technique was simply reaction of a low concentration (ca. 2 mM) of HgCl2 with (at least three times the Hg concentration) Na2S solution, which precipitates HgS as a colloid. HgS deposited slowly over many hours. It is interesting that while some film deposition is expected by this second method, thicknesses of at least 0.7 m were obtained—much thicker than would intuitively be expected. The temperature dependence of the growth (at least for the thiourea method) was complicated and depended on other parameters. Optimum temperatures were around room temperature; temperatures higher than 25°C resulted in rough films. The absorption spectra of both types of films were rather strange—a gradual absorption onset at somewhat less than 700 nm (ca. 1.8 eV) and a sharp onset at 400 nm (3.1 eV). HgS occurs in (at least) two forms: red (distorted rocksalt, cinnabar—the stable form at normal temperatures) with a bandgap of ca. 2.0 eV and a black form

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(zincblende) with a bandgap of ca. 0.5 eV (although large variations of this value have been reported—see Hg ternaries in Chap. 8). On precipitation of a solution of Hg2 ions with sulphide solution, a black precipitate is formed (it is not clear if this is the same as the zincblende form), which eventually should convert to the more stable red form, although this change might take a long time. From the absorption spectra of the films given in this study, with no absorption at wavelengths greater than 700 nm, the films were clearly not of the black form. The sharp onset at 3.1 eV could conceivably be due to absorption by very small crystals However, it is more likely that it is a higher transition of HgS, since HgS that is not quantized also shows this feature (see later). The gradual onset from ca. 1.8 nm could be due to absorption in a distribution of larger sizes also containing a low concentration of “black” heavily quantized HgS (low concentration since the absorption below 2 eV is very weak). It must be stated that no structural characterization of these films was made; however, it is unlikely that anything other than HgS was formed under the deposition conditions. The other studies used thiosulphate as a sulphur source. One was carried out in a simple mixture of HgCl2 and Na2S2O3, presumably under acidic conditions [152]. Crystal size measured by XRD was reported to range from 3 nm to 8 nm over a temperature range of 0–85°C (film thickness also varied over this temperature range from 50 to 180 nm), and this size regime was confirmed by TEM. The bandgap estimated from optical spectra varied monotonically as a function of deposition temperature from 2.3–2.4 eV at 0°C to ca. 1.9 eV at 85°C, due to size quantization. The resistivity of the films varied from 104 to 103 -cm over the same temperature range and generally decreased strongly with increasing measurement temperature to low values (very approximately 100 -cm) at temperatures over 150°C. A similar deposition, only carried out under alkaline conditions (with added ammonia) at pH 11 and at 65°C, was described [153]. According to this study, simple mixing of a mercuric salt solution with thiosulphate results in immediate formation of a black precipitate of HgS. By first treating the Hg2 solution with aqueous ammonia, a white precipitate formed by the following reaction: Hg(NO3)2  2NH3 → (NH2Hg)NO3  NH4NO3

(4.9)

This precipitate dissolves in thiosulphate to form a thiosulphate complex, which, in common with other metal thiosulphate complexes, decomposes when heated to the metal sulphide (see Sec. 3.3.3). Besides direct decomposition of the thiosulphate complex, another possibility suggested in this study is formation of sulphide ion by alkaline hydrolysis of thiosulphate [Eqs. (3.20) to (3.24)] and reaction with Hg2 to form HgS. The substrates were glass precoated with a very thin film of CD PbS (presumably this improved adherence and/or homogeneity). The films deposited by this process were golden yellow if thin (ca. 100 nm) and became red (the normal cinnabar color) if thick (ca. 500 nm). The terminal

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thickness of the films was 80–90 nm; thicker films were obtained by repeated depositions. X-ray diffraction showed that the deposit (both as a film and the precipitated material) was cinnabar HgS. The diffraction peaks were sharp (crystal size at least 20 nm, possibly much greater). Optical transmission spectra showed a sharp absorption at ca. 400 nm (3.1 eV) and a weaker one at ⬇600 nm (2.0–2.1 eV), the latter probably corresponding to the HgS bandgap, since little or no size quantization should be observed in these films due to the relatively large crystal size. However, there was a gradual loss in transmission to beyond 800 nm, which was more pronounced for thicker films. It is not clear to what extent this is due to scattering (the thicker films were reported to be less reflecting than the thinner ones, implying increased scattering) or to absorption. The resistivity of the films was greater than 104 -cm (calculated from a reported sheet resistivity of 109  and an assumed thickness of 100 nm).

4.7 HgSe Three studies in total have been made on CD HgSe. In the earliest, selenosemicarbazide was used as a Se source and the Hg was complexed with iodide [99]. The deposition was carried out at 20°C. No details of the films were given other than that they were specular. In the other two studies, selenosulphate was used. In one, a formamide complex of Hg, made by dissolving HgO in formamide, was used [154]. The solution was made ca. 0.5 M in NaOH, and a trace of polyvinyl pyrollidone was added. The deposition was carried out at room temperature. The polyvinyl pyrollidone slowed the deposition somewhat and apparently improved film uniformity and adherence as well as slightly increased terminal thickness (500 nm). It was noted that films were not obtained with the usual complexants, such as ammonia, triethanolamine, and cyanide. It is not mentioned in which way these complexants were unsuitable; ammonia and triethanolamine might be too weak, resulting in immediate precipitation in solution. Also, addition of ammonia to some mercuric salts tends to lead to precipitation of insoluble products. Cyanide, however, is a very strong complexant and would be expected to control such bulk precipitation better than formamide. Iodide, a strong complex for Hg2 (and successfully used to deposit HgS, as described earlier), resulted in film deposition but with poor reproducibility. No XRD pattern was found for the films, and on this basis they were believed to consist of amorphous HgSe. Based on more recent XRD studies of nanocrystalline materials, the lack of an XRD pattern was likely due to very small crystal size (supported by the increased bandgap; see later). Annealing at 200°C “crystallized” the HgSe to an extent that it was clearly identified by XRD. Optical spectroscopy gave a bandgap value of 1.42 eV. Bulk HgSe is a semimetal with

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a bandgap of -0.15 eV (i.e., overlapping valence and conduction bands). This suggests very strong size quantization (see HgSe in Chap. 10). The room-temperature resistivity was ca. 300 -cm. In the other study [155], ammonia-complexed Hg(NO3)2 was mixed with the selenosulphate solution. As for the corresponding HgS deposition, a white precipitate formed on addition of ammonia to the Hg(NO3)2 [Eq. (4.9)]. This precipitate dissolved partly in the excess ammonia used, due to formation of various ammine complexes, and completely when the selenosulphate solution was added, due to additional formation of selenosulphate (and maybe sulphite from the excess sulphite in the selenosulphite solution) complexes. It is likely that mixed ammineselenosulphate/sulphite complexes were formed. The deposition was carried out on polyester substrates (the transparencies used in overhead projectors) at 10°C. Deposition occurred over several hours to a terminal thickness of ca. 250 nm. Bulk precipitation occurred in parallel with the deposition, suggesting that the cluster mechanism was dominant. X-ray diffraction of the powder precipitated in solution confirmed it to be HgSe (sphalerite). The spectrum of the film showed a strong (111) peak and virtually no other reflection, suggesting a high degree of texture of these films. From the peak broadening, a crystal size of 7.7 nm was calculated. From optical spectra, a bandgap of 2.5 eV was estimated (based on an indirect gap), and this increase from the negative bandgap of bulk HgSe (see earlier) was attributed to size quantization. The sheet resistivity was measured to be 13 k-cm2. Although the film thickness was not given, based on a thickness of 250 nm, this translates into a specific resistivity of 1 -cm. Annealing the films at low temperatures (ca. 100°C) results in a decrease in resistivity up to as much as an order of magnitude (the crystal size, as measured by XRD, increases only slightly via this treatment).

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PK Nair, MTS Nair, J Campos. In: CM Lampert, ed. Proc. SPIE, San Diego, CA, 1987, Vol. 823, p 256. PK Nair, MTS Nair, J Campos. Sol. Energy Mater. 15:441, 1987. IO Oladeji, L Chow, JR Liu, WK Chu, ANP Bustamante, C Fredricksen, AF Schulte. Thin Solid Films 359:154, 2000. L Pintilie, E Pentia, I Pintilie, D Petre. Mater. Sci. Eng. B 44:403, 1997. M Froment, MC Bernard, R Cortes, B Mokili, D Lincot. J. Electrochem. Soc. 142:2642, 1995. AG Shikalgar, SH Pawar. Sol. State Commun. 32:361, 1979. AG Shikalgar, SH Pawar. Thin Solid Films 61:313, 1979. CD Lokhande, SH Pawar. Sol. State Commun. 44:1137, 1982. JA Akintunde. J. Mater. Sci.-Mater. Electron. 11:503, 2000. AI Oliva, R Castro-Rodriguez, O Ceh, P Bartolo-Perez, F Caballero-Briones, V Sosa. Appl. Surf. Sci. 148:42, 1999. D Bhattacharyya, MJ Carter. Thin Solid Films 288:176, 1996. D Lincot, R Ortega-Borges, M Froment. Appl. Phys. Lett. 64:569, 1994. R Cortes, M Froment, B Mokili, D Lincot. Philos. Mag. Letts. 73:209, 1996. D Lincot, B Mokili, R Cortes, M Froment. Microsc. Microanal. Microstruct. 7:217, 1996. MJ Furlong, D Lincot, M Froment, R Cortés, AN Tiwari, M Krejci, H Zogg. In: 14th ECPV Solar Energy Conf. Barcelona, Spain, 1997, p 1291. MJ Furlong, M Froment, MC Bernard, R Cortés, AN Tiwari, M Krejci, H Zogg, D Lincot. J. Crystal Growth 193:114, 1998. YK Hwang, SY Woo, JH Lee, DY Jung, YU Kwon. Chem. Mat. 12:2059, 2000. RL Call, NK Jaber, K Seshan, JR Whyte Jr. Sol. Energy Mater. 2:373, 1980. GA Kitaev, SG Mokrushin, AA Uritskaya. Kolloid Zh. 27:51, 1965. R Ortega-Borges, D Lincot. J. Electrochem. Soc. 140:3464, 1993. V Polpescu, EM Pica, I Pop, R Grecu. Thin Solid Films 349:67, 1999. K Yamaguchi, T Yoshida, T Sugiura, H Minoura. J. Phys. Chem. B 102:9677, 1998. K Yamaguchi, T Yoshida, N Yasufuku, T Sugiura, H Minoura. Electrochem. 67:1168, 1999. PK Nair, MTS Nair. Sol. Energy Mater. 15:441, 1987. V Weiss, AA Friesem, A Peled. Thin Solid Films 218:193, 1992. F Goto, M Ichimura, E Arai. Jpn. J. Appl. Phys. 2 36:L1146, 1997. M Ichimura, F Goto, E Arai. J. Electrochem. Soc. 146:1028, 1999. JGea Vázquez-Luna. J. Crystal Growth 187:380, 1998. M Gros-Jean, R Herino, D Lincot. J. Electrochem. Soc. 145:2448, 1998. DS Boyle, A Bayer, MR Heinrich, O Robbe, P O’Brien. Thin Solid Films 361:150, 2000. D Hariskos, M Powalla, N Chevaldonnet, D Lincot, A Schindler, B Dimmler. Thin Solid Films 387:179, 2001. A Kylner, J Lindgren, L Stolt. J. Electrochem. Soc. 143:2662, 1996. A Kylner. J. Electrochem. Soc. 146:1816, 1999. A Kylner, M Wirde. Jpn. J. Appl. Phys. 1 36:2167, 1997. DW Niles, G Herdt, M AlJassim. J. Appl. Phys. 81:1978, 1997.

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80. Y Hashimoto, T Nakanishi, T Andoh, K Ito. Jpn. J. Appl. Phys. 2 34:L382, 1995. 81. T Nakada, H Fukuda, A Kunioka, S Niki. In: 13th ECPV Solar Energy Conf. Nice, France, 1995, p 1597. 82. R Jayakrishnan, JP Nair, BA Kuruvilla, SK Kulkarni, RK Pandey. J. Mater. Sci. 7:193, 1996. 83. GA Kitaev, TS Terekhova. Zh. Neorg. Khim. 15:48, 1970. 84. RC Kainthla, DK Pandya, KL Chopra. J. Electrochem. Soc. 127:277, 1980. 85. RA Boudreau, RD Rauh. J. Electrochem. Soc. 130:513, 1983. 86. GA Kitaev, AA Uritskaya, LE Yatlova, TS Terekhova, TI Dzyuba. Zhurn. Neorg. Khim. 35:3072, 1990. 87. G Hodes, A Albu-Yaron, F Decker, P Motisuke. Phys. Rev. B 36:4215, 1987. 88. S Gorer, G Hodes, Y Sorek, R Reisfeld. Mater. Lett. 31:209, 1997. 89. F Trojanek, R Cingolani, D Cannoletta, D Mikes, P Nemec, E Uhlirova, J Rohovec, P Maly. J. Cryst. Growth 209:695, 2000. 90. G Hodes. Isr. J. Chem. 33:95, 1993. 91. R Garuthara, G Levine. J. Appl. Phys. 80:401, 1996. 92. P Nemec, D Mikes, J Rohovec, E Uhlirova, F Trojanek, P Maly. Mater. Sci. Eng. B 69:500, 2000. 93. M Perakh, A Peled. Thin Solid Films 50:293, 1978. 94. H Cachet, H Essaaidi, M Froment, G Maurin. J. Electroanal. Chem. 396:175, 1995. 95. KC Mandal, O Savadogo. J. Mater. Sci. Lett. 10:1446, 1991. 96. R Lozada-Morales, M Rubin-Falfan, O Portillo-Moreno, et al. J. Electrochem. Soc. 146:2546, 1999. 97. MTS Nair, PK Nair, H Pathirana, RA Zingaro, EA Meyers. J. Electrochem. Soc. 140:2987, 1993. 98. BK Rai, HD Bist, RS Katiyar, MTS Nair, PK Nair, A Mannivannan. J. Appl. Phys. 82:1310, 1997. 99. AA Velykanov, EK Ostrovskaya, NP Garina, VA Turacova, AA Tchurkan. Ukr. Chim. Zh. 49:764, 1983. 100. H Cachet, R Cortés, M Froment, G Maurin, N Shramchenko. J. Electrochem. Soc. 144:3583, 1997. 101. E Lifshitz, I Dag, I Litvin, G Hodes, S Gorer, R Reisfeld, M Zelner, H Minti. Chem. Phys. Lett. 188, 1998. 102. XC Ai, R Jin, CB Ge, JJ Wang, YH Zou, XW Zhou, XR Xiao. J. Chem. Phys. 106:3387, 1997. 103. P Maly, J Kudrna, F Trojanek, D Mikes, P Nemec, AC Maciel, JF Ryan. Appl. Phys. Lett. 77:2352, 2000. 104. K Rajeshwar, L Thompson, P Singh, RC Kainthla, KL Chopra. J. Electrochem. Soc. 128:1744, 1981. 105. RC Kainthla, DK Pandya, KL Chopra. Sol. State Electron. 25:73, 1982. 106. VM Garcia, MTS Nair, PK Nair, RA Zingaro. Semicond. Sci. Technol. 11:427, 1996. 107. MTS Nair, PK Nair, RA Zingaro, EA Meyers. J. Appl. Phys. 74:1879, 1993. 108. G Hodes, IDJ Howell, LM Peter. In: Photochemical and Photoelectrochemical Conversion and Storage of Solar Energy. ZW Tian, Y Cao, eds. Int. Acad. Publishers, Beijing, China. 1993, p 331.

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L Kronik, L Burstein, M Leibovitch, Y Shapira, D Gal, E Moons, J Beier, G Hodes, D Cahen, D Hariskos, R Klenk, HW Schock. Appl. Phys. Lett. 67:1405, 1995. E Lifshitz, I Dag, I Litvin, G Hodes. J. Phys. Chem. B 102:9245, 1998. B Alperson, I Rubinstein, G Hodes. Phys. Rev. B 6308:1303, 2001. GK Padam, SK Gupta. Appl. Phys. Lett. 53:865, 1988. RW Buckley. In: 11th ECPV Solar Energy Conf. Montreux, Switzerland, 1992, p 962. VB Patil, PD More, DS Sutrave, GS Shahane, RN Mulik, LP Deshmukh. Mater. Chem. Phys. 65:282, 2000. AV Kalyakina, RI Pelyukpashidi. Tr. Khim. Met. Inst. Akad. Nauk. Kaz. SSR 17:114, 1973. M Sotelo-Lerma, RA Zingaro, SJ Castillo. J. Organomet. Chem. 623:81, 2001. P O’Brien, J McAleese. J. Mater. Chem. 8:2309, 1998. B Mokili, M Froment, D Lincot. J. de Phys. IV 5:261, 1995. P O’Brien, DJ Otway, D Smith-Boyle. Thin Solid Films 361:17, 2000. JM Doña, J Herrero. J. Electrochem. Soc. 141:205, 1994. IO Oladeji, L Chow. Thin Solid Films 339:148, 1999. B Mokili, Y Charreire, R Cortes, D Lincot. Thin Solid Films 288:21, 1996. J McAleese, P O’Brien. In Mater. Res. Soc. Symp. Proc., 1998; Vol. 485; p 255. IC Ndukwe. Sol. Energy Mater. Sol. Cells 40:123, 1996. S Biswas, P Pramanik, PK Basu. Mater. Lett. 4:81, 1986. R Ortega Borges, D Lincot, J Vedel. In: 11th ECPV Solar Energy Conf., Montreux, Switzerland, 1992, p 862. PK Nair, MTS Nair. Semicond. Sci. Technol. 7:239, 1992. PK Nair, MTS Nair, O Gomezdaza, RA Zingaro. J. Electrochem. Soc. 140:1085, 1993. OL Arenas, MTS Nair, PK Nair. Semicond. Sci. Technol. 12:1323, 1997. K Yamaguchi, T Yoshida, T Sugiura, H Minoura. J. Mater. Res. 13:917, 1998. CD Lokhande. Mater. Chem. Phys. 28:145, 1991. M Ichimura, F Goto, Y Ono, E Arai. J. Cryst. Growth 199:308, 1999. GA Kitaev, AA Uritskaya, LE Yatlova, VR Mirolyubov. Russ. J. Appl. Chem. 67:1415, 1994. QW Chen, YT Qian, ZY Chen, L Shi, XG Li, GE Zhou, YH Zhang. Thin Solid Films 272:1, 1996. QW Chen, XG Li, Y Qian, JS Zhu, G Zhou, WP Zhang, YH Zhang. Appl. Phys. Lett. 68:3582, 1996. P Pramanik, S Biswas. J. Electrochem. Soc. 133:350, 1986. RK Nkum, AA Adimado, H Totoe. Mater. Sci. Eng. B 55:102, 1998. GA Kitaev, TP Sokolova. Russ. J. Inorg. Chem. 15:167, 1970. GA Kitaev, TP Sokolova. USSR Patent 356, 319 1972. CD Lokhande, PS Patil, A Ennaoui, H Tributsch. Appl. Surf. Sci. 123:294, 1998. JM Doña, J Herrero. J. Electrochem. Soc. 142:764, 1995. AM Chaparro, MA Martinez, C Guillen, R Bayon, MT Gutierrez, J Herrero. Thin Solid Films 361:177, 2000. A Ennaoui, M Weber, M Saad, W Harneit, MC Lux-Steiner, F Karg. Thin Solid Films 361:450, 2000.

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143. 144. 145. 146. 147. 148. 149. 150. 151. 152. 153. 154. 155.

J Herrero, MT Gutierrez, C Guillen, JM Dona, MA Martinez, AM Chaparro, R Bayon. Thin Solid Films 361:28, 2000. CA Estrada, PK Nair, MTS Nair, RA Zingaro, EA Meyers. J. Electrochem. Soc. 141:802, 1994. C Wang, XF Qian, WX Zhang, XM Zhang, Y Xie, YT Qian. Mat. Res. Bull. 34:1637, 1999. GN Chaudhari, SD Sathaye, P Singh, VJ Rao, V Manorama, RS Bhide, SV Bhoraskar. J. Mater. Sci. Lett. 11:1097, 1992. GN Chaudhari, S Manorama, VJ Rao. J. Phys. D:Appl. Phys. 25:862, 1992. NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Mat. Res. Bull. 11:1109, 1976. NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Thin Solid Films 42:383, 1977. M Skyllas-Kazacos, JF McCann, R Arruzza. Appl. Surf. Sci. 22/23:1091, 1985. M Perakh, H Ginsburg. Thin Solid Films 52:195, 1978. SS Kale, CD Lokhande. Mater. Chem. Phys. 59:242, 1999. M Najdoski, I Grozdanov, SK Dey, B Siracevska. J. Mater. Chem. 8:2213, 1998. P Pramanik, S Bhattacharya. Mat. Res. Bull. 24:945, 1989. B Pejova, M Najdoski, I Grozdanov, SK Dey. J. Mater. Chem. 11:2889, 1999.

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5 PbS and PbSe

5.1 HISTORICAL BACKGROUND PbS holds the honor of being the first reported compound to be deposited by CD. In 1869, Puscher described a “new and cheap process, without using dyes, to coat various metals with splendid lustrous colors” [1]. This involved deposition from a thiosulphate solution of lead acetate (and also, in the same paper, from Cu and Sb salts to give presumably corresponding sulphides). These shiny, colored coatings prompted further studies in this process, both to expand the process to other metal sulphides and to understand the process. These studies are discussed in Section 5.2.1. The common thiourea process for CD was also first used for PbS [2]. The thiourea method became the preferred one for the emerging development of PbS photoconductive cells as infrared (IR) detectors during the Second World War. Obviously much of this early work, carried out by German groups for military purposes, was secretive and was not published at the time. Later, photoconductive cells using CD PbS (which gave better cells than did the more conventional, evaporated PbS films [3]) became commercial as IR detectors, and, together with CD PbSe, they remain so until this day. An early description of such cells is given by Kicinski [4]. This application then became the driving force for CD studies, which were limited almost entirely to PbS and PbSe up to the start of the 1960s. Section 5.2.6 gives a brief overview of the operation of photoconductive cells in general and lead salt cells in particular.

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Since the expansion of CD to II–VI compounds in the 1960s and subsequently to other compounds, work on the IV–VI materials has not kept pace with this expansion, and there has been only a relatively limited amount of work since then, mainly on the use of PbS for solar control coatings. This chapter does not aspire to comprehensively cover the older literature, although it is intended that the important aspects of these studies be included. Previous reviews that deal with these materials (mostly reviews on infrared detectors) include Refs. 3 and 5–8.

5.2 PbS 5.2.1 Deposition Using Thiosulphate The original CD process involves decomposition of a thiosulphate solution of PbAc2 [1]. From the outset, it must be noted that the mechanism of this deposition has not been unambiguously elucidated up to the present time. There are two main possibilities for the reaction. 1. Decomposition of a thiosulphate complex of Pb, a possible reaction being 2 Pb(S2O3) 2 2  H2O → PbS  S2O 3  H2SO4

(5.1)

S2O 2 3

2. Decomposition of to give free sulphide ion, which then reacts with free Pb2. It has been suggested [9,10] that this can occur by reduction of elemental S, which forms in acidic solutions of S2O 2 3   S2O 2 3  H → S  HSO 3

(5.2)

by electrons formed in the half cell reaction:  2 2S2O 2 3 → S4O 6  2e

(5.3)

(An internal electrochemical mechanism was proposed long ago for deposition on certain metal substrates, since the rate of deposition sometimes depended on the nature of the substrate [11].) The standard potential of Reaction (5.3) is 0.08 V, considerably more positive than the reduction potential of S to S2 (0.45 V). Free sulphide, if formed, would be in a very low concentration, since it will be removed continually by precipitation of PbS; this will move the S reduction potential strongly positive according to the Nernst equation [Eq. (1.32)]. This positive shift will be even greater than normal because of the non-Nernstian behavior of the S2/S couple when [S]  [S2] (at least in alkaline solution) [12]. In opposition to this, the solubility of S in the (slightly acidic) aqueous solutions is very low, which will move the potential in the opposite direction. Add to this the very small concentration of S2 in acid solution [Eq. (1.15)], and it becomes clear that it is not trivial to estimate the feasibility of the formation of PbS by free sulphide. The non-Nernstian behavior of the sulphur-rich S2/S couple and the lack of knowledge of the solubility of free S in the deposition solution are the two factors that complicate what would have been a tractable thermodynamic calculation.

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The decomposition of metal thiosulphates to the corresponding sulphide is well known and is generally assumed to occur by breaking of the S–S bond in thiosulphate. Thus, it has been shown, by a radioactive tracer method, that if thiosulphate is prepared by dissolving radioactive S in sulphite solution: 2 S*  SO 2 3 D S*MSO 3

(5.4) 

and this solution is reacted with Ag to form a precipitate of Ag2S2O3, then the AgS formed on boiling this precipitate in water contains only labeled S, showing that the metal bonds only to the S which is not attached to oxygen [13]. Admittedly, even if the AgS was formed by free sulphide, the same result would be probable due to the lability of the SMS bond; however, such reactions have been studied for many metal cations, and it is invariably assumed (and this does not necessarily make the assumption correct) that the reaction occurs by simple fission of the SMS bond in the thiosulphate. Additionally, since this reaction occurs most readily for metals with a high affinity for S (Ag, Cu(I), Hg and Pb), there is no reason to expect that the MMS bond present in the thiosulphate (or complex thiosulphate) would be broken, as would be the case were the formation of metal sulphide to occur by free sulphide, although, as has already been pointed out, these same metal sulphides have a very low-solubility product and therefore very little sulphide is required to precipitate them. One potentially useful piece of information that can be explained more readily based on a free-sulphide generation comes from an early study on the formation of PbS by boiling Pb2 and thiosulphate in water, when it was found that PbS formed much more readily when excess thiosulphate was present [14]. To sum up, the mechanism of formation of PbS using thiosulphate is still not defined, either in general or even for any specific case. Some dedicated research to solve this question is clearly required. While a number of early studies described PbS (and other sulphides) formation from thiosulphate solutions [1,11,15,16], these early studies provided little characterization of the films other than the interference colors obtained (due to different thicknesses of the films). There appear to be only two modern studies that provide some general characterization of these PbS films. In one [10], films were deposited on glass at 80°C from a solution of PbAc2 and Na2S2O3 at a pH between 5 and 6. X-ray diffraction (XRD) showed sharp peaks of polycrystalline PbS. Although the thickness of the films was not given, they were clearly thick by CD standards, since the optical transmission over the range from 400 to 2000 nm was less than 1%. Electrical resistivity was measured to be ca. 105–106 -cm. In another study [17] (also by the same group on Cu2S-coated glass [18]), the conditions were similar (the pH range was between 4 and 6, and acetic acid was added, if necessary, to reach this pH range; the bath temperature was slightly lower—ca. 70°C—and the reactant concentrations were lower). The PbS films were thinner (0.1 m), as could be seen by the greater optical transmission (60% at 2.5 m,

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dropping gradually to ca. 7% at 500 nm and more rapidly at shorter wavelengths). Such semitransparent films are considered useful for solar control coatings (see Sec. 2.13). The films were p-type with a very low resistance (4 k/sq, equivalent to 4  102 -cm). The large difference between the resistivities in the two studies seems worthy of further investigation.

5.2.2 Deposition Using Thiourea With the exception of the few cases just described, all the modern (and most of the earlier) studies on PbS deposition were carried out using thiourea as a source of sulphide, as first described by Emerson-Reynolds in 1884 [2]. In the original study, the lead was present as a strongly alkaline solution of lead tartrate (probably a mixed tartrate–hydroxide complex of lead). When this solution was heated with a solution of thiourea, “at about 50° a specular layer forms at the bottom and sides of the vessel. When the beaker is thoroughly clean in the first instance, the adhesion is uniform and strong.” Most of the PbS formed as a precipitate. The competing effects of temperature were noted: At room temperature the deposition was very slow (days) and the film tended to be less even (more heterogeneous to the eye), while at higher temperatures thin (due to excessive precipitation) brownish films were obtained. Deposition was reported on a number of different materials (including porcelain and ebonite as well as iron and steel, although films on the metals had poorer adhesion). The same basic technique was used to study the deposition of PbS on different types of glass substrates. The quality of the film varied greatly from very poor, partially formed films that were not adherent (quartz and borosilicate glass), through occasionally good but irreproducible films (on sodalime glass), to homogeneous, adherent films on flint (lead containing) and Zn-containing (crown) glasses [19]. The good adherence on the Pb and Zn glasses was explained by formation of insoluble sulphides by these metals; such sulphides would form a good binding site for further deposition of PbS (presumably also for other compounds). Lead acetate was later usually employed as the Pb salt (e.g., Refs. 4 and 20). The thiourea acts not only as a source of S but also as a complexing agent, as shown by Kicinski. In the absence of thiourea, hydroxide can also complex Pb2, but a considerable excess is usually required in practice. Small amounts of lead oxide and hydroxide were detected in these films by XRD [4]. There does not seem to be a clear consensus as to the mechanism of PbS formation in these and similar studies. It is often stated that the formation of PbS occurs via decomposition of a Pb–thiourea complex species [4,21,22]. This was often based on the absence of any measurable concentration of sulphide on alkaline hydrolysis of thiourea. However, as discussed in Section 3.3.3.1, this is not a valid criterion for the absence of a sulphide-mediated reaction. Even today, it cannot be stated categorically which mechanism is operative or even dominant.

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A microscopic study of the PbS formed in these films deposited on glass at 50°C showed the formation of large (0.2–1 m) cubic crystals from relatively concentrated solutions (the films were specularly reflecting) and somewhat larger, but less well-defined crystals from more dilute solutions (resulting in more highly scattering films) [23]. Addition of very small amounts of CuSO4 to the dilute deposition solution resulted in specularly reflecting films with smaller, more evenly sized crystals. From XRD line broadening, the crystal size of PbS deposited by the method described by Kiciniski was estimated at ca. 70 nm [24]. Another report using films prepared by Kicinski on pyrex described a preferred orientation with the (001) faces parallel to the glass surface and a crystal size larger than 15 nm and in some cases ca. 50 nm [25].

5.2.3 Variations in Deposition This section deals only with thiourea-based baths. There is little variation in the thiosulphate baths that have been reported.

5.2.3.1 NH4OH (in place of NaOH or KOH) Ammonium hydroxide has also been used in place of NaOH or KOH [20,24,26]. In reported contrast to films deposited from alkali metal hydroxide, these films, prepared at or slightly above room temperature, were photoconductive (photosensitivity ca. 10) as deposited without need for air-annealing [26]. The crystal size of films deposited at different temperatures was measured (XRD line broadening) to be 10–15 nm (30°C), 17 nm (40°C), and 39 nm (50°C) [24,26]. The presence of strain in the crystals was inferred from the same XRD measurements [24].

5.2.3.2 Addition of Halides A comprehensive study has been made on the effect of added ammonium halides to deposition from solutions of citrate-complexed Pb2 containing NH4OH [27]. The effects of the ammonium halides were both strong and varied. The deposition rate decreased with increased halide concentration. This is expected, if only due to the lower pH of the buffered solution. Also, the retardation effect of the ammonium halide increased from Cl to I. This could be due to partial removal of Pb by the sparingly soluble halides (the iodide is the least soluble and therefore will most effectively remove the Pb). It is also possible that the more strongly adsorbed iodide retards growth due to adsorption and surface capping. This is supported by the gradual decrease in crystal size (measured by electron microscopy) from ca. 0.8 m (no halide) through 0.5 m (Cl), 0.3 m (Br), and 0.2 m (I), which parallels the increase in adsorption power of the halide ions from Cl to I. On the other hand, substitution of NH4I by KI or NaI does not have the same effect (the crystal size remains ca. 0.8 m). The films were preferentially (111) textured in the absence of ammonium halide, and this texture decreased as the ammonium halide

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concentration, as well as its atomic weight, increased, again as expected based on adsorption of halide; but again this texture-reducing effect did not occur (at least not to the same extent) with KI or NaI. Other effects of the different halides, given only for the ammonium halides (i.e., it was not known whether NaI or KI had a similar effect or not), were change in shape of the crystals from cubic (Cl) to spherical (I) (which could also be explained by strong adsorption of iodide on all crystal faces) and an energy of activation of the deposition varying from 38 (Cl) to 67 (I) kJ/mole; Both are characteristic of a chemical rate-determining step, but it is clear that there is an important difference in the two mechanisms. It seems that, while anion adsorption may play a role in these effects, the main role, at least in those effects where ammonium and alkali metal halides were compared, has another explanation. A possible one, in particular for the effect on deposition rate, is the lower pH of the ammonium-buffered solutions. Another property of the PbS films deposited from the ammonium-buffered solutions was the relatively high photosensitivity obtained for the as-deposited films, particularly from the Cl bath. Increases in photosensitivity of close to two orders of magnitude were obtained, compared to films deposited from solutions containing no ammonium halide. It was suggested that the halide compensated the PbS (low-temperature measurements showed quasi-intrinsic conduction in these films), resulting in high photosensitivity. Additionally, there is a general inverse relation between the photosensitivity and the grain size/degree of texturing. This relationship between grain size and photosensitivity is often seen and implies a major role of the grain surfaces in the photoconduction mechanism. Note that this is in contrast to photoconductivity in the II–VI semiconductors, where there was no obvious correlation between sensitivity and grain size (see Chap. 4). A related study, using KBr as an additive [28], found a modest increase in crystal size with increasing KBr concentration together with an increase in crystal height (up to 0.4 mM Br concentration followed by a subsequent decrease with further increase in KBr concentration) and a slight (200) preferred orientation. It was suggested that this crystal growth was due to retarded growth of small nuclei due to the complexing power of the bromide. Note that the pH of this bath was ca. 12, probably higher than the buffered ammonium bath described earlier. As with the ammonium halide baths, the photosensitivity of the resulting PbS films was found to increase with added KBr in the bath (no Br was found in the layers themselves, although very small “doping” concentrations might not be detected in the analyses), up to a maximum of 0.4 mM KBr, and then decreased, which correlates with the crystal height. The increase in photosensitivity was attributed to disorder at the grain boundaries. The study using ammonium halides also attributed a major role of the grain surface to the photoconductivity, but the relation between grain size and sensitivity is very different in the two studies.

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5.2.3.3 Oxidizing Agents in the Bath PbS films deposited from most basic baths are only weakly photoconducting. For maximizing photoconductivity response, air-annealing is usually carried out. By employing an oxidizing agent in the bath, films can be deposited that have relatively high response as deposited (explained by introduction of sensitizing centers in the PbS; see Sec. 5.2.6). One study [29] has described the affect of an oxidizing component in the bath on the sensitivity of the resulting films (ca. 1 m thick). Unfortunately, neither the bath conditions (probably a Pb2/hydroxide/thiourea bath) nor the identity of the oxidizing agent was revealed, other than that they were “variations of a commercially available material.” The films were all p-type with carrier density between 1016 and 1017/cc. The detectivity of a film deposited without the oxidant was lower (by orders of magnitude at room temperature, less at low temperatures) than for a film prepared with an excess of oxidant or a standard (presumed commercial) film. The absorption spectra of the films deposited with and without oxidant were also quite different (see Sec. 5.2.5). It is notable that this difference in absorption was not reflected in the photoconductivity response shape near the response onset; the samples behaved similarly with a photoconductivity onset of ca. 0.4 eV, although at higher photon energies the response of the oxidantfree film decreased strongly with increasing photon energy, in contrast to a modest decrease for the oxidant samples. It was also noted that the film with a high oxidant concentration appeared more porous (in SEM micrographs) than the oxidant-free film and that the particle size was ca. 70 nm (somewhat less in the oxidant-containing films than in the oxidant-free ones). A similar study on the effect of (the as-usual unspecified) oxidant on PbS films deposited from a thiourea/hydroxide bath found similar effects to the foregoing study, namely, better sensitivity, smaller particle size (from almost a micron in the absence of oxidant dropping to 0.2 m with a high oxidant concentration) and more porous with increasing oxidant concentration [30]. In this case, the sensitivity peaked at a certain oxidant concentration, followed by a decrease. Also, while the oxidant films were p-type, those deposited without the oxidant were n-type. The oxygen content was constant and independent of the oxidant concentration (as found also in the previous study [29]), and it was concluded that the origin of the active sensitizing centers was related not to the oxidizing agent itself but to the structure of the film. Apart from considerations of photoconductivity, one oxidant, H2O2, was shown to exert considerable influence on the crystallographic texture of PbS deposited from a PbAc2/NaOH solution with thiourea at close to room temperature [31]. While the (111) reflection was fairly dominant in the absence of H2O2, the (200) reflection became very dominant in its presence. The H2O2 was added, not with the other reactants, but some time after film formation became visible, and the degree of texturing reached a maximum at a certain time, after which it again decreased. No explanation for these effects was given.

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5.2.3.4 Miscellaneous Variations in Deposition Addition of Inhibitors. Addition of salts of Sb, Sn, or As to a basic PbAc2 /thiourea/OH bath resulted in a slower deposition, typically by 2–4 times [21]. These metal ions all form complex sulphides in aqueous solution, and it was believed that they caused “peptization” of growing nuclei; peptization converts large particles into colloidal ones. Probably what is meant is some sort of capping action due to adsorption of the retarding ions on the growing PbS crystals, preventing (more correctly, retarding) further growth. Illumination-Induced Growth. As for CdS, illumination has been found to increase the growth rate of CD PbS films [32]. The rate increase was attributed to a combination of heating by light absorption and the formation of charge carriers in the PbS film, the latter resulting in activation of the deposition. As for CD of CdSe under illumination, described in Chapter 4, such activation may result from either a photochemical reaction by the photogenerated electrons and/or holes on the PbS surface, as suggested earlier (the most likely scenario), and/or by a change in the electrical double layer at the surface of the PbS particles (adsorbed on the surface or as a colloid in the solution near the substrate), which might lower a potential barrier to adsorption and coalescence of the PbS colloids. Alkali-Metal-Free Solutions. Films of CD PbS are usually p-type as deposited. One early suggestion to explain this was that the alkali metal ions used in the deposition solution (as NaOH or KOH) act as a p-type dopant [33]. Based on this supposition, Bloem deposited PbS from a solution of PbAc2, hydrazine hydrate, and thiourea (free of Na or K). The as-deposited films were initially ntype but changed to p-type on exposure to air. Attempts to dope the films permanently n-type by adding trivalent ions to the deposition solution were unsuccessful. However, by depositing the films on a substrate coated with trivalent ions (such as Al, In, Ga), n-type behavior could be maintained for a considerable time. PbS p-n junctions were fabricated using this approach (see Chap. 9).

5.2.4 Substrate Effects (See also Sec. 5.2.2) 5.2.4.1 Epitaxial Deposition There have been a few reports on epitaxial deposition of PbS on various singlecrystal substrates. PbS (n-type) was epitaxially deposited on (111) Ge (5.4% mismatch) from a Pb(NO3)2 /KOH/thiourea solution at room temperature with (111) orientation [34] (although another study using apparently the same conditions found the deposit to be p-type and polycrystalline with some (100) preferred orientation [35]). From a similar solution (with addition of some ethanol), PbS was deposited on single-crystal CdS (ca. 6.6% mismatch) with varying degrees of epitaxy [36]. On the (0001) faces of CdS, the growth was (111) [(111) cubic corre-

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sponds to (0001) hexagonal] and on the (112 0) face, it was (220) oriented. The degree of epitaxy was moderate on the Cd (0001) face, better on the S (0001 ) face, and very high on the (112 0) face. Using the same reactants (but at higher concentrations), also at room temperature, epitaxial PbS growth was observed on (100) InP, (100) Ge, and also (although with rougher morphology) on (111) Ge [37]. In these experiments, the films, ca. 300 nm thick, were all p-type. The energy of activation for film formation in this study was 65 kJ/mole—similar to CdS deposition from thiourea solutions but different from the value of activation energy measured for ZnS deposition from a thiourea bath (see Sec. 4.4.1). Homogeneous precipitation also occurred gradually, but the films were removed before it interfered by adhesion of clusters (presumably before it became excessive). PbS films grown on single-crystal GaAs (lattice mismatch ca. 5%) were polycrystalline or only somewhat oriented [38], and those on Si (9% mismatch) were polycrystalline [39]. Some of these heterojunctions with CD PbS are discussed in Chapter 9.

5.2.4.2

Deposition on Monolayers

PbS has been deposited from hydroxide/thiourea solutions onto Au and thiol monolayer–coated Au substrates [40]. The quality and texturing of the films varied according to the specific substrate used and also on the hydroxide concentration. For dilute OH solutions (0.1 M OH; 0.01 M Pb2), good deposition was obtained only on some surfaces, while for more concentrated OH (0.25 M OH; 0.01 M Pb2), the deposition was not very surface dependent. Various degrees of (111) or (100) texture were obtained, depending on the substrate. On bare (111) Au, a (100) texture was obtained. It was noted that a high degree of texture was obtained on some of the monolayers, which themselves were poorly ordered. The crystal size varied according to both film thickness and other deposition conditions. For very thin films, very small crystal sizes were obtained (from a few nanometers up). For thick films, relatively large, more or less well-defined cubic crystals of between 50 and several hundred nanometers were obtained. In general, the depositions from the more concentrated OH solutions gave larger, betterdefined crystals. The differences were rationalized in terms of hydroxide and ionby-ion mechanisms predominating in the low- and high-concentration OH solutions, respectively. The ability of certain monolayer-coated surfaces to enhance or retard film growth was exploited to pattern CD PbS films (see also Secs. 2.8 and 4.1.5.3). This patterning was based on the UV-induced oxidation of a thiol linkage to Au (a strong bond) to a weakly bound sulphonate group that could be rinsed away [41]. Figure 5.1 shows the processes involved in this patterning. A long-chain mercapto-carboxylic acid (16-mercaptohexadecanoic acid) was self-assembled on a Au substrate and exposed to short-wavelength UV radiation through a patterned

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FIG. 5.1 Scheme for patterned deposition of PbS. (a) Self-assembled monolayer (SAM) of long-chain mercapto-carboxylic acid. On exposure to UV radiation through a mask, the exposed thiol group is oxidized to a sulphonate group (b) that is weakly bound to the Au substrate and can be easily rinsed away (c). Subsequent formation of a long-chain alkyl thiol SAM occurs only on the exposed Au (d). CD of PbS occurs only on hydrophilic carboxylate endgroups and not on hydrophobic methyl groups (e). (See Ref. 41).

mask, in this case a TEM grid (a). The parts of the monolayer exposed to the UV were oxidized to sulphonates (b) that were then rinsed away, leaving the substrate patterned with mercapto-carboxylates and exposed Au (c). A long-chain alkyl thiol (16-mercaptohexadecane) was then self-assembled onto the (exposed Au portions of the) substrate (d). Chemical deposition of PbS (Pb2, NaOH, thiourea) onto this patterned substrate resulted in PbS deposition on the hydrophilic carboxylate endgroups and almost not at all on the hydrophobic alkyl endgroups (e). Figures 5.2a and b show the resulting deposit: A dense PbS deposit formed on the parts that were unexposed to the UV radiation (dark grid areas in 5.2a and lower part of 5.2b), while only scattered particles were found on the exposed areas.

5.2.4.3

Ferroelectric Substrate

Deposition of PbS onto a poled ferroelectric substrate (a complicated oxide of mainly Pb, Zr, and Ti) from a Pb(NO3)2/NaOH/thiourea bath (containing also hy-

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FIG. 5.2 SEM image of (a) PbS (dark areas) growing on patterned carboxylate-terminated regions of substrate and (b) boundary between PbS-covered carboxylate region (lower part of micrograph) and almost bare thiol-terminated region. (From Ref. 41 with permission from Elsevier Science).

droxylamine hydrochloride as a reductant and a trace of Bi3; the reason for these additives was not given) resulted in films with larger particle size ( 1 m) compared to that of films deposited onto an unpoled ferroelectric or glass (particle size ca. 0.3 m) [42]. The larger particle size was explained by the electric field present at the substrate surface, which attracts ions from the solution and increases the growth rate. It was not stated whether the film thickness was greater for the poled substrates (as should be for a faster growth rate). In any case, a faster growth rate does not necessarily mean a larger crystal size; the opposite is often true. The electrical properties of the films also depended on the poled state of the substrate. The resistivity of the PbS on the poled surfaces (10 -cm on the positively poled face and 20 -cm on the negatively poled face) was overall lower than that on the unpoled ceramic (varied between 10 and 100 -cm) or on glass (50 cm). The photoconductivity response varied over almost an order of magnitude, depending on substrate and temperature. The room-temperature response was highest on glass and lowest on the poled substrate, but this order was reversed at liquid N2 temperatures. The films were all p-type except for that deposited on the negatively poled face, which was n-type. All these results were explained by the effects of the electric field and the surface charge on the depositing film. The temperature dependence of the photoconductivity effects was attributed to trapping dominating the photoconductivity at lower temperatures. It should also be noted that the particle size may also affect the electrical properties; e.g., fewer grain boundaries would result in a lower resistivity (assuming identical bulk properties).

5.2.4.4

Deposition on Liquid

Although not CD in the usual sense of the technique, it is worth mentioning a report of formation of PbS films by passing H2S over PbAc2 aqueous solution slightly acidified with acetic acid [25]. The films were picked up on a gauze and

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were analyzed by electron diffraction. The (001) faces were found to be more or less parallel to the solution surface, and the crystal size was ca. 25 nm. It is clear that this technique should be applicable to almost all metal sulphides (also selenides—see Sec. 5.3—and tellurides). Deposition from colloidal sol, a related deposition technique (although not depending on a particular substrate surface) also resulted in PbS films using gaseous H2S. A very thin PbS film was formed when a quartz plate was immersed overnight in a solution of Pb(ClO4)2 and poly(vinyl alcohol) through which H2S had been bubbled [43]. The absorption spectrum of this film was similar to that of the PbS sol and consisted of several absorption peaks with an absorption onset of ca. 630 nm (strongly blue shifted from the PbS bulk bandgap). The XRD crystal size of the precipitate was ca. 3 nm (see Chap. 10 for more details).

5.2.5 Optical Absorption/Transmission/Reflection of PbS Films Some general remarks on the optical absorption of lead chalcogenides are in order here, since transmission spectra of thin films of these compounds are open to misinterpretation. If a transmission spectrum of, say, a thin (⬇100 nm) PbS film is taken, a sharp drop in transmission will be seen in the red–near IR region (usually between 600 and 800 nm). It is easy to translate this into the bandgap of PbS, and this has been done, even in careful studies. Thus, one thorough study of films of the Pb chalcogenides deposited both by CD and by evaporation (for PbTe, only by evaporation) has reevaluated their bandgaps upwards (e.g., 1.3 eV for PbS instead of 0.4 eV) [44]. It was later clarified that the bandgaps were indeed much lower than originally believed (e.g., Ref. 45). These materials have an absorption coefficient in the region of their bandgap, and for considerably higher photon energies, that is only moderate (ca. 104 cm1), and only at much higher photon energies does this absorption coefficient increase substantially (by about an order of magnitude), giving the apparent bandgap onset measured in thin films. In thick films (microns and up), the absorption close to the bandgap is high. However, for much thinner films, this absorption is weak. Furthermore, it is often masked by reflection. It is particularly important when dealing with films that absorb weakly and that have high reflectivity (such as PbS) to stress that transmission measurements that are uncorrected for reflection should not be directly converted to absorption, since the differences in transmission with change in wavelength may well be due to differences in reflection rather than in absorption. Since the lead chalcogenides have high dielectric constants (therefore high refractive indices and high reflection), masking of weak or even moderate absorption by reflection is probable. Although many studies of the optical properties of CD PbS films have been made, most of them do not extend to wavelengths corresponding to the bulk PbS bandgap (ca. 3 m). The study noted by Gibson [44], which was corrected for re-

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flection, showed a strong absorption beginning at ca. 800 nm and a relatively weak absorption extending out to at least 6 m. In another study that does show the extended IR region for 1-m-thick films, the absorption edge is very dependent on the presence or absence of an oxidizing agent in the bath (the absorption coefficients at energies well above the onset were similar) [29]. The films prepared with oxidant had a very gradual onset starting well below the normal bandgap of PbS (ca. 0.4 eV), while that deposited without oxidant exhibited a sharp onset at ca. 0.52 eV, considerably above the normal bandgap. Various possibilities might explain these differences. The lower onset of the “oxidant” films might be caused by bandgap tailing due to a high concentration of states in the gap near the band edges or to electric field effects on the bandgap due to these or other states. The larger bandgap of the nonoxidant films could be due to size quantization (the crystal size—as opposed to particle size— was not measured), band filling (Burstein–Moss shift), or, as suggested in the paper, a high concentration of oxygen in the PbS (although, as pointed out in Sec. 5.2.3.3, the use of oxidant in the deposition solution did not affect the oxygen concentration in the film). Most other studies show the optical spectra of thin (often 100-nm) films, with emphasis on their solar control properties, and limited to a long-wavelength limit of 2.5 m [46–48] (Ref. 49 shows reflection spectra to longer wavelengths). These films usually have an apparent absorption onset in the region of 600–800 nm, the longer wavelengths characteristic of thicker films. The spectra at longer wavelengths are typically dominated by reflection rather than by absorption. This would suggest bandgaps of between 2 and 1.5 eV. Figure 5.3 shows typical examples of the transmittance and specular reflectance of PbS films deposited from

FIG. 5.3 Optical transmittance (a) and near-normal specular reflectance (b) of CD PbS films of different thicknesses. The thickness increases from A to F over an estimated range of ca. 50 nm to 200 nm. (Adapted from Ref. 46 with permission from IOP Publishing Ltd.).

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a Pb(Ac)2/triethanolamine/NaOH/thiourea bath; the absence of absorption in the near-IR region, where PbS would normally absorb, is evident from a comparison of the transmission and reflection data. In fact, observation of various CD PbS (from thiourea baths) spectra shows that there is absorption starting around 1.5–2 m, but it is usually weak. Partly this is due to the fact that much of the literature deals with thin films (often 100 nm), but the absorption does seem to be smaller than expected. In the study by Pop et al. [49], moderately strong absorption still occurs in multilayer films (some hundreds of nanometers thick) at wavelengths longer than 2 m. The small amount of information on films deposited from thiosulphate baths suggest that the absorption of these films in the near-IR range may be higher than those deposited from thiourea baths. Gadave et al. measured a transmission under 1% over the entire measured range from 2 m to 400 nm (in fact the transmission increased slightly towards shorter wavelengths) [10]. No reflectivity or thickness data were given: however, the films must have been very thick (by CD standards) to give such a low transmission at 2 m. Reference 17 shows a gradually decreasing transmission from the maximum wavelength measured (2.5 m) for a film of unknown thickness, with a transmission of ca. 20% at 800 nm, while a 60nm film deposited by a similar process on ultrathin Cu2S (which, according to the data of Ref. 17, did not much affect the spectrum) showed a transmission of 11% at 800 nm (lower than most layers of comparable thickness deposited from a thiourea bath) but with a poorly defined long-wavelength onset. To sum up, while there is too little information available to draw any firm conclusions, it appears that films deposited from most thiourea baths are weakly absorbing in the near-IR region and that films deposited from thiosulphate solutions (which are mildly acidic) may possess different optical properties in general than those deposited from (alkaline) thiourea baths. In this respect, and if this difference is real, it would be interesting to deposit PbS from thioacetamide baths, which can be both acidic and alkaline.

5.2.6 Photoconductivity of CD PbS (and PbSe): General Considerations Specific photoconductive properties of PbS films have already been treated. This section deals with more general aspects of CD PbS (for the most part, also relevant for PbSe) films. The bandgap of PbS at room temperature is 0.4 eV, corresponding to a wavelength of 3.1 m, which more or less gives the long-wavelength detection limit. The bandgap increases with increasing temperature, in contrast to the normal semiconductor bandgap dependence on temperature. Therefore the longwavelength detection limit of detectors made using these films shifts to shorter wavelengths with increasing temperature. A low temperature of operation there-

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fore not only increases the sensitivity of the detectors (reduction of noise) but also widens the spectral response. The air-annealing (typically between 300 and 400°C) of PbS and PbSe used, in most cases, to maximize the photoconductive response results in some surface oxidation, probably to PbSO3 and PbSO4. Such compounds have been detected in PbS precipitate, prepared from a typical CD process, after air-annealing at 300°C or higher [50]. Oxidizing agents are often added to the CD bath to give photoconductive films in the as-deposited state [29,30]. In fact, commercial photoconductive films used a proprietary oxidant as a matter of course. These commercial films, prepared from alkaline solutions of a lead salt and thiourea, were made up of several layers of CD PbS (total thickness ca. 1 m) to maximize the photoresponse, in particular the long-wavelength response. The cause of photoconductivity in these films has been thoroughly studied. Here only a very condensed account of the theory is given. References 51 and 52 give more detail. Electrical conductivity, , is given by the product of free-carrier concentration, n, carrier mobility, , and carrier charge, e: 
ne

(5.5)

Therefore an increase in conductivity upon illumination (photoconductivity) can be due to either an increase in carrier concentration and/or an increase in mobility. In general, it is believed that an increase in carrier (hole) concentration is the dominant cause for room-temperature photoconductivity for the lead chalcogenides and that an increase in mobility becomes increasingly important at low temperatures. The dark conductivity of films deposited with or without added oxidant were similar; the difference in photoconductivity between them was ascribed to the formation of sensitizing centers (interband states) due to the oxidant. Finally, it is worth mentioning a comment made in a paper describing junctions between Ge and CD PbS [34]. It was noted that evaporated epitaxial PbS films were poorly, if at all, photoconducting, while CD films, with mobility lower by two orders of magnitude and much “poorer” structure, were much superior in this respect. In a way, this should not be surprising since, for good photoconductivity, low dark conductivity (and therefore either low mobility and/or low carrier concentration) is necessary.

5.3 PbSe PbSe has a considerably more recent history than does PbS; having been deposited by CD only about 60 years ago, instead of more than 130 years for PbS. The earliest work appears to have been carried out by the Germans during World War II and, as for PbS, was shrouded in secrecy. While PbSe has been deposited by the Se analogues of thiosulphate and thiourea, as for PbS, in contrast to the history of

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PbS deposition, the early work used mostly selenourea (or selenourea derivatives), while the more recent studies used the selenosulphate baths. As for PbS, these two baths will be treated separately. Optical and photoconducting properties will not be treated separately as for PbS, for which there is a larger body of results, but are included with the description of the various deposition baths.

5.3.1 Selenourea-Based Baths The first apparent report in the open literature of CD PbSe for photoconductive detectors was in 1949 [53]. The PbSe was deposited from a solution of PbAc2 and selenourea onto a predeposited (from PbAc2 and thiourea) layer of PbS. The PbS layer acted as a seed layer, presumably to obtain faster deposition (it was noted that the PbSe deposition was much slower than that of PbS). The photoconductivity of this film exhibited a broad maximum between 3 and 4 m, giving a reasonable response out to beyond 4.5 m (PbS drops off at 3 m). A detailed description of PbSe deposition using N,N-dimethylselenourea (DMS) was presented in 1964 [54]. DMS was used instead of selenourea because of its greater stability. Even so, Na2SO3 was added to inhibit decomposition to Se. The Pb (as nitrate) was complexed with citrate to keep it in solution in the alkaline conditions used (pH
9.8 using ammonia). This pH was fairly critical: At 9.5 (and at 25°C) the deposition was much slower, and at 10.1 rapid bulk precipitation occurred. Counter intuitively (and contrary to the case using selenosulphate), the reaction rate was faster for an aged DMS solution than for a fresh one (after aging for 15 hr, little additional increase in rate occurred on further aging). This was explained by inhibition of the DMS decomposition to selenide by the sulphite (Sec. 3.2.2.2). It was noted that PbSe formation occurred more rapidly in the absence of sulphite and that adding fresh sulphite to an aged DMS solution reduced the reaction rate. The best films were obtained if a fresh solution was used to deposit PbSe on glass slides that had been precoated with a thin PbSe layer. In this case, film growth started almost immediately rather than after an induction period. The resistivity of the films varied from 106 to 107 . Since the thicknesses of the films were not given, it is difficult to convert these values to a specific resistivity. However, it does appear that the variations in resistivity were due mainly to film thickness and an upper value of 100 -cm seems reasonable. A detailed study of both the mechanism [55] and the kinetics [56] of PbSe deposition from selenourea baths has been carried out. The Pb was complexed with citrate, and hydrazine was used to control alkalinity (remember that hydrazine can function also as reducing agent and as complexant). As with the foregoing study, immediate deposition, with no induction time, occurred on glass on which PbSe had been precoated (or on glass that was sensitized with SnCl2 solution, resulting in formation of tin hydroxide/oxide nuclei). On untreated glass, an induction time for film deposition, which paralleled that for homogeneous pre-

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cipitation in solution, was observed. The pH was somewhat lower than that of the preceding study (typically ca. 8.8) and evidently less critical, although the deposition rate increased with increasing OH concentration. The rate was also dependent on the selenourea concentration but independent of the Pb or citrate concentration. Based on these observations, it was suggested that the rate-determining step was decomposition of selenourea by OH to selenide ion: (NH2)2CSe  OH → HSe  CH2N2  H2O

(5.6)

The selenide could then react with whatever Pb species were present, either in solution or as a solid phase, if present. The activation energy of this process, 60 kJ/mole, was consistent with a chemical rate-determining step. Films deposited from a selenourea bath (details not provided) were annealed at 350°C in air and changes in their resistivity and photoconductivity measured [57]. The resistivity of the films (between 1 and 4 m thick) dropped from a few hundred k (assumed lateral conduction across the film) to about 1 k after heating for two minutes, with a peak in the resistivity after heating for about a minute. The photoconductivity maximum shifted from 1.5–1.8 m for the as-deposited film to 3.6 m for the annealed one. These phenomena were attributed to crystallization of a matrix of amorphous or nanocrystalline PbSe surrounding larger crystals, which crystallize or grow with annealing time. The crystal size of as-deposited films in this and other studies by the same group was typically some tens of nanometers.

5.3.2 Selenosulphate-Based Baths Sodium selenosulphate was used to precipitate PbSe from Pb2 solutions by mixing PbAc2 and Na2SeSO3 solutions [58]. The precipitation was rapid, and film formation did not occur (at least to any visible extent). Complexation was required to slow the reaction. Films of PbSe were first deposited using selenosulphate from a hydroxide-complexed (plumbite) bath [59]. By complexing the Pb [with citrate, Rochelle salt (tartrate), or oxalate] and adding alkali (NaOH, KOH, or NH4OH) to an optimum pH of 11.05, mirror films of PbSe were obtained [60]. The mechanism proposed, for both precipitate [58] and films using carboxylic acids [60], was formation of PbSeSO3: PbAc2  Na2SeSO3 → PbSeSO3  2NaAc

(5.7)

followed by decomposition of the PbSeSO3 to PbSe by water: PbSeSO3  H2O → PbSe  H2SO4

(5.8)

No evidence was presented for this mechanism, although it is a reasonable route, nor for the assumption of an identical mechanism for precipitate and film formation (although that too is likely. In fact, in another study, based on a kinetic study

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of PbSe formation from a similar bath employing glass powder (particle size ca. 80 m) as a “substrate,” it was concluded that the same mechanism was operative both for PbSe formation as a precipitate in solution and for that deposited on the glass surface [61]). On the other hand, Fofanov and Kitaev assumed a mechanism whereby selenide ion, formed by hydrolysis of selenosulphate, reacted with Pb2 via the ion-by-ion mechanism [59]. The thickness of the films using the carboxylate complex was not dependent on the Pb concentration (at least from 10 mM up). It was, however, very dependent on the pH value, with a sharp maximum, the position of which was dependent on the anion of the lead salt and, more so, on the hydroxide used (Na, K, NH4). The thickness varied from 50 nm to 300 nm. The film thickness was also moderately dependent on deposition temperature, increasing with increase in temperature (in contrast to the more usual decrease in terminal thickness with increasing temperature). It is not certain, however, whether these thicknesses (and others measured in this study) were actually terminal thicknesses or only the thickness measured after a certain time. The bandgap of these films was measured to be 0.28 eV—the same as the literature value. Films were also deposited from similar solutions, only using selenourea instead of selenosulphate and at somewhat lower pH (⬇9). The quality (of the mirror surface) and thickness of the films deposited from selenosulphate solutions were similar to those deposited from the (at that time) conventional selenourea ones. Kainthla et al. carried out an investigation on the parameters that affected deposition rate [62]. The rate increased, as expected, with increase in temperature and selenosulphate concentration. However, it decreased with increase in pH. This was explained on the basis of the expected dominant hydroxy-citrate complex of Pb, [Pb(OH)C6H5O7]2. The equilibrium involving this complex is [Pb(OH)C6H5O7]2 D Pb2  OH  C6H5O 3 7

(5.9)

A greater hydroxide concentration will shift the equilibrium to the left, decreasing the free-Pb2 concentration. This explanation means that the rate is dependent on the lead concentration, in contrast to the previous studies discussed earlier, which found the rate to be essentially independent of this parameter. It was also noted that the optimal pH was temperature dependent: At low deposition temperature it was ⬇9, while at high temperatures it was ⬇10. This follows from the inverse dependence of rate on pH. The optimal pH is a balance between slow-enough formation of PbSe (to prevent precipitation in solution) but not too slow to prevent formation of a film in a reasonable time. At higher temperatures the rate is faster, and therefore the optimal pH should be lower. Films (between 0.4 and 2.5 m thick) deposited from selenosulphate solution were characterized by optical spectroscopy and (photo)conductivity [63].

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There appeared to be a very large variation in optical properties from film to film; in one case, little absorption (corrected for specular reflection) was measured from 2 m toward increasing wavelength, while another sample showed a sharp absorption onset close to 5 m. The films had a resistivity of between 3.103 and 3.104 -cm and a photoconductivity maximum varying between 1.2 and 1.9 m. An investigation of PbSe deposited from selenosulphate baths was carried out using various complexants and conditions and with emphasis on the mechanisms and, in particular, on the crystal size and morphology of the deposits [64]. Citrate, nitrilotriacetate (NTA), and hydroxide (in order of increasing complexing strength with Pb) were used under conditions where a solid hydrated oxide phase was either present in solution (cluster) or absent (ion-by-ion). The presence of this phase, and a semiquantitative estimation of its relative concentration, was measured by UV absorption and light-scattering measurements. As for CdSe, the morphologies of the two types of films were very different, with the ion-by-ion films having a larger crystal size. However, the difference depended on the complexant used as well as on other deposition conditions. A wide variety of crystal sizes was obtained. Figure 5.4 shows some TEM micrographs of various PbSe films. The conditions of deposition for each film (the important factors are given in the figure legend) are not important here; the purpose is to demonstrate the wide range of crystal sizes and morphologies in PbSe films obtained using different complexants, solution compositions, and deposition temperatures. Under certain conditions, relatively large and small crystals coexisted in the same films. For example, in “cluster” deposition from citrate complex, both small (ca. 4 nm) and medium-sized (6–12 nm) crystals were formed, although only the small crystals formed at the beginning of the deposition, and the larger ones appeared in thicker films, again, an indication of both mechanisms operating (ionby-ion growth normally is slower than cluster growth, therefore it takes longer for the larger crystals to appear). Additionally, the small and larger crystals were not deposited together but in separate regions; i.e., regions only of larger crystals were formed surrounded by (most of the deposit) only small crystals. It was suggested that the larger crystals formed in regions when no previous small-crystal deposition had occurred (i.e., in pinholes in the originally thin film). For low selenosulphate concentrations, only the small crystals were formed, even in thicker films, and this was rationalized by the lower steady-state selenide concentration, which would favor cluster growth over ion-by-ion formation (the product of free lead and selenide ions needs to be larger than the solubility product of PbSe for ion-by-ion deposition to occur). An important difference between the citrate depositions and the NTA or hydroxide ones is that, even in the “ionby-ion” citrate deposition, some low concentration of colloidal hydrated oxide was present, due to the relatively low complexing strength of citrate. The pH of the hydroxide baths (13) was much higher than that of the citrate or NTA baths (10.8).

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FIG. 5.4 TEM micrographs of PbSe films deposited under different conditions. A: from citrate/0°C/hydroxide mechanism. B: nitrilotriacetate/60°C/ion-by-ion mechanism. C: hydroxide/60°C/ion-by-ion mechanism. D: hydroxide/80°C/ion-by-ion mechanism. All films are as deposited.

In common with CdSe deposition from selenosulphate baths, cluster deposition of PbSe normally resulted in specular films, while ion-by-ion films were initially highly scattering as thin films but eventually (usually) became specularly reflecting with increase in thickness. As for CdSe, the development of specularity with thickness of ion-by-ion films could be explained by filling in of voids between the large crystals. The wide range of very small crystal sizes in these films gives rise to strong blue shifts in their optical spectra due to size quantization [65]. This aspect of these films is dealt with in detail in Chapter 10. Other complexants have been used for PbSe deposition. Triethanolamine was used in one study [66]. While deposition occurred over a wide range of temperatures, optimum results (in terms of rate of deposition and film thickness) were obtained at a deposition temperature of 75°C. In another study, lead nitrate was dissolved in an excess of hydroxide and excess selenosulphate was also used as an additional complexant [67]. The pH was 10 (adjusted with acetic acid), and depo-

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sition was carried out both at room temperature and at 70°C. The substrate (glass or polyester) was pretreated with SnCl2 solution—the standard tin sensitization; presumably this improved the adherence and/or homogeneity of the films. The resistivity of these films was ca. 104 -cm as deposited. Annealing in air at 100°C for 24 hr decreased this resistance to 3  103 -cm, a drastic change for such a mild treatment and one that suggests that the electrical properties of these as-deposited films may change greatly with time, even under ambient conditions.

5.3.3 Variations in PbSe Deposition X-ray amorphous PbSe has been deposited by addition of Na2S2O3 (at least half the selenosulphate concentration) to a room-temperature selenosulphate/citrate bath [68]. The thiosulphate increased the induction time for PbSe formation in the solution, although eventually film growth was even faster than in the absence of thiosuphate. It was noted that the initial, very thin deposit was yellow. A clear XRD pattern was obtained (for thicker films) after annealing at 350°C for one minute, giving a crystal size of 13 nm (which grew with continuing annealing). It is interesting to note that, in contrast to films deposited in the absence of thiosulphate, those made using thiosulphate did not peel off, even after remaining a long time in the depositing solution. Chemically deposited films will often peel off if left too long in solution, probably due to stress in the thickening film, and this is expected to be less or absent in an amorphous film. The amorphous structure was attributed mainly to the increased deposition rate caused by the thiosulphate. Since the growth rate was apparently in the range of hours, this by itself would not be expected to result in an amorphous film, although it might contribute to the effect. A thorough investigation of this deposition would be desirable, including TEM imaging (to see whether the film is, indeed, amorphous, since XRD by itself is not a good enough verification of this) and, particularly important, compositional analysis (to see if there is an appreciable amount of S in the film, thiosulphate itself forming PbS). From the optical spectrum, an approximate bandgap of 1.5 eV could be estimated, although the interpretation of the spectrum is open to ambiguity. A photoconductivity maximum at 1.1 m (1.1 eV) was measured. These anomalously high energy values were attributed to size quantization (Chap. 10). Transmission in the mid-IR was also measured for these films. Although there were some absorption bands, probably due to adsorbed species from the deposition bath, the films were essentially transparent at wavelengths longer than ca. 5 m, the transmission decreasing gradually at shorter wavelengths. While annealing at 350°C did change the IR spectral shape, particularly the intensity of some of the absorption bands, overall the transparency beyond 8 m was maintained (in some regions, even increased). This implies that the films, both as deposited and annealed, were very intrinsic, since free carriers would result in absorption in this spectral region.

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The (presumed lateral) resistance of the films was ca. 108 , which decreased 4–5 orders of magnitude after annealing at 350°C. In analogy with PbS, a film of PbSe has been formed by passing H2Se over PbAc2 solution acidified with acetic acid, giving specularly reflecting films that could be picked up from the surface of the solution [25]. Electron diffraction of these films showed a tendency to grow with (001) planes parallel to the solution surface. The average crystal size of these films was 25 nm. A variation of the CD process for PbSe involved deposition of a basic lead carbonate followed by selenization of this film with selenosulphate [64]. White films of what was identified by XRD as 6PbCO33Pb(OH)2PbO (denoted here as PbMOHMC) were slowly formed over a few days from selenosulphate-free solutions that contained a colloidal phase and that were open to air (they did not form in closed, degassed solutions). CO2 was necessary for film formation—other than sparse deposits, no film formation occurred of hydrated lead oxide under any conditions attempted in this study. Treatment of these films with selenosulphate solution resulted in complete conversion to PbSe at room temperature after 6 min. The selenization process of this film was followed by XRD, and it was seen to proceed by a breakdown of the large PbMOHMC crystals to an essentially amorphous phase of PbSe with crystallization of this phase to give finally large (ca. 200 nm) PbSe crystals covered with smaller (15–20 nm) ones as well as some amorphous material. Hydrazine added to a selenourea//Pb2 bath was found to strongly affect the electrical and photoconductive properties of the PbSe films [69]. As the hydrazine concentration increased, the dark resistivity increased in a very non-monotonic way, from ca. 0.1 -cm (low hydrazine concentration) to 104 -cm. Photoconductivity was observed, as might be expected, only for the high-resistivity films. The effect of the hydrazine was attributed to an increase in the deposition rate by an increase in pH; however, the pH values of the solutions were not reported, and hydrazine would be expected to affect the reaction as a reducing agent apart from pH considerations. As an interesting aside, it was noted that only p-type films could be deposited by this technique (and this is the case for most CD PbS), while evaporated PbSe films were normally n-type (due to Se vacancies). The absorption edge of the films was ca. 4.2 m, close to to the literature bandgap of 0.28 eV.

5.3.4 Comparison of Films Deposited from Selenourea and Selenosulphate Baths A series of studies comparing films deposited from selenourea and selenosulphate baths were carried out, with emphasis on photoconductivity behavior and the effects of annealing. A broad photoconductivity maximum occurred at ca. 1.6 m for the selenosulphate film, decreasing strongly beyond 2 m. A similar spectrum,

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but shifted 0.2–0.4 m to longer wavelengths, was observed for the selenourea film. Annealing at 350°C in air broadened both spectra to ca. 5 m (approximately the bulk bandgap of PbSe), although the the photoconductivity response of the selenosulphate films was initially degraded (mainly due to increased noise), and a long (ca. 50 hr) annealing time was required for good response [70]. Photoluminescence spectra of the films were measured (77 K) and compared with an epitaxial PbSe layer in the same study. Blue shifts in the spectra (greater for the selenourea films) were attributed to quantum size effects (see Chap. 10). The crystal size was reported to be 40–60 nm (the selenourea ones being somewhat smaller than the selenosulphate ones), growing to 100–150 nm after annealing. The changes in resistivity with annealing of films deposited from selenourea and selenosulphate baths, as well as evaporated films, were compared [71,72]. Although there were small differences between the various films, no major difference was found. Additionally, the resistivity of as-deposited films, deposited from both selenourea and selenosulphate baths, does not change with time over a period of months in air. However, after annealing in air at 350°C when the resistivity increases, there is a gradual decrease in room-temperature resistivity (and also in photoconductivity response) with time [73]. These variations were related to formation of PbSeO3 and adsorbed oxygen on the surface of the annealed crystals.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

C Puscher. Dingl. J. 190:421, 1869. J Emerson-Reynolds. J. Chem. Soc. 45:162, 1884. DE Bode. In: Physics of Thin Films. Vol. 3, Academic Press, New York, NY, 1996, p 275. F Kicinski. Chem. Ind. 54, 1948. TS Moss. Proc. IRE 43:1869, 1955. WN Arnquist. Proc. IRE 47:1420, 1959. DJ Lovell. Am. J. Phys. 37:467, 1969. TH Johnson. Proc. SPIE 60, 1983. CD Lokhande. Mater. Chem. Phys. 28:1, 1991. KM Gadave, SA Jodgudri, CD Lokhande. Thin Solid Films 245:7, 1994. E Beutel, A Kutzelnigg. Z. Elektrochem. 36:523, 1930. PL Allen, A Hickling. Chem. Ind. 51:1558, 1954. EB Andersen. Z. Phys. Chem. B. 32:237, 1936. WH Perkins, AT King. J. Chem. Soc. 103:301, 1913. E Beutel. Z. Angew. Chem. 26:700, 1913. E Beutel, A Kutzelnigg. Monats. 58:295, 1931. I Grozdanov. Semicond. Sci. Tech. 9:1234, 1994. M Najdoski, B Mincevasukarova, A Drake, I Grozdanov, CJ Chunnilall. J. Mol. Struct. 349:85, 1995. HL Smith. J. Sci. Instrum. 4:115, 1927.

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20. O Hauser, E Biesalski. Chem. Ztg. 1078, 1910. 21. G Bruckman. Kolloid Z. 65:1, 1933. 22. MK Norr. J. Phys. Chem. 65:1278, 1961. 23. H Pick. Zeit. Phys. 126:12, 1949. 24. HN Acharya, NK Misra. J. Phys. D: Appl. Phys. 4:1968, 1971. 25. H Wilman. Proc. Phys. Soc. 60:117, 1948. 26. HN Acharya, HN Bose. Phys. Status Solidi (a) 6:K43, 1971. 27. VF Markov, LN Maskaeva, GA Kitaev. Inorg. Mater. 36:657, 2000. 28. EM Larramendi, O Calzadilla, A Gonzalez-Arias, E Hernandez, J Ruiz-Garcia. Thin Solid Films 389:301, 2001. 29. GH Blount, PJ Schreiber, DK Smith, RT Yamada. J. Appl. Phys. 44:978, 1973. 30. GP Kothiyal, B Ghosh, RY Deshpande. J. Phys. D: Appl. Phys. 13:869, 1980. 31. IC Torriani, M Tomyiama, S Bilac, GB Rego, JI Cisneros, ZP Argüello. Thin Solid Films 77:347, 1981. 32. PK Nair, VM Garcia, AB Hernandez, MTS Nair. J. Phys. D: Appl Phys. 24:1466, 1991. 33. J Bloem. Appl. Sci. Res. Section B 6:92, 1956. 34. JL Davis, MK Norr. J. Appl. Phys. 37:1670, 1966. 35. G Guizzetti, F Filippini, E Reguzzoni, G Samoggia. Phys. Status Solidi (a) 6:605, 1971. 36. S Watanabe, Y Mita. J. Electrochem. Soc. 116:989, 1969. 37. M Isshiki, T Endo, K Masumoto, Y Usui. J. Electrochem. Soc. 137:2697, 1990. 38. BL Sharma, SN Mukerjee. Phys. Status Solidi (a) 2:K21, 1970. 39. H Sigmund, K Berchtold. Phys. Status Solidi 20:255, 1967. 40. FC Meldrum, J Flath, W Knoll. J. Mater. Chem. 9:711, 1999. 41. FC Meldrum, J Flath, W Knoll. Thin Solid Films 348:188, 1999. 42. I Pintilie, E Pentia, L Pintilie, D Petre, C Constantin, T Botila. J. Appl. Phys. 78:1713, 1995. 43. S Gallardo, M Gutierrez, A Henglein, E Janata. Ber. Bungenges. Phys. Chem. 93:1080, 1989. 44. AF Gibson. Proc. Phys. Soc. B 63:756, 1950. 45. JGN Braithwaite. J. Sci. Instr. 32:10, 1955. 46. PK Nair, MTS Nair, A Fernandez, M Ocampo. J. Phys. D: Appl. Phys. 22:829, 1989. 47. PK Nair, MTS Nair. J. Phys. D: Appl. Phys. 23:150, 1990. 48. PK Nair, M Ocampo, A Fernandez, MTS Nair. Sol. Energy Mater. 20:235, 1990. 49. I Pop, C Nascu, V Ionescu, E Indrea, I Bratu. Thin Solid Films 240: 1997. 50. GA Kitaev, LG Protasova, VG Kosenko, VF Markov, LV Dyadyuk, GV Stepanova. Inorg. Mater. 29:1182, 1993. 51. S Espevik, C-H Wu, RH Bube. J. Appl. Phys. 42:3513, 1971. 52. E-H Lee, RH Bube. J. Appl. Phys. 43:4259, 1972. 53. CJ Milner, BN Watts. Nature 163:322, 1949. 54. RA Zingaro, DO Skovlin. J. Electrochem. Soc. 111:42, 1964. 55. AB Lundin, GA Kitaev. Inorg. Mater 1:1900, 1965. 56. AB Lundin, GA Kitaev. Inorg. Mater. 1:1905, 1965. 57. LP Biró, A Darabont, P Fitori. Europhys. Lett. 4:691, 1987. 58. NI Glistenko, AA Eremina. Zh. Neorg. Khim. 5:1003, 1960.

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59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73.

GM Fofanov, GA Kitaev. Russ. J. Inorg. Chem. 14:322, 1969. NI Glistenko, IN Shramchenko, AN Kosilova. Zhurn. Prikl. Khim. 45:1356, 1972. GA Kitaev, AZ Khvorenkova. Russ. J. Appl. Chem. 72:1520, 1999. RC Kainthla, DK Pandya, KL Chopra. J. Electrochem. Soc. 127:277, 1980. RM Candea, N Dadarlat, R Turcu, E Indrea. Phys. Stat. Sol. (a) 90:K91, 1985. S Gorer, A Albu-Yaron, G Hodes. Chem. Mater. 7:1243, 1995. S Gorer, A Albu-Yaron, G Hodes. J. Phys. Chem. 99:16442, 1995. P Pramanik, S Biswas, PK Basu, A Mondal. J. Mater. Sci. Lett. 9:1120, 1990. I Grozdanov, M Najdoski, SK Dey. Mater. Lett. 38:28, 1999. LP Biro, RM Candea, G Borodi, A Darabont, P Fitori, I Bratu, D Dadarlat. Thin Solid Films 165:303, 1988. DH Roberts, JE Baines. J. Phys. Chem. Solids 6:184, 1958. D Dadarlat, RM Candea, R Turcu, LP Biro, Zasavitskii, II, MV Valeiko, AP Shotov. Phys. Status Solidi (a). 108:637, 1988. RM Candea, R Turcu, P Margineanu, N Dadarlat. Phys. Status Solidi (a) 96:337, 1986. RM Candea, R Turcu, G Borodi, I Bratu. Phys. Status Solidi (a) 100:149, 1987. RM Candea, LP Biro, N Dadarlat, G Borodi, A Darabont, P Fitori, R Turcu. Phys. Status Solidi (a). 108:233, 1988.

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6 Other Sulphides and Selenides

6.1 INTRODUCTION This chapter covers sulphides and selenides not included in Chapters 4 and 5, i.e., all metals except for Zn, Cd, Hg, and Pb. Some of these materials, e.g., the sulphides of Bi, Cu, and Ag and Cu-Se, have been the subject of many investigations. There are others, however, on which as little as one paper has been published altogether. In order to minimize, as much as possible, making this chapter into a list of recipes and properties, its layout will be somewhat different than that used up to now. The chapter is divided into sections alphabetically (by English name rather than by chemical symbol), designated by the name of the metal. Basic data for those compounds for which there are at least several papers will be listed in tabular form together with references. The data in the first column of the table (the deposition bath composition) will be given as sulphide or selenide source/complex/temperature and pH. The specific metal ion is not given in most cases, because this will be clear from the location of the table and the metal salt used will normally be found in the specific description of the study in question. It should be kept in mind that the chalcogen source itself is a complexant, often fairly weak but sometimes strong. Resistivity is given as specific resistivity; if not defined, then the resistance reported is given in parentheses. Specific details outside of these basic data will be treated separately in the order in which the particular study appears in the table and,

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if connected with a specific datum in the relevant table, will be signified by an asterix after this datum. Each discussion of this type will include the reference number to identify it. Abbreviations used in the “Solution” column are: DMSeU—N,Ndimethyl selenourea; EDTA—ethylenediaminetetra-acetate; S2O3—thiosulphate; SeSO3—selenosulphate; TA—thioacetamide; TEA—triethanolamine; TU— thiourea; RT—room temperature. The abbreviations in the “Bandgap” column are: dir.—direct; indir.—indirect. Before discussing the specific studies, a general overview will be given for each metal. This overview relates only to information of direct relevance to the subject and will be relatively lengthy in some cases and very brief or even nonexistent in others, according to need.

6.2 ANTIMONY 6.2.1 Sb2S3 (See Table 6.1) All reported cases of CD antimony sulphide involve the trisulphide, Sb2S3. Sb2S3 is soluble in hydroxide to give antimonates and in excess sulphide to give thiocomplexes. The latter is not a problem in CD since free sulphide, if it exists, does so in a very low concentration and is rapidly taken up to precipitate Sb2S3. However, the solubility in alkaline solutions limits the pH of the deposition solutions. Sb2S3 exists in two forms: so-called amorphous Sb2S3, which varies in color from

TABLE 6.1 Antimony Sulphide Solution TA/NH3/TEA/RT Annealed, 300C TA/NH3/TEA/RT Annealed, 300C TA/tartrate/RT/pH 9.5 TA/SbCl3 in CH3COOH (nonaqueous)/RT*/pH 1.5 Annealed 200C TU/RT/pH 1–1.2 Annealed  200C S2O3/SbCl3/EDTA/RT/pH 2–3 S2O3SbCl3 in CH3 COOH Annealed 170–200C S2O3/SbCl3 in acetone/10C/pH 5 Annealed 250C

Bandgap (eV)

Resistivity (-cm)

Conduct. type

1.85 ind. 1.74 ind. 1.86 ind. 1.74 ind. 1.62 ind. 1.75 dir.

5  104 250 4  108 5  106 (5  108 ) 107 5 7  10 (130C)

n-

3

n-

4 5 6 7

1.62 dir.

2  105 (130C)

n-

n-

Ref.

8 1.82 2.48* dir. 1.76 dir. 2.21 dir. 1.79 dir.

108 104–105
107
105
107

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nn-

9 10, 11 12

yellow to red, and black crystalline Sb2S3, with a bandgap between 1.7 and 1.8 eV. Thioacetamide is the most commonly used S source for this compound. Antimony sulphide deposition onto metallic substrates, together with PbS and Cu-S, was first reported in the original paper of Puscher [1] using thiosulphate and antimony tartrate. No characterization of this film was carried out nor properties given. The same method was also described “recently,” in 1931, using a number of different metals as substrates [2]. It was noted that SbCl3, when mixed with thiosulphate, reacted too rapidly, hence the use of the tartrate (tartaric acid is a complexant). Again, no characterization of the films was made.

6.2.1.1

Thioacetamide

In Refs. 3–5, potassium antimonyl tartrate was used as a source of Sb. While the deposition conditions appeared to be almost identical in all cases (other than the silicotungstic acid added in some cases in the latter two studies), there were some unexplained differences in structural and electrical properties (the latter shown in the table). The lower resistivities obtained in the first study were mirrored by higher carrier concentration (2  1015 cm3 vs. 1012 cm3) and slightly higher mobility (14 vs. 10 cm2 V1 sec1). The as-deposited film gave either no peaks or very broad ones in the XRD pattern, from which it is difficult to extract a crystal size, since there are so many close peaks in the spectrum. After annealing, well-defined peaks were observed. The films in the first study (deposited on glass) had a grain size (measured by SEM, not necessarily crystal size) of 0.12 m as deposited, growing to nearly 4 m after annealing. From Ref. 5, the grain size of the annealed films on SnO2coated glass or steel (both with well-defined crystal surfaces) was close to 1 m, but for films deposited on glass it was 0.08 m. The films in the first study were reported to be close to stoichiometric, while those in the other studies were more or less S rich. Addition of silicotungstic acid to the bath [4,5] reportedly improved the stoichiometry and increased the grain size somewhat (as well as introducing a separate WO3 phase) of the annealed films. There was no major effect of the STA on the electrical resistance. The resistivity of the annealed (not the as-deposited) films was also very substrate dependent [5]: ca. 5  108 -cm for films on glass or steel and two orders of magnitude less on SnO2-coated glass. Therefore there is no direct relationship between the resistivity and the grain size. Photoconductivity in these layers was studied [5]. The ratio between dark and light resistivities was as high as 104 (for annealed films deposited with silicotungstic acid), but this ratio decreased, and the decay time increased, with exposure to air. The ratio for as-deposited films was ca. 10. Photoconductivity decay times varied from 1 sec to several seconds. In a similar study [6], SbCl3 was complexed with tartaric acid. Also, NaOH was used instead of ammonia, and the solutions were more concentrated

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than the previous one. The films were close to stoichiometric. A lower value of indirect bandgap was measured for this study (1.62 vs. 1.85 eV). It is noteworthy that while the films exhibited no XRD pattern, the precipitate in solution gave sharp XRD peaks. Whether this is due to a different mechanism for the two routes or an effect of the glass substrate, two suggestions made by the authors, is an open question. A nonaqueous (acetic acid) bath is described in Ref. 7. While the deposition was carried out mostly at room temperature, the bath was initially heated to 80°C for 10 min to speed up the deposition. Unlike other films of Sb2S3, the asdeposited films showed an XRD pattern of well-defined peaks (although these became more numerous and with mostly different reflections on annealing). It is interesting to note, in possible connection with this observation, that the as-deposited films were described as pink, which is not typical of any form of Sb-S and suggests a mixture of a white material with a reddish one.

6.2.1.2

Thiourea

Reference 8 is the only study in which thiourea has been used to deposit Sb2S3. A methanolic solution of SbOCl at a very low pH (ca. 1) was used. This seems to be a unique case of deposition using thiourea at low pH and suggests that the reaction may proceed through some complex-decomposition reaction, since free sulphide is not expected to form under such conditions. The as-deposited films were highly scattering, white, and nonadherent, but converted to adherent, still scattering films with the typical orange color of as-deposited Sb2S3 after heating at 120°C. X-ray diffraction of this film showed no pattern; annealing at over 200°C converted the film to gray Sb2S3 with a well-defined XRD spectrum. While optical spectra were given, it is difficult to interpret them, due to the large degree of scattering.

6.2.1.3

Thiosulphate

Reference 9 involved the first thiosulphate deposition after the original reports of Refs. 1 and 2 and gives some properties of the films but no structural characterization. In Refs. 10 and 11, aqueous Na2S2O3 was added to SbCl3 in glacial acetic acid (SbCl3 hydrolyzes in water unless complexed or the solution is moderately acidic or strongly alkaline). A pH of ca. 3 was optimum; below 2.5, adhesion was poor; above 4, basic antimony salts precipitated. The solution was kept below room temperature to prevent rapid bulk precipitation. No XRD pattern was found for the as-deposited film, which was presumed to be amorphous. Annealing at 170°C crystallized the film, at least partly. The bandgap of the as-deposited film was reported to be 2.48 eV and that of the annealed film 1.76 eV. Photoconductivity was exhibited by the annealed film but not by the as-deposited one.

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A similar thiosulphate bath, only using acetone instead of acetic acid to dissolve the SbCl3 was described in Ref. 12. It was reported that the films made using acetic acid tended to flake off, and use of acetone prevented this. Photoconductivity studies showed that the photosensitivity was poor for as-deposited films (ca. a factor of 2 increase in conductivity) but increased to between 102 and 103 after optimum annealing at 300°C in N2 (annealing at 250°C or less had little effect on the photosensitivity).

6.2.2 Sb2Se3 Only one group has reported CD of Sb2Se3. The solution used was potassium antimonyl tartrate, complexed with triethanolamine and ammonia. Selenosulphate was used as the Se source. No XRD pattern was found, as for the sulphide deposited under equivalent conditions. The bandgap was 1.88 eV, and resistivity ⬇107 -cm [13,14]. Continued study of this deposition showed the effect of various parameters on deposition rate and film thickness (the latter typically reaching 1 m) [15]. This study also described some photoelectrochemical behavior of these films (Chap. 9).

6.3 ARSENIC There are at least three sulphides of As: As4S4, As4S6, and As4S10. The dimeric arsenous (III) sulphide, As4S6 (often given as As2S3) and arsenic (V) pentasulphide (As4S10) can be precipitated by H2S as yellow solids from acidic solutions of the respective As salts. These sulphides are soluble in alkaline solution and in (poly)sulphide solutions and must therefore be deposited from acidic or at most neutral solution. The pentasulphide is not a very stable compound; it is hydrolyzed by boiling water to arsenious acid and also is unstable as a solid in air above ca. 100°C. Only two papers on CD of an arsenic chalcogenide (arsenic sulphide) were found. Films were obtained using thiosulphate in an EDTA solution of As2O3 at room temperature and a pH of 2–3 [9]. No compositional information was given. A bandgap of 2.0 eV and a resistivity of 104–105 -cm were measured. As2S3 was deposited at room temperature (27°C) from an acidic (pH ⬇ 2) thioacetamide bath containing As2O3 dissolved in concentrated HCl (and in some cases complexed with EDTA) [16]. The terminal thickness (which reached a maximum and then decreased with time) was studied as a function of various deposition parameters. Well-defined XRD peaks were obtained showing the monoclinic structure (notable since this compound has a tendency to be amorphous or nearly so as deposited). A direct bandgap of 2.42 eV (similar to the standard value for As2S3) was estimated from the optical spectrum. The resistivity was ca. 105 -cm.

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6.4 BISMUTH 6.4.1 Bi2S3 (See Table 6.2) Bismuth sulphide, Bi2S3, has been rather extensively investigated. Bi3 is readily hydrolyzed in aqueous solution and is either used in acid solution or strongly complexed. Note the very low-solubility product of this compound—1098 (Table 1.1). The very low value is due largely to the large number of ions involved (five). However, even apart from this, the solubility (given as the concentration of free Bi3 and S2 ions) is low, and the solid is very readily precipitated. A glance at the various values of bandgap shown in the table is enough to see how wide this range is. These differences are often attributed either to an amorphous structure or to size quantization. However, as discussed in more detail in Chapter 10, many of these values are not dependable, and the values given in this chapter are, in many cases, more a measure of the shape of the transmission (absorption) curve than of any specific bandgap value. Deposition of Bi2S3 was first reported in 1931 from a thiosulphate bath onto metal substrates, although no details of the deposit properties were provided [2].

6.4.1.1

Thiosulphate

In Ref. 18, the initial deposition temperature was 60°C, which was then lowered to 27°C. The properties of the films were found to vary considerably with film thickness. The crystal sizes and bandgaps are discussed in Chapter 10. Films 220

TABLE 6.2 Bismuth Sulphide Solution S2O3 /EDTA/60–27C*/pH 2 S2O3 /HCHO/CH3COOH/ 7C/pH 1.4 S2O3 /HCHO/CH3COOH/ 25C/pH 1.5 TA/TEA/NH3 TA/TEA/25 or 50C/pH 8.5 TA/(EDTA)/pH1–2 TU/TEA/NH3 / 100C-RT*/pH 8–10 TU/TEA/NH3 /95C/pH 9.5 TU/NH3 /up to 90C

Bandgap (eV)

Resistivity (-cm)

Conduct. type

Ref.

2.22–1.62* dir. 1.9 dir.

5.103–106 106

n-

9, 17, 18 19

1.27 dir. 1.7 indir. 1.9 1. 6 (illum.) 1.84 dir. ca. 1.5* 1.6 1.76

105 103 ca. 106*

20 p-

105 107 -cm

n-

105–106

n-

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21 22, 23 24 25, 26 27 28

nm thick had a resistivity of 5  103 -cm, compared to 53-nm-thick films with a value of ca. 3  104 -cm. Reference 19 described a nonaqueous deposition using a mixture of acetic acid (in which the Bi(NO3)3 was dissolved) and formaldehyde (in which the sodium thiosulphate, which was insoluble in acetic acid, was dissolved). Deposition was carried out at 7°C because the resulting films were more homogeneous than at higher temperatures. No clear XRD pattern was found for the films, although one was for the precipitated powder. A more recent investigation of the previous deposition [24] showed an XRD pattern of moderately sharp peaks similar (in peak widths) to the precipitated powder but weaker (presumably due to less material) and with only some of the reflections (suggesting texturing of the film).

6.4.1.2

Thioacetamide

The majority of the various studies on Bi2S3 used thioacetamide as the source of sulphur. In the first of these studies [20], very broad peaks were observed in the XRD spectrum and the film was classified as amorphous. (A radial distribution function analysis of these films allowed a structure to be proposed [29]). Much sharper XRD peaks were obtained after a mild annealing at 150°C (for 6 hr). The broad peaks in the as-deposited film seem, for the most part, to have different values of 2 from those of the annealed film. In Ref. 21, it was noted that use of ammonia (or NaOH) resulted in films containing particulate deposits and incorporation of Bi(OH)3 into the films. Therefore the deposition solution was ammonia free, and, although alkaline, the pH was relatively low. It was also noted that use of thiourea instead of thioacetamide gave very thin films. No difference in either optical or electrical properties was found for films deposited at 25 or 50°C. The dark resistivity dropped from ca. 106 -cm to as low as 0.3 -cm (this value was apparently very variable from sample to sample) after air-annealing at 200°C, although it increased steeply again at higher annealing temperatures. The photoconductivity of these films was studied. The sensitivity was ca. 20 for the as-deposited films, increasing to ca. 100 or more after (optimum) annealing at 150°C. This temperature corresponded to the temperature at which the dark resistivity began to decrease strongly; above this temperature, the dark resistance dropped and therefore so did the photoconductivity response. The effect of various annealing treatments on the films was studied in more detail in a separate work [30]. Here, XRD showed no peaks whatsoever for the as-deposited films but strong and sharp peaks after annealing at 200°C. Specific studies of the effect of annealing on the electrical and photoconducting properties of these films annealed in argon or hydrogen [31] and in oxygen [32] have been described but will not be discussed further here. The rate of deposition of these films increased on exposure to light. This phenomenon is often observed in CD films and is believed to be due to photogen-

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erated charge formation in the growing particles, leading to an electrochemical deposition process in parallel with the CD one. This phenomenon has been exploited to generate photographic images by shining light through a mask (in this case, a photographic negative) onto the depositing film [22]. The films in this case were deposited on a predeposit of ZnS to improve adhesion. The contrast (mainly in transmission but also seen in reflection) between the thinner and thicker regions of the film reproduced the image on the negative. Changes in bandgap due to size quantization may also have contributed to the contrast (see Chap. 10). It was mentioned that better adhesion was obtained on glass if ZnS was predeposited first [22,33]. Improved adhesion was also obtained from films deposited from this bath if the glass substrates were first treated with an organosilane [22]. The silane binds to the glass surface, and the growing film anchors to terminal amino or thiol groups, both of which (the thiol in particular) bind strongly to metal chalcogenides. A fairly strongly acidic thioacetamide bath (pH between 1 and 2) was described in Ref. 23, both with and without EDTA (sodium salt) as a complexant While some structural differences as well as variations in film thickness were noted between the films deposited from baths with and without EDTA, they were not highly significant. As with most other acidic baths, the films were clearly crystalline, showing defined XRD peaks that allowed a crystal size of the order of 10 nm to be estimated.

6.4.1.3

Thiourea

Triethanolamine was used in Ref. 25 as complexant together with ammonia, the latter to slow the reaction, and it also apparently improved adherence. The solution was heated to boiling for 40 min and then left at room temperature for 4 hr. As for other depositions using this strategy, no rationale was given for this regime, although some hints as to its reason might be gleaned from the follow-up paper ([26]; see next paragraph): It can be presumed that it provided better films than those obtained by simply depositing at one temperature. No XRD pattern was obtained as for other alkaline bath depositions. In a follow-up investigation [26], the initial pH was 10.17, and this dropped somewhat during the deposition. The deposition was dependent on the age of the Bi(NO3)2/triethanolamine solution. A solution aged for ca. 5 hr before making up the deposition solution gave films about twice as thick (0.1–0.2 m) as a freshly made one. However, if aged for 24 hr, the color of the heated solution, instead of turning brown, became white, and no film formed. This was explained by increased hydrolysis of the Bi/triethanolamine solution on aging; it seems that some hydrolysis is good for the process, but too much prevents adherence and therefore film formation. A solution similar to the previous ones was used in Ref. 27. One major difference between the films obtained in this study and the previous ones is the film

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thickness—ca. 1 m thick after 30 min at 95°C, compared to an order of magnitude less after 40 min in a boiling solution, then 4 hr at room temperature. The most likely difference, in the absence of a knowledge of the triethanolamine concentration used in the present study, is that this triethanolamine concentration was lower in the present study and therefore thicker films could be obtained in a shorter time, since the the Bi would be less complexed. This is supported by the fact that the Bi(NO3)3 was triturated with the triethanolamine, which suggests that only enough triethanolamine was added to dissolve the Bi(NO3)3.

6.4.2 Bi2Se3 (See Table 6.3) There are only a few papers dealing with Bi2Se3, and therefore they will all be treated in one section. The first report was based on the early Bi2S3 depositions, using a triethanolamine/ammonia bath and selenosulphate [34]. Films were deposited both with and without hydrazine. The deposition was faster with hydrazine, as may be expected, although in this case the films lost adherence if left in the solution more than ca. 30 min. Also, while the films deposited from the hydrazine bath were single phase, traces of elemental Se were found in the hydrazine-free bath. The large difference in resistivity between films deposited from the two baths is interesting, although no reason for this difference was suggested. In contrast to the corresponding deposition of Bi2S3, where the age of the Bi/triethanolamine solution was important (see earlier), the age of the Bi solution is not critical to obtaining a good film; in fact, thicker films were obtained using even strongly aged Bi solutions [26]. This was ascribed to the low temperature used in the formation of the selenide compared to the sulphide. Films were also made using N,N-dimethylselenourea as Se source [35]. Na2SO3 was added, as usual for selenoureas, to minimize oxidation of the selenourea. X-ray diffraction showed only a very broad and ill-defined spectrum of the as-deposited film. As for Bi2S3, annealing at a relatively low temperature (200°C) was sufficient to crystallize the film and show well-defined peaks. The

TABLE 6.3 Bismuth Selenide Solution Na2SeSO3 /TEA-NH3 / 30C/pH 9.9  0.1  hydrazine Na2SeSO3 /NH3 /RT DMSeU*/TEA/RT—40C Annealed 200C

Bandgap (eV)

Resistivity (-cm)

Type

1.15*

ca. 104–105

n-

1.03* 1.42 1.7–1.4* 1.57–1.07

5  102 107 0.1

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Ref. 26, 34

28 35 n- dir.

crystal size in the annealed film, estimated from the peak widths, was ca. 12 nm. The properties, in particular the bandgap calculated from the optical spectra, were dependent on the film thickness. Thus the range of bandgaps given in the table vary from a film 0.09 m thick (higher bandgap) to 0.15 m (lower bandgap). This is a very large range for such a small change in thickness. Since the variations are probably due to size quantization, a measure of the XRD peak widths for the different thicknesses (of minimally annealed films) would be of interest here. The as-deposited films were very photosensitive (light-to-dark ratio up to 70), but the annealed films were much less so.

6.5 COBALT In the early work by Beutel and Kutzelnigg on film formation on various metals from hot thiosulphate solutions of various metal compounds, one of those that resulted in apparent CD was Co [2]. Since no characterization of the films was given and it is possible that the coloration was due to reaction between the thiosulphate and the substrate metals, this author made a single experiment, mixing Co2 with excess thiosulphate and heating in a glass vial. A black film formed on the glass (this was also carried out for Ni and Fe; see later). Although no characterization was made of this film, it indicates that Co-S was indeed formed in the experiment of Ref. 2. CoS was deposited at room temperature from a triethanolamine/ammoniacomplexed solution of CoCl2 using thioacetamide as sulphur source [36]. Both compositional analysis (CoS1.035) and XRD analysis showed the formation of CoS. From the optical spectrum, a direct bandgap of 0.62 eV was found. The films were p-type with a resistivity of ca. 106 -cm. Metallic grey-brown films of Co3S4 were deposited from a CoCl2/NH3 solution using thiourea at temperatures between room temperature and 50°C [36a]. The film resistivity was approximately 105 -cm. CoSe was deposited using a similar composition to that for CoS, except that selenosulphate was used in place of thioacetamide, NaOH and hydrazine were also added to the solution, and the deposition was carried out close to 100°C [37]. In contrast to CoS, no XRD pattern was observed for the as-deposited films; annealing at 280°C crystallized the films to give a defined pattern corresponding to CoSe. A direct bandgap of 0.45 eV was estimated from the optical spectrum. The films were p-type and the resistivity 104 -cm.

6.6 COPPER Copper chalcogenides can be readily deposited by CD. There is a strong affinity between Cu and S or Se; metallic Cu exposed to elemental S dissolved in a solvent (e.g., dimethyl sulphoxide) will quickly turn black due to formation of cop-

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per sulphide. This strong affinity between the elements is manifested as a low-solubility product of the various Cu chalcogenides. A complicating factor in this deposition, however, is the large number of different phases and stoichiometries that can exist. Four well-defined room-temperature phases of the CuxS system are known: chalcocite (x ⬇ 2); djurleite (x ⬇ 1.96); digenite (x ⬇ 1.8), and covellite (x
1). Mixed phases can also occur, of course. Since only CuS gives an identifiable XRD pattern in the CD Cu-S films, a knowledge of which phase(s), other than CuS, exists is unknown, although the average composition has been measured in many cases. The various phases can, in some cases, be changed from one to another after deposition. This has been shown for Cu-Se using electrochemical polarization [38] and by aging and then reversed by heating [39]. Treatment of CuxS film with gaseous H2S apparently changes the composition (possibly also the phase) toward a more S-rich film [40]. Cu-S shares with PbS and Sb-S the distinction of being the first published CD compound [1]. This and (for many decades) subsequent reports involving CD Cu-S described decomposition of thiosulphate solutions of Cu salts to give Cu-S films. These (and other, mainly PbS) films were known as lüsterfarben (lustrous colors) due to the varied interference colors obtained on metal substrates by deposition of PbS or Cu-S (see Sec. 2.1 and Sec. 5.2 for more details of the history of these lüsterfarben). As for PbS, very little characterization of these deposits was reported in those early papers apart from their various colors. If this early work on CD of Cu-S was driven by the attractive colors they imparted to metallic substrates, more recent studies were driven initially by their potential use in Cu2S/CdS photovoltaic cells (these cells are no longer studied to any extent due to their perceived instability, although, with what has been learned about Cu-containing chalcopyrite, such as CuInSe2, thin-film cells over the past couple of decades, it would not be surprising if such studies were again pursued) and later in solar control coatings (see Sec. 2.13). Thiosulphate and sulphite are sufficiently reducing to reduce Cu2 to Cu. Therefore the Cu in solutions of Cu2 containing sufficient thiosulphate, selenosulphate, or sulphite should be predominantly in the monovalent form. This would lead to the expectation that the main product will be something close to Cu2S(e). While this is often the case, CuS(e) is deposited in some cases. However, it is arguable whether this reduction of Cu2 is, in fact, important in practice. The reason is based on an XPS study that showed that Cu in its compounds with S, Se, and Te is normally in the monovalent state; it is the chalcogenide ion (or polyion) that is believed to change oxidation states in these compounds [41]. An interesting characteristic of CD Cu-S films deposited from thiosulphate solution is the range of compositions that can be obtained by varying the deposition conditions [40]. Elemental analyses of the precipitated CuxS powders obtained by heating a solution 0.1 M each in CuSO4 and Na2S2O3 showed that the composition varied from x
1.7 to x
1.0, with longer reaction times and higher

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temperature giving lower values of x. X-ray diffraction of these end-members was consistent with a phase where 1.86  x  1.96 (therefore the Cu1.7S was a mixture of predominantly this phase, presumably djurleite, and other phases) and x
1.0, respectively. XPS of films deposited from the same solutions showed the sulphur of the CuS to have a higher binding energy (more oxidized) than that of the Cu1.7S, in agreement with a variable valence state of the chalcogen in these compounds. Variation of the Cu:thiosulphate ratio also affected the composition. Compositions corresponding to Cu2S (Cu:thiosulphate 1:2); Cu1.8S (1:1); Cu1.4S (1:2.5), and CuS (1: 3) were obtained, although in film form, only CuS gave an XRD pattern [42,43]. The variation in composition is not a monotonic function of the Cu:thiosulphate ratio. For equimolar Cu and thiosulphate, there is not enough thiosulphate to both reduce all the Cu2 and form the sulphide, hence a partially reduced Cu2 is formed; while for a ratio of 1:2, there is enough thiosulphate to both reduce all the Cu2 to Cu and to form the sulphide. More difficult to understand at first sight is why an increasing concentration of thiosulphate apparently oxidizes the Cu-S to an increasingly greater amount of CuS, until at a ratio of 1:3, only CuS is formed. However, it must be kept in mind that CuS is not simply composed of divalent Cu, but is either 2CuS 2 2 or a mixed-valence compound with both S2 and S 2 2 groups. Thus it may be more useful to consider the effect of the thiosulphate on the sulphur species. A possible hypothesis is that excess thiosulphate results in the formation of elemental S, which can react with S2 to give the polysulphide ion, S 2 2 , which exists in CuS. More generally, and regardless of the specific mechanism, a higher concentration of S relative to Cu can be expected to favor more S in the final product. The various Cu-S and Cu-Se films generally exhibit similar optical spectra for comparable thicknesses. Figure 6.1 shows some such spectra. The peak at ca. 0.6 m is characteristic of these films, and the drop in transmission at longer

FIG. 6.1 Optical transmission spectra of various CuxS films. (a) Cu2S; (b) Cu1.8S; (c) Cu1.4S; (d) CuS. (Adapted from Ref. 43, with permission from Elsevier Science (USA)).

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wavelengths is presumably due to absorption and reflection by free holes (CuxS(e), as for the other Cu-S(e) compounds, is p-type and in most cases highly degenerate, therefore relatively highly conductive). The CuS (covellite) phase tends to be somewhat more conductive than the other sulphide phases. CD Cu-S(e) films have been proposed for a number of different potential applications. Solar control coatings, where the visible and IR transmission and reflectivity can be varied, is probably the most studied, e.g., Refs. 44 and 45. The relatively high conductivity and the partial transmittance in the visible spectrum are useful for transparent conductors [46]. Other possible applications are for Cu2 sensor electrodes and electrical contacts for ceramic devices [46]. In the tables for both Cu-S and Cu-Se (Tables 6.4 and 6.5), the column denoting conductivity type has been deleted (these semiconductors are always ptype), and, in its place, the phase (and/or composition) has been given. In some cases, particularly for the sulphides, where no XRD pattern was seen (except for CuS), no phase (or composition) was proven and therefore no entry is given in the table. This was not a problem for Cu-Se, since XRD spectra were always clear and definitive.

6.6.1 Cu-S (See Table 6.4) In Ref. 9, no structural or compositional characterization was given; Cu2S was assumed. The films were not uniform and did not adhere to the substrates. The high values of resistivity are unusual for this material: it may be that 103–104 was intended. TABLE 6.4 Copper Sulphide Solution

Bandgap (eV)

S2O3 /RT/pH ca. 2 S2O3 /60C/pH 0.5 S2O3 /60 or 70C/pH 2–5 S2O3 /50C/pH 5 (AcH)

1.2 2.4 dir.

S2O3-dimethylthiourea Anneal 300C Anneal 400C TU/NH3 /30°C TU/TEA-acetate-NH3 / 40C/pH 9.4 CuCl/TU/EDTA/NH2OH/ RT—80C/pH 8.5–11.5 TU/TEA-NH3 /RT

1.55 indir. 1.55 indir. 1.4 indir. ca. 2.26 ind. 2.58 dir. and ca. 1.7* 1.45

1.7–2.0

ca. 1.5 indir.

Resistivity (-cm) 104–105* 104 2 10 –104* 103 (CuS) 3  103 (Cu2S) 8  103 8  104 103 0.25 3  103 103–104 0.004  1

Phase or composition

Varied Varied CuS Cu1.8S Cu1.96S Cu1.86S CuxS 1.83  x  1.85

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Ref. 9 47 40 42, 43 48

49 50 51 52

In Ref. 47, as before, no structural or compositional characterization was given, and the films were classified as CuxS. Average grain size, measured by scanning electron microscopy (SEM), was ca. 70 nm. In Ref. 40, films were deposited on various polymer substrates that in most cases required pretreatment with either various organosilanes or poly(ethyleneimine) in order for good film formation to occur. This technique appears to constitute a case of adsorption of colloids from the solution rather than formation of Cu-S directly on the substrate. For example, MNH groups on the iminecoated substrate become protonated in the acidic solution and attract the apparently negatively-charged Cu-S particles. The composition varied with temperature and time of deposition (see the previous general discussion). The resistivity of the films varied with the deposition conditions and substrate; CuS was more conducting than compositions closer to Cu2S. In Refs. 42 and 43, composition depended on the Cu:thiosulphate ratio (see the earlier general discussion). Only CuS gave a measurable XRD pattern. CuS was less transmitting in the IR region than CuxS (x
1.4–2.0), and for CuS, greater thiosulphate concentration (constant [Cu]) resulted in less transparency in the IR, although film thickness was fairly constant. This is expected from the data of Ref. 40, where greater thiosulphate concentration resulted in a deposit closer to CuS in composition and with higher electrical conductivity (therefore less transparent in the IR). A combination of thiosulphate and dimethylthiourea was used in Ref. 48. Using only thiosulphate resulted in slower deposition and, more importantly, poorly adherent films. The resistivity of the as-deposited CuS dropped to 3  104 -cm on annealing at 200°C (in N2) without a change in phase or composition. Annealing at 300 and 400°C resulted in loss of S and phase changes. The crystal size was 11 nm (200°C anneal), 19 nm (300°C), and 20 nm (400°C). Presumably no clear XRD pattern was obtained for the as-deposited film of CuS. In Ref. 49, the composition of the films was given as Cu1.8S. However, no XRD pattern was found and no compositional analysis given, and therefore it is unclear just what the actual composition and phase were. A potentiometric technique was used in Ref. 50 to measure composition, found to be Cu1.86S. In this deposition, stirring the deposition solution resulted in nonuniform and poor-quality films, while good films were obtained in unstirred films. The bandgap was measured to be direct, with a value of 2.58 eV. This is a particularly high value. Examination of the transmittance spectrum showed a sharp drop in transmission at ca. 1.7 eV, which is more likely to be the true bandgap. In Ref. 51, the composition was CuxS, with 1.83  x  1.85. The films were thicker than most others—several microns. The terminal thickness was very pH dependent; at pH 8.5, it was ca. 0.5 m, and at pH 11, an order of magnitude

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greater. A pH of 10 was optimum in terms of maximal thickness without excessive bulk precipitation. No structural or compositional data was given in Ref. 52. Although a value of bandgap was not given, from the optical spectra, an indirect bandgap of between 1.5 and 1.6 eV could be estimated. The specific resistivity dropped with increasing film thickness by more than two orders of magnitude, ⬇ 1 -cm for a thickness of 0.15 m to 4  103 -cm for 0.35 m. It was suggested that this was due to increasing nonstoichiometry as deposition proceeded. If a second deposition was carried out on a previous one, the resistivity (measured across the film) was characteristic of a single layer, suggesting that a relatively insulating layer was deposited on the first layer in the early stages of the second deposition. This group published a number of other papers on CuxS, using the same basic deposition solution, with emphasis on variations of their spectral properties for possible use in solar control coatings. Some examples are given in Refs. 44, 45, and 53 (the last deals with films deposited on Kapton foil). This last showed some weak CuS (covellite) peaks in the XRD spectrum; no inference of whether this was the major phase or not could be made.

6.6.2 Cu-Se (See Table 6.5) The three main phases encountered in CD Cu-Se films are Berzelianite (Cu2xSe, where x is typically ca. 0.2); Umangite (often written as Cu3Se2 but may be considerably lower in Cu) and Klockmannite (CuSe). In Ref. 54, XRD showed the deposit to be hexagonal CuSe. Analysis of the absorption spectrum gave a direct bandgap of 2.02 eV. As commonly seen for these compounds, there was still strong absorption at lower energies (e.g., at 1.9 eV, the absorption coefficient was 7  104 cm2), possibly due to an indirect transition but likely due, at least in part, to free-carrier absorption. From Hall measurements, the doping (acceptor) density was ca. 1022 cm2 (heavily degenerate) and the mobility ca. 1 cm2V1sec1. The dependence of film thickness and deposition rate on the deposition parameters has been studied in a separate paper [62]. Nitrilotriacetate was used as complexant in the deposition in Ref. 55. Cu-Se could be both electrodeposited and chemically deposited from this solution. The electrodeposited film was Cu1.8Se with the berzelianite structure, while the CD one was Cu1.2Se with the umangite structure. The XRD pattern of the CD films showed sharp peaks (instrument broadening) with no preferential texture. Electron microscopy of these films (Fig. 6.2) shows large (micron scale) particles that, from their faceted shape and together with the sharp XRD peaks, appear to be single crystals. This is a particularly large crystal size for a CD film; from this and

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TABLE 6.5 Copper Selenide Solution SeSO3/TEA  NH3/RT SeSO3/NTA*/RT/pH
9 SeSO3/TEA  NH3/
95C/pH ca. 10 SeSO3/NH3/
45C/ pH
10 SeSO3/pH
10 SeSO3/TEA  NH3 /75C SeSO3/NH3/RT SeSO3/citrate/5–27C/ pH
7 SeSO3/citrate/60C pH 9 pH 12 DMSeU/tartrate/50C DMSeU/tartrate/50C

Bandgap (eV)

Resistivity (-cm)

ca. 2.0 dir.*

103

1.20 dir.


2.36 dir. 1.9 indir.* See text

2.37 2.0 ca. 2.13

Phase

Ref.

103–2  102

CuSe Cu1.2Se umangite Cu1.86Se berzelianite

54 55 56

ca. 10 -sq.

Cu2Se

57

3  103–104* 2  104

Cu2Se or CuSe* Cu2xSe berzelianite* Cu2xSe berzelianite

58 38 59

ca. 2  103 (for all films)

Cu2xSe berzelianite Cu3Se2 umangite

39

60

103–104* 2  104

Cu3Se2 umangite CuSe klockmannite CuSe CuSe klockmannite

61 59

the other studies on CD Cu-Se, it can be seen that Cu-Se has a tendency to form relatively large crystals. This may be due, at least in part, to the high mobility of Cu (although the tendency is much less for Cu-S) and the relatively low melting points of the Cu-Se compounds in general.

FIG. 6.2 Scanning electron micrograph of a Cu1.2Se film deposited from a selenosulphate solution of Cu2 complexed with nitrilotriacetate.

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The main difference between Ref. 56 and Ref. 54 is the higher deposition temperature of the former. Elemental analysis gave a composition of 65% Cu and 35% Se (at.%) corresponding to Cu1.86Se. X-ray diffraction confirmed the cubic berzelianite phase. The deposition in Ref. 57 appears similar to those of Refs. 54 and 56, except that only ammonia was used to complex the CuSO4 and the deposition temperature was ca. 45°C. The films were deposited on polyester (overhead transparency) films. From XRD, a composition of Cu2Se (no specific phase given) was assigned. The transmission spectrum was similar to those of Cu-S in general, with a strong free-carrier absorption beginning in the near IR and a strong absorption (presumably bandgap) onset at ca. 700 nm. Reference 58 involved a similar deposition as the previous case, but without ammonia. The composition of the Cu-Se was Cu2Se (using a Cu:selenosulphate ratio of 1:1) and CuSe (1:5). (This is the same trend as found for sulphides deposited from thiosulphate solutions.) The transmission spectrum of the Cu2Se was similar to that in the previous study, while that of the CuSe showed an absorption onset at ca. 600 nm. Both films showed strong apparent free-carrier absorption starting in the near IR (for the CuSe, even at somewhat shorter wavelengths). However this absorption appeared to be stronger for the Cu2Se than for the CuSe, although the resistivity of the former (3  103 -cm) was higher than that of the CuSe (104 -cm). Copper acetate was used in Ref. 38; it was noted that if chloride was used instead of acetate, no deposition occurred, and this was attributed to adsorption of chloride on the substrate (Pt). The berzelianite phase with a small amount of umangite impurity was obtained. The composition and phase of the film could be altered by electrochemical cathodic polarization (in an aqueous K2SO4 solution). Initially, there occurred an increase in lattice parameters and decrease in x (Cu2xSe). With continued polarization, a phase change occurred until eventually only orthorhombic Cu2xSe was present in the film. The umangite phase also disappeared, and it was believed that this impurity phase catalyzed the phase transformation. The change in composition during cathodic polarization was attributed to reduction of zerovalent Se to Se2, which was dissolved in the solution. Based on the study of Folmer and Jellinek [41] discussed earlier, this explanation can be 2 interpreted as reduction of Se 2 (divalent Se). 2 (“monovalent” Se) to Se Reference 59 provides a comprehensive explanation of the optical spectra and extracted bandgaps. The direct bandgap of ca. 2.36 eV is compared to the literature value of ca. 2.2 eV and explained by size quantization in the fairly small (20 nm) crystals. An indirect bandgap of 1.9 eV was measured (literature value  1.4 eV), but it was stressed that this provided an upper limit only, since the absorption in this region was dominated by free-carrier absorption, which masked the indirect absorption. Annealing decreased the conductivity and the free-carrier absorption and changed the indirect gap to ⬇1.3 eV.

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In Ref. 39, citrate was used a complexant and the pH was lower than in other similar studies. The depositions were carried out at around room temperature or, if the deposition was too fast, at lower temperatures (no difference in the nature of the films was found with different temperatures). The composition and structure of the deposit was found to be pH dependent: At a pH ca. 7 or lower, berzelianite Cu2xSe was deposited, while at a slightly higher pH (7.8), the product was umangite (ca. Cu3Se2). The exact compositions varied with change in ratio between Cu and selenosulphate concentrations. As with the other Cu-Se and Cu-S films, bandgap determination was complicated by the strong free-carrier absorption. Direct bandgaps of 2.2 eV (Cu2xSe) and 2.8 eV (Cu3Se2) were measured from the transmission spectra. However, from examination of these spectra, it can be inferred that a strong absorption, not arising solely from free carriers, occurred at lower energies. An approximate reanalysis of the transmission spectra, taking into account free-carrier absorption, allowed estimation of indirect bandgaps of 1.5–1.6 eV (Cu2xSe) and 2.0–2.1 eV (Cu3Se2). The berzelianite phase was subsequently found to slowly transform to the umangite one under ambient conditions [63]. By heating at 140°C in air, this phase transformation could be reversed. These phase changes could be repeated in a cyclical manner. In Ref. 60, the differences in the two solutions giving Cu3Se2 and CuSe were the lower pH (9) and higher Cu and citrate concentrations (6 mM) for Cu3Se2, compared to pH
12 and Cu (and citrate) concentrations of 4 mM (constant selenosulphate concentration of 30 mM in both cases). The films were deposited on flexible polyester substrates. It was noted that the deposition was paralleled for the most part by bulk precipitation. The average crystal size of both Cu3Se2 and CuSe was 42 nm. In Ref. 61, N,N-dimethylselenourea was used (together with CuCl2 instead of CuSO4 commonly used in the selenosulphate depositions). Film (specific) resistivity dropped as thickness increased. Reference 59 is similar to the previous study. Values for film resistivity were not given, but it was noted that the films were less conductive than Cu2xSe films (2  104 -cm) made in the same study using selenosulphate instead of the selenourea.

6.7 INDIUM Indium is very readily hydrolyzable, with a pKa of 4.0, and forms the hydroxide even in moderately acidic solutions (see Sec. 1.1.2). This means that unless deposition is carried out in strongly acidic solution, some hydroxide is likely to be present in any chalcogenide formed by CD. This is indeed the case in most studies of In2S3 deposition reported up to now. Films of In2S3 on glass deposited from a solution of InCl3 and thioacetamide were described as early as 1976 by the Kitaev group [64]. Broad XRD Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

peaks corresponding to -In2S3 with a crystal size of ca. 5 nm were obtained. A bandgap of 2.45 eV could be estimated from the optical spectra. The resistivity was very high—ca. 1012 -cm at room temperature—with an activation energy of 1.2 eV (implying a midgap Fermi level and highly intrinsic material). Films annealed at 250°C were photosensitive, with a photoconductivity maximum at ca. 500 nm (2.5 eV). It was noted that this was blue-shifted by 0.13 eV, compared to single-crystal In2S3. The small size of the crystals suggests that size quantization occurs here. In2S3, or, as shown later more probably In(OH,S), was deposited by CD for use as a buffer layer in photovoltaic cells [65,66]. The deposition bath was again InCl3 and thioacetamide operating at a temperature of 70–80°C. Analysis of these films by XPS showed that oxygen was present in the films, presumably as hydroxide [67]. Importantly, the results were inconsistent with a mixture of sulphide and hydroxide (which might be expected from this bath) and suggested rather some compound formation. (Details of photovoltaic cells using these films are given in Chap. 9.) In the first of a series of studies, the same basic bath as previously, but using acetic acid to adjust the pH (to ca. 3, probably somewhat lower—see following reference) was used [68]. The main purpose of the acetic acid is probably to lower the pH and therefore to reduce the In3 hydrolysis, although its (weak) complexing ability with In3 might also play some role in minimizing this hydrolysis. The color of the film (and homogeneous precipitate) varied from whitish yellow to yellow (the bandgap of In2S3 is ca. 2.4 eV), and XPS analysis showed that the S:In ratio was always less than that expected for stoichiometric In2S3 (1.5), although it was higher when the thioacetamide and acetic acid concentrations were increased. Higher thioacetamide concentration increased the concentration of sulphide formed, while more acetic acid decreased the hydrolysis to hydroxide. The resulting films were therefore believed to be composed of both sulphide and hydroxide, designated as Inx(OH)ySz. The bandgaps (indirect) varied between 2.0 eV (highest S content) to 2.5 eV. For the film with the highest S content, the value of 2.0 eV is probably an underestimation, particularly taking into account the yellow color typical of a film with a bandgap of at least 2.3 eV. The films were highly resistive, between 107 and 108 -cm. An XPS investigation of these films was carried out [69]. The pH was more accurately measured to be between 2.2 and 2.5. Also it was noted that, although higher concentrations of acetic acid minimized the codeposition of hydroxide, above 0.1 M acetic acid, the films were not homogeneous and poorly adherent. From the XPS spectra, it was concluded the films were of the composition In(OH)S, with small variations in the S:OH ratio. Sulphate, probably as surface oxidized In-S, was also present. Structural (XRD and microscopic) studies of the films allowed more definite assignments to be made as to the identity of the films [70]. The structure was dependent on the thioacetamide and acetic acid concentrations. At low concentraCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

tions a composition identified as In5S4 was obtained, while at higher concentrations a mixture of cubic - and -phases resulted. The former possessed a stringlike morphology, while the latter was typically composed of more or less spherical granules. The XRD peaks of the In5S4 deposit were broad, equivalent to a crystal size of ca. 5 nm, verified also by TEM. The In3S3 deposits were apparently of somewhat larger crystal size (10 nm). A study of the species present in these solutions and the mechanism of the deposition has been presented [71]. Under the conditions of the depositions, the main solution indium species (in the absence of thioacetamide) are In-Cl (mainly [InCl2]) complex species. Only ca. 1% of the total In content is present as free In3. No In(OH)3 or hydroxy-complexes were calculated to be present if acetic acid was present (in the absence of acetic acid, the hydroxide could form). From a kinetic analysis of the deposition reaction, it was concluded that the deposition occurred by direct reaction between the thioacetamide and the chloro-indium complexes. It was noted that thioacetic acid was the main by-product and that no acetamide was detected (see Sec. 3.2.1.3 for a description of the possible mechanisms and by-products of thioacetamide hydrolysis). Acetonitrile (CH3CN), a less common by-product, was also detected at the higher pH values (these depositions took place between a pH of 2 and 3) but not at the lower ones. A different study, using essentially the same deposition solution (InCl3  thioacetamide) at a pH of 3.1 or lower, has been described [72]. The films, deposited on ITO/glass adhered to the ITO side but not well to the glass. In this case, compositional analyses showed the films, which gave electron diffraction patterns and XRD spectra characteristics of In2S3 (mainly the -phase, possibly together with the -phase) to be slightly S rich (the precipitate formed in solution was more or less stoichiometric). Thus, although there was no evident difference in the deposition parameters, these films appear different than the mixed sulphide-hydroxide ones described previously. Microscopic investigation showed the films to consist of a mixture of round particles and needles, forming a porous, spongelike morphology. The films exhibited an increase in bandgap, together with decrease in crystal size with decreasing deposition temperature, due to size quantization (see Chap. 10 for more details). Decrease in solution pH also resulted in a decrease in the bandgap. The bandgap varied between ca. 2.3 and 2.7 eV.

6.8 IRON No clear-cut example of Fe-S has been described. The closest is in the report of Ref. 2, where, among other metal salts, a boiling aqueous solution of iron thiosulphate imparted coloration to an iron substrate [2]. No confirmation was given as to the composition or nature of this coloration, and a few attempts by the author to deposit Fe-S on glass by heating thiosulphate solutions of iron salts were unsuccessful (unlike the corresponding Co and Ni cases). Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

6.9 MANGANESE The stoichiometry of Mn-S precipitated from solution is normally MnS. The stable form of MnS is green -MnS, which has the rocksalt structure. However, the pink form, which is the form that usually precipitates from solution, is a mixture of -MnS (zincblende) and -MnS (wurtzite), both of which are metastable. MnSe behaves analogously. MnS was deposited from a room-temperature solution of Mn(II) acetate complexed with triethanolamine and buffered with NH4Cl [73]. Thioacetamide was used as a sulphur source, and hydrazine was also used (it was not specified whether the reaction proceeded in its absence). No XRD pattern was seen in the as-deposited (grey-pink) film; annealing at 500°C in an inert atmosphere gave a pattern corresponding to MnS. A bandgap (indirect) of 3.25 was measured from the optical spectrum. The film was p-type with a resistivity of ca. 105 -cm. Optimization of the film growth from the foregoing bath was carried out [74]. In contrast to many other CD reactions, the growth rate decreased slightly with increasing temperature (also, the terminal thickness was greater at lower temperature—a common occurrence due to reduced bulk precipitation). It may be that bulk precipitation was so rapid at higher temperatures that the thickness of the film deposited at the higher temperatures was less by the time of the first thickness measurement (10 min). While, as with the previous study, no XRD pattern was seen in the as-deposited film on glass, a clear pattern was observed for films deposited on SnO2-conducting glass, showing a mixture of the cubic - and hexagonal -phases. However, TEM/ED showed the presence of 3- to 4-nm-sized MnS crystals on glass. The optical bandgap (for the film on glass) was estimated to be ca. 3.0 eV and direct. There does not seem to be a clear-cut value in the literature for the bandgap of MnS, but it is has been given as 3.0  0.2 eV. The films showed only weak photoconductivity. MnS has also been deposited from an alkaline (pH 9.7–9.8) thiosulphate bath, using MnCl2 [75]. The deposition was carried out at room temperature after initial heating at 70°C (this initial heating step was noted to be essential, although no explanation for this was given). The XRD spectrum was barely indistinguishable from the noise; there was a possible correlation with the spectrum of -MnS. Optical studies showed a direct bandgap of 3.1 eV, and the resistivity was measured to be between 107 and 108 -cm.

6.10 MOLYBDENUM There are two main sulphides of Mo. The stable form is the black, layered MoS2, commonly used as a lubricant. Precipitation from (acidic) solution normally gives the amorphous MoS3, which converts to MoS2 on heating. An important issue when using molybdates as a source of Mo is that solutions of molybdates do not precipitate the sulphide (selenide) when sulphide (selenide)—either as H2S(Se) or Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

as an alkaline sulphide (selenide)—is reacted with the molybdate solution, but rather form soluble thio(seleno)molybdate ions, such as MoS 2 4 . Chemical deposition of both MoS2 and MoSe2 has been reported from ammonium molybdate solution [76,77]. For the sulphur and selenium sources, thioacetamide and selenosulphate were used, respectively. Ammonium hydroxide was added to the sulphide solution, while an acetic acid/ammonium acetate buffer was used with the selenide solution (pH values were not given). Reducing agents (either hydrazine [76] or sodium dithionite [77]) were added to the baths. Deposition was started at 90–100°C, followed by lowering to room temperature. No XRD pattern was observed for the Mo-S deposit, but after heating (apparently in the deposition solution) in an autoclave at 300°C, the XRD pattern of MoS2 was obtained. The XRD pattern of MoSe2 was obtained for the as-deposited film. It is possible that the as-deposited Mo-S was MoS3, which is often obtained in an amorphous form from solution reactions at relatively low temperatures and converts to crystalline MoS2 on annealing. The estimated bandgaps for the two materials were 1.17 eV indirect (MoS2) and 1.14 eV direct (MoSe2). The latter is unusual, since this is the approximate value of the indirect gap of MoSe2; the direct gap is substantially higher. While the role of the reducing agents (hydrazine and dithionite) was not explicitly discussed, it must be assumed that they play an essential role in forming the Mo chalcogenides rather than the soluble thio(seleno)molybdate ion. MoSe2 was deposited from a Mo(VI) (the source used was not specified) solution complexed with ammonia to give a hexammine complex and mixed with hydrazine and selenosulphate at 40°C [78]. The as-deposited films were XRD amorphous but converted to crystalline MoSe2 after annealing in N2 at 380°C. Elemental analysis showed the as-deposited films to be nearly stoichiometric MoSe2. A direct bandgap of 1.48 eV (1.36 after annealing) was measured. The films were n-type with a resistivity of ca. 4  103 -cm (ca. 1 -cm after annealing).

6.11 NICKEL Ni-S behaves rather similarly to Co-S (see Sec. 6.5, Cobalt). Note that the freshly precipitated monosulphides of both metals transform in solution to a more insoluble form–possibly M(OH)S. As for cobalt (see earlier), an early study of the coloration of metals by immersion in boiling metal salt–thiosulphate solutions resulted in coloration of the metals [2]. Also as for cobalt, a single experiment by the author repeating (similar, not necessarily identical conditions) this experiment, only on glass, rather than on metals which might be colored by the thiosulphate alone, resulted in a black film on the glass. Other than the foregoing, only one other paper was found dealing with NiS (and also NiSe) [79]. The baths were based on NiSO4, triethanolamine, and am-

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monia. For the sulphide, thioacetamide was used, while for the selenide, selenosulphate was the Se source, and NaOH and hydrazine were added to the bath. Depositions were carried out at room temperature. XRD confirmed the formation of NiS and NiSe films. The bandgaps (direct) were 0.35 eV (NiS) and ) 0.23 eV (NiSe). The films were both p-type, with resistivities of 10 -cm (NiS) and 0.1 -cm (NiSe).

6.12 SILVER 6.12.1 Ag2S (see Table 6.6) The Ag ion forms strong complexes with thiourea (log K
12.7 for the Ag(thiourea) 3 complex and with thiosulphate (log K ⬇ 13). The strong binding of Ag to thiosulphate is exploited in the use of Na2S2O3 solution to remove excess Ag during the developing of photographic films. For this reason, Ag can

TABLE 6.6 Silver Sulphide Solution S2O3 /RT/pH 2–3 S2O3 /RT/pH 2.6 (2.2)*  EDTA S2O3 /NH3 /50C/pH 9–11 TU/S2O 32/RT/pH 10.1* TU/8–55C/pH 8–10 (NH3)  EDTA TU/NH3, kinetic study TU/NH3, thermodynamic analysis TU/NH3 /NH 4 /pH ca. 11, structural study TU/40–80C/pH 9, study of mechanism TU/40–80C/pH 9 TU (NH4OH) Hg2 doped TA/RT/8–55C  EDTA Dip technique Ag/S2O3 /RT  TU/80C/pH 8–11*

Bandgap (eV)

Resistivity (-cm)

1.2 0.95 dir. 0.73 dir.
2.2 2.3 dir. — ca. 0.8 dir.

106 –107 103–104 104–105
10 102 103–105* 2  104

Conduct. type

Ref. 9 80

pn-

11 81 82, 83 84 85 86 87

ca. 1.0* 0.8 (from photocond.) 0.95 dir. 0.83 indir.

ca. 103 (dark)

n-

88 89, 90

ca. 105 ca. 104 (100 )

n-

91

n-

92

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be kept in solution under alkaline conditions without the need for another complexant. The strong binding of Ag to sulphur in thiourea and thiosulphate suggests that the mechanism of Ag2S formation may be of the complex-decomposition type rather than through formation of free sulphide. Thus any bond breaking involving the Ag-complex is intuitively expected to occur at the SMC or SMS bonds of the thiourea or thiosulphate, respectively (see Sec. 3.3.3.1, Eqn. 3.5.5 and discussion following for more details on this topic). Two forms of Ag2S exist—-Ag2S (acanthite), a monoclinic form, and Ag2S (argentite), which is cubic. -Ag2S is the form that is stable at room temperature and is invariably the one that occurs in CD films. Thiosulphate was used in an acid bath in Ref. 9. This study covered many sulphides deposited using thiosulphate, and little detail was given on the deposition or on the films themselves. Equimolar quantities of Ag and thiosulphate were used in Ref. 80 (as for the previous study), so the complexation of the Ag by thiosulphate was not as strong as it would have been in an excess of thiosulphate. A suggested mechanism for the deposition was reduction of elemental S to sulphide, formed in the acidic thiosulphate solution, by the moderately reducing thiosulphate. It was stressed that the thiosulphate was slowly added to the AgNO3 solution with heavy stirring, with the implication that otherwise the thiosulphate would be oxidized. The films were reported to be rough rather than the smooth specular films often obtained by CD. Differences in properties were obtained if EDTA was added as an additional complexant. The films with EDTA were somewhat thinner (0.14 instead of 0.19 m), and XRD of the EDTA-free films gave sharp peaks, while those deposited from an EDTA-containing bath apparently showed no pattern, hence were probably either very small crystalline or amorphous. As seen from Table 6.6, both apparent bandgap and resistivity changed on addition of EDTA to the bath. The higher resistivity with EDTA was explained by the smaller crystal size. The lower bandgap is less obvious; very small crystal size would increase the bandgap, and an amorphous semiconductor has often (although not always) a higher bandgap than the crystalline form. The deposition in Ref. 11, from an alkaline thiosulphate bath, was reported in the context of a general description of deposition of various materials by CD, and only a little characterization was reported. X-ray diffraction showed some Ag2S peaks. Optical spectroscopy showed a gradual decrease in transmission over a wide spectral range, and it would be difficult to extract a reliable value for the bandgap from the spectrum. Note that the deposition in Ref. 81 used thiosulphate as a complexing agent and not ostensibly as a source of S. The thiourea concentration is critical. Thiourea is added to the Ag/S2O3 just until some solid Ag2O is formed. Too little thiourea results in thin, scattering films. Too much results in films that in their as-deposited state are good, but after annealing (300°C), voids form. The pH is also critical:

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Under 10 the deposition is slow, while above 10.2 it is too fast and bulk precipitation dominates. In Refs. 82 and 83, this deposition was employed both with and without EDTA, a strong complexant. The deposition was studied at various deposition temperatures. Better adhesion was obtained at low temperature (8°C). The resistivity of the films was dependent on deposition temperature: 2.103 -cm (8°C) and 1.5  105 -cm at 25°C. For a film deposited from an EDTA-containing bath at 8°C, it was ca. 2  104 -cm. The reaction kinetics were analyzed in an early study [84]. Of particular note is the unusually high activation energy (160 kJ/mole)—about twice the normal value for reaction-controlled CD processes. This contrasts with the much smaller value (20.4 kJ/mole) measured in Ref. 87, although there were several differences in the deposition: The present solution contained ammonia, the pH was higher—probably between 11 and 12 (compared to the borate buffered solution with pH
9 of Ref. 87), and the activation energies were measured at lower temperatures. It is interesting that the pH could not be measured directly in this study using a pH meter since the pH electrode was apparently rapidly coated with Ag2S. Based on the kinetic study, the overall reaction proposed was  2Ag(N2H4CS) 3  2OH → Ag2S  5N2H4CS  NH2CN  2H2O (6.1)

While both ammonia and thiourea were present, it is probable that with sufficient thiourea present, the main complexant was thiourea (see the next paragraph). This would also explain the observed dependence of deposition rate on thiourea concentration: Initially the deposition increased (since thiourea is the source of S), reached a maximum, and then decreased (due to increasing complexation). Reference 85 presents the thermodynamic side of the previous paper. It is pointed out that although both ammonia and thiourea are present in the solution, because of the much higher stability constant of the Ag-thiourea complexes compared to the Ag-ammines, essentially all the Ag will be present as a thiourea complex. In this case, it can be assumed that the role of ammonia is only to control pH. An interesting observation in Ref. 86 was that the density of nuclei formed in the early stages of film deposition did not change with time. The film developed by growth of the initially relatively small (ca. 20 nm) nuclei. This suggests an ionby-ion type of growth rather than a cluster one. Reference 87 is a mechanistic study of Ag2S deposition from a thiourea bath (buffered to pH 9 with a borate buffer). There are some unusual properties of this deposition. One, the unusually strong effect of stirring on the deposition rate, has already been dealt with (Sec. 3.7) and, together with the measured activation energy of 20.4 kJ/mole, suggests a rate-determining diffusion step in the deposition. Another observation is that the films reach a maximum thickness (between 0.5 and 1.6 m) after 30 min of deposition (the thickness increases with temperature but the 30 min is, surprisingly, temperature independent) and then become thinner

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with further time in the bath, presumably because of loss of adhesion of the film due to increasing stresses in the film. The pH is fairly critical: 8.8 slows the growth down greatly, while 9.4 results in a fast homogeneous precipitation in solution (such precipitation occurs, even at the optimum pH of 9, in parallel with film formation). The properties of the foregoing films (only without the borate buffer) were described in Ref. 88. The XRD peaks of the as-deposited films were narrow and sharp, evidencing relatively large crystal size. The bandgap of the as-deposited film, measured from transmission spectra, was ca. 1.0 eV but varied somewhat with deposition temperature: At 40°C it was 0.91 eV and reached a maximum of 1.02 at 60°C and then decreased slightly at still higher temperatures. The absorption onset was sharper at higher temperatures, which was interpreted as being due to denser films. The refractive index was also slightly higher for higher deposition temperature, again explainable by the same rationale. The films, annealed in N2 at 250°C, were photoconducting, and the photoconductivity spectrum was similar to the absorption spectrum. Time-resolved microwave conductivity measurements were carried out on the films. Fast decay times and moderately good mobilities were found from these measurements. In particular, the mobility was very temperature dependent, and the highest value (5.3 cm2V1sec1) was obtained for annealed films that had been deposited at 60°C; both higher and lower deposition temperatures gave much lower mobilities. While not understood, this dependence emphasizes the need to optimize these films specifically for any particular application. A study of the photoconductivity of Ag2S doped with Hg2 or Au3 was described in Ref. 89. Illumination decreased the resistivity typically 2–3 times. Photoconductivity spectra showed best results for Hg doping; Au doping gave a higher peak sensitivity but a narrower spectrum, with lower sensitivity at shorter wavelengths compared even to undoped films [90]. Reference 91 involved an acid bath (although pH was not given) using thioacetamide as the source of S. The terminal film thickness was greater for lower-temperature deposition; films thicker than 3 m were obtained at 8°C. In contrast to the acid thiosulphate bath, the use of EDTA decreased the resistivity, as did deposition at lower temperatures. Photoelectrochemical activity was found for these films (see Chap. 9). Reference 92 describes not a normal CD process, but one closer to the SILAR technique described in Sec. 2.11.1. However, while the SILAR method involves dipping the substrate in a solution of one ion (e.g., sulphide), rinsing to remove all but (ideally) a monolayer of adsorbed ions and then dipping in a solution of the other ion (e.g., Ag), the present technique omits the intermediate rinsing step. This means that a relatively large amount of solution can remain on the substrate between dips, and layer formation proceeds much more rapidly than for SILAR, albeit with less control. A typical rate was 4 nm/dip cycle. In this case, a visible layer of Ag2S formed after several dips. Since interference colors were ob-

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tained, the films were smooth, because such colors are not seen on rough, highly scattering films. The film thickness/number of dips increased with increasing pH; however, the best films were obtained at pH
9.

6.12.2 Ag2Se Ag2Se films were first deposited from a bath using selenosemicarbazide as a Se source and thiourea to complex the Ag at 20°C [93]. The films were specular and had a resistivity of between 2 and 20 -cm. The only other true CD of Ag2Se describes films deposited on polyester substrates from an ammoniacal AgNO3 solution with selenosulphate at 0°C and a pH of 10–11 [94]. They were strongly (111) textured, with a crystal size of 9 nm. The optical bandgap (direct) was estimated to be 1.8 eV, compared to the normal value of ca. 1.3 eV. This was attributed to size quantization. The absorption spectra showed considerable absorption (scattering?) at longer wavelengths, which could be due to a lower, indirect bandgap if not to scattering. The resistivity of the films was ca. 2  103 -cm (200 -sq). A dip technique in which metallic Ag films were converted into Ag2Se was described [95]. The Ag film was made by successive dipping of glass substrates in a AgNO3 solution, followed by dipping in a solution of formaldehyde, and was converted to the sulphide by treatment with a solution of SeO2. The films were rough and apparently poorly adherent. The resistivity of the films was ca. 103 -cm.

6.13 THALLIUM The few cases reported for CD sulphides and selenides of Tl all reported the monosulphide (selenide)—TlS or TlSe. Tl can be monovalent or trivalent, and these apparently divalent compounds are believed to be mixed-valence compounds, with both Tl(I) and Tl(III) present.

6.13.1 TlS TlS was deposited from a solution of TlNO3, ammonia, and thiourea at room temperature (26°C) [96]. X-ray diffraction showed the formation of TlS. Optical spectroscopy (both transmission and diffuse reflection) allowed an approximate bandgap of 1.0 eV to be estimated. The films were p-type, with resistivity of ca. 2  103 -cm. Photoconductivity was measured (although not quantified) with a peak at ca. 1.2 eV (ca. 1 m) and extending from 1.0 eV to beyond 1.4 eV.

6.13.2 TlSe TlSe was first deposited from a solution of thallium(I) acetate and selenosulphate with added NaOH and hydrazine at room temperature [97]. The initial films were

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mirror-like but became thick (ca. 3–5 m) and matte black with increasing deposition. X-ray diffraction confirmed the film to be tetragonal TlSe. The main purpose for making these films was to study their photoconductivity. The films were p-type, with a resistivity of 500 -cm, which decreased by a factor of 2 after annealing (150°C in air). Indium doping (by adding In3 to the deposition bath) increased the resistivity (of the annealed film) by an order of magnitude, probably by introduction of compensating donors. The films, particularly the annealed, doped ones, were highly photoconducting, with a maximum photosensitivity (change in conductivity on illumination/dark conductivity) of 107 at a wavelength of 1.1 m. The response extended to ca. 1.5 m (0.83 eV) (low-energy side) and ca. 0.7 m (high-energy side). TlSe was also deposited from a solution of Tl2SO4 complexed with triethanolamine and ammonia and selenosulphate at 30°C [49]. Tetragonal TlSe was identified by XRD. The bandgap was estimated at 1.12 eV; however, the absorption spectrum appears to show two transitions—one (possibly indirect) at ⬇0.9 eV and another at ⬇1.3 eV. The films were p-type, with a resistivity of 105 -cm. Considering the high carrier concentration measured (almost 1020 cm3), this resistivity value appears unusually high.

6.14 TIN There are three sulphides of tin: SnS (grey, metallic; usually nonstoichiometric), Sn2S3 (black), and the layered, yellow SnS2. SnS and SnS2 are formed when hydrogen sulphide is passed into solutions of Sn(II) and Sn(IV), respectively. The analogous selenides also exist, although the existence of Sn2Se3 is apparently in some doubt. Tin forms soluble thio(seleno)anions. The sulphides tend to be soluble in very alkaline solutions.

6.14.1 Sn-S (See Table 6.7) Deposition of Sn-S (from a thiosulphate bath) was claimed as far back as 1870 [98]. This was based on deposition from a boiling solution of “zinnsalz” [probably tin chloride, but not clear whether Sn(II) or Sn(IV)] complexed with tartrate and using thiosulphate as the source of sulphur. Unfortunately, the substrate was brass; since brass will slowly convert to a dark-colored sulphide upon immersion in boiling thiosulphate solution and no characterization of the film was made other than its color (which would vary initially according to thickness, due to interference effects), there is no evidence that the films were really a sulphide of Sn. In fact, an attempt to reproduce these results concluded that there was no Sn in the layer [2]. On the other hand, SnS2 films were later obtained using a thiosulphate bath, although the solution composition was different (see later discussion of Ref. 99.)

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TABLE 6.7 Tin Sulphide Solution

Bandgap (eV)

Resistivity (-cm)

Composition

Ref.

S in RCOOH/SnCl2/90*

See text

100

2.35 1.51 indir. 1.3*

103–104 (n-) 2.5  1010 (n-) 2  104–107 (p-)

SnS (p-) Sn2S3; SnS 2* SnS2 SnS (anneal) SnS

S2O3/SnCl4/RT/pH 1–2 TA/SnCl2/AcH/TEA  NH3/RT TA/SnCl2·2H2O in acetone TEA  NH3/RT—75C TA/SnCl4/EDTA  NH3 hydrazine/RT/pH ca. 10

99 101 102

2.3 dir.

1.2 (n-)

SnS*2

103

Reference 100 described a different technique than usual, in that it used elemental S dissolved in a carboxylic acid. Propionic acid was most often used, although other carboxylic acids could also be employed. It was noted that best results were obtained when the SnCl2 was added as a powder to a freshly prepared S solution; aged solutions reacted much more slowly, if at all. It was surmised that this was due to changes in the nature of the dissolved Sn2 with time, such as tin oxide, or propionate formation and loss of HCl. The stoichiometry of the deposit was dependent on the water content of the solution and on the presence of a complexant. In anhydrous solution and without complexant, a uniform, brown film with the approximate composition Sn2S3 was formed. Addition of ca. 1% of water resulted in a uniform, slate-gray film of Sn1xS. This was explained by the increased ionization of Sn species in water and therefore increase in the concentration of Sn2. A substantial increase in the water concentration resulted in patchy films of an irreproducible nature, with both brown and gray regions forming. Complexation of the Sn reduced the reaction rate and allowed more time for the growing films to react with dissolved S, resulting in films of approximate composition SnS2 (yellow) together with some Sn2S3. In this case, some water was needed to form good films. An interesting and unusual characteristic of this deposition method is that if the deposition is allowed to proceed for a long time, the amount of precipitate formed in the reaction is decreased and the films grow thicker. This was explained through dissolution and reprecipitation of Sn-S solid phases via soluble thiostannate species. Dissolution was assumed to occur preferentially in the bulk precipitate due to its greater accessible surface area compared to that of the film. Another unique characteristic of this process is that a band of Sn-S above the level of the deposition solution was frequently observed. This above-solution film was ascribed to reaction of volatile SnCl4 and H2S. The presence of the for-

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mer could be explained by reactions of the type 2SnCl2  S → SnCl4  SnS

(6.2)

while the H2S (which could be detected separately) might be formed by SnCl2-reduction of S. X-ray diffraction showed the gray (low-water-concentration) films to have the pattern of SnS, while compositional analysis showed them to be Sn deficient by typically 10%. Crystals were large (micron-sized or larger). The other films tended to be mixtures of approximate compositions Sn2S3 and SnS2, with a needle-like morphology (typically 1 m long by 0.1 m wide). The SnS had an indirect bandgap of 1.0–1.3 eV and was p-type. It was more difficult to estimate the bandgaps of the other films due to their mixed nature. However, approximate bandgaps of 1.8 eV (Sn2S3) and 2.4 eV (SnS2) could be estimated from the optical spectra. The possible mechanisms of this unique deposition are not considered in Chapter 3 and therefore will be done so here. The reaction of metal salts with elemental S in nonaqueous solvents in which S dissolves is known, even if the mechanism is not clear. In the present case, two plausible mechanisms can be given. The Sn(II)/Sn(IV) redox potential is relatively negative (0.15 V vs. SHE in aqueous solution; much more negative in alkaline solution, although this is not relevant in the acidic conditions used here). Since only Sn(II) was added to the solution, any Sn(IV) present will be formed in the solution and is likely to occur in low concentration under most conditions. This means that, from the Nernst equation (see Chap. 1), the potential of the solution will be more negative than the standard potential, possibly by a large amount. Another, more negative potential that may be relevant is the Sn2/Sn0 (0.14 V; see later). The S/S2 standard potential is 0.45 V. Since only the oxidized form of this couple was added to the solution, the redox potential (which for our purposes means the potential where some appreciable concentration of S2 will be formed) will be considerably positive of this value. Add to this the known non-Nernstian behavior of the S/S2 couple when [S]  [S2], when the potential shifts strongly positive to a greater extent than expected from the Nernst equation [104], and it is feasible that the Sn(II) may reduce S to S2 in sufficient concentration to form Sn-S. By esentially the same reasoning, it could be argued that the Sn(II) might reduce itself (disproportionate) into elemental Sn (at low concentration). As just noted, the standard potential of Sn(II)/Sn(0) is 0.14 V. It might be argued that thermodynamically, this is more likely than the reduction of S (although the nonNernstian behavior of S/S2 will at least reduce this difference). Metals immersed in nonaqueous solutions of S can react to form a layer of the metal sulphide (the rate depending on the metal and on the temperature—e.g., Cu will readily sulphide at room temperature, while sulphidization of Zn will proceed slowly even at high temperature). The small Sn0 nucleii that may be formed in the disproportionation

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of Sn(II) would be chemically very active and more likely to react with the S in solution. In connection with this mechanism, it has been reported that the reaction between elemental S and Sn in alkylammonium compounds or amines under hydrothermal conditions gives various organic tin-suphide species and even, under some conditions (low pH), SnS2 [105]. A possible mechanism for this process was proposed based on nucleophilic attack of the basic amine or hydroxide on the S8 chain: 

S8  RNH2 → SMS哹SMNH2R

(6.3)

resulting in formation of polysulphide ions. This mechanism is unlikely in the acidic solutions used for the CD process. However, according to the principles just discussed based on the Nernst equation, elemental S in solution may be expected to contain a (very low) equilibrium concentration of (poly)sulphide. If the sulphide is removed by reaction (with Sn2, in this case), then even this very low concentration may be enough to sustain the formation of a Sn-S solid phase. Clearly, the lower the solubility product of the metal sulphide, the more likely this process is to occur. Considering that homogeneous precipitation of metal chalcogenides (mainly sulphides) by reaction between metal ions and dissolved chalcogen is well established, the main difference between this deposition and similar reactions seems to be that the products adhere to a substrate to give a visible film (in this case) rather than only precipitate. Whether this is connected with the redissolution/redeposition process that occurs with the Sn-S system or has some other explanation is important. If the former, it may be limited to only those systems that behave similarly. Otherwise it is not unreasonable to expect that other metal sulphides and selenides (possibly also tellurides, although tellurium tends to be much less soluble, if at all, in such solvents) may be deposited as films in this manner. In Ref. 99, yellow-gold films were obtained that gave no XRD pattern, but the chemical composition (as well as color) was consistent with SnS2. Increase in pH (more than 2) reduced the deposition rate, while increase in temperature led to precipitation in solution and therefore thinner films. In Ref. 101, the films in the deposition were deep brown. No XRD pattern was observed, but after annealing in an inert atmosphere at 410°C the pattern of SnS was obtained and the stoichiometry confirmed by elemental analysis. The bandgap (1.51 eV, indirect transition) was higher than the literature value (1.3 eV), and this was rationalized as resulting from the apparent amorphous structure. The room-temperature conductivity (4.109 S-cm1
a resistivity of 2.5  108 -cm) is low for a relatively low-bandgap material, suggesting either a very stoichiometric and intrinsic material or high grain-boundary resistance. The films were photoconductive (the sensitivity was not given), with a spectral range from 550 to 1050 nm and a peak at ca. 850 nm.

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In Ref. 102, the solution should be turbid; if too much triethanolamine was added and the solution was clear, no deposition occurred. Too much NH3 (pH not given) led to incorporation of hydroxide. The bandgap was given as ca. 1.3 eV; however, from the optical spectra, there is a variation in the bandgaps, depending on the deposition conditions and film thickness. In particular, the lower the temperature of deposition, the larger the bandgap (which analyses of the spectra show to be indirect and to vary from 1.45 to 1.2 eV). This behavior is typical of size quantization. The resistivity varied strongly with film thickness; the values shown in Table 6.7 are for 0.35 m (107 -cm) and 1.2 m (2  104 -cm). The films were mildly photoconductive (maximum sensitivity  10). These SnS films (in one study, propylene glycol was used instead of acetone to dissolve the SnCl2), coated with CD CuxS, were shown to possess spectral characteristics favorable for various solar control purposes [106,107]. Depending on the thicknesses of the two layers, films with varying absorption and reflection could be obtained that might be suitable for solar collectors or for window glazing. In Ref. 103, the pH of the bath was rather critical; films deposited at pH  10.5 were powdery and poorly adhering, while at pH  9.5, no deposition was observed. No structural or compositional characterization was given, but from the transmission spectrum it could be assumed that the film was SnS2 (and also highly scattering). Films designated as Sn(O,S) were deposited from a solution containing tin(IV) acetate, HCl, and thioacetamide (the concentration of the latter two components determining the S:O ratio) [108]. These films were prepared as buffer layers for photovoltaic cells (see Chap. 9), and little characterization of the films themselves, other than some XPS, was reported. The XPS results suggested that the films were a mixture of SnO2 and some Sn-S species. While not strictly CD, SnS has been deposited by an immersion technique whereby a glass substrate was immersed in a cold sulphide solution, followed, without rinsing, by immersion in a hot SnCl2 solution, and this cycle was repeated to increase the film thickness [109]. The film properties, in particular the electrical resistivity, were very dependent on the pH of the SnCl2 solution.

6.14.2 Sn-Se Only one example of Sn-Se has been reported [110]. Films were deposited from a room-temperature selenosulphate solution of SnCl2 complexed with triethanolamine and added NaOH. Polyvinylpyrollidone (PVP) was also added and in general slowed down the deposition. At an optimum concentration of PVP, a maximum terminal thickness was obtained (although no comparison with films deposited from PVP-free solutions was given). No XRD pattern was observed for the as-deposited films; heating in an inert atmophere at ca. 330°C gave the pattern of SnSe. The bandgap was 0.95 eV (indirect). The films were n-type, with a resistivity of ca. 10 -cm

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69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103.

R Bayón, C Mafftiotte, J Herrero. Thin Solid Films 353:100, 1999. R Bayón, J Herrero. Appl. Surf. Sci. 158:49, 2000. R Bayón, J Herrero. Thin Solid Films 387:111, 2001. T Yoshida, K Yamaguchi, H Toyoda, K Akao, T Sugiura, H Minoura. Proc. Electrochem. Soc. 97–20:37, 1997. P Pramanik, MA Akhter, PK Basu. Thin Solid Films 158:271, 1988. CD Lokhande, A Ennaoui, PS Patil, M Giersig, M Muller, K Diesner, H Tributsch. Thin Solid Films 330:70, 1998. CD Lokhande, KM Gadave. Turkish J. Phys. 18:83, 1994. P Pramanik, RN Bhattacharya. J. Mater. Sci. Lett. 8:781, 1989. P Pramanik, S Bhattacharya. Mater. Res. Bull. 25:15, 1990. KC Mandal, O Savadogo. J. Mater. Chem. 1:301, 1991. P Pramanik, S Biswas. J. Solid State Chem. 65:145, 1986. SS Dhumure, CD Lokhande. Thin Solid Films 240:1, 1994. AJ Varkey. Sol. Energy Mater. 21:291, 1991. SS Dhumure, CD Lokhande. Mater. Chem. Phys. 27:321, 1991. SS Dhumure, CD Lokhande. Sol. Energy Mater. Sol. Cells 28:159, 1992. GA Kitaev, TP Bol’shchikovsa. Izv. Akad. SSSR, Neorg. Mater. 2:65, 1966. GA Kitaev, TP Bol’shchikovsa, TA Ust’yantseva. Izv. Akad. SSR, Neorg. Mater. 3:1080, 1967. TP Bol’shchikovsa, GA Kitaev, VI Dvoinin, MV Degtyarev, LM Dvoskina. Izv. Akad. SSSR, Neorg. Mater. 16:387, 1980. H Meherzi-Maghraoui, P Cowache, D Lincot, M Dachraoui. J. Chim. Phys. 96:259, 1999. H Meherzi-Maghraoui, M Dachraoui, S Belgacem, KD Buhre, R Kunst, P Cowache, D Lincot. Thin Solid Films 288:217, 1996. MJ Mangalam, KN. Rao, N Rangarajan, CV Suryanarayana. Brit. J. Appl. Phys. (J. Phys. D), Ser. 22:1643, 1969. MJ Mangalam, KN. Rao, N Rangarajan, CV Suryanarayana. Ind. J. Pure and Appl. Phys. 7:628, 1969. SS Dhumure, CD Lokhande. Mater. Chem. Phys. 28:141, 1991. M Ristova, P Toshev. Thin Solid Films 216:274, 1992. AA Velykanov, EK Ostrovskaya, NP Garina, VA Turacova, AA Tchurkan. Ukr. Chim. Zh. 49:764, 1983. B Pejova, M Najdoski, I Grozdanov, SK Dey. Mater. Lett. 43:269, 2000. AB Kulkarni, MD Uplane, CD Lokhande. Thin Solid Films 260:14, 1995. A Mondal, P Pramanik. Thin Solid Films 110:65, 1983. MJ Mangalam, KN Rao, N Rangarajan, CV Suryanarayana. Jpn. J. Appl. Phys. 8:1258, 1969. C Puscher. Dingl. J. 195:375, 1870. CD Lokhande. J. Phys. D: Appl. Phys. 23:1703, 1990. RD Engelken, HE McCloud, C Lee, M Slayton, H Ghoreishi. J. Electrochem. Soc. 134:2696, 1987. P Pramanik, PK Basu, S Biswas. Thin Solid Films 150:269, 1987. MTS Nair, PK Nair. Semicond. Sci. Tech. 6:132, 1991. AJ Varkey. Int. J. Mater. Prod. Technol. 12:490, 1997.

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104. 105. 106. 107. 108.

PL Allen, A Hickling. Chem. Ind. 51:1558, 1954. T Jiang, GA Ozin, RL Bedard. Adv. Mater. 6:860, 1994. PK Nair, MTS Nair. J. Phys. D:Appl. Phys. 24:83, 1991. MTS Nair, PK Nair. J. Phys. D:Appl. Phys. 24:450, 1991. D Hariskos, R Heberholz, M Ruckh, U Ruhle, R Schäffler, HW Schock. In: 13th ECPV Solar Energy Conf., Nice, France, 1995, p 1995. 109. M Ristov, G Sinadinovski, I Grozdanov, M Mitreski. Thin Solid Films 173:53, 1989. 110. P Pramanik, RN Bhattacharya. J. Mater. Sci. Lett. 7:1305, 1988.

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7 Oxides and Other Semiconductors

Most of the compounds deposited by CD have been sulphides and selenides. Apart from a very few examples of tellurides (and some related telluride experiments) and with a very few exceptions, discussed at the end of this chapter, what is left is confined to oxides (including hydrated oxides and hydroxides and two examples of basic carbonates.) This chapter deals mainly with these oxides. In addition, as noted in Chapter 3, there are a number of slow precipitations that result in precipitates, rather than films, of various other compounds, not necessarily semiconductors in the conventional sense. These potential CD reactions, briefly discussed in Chapter 3, will be somewhat expanded on in this chapter. Oxide films are often deposited because of their electrical (resistance) and optical properties. A selection of such properties of CD oxides is given in Table 2.2. The reader is strongly urged to read Section 3.2.4 (precursors for oxide deposition) before reading this chapter or at least to refer back to it when necessary.

7.1 GENERAL CONSIDERATIONS FOR OXIDE/HYDRATED OXIDE/HYDROXIDE DEPOSITION The old analytical chemistry literature is rich with methods involving homogeneous precipitation from solution, the purpose being to obtain dense (therefore easily filterable), contamination-free precipitates for purposes of analyses. The

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urea method, in particular, has been extensively used in the past to form precipitates of oxides and basic salts for analytical purposes. Urea hydrolyzes to ammonia and (bi)carbonate, and the ammonia hydrolyzes further to give OH ions, with a subsequent increase in pH. This leads to the formation of hydroxides, hydrated oxides, carbonates, and basic salts. The formation of films has often been noted in precipitations using urea. Thus, Gordon [1] noted: “The precipitation of basic salts with urea is characterized by the formation of thin transparent films of precipitate which strongly adhere to glass surfaces.” Also, in Ref. 2 Gordon et al. wrote (p. 39): “Basic thorium formate adheres tenaciously to glass surfaces in the manner characteristic of the basic salts precipitated by the urea method.” Basic sulphates of Al [3] and Ga [4], which under suitable conditions contain very little sulphate (and are probably oxides or hydroxides), have been observed to form on glass using urea precipitation. This film formation was an undesirable side effect—for accurate analysis, the film needed to be removed and added to the precipitate. However, as pointed out by Gordon [1], “To those who have worked with the urea method, the existence of these films will always be a reminder of Willard’s* fond hope that he will someday find a way of making all the precipitate adhere to the beaker so that it will only be necessary to dry and weigh the beaker after discarding the solution.” Film formation is, in retrospect, not surprising, since the slow reaction characteristic of many homogeneous precipitation reactions is normally required (among other factors) for appreciable film formation to occur. Furthermore, the films tended to be very adherent. In their book Precipitation from Homogeneous Solutions [2], Gordon et al. write on the films formed by precipitation of basic tin sulphate: “The films cannot be removed by scraping with a policeman. However, by adding a few milliliters of hydrochloric acid . . . the films are easily dissolved.” (For those readers who, like the author, found this first sentence evoking amusing mental images, a policeman is (or was) a glass rod with a piece of rubber attached to the end, used to scrape deposits out of reaction vessels.) A notable exception to this “easy” dissolution was deposition of “basic stannic sulfate” using urea. This material formed such an adherent film that “it poses a difficult removal problem” [2,4a]. The films could be peeled off in relatively large, transparent sheets by warming for 30 min in a solution of (NH4)2SO4 and NaOH at pH of 9  0.5, a fact that suggests using CD for the preparation of thin, self-supporting films. (This also meant that the film particles adhered strongly not only to the substrate, but also to each other.) Actually, though not analytically well defined, this “basic sulfate” actually contained very little sulphate and was probably mostly SnO2 (see the beginning Sec. 7.2.14 on SnO2). It should be remembered that most of these products form white precipitates; therefore, if nonscattering, the films may not even be

* H. H. Willard was one of the pioneers of homogeneous precipitation for chemical analysis.

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visible. Thus film formation may have occurred in cases where it was not reported or even recognized. As a general point, it might be expected that the product of reaction between metal ions and hydroxide is a hydroxide (or basic salt) rather than an oxide. In many reported cases, oxides are formed directly. This is probably due to two factors. Many of the metal ions used (e.g., Pb, Sn, Tl, Ti, Zr, Si) do not readily, if at all, form simple hydroxides; most of these cations have a greater tendency to form what may be called oxide polymers, involving condensation to chains of OMMMOMMMO species at the reaction pH. Some hydroxides (e.g., Ag, Cu, Mn) are not very stable and are quite readily converted to the oxide, even in aqueous solutions. In some cases, simple hydroxides do form and need to be heated to dehydrate to the oxide. Very acidic (high valent) cations will readily hydrolyse in aqueous solution, often even at low pH. These cations tend to form the polymeric metal oxide chains mentioned previously. This hydrolysis can be controlled by addition of boric acid (see Sec. 3.2.4.4) and forms the basis of a technique referred to as liquid phase deposition. This method can be reasonably included in the more general term of chemical solution deposition, and is treated, although not comprehensively, in this book. Ref. 5 deals more thoroughly with this technique and describes many cases of SiO2 as well as some examples of several other oxides not covered in this chapter.

7.2 SPECIFIC OXIDES AND HYDROXIDES 7.2.1 Antimony Oxide (Sb2O3) Films of Sb2O3 (more strictly, Sb4O6) have been deposited from a room-temperature solution of potassium antimonyl tartrate and sodium selenosulphate [6]. The films showed a clear XRD pattern, and compositional analysis confirmed the composition. It is interesting that the selenide was not formed from this solution (it was formed if the Sb solution was mixed with triethanolamine and ammonia before adding the selenosulphate; see antimony selenide in Chap. 6). The most likely explanation for this is that the more alkaline solution, containing triethanolamine and ammonia, keeps the oxide, which might tend to form, in solution, both because Sb4O6 is soluble in sufficiently strong alkali and because of complexation by the triethanolamine and ammonia. Hydroxide would be present in much larger concentrations than selenide, even under mildly alkaline conditions. The resistivity of the films was of the order of 109 -cm.

7.2.2 Cadmium (Hydr)oxide (Cd(OH)2, CdO) Films of what was presumably Cd(OH)2 were deposited by heating to 80°C an alkaline cyanide solution of Cd2 containing H2O2 [7]. After heating the white as-

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deposited films at 250–300°C, they turned brown to give CdO with resistivities of typically a few k/sq. (CdO is normally a degenerate semiconductor with a low resistivity.) H2O2 was used in two other studies to deposit cadmium hydroxides or hydrated oxides. A mixed CdMOMOH film was deposited from a Cd-ammine solution at pH
10 and at various temperatures onto glass and quartz [8]. The deposited films exhibited clear XRD peaks that were identified with Cd(O2)0.88(OH)0.24 and were (111) textured. However, if KBr was added to the deposition solution, the preferred texture became (200) and the rate of film growth, as well as the terminal thickness, decreased. Annealing the films at ca. 200°C converted them into CdO with the same texture present as in the as-deposited films. From the optical spectra, a bandgap of ca. 2.6 eV was obtained (slightly higher for the Br-free bath and slightly lower for the Br-containing bath). The resistivity of the films varied from 2 to 20 k/sq ( 2  102 -cm). In this case, thicker films had a higher resistivity, and this was ascribed to cracking of the thicker films, observed in optical microscopy studies. The second study used a similar solution and obtained films, identified by XRD as CdO2, with a high resistivity (106–107 cm) that converted to low-resistivity (103 -cm) CdO on annealing in air at 450°C with a direct bandgap of 2.3 eV [9]. Deposition without H2O2 has also been described, using ammonia-complexed Cd. In one study [10,11], deposition was carried out at either room temperature or 50°C. The as-deposited Cd(OH)2 films on glass were annealed at 400°C in either air or an inert atmosphere to convert them to CdO. The as-deposited film was X-ray amorphous, while the annealed film was polycrystalline CdO. A bandgap of 2.2 eV (a little lower than the standard value of 2.4 eV) was obtained from the optical spectra of the CdO films, and the transmission in the nonabsorbing region was high (up to 90%). The resistivity of the as-deposited Cd(OH)2, as expected, was high (107 -cm), while that of the CdO was 103 -cm. In the other study, a higher ammonia concentration was used and deposition was carried out at room temperature [12]. It was noted that the uniformity of the films was better than when deposited at higher temperature. The Cd(OH)2 was heated in air at 150°C to convert it to CdO (the previous films required heating to ca. 275°C to convert the Cd(OH)2 to CdO [11]). (This is a not insignificant difference for what is assumed to be identical material.) The bandgap was 2.3 eV, and, as in the previous study, the films were very transparent to photons with less energy than the bandgap. The resistivity of the films was between 2  102 and 5  102 -cm.

7.2.3 Cobalt (Hydr)oxide and Hydroxy-Oxide (Co(O)OH, Co(OH)2, CoO, Co3O4) Co(O)OH was deposited from an ammonia-complexed solution of CoCl2 [13]. The Co(II) ammonia complex was allowed to oxidize for two days in air to the

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more stable trivalent cobaltic (III)–ammine complex. Heating the cobaltic–ammine complex to 65–70°C resulted in deposition of ca. 0.1 m of an adherent, brown CoO(OH) deposit on a glass substrate after 4 hr. Spectral measurements allowed an estimation for the bandgap of this film of ca. 2.4 eV. The Co(O)OH was oxidized to an adherent Co3O4 film by heating over 300°C in air. Films of CoO were deposited from a somewhat similar bath (but not left to oxidize) after annealing in O2 at 350°C; Co3O4 started to appear only at 500°C [13a]. Hydrolysis of urea to increase the pH by formation of ammonia was used to deposit Co(OH)2 by heating a solution of Co2 with urea at 100°C [14]. The pink film was shown by XRD to be a mixture of phases of Co(OH)2. Heating the films at 350–400°C converted the hydroxide to Co3O4.

7.2.4 Copper Oxide (Cu2O) Cu2O films were deposited by treating a thiosulphate-complexed solution of Cu(NO3)2 with NaOH [15]. This was based on an early study where a glass substrate was alternately and repeatedly dipped in NaOH and then Cu-thiosulphate solutions [16]. The thiosulphate (S2O2 3 ) both reduced the Cu(II) to Cu(I) and acted as a complexing agent. The films were deposited at 60–70°C, resulting in a thickness of ca. 0.3 m in one hour. The substrates—glass slides or polyester film—were precoated with very thin CuxS films by immersion in copper thiosulphate solution at 40°C. While the role of this prelayer was not clear, it was implied that in its absence, the Cu2O films were not uniform. While the mechanism for the deposition was not discussed, the instability of the copper hydroxides (the hydroxide of Cu(I) probably does not even exist) toward dehydration, together with the reducing action of the thiosulphate, leads to the expectation that Cu2O will be the product of the hydrolysis of Cu(I) in alkaline solution. It should be noted, however, that the Cu-thiosulphate solution itself is not very stable and apparently forms predominantly CuxS in the absence of NaOH. X-ray diffraction showed the film to be Cu2O, with no detectable amount of CuO and with a crystal size, estimated from the peak widths, of
20 nm. Optical transmission measurements of the films gave a value of (indirect) bandgap of 2.28 eV (literature room-temperature bandgap
2.1 eV but is rather variable). The electrical resistivity of the films (between 0.1 and 0.25 m thick), measured through Au contacts, was ca. 2.5 k/sq (ca. 5  102 -cm). This value increased with air-annealing (250°C, 20 min) up to 60 k/sq. The relatively low resistivity was attributed to incorporation of S, either from the CuxS prelayer or by 2 hydrolytic decomposition of S2O2 3 to S . Treatment of the films with Na2S solution decreased the resistivity by nearly two orders of magnitude, and S was found in the films. It is likely that the surface of the Cu2O crystals was partially

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converted to CuxS, and surface conduction via this surface layer was responsible for the enhanced conductivity.

7.2.5 Indium Oxide (In2O3) In2O3 was deposited from a solution of InCl3 that was slowly hydrolyzed to form the hydroxide [17]. The rate of hydrolysis was slowed down sufficiently to prevent rapid bulk precipitation by a combination of a freezing agent (sodium citrate) and a relatively low pH (7.5). Ag (as AgNO3) was added, supposedly as a catalyst (although it is not clear what needed to be catalyzed) and to improve adherence. The citrate will also act as a complexant for the In3, which may be an important factor. Addition of SnCl4 to the deposition solution allowed doping with Sn [tin-doped indium oxide or indium tin oxide (ITO)]. The film grew to a terminal thickness of ca. 360 nm in 30 min. At higher solution pH, the terminal thickness decreased. Heating at 200°C for 2 hr in a vacuum resulted in conversion of the hydroxide(s) to crystalline oxide, with an average grain size of ca. 25 nm (In2O3) and ca. 54 nm (ITO). Figure 7.1 shows the optical transmission and reflectance spectra of the two films. The main difference is an increase in the mid-IR reflectance of the doped film compared with the undoped one, due to the high free electron concentration

FIG. 7.1 Transmittance and reflectance spectra of In2O3 films. Broken lines: undoped In2O3; solid lines: ITO (10% Sn). (Adapted from Ref. 17.)

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in the conduction band. The bandgaps measured from the spectra were ca. 3.5 eV (undoped) and ca. 3.6 eV (ITO). Electrical measurements of the In2O3 (ITO) films gave temperature-independent resistivities ( ) of 2  102 (103) -cm, carrier concentrations (Nd) of 2  1020 (1021) cm3, and mobilities () of 3 (17) cm2V1sec1. The temperature independence of the resistivity indicated that the films, even the nominally undoped ones, were degenerate semiconductors. In2O3 has been deposited on Sn/Pd-activated glass by first depositing a film of In(OH)3 and then heating in air at a temperature of 200°C or more [18]. The In(OH)3 was deposited using a solution of dimethylamineborane and indium nitrate maintained at 60°C. The deposition rate was dependent on the borane concentration up to a limiting concentration of 0.03 M, and the film thickness was proportional to deposition time, with final thicknesses of ca. 1 m. X-ray diffraction showed mainly one sharp peak corresponding to the (210) plane of In(OH)3, which, after annealing, converted to polycrystalline In2O3. Optical transmission spectroscopy of the In2O3 film showed a high transmission at 800 nm, gradually decreasing with decreasing wavelength, characteristic of a somewhat scattering film. The bandgap was estimated from the spectrum to be 3.6 eV. Electrical conductivity measurements of the as-deposited In(OH)3 showed an expectedly high resistivity of ca. 109 -cm. That of the annealed oxide film decreased to 33 -cm (carrier concentration
1.85  1016 cm3; mobility
10 cm2V1sec1). The resistivity is high compared to many other In2O3 films (which are often used as transparent conductors), mainly due to the low carrier concentration, implying a high degree of stoichiometry. Finally, although no attempt was made to convert the film to oxide, In(OH)3, for use as a buffer layer on PV cells (see Chap. 9), was deposited from a thiourea-based solution of InCl3 at a pH of 3.3 [19]. Apparently no sulphide was formed, possibly due to the relatively high (for In) pH, which favored hydroxide formation.

7.2.6 Iron Oxides and Hydroxy-Oxide The instability of Fe(III) compounds toward hydrolysis has been exploited to form Fe(O)OH films [20]. The substrate in this study was a sulphonated-vinyl terminated self-assembled monolayer (SAM). Deposition was accomplished by heating Fe(NO3)3 solutions. The pH of the solution was rather critical; a pH of 2.0 or slightly higher was necessary. At lower values of pH, hydrolysis did not occur; at appreciably higher values, rapid hydrolysis occurred, resulting in precipitation rather than film deposition [e.g., at a pH of ca. 3, only very thin films (ca. 5 nm thick) of colloidal Fe(O)OH particles formed]. The films were columnar, with a column diameter of ca. 20 nm, and the columns were composed of lamellae ca. 2

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nm in size. Nucleation of the (noncolloidal) film occurred by binding of Fe species to the sulphonate endgroups (one Fe to two sulphonate groups, measured by XPS) [21,22] -Fe2O3 was deposited on Si (111) or Si (100) by using hydrolysis of urea at high temperatures [23]. An aqueous solution of Fe(NO3)3 and urea (pH between 5 and 6), together with the Si substrate was heated to between 100 and 200°C (presumably in an autoclave) for 4–24 hr. X-ray diffraction showed the formation of -Fe2O3 with some (101) texture. From SEM measurements, the films were ca. 100 nm in thickness, with a morphology depending on the crystal face of the Si. For (111) Si, the grains were spherical and 10 nm in size; for (100) Si, columnar grains, 30  5 nm, were obtained. If reduced Fe powder was added to the preceding solution at an optimum pH of 6–7, magnetite (Fe3O4) was deposited onto Si (100) or -Al2O3 at ca. 140°C over several hours [24]. No other phase was found in the XRD spectrum. It was suggested that the Fe3O4 formed by reaction between Fe(OH)3 (presumably formed by hydrolysis of the ferric nitrate) and Fe(OH)2 formed by hydrothermal oxidation of the Fe powder. Particle sizes of 150 nm (on Al2O3) and 50 nm (on Si) were measured by SEM. By adding Co(Cl)2 to the deposition solution and heating the resulting films in air (2 hr at 400°C) conversion of the magnetite to -Fe2O3, doped with Co, occurred [25]. The grains were needle-shaped (50  10 nm). The films exhibited good magnetic properties. A variation of the foregoing urea method was used to deposit -Fe2O3 on SnCl2-sensitized glass substrates [26]. A solution of FeCl2 and urea at pH
3 was heated at 90°C for 2 hr. The as-deposited film was probably FeO(OH) (the hydroxide group was seen in FTIR studies). On annealing at 350–400°C (presumably in air), -Fe2O3 was formed, with a crystal size of 22 nm. Optical spectroscopy of the as-deposited film showed a direct bandgap of 3.2 eV and a weak (possibly indirect) absorption starting at ca. 2.2 eV. The bandgap of the oxidized films was 2.0 eV. Resistivity was ca. 2 -cm, which dropped by a factor of up to four when exposed to high humidity, suggesting possible use as a sensor for water vapor. Magnetite was also deposited on glass by the dimethylamine borane technique described for In2O3, using a solution of Fe(NO3)3 and dimethylamine borane with a pH
3.5 at 20°C [27]. At higher deposition temperatures, Fe(O)OH and Fe2O3 were apparently also formed, and the magnetite content decreased, until, at 60°C, no magnetite was observed and the films were yellow (the magnetite films were black). The formation of pure Fe(III) oxide (hydroxide) at higher temperatures, compared to the mixed Fe(II)/Fe(III) magnetite at lower temperatures, is likely due to more facile oxidation of Fe(II) to Fe(III) at higher temperature. The borane is a reducing agent, and this is presumably the reason that the mixed va-

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lency magnetite, which can be envisaged as Fe(III)2O3Fe(II)O, is formed at lower temperature. The resistivity of the magnetite was ca. 2 k-cm—much higher than usual for magnetite. This was explained by a lower concentration of Fe(III) than expected from stoichiometry, which was obtained in these films. Magnetic properties of the films were described.

7.2.7 Lead Oxide (PbO2) A solution of Pb2 ions can be oxidized to PbO2 by persulphate [28]: 2  Pb2  S2O2 8  2H2O → PbO2  2SO 4  4H

(7.1)

This reaction normally resulted in a precipitate of PbO2. However, this process was subsequently modified to give films of PbO2 without precipitation in solution, apparently based on the observation of the authors that films of PbO2 were sometimes observed to form on the walls of the glass beaker using this reaction [29]. The film deposition was carried out at room temperature from an aqueous solution of plumbous acetate, ammonium acetate, and ammonium persulphate, using NH4OH to bring the pH to 6. A trace of AgNO3 was added as a catalyst for reaction 7.4 [29]. A film of PbO2 ca. 50 nm thick was formed in an hour. Once this initial film was deposited, thicker films could be built up, usually at a somewhat higher pH, in the absence of the AgNO3. The initial film formation appears to be a pure CD reaction. However, electrochemical studies of further film buildup showed that an electroless deposition mechanism, involving two partial electrochemical reactions, was responsible for film formation. A few words on the difference between CD and electroless deposition are in order here. Electroless deposition is related to electrodeposition, except that instead of charge being supplied by an external power supply, it is supplied internally by oxidation (reduction) of a strong reducing (oxidizing) agent; the two partial electrochemical reactions occur at different sites on the substrate (the substrate is initially a sensitized solid and subsequently the deposited material itself). This implies a reasonable electrical conductivity of the material to be deposited. For this reason, electroless deposition is used mainly for metals, but can be used for electrically conducting compounds, of which PbO2 is an example. Since charge transfer is involved, a change in valence state of the metal cation normally occurs between solution and film. Mindt [29] described some properties of these films (thicker electroless films, not the initial purely CD ones). Electron diffraction showed that the film was -PbO2. The crystal (more correctly the particle) size was found, by electron microscopy, to be ca. 200 nm. The carrier density, measured by the Hall effect, was ca. 1021 cm3. The resistivity was somewhat dependent on the pH of deposi-

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tion, varying from 2  103 (pH 7) to 3  102 -cm (pH 10), although the carrier density was not found to vary appreciably with pH (implying that the mobility did vary with pH). In a similar study, the deposition conditions were modified [pH of 8 (by ammonia), no Ag catalyst, and a deposition temperature of 80°C] [30]. A primary thin film was deposited, followed by a second deposition, resulting in films several microns thick. Optical absorption spectroscopy gave a bandgap of 1.7 eV. The film resistivity was 1.3  103 -cm (carrier density
8  1019 cm3; mobility
50 cm2V1sec1). White films of 6PbCO33Pb(OH)2PbO (from XRD analysis) were slowly formed over a few days from alkaline-complexed Pb2 solutions that contained a colloidal hydrated oxide phase and that were exposed to air [31]. This was due to reaction with CO2 in the air (see Sec. 5.3.3).

7.2.8 Manganese Oxide (Mn2O3, MnO2) Aqueous solutions of permanganate will slowly oxidize, forming a brown film [not clear if this is Mn2O3, MnO2, or Mn(O)OH] on the walls of the vessel in which it is stored. Increase of either acidity or alkalinity of the solution can accelerate this decomposition reaction. This reaction has been used to treat polymer substrates prior to chemical deposition, for improved adhesion [32]. In that work, the MnMO film was dissolved before the CD process, and the improved adhesion was probably due to some increase in the hydrophilic character of the mostly hydrophobic polymers. Mn2O3 films have been deposited on glass from an ammoniacal solution of Mn2 [33]. Ammonium chloride was added to decrease the pH and slow down the rate of hydrolysis. The initial product was believed to be Mn(OH)2, which oxidized in air to Mn2O3. The persulphate technique used for PbO2 described earlier was extended to MnO2 deposited on glass, using manganous acetate in place of lead acetate, together with ammonium persulphate and AgNO3 as a catalyst [34]. Adherent films up to 0.5 m could be obtained. No XRD pattern was found for the films, implying that the deposit was amorphous or made up of very tiny nanocrystals (nondefected crystal size of 2 nm or more can usually be detected by careful XRD). The resistivity of the films was very dependent on the solution pH, with values of 4  102 -cm (pH 8) and 2  104 -cm (pH 6.3).

7.2.9 Molybdenum Oxide An early attempt to deposit Mo—S on various metal substrates using ammonium molybdate and thiosulphate resulted in films that were found to be sulphur free and believed to be an oxide, although this was not investigated further [35].

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7.2.10 Nickel Oxide (NiO) A brief communication on the deposition of NiO using persulphate was described [36]. An ammoniacal solution containing NiSO4 and potassium persulphate was used to deposit black NiO on glass at room temperature. It was suggested, based on the importance of the NH4OH (NaOH or KOH did not give NiO; other alkaline reducing species, such as amines, did) that a mixture of NiO and Ni2O3 formed and that the higher-valent Ni2O3 was reduced by the ammonia The NiO was confirmed by XRD. Optical absorption spectroscopy was used to estimate a direct bandgap of 1.75 eV (see later). The films were p-type (thermoelectric probe), with a resistivity of 105 -cm. In another study, an ammoniacal solution of Ni2 was heated at 60–80°C to deposit a green-gray film of what was reported to be NiO [37]. This was assumed to form via the hydroxide, although no structural or compositional characterization of this deposit (or of the annealed film) was given. The deposition was carried out in a beaker, and no deposition occurred at room temperature; this suggests that deposition occurred by loss of ammonia. Heating this film in air at 280°C formed Ni(O)OH. From the optical transmission and reflection spectra of the NiO and Ni(O)OH, it appears that their absorption spectra were very similar, with a weak absorption (possibly also scattering) in the visible and strong absorption in the near-UV region. The resistivity of the Ni(O)OH was 800 -cm, while that of the as-deposited NiO was apparently too high to measure. Ni(OH)2 was deposited from a solution of urea and Ni2 ions at an initial pH of 6 and a temperature of 100°C. The hydroxide was then annealed in air at 350–400°C to convert it to NiO [38]. Films almost 1 m thick were obtained after ca. 2 hr, with an average crystal size (from XRD) of 13 nm. Optical absorption spectroscopy of the annealed films gave a direct bandgap of 3.6 eV, somewhat lower than the rather variable literature values of 3.7–4.0 eV. It should be mentioned that it is not as simple to correlate the band structure of NiO with its optical and electrical properties as it is for most of the other semiconductors dealt with here. This feature, common to many transition metal compounds, is a consequence of electrons in (often-narrow) d-bands that are relatively localized by electron–electron repulsions. Thus, although the 3d band in NiO (which in the pure state is green, like the hydrated Ni2 ion) is only partially filled, pure NiO is insulating due to the localized 3d electrons. The conductivity (and black color) of NiO as it is normally obtained is due to nonstoichiometry leading to doping. The room-temperature resistivity of the annealed NiO in this case was several M/sq. (several hundred -cm).

7.2.11 Silicon Oxide (SiO2) SiO2 films were deposited on soda lime glass from a silica gel–saturated solution of hydrofluorosilicic acid (H2SiF6) and boric acid [39]. The boric acid reacts with

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the H2SiF6 to form SiO2 by removal of HF formed in the H2SiF6 hydrolysis equilibrium: H2SiF6  2H2O D 6HF  SiO2  H3BO3  4HF D BF 4  H3O  2H2O

(7.2) (7.3)

The SiO2 formed gradually deposited on the substrate. The deposition rate was typically 10–20 nm/ hr, depending on the boric acid concentration and solution temperature. Film thickness was ca. 100 nm. Infrared measurements showed that the films had a higher concentration of SiMOMSi bonds than some other silica films made by different techniques. This was interpreted to mean that the silica network of the films was more orderly, a property that was evidenced by greater stability of the films against chemical etchning and good blocking properties to sodium diffusion from the soda glass, compared to many other silica films. This is the first of many SiO2 depositions using essentially the same technique. The others are given in Ref. 5.

7.2.12 Silver Oxide (AgO and Ag2O) Oxides of Ag have been deposited by deposition from a triethanolamine-complexed Ag solution at a pH  11.5 [40]. At room temperature, a black deposit formed over some hours that converted into a brown film when air-annealed at 150°C. Based presumably on the color of these films, the as-deposited film was assumed to be AgO, which turned into Ag2O on annealing (no structural characterization of the films was reported). From the optical spectra, a bandgap of 2.25 eV was estimated for the annealed films ([literature (direct) bandgap of Ag2O
1.2 eV; a pure semiconductor with a direct bandgap of 2.25 eV should be orange]. Both films were insulating (no values given) but became much more conducting when high voltages (1 kV) were applied to two laterally spaced Ag electrodes on the (ca. 500-nm-thick) films. Using a similar solution, films of either Ag or AgO were deposited on both glass and polyester film [41]. Addition of triethanolamine to a Ag solution caused initial precipitation (silver oxide or hydroxide), which redissolved in excess triethanolamine. Deposition from a solution where some precipitate remained resulted in AgO (possibly with some Ag2O), while a solution where this precipitate was completely redissolved gave metallic Ag. The reducing action of the free triethanolamine present in the latter case may be the cause of the formation of metallic Ag.

7.2.13 Thallium Oxide (Tl2O3) The same persulphate technique described earlier for PbO2 and MnO2 was also used to deposit Tl2O3, with thallous acetate in place of the other metal acetates

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[34]. Adherent films could be deposited up to a thickness of 10 m—very thick films for the CD method. X-ray diffraction confirmed that the films were cubic Tl2O3. The film resistivity was 5  104 -cm. The modifications employed by Bhattacharya and Pramanik for the PbO2 deposition (higher pH and temperature, no Ag) were also used for Tl2O3 [30]. Film resistivity was similar to that of the previous study (3.7  104 -cm), with a carrier concentration of 4  1020 cm3 and a mobility of ca. 50 cm2V1sec1.

7.2.14 Tin Oxide (SnO2) Very strongly adhering films of “basic stannic sulfate,” which was probably SnO2 (the Sn:SO 2 4 ratio was 44:1) were reported using an aqueous solution of SnCl4, urea, H2SO4, (NH4)2SO4, and HCl at pH 0.5 [4a]. The films adhered to the walls of the glass deposition vessel so strongly that HF was one of the few reagents available to remove them. Films which are mainly (hydr)oxides of Al, Ga and Th have also been reported in the early literature using the urea method [2]. Tin salts, in particular those of Sn(IV), are readily hydrolyzed, and the stable product is SnO2. This has been exploited in a number of studies, most with minor differences between them, to deposit SnO2 films. SnO2 films were deposited using SnCl4 and NH4F (the latter apparently as a complexing agent to slow down hydrolysis of the Sn4 by the alkaline solution) [33]. The resistivity of the as-deposited film was 200 -cm (102 -cm for In doped). A similar method was described a few years later in which the NH4F was used as a freezing agent to slow down the rate of hydrolysis [42]. AgNO3 was added as a catalyst(?) and to improve the film adherence, although it is not clear why a catalyst was needed or even desirable, since the objective was to slow down the reaction. NaOH was used instead of ammonia to adjust solution pH to between 7.5 and 8.5. The growth rate was linear and decreased with decreasing pH, with a limiting thickness that increased with decreasing pH. The growth rates and terminal thicknesses were similar to those for ZnO deposited by the same technique (see Fig. 7.2 in Sec. 7.2.17), only the rates were two to four times slower. These relationships were explained in the same way as for ZnO. The films were found to be SnO2 with the rutile structure (by XRD), with a grain size of 20–30 nm (by TEM). Optical transmission and reflectance spectroscopy showed that the films were close to 80% transmitting up to ca. 1 m and highly reflecting in the mid-IR. These spectra were similar to those of ZnO shown in Figure 7.3, (Sec. 7.2.17), except for the absorption in the UV of the undoped ZnO due to the lower bandgap. The bandgap (direct) was 3.56 eV. Electrical measurements gave a resistivity ( ) of 0.1 -cm, carrier concentration (Nd) of 1019

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cm3, and mobility () of ca. 6 cm2V1sec1. Annealing in vacuum decreased to a minimum of 2  103 -cm at 375°C. Antimony-doped SnO2 films were deposited by adding SbCl3 to the deposition solution. Sb is a well-known n-type dopant used to increase the conductivity of SnO2 films. The Sb concentration in the films increased linearly with that in the deposition solution and was somewhat less than the solution concentration (e.g., 6% Sb in solution gave ca. 4% in the film). The Sb doping increased both the visible/near-IR transmission and mid-IR reflectance of the films, compared to the undoped films. These spectra are similar to those for doped ZnO (Fig. 7.3), and the effect of doping can be explained in the same way. The bandgap increased to 4.1 eV, compared to 3.56 eV for the undoped film, explained through band filling by free electrons. The conductivity, Nd, and  all increased with increasing Sb concentration up to 5% and then decreased again with increasing Sb (Nd leveled off). This was due to the obvious increase in Nd with increased doping, the measured increase in grain size with doping (30–65 nm), resulting in increased  followed by segregation of dopant at the grain boundaries at greater Sb concentrations, which again decreased . The resistivity of the 5% Sb films decreased from 103 -cm as deposited to almost as low as 104 -cm on annealing at 375°C. In another preparation [43], ammonia gas was passed through a solution of SnCl4, the precipitate rinsed well (to remove Cl, which caused the final films to be porous—a useful observation since porous films are sometimes preferred over compact ones), and redispersed in concentrated HNO3 to give a semitransparent sol at a pH of 5–7. Films were deposited from this solution onto Si (100) by heating at 60–100°C for 4 hr, followed by 100–200°C for 6–12 hr. The initial lower-temperature step was necessary to obtain nucleation; if omitted, no film was deposited. Films ca. 200 nm thick were obtained with a crystal size of 3.5–4 nm. As with the original deposition of “basic stannic sulphate,” urea was used in another deposition to slowly increase the pH of the initially strongly acidic (by HCl) solution, thereby hydrolyzing the SnCl4 [44]. Besides slowing down the hydrolysis of the SnCl2 due to increased acidity, the HCl also complexed the tin as hexachlorostannate, (SnCl6)2, further slowing down the hydrolysis. From the optical spectrum, a bandgap of 4.0  0.1 eV was estimated. These films were used as buffer layers on CuInSe2 solar cells (see Chap. 9). SnO2 was deposited on hydrolyzed Si and on Si coated with sulphonateterminated self-assembled monolayers from a solution of SnCl4 in dilute HCl at 80° [45]. The films, up to 65 nm thick and consisting of a dense-packed aggregate of SnO2 nanocrystals (5–10 nm) together with some amorphous basic tin oxide, contained ca. 3 at.% Cl. They were adherent on all substrates, although the adherence and homogeneity on Si was less reproducible than on the monolayer-coated Si. Films were also deposited using a continuous flow system. The films were sim-

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ilar to those using a static solution, but the deposition rate was considerably faster (20 nm hr1 compared to an average of several nm hr1 for the static solution), and thicker films (160 nm) could be obtained. Films of SnO2 up to 0.5 m thick were grown by hydrolysis of aqueous SnF2 (a change from the usual SnCl4) solutions, optimally at ca. 60°C on various glasses and on Si [46]. The as-deposited films contained 6–16 mole% F, an n-type dopant in SnO2. Films annealed at 300°C in air contained almost no F and were crystalline SnO2 (cassiterite). It was suggested that the as-deposited material possessed a tin–oxygen polymer structure, with tin–fluorine bonds substituting for some of the tin–oxygen bonds. Conductivities of the order of 102 -cm were measured for films annealed at 500°C. Deposition of SnO2 onto sulphonated polystyrene at 40°C has been briefly described in a review by Bunker et al., although without experimental details [47]. The point was made that the deposition of SnO2 involves hydrolysis and condensation reactions involving poorly characterized species as opposed to precipitation of an ionic salt. Also, the deposition is very dependent on pH—a change in pH of one unit can change the solubility of the products by four orders of magnitude. In spite of these factors, which imply difficulty in depositing films in the absence of bulk precipitation in solution, they apparently succeeded in achieving preferential nucleation of SnO2 on the sulphonated polystyrene. The films were dense and composed of cassiterite, with a grain size of 4 nm.

7.2.15 Titanium Oxide (TiO2) Sulphonate-terminated self-assembling monolayers on Si were used as substrates for TiO2 deposition [48]. The deposition solution consisted of TiCl4 in 6M HCl at 80°C (considerably more dilute HCl solutions resulted in immediate bulk precipitation, while much stronger solutions were stable against hydrolysis and therefore no deposition occurred.) The films, ca. 50 nm thick, comprised small (2–4 nm) nanocrystals of anatase TiO2, possibly in an amorphous matrix, and were uniform, adherent, and pore free. In contrast, only a small amount of irregular deposit was formed on bare Si. The role of the sulphonate endgroups was believed to promote nucleation of the nanocrystals and/or facilitate attachment of TiO2 clusters in solution to the substrate. Hydrolysis of TiCl4 proceeds through various titanium hydroxy and chloro-hydroxy complex cations. The anionic sulphonate groups could thus promote attachment and nucleation of these cationic complexes. Annealing in air increased the crystal size (up to 25 nm at 600°C) without damage to the film while retaining the anatase structure (the rutile structure was barely noticeable after 2 hr at 600°C). Using this deposition technique, TiO2 was deposited onto patterned selfassembled monolayers [49]. Thioacetate-terminated trichlorosilane monolayers

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were self-assembled onto oxidized Si substrates. Illumination through a grid-patterned mask with grid openings of ca. 10 m resulted in photolysis of the exposed, somewhat hydrophobic thioacetate groups to hydrophilic sulphonate groups. TiO2 was then deposited from the TiCl4 solution onto the irradiated hydrophilic regions of the substrate (see Figs. 2.5 and 2.6). The resistivity of this TiO 2 was ca. 109 -cm. TiO2 was deposited on (100) Si from a solution of TiO 2 4 prepared by dissolving Ti metal in an aqueous solution of ammonia and H2O2 [50]. An initial lowtemperature–final high-temperature regime similar to that described earlier for SnO2 by the same group was used at a pH between 6 and 7. As before, the lowtemperature stage was necessary for film formation. A film formed in the lowtemperature stage but was X-ray amorphous. After the high-temperature stage, pure anatase-phase TiO2 was obtained in the form of square platelets ca. 10  10 nm in size. The film was highly (112) textured.

7.2.16 Vanadium Oxide Vanadium oxide films were deposited by dissolving V2O5 in aqueous HF and immersing an Al plate in the solution (the Al acts as a scavenger for F in place of the more commonly-used boric acid) [50a]. The brown film was X-ray amorphous but crystallized on heating in air through a mixed V(IV)- V(V) oxide to V2O5 (in an inert atmosphere, VO2 was formed on annealing). The as-deposited material was believed to contain mainly V(IV).

7.2.17 Yttrium Oxide (Y2O3) Basic yttrium carbonate [Y(OH)CO3] was deposited by CD and subsequently annealed in air at 600°C to Y2O3 [51]. Si wafers and self-assembled monolayers with sulphonate endgroups were used as substrates. An aqueous solution of YNO3 and urea was heated at 80°C in sealed vials. The increase in pH, together with generation of carbonate from hydrolysis of urea (Sec. 3.2.4.1), resulted in formation of the basic carbonate. The film thickness was 35 nm. Various analytical techniques were used to confirm that the deposit was amorphous Y(OH)CO3. Annealing at 600°C was necessary to convert the film to crystalline Y2O3 (amorphous oxide was formed between 300 and 400°C), with a film thickness of 25 nm and crystal size of ca. 20 nm.

7.2.18 Zinc Oxide (ZnO) ZnO is the most studied of all the oxides deposited by CD. This is largely due to its use as a transparent, electrically conducting layer. The first description of CD ZnO arose from the observation that deposition of mixed (Cd,Zn)S resulted in large amounts of ZnO, and this led to development

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of a technique for depositing pure ZnO [7]. The method is similar to that used for CdO—heating an alkaline cyanide solution of a zinc salt to 80–90°C (although, in contrast to the Cd case, no H2O2 was added). The films, a few hundred nanometers in thickness, were shown by XRD to be ZnO. Sheet resistivities were ca. 108 /sq, which dropped three orders of magnitude after heating at 350°C for 10 min in forming gas. ZnO was deposited by CD using the freezing technique described previously for In2O3 [52]. The substrates (glass or quartz) were immersed in an aqueous solution containing ZnCl2, NH4F (as freezing agent—probably also acting as a complexing agent for the Zn2), and Ag catalyst (as for In-O and SnO2, it is not clear for what purpose), and NaOH was then added to a pH of between 7.5 and 8.5. The resulting Zn(OH)2 films were then annealed for a few hours in air or vacuum at 180–200°C. The rate of deposition and terminal thickness of the ZnO films are shown in Figure 7.2. The films grow faster at higher pH but with a lower terminal thickness than at lower pH. At a pH substantially lower than 7.5, no growth occurs; the OH concentration is too low to precipitate Zn(OH)2. As the pH is increased, the rate of formation of Zn(OH)2 increases. This results in increasing homogeneous precipitation in the solution, leading to loss of reactant and thinner films, until, at a pH substantially greater than 8.5, essentially all the Zn is precipitated and no (or, more likely, only an ultrathin) film formation occurs. Al-doped ZnO films were also deposited by adding AlCl3 to the deposition solution. The amount of Al in the films (given as at.% with respect to the Zn concentration) was somewhat smaller than that in the deposition solution but was proportional to the concentration in solution (up to the maximum measured concentration in the films of 5.5%). The films (both ZnO and ZnO:Al) were wurtzite structure with a preferential texturing (c-axis ⊥ substrate). No Al2O3 was found in the XRD spectra, suggesting either dispersal of the Al in the ZnO matrix or its presence as very tiny crystals of (hydr)oxide on the ZnO surface. TEM measurements showed an average grain size of 25 nm (ZnO) and 45 nm (ZnO:Al). Due to the possible application of these films for transparent electrically conducting or infrared-reflecting purposes, the optical and electrical properties of the films were the subject of careful study. Figure 7.3 from 268 shows the optical transmittance and reflection spectra of both undoped and Al-doped films. The doped films have a higher visible/near-IR transmittance. The transmittance cutoff is blue-shifted for the doped film due to conduction band filling by electrons from the Al dopant; this results in an increase in the effective optical bandgap (the Burstein–Moss shift) and therefore an absorption blue shift. (The bandgaps measured from the spectra were 3.40 and 3.98 eV for the undoped and doped films, respectively. The value of 3.4 eV for the undoped sample is considerably higher than the usual value of 3.2 eV; these films were already quite highly conducting— see the electrical measurements later and Table 7.1). The mid-IR reflectance of the doped film also increased due to free-electron reflection.

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FIG. 7.2 ZnO film thickness vs. deposition time for different values of solution pH. (Adapted from Ref. 52.)

FIG. 7.3 Transmittance and reflectance spectra of ZnO films. Broken lines: ZnO; solid lines: ZnO:A1 (4 at.%). (Adapted from Ref. 52.)

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TABLE 7.1 Variation of Electrical Parameters of ZnO Films with Al Doping and Annealing Conditions 200°C

225°C oxygen

350°C vacuum

at.% Al

N





N





N





0 4 5.5

0.67 9.0 9.4

7.7 11.9 1.5

12 0.60 4

2.5 19

10.4 16.7

2.5 0.21

2.4 18

10.2 16.7

2.5 0.20

200C is the basic anneal that converts the hydroxide into the oxide. The films were then reannealed. The two conditions given here are those that give the most highly conducting films. Higher temperatures in oxygen decrease the conductivity, while higher temperatures in vacuum have no further effect. N: free electron concentration ( 1020 cm3); : mobility (cm 2V1s1); : resistivity  103 -cm. (From Ref. 52.)

The resistivity of the films decreased with increasing A1 content up to 4% and then increased again to the maximum measured A1 content of 5.5%. This is due to a combination of increase of carrier concentration at 4% A1, which then levels off, and increase in mobility at 4% A1 followed by a sharp decrease at 5.5% A1. The increase in carrier concentration is clearly due to the doping. Increase in mobility was linked to the increased grain size of the doped films, while the decrease in mobility (and conductivity) at doping levels  4% A1 was attributed to grain boundary segregation, resulting in higher intergrain barriers. The relevant electrical parameters are listed in Table 7.1. Further decrease in resistivity was obtained upon annealing (see Table 7.1). In all cases where maximum conductivity was obtained (350°C in vacuum or 225°C in oxygen), the decreased resistivity was due mainly to increased carrier concentration, although a moderate increase in mobility was also measured. For vacuum-annealing, these two effects were explained by an increase in oxygen vacancy concentration and desorption of oxygen from grain boundaries (therefore decrease in grain boundary barrier), respectively. The reason for the increase in carrier concentration and mobility after annealing in oxygen at 225°C is not clear, although the decrease in these parameters on annealing at higher temperatures in oxygen follows naturally from the foregoing explanation. Hexagonal ZnO films were deposited on glass or SnO2 /glass using a solution of zinc acetate complexed with ethylenediamine and with the pH adjusted to between 10.5 and 11.0 with NaOH [53,54]. The formation of the films occurred only under conditions where Zn(OH)2 was calculated to be present in the deposition solution. Good-quality, adherent films were obtained only on glass and within narrow pH and composition ranges. Films could be obtained outside these conditions, but they were then poorly adherent. In general, the adherent films formed under conditions where the deposition was slow (relatively low pH, relatively high complex:Zn ratio).

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It was proposed that deposition of adherent films proceeded on Zn clusters bound via OH groups to the glass surface. The poorer adherence on a SnO2 surface was explained by a lower concentration of hydroxy groups on this surface. In a related study, McAleese and O’Brien showed how ZnO nucleated on glass and SnO 2/glass from solutions of zinc acetate, ammonia, and thiourea (and sometimes also hydrazine) [55]. Such a deposition solution is normally used to deposit ZnS, and this study showed that ZnO can form together with, or even instead of, ZnS. From the optical spectra of the films deposited from the ethylenediaminecomplexed solutions, a bandgap of 3.15 eV was calculated (literature value 3.2 eV). ZnO was deposited on Si (100) by heating an ammoniacal solution (at pH
7) of ZnAc2 [56]. The heating regime was important: First the solution was heated to between 60 and 100°C for 6 hr and then the temperature was raised to between 100 and 200°C for 6–12 hr (in an autoclave). The initial lower temperature was necessary to obtain deposition, and it was suggested that nucleation of ZnO on the substrate occurred during this step. No XRD pattern was obtained after this lower-temperature stage; wurtzite ZnO was clearly seen by XRD after the high-temperature stage. The films, ca. 65 nm thick, were smooth, dense, and homogeneous, with some (10.0) texture. ZnO films for use as buffer layers in photovoltaic cells (see Chap. 9) have been chemically deposited from aqueous solutions of ZnSO4 and ammonia [57]. The solution was heated to 65°C, and adherent, compact Zn(OH)2  ZnO films were formed after one hour. Low-temperature annealing converted the hydroxide to oxide. The solution composition will be important in this deposition. On one hand, increased ammonia concentration will increase the pH and therefore the homogeneous Zn(OH)2 precipitation in solution. However, further increase in ammonia concentration will redissolve the hydroxide as the ammine complex. There will clearly be an optimum ammonia (and zinc) concentration where Zn(OH)2 does form, but slowly enough to prevent massive homogeneous precipitation. The use of ammonia in (hydr)oxide deposition derives, in part at least, from its gradual loss by evaporation if the system is not closed [58]. Any open solution of an ammonia-complexed metal ion (which forms an insoluble hydroxide or hydrated oxide) should eventually precipitate the (hydr)oxide for this reason alone. The borane technique, described earlier for In2O3 preparation, was also used to deposit ZnO on Sn/Pd-activated glass using a solution of Zn(NO3)2 and dimethylamine borane at 50°C [59]. The films were randomly oriented polycrystalline ZnO, and the crystals were hexagonal shaped, with a typical size of 0.2 m, at moderately high dimethylamine borane concentrations; at low concentrations, the grains were smaller and more irregular. Optical transmission spectra gave an estimated bandgap of 3.3 eV. From the spectra, the films showed some scattering, with the most transparent films having an approximate integrated transmission over the visible region of 70%, obtained from a solution containing 0.05 M/l dimethylamine borane. This correlated with

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the most regular morphology measured by SEM. A more recent study by the same group, using higher Zn concentrations and slightly higher deposition temperature (60°C), showed increased transmission up to nearly 80% over the visible range with a borane concentration of 0.1 M [60]. Electrical measurements of films as a function of boron doping were also reported in this study. The amount of boron in the films varied over only a small range, from 0.01% at very low borane concentrations in solution (a minimum concentration was needed for the deposition to proceed) to 0.02% at a borane concentration in solution of 0.1 M. The resistivity of the films was high in all cases, varying from ca. 20 k-cm to 0.4 k-cm with increase in boron content. Mobility and carrier concentration measurements showed low values of both (maximum values: 
1 cm2V1 sec1, N
1.8  1016 cm3). The low carrier concentration implies that these films are highly stoichiometric. While unfavorable for transparent conducting purposes (although annealing would probably improve the conductivity), the ability to make relatively insulating ZnO may be advantageous for other purposes.

7.2.19 Zirconium Oxide (ZrO2) A process similar to that described earlier for TiO2 was used to deposit ZrO2 films [61]. Self-assembled monolayers with terminal sulphonate or methyl groups on Si were used as substrates; no film growth occurred on bare Si. The deposition solution was Zr(SO4)2 dissolved in an aqueous HCl solution at 70°C. The film growth was slow—15 nm after 4 hr and with a limiting thickness of 40 nm after ca. 20 hr, although it was faster at the beginning of the deposition. The films were composed of a mixture of crystalline tetragonal ZrO2 and amorphous material—probably a basic zirconium sulphate. The (thicker) films varied in their thickness, from up to 10-nm-sized crystals near the substrate to 2- to 3-nm crystals, together with a greater proportion of amorphous basic sulphate, toward the film surface. It was suggested that both electrostatic forces (between the negatively charged sulphonate surface groups and positively charged zirconium oxide and basic sulphate colloids) and van der Waals attractive forces cause the observed good adhesion between the films and sulphonate monolayers while the inferior adhesion to the uncharged methyl-terminated monolayers was due solely to van der Waals forces. Annealing the films for 2 hr at 500°C, though causing pyrolysis of the monolayer, did not damage the film (or its adhesion to the substrate), which remained tetragonal ZrO2 with a crystal size ca. 10 nm and with some sulphate, the latter disappearing after prolonged annealing. Annealing at 600°C or higher resulted in a change to the monoclinic phase of ZrO2 Yttrium-doped ZrO2 was deposited by adding Y2(SO4)3 and urea and depositing at 80°C (see deposition of Y2O3 described earlier). A higher pH (2.5–3.0)

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was needed to cause coprecipitation of Y2O3, and the Y concentration in the film was ca. 25% of that in the deposition solution. Basic carbonates and sulphates were also present in the films. Annealing in air at 500°C for 2 hr resulted in complete crystallization of the films to yttrium-stabilized zirconia.

7.2.20 Ternary Oxides Deposition of two ternary oxides—(Cd,Sn)O and (Zn,Cd)O—will be mentioned very briefly next. These will be treated more fully in Chapter 8. Also, a brief description of a related technique (SILAR; see Sec. 2.11.1) that has been used for complex oxides and related compounds will be given.

7.2.20.1 Cadmium Stannate (Cd2SnO4) The ammonium fluoride technique used by Raviendra and Sharma for ZnO (described earlier) has also been used by them to deposit cadmium stannate using a mixture of CdCl2 and SnCl4 [52]. After annealing at over 200°C, Cd2SnO4 was obtained. Optical and electrical properties of these films are described in Chapter 8.

7.2.20.2 Zinc Cadmium Oxide (Zn,Cd)O Films of ZnxCd1xO, with varying values of x, were deposited from ammoniacal solutions of Cd and Zn chlorides containing hydrogen peroxide at 45°C followed by annealing in air at 500°C, presumably to convert the hydroxides to oxide [62]. The optical and electrical properties of these films are described in Chapter 8.

7.2.20.3 SILAR Deposition of Metal Oxides, Hydroxides, and Peroxides In SILAR (successive ionic-layer and reaction) deposition, discussed in Section 2.11.1, successive compound layers are built up from reaction between adsorbed (ideally mono-) layers and a reactive solution. This technique has been applied to sulphides, selenides, and oxides, including hydroxides, peroxides, and ternary oxides. As an example, films of LaxNbOy were deposited [63]. This technique was used to deposit many other oxides, hydrated oxides, and peroxides (see references in Ref. 63). In view of the uncertain purpose of H2O2, often used in CD of oxides, the role of this chemical as explained in this work is of interest. Except at a pH below ca. 2, the surface of oxidized Si (Si was used as a substrate in these experiments) in aqueous solution, and also glass, is composed of SiMO groups [see Eq. (2.16)]. These groups can attract metal cations to give an uncharged (for monovalent cation) or positively charged (for a higher-valent cation) surface. Thus, for a divalent cation: SiMO  M2 D SiMOMM

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(7.4)

This surface reacts with H2O2 to give SiMOMM  H2O2 D SiMOMMMOOH  H

(7.5)

which, in neutral or alkaline solution, dissociates to give a negatively charged surface again: SiMOMMMOOH D SiMOMMMOO  H

(7.6)

Additionally, the ability of H2O2 to oxidize a metal to a higher-valent state, resulting in a more insoluble hydroxide (higher-valent metal hydroxides are more insoluble at a particular pH than the hydroxide of the metal in a lower-valent state) has been pointed out in this study.

7.3 OTHER SEMICONDUCTORS: Se AND SILVER HALIDES The remaining semiconductors (apart from ternaries, which are treated in Chap. 8) that have been deposited by CD are elemental Se and silver halides. The little done on these materials will be discussed here.

7.3.1 Se Elemental Se exists in several forms, including amorphous (red), monoclinic (red), and gray (hexagonal). Gray Se is the most stable form. There are several reports on CD Se. The earliest is based on the fact that selenosulphate is unstable in acidic solution and, if made acid, will immediately precipitate red Se (note that its analogue, thiosulphate, behaves similarly but that a much lower pH is required to precipitate elemental S). Se was deposited by slightly acidifying dilute (10–50 mM) selenosulphate solution [64]. Films of amorphous, red Se were obtained at 10–15°C, while gray Se was obtained at 30°C or higher. Using weak acids (citric or ascorbic), amorphous Se films were likewise deposited at 0°C [65]. These conditions slow the formation of Se enough to allow films (ca. 50 nm thick) to form. A direct bandgap of 2.0 eV was measured for these films. Heating at 85°C transformed the films to gray, hexagonal Se. Another technique, which is not strictly true CD but close enough to warrant inclusion, is photodeposition from an amorphous Se colloid [66–68]. The Se colloid was prepared by reduction of a solution of SeO2. Illumination of a substrate in this solution with light that was absorbed by the Se (bandgap 2.05 eV) resulted in film formation on a substrate. Film formation occurred in the absence of illumination but was extremely slow, particularly at lower temperatures (at temperatures above 25°C, gray Se began to be formed). No XRD structure was found for the films (deposited below 25°C). Raman spectroscopy revealed the presence of Se chains and rings in the film. The mechanism of the deposition was not completely understood, but it was clearly connected with the photogeneration of

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charges in the Se, since only superbandgap illumination was effective. The bandgap of the films was measured, from the optical spectra, to be 2.05 eV [69].

7.3.2 Silver Halides (AgI, AgBr, AgCl) Section 3.2.6 begins with the sentence “Although they do not appear to have been used (at least deliberately) to form films, there are other slow anion-generating reactions.” And near the end of this section is written “It should be stressed that these reactions were used to form precipitates and not films. There is no guarantee that films can be formed using these reactions. However, it is reasonable to expect that, under the right conditions, it may be possible to produce films of these compounds. It is left as an exercise for the curious reader to find these ‘right’ conditions.” For most of the period while this book was being written, the halides were included in that section of Chapter 3. However, the very act of writing down that it should be possible to produce films of compounds containing these anions led me (after a couple of years, when the opportunity arose to spend an extended period of time in the laboratory during a visit to the university of Bern) actually to try to do this. For a number of reasons, halides were the anion of choice, specifically silver halides. The slow generation of halides is discussed in Section 3.2.5 and is based on the slow hydrolysis of haloalcohols. AgI was deposited by hydrolysis of 2-iodoethanol (ca. 50 mM): ICH2CH2OH  H2O D I  H  HOCH2CH2OH

(7.7)

in an aqueous solution containing AgNO3 (ca. 10 mM) and a small amount of triethanolamine (ca. 0.5 mM). The triethanolamine, originally added as a complexant, reduced the Ag to Ag if used in high concentration. However, small amounts were found to give a more homogeneous film than if no triethanolamine was used. The deposition works best at room temperature—heating results in excessive precipitation in solution. The iodopropanol usually contains some free iodide, the amount of which increases with age, and this can deleteriously affect the deposition if present in too high quantities. The yellow AgI films exhibited sharp XRD peaks (no line broadening), showing them to be a mixture of wurtzite and sphalerite AgI. The films (other than very thin ones) scatter light moderately strongly, and transmission spectra were taken using an integrating sphere. Figure 7.4 shows the spectrum of such a film. The strong absorption onset at 440–450 nm is due to the direct bandgap of AgI (ca. 2.8 eV). AgI also has an indirect absorption at longer wavelengths, and the decrease in transmission over this region is due partly to this absorption and probably partly to scattered radiation not collected in the integrating sphere. The deposition occurs in parallel with homogeneous precipitation, suggesting that film formation is due to adhesion of crystals from the solution. This is supported by SEM pictures that show scattered crystal formation, with gradual denCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

FIG. 7.4 Total transmission spectrum (measured with an integrating sphere) of a CD AgI film on glass.

sification of the crystal structure as the deposition proceeds. Figure 7.5 A shows such a micrograph of a AgI film on glass. The average crystal size is a few hundred nanometers, and, as might be expected from such a morphology, the film adhesion is poor—the films can be wiped off with a tissue, although they will usually stand up to cleaning in an ultrasonic bath. Adhesion is much better on

FIG. 7.5 TEM micrographs of (A) a CD AgI film on glass and (B) a similar film on SnO2-coated glass. Copyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

SnO2-coated glass; the films are not removed by rubbing with a tissue. The film morphology is different than for the films on glass; the crystals are somewhat smaller and tightly aggregated (Fig. 7.5B). AgBr films could be made in the same manner using various bromo-alcohols. AgBr is colorless but is usually slightly brownish due to photolytic formation of small particles of metallic silver (this occurs much more strongly if deposition is carried out in room light, but formation of some brown coloration is still noticeable even if deposition is carried out in the dark). This coloration masked the optical absorption, which occurs mainly in the UV for AgBr. For bromides, triethanolamine was not needed. AgCl was not so readily deposited, and while occasionally some film formation did occur, it was nonhomogeneous, thin, and irreproducible. AgCl could be deposited by immersing a substrate in an aqueous solution of AgNO3 to which NaCl solution was added, resulting in a cloudy precipitate/sol of AgCl. The concentrations of both reactants were important: ca. 20 mM each, preferably with a slight excess of Ag. Much lower concentrations resulted in little deposition (at least within a reasonable time), while even twice that concentration resulted in immediate coagulation of the AgCl and no film formation. What was unique about this deposition was that on two occasions (among five or six experiments in total), visible film formation occurred virtually immediately on mixing the solutions. This is the only example known to the author where visible film formation occurs in a rapid precipitation and contradicts the “conventional wisdom,” which otherwise seems to be valid, that the reaction leading to formation of the compound must be slow in order for appreciable film formation to occur. More details on this work are given in Ref. 70.

7.4 EXTENSION OF CD TO OTHER MATERIALS? To conclude this chapter, we look back at the earlier literature in hopes of widening both the potential deposition methods and the materials that can be deposited. As well as oxides and related compounds, other anions are considered. The resulting compounds do not necessarily fall under the common heading of semiconductors, but they are relevant in the hope of expanding the scope of chemical deposition. Table 7.2 summarizes a range of homogeneous precipitation reactions. Details of all these reactions can be found in Ref. 2 (this book, in spite of its age, is required reading for anyone wishing to pursue this line; more recent books may exist, but will probably not reduce its value). It should first be stressed that the material in this final section relates to precipitates rather than to films. However, with some effort (in some cases only a little or none, as seen from the common film formation occurring in urea precipitations), it is reasonable to expect extension to form films of the same materials in at least some cases.

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TABLE 7.2 Homogeneous Precipitation Reactions Precipitant Hydroxide

Phosphate

Oxalate

Sulphate

Sulphide Iodate

Carbonate Chromate

Chloride Arsenate Fluoride

Reagent Urea Acetamide Hexamethylenetetramine Metal chelate  H2O2 Triethyl phosphate Trimethyl phosphate Metaphosphoric acid Dimethyl oxalate Diethyl oxalate Urea and an oxalate Dimethyl sulphate Sulphamic acid Potassium methyl sulphate Ammonium persulphate Metal chelate  persulphate Thioacetamide Iodine  chlorate Periodate  ethylene diacetate (or -hydroxyethyl acetate) Trichloroacetate Urea  dichromate Potassium cyanate  dichromate Cr(III)  bromate Ag–ammonia complex  chloride  hydroxyethylacetate Arsenite  nitric acid Fluoboric acid

Elements precipitated Al, Ga, Th, Fe(III), Sn, Zr Ti Th Fe(III) Zr, Hf Zr Zr Th, Ca, Am, Ac, rare earths Mg, Zn, Ca Ca Ba, Ca, Sr, Pb Ba, Pb, Ra Ba Ba Ba Pb, Sb, Bi, Mo, Cu, As, Cd, Sn, Hg, Mn Th, Zr Th, Fe(III) Rare earths, Ba, Ra Ba, Ra Ba, Ra Pb Ag Zr La

Source: Modified from: I. M. Kolthoff and P. J. Elving, eds. Treatise on Analytical Chemistry, Part 1, Vol. 1. New York: Interscience Encyclopedia, 1959, p 741.

The list in Table 7.2 may appear incomplete to the modern chemist utilizing or studying chemical deposition; e.g., only thioacetamide is noted as a sulphide source and selenides are not included. However, when we reflect that the vast bulk of the work carried out on CD concerned just sulphides, selenides and oxides, this “old” table might point the way to a major expansion of the CD technique, both for semiconductors and for other compounds. Further processing may be expected to extend the types of material even further. For example, arsenates and phosphates may be reducible in some cases to the better-known (to the semiconductor community) arsenides and phosphides.

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REFERENCES 1. 2. 3. 4. 4a. 5. 6. 7. 8. 9. 10. 11. 12. 13a. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35.

L Gordon. Anal. Chem. 24:459, 1952. L Gordon, ML Salutsky, HH Willard. Precipitation from Homogeneous Solutions. New York: Wiley, 1959. HH Willard, NK Tang. Ind. Eng. Chem. Anal. Ed. 9:357, 1937. HH Willard, HC Fogg. J. Am. Chem. Soc. 59:40, 1937. HH Willard, L Gordon. Anal. Chem. 25:170, 1953. TP Niesen, MR De Guire. J. Electroceram. 6:169, 2001. RN Bhattacharya, P Pramanik. J. Electrochem. Soc. 129:1642, 1982. RL Call, NK Jaber, K Seshan, JR Whyte Jr. Sol. Energy Mater. 2:373, 1980. M Najdoski, I Grozdanov, B Minceva. J. Mater. Chem. 6:761, 1996. M Ortega, G Santana, A Morales-Acevedo. Solid State Electron. 44:1765, 2000. M Ocampo, AM Fernandez, PJ Sebastian. Semicond. Sci. Technol. 8:750, 1993. M Ocampo, PJ Sebastian, J Campos. Phys. Status Solidi (a) 143:K29, 1994. AJ Varkey, AF Fort. Thin Solid Films 239:211, 1994. FC Eze. J. Phys. D: Appl. Phys. 32:533, 1999. AJ Varkey, AF Fort. Sol. Energy Mater. Sol. Cells 31:277, 1993. B Pejova, A Isahi, M Najdoski, I Grozdanov. Mater. Res. Bull. 36:161, 2001. I Grozdanov. Mater. Lett. 19:281, 1994. M Ristov, G Sinadinovski, I Grozdanov. Thin Solid Films 123:63, 1985. RP Goyal, D Raviendra, BRK Gupta. Phys. Status Solidi (a) 87:79, 1985. M Izaki. Electrochem. Solid State Lett. 1:215, 1998. KO Velthaus, J Kessler, M Ruckh, D Hariskos, D Schmid, HW Schock. In: 11th ECPV Solar Energy Conf., Montreux, Switzerland, 1992, p 842. PC Rieke, BD Marsh, LL Wood, BJ Tarasevich, J Liu, L Song, GE Fryxell. Langmuir 11:318, 1995. BJ Tarasevich, PC Rieke, J Liu. Chem. Mat. 8:292, 1996. PC Rieke, R Wiecek, BD Marsh, LL Wood, J Liu, L Song, GE Fryxell, BJ Tarasevich. Langmuir 12:4266, 1996. QW Chen, YT Qian, H Qian, ZY Chen, WB Wu, YH Zhang. Mater. Res. Bull. 30:443, 1995. QW Chen, YT Qian, ZY Chen, Y Xie, GE Zhou, YH Zhang. Mater. Lett. 24:85, 1995. QW Chen, XG Li, YT Qian, YH Zhang. Mater. Lett. 31:247, 1997. B Pejova, M Najdoski, I Grozdanov, A Isahi. J. Mater. Sci.-Mater. Electron. 11:405, 2000. M Izaki, O Shinoura. Adv. Mater. 13:142, 2001. P Ruëtschi, B Cahan. J. Electrochem. Soc. 104:406, 1957. W Mindt. J. Electrochem. Soc. 117:615, 1970. RN Bhattacharya, P Pramanik. Bull. Mater. Sci. 2:287, 1980. S Gorer, A Albu-Yaron, G Hodes. Chem. Mater. 7:1243, 1995. P Pramanik, S Bhattacharya. J. Mater. Sci. Lett. 6:1105, 1987. KL Chopra, RC Kainthla, DK Pandya, AP Thakoor. In: Physics of Thin Films, Vol. 12, Academic Press, New York and London, 1982, p 167. W Mindt J. Electrochem. Soc. 118:93, 1971. E Beutel, A Kutzelnigg. Monats. 58:295, 1931.

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36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 50a. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70.

P Pramanik, S Bhattacharya. J. Electrochem. Soc. 137:3869, 1990. AJ Varkey, AF Fort. Thin Solid Films 235:47, 1993. B Pejova, T Kocareva, M Najdoski, I Grozdanov. Appl. Surf. Sci. 165:271, 2000. H Nagayama, H Honda, H Kawahara. J. Electrochem. Soc. 138:2013, 1988. AJ Varkey, AF Fort. Sol. Energy Mater Sol. Cell 29:253, 1993. T Kocareva, I Grozdanov, B Pejova. Mater. Lett. 47:319, 2001. D Raviendra, JK Sharma. J. Phys. Chem. Solids 46:945, 1985. QW Chen, YT Qian, ZY Chen, GE Zhou, YH Zhang. Thin Solid Films 264:25, 1995. D Hariskos, R Heberholz, M Ruckh, U Ruhle, R Shäffler, HW Schock. In: 13th ECPV Solar Energy Conf., Nice, France, 1995, p 1995. S Supothina, MR De Guire. Thin Solid Films 371:1, 2000. K Tsukuma, T Akiyama, H Imai. J. Non-Cryst. Sol. 210:48, 1997. BC Bunker, PC Rieke, BJ Tarasevich, AA Campbell, GE Fryxell, GL Graff, L Song, J Liu, JW Virden, GL McVay. Science 264:48, 1994. H Shin, RJ Collins, MR Deguire, AH Heuer, CN Sukenik. J. Mater. Res. 10:692, 1995. RJ Collins, H Shin, MR DeGuire, AH Heuer, CN Sukenik. Appl. Phys. Lett. 69:860, 1996. QW Chen, YT Qian, ZY Chen, WB Wu, ZW Chen, GE Zhou, YH Zhang. Appl. Phys. Lett. 66:1608, 1995. S Deki, Y Aoi, Y Miyake, A Gotoh and A Kajinami. Mater. Res. Bull. 31:1399, 1996. M Agarwal, MR DeGuire, AH Heuer. Appl. Phys. Lett. 71:891, 1997. D Raviendra, JK Sharma. J. Appl. Phys. 58:838, 1985. T Saeed, P O’Brien. Thin Solid Films 271:35, 1995. P O’Brien, T Saeed, J Knowles. J. Mater. Chem. 6:1135, 1996. J McAleese, P O’Brien. Mater. Res. Soc. Symp. Proc. Vol. 485. 1998, p 255. QW Chen, YT Qian, ZY Chen, YH Zhang. Mater. Lett. 22:93, 1995. A Ennaoui, M Weber, R Scheer, HJ Lewerenz. Sol. Energy Mater. Sol. Cells 54:277, 1998. AJ Varkey. Int. J. Mater. Prod. Technol. 10:94, 1995. M Izaki, T Omi. J. Electrochem. Soc. 144:L3, 1997. M Izaki, J Katayama. J. Electrochem. Soc. 147:210, 2000. M Agarwal, MR De Guire, AH Heuer. J. Am. Ceram. Soc. 80:2967, 1997. G Contreras-Puente, O Vigil, M Ortega-Lopez, A Morales-Acevedo, J Vidal, ML Albor-Aguilera. Thin Solid Films 361:378, 2000. VP Tolstoy. Thin Solid Films 307:10, 1997. GA Kitaev, GM Fofanoy. Zh. Prikl. Khim. 43:1694, 1970. B Pejova, I Grozdanov. Appl. Surf. Sci. 177:152, 2001. M Perakh, A Peled, Z Feit. Thin Solid Films 50:273, 1978. M Perakh, A Peled. Thin Solid Films 50:283, 1978. M Perakh, A Peled. Thin Solid Films 50:293, 1978. A Peled. Philosoph. Mag. B 53:171, 1986. G Hodes, G Calzaferri. Adv. Funct. Mater. in press.

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8 Ternary Semiconductors

8.1 GENERAL CONSIDERATIONS FOR PRECIPITATION AND CHEMICAL DEPOSITION OF TERNARY COMPOUNDS Mixed compositions are of interest mainly because they allow tuning of the semiconductor properties (most commonly bandgap and, therefore, spectral sensitivity). This is useful for various device applications. Photoconductive detectors, where a certain spectral sensitivity range is desired, is probably the main application that drove many studies on CD of ternary semiconductors. Mixed metal chalcogenides have been deposited by CD. According to simple fundamental considerations, the deposition should proceed according to the solubility products of the two separate metal chalcogenides; the one with a smaller Ksp should precipitate first, and only after, when the concentration of free (first) metal ion was low enough, would the other chalcogenide precipitate, assuming a sufficient supply of chalcogenide ions. If the difference in Ksp was large, then the solution would be almost entirely depleted of the low-Ksp chalcogenide before precipitation of the second would start. Note that this discussion relates to CD where deposition is slow; for rapid precipitation, kinetic factors might be more important, and differences in concentrations of the two cations are likely to play a more dominant role.

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This picture is in many (probably most) cases oversimplified. There are other factors that may become important in such mixed products. The literature is full of examples of precipitation of mixtures/compounds, and we will consider some of these, keeping in mind that they refer to precipitation and not CD. One is coadsorption of one cation on the precipitated compound of the other. A wellknown example of this phenomenon is adsorption of Cr 3 on ferric hydroxide precipitated from solutions containing Fe3 and small concentrations of Cr 3. The chromium is precipitated together with the Fe(OH)3 at a pH where chromium hydroxide itself would not readily precipitate [1]. This is due to strong adsorption of the “impurity” cation (Cr 3) on the precipitating Fe(OH)3. The ability of some precipitates to either cause coprecipitation of other metal salts that would, by themselves, not precipitate under the same conditions—an effect that was known as induced precipitation—or often even incorporate the “soluble” metal salt if the freshly precipitated insoluble salt was exposed to a solution containing the more soluble one—has been known for a long time; many examples from the old literature are given in a review by Kolthoff and Moltzau [2]. Another, and on the face of it, rather different example, is the coprecipitation of solid solution compounds, such as CuInS2 and CuInSe2—semiconductors of particular interest due mainly to their applicability for photovoltaic cells. It was shown, by X-ray diffraction, that the precipitate resulting from reaction between H2S and an aqueous solution containing both Cu and In3 ions was, at least in part (depending on the concentrations of the cations), single-phase CuInS2 [3]. Two factors were found to be necessary for this compound formation: (1) the presence of sulphide on the surface of the initially precipitated colloidal solid metal sulphide and (2) one of the cations being acidic and the other basic. The monovalent Cu cation is relatively basic, while the trivalent In3 cation is relatively acidic. It is not clear what the physical reason is for this latter requirement. A difference in practice between acidic and basic cations is that, in an aqueous solution of both cations, the acidic cation is more likely to be in the form of some hydroxy species (not to be confused with hydrated cations), while the basic cation is more likely to exist as the free cation. In a subsequent study of AgMGaMS precipitation by the same group, a mechanism for solid solution formation was proposed [4]. Colloidal Ag2S, with its lower solubility product, formed initially. The surface of the colloid adsorbed both sulphide and Ga (or Ga-hydroxy; see earlier) ions to form a gallium sulphide layer. This then would be similar to induced precipitation. It was suggested that this GaMS then diffused into the Ag2S, where solid solution formation occurred. In view of the high mobility of Ag ions, it seems more likely that Ag diffused outward rather than the Ga-species diffusing inward. Of course, precipitation is a highly nonequilibrium process, while CD, depending on which mechanism is operative, is closer to an equilibrium process (an ion-by-ion deposition occurs close to equilibrium, while the initial hydroxide formation in a hydroxide mechanism is

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more like a precipitation reaction; the subsequent chalcogenide exchange is again closer to equilibrium). Thus, extension of precipitation reactions to CD processes may be useful, but the comparison should be made cautiously. Various possibilities for precipitation of two cations, M1 and M 2, by sulphide are shown in Figure 8.1. The case where both metal sulphides have similar solubility products (Cd,PbS would be an example of this) is shown in the left-hand process, where a particle of M1M2S is formed (to simplify things, no information on stoichiometry is implied here). This particle may be either a mixed phase or a solid solution, depending on the miscibility of the two sulphides and the kinetics of the precipitation. The middle process shows the case for two sulphides with different solubility products. The metal sulphide with the lower-solubility product (assumed to be M1) precipitates as a separate phase, which adsorbs both sulphide ions and M2 ions (the latter either onto adsorbed S or onto lattice S). Eventually a shell of M 2S (probably containing M1S) will form. Depending on the driving forces involved, this core-shell structure may remain in that state, or diffusion into a mixed two-phase particle or single-phase solid solution may occur. A third possibility, shown on the right-hand side, is that only M1S is formed. If the M1 is sufficiently depleted by precipitation of M1S, and if sulphide formation continues, then M2S will eventually precipitate. These processes have been shown for free metal ions. However, if a cluster mechanism based on metal hydroxide colloids is involved, they are equally applicable to the formation of the solid hydroxide species. The degree of conversion of

FIG. 8.1 Scheme of various possibilities for coprecipitation of two metal (M1 and M2) sulphides.

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the hydroxides to sulphide will then depend on the differences between the solubility products of the hydroxides and sulphides as well as the hydroxide concentration in solution (pH and temperature). If a complex-decomposition mechanism is involved, then the most important factor is probably the differences between the strengths of the metal–sulphur (keeping sulphide as our example) bonds, and this is similar to the differences in solubility products of the sulphides. These processes have been described for rapid precipitation reactions. However, they should also be valid in general for slow precipitation—i.e., for CD— with possible differences due to the very different kinetics involved. Thus, if free sulphide is involved, since it is always present in very low concentration, the lower-solubility product metal sulphide is more likely to deposit first, compared to rapid precipitation. Solid-state diffusion processes have much more time to occur in CD (although they may occur in rapid precipitation after the precipitation itself), increasing the probability of solid solution formation. Probably the most basic question to be asked when depositing ternaries is: Is the as-deposited material a true single-phase solid solution? A look at Table 2.3 shows that this is often not the case. In some studies, clear structural characterization (usually XRD) shows that the film is (at least predominantly) a single-phase solid solution. Others state that the product is not a solid solution or contains a large component of other phases. However, some studies claim that the films are solid solution, without presenting clear evidence. In most cases, these are older studies, based on XRD spectra which, by today’s standards, are not clear. The XRD spectrometers of today are a lot better than those of not so long ago. Additionally, nowadays it is recognized that nanocrystalline films (as usually deposited by CD) often require more care in sampling, to avoid the common danger of incorrectly pronouncing them “amorphous” or “poorly crystalline.” Films that are not a true solid solution as deposited may often be converted to a solid solution by annealing. Similarly, solid solution formation may sometimes occur if two separate films are deposited, one on top of the other, and then annealed to effect interdiffusion. While the original intention of this chapter was to confine the contents to genuine solid solutions (or at least those that might be solid solutions), in some cases ternary films that clearly are not solid solutions are included, and this is then made clear. Another consideration is whether the deposited film is homogeneous throughout its thickness. If the composition is a function of the relative solubility products of the individual binary compounds, then the metal ion that has the lower-solubility product with the anion will deposit preferentially at first, but, due to depletion, it may become lower in concentration in the film as deposition proceeds. Thus for complete characterization of these films, compositional analysis should ideally be made as a function of spatial position in all three dimensions; this is rarely carried out in practice. The semiconductors described in this chapter are divided into two types: those composed of two different metal cations (most of the studies) and those with two different anions. In their 1982 review [5], Chopra et al. give a list of eight difCopyright © 2002 by Marcel Dekker, Inc. All Rights Reserved.

ferent ternary compounds that had been deposited by CD at that time. The work on some of those compounds has apparently not been published, other than in theses, and details of their preparation and characterization are not given; therefore those studies are not discussed here. These include the sulphides of Cd,Zn and Cd,Hg and the selenides of Cd,Hg, Cd,Pb, and Pb,Hg. Additionally, two papers on bilayer formation are treated separately in Section 8.4.

8.2 MIXED METAL COMPOUNDS 8.2.1 (Cd,Zn)O (Cd,Zn)O films were deposited on glass at 45°C from solutions of Cd and Zn chlorides to which ammonia and then H2O2 were added (the purpose of the H2O2 was not given); they were then annealed at a final temperature of 500°C [6]. (No description of the as-deposited films was given; they were presumably mixed hydroxides.) Optical spectroscopy showed that the optical bandgap shifted strongly to the red with increasing Cd concentration (at fixed Zn concentration) for small Cd:Zn ratios in the deposition bath, but it was affected only to a small extent by variations in this ratio when the Cd:Zn ratio was greater than ca. unity. From this it can be inferred that Cd(OH)2 was preferentially deposited, even though the solubility product of Zn(OH)2 is lower (by ca. 50 times). This could be explained by the greater strength of the Zn-ammine complex compared to the corresponding Cd complex (two orders of magnitude higher), resulting in a hundred times lower free-Zn2 concentration compared to Cd (for the same total concentration of each). This more than offsets the lower-solubility product of the Zn(OH)2, although the difference is not large, resulting in preferential deposition of Cd(OH)2. This provides a good example of the need to consider all the relevant parameters when trying to understand the specifics of the depositions. The resistivity of pure CdO was 3  103 -cm (CdO is normally a degenerate n-type semiconductor), which increased approximately linearly (on a semilog scale) with increasing solution Zn content up to 107 -cm (at 60% Zn) and then tailed off to a value of ca. 108 -cm for very Zn-rich films.

8.2.2 (Cd,Zn)S By far the greatest interest and effort in the CD of ternary semiconductors has been focused on cadmium zinc sulphide (Cd,Zn)S. This interest has been driven by the expected improvement in performance of thin-film photovoltaic cells (CdTe- and CIS (CuInSe2-based cells) using (Cd,Zn)S rather than the presently used CdS. This expectation arises mainly from the increased bandgap of the Zn-containing solid solution, resulting in increased transparency to shorter wavelengths of light (see Sec. 9.1.4.5 for more details). Another consideration for heterojunction for-

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mation is the decrease in electron affinity of the semiconductor with increase in Zn content (ZnS has a smaller electron affinity than CdS). The electron affinity of a semiconductor is a measure of the position of the conduction band with respect to the vacuum energy level; a lower value means a higher conduction band. Thus the alignment of the conduction band of the (Cd,Zn)S with that of the second semiconductor can be controlled to a large extent by varying the film composition. In spite of the overall chemical similarity of Cd and Zn, however, it has not proven simple to deposit true solid solutions of the sulphides. There are a number of reasons for this, some of which have been treated in detail in Chapter 4 in the discussion on ZnS deposition. We sum them up here. While CdS is less soluble than ZnS, Cd(OH)2 is more soluble than Zn(OH)2. For this reason, ZnS is more difficult to deposit than CdS, since Zn(OH)2 tends to form instead of, or together with, ZnS. (Although the solubility product of ZnS is lower than that of Zn(OH)2, the concentration of hydroxide in any typical aqueous solution will be much higher than that of sulphide). In an alkaline solution (the most common medium for CD), CdS deposition will be preferred over ZnS. The concentration of free Cd 2 should be much lower than that of Zn2 in order for ZnS to deposit according to simple solubility product considerations. However, the strength of complexation of most ligands is comparable for both Cd and Zn (ammonia and hydroxide give stronger complexes with Zn). We know of no ligand that will complex Cd enough to bring its free-ion concentration in solution the orders of magnitude lower than that of Zn in the same solution that is required (cyanide is maybe the closest to this ideal, but the difference is still not enough, and cyanide is such a strong complexant that deposition might be relatively difficult from solutions containing it in large amounts). Because of the lower solubility of Zn(OH)2, it should be possible to adjust the complexant concentration so that Zn(OH)2 is present in solution but Cd(OH)2 is not, ignoring the possibility of induced coprecipitation. CdS would therefore be formed by the (usually slower) ion-by-ion mechanism, while ZnS might be formed by the (usually faster) hydroxide cluster mechanism. This is probably not as ideal as it may sound, both because Zn(OH)2 does not readily methasize to ZnS (due to the much higher hydroxide than sulphide concentration in the solution) and because, if it did occur, it is more likely that separate phases would be formed. One point in favor of a single-phase solid solution deposition is that CdS and ZnS do readily form solid solutions in general. Thus, if the two sulphides can be simultaneously deposited, there is a good chance that they will form a solid solution if the temperature is high enough. Put another way, annealing of a well-mixed two-phase mixture of CdS and ZnS will form a solid solution if the temperature is high enough. For finely divided precipitates (as normally occurs in CD), this temperature is expected to be relatively low. There are a number of reports on CdMZnMS deposition where the CD film was either clearly shown to be mixture of phases [7] or there was insufficient ev-

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idence to support solid solution formation [8,9]. The first well-characterized deposition of a true (Cd,Zn)S alloy film was described by Padam et al. [10]. Cd and Zn acetates were used in various ratios complexed with ammonia and triethanolamine and thiourea at 90–95°C to deposit (Cd,Zn)S over the complete composition range onto glass substrates. Interestingly, the Zn will be more heavily complexed than the Cd in this solution, which shows that the mechanism of deposition is not one based solely on solubility products of the sulphides. In fact, from the crystal size measurements of a similar deposition described in Ref. 11 (see later), it is possible that the deposition mechanism is different for the two cations (the small crystals of pure ZnS and the larger ones of pure CdS suggest a cluster mechanism and an ion-by-ion mechanism, respectively). The films were characterized by a variety of techniques. Elemental analysis (EDS) showed that the Zn:Cd ratio in the film was almost equal (slightly less) than that in the deposition solution. X-ray diffraction and ED were used for phase and compositional analyses. All the compositions up to 80% Zn were wurtzite structure, while pure ZnS was sphalerite. Interestingly, while most of the films gave ring ED patterns showing nonoriented growth, some showed a degree of orientation, in spite of the glass substrate. The bandgap, calculated from optical absorption spectroscopy, varied almost linearly with composition between that of CdS (2.4 eV) and ZnS (3.6 eV), providing further evidence for solid solution formation. The films were all n-type (hot probe) with resisitivity that varied linearly (on a log scale) from 109 -cm (CdS) to 1012 -cm (ZnS). Doping by In (as InCl3 in the deposition solution) reduced ; e.g., for a Cd0.8Zn0.2S film, dropped linearly (on a log scale) with In content from ca. 1010 -cm (undoped) to ca. 105 -cm (1.5% In—the In ratio in the film was similar to that in the solution). At high In ratios, increased, explained by a decrease in mobility due to scattering by In. Annealing in H2 at 200°C also decreased . For example, a Cd0.8Zn0.2S:1.5% In film showed a minimum value for of ca. 10 -cm, presumably due to loss of S. This same method was more recently repeated with very similar results [11]. It was additionally found that the films were strongly textured (only one XRD peak—either (0001) wurtzite or (111) sphalerite), although this texture was lost if a subsequent layer was deposited to produce thicker films. The crystal size (measured from XRD peak width) varied from 20 nm (CdS) to 9 nm (ZnS). The bandgap varied between the same limits as found in the previous study, but changed more rapidly for high Zn content. The resistivity of the films varied (linearly on a log scale) from 109 -cm (CdS) to 1014 -cm (ZnS)— the latter higher than the value measured by Padam et al. Boron doping (addition of boric acid to the solution) decreased the resistivity of CdS by three orders of magnitude. Using ammonia-complexed metal iodides and thiourea at pH 10, films were formed whose properties depended on the temperature–time regime of the depo-

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sition solution [12]. If the reagents were mixed at room temperature and deposition occurred while heating the solution to 80°C, only CdS was deposited from various mixtures of Cd and Zn iodides, even where the Zn was in large excess. If, however, the reagents were mixed at a higher temperature (60°C), followed by heating the bath to 80°C during deposition, Zn was incorporated into the films, although the Zn content in the films was much lower than in the solutions; the Zn content increased slowly up to ca. 80% Zn in solution and then rapidly at higher Zn concentrations. This was seen by both optical spectra and XRD, the latter in particular supporting solid solution formation. A reason for this dependence on temperature programming was not suggested. A possible clue may be obtained from a consideration of the temperatures at which solid Cd(OH)2 and Zn(OH)2 are formed (from thermodynamic calculations based on the equilibrium constants of the Cd- and Zn-ammine complexes, the solubility products of the hydroxides, and dependence of hydroxide concentration in water on temperature). At room temperature, no solid hydroxide phase is calculated to be present in the various Cd/Zn solutions. At approximately 40°C, hydroxide will form (slightly lower temperature for Cd, slightly higher for Zn, although the difference is not large, and, in view of the approximations used in these calculations, as well as kinetic factors, it is not certain that Cd(OH)2 will, indeed, form first). Also, it has been shown that Cd(OH)2 can form on the substrate in some cases before it forms in solution (discussed in Chap. 3), but the equivalent experiment has not been done for Zn. However, if we assume that Cd(OH)2 will form before Zn(OH)2 as the temperature is raised, then this might explain, at least in part, the formation of only CdS. For the case where deposition was started at 60°C, then both hydroxides are present to begin with and it is more likely that Zn will be incorporated into the final films. This reasoning is based on the expectation that cluster deposition will be much faster than ion-by-ion deposition. The resistivity of the films decreased with increase in Zn content (from 1010 -cm for CdS (a very high value for CdS) to ca. 106 -cm for 90% (solution concentration) Zn and then increased to ca. 109 -cm for pure ZnS. No explanation for this effect was given. The films were photoconductive, with the resistivity decreasing in a somewhat sporadic manner as a function of composition, up to a maximum dark:light ratio of 5  103 for the 90% Zn films. Similar conditions, but at a lower pH of 8.4, were also used by the same group [13]. The ammonia concentration was reported to be important in forming the solid solutions, although this concentration was not given. Only Zn-rich solutions were described in this study (between 80 and 99% Zn in solution). As for the previous study at higher pH, the Zn concentrations in the films were quite different from those in the solution (except for very Zn-rich solutions); the films were richer in Cd up to at least 92% Zn in solution, and then the Zn concentration in the films increased rapidly with further increase in solution Zn concentration up to 99% Zn, from which deposited films with very little Cd.

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X-ray diffraction showed well-defined peaks that shifted in position with change in composition, while the optical absorption spectra gave values of Eg that also varied gradually with composition. The well-defined shift in XRD peaks and, to a somewhat lesser extent, the gradual change in estimated bandgap with composition provide good evidence for true solid solution formation. Unlike other studies as well as the higher-pH studies by the same group, the electrical resistivity did not vary much with composition, being ca. 5  107 -cm. An overview of the variation of resistivity of some of these (Cd,Zn)S films with composition is given in Figure 8.2. There is quite a large variation, both in resistivity values and in their compositional dependence. The question must be asked: Why are solid solutions formed in some cases and not in others? A common denominator in the successful films and their difference from the unsuccessful ones (success being defined as formation of a solid solution) is the higher temperature used in the former (80–95°C). Higher temperature will facilitate intermixing of the codeposited CdS and ZnS.

8.2.3 (Cd,Zn)Se Two selenosulphate baths have been described for (Cd,Zn)Se. In the first [14], solid solution formation was claimed, at least for annealed (300°C) samples, al-

FIG. 8.2 Resistivity data for some (Cd,Zn)S films. The data from the two Yamaguchi papers were modified to show resistivity values as a function of approximate film composition, rather than solution composition as given in the original papers.

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though XRD results of the (annealed) samples were difficult to understand (the “a” lattice parameter increased as Zn concentration increased). The effects of composition and annealing on the bandgaps and electrical conductivity were described. In the other [15], solid solution formation over part of the composition range was claimed. However, the XRD spectra of the films (relatively sharp peaks characteristic of ZnO), together with the visual description of the pure ZnSe films (white with a slight greyish tinge), suggest mixed-phase formation containing ZnO.

8.2.4 (Cd,Hg)S Solid mixtures of CdS and HgS have been shown to form solid solutions after treatment with certain solutions, such as concentrated ammonium sulphide [16]. This may be due partly to the very similar ionic radii of the two cations and (maybe more important) the ability of Hg to diffuse readily in solids. Therefore it is probable that solid solutions can readily form in this system. (Cd,Hg)S was deposited from a solution containing CdCl2 and HgCl2 complexed with a low concentration of cyanide (15 mM CN to 50 mM CdCl2—the Cd concentration was fixed and the Hg was varied), thiourea, and ammonia together with KOH at 80–85°C onto Ti substrates [17]. The mole fraction of Hg in the films was ca. four times its concentration in the solution. This was expected based on the lower-solubility product of HgS compared to CdS. The maximum mole fraction of Hg in the films (composition measured by atomic absorption spectroscopy) was 0.18; attempts to increase this value by adding more HgCl2 resulted in rapid precipitation in the solution and little film formation. Increased Hg concentration in the films could probably be obtained by optimization of the conditions, e.g., by reducing the Cd concentration and/or by using a specific complex for Hg, such as iodide. The bandgap of the films (annealed at 320°C for 3 hr in air), measured by photoelectrochemical photocurrent spectroscopy, decreased with increasing Hg content down to 1.8 eV for 0.18 mole fraction Hg. The shape of these spectra suggested that the (annealed) films were solid solutions, although no structural characterization was made. The main purpose for making these films seems to have been to study their photoelectrochemical properties, which are described in Chapter 9. Triethanolamine was also used as a complexant to deposit these films from thiourea baths [18]. As with the previous study, there was a maximum Hg content in the bath (0.05 mole fraction—absolute concentrations were not given), which led to a 0.18 Hg mole fraction in the films, above which, although films were formed, the Hg content decreased, also explained by rapid precipitation of HgS in the solution. X-ray diffraction showed the formation of a single phase, up to a Hg content (in the bath) of 0.15, and two-phase formation at higher concentrations. The optical bandgap dropped from 2.4 eV (pure CdS) to 1.76 eV (0.05 Hg in bath,

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0.18 in film) and then slowly increased again with increasing Hg in the bath to ca. 1.9 eV.

8.2.5 (Cd,Pb)S Solid solutions of Cd and Pb sulphides have been popular in ternary CD. The crystal structures of the individual sulphides are different: PbS crystallizes in the rocksalt structure, while CdS forms tetrahedrally bonded sphalerite or wurtzite structure (a rocksalt form of CdS exists but normally only at high pressures). This suggest that the solubility range of the alloys will be limited. Acetates of Cd and Pb were mixed with ammonia and thiourea and films deposited on glass at room temperature [19]. The concentration of Cd in the films was a little higher than in the solution. The lattice constant increased with increase in Cd but was greater than that of PbS (which has a slightly larger lattice constant than CdS) at all levels measured. This is not typical of solid solution formation, although it does imply a single phase. It was suggested that lattice expansion occurred due to PbS entering as an interstitial into the CdS. It seems more probable that, if interstitial expansion is the correct explanation, Cd2, which is the lowerconcentration component, will go into interstitial sites in the PbS. The morphology of these layers was investigated as a function of composition [20]. In a similar study, (Cd,Pb)S was deposited on Ti at ca. 75°C [17]. The ratio of Cd to Pb in the films was found (by atomic absorption spectroscopy) to be very similar to that in the solution, a consequence of the similar values for the solubility product of the two sulphides. Since the main purpose for investigating these films was to study their photoelectrochemical properties (see Chap. 9 for details), little no characterization, other than compositional and photoelectrochemical, was made. Photoelectrochemical spectroscopy (of films annealed at 460°C in air for 4 hr) showed a decrease in bandgap with increasing Pb content down to 1.6–1.7 eV for Cd0.82Pb0.18S, although very nonlinearly—a strong drop in bandgap, of 0.5–0.6 eV, occurred between 0.1 and 0.18 mole fraction Pb. Films deposited from mixed Cd/Pb solutions complexed with ammonia (for Cd) and hydroxide (for Pb), both in a minimum amount to effect dissolution, at pH values between 10 and 13 and deposition temperatures between 60 and 80°C were concluded, from consideration of the XRD, TEM (which showed two different crystal sizes), and optical spectra, to be mixtures of the two sulphides rather than a solid solution [21]. A study of these films for solar-selective surfaces was carried out. A study of variation of the composition of (Cd,Pb)S films as a function of the free-[Cd2]:[Pb2] ratio (i.e., the uncomplexed metal ions, which could be calculated from the concentrations of total metal ions, concentration of complex, and the respective stability constants of the metal complexes) showed a linear increase in Cd content of the films up to a ratio value of 10, followed by a sudden decrease

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in Cd content above this value (the details of the solution composition were not given) [22]. The maximum amount of Cd that could be incorporated as a solid solution with PbS was 15%. It was suggested that for [Cd2] greater than this, the rate of CdS formation was so great that it became more favorable for the CdS to form a separate phase than a solid solution. The crystallographic texture of the films was dependent on the Cd content. Up to 3 at.%, the films were (111) textured, while for higher Cd concentrations they became (200) textured. The crystal size (measured from electron microscopy) was of the order of some hundreds of nanometers (somewhat smaller for larger Cd content) but increased again to ca. 1 m for maximum Cd content just before phase separation. The resistivity of the films increased from 106 -cm for very low Cd content to a maximum of ca. 108 -cm at 6% Cd and then slowly dropped again with increasing Cd. This maximum correlates with the minimum crystal size, suggesting a dominant role of grain boundaries in the conduction mechanism. The spectral response of the photoconductivity blue-shifted with increase Cd content up to a peak response at 1.35 m for 8.4% Cd. The presence of cyanamides of Cd and Pb in films of (Cd,Pb)S was confirmed by thermal desorption mass spectrometry [23]. Cyanamide (H2CN2) is a product of the decomposition of thiourea and forms sparingly soluble metal salts. The metal cyanamide content of the film varied from ca. 5% up to ca. 20% (by weight). The presence of the cyanamides decreased the intensity of the XRD reflections, presumably due to poorer crystallization of the sulphides. Interestingly, the photosensitivity of the films increased with higher metal cyanamide content, although whether this was due specifically to the presence of the cyanamide or to its effect on the crystal growth was not known. The same group also deposited (Cd,Pb)S using a flow system [24]. In this case, metal cyanamides were not detected by XRD, presumably because the flow system removed the cyanamide. The rate of flow affected the crystal size: Larger flow rate resulted in finer-grained deposits. Elemental analysis and XRD showed the incorporation of Cd in the films, again up to ca. 10%, as a solid solution.

8.2.6 (Hg,Pb)S Films of (Hg,Pb)S were grown on glass at 30°C from a solution of PbAc2, HgCl2, thiourea, and NaOH at pH
10 [25,26]. It was noted that the pH had to be critically controlled to obtain good-quality films. Additionally, the order of mixing of the solutions was unusual (it was not stated if this was critical or not) in that the PbAc2 and thiourea were first mixed, the NaOH was then added until a light brown color appeared in the solution, and only then was the HgCl2 solution added, followed by adjustment of pH to 10.

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For films deposited at temperatures up to 45°C, ED showed the presence of a single phase, with increasing lattice spacing with increased Hg content up to 4% Hg. Above this concentration, aggregates of -HgS (metacinnabar) were found in the deposit, which originated from the colloidal solution. While not clearly defined, the transmission spectra shifted to the blue (increasing bandgap) with increase in Hg concentration. This increase in bandgap, together with the ED data, suggests that the films are alloys of PbS with -HgS (cinnabar, bandgap ca. 2 eV). For deposition temperatures greater than 45°C, the lattice parameters decreased with increasing Hg, and the -HgS phase was formed (the lattice parameters of HgS are ca. 1.5% smaller than those of PbS). If FeCl3 was added to the HgCl2 solution, the film properties, even for those deposited at low temperatures, were similar to those grown without FeCl3 at higher temperatures. Single-phase films were obtained up to 50 at.% Hg with the -HgS phase, and the lattice spacing decreased with increasing Hg content. The transmission spectra, while again mostly not well defined, shifted to the red with addition of Hg. The bandgap of -HgS was taken to be ca. 0.1 eV in these studies; other values of ca. 0.5 eV have also been measured for this phase. The FeCl3 was believed to stabilize the -phase. A specific study of the optical and electrical properties of these Pb1xHgxS films was carried out with an emphasis on the difference between the - and -phase alloys deposited at 30°C as described earlier [27]. A linear increase in bandgap (up to 0.9 eV for x
0.33) for the -phase alloy and a linear decrease of the -phase down to ca. 0.18 eV for x
0.33 was measured. The resistivities of the alloys (x
0.14) were higher than for the pure PbS (ca. 10 -cm) by a factor of ca. 5 (-phase) and of 10 (-phase). Both the photoconductivity response and the thermoelectric power of the alloys were greater than for the pure PbS. The electrical properties were believed to be controlled mainly by intergrain barriers. Using the FeCl3-containing solution, epitaxial films of -Pb1xHgxS were grown on (111) Si or Ge single crystals, where x varied between 0 and 0.33 [28]. The conditions to obtain epitaxy were low temperature (20°C), relatively dilute solution (concentrations not given, but the typical concentrations were high— metal concentration probably several hundred mM), and relatively thin films (80 nm; above this thickness, -phase deposition occurred). The films on Ge 11) oriented, while those on Si were (1 12). The requirement for low temwere (1 perature and relatively low concentration of reactants, both of which slow the deposition process, suggest that the epitaxy occurs if enough time is allowed for crystal growth to occur. Decreased temperature will decrease both crystal growth and the rate of attainment of epitaxy (probably a surface diffusion process); for epitaxy to be preferred at lower temperatures implies that the effect of temperature on the former is greater than on the latter. The attainment of epitaxial growth is strong evidence for an ion-by-ion mechanism, even though parallel homogeneous precipitation occurred in the solution.

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Terminal thickness usually decreases with increase in temperature, due to faster homogeneous precipitation in the solution in parallel with film deposition. In this study, the terminal thickness increased with increasing temperature. This trend also suggests that ion-by-ion growth dominates, since homogeneous precipitation is less likely in an ion-by-ion process than in a cluster one. Although such homogeneous precipitation was observed, it seems likely that it occurred to a lesser extent than it would in a cluster process.

8.2.7 (Cd,Bi)S Films of BiMCdMS were deposited from triethanolamine/ammonia-complexed nitrates of Cd and Bi using thiourea [29]. The films were deposited at a pH of ca. 10 on glass or Si (111) at 95°C (for 90 min), followed by cooling to room temperature and continuing deposition for 24 hr. Rutherford Backscattering (RBS) analysis showed that both Cd and Bi were incorporated into the films very roughly in proportions similar to those present in the solution. No structural data was given to support solid solution formation. The gradual variation of bandgap with composition of annealed films (450°C in argon for 2hr), from 1.65 eV (for pure Bi2S3) to 2.43 eV (for pure CdS), suggested solid solution formation, at least for the annealed films.

8.2.8 (Cu,Pb)S Two-phase films of PbSMCuS were deposited on glass from a triethanolamine/ thiourea bath at room temperature [30]. As deposition proceeded, the films became Pb-rich as Cu was depleted by more rapid formation of sulphide. The resistivity of the films was
10 -cm (1 m/sq. for a film thickness of ca. 0.1 m).

8.2.9 (Cu,Bi)S (and (Cu,Sb)S) Films were deposited from a triethanolamine/thiourea bath containing CuSO4 and Bi(NO3)3 [31]. No compositional or structural characterization was given; therefore there is no evidence that this was a solid solution or even a mixture of Cu and Bi sulphides. An example of solid solution formation by separate deposition of binary layers followed by annealing to interdiffuse the two layers is given for Cu3BiS3 deposition [32]. Bi2S3 (film thickness ca. 90 nm) was deposited at room temperature from a Bi(NO3)3/triethanolamine/thioacetamide bath onto glass slides. CuS (300–600 nm thick) was then deposited on this film from a CuCl2/triethanolamine/ammonia/NaOH/thiourea bath at room temperature. The films were annealed at 250°C for 1 hr. Formation of the Cu3BiS3 phase could be seen from the XRD pattern. Measurement of precipitated powders (prepared by putting the Bi2S3 precipitated in the first deposition in the CuS deposition solution) annealed at 300°C showed more clearly the formation of the solid solution.

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Transmission spectra of the annealed films showed an approximate bandgap of 1.9 eV. The films were p-type semiconductors with fairly low resistance (depending on annealing conditions; even the as-deposited films had a sheet resistance of 7 k/sq., which dropped to ca. 100 /sq. on mild annealing at 150°C). In a similar manner, Cu3SbS4 was formed by depositing a layer of CuS (as earlier) onto a previously deposited film of Sb2S3 (from a thiosulpahte bath) and annealing in N2 at 250°C [33]. These films (typically between 0.1 and 0.3 m thick) were highly conducting (some tens of /sq). The films were evaluated for solar-control purposes and exhibited good IR reflectivity/low IR transmittance with sufficient visible transmittance.

8.2.10 (Pb,Sn)Se Films were deposited from solutions of lead and tin salts (the salts used were not specified) with ammonium acetate, ethylenediamine, and selenourea at a pH  9 (probably at least 11) [34]. To obtain thicker films, deposition was repeated a number of times and the films were annealed; therefore it is not known if solid solution formation occurred in as-deposited films. In annealed films, Pb1xSnxSe solid solutions with x up to 0.11 were verified by XRD. The spectral response of the photoconductivity of the (annealed—as-deposited films were not photosensitive) films shifted from a peak at ca. 4 m (pure PbSe) to ca. 7.5 m (11% Sn), supporting solid solution formation of the annealed films. The room-temperature, dark resistance of the (probably annealed, but not certain) films varied from 1 to 300 k, depending on deposition conditions.

8.2.11 (Pb,Bi)S From solutions of Bi and Pb nitrates, complexed with triethanolamine and ammonia, mixed sulphides were deposited with thiourea on glass at pH values between 9.5 and 11 and at 100°C (initially) followed by slow cooling in the solution [35]. Elemental analyses showed the presence of both metals in the films. It is not clear whether solid solution formation occurred in the as-deposited films, although the lattice parameters did vary non-monotonically, depending on composition.

8.2.12 CuInS2 and CuInSe2 CuInS2 (and, even more, CuInSe2) are strong candidates for thin-film photovoltaic cells. For this purpose, the chalcopyrite structure (which is an ordered lattice) is preferred over the disordered, zincblende form. Due to the large absorption coefficients of these materials, a 1-m-thick film is more than enough to absorb almost all the suprabandgap radiation. Somewhat thicker films are generally used, due to problems of pinholes, which commonly occur in thinner films. A number of methods have been used to deposit these films. Surprisingly, very few (published) attempts have been made to deposit them by CD.

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A note of caution is necessary when dealing with these materials. It is not trivial to distinguish between CuInS(Se)2 and some phases of CuMS(Se). Diffraction and optical properties may be similar. Elemental analysis is particularly important to verify inclusion of indium in the films and in the correct ratio. A fingerprint of the chalcopyrite XRD is the presence of a weak peak at 2
17–18°, corresponding to the (101) chalcopyrite reflection. This is often not seen, although this could be either because the deposit is not chalcopyrite or because weak peaks are usually not seen in nanocrystalline materials with particularly small crystal size.

8.2.12.1 (Cu,In)S Deposition of single-phase, chalcopyrite CuInS2 was claimed using a solution of CuCl2, InCl3, complexed by triethanolamine and ammonia, with thiourea [36]. (Note that this claim was contested, based on the diffraction and compositional data [37,38].) The best films were somewhat Cu rich (typically Cu1.08In1S1.5) and, more importantly, very S deficient. The excess Cu is not surprising, considering (1) the greater concentration of Cu in the solution compared to In and (2) the lower solubility of CuMS compounds compared to InMS ones (see, however, the discussion on CuMInMS compound precipitation in Sec. 8.1). The fact that films with excess In can be obtained with In:Cu ratios in solution of less than 1 suggests that the coprecipitation is more complicated than expected based solely on solubility products or even taking into consideration adsorption of In on CuMS as inferred from early studies on this and similar systems described in Section 8.1 [3,4]. Overall, it appears likely that the films contained chalcopyrite CuInS2 mixed with other phases with similar diffraction patterns. Separate microstructural characterization (EDS) of films with varying composition (ca. 10% excess Cu or In) showed the formation of separate phases of Cu2S and In2S3, respectively, along with the CuInS2 [39]. The best films were obtained at high deposition temperatures (80°C) and with stirring. Lower deposition temperature resulted in poorer stoichiometry (less S), and stirring improved film uniformity. Grain size, measured by TEM (which does not necessarily show crystal size) was 100–400 nm. From the optical transmission spectra, a bandgap of 1.50 eV was found for the most stoichiometric (in terms of Cu:In
1; the S content was always found to be low) films. This value dropped slightly for nonstoichiometric films [39]. Resistivities varied with composition, from ca. 50 -cm for In-rich films down to ca. 0.1 -cm for very Cu-rich ones [36].

8.2.12.2 (Cu,In)Se Three studies used essentially the same baths—ammonia and triethanolaminecomplexed Cu and In salts—and selenosulphate as Se source [40–42]. In all cases,

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a Cu:In ratio close to unity in the solution resulted in an optimum (and ca. unity) ratio in solution. In fact, the ratio in the film was not strongly dependent on the ratio in solution (at least, over the narrow range measured) [42]. This is not obvious based on solubility product considerations and suggests some form of compound formation, as described in the related precipitation experiments of Rudnev et al. [3,4]. In the earliest study, the deposition was carried out at room temperature; no elemental analysis was made, and the diffraction data do not show the presence of chalcopyrite CuInSe2, although the sphalerite phase might be present. The other studies used higher temperatures, a parameter that appears to be important. More stoichiometric films were found at 90°C than at 50°C [41], and in one case, good, adherent films were obtained only at 85°C; at lower temperatures, powdery, nonadherent films were obtained [42]. Both these studies reported chalcopyrite CuInSe2 as deposited. A bandgap of 1.08 eV was calculated from the optical spectrum for the low-temperature deposition of Ref. 40 and one of 0.9 eV for the hightemperature (90°C) deposition of Ref. 41. All films that were characterized for conductivity type were found to be p-type, with resistivities that varied between 0.08 and 500 -cm, depending mainly on the excess Cu content, which resulted in low-resistivity films [41,42]. These are low values and suggest appreciable Cu excess, although elemental analysis showed some of the films to be close to stoichiometric. A similar deposition, using ammonia but with citrate complexant instead of triethanolamine and at 40°C, was also reported [43]. From XRD measurements (not shown in the study), predominantly chalcopyrite CuInSe2 was reported if the Cu:In:Se ratio in solution was ca. 1:1:2. EDS analysis confirmed this approximate ratio in the films. From the absorption spectra, a bandgap of 1.4 eV was measured, which decreased to 1.15 eV on annealing to 520°C (the literature bandgap of CuInSe2 is ca. 1.0 eV). Resistivity and Hall measurements of these films as a function of composition are interesting. The resistivity increased to a sharp maximum of ca. 108 -cm at a Cu:In ratio of 1.5; lower values of Cu:In resulted in lower-resistance p-type films, while higher values (more Cu) gave low-resistance n-type films (Fig. 8.3). This is unexpected in that (a) In-rich films of CuInSe2 are normally n-type, while Cu-rich films are p-type and (b) the highest resistivity would be expected for Cu:In  1 if the material is CuInSe2. These results, together with the higher-thanusual bandgap of the (mainly annealed) films, suggest that the films are not simply single-phase CuInSe2, but either a mixture of phases or a different composition. The most recent investigation is closest to the previous one, but using only citrate as a complexant (no ammonia), a lower pH than the other studies (8 instead of ca. 10), and deposited at room temperature [44]. The films from this deposition were not adherent as deposited and required annealing (300°C) to become adher-

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FIG. 8.3 Resistivity vs. Cu:In ratio for CuMInMSe films (annealed at 520°C). (Adapted from Ref. 43 with permission from Elsevier Science.)

ent. The films again appeared, from XRD, to be chalcopyrite. The XRD peaks of the as-deposited films were rather broad (crystal size of ca. 14 nm). As with the previous study, the bandgap of the as-deposited films was anomalously high (1.3 eV), which, considering the small crystal size, may be a quantum size effect. After annealing, the crystal size almost doubled and the bandgap dropped to the normal CIS value of 1.02 eV. The films were p-type and highly conducting: 103 cm as deposited and ca. 2 -cm after annealing. Again, these low resistivities suggest that free CuMSe species were present.

8.2.13 (Cd,Sn)O Cadmium stannate (Cd2SnO4) was deposited from a solution of CdCl2 and SnCl4 using NH4F, ostensibly as a freezing agent, although it is probable that it also functioned as a complexant (see this technique for SnO2 deposition in Chap. 7) along with a small amount of AgNO3 as catalyst (not clear for what) and NaOH to adjust pH to between 7.5 and 8.5 [45]. The film grew to a maximum thickness of 0.8 m in 40 min at a pH of 7.5 (faster deposition but lower terminal thickness at higher pH). X-ray diffraction of films annealed at 200°C or above showed them to be Cd2SnO4 with a grain size of 25 nm (20 nm before annealing).

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FIG. 8.4 Transmittance and reflectance spectra of CdMSnMO (Cd2SnO4) films: as deposited, annealed in H2 and in vacuum, both at ca. 200°C. (Adapted from Ref. 45.)

The optical transmittance and reflectance spectra of the as-deposited film and films annealed in H2 and in vacuum at ca. 200°C are shown in Figure 8.4. All the films show high transmission in the visible/near-IR region, in particular the annealed ones. A high transmittance is important if the films are to be used for their transparent, conducting properties. The blue shift of the absorption after annealing was explained by the increase in bandgap due to a high free-carrier concentration (the Moss–Burstein shift caused by filling of the lower part of the conduction band by free carriers). The increased carrier concentration after annealing is also the cause of the shift in the reflectance spectrum to shorter wavelengths upon annealing; annealing in H2 results in a larger free-carrier concentration due to removal of more oxygen (heavier n-type doping). The bandgaps, calculated from the transmission spectra, were all indirect and were 2.7 eV (as deposited), 3.1 eV (vacuum-annealed), and 3.2 eV (H2-annealed). The resistivities of the films were 4  101 (as deposited), 102 (vacuumannealed), and 4  103 -cm (H2-annealed). The decrease in resistivity was due mainly to increase in free-electron concentration (2  1018, 5  1019, 1020 cm3) for the three films; the mobility increased by a factor of two between the as-de-

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posited and H2-annealed films. The increase in electron concentration was due to oxygen removal, as mentioned earlier. The increased mobility was possibly due to desorption of surface-adsorbed oxygen.

8.3 MIXED CHALCOGEN COMPOUNDS A number of mixed sulphide/hydroxides have been deposited, mainly in the search for improved window layers for photovoltaic cells (Chap. 9). These are mostly probably mixed-phase films, although in one case, In(OH)S, experimental evidence suggests true compound formation [46]. Most of these films have been dealt with in previous chapters (see Chap. 4 under ZnS and Chap. 6 under In and Sn sulphides). One study (described from the viewpoint of its properties in photovoltaic cells in Chap. 9) has not been described previously and will be mentioned briefly here. This deals with Zn(O,OH) and Zn(O,OH,S) deposited from Zn-ammine solutions, the latter film from solutions also containing thiourea [47]. It is of interest to note that the Zn(O,OH) films did not deposit on glass but did on both ZnO- and CuInSe2-type substrates. Even after annealing at 300°C, hydroxide groups were still present in those films.

8.3.1 Cd(S,Se) In two of the studies made of Cd(S,Se) deposition [48,49], the solutions and conditions of deposition were similar. In both cases, Cd was complexed with ammonia, a mixture of thiourea and selenosulphate was used, and the deposition was carried out on glass at 75°C. From the study of Kainthla et al. [48], XRD of the films showed clearly that solid solution formation occurred; the (predominantly sphalerite) diffraction peaks shifted with change in composition. For compositions with S concentration  60%, only zincblende structure formed; the amount of wurtzite increased with increasing S content but was always low. The concentration of S in the films was somewhat greater than that in the deposition solution; i.e., S deposited preferentially. This is not surprising since CdS deposition is normally faster than that of CdSe. The concentration of ammonia was increased as the thiourea:selenosulphate ratio increased, ostensibly to slow down the rate of formation of CdS through decreased Cd2 concentration (although the rate of CdSe formation is also dependent on this same factor). Optical spectra showed a gradual shift of the onset with composition, as expected for a true solid solution. Additionally, the refractive index (and therefore the dielectric constant) increased gradually as the Se content increased, mirroring the larger dielectric constant of CdSe compared with CdS. The study of Ref. 49 gave similar results, although XRD peaks of the films were very weak and difficult to interpret. In addition, the electrical resistivity of

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the films was measured. The resistance of the S-rich and slightly Se-rich films was typically 1–3  106 -cm, while for high Se concentrations the resistivity fell to a value of 7  104 -cm for pure CdSe. Another study of Cd(S,Se) deposition was similar to the preceding ones, with the differences that triethanolamine was used instead of ammonia (solution pH
10.4) and the deposition was carried out at 55°C [50]. The main difference in the films was that a true single-phase solid solution formed over only part of the composition range (for values of x in CdSexS1x between 0 and 0.4 and between 0.85 and 1). This was paralleled by sharp absorption onsets for the solid solutions and more gradual ones for the mixed-phase systems. Crystal sizes varied in a nonmonotonic manner from 27 nm (CdS) to 9 nm (CdSe). The resistivities were of the same order (for CdS, slightly less) as those of the previous study, although the variation of resistivity with composition was somewhat different. However, since the variation in resistivity over the entire composition range was only a little more than an order of magnitude, which is not particularly large, such differences need not be very meaningful. The same group also studied In-doped (by addition of InCl3) CdS1xSex [51]. As the In concentration increased, the degree of crystallinity (measured by the height of the XRD peaks) and crystal size increased, reached a peak, and then decreased. These structural changes correlated with other properties: The bandgap and resistivity were minimum when the crystallinity and crystal size were maximum. The concentration of In in the films was much higher than that in solution (2% of the Cd concentration, compared to 0.1% in the solution) [52].

8.3.2 Zn(S,Se) Zn(S,Se) has been deposited on both glass and on single-crystal GaAs (110) from a hydroxide-complexed solution of Zn2 using, as for Cd(S,Se), a mixture of thiourea and selenosulphate [53,54]. Apparently conditions were chosen to give the composition ZnS0.056Se0.944 because of its perfect lattice match with the GaAs substrate. The composition did not appear to be dependent on the deposition temperature. Room-temperature deposition resulted in films with very broad peaks, which sharpened considerably with increasing deposition temperature to give a crystal size of ca. 20 nm at a deposition temperature of 90°C. The high-temperature films on GaAs exhibited a fairly high degree of epitaxy, as seen by the spots in the electron diffraction pattern.

8.3.3 Pb(S,Se) Pb(S,Se) was deposited from a hydroxide-complexed solution of Pb(NO3)2 using, as before, a mixture of thiourea and selenosulphate [55]. The films were

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confirmed to be single phase by XRD (Debye–Scherrer photographs). In contrast to the case for Cd and Zn, Se was preferentially deposited. (Small selenosulphate concentrations in solution resulted in much larger Se concentrations in the film. For example, from a solution containing 0.5 mM selenosulphate and 500 mM thiourea, the film composition was ca. PbS0.75Se0.25.) This was explained by the large difference in solubility products of PbS and PbSe (nearly 10 orders of magnitude; see Table 1.1). For Cd and Zn, this difference is some orders of magnitude less. Thus, while for Cd and Zn it seems that the faster decomposition of thiourea compared to selenosulphate more than compensates for the lower-solubility products of the selenides, for Pb the difference in solubility products between sulphide and selenide becomes the main composition-determining factor. From XRD line widths, it was noted that the crystal size increased as the S concentration increased, although values were not given. A simple but useful characterization technique to quantify film thickness homogeneity was used here. The transmittance of a focused light spot scanned across the sample showed excellent homogeneity.

8.3.4 Bi2(S,Se)3 In this case, thioacetamide was used as the sulphur source, instead of thiourea as for the previous mixed sulphides-selenides (selenosulphate, as before, was used as the Se source) [56]. Bi(NO3)3 was complexed with triethanolamine and the pH adjusted with ammonia to 8.2. The deposition was carried out at 55°C. The composition was varied by varying the thioacetamide/selenosulphate ratio. Although it is not clear what the elemental compositions of the various films were, from the limited XRD data given, it seems that solid solution did occur. The crystal sizes increased from 6 nm (pure sulphide) to 13 nm (pure selenide), and bandgap values decreased over the same range from ca. 1.9 to 1.0 eV.

8.4 BILAYERS OF DIFFERENT SEMICONDUCTORS This chapter has dealt with true ternary compounds, with the underlying implication that deposition of separate phases is undesirable. However, it needs to be stressed that what is undesirable for one purpose may be preferable for another (examples being small crystal size and a large amount of scattering). So, too, a composite of different phases may be the goal of a particular deposition. This issue does not appear to have been dealt with in CD. Taking the principle of separation of phases one step further, separate layers may also be deposited, one on top of another. This has been done in a number of cases and should present no problem (taking into consideration that there may be some cases where deposition of the second layer will destroy or change in some

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way the first one). There are, however, two examples in the literature of deposition of a bilayer from a single deposition solution. In the first [57], a bilayer of CdS/ZnS was formed using the electrochemically assisted technique described for CdS (Sec. 4.1.6.6). In this technique, a mixture of CdCl2 and ZnCl2 was mixed with thioacetamide at a pH of 2.45 (added HCl) and deposited on ITO/glass at 70°C. In the absence of electrochemical polarization, only ZnS, with a low percentage of CdS, deposited after an induction period. CdS does not deposit by itself under these acidic conditions. When subjected to cathodic polarization (0.65 V vs. S.C.E.), CdS preferentially deposited due to local increase in pH at the cathode (by water electrolysis), while eventually, as the Cd concentration dropped, ZnS formation became more favorable. Depth analysis of the films showed that Zn was formed preferentially at the surface while CdS formed preferentially near the substrate. This order is expected, considering the lower-solubility product of CdS. The second example, which was a pure CD process, produced ZnO on top of CdS [58]. The principle is based on the facts that CdS deposits much more readily than does ZnS (see Sec. 4.4.1) and that ZnO (or Zn(OH)2, which readily converts to ZnO) tends to deposit readily, more so than ZnS unless under conditions of high active sulphur concentration (whether sulphide ion or sulphur-containing complex) and low pH. The solution contained Cd and Zn ions (the latter in excess) complexed with ethanolamine and ammonia (therefore at least fairly high pH) and thiourea. Cross-sectional microprobe analysis showed that the film contained the more readily deposited CdS at the substrate (ca. 0.3 m thick), covered with a 2m ZnO layer, which formed as the solution conditions (probably mainly the drop in Cd concentration) favored the ZnO deposition. The films were pale yellow, and optical spectroscopy showed two transitions—one at the CdS bandgap (ca. 2.6 eV, greater than the bulk bandgap, suggesting that the CdS crystal size was small—ca. 4–5 nm) and the other corresponding to ZnO at 3.2 eV. The films were photosensitive, about an order of magnitude more sensitive than ZnO deposited by itself. A different form of bilayer can be formed using topotactic exchange reactions. This type of exchange is well known, e.g., for the conversion of CdS into Cu2S by immersing in a hot CuCl solution, used in the past for fabricating CdS/Cu2S photovoltaic cells (see Sec. 9.1.2). It has been used more recently to convert CD films of one semiconductor into another, e.g., CdS and CdSe into Ag2S and Ag2Se [59] and SnS2 into Ag2S [60]. While these studies describe conversion of one semiconductor into another, it is clear that, if carried out in a controlled fashion, partial exchange can occur, leading to the expected formation of a shell of the exchanged semiconductor around a core of the original semiconductor for each individual crystal in the film (assuming the film to be at least somewhat porous, as it invariably will be). This process therefore can lead to films of core/shell nanocrystals.

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8.5 CONCLUSION It is fair to state that the understanding of deposition of ternary compounds lags behind that of binaries. A better understanding of the factors that control codeposition, as well as solid solution formation, is needed. However, it is also clear that there is scope for deposition of a wide range of compounds, not only ternaries, but quaternaries and even higher-multinary materials. Additionally, the scope for deposition of mixed-phase films, either as consecutive layers (as shown earlier) or as composites, is great, and this aspect of CD will undoubtedly be pursued.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

CS Brooks. Metal Finishing 88:21, 1990. IM Kolthoff, DR Moltzau. Chem. Rev. 17:293, 1935. NA Rudnev, GI Malofeyeva. Talanta 11:531, 1964. NA Rudnev, IV Melikhov, AM Tuzova. Zh. Anal. Khim. 28:635, 1973. KL Chopra, RC Kainthla, DK Pandya, AP Thakoor. In: Physics of Thin Films. Vol. 12, Academic Press, New York and London, 1982, p 167. G Contreras-Puente, O Vigil, M Ortega-Lopez, A Morales-Acevedo, J Vidal, ML Albor-Aguilera. Thin Solid Films 361:378, 2000. JM Doña, J Herrero. Thin Solid Films 268:5, 1995. NC Sharma, RC Kainthla, DK Pandya, KL Chopra. Thin Solid Films 60:55, 1979. T Nakazawa, S Kuranouchi, A Ashida, N Yamamoto. In: 12th EC Photovolt. Solar Energy Conf. Amsterdam, the Netherlands, 1994, p 601. GK Padam, GL Malhotra, SUM Rao. J. Appl. Phys. 63:770, 1988. SA Al Kuhaimi, Z Tulbah. J. Electrochem. Soc. 147:214, 2000. T Yamaguchi, Y Yamamoto, T Tanaka, Y Demizu, A Yoshida. Thin Solid Films 281–282:375, 1996. T Yamaguchi, Y Yamamoto, T Tanaka, A Yoshida. Thin Solid Films 344:516, 1999. KC Sharma, JC Garg. J. Phys. D: Appl. Phys. 23:1411, 1990. DS Sutrave, GS Shahane, VB Patil, LP Deshmukh. Mater. Chem. Phys. 65:298, 2000. ES Rittner, JH Schulman. J. Phys. Chem. 47:537, 1943. M Skyllas-Kazacos, JF McCann, R Arruzza. Appl. Surf. Sci. 22/23:1091, 1985. LP Deshmukh, KM Garadkar, DS Sutrave. Mater. Chem. Phys. 55:30, 1998. BB Nayak, HN Acharya, GB Mitra. Bull. Mater. Sci. 3:317, 1981. BB Nayak, HN Acharya. J. Mater. Sci. Lett. 4:651, 1985. GB Reddy, DK Pandya, KL Chopra. Sol. Energy Mater. 15:383, 1987. GA Kitaev, VF Markov, LN Maskaeva, LE Vasyunina, IV Shilova. Inorg. Mater. 26:202, 1990. GA Kitaev, LN Maskaeva, VF Markov, AY Kurkin, LE Vasyunina. Inorg. Mater. 25:1065, 1989. LN Maskaeva, VF Markov, GA Kitaev. Russ. J. Appl. Chem. 73:751, 2000. NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Mat. Res. Bull. 11:1109, 1976. NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Thin Solid Films 42:383, 1977.

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27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60.

NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Thin Solid Films 62:97, 1979. NC Sharma, DK Pandya, HK Sehgal, KL Chopra. Thin Solid Films 59:157, 1979. S Misra, HC Padhi. J. Appl. Phys. 75:4576, 1994. R Suarez, PK Nair. J. Solid State Chem. 123:296, 1996. LP Deshmukh, DS Sutrave, BM More, CB Rotti, KM Garadkar. Semicond. Devices 421, 1996. PK Nair, L Huang, MTS Nair, H Hu, EA Meyers, RA Zingaro. J. Mater. Res. 12:651, 1997. MTS Nair, Y Pena, J Campos, VM García, PK Nair. J. Electrochem. Soc. 145:2113, 1998. VM Markov, LN Maskaeva, LD Loshkareva, SN Uimin, GA Kitaev. Inorg. Mater. 33:555, 1997. N Parhi, BB Nayak, BS Acharya. Thin Solid Films 254:47, 1995. GK Padam, SUM Rao. Sol. Energy Mater. 13:297, 1986. D Cahen. Sol. Energy Mater. 15:225, 1987. GK Padam, SUM Rao. Sol. Energy Mater. 15:227, 1987. GK Padam, GL Malhotra, SUM Rao. Phys. Status Solidi (a): 109:K45, 1988. RN Bhattacharya. J. Electrochem. Soc. 130:2040, 1983. GK Padam. Mat. Res. Bull. 22:789, 1987. KR Murali. Thin Solid Films 167:L19, 1988. JC Garg, RP Sharma, KC Sharma. Thin Solid Films 164:269, 1988. PK Vidyadharan Pillai, KP Vijayakumar, PS Mukherjee. J. Mater. Sci. Lett. 13:1725, 1994. D Raviendra, JK Sharma. J. Appl. Phys. 58:838, 1985. D Hariskos, M Ruckh, U Ruhle, T Walter, HW Schock, J Hedstrom, L Stolt. Sol. Energy Mater. Sol. Cells 41–2:345, 1996. K Kushiya, T Nii, I Sugiyama, Y Sato, Y Inamori, H Takeshita. Jpn. J. Appl. Phys. 35:4383, 1996. RC Kainthla, DK Pandya, KL Chopra. J. Electrochem. Soc. 129:99, 1982. RS Mane, CD Lokhande. Thin Solid Films 304:56, 1997. GS Shahane, BM More, CB Rotti, LP Deshmukh. Mater. Chem. Phys. 47:263, 1997. GS Shahane, KM Garadkar, LP Deshmukh. Mater. Chem. Phys. 51:246, 1997. GS Shahane, LP Deshmukh. Mater. Chem. Phys. 70:112, 2001. GN Chaudhari, S Manorama, VJ Rao. Thin Solid Films 208:243, 1992. GN Chaudhari, SN Sardesai, SD Sathaye, VJ Rao. J. Mater Sci. 27:4647, 1992. YS Sarma, HN Acharya, NK Misra. Thin Solid Films 90:L43, 1982. AR Patil, VN Patil, PN Bhosale, LP Deshmukh. Mater. Chem. Phys. 65:266, 2000. K Yamaguchi, T Yoshida, T Sugiura, H Minoura. J. Mater. Res. 13:917, 1998. SJ Castillo, M Sotelo-Lerma, RA Zingaro, R Ramirez-Bon, FJ Espinoza-Beltran, R Guillemette, MA Dominguez. J. Phys. Chem. Solids 62:1069, 2001. CD Lokhande, KM Gadave. Mater. Chem. Phys. 36:119, 1993. CD Lokhande, VV Bhad, SS Dhumure. J. Phys. D: Appl. Phys. 25:315, 1992.

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9 Photovoltaic and Photoelectrochemical Properties

A large number of studies on CD have been driven by two related potential uses: photoelectrochemical (PEC) cells, mostly the earlier studies, and, more recently, photovoltaic (PV) cells. This chapter is devoted to these two topics where CD films have been used.

9.1 PHOTOVOLTAIC CELLS 9.1.1 Introduction Since almost all thin-film (CdTe- and CuInSe2-type) cells today use CD films (mostly CdS), there is no attempt here to be comprehensive regarding the literature. Rather, studies that emphasize the CD film itself are discussed. Before discussing the role of the CD layer specifically, a brief overview of the relevant cells will be given. There are three main thin-film PV cells under development at present: amorphous Si, CdTe/CdS, and CI(G)S/CdS [CI(G)S refers to copper indium (gallium) selenide]. Of these, the last two are polycrystalline (as opposed to amorphous), and both normally employ CD CdS. Crystalline Si cells are not thin films, being at least tens and usually hundreds of microns in thickness, compared to the few microns of active thickness of the thin-film cells. Schematic diagrams of the CdTe and CI(G)S cells are shown in Figure 9.1. The main difference in their construction is that the CdTe cell is a superstrate

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(backwall) cell (illuminated through the conducting glass substrate), while the CI(G)S cell is a substrate (frontwall) cell (illuminated through the front surface). The CI(G)S cell is a development of the original CuInSe2 /CdS (CIS/CdS) cell, with Ga added to increase the bandgap. Pure CuGaSe2 /CdS has also been investigated, although considerably less than CI(G)S, due to its (at present) considerably lower conversion efficiency. In the following, CIS will be used where CuInSe2 is intended, while CI(G)S refers to CuInxGa1xSe2, with 1  x typically 0.2  0.1. Some studies have also been made on CuInS2, which has a higher bandgap than CIS and in principle should give a better cell (in practice it is inferior, although somewhat better than CGS). In both cells, the absorber layer (CdTe or CI(G)S) is a few microns thick, while the CD CdS (or other CD layer) is typically 50–100 nm thick. The CD layer is often called the “buffer” layer, a term that serves to show the lack of understanding of its role. Nominally, the CdS is the n-type part of the p-n junction. The basic mode of action of a p-n PV cell is shown in Figure 9.2. The short-circuit current (denoted in this chapter as ISC) is the current flowing in the illuminated cell when the two sides (terminals) of the cell are shorted (full band bending, left figure) while the open-circuit voltage (denoted as VOC) is the voltage generated between the terminals at open circuit when no external current flows (right figure). The absorber CdTe and CI(G)S are always p-type in these cells; n-type absorbers have been little investigated, mainly because suitable high-bandgap p-type materials are not readily available. A high bandgap of the buffer layer is necessary because light passes through this layer on the way to the absorber, and some of the light absorbed in the buffer layer is lost for current generation. The superiority of CD CdS over evaporated CdS in both types of cell suggests that something other than loss of photons absorbed in the buffer layer is involved. The factors believed to contribute to this superiority, in particular the effects of the deposition solution on the absorber, are discussed later.

FIG. 9.1 Schematic cross section of CI(G)S/CdS and CdTe/CdS PV cells.

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FIG. 9.2 Band diagrams of p-n cell in the dark (or under illumination at short circuit) and under illumination at open circuit.

The foregoing explanation of the operation of these cells, while very basic, will be almost sufficient for our purposes (there are many sources explaining the mode of action of PV cells in more detail). One other process, which plays an important role in PV cells in general, should be described: electron–hole recombination. This is central to PV cell operation. Photogenerated electrons and holes are ideally separated at the p-n junction and flow in opposite directions to give an external current. However, there are many pitfalls awaiting these charges on their way to the terminals where at least one of them can be extracted. These pitfalls, which cause the electrons and holes to recombine before external current flows, are various forms of recombination centers. They can occur at the interface between the p- and n-type semiconductors, at grain boundaries, or in the bulk of the semiconductor crystals. A major part of PV cell research is devoted to minimizing such recombination centers.

9.1.2 CuxS/CdS Cells Before describing studies on the CIS and CdTe cells, there are two CD-related papers on the Cu2S/CdS cell, which was intensively investigated around 20 years ago and was eventually abandoned because of perceived insoluble stability issues, a perception that, it should be noted, while widely held, is not undisputed. Should this cell make a comeback, CD is likely to be a method that will be considered for either of the two semiconductors or even for both. One study utilized CD CdS, built up from several layers (probably a total thickness of the order of a micron) and annealed [1]. The CuxS layer was formed by the usual (for this type of cell) topotactic reaction between a CuCl solution and the CdS substrate. While the cell performance varied over a wide range, depend-

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ing on the CD process, the maximum efficiency obtained was 0.13%; all parameters of the cell were very poor. In the other study, CuxS was chemically deposited on (presumably evaporated) CdS films from a triethanolamine/ammonia/thiourea bath (see Chap. 6, copper sulphides) [2]. Very low currents and poor fill factor were obtained, although the VOC was reasonable (ca. 0.5 V), with an efficiency of ca. 0.5%. The stoichiometry of the Cu-S was not ideal for PV cell use, although this could be varied to an extent by electrochemical treatment. In view of the sensitivity of the Cu2S/CdS cell to the nature and phase of the Cu-S, it is likely that much better performance can be obtained if an effort is made to do so.

9.1.3 CdTe/CdS Cells While CD CdS is commonly used for CdTe-absorber cells, there is relatively little work that emphasizes the CdS film. A 13.4% efficient cell fabricated by close space sublimation of CdTe on CD CdS was reported in 1991 [3], followed by a 14.5% cell a year later by the same group [4]. The CdS thickness was between 50 and 150 nm. The cells were illuminated through the tin oxide/glass, which was used as the substrate for the CdS deposition, and this geometry has been used ever since for these cells. The most comprehensive study of the effect of the CdS deposition parameters on the resulting CdTe/CdS appears to have been made for electrodeposited CdTe [5]. The most simple variable is CdS film thickness. Clearly a minimum thickness is required for junction formation and to prevent shunting from the CdTe through the CdS to the substrate (usually conducting glass). On the other hand, an increase in this thickness leads to a decrease in ISC due to light absorption in the CdS, which is clearly seen as a decrease in the short-wavelength response of the cell. The optimum CdS thickness was found to be ca. 70 nm, although good cells were also made with more than twice this thickness. Other deposition parameters affected mainly VOC and fill factor rather than ISC. These included an increase in thiourea concentration and the use of buffered (ammonium ion, lower pH) solutions; both these factors resulted in higher S:Cd ratios, therefore more stoichiometric films (CD CdS films are often Cd-rich; this does not necessarily mean n-type doped but is more likely due to the presence of other Cd species, e.g., Cd(OH)2). The use of chloride as an anion in solution rather than sulphate also gave better cells. It was believed that all these factors influenced the nature of both the CdS and the CdTe films after annealing the cell. Specifically, it was thought that small grain size and high defect density (the CdS was believed to be polytype with a large density of stacking faults [6]) in the cubic CdS film was beneficial for the resulting recrystallization process and for intermixing between the CdS and CdTe during the recrystallization and phase change (to

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hexagonal phase) of the CdS during the annealing step. In this respect, it was noted that if the CdS was recrystallized prior to CdTe deposition (and the cell then reannealed as usual), the resulting cells were very inferior in all output parameters to normal cells. The VOC of electrodeposited CdTe on CD CdS cells was studied as a function of the CdS deposition parameters [7]. While there were a number of different variables involved, it was clear that conditions leading to thicker films (ca. 180 nm), such as lower pH, high thiourea:Cd ratio (or possibly higher thiourea concentration), and repeated deposition of the films, resulted in the highest voltage (ca. 0.7 V). The short-wavelength response was poorer for thicker films, not only due to absorption in the CdS, but also at wavelengths longer than the CdS absorption onset, suggesting recombination at interface states. A clear effect of CD CdS on heterojunction formation has been shown for CdTe that was electrodeposited onto CdS films on single crystal (1¯1¯1¯) InP [8]. CdTe electrodeposited directly onto the (1¯1¯1¯) InP shows some degree of epitaxy but also considerable polycrystallinity (the latter not surprising, considering the 9.5% lattice mismatch between the two materials). If, however, a thin (20–30 nm) CdS film was chemically deposited onto the same InP surface, the epitaxy of CdTe electrodeposited onto the CdS/InP was found to be very good. The CdS was largely epitaxial with the InP (the lattice parameters of CdS and InP are very close), with ca. 15% polycrystallinity. Interestingly, the CdTe deposited on the CdS exhibited an even higher degree of epitaxy than that of the CdS itself, showing that the small but appreciable amount of polycrystalline CdS did not substantially degrade the epitaxy of the CdTe. It was suggested that the improvement in epitaxy due to the CdS was caused by a graded interface. Similarly, an XRD comparison of CdTe electrodeposited onto SnO2/glass (activated by a cathodic treatment) or onto CdS chemically deposited on SnO2/glass showed better crystallinity (narrower XRD peak) for the CdTe deposited on the CdS and also better texture (only the (111) reflection was seen, while a small additional (220) peak was evident for the CdTe deposited directly on the SnO2) [9].

9.1.4 CdS/CI(G)S Cells 9.1.4.1

General Considerations

A considerably greater body of work with more emphasis on the CD buffer layer exists for this cell. Much of this involves the specific effects of the CD process at the interface, and this will be discussed in a later section. An experimental measurement of the band lineup between the CdS and solar-grade polycrystalline CIGS has been made using contact potential difference (Kelvin probe) measurements in air [10]. This lineup is shown in Figure 9.3. In particular, it shows that no spike was found in the conduction band. The presence of such a spike (believed to occur from previous studies either on single crystals

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FIG. 9.3 Band offset between CIGS and CdS. (After Ref. 10.)

or on polycrystalline films in vacuum) would constitute a barrier to electron flow over the interface. From the practical point of view, reuse of the deposition solution after filtering out the precipitated CdS and addition of fresh reagent was shown to have no effect on the device properties of CdS/CIGS cells (see Sec. 4.1.6.10 for more details of the deposition) [11].

9.1.4.2 Chemical Reaction and Diffusion at the CdS/CI(G)S Interface There are a number of studies on the effects of the CD process on the surface of the CIS or CIGS. In many of these studies, the absorber surface was treated with partial CD solutions, in particular, ammonia or ammonia  Cd2 [12–19]. There are several reactions that occur during these treatments, and these will be discussed in general before specific results from the different studies are treated. Aqueous ammonia removes surface oxides from the CI(G)S, in particular, indium (possibly due to the tendency of the CIS surface to be enriched in In-O). In this respect, it can be considered as an etchant, although the etching is limited to the near surface region and does not continue (although in one case, it was found that, if oxygen is present, etching can continue, presumably due to continual oxidation of the freshly exposed surface [13]).

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Cd2 dissolved in ammonia has a particularly strong effect on the absorber surface. The effect of ammonia  Cd with respect to the In and Cu concentrations (decrease in Cu) is opposite to that of ammonia by itself (decrease of In), and it was noted that the surface reaction with all components present in the CD bath would depend on the relative kinetics of the various partial reactions. Probably more important, however, is that Cd rapidly substitutes for Cu for between several nanometers and 20 nm into the absorber. This has been found to occur on single crystals of CIS [15,16] as well as on polycrystalline films, and therefore is not simply grain boundary diffusion, as might at first be suspected. In contrast to these results, one study found no Cd substitution for Cu in CIS single crystals and related this to the supposed absence of a Cu-deficient layer in the single crystals, compared to films [19]. It is not clear whether Cd indiffusion promotes Cu outdiffusion or vice versa. It should be noted that Cu is readily complexed by the ammonia at the surface and therefore is easily removed. Also, the ionic radii of Cd 2 and Cu are almost identical, facilitating exchange. It is believed that the indiffused Cd type converts CIS to n-type and that the junction is a buried n-p one rather than located at the CdS/CIS interface. There is also the question of how abrupt the junction is; the Cu-poor surface region has been considered a separate phase— CuIn3Se5—known as an ordered vacancy compound (OVC). Exchange of Cu and Cd was shown to be easier for single-crystal CuIn3Se5 than for CIS [16]. The effect of this Cd/NH3 treatment on the PV properties are very marked. While cells fabricated without a buffer layer [ZnO sputtered directly on the CI(G)S] are very poor, with all parameters very low, the same cells, but subjected to the Cd/NH3 treatment before ZnO deposition, are very much better, and in fact the efficiencies are only a little lower than CD CdS cells, due to lower VOC (ISC is actually often higher due to the better blue response in the absence of CdS). This is a particularly important result since it shows that the main role of the buffer layer is not related to the specific properties of the CdS itself, but rather to nearsurface modification of the CI(G)S. Substitution of Zn for Cd in the Cd/NH3 treatment gave comparable results [15]. This is in contrast to the use of CD ZnS, which was inferior to that of CdS, although not necessarily by much (see Section 9.1.4.5). However, the presence of the CD CdS is still required in order to obtain optimal efficiency, and therefore the CdS itself does play some role, possibly to prevent sputter damage to the absorber (although Cd/NH3-treated absorbers, which then have evaporated CdS deposited, are still not as good as CD CdS cells). The conformal coverage of the irregular absorber surface by the CD CdS is another factor often invoked. Incomplete coverage (e.g., pinholes) could lead to some shunting between the ZnO and CI(G)S, either due to the lack of CdS at the pinholes and/or because of sputter damage at the pinholes. Also in connection with the nature of the CdS itself, studies on the effects of CdS thickness and impurity content on CdS/CIGS cell parameters have been made [20]. The impurity content

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was controlled by the concentration of thiourea in the bath (see Sect. 4.1.7 for details) while maintaining the CdS thickness constant. After a very thin CdS film had been deposited, the main effect of increasing both thickness and impurity concentration was to increase the VOC. This suggests that these effects are related to the “bulk” properties of the CdS rather than to effects of the different solutions on the absorber surface. It should be kept in mind, however, that the predeposited ultrathin layer either may not have totally covered the CIGS or may have been porous enough to allow contact between the second deposition solution and the CIGS. In connection with impurity considerations, a general property of CD CdS is its relatively large oxygen content. While it is not clear whether this is of any importance for its use in PV cells, it is often studied in this respect. The subject of oxygen impurities in CdS is treated in some detail in Section 4.1.7. Photoluminescence (PL) measurements to monitor the changes due to the Cd/NH3 treatment of CIGS have been carried out by two groups. Both show major, though very different, effects of the treatment on the PL spectra. In one case [14], two shallow subbandgap peaks, attributed to donor–acceptor pair recombination, were found in the nontreated CIGS. After treatment, the higher-energy peak was quenched and a strong new, lower-energy peak appeared. This could be interpreted as the removal of states that caused interface recombination, although the nature of the strong lower-energy luminescence was not understood. Notably, similar behavior was obtained when CdS was deposited, providing further evidence for the dominating role of the Cd/NH3 treatment in the CdS buffer layer formation. In the second study, the bandgap luminescence was found to increase 15 times after the Cd/NH3 treatment [16]. This could be interpreted as a passivation of recombination centers. The luminescence from CD CdS/CIGS was stronger (by ca. two times) than from CdS evaporated on Cd/NH3-treated CIGS and much stronger (nearly 20 times) than from CdS evaporated on nontreated CIGS. From these results, it was believed that both the Cd/NH3 treatment and the CdS deposition were important, although, again, it appears that the major effect is from the Cd/NH3 treatment. On an In-rich CIS surface (one from which good-quality cells could be made), the change in surface composition was followed upon deposition of CdS from a complete bath [12]. An important observation was that initially, while Cd was found at the surface, the S concentration was low, and stoichiometric CdS only formed later in the deposition. It was stressed, however, that since the deposition was begun at room temperature and then gradually heated to 60°C, it might be that the Cd exchange process occurred preferentially at low temperature, where CdS deposition was still very slow, while the thiourea decomposition, which was strongly temperature dependent, occurred more readily as the deposition temperature increased. The initial Cd species were believed to act as nucleation centers for CdS formation. This heating regime also slowed down precipitation in solution and resulted in a lower terminal thickness (ca. 75 nm compared to 100 nm

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for a preheated solution [13]), which would make the deposition time to obtain the required thickness (ca. 50 nm) less critical. An investigation of CdS deposition on CIGS (In  Ga rich—good PV cell quality) substrates was carried out using two different baths: a “standard” bath and a low-thiourea (3 mM) one [17]. Cd was preferentially incorporated into (onto) the substrate compared to S in the initial stages of the deposition, in agreement with the previous study. In contrast to the previous study was the conclusion that no major preferential removal of any of the substrate constituents occurred (In, Ga, and Se native oxides were removed). In addition, no clear evidence of compound formation between the Cd and substrate was found, although it is clear that Cd was incorporated into the substrate. Although the low-thiourea bath resulted in CdS with lower impurity (mainly N compounds) levels and more Cd incorporation into the CIGS, no improvement in cell performance resulted from such baths. While the nature of the Cd incorporation in the early stages of the deposition could not be unambiguously defined, it was suggested that this Cd was in the form of Cd(OH)2, which converted to CdS, either as the thiourea slowly hydrolyzed to sulphide or as decomposition of a Cd(OH)2–thiourea complex occurred, both of which are very temperature dependent. An explanation for the beneficial effect of the CD process on CI(G)S cells was suggested based on the known effects of oxygen treatment on these cells [21]. Annealing in oxygen removes Se vacancies, which in turn decreases recombination at grain boundaries, surfaces and at the CdS/CI(G)S interface. In opposition to this beneficial effect, the oxygen treatment also has been shown to reduce positive charge, and therefore band bending at the interface, and to increase Cu diffusion into the bulk of the CI(G)S, thus reducing the acceptor doping; both these effects are detrimental to cell performance. The CD deposition was believed to restore the positive charge to the interface (with which the deposition solution was in contact) but not to the grain boundaries (where it did not reach) by creation of Cd on Cu vacancies (CdCu) and possibly also removal of oxygen on Se vacancies (OSe). This removes the detrimental effect of the oxygen treatment at the interface but not the beneficial effect at the grain boundaries.

9.1.4.3

Epitaxy of CdS on CIS

The surface cleaning of the CIS also affected the mode of deposition of the CdS. The CdS was found to grow to a greater or lesser extent of epitaxy on single-crystal (heteroepitaxial layer) CIS [22]. Very good epitaxy of cubic CdS was found for cyanide-treated CIS; somewhat lower epitaxy was found for ammonia-treated surfaces and poorer epitaxy obtained for untreated surfaces that contained considerable oxides. Additionally, the epitaxy was only obtained at higher deposition temperatures (70°C); at lower temperatures, the growth was polycrystalline. Epitaxial growth of cubic CdS {111} on the CIGS {112} was also found to occur using a room-temperature bath, with gradual heating to 80°C 18,19]. Again,

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the epitaxy was explained by ammonia cleaning of the CIGS followed by ion-byion growth.

9.1.4.4

CuGaSe2-based cells

CuGaSe2 (CGS) has also been studied as a PV material, although efficiencies of cells based on this semiconductor are, at present, much lower than those using CIS or CIGS. There is one study on the specific interaction between CD CdS and CGS [23]. Two different CD baths were used to deposit CdS in this investigation: one at 60°C and the other at 80°C (the latter with a higher ammonia concentration to slow down the deposition). Several pronounced differences were found between the two baths, in spite of the relatively small difference between them. For Cu-rich CGS, Cu-S inclusions in the CdS were formed in the high temperature bath, due to interaction between Cu in Cu-rich CGS and CdS, but not in the low-temperature bath. Such inclusions could lead to shunting. For Ga,In-rich films (which gave better cells), the 80°C deposition resulted in Se being found in the CdS layer. Additionally, the higher-temperature CdS was less defected and formed a less defected interface with the CGS. The 80°C deposition gave better PV properties for Ga-rich films (up to 9.3% efficiency). The 60°C deposition, however, was better for the (poorer cell quality) Cu-rich films, which could be explained by shunting through the Cu-S inclusions in the 80°C Cu-rich CGS devices. This investigation clearly shows the necessity for optimization of the CD process, not only for every specific absorber material, but even for different types of the same absorber.

9.1.4.5

Cd-Free Buffer Layers

While most of the reported studies on CD buffer layers deal with CdS, there have been a number of attempts to chemically deposit other materials. There are several reasons for this. One is the desire, particularly prominent in Europe, to find a Cd-free cell, for obvious environmental reasons (this clearly is relevant for the CIS and not for the CdTe cell). Another reason is that part of the light absorbed in the CdS is lost for current generation (a point that does not seem to have been rigorously investigated but that probably is due to a high recombination rate for holes photogenerated in the CdS). A higher-bandgap material will therefore allow more short-wavelength photons to reach the CIS absorber and generate photocurrent. Since the CD process is believed to be largely responsible for the beneficial properties of the CdS films, it has been anticipated that other CD layers would behave likewise. Other possible factors are the band lineup between CIS and CD layer and (photo)conductivity (or maybe better, resistivity, of the buffer layer, since it appears that a high resistivity is beneficial). ZnS. ZnS, with a high bandgap of 3.7 eV, is an obvious choice for alternate buffer layers. ZnS films, prepared using thioacetamide from an ammonia-free

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bath (see Sec. 4.4.2 for a description of the films) were deposited onto CIS films [24]. The higher short-circuit current expected for the higher-bandgap ZnS was obtained, although the overall efficiencies were lower than for CdS deposited on similar substrates, due to lower photovoltages. Efficiencies as high as 9% were obtained. Most ZnS baths probably give films containing some hydroxide (at higher temperature, maybe oxide), as discussed in Section 4.4.1. The previous bath may be an exception, since it was presumably carried out at relatively low pH. Using Zn(O,OH,S), efficiencies as high as 12.8% were obtained on CIGS substrates [25]. This efficiency was obtained after ca, 60-mn illumination and reversibly decreased when kept in the dark. Spectral response measurements showed the expected increase in short-wavelength response. Such an efficiency can be compared with the best reported value over 16% obtained at that time; however, since the highest reported values are considerably higher than those more routinely obtained, then the value of 12.8%, while lower than that obtainable using CdS, is not too far below it. Ideally, such experiments should be compared with state-of-theart CdS deposited on (as closely as possible) identical substrates. A somewhat more recent study reported 15.1% for ZnS/CIGS (compare with 17.0% state-of-the art CdS on the same CIGS substrates) [26]. This study showed the differences in optimization of the cell, depending on whether CdS or ZnS was used. Thus, while a 40 nm-thick ZnS layer led to a large improvement in all cell parameters, as was also the case for CdS, in contrast to CdS, increasing thickness of the ZnS to 90 nm reduced the cell performance, giving in particular a very low fill factor. This was explained by the higher resistivity of the ZnS (by two orders of magnitude), compared to CdS. In addition, while a high-resistivity sputtered ZnO layer was deposited on the CdS prior to the conducting ZnO:A1 to obtain maximum VOC, such a layer degraded the ZnS-buffer cell. Since this ZnO was sputtered in an oxygen-containing atmosphere, and no such degradation occurred for the conducting ZnO sputtered in an oxygen-free environment, it was believed that the degradation was caused by negative oxygen ions or energetic neutral particles. Finally, while the short-wavelength response of the ZnS cell was better than that of the CdS cell, this improvement did not compensate for the lower quantum efficiency at longer wavelengths for the ZnS cell. ZnSe. ZnSe, with a bandgap of ca. 2.7 eV, is another obvious substitute for CdSe. As with ZnS, there is a tendency for CD ZnSe to contain hydroxide. Zn(Se,OH) deposited from a selenourea bath was deposited on CIGS [27] (see Sec. 4.5.2 for details of the Zn-Se films). Efficiencies up to 13.7% (highest literature value 17% using CdS) were found. Spectral response measurements showed the expected improvement in short-wavelength response. The optimal thickness of the CD layer was 7–8 nm; a layer ca. twice that thickness resulted in a drop in efficiency (to 10.4%), mainly due to a drop in fill factor, probably be-

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FIG. 9.4 Band diagram of the CIGS/Zn(Se, OH)/ZnO cell. (After Ref. 27.)

cause of the higher resistivity of the thicker film. From XPS measurements, the band alignment could be estimated; this is shown in Figure 9.4. Tunneling of electrons through the large barrier at the junction between the CIGS and Zn(Se,OH) conduction bands was presumed to occur readily for the thin films but much less so for the thicker ones. ZnSe, deposited by the same method, was also used as a buffer layer for CuInS2 cells [28]. Higher currents and voltages but lower fill factor were obtained, compared to CdS, with a slightly lower overall efficiency. The band lineup for the ZnSe/CuInS2 junction was also measured by XPS for this system. ZnO. ZnO, which is normally used as the conducting window layer on the CIS-type cell, has also been used as a buffer layer for CuInS2 cells by annealing CD Zn(OH)2 (deposited from Zn2/ammonia solutions) [29] (see Sec. 7.2.18, ZnO). The efficiency was much lower (3.8%) compared with comparable CdS buffer layers (8.6%). The difference was due to much lower open-circuit voltage and resulting lower fill factor; the photocurrents were similar using both buffer layers. The increase in photocurrent at short (500 nm) wavelengths due to the higher-bandgap ZnO was offset by lower photocurrents over the rest of the active

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spectrum. While the thickness of the ZnO was not given, it was noted that the In XPS signal was still visible after the ZnO deposition. Comparing this with the previous study by the same group using Zn(Se,OH), where the In peak disappeared when the CD layer was 7 nm and was already weak at ca. 4 nm, suggests either that the ZnO layer was thinner than optimal or that the coverage was not homogeneous. In view of the previous experiments on CI(G)S, which showed that a Zn2/ammonia treatment was as effective as a Cd2/NH3 treatment and both not much less effective than optimally deposited CdS, the Zn2/NH3 bath used here was considerably inferior to CdS buffer layers, even assuming an overly thin film deposited. Whether this difference is due to the different substrate (CuInS2 instead of CI(G)S), to the annealing treatment, or to some other reason remains to be investigated. In(S,OH). Various compounds of In have been used, with some success, for buffer layers. In(OH)3 was grown on CIS (In-rich) films from a solution of InCl3 with thiourea (which possibly acted to gradually increase pH rather than as a source of S) [30]. In spite of the higher blue response compared to a control CdS cell [In(OH)3 is colorless as a film], the red response was poorer, leading to a somewhat reduced overall photocurrent. The fill factor was also less. A best efficiency of 9.5% was obtained, compared to 11.9% for the control using the same batch of substrates. In the same paper, the deposition of In2S3 [based on later studies, this may have been In(S,OH)] from a thioacetamide bath at a pH between 1 and 2 was described, but with considerably poorer results (both photocurrent and fill factor were much lower). Using essentially the same method for depositing In(S,OH) on CuInS2 films, the same group found slightly better performance (11.4%) with this buffer layer than for CdS (10.8%), due to increase in photocurrent and also in open-circuit voltage [31]. Since CuInS2 has a higher bandgap (1.5 eV) than CIS or CIGS, the fractional increase in photocurrent due to the improved blue response is larger. These same In(S,OH) films were also investigated on CIGS substrates [32]. XPS measurements (on CuGaSe2 used to prevent interference by In from the CIGS) showed that the deposited film was not a mixture of In2S3 and In(OH)3, as might be reasonably expected, but was a compound containing In, O, and S. Higher thioacetamide concentrations resulted in better device performance, while the In concentration was not found to influence the performance in any reproducible way. The completed cells (with ZnO window layer) required an anneal (2 min at 200°C) and light soak; the performance without this anneal was much poorer. The best cell gave an efficiency of 15.7%. In comparison with CdS buffers, the VOC increased, the ISC decreased, and there was no change in the fill factor. Overall, the efficiencies were only slightly less than for cells using CdS. For CIGS substrates deposited on Corning “Pyrex-type” glass (good cells are usually deposited on soda-lime glass and are beneficially affected by Na diffusion through the film), the In(S,OH) cells

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were actually better than the CdS ones. In fact, in contrast to CdS cells, the efficiency of the In(S,OH) ones were independent of the type of glass used for the substrate. Capacitance–voltage measurements indicated a 10 times increase in acceptor concentration of the absorber layer and narrowing of the space charge layer, compared to CdS. This explained the decrease observed in the red response (the blue response was better) and the increase in the VOC, and also the independence on the glass, since acceptor doping occurred by the CD process rather than by Na diffusion. Since the CD process did not change the bulk of the CIGS, it was assumed that the doping changes were due to surface (and grain boundary) effects. A list of cells made on CIS, CuInS2, and different CIGS substrates using In(OH,S) buffer layers is given by Hariskos et al. [32]. They note that the fill factor of these cells drops with time when the cells are kept in the dark but that illumination (by light absorbed by the In(OH,S) film) reverses the degradation. This behavior suggests that the degradation is due to adsorption of some species (oxygen, water?) that reacts with photogenerated electrons and/or holes. Sn(O,S). In the same paper, films of SnO2 and Sn(O,S)2—the latter apparently a mixture of oxide and sulphide—were used as buffer layers on CIS and CIGS but with all cell parameters lower than when using CdS and efficiencies lower than those obtained using In(OH,S)

9.1.5 Other Cells and Related Studies 9.1.5.1

Other Heterojunctions and Devices

There are several studies on heterojunctions, other than those already described, formed with at least one CD semiconductor. These will be described briefly here. The earliest of these studies was on PbS. PbS can have either p- or n-type conductivity, although CD PbS is usually p-type. Based on the belief that the ptype conductivity may be due to alkali metal cations from the deposition solution, an alkali metal—free deposition, using lead acetate, thiourea, and hydrazine hydrate was used [33]. While initially n-type, the film converted to p-type in air. Attempts to stabilize the p-type material by adding trivalent cations to the deposition solution were unsuccessful. However, deposition of the PbS on a trivalent metal, such as Al, did stabilize the n-PbS, at least for a time. In this way, p-n junctions were made (the PbS close to the trivalent metal was n-type, while the rest of the film was p-type). Photovoltages up to 100 mV were obtained from these junctions at room temperature and almost 300 mV at low temperatures (90 K). PbS was also deposited on single-crystal n-and p-type Ge [34]. The PbS was epitaxial (111) with the (111) Ge (Ge has a 5% smaller lattice spacing than PbS). A photovoltage was measured from the junctions. The photoresponse extended to 1.75 m for the junctions on p-type or intrinsic Ge and to 3.35 m on n-type Ge. The difference could not be explained, although it can be noted that these values

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correspond approximately to the bandgaps of Ge and PbS, respectively. Hot probe measurements indicated that the PbS layers were n-type; however, it was mentioned that this measurement may be affected by the junctions and that PbS chemically deposited onto glass by the same method normally gives p-type layers. Investigations of junctions formed by CD of PbS on n-Si [35] and n-Sb2S3 on p-Si [36] and on p-Ge [37] have been made. In particular, the Sb2S3 junctions were found to be much more PV active if a small amount of silicotungstic acid (STA) was added to the deposition solution. Conversion efficiencies of 5.2% on Si and ca. 4% on Ge could be calculated from the photocurrent–voltage characteristics. The STA resulted in formation of WO3 in the Sb2S3 film, and it was believed that this may, at least in part, be responsible for the improvement. Of particular note was the relatively large open-circuit voltage (nearly 0.7 V) obtained from the junction with Ge; this value is almost as large as the Ge bandgap. It is tempting to wonder if this junction was not closer to a Schottky (metal–semiconductor) junction, where the Ge behaves as a metal, in which case, the maximum photovoltage is limited by the Sb2S3 bandgap (ca. 1.7 eV). The p-Ge was highly doped (4  1018 cm3), although less than might be expected if it were to behave as a metal. CuInSe2, deposited by CD (see Chap. 8), has been screened for photovoltaic activity. In one study, CdS was evaporated on CIS that was chemically deposited onto conducting glass [38], while in the other the CIS was chemically deposited onto single-crystal Si [39]. The cells gave low activity, although the CdS/CIS cell gave a short-current photocurrent of 4 mAcm2 (AM1 illumination), which is quite appreciable, if still low, and suggests further studies in this direction might be fruitful, in particular also using CD CdS. CdO, a degenerate n-type semiconductor, was chemically deposited on single-crystal p-type Si [40]. The junction showed clear diode behavior, and, although no photovoltaic effect was observed, photocurrent was generated under reverse bias. From the spectral response of the photocurrent, almost all of the current generation occurred in the Si. CdS/SnxS PV cells have been fabricated where the CdS was deposited by CD and the SnS deposited by a variant of CD where the substrate is dipped first in a solution of one of the ions and then in the other without rinsing in between, as would be the procedure for SILAR (see Sec. 2.11.1) [41]. While the cells showed very low conversion efficiencies, the main emphasis was on Ag-doping of the CdS in order to increase the conductivity and the effect of this doping on the PV cells. An increase in efficiency from 0.03% to 0.08%, mainly as a result of an increase in short-circuit current, was obtained by doping the CdS with Ag. The doping was carried out by an ion exchange process whereby the undoped CdS film was immersed in a solution containing Ag complexed by thiosulphate. A heterojunction between two different CD semiconductors—n-Ag2S deposited on top of p-PbS—was fabricated [42]. Some photoconductivity was found at wavelengths longer than that corresponding to the Ag2S bandgap (ca. 0.9 eV).

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However, at shorter wavelengths, where light could be absorbed in the Ag2S, negative photoconductivity (increase in resistivity with light) was found. This was explained by a combined Dember effect/electrostatic attraction of electrons by Ag ions. Metal–CD semiconductor Schottky junctions have been examined as solar cells. As for the p-n junctions described earlier, the addition of STA to both Sb2S3 and CdSe improved the cell parameters greatly [43]. AM1 efficiencies (for annealed films) of 7.2% and 5.5% were obtained using Pt contacts on annealed CdSe and Sb2S3, respectively. Again, the improvement due to the STA was attributed to the presence of WO3. In this study, it was suggested that the WO3 might introduce favorable interface states in the devices. Whatever the reason for the improvement, the strong effect of the STA treatment and its applicability to two different semiconductors (as well as to both solid-state and liquid junction cells) warrants further investigation. A similar study, using CD CdSe (with STA) deposited on poly(3methylthiophene), the latter electropolymerized (or the polythiophene electrodeposited onto CdSe), was carried out [44]. Poly(3-methylthiophene) can be prepared either p-semiconducting or doped to a metallic conductivity. While the undoped pn junctions gave poor photoresponse, the Schottky-type doped thiophene-CdSe junction gave conversion efficiencies of 2.7%, which were stable for at least 72 hr of illumination. CdSe deposited without STA resulted in lower efficiencies (1.3%). Thin-film transistors have been fabricated by depositing 50 nm of CdS onto SiO2-covered n Si and evaporating two A1 electrodes (source and drain) onto the CdS [45]. Similar devices were also made using CdS deposited on polyimide substrates with three (source, drain, and gate) evaporated metal electrodes and various sputtered insulator layers for the gate electrode.

9.1.5.2

Passivation Studies

Various sulphiding treatments have been known to passivate the surfaces of III–V semiconductors like GaAs and InP. Similar effects have been found with CD of very thin CdS films on InP surfaces. By deliberately oxidizing InP surfaces to produce In and P oxides, it was shown that a standard CdS solution (only with a higherthan-usual concentration of ammonia) removed these oxides and prevented oxide regrowth [46]. The same treatment also removed P vacancies deliberately introduced by annealing in H2, as seen by photoluminescence studies, presumed due to filling of P vacancies by S. Deposition of CdS (5–7 nm) onto InP both improved the C–V behavior and lowered the interface state density of SiO2/InP junctions by an order of magnitude. Further studies found that a pretreatment of the InP surface by a solution of NH4OH/thiourea improved the passivation but that CdS deposition (3–4 nm) was still important [47]. The NH4OH/thiourea treatment was believed to produce a stable In-S terminated InP surface [48]. This latter study also reported that the SiO2/InP junction quality, measured mainly by the interface state density, was maximum at a very low CdS thickness (ca. 1 nm) and that this quality degraded

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gradually as the CdS thickness increased, presumably due to interface states introduced by the thicker CdS. It appears that the NH4OH/thiourea treatment passivates the surface, while the very thin CdS protects it from reoxidation. This treatment was extended to various device structures based on ternary semiconductors (InA1As and InGaAs) with similar improvements due to removal of interface states [49]. These studies are clearly related to similar ones on CI(G)S. CD ZnSe has also been demonstrated to passivate surface states, 0.92 eV below the conduction band edge (measured by thermally stimulated exoelectron emission) on single crystal GaAs. This passivation resulted in bandgap luminescence from the originally non-luminescent GaAs [49a].

9.2 PHOTOELECTROCHEMICAL CELLS (PECs) 9.2.1 Introduction and Background As noted earlier, there are numerous studies on the photoelectrochemical (PEC) properties of CD films. Many, if not most, of these studies describe the preparation of the films and some PEC properties. In such cases, rather than describe each study separately, it is more useful and efficient to tabulate the results, providing important cell parameters together with the reference. Additional relevant information will be given separately for each individual reference. However, no attempt is made to cover all the individual studies in any detail, but rather to give enough information to allow the reader to decide whether it may be worthwhile to refer to the original reference. Some studies that treat the PEC properties of these films in a more fundamental way will be discussed separately. An important issue is whether the films have been annealed or not and under what conditions; annealed films usually give better performance (normally much better) than nonannealed ones.

FIG. 9.5 Schematic diagram of a PEC.

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We start with a brief introduction to PECs. Figure 9.5 presents a schematic diagram of a PEC showing a semiconductor film on a substrate—the photoelectrode—connected through an external meter and/or load to a second electrode (the counterelectrode). The two electrodes are immersed in an electrolyte, and the semiconductor film is exposed to illumination. If the substrate is transparent, conducting glass, the light can pass first through the substrate and then to the semiconductor (and the glass can also function as the window of the cell); this configuration is know as a backwall cell, in contrast to the normal frontwall cell (shown in Fig. 9.5), where light is incident directly on the nonsubstrate side of the semiconductor film. The following discussion assumes that the semiconductor crystal size is large enough so that charge transport is dominated by a space charge layer in the semiconductor. This is typically the case when chalcogenide films have been annealed at temperatures of ca. 400°C or more, where the crystal size is typically of the order of hundreds of nanometers. This assumption is usually not valid for asdeposited CD films or for those annealed at low temperatures (e.g., 250°C or less). The mode of operation of such films is treated in Section 9.2.2.2. Figure 9.6a shows the band diagram of a semiconductor–electrolyte junction (with an n-type semiconductor; most studies described here have used films that give n-type response). The band bending represents the situation in the dark at equilibrium (flat Fermi level) or under illumination at short circuit (i.e., the photoelectrode (or photoanode) is short-circuited to the counterelectrode). (We ignore complications of quasi-Fermi levels due to any nonequilibrium situation in the light; it does not change the simple picture for our purposes). Photogenerated electrons (in the conduction band) and holes (in the valence band) are spatially separated by the space charge layer and flow in opposite directions. The holes flow to the semiconductor/electrolyte junction and oxidize some species in the electrolyte. The electrons flow through the back (ohmic) contact to the external circuit. Electrons thus flow to the counterelectrode and reduce some electrolyte species. (For a p-type semiconductor, a photocathode, the direction of charge flow is opposite: Electrons flow to the semiconductor/electrolyte junction and holes to the back contact.) The resulting current flowing is the short-circuit current (ISC). For a regenerative PEC, the electrolyte species that are oxidized and reduced form a single redox couple. A common example is the polysulphide redox solution. The reactions occurring at the photoanode are: S2  2h → S;

S  S2 → S2 2

etc. up to approximately S2 (9.1) 5

while at the counterelectrode, the reverse occurs:  2 S2 2  2e → 2S

or

S  2e → S2

(9.2)

The result is no net change in the PEC (hence the term regenerative) and electricity is produced. The regenerative PEC has been the type predominantly studied using CD

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A

B FIG. 9.6 Band bending in a PEC (a) in the dark or under illumination at short circuit and (b) under illumination at open circuit.

films. However, there is another important type, where the anodic and cathodic reactions occur with different redox species. The holy grail for this type of cell has been to photoelectrolyze water to hydrogen and oxygen. While this goal has attained only limited success, the search has led to very decided success in other, related directions, usually connected with photo-oxidizing adsorbed layers of pollutants or bacteria on TiO2 [50]. If the photoanode is not connected to the counterelectrode, then current cannot flow, and instead the potential of the photoanode (the Fermi level) rises until balanced by recombination of the photogenerated charges (Fig. 9.6b). The difference between this potential under illumination and the original potential (more correctly, the potential of the counterelectrode, which ideally is equal to that of the nonilluminated photoelectrode) is the open-circuit voltage (VOC) of the cell.

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In practice, in order to generate electrical power, the cell must operate under conditions where both current and voltage are generated, as with the photovoltaic cells described earlier. This situation is shown in Figure 9.7, which gives the photocurrent/photovoltage characteristics of the cell. The maximum power (Pmax) is generated when the load is such that the product of the current (IP) and voltage (VP) is a maximum. The shape of the photocurrent/photovoltage characteristic, which determines Pmax, is quantified as the fill factor (FF), which is defined as the ratio between Pmax and the product of ISC and VOC, i.e., Pmax FF
 ISC  VOC

(9.3)

and is given either as a fraction or, commonly, as a percentage. An important feature of photoanodes is that the photogenerated holes, which are normally very strongly oxidizing, may oxidize the semiconductor instead of, or as well as, the electrolyte species. This phenomenon is known as photocorrosion. For the purposes of the limited explanation of PECs given here, it is enough to note that continuous photocorrosion will destroy the photoelectrode. (This need not necessarily occur if corrosion is confined to the semiconductor surface.) The substrate on which the semiconducting photoelectrode is deposited is important, not just as an ohmic contact to extract charge (usually electrons, since most photoelectrodes are photoanodes), but also because the substrate should be electrocatalytically as inactive as possible toward the electrolyte species. The reason for this is explained in Figure 9.8. The heavy solid line shows a typical pho-

FIG. 9.7 Current–voltage characteristic of a photoanode showing maximum power point (Pmax). The fill factor is given by the product VP  IP divided by the product of VOC and ISC.

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FIG. 9.8 The effect of exposed substrate on a PEC. The thick solid line gives the lightinduced current–voltage characteristic. The thin solid line gives the net current–voltage characteristic when the substrate has the electrocatalytic activity given by the broken line (fair electrocatalyst). If a substrate with poor electrocatalytic activity is used, there is little effect of the substrate on the photocharacteristics in the fourth quadrant.

toanode response; it is assumed that dark currents are negligible. We now consider the dark (or under illumination—there will be no difference) current–voltage characteristic of any exposed substrate. This is shown for a poor electrocatalyst and a better (still not good) electrocatalyst. The higher the electrocatalytic activity of an electrode, the greater the current at a given bias. For a photoanode, the current in the fourth quadrant is the most important (for a photocathode, it is that in the second quadrant). If the positive current of the (no dark current) photoanode and the negative current of the exposed substrate are now added, the effect of the poor electrocatalyst is negligible while that of the better electrocatalyst is to strongly reduce both the VOC (from VOC to VOC(1)) and the fill factor of the PEC. This is the reason that Ti, a poor electrocatalyst for most redox systems, is used so commonly as a substrate for photoanodes. Another substrate sometimes used is Cr (also a poor electrocatalyst in general), e.g., for Bi2S3 [51] and Sb2Se3 [52]. This is in addition to the fact that these metals tend to form satisfactory ohmic contacts to n-type semiconductors. Note that for nanocrystalline films where a space charge layer may not exist, the concept of ohmic or nonohmic contact is not necessarily the same as for bulk semiconductors (see Sec. 9.2.2.2).

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9.2.2 CdSe 9.2.2.1 Annealed Films CdSe has been the most extensively studied semiconductor for PEC purposes. This is due to its fairly favorable bandgap for solar cell use (1.73 eV), compared to the higher-bandgap CdS, although there are many, usually less detailed, PEC studies on CdS as well. Cell details are given in Table 9.1, specific comments follow. In Ref. 53, the main purpose of the study was to investigate PEC properties of CD CdSe films rather than to optimize the actual solar cell (e.g., CdSe annealing was carried out at the relatively low temperature of 280°C and the film thickness was only ca. 0.9 m), hence the cell parameters are not as high as they could be. The CdSe/polysulphide junction was characterized by a number of techniques. The effects of an surface layer of CdS, due to exchange of Se by S from the polysulphide, were considered. Preliminary PEC results in Ref. 54 were previously described [60]. This study was directed to optimization of the PEC parameters. The deposition was based on the ammonia/selenosulphate bath. The Ti substrates were treated with a suspension of Cd(OH)2 and the deposition carried out in a sealed tube to prevent loss of ammonia and thereby to improve reproducibility. Several layers (about five) were deposited to give a total thickness of ca. 2.5 m, which was found to be optimum. The first layer was annealed at 500°C in air (to improve film adherence) and the final film at 550°C, also in air. The annealed films were etched (9M HCl) and treated with a ZnCl2 solution. The maximum efficiency was obtained at 50 mWcm2 illumination (6.8%). At higher illumination intensities, the effi-

TABLE 9.1 PEC Parameters of CD CdSe Efficiency (%) Annealed 0.15 6.3 4.9a 4.4 11.7 Not annealed 0.1 a

Isc (mAcm2)

Voc (V)

FF

Electrolyte

Illumination (mWcm2)

Refs.

1.92 15.3 16.0 8.4 11.8

0.23 0.66 0.66 0.85 0.59

0.34 ca. 0.42 0.46 0.62 0.67

Polysulphide Polysulphide Polysulphide Polysulphide Selenosulphate

100 67 100 (AM1) 100 40

53 54 55 56 57, 58

0.06

0.1

Polysulphide

Not defined

59

0.28

Higher efficiency (5.5%) was found for lower solar illumination (42.5 mWcm2).

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ciency dropped, due to a sublinear increase in ISC with illumination at higher intensities. The effects of annealing and film thickness on the PEC properties were investigated in Ref. 55. The optimum annealing conditions were 470°C for one hour in air. ISC in particular was found to drop strongly both at lower and higher annealing temperatures. ISC and VOC, and therefore efficiency, increased strongly with film thickness up to ca. 1.5 m and then more gradually up to ca. 2 m, after which no further change was observed. Such thicknesses result from a number of successive depositions. Larger-area photoelectrodes of 18 cm2 were also made; the efficiency of these cells dropped to 3.2% compared with small cells, mainly due to a drop in fill factor and ISC. A triethanolamine/ammonia/selenosulphate bath was used in the experiments of Ref. 56. Three depositions were employed to give a final film thickness of 4 m. The films were annealed at 500°C in N2 and etched in dilute aqua regia. The efficiency at this stage in a polyselenide electrolyte was ca. 0.5%. This study concentrated on the effects of a treatment consisting of a 50% HCl etch, which resulted in a black matte surface (indicating a high surface roughness), followed by a dip in an acidified ZnCl2 solution (a Zn2 dip has been previously been used beneficially for CdSe photoelectrodes [61]). This surface treatment improved all cell parameters. Using Kelvin probe studies, it was shown that the surface potential (measured against a Pt vibrating probe in nitrogen ambient) changed by 0.31 V after the treatment, interpreted as an increase in band bending due to change in the surface charge. This change was the same as the increase in VOC of the cell. Analysis of the dark current–voltage characteristics of the cell showed a decrease in both the reverse saturation current and the ideality factor after the surface treatment. That Zn was present in the samples was verified by separate XPS measurements [62]. A spectral response study of these films was carried out, with an emphasis on subbandgap response, which might be related to surface states [63]. The surface treatment decreased the subbandgap response (although it should be remarked that the subbandgap response of the untreated film was unusually high, extending out to 1 m and with relatively high quantum yields for a subbandgap response). It also preferentially increased the short-wavelength response, interpreted as a reduction in surface recombination. The beneficial effect of the surface treatment was found to occur not just in polyselenide electrolyte, but also in a ferrocyanide electrolyte. This suggests that the effect of the Zn is not related to formation of ZnSe at the surface, as might be thought. On the other hand, a large part of the improvement in polyselenide electrolyte was due to the HCl matte etch, and it is not clear how much of the improvement in the ferrocyanide electrolyte was caused by this part of the treatment. In Ref. 57, a small concentration (
105 M) of silicotungstic acid (H4SiW12O40) was added to a triethanolamine/ammonia/selenosulphate bath,

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deposition carried out at 40°C and annealing in air at 430°C (the CdSe remained cubic, in contrast to the hexagonal form usually obtained after annealing at such temperatures). A large increase in the PEC performance was obtained, compared to photoelectrodes deposited without the silicotungstic acid. All cell parameters were increased, but the major effect was on ISC. The stability of the PEC also improved: It was stable (at 40 mW-cm2 illumination) for greater than 2800 C/cm2, although the selenosulphate electrolyte is unlikely to be stable in a PEC over the long term. A follow-up study of these films [58] showed the presence of WO3 in the films. From the dark current–voltage characteristics, the reverse saturation current, ideality factor, and series resistance all decreased as a result of the modification. Two possibilities were put forward to explain the effect of the silicotungstic modification. One was that formation of a CdSe/WO3 heterojunction occurred, improving the charge transfer at the semiconductor/electrolyte interface. Another possibility, if a true heterojunction exists, is reduction in majority carrier (electron) transfer to the electrolyte, as in a metal/insulator/semiconductor device. In another study, it was suggested that the improved PEC response was due to charge transfer catalysis by the W-containing groups adsorbed at the CdSe crystal surfaces [64]. It should also be noted that the silicotungstic acid modification also improved the performance of solid-state cells (see Sec. 9.1.5.1) where no charge transfer to an electrolyte is involved, suggesting (although not proving) that electrocatalysis is not the reason for the improved PEC behavior.

9.2.2.2 Nonannealed Films In Ref. 59, films 0.8 m thick were deposited from an ammonia/selenosulphate bath. Various configurations of PECs were studied. A study of the mode of operation of nanocrystalline CdSe photoelectrodes was carried out [65]. This was based on the expectation that nanocrystals, being (usually) much smaller than typical space charge layer widths, would not support such a space charge layer and therefore that some other mechanism for charge separation should be considered. It was originally reported that nanocrystalline CD CdS [66] and CdSe [67] photoelectrodes, which normally gave photocurrent–voltage behavior characteristic of an n-type semiconductor, gave ptype behavior after etching in dilute HCl, an example of which is shown in Figure 9.9. The study by Kronik et al. [65] used surface photovoltage (SPV) spectroscopy and X-ray photoelectron spectroscopy (XPS) to investigate this effect. It was concluded that the CdSe itself was close to intrinsic, as might be expected for very small nanocrystals [the crystal size in this study was 4–5 nm; no such apparent reversal of semiconductor type was found for considerably larger (ca. 20 nm) crystal size]. The direction of photocurrent flow, rather than being determined by an electric field (the space charge layer) in the semiconductor, was determined by trapping of charge carriers at the individual nanocrystal surfaces. Both electrons and holes could be trapped at the surface; however, the

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FIG. 9.9 Current–voltage characteristics under chopped illumination (
AM1) of a nanocrystalline (4–5 nm) CdSe film, deposited by CD, in a polysulphide electrolyte. The two characteristics are for as-deposited CdSe (top) and after etching with dilute HCl (bottom). (After Ref. 67.)

charge that was preferentially trapped (longer lifetime of trapped charge, which usually correlated with deeper trap sites) would preferentially be transferred to the electrolyte. Such a scenario is reasonable due both to the longer residence time of the charge at the surface as well as better overlap of a deeply trapped carrier with the redox species, which, for the polysulphide system used in these experiments, is located deep in the bandgap. The presence of such charge traps and the effect of water on them was further substantiated in later studies using photoluminescence [68,69] and scanning probe spectroscopies [70]. See Section 4.2.7.1 for a description of surface trapping in CD CdSe measured by various photoluminescence studies. The photovoltage of such a cell comes not from neutralization of a built-in space charge layer, but from change in the Fermi level in the almost intrinsic nanocrystal upon illumination. This change in Fermi level is also determined by the relative trapping of electrons and holes. If no trapping occurred and no charge extraction took place, then illumination would not change the potential of the semiconductor appreciably, since both electron and hole Fermi levels would change by (close to) the same amount and in opposite directions. However, to take a simple example, if the holes are deeply trapped and the electrons are not trapped at all, then the shift in the overall Fermi level is determined almost completely by

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the electron concentration; i.e., the Fermi level rises on illumination, similar to what happens in an n-type semiconductor. The opposite will occur if electrons are trapped and holes are free. In practice, both charges are usually trapped to a greater or lesser extent, and the shift in Fermi level, and therefore the photovoltage direction and magnitude, will be dependent mainly on the relative trapping depths of the two carriers. While it has not been studied, it is probable that there exists a range of crystal sizes, possibly quite wide, where both surface trapping and space charge layer effects contribute to the PEC functioning of the photoelectrode. While the cell efficiencies of these films were not specifically investigated, best parameters of 2 mAcm2 (ca. AM1 illumination; quantum efficiencies increased with decreasing illumination intensity due to diffusion limitations in the nanoporous film); 0.5 V and ca. 50% fill factor were obtained. However, great variation in these parameters were obtained; one reason for this can be seen from a consideration of Figure 9.9. If a CdSe film is etched, but less than optimally (shorter time, more dilute HCl), it is clear that after a certain, unique etch treatment, zero net photocurrent will be obtained. The actual photocurrent (and other output parameters) of the film is a balance between photoanodic and photocathodic currents. A word on “ohmic” contacts to space charge layer–free nanocrystalline films. The ability of a contact to function as an ohmic contact to such films is determined by the offset between the metal contact Fermi level and the semiconductor energy bands, rather than by a potential barrier in the form of an electric field in the semiconductor. It seems that, in practice, most conductors act as good sinks for photogenerated charges in nanocrystalline semiconductors. One reason for this may be that accumulation of one type of charge (the other is usually removed rapidly by the electrolyte at the very high contact area between nanocrystals and electrolyte) will charge the particle and raise the energy levels until the accumulated charge can flow to the sink.

9.2.3 CdS In contrast to CdSe, most studies on CdS involved either nonannealed films or films annealed at relatively low temperatures. Relative efficiencies (i.e., taking into account that the efficiency of the higher-bandgap CdS will be lower than that of CdSe due to its lower light absorption) are therefore low. Additionally, most PEC studies on CD CdS involve doped films, where ions of the dopant were added to the deposition solution. In Ref. 71, Al doping (0.1 wt% Al added to the deposition bath; not necessarily the same percentage in the films) improved the PEC properties. Annealing (in H2) at 200°C, which removes S and should make the films more n-type, improved the cell parameters of both doped and undoped films [72,76]. Stability data

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TABLE 9.2 PEC Parameters of CD CdS Efficiency (%) 0.003 0.0075 (Al-doped) 0.04 (Al-annealed) 0.01 0.05 (In-doped) 0.022 0.05 (As-doped) —

Isc (mAcm2)

Voc (V)

0.06 0.125 0.5 0.17 0.4 0.3 0.5 ca. 0.45

ca. 0.11 ca. 0.11 0.24 0.17 0.23 0.175 0.22 0.58–0.6

FF

Electrolyte

0.33 0.30 0.35 0.44 0.53 0.42 0.44

Polysulphide Polysulphide Polysulphide Polysulphide Polysulphide Polysulphide Polysulphide Polysulphide or sulphite

Illumination (mWcm2) 100 100 100 100 100 100 100 60

Ref. 71 72 73 73 74 74 75

in polysulphide electrolytes were presented. In doping (optimum 0.01% In) also improved the PEC characteristics [73]. The study in Ref. 74 followed essentially the same deposition procedure as before, but on Cr-plated steel (compared with steel in the previous studies; see Sec. 9.2.1 for an explanation of the effect of Cr plating based on the poor electrocatalytic activity of Cr) and with a CuCl:KCN etch at 90°C (the reason for this specific etch was not explained). These modifications led to improved PEC response, which was further improved using As-doped CdS, deposited by adding AsCl3 to the deposition bath. In general, there is an optimum doping concentration, which varied from dopant to dopant. The dopants studied here were donors (increased n-type CdS). Increase in doping density can increase all parameters; but if too much dopant is present, the parameters can degrade due to a narrower space charge layer (poorer ISC due to decreased collection efficiency in the red) and increased recombination due to impurity centers. For nanocrystalline nonannealed (or low-temperature-annealed) films, where there may be no space charge layer as discussed previously, the effect of recombination centers will still be valid—maybe even more important due to the lack of a space charge layer to separate charges. Cu doping of CdS has been investigated [77]. Since the light intensity was not specified, the cell parameters are not given in Table 9.2. The doping caused a small increase in ISC and an equally small decrease in VOC, with no appreciable change in efficiency, although it is arguable if these changes are significant. These electrodes were also used with a Ag2S storage electrode in a photoelectrochemical storage cell. Photoelectrochemical characterization was also carried out on CdS films using different sizes of CdS nanocrystals [75]. VOC increased with decreasing crystal size from 0.58V (75 nm) to 0.68 V (5 nm). Surprisingly, ISC was not dependent

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on size, as would be expected due to the increasing bandgap and therefore lowered light absorption with decreasing crystal size. Besides evaluating photoelectrodes for use in PECs, photoelectrochemical characterization can be used for other purposes. For example, photocurrent spectra of CD CdS has been used to measure the semiconductor bandgap (as I2 vs. h), and agreement between the bandgap values measured by this method and by absorption spectroscopy for as-deposited and annealed films was found [78].

9.2.4 Other Photoelectrodes A number of other CD semiconducting materials as photoelectrodes have been reported, the basic PEC characteristics are given in Table 9.3. Further details can be found in the original references. Most semiconductors described in this chapter gave n-type response (as already shown, for nanocrystalline semiconductors this does not necessarily mean that the semiconductors are actually n-type but rather that the net photogenerated hole current is to the electrolyte while that of the electrons is to the substrate). One example of a “p-type” photoelectrode is nanocrystalline CD PbSe [83]. Figure 9.10 shows the response of such a film. The crystal size of the PbSe in this film is ca. 4 nm, with regions of crystals two to three times larger (the crystal size distribution is bimodal—see Ref. 88). Interestingly, a single crystal of PbSe was not found to give any photoresponse at all under the same conditions, and the current–voltage behavior of the single crystal was essentially ohmic rather than the clearly asym-

FIG. 9.10 Chopped illumination (100 mW-cm2) current–voltage characteristics of a nanocrystalline PbSe film deposited from a citrate/selenosulphate bath at 60°C. The electrolyte is the original solution from which the film was deposited. (After Ref. 83.)

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TABLE 9.3 PEC Parameters of Miscellaneous Semiconductor Films Efficiency (%) Sb2S3 0.8 (as-dep.) 2.0 (ann.)a 3.9 (STA)b 0.008c Sb2Se3 (nonannealed) 0.06 Bi2S3 (nonannealed) 0.12 Ag2S (nonannealed) –– 0.088 PbSe (nonannealed) p-type response 0.029 SnS2 (nonannealed) 0.006 (Cd, Pb) S 0.14 ann.f 0.83g no ann. (Cd, Zn) S (nonannealed)h 0.04 (Cd, Hg)S (annealed at 320C in air)i 0.36

Isc (mAcm2)

Voc (V)

FF

Illumination (mWcm2)

Ref.

2.0 4.0 5.6 0.14

0.44 0.48 0.54 0.155

0.37 0.44 0.52 0.41

Polyiodide

40

79

Ferrocened

100

80

0.45

0.37

ca. 0.3

Polyiodide

80 (IR filter)

52

0.34

0.39

0.55

Polysulphide

60 (water filter)

51

0.1 0.25

0.1 0.09

–– ca. 0.25

Polysulphide Polysulphide

100 200

81 82

1.8

0.08

0.4

Pb2/SeSO32

100

83

0.019

0.45

0.43

KCl

100

84

1.4 ca. 2.8

0.2 ca. 0.5

0.39 0.3

Polysulphide Polysulphide

75 50

85 86

0.065

0.31

0.43

Polysulphide

22

87

1.6

0.51

0.33

Polysulphide

75

85

Electrolyte

e

a

Annealed at 300C in N2. With silicotungstic acid and annealed at 300C in N2. c As deposited. d In dimethylsulphoxide solution. e The solution from which the PbSe was deposited. f Annealed at 320C in air; 0.11 mole fraction Pb––efficiency decreased as Pb increased. g For Cd0.925Pb0.075S. h For Cd0.7Zn0.3S. i For Cd0.82Hg0.18S. b

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metric diode behavior of the nanocrystalline film. Note also the p-type response of etched CdSe (also CdS) nanocrystalline photoelectrodes discussed in Section 9.2.2.2. Photoresponse spectra of PbSe films of different crystal sizes, reflecting the varying bandgap due to size quantization, are shown in Figure 10.6. Some ternary compounds have also been used as photoelectrodes, with optimum efficiencies reported in Table 9.3. For Cd1xPbxS, the efficiency increased from 0.3% (pure CdS) to 0.83% (x
0.075) and then decreased sharply to give an even lower efficiency than for pure CdS (at x
0.1), continuing to drop more slowly at higher Pb fractions [86], although another study on this material found the efficiency of all alloys to be less than that of the pure CdS [85]. A similar trend of initial increase in efficiency was found for Cd1xZnxS, with an increase in efficiency from 0.01% (CdS) to a maximum of 0.04 (Cd0.7Zn0.3S), dropping to the CdS value at x
0.4 and continuing to drop slowly at higher Zn values [87]. The initial increase in efficiency for the (Cd,Pb)S film with increasing Pb content was probably due to the lower bandgap and therefore increased absorption. The reason for the subsequent decrease in efficiency with further increase in Pb cannot be explained so simply, although it may be noted that CD PbS is not very photoactive, possibly due to its relatively high conductivity. For the (Cd,Zn)S electrodes, since the bandgap of ZnS is much higher than that of CdS, the initial increase in efficiency (ISC also increases) is unexpected, although a possible reason is better doping characteristics, resulting in an optimum resistivity; The series resistance of these photoelectrodes was a minimum at x
0.3. For (Cd,Hg)S films, while the bandgap decreased with addition of Hg, the efficiency of films with a little (0.04 mole fraction) Hg was much lower than those of pure CdS but increased with further addition of Hg until efficiencies close to those of pure CdS were reached at 0.18 mole fraction Hg [85].

9.2.5 Coupled Photoelectrodes Some studies involving coupled photoelectrodes of two or even three semiconductors, with at least one deposited by CD, have been reported. A relatively popular subject is sensitization of nanocrystalline wide-bandgap semiconductors with lower-bandgap semiconductors. This is due to the wide interest in dye-sensitized TiO2 solar cells [89]. A range of semiconductors has been deposited on widebandgap semiconductors using a variety of techniques, most commonly by dipping the wide-bandgap semiconductor film (usually TiO2) into solutions of a metal salt and then a chalcogenide solution, but also using electrodeposition and chemical vapor deposition (Refs. 90 and 91 give some examples). Charge separation efficiency often is improved compared to the absorbing semiconductor itself, usually due to electron injection into the conduction band of the large-bandgap semiconductor, which reduces electron–hole recombination. The stability of the semiconductor is also often improved.

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In such a coupled system, CdSe has been chemically deposited onto TiO2 (the latter prepared by both spray-painting and by screen printing) from a triethanolamine/ammonia/selenosulphate bath [92]. The films were annealed at 400°C in air. The quantum efficiencies of such films using spray-painted TiO2 and polysulphide electrolyte, were found to be ca. 20 times higher than films of CdSe by itself (maximum quantum efficiency was reported to be over 0.6, and light-toelectricity conversion was 1%). Much lower values were found for the CdSe-sensitized (lower-surface-area) screen-printed TiO2 films, for nonannealed films, and in ferro/ferricyanide electrolyte. The spectral response of the CdSe/screen-printed TiO2 electrode exhibited a pronounced response beyond the bulk bandgap of CdSe. This subbandgap response was absent in the CdSe electrode. It also was not apparent in the absorption spectrum of the coupled electrode, although this absorption spectrum was not clearly defined (it was obtained from diffuse reflectance measurements). CdSe was deposited on CdS (both deposited by CD) and subjected to different annealing temperatures [93,94]. The purpose was to see if the CdS/CdSe heterojunction affected the PEC properties. The main effect of the coupled system compared to only CdSe was to improve the stability of the photoelectrode in ferro/ferricyanide electrolyte (a partially stabilizing electrolyte for CdSe). The spectral response of the coupled system (measured, as usual, at low light intensities) was closer to pure CdS than to CdSe, although the values of ISC were similar for both at solar intensities, indicating an illumination-dependent behavior. Deposition of ZnO (by dipping in ZnAc2/methanol and annealing at 400°C) improved the stability of the system greatly, although with a decrease in ISC, so it is difficult to know how much of the increase in stability was due to the lower ISC and how much to other factors (stability normally will increase at lower currents). Changes in surface morphology, related to formation of Cd2Fe(CN)6, were measured on all samples. Various hypotheses were put forward to explain the effects of the coupled system and of the ZnO on the stability.

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66. 67.

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G Hodes, A Albu-Yaron. Proc. Electrochem. Soc. 88-14:298, 1988. G Hodes, IDJ Howell, LM Peter. In: Photochemical and Photoelectrochemical Conversion and Storage of Solar Energy. Tian, Z. W., Cao, Y., eds. Int. Acad. Publishers, Beijing, China, 1993, p 331. E Lifshitz, I Dag, I Litvin, G Hodes, S Gorer, R Reisfeld, M Zelner, H Minti. Chem. Phys. Lett. 188, 1998. E Lifshitz, I Dag, I Litvin, G Hodes. J. Phys. Chem. B 102:9245, 1998. B Alperson, I Rubinstein, G Hodes. Phys. Rev. B 6308:1303, 2001. CD Lokhande, SH Pawar. Mat. Res. Bull. 18:1295, 1983. CD Lokhande, SH Pawar. Sol. State Commun. 49:765, 1984. SH Pawar, LP Deshmukh. Mater. Chem. Phys. 10:83, 1984. LP Deshmukh, AB Palwe, VS Sawant. Solar Cells 28:1, 1990. KK Nanda, SN Sarangi, S Mohanty, SN Sahu. Thin Solid Films 322:21, 1998. CD Lokhande, MD Uplane, SH Pawar. Ind. J. Pure & Appl. Phys. 21:78, 1983. CD Lokhande, MD Uplane, SH Pawar. Sol. State Commun. 43:623, 1982. ME Özsan, DR Johnson, M Sadeghi, D Sivapathasundaram, G Goodlet, MJ Furlong, LM Peter, A Shingleton. J. Mater. Sci. Mater. Electron. 7:119, 1996. O Savadogo, KC Mandal. J. Electrochem. Soc. 139:L16, 1992. LP Deshmukh, SG Holikatti, BP Rane, BM More, PP Hankare. J. Electrochem. Soc. 141:1779, 1994. SS Dhumure, CD Lokhande. Mater. Chem. Phys. 28:141, 1991. SS Dhumure, CD Lokhande. Sol. Energy Mater. Sol. Cells 28:159, 1992. S Gorer. PhD dissertation, Weizmann Institute of Science, Rehovot, Israel, 1996. CD Lokhande. J. Phys. D: Appl. Phys. 23:1703, 1990. M Skyllas-Kazacos, JF McCann, R Arruzza. Appl. Surf. Sci. 22/23:1091, 1985. LP Deshmukh, BM More, CB Rotti, GS Shahane. Mater. Chem. Phys. 45:145, 1996. LP Deshmukh, CB Rotti, KM Garadkar. Mater. Chem. Phys. 50:45, 1997. S Gorer, A Albu-Yaron, G Hodes. J. Phys. Chem. 99:16442, 1995. B O’Regan, M Gratzel. Nature 353:737, 1991. D Liu, PV Kamat. J. Phys. Chem. 97:10769, 1993. R Vogel, P Hoyer, H Weller. J. Phys. Chem. 98:3183, 1994. ME Rincon, O Gomez-Daza, C Corripio, A Orihuela. Thin Solid Films 389:91, 2001. ME Rincon, M Sanchez, A Olea, I Ayala, PK Nair. Sol. Energy Mater. Sol. Cells 52:399, 1998. ME Rincon, M Sanchez, J Ruiz-Garcia. J. Electrochem. Soc. 145:3535, 1998.

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10 Nanocrystallinity and Size Quantization in Chemical Deposited Semiconductor Films

Chemical deposition is a low-temperature technique compared with most other semiconductor film deposition methods. This has both advantages and disadvantages. An obvious advantage is the simple processing often inherent in a low-temperature technique. What may be a more important advantage, however, is that low-temperature deposition techniques usually (although not always) result in small crystal size. As recently as a decade ago, this would have been considered a decided disadvantage—large crystal size was almost always desired then, e.g., for photovoltaic cells in order to minimize grain boundary recombination. However, with the increasing emphasis on nanostructured materials over the past decade, this characteristic of CD films is now often considered an advantage. This chapter deals mainly with quantum size effects in CD nanocrystalline films. However, another, quite separate property of such films is related to the large percentage of atoms located on the surface of the nanocrystals of these films, e.g.
50% for a crystal size of a few nm; this is the effect of adsorption of molecular and ionic species on the nanocrystal surfaces. This aspect has been dealt with much less than has size quantization; therefore, it constitutes only a very small part of this chapter, mainly Section 10.2.3, which discusses the effect of adsorbed water on CD CdSe films. Section 9.2.2.2 deals in somewhat more detail with this particular issue.

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10.1 THE QUANTUM SIZE EFFECT: BACKGROUND Quantum size effects in semiconductor nanocrystals became an important field of research in the 1980s, when a number of groups, notably those of Brus at Bell Labs and Henglein at the Hahn Meitner Institute, published seminal papers on the effects of the size of semiconductor colloids on their optical properties and correlated crystal size with changes in electronic band structure. Quantum size effects in semiconductor nanocrystals have been seen before, although the effect presumably was not realized. Early references to precipitates formed when alkali, cyanide-containing selenide solution was added to ammoniacal, cyanide-containing Cd2 described them as orange-yellow when precipitated in the cold and changing to red-brown when heated, and also noted that finely divided Cd-selenide varies in color from yellow to red-brown [1–3]. As is described later, these color changes from normal dark-brown or black CdSe are the most obvious and visual manifestations of the quantum size effect (or size quantization). The terms nanocrystals and quantum dots are often used interchangeably. Quantum dots, as used here, are invariably nanocrystals (amorphous materials could, in principle, also exhibit quantum size effects as long as some electronic separation between different particles occurs) that show quantum effects, while nanocrystals may or may not be small enough to exhibit such effects. Three-dimensional size quantization is due to localization of electrons and holes in a confined volume—e.g., a semiconductor nanocrystal—resulting in a change of the energy band structure. As the crystal size decreases below a certain limiting size, associated with its exciton Bohr diameter, the spacing between levels in the bands becomes larger [the energy structure changes from a quasi-continuum (band) to discrete, quantized levels] and the bandgap increases. This latter change is manifested as a blue shift in the optical spectrum of semiconductor quantum dots (the “quantum size effect”). A simple way to visualize this effect is to consider a silicon atom (silicon being the best-known and most common semiconductor), with two electrons in the 3s level and two in the 3p levels (Fig. 10.1). sp3 hybridization of the single s and three p levels leads to formation of four degenerate (i.e., equal-energy) sp3 levels, each containing one electron. Interaction of these levels with neighboring atoms results in splitting into bonding () and antibonding (*) orbitals, with all electrons in the bonding orbitals and none in the antibonding ones. Up to here, the situation in Figure 10.1 is shown for a single atom. However, as additional atoms are added, each atom contributes its own orbitals and the localized orbitals of the single atom are gradually broadened into a range of molecular orbitals and, eventually, when the number of atoms becomes very large, into the familiar filled valence band and empty conduction band characteristic of a semiconductor. If we consider the process from the other direction

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FIG. 10.1 Scheme showing (from left to right) how the relevant energy levels of silicon hybridize, interact with other atoms, split in a cluster, and eventually broaden into bands. (Adapted from Fig. 4 in LE Brus. Nouv. J. de Chem. 11:23, 1987.)

(right to left), then the bandgap increases and the levels within the bands, which in the bulk semiconductor are extremely (almost infinitesimally) close to neighboring levels, open up into discrete levels, as described previously. This picture is reasonably valid for covalent silicon but rather simplistic for many of the semiconductors common in CD, which are usually mixed covalent and ionic. However, it serves to give a feeling for size quantization. For those readers who would prefer a more realistic interpretation for semiconductors with considerable ionic character, it is suggested that they construct a similar scheme for purely ionic materials and then “imagine” the required combination of ionic and covalent character. There are many theoretical models to correlate the increase in semiconductor bandgap with crystal size. However, for our purposes we will show only the original model, known as the effective mass model, since this is the easiest to understand, in spite of its limited accuracy. The effective mass model is based on the energy of the lowest-energy particle-in-a-box configuration, taking into account that the relevant mass term is given by a reduced effective mass, , where  is given by 1 1 1 
    me mh

(10.1)

The effective masses of electrons (me) and holes (mh) represent the masses that these charges appear to have when moving in the solid rather than in free space, and these vary from material to material. (In the size quantized regime, they can also vary with crystal size, particularly for small quantum dots, hence the limitations of the effective mass model).

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The increase in bandgap, E, of a semiconductor due to size quantization is then given by 2 2 1.79e2 E
2   R 2R

(10.2)

where the first term on the right-hand side is the localization energy [the particlein-a-box energy of the charges in a box (more correctly, in this case, a sphere) of radius R, modified by the reduced effective mass term) and the second term represents a reduction in the energy increase due to coulombic interaction between the electron and the hole and is a function of the dielectric constant of the semiconductor, . The increase in bandgap is inversely dependent on both the reduced effective mass and on the square of the radius. The bandgap should therefore increase as a parabolic function of the decrease in size. In practice, the rate of increase is less than this, and an exponent considerably lower than 2 gives a better fit of the bandgap increase with decrease in crystal size. An obvious importance of this size quantization is that a single semiconductor can possess a range (sometimes a wide range) of bandgaps, which can be controlled if the semiconductor crystal size is controlled. This “bandgap tailoring” allows (ideally) control of all properties that depend on the bandgap. The most obvious is the optical transmission (absorption) spectrum, and this is the property most often measured in quantifying size quantization, since the bandgap can be estimated from this spectrum. In this chapter, size quantization effects in CD films are described. Since the majority of reports on size quantization in CD films mention the effect but do not go into detail on this aspect, as with many other chapters in this book, it will be more efficient to tabulate the relevant literature and to deal with individual studies that provide additional results of interest outside of what is included in the table or require further discussion. CdSe and PbSe will be dealt with in a more integrated manner, since films of these materials, in particular CdSe, have been the most intensively studied from the viewpoint of their nanocrystallinity and quantum size effects. Some words of caution in interpreting optical transmission (absorption) spectra (see also Sec. 1.4.2). Since the energy structure of nanocrystals in the quantum size domain is more like an atomic structure, with separate levels, than the band structure of bulk semiconductors, the derivation of bandgap from the absorption spectra using the (h)n vs. h plot, which is based on the density of states (band structure) of the bulk semiconductor, is not really valid. Nonetheless, it is often used to give an approximation. Since there is always a distribution of crystal sizes (and therefore of bandgaps) in these size-quantized films, there is in any case no single “correct” bandgap, and any value measured will be an approximation. Another area where nanocrystallinity is important is that of surface effects due to the high real surface areas of such films (often tens of percent of all atoms

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are located at a surface). Since these films are invariably porous to a greater or lesser extent, much of this surface is accessible to modification and treatment by liquids and gases. This aspect has been less dealt with for CD films than quantum size effects, but some examples do exist and are briefly discussed in this chapter, with references to other sections of this book where they are of specific relevance. The factors that influence crystal size in CD films are discussed in Section 10.2 on CdSe, since they have been most studied for that material. However, the principles involved are general for all semiconductors.

10.2 CdSe 10.2.1 Historical Background This chapter is the one closest to the interests of the author. These interests, both in chemical deposition and in size quantization, were a direct result of a single serendipitous observation. As a historical and personal interlude, this “experiment,” involving CD CdSe, will be briefly described and its rationale explained (in reverse order). Accepting an invitation to spend some time in Campinas, Brazil, I planned to utilize my experience in electrodeposition of CdSe films. Shortly before this trip, a paper appeared describing a method of electrodeposition of CdSe, based on the use of selenosulphate instead of the commonly used SeO2 (Ref. 4, which itself was based on an earlier study [5]). The trip to Campinas and the resulting ability to spend a lot of time in the lab seemed a good opportunity to try this method. The method proved to be simple to reproduce. More relevantly, however, in keeping with my overall philosophy on life, instead of clearing away the “finished” experiment, the beaker containing the electrodeposition solution was left sitting over the ensuing weekend on the laboratory table to be taken care of at a more suitable (i.e., later) time. On returning to the lab after the weekend, I found the beaker in question to contain a bright red precipitate and also a similarly colored, transparent film on the inside of the glass walls. My first thought was that this was elemental Se (which is usually bright red in the freshly precipitated state), based on my understanding that selenosulphate is not very stable and can readily form elemental Se under certain conditions. A simple chemical test (treatment with Na2S solution, in which Se dissolves) showed that the precipitate was not Se. To make a fairly long story short, the next two ideas (occurring more or less at the same time if my memory does not fail me) were either that I had discovered a new amorphous form of CdSe or that the CdSe was size quantized, and some further experiments were all that were needed to verify the latter hypothesis. Had I been more familiar with chemical deposition at that time, I might well have ignored the “red” CdSe as something obvious; an earlier study on CD CdSe had noted that the as-deposited films were red [6].

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Just to complete this history, it should be mentioned that the weekend in question occurred shortly after I had spent several days in Rio de Janeiro during Carnival. I do not make any (overt) claims that this experience affected my ability to interpret the experiment in any way. The results of this experiment and the subsequent investigation were published in Ref. 7. At the time, it was not obvious that quantum size effects would be seen in strongly aggregated nanocrystals. Subsequently, this was found to occur quite commonly (as seen in this chapter). While some degree of electronic isolation between crystals is needed for size quantization to be exhibited, this isolation need by no means be absolute. CdSe is a particularly attractive material to show the phenomenon of size quantization, since its color can vary from very deep red (even black in powder form) to yellow (almost white for very tiny crystals). To give a feeling for the size dependence of the color, crystals of 6 nm will be red, 4 nm orange, and 3 nm yellow, a wide range in color for a small range in crystal size. This variation in color can be seen from the transmission spectra shown in Figure 2.9. The spectral changes show how the absorption onset (equal to the start of the transmission decrease) moves to the red as the deposition temperature increases and also as the mechanism changes from a cluster mechanism to an ionby-ion one (see later). The bandgap range, from ca. 2.3 eV to 1.8 eV, parallels a change in crystal size from ca. 3 nm to over 12 nm (the rightmost spectrum is for a film with crystal size ca. 20 nm, but this spectrum is reached by the time the crystal size is close to 12 nm, where it can be considered to be bulk from the point of view of size quantization).

10.2.2 What Determines the Size of Nanocrystals in CD Films? The factors that determine crystal size, a topic of particular relevance to this chapter, have been discussed to some extent in Section 3.4. There are two main factors that generally affect crystal size for any particular material: the deposition mechanism and the deposition temperature. The hydroxide cluster mechanism is expected to give a crystal size similar to that of the original hydroxide cluster (with some growth possible as deposition proceeds). That size depends mainly on temperature, both because higher temperatures allow more grain growth and, possibly more important, lower temperatures kinetically stabilize very small nuclei in solution that are thermodynamically unstable. For example, in the hydroxide cluster mechanism, where crystal size is believed to be controlled mainly by the size of the Cd(OH)2 colloids, the relevant equilibria are Cd2  2OH D Cd(OH)2

(10.3)

A number n of these molecules can form a cluster, [Cd(OH)2]n. This cluster can

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continue to grow by collecting a variety of solution species, one possibility being [Cd(OH)2]n  2OH D [Cd(OH)2]n.2OH [Cd(OH)2]n.2OH  Cd2 D [Cd(OH)2]n1

(10.4) (10.5)

This may continue until eventually the cluster is large enough to be thermodynamically stable (i.e., will not redissolve). However, if the cluster is smaller than the critical nucleus size, then there is the possibility that the nucleus will redissolve. The lifetime of the nucleus will then depend on its size and also on the temperature; lower temperatures will slow the redissolution step. Thus lower temperature increases the chance that a subcritical nucleus will eventually grow to a stable size rather than redissolve. This kinetic stabilization of small nuclei results in a greater total density of nuclei and therefore smaller crystal size, since the total quantities of reactants are fixed. For the ion-by-ion reaction, nucleation is generally slower and the density of nuclei smaller. Additionally, growth occurs (ideally) only at a solid surface; therefore nucleation is confined to two dimensions, in contrast to three dimensions for the cluster mechanism. The crystal growth may terminate when adjacent crystals touch each other or by some other termination mechanism, e.g., adsorption of a surface-active species. These factors should be valid regardless of whether the mechanism proceeds via free chalcogenide ions or by a complex-decomposition mechanism. The effect of temperature and mechanism on the optical spectra, through the crystal size, is clearly seen in Figure 2.9. In particular, the difference in crystal size between the two rightmost spectra, both deposited at 80°C but one through the cluster mechanism and the other (HC) through the ion-by-ion mechanism, is relatively large: 8.5 nm for the film deposited by the cluster mechanism and 20 nm for that deposited by the ion-by-ion mechanism. While the effect of temperature is gradual, that of mechanism is sudden. It is determined by the conditions that separate the formation of metal hydroxide colloids from a solution with no metal hydroxide phase. Decrease in the complex:metal concentration ratio and increase in temperature and pH will all favor hydroxide formation. This sudden transition on varying the complex:metal ratio is shown, for two different solution temperatures, in Figure 10.2. The spectra are independent of the NTA:Cd ratio (NTA, nitrilotriacetate, the complex used) until the transition between hydroxide-containing and hydroxide-free solutions is reached, whereupon they suddenly undergo a red shift (increase in crystal size, decrease in bandgap) and then no further change as the NTA:Cd ratio increases further. The crystal sizes in these films were also shown, by XRD, to change only at the NTA:Cd ratio where the spectrum changes and the experimental results agreed with thermodynamic calculations on the region of

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FIG. 10.2 Optical transmission spectra of CdSe films deposited at 10°C (top) and 80°C (bottom) with varying NTA:Cd molar ratios (shown in the figures). The bandgaps estimated from the spectra are indicated by the thin vertical lines. (Adapted from Ref. 8.)

existence of Cd(OH)2 as a function of temperature, pH, and solution composition [8]. Illumination, by light that is absorbed by the growing semiconductor crystals, has been shown to increase the crystal size somewhat, seen as a red shift in the optical spectrum and decrease in bandgap by as much as 0.2 eV [9–11] (see Fig 4.3). This is probably due to photoelectrochemical reactions taking place at the crystal surface. The chemical deposition solution can also be used to electrodeposit CdSe. Electrons (either from an external source, as in the case of elec-

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trodeposition, or generated by illumination during CD) can reduce selenosulphate to selenide, which reacts with Cd ions, causing (photo)electrodeposited CdSe to form on the crystals and thereby increasing the crystal size. Only light that is absorbed by the semiconductor can cause this effect, as expected based on the photoelectrochemical mechanism. In addition, a certain minimum intensity is needed to cause a measurable red shift, but the effect saturates at high light intensities. At low intensities, recombination was assumed to remove the photogenerated electron before it had time to reduce selenosulphate, hence the threshold. As intensity increases, there is an increasing likelihood of photoelectrochemical reduction. The saturation was explained by assuming that only one electron/hole pair was effective; further increase in intensity has no further effect [10,11] Since size-quantized CdSe can undergo visible color changes with change in crystal size, this effect of illumination can clearly be used to form patterns on the film by illuminating the film through a mask. We have even observed interference fringes at the edges of these patterns, alternating between orange (nonilluminated) and brown-red (illuminated), corresponding to destructive and constructive interference, respectively, at the pattern edges. It is a general observation when quantum size effects are observed in CD films that the blue shift is reduced somewhat as the film thickness is increased, and this has been shown clearly for CdSe [9,11–13]. From absorption spectra, a difference of ca. 0.08 eV (a crystal size difference of 0.26 nm for a crystal size of ca. 4 nm) was shown to occur between thin and thick films (growth time between 10 and 190 hr) [11]. In another study, using the same basic deposition solution (NTA/selenosulphate), red shifts in the photoluminescence spectra could be correlated with a change in crystal size from ca. 4.5 nm (2.1 eV) to ca. 8 nm (1.85 eV) [12].* The widths of the photoluminescence peaks increased as the deposition time (and therefore crystal size) increased. This was explained as an increase in size distribution of the crystals as deposition proceeded. Since the change in bandgap with crystal size increases as crystal size decreases, wider peaks in the photoluminescence spectra, if due to increase in size distribution, means that the size distribution increases greatly with increasing deposition time. If no change in the size distribution were to occur, the peak widths should actually decrease with increase in average crystal size. While photoluminescence peak widths of sizequantized films may certainly be influenced by size distribution, other factors, in particular recombination from surface sites with various spatial separation, will also affect the width. A microscopic study of such films would be the most dependable way to measure size distribution.

* The values of crystal size reported in this study varied between 6.6 and 10.6 nm. These values were based on the effective mass approximation, which overestimates the bandgap increase with decreasing crystal size, particularly for small sizes.

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Increase in size distribution with increasing film thickness (or deposition time) is expected for a number of reasons. One is the obvious one that, since film growth can involve both crystal growth and new nucleation, the chances of any particular crystal growing increases with time (film thickness), due to deposition of new material on the crystal. Also, as deposition proceeds, the solution composition changes, and this can lead to changes in crystal size. There are two main causes of this. One is that, while the Cd concentration decreases, that of the complex does not change (this may not be the case where a volatile complexant, like ammonia, is used in an open bath). Therefore the complex:Cd ratio increases during deposition, and at some point the mechanism may change from hydroxide cluster to ion-by-ion; the latter normally gives larger crystals and may also occur on already existing crystals. The second cause is that, even assuming no change in deposition mechanism, crystal size grows slightly with decrease in reactant concentrations; reducing the concentration of all reactants to one-half the original concentration resulted in a bandgap increase of between 0.05 and 0.10 eV, corresponding to a size increase of ca. 10% [9]. This can be explained based on the concept of kinetic stabilization of small nucleii. As described earlier, small nucleii are stabilized at lower temperatures, thus providing a greater chance of growth to a size where the crystal will be thermodynamically stable and resulting in a smaller crystal size. However, the lower the concentration of reactants (selenosulphate and free Cd2, the latter of which will decrease due to both decrease in total Cd concentration and increase in complex:Cd ratio), the slower will be this growth and therefore, as with increased temperature, the greater likelihood of redissolution of the small nuclei, resulting in a larger final crystal size. Of course, this concentration effect is important not only in the context of varying film thickness (deposition time) but also as a means for further control of crystal size. Thicker films with crystal size more characteristic of thin films could be formed by depositing on an existing thin film from a new solution and repeating this to the desired thickness. This suggests that the main reason for increase in crystal size with continuing deposition is not simply because of deposition on already-deposited crystals, but because of changes in the composition of the deposition solution. This is also borne out by a comparison of modulated electrotransmission (ET) and electroreflection (ER) spectroscopies [where modulation of the potential of the film in an electrolyte results in corresponding changes in the absorption and reflection) of a CD CdSe film (Fig. 10.3)]. The bandgap of the film is ca. 2.1 eV, but the ET spectrum is broader and shifted mainly to the high-energy side compared to the ER spectrum. Since the ER spectrum samples the near-surface region of the film and the ET spectrum the total film, this difference suggests that the film is composed of smaller crystals close to the substrate and larger ones toward the surface. The effect of film thickness was very evident for films deposited from baths based on N,N-dimethylselenourea and complexed with both ammonia and either

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FIG. 10.3 Modulated electrotransmission (ET) and electroreflectance (ER) spectra of a CD CdSe film deposited from an NTA/selenosulphate bath at 30°C. The experiments were carried out in an electrolyte containing sulphide and sulphite (the latter to prevent formation of colored polysulphide) at a pH of 10 (buffered with NaH2PO4).

citrate or tartrate [14]. Bandgaps estimated from the transmission spectra varied by at least 0.2 eV over a thickness range from under 100 nm to a few hundred nanometers. (Similar thickness effects were seen using photoluminescence in similarly prepared films [13].) Films from the tartrate bath were deposited more rapidly than those from the citrate bath [15], and the higher bandgap found for the citrate-based films in this study was probably due to the fact that those films were thinner than the tartrate ones. For comparable film thickness, the bandgaps were similar (derived from a comparison of deposition rate [15] and spectra as a function of deposition time [14]). No XRD patterns were found for these films; therefore crystal size was not directly measured. It was noted that the deposition solution color changed during deposition from colorless through turbid yellow, orange, to orange-red. This color change is typical for low-temperature, hydroxide-cluster-mechanism CD of CdSe in general when carried out at relatively low temperatures. Annealed films deposited from a N,N-dimethylselenourea/citrate/ammonia bath were shown to exhibit a (0001) XRD reflection at 2
13°, a reflection normally forbidden in hexagonal CdSe [the (0002) reflection is the one normally seen] [13]. This was explained by a breaking of the selection rules due to the small crystal size. Interestingly, this peak was very weak in thin films and prominent in

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thicker ones. A broad Raman band at 250 cm1, which is not observed normally in CdSe and disappeared after annealing, was also observed in these films and attributed to a surface optical mode.

10.2.3 Photoluminescence (See also Sec. 4.2.7.1) Photoluminescence was mentioned earlier in connection with studies in change of crystal size with film thickness. On a more general note, a number of photoluminescence studies have been reported on size-quantized CD CdSe films. Spectra showing both bandgap luminescence (this is probably not true bandgap but recombination from shallow surface states) [7,12,16] and dominant deep-trap luminescence [10,11,17,18] have been reported. One study, using dimethylselenourea instead of the more common selenosulphate, found both “bandgap” luminescence (which red-shifted on increasing film thickness, explained by increasing crystal size) and a lower-energy peak at ca. 1.75 eV, attributed to larger, weakly quantumconfined crystals. It is very possible that this low-energy peak arises from surface states, since there was no evidence of a bimodal size distribution that would lead to two separate peaks The role of water adsorbed on the surface of these nanocrystals in passivating surface states was discussed in Section 9.2.2.2. This is seen in luminescence by change in the spectrum from deep-trap-dominated in a dry ambient to nearbandgap-dominated in a humid ambient [17]. A study of the deep-trap luminescence showed that the luminescence originated from recombination of trapped charges [17] (see also surface photovoltage spectroscopy measurements in Ref. 18). From a consideration of the optically detected magnetic resonance (ODMR) signals, it was shown that the recombination occurred from sites of low symmetry, i.e., at the surface of the nanocrystals [19]. However, based on time-resolved photoluminescence and transient absorption measurements, even the near-bandgap luminescence was believed to result from shallow-trapped carriers. The dynamics of the photogenerated charge trapping for these films showed, using transient absorption measurements, fast (subpicosecond) electron trapping to shallow surface states [10,20] and slower, but still relatively fast (ca. 50 ps) emptying of these shallow traps, either to the ground state or, more likely, to deeper traps [20]. Even taking into account this strong effect of water adsorption, in the author’s experience the luminescence from these CdSe films can be highly variable, not only in whether it is dominated by near-bandgap or deep-trap recombination but, maybe even more so, in the intensity of the luminescence, varying from relatively strong luminescence (visible to the eye) to no measurable signal at all. While the reason for this large variation is unknown, it is reasonable to assume that it is related to the nature of the surface of the individual nanocrystals. A collection of studies on size quantization in CD CdSe films, together with some relevant data, is presented in Table 10.1.

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TABLE 10.1 Possible Quantum Size Effects in CD CdSe Max. Eg (eV) 2.3 2.3

Min. crystal size (nm) ca. 3.5

2.1 2.28 2.06 2.06 ca. 2.1 ca. 2.05c ca. 2.3 2.3 2.28

ca. 4.5 (6.5)b 6.4d 4.25 4.34

Miscellaneous

Refs.

NTA/selenosulphate Citrate or tartrate/ammonia/ DMSea Triethanolamine/ammonia/ selenosulphate Citrate/ammonia/DMSe Tartrate/ammonia/DMSe Citrate/selenosulphate NTA/selenosulphate NTA/selenosulphate Citrate/ammonia/DMSe NTA/selenosulphate NTA/selenosulphate

7, 8, 9, 17, 19 14 21 15 15 15 12 20 13 11 10

Bulk bandgap
1.73 eV (wurtzite), ca. 1.8 eV (zincblende). All bandgap values given in this chapter are room-temperature values. a N,N-Dimethylselenourea. b This reported size is probably overestimated; simple effective mass approximation used in its estimation. c The optical absorption spectrum consists of a reasonably sharp onset corresponding to a bandgap of ca. 2.05 eV and a broad tail to longer wavelengths (onset of between 1.8 and 1.9 eV). d The films in this study were thick (ca 1 m), and therefore the average crystal size may be expected to be larger than most films, which are usually in the range of 100–200 nm thick.

10.2.4 Annealed Films It appears that most reported CD CdSe films are size-quantized, with crystal sizes of 10 nm. There are some exceptions. Films deposited via an ion-by-ion mechanism at high temperature possess larger crystal size and show no size effects [8]. Films deposited from an ammonia/selenosulphate bath were reported with a bandgap, measured from the absorption spectrum, typical of bulk CdSe (ca. 1.74 eV) [22]. The CdSe in this study was grown at 80°C from a solution containing 48 mM Cd and 2.1 M NH3. Taking into account the additional complexing power of the selenosulphate, such a solution may be close to the transition between a hydroxide mechanism and an ion-by-ion one. One technique commonly used to illustrate size quantization in these films is annealing them to increase crystal size. This results in a gradual red shift in the spectra until eventually “bulk” CdSe (ca. 11 nm in size) is reached, after which no further shift, at least not one due to size, is seen (see Refs. 7, 13, 14, 15, and 23). The increase in crystal size with annealing depends on temperature and time of an-

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FIG. 10.4 TEM micrographs of a nanocrystalline CdSe film as deposited (upper left) and after air-annealing for 20 mn at increasing temperatures up to 500°C (lower right). The 50-nm scale is the same for all micrographs except for the 450°C and 500°C ones, which are marked by the 100-nm scale. (From S. Gorer and G. Hodes, unpublished results.)

nealing as well as on the material annealed and on the annealing atmosphere. For CdSe and CdS, which are usually annealed in air or sometimes in an inert atmosphere, as a rule of thumb, increase in crystal size is slow up to a temperature of ca. 300°C and increases greatly at a temperature somewhere between 300 and 400°C (with further growth at higher temperatures), together with a phase transformation (if the original film is sphalerite) to wurtzite structure. This can be seen in Figure 10.4, which shows a CdSe CD film as deposited and after sequential annealing treatments. (See also Fig. 1a in Ref. 24, which shows XRD spectra of essentially the same process). Although the size increase at low annealing temperatures is small, the red spectral shift occurs mainly in this region, since only a small increase in crystal size is necessary to obtain an appreciable red shift.

10.3 CdS Films described in Table 10.2 are assumed to be deposited using a standard bath (Cd2/ammonia/thiourea/relatively high temperature) unless described otherwise.

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The first report of size quantization in a CD semiconductor film was in 1981 for very thin films of CdS deposited from CdSO4 and thioacetamide [25]. Quantum shifts were measured by photoluminescence (absorption spectra in the bandgap region could not be measured, probably because the films were extremely thin). Compared to an exciton luminescence peak at ca. 507 nm (2.45 eV) for bulk material, a peak at 468 nm (2.65 eV) was obtained for ultrathin films (particles of ca. 10 nm separated from each other by ca. 20 nm) and between 470 and 495 nm (2.64–2.51 eV) for thicker films (thickness not defined). The peak of the ultrathin films did not change if the film was heated at 800°C in H2S, a treatment that would almost certainly result in crystal growth well beyond the size quantization limit for thicker films. The 10-nm particle size is at least twice the size needed to see the observed blue shift in the spectra, suggesting that either the particles were composed of a few aggregated crystals or the vertical dimensions of the particles were much smaller than the measured lateral ones. Blue shifts in very thin films (3 nm average thickness) were also more recently measured for CD CdS films deposited in an ultrasonic bath [34]. In general, very thin films would be expected to exhibit size quantization if the film thickness is less than the Bohr diameter, since quantization will occur in at least one dimension. However, in most cases, such very thin CD films can be

TABLE 10.2 Possible Quantum Size Effects in CD CdS Max. Eg (eV)

Min. crystal size (nm)

2.66 2.63 2.5 2.64 ca. 2.7; see text 2.60 2.5 2.69

See text ca. 5 4.8 20 nm 4.1

2.58 3.2 5.0

Miscellaneous

Refs.

From thioacetamide bath, ultrathin layer CdI2 Eg larger with increase in thiourea concentration Citrate/ammonia bath NTA bath; cluster mechanism Standard bath 65–85°C Cd acetate Electrochemical-induced and thiol capping agent Citrate/ammonia bath Ultrathin layer (2–3 nm), deposited in ultrasonic bath —

25

Bulk bandgap
2.43 eV.

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26 27 28 8 29 30 31 32 33 34 35

expected to be composed of islands, as occurs in the films described earlier, and size quantization is not then expected if the individual crystal size is greater than the Bohr diameter. In Ref. 26, optical transmission spectra of films deposited using CdI2 in a NH3/NH 4 /thiourea bath were substantially blue-shifted (to 2.63 eV) from the normal CdS absorption. While no crystal size was given, from the XRD spectrum given in the paper it appears that the crystal size was considerably larger than 10 nm, and therefore the cause of the blue shift is not clear. No such shift was seen for films deposited under similar conditions but using CdCl2 instead of the iodide. In Ref. 27, small increases in bandgap were found when high thiourea concentrations were used. This was explained, in general terms, as a decrease in grain size with increase in deposition rate. In Ref. 28, using a citrate/ammonia bath, clear increases in bandgap (up to 2.64 eV) were observed. A crystal size of ca. 4 nm was calculated from the XRD data, and this size would be consistent with the observed optical spectra. Most of the XRD peaks, while riding on a strong and noisy background, do appear to be much sharper than would be expected for a 4-nm size. On the other hand, one higher-angle peak [(11.0)] is considerably wider, although it remains wide even after annealing at 450°C for one hour, a treatment that normally will cause crystal growth into the size range of hundreds of nanometers. The main differences between this bath and many standard baths are the use of citrate together with ammonia (citrate is a weaker complexant for Cd than is ammonia, although it might act as a surface blocking agent through adsorbed carboxylate groups) and the high thiourea:Cd ratio (ca. 17). In Ref. 8, crystals ca. 5 nm in size were deposited from a nitrilotriacetate (NTA)-complexed bath (no ammonia) at 40°C (a lower temperature than most CdS depositions). The composition of the bath was such that Cd(OH)2 was present as a colloidal phase (cluster mechanism–see Chap. 3). Under conditions where no hydroxide phase was present and the reaction proceeded via an ion-byion mechanism, much larger crystals (70 nm) and a red-shifted spectrum were found. See Section 10.2.2 for more detail on the dependence of crystal size on the deposition mechanism. In Ref. 29, bandgaps up to 3 eV were reported. From the optical spectra, it is more probable that the maximum bandgap value is ca. 2.7 eV. The bandgap decreased strongly with increase in film thickness up to 250 nm, at which stage the bandgap was close to the bulk value. While most films were deposited at a pH of 11.7, it was noted that smaller particle size was obtained at pH12 and larger ones at pH  9. In Ref. 30, small increases in bandgap (up to a bandgap of 2.6 eV) were obtained for CdS deposited in the presence of a magnetic field. The main factor determining the higher bandgap was a lower deposition temperature (65°C). However, since the minimum crystal size, measured by XRD, was at least 20 nm, and

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bandgap increase due to size quantization in CdS will only become apparent well below 10 nm, there is probably another reason for the apparent increase in bandgap. In Ref. 31, using CdAc2, a blue shift of nearly 0.1 eV (to 2.50 eV) was obtained, compared to films using CdCl2. In contrast to the CdCl2 deposition, where XRD showed sharp peaks, no pattern was observed for the acetate films (amorphous or small crystal size). In Ref. 32, using a process of electrochemically induced CD of CdS (see Sec. 4.1.6.6), nanocrystalline CdS films were deposited by using 2-mercaptoethanol as a strongly adsorbing growth-termination (capping) agent. By increasing the concentration of mercaptoethanol, crystal size was reduced to the region where quantum size effects were observed. Crystal sizes down to 4.1 nm (bandgaps up to 2.69 eV) were obtained. Above ca. 10 mM mercaptoethanol concentration, film formation was prevented. It is particularly interesting (and fortunate) that, up to this concentration, film thickness was not seriously affected by the mercaptoethanol. A strongly adsorbed surface layer on the crystals might be expected to prevent adhesion between crystals and between crystals and substrate, necessary to form a film (and apparently does so at high enough concentration). This technique of surface capping by strong adsorbents, well known for semiconductor colloids, can therefore be applied—albeit with greater limitations expected on the maximum adsorbent concentrations—to control particle size in CD nanocrystalline films. In Ref. 33, a bandgap of 2.58 eV was obtained from a citrate/ammonia bath (deposition solution similar to that used in Ref. 28, only with ten times the Cd concentration). No crystal size measurement was given. A transient optical absorption study of charge recombination dynamics was recently reported on CD CdS with a crystal size of 5 nm (no details of the deposition were provided) [35]. Three different time constants for the decay of the transient absorption bleaching were measured (0.8 psec, 17 psec, and 800 psec) and were attributed to three groups of nanocrystals with specific defects. Trioctylphosphine oxide (TOPO), a well-known passivating agent for CdSe nanocrystals that acts by binding to Cd, was found to increase the relative contribution of the 800-psec recombination, which was therefore attributed to volume recombination, since the TOPO is expected to reduce the relative contribution of surface recombination (represented by the two shorter time constants).

10.4 PbSe The other semiconductor, apart from CdSe, that has been studied with deliberate emphasis on quantum size effects, is PbSe (see Table 10.3). There are several reasons for this. One is the long-known use of CD to deposit PbS and later PbSe. Second, and of particular importance, the electron/hole effective mass in PbSe is very

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low (0.034 times the free-electron mass at room temperature); from Eq. (10.2), this means that the quantum size effect will be particularly large, i.e., the bandgap will increase noticeably even for relatively large crystals and the increase for small crystals will be very large. Third, in keeping with the usual small crystal size obtained by CD, TEM images of CD PbSe from an earlier study showed crystals of the order of 10 nm in size [36]. Motivated by these considerations, a study of CD PbSe films was carried out with an emphasis on crystal size, the control of this size, and size quantization effects in the films [9,37–39]. Much of this study parallels that of CdSe; the reader is referred to Section 5.3.2 for a discussion of these films. Summing up, crystal size can be controlled by deposition temperature (lower temperature gives a smaller crystal size) and by the dominant deposition mechanism (ion-by-ion results in larger size than cluster growth), for the same reasons as described for CdSe. Annealing can also be used to increase crystals size, which results in red shifts of the absorption spectra [39]. However, there are also some differences between PbSe and CdSe films. PbSe sometimes grows with a bimodal size distribution, i.e., areas of small crystals and others of considerably larger ones. Also, PbSe forms very large (ca. 1 m) crystals under conditions of high deposition temperature and ion-by-ion growth, in contrast to CdSe, where the crystal size is typically limited to tens of nanometers. The bulk bandgap of PbSe is 0.28 eV. Values of bandgap up to 1.5 eV for a crystal size of ca. 4 nm have been measured [39]. Figure 10.5 shows experimental values for PbSe films of bandgaps, measured from absorption spectra against crystal size (horizontal bars that show the size distribution), as well as theoretical calcu-

TABLE 10.3 Possible Quantum Size Effects in CD PbS and PbSe Max. Eg (eV) PbS 1.97? PbSe ca. 0.67b; 0.35 c 1.5 ⬇0.8

Min. crystal size (nm)

Miscellaneous

Refs.

3a

Pb2  H2S

47

ca. 30

Selenourea or selenosulphate

43, 44

3.5

Selenosulphate and various complexants Selenosulphate  thiosulphate

9, 37, 39

Amorphous

46

PbSe bulk bankgap
0.28 eV. a Size of coprecipitated PbS. b Photoconductivity maximum; see text. c From low-temperature photoluminescence spectrum and converted to room temperature.

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FIG. 10.5 Experimentally measured values of bandgap of PbSe films (horizontal bars: The length gives the experimental uncertainty in size, mainly due to the size distribution). The broken curve gives the theoretical relationship between bandgap and crystal size based on the hyperbolic band approximation used for PbS in Ref. 40. The room-temperature reduced effective mass (0.034) was calculated from the low-temperature value (0.022) (R. Dalven, Infrared Phys. 9:141, 1969.) according to the temperature dependence given in H. Preier, Appl. Phys. 20:189, 1979. The dotted curve is a more recent calculation based on an envelope function calculation [41].

lations of the bandgap–size relationship based on a hyperbolic band model [40] and an envelope function calculation [41]. The results agree for the most part with the former model. Such agreements between theory and experiment for size quantization should be treated with some caution. Thus, more recent experiments on PbSe nanocrystals in a phosphate glass matrix suggest a smaller increase in bandgap with decreasing crystal size [42]. For example, a crystal size of 3.5 nm, which exhibited an apparent bandgap of ca. 1.5 eV in the CD films, showed one of ca. 1.0 eV in the glass samples. Several reasons could be put forward to explain this discrepancy. The surrounding glass matrix could reduce the confinement in the glass samples, although it might be argued, with equal validity, that the contact between the nanocrystals in the film does the same thing. Additionally, the lead chalcogenides usually exhibit weak absorption near (and even quite far from) the bandgaps, and it is sometimes difficult to unambiguously determine the bandgap in a thin film.

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FIG. 10.6 Photoelectrochemical photocurrent spectra (in a selenosulphate electrolyte) of CD PbSe films deposited under various conditions with four different crystal sizes (see text). The cutoff at 1400 nm is due to absorption by the water. (From S. Gorer, PhD dissertation, Weizmann Institute of Science, Rehovot Israel, 1996).

These films were active as electrodes in a photoelectrochemical cell (see Sec. 9.2.4). Figure 10.6 shows the photocurrent spectral response of four different PbSe films with different crystal sizes (the long-wavelength cutoff at 1.4 m is due to absorption by water from the aqueous solution used). Spectrum (a) is for very small (4 nm) crystals with a bandgap of ca. 1.5 eV. Spectrum (b) is for a film with mainly small crystals and also areas of larger crystals (bimodal distribution). The longer-wavelength feature is due to the larger crystals. Spectrum (c) represents a mixture of moderately small and moderately large crystals, while (d) contains only fairly large crystals. The important point to be made is the control over the shape of the spectral response (the details of the crystals sizes and preparation methods are less important and are not dealt with here). These spectral responses can be translated, in principle, into spectral responses for detectors based on these films (in practice, these films, used as photoelectrodes in the present setup, would probably be neither efficient nor very stable as detectors). Since PbSe photodetectors peak in the infrared and their sensitivity decreases strongly toward the visible, the ability to control both the crystal size (
bandgap) and size distribution (range of spectral response) allows extensive tailoring of the spectral response of a detector based on these films. Response (c) in Figure 10.6 is particularly interesting, for it indicates the possibility of a flat response over a wide range. Blue shifts in photoconductivity spectra, attributed to quantum size effects, were noted in PbSe films deposited using either selenourea or selenosulphate

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[43,44]. The peak in the photoconductivity spectrum shifted from ca. 1.7 m for the as-deposited films to 3.6 m after annealing. These peak values do not give a value of the bandgap directly, but an approximate value of 0.67 eV for the as-deposited film may be estimated, assuming that the shapes near the photocurrent onset of the as-deposited and annealed photoconductivity spectra are similar. Crystal size in similar films was measured to be ca. 30 nm [43,45], slightly less for the selenourea samples and slightly more for the selenosulphate ones. Only small bandgap shifts (0.1 eV) are expected for this crystal size. Shifts in photoluminescence peaks of ca. 0.02 eV (selenosulphate bath) and 0.07 eV (selenourea bath) were more in line with the measured crystal size, particularly taking into account that the crystal size of the former was somewhat larger than that of the latter [44]. In comparing these results (using the selenosulphate bath) with those earlier, which give larger quantum effects, two differences are relevant. One is that the former are considerably thicker (by at least several times) than the latter. This increases the crystal size in general and particularly so for many of the PbSe films. The second is that without further information, neither photoconductivity nor photoluminescence spectroscopy (in particular when peaks, rather than onsets, are used) are accurate indicators of bandgap, although they can provide indication of quantum size effects. When thiosulphate was added to the selenosulphate deposition bath, no XRD pattern was observed in the films, and they were presumed to be amorphous [46]. Annealing at 350°C induced crystallization. Just one minute at this temperature was enough to give a definite XRD pattern. The crystal size after this treatment was 13 nm, meaning that the size before annealing, if any crystallinity did exist, was considerably less than this. A strong blue shift in the absorption spectrum of the as-deposited film showed an onset of ca. 0.8 eV. Finally, two related studies can be mentioned here. It was noted that when a quartz plate was immersed overnight in a solution of Pb(ClO4)2 and poly(vinyl alcohol) through which H2S had been bubbled, a film formed on the plate parallel with formation of a colloidal PbS sol [47]. The film was extremely thin (maximum absorbance of 0.015 at 400 nm). The absorption spectrum of this film was similar to that of the PbS sol and consisted of several absorption peaks with an absorption onset of ca. 630 nm (1.97 eV). It is not clear that this is the true bandgap onset, for the same reasons as discussed previously (weak absorption close to the bandgap). The XRD crystal size of the precipitate was ca. 3 nm. It is interesting that, apart from this study, quantum size effects have not been described in CD PbS films, in contrast to PbSe ones. Although PbS does show weaker quantum effects than does PbSe (because of its larger effective mass), it still should show strong quantum size effects—greater than CdSe, for example. For some reason, PbS seems to grow with larger crystal size than many other semiconductors. However, there is no a priori reason to indicate that size-quantized PbS could not be deposited by CD, and it is likely that an effort to do so would bear fruit.

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The second study of possible relevance reported that PbSe, precipitated from selenosulphate solution (not in the form of a film), was found to have an (electrical) bandgap, measured by temperature-dependent resistivity, of 0.4 eV [48]. In the same study, samples prepared by reaction of solid lead tartrate with H2Se exhibited an electrical bandgap of 0.92 eV. These results suggest the occurrence of size quantization.

10.5 OTHER SEMICONDUCTORS Quantum size effects have been noted, either explicitly or implicitly (in some cases, not noted in the study but inferable from the optical spectra), in a number of other CD semiconductor films. In most cases, the quantum size effects are a side issue and are only briefly mentioned. For that reason, the results will be given in Tables 10.4 and 10.5, with additional information given for specific studies where available.

10.5.1 ZnS In Ref. 49, the bandgaps of these ZnS films were dependent on the deposition temperature and varied from 4.05 eV (3°C) to 3.88 eV (90°C) (bulk cubic bandgap
3.6 eV). The crystal size measured by TEM for the 90°C deposition was 6–8 nm, in the range, if somewhat large, for the observed bandgap increase in ZnS. For an acidic bath (no ammonia and higher Zn concentration), the bandgap was the 3.6 eV of bulk cubic ZnS. No mention of quantum size effects was made in the study in Ref. 50, and in fact the measured bandgap (3.76 eV) was reported to agree with the literature value. The literature value meant was probably that of the hexagonal form of ZnS, which has a somewhat higher bandgap (ca. 3.8 eV) than the cubic form. However, electron diffraction showed the cubic form was deposited. Therefore this value is indicative of size quantization. In Ref. 51, the films were reported to contain significant amounts of hydroxide or oxide. This is a common property of CD ZnS films deposited from alkaline solution (see Chap. 3). In Ref. 52, no XRD pattern was observed in the as-deposited films. Annealing at 450°C resulted in a bandgap of 3.7 eV (and conversion of some of the ZnS to ZnO). This value suggests that hexagonal ZnS is formed (at least, after annealing).

10.5.2 ZnSe In Ref. 53, from the optical absorption spectrum, a slight blue shift (2.63 eV compared to bulk 2.58 eV) was explained by size quantization. Note that XRD of the precipitated powder suggests a crystal size of ca. 3 nm (no pattern was discerned

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TABLE 10.4 Possible Quantum Size Effects in other CD II-VI Semiconductors. Max. Eg (eV) ZnS 4.05 (3°C) 3.88 (90°C) 3.76 3.85 (85°C) 3.9 ZnSe 2.63 2.7 2.9 HgS ca. 2.3 (0°C) 1.9 (85°C) see text HgSe 1.42 2.5

Min. crystal size (nm) —

Miscellaneous Thioacetamide/ammonia

6–8 3

ca. 3 nma ⬇10 2–2.5

3–4 8 20 No XRD seen 7.7

Thiourea/ammonia/hydrazine Thiourea/ammonia/amines Thioacetamide/triethanolamine/ ammonia DMSb/citrate/ammonia/50°C Selenourea/ammonia/ hydrazine/60°C Selenourea/ammonia/ hydrazine/70°C

Refs. 49 49 50 51 52

53 54 55,56

Thiosulphate

57

Thiosulphate/ammonia

58

Selenosulphate/Hg0formamide-NaOH Selenosulphate/ammonia/10°C

59 60

ZnS (bandgap 3.8 eV hexagonal and 3.6 eV cubic), ZnSe (2.58 eV), HgS (2.1 eV cinnabar; zero-gap metacinnabar), and HgSe (0.15 eV) a From precipitate. b N,N-dimethylselenourea.

in the film, possibly due to the small amount of material together with the small crystal size). Such a small size would be expected to result in a larger bandgap. No reference to size quantization was made in Ref. 54. However, the bandgap is somewhat larger than for a nonquantized sample. The crystal size (ca. 10 nm from both electron microscopy and X-ray diffraction) is a bit large to show size effects in the spectra. In Table 10.4, there is some degree of inconsistency between the various bandgaps and crystal sizes, if it is assumed that size quantization determines the bandgap. It may be that other factors influence the bandgap here. In Refs. 55 and 56, the Zn:Se atomic ratio of 3:2 found in the films, together with the TEM micrographs, which show very small ZnSe crystals, suggest that size-quantized ZnSe is formed together with ZnO [or Zn(OH)2].

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TABLE 10.5 Possible Quantum Size Effects in Miscellaneous CD Semiconductors Max. Eg (eV) Sb2S3 2.48

Min. crystal size (nm) —

Miscellaneous

Refs. 67

1.7 2.21

⬇20 —

1.79 Cu-S CuS—2.0 CuS—1.55 Cu1.8S—1.55 Cu1.96S—1.4 Cu-Se 2.38(d) 1.9(i) 2.20(d) 1.3(i) In2S3 2.7 (50°C) Ag2Se 1.8 direct CuInSe2 1.4 (direct)

⬇20

Thiosulphate/SbCl3 in CH3COOH Annealed at 170°C Thiosulphate/SbCl3 in acetone/ 10°C/pH 5 Annealed at 250°C

⬇13 11 (200°C) 19 20

Thiosulphate/pH 5/50°C Thiosulphate/dimethylthiourea Annealed at 300°C Annealed at 400°C

69 70

20 29

Selenosulphate/ammonia/25°C Annealed at 300°C (Cu1.85Se)

71

ca. 6

Thioacetamide/pH 1–3.1 (HCl)

75

Selenosulphate/ammonia

77

Selenosulphate/ammonia/ citrate/40°C

79

9 —

1.15 (520°C) 1.3 direct

20 14

1.02 (300°C)

23

Selenosulphate/citrate/ pH 8/25°C

67 68 68

79 80 80

10.5.3 HgS In Ref. 57, since the film thickness increased with deposition temperature [from 50 nm (0°C) to 180 nm (85°C)], the decrease in bandgap with increase in deposition temperature may also be due, to some extent, to this factor. From the weak XRD pattern (shown for a deposition at 28°C), the HgS appears to be the cinnabar crystal form with a bulk bandgap of 2.1 eV. However, this would not explain the bandgap of 1.9 eV measured for the 85°C deposition. A similar deposition, only with the addition of ammonia, was carried out [58]. This is a major change (e.g., alkaline instead of acidic conditions). While a high bandgap (3.1 eV) was measured from the sharp drop at 400 nm in the optical

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transmission, this is not likely to be a bandgap transition, for several reasons. One is the color of the films—golden yellow (100 nm thick) to red (500 nm). This color was seen also as a more gradual drop in the transmission beginning at ca. 600 nm. (An even more gradual reduction in transmission occurred to longer wavelengths and is probably related to small amounts of lower-bandgap modifications found in the films). Second, the XRD spectrum of the films indicated a crystal size of at least 20 nm, making it unlikely that the transition at 3.1 eV was due to size quantization of cinnabar with a bulk bandgap of 2.1 eV.

10.5.4 HgSe HgSe is normally a semimetal with a negative bandgap. However, from the two studies noted next, either the bulk bandgap of these films is considerably larger, the quantum size effect is huge, or the measured bandgaps in all cases were higher transitions rather than the primary bandgap. More studies with careful investigation of the optical properties are needed here. In Ref. 59, no XRD pattern was observed for the as-deposited films, which were thus assumed to be amorphous. Annealing at 200°C was enough to give an XRD pattern (of hexagonal HgSe), although no crystal size details were given. In Ref. 60, the HgSe was probably in the form of the cubic phase: The precipitated powder was cubic, while the film was highly textured (111) [which could also belong to the hexagonal (0002) reflection]. The very large value of bandgap found (2.5 eV) was compared with the value of ca. 3.2 eV previously found for 2to 3-nm HgSe colloids [61].

10.5.5 Bi2S3 There are many reports on CD Bi2S3 where larger-than-normal bulk value (1.3 eV) bandgaps are obtained. The section on Bi2S3 in Chapter 6 should be read for those interested in quantum size effects in this material. Table 6.2, on Bi2S3, gives values of reported bandgaps, and values of between 1.6 and 1.9 are particularly common. However, as pointed out in that chapter, such values should be treated with caution; most of the bandgap values are calculated from an extrapolation of the plot of (h)2 vs. h to (h)2
0 (see Sec. 1.4.2), and consideration of the raw transmission spectra, where they are given, strongly suggests that the absorption onsets are often considerably red-shifted from this calculated value. This is particularly true for films deposited from acidic solutions, which have a larger crystal size than those deposited from alkaline solution (see later). X-ray diffraction patterns, with peak widths characteristic of a crystal size of at least 10 nm and usually 20 nm or more, have been obtained from Bi2S3 films deposited from acidic solutions. In contrast, XRD has not shown any peaks for asdeposited films from alkaline baths, but annealing at low temperatures (150–200°C) was sufficient to give fairly sharp (equivalent to at least 20-nm crys-

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tal size) XRD peaks. However, the optical absorption spectra changed only little, if at all, after this annealing (see, e.g., Refs. 62 and 63). This provides further evidence that the optical spectra do not reflect size quantization in general. Taking all this into account, there seems to be no merit in including Bi2S3 in the present tables. Reported bandgap values have already been given in Table 6.2, and crystal sizes have been summed up in the previous paragraph. Instead, some remarks on specific studies will be made. In Ref. 64, red spectral shifts were obtained upon illumination of the films during deposition. It was thought that this was due to increasing film thickness as a result of the illumination. However, consideration of the raw transmission spectra in this study does indicate a change in bandgap with varying illumination conditions (intensity and particularly duration). Estimation (by this author) of the bandgap from the transmission spectra gives ca. 1.9 eV (no illumination) to 1.5 or less eV under full solar illumination during the entire deposition. Similar illumination-induced crystal growth with resulting increased bandgaps has been discussed for CdSe earlier in this chapter. Reference 65 does relate in some detail to size quantization in the Bi2S3 films. The crystal size was reported to increase and the bandgap to decrease with increasing film thickness. The deposition was from acidic solution, and consideration of the XRD pattern shown indicates a crystal size of at least 20 nm. Slow-scan XRD was interpreted to give a range of crystal sizes from 5.2 nm (for a 53-nm-thick film) to 8 nm (220-nm-thick), which is not in agreement with the (normal scan speed) spectra shown. Bandgaps as high as 2.22 eV were calculated from the (h)2 vs. h plots, but these do not match the values expected from the optical absorption spectra, which indicate much lower values. A re-estimation from these (h)2 vs. h plots, but at lower values of absorbance (
104 cm1 rather than the range closer to 105 cm1 that was used), gives better linearity and a gap of ca. 1.4 eV, fairly independent of film thickness and quite close to the literature value.

10.5.6 Bi2Se3 There are very few studies on CD Bi2Se3 (see Chap. 6) and only one that provides a reasonably in-depth study of the optical properties [66]. In this work, N,Ndimethylselenourea was used as the Se source in an alkaline triethanolamine bath. The spectrum of a thin film (90 nm thick) was quite different than the spectra of thicker ones (150 or 250 nm, which were similar). The 90-nm film exhibited a clear onset with a direct bandgap, calculated from the (h)2 vs. h plot, of 1.7 eV. (The literature room-temperature value is ca. 0.12 eV.) Annealing at 200°C caused a red shift in the spectrum to a calculated bandgap of 1.57 eV. For the thicker films, the as-deposited direct bandgap was 1.41 eV and after annealing, 1.07 eV. For these thicker films, however, there was also loss of transmission at lower energies. Both scattering and an indirect bandgap were suggested as possi-

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ble causes of this. Based on the latter possibility, indirect bandgaps between 0.2 and 0.35 eV were calculated for this region of the spectra. X-ray diffraction of these films showed some sign of crystallinity and a very rough estimation of
10 nm could be made from this spectrum. Annealing at 200°C increased the crystal size by only a few nanometers. A rough estimation of the increase in bandgap expected for a 12-nm crystal size (based on a literature value of the reduced effective mass of
0.06) is ca. 0.15 eV. This suggests that the apparently indirect gap measured for the thicker films is somewhat quantized and that the direct gap measurement is a higher transition. The higher value for the thin film may be a large size quantization effect, but information on crystal size is needed.

10.5.7 Sb2S3 In Ref. 67, the increase in bandgap of the as-deposited film was attributed to a mixed amorphous-polycrystalline structure (apparently no XRD pattern was found for the as-deposited film). The onset of absorption in the transmission spectrum was sharp for the as-deposited film, and any polycrystalline (nonquantized) material would be expected to give some absorption at lower energies, even if the amorphous phase possessed a higher bandgap. Therefore size quantization seems to be a more reasonable explanation for the high bandgap of the as-deposited film. In Ref. 68, the higher bandgap of the as-deposited films (2.21 eV vs. 1.79 eV for annealed films) was explicitly explained by size quantization.

10.5.8 Cu-S Bandgap measurements for Cu sulphides and selenides are complicated by the fact that these semiconductors are normally degenerate, with high free-carrier absorption in the near-infrared and possible Moss–Burstein shifts (due to saturation of the top of the valence band by holes) in the optical gap. It is quite possible that variations in bandgaps in these materials are due to differences in stoichiometry, phase, and doping rather than to any quantum size effect. Only studies where crystal size can be estimated and are possibly in the quantum size range are given here. In Ref. 69, various compositions were obtained by varying the Cu:thiosulphate ratio. Only CuS gave an XRD pattern that allowed an estimation of crystal size; other compositions and phases (Cu1.8S, Cu1.4S, Cu2S) showed no XRD patterns. The (indirect) bandgaps found for these films were 2.0, 2.0, and 1.7 eV, respectively. In Ref. 70, small increases in bandgap were found for Cu1.8S (1.55 eV; literature value 1.2) and Cu1.96S (1.4 eV; literature value ca. 1.3). The apparent discrepancy between the bandgaps of CuS (both indirect bandgap values) found from these two studies should be noted. The former found a considerably larger value, even though the crystal size was slightly larger and the film was not annealed. A comparison of the two transmission spectra does not

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suggest such a large difference, and the difference may be due simply to the problems in measuring bandgaps of Cu-chalcogenides discussed earlier.

10.5.9 Cu-Se Reference 71 is the only study that explicitly invokes size quantization for Cu-Se films. This was based mainly on the decrease in bandgap found upon annealing. In the absence of effective mass values for the various compounds, it can only be said that the 20-nm size measured for the as-deposited films is a bit large in general to explain the observed spectral shifts, but not decisively so. It is important to note that the electrical resistivity of the films did not change after annealing; this implies that the decrease in bandgap on annealing is not due to a change in the Moss–Burstein band filling, since this would be expected to be paralleled by an increase in resistivity. The same study also described CuSe films, which were obtained when N,Ndimethylselenourea was used instead of selenosulphate. The crystal size in this case was slightly larger ( 23 nm). The bandgap was measured to be ca. 2.13 eV. In view of the difficulty in measuring bandgaps in these materials, this is close to the value of 2.0 eV reported for CD CuSe films with a crystal size of 40 nm, where no quantum size effects are expected [72]. Other studies on Cu-Se where small crystal sizes were measured and therefore are potential candidates to exhibit size quantization give crystal sizes between 10 and 20 nm and bandgaps (direct) of ca. 2.2 eV for Cu2xSe and ca. 2.8 eV for Cu3Se2 [73,74]; a value of 2.37 eV was measured for CD Cu3Se2 with a 40-nm crystal size [72].

10.5.10 In2S3 In Ref. 75, the bandgap increased gradually from 2.32 eV to 2.7 eV as the deposition temperature varied from 90°C to 50°C. The pH also had an effect on the crystal size (at 70°C, Eg was 2.56 eV at a pH of 3.1 and 2.41 at a pH of either 1.3 or 1.5). This could be due to a change of mechanism [In(OH)3 is very acidic and is likely to be present at the higher pH but not at the lower one). Transmission electron microscopy showed the films to be composed of both needles and round particles. At 90°C, the round particles were 8–10 nm in size and the needles 20–30  2 nm. At 50°C, there were fewer needles and the round particles were ca. 6 nm. X-ray diffraction showed a larger particle size: 15 nm (90°C), 13 nm (70°C), and 10.6 nm (50°C). The large apparent size estimated from XRD, and the fact that annealing, which reduced the bandgap (see later), did not narrow the XRD peaks, suggested that the XRD spectrum was weighted toward a small fraction of larger crystals. Eg decreased gradually on annealing from 2.7 eV (50°C sample) to 2.53 eV (350°C anneal). The increase in Eg was ascribed to a combination of size quantization and excess S (measured by EDX analysis).

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The original CD of In2S3, by the same method as earlier [76] noted a crystal size that was probably no larger than 5 nm, although the bandgap was measured to be 2.45 eV, similar to the direct bulk gap of In2S3. A lower energy weak absorption which was seen in single-crystal material was absent in the films. In addition, the photoconductivity spectral peak (of a film annealed at 250°C) was blueshifted 0.13 eV, compared to single-crystal material.

10.5.11 Ag2Se In Ref. 77, the (direct) bandgap of 1.8 eV was compared to the bulk value of ca. 1.3 eV bulk and the difference ascribed to size quantization. The absorption spectrum shows considerable absorption to beyond 820 nm (the maximum wavelength measured), and it is not clear whether this is due to scattering or to a lower-energy (possibly indirect) absorption.

10.5.12 CuInSe2 Two cases of what are probably size quantization can be found in CD CuInSe2 films. In neither case was size quantization used to explain the larger-than-normal (1.02 eV) bandgaps. It is notable that these examples appear to be the only ones showing this phenomenon for CuInSe2 in any form. One example of an anomalously high bandgap—1.38 eV—was found for CuInSe2 films sputtered onto glass at 77 K and was attributed to the amorphous structure; i.e., no XRD pattern was observed [78]. This could also be due to size quantization. The bandgap of these sputtered films was calculated from a plot of (h)n vs. h, where n
1 (amorphous semiconductor). If replotted for n
2 (direct bandgap, crystalline semiconductor), an even higher bandgap would result. In Ref. 79, as-deposited films exhibited a bandgap of 1.4 eV, which dropped on annealing at 110°C to 1.225 and then fell more gradually to a final value of 1.15 eV at 520°C. Crystal size measured by XRD was given only for the 520°C annealed film (20 nm; very small for that temperature of annealing), which means the crystal size of the as-deposited film was probably much smaller. Taken together with the following study, this seems to be a clear case of size quantization. The main differences in the deposition procedure between the study in Ref. 80 and the previous study are the somewhat lower temperature, the (almost certainly) lower pH (while not given; the previous deposition, using ammonia, was probably carried out at a minimum pH of 10), and the lack of ammonia in this study.

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