Solution Thermodynamics and Its Application to Aqueous Solutions: A Differential Approach

PREFACE

At the risk of being overly simplistic, the last century was the century of quantum mechanics. A handful of geniuses, driven just by curiosity, have learned how atoms and electrons behave, how an atom is made up, how a nucleus is held together, and so on, and constructed what we know as quantum mechanics. While it was developed purely for the fulfillment of human curiosity, it has led to its application to various technologies to the extent unimaginable to early mankind. Examples include the usage of nuclear power, the invention of transistor and various electronic devices, and drug design, to name but a few. How many of these geniuses who served an essential role in bringing in modern “conveniences” (and hence huge businesses) would have dreamed of the results of their endeavor? A philosophical spin-off of quantum mechanics, on the other hand, is that human beings have learned to look into details at the nanoscopic level of constituent elements. After all, we live in a macroscopic world which consists of components, sub-components, sub-sub-components and so on down to the elementary particles. And we are now good at, or at least comfortable in, studying the nature of what we believe to be the smallest components in detail. However, we all know that the nature of an assembly is not just the sum of the properties of all the constituent elements. Rather, each assembly has the specific extra characteristics due to the interactions among elements – the so-called “Many-Body Problem”. Recognition of such characteristics of assemblies is more common in social sciences, a theory in ecology based on Holism (Merchant, 1992), for example. In science, there is an effort to extract hitherto unknown laws purely by observing the peculiar characteristics of various many-body problems including biological systems and “life” itself. (Makishima, 2001). Thermodynamics, developed in the 19th century, on the other hand, deals with the macroscopic world. A number of geniuses, driven equally by curiosity, extracted the laws of thermodynamics. Since it deals exclusively with macroscopic phenomena, recognition of the many-body nature is inherent, as typified by the common use of the excess thermodynamic functions in solution thermodynamics, excess over the sum of the thermodynamic functions of all constituents. At this point in time, therefore, I believe it is worthwhile to revisit thermodynamics in order to tackle the many-body problems, since quantum mechanics is not so well designed to do so. However, conventional usage of thermodynamics is basically phenomenological in nature and generally lacks the power to elucidate underlying mechanisms. For the last 20 years or so, we have been making an attempt at improving this point with some success. Our approach is purely experimental.

xii

In contrast to conventional thermodynamic studies, we make an effort to obtain thermodynamic quantities that are proportional to at least the second and possibly the third and higher derivatives of Gibbs function. We introduce the interaction functions and the partial molar fluctuations, both being the third derivatives. Without resorting to any model system, we learn the detailed thermodynamic behavior of the system in question, and hence the underlying mechanisms, directly from the composition or the temperature dependence of these higher order derivative quantities. With this methodology, we learn the mixing schemes or the “solution structures” in aqueous solutions in more detail than hitherto possible. In some cases, the findings by the present novel approach to solution thermodynamics more than complement those by other techniques including modern spectroscopic studies. In particular, the power of the present methodology in elucidating the molecular processes in dilute aqueous solutions is remarkable, as will be demonstrated throughout this book. This book has emerged from notes given to new students in our research group. All of them have had some exposure to the first formal introduction to thermodynamics in their undergraduate courses. Thus, it starts off with a brief review of the basics, followed by the development suitable for our approach. Then, what we learned about aqueous solutions using our thermodynamic methodology is reviewed. As appropriate, the findings about the mixing scheme or the “solution structures” by other methods are compared. Thus, the main part of this book is still growing and may necessarily be revised later. I wish to dedicate this book to my family of the past, the present and the future. For me to reach the present point in my career, I owe a great deal to my mentors, collaborators, postdoctoral fellows and students. I list their names on the following pages in an approximately chronological order. I wish to thank each and all sincerely. The work has been mainly supported by the Natural Science and Engineering Research Council (NSERC) of Canada. Although there has been a continuous bureaucratic reluctance, I thank NSERC all the same. Vancouver January, 2007 Yoshikata Koga

ACKNOWLEDGEMENTS

I thank:Shoji Makishima Charles A. McDowell Hiroshi Suga Akira Inaba Loren G. Hepler Chris G. Vonk Robert M. Miura Thierry Michael Rita Kasza Adam Chong Margaret Slater L. Le Joan S. C. Kam Frankie Lau Catherine Hui Christa Trandum Kathy Puhacz Mark Nitz Mark Klippenstein Yvonne Yau Betty Lam Anu Saini Ichiro Fujihara Candy Cheung Ami Hui Shigeomi Takai Daniel Chen Jesper Hansen Hiroshi Matsuo Denise Wong Pui Ming Chu

Eiji Ootsuka Hideaki Chihara Tooru Atake G. C. Benson Peter Westh Ichiro Arakawa Sachio Murakami William W. Y. Siu Lisa C. F. Chao Kirby Wong-Moon Mahiko Nagao Raymond Huie Yingnian Xu Stephen Cheng Heidi Lam Hikari I. Yoshihara Alice Ho Steven Tanaka Ken Suh Jerri Lai Sandy Hundall Paul Yenson Ole G. Mouritsen Yet Chow H. Kawaji K. Kawasaki Eric To Yoshihiro Taniguchi Jianhua Hu Wesley Chiang Yasutoshi Kasahara

Lionel G. Harrison James A. Morrison Keiko Nishikawa Aase Hvidt Jesper Kristiansen R. Troy Lassau Katsutoshi Tamura Terrance Y. H. Wong John T. W. Lai David Fraser Anselm Wong Laurence Beach Jenny Chuang Johny Shim Damon Robb Susan Chan James V. Davies Virginia Loo Arezoo Nasiry Paul Cheung Rena Kare Kazuko Mizuno Kenneth Chong George Makris T. Tojo Y. Ootsuka Christian Svane Seiji Sawamura Charles A. Haynes Jason Sze Matthew T. Parsons

xiv

Kim B. Anderson Jonathan Lau Mark Chen Kitty Chan Hideki Katayanagi Mette N. Kristensen Naoki Toshima Aurelien Perera Ken-ichi Tozaki David Siu Pang Hiew

Eric Yee Kevin Ralloff Amy Wu Harris Huang Hideki Shimozaki Hitoshi Kato Yukihide Shiraishi Franjo Sokolic Yasutaka Kido Lion Kit

Allen Liu Karin Liltorp Andrew Racaza Takehiko Ichioka Kumiko Miki Fumie Sebe Masao Kanemaru Laszlo Almasy Lori Leong Kenneth Wong

Chapter 0

INTRODUCTION

[0-1] INTRODUCTION As the title implies, this book describes an application of thermodynamics to solutions – multi-component systems. The original formulation of thermodynamics for multi-component systems was made by Gibbs (1993) and we basically follow his derivation here. We then emphasize the importance and usefulness of higher order derivatives of Gibbs energy, G, than have hitherto been utilized. In particular, we show that the second and third derivatives of G, if determined within reasonable accuracy, are powerful in providing information towards molecular level understanding of the “mixing schemes” or the “solution structures” of aqueous solutions. At this point in time, the response functions (heat capacity, compressibility, thermal expansivity), the second derivative quantities, can be experimentally determined with a high accuracy. At the risk of over-generalization, however, it is common practice to integrate these measured second derivative quantities. The resulting very accurate thermodynamic quantities that are first derivatives of G are then compared with theoretical predictions. We argue that this practice is a waste of the qualitative information contained in the second derivative quantities themselves. For example, a heat capacity, Cp , provides information regarding the entropy fluctuation of the system, while its integrated quantity, the entropy, gives the global average. We suggest that the information contained in the second derivative quantities be taken advantage of by itself. Our originality is that we further take one more derivative of the measured second derivative quantities graphically without resorting to any fitting function. We have thus model-free, experimentally accessible third derivative quantities. Of course, with this treatment the resulting third derivatives are not without increased uncertainties. As shown below, they could amount to several per cent with the measured second derivative data determined within 0.1%. However, this quantitative disadvantage is well compensated for by the gain in higher quality information which is powerful in elucidating the molecular processes in aqueous solutions. We will demonstrate this throughout the book. The importance of H2 O and aqueous solutions on this planet earth requires no emphasis. A vast amount of effort has been invested in understanding the nature of aqueous solutions, yet we are still far from a complete understanding, if such is at all possible. We believe this new approach to solution thermodynamics adds, and will continue to add, some new insights into this perpetual human question.

2

Fig. 0-1

[0-2] QUALITATIVE JUDGMENTS AND SCIENCES Natural sciences are believed to be quantitative. Or are they? Let us consider the process of determining the mass of a sample, the most fundamental quantity in chemistry. We are given a balance and a set of standard weights. We place a sample on the lefthand side of the balance, and we discover that the left side is heavier, as shown in Fig. 0-1(a). A 10 g standard weight is placed on the right-hand pan, and we judge that a 10 g standard is too heavy for the sample, Fig. 0-1(b). We replace the 10 g standard with two 1 g weights, and realize the sample is heavier than 2 g, Fig. 0-1(c). We place another 1 g standard on the right pan, and learn that 3 g is too heavy, Fig. 0-1(d). Thus we conclude that the mass of the sample is between 2 and 3 g. These processes are repeated m times using 0.1 g standards until we can conclude that the mass of the sample is 2 + m − 1 × 01 g < mass < 2 + m × 01 g. We continue until a required uncertainty within which we wish to know the mass of the sample. As this example hints, the mass is determined within the required accuracy by a series of qualitative judgments as to

(b)

(a)

10 g

(c)

2g

(d)

3g

Fig. 0-1. Weighing a sample using a balance and standard weights, 10 g, 1 g, 0.1 g, 0.01 g    .. Weighing continues while a qualitative judgment is made as to which way the balance tips until we know the range of the mass within the required accuracy. See the text, Section [0-2].

3

which way the balance tips. Here we assume that the balance is sensitive enough for the required accuracy and that a set of standard weights are available down to the required sensitivity. As we see below, qualitative judgments are prevalent throughout this book.

[0-3] OUTLINE OF THIS BOOK In Part A, we discuss a differential approach to solution thermodynamics. It is new only because the second and third order derivatives of G are used as the basis of qualitative judgment as to what the thermodynamic quantities imply. At first, we briefly review the foundation of thermodynamics. To this end, we basically take the same attitude as Gibbs (1993), Hashizume (1981) and Callen (1985), whereby the existence of entropy and its nature are taken for granted, and build thermodynamics relations analytically. However, since a much smoother and more straightforward introduction of entropy is available by the “order-theoretical” approach (Giles, 1964; Lieb & Yngvason, 1999), we very briefly introduce Giles’ approach for entropy, sacrificing the rigor of the original. We then follow in Chapter II the derivation of the second derivatives of G – partial molar quantities, and the third derivatives – intermolecular interaction functions. Their physical meanings are briefly discussed in Part A, but their actual power in elucidating the nature of mixing schemes or “solution structures” is demonstrated in Part B, where these quantities are evaluated for actual aqueous solutions. For some unknown reasons, the experimental determination of the partial molar quantities, the other second derivative quantities than the response functions mentioned above, has been largely ignored, though the technical capability exists. Hence we devote Chapter III to the experimental methods of determining the partial molar quantities. Also included in the same chapter is a graphical differentiation to obtain the interaction functions – the third derivatives. A computer graphic technique, the “cubic B-spline”, could be used for this purpose. The present state of the cubic B-spline technique is discussed in Appendix A. Chapter IV deals with other second derivatives – response functions. We introduce the fluctuation functions that could provide qualitative information about the intensity (amplitude) as well as the extensity (wavelength) of fluctuations. We take one more derivative and introduce the partial molar fluctuation functions, the effect of additional solute on the fluctuation characteristics of the system. Part B describes our studies up to the present time on aqueous solutions using the second and third derivative quantities, and thus demonstrates their power in elucidating the mixing schemes. Chapter V deals with the aqueous mono-ols to which we have applied our methodology so far, while Chapter VI is about the aqueous solutions of some other non-electrolytes. In these two chapters the thermodynamic signatures and hence the mixing schemes of hydrophobic and hydrophilic solutes are discussed. These understandings, however, become deeper with our “three component 1-propanol probing methodology” described in Chapter VII and Chapter VIII. The concept of partial molar quantities is most useful in multi-component systems in that they refer only to a target component even in the presence of many other components. Namely, it is only a

4

target component that is required to be perturbed. Furthermore, the special nature of H2 O dictates that the solution in the H2 O–rich composition range is governed by the interactions of each solute via bulk H2 O, rather than by direct contacts among solutes themselves. This led to our “three component 1P-probing methodology” in which we use thermodynamic properties (in terms of third derivative quantities) of 1-propanol (1P) in a mixed aqueous solution including a third component in the 1P – a third component – H2 O system. We use the induced changes in the third derivative quantity of 1P by adding a third component as a probe to learn the effect of the third component on the molecular organization of H2 O. We apply this methodology to elucidate the effects of amphiphiles on H2 O. We learn that the hydrophobic and the hydrophilic moieties of an amphiphile compete or cooperate towards interaction with H2 O. The net result is that the effects of each moiety manifest themselves competitively or cooperatively in the behavior of the third derivative quantities. We thus establish a hydrophobicity/hydrophilicity scale within the methodology. In Chapter VIII, we apply the same 1P probing methodology to some electrolyte salts, and learn the effects of cations and anions on the molecular organization of H2 O. The effects of ions on the structure and the function of biopolymers in aqueous solutions have been studied extensively. As a result, ions have been ranked as the Hofmeister series (Hofmeister, 1887) according to the propensity and the strength of the effect of each ion. The generality of this ion ranking for a wide range of processes has sparked a massive interest in the possible underlying mechanism. However, the molecular level understanding of the Hofmeister effect is still fragmentary. We make an attempt at sorting out the ranking in terms of the hyrophobicity/hydrophilicity scale within our methodology. In addition, we show that our method could identify some ions as simple hydration centers without altering the nature of bulk H2 O. The hydration number of these ions could be evaluated by this method. We show for such ions that the Hofmeister ranking is consistent with the Hofmeister’s original idea, “water withdrawing power” (Hofmeister, 1887). In Chapter IX, we discuss the treatment by the present methodology of the three component systems in general. We introduce the heterogeneous interaction functions between two different species in the second and third derivative levels. We show that the heterogeneous interactions also occur via bulk H2 O competitively or cooperatively but not linearly additively. We then provide the executive summary of our findings in the last chapter, Chapter X. We had to omit from this book our recent unfinished works using the 1-propanol probing methodology on the nature of aqueous solutions of “ionic liquids” and that of aqueous “osmolytes”. Both works will be dealt with in a future revised edition. As appropriate, we review and compare recent findings about the mixing schemes, or the solution structures, in aqueous solutions by other techniques; spectroscopic, scattering, molecular dynamics and modern ultra fast techniques. We demonstrate that the present thermodynamic methodology is just as powerful as, if not more than, those modern techniques. Throughout this book, all the figures, tables and equations are numbered in Arabic numerals preceded by a Roman number denoting the chapter. Thus, eq. (III-21) is the 21st equation in Chapter III. However, when cited within the same chapter, the first Roman

5

numerals are omitted. We make frequent references to equation and figure numbers in the text. For the convenience of a reader, we make entries of equation and figure numbers on the top of each page. When the table is referred, its page number is given. As for any other thermodynamic discussions, we use partial differentiations frequently. When obvious, the lists of variables that remain constant are often omitted. Readers are asked to interpret that all the variables other than those of differentiation are kept constant. Indeed, we exclusively dwell in the Gibbs variable system (p, T , ni ), and it should be obvious which are kept constant.

Eq. (I-1) to (I-3)

Chapter I

BASICS OF THERMODYNAMICS – DERIVATIVES OF GIBBS ENERGY, G

[I-1] SYSTEM, STATE, FIRST AND SECOND LAWS OF THERMODYNAMICS Consider a “system” surrounded by a “surrounding”. A system can be “open” or “closed” depending on whether it can exchange matter with the surroundings or not. An “isolated” system cannot exchange “energy” with the surroundings. Here we accept “energy” as one of the fundamental quantities in describing the “state” of a system. Historically, however, careful observation and precise logical deduction by a number of geniuses on macroscopic phenomena regarding heat and work acting on a closed but not isolated system have led to the discovery of the energy or the internal energy, E, described by equation, dE = q − w = qrev − wrev 

(I-1)

This is the well known first law of thermodynamics. Eq. (1) is the result of the realization that the heat received by the system, q, and the work done on the surroundings by the system, w, are equivalent and that the combination of the two determines the change in the internal energy, E, which is a state function. Among an infinite number of possible combinations of the values of q and w resulting in the same value of dE, there is the unique value, qrev and wrev respectively, such that qrev  ≥ q and

wrev  ≥ w

(I-2)

Namely, the absolute value of qrev takes the maximum, and so does that of wrev under the special situation called the reversible, or the quasi-static process. This unique value of qrev defines another thermodynamic state function, entropy, S, by eq. (I-3) – the second law of thermodynamics. dS =

qrev ≥ 0 T

(I-3)

In the advent of statistical mechanics, we learned that S represents the randomness of the system. While this description based on statistical mechanics is easier to understand

10

Fig. I-1

Surrounding

System closed, not isolated

w work out

dE q heat in

dE = q - W = q i - Wi etc. = qrev - Wrev

Fig. I-1. A closed but not isolated system receives heat q from, and gives work w to the surrounding. As a result, the internal energy of the system, E, changes by dE = q − w. For a given value of dE, there are countless combinations of the values of q and w. Out of these combinations, there is the unique combination, qrev and wrev , such that the absolute value of qrev or wrev is the largest, pertaining to the reversible process, i.e. the interaction between the system and the surrounding takes place “reversibly”.

than the original derivation by Clausius leading to eq. (3) (Kestin, 1976), thermodynamics must be self-contained within a macroscopic description. Below, we very briefly introduce Giles’ method (1964) of introducing entropy, S, which we find is much more straightforward than Clausius’ original derivation (Kestin, 1976), without extensive use of gedanken experiments on heat engines. Indeed, in textbooks and lectures on thermodynamics, chemistry students are exposed for the first time to careful consideration about heat engines only at the beginning. Once the entropy is introduced, heat engines are completely forgotten and the difference in thermodynamic functions between two states then takes center stage throughout. For example, the Gibbs free energy difference is emphasized between 2 mole of H2 and 1 mole of O2 gas put together without mixing them and their gaseous mixture, or between the latter gaseous mixture and 2 mole of H2 O. Using Giles’ approach we avoid discussion on heat engines, as shown below. The second inequality of eq. (3) means that the system completely isolated can only become more disorganized spontaneously; otherwise, dS = 0, the system is in equilibrium. This will lead to the stability criteria of various thermodynamic functions discussed in Section [I-4]. If we limit our attention to the case where work is done only by volume expansion, wrev = p dV

(I-4)

Eq. (I-4) to (I-6d)

11

where p is the pressure of the surroundings and V is the volume of the system. Equations (1), the first half of (3) and (4) yield the first law of thermodynamics, dE = T dS − p dV

(I-5)

This is the most fundamental relation in thermodynamics and is the starting point for further discussion.

[I-2] GILES’ DERIVATION OF ENTROPY [I-2-1] State, process, and irreversibility of process We follow primarily the route of the “order–theoretical” approach, introduced by Giles (1964), but later more rigorously treated by Lieb and Yngvason (1999). Consider a closed and isolated system, A, that could take a state labeled as a, b, c    . Consider, for example, a system consisting of 2 mole of H and 1 mole of O atoms. The system could be in a state labeled as a, which consists of 1 mole of H2 and 1/2 mole of O2 . The same system could be filled with a mole of H2 O vapor, which is called state b. Since the system is isolated and closed, the total amount of H and O remains the same as well as the internal energy E. To ensure that there is no work done on the system in addition to no heat exchange, the volume of the system should be the same, V . We now consider a process a→b. Suppose we have some means of judging whether such a process indeed occurred or not. We do not care what happened during the course of the process, as long as E and V and the total amount of H and O are all kept the same at the beginning and the end of the process. We could thus add a pinch of catalysts into the system and remove them at the end without disturbing the above conditions. We can summarize the results of the observation by the following statements: a reversible

if a→b

and

b→a

b natural irreversible

if a→b

and

b −→ a

NOT

c anti-natural irreversible if a −→ b and d impossible

NOT

if a −→ b and NOT

NOT

b→a NOT

b −→ a

(I-6a) (I-6b) (I-6c) (I-6d)

Here → denotes the process that actually occurred, while −→ means never materialized. We now consider the association of two systems, A and X, with the resulting compound system A + X, and they take states, a, x, and a + x. The two systems are sitting side by side, but not in direct contact with each other, within the same surroundings. They do not exchange E, V , or matter between them. We can imagine x being the state

12

in which 1 mole of H2 and 1 mole of S are enclosed in system X, and in state y 1 mole of H2 S is present. Suppose we know that a→b, and x→y are both natural irreversible, NOT NOT namely, b −→ a, and y −→ x separately. For convenience of further discussion, we define process  ≡ (a, b), and  ≡ (x, y). “≡” signifies the definition. Assume that both  and  are natural irreversible independently. Let us now consider a new process  −  ≡ (a + y, b + x) for the combined system A + X. If such a process  − . is observed to be natural irreversible, we then conclude that  is more irreversible than . If we assign a value to each process representing its “irreversibility” as, I, I etc, we can write I > I. If, on the other hand,  −  is found to be anti-natural irreversible, then  is less irreversible than  and I < I. We note in passing by comparing with eq. (6) above that I = 0 for  reversible, I > 0 for natural irreversible, and I< 0 for anti-natural irreversible. With positive integers k and m, we may consider a process k − m, namely, we associate k systems of A and m systems of X, and connect states ka + my and kb + mx. If it is found that k −m. is natural irreversible, then kI> mI, or I > (m/k)I. If we add one more . reversed, i.e. we try the process k − (m + 1), and the result turns out to be anti-natural irreversible, we conclude that (m/k)I < I < {(m + 1)/k}I. With sufficiently large integers of k and m, we could sandwich the value of I with a sufficiently small interval in reference to the irreversibility of a “standard” process .

[I-2-2] Entropy, and potentials – defining T, and p Since  connects two states a and b, and since I is determined within a required accuracy relative to a standard I, we define a state function S for each state such that, I ≡ Sb − Sa

(I-7)

Hence, following (6), we make the following statements as, Sb = Sa if  is reversible

(I-8a)

Sb > Sa if  is natural irreversible

(I-8b)

For an infinitesimally small process, i.e. a ≈ b, we may rewrite eq. (8) as, S = 0 if reversible equilibrium

(I-9a)

S > 0

(I-9b)

if irreversible non-equilibrium

For the above two statements, anti-natural and impossible cases are omitted for obvious reasons. The definition of an equilibrium state is given by eq. (9). Namely, if an

Eq. (I-7) to (I-13)

13

infinitesimal deviation from a state could naturally bring back to the original state, such a state is said to be in equilibrium. We now have a state function S, in addition to E, and V considered above. There is obviously another variable describing the amount of each component in the system, ni . At present, however, we fix ni and ignore its variation. Hence we could consider S to be a function, a differentiable function, of E and V , and write as, S = SE V 

(I-10)

Then its perfect differential is written as,  dS =

S

E



 dE + V

S

V

 dV

(I-11)

E

We define “potentials” following Giles (1964) as,  

S

E

S

V

 ≡ 1/T

(I-12)

≡ p/T

(I-13)

V

 E

We note in passing that if ni is kept as a variable, we could have defined the chemical potential of the i-th component as S/ni  ≡ i /T . Here we are keeping ni constant. Later we will introduce chemical potential, by introducing the variation in ni . With these definitions (12) and (13), eq. (11) is rewritten as, dE = T dS − p dV

(I-5)

Thus we recover eq. (5). The original derivation by Giles (1964) is much more rigorous than the present brief description. However, this brief account will provide the essence of the “ordertheoretical” method of introducing entropy without resorting to the use of heat engines including an unreal Carnot cycle. These heat engines are not familiar to chemistry students, and in most textbooks of thermodynamics, such heat engines are used only once at the beginning for introducing the entropy; in all the remaining applications to chemistry no heat engines are mentioned. Giles’ treatment is not without a problem. The connection of two states via → and the differentiability of S are taken for granted. Lieb & Yngvason (1999) claim, however, that both premises are provable.

14

[I-3] LOGICAL (MATHEMATICAL) DEDUCTION Since eq. (5) can be rewritten as,  dE =

E

S





E dS +

V V

 dV

(I-5’)

S

it follows that  T=

E

S

 

(I-14)

V

which recovers eq. (12) above and 

E p= −

V

 

(I-15)

S

If we invent enthalpy, H, following Gibbs as, H = E + PV 

(I-16)

dH = dE + p dV + V dp = T dS + V dp

(I-17)

it follows using eq. (5) that

The motivation for adding the pV term in eq. (16) is to change the independent variable from V to p, as is evident in eq. (17). Namely, in eq. (5), we change V (together with S and seek the induced change in terms of E, while in (17) we change p infinitesimally. The choice of the independent variable, p or V , is a matter of convenience. For theoretical works, V is more convenient conceptually, while p is the choice of experimentalists. Eq. (17) then implies that  T=

H

S

 

(I-18)



(I-19)

p

and  V=

H

p

 S

We have Helmholtz energy, A, and Gibbs energy, G, thanks to these geniuses. A = E − TS

(I-20)

Eq. (I-14) to (I-25)

15

and G = H − TS

(I-21)

Here again, the TS term is introduced to convert the independent variable from S to T , the convenience of which is immediately obvious. For Gibbs function, G, eqs. (5), (16) and (21) yield, dG = −S dT + V dp

(I-22)

Eq. (22) demonstrates the experimental convenience of G, since its variation is determined by changing T and p, both of which are controllable experimentally. The discussion of the stability criteria in Section [I-4] below shows its further convenience in the (p, T ) variable system. It follows from eq. (22) that, 

G S=−

T

 

(I-23)



(I-24)

p

and  V=

G

p

 T

We see that  H = G + TS = G − T

G

T



 = −T p

2

G/T

T

 

(I-25)

p

We note in eqs. (22)–(24) that G is a convenient function in the system of the independent variables (p T ), and that S and V are the first derivative of G with respect to T and p. Its convenience becomes more convincing below when we discuss the stability criteria in Section [I-4]. We point out from eq. (25) that H is not strictly the first derivative of G. However, it contains the first derivative of G also. Equivalent relations using A, eq. (20), instead of G can be obtained for the variable system, (V T ). For the convenience of experimentalists, however, we limit our attention to the (p T ) variable system. p and T are experimentally controllable and are also both intensive quantities. In Section [I-5], we expand our discussion to the (p T ni ) variable system, where ni is the amount of the i-th component of the mixture – the main subject of this book, the solution thermodynamics. We note that ni is also experimentally controllable.

16

[I-4] STABILITY CRITERIA The inequality of the second law, eq. (3) or (9), above dictates that for an isolated and closed system any spontaneous process produces entropy, dS ≥ 0

(I-26)

This is unconditionally true as long as the system under consideration is isolated and closed. Thus, if a process were to be associated with dS < 0, such a process will never be materialized. When dS = 0 the system is then in equilibrium and nothing apparently changes. If dS > 0, the system then undergoes a spontaneous change. Thus, eq. (26) shows the spontaneity of the process, or the stability of its final state, and is called a stability criterion. It would also be useful if we have some indication of spontaneity in terms of other state functions, E, H, V, and G For this purpose, we combine eqs. (1), (3) and (4) and obtain, dE ≤ T dS − p dV

(I-27)

Therefore, dE ≤ 0 if

dS = dV = 0

(I-28)

dS ≥ 0 if

dE = dV = 0

(I-29)

Since H = E + pV, eq. (27) is rewritten as, dH ≤ T dS + V dp

(I-30)

Therefore, dH ≤ 0 if

dS = dp = 0

(I-31)

dS ≥ 0

dH = dp = 0

(I-32)

if

Since G = H – TS, eq. (30) is rewritten as, dG ≤ −SdT + V dp

(I-33)

Therefore, dG ≤ 0 if

dT = dp = 0

(I-34)

Eq. (I-26) to (I-38)

17

Eq. (34) shows the usefulness of G in the (p T ) variable system, as pointed out above. Namely, we can control p and T to be constant and seek a change in state such that dG ≤ 0. We know then that such a state is stable to any disturbances. As discussed for dS above, the equal sign indicates the equilibrium situation. We point out in passing that the stability criterion for dV cannot be drawn in a meaningful manner. Readers might like to check this out.

[I-5] MULTI-COMPONENT SYSTEM – PARTIAL MOLAR QUANTITIES We now consider a multi-component system (Koga, 2004b). Up to this point the amount of material in the system was left implicit and fixed. Now the independent variables are (p T ni , where ni is the amount of the i-th component, and i = 1, 2, 3,    k, for a k-component system. Thus, we have a convenient Gibbs function, G, as a function of all these variables. G = Gp T ni 

(I-35)

When the independent variables are perturbed, the resulting total differential of G is written as,  dG =

G

T



 dT + pni

G

p



   G dp + dn 

ni pTnj=i i Tni i

(I-36)

The first two partial derivatives on the right of eq. (36) are given in (23) and (24). The third and onward derivatives are called by Gibbs “the chemical potential of the i-th component”.  i =

G

ni

 

(I-37)

pTnj=i

where the partial derivative is taken with respect to ni keeping p T and nj  j = i) constant. The operational meaning of i is the response of the entire system in terms of G when the system is perturbed in ni . Thus, i implies the actual contribution of the i-th component towards the entire system in terms of G or the actual situation of the i-th component in terms of G in the mixture. As is obvious from eq. (37), i is the first derivative of G. By differentiating both sides of eqs. (23)–(25) with respect to ni and using eq. (37), we have  Si ≡

S

ni



 =−

nj=i

i

T

  pnj=i

(I-38)

18

 Vi ≡  Hi ≡



V

ni

 =

nj=i

H

ni

i

p





 = −T 2 nj=i



(I-39)

Tnj=i

 i /T

T

 

(I-40)

pnj=i

They are called, respectively, the partial molar entropy, volume and enthalpy of the i-th component. Similar to i , they show the actual situation of the i-th component in the mixture in terms of entropy, volume and enthalpy, respectively.

[I-6] EXCESS QUANTITIES The property of a mixture is not just the sum of those of constituent components. Aside from the mixing entropy discussed below, it is only an ideal, and hence a non-real situation, when the sum of the properties of the constituent components fully describes the property of the resulting mixture. For real cases, an assembly of many components often shows some characteristics unimaginable from the properties of each component. Such a special characteristic manifests itself through strong interactions among the constituents – the so-called “many-body problem”. In terms of enthalpy, for example, the total enthalpy of a mixture, H, is written as, H=



ni Hi∗ + H E 

(I-41)

ni

where Hi∗ is the enthalpy per unit amount of the i-th component in its pure state. For an ideal, and hence a non-real mixture, the total enthalpy is just the sum of those of each component, i.e. H E = 0. In reality, however, H E = 0, as a manifestation of intermolecular interactions among constituents. Equivalently, for other extensive quantities, V=



ni Vi∗ + V E 

(I-42)

ni Si∗ − R lnni /N  + S E 

(I-43)

ni G∗i + RT lnni /N  + GE 

(I-44)

i

S=

 i

G=

 i

The second term on the right of the last two equations R ln(ni /N ), is the mixing entropy arising from the fact that each component is distinguishable from the others and hence their local accessibility increases by mixing even in the ideal solution. We introduce the mixing entropy in consideration of the process of mixing below in [II-5]. Here, N = ni .

Eq. (I-39) to (I-51)

19

The key concept in understanding the nature of a mixture is the intermolecular interactions. Thus, the solution thermodynamics deals mostly with the excess functions. Therefore, instead of eqs. (37)–(40), we use the following excess partial molar quantities of the i-th component more often.  Ei ≡  SiE

≡ 

ViE ≡  HiE

=

GE

ni

S E

ni

 



V E

ni

(I-45)

pTnj=i



Ei =−

T

pTnj=i



 =

pTnj

Ei

p



H E

ni



(I-46)

pnj=i

 

(I-47)

Tnj=i

 = −T



2

pTnj=i

 Ei /T

T

 

(I-48)

pnj=i

Equations (45)–(48) indicate Ei is the first derivative and SiE , ViE , HiE are the second derivative of GE .

[I-7] RESPONSE FUNCTIONS There are three more second derivatives of G, in the (p T ni  variable system. They are called collectively the response functions: heat capacity, Cp , isothermal compressibility, T , and isobaric thermal expansivity, p .  Cp =

H

T

1 T = − V p =

1 V







 =T

p



V

p

V

T

T

 = T

S

T

 p

1 =− V 1 V



 = −T 



T

2 G

p2

  p

2 G

T 2

 

(I-49)

p

 

(I-50)

T

G

p

 

(I-51)

T

In all the above derivatives, ni are kept constant. From our experience, we know that Cp > 0, and T > 0. We have not encountered a system in which the temperature decreases on addition of heat or the volume increases on applying pressure. The nonnegativity of Cp and T is called the mechanical stability criterion. As will be discussed

20

in Chapter IV, Cp is proportional to the mean square (hence non-negative) fluctuation of entropy and T to that of volume. Hence they cannot be negative. [I-8] THERMODYNAMIC QUANTITIES – ORDER OF DERIVATIVE So far we have seen thermodynamic quantities that contain the first and second derivatives of G with respect to independent variables in the Gibbs ensemble, the (p T ni  variable system. We summarized them in Table I-1. They are classified by the highest order of derivative that each thermodynamic quantity contains. The first column shows the order of derivative of the leading term and the number of the quantity in each class is shown in parentheses. The second column is the notation. Those in the upper portion of the table are commonly accepted notations, while towards the bottom of the table new notations are indicated. The list in { } indicates the variables of differentiation. Thus, Cp , Table I-1. Thermodynamic quantities in p T ni  variable system. 0th

G

1st (4)

H  T a S  T  V  p i  ni 

H = G − T G/ T S = − G/ T V =  G/ p i =  G/ ni 

H E = GE − T GE / T S E = − GE / T V E =  GE / p Ei =  GE / ni 

2nd (7)

Hi  T ni  Si  T ni  Vi  p ni  Cp  T T  T  p p p  T p i−j  ni  nj 

Hi =  H/ ni  Si =  S/ ni  Vi =  V/ ni  Cp =  H/ T = T S/ T T = − V/ p/V p =  V/ T/V i−j ≡ N i / nj 

HiE =  H E / ni  SiE =  S E / ni  ViE =  V E / ni  S  ≡ Cp /V , S  = SR/V V  ≡ TT , V  = VR/V SV  ≡ Tp , SV  = SVR/V i−j E ≡ N Ei / nj 

3rd (11)

Hi−j  T ni  nj  Si−j  T ni  nj  Vi−j  p ni  nj  S i , S i V i , V i SV i , SV i

Hi−j ≡ N Hi / nj  Si−j ≡ N Si / nj  Vi−j ≡ N Vi / nj  S i ≡ N S / ni , S i ≡ N S / ni  V i ≡ N V / ni , V i ≡ N V / ni  SV i ≡ N SV / ni , SV i ≡ N SV / ni   Cp / T  Cp / p =  p / T  p / p = − T / T  T / p  i−j / nk 

Hi−j E ≡ N HiE / nj  Si−j E ≡ N SiE / nj  Vi−j E ≡ N ViE / nj 

4th (16) a

               ..

The variable in { } is the variable of differentiation.

Eq. (I-52)

21

for example, contains two T , indicating G is differentiated with respect to T twice to obtain Cp . The third column contains the definition equality. In this book, we more often use the quantities in the fourth column, the excess functions, for the reasons discussed above. Furthermore, instead of using the response functions, Cp , T , and p , directly we introduce (Koga, 1999; Koga & Tamura, 2000) and use q  and q (q = S, V , or SV). q  signifies the amplitude or the intensity of the mean square fluctuation in quantity q, while q  contains qualitative information about the wavelength or the extensity of fluctuation in q. For q = SV, both fluctuation functions, SV  and SV , show the cross fluctuation between S and V Unlike normal liquids, H2 O has a negative contribution in the SV cross fluctuation, due to a putative formation/destruction of ice-like patches. Hence SV  and SV  are important quantities in studying aqueous solutions. The process of modification of the molecular organization of H2 O by solutes could be monitored by these quantities. We discuss fluctuation functions in more detail in Chapter IV. Having tabulated the thermodynamic quantities up to the second derivatives, it is natural to expand the table to include higher order derivatives, as shown in Table 1 (p. 20). In the subsequent discussion of the study of aqueous solutions, we fully utilize the quantities listed in the second column of the third derivative class. We have not yet exploited the bottom five third derivative quantities and thus no new notations are given. It is our challenge to experimentally determine these and even higher order derivatives in the future. G, the zeroth order derivative, dictates the fate of an equilibrium state, by the stability criterion discussed above, eq. (34). The first derivatives, H and S indicate the ingredient of G. A second derivative, Cp for example, would indicate the degree of fluctuation in H or S Thus, the higher the order of the derivative, the more detailed information it provides. Our approach is, therefore, to obtain quantities of the second derivative by direct measurements. As much as the situation permits, we take one more derivative graphically without resorting to any fitting function. We then learn what the nature tells us directly from the data of the third derivative quantities that are model-free and experimentally accessible.

[I-9] INTERACTION FUNCTIONS – THIRD DERIVATIVES For an extensive excess function, F E (=GE  H E  S E or V E , the excess partial molar quantity of the i-th component, FiE , is written following eqs. (45)–(48) as,  FiE

=

F E

ni

 

(I-52)

pTnj=i

For historical reasons (since Gibbs used the notation), FiE = Ei for F E = GE ; otherwise FiE = HiE , SiE , or ViE . The physical meaning of the excess partial molar quantities of the i-th component is the actual situation that the i-th component experiences in terms

22

of GE , H E , S E , and V E respectively within the mixture. We now define the following derivative.  E

Fi E Fi−j ≡N  (I-53)

nj pTn k=j

This quantity, which we call the interaction function, is the change in FiE due to a unit E perturbation in nj , keeping all the other independent variables constant. Therefore, Fi−j signifies the effect of the j-th component on the actual situation of the i-th component in terms of F E . Thus, Ei−j is the i-j interaction in terms of chemical potential. It is clear E from the definition that Ei−j is the second derivative of G (see Table 1 (p. 20)). Hi−j , E E Si−j , and Vi−j are, on the other hand, the third derivatives of G, and they are indicative of the i-j interactions in terms of H E  S E , and V E respectively. E According to the stability criteria, if Ei−j or Hi−j is negative, the j-th component makes the chemical potential or the enthalpic situation of the i-th component more favorable. Therefore we may say that the i-j interaction is favorable or “attractive” in terms of chemical potential or enthalpy. If it is positive, then the i-j interaction is said to be unfavorable or “repulsive”. Of course, i and j are interchangeable and j could be equal to i also. Since the stability criterion for entropy is the opposite of enthalpy E or Gibbs function, a positive value of Si−j is indicative of a favorable (or attractive) i-j interaction in terms of excess entropy. For volume, there is no stability criterion. Thus, E Vi−j provides simply the degree of the effect of the j-th component on the volumetric situation of the i-th component in the mixture. We note in passing,  

i i−i ≡ N  (I-54)

ni is inversely related to the concentration fluctuation accessible by scattering experiments. (Note this is not the excess function, Ei−i , eq. (53) with F = G with i = j). As will be discussed in Chapter IV, this is inversely proportional to the mean square (hence non-negative) concentration fluctuation. The non-negativity of i−i , or the right-hand side of eq. (54), is known as the diffusional stability criterion. It is also related to the Kirkwood-Buff (1951) integrals. If the excess partial molar quantity, FiE , is determined accurately and in small increE ments in nj , then it is possible to evaluate Fi−j graphically within an acceptable accuracy, E , giving several per cent. (See [III-4], for graphical differentiation). Thus, we have Fi−j important information about the i-j interaction, purely from experiment without resorting to any model. This information is crucial in understanding the nature of a non-ideal, and hence a real mixture. We will make extensive use of this model-free, experimentally accessible information about the i-j interactions in order to study the molecular processes in aqueous solutions.

Eq. (I-53) to (II-6)

Chapter II

SOLUTION THERMODYNAMICS – USE OF THE SECOND AND THIRD DERIVATIVES OF G

[II-1] MIXTURE Consider  a gas mixture consisting of ni of the i-th component, i = 1 2 3    k, with N = ni . The mixture occupies the volume, V , under the total pressure, p. We define i

the mole fraction of the i-th component, xi , as, ni  N

xi =

(II-1)

The partial pressure of the i-th component, pi , is defined as, pi = xi p

(II-2)

and hence, p=



pi 

(II-3)

i

A perfect gas by definition obeys the Boyle-Charles law, a special form of the state function, V = fp T ni  V=

NRT  p

(II-4)

V=

ni RT  pi

(II-5)

Using eqs. (1), (2) and (3),

Differentiating eq. (4) with respect to ni , keeping p, T, and nj=i constant and using eq. (I-39), we obtain,   V V RT =  (II-6) = Vi = ni pTnj=i p N

24

Recalling the second equality of eq. (I-39), and using eq. (6),   i RT  = Vi = p Tnj=i p

(II-7)

Integration of eq. (7) yields, i =

0i + RT

pi 1 p0

p

dp

(II-8)

Thus,  pi  ln p0 

i =

0i + RT

(II-9)

where p0 is the pressure of a standard state, which is normally taken to be 1 atm, and 0i is the value of i at p = p0 . We note in passing that for non-perfect gases, the state function, eq. (4), will have virial correction terms. In such a case, eq. (9) should be written as,   f (II-10) i = 0i + RT ln i  f0 where fi and f0 are fugacities. By virtue of eq. (2), eq. (9) is rewritten as, i = ∗i + RT ln xi 

(II-11)

where ∗i is the chemical potential of the i-th component in its pure state, i. e. xi = 1, with  ∗ pi ∗ 0 i = i + RT ln p0 where pi∗ is the vapor pressure of pure liquid i at the same temperature. Eq. (11) is valid strictly for a perfect gas mixture or an ideal mixture. A perfect gas mixture is characterized by the total lack of intermolecular interaction among constituent molecules. This situation is not at all real. For a real non-perfect gas mixture, an artificial introduction of fugacity, fi , eq. (10), instead of measurable pi is one way to cope with the situation. Another is to introduce an extra term, the excess chemical potential, Ei , of i to the right of eq. (11), as shown in eq. (12) below. We take the latter route. We then experimentally determine Ei . i = ∗i + RT ln xi + Ei 

(II-12)

Eq. (II-7) to (II-17)

25

Ei is called the excess chemical potential of the i-th component, and takes a non-zero value reflecting the real situation in a mixture in which intermolecular interactions are prevalent. We note that by the present choice of ∗i , i must be equal to ∗i and hence Ei = 0 at xi = 1. This is equivalent to choosing the reference state, where Ei = 0, to be the pure state for all the components. Hence this is often called the symmetric reference system, since all the components are treated equally. Any other choices are possible, but the dilute ideal solution is often taken as the reference. Namely, Ei is taken as zero at the infinite dilution, i. e. all the mole fractions of the solutes are zero. This is a convenient choice for studying exclusively on dilute solutions. Here we choose the symmetric reference system, since we are interested in the mixing schemes in the entire composition range. Instead of eq. (12), a more conventional way is to write i = ∗i + RT ln ai 

(II-13)

where ai is the activity of the i-th component in the mixture and a i = x i i 

(II-14)

i is called the activity coefficient. Comparing eq. (12) with (13) and (14), Ei = RT ln i 

(II-15)

We use eq. (12) throughout this book, instead of the more common eq. (13), and determine Ei experimentally. The usage of ai seems to us to have some apologetic connotation; apologetic in the fact that the reality does not follow eq. (11).

[II-2] GIBBS-DUHEM RELATION From eqs. (I-23), (I-24), (I-36) and (I-37), dG = −S dT + V dp +



i dni 

(II-16)

i

With T and p being kept constant, dG =

 i

i dni 

(II-17)

26

Since G is an extensive quantity, the value of G is proportional to the amount of the system. Namely, if we increase ni to rni i = 1 2 3    k, where r is a positive constant, then the value of the Gibbs function of the r-fold system, G , will be equal to rG. Or, G rni  = rGni 

(II-18)

The chemical potential, on the other hand, is an intensive variable and independent of the size of the system, as long as xi ’s are constant. To show this, we calculate the chemical potential of the i-th component in the r-fold system, i , as, 

i

G = rni 





rG = rni 

nj



 = nj

G ni

 = i 

Thus, the chemical potential of the i-th component is invariant to the change from ni to rni , i.e. i is an intensive quantity. The same argument leads also to the fact that other partial molar quantities, eq. (I-38) to eq. (I-40), are all intensive quantities. We now integrate eq. (17) from r = 0 to r = 1. The result is written with the zero integration constant as, G=



n i i 

(II-19)

i

A similar deduction leads to the relation for the excess Gibbs function as, GE =



ni Ei 

(II-20)

i

The same arguments as those above also lead to similar expressions, HE =



ni HiE 

(II-21)

ni SiE 

(II-22)

ni ViE 

(II-23)

i

SE =

 i

VE =

 i

We now take the differential of eq. (19) mathematically. Namely, dG =

 i

i dni +

 i

ni di 

(II-24)

Eq. (II-18) to (II-30)

27

For this equation to be physically consistent with the original eq. (17), the second term on the right of eq. (24) must be identically equal to zero. 

ni di = 0

(II-25)

i

This is the Gibbs-Duhem relation. We make use of this relation later in analyzing vapor pressure data to calculate the excess chemical potentials. Similar deduction for the other excess functions yields, 

ni dEi = 0

(II-26)

ni dHiE = 0

(II-27)

ni dSiE = 0

(II-28)

ni dViE = 0

(II-29)

i

 i

 i

 i

Thus, the changes in (excess) partial molar quantities are not completely independent. These relations are useful in checking the self-consistency in experimental data. For example, the excess partial molar enthalpies of each component in a binary mixture E and HWE , must be related thanks to of ethanol (abbreviated as ET) and H2 O (W), HET eq. (27) as, E xET HET + xW HWE = 0

(II-30)

where xET and xW are the mole fraction of ET and W respectively. Fig. (II-1) shows the E , and that of W, HWE , actually measured by the excess partial molar enthalpy of ET, HET method described in Chapter III (Tanaka et al. 1996; Liltorp et al., 2005). Thus the slopes E and HWE , are related by eq. (30). The at a given mole fraction shown in the figure, HET thermodynamic consistency of these data is therefore checked by eq. (30). The results of the left-hand side of eq. (30) are listed in Table II-1. As is evident, the left-hand sides of eq. (30) are indeed zero within ±002 kJ mol−1 , which is comparable with the E and HWE , ±003 kJ mol−1 . Thus, this set of data is thermodynamically uncertainty of HET E available, eq. (30) consistent. If, on the other hand, there are only the data for HET E E provides the means to estimate HW . Since HW is identically zero at xW = 1, or xET = 0, the increase in HWE  HWE for a small increment in xET  xET is given by eq. (30) knowing E E  HET . Thus, HWE can be calculated stepwise from xET = 0 to the the increment of HET entire composition range.

28

Fig. II-1

2

H ETE

H ETE or H WE / kJ mol-1

0

-2

H WE

-4

-6

-8

-10

-12 0.0

0.2

0.4

0.6

0.8

1.0

xET

Fig. II-1. Excess partial molar enthalpies of ethanol (ET) and H2 O(W) in ET-H2 O at 25  C. The slopes of the tangents (shown by straight lines) at a fixed mole fraction of ET, xET , (indicated by E the dotted line) are related by the Gibbs-Duhem relation, xET HET + xW HWE = 0. The data from Tanaka et al. (1996). Table II-1. Gibbs–Duhem for the excess partial molar enthalpy of ethanol (ET) and water (W) in aqueous ethanol at 25 C. xET

E HET kJ/mol

HWE kJ/mol

Error 0 001 002 003 004 005 006 007 008 009 01 011 012 013

003 −1011 −972 −926 −871 −807 −734 −652 −568 −486 −41 −337 −273 −22 −174

003 0 0 −001 −003 −007 −01 −013 −019 −024 −031 −039 −047 −054 −062

xET

Gibbs-Duhem Sum kJ/mol

0005 0015 0025 0035 0045 0055 0065 0075 0085 0095 0105 0115 0125 0135

0002 −0003 −0006 −0016 0004 0017 −0002 0015 0001 −0003 −0004 −0001 −0013 −0011

29

Table II-1. (Continued) xET

E HET kJ/mol

HWE kJ/mol

xET

Gibbs-Duhem Sum kJ/mol

014 015 016 017 018 019 02 022 024 026 028 03 032 034 036 038 04 042 044 046 048 05 052 054 056 058 06 062 064 066 068 07 072 074 076 078 08 082 084 086

−137 −104 −078 −054 −035 −02 −008 006 015 019 023 023 023 023 022 021 02 018 016 013 0105 008 006 004 002 0 −002 −004 −006 −008 −009 −01 −01 −01 −01 −01 −01 −01 −009 −008

−069 −074 −08 −083 −0875 −09 −093 −097 −101 −101 −101 −101 −101 −101 −1 −099 −098 −096 −094 −092 −09 −089 −087 −085 −082 −08 −078 −076 −073 −07 −068 −066 −063 −06 −059 −059 −059 −06 −062 −065

0145 0155 0165 0175 0185 0195 021 023 025 027 029 031 033 035 037 039 041 043 045 047 049 051 053 055 057 059 061 063 065 067 069 071 073 075 077 079 081 083 085 087

0005 −0010 0015 −0004 0007 −0001 −0002 −0010 0010 0011 0000 0000 0000 0003 0003 0002 0004 0003 −0003 −0001 −0007 0000 −0001 0003 −0003 −0004 −0004 −0001 −0002 0000 −0001 0009 0008 0003 0000 0000 −0002 0005 0004 0002 (Continued)

30

Fig. II-2

Table II-1. (Continued) xET 088 09 092 094 096 098 1

E HET kJ/mol

−007 −0055 −004 −003 −002 −001 0

HWE kJ/mol

xET

−07 −077 −085 −095 −11 −126 −148

089 091 093 095 097 099

Gibbs-Duhem Sum kJ/mol 0006 0006 0002 0002 0005 0008

[II-3] VAPOR PRESSURES Consider the equilibrium between a liquid and a gas mixture with p and T kept constant. At the equilibrium, i gas = i liquid

(II-31)

This is shown by the following argument. Transfer an infinitesimal amount ni from the gas phase to the liquid phase, as shown in Fig. II-2. If the entire system is in equilibrium, the change in the total Gibbs function is zero,

Gtotal = 0 = Ggas + Gliquid

(II-32)

Ggas = − Gliquid

(II-33)

Therefore,

Since,

Ggas = i gas ni  and,

Gliquid = i liquid− ni  we immediately recover eq. (31). We point out in passing that eq. (31) leads to the Gibbs phase rule. In the system consisting of C-components in equilibrium with P-phases with no chemical reactions among components in all phases, the variance, or the number of degrees of freedom of variables, f , is written as, f = C − P + 2

(II-34)

Eq. (II-31) to (II-37)

31

Gas phase

p i (i = 1, 2, 3,...,k) δn i

ni (i = 1, 2, 3,...,k)

Liquid phase

Fig. II-2. A k-component mixture is in a closed and isolated system consisting of two phases: gas and liquid. Each phase is open and non-isolated to the other. They can exchange material and energy. pi is the partial pressure of the i-th component in gas phase and ni the amount of the i-th component in liquid mixture.

The deduction is straightforward and most elementary textbooks contain a paragraph devoted to this discussion. i gas and i liquid are rewritten by virtue of eqs. (9) and (12) as,   pi 0  (II-35) i gas = i + RT ln p0 and, i liquid = ∗i + RT ln xi + Ei 

(II-36)

with the condition for the equilibrium between the gas and liquid phases, i gas = i liquid

(II-37)

Here, we assume that the gas phase mixture can be treated as an ideal mixture in comparison with the non-ideality of liquid mixture. Since the density of gas is generally 1000-fold smaller than that of liquid, the intermolecular interactions in gas could be

32

Fig. II-3

negligibly small in comparison with those in liquid. This assumption is shown below to be acceptable. Consider now separately the equilibrium between the pure liquid and the pure gas of the i-th component at the same temperature. Instead of eqs. (35) and (36), we have for gas,  i pure gas = 0i + RT ln

 pi∗  p0

(II-38)

where pi∗ is the vapor pressure of the pure liquid at the same temperature. For liquid, i pure liquid = ∗i 

(II-39)

with the condition for equilibrium, i pure gas = i pure liquid

(II-40)

Combination of eqs. (35) to (40) yields,  pi  = RT ln xi pi∗ 

Ei

(II-41)

This provides the method of evaluating Ei in the liquid mixture by measuring the equilibrium partial pressure of the i-th component, pi . When non-ideality of the gas phase mixture must be taken into account, eq. (41) needs another term with the virial coefficients of all the components, Bii , together with all the cross terms, Bij in the gas phase. Thus, instead of eq. (41),  pi + Ci  = RT ln xi pi∗ 

Ei

(II-42)

with Ci = Bii − Vi∗ p − pi  + p1 − pi /pz

(II-43)

where z = 2B12 − B11 − B22 for a binary mixture consisting of components 1 and 2. However, for the aqueous solutions of our main interest, the first term on the right-hand side of eq. (42) gives a large non-zero value and the additional virial correction term is negligible. We therefore use eq. (41) for most cases. For aqueous tert-butanol, for example, the excess chemical potential of tert-butanol, ETBA calculated by eq. (41) was about 6000 J/mol in the H2 O-rich region, while the virial correction term in eq. (42) amounted only to several J/mol. (Koga et al., 1990a; Koga et al., 1990c).

Eq. (II-38) to (II-45)

33

[II-4] RAOULT’S LAW AND HENRY’S LAW For the very special case of an ideal mixture which does not exist in reality, Ei = 0. Then from eq. (41), pi = xi pi∗ 

(II-44)

This is Raoult’s law, which says that the other components work only as diluents on the vapor pressure of the i-th component as a result of no intermolecular interaction. For such a binary mixture of component 1 and 2, the total pressure is written as, p = p1 + p2 = x1 p1∗ + x2 p2∗ 

(II-45)

Fig. II-3 shows the vapor pressure of an ideal, and hence a non-real, mixture, following eq. (44) and eq. (45). In reality, however, the vapor pressures depend on the mole fraction in a much more complicated manner. See Fig. II-4, for the mole fraction dependence of the vapor pressures for tert-butanol (TBA) – H2 O at 25  C. pTBA and pW are the partial pressures of tert-butanol (TBA) and H2 O(W), respectively. (Koga et al., 1990a) This complex mole fraction dependence is the manifestation of the intermolecular interactions

p 2*

p, p1, or p 2

Ideal mixture 1 - 2

p = p1 + p 2

p2

p1*

p1

0

1

x2

Fig. II-3. Total pressure, p, and the partial pressures over the ideal mixture 1–2. p1 and p2 are the partial pressures of 1 and 2. p1∗ and p2∗ are the vapor pressures of pure 1 and pure 2, respectively, at the same temperature.

34

Fig. II-4

60

50

p, p TBA, or p w / Torr

p = p TBA + p w 40

p TBA

30

20

pW

10

0 0.0

0.2

0.4

0.6

0.8

1.0

x TBA

Fig. II-4. Vapor pressures of a real mixture, for example, tert-butanol (TBA)−H2 O(W) at 25  C. pTBA , and pW are the partial pressures of TBA and W. p is the total pressure. The broken lines would be the partial pressures of TBA and W if the mixture were to be ideal. Note that pTBA is close to the partial pressure of the ideal solution near xTBA = 1, and pW also near xW = 1, or xTBA = 0. The data from Koga et al. (1990a).

occurring in the mixture and is the very key to understanding the molecular mechanisms, as will be discussed throughout this book. We note in Fig. II-4, however, that the partial pressure of H2 O, pW , approaches the broken line in the figure, the hypothetical Raoult’s law vapor pressure, as xTBA → 0. Namely, the other component, TBA, is so dilute that the interactions from TBA do not affect H2 O much, and thus TBA molecules are working just as diluents for H2 O. Thus, the partial pressure of H2 O, pW , in this region is approximated as, ∗  pW ≈ 1 − xTBA pW

(II-46)

This approximation will be used to calculate the partial pressures from the total vapor pressure data for a binary system. Whether eq. (46) becomes strictly the equality, ∗  pW = 1 − xTBA pW

(II-47)

in a very dilute but a finite concentration range is an unresolved issue. (Koga, 1995a; Westh et al, 1998). If eq. (47) holds at all in a finite range or even in an asymptotic

Eq. (II-46) to (II-54)

35

condition of xTBA → 0 and xW → 1, then due to the Gibbs-Duhem relation, the partial pressure of TBA must also be proportional to xTBA in the same range as, pTBA = kH xTBA 

(II-48)

kH is called Henry’s constant. Observe in Fig. II-4 that for xTBA → 0, pTBA is almost on a straight line against xTBA . The same argument applies for pW in the region of xW → 0, or xTBA → 1.

[II-5] PROCESS OF MIXING – MIXING ENTROPY Let us consider mixing nB mol of B and nW mol of W resulting in the total N= nB + nW  mol of the mixture with T and p kept constant. Namely, nB mol of pure B + nW mol of pure W → Mixture N = nB + nW 

(II-49)

The change in the Gibbs energy associated with this process, mix G, is the difference in the total Gibbs energy between the right and the left of the process (49). Hence,

mix G = GnB  nW  − G∗ nB  + G∗ nW 

(II-50)

where GnB  nW  is the Gibbs energy of the mixture, while G∗ nB  and G∗ nW  those of nB of pure B and nW of pure W, respectively. According to eq. (19), GnB  nW  is written as, GnB  nW  = nB B + nW W 

(II-51)

G∗ nB  = nB B ∗ 

(II-52a)

G∗ nW  = nW W ∗ 

(II-52b)

and

The chemical potentials are written due to eq. (12) as, B = B ∗ + RT lnxB  + EB 

(II-53a)

W = W ∗ + RT lnxW  + EW 

(II-53b)

mix G = nB EB + nW EW  + RT nB lnxB  + nW lnxW 

(II-54)

Thus,

36

The above process, eq. (49), is nothing but the defining process of the Gibbs energy of the mixture at p, T, nB and nW , if we choose the pure states of each constituent as the reference, i.e. G∗ nB  = G∗ nW  = 0. Hence, Gp T nB  nW  = mix G = GE + GIdeal

(II-55)

with GE = nB EB + nW EW 

(II-56a)

GIdeal = RT nB lnxB  + nW lnxW 

(II-56b)

The superscript “Ideal” on the left of eq. (56b) is due to the fact that this term remains non-zero, even if the mixture is ideal with EB = EW = 0. By differentiating with respect to T , and recalling eq. (I-23), we obtain,   G = S E + S Ideal  (II-57) Sp T nB  nW  = − T with, S E = nB SBE + nW SWE 

(II-58a)

S Ideal = −RnB lnxB  + nW lnxW 

(II-58b)

where  SBE = −

EB T



  and SWE = −

EW T

 

(II-59)

from eq. (I-46). By virtue of the definition, G = H–TS, Hp T nB  nW  = H E + H Ideal 

(II-60)

H E = nB HBE + nW HWE 

(II-61a)

H Ideal = 0

(II-61b)

with

where  HBE

= −T

2

   EB /T EW /T E 2  and HW = −T  T T

(II-62)

Eq. (II-55) to (II-65)

37

from eq. (I-48). From the relation, V = G/p, eq. (I-24), V = V E + V Ideal 

(II-63)

V E = nB VBE + nW VWE 

(II-64a)

V Ideal = nB VB∗ + nW VW∗ 

(II-64b)

with

where  VBE

=

EB p



  and

VWE

=

EW p

 

(II-65)

from eq. (I-47). We note that S Ideal , eq. (58b), is always non-zero for any mixture. Even if the solution is free from interactions and S E is zero, this term is positive since xB and xW are both less than unity. This is called the mixing entropy, arising from the fact that species B and W are distinguishable. Namely when B and W are added together, they spontaneously tend to mix randomly, even if there is no interaction between them. This brings about the finite negative value of GIdeal , eq. (56b); the spontaneity of the process of mixing B and W even if there is no interaction between them. In Statistical Mechanics, the mixing entropy, eq. (58b) is nothing but the increase in accessibility of distinguishable B and W species in a completely random mixture. Due to the choice of the reference states, H Ideal , eq. (61b) is apparently zero. Namely, for an ideal solution lacking in interaction there is no enthalpy of mixing. V Ideal , eq. (64b), on the other hand, is non-zero and is the sum of the volumes of nB of B and nW of W in their pure states. This apparent difference simply comes from the fact that volumes have an absolute scale while energies and enthalpies have not. In reality, the excess functions, GE , eq. (56a), S E , eq. (58a), H E , eq. (61a), and V E , eq. (64a) are nonzero due to interactions among the constituents of the mixture. Thus, the experimental determination of these excess functions and their higher order derivatives is crucial in obtaining information about the interactions in order to elucidate the mixing schemes or the solution structures.

[II-6] CONVERSION FROM nB  nW  TO xB  N VARIABLE SYSTEMS The excess partial molar quantities, eq. (I-45) to eq. (I-48), are defined in the (p, T , nB , nW ) variable system for a binary mixture of B and W. For experimentalists, however, it is more convenient to use the (p, T , xB , N ) variable system, in order to present the results in a graphical form. Otherwise, the entire composition range from the pure B to

38

the pure W end cannot be realistically covered in the (nB , nW ) scale. Consider now any excess thermodynamic function, F E , in both variable systems. Thus, keeping p, and T constant, F E nB  nW  = F E xB  N

(II-66)

Recalling eq. (1), xB =

nB n = B n B + nW N

(II-67)

The total differential of F E is identical in each variable system, as long as the condition eq. (67) holds. Hence, dF E nB  nW  = dF E xB  N

(II-68)

and  dF nB  nW  = E

 dF E xB  N =

F E nB

F E xB





F E dnB + nW

nW



 dxB +

N

F E N

 dnW 

(II-69)

nB

 dN

(II-70)

xB

Inserting eq. (69) and eq. (70) into eq. (68) and dividing both side of eq. (68) by nB keeping nW constant, i.e. nW = 0, we obtain, 

F E nB



 Since from eq. (67), 

 = nW

xB nB

F E xB



F E nB

nW

  N

xB nB



 +

nW

1 − xB  , and = N

 = nW

1 − xB  N





F E xB

F E N

N nB

  xB

N nB

 

(II-71)

nW

 = 1, eq. (71) is rewritten as, nW



 + N

F E N

 

(II-72)

xB

If F E is an extensive quantity, such as GE  H E  S E , and V E , the second term on the right of eq. (72) is identically equal to F E /N which is the molar excess quantity, FmE . However, if F E is an intensive quantity as Ei , HiE , SiE , and ViE , then the second term of eq. (72) is identically zero. Hence, for an extensive quantity, 

F E nB



 = 1 − xB  nW

FmE xB

 + FmE 

(II-73)

Eq. (II-66) to (II-76d)

39

and for an intensive quantity, 

F E nB

 nW

1 − xB  = N



F E xB

 

(II-74)

These equations, eq. (73) and eq. (74) provide the means to raise the order of derivative of a thermodynamic function by one step in a convenient xB  N variable system. For F ≡ H, for example, eq. (73) is rewritten as,  HBE = 1 − xB 

HmE xB

 + HmE 

(II-75)

Hence if HmE = H E /N is experimentally determined as a function of xB , HBE should be calculated by eq. (75). As will be discussed in Chapter III, this route of obtaining HBE , a second derivative of G, is problematic and it is much more advantageous to determine HBE experimentally. Recall eq. (I-53), in which we introduced the thermodynamic interaction functions. Eq. (74) above provides the basis for calculating them in the xB  N variable system. For F E = EB , HBE , SBE , or VBE , we redefine the interaction functions in terms of G, H, S, or V as,  E  E B B E  (II-76a) = 1 − xB  B−B ≡ N nB xB  E HB−B

≡N 

E SB−B

≡N 

E VB−B

≡N

HBE nB

SBE nB





VBE nB



HBE = 1 − xB  xB 

SBE = 1 − xB  xB





 

(II-76b)



VBE = 1 − xB  xB



(II-76c)

 

(II-76d)

As the definitions imply, these interaction functions signify the effect of a unit increase of B on EB , HBE , SBE , or VBE , the actual situation of existing B in the mixture. Hence they represent the B-B interaction. Due to the stability criteria discussed in [I-4], the sign of the interaction function, eq. (76), indicates that an addition of B results in a favorable or a non-favorable situation of existing B in terms of each quantity. We call the B-B interaction “attractive” or “repulsive” for short. Thus, the B-B interaction is E are negative (or positive) in terms of chemical attractive (or repulsive) if EB−B and HB−B E except that the sign is potential and enthalpy, respectively. The statements hold for SB−B E reversed. Since there is no stability criterion for volume, VB−B simply shows the effect of an additional B on the volumetric situation of B.

40

We use the word “interaction” in accordance with the sense conveyed by eq. (76) above. Hence, it is not necessarily limited to such conventional meanings as the potential or the force field. The latter and many other effects are included in a holistic manner in our usage. So far, we have dealt only with a binary mixture (B, W). Generalization of the above discussion to a multi-component system 1 2 3     k is straightforward following the route starting at eq. (66) and keeping track of which variables are kept constant. Thus an equivalent expression for eq. (72) is written as, 

F E n1

 n2 n3     nk nk+1

1 − x1  = N  +

F E N



F E x1



  k  xm F E − xm xj=m x2 x3     xk m=2 N

 

(II-77)

xi

[II-7] INTERACTION FUNCTIONS DUE TO THE IDEAL MIXING ENTROPY As we have discussed so far, the excess functions, GE  H E  S E , and V E contain the information about the solute-solute interaction in real mixtures crucial for understanding the underlying molecular processes – mixing schemes. However, there is the ideal mixing entropy that dictates the complete random mixing even in the absence of the excess entropy term, S E . The ideal mixing entropy, S Ideal , is written due to eq. (58b) as, S Ideal = −RnB lnxB  + nW lnxW 

(II-78)

and its molar quantity, Smid , is Smid = −RxB lnxB  + xW lnxW 

(II-79)

Then the partial molar ideal mixing entropy of B, SBid , can be written following eq. (72) as,  Smid + Smid = 1 − xB  xB 

SBid

= −RlnxB 

(II-80)

id Thus, the B-B interaction in terms of the ideal mixing entropy, SB−B could be calculated by eq. (74) as,

id = −R SB−B

1 − xB   xB

(II-81)

Eq. (II-77) to (II-83)

41

id id This SB−B remains negative in the entire composition range. Namely, SB−B drives B molecules apart though its degree is small except for the region sufficiently near xB = 0. Similar arguments lead to the B-B interaction in terms of the ideal chemical potential, id B−B , written as,

id B−B = RT

1 − xB   xB

(II-82)

which is always positive, indicating repulsion in terms of ideal chemical potential. This effect is dominant only at a very dilute region over that by the excess functions. The id E E contributions from SB−B and/or id B−B over corresponding SB−B and/or B−B will be discussed when we determine the latter interaction functions for actual aqueous solutions of non-electrolytes in Part B below.

[II-8] PHASE SEPARATION – CRITICAL POINT (UCST OR LCST) In this and the next section we discuss the topics that will be useful for Part B. At constant pressure and temperature, the thermodynamic stability criterion for G is dG ≤ 0, eq. (I-34) As will be discussed in Chapter IV, the heat capacity, Cp , and the isothermal compressibility, T , can never be negative. This is completely consistent with human experience and is often called the mechanical stability criteria. Another inequality, called the stability criterion, is applicable in a binary mixture (B, W) such that  diffusional

B xB must never be negative. This together with the mechanical stability criteria are supported by the fact that these quantities are related to the mean square fluctuations, as discussed in Chapter IV, i.e. they are never negative, being squared quantities. Namely, 

B xB

 ≥ 0

(II-83)

Note, however, that the xB derivative of the excess chemical potential of B, EB , could be either positive or negative. Indeed, this is related to the B-B interaction in terms of excess chemical potential, EB−B , as shown in eq. (76a). The sign of EB−B is an important piece of information for understanding that the B-B interaction is favorable or not in terms of excess chemical potential, as will be used repeatedly throughout this book. Eq. (83) is concerned with the sign of the derivative to the total chemical potential, not the excess chemical potential. The difference is, of course, the term due to the mixing entropy, as discussed above in Section [II-7]. When the interaction among the constituents becomes sufficiently strong in a binary (B, W) mixture, we may observe a phase separation consisting of two liquid phases; one rich in B and the other in W. Under a constant pressure condition, the region of phase separation is depicted in the temperature-mole fraction field as shown in Fig. II-5. Two typical cases are given in Fig. 5(a) and 5(b). The former, Fig. 5(a), is

42

Fig. II-5(a) to II-5(c)

(a) Single phase mixture

UCST

T

Two Liquids

XB

Fig. II-5(a). Phase separation with the upper critical solution temperature (UCST). The necessary E E but not sufficient conditions for UCST to occur are HB−B < 0, and SB−B < 0. (See text; see also Figs. II – 7(a) and 7(b).)

(b)

Two Liquids

T LCST

Single Phase Mixture

XB

Fig. II-5(b). Phase separation with the lower critical solution temperature (LCST). The necessary E E but not sufficient conditions for LCST to occur are HB−B > 0 and SB−B > 0. (See also Figs. II – 8(a) and 8(b).)

the case in which phase separation occurs at low temperatures with a critical point. The temperature above which the mixture forms a single phase is called the upper critical solution temperature (UCST). This case is easily understandable in that a strong attraction between like species in terms of enthalpy would bring a phase separation at

Eq. (II-84)

43

70

(c)

60

temperature / °C

50

2-butanone - H2O

40 30

Two liquids 20 10 0

0.0

0.2

0.4

0.6

0.8

1.0

x BUT

Fig. II-5(c). Liquid phase diagram for 2-butanone (BUT) – H2 O. Phase boundaries are due to Felino et al. (1983). Freezing prevents observation of LCST, which is estimated from −22 to −6  C (Francis, 1961). (See also Wong et al. (1992).)

low temperatures, while at higher temperatures the total entropy effect drives the system to a random mixture. Fig. 5(b) shows the case in which phase separation occurs at high temperatures with the lower critical solution temperature (LCST). This case is not so easily explained as the UCST case above. However, the following argument using the diffusional stability criterion, eq. (83), provides a general understanding for both the UCST and the LCST cases. Since B = B ∗ + RT lnxB  + EB , eq. (12), eq. (83) is rewritten as,   E  RT B B = ≥ 0 (II-84) + xB xB xB The first term on the right of eq. (84) is always positive. If the second term is positive, i.e. EB /xB  > 0, then the inequality eq. (83) or eq. (84) is always satisfied and the system is in a single phase. However, if the second term is sufficiently negative, i.e. EB /xB  << 0, then it is possible that the diffusional stability, eq. (83) is violated in a certain range, as shown in Fig. II-6. Namely, by adding the RT lnxB  and EB − B ∗  terms shown in Fig. 6(a), the stability criterion, eq. (83) or eq. (84), is violated in the region between B and C in Fig. 6(b). The system then phase-separates in the region between A and D, keeping B /xB  = 0, saving the system from violating eq. (83). The region between A and B and that between C and D are meta-stable. While eq. (83) is satisfied in these regions, the value of G = nB B + nW W can be shown to be larger than that for the sum of two separate phases. Similar to the Maxwell construction for

44

Fig. II-6(a) to II-6(b)

(µ BE - µ B*) or RT ln(x B)

µ BE - µ B* 0

RT ln(x B)

(a)

0

1

xB

Fig. II-6(a). The mole fraction xB  dependence of the excess chemical potential of B and that from the ideal mixing entropy RT lnxB . (See text.) 0

T

B M

µ B - µ B*

A

D C

TC Q

(b)

0

1

xB

Fig. II-6(b). The mole fraction xB  dependence of the chemical potential of B for phase separating B-W system at temperature T . (See [II-8] for A, B, C, D, and M in the figure.) At Section

B temperature, TC , a critical point, d ≥ 0, and is zero at point Q. If TC > T , the point Q is dxB UCST, and if TC < T , point Q is LCST. The scale of ordinate is arbitrary.

45

a van der Waals fluid (Chapter XVI in Prigogine & Defay, 1954), the area encircled by AMBA and that by DMCD are the same. We note that points B and C, the edge of meta-stability, are called the spinodal points. We also note that at a certain condition (temperature), eq. (83) is always satisfied, but at point Q in Fig. 6(b), B /xB  = 0. This is the critical point, the end point of phase separation. Thus, a necessary condition for phase separation to occur is that EB /xB  be sufficiently negative. For this to occur, there are two simple possibilities depicted in Figs. II-7 and II-8. For Case I, Fig. 7(a), HBE > TSBE with both HBE /xB  and SBE /xB  are negative, hence B /xB  < 0 with the possibility of phase separation. However, if temperature increases sufficiently, then a situation may arise such that HBE < TSBE as shown in Fig. 7(b). Hence EB /xB  > 0 and there is no phase separation. Namely, phase separation could occur at low temperatures with a UCST. In this case, the B-B interactions introduced by eq. (76) are attractive in terms of enthalpy, but repulsive in entropy. This is consistent with the intuitive discussion above in that the attractive enthalpic B-B interaction dominates at low temperatures and provides the possibility for phase separation. At high temperatures the repulsive entropic B-B interaction coming from the excess entropy, together with that from the ideal mixing entropy, dictates that the system mixes well. Case II, Fig. 8, shows the opposite situation, whereby the B-B interaction is repulsive in terms of enthalpy but attractive entropy-wise. At high temperatures, HBE > TSBE and hence EB /xB  < 0, as shown in Fig. 8(a), providing the possibility of phase separation. At sufficiently low temperatures, however, HBE < TSBE resulting in EB /xB  > 0 with no chance for phase separation. Thus, if phase separation is to occur, it must be with an LCST, as shown in Fig.5(b). E In summary, the system with a UCST, Fig. 5(a), is consistent with the case HB−B <0 E and SB−B < 0. We note that the phase boundary in the low xB side slants with the slope dT/dxB > 0, with dHBE /dxB < 0, and dTSBE /dxB < 0. The system with an LCST, Fig. 5(b), E E on the other hand, has the situation HB−B > 0 and SB−B > 0 with the phase boundary E slanting with a negative slope, dT/dxB < 0, with dHB /dxB > 0, and dTSBE /dxB > 0. There are many binary systems with various types of phase diagrams and complete generalization is not possible. There are cases including aqueous 2-butanone. As shown in Fig. II-5(c), there appears to be an LCST, but it is hidden by the formation of solid phases. In this and more general cases, it is possible to draw the relationships between E E the sign of HB−B and SB−B in a single phase of small xB side with the sign of the slope of the phase boundary, dT/dxB , with the two phases on the larger xB side. The results E E are consistent with the discussion above; namely, if HB−B > 0 and SB−B > 0 the phase E E < 0, then boundary slants with a negative slope, dT/dxB < 0, but if HB−B < 0 and SB−B the phase boundary has a positive slope, dT/dxB > 0. Note, however, the inverse is not necessarily true. We will not repeat the discussion here; readers may refer to Wong et al. (1992). The UCST and the LCST are the critical points and have specific characteristics of the so-called critical phenomena (Stanley, 1971). As with any other critical points at the super critical region in a single phase domain, the solution shows a large fluctuation

46

Fig. II-7(a) to II-8(b)

H BE and TS BE

H BE

(a) Low temperature Possible phase separation

TS BE

µ BE 0 0

1

xB

Fig. II-7(a). The case in which HBE > TSBE at low temperature, resulting in the required behavior E E of EB for possible phase separation. Note that HB−B < 0 and TSB−B < 0.

H BE and TS BE

TS BE

(b) High temperature No phase seperation

H BE

0

µ BE 0

1

xB

Fig. II-7(b). The case in which HBE < TSBE at high temperature resulting in the behavior of EB E E which gives no chance for phase separation. Note HB−B < 0, and SB−B < 0.

47

µ BE, H BE, and TS BE

µ BE 0

H BE

TS BE (a) High temperature Possible phase separation

0

1

xB

Fig. II-8(a). The case in which HBE > TSBE at high temperature resulting in the required behavior E E of EB for phase separation. Note that HB−B > 0 and SB−B > 0. 0

µ B E, H BE, and TS BE

µ BE

TS BE

H BE (b)

Low temperature No phase separation

0

1

xB

Fig. II-8(b). The case in which TSBE > HBE resulting in the behavior of EB which gives no chance E E for phase separation. Note HB−B > 0, and SB−B > 0.

48

Fig. II-9

in xB which diverges to infinity at the critical point. Namely the solution in the super critical region consists of two kinds of clusters; one rich in B and the other in W. These clusters are nanoscopic in size but their size grows as the condition approaches that of the critical point and eventually diverges to infinity at the critical point. Within the two-phase region then, these clusters acquire macroscopic size and the system becomes a two-phase liquid mixture. Indeed, scattering experiments reveal dramatically enhanced concentration fluctuations in the super critical tert-butanol−H2 O (with a possible LCST) (Koga, 1984; Nishikawa et al., 1987 and 1989) and acetonitrile−H2 O with the UCST (Nishikawa et al., 2002).

[II-9] AZEOTROPE The azeotrope occurs when the liquid phase composition and the gas phase composition are identical at the equilibrium, or the total vapor pressure as a function of xB shows an extremum, as depicted in Fig.II-9, i. e. dp/dxB = 0 with p = pB + pW . It follows then that, 

dp dxB



 =

   dpB dpW + = 0 dxB dxB

(II-85)

p

Azeotrope

xB

Fig. II-9. The azeotrope, where the total pressure has the extremum, dp/dxB = 0. It can be shown that the composition of the vapor phase and that of the liquid phase are identical, i. e. xB /xW = pB /pW .

Eq. (II-85) to (II-87)

49

Taking into account the Gibbs-Duhem relation. eq. (26) and the expression for the chemical potential using the partial pressure, eq. (41), eq. (85) is rewritten as, 

dp dxB



 =

dpB dxB



xB xW − pB pW

 = 0

(II-86)

Recall that at a critical point dB /dxB  = 0 and hence dpB /dxB  = 0. Thus, unless the system is at a critical point, the following relation holds at an azeotrope, xB p = B xW pW

(II-87)

This occurs quite commonly in aqueous alcohols. As discussed in Chapter III, this situation is problematic in analyzing the vapor pressure data to obtain chemical potentials by the Boissonnas method (1939). We will circumvent this difficulty by using the concentration fluctuation data accessible from scattering experiments.

Chapter III

DETERMINATION OF THE PARTIAL MOLAR QUANTITIES

[III-1] INTRODUCTION While a vast amount of the excess (integral) molar enthalpy data of mixtures exists in the literature, the excess partial molar enthalpy of a component has been rather rarely used in thermodynamic studies, despite its obvious power of providing more detailed information of intermolecular interactions. Arnett et al. (1965a, 1965b, 1966 and 1969) determined the differential molar heats of solution of various alcohols and amines, which are the same as the excess partial molar enthalpies of the solute at the infinite dilution. From their temperature dependence, the heat capacities of solution were calculated at the infinite dilution. This quantity is a third derivative of G and has already been shown to be powerful in providing evidence for the so-called “iceberg” formation. Unfortunately, however, the determination was limited to the infinite dilution. Bertrand et al. (1966), on the other hand, determined the excess partial molar enthalpies of methanol and ethanol in the respective binary aqueous solution for the entire composition range. They used the data for calculating the heats of neutralization of H+ + OH− → H2 O in the mixed solvent of aqueous alcohol. The potential usefulness of these partial molar enthalpy data for understanding the molecular processes in aqueous solutions was not explored until the late 1980s (Koga, 1986; Anderson and Olofsson, 1988). Thus, it may be of some use to review in this chapter the method of determining partial molar enthalpy, together with partial molar entropy and partial molar volume. First, however, we discuss the relationship between the excess (integral) molar enthalpy, HmE , that has been commonly measured, and the excess partial molar enthalpy of the i-th component, HiE , in the way of conversion of the variable system.

[III-2] CALCULATION OF HiE FROM HmE Fig. III-1 shows the plots of the excess (integral) molar enthalpy, HmE , data for ethanol(ET)−H2 O by Costigan et al.(1980). Recall eq. (II-75), which provides the E relationship between the excess partial molar enthalpy of ethanol (or H2 O), HET E E (or HW ), and Hm .

52

Fig. III-1 to III-2

Ethanol(ET) - H2 O at 25 °C

0

H mE

R

P

-200

HET E

H wE

-400

-600

Q

-800

-1000 0.0

0.2

0.4

0.6

0.8

1.0

xET

Fig. III-1. The excess molar enthalpy, HmE , of ethanol (ET)−H2 O(W) at 25  C, according to E Costigan et al., (1980). The excess partial molar enthalpy of ET, HET , at point P is the value of the intercept at xET = 1 of the tangent at point P. The excess partial molar enthalpy of W, HWE , is the value of the intercept at xET = 0 of the tangent at point P. Reproduced with permission from the Australian Journal of Chemistry 33(10):2103–2119 (Costigan, MJ et al.). Copyright CSIRO 1980. Published by CSIRO PUBLISHING, Melbourne, Australia – http://publish.csiro.au/ nid/52/issue/3652.htm.

 E HET = HmE + 1 − xET 



HWE

=

HmE

HmE xET

HmE + 1 − xW  xW

 

(III-1a)



 =

HmE

− xET

HmE xET

 (III-1b)

E What eq. (1) indicates graphically is that the value of HET at point P in the figure is the E value at the intercept at the xET = 1 axis, R, while HW is that at the xET = 0 axis, Q. The line QPR in the figure is the tangent at point P. If we know the function form of HmE , it is E theoretically possible to obtain HET analytically. Unfortunately, however, it is generally difficult to find the exact and true function form covering the entire composition range. The goodness of curve-fitting with a trial function form is judged by the sum of the square deviations between the experimental values and the fitted values. When interest lies in a dilute, H2 O-rich region, for example, the values of HmE are close to zero. For these data points the absolute deviation could be small in comparison with the middle composition region, but the relative deviations could be systematic and substantial. This difficulty in the dilute region could give a devastating error when differentiated

Eq. (III-1a) to (III-2)

53

analytically. Ott et al. (1986) devised the following fitting function to the HmE data for ET−H2 O determined by Costigan et al. (1980), shown in Fig. III-1. HmE = x1 − xexp−axbj 1 − xj + 1 − exp−ax ck 1 − xk 

(III-2)

This function turned out to be quite successful in reproducing the excess partial molar E E . By virtue of eq. (1) and eq. (2), the calculated values of HET , enthalpy of ethanol, HET with j = 0 to 2, and k = 0 to 4, are shown in Fig. III-2 together with our data (Tanaka et al. 1996) measured directly by the method described in the next section. As is evident from Fig. 2, the match is satisfactory. Such success in finding the correct function for HmE in the entire composition range is quite rare. We note that eq. (2) contains the exponential term, which seems to work as a switching function from the first term to the second term dominance. We have applied a similar equation to eq. (2) using the exponential term to the HmE data of the 2-butoxyethanol BE−H2 O carefully determined by Davis et al. (1988) without success. It appears that a much stronger switching function is required for BE−H2 O. Indeed, Davis (1983, 1984, 1985) devised the four segment model in which four different fitting functions are used in each of four mole fraction regions. Douheret et al. (1989) thus fitted a total of 13 parameters including the segment boundaries with reasonable success, 2

0

HET E / kJ mol-1

-2

-4

-6

-8

Measured -10

-12 0.0

0.2

0.4

0.6

0.8

1.0

xET

Fig. III-2. The excess partial molar enthalpy of ethanol (ET) in ET−H2 O at 25  C. The hollow circles are the measured by Tanaka et al. (1996). The line is calculated by curve-fitting by Ott et al. (1986) of the HmE data by Costigan et al. (1980). A rare case of successful curve-fitting to E obtain reasonable values of HET in the entire composition range (see text). The measured data from Tanaka et al. (1996). Reproduced with permission from the Canadian Journal of Chemistry 74, 713, by Tanaka et al., Copyright (1996), with permission from NRC Research Press.

54

Fig. III-3 to III-4

100

BE - H2 O at 25 °C 0

Measured Line - 4 segment fit Broken line - fit by eq(III-3)

HmE / J mol-1

-100

-200

-300

-400

-500 0.0

0.2

0.4

0.6

0.8

1.0

xBE

Fig. III-3. The excess molar enthalpy, HmE , of 2-butoxyethanol BE−H2 O against mole fraction of BE, xBE , at 25  C. The open circles indicate the data by Davis et al. (1988). The line is the four-segment fit by Douheret et al. (1989). Four different fitting functions were used bounded at about the arrows to cover the entire composition range. The broken line is fitted by eq. (III-3). (See Koga(2004a) for details.) Reproduced with permission from 3.1.3 Determination of Partial Molar Quantities, (Koga, Y) in Comprehensive Handbook of Calorimetry and Thermal Analysis, (Sorai, M. Editor-in-chief), Copyright (2005), John Wiley & Sons.

as shown in Fig. III-3. There are, however, small but definite systematic errors in the segE . This effect ment boundaries. These systematic errors will lead to incorrect values of HBE is shown in Fig. III-6 in the next section. We have also tried another fitting function of the all-purpose Abbott-Van Ness type (Van Ness & Abbott, 1982) for the BE−H2 O data,  i HmE ax (III-3) =  i j bj x x1 − x E The resulting HBE data were not satisfactory in the H2 O-rich region, as shown in Fig. III-6 below (see Koga (2004a) for details). Thus, it would be much simpler to determine the excess partial molar enthalpy experimentally.

[III-3] EXPERIMENTAL DETERMINATION OF EXCESS PARTIAL MOLAR ENTHALPY The excess partial molar enthalpy of the i-th component is, by definition, eq. (I-48), the slope of the tangent drawn on the curve H E vs. ni at point P in Fig. III–4. Experimentally,

Eq. (III-3)

55

n j ( j other than i) are fixed

R

HE

δH E/δn i H iE P Q

δn i / 2

δni /2 ni

 E Fig. III-4. Approximation of the slope of tangent at point P, HiE = H , by the slope of ni n j=i  E (see text). Reproduced with permission from 3.1.3 Determination of Partial the cord QR, H ni Molar Quantities, (Koga, Y) in Comprehensive Handbook of Calorimetry and Thermal Analysis, (Sorai, M. Editor-in-chief), Copyright (2005), John Wiley & Sons.

however, the amount of the i-th component, ni , must be perturbed by a small but finite amount ni and the resulting change in the excess enthalpy, H E is determined at the constant pressure. This corresponds to approximating the slope of the tangent at P with the slope of the cord QR in Fig. 4. The goodness of this approximation could be checked by reducing the size of ni and seeing if the resulting quotient H E / ni converges to a constant value. This converged value could be safely assumed to be the desired slope of the tangent. Fig. III-5 is a sketch of the calorimeter for this purpose according to Wadso (1966). It is commonly called an isoperibolic titration calorimeter. It is operated under a semi-adiabatic condition. The air space separating the cell and the water bath, the temperature of which is controlled within ±0001 C, serves as a semi-adiabatic shield. A small amount of heat leak, however, is monitored by a thermal sensor and corrected for. The target component is titrated into the cell containing the mixture. Generally, for non-electrolyte−H2 O systems, a titre of less than 0.02 mol is sufficiently small for the above approximation if the mixture in the cell is about 5 mol. There are a number of commercial titration calorimeters that use a similar or smaller titrant to titrand ratio. The thermal response of the cell on titration is picked up by a thermal sensor which is separately calibrated by a known amount of electric energy. For the calorimeter shown in Fig. 5 a thermistor of 2 k is satisfactory. The process of measurement to titrate nB of B into a solution made of nB of B and nW of W can be described as, nB of B + nW of W +  nB of B → nB + nB  of B + nW of W

56

Fig. III-5 to III-6

Dry nitrogen

Buret

Heat exchanger

Bath Heater

Thermistor

Bath Air space

Fig. III-5. An isoperibolic titration calorimeter. Cell contains about 100 mL of titrand into which about 0.5 mL of a titrant is delivered. A thermistor of about 2 k and a standard heater calibrate the thermal effect of titrant.

The thermal effect qp associated with the process is written as, qp = HnB + nB  nW −HnB  nW  − nB HB ∗ 

(III-4)

where H is the total enthalpy and HB ∗ is the molar enthalpy of pure B. It follows then, qp HnB + nB  nW  − HnB  nW  = − HB ∗  nB nB

(III-5)

and for a sufficiently small nB , qp = nB



H nB

 − HB ∗ = HBE  nW

(III-6)

Eq. (III-4) to (III-6)

57

0 -2

HBE E / kJ mol-1

-4 -6

(a)

-8 -10

Measured

-12

Line - calc. by 4 segment fit -14

Broken line - Abbott / vanNess fit

-16 -18 0.00

0.02

0.04

0.06

0.08

0.10

0.8

1.0

xBE 1

HBE E / kJ mol-1

0

-1

(b)

-2

Measured Line - calc by 4-segment fit

Broken line - calc. by eq.(III-3) -3

-4 0.0

0.2

0.4

0.6

xBE

Fig. III-6. The excess partial molar enthalpy of 2-butoxyethanol (BE) in BE−H2 O at 25  C. The open circles are the measured data by Siu & Koga (1989). The line is calculated by the foursegment fit by Douheret et al. (1989). Broken line is calculated using the Abbott-van Ness fitting, E eq. (III-3). Clearly both fittings are not satisfactory for calculating HBE , although the fit for HmE looks almost acceptable (see Fig. III – 3).

If we titrate W instead of B, we obtain HWE . Fig. III-6 shows the results of such E measurements for 2-butoxyethanolBE−H2 O at 25  C, HBE , the excess partial molar E enthalpy of BE. For comparison, the calculated values of HBE (Douheret et al. 1989) E E using the four segment fit for Hm (Davis et al. 1988) and the values of HBE calculated using eq. (3) (Koga, 2004a) are also shown.

58

Fig. III-7

[III-4] EXPERIMENTAL DETERMINATION OF EXCESS PARTIAL MOLAR VOLUME The method described for determining the excess partial molar enthalpy is completely applicable to volume. Thus, the volume change V is determined on perturbing the amount of the i-th component. Namely, for the process, nB of B + nW of W +  nB of B → nB + nB  of B + nW of W  the volume change V is written equivalently to eq. (6) above as, V = nB



V nB

 − VB ∗ = VBE 

(III-7)

nW

where VB ∗ is the molar volume of pure B. V could be determined directly by dilatometer (Davies et al., 1994). Alternatively, if the density data of the solution are determined accurately with five to six significant figures, using for example an automatic vibrating tube densimeter (Dethlefsen, 1984; Dethlefsen et al., 1984), then the excess partial molar volume can be determined by using the relation, 

VBE

=

VmE

VmE + 1 − xB  xB

 

(III-8)

The derivative of the second term on the right is taken by graphical differentiation (Koga et al., 1993). The pros and cons of graphical and manual differentiation as opposed to the computer garphics technique using B-spline is discussed in Appendix A.

[III-5] EXCESS PARTIAL MOLAR ENTROPY – EXCESS CHEMICAL POTENTIAL In theory, the excess partial molar entropy of the i-th component can be determined by measuring the resulting entropy change of the system on perturbing the amount of the target component i. In practise, however, the direct experimental determination of entropy has been difficult, if not impossible. Instead, we determine the excess chemical potential of the i-th component, Ei , and take advantage of the relation, Ei = HiE − TSiE 

(III-9)

to calculate SiE . Eq. (II-41) shows the way to obtain the chemical potential of a component from its equilibrium partial pressure, pi , i.e.

Eq. (III-7) to (III-10)

59

 pi  = RT ln xi pi∗ 

Ei

(III-10)

where pi∗ is the vapor pressure of pure i at the same temperature. The first step that we take is to determine the total vapor pressure of the head space gas mixture over a given liquid mixture directly. For this purpose capacitance manometers can be used to advantage. A typical sensitivity is 1 mTorr with a full scale 100 Torr. The temperature variation of the sample can be corrected for by the Gibbs-Konovalov correction, shown in Appendix B. Thus, the total pressures are obtained at a fixed temperature, 298.15 K, for example. A simple vacuum system capable of pumping down to 10−6 Torr made of Pyrex glass or stainless steal is sufficient. A typical vapor pressure apparatus is shown in Fig. III-7. The region encircled by a broken line is an air bath, the temperature of which is kept higher than the temperature of the cells, in order to avoid gas condensation. For the same purpose, tape heaters are wound around the tubes connecting cells to the measuring system, in order to keep the temperature of the connecting tubes higher than that of the cells. An amount of each component can be measured in the gas handling manifold with the known volume, temperature and pressure, then transferred Vaccum line Capacitance Manometer Air Bath 5 L bulb 0.5 L

Sample 3

Sample 1

Sample 2 Band Heater

Bellow Valves Cell Water bath Up/Down

Fig. III-7. Schematic diagram of vapour pressure apparatus (see text).

60

to the cell by cooling the latter with liquid nitrogen. During this operation the water bath is lowered. The bath is then raised; the liquid mixture is thawed and mixed with a magnetic spin bar. After thermal equilibration, the pressure is measured by a capacitance manometer. After determining the equilibrium pressure, the gas phase sample may be transferred back to the cell by freezing the cell with liquid nitrogen, or may be sent to gas chromatographic analysis to determine the gas phase composition. For the latter operation, the amount of gas phase lost must be recorded and corrected for the subsequent determination of the mole fraction. For binary systems, numerical analysis based on the Gibbs-Duhem relation is possible due to Boissonnas (1939) without gas analysis. For some ternary systems a similar numerical analysis is also possible. The total pressures of the ternary system must behave monotonously as a function of all the composition variables without showing any extremum. We will show in Chapter IX such an analysis for a tert-butanol−dimethyl sulfoxide−H2 O system (Trandum et al., 1998).

[III-6] BOISSONNAS ANALYSIS – EXCESS CHEMICAL POTENTIAL Table III-1 lists the vapor pressure data at 293.15 K for acetone(AC)−H2 O, determined by Perera et al. (2005). The total pressure, p, increases monotonously as a function of the mole fraction of AC, xAC . This is an ideal situation for the following analysis, as will become evident shortly. Our task is to evaluate the partial pressures of AC and H2 O, pAC and pW , from the data of p against xAC . The Gibbs-Duhem relation, eq. (II-26), is described for the present binary system as, xAC EAC + xW EW = 0

(III-11)

Using partial pressures, eq. (11) is rewritten as,  xAC

pAC pAC



 + xW

pW pW

 = 0

(III-12)

With p = pAC + pW , and xAC + xW = 1, eq. (12) is rewritten by eliminating pW as, pAC = pAC 

(III-13)

where AC is what we call the Boissonnas coefficient and is defined as, AC = 1/1 − pW /pAC xAC /xW  

(III-14)

If we know the values of pAC , pW , xAC and xW at one data point, then the increment in pAC for the next point is calculated by eq. (13) by measuring the change in the total pressure, p. We repeat the process stepwise until all the data points are solved. To start this iteration, however, we have to assume that at the most dilute data point, the

Eq. (III-11) to (III-14)

61

Table III-1. Vapor pressure analysis by the Boissonnas method Acetone (AC) - H2O (W) at 20.00 degC Perera et al. (2005). Starting at xAC = 0 xAC units

p(meas) Torr

pW Torr

pAC Torr

0 000791 001491 002173 002883 003604 00433 005147 006048 007033 008089 009192 010308 01134 012346 013177 015353 018451 022605 027538 032203 035752 041168 050181 059569 06981 075982 081295 086031 090702 09421 097658 099336 099778 1

17527 25497 32293 38675 45012 51256 57251 63592 70191 76932 83671 89915 95725 100631 102053 105112 112145 120051 128446 135746 140955 143956 148237 153642 158513 16372 16713 17043 173883 177659 180658 183588 185 18534 185474

17527 17388 17270 17157 17043 16928 16815 16694 16564 16427 16286 16150 16018 15902 15867 15786 15592 15354 15065 14771 14517 14344 14063 13635 13106 12308 11519 10510 9129 7173 5085 2474 0605 0277 0

0000 8109 15023 21518 27969 34328 40436 46898 53627 60505 67385 73765 79707 84729 86186 89326 96553 104697 113381 120975 126438 129612 134174 140007 145407 151412 155611 159920 164754 170486 175573 181114 184395 185063 185474

Starting at xW = 0 fai(AC)

101739 10177 101803 101842 101878 101918 10197 102029 102097 102173 102267 102364 10246 102662 102756 103017 103432 104038 104866 105768 106563 107915 110877 115313 123149 130579 139981 151803 169613 189133 232334 196534

xW 1 099209 098509 097827 097117 096396 09567 094853 093952 092967 091911 090808 089692 08866 087654 086823 084647 081549 077395 072462 067797 064248 058832 049819 040431 03019 024018 018705 013969 009298 00579 002342 000664 000222 0

pW Torr −3850 −3839 −3826 −3811 −3795 −3779 −3760 −3738 −3715 −3690 −3665 −3641 −3619 −3612 −3597 −3558 −3507 −3444 −3377 −3321 −3284 −3221 −3115 −2986 −2792 −2632 −2441 −2195 −1850 −1497 6634 0984 0278 0000

pAC Torr

fai(W)

29347 000161 36132 0002 42501 000231 48823 000257 55051 000279 61030 000302 67352 000324 73929 000347 80647 00037 87361 000395 93580 000419 99366 000442 104250 000479 105665 0005 108709 000555 115703 000638 123558 000757 131890 000914 139123 001081 144276 001226 147240 001466 151458 001962 156757 002652 161499 003732 166512 004676 169762 005783 172871 00713 176078 009136 179509 011792 182155 277502 176954 −40016 184016 −207583 185062 185474

62

Fig. III-8 to III-9

second point (xAC = 0007907 for the AC−H2 O vapor pressure data listed in Table 1), the system is in the Henry’s law region. Hence, the major component W must obey ∗ ∗ 1 − xAC , where pW is Raoult’s law, as discussed in Section [II-4]. Namely, pW = pW ∗  the vapor pressure of pure H2 O at 20 C, i, e. pW = 17527 Torr. Then, pAC = p − pW . Thus the second data point is now solved. Using eq. (13) the third and the subsequent data points are solved point by point. The results are shown in the third and fourth columns of Table 1, and plotted in Fig. III-8. For the last point at xAC = 09978, the table shows pW = 0277 Torr and pAC = 185063 Torr, which are favorably compared with pW = 0278 Torr and 185.062 Torr calculated by assuming Henry’s law in this most AC-rich point, since the sensitivity of the pressure determination of the gauge used is ±0001 Torr. Hence, we conclude the analysis was satisfactory. We note that the values of AC in column 5 are about 1. Hence the measured increment in p is carried almost directly to the increment of pAC . This means that the uncertainties in pAC are

200

150

p / Torr

Total p pW

100

p AC

50

0 0.0

0.2

0.4

0.6

0.8

1.0

xAC

Fig. III-8. Vapor pressures of acetone (AC)−H2 O at 20  C. The total pressure (measured) increases monotonously as the mole fraction of AC, xAC , increases. Thus, the Boissonnas analysis is possible for the entire composition range. The data are according to Perera et al. (2005). (see text.) Reproduced with permission from the Journal of Chemical Physics 123, 024503 (Perera, A. et al.), Copyright (2005), American Institute of Physics.

63

about the same as those in p. The excess chemical potentials for AC and H2 O, EAC and EW are calculated by eq. (10) and plotted in Fig. III-9. This method is due to Boissonnas (1939). There are many other methods of solving the total vapor pressure data; see Prausnitz et al (1967), for example. However, most of these methods assume a fitting function for the excess Gibbs function, GE , a priori, and hence there is a danger of artificially forcing the thermodynamic property of the system in question to behave in the form of the chosen fitting function. The Boissonnas method, on the other hand, being purely a numerical analysis based on the first principle, the Gibbs-Duhem relation, is free from such a danger. The Boissonnas method, however, starts with the assumption that the most dilute point is in the Henry’s law region. Is this assumption correct? This issue was dealt with by us and it was concluded that the Henry’s law region does not exist even at the mole fraction of a solute as low as 10−5 . (Koga, 1995a; Westh et al., 1998). This still leaves the possibility that Henry’s law is obeyed below 10−5 mole fraction. It is, however, very impractical for determining the vapor pressure at this low mole fraction. About 10−3 is a convenient mole fraction at which to measure the most dilute solution, as is the case for the above AC-H2 O, as well as the alcohol−H2 O (Hu et al., 2003). Thus, the partial

2.0

μ ACE/RT or μ WE/RT

1.5

μ WE/RT 1.0

μ ACE/RT 0.5

0.0

0.0

0.2

0.4

0.6

0.8

1.0

xAC

Fig. III-9. The excess chemical potential of acetone (AC), EAC , and H2 O (W), EW , in AC-H2 O at 20  C due to Perera et al. (2005). They are calculated by eq. (III-10), using the partial pressure data obtained by the Boissonnas method (see text). Reproduced with permission from the Journal of Chemical Physics 123, 024503 (Perera, A. et al.), Copyright (2005), American Institute of Physics.

64

Fig. III-10

pressures at the point determined on the assumption of Henry’s law are in fact in error. This error may be at first sight carried systematically through the entire process of the above analysis. It turned out, however, that eq. (13) or eq. (14) contain a self-correcting feature. (Koga, 1995a). Namely, if the value of pAC is erroneously large, the value of AC , eq. (14), will be unduly small. It follows then that the increment for pAC , pAC , calculated by eq. (13) becomes smaller than the proper value. Thus the value of pAC for the next point will be closer to the real value. If this correction is not enough and the value of pAC is still too high, then the next one could be more correct, and so on. The same corrective feature applies for the case of too low a value of pAC . Depending on how dilute the first point is and also on the size of increments, several points appear to be sufficient to achieve stable (and hence most likely correct) results (Koga, 1995a). For the present AC−H2 O, the first five points of EAC below xAC <0004 which seem to deviate downward could be due to this effect. In deriving eq. (13) from eq. (12) above, pAC could be eliminated instead of pW resulting in, pW = p W 

(III-15)

W = 1/1 − pAC /pW xW /xAC  

(III-16)

The Boissonnas analysis could equally start from the xW = 0 end in the order of descending total pressure. Columns 6 to 9 of Table III-1 (p. 61) show the results of such an effort. As is evident from the table, the value of W varies wildly, resulting in negative partial pressures. This is due to the fact that W is mostly almost zero and the change in p is not numerically conveyed to that in pW . It is therefore mandatory to start the Boissonnas analysis at the low total pressure end and calculate stepwise in the order of ascending total pressure. That the AC values are close to unity can be understood qualitatively in the following manner. For this choice of starting point, the vapor pressures of pure components are such that pAC ∗ is much larger than pW ∗ . If we approximate the xAC -dependence of pAC and xW -dependence of pW as, pAC = pAC ∗ xAC + higher order in xAC  pW = pW ∗ xW + higher order in xW 

(III-17) (III-18)

then the denominator of eq. (13) becomes about 1 − pW ∗ /pAC ∗ +     , and since pW ∗ /pAC ∗ <<1 AC is of the order of unity. In reality, however, a binary aqueous system with a monotonous increasing/decreasing total pressure is rather rare. Besides the AC−H2 O system dealt with above, the methanol−H2 O system is another example (Hu et al., 2003). For dimethyl sulfoxide (DMSO)−H2 O, the total pressure decreases monotonously as the concentration of DMSO increases. Therefore, the partial pressures were calculated from the DMSO rich end

Eq. (III-15) to (III-19)

65

(Lai et al., 1995). Most aqueous solutions have an azeotrope where the total pressure shows the extremum and the mole fraction in vapor and that in liquid phase are identical, as discussed in Section [II-9]. For example, the total vapor pressure data for 1-propanol (1P)−H2 O at 25  C (Hu et al., 2003) are shown in Fig. III-10. It is evident that there is an azeotropy at about x1P = 035. At the azeotropy, pW /p1P x1P /xW  = 1

(III-19)

and the value of 1P and that of W become infinity and p = 0. Thus, eq. (13) or (15) is indefinite and the Boissonnas method no longer works. For such a case we start the Boissonnas method from both ends of mole fraction in ascending order of the total pressure to the maximum. Near the maximum, the value of 1P or W becomes large, of the order or 10 or more, then the uncertainty in p1P or pW is magnified systematically. We then use the knowledge that at the azeotrope the liquid phase mole fraction is equal to the gas phase mole fraction y1P = p1P /p, and interpolate y1P values graphically between the points where the value of the Boissonnas coefficient, 1P or W , is larger than 2, say. The details of this are discussed in the next section.

40

35

azeotropy

p / Torr

30

25

1-propanol - H2O at 25 °C

20

15

0.0

0.2

0.4

0.6

0.8

1.0

x1P

Fig. III-10. The total vapor pressure of 1-propanol (1P)−H2 O at 25  C. (Hu et al., 2003). There is an azeotropy, which was determined to be at x1P = 0390 as discussed in Section [III-7]. Reproduced with permission from the Canadian Journal of Chemsitry 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press.

66

Fig. III-11 to III-12

[III-7] PARTIAL PRESSURES OF 1-PROPANOL (1P)−H2 O The total vapor pressure, p, of 1-propanol (1P)−H2 O, Fig. III-10, shows a rather flat plateau in the mid range of the mole fraction of 1P, x1P . Though there must be an azeotropy, its exact locus is not immediately obvious. The Boissonnas method was applied from both ends in the ascending order of p. The values of 1P and W used to solve partial pressures are plotted in Fig. III-11. In the mid range, 03 < x1P < 06, the values of 1P and W become significantly larger than 1, and the resulting partial pressures would carry some systematic errors in this range. The hollow symbols in the figure indicate these unsatisfactory points. Fig. III–12 shows the plots of y1P vs. x1P . Thus, the data points in this mid range, the points with hollow symbols in Fig. 12, are not reliable. To correct for this, we draw a smooth curve through all the reliable data points with filled symbols by the aid of a flexible ruler, as shown in the figure. We then read the value of y1P off the smooth curve drawn. Using the value of y1P , the value of p1P is recalculated. From the point of intersection of the curve and the line of y1P = x1P , the azeotrope was concluded to be at x1P = 0390 ± 0005 (see Hu et al., (2003) for details). In this manner, the results are obtained without using any fitting function, and free from

10 8 6

φ 1P or φW

4 2 0 -2

φ 1P OK

-4

φ W OK φ 1P > 2

-6

φW > 2

-8 -10 0.0

0.2

0.4

0.6

0.8

1.0

x1P

Fig. III-11. The Boissonnas coefficients, 1P and W , for 1-propanol (1P)−H2 O at 25  C. The analyses with 1P , W > 2 in the range, 03 < x1P < 06, are not reliable. Hollow symbols indicate non-reliable analyses (see text). The data from Hu et al. (2003). Reproduced with permission from the Canadian Journal of Chemsitry 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press.

Eq. (III-20) to (III-21)

67

1.0

0.8

y1P

0.6

0.4

0.2

0.0 0.0

0.2

0.4

0.6

0.8

1.0

x1P

Fig. III-12. The mole fractions of 1-propanol (1P) in 1P−H2 O at 25  C. x1P is that in liquid phase and y1P in the equilibrium gas phase. The data of hollow symbols in the figure are not reliable since the Boissonnas coefficients are larger than 2 (see Fig. III–11 and text). To correct for this, a smooth curve is drawn through all the good points and the value of y1P is read off the smooth curve drawn. From the intersect of the curve and the line of y1P = x1P , we determine the azeotropy to be at x1P = 0390 ± 0005 (See Hu et al. (2003) for details). Reproduced with permission from the Canadian Journal of Chemsitry 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press.

any model. Here again, we prefer this manual graphing method to the computer graphic technique, B-spline, as discussed in Appendix A. The resulting p1P data were used to calculate the excess chemical potential of 1P following eq. (10) as, 

E1P

p1P = RT ln ∗ p1P x1P

 

(III-20)

The total chemical potential of 1P is then, 1P = ∗1P + RT lnx1P  + E1P 

(III-21)

The resulting chemical potential is now compared with the result obtained by a small angle x-ray scattering (SAXS) study (Hayashi et al., 1990). The concentration fluctuation

68

Fig. III-13

5

SAXS, Hayashi et al.(1990) From vapour pressure

RT(1-x1P)/(dμ 1P /dx1P)

4

3

2

1

0

0.0

0.2

0.4

0.6

0.8

1.0

x1P 1P  Fig. III-13. The ordinate quantity, RT 1−x 1P , is the concentration fluctuation that can be deterx1P

mined from the forward scattering intensity of small angle x-ray scattering (SAXS). The hollow square symbols in the figure are the result by Hayashi et al. (1990). The filled circles are calculated from E1P data. See text [III-7] for details. Reproduced with permission from the Canadian Journal of Chemsitry 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press.

mentioned in Section [I-9], is equal to the so-called partial structure factor, Sxx(0) (Bhatia & Thornton, 1970), which is directly obtained from the forward scattering intensity. Thus, 1 − x    Sxx0 = RT  1P 1P x1P

(III-22)

Fig. III–13 shows the values of the left of eq. (23) from SAXS measurements (Hayashi et al., 1990) and the right of eq. (23) obtained from the above vapor pressure analysis (Hu et al., 2003). The match between the two sets of data is satisfactory including the mid region of the above y1P correction, 03 < x1P < 06, assuring the soundness of both sets of determination.

Eq. (III-22) to (IV-3)

Chapter IV

FLUCTUATION AND PARTIAL MOLAR FLUCTUATION – UNDERSTANDING H2 O

[IV-1] INTRODUCTION So far we have dealt with the excess partial molar quantities, Ei  HiE  SiE , and ViE . While Ei is the first derivative of G, the rest are the second derivatives (see Table I-1 (p. 20)). We then took one more compositional derivative and obtained the interaction E E E E E  Si−j , and Vi−j . Of course, Ei−j is the second and Hi−j  Si−j and functions, Ei−j  Hi−j E are the third derivatives. As shown in Table I-1 (p. 20), there are three more second Vi−j derivative quantities, i.e. the heat capacity, Cp , the isothermal compressibility, T , and the isobaric expansivity, p . They are collectively called the response functions, and defined as,    2  S  G Cp ≡ T  (IV-1) = −T T T 2 1 T ≡ − V 1 p ≡ V





V p

V T





1 =− V

1 = V





 2 G  p2

2 G pT

(IV-2)

 

(IV-3)

The partial derivatives are taken keeping all the other variables constant in the p T ni variable system. There is another second derivative, i−j , in Table I-1 (p. 20). This quantity is inversely proportional to the concentration fluctuations and is related to the forward scattering intensity in light, x-ray and neutron scattering experiments. This is dealt with in Section [IV-5] below.

[IV-2] FLUCTUATION FUNCTIONS – COARSE GRAIN The response functions, eq. (1) to eq. (3), are related to the mean-square fluctuation functions in V and S through standard statistical thermodynamics. (Landau & Lifshits, 1980; Debenedetti, 1986; Panagiotopoulos & Reid, 1986; Yan & Chen, 1992). For convenience

70

in dealing with fluctuations, we impose the conditions such that the independent variables p T ni are fixed and their conjugates, S = −G/T and V = G/p are allowed to fluctuate. The mean-square fluctuation, for example, is evaluated by taking the average of the square of the variation of a local/instantaneous value from the global average. Since S and V are extensive quantities, a local/instantaneous value must be evaluated in a coarse grain consisting of an equal amount of each component nzi i = 1 2  m,  for m-component system). The total size of the coarse grain, N z = nzi , must be large enough that the thermodynamic quantities S and V can be defined, but small enough that the fluctuation can be detected, as pointed out by Debenedetti (1986) and Panagiotopoulos and Reid (1986). The fundamental problem at hand is that there is no a priori knowledge about N z . It may not even be universal. For now, however, we use an unspecified N z for the size of coarse grain and rewrite the thermodynamic relations between fluctuations and response functions, taking account of N z explicitly (Koga et al., 1996; Koga, 1999; Koga & Tamura, 2000; Siu & Koga 2005). Namely,  <V > = kT 2

 <S > = kT 2

 <S V > = kT

V p S T V T

 = kT VN Z T 

(IV-4)

= kCpN Z 

(IV-5)

= kT VN Z p 

(IV-6)

 

 signifies the variation of the quantity in a coarse grain from its global average; k is the Boltzmann constant; VN Z is the average volume of the coarse grain; and CpN Z is the average heat capacity of the coarse grain containing N z . We stress that we have no way of knowing N z and that N z may not be universal and may depend on the conditions. For example, as a demixing critical point is approached from a super-critical region, fluctuation becomes progressively vigorous not only in its intensity (amplitude) but also its extent (wavelength). Eventually at the critical point, the extensity (wavelength) reaches a macroscopic size, and hence N z must diverge at the critical point. N z could also be different from material to material. Indeed, it has been pointed out that in liquid H2 O a large number of H2 O molecules fluctuate collectively due to hydrogen bonding (Tanaka & Ohmine, 1987; Ohmine, 1995; Iijima & Nishikawa, 1994). Hence, it is expected that the size of the coarse grain in H2 O is necessarily larger than that in a normal liquid without a hydrogen bond network. Thus, when the values of the mean square fluctuations, eq. (4) to eq. (6) are quoted, the value of N z , hence those of VN Z and CpN Z , must be specified. To circumvent this difficulty, it is common practice (though not explicitly explained) to divide both sides of eq. (4) to eq. (6) by the average volume of a coarse grain, = VN Z , and we obtain

Eq. (IV-4) to (IV-13)

71

V

≡ <V 2 >/k = T T 

(IV-7)

S

≡ <S 2 >/k = CpN Z /VN Z = Cp m /Vm 

(IV-8)

≡ <S V >/k = Tp 

(IV-9)

SV

(for q = V S, or SV) is the notation for convenience. Cp m and Vm are the molar heat capacity and the molar volume. We call these q ’s, eq. (7) to eq. (9), the volume, the entropy and the entropy-volume cross “fluctuation density”, respectively. It is immediately evident that these quantities, q ’s, are devoid of information about N z , and hence they may be indicative of the intensity or the amplitude of fluctuation. q ’s are second derivatives of G. We then define third derivatives, the partial molar fluctuation density for a binary mixture as,  q   q    q = 1 − xi  (IV-10) i ≡ N xi ni

q

which signifies the effect of the i-th component on the respective fluctuation functions. When the fluctuation propensities are to be compared between aqueous and nonaqueous solutions, the extensity of fluctuation could be important. As mentioned above, molecular dynamic studies (Tanaka & Ohmine, 1987; Ohmine, 1995) and an x-ray scattering investigation (Iijima & Nishikawa, 1994) indicated that in H2 O collective fluctuations involving a large number of molecules are prevalent. This suggests that H2 O has a characteristically larger extensity of fluctuation than normal liquids. It is therefore expected, though not rigorously proven, that the size of coarse grain is necessarily larger for H2 O than for normal liquids. Is there any way of identifying this effect of H2 O? It turns out that the mean square normalized fluctuations, q , defined below by eq.(14) to eq.(16) serve this purpose though qualitatively (Koga, 1995b; Koga et al., 1996; Koga and Westh, 1996; Tamura et al., 1999; Koga, 1999; Koga & Tamura; 2000; Siu & Koga, 2005). First, we normalize the deviation of a local/instantaneous value in a coarse grain from the global average by the mean volume of coarse grain. It could very well be more meaningful to normalize by the actual volume of coarse grain. However, for mathematical simplicity, here we use the average volume to normalize with. Thus,   V 2 > = kT T / = kT T /N z Vm  (IV-11) < 

 S 2 < > = k/2 = kCp m /Vm /N z Vm     S V < > = kT p / = kT p /N z Vm 

(IV-12)

(IV-13)

72

The second equalities use the relation, = VN Z = N z Vm , and = CpN Z = N z Cp m . On the other hand, we define the normalized fluctuation functions as, V

 ≡ RT T /Vm 

(IV-14)

S

 ≡ RCp m /Vm 2 

(IV-15)

SV

 ≡ RT p /Vm 

(IV-16)

and the partial molar normalized fluctuation functions as,  q

i ≡ 1 − xi

 q   xi

(IV-17)

We now wish to compare two systems, A and B. For this we consider the ratio of the left-hand sides of eqs. (11) to (13). Using eqs. (14) to (16), the ratios between A and B are written as, 

V <  <  <

S S

2



V > /<

2

A



2 > A /< 

S

B

>

 V  A  V  B   = /  z N Nz

(IV-18)

 S  A  S  B   /  Nz Nz

(IV-19)

2 > B =

     SV  A  SV  B V V   S > A /< > B = /  z N Nz (IV-20)

Here superscripts [A] and [B] designate each system. If we assume that the left-hand sides of eq. (11) to eq. (13) are universal and independent of the system under consideration, then the left-hand sides of eq. (18) to eq. (20) are identically equal to unity. If so, N z A > N z B if q  A > q  B 

(IV-21)

and the converse is also true. Namely, if we see q  A > q  B when calculated by eq. (14) to eq. (16), then we conclude N z A > N z B , and hence the extensity or the wavelength of fluctuation is larger for sample [A] than for [B]. If we relax the above assumption and allow the ratios of the left-hand sides of eq. (18) to eq. (20) to take a positive value less than unity (for the case where the ratio is larger than 1, the labels A and B can be reversed), it follows that the above statement is partly true. Namely, if we find q A  > q  B , then it is necessary but not sufficient that N z A > N z B . Thus, the relative size of q  between the two systems in comparison provides a qualitative indication as

Eq. (IV-14) to (IV-21)

73

to the relative size of coarse grain, which in turn could be parallel to the wavelength of fluctuation of each system. We used both q ’s and q ’s and their temperature and pressure dependences to contrast the nature of H2 O from that of n-hexane (Koga et al.; 1996; Koga & Tamura, 2000). Within the limited applicability of q  discussed above, we learned about the collective nature of fluctuation in H2 O in contrast to a normal liquid, n-hexane, as described below in Section [IV-3].

[IV-3] H2 O vs. n-HEXANE It is well known that H2 O is unique in comparison with normal liquids, its uniqueness being due to its hydrogen bonding capability. (Eisenberg & Kauzmann, 1969; Franks, 1972; Suzuki, 1980). A number of models have been devised for understanding this peculiar nature of H2 O in the past. They may be crudely classified into two categories represented by the mixture model (Roentgen, 1892) and the bent hydrogen bond model (Bernal & Fowler, 1933). The former postulates that liquid H2 O is a mixture of two basic components: One is a bulky ice-like and the other a dense normal liquid. This model has been developed further by many others, recently by Robinson et al. (Bassey et al., 1987; Vedamuthu et al., 1994 & 1995). In the bent hydrogen bond model, liquid H2 O is regarded as forming a more or less completely hydrogen bonded network with a wide distribution in the hydrogen bond strength due to bending – fluctuating widely. This concept has also been extended, an example being a series of work by Sceates et al. aided by spectroscopic data (Sceates et al., 1980). Here we compare H2 O and n-hexane, and see the differences in terms of fluctuation functions developed above. Figs. IV-1, IV-2 and IV-3 are plotted using the data in our earlier paper (Koga & Tamura, 2000). The raw data were taken from the literature (Pruzan, 1991; TerMinassian et al., 1981; Grindley & Lind, 1971; Kell, 1975). If n-hexane, a typical van der Waals liquid, is regarded as an assembly of harmonic oscillators, the mean square volume fluctuation is proportional to the absolute temperature and inversely proportional to the force constant of the harmonic oscillator (Landau & Lifshits, 1980; Yan & Chen, 1992). Hence, the entropy and the entropy-volume cross fluctuation are zero for strictly harmonic oscillators. The volume fluctuations for n-hexane, V and V , in Fig. 1 seem to approach the temperature axis at −273  C, though curling up towards high temperature. The latter curling up effect is no doubt due to non-harmonicity. Hence entropy, Fig. 2, and entropy-volume cross fluctuations, Fig. 3, take small non-zero values. As pressure increases, the force constant increases, and hence V and V  decreases as shown in Fig. 1. Fig. 1(a) suggests that the amplitude of volume fluctuation density, V , for H2 O is in fact smaller than that of n-hexane. At high pressures, 500 MPa for example, the values of V are almost the same and show almost the same temperature-dependence between H2 O and n-hexane, consistent with earlier suggestions that H2 O becomes the same as normal liquids at high pressures (Stillinger & Rahman, 1974; Jonas et al., 1976; Wilbur et al., 1976). V  in Fig. 1(b), on the other hand, shows that the value of V 

74

Fig. IV-1(a) to IV-2(b)

1.2

hollow: n-hexane filled: H2O

1.0

0.1 MPa 50 MPa 200 MPa 500 MPa

0.8

0.6

V

δ / K MPa-1

(a)

0.4

0.2

0.0 -40

-20

0

20

40

60

80

100

120

140

60

70

temp / °C hollow: n-hexane, filled: H2O 0.1 MPa 50 MPa 100 MPa 200 MPa 300 MPa 500 MPa

0.12

0.10

V

Δ

0.08

(b)

0.06

0.04

0.02

0.00

-20

-10

0

10

20

30

40

50

temp / °C

Fig. IV-1. (a) The volume fluctuation density, V , (the amplitude of volume fluctuation) and (b) the normalized volume fluctuation, V , (the amplitude and the qualitative wavelength of volume fluctuation), for H2 O (filled symbols) and n-hexane (hollow symbols). The dashed line is a presumed contribution of the normal liquid-like part of H2 O with no ice-like patches. Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis) 27, 195, (Koga, Y. and Tamura K.), Copyright (2000), Japan Society of Calorimetry and Thermal Analysis.

75

5

filled: H2O, hollow: n-hexane

0.1 MPa 200 MPa 500 MPa

3

S

δ / JK-1 cm-3

4

(a)

2

1

-40

-20

0

20

40

60

80

100

120

140

120

140

temp / °C

3

2

0.1 MPa

S

Δ / J2K-2 cm-6

filled: H2O, hollow: n-hexane

200 MPa

1

500 MPa

(b)

0 -40

-20

0

20

40

60

80

100

temp / °C

Fig. IV-2. (a) The entropy fluctuation density, S , (the amplitude of entropy fluctuation), and (b) the normalized entropy fluctuation, S , (the amplitude and the qualitative wavelength of entropy fluctuation), for H2 O (filled symbols) and n-hexane (hollow symbols). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis) 27, 195, (Koga, Y. and Tamura K.), Copyright (2000), Japan Society of Calorimetry and Thermal Analysis.

76

Fig. IV-3

filled: H2O, hollow: n-hexane 0.8

0.1 MPa 50 MPa

(a)

100 MPa

0.6

SV

δ

500 MPa 0.4

0.2

0.0

-40

-20

0

20

40

60

80

100

120

140

120

140

temp / °C

filled: H2O, hollow: n-hexane 0.1 MPa 50 MPa 100 MPa 200 MPa 300 MPa 500 MPa

SV

Δ / J K-1 cm-3

0.15

0.10

(b) 0.05

0.00

-40

-20

0

20

40

60

80

100

temp / °C

Fig. IV-3. (a) The entropy-volume cross fluctuation density, SV , (the amplitude of the entropyvolume cross fluctuation), and (b) the normalized entropy-volume cross fluctuation, SV , (the amplitude and the qualitative wavelength of the entropy-volume fluctuation) for H2 O (filled symbols) and n-hexane (hollow symbols) The dashed line indicates a presumed contribution of the normal liquid-like part of H2 O with no icebergs. Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis) 27, 195, (Koga, Y. and Tamura K.), Copyright (2000), Japan Society of Calorimetry and Thermal Analysis.

77

for H2 O becomes generally larger than that of n-hexane, indicating that the wavelength of volume fluctuation, or the size of coarse grain, is qualitatively larger for H2 O, as discussed above in Section [IV-2]. At low pressures and temperatures (less than 200 MPa and 50  C), both V and V  for H2 O show an anomalous initial decrease as temperature increases. Dashed lines in Fig. 1 are drawn as a guide for the presumed normal behavior for 0.1 MPa. Thus, the anomalous positive contribution from the dashed line that decreases on increasing temperature is consistent with the putative formation/destruction of ice-like patches of the mixture model of liquid H2 O. At higher temperature the probability of hydrogen bond decreases and hence the formation of ice-like patches is retarded. At pressures higher than 300 MPa, this anomalous positive contribution to V and V  is no longer present. As will be discussed below regarding entropy fluctuation, this is due rather to retardation of forming ice-like patches according to the Le Chatelier principle than a direct reduction of hydrogen bond probability by pressure. Fig. 2 indicates that the entropy fluctuations for H2 O are larger by 3-fold for S and by about 10-fold for S . The entropy fluctuation is associated with the enthalpy fluctuation, i.e. <H 2 > = T 2 <S 2 >, which in turn is indicative of the bond strength fluctuation. Accordingly, this observation suggests that for H2 O the bond strength fluctuates more widely not only in the amplitude but also in the wavelength than that for n-hexane, consistent with the “bent hydrogen bond model”. A high resolution inelastic neutron scattering study suggested that there are two kinds of hydrogen bond strength with the force constant ratio of 2:1 even in ice Ih (Li & Ross, 1993; Li et al., 1994a and 1994b). While criticisms were raised of this “exotic” model by some (Kuo et al., 2005; Tse & Klug, 1995), it seems to be agreed by all that the proton position in ice Ih is highly disordered. For liquid, therefore, an even wider distribution of the hydrogen bond strength is expected, i.e. the bent hydrogen bond model. Both S and S  for H2 O decrease on increasing temperature, in contrast for n-hexane. This is due to a decrease in the probability of intact hydrogen bond for H2 O on temperature increase. Namely, as the temperature increases, the bent hydrogen bonds are bent excessively to be broken. The molecules without hydrogen bonds contribute very little to S and S  as n-hexane. On the other hand, the pressure generally increases S and S  for both H2 O and n-hexane, except for the low pressure and temperature region, below 200 MPa and 50  C for H2 O. The latter region is where the ice-like patches are prevalent as discussed above. Thus, the pressure alters the force constant in a similar manner for both H2 O and n-hexane, and may not reduce the hydrogen bond probability directly. Otherwise, S and S  should decrease for H2 O on pressure increase, as in the case for the temperature increase discussed above. For such normal liquids as n-hexane, when fluctuation causes an increase in volume of a coarse grain, then inevitably its entropy increases. Hence the entropy-volume cross fluctuation is always positive. Since both volume and entropy fluctuations, Figs. 1 and 2, increase on temperature increase, the cross fluctuation should also increase on temperature increase, which is indeed the case for n-hexane, Fig. 3. The effect of pressure on volume, Fig. 1, and the entropy, Fig. 2, fluctuations are the opposite of each other.

78

Fig. IV-4 to IV-5

The resulting cross fluctuations decrease on pressure increase, Fig. 3. Apparently the pressure effect on volume fluctuation is stronger. The peculiar behavior for H2 O in Fig. 3 is naturally interpreted as being due to the putative formation/destruction of ice-like patches at the low pressure and temperature region. Namely, unlike n-hexane, a positive volume fluctuation may be partly due to the formation of ice-like patches which is associated with an entropy decrease, and vice versa. Hence the cross term <S V > will have a negative contribution. This negative contribution is prevalent in the low temperature and pressure region where ice-like patches are putatively formed and destroyed. As the temperature and the pressure increase above 50  C and 200 MPa, this negative contribution diminishes therefore rapidly. At 500 MPa, SV and SV  behave in a similar manner as those for n-hexane. Comparison between Fig 3(a) and Fig 3(b) suggests that the fluctuation wavelength is qualitatively larger for H2 O than n-hexane. We have thus demonstrated the usefulness of the fluctuation functions introduced above in recovering the general understanding of liquid H2 O (Eisenberg & Kauzmann, 1969; Franks, 1974; Suzuki, 1980), and that the mixture model and the bent hydrogen bond model are only two sides of a coin (Stanley & Teixeira, 1980; Blumberg et al., 1984). Furthermore, we have pointed out the importance of distinguishing the amplitude and the wavelength of fluctuation.

[IV-4] SITE-CORRELATED PERCOLATION MODEL OF H2 O As we saw in the entropy fluctuation functions, S and S , in [IV-3], H2 O is characterized by a large fluctuation in hydrogen bond strength in terms of the intensity (amplitude) as well as the extensity (wavelength) in comparison with a normal liquid, n-hexane. This could be understood to be due to the molecular structure of H2 O, together with its hydrogen bonding capability. As shown in Fig. IV-4, the H−O−H angle of

H 104.5°

O

H

Fig. IV-4. H2 O molecular structure. The H-O-H angle is known to be 1045 in the gaseous state, indicating that the sp3 configuration is not perfect. In solid and liquid states, this angle does not expand much (see text and Fig. IV-5). Reproduced with permission from Water and Aqueous Solutions (Suzuki, K.), Copy right (1980), Kyouritsu Shuppan, Tokyo.

79

a gaseous molecule is 1045 , which does not increase much in the condensed phase. (Suzuki, 1980, p. 53). In ice Ih, the bend angle is cited as 107 ± 1 (Kuhs & Lehman, 1986), 1066 ± 15 (Petrenko & Whiteworth, 1999), 1063 ± 04 (Kuo et al., 2005) and 1081 ± 08 (Fanourgakis & Xantheas, 2006). When H2 O freezes to ice Ih, O atoms are organized in a tridymite type structure with four hydrogen bonds to the adjacent O atoms as sketched in Fig. IV-5. The lattice arrangement dictates that the O−O−O angle must be 1095 . The discrepancy by a few degrees between the O−O−O angle and the molecular H−O−H angle is the key for a large fluctuation in S and S . Namely, if one of the two H’s belonging (momentarily) to the central O is on line of the adjacent O−O direction to form a proper hydrogen bond, the other becomes inevitably off line by a few degrees, resulting in a bent and weak hydrogen bond. Indeed, high resolution inelastic neutron scattering experiments showed that there are two kinds of hydrogen bonds in ice; one is weaker than the other. (Li and Ross, 1993; Li et al., 1994a and 1994b). Of course, this distribution occurs momentarily, and it switches back and forth, i.e. the hydrogen bond strength is fluctuating even in ice. While this “exotic” model has been criticized (Kuo et al., 2005; Tse & Klug, 1995), there is no doubt that the location of proton in ice Ih is highly disordered. In liquid H2 O, the lattice structure of O atoms is disorganized. In addition, the bend angle in liquid is shown to be 1061 ± 18 by neutron scattering (Ichikawa et al., 1991) and 1063 ± 49 by simulation (Fanougakis & Xantheas, 2006). The pair correlation function obtained by x-ray scattering experiments

H 109°

106° H

O

Fig. IV-5. The local molecular arrangement of ice Ih. The O atom arrangement is that of tridymite and the O−O−O angle is 109 . The location of the H atom is highly disordered and the average H-O-H angle remains less than 109 . Another way to describe the highly disordered H arrangement is that when one H aligns with one of the four nearest O then the second H will be off-aligned, and hence there are two kinds of hydrogen bonding, one strong and the other weak (see text).

80

Fig. IV-6(a) to IV-6(b)

shows the first nearest population is at about 28 A and the second at 45 A with an interstitial population at 35 A . The total number of O atoms in these configurations is now 4.4, as opposed to 4 in ice. (Eisenberg & Kauzmann, 1980; Suzuki, 1980). This suggests that there is a substantial hydrogen bond network still intact in liquid, with a wide distribution in the hydrogen bond strength, consistent with the bent hydrogen bond model of H2 O (Bernal & Fowler, 1933; Sceates & Rice, 1980). This results in large S and S  for H2 O as we saw in Section [IV-3]. The behaviors of the volume fluctuation functions, V and V , reflect the mixture model of H2 O (Roentgen, 1892; Bassey et al., 1987; Vedamuthu et al., 1994 and 1995), as we saw above in Section [IV-3]. Thus, both models represent the two possible and simplified views of the reality, which of course is the very purpose of devising a model. Stanley et al., on the other hand, have paid attention to the connectivity nature of the hydrogen bond network of H2 O (Stanley & Texeira, 1980; Blumberg et al., 1984). They suggested that for any reasonable definition of the “intact” hydrogen bond, the hydrogen bond network at normal temperatures is bond percolated. Namely, a snapshot of a given moment would show that the hydrogen bond network is connected throughout the entire bulk of liquid H2 O. As the global probability of the “intact” hydrogen bond decreases as temperature increases (or certain solutes are added as is the main theme of this book), a threshold, the bond percolation threshold, could be reached whereby the network connectivity is lost. This connectivity idea turned out to be useful in understanding the effect of a non-electrolyte solute on H2 O in the present studies using the second and the third derivatives of G, as will become evident in Part B. We now follow Stanley and Texeira (1980) and study the bond percolation nature using a simplified two-dimensional square lattice. First they assume a certain (reasonable) cutoff in the energies of hydrogen bonds, so that all possible hydrogen bonds are classified as “intact” or “broken”. Then the global average of the “intact” hydrogen bond can be calculated. For a given hydrogen bond probability, the network connectivity is shown in Fig. IV-6 as a snapshot of a given instance. At every corner of the square lattice an O atom is placed. With a given probability, 25% say, a line is drawn at random for each of four possible directions from each O atom. The results are shown in Fig. IV-6(a). In addition, the O atom which has all four hydrogen bonds intact is marked with an asterisk. As is evident in Fig. IV-6(a), the hydrogen bonds are not connected throughout the entire square, and we cannot walk on bonds through from the top to the bottom of the square. This says that the (hydrogen) bond percolation is not yet established. Figs. IV-6 (b), (c) and (d) are the results of the same exercise for the hydrogen bond probability of 45, 55 and 75%, respectively. We see that the bond percolation is not yet established for (b), but in (c) we can walk through the square lattice from one end to the other, i e. the bond percolation is on. Therefore, the threshold value of the probability above which the bond percolation is established must be between 45 and 55%. It is known that the bond percolation threshold is 50% for the two-dimensional square lattice with four bonds. (Parsonage & Staveley, 1978) Furthermore we note in Fig. IV-6(c) and (d) that the asterisks, i.e. the O atom with four hydrogen bonds, show a clustering tendency, while areas without asterisks are also present. These clusters of O atoms with four bonds can

81

(a) Bond probability 0.25 | o

o | o -- o -- o

o | o -- o -- o o o o

o

o

o -- o o | o o -- o | o o o o | o

o o | o -- o

o | o

o -- o | o -- o

o o | | o o | o -- o | o o o

o

o o

o -- o -- o

o | o -- o | o o | o o

o

o -- o -- o

o

o

o

o | o

o -- o

o o o

o -- o

o o | o -- o |

o

o

o -- o -- o | o o -- o

o -- o

o -- o | o o |

o | o

| o -- o -- o | o -- o o | o o o

o o

o | o

o | o

| o

| o -- o

| o

| o -- o

o

o

o

o

o

o

o | o

o

o | o | o | o --

o -- o o | o

o

o

o -- o

o

o

o

o o | | o -- o -- o -- o -- * -- o | o o o o o o | o -- o o o -- o o | | o o -- o -- o -- o o | o o -- o o o -- o | | o o o o o o | o -- o o o -- o -- o | | o -- o -- o o o o | o -- o o o o o |

-- o

o

o o

o o | o

o

o

o

o

o -- o | o -- o

o o --

o -- o | -- o o | o o --

o | o

o

o

o

o --

o

o

o --

-- o

o o -| o -- o | o o | o o -| o -- o | o -- o

o

o -- o o -- o o | -- o | o

o | o o

o -o o o o | o

Fig. IV-6(a). The bond percolation of 256 molecules of the two dimensional square “H2 O” with 25% hydrogen bond probability. o is the oxygen atom of H2 O. -- shows the intact lateral hydrogen bond.  shows the vertical intact hydrogen bond. ∗ is the oxygen of H2 O which has all four hydrogen bonds intact. Note there is only one ∗ . (b) Bond probability 0.45 o -- o -- o o | | | | o o -- o o | -- o o o -- o | | o o -- o o | o -- o -- o o | | | -- o o -- o -- o | o o o o | o o o -- o | | -- * -- o o o | | -- o -- o o -- o | | o o o -- o | | | -- o o o o | | -- * -- o -- o -- o | o -- o -- o o | o o o -- o | | o o -- * -- o | | |

o -- o o -| | o o -- o | | | o o -- o | | -- o o o | | o o -- o -| | o -- o -- o --

| o | o -- o | o -- o o

| | o -- o o | | o -- o o | o o -- o | | o -- o o -- o | o o o -- o | o -- o -- o -- o | o -- o o -- o | | | o o o -- o | | * -- o o -- o | | | * -- o -- o o | | | | o o o o

o o | | o -- o | o o --

o -- o o | o o o | | o -- o o | | -- o -- o -- o o o o | | | | -- o -- o o -- o o -- o | | | o o o -- o -- o o | | o -- o o o o -- o | o o o -- o o o | o o o -- o o -- o | -- o o -- o -- o o o | | | o -- o -- o -- * -- o -- o | | -- o o o o o o | | | -- o -- o o -- o -- o -- o |

---

---

o o | o -- o

o | o | o -- o

| | -- o -- o -o

o

-- o -- o | o o -| o -- o | | o -- o o o -| -- o o | -- o -- o -- o o | | o -- o

o -- o -- o | | o -- o | -- o o -- o o -| -- o -- o -- o o -- o o -| | | -- o o -- o -- o o -- o | | o

Fig. IV-6(b). The bond percolation of 256 molecules of the two dimensional square “H2 O” with 45% hydrogen bond probability. o is the oxygen atom of H2 O. -- shows the intact lateral hydrogen bond.  shows the vertical intact hydrogen bond. ∗ is the oxygen of H2 O which has all 4 hydrogen bonds intact. Note there are 6 ∗ out of which two of them are next to each other. The bond percolation is not yet established hence one cannot walk through the lattice from the top to the bottom on the bonded path.

82

Fig. IV-6(c) to IV-7

(c) Bond probability 0.55

Fig. IV-6(c). The bond percolation of 256 molecules of the two dimensional square “H2 O” with 55% hydrogen bond probability. o is the oxygen atom of H2 O. -- shows the intact lateral hydrogen bond.  shows the vertical intact hydrogen bond. ∗ is the oxygen of H2 O which has all 4 hydrogen bonds intact. Note the bond percolation is now established; one can walk through the lattice from the top to the bottom. Furthermore there is a clear tendency that oxygen molecules with four bonds cluster together. The bond percolation threshold for the two dimensional square lattice with 4 possible bonds is 0.50. (d) Bond probability 0.75

Fig. IV-6(d). The bond percolation of 256 molecules of the two dimensional square “H2 O” with 75% hydrogen bond probability. o is the oxygen atom of H2 O. -- shows the intact lateral hydrogen bond.  shows the vertical intact hydrogen bond. ∗ is the oxygen of H2 O which has all 4 hydrogen bonds intact. Not only bond percolation is intact but also the clusters of ∗ -oxygens are almost connected, i.e. the situation is close to forming ice. These four pictures, (a), (b) (c) and (d), are generated by a Macro in MSExcel created by Matthew T. Parsons following the idea by Stanley and Teixeira (1980). I thank Matthew T. Parsons for kindly allowing me to use the Macro. Reproduced with permission from the Journal of Chemical Physics 73, 3404, (Stanley, H. E. and Teixeira, J.), Copyright (1980), American Institute of Physics.

83

naturally be regarded as putative ice-like patches that are continually forming/breaking here and there. This is consistent with the mixture model. This clustering tendency of four-bonded O atoms occurs from the topological requirement, in spite of the fact that intact hydrogen bonds are placed randomly with a given probability. In reality, once the hydrogen bond is formed from an O atom to one direction donating H to the adjacent O, then a lone pair of the same O atom has a preference for accepting H from the other direction, i.e. the cooperativity of hydrogen bonding in H2 O (Frank and Wen, 1957; Frank, 1958). Thus, the clustering tendency is stronger in reality than is depicted in Fig. IV-6. The topological connectivity of real H2 O is, of course, not a two-dimensional square lattice, but is of the three-dimensional tridymite type. For this, the percolation threshold is known to be 39% bond probability. (Parsonage & Stavely, 1978). Fig. IV-7 shows a number of estimates of the temperature dependence of the hydrogen bond probability. These estimates are based on a variety of principles under a variety of assumptions. While there is no priority as to which estimate is closer to reality, we note that the estimate by Stanley using the density data (Stanley & Texeira, 1980) reaches 39% at about 80  C. We show below in Section [V-1] that the hydrogen bond probability of bulk 1.0 1

Hydrogen bond probability

0.8

2 3 4

0.6

0.4

65 7 8 9

0.2

0.0 0

20

40

60

80

100

120

Temperature / °C

Fig. IV-7. The temperature dependence of the hydrogen bond probability in H2 O, except for the broken lines which are for D2 O. 1. Haggis et al. (1952). 2, and 3. Stanley & Teixeira (1980). 4. Walrafen (1966). 5. Nemethy & Scheraga (1964). 6. Nemethy & Scheraga (1962). 7 and 8. Marcus & Ben-Naim (1985). 9. Grjotheim & Krogh-Moe (1954). Reproduced with permission from The structure and properties of water, (Eisenberg, D. and Kauzmann, W.), Copy right (1969), Oxford University Press, Oxford.

84

H2 O away from solutes is reduced on addition of 2-butoxyethanol and at a threshold concentration, the behavior of the third derivative thermodynamic quantities suggests the breakdown of the hydrogen bond percolation. This concentration threshold decreases on increasing temperature and it extrapolates to the zero solute concentration (pure H2 O) at about 83  C. This coincidence of about 80  C between the two sets of observations is remarkable and is fully discussed in Chapter V.

[IV-5] CONCENTRATION FLUCTUATIONS AND KIRKWOOD-BUFF INTEGRALS In considering the concentration fluctuations, we choose a coarse grain such that nzj j = i are fixed and the amount of the target component ni only is allowed to fluctuate. As  discussed above, the size of the coarse grain, N z = nzj , must be large enough for thermodynamic quantity such as i to be defined and small enough so that the fluctuation can be detected. However, its value remains unspecified. Thus, the mean square variation of ni  ni , can be written as (Debenedetti, 1986),  i  <ni > = RT/ ni 

2

(IV-22)

The partial derivative is taken keeping p T , and nj=i constant. When nj for j = i are kept constant, xi = ni 1 − xi /N and 

i ni



 i = 1 − xi /N   xi 

Hence eq. (22) is rewritten as,  i N <xi > = RT1 − xi /  xi 

2

(IV-23)

The left-hand side of eq. (23) is called the concentration fluctuation of the i-th component and is equal to the concentration structure factor at the long wavelength limit, Sxx(0), which is accessible by light, x-ray, or neutron scattering experiments. Namely, it is calculated by using the forward scattering intensity, the scattering amplitude, the partial molar volume of each component, the total volume and the compressibility of the mixture (Bhatia & Thornton, 1970; Koga, 1984; Nishikawa et al., 1987; Perera et al., 2005). The right-hand side of eq. (23) can be calculated using the thermodynamic data of Ei as discussed in Section [III-7] for 1-propanol 1P −H2 O. Thus, in

Eq. (IV-22) to (IV-25c)

85

determining the concentration fluctuation, the route via scattering experiments and that from the thermodynamic determination of the chemical potential are complementary. In particular, consideration of relative errors reveals immediately that when the value of the concentration fluctuation is large, scattering experiment is advantageous in that the relative uncertainty is small; about the same as the uncertainty of the scattering intensity measurements. From the thermodynamic route, on the other hand, the denominator, i /xi , is small near the maximum of the concentration fluctuation and its relative uncertainty inevitably becomes large. The concentration fluctuation is closely related to the Kirkwood-Buff integral, Gij , defined as, Gij  ≡



4r 2 gij r − 1dr

(IV-24)

0

where gij r is the pair correlation function between i-j pair of molecules with a distance r apart. On the other hand, Gij for a two component system consisting 1 and 2 can be calculated by using thermodynamic quantities as (Kirkwood & Buff, 1951; Bhatia & Thornton, 1970; Koga, 1984; Nishikawa et al., 1987; Perera et al., 2005), G12 = RTT −  Gii = G12 +

V1 V 2  Vm D

 Vj − Vm /xi  D

(IV-25a)

(IV-25b)

where D is defined as, D≡

x 1 x2  N <x2 2 >

(IV-25c)

and Vi is the partial molar volume of the i-th component (i = 1, or 2, and j = 2, or 1), Vm the total molar volume and T the isothermal compressibility. The denominator on the right of eq. (25c) is the concentration fluctuation given by eq. (23). We note that all these thermodynamic quantities contain the second derivatives of the Gibbs energy, G, as seen in Table I-1 (p. 20). Eq. (25) is a remarkable relation. The right-hand side is calculated from the macroscopic thermodynamic quantities while the left contains the microscopic pair correlation function, albeit in the integral form. While there is some skepticism about the power of Gij because of the integration, (Donkersloot, 1979; Zaitsev et al., 1989), eq. (25) in turn suggests that if we could manage to take a derivative of both sides of eq. (25), then we could have third derivative quantities accessible from thermodynamic measurements on the right-hand side. The “derivative” of Gij on the left might provide some microscopic information of the level of the pair correlation function. This is the main theme of this

86

study as briefly mentioned in Section [I-9] (see eq. (I-53) and the discussion immediately following it). Throughout the remainder of this book, we obtain various third derivatives and demonstrate their usefulness in elucidating the molecular processes in aqueous solutions.

Chapter V

MIXING SCHEMES IN AQUEOUS MONO-OLS

[V-1] MIXING SCHEMES IN 2-BUTOXYETHANOL (BE)–H2 O To describe molecular processes in aqueous (or any other) solutions, the phrase “solution structures” is often used. The word “structure”, however, implies an ordered solidlike structure, so instead we use the term “mixing scheme” throughout this book. 2-Butoxyethanol (abbreviated as BE), C-C-C-C-O-C-C-OH, is the largest mono-ol that mixes with H2 O in the entire composition range at room temperature, although at temperatures higher than 50 C a phase separation occurs with the lower critical solution point (LCST) (Ellis, 1967; Ito et al., 1983; Quirion et al., 1990). At low temperatures, the mixture freezes into ice and the addition compound, BEH2 O38 , or a solid solution consisting of the addition compound and perhaps the solid BE, depending on the composition range (Koga et al., 1994) The phase diagram is shown in Fig. V-1. Our conclusion at the end of this section [V-1] is that there are three composition regions labeled as I, II, and III within a single liquid phase domain in the phase diagram, in each of which the mixing scheme is qualitatively different from those in the other two regions. The hint for this finding came immediately from the composition dependence of the excess E E E partial molar enthalpies and entropies (times T ) of BE and H2 O TSBE HBE  HWE  TSBE , E and TSW . These quantities are determined experimentally as discussed in Chapter III, and are the second derivatives of G. Prior to these second derivative quantities, what were available as thermodynamic quantities are the excess molar (integral) enthalpy and E entropy, HmE and SmE , the first derivative of G  HmE = xBE HBE + xW HWE , for example (see  Table I-1 (p. 20) in Section [I-8]). These values at 25 C are shown in Fig. V-2, together with GEm = HmE − TSmE . Evidently, on mixing BE and H2 O there is a modest enthalpy gain, but the loss of entropy surpasses the enthalpy gain resulting in a positive GEm . Note that in the figure T is multiplied to SmE and its units are the same as HmE and GEm . This practice of using the product of T and entropy is adopted throughout this book. Thus, we could conclude that the BE−H2 O system is rather hard to mix. The degree of difficulty of mixing, i.e. the positive value of GEm , is not large enough, so that the ideal mixing entropy term, RT xBE ln xBE + xW ln xW , is larger in its absolute value than GEm , and the resulting Gm = GEm + RT xBE ln xBE + xW ln xW  is negative. Thus, the system remains mixed in the entire composition. Is there anything more that Fig. 2 suggests? Theoretically, all the thermodynamical information, at least its mole fraction dependence, should be contained in Fig. 2. Being human, however, we have difficulty in extracting all the information from

90

Fig. V-1 to V-3

100 80

Two Liquids

temp. / °C

60 40

III

I

20

II

0

Glass ?

Two Solids -20

Solid Solution 0.0 S

0.1

0.2

0.3

0.4

0.5

0.6

xBE

Fig. V-1. Phase diagram for 2-butoxyethanol (BE)−H2 O. S is an addition compound, possibly a clathrate, BEH2 O38 . The dotted lines indicate the boundary between Mixing Schemes I and II, and that between II and III. The details of Mixing Schemes I, II, and III are discussed throughout this section [V-1]. The data for phase boundaries are from Koga et al. (1994), Ellis (1967), Ito et al. (1983) and Quirion et al. (1990). Reproduced with permission from the Journal of Physical Chemistry, 100, 5172, (Koga Y.), Copyright (1996), American Chemical Society.

BE - H2O, 25 °C

GmE, HmE, or TSmE / kJ mol -1

1.0

G mE 0.5 0.0

H mE

-0.5

TS mE

-1.0 -1.5 -2.0

RT {x BEln(x BE) + x Wln(x W)}

-2.5 0.0

0.2

0.4

0.6

0.8

1.0

x BE

Fig. V-2. The excess molar functions, GEm  HmE , and TSmE for 2-butoxyethanol (BE)−H2 O at 25  C, together with the ideal mixing free energy, RT xBE lnxBE  + xW lnxW .

91

Fig. 2 alone. Often, a model system is constructed and GEm  HmE or TSmE are calculated accordingly. The resulting GEm etc. is then compared with those determined experimentally. We do not follow this route: rather, we determine the second and third derivative quantities directly, without resorting to any model system, and learn from the data. E E We thus show HBE and TSBE in Fig. V-3, and HWE and TSWE , in Fig. V-4, experimentally  determined at 25 C (Koga, 1995c;, 1996;, 2003a). It becomes immediately obvious here that the thermodynamic behaviors and hence the mixing schemes are qualitatively different in three composition regions approximately bounded by arrows in Fig. 3. We label these three regions I, II, and III from the H2 O-rich side, and call mixing schemes operating in each region Mixing Scheme I, II, and III, respectively. In the most BE-rich region, the BE mole fraction, xBE , larger than about 0.6, the E E values of HBE and TSBE are almost zero in comparison with the H2 O-rich region. This

0

H BEE or TS BEE (kJ mol -1)

-10

H BEE

-20

TS BEE

(a)

-30 0.0

0.2

0.4

0.6

0.8

1.0

0

-10

-20

(b) -30 0.00

0.05

0.10

0.15

x BE E Fig. V-3. The measured excess partial molar enthalpy and entropy (times T ) of BE, HBE and E  − TSBE for 2-butoxyethanol (BE) H2 O at 25 C; (a) for the entire composition range, and (b) for H2 O-rich region. See text for arrows. Data from Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 5172, (Koga Y.), Copyright (1996), American Chemical Society.

Fig. V-4

H WE or TS WE (kJ mol -1)

92

H WE

1

TS WE

0

-1

-2

-3 0.0

0.2

0.4

0.6

0.8

1.0

x BE E Fig. V-4. The measured excess partial molar enthalpy and entropy (times T ) of H2 O HWE and TSW  for 2-butoxyethanol (BE)−H2 O at 25 C. See text for arrows. Data from Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 5172, (Koga Y.), Copyright (1996), American Chemical Society.

suggests that BE molecules in the mixture are in almost the same environment locally as in the pure state, and hence BE molecules must exist as clusters of themselves, perhaps in a micellar form. Fig. 4 shows in the same range that the value of HWE is small and constant, while TSWE shows modest composition dependence but the values themselves are also small noting the scale of the ordinate in Fig. 4. These observations suggest that H2 O molecules break away from liquid H2 O and interact with the existing BE clusters with very small enthalpy and entropy changes. If the BE clusters discussed above are of micelle type with −OH group pointing outward, then it could be probable that H2 O molecules interact with −OH’s of the BE clusters one by one, without much influence among H2 O molecules themselves. We call this mode of mixing Mixing Scheme III for BE-H2 O. In the intermediate region, BE-H2 O has a phase separation with the LCST at 50  C and at xBE = 006. At 25  C, therefore, this composition region is super-critical. Indeed, the intensities of light scattering (Ito et al., 1981; 1983) and small angle x-ray scattering (Hayashi and Udagawa, 1992b) show a sharp maximum towards the composition of the LCST at room temperature. On increasing temperature to that of the LCST, the maximum intensity grows rapidly. There is thus no doubt that clusters are present in this composition-temperature range, which grow in size and eventually reach the macroscopic E E size at the LCST. Furthermore, that HBE and TSBE have positive slopes as xBE increases in this region (though much less sharp in comparison with those in the H2 O-rich region) is consistent with the system having a potential of phase separation with an LCST, as discussed in detail in Section [II-8]. Thus, we conclude that BE-H2 O in this region is a mixture consisting of two kinds of clusters: One rich in H2 O and the other in BE. Below

93

50  C these clusters are still sub-macroscopic in size, but grow to the macroscopic scale on reaching the LCST. Above this, two phases exist with compositions at each end of the phase boundary on the isothermal tie line. We call this Mixing Scheme II. The BE-rich clusters are no doubt reminiscent of the clusters in Mixing Scheme III discussed above. E E We now turn to the H2 O-rich region, where both HBE and TSBE change sharply. We E recall that measurements of HBE are made by titrating a minute amount of BE into a large quantity of a given mixture of BE and H2 O. Thus, at the infinite dilution, the value E of HBE is the enthalpic effect of the process in which one mole of BE breaks away from its pure state and mixes into the infinite amount of H2 O. Its value is evidently E is −26 kJ mol−1 . Thus the above process is associated −17 kJ mol−1 , but that of TSBE with a large enthalpy gain, but a larger entropy (time T ) loss, hinting that some ordering occurs in the resulting mixture. This magnitude is only discernable if we consider a strong organization caused in H2 O by BE. Thus, the enhancement of the hydrogen bond network of H2 O in the immediate vicinity of BE, i.e. the “iceberg” formation, is a reasonable interpretation. The “iceberg” concept was formulated by Frank and Evans (1945) using the partial molar entropy (a second derivative of G!) data at the infinite dilution available at the time. It is therefore applicable only in dilute aqueous solutions. Nonetheless, this concept has been taken overly seriously and each time an experimental observation has been made contrary to this concept, it has been criticized or challenged, irrespective of the concentration of solute (for example, Dixit et al., 2002a). Indeed, we suggest throughout this section that it is only in this H2 O-rich region that the “iceberg” concept is partially true. Another confusion in the literature arises regarding the size of the solute in question. The original suggestion was based on the data of non-electrolyte solutes with the molecular weight 100 at most (Frank and Evans, 1945). Thus, the blind application of this concept for functions and structures of large bio-molecules in H2 O is an over-interpretation. Indeed, there is a theoretical work suggesting that the “iceberg” concept is applicable for small hydrophobic moieties. For solutes with large surface areas and curvatures, the hydrogen bonds are broken at all temperatures (Southall and Dill, 2000). There are mixed conclusions regarding the structure of H2 O in the vicinity of flat hydrophobic surfaces (infinite curvatures) (Ball, 2004). See Koga et al. (2004a) for a more detailed discussion. Fig. V-5 shows the isotope effect of the excess partial molar enthalpy of BE (Siu et al., 1992), which supports indirectly the “iceberg” formation. We titrated either deuterated C-C-C-C-O-C-C-OD (BD) or non-deuterated BE into either H2 O or D2 O. The four sets of measurements were carried out at 30  C. The results indicate clearly that it is the solvent (H2 O or D2 O) that dictates the value of the excess partial molar enthalpy of BE or BD, while it makes no difference whether the titrant is BE or BD up to the boundary to Mixing Scheme II. It is well known that the average hydrogen bond strength is higher for D2 O than for H2 O (Eisenberg and Kauzmann, 1969; Suzuki, 1980). The fact that the E E and HBD are larger in D2 O than in H2 O indicates that the main absolute values of HBE thermodynamic effect in Mixing Scheme I is occurring in solvent H2 O or D2 O. This indirectly supports the “iceberg” formation. In Mixing II, on the other hand, whether

94

Fig. V-5 to V-6

H BDE or H BHE / kJ mol -1

0 -2

BH into D2O

-4

BD into D2O

-6 -8 -10 -12 -14

BH into H2O

-16

BD into H2O

-18 0.00

0.01

0.02

0.03

x BD or x BH

Fig. V-5. The measured isotope effect of the excess partial molar enthalpy of 2-butoxyethanol in aqueous solution at 30  C. The deuterated 2-butoxyethanol, C-C-C-C-O-C-C-OD, (BD), or normal 2-butoxyethanol, C-C-C-C-O-C-C-OH, (BH), was titrated into D2 O or H2 O. See text for implication of the results. Data from Siu et al. (1992). Reproduced from the Journal of Chemical Thermodynamics, 24, 159, (Siu, W. W. Y. et al.), Copyright (1992), with permission of Elsevier.

the titrant is BD or BE, and not whether the titrand is D2 O or H2 O, seems to be more E E or HBD . This suggests the onset of BE or BD clustering, important in the values of HBE pertaining to the BE-rich clusters putatively present in this intermediate region. On increasing xBE further, the latter clusters dominate, or rather the other H2 O-rich clusters disappear, and Mixing Scheme III sets in. E at the infinite dilution decreases, On increasing temperature, the absolute value of HBE and the mixing scheme boundary to II seems to occur at a progressively smaller value of xBE , as shown in Fig. V-6 (Koga et al., 1990b). Since the hydrogen bond probability decreases on temperature increase, this observation also suggests that it is H2 O that dictates the thermodynamics in Mixing Scheme I. Returning to the very H2 O-rich region, it is striking that the thermodynamic situation is rapidly changing as xBE increases, as is evident in Figs. 3 and 5; namely the absolute E is rapidly decreasing as xBE increases. If the first bunch of BE molecules value of HBE has gained a large enthalpy, then the second bunch must settle in the system far away from the first so as to gain enthalpy equally, i.e. enthalpy repulsion. Entropy-wise, however, since the first bunch has lost a large entropy (times T ), the second tends to settle in the vicinity so as not to lose as much entropy as the first, i.e. entropy attraction. If these two effects compensate completely, then the ideal mixing entropy dictates a random distribution of BE. However, the entropy attraction is larger than enthalpy repulsion, hence BE molecules tend to attract each other at this composition region, except for the extremely dilute region where the ideal mixing entropy dictates repulsion, as discussed in Section [II-7].

Eq. (V-1a) to (V-1b)

95

0 -2

HBEE / kJ mol-1

-4 -6 -8

29.90 °C 40.00 °C 50.09 °C 56.15 °C 60.13 °C

-10 -12 -14 -16 -18 0.000

0.005

0.010

0.015

0.020

0.025

0.030

xBE

Fig. V-6. The temperature dependence of the excess partial molar enthalpy of 2-butoxyethanol (BE) in BE−H2 O. Koga et al. (1990b). Reproduced with permission from the Journal of Physical Chemistry, 94, 3879, (Koga, Y. et al.), Copyright 1990, American Chemical Society.

Fig. V-7 and Fig. V-8 show the BE-BE interaction functions, the third derivatives, E E HBE−BE and TSBE−BE defined as (see eq. (I-53), eq. (II-76b), and eq. (II-76c)),  E   E  HBE HBE E HBE−BE ≡ N = 1 − xBE   (V-1a) nBE xBE  E TSBE−BE

≡ NT

E SBE nBE





E SBE = T1 − xBE  xBE

 

(V-1b)

E E Both HBE−BE and TSBE−BE are positive, and hence the BE-BE interaction is repulsive (unfavorable) in enthalpy but attractive (favorable) entropy-wise. The value of the latter is larger, and hence the entropy attraction surpasses the enthalpy repulsion. We note, however, at the very dilute solution, the BE-BE entropic repulsion coming from the ideal mixing entropy is large, as discussed in Section [II-7]. The entropic attraction by the excess entropy, as calculated by eq. (1b), compensates the repulsion from the ideal mixing entropy at about xBE = 00065, as indicated in Fig. 8(b). The most striking feature in Figs. 7 and 8 is the sharp initial increase to point X and the E E and TSBE−BE , third derivatives of G. This equally sharp drop to point Y in both HBE−BE resembles the heat capacity anomaly associated with the typical second order phase transition. See Appendix C for various types of anomalies in heat capacity associated with phase transitions and the commentary about the “order” of phase transition. Desnoyers et al. (Roux et al., 1978a & b) observed that the partial molar heat capacity of BE (a third derivative of G) show a sharp peak for BE-H2 O. They called this phenomenon in the

96

Fig. V-7(a) to V-8(b)

1200

25.00 °C 29.90 °C 40.00 °C 50.09 °C 60.13 °C

X 1000

HBE-BEE / kJ mol-1

Error 800 600

(a)

400 200

II I

Y

0

0.00

0.01

0.02

0.03

0.04

0.05

0.06

xBE E Fig. V-7(a). The enthalpic interaction between solutes in 2-butoxyethanol (BE)−H2 O HBE−BE , at various temperatures. The lines drawn are a guide for the eye; they are used to locate point X, the presumed maximum, which is equivalent of the inflection point in the H2 O-rich region for the E HBE vs. xBE curves shown in Figs. 3, 5, and 6. Point X marks the onset of transition from Mixing Scheme I to II, and point Y is the end point of transition. Mixing Scheme II is operative to the right of point Y. Mixing Scheme III is apparent in Fig. 7(b). Koga et al. (1990b). Reproduced with permission from the Journal of Physical Chemistry, 94, 3879, (Koga, Y. et al.), Copyright 1990, American Chemical Society.

50

(b) 25 °C

HBE-BEE / kJ mol-1

40

30

II

20

III

10

0

-10 0.0

0.1

0.2

0.3

0.4

0.5

0.6

xBE E Fig. V-7(b). The same as Fig. 7(a) but an expanded view for a small HBE−BE and a large xBE region  for 25 C only. The boundary between Mixing Schemes II and III is now apparent. Tanaka et al. (1996). Reproduced with permission from the Canadian Journal of Chemistry, 74, 713, (Tanaka, S. H. et al.), Copyright (1996), NRC Research Press, National Research Council of Canada, Ottawa.

97

1400

X

1200

(a)

TSBE-BEE / kJ mol-1

1000 800 600 400

II

I

200

Y 0 0.00

0.01

0.02

0.03

0.04

0.05

xBE

Fig. V-8(a). The entropic interaction (times T ) between solutes in 2-butoxyethanol E E E (BE)−H2 O TSBE−BE , at 25  C. Note that the values of TSBE−BE is larger than those of HBE−BE shown in Fig. 7. Namely, the entropic BE−BE attraction is stronger than the enthalpic repulsion, E in terms of excess quantity. Calculated using TSBE data (Koga, 1991).

TSBE-BEid or -TS BE-BEE / kJ mol-1

0 -200 -400

(b) -600 -800 -1000

TS BE-BEid

-1200

-TSBE-BEE

-1400 0.00

0.01

0.02

0.03

0.04

0.05

xBE

Fig. V-8(b). Comparison between the excess entropic attraction between solutes in 2E id butoxyethanol (BE)−H2 O TSBE−BE , and the ideal mixing entropic repulsion, TSBE−BE . See text for detail.

98

Fig. V-8(c) to V-9

100

(c)

TSBE-BEE / kJ mol-1

80

60

II

40

III

20

0 0.0

0.2

0.4

0.6

0.8

1.0

xBE E Fig. V-8(c). An expanded version of Fig. 8(a) in small TSBE−BE and large xBE range. The boundary E between Mixing Schemes II and III is apparent. Calculated using TSBE data (Koga, 1991).

BE-H2 O the “pseudo-phase” transition, by noting its resemblance to the anomaly of heat capacity (a second derivative quantity) for a phase transition. Here we suggest that the anomalies of the third derivatives of G shown in Figs. 7 and 8 are associated with the transition of mixing scheme, the change-over from Mixing Scheme I to II. The “phase” is a macroscopic entity and its transition is understood as the change in some long-range order parameters, which is accompanied by anomalies of the second derivatives of G. Here, the “mixing scheme” is a molecular arrangement of the mixture within a single liquid phase, and the anomalies are observed in the third derivatives. Hence it is more subtle, perhaps associated with some short- or medium-range order parameters. Thus, the boundary from Mixing Scheme I to II in the BE−H2 O is associated with a peak type anomaly in third derivatives of G. E E and TSBE−BE extrapolate Another point to note in Figs. 7 and 8 is that both HBE−BE E E to non-zero values at xBE = 0; or equivalently, HBE and TSBE do not show zero slope towards xBE = 0 in Fig. 3. If there were a finite range where their slopes were zero, it would be the Henry’s law region as discussed in Section [II-4]. Since we did not observe this at least down to the first measurement at xBE ≈ 0001, we expanded the determination E to xBE ≈ 1 × 10−5 without any sign of Henry’s law (Westh et al., 1998). This of HBE means that the BE-BE interaction is of extremely long range, or a BE molecule interacts with another BE in the presence of 2 × 105 molecules of H2 O. Therefore, if we attempt to understand Mixing Scheme I by means of computer simulation, the size of a coarse grain ought to be at least an order of magnitude larger than this number, i.e. at least 106 molecules of H2 O are required. As discussed in Section [IV-4] above, liquid H2 O has its characteristic bond-percolated hydrogen bond network, and such a long-range BE-BE

99

Hydrogen Bond Probability

interaction might be occurring via the hydrogen bond network. This is indeed the case as will be discussed below in Section [V-3]. We now make an attempt to understand the details of molecular processes in Mixing E E Scheme I. Returning to Fig. 3, the absolute values of HBE and TSBE decrease sharply as xBE increases. The first bunch of BE alters the molecular organization of H2 O by enhancing the hydrogen bond network of H2 O in the immediate vicinity, but there must be some additional effects. One possibility for such a long-range interaction could be that the hydrogen bond probability profile may be established such as sketched in Fig. V-9. Namely, the “first BE” enhances the probability to unity in its immediate vicinity. At the same time, it affects the probability profile to a long-range as depicted in Fig. 9 over the average probability for pure H2 O. Then the second needs to exert a lesser effect on the hydrogen bond probability profile. If this process continues as xBE increases, when it completes its course, the icebergs may completely fill the entire system. At this point there is no room to accommodate BE any longer, marking the end of Mixing Scheme I. More BE introduced in the system now must cluster together pertaining to the BE-rich cluster in Mixing Scheme II. If this picture is correct and the ice-like structure fills the entire space, then the ionic conductivity of H+ and OH− should increase in the BE-H2 O as xBE increases up to the boundary between Mixing Schemes I and II. Such a structure would facilitate H+ hopping along the completed hydrogen bond network. Contrary to this expectation, the ionic conductivity in fact decreased (Koga et al., 1995a), as shown in Fig. V-10. The ionic conductivity of H+ OH− was calculated from those of KCl, KOH, and HCl at the infinite dilution. This result was the first hint that the iceberg does

1.0

0.5

Average for pure H2O

1st BE

2nd BE

0.0

Space

Fig. V-9. A presumed snapshot (Davies et al., 1994) of the hydrogen bond probability profile for Mixing Scheme I in 2-butoxyethanol (BE)−H2 O. As will become evident, this picture is wrong (see text [V-1]). See Fig. 14 for a more correct picture at the present stage. Reproduced from the Journal of Solution Chemistry, 23, 339, (Davies, J. V. et al.), Copyright (1994), with kind permission from Springer Science and Business Media.

100

Fig. V-10 to V-11(b)

35 30

0Λ' j

/ Ω-1 cm-1

25

j=

20

H+OH-

15

H+Cl10

K+OHK+Cl-

5

I 0 0.00

II

0.01

0.02

0.03

0.04

0.05

xBE

Fig. V-10. The mole fraction conductivity of the j-th species, 0 j , in 2-butoxyethanol (BE)−H2 O at 25  C. j = HOH, HCl, KOH, or KCl (Koga et al., 1995a). See text. Reproduced with permission from the Canadian Journal of Chemistry, 73, 1294, (Koga Y. et al.), Copyright (1995), NRC Research Press, National Research Council of Canada, Ottwa.

not fill the entire system, but rather the hydrogen bond probability of bulk H2 O away from iceberg-clad BE may in fact be decreasing as xBE increases. The remainder of this section is devoted to a discussion of this issue. Fig. V-11 shows another second derivative of G, the excess partial molar volumes of E (Koga et al., 1993). As is evident in Fig. 11(a), VWE increases H2 O and BE, VWE and VBE at first, supporting the iceberg formation, but then it decreases before the I-II boundary. Furthermore, if the iceberg would fill the entire system at the boundary as argued above, −1 the value of VWE would be close to that of ice, about 15 cm3 mol ! Fig. 11(a) shows an −1 increase of the order of 0001 cm3 mol , some 1500 times smaller. There must therefore be a different kind of additional effect from what we argued above operating to reduce VWE progressively and to reduce the ionic conductivity of H+ OH− . Before settling this issue, we now turn to the temperature dependence of the boundary between Mixing Schemes I and II. As we argued above, the anomaly of the third derivaE E and TSBE−BE , represents the change-over tive quantity, in particular, point X of HBE−BE point. As introduced in Section [I-8], we label these third derivatives by listing the E ; independent variables of differentiation as T nBE  nBE . Fig. V-12 shows VBE−BE p nBE  nBE  (Koga, 1992) defined as,  E VBE−BE

≡N

E VBE nBE





E VBE = 1 − xBE  xBE

 

(V-2)

Eq. (V-2)

101

0.02 0.00

VWE / cm3 mol-1

-0.02 -0.04 -0.06 -0.08 -0.10

(a)

-0.12 -0.14 0.00

0.01

0.02

0.03

0.04

xBE

Fig. V-11(a). Excess partial molar volume of H2 O in 2-butoxyethanol (BE)−H2 O VWE , at 25  C (Koga et al., 1993). The arrow indicates the transition of mixing scheme from I to II. Reproduced from the Journal of Chemical Thermodynamics, 25, 51, (Koga et al.), Copyright (1993), with permission from Elsevier.

-3 -4

VBEE / cm3 mol-1

-5 -6 -7 -8

(b)

-9 -10 -11 0.00

0.01

0.02

0.03

0.04

xBE E Fig. V-11(b). Excess partial molar volume of 2-butoxyethanol (BE) in BE−H2 O VBE , at 25  C (Koga et al., 1993). The arrow indicates the transition of mixing scheme from I to II. Reproduced from the Journal of Chemical Thermodynamics, 25, 51, (Koga et al.), Copyright (1993), with permission from Elsevier.

102

Fig. V-12 to V-13

800

15 °C 25 °C

VBE-BEE / cm3 mol-1

600

400

200

0

-200

-400 0.00

0.01

0.02

0.03

0.04

xBE E Fig. V-12. The volumetric interaction between solutes in 2-butoxyethanol (BE)−H2 O VBE−BE , at  15 and 25 C (Koga, 1992). Reproduced with permission from the Journal of Physical Chemistry, 96, 10466, (Koga, Y.), Copyright (1993), American Chemical Society.

E An anomaly similar to that of HBE−BE is apparent with point X being at about the same locus of xBE . We calculated from the second derivative quantities a variety of third derivatives at various temperatures. They include the nW -derivative of the excess partial molar volume of W, p nW  nW  (Koga, 1992); the nBE -derivative of the excess partial molar heat capacity of BE, T T nBE  (Roux et al., 1978a & b); the nBE -derivative of the compressibility, p p nBE  (Koga et al., 1995b); the nBE -derivative of the thermal expansivity, p T nBE  (Davies et al., 1992); the T -derivative of the heat capacity at E a given xBE  T T T  (Westh et al., 1994; Atake and Koga, 1994), and HBE−BE at various temperatures, T nBE  nBE  (Koga et al., 1990b). They all show similar peak anomalies except for T T T  which showed a step type of anomaly. The loci of anomalies (point X) of all these third derivatives of G are plotted in Fig. V-13 (Koga et al, 2004a) together with the phase diagram for BE-H2 O (Koga et al., 1994; Ellis, 1967). The plots seem to form a single line that separates the xBE -T field into the regions of Mixing Schemes I and II. This line is sometimes called the Koga-line (Hvidt, 1993; Perera, 2002; Nishikawa et al., 2003). In the region above and to the right of the Koga-line (beyond point Y), Mixing Scheme II is operative in which two kinds of clusters are mixed with each other, one rich in H2 O and the other in BE. In terms of the percolation nature of the hydrogen bond network of H2 O, Mixing Scheme II is devoid of the hydrogen bond percolation. This hints that the Koga-line marks the end of the hydrogen bond percolation nature that exists in Mixing Scheme I all the way down to xBE = 0, the pure H2 O. We recall the site-correlated percolation model of pure H2 O by Stanley & Teixeira (1980) discussed in Section [IV-4]. We also recall that their estimate for the average hydrogen bond

103

100

L+L

80 60 40

Temperature (°C)

20

I

II

I

II

(a)

0 -20 20

10

(b)

0

tP -10

-20

Solid Sol.

Ice + S

-30 0.00

0.02

0.04

S

0.06

0.08

0.10

xBE

Fig. V-13. Mixing scheme and phase diagram for 2-butoxyethanol (BE)−H2 O (Koga et al., 2004a). S is an addition compound of BEH2 O38 . tp is its incongruent melting point. Reproduced with permission from the Journal of Physical Chemistry A, 108, 3873, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

probability decreases on increasing temperature, and reaches 39%, the bond percolation threshold for ice Ih (or tridymite) type bond connectivity at about 80  C. Fig. 13 indicates that the Koga-line seems to cut the xBE = 0 axis at about 83  C. This general coincidence is a circumstantial support for the above conjecture that the Koga-line is the hydrogen bond percolation threshold in BE-H2 O. Thus, Mixing Scheme I is such that BE enhances the hydrogen bond network of H2 O in the immediate vicinity of BE, i.e. the iceberg formation. At the same time the hydrogen bond probability of bulk H2 O away from the iceberg-clad BE decreases concomitantly as if the global temperature increases. As xBE increases, and more icebergs form, the hydrogen bond probability of bulk H2 O away from BE decreases progressively to 39%, whereupon the bulk H2 O loses its integrity as the liquid H2 O and Mixing Scheme II sets in. This will explain the decrease in the H+ OH− conductivity (Fig. 10) and the minute initial increase in VWE in Fig. 11(a). E On the other hand, the initial decrease in VBE (Fig. 11(b)) is the consequence of the E initial increase in VW (Fig. 11(a)), according to the Gibbs-Duhem relation, but could be

104

Fig. V-14 to V-15(a)

Hydrogen Bond Probability

explained at the molecular level as follows. The molecular size of BE is much larger by about 8-fold than that of H2 O. Hence “sands filling the gaps of pebbles” analogy would result in a negative excess partial molar volume of BE. However, H2 O is not just a small molecule, but is bulkier due to the hydrogen bond network with putative ice-like E and the net result is the actual patches. This bulkiness contributes positively to VBE E value of VBE at xBE = 0. As xBE increases, the hydrogen bond probability of bulk H2 O E deceases decreases where incoming BE settles. Hence the positive contribution to VBE E E progressively, resulting in the initial decrease in VBE . The sharp increase in HBE−BE and E must then be coming from the competition between the short-range effect of TSBE−BE iceberg formation and the long-range effect from the reduction of the hydrogen bond probability of bulk H2 O. Thus, the hydrogen bond probability profile (Fig. V-14) may be more appropriate in place of the scenario of the iceberg filling the entire system (Fig. 9). Another important point to note in Fig. 13 is that the Koga-line ends at the incongruent melting point, −37  C, of an addition compound labeled as “S” in the phase diagram. The composition of S translates to BEH2 O38 , and 38 H2 O molecules are just enough to cover the surface of BE, using the respective molar volumes and assuming both are spherical. This suggests that at the ambient condition within Mixing Scheme I the solution is preparing to form the addition compound on freezing, providing more circumstantial evidence for the existence of the icebergs with BE in the middle in solution. On cooling BE-H2 O in Mixing Scheme I, H2 O in the bulk away from the icebergs precipitates as ice at the liquidus curve, and eventually the addition compound BEH2 O38 forms

1.0

Average for pure H2O

0.5

Percolation Threshold

1st BE

2nd BE

0.0

Space

Fig. V-14. A snapshot of the hydrogen bond probability profile for Mixing Scheme I of 2butoxyethanol (BE)−H2 O. In the immediate vicinity of BE the probability is one, but in the bulk H2 O away from solutes the probability is reduced from the average for pure H2 O. When the bulk average reaches the hydrogen bond percolation threshold, the hydrogen bond network loses its infinite connectivity and the mixture consists of two kinds of clusters, one rich in H2 O and the other in BE, i.e. Mixing Scheme II sets in.

105

at −37  C. Since the hydrogen bond probability of bulk H2 O is reduced by the presence of the iceberg-clad BE, it requires a lower temperature than 0  C to form ice. This is a molecular level understanding of the freezing point depression normally deduced by thermodynamic reasoning alone; see Appendix D for thermodynamic arguments for freezing point depression. In summary, Mixing Scheme I is such that BE molecules are surrounded by a number of H2 O molecules whose hydrogen bond probability is enhanced, i.e. iceberg formation. Concomitantly, however, the hydrogen bond probability of bulk H2 O away from the iceberg-clad BE is reduced, as shown in Fig. 14 as a snapshot of the hydrogen bond probability profile. The probability is still high enough initially until it reaches the bond percolation threshold. Hence, the hydrogen bond percolation is intact up to the threshold composition. In other words, H2 O protects its integrity as the liquid H2 O though sacrificing the hydrogen bond probability of bulk H2 O against invading BE. When xBE reaches the threshold value, the hydrogen bond percolation is broken and Mixing Scheme II sets in. As will be shown in Setion [V-3] below, the H2 O molecules in the iceberg do not participate in the characteristic hydrogen bond strength fluctuation; namely, icebergs are solid units that would eventually form the addition compound. Fig. V-15 shows a sketch of the Mixing Schemes I, II and III, by representing H2 O as a two-dimensional square lattice shown in Fig. 6. One remaining point to discuss about the boundary between Mixing Schemes I and E E II is that the anomaly in HBE−BE (and TSBE−BE ) starts at point X and seems to end at point Y as seen in Figs. 7(a) and 8(a). Why does the boundary seem to have this width? Our conjecture is that the width reflects the fluctuating hydrogen bond strength distribution. Point X in Figs. 7(a) and 8(a) is where the strongest hydrogen bond in distribution reaches the bond percolation threshold, while at point Y, the weakest and hence all the hydrogen bonds lose the network connectivity. At present there is no experimental evidence to support this, and until such evidence becomes available, it remains just a conjecture. The boundary from Mixing Scheme II to III, on the other hand, is not as conspicuous as that between I and II. As shown in Figs. 7(b) and 8(c), Mixing Scheme I:

Fig. V-15(a). A snapshot of Mixing Scheme I (Parsons, 2001). H2 O is represented by the twodimensional square lattice shown in Fig. IV-6. Large filled circles represent BE molecules. In their immediate vicinity the hydrogen bond network is enhanced (iceberg formation). The bulk H2 O away from BE has a lower hydrogen bond probability than in the pure state, but the hydrogen bond percolation is still intact. The BE−BE interaction is occurring via the percolated hydrogen bond network. See text.

106

Fig. V-15(b) to V-17

Mixing Scheme II

H2O rich cluster

BE rich cluster

Fig. V-15(b). Mixing Scheme II is such that the system is a mixture of two kinds of clusters, one rich in H2 O, and the other in BE. The H2 O-rich clusters are reminiscent of the icebergs in Mixing Scheme I, and the BE-rich clusters are predominant in Mixing Scheme III. Parsons (2001).

Mixing Scheme III:

Fig. V-15(c). Mixing Scheme III for 2-butoxyethanol (BE)−H2 O (Parsons, 2001). BE molecules form clusters of its own kind and H2 O interacts with the BE clusters as a single molecule. For BE, the cluster could be of micelle type with OH groups pointing outward. H2 O molecules interact predominantly with OH group one by one, without lateral interaction with other H2 O. The large circle represents OH group, small circle H2 O molecule and wiggled line the alkyl chain in BE.

the boundary is located as the intersection of two straight lines crossing with a shallow E angle. Hence locating the II-III boundary using the xBE dependence of HBE−BE and/or E E E TSBE−BE is problematic. This is due to the fact that the values of HBE and TSBE are small in Mixing Scheme III. (This fact was instrumental in arriving at the interpretation of Mixing Scheme III.) This difficulty is circumvented below in Section [V-5].

[V-2] MIXING SCHEMES IN OTHER MONO-OLS (AL)–H2 O E Figs. V-16, 17 and 18 show the excess partial molar enthalpy, HAL , entropy (times T ), E E TSAL , and volume, VAL , of alcohol (AL) in the binary AL−H2 O at 25  C, respectively. Alcohols studied are methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P) and tert-butanol (TBA). Together with 2-butoxyethanol (BE) discussed above in Section [V–1], these alcohols mix with H2 O in the entire composition range at about room

107

5

HALE / kJ mol-1

0

-5

ME ET

-10

1P 2P

-15

TBA

-20 0.0

0.2

0.4

0.6

0.8

1.0

xAL (AL = ME, ET, 1P, 2P, or TBA) E . Alcohols Fig. V-16. The excess partial molar enthalpy of alcohol (AL) in AL−H2 O at 25  C HAL are methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), and tert-butanol (TBA). Data E are from Tanaka et al. (1996), except for TBA. The values of HTBA are extrapolated to 25  C using E the data at 2690  C (Koga, 1986; 1988a) and the T -dependence of HTBA . See Koga et al. (1990a).

5

T SALE / kJ mol-1

0

-5

ME -10

ET 1P

-15

2P -20

TBA

-25

0.0

0.2

0.4

0.6

0.8

1.0

xAL (AL = ME, ET, 1P, 2P, or TBA) E Fig. V-17. The excess partial molar entropy (times T ) of alcohol (AL) in AL−H2 O at 25  C TSAL . The data are taken from Koga et al. (2003) except for TBA (Koga et al., 1990a). Reproduced with permission from the Canadian Journal of Chemistry, 81, 150, (Koga, Y. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

108

Fig. V-18 to V-19

2

VALE / cm3 mol-1

0 -2 -4

ME, 25 °C ET, 25 °C 1P, 25 °C 2P, 25 °C TBA, 26 °C

-6 -8 -10 -12 0.0

0.2

0.4

0.6

0.8

1.0

xAL (AL = ME, ET, 1P, 2P, or TBA) E Fig. V-18. The excess partial molar volume of alcohol (AL) in AL−H2 O VAL . The density data for methanol (ME), ethanol (ET), 1-propanol (1P) are taken from Benson and Kiyohara (1980). The density data for 2-propanol (2P) are from Sakurai (1988), and those for tert-butanol (TBA) are from Sakurai (1987).

temperature. There are no other mono-ols that mix with H2 O in the entire composition. As is immediately evident from the figures, their thermodynamic behaviors are qualitaE E tively the same as that for BE−H2 O: Namely, HAL and TSAL (Figs. 16 and 17) show a sharp increase in the H2 O-rich region, while both stay almost constant and close to zero E in the AL-rich region, xAL >06. The excess partial molar volume of AL, VAL (Fig. 18) shows the characteristic initial decrease as for BE−H2 O (Fig. 11(b)). It is thus likely that Mixing Schemes I and III are equally operative in all AL−H2 O systems. As for the E E intermediate region, while TSAL shows a slow increase for all AL as for BE-H2 O HAL decreases for all AL but ME. However, as Fig. V-19 indicates, the excess chemical potential of AL is a decreasing function of xAL for all AL including BE, which is a requirement for a phase separation to occur. Furthermore, as will be discussed below in Section [V-4], there are a number of light or x-ray scattering works showing a large scattering intensity in this intermediate region for all AL including BE, suggesting the presence of clusters. Thus, we tentatively suggest that Mixing Schemes I, II, and III are equally operative in all AL that mix with H2 O at room temperature, the same as in BE−H2 O. We discuss Mixing Scheme II in Section [V-4] and III in Section [V-5] in some more detail, and concentrate on Mixing Scheme I in this section. E E E Table V-1 lists the values of EAL  HAL  TSAL and VAL at the infinite dilution together with limited existing data for comparison. (The other existing data are mostly limited at E the infinite dilution.) As for BE-H2 O, the absolute value of TSAL is always larger than E E that of HAL , which is the consequence of AL to be positive. The latter is a requirement for the mixture to have the possibility of phase separation as discussed above and in

109

10

ME ET 1P 2P TBA

μALE / kJ mol-1

8

6

BE

4

2

0

-2

0.0

0.2

0.4

0.6

0.8

1.0

xAL (AL = ME, ET, 1P, 2P, TBA or BE )

Fig. V-19. The excess chemical potential of solute in alcohol−H2 O EAL , at 25  C. The data for methanol (ME), ethanol (ET), 1-propanol (1P), and 2-propanol (2P) are taken from Hu et al. (2003), those for tert-butanol (TBA) from Koga et al. (1990a), and those for 2-butoxyethanol (BE) from Koga (1991). Reproduced with permission from the Canadian Journal of Chemistry, 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

Section [II-8]. Also shown in the list are the data for 1-butanol (1B), 1-pentanol (1Pe), 1-hexanol (1Hx), 1-Octanol (1Oc), and 1-Decanol (1De). All these latter alcohols have de-mixing gaps with H2 O, which are listed in Table V-2. For 1B−H2 O, it was shown that within the narrow solubility limit in the H2 O-rich region, x1B <0019, Mixing Scheme I is operative right up to the phase separation. (Tamura et al, 2000) The de-mixing is after all the macroscopic manifestation of Mixing Scheme II. We have not studied the H2 O-rich region of other large 1-mono-ols due obviously to the narrowness of the one phase range in these alcohols (see Table 2). Turning back to the data shown in Table 1, since the only variation in the series of AL listed is in the size of the alkyl group, the quantitative differences shown in the table reflect the effect of the hydrophobic moiety on H2 O. Hence it is tempting to rank the hydrophobicity of each AL by the data in the E E or TSAL at xAL = 0 table. As is evident, the ranking by the value of EAL , and that by HAL E E E do not provide a consistent result. The values of AL  HAL and TSAL at the infinite dilution indicate the thermodynamic effect of one mole of AL (a) breaking away from its pure environment and (b) producing a void in H2 O of the size of AL, and then (c) letting H2 O accommodate AL by altering the molecular organization of H2 O. The net results are of the order of −10 kJ mol−1 as shown in the table, while the absolute values of each competing steps could be an order of magnitude larger. Hence a subtle difference among AL in each step could bring about a relatively sizable difference in the net result. Thus, it is safer to use the relative difference within a given alcohol. As is our basic

110

Fig. V-20

Table V-1. The excess partial molar quantities of alcohols (AL) at the infinite dilution in AL−H2 O at 25 C. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa. Mono-ol (abbr.)

EAL (ref) kJ mol−1

E HAL (ref) kJ mol−1

E TSAL (ref) kJ mol−1

E VAL (ref) −1 cm3 mol

Methanol (ME)

1.3(a) 1.24(e) 1.21(g)

−87 (c)

−22 (d) −253(y)

Ethanol (ET)

3.3(a) 3.38(e) 3.37(k)

−135(c) −133 ( j)

−35(d) −352(y)

1-Propanol (1P)

6.6(a) 6.60(e) 6.50(n)

−74 (b) −735(f) −704((h) −729 (i) −102(b) −996(f) −1017(h) −1018(l) −1015(i) −104(b) −1012(l) −1012(i)

−170(c) −167( j)

−44(m) −422(y)

2-Propanol(2P)

5.0(a)

−181(c)

−41(o) −515(x)

1-Butanol(1B)

−193(p) −195( j)

−54(m)

tert-Butanol(TBA)

10.0(p) 9.77(e) 9.71(n) 9.75(z) 6.2(q)

−131(b) −1309(i) −1307(l) −93(p) −932(i) −94(l)

−236(q)

1-Pentanol (1Pe)

13.5( j)

−174(q) −1746(l) −78(l) −799(i)

−213( j)

−63(r) −722(s) −61(m)

2-butoxyethanol(BE)

9.2(t)

−170(u) −170(w) −1719(x)

−262(t)

−82(v)

(Ref.) (a) Hu et al. (2003), (b) Tanaka et al. (1996), (c) Koga et al. (2003), (d) Dethelefson (1984), (e) Vrbka et al. (2005), (f) Bertrand et al. (1966), (g) Christian et al. (1981), (h) Dohnal et al. (1994), (i) Hallen et al. (1986), (j) Tanford (1973), (k) Nord et al. (1984), (l) Arnett et al. (1969), (m) Vikingstad (1979), (n) Larkin & Pemberton (1973), (o) Akers & Gabler (1991), (p) Tamura et al. (2000), (q) Koga et al. (1990a), (r) Hvidt et al. (1978), (s) Sakurai (1987), (t) Koga (1991), (u) Siu & Koga (1989), (v) Koga et al. (1993), (w) Kusano et al. (1973), (x) Sakurai (1988), (y) Benson & Kiyohara (1980) (z) Hovorka et al. (2002). E attitude in this book, we use the xAL dependence, i.e. the interaction functions, HAL−AL E and TSAL−AL , for this purpose. E E and TSAL−AL , to A qualitative similarity of the third derivative quantities, HAL−AL those for BE−H2 O is apparent in Figs. V-20 to 23. Peak type anomalies similar to that

111

Table V-2. The liquid–liquid phase boundaries of some alcohol (AL)−H2 O at 25 C. The phase separation occurs between the H2 O-rich and the AL-rich boundaries given as the mole fraction of AL, xAL . Data from Sorensen & Arlt (1979). Alcohol (abbrev.)

H2 O-rich xAL

AL-rich xAL

1-Butanol (1B) 1-Pentanol(1Pe) 1-Hexanol (1Hx) 1-Octanol (1Oc) 1-Decanol (1De)

00192 000375 000104 00000703 00000162

0488 0658 0710 0793 0736

400

HAL-ALE / kJ mol-1

350 300

ME, 25 °C

250

ET, 25 °C 1P, 25 °C

200

2P, 25 °C

150

TBA, 26.9 °C

100 50 0 -50 0.0

0.1

0.2

0.3

0.4

0.5

xAL (AL = ME, ET, 1P, 2P, or TBA) E . The data Fig. V-20. The enthalpic interaction between solute alcohol (AL) in AL−H2 O HAL−AL for methanol (ME), ethanol (ET), 1-propanol (1P) and 2-propanol (2P) are taken from Tanaka et al. (1996), and those for TBA from Koga (1988a). Reproduced with permission from the Canadian Journal of Chemistry, 74, 713, (Tanaka S. H. et al.), Copyright (1996), NRC Research Press, National Research Council of Canada, Ottawa.

E E for BE−H2 O (Figs. 7(a) and 8(a)), are evident for both HAL−AL and TSAL−AL with point E E X and point Y. The difference among AL lies in the value of HAL−AL or TSAL−AL at E point X and its xAL locus. These values are listed in Table V-3. The value of TSAL−AL E , as in the case for BE−H2 O. We conclude that at point X is larger than that of HAL−AL the iceberg formation with a concomitant reduction in the hydrogen bond probability of bulk H2 O away from icebergs is also true in all AL−H2 O systems studied. The AL − AL interaction function increases rapidly to the breaking point, point X, with a E E larger value of HAL−AL or TSAL−AL for a larger AL. In addition, a larger alcohol requires

112

Fig. V-21 to V-23

50

H ME-MEE / kJ mol-1

40

X 30

20

III

II I

10

Y

0

0.0

0.2

0.4

0.6

0.8

xME E , at 25  C Fig. V-21. The enthalpic solute–solute interaction in methanol(ME)−H2 O HME−ME E (Tanaka et al., 1996). Point X corresponds to the inflection point of HME vs. xME and it marks the onset of transition from Mixing Scheme I to II. Point Y is the end point of transition and beyond this point Mixing Scheme II is operative up to about xME = 045. Mixing Scheme III is in the range xME >045. See text. Reproduced with permission from the Canadian Journal of Chemistry, 74, 713, (Tanaka S. H. et al.), Copyright (1996), NRC Research Press, National Research Council of Canada, Ottawa.

400 350

TS AL-ALE / kJ mol-1

300

ME

250

ET 1P

200

2P TBA

150 100 50 0 -50

0.0

0.1

0.2

0.3

0.4

0.5

xAL (AL = ME, ET, 1P, 2P, or TBA) E Fig. V-22. The entropic interaction (times T ) between solute alcohols (AL), TSAL−AL , at 25  C. The data are taken from Koga et al. (2003) for all but TBA. The latter data are calculated using E TSTBA data at 25  C from Koga et al.(1990a). Reproduced with permission from the Canadian Journal of Chemistry, 81, 150, (Koga, Y. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

Eq. (V-3)

113

50

40

TS ME-MEE / kJ mol-1

X 30

20

III

II

I 10

Y 0 0.0

0.2

0.4

0.6

0.8

1.0

xME E Fig. V-23. The entropic interaction between solutes in methanol (ME)−H2 O, TSME−ME , at 25  C (Koga et al., 2003). The boundary between Mixing Schemes II and III is less conspicuous. Reproduced with permission from the Canadian Journal of Chemistry, 81, 150, (Koga, Y. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

a smaller value of xAL to bring the system to point X as well as to point Y. We may rank the hydrophobicity of each alcohol by these sets of data, which give the consistent ranking as, ME<ET<2P<1P
(V-3)

The equivalent interaction in terms of chemical potential, EAL−AL (the second derivative of G) shows a much less conspicuous xAL dependence feature due to the entropy– enthalpy compensation effect prevalent in aqueous solution. (Lumry & Rajender, 1970; Lumry, 2003). We stress that this ranking is the effect of each AL on H2 O in the H2 O-rich environment. The peak anomalies in this H2 O-rich region are very prominent, as shown in Figs. 20 and 22, and point X and point Y can be located with much confidence, ±0005 in xAL locus. Figs. 21 and 23 are for ME−H2 O with the smallest anomaly in the series, which still show the peak clearly. As discussed for BE−H2 O in Section [V-1], estimation of the boundary between Mixing Schemes II and III is problematic, since we have to seek an intersection of two lines crossing with a shallow angle. The same is true for all AL−H2 O, as is evident in Figs. 21 and 23. The uncertainty in the II–III boundary by E E or TSAL−AL may be ±01 in xAL . They are listed in Table 3. locating a break in HAL−AL The II–III boundaries will be reexamined below in Section [V-5] by using anomalies in HWE −W or TSWE −W , and more confident II–III boundary values will be given. E E and HTBA−TBA , which Figs. V-24 and 25 shows the temperature dependence of H1P−1P E E are similar to that of HBE−BE (Fig. 7(a)). Namely, the value of HAL−AL starts with the

114

E E Table V-3. The summary of behaviors of the alcohol–alcohol interactions, HAL−AL and TSAL−AL , for AL−H2 O at 25 C. See Koga et al. (2003a), with revision regarding II-III boundary for TBA. See footnote below for detail.

AL

ME ET 2P 1P TBA BE a

b

c

d

E From HAL−AL E HAL−AL at xi = 0

24 ± 3 35 ± 5 45 ± 7 50 ± 7 90 ± 20 350 ± 30

at point X

32 ± 3 73 ± 5 180 ± 10 210 ± 10 350 ± 20 1100 ± 50

Loci of X ±0005 0.07 0.06c 0.05 0.049 0.045 0.017

Y

0.25 0.19d 0.11 0.09 0.065 0.021

II-IIIa boundary ±01 0.45 0.35 0.16 0.27 ---b 0.41

E From TSAL−AL E TSAL−AL at xi = 0

25 ± 3 40 ± 5 60 ± 8 60 ± 8 150 ± 20 470 ± 35

at point X

34 ± 3 81 ± 5 185 ± 10 225 ± 10 360 ± 20 1150 ± 50

Loci of X ±0005 0.065 0.055 0.05 0.049 0.045 0.017

Y

0.25 0.19 0.11 0.09 0.065 0.021

II-IIIa boundary ±01 0.54 0.38 0.19 0.30 ---b 0.43

Fig. V-24 to V-25

E E The mixing scheme boundary between II and III is difficult to evaluate from the behavior of HAL−AL and TSAL−AL , due to the smallness of their values and E the subtlety of their variations. In Section [V-5] we use HWE −W and TSW −W instead which provide more reliable loci of the II-III boundary. E The locus of the II-III boundary was given in 1990 as xTBA = 062 from a sharp bend in TSW −W (Koga et al., 1990). However, the vapor pressure analysis in the 1990 paper did not circumvent correctly the difficulty associated with the azeotrope at xTBA = 06. Subsequently the correct procedure discussed in Section [III-9] was applied and no break was observed in HWE −W (see Section [V-5]). It could be interesting to note that the mole fraction at point X for ET−H2 O translates to about 18 volume % which is the highest alcohol content in classical Japanese sake and classical ice-wines. This seems to suggest that yeast ceases to function at point X. It could also be interesting to note that the mole fraction of point Y is about 40 volume %, which is the alcohol content of the strongest spirit on the market. In an unpublished observation, we saw that the vesicles of dipalmitoyl phosphatidylcholine(DPPC), a model cell membrane, are dissolved completely and the colloidal dispersion became clear when the mixed solvent of ET−H2 O reached point Y.

115

250

X X

H1P-1PE / kJ mol-1

200

25 °C

150

35 °C

100

50

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. V-24. The temperature dependence of the 1-propanol (1P) enthalpic interaction, H1P−1P , in 1P−H2 O. Unpublished data (Koga).

500

HTBA-TBAE / kJ mol-1

400

X

X 300

temp. °C 26.90 30.00

Error

45.17

X

59.49

200

Y

100

0 0.00

0.02

0.04

0.06

0.08

0.10

xTBA E Fig. V-25. The temperature dependence of the tert-butanol enthalpic interaction, HTBA−TBA , for TBA−H2 O. Data from Koga (1988a). Reproduced with permission from the Canadian Journal of Chemistry, 66, 1187, (Koga, Y.), Copyright (1988), NRC Research Press, National Research Council of Canada, Ottawa.

same value at xAL = 0 independent of temperature. It increases along the same line as E reaches point X at smaller values of xAL increases. For a higher temperature, HAL−AL E HAL−AL and xAL , and makes a nose-dive to point Y. For TBA−H2 O, the loci of point X in Fig. 25 and anomalies in other third derivative quantities are collected and plotted

116

Fig. V-26

100 80 60

II

40

I Temperature (°C)

20 0

20

10

II

I

0

tP -10

0.00

S1 + S2

Ice + S1 0.05

0.10

S1

0.15

xTBA

Fig. V-26. The mixing scheme boundary between Mixing Schemes I and II – the Koga line (Koga et al., 2004a). The phase boundaries are taken from Ott et al. (1979), Mootz & Staeben (1993) E and Takaizumi (2000). The mixing scheme boundary is the loci of anomalies in HTBA−TBA (Koga, E 1988a), TSTBA−TBA (Koga et al.,1990c), both are expressed following the labelling of Table I-1, as T nTBA  nTBA . Others are p p nTBA  and p T nTBA  evaluated by using Sakurai’s density data (1987), and T T nTBA  using the heat capacity data (DeVisser et al. (1977)). Reproduced with permission from the Journal of Physical Chemistry A, 108, 3873, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

in the temperature–mole fraction field together with the phase diagram in Fig. V-26. Similar to Fig. 13, they form a line, the Koga line, that separates Mixing Scheme I and II fields. The boundary seems to point also to 80 to 85  C at the xTBA = 0 axis. For the low temperature side, the Koga line seems to point to the incongruent melting point of the addition compound S1 of the type TBAH2 Om . The value of m remains controversial: m = 6 (Ott et al., 1979), 7 (Mootz & Staeben, 1993) and 7.7 (Takaizumi, 2000) have all been suggested. In any case, the fact that this addition compound and ice form on cooling the TBA−H2 O in Mixing Scheme I also provides support for the “iceberg” formation concept in the H2 O-rich region.

Eq. (V-4) to (V-11)

117

[V-3] FLUCTUATION FUNCTIONS – MORE ABOUT MIXING SCHEME I So far we have learned about Mixing Scheme I (and II and III to some extent) in AL−H2 O systems from the behavior of the second and third derivative quantities. In particular, deeper insights into the mixing schemes are gained by the temperature and the mole E E E E E E  TSAL  VAL  HAL−AL  TSAL−AL , and VAL−AL . In this section fraction dependences of HAL we seek for additional information from other second and third derivative quantities, i.e. the fluctuation functions introduced in Chapter IV. The mean square fluctuation densities, q , where q = V, S, or SV, are rewritten as, recalling eqs. (IV-7) to (IV-9), V

≡ T T 

(V-4)

S

≡ Cp m /Vm 

(V-5)

≡ T p 

(V-6)

SV

They are the volume, the entropy, and the entropy-volume cross fluctuation densities respectively, and show the mean square amplitudes (or intensities) of the respective fluctuations, as discussed at some length in Chapter IV. Their partial molar quantities of the i-th solute, q i , are defined as,  q  q  (V-7) i ≡ 1 − xi  xi which signifies the effect of the i-th component on the respective fluctuation density. We also introduced, recalling eqs. (IV-14) to (IV-16), the normalized fluctuation functions, q , as V

 ≡ RT T /Vm 

(V-8)

S

 ≡ RCp m /Vm2 

(V-9)

SV

 ≡ RTp /Vm 

(V-10)

They qualitatively indicate the amplitude and the wavelength (extensity) of the respective fluctuation, as discussed in Section [IV-2]. Their partial molar quantities of the i-th component are defined as,  q   q  (V-11) i ≡ 1 − xi  xi which signifies the effect of the i-th component on the respective fluctuation functions. We calculated all these fluctuation functions using the available literature data. For methanol (ME)−, ethanol (ET)−, and 1-propanol (1P)−H2 O, the raw data by Benson

118

Fig. V-27(a) to V-27(b)

et al. (Benson & Kiyohara, 1980; Benson et al., 1980; Kiyohara & Benson, 1981) are interpolated to the fixed mole fraction intervals graphically as detailed in Koga (1999). For TBA−H2 O, the literature data (Tamura et al., 1999; Sakurai, 1987; Koga & Tamura, 2000) are treated in a similar manner (Koga & Tamura, 2000). For BE−H2 O, the necessary data (Koga et al., 1995b; 1993; Roux et al., 1978a; Davies et al., 1992) are also treated by graphical interpolation (Koga & Westh, 1996; Koga & Tamura, 2000). The results are shown in Figs. V-27 to V-36. Fig. 27 shows the volume fluctuations, V and V . Both decrease on increasing xAL within Mixing Scheme I. We recall that in evaluating fluctuation functions we fixed all the independent variables T p nAL  nW , where nAL and nW are the amount of AL and H2 O. Hence, each coarse grain, the volume of which is allowed to fluctuate, should contain the same amount of AL and H2 O, and the mole fraction xAL in a coarse grain is the same as the global average. If the “iceberg” formation occurs as a part of fluctuation, it is expected that the volume fluctuation should increase as xAL increases. The present observation that both V (amplitude) and V  (amplitude and wavelength) in fact decrease on increasing xAL may indicate that the “iceberg” parts do not participate in fluctuation, and the “icebergs” are rigid entities.∗ Thus, only the bulk H2 O away from “icebergs” fluctuates with a lesser intensity (amplitude) and extensity (wavelength) due to a reduced hydrogen bond probability with a reduced degree of formation/destruction of ice-like patches. The initial decreases in both V and V  turns around to increase at about the mixing scheme boundary for each AL. The third derivative quantities, V AL and V AL , in Fig. V-28 provide more conspicuous xAL dependences. In particular, the arrow in Fig. 28 (b) shows the xAL locus of point Y for each AL taken from Table 3 (p. 114). Thus, the initial negative effect of AL on V  decreases on increasing xAL and then reaches about zero at point Y, thereupon Mixing Scheme II starts; namely the effect of incoming AL is to reduce the amplitude as well as the wavelength of volume fluctuation, by reducing the hydrogen bond probability of bulk H2 O away from icebergs. When such an effect is no longer operative, Mixing Scheme II sets in. Fig. V-29 shows the entropy fluctuations. Fig 29(a) indicates that S (amplitude) increases on increasing xAL initially, with a lesser increase for smaller AL. This general trend is reversed, however, from TBA to BE. While we note the total number of carbon in BE is 6 as opposed to 4 for TBA, there is an extra −O− in ether type for BE. The issue of relative hydrophilicity between the hydroxyl group, −OH, and the ether −O−, will be dealt with in Chapter VII. On the other hand, S  (amplitude and wavelength) decreases as xAL increases as shown in Fig. 29(b). The mean square entropy fluctuation, <S2 >, can be translated to the fluctuation of hydrogen bond strength, since T 2 <S2 > = <H2 >, for constant p. In terms of the bent hydrogen bond model, S (amplitude) reflects how widely the hydrogen bonds are bending back and

∗ In this connection, we recall the recent femto-second pump-probe spectroscopic study on the dynamics of H2 O in the hydration shell of anions (Omta et al., 2003; Bakker et al., 2004; Kropman & Bakker, 2001). They showed that the dynamics of H2 O in the hydration shell is much slower than that in the bulk H2 O away from the hydration shell.

119

ME ET 1P TBA BE

(a)

0.15

V

δ / K MPa-1

0.20

0.10

0.0

0.1

0.2

0.3

xAL (AL = ME, ET, 1P, TBA, or BE)

Fig. V-27(a). The volume fluctuation density in alcohol (AL)−H2 O V , at 25  C. Alcohols (AL) are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA) and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

0.08

0.07

ME ET 1P TBA BE

(b)

V

Δ

0.06

0.05

0.04

0.03

0.02 0.0

0.1

0.2

0.3

xAL (AL = ME, ET, 1P, TBA, or BE)

Fig. V-27(b). The normalized volume fluctuation, V , for alcohol (AL)−H2 O at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA), and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

120

Fig. V-28(a) to V-29(b)

1.0

0.5

0.0

V

δAL / K MPa-1

(a)

ME ET 1P TBA BE

-0.5

-1.0 0.0

0.1

0.2

0.3

xAL (AL = ME, ET, 1P, TBA, or BE)

Fig. V-28(a). The partial molar volume fluctuation density of alcohol (AL) in AL−H2 O V AL , at 25  C. Alcohols (AL) are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA), and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

0.2

(b)

TBA BE

0.0

ME

1P

V

Δ AL

ET -0.2

ME ET 1P TBA BE

-0.4

-0.6

-0.8 0.0

0.1

0.2

0.3

xAL (AL=ME, ET, 1P, TBA, or BE)

Fig. V-28(b). The partial molar volume normalized fluctuation of alcohol (AL) in AL−H2 O V AL , at 25  C. The arrows indicate the loci of point Y, the end of the transition from Mixing Scheme I to II, for each alcohol. Alcohols are methanol (ME), ethanol (ET), and 1propanol (1P) (Koga, 1999), and tert-butanol (TBA), and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

121

4

ME ET 1P TBA TBA

S

δ / J K-1 cm-3

(a)

3

0.0

0.1

0.2

0.3

xAL (AL=ME, ET, 1P, TBA, or BE)

Fig. V-29(a). The entropy fluctuation density of alcohol (AL)−H2 O S , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA) and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

2.0

1.5

ME ET 1P TBA BE

1.0

S

Δ / J2 K-2 cm-6

(b)

0.5

0.0

0.1

0.2

0.3

xAL (AL=ME, ET, 1P, TBA, or BE)

Fig. V-29(b). The normalized entropy fluctuation in alcohol (AL)−H2 O S , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA), and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

122

Fig. V-30(a) to V-30(b)

forth. The observed facts in Fig. 29(a) and (b) may suggest that the H2 O molecules in the vicinity of “icebergs” are forced to bend more. At the same time, however, the wavelength of fluctuation diminishes more rapidly in the bulk H2 O away from “icebergs” to compensate the initial increase in the amplitude of entropy fluctuation. Fig. V-30 E and shows sharp drops in S AL and S AL corresponding to points X and Y in HAL−AL E TSAL−AL . For smaller alcohols the degree of these anomalies are less conspicuous. The unique nature of H2 O manifests itself in the entropy volume cross fluctuation in that there is a negative contribution to the product SV unlike normal liquids; namely, the volume increase, V >0, during the course of fluctuation contains an increase due to the putative formation of ice-like patches, which is associated with a negative S, and vice versa. Thus, in Fig. V-31, the initial increases in SV and SV  could be understood as a reduction of this negative contribution due to formation/destruction of ice-like patches as xAL increases. As a guide for the eye, the dotted line is drawn for “a normal behavior” of 1P−H2 O. The difference from this dotted line can be regarded at least semi-quantitatively as the negative contribution. Thus the trace of SV and SV  reflects the manner in which the added AL reduces the characteristics of the H2 Olikeness of the bulk H2 O away from icebergs. We note that the xAL dependence of SV E or SV  is sigmoidal and resembles that of HAL shown in Figs. 3 and 16 in the H2 O−rich SV E region. Indeed, the next derivatives, AL and SV AL , and HAL−AL show similar peak anomalies. The similarity is striking as shown in Figs. V-32 to 36. In particular, the xAL dependence of SV AL matches almost exactly for 1P−, TBA−, and BE−H2 O on scaling the ordinate, as shown in Figs. 34(b), 35(b) and 36(b). Considering the uncertainties in both third derivative quantities, it may be safe to conclude that xAL -dependence is the E and SV AL . This means that the AL − AL enthalpic interaction same between HAL−AL and the AL’s effect on the S−V cross fluctuation share the same fundamental cause. We suggest therefore that the AL−AL interaction is operative via the bulk H2 O away from the “icebergs” in solution, where putative formation/destruction of ice-like patches causes the negative contribution to SV AL . Hence, the AL−AL interaction is bulkH2 O mediated and of long range via the hydrogen bond network of H2 O. There are E mismatches between SV AL and HAL−AL for the portion from Point X to Y systematically, as shown in Figs. 34(a), 35(a) and 36(a). This observation together with good matches E may hint at the importance of the extensity (wavelength) of between SV AL and HAL−AL the fluctuation nature of H2 O and of aqueous solutions. The relatively poor matches for ME−, and ET−H2 O shown in Figs. 32 and 33 may be attributed to some specific xAL dependence of the density data for ME−, and ET−H2 O, when p is calculated from the density data at 15, 25, and 35  C as, p =

V35 − V15  V25 35 − 15

(V-12)

where V35 etc. are the molar volume at 35  C calculated from the measured density. More direct measurements, such as dilatometry (Davies et al., 1992), are awaited.

Eq. (V-12)

123

ME ET 1P TBA BE

S

δAL / J K-1 cm-3

5

0

-5

(a) -10 0.0

0.1

0.2

0.3

xAL (AL=ME, ET, 1P, TBA, or BE)

Fig. V-30(a). The partial molar entropy fluctuation density of alcohol (AL) in AL−H2 O S AL , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tertbutanol (TBA), and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

(b)

-5

ME ET 1P TBA BE

S

ΔAL / J2 K-2 cm-6

0

-10

0.0

0.1

0.2

0.3

xAL (AL=ME, ET, 1P, TBA, or BE)

Fig. V-30(b). The partial molar normalized entropy fluctuation of alcohol (AL) in AL−H2 O S AL , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA) and 2-butoxyethanol (BE) (Koga & Tamura, 2000). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

124

Fig. V-31(a) to V-32(b)

ME ET 1P TBA BE

0.35 0.30

SV

δ

0.25 0.20 0.15

(a) 0.10 0.05 0.0

0.1

0.2

0.3

xAL (AL = ME, ET, 1P, TBA, or BE)

Fig. V-31(a). The entropy volume cross fluctuation density in alcohol (AL)−H2 O SV , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA) and 2-butoxyethanol (BE) (Koga & Tamura, 2000). The dotted line shows a presumed “normal behaviour” of SV in 1P−H2 O. The difference from this line is regarded as the negative contribution to SV , which diminishes as x1P increases from x1P = 0. When this negative contribution becomes zero at point Y, Mixing Scheme II is operative. Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa. 0.10

(b)

SV

Δ / J K-1 cm-3

0.09 0.08 0.07 0.06

ME ET 1P TBA BE

0.05 0.04 0.03 0.0

0.1

0.2

0.3

0.4

xAL (AL = ME, ET, 1P, TBA, or BE)

Fig. V-31(b). The normalized entropy volume cross fluctuation of alcohol (AL)−H2 O SV , at 25  C. Alcohols are methanol (ME), ethanol (ET), and 1-propanol (1P) (Koga, 1999), and tert-butanol (TBA) and 2-butoxyethanol (BE) (Koga & Tamura, 2000). The dotted line shows a presumed “normal behaviour” of SV  in 1P−H2 O. The difference from this line is regarded as the negative contribution to SV , which diminishes as x1P increases from x1P = 0. When this negative contribution becomes zero at point Y, Mixing Scheme II is operative. Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

125

0.8

(a)

δME

SV

0.4

H ME-MEE

SV

δ ME

0.6

0.2

0.0

0.0

0.2

0.4

0.6

0.8

xMe

Fig. V-32(a). The partial molar entropy volume cross fluctuation density of methanol (ME), E SV ME , in ME−H2 O at 25  C. Also plotted are the ME−ME enthalpic interactions, HME−ME , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

0.20

XMe

(b)

0.10

SV

Δ ME

H ME-MEE

0.05

SV

ΔME / J K-1 cm-3

0.15

YMe 0.00

I

II

-0.05 0.0

0.2

0.4

0.6

0.8

xMe

Fig. V-32(b). The partial molar normalized entropy volume cross fluctuation of methanol (ME), SV E ME , in ME−H2 O at 25  C. Also plotted are the ME−ME enthalpic interaction, HME−ME , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

126

Fig. V-33(a) to V-34(b)

1.4 1.2

(a)

1.0 0.8

δET

SV

δ ET

SV

0.6

H ET-ETE

0.4 0.2 0.0 0.0

0.1

0.2

0.3

0.4

0.5

xET

Fig. V-33(a). The partial molar entropy volume cross fluctuation density of ethanol (ET) in E ET−H2 O SV ET , at 25  C. Also plotted are the ET−ET enthalpic interaction, HTE−ET , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

0.5

XEt

SV

ΔET / J K-1 cm-3

0.4

(b)

0.3

Δ ET

SV

0.2

H ET-ETE 0.1

YEt

0.0

I

II

-0.1 0.0

0.1

0.2

0.3

0.4

0.5

xET

Fig. V-33(b). The partial molar normalized entropy volume cross fluctuation of ethanol (ET) E in ET−H2 O SV ET , at 25  C. Also plotted are the ET−ET enthalpic interaction, HTE−ET , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

127

2.5

(a)

2.0

1.5 SV

δ 1P

SV

δ1P

H 1P-1PE

1.0

0.5

0.0 0.0

0.1

0.2

0.3

x1P

Fig. V-34(a). The partial molar entropy volume cross fluctuation density of 1-propanol (1P) in E 1P−H2 O SV 1P , at 25  C. Also plotted are the 1P−1P enthalpic interaction H1P−1P , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

1.0

I II

0.8

(b)

SV

Δ1P / J K-1 cm-3

X1P 0.6 SV

0.4

Δ1P

H 1P-1PE 0.2

0.0

-0.2 0.00

Y1P

0.05

0.10

0.15

0.20

0.25

0.30

x1P

Fig. V-34(b). The partial molar normalized entropy volume cross fluctuation of 1-propanol (1P) in E 1P−H2 O SV 1P , at 25  C. Also plotted are the 1P−1P enthalpic interactions, H1P−1P , with arbitrary units. (Koga, 1999). Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

128

Fig. V-35(a) to V-36(b)

4

(a) 3

SV

δ TBA

2

SV

δ TBA

H TBA-TBAE

1

0

0.00

0.05

0.10

0.15

x TBA

Fig. V-35(a). The partial molar entropy volume cross fluctuation density of tert-butanol (TBA) in E TBA−H2 O SV TBA , at 25  C. Also plotted are the TBA−TBA enthalpic interactions, HTBA−TBA , with arbitrary units. (Koga and Tamura, 2000). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 27, 195, Copyright (2000), the Japan Society of Calorimetry and Thermal Analysis.

1.5

Δ TBA

SV

XTBA

0.5

II

SV

Δ TBA / J K-1 cm-3

(b) 1.0

H TBA-TBAE

0.0

0.00

I

YTBA

0.05

0.10

0.15

x TBA

Fig. V-35(b). The partial molar normalized entropy volume cross fluctuation of tert-butanol (TBA) E in TBA−H2 O SV TBA , at 25  C. Also plotted are the TBA−TBA enthalpic interactions, HTBA−TBA , with arbitrary units. (Koga & Tamura, 2000). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 27, 195, Copyright (2000), the Japan Society of Calorimetry and Thermal Analysis.

129

5

δBE

SV

(a)

SV

δBE / J K-1 cm-3

4

H BE-BEE

3 2 1 0

0.00

0.02

0.04

0.06

xBE

Fig. V-36(a). The partial molar entropy volume cross fluctuation density of 2-butoxyethanol (BE) E in BE−H2 O SV BE , at 25  C. Also plotted are the BE−BE interactions, HBE−BE , with arbitrary units. (Koga & Tamura, 2000). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 27, 195, Copyright (2000), the Japan Society of Calorimetry and Thermal Analysis.

2.0

(b) XBE

ΔBE

SV

1.0

H BE-BEE 0.5

SV

Δ BE / J K-1 cm-3

1.5

0.0

I

YBE II

0.00

0.02

0.04

0.06

xBE

Fig. V-36(b). The partial molar normalized entropy volume cross fluctuation of 2-butoxyethanol E (BE) in BE−H2 O SV BE , at 25  C. Also plotted are the BE−BE enthalpic interactions, HBE−BE , with arbitrary units (Koga & Tamura, 2000). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 27, 195, Copyright (2000), the Japan Society of Calorimetry and Thermal Analysis.

130

Fig. V-37(a) to V-37(b)

We now list additional information gained in this section regarding Mixing Scheme I. (1) H2 O molecules within “icebergs” do not participate in fluctuation. Namely, “icebergs” are rigid entities and H2 O molecules in icebergs show much slower dynamics than those in bulk H2 O away from icebergs. (2) The intensity (amplitude) of entropy fluctuation increases but the extensity (wavelength) decreases on increasing xAL . (3) The AL − AL interaction is H2 O mediated and is therefore of a long range within Mixing Scheme I.

[V-4] CONCENTRATION FLUCTUATIONS – MIXING SCHEME II As discussed in Section [V-1] for 2-butoxyethanol (BE)−H2 O, Mixing Scheme II is operative in the intermediate composition region, whereby two kinds of clusters mix; one rich in H2 O and the other in BE. The fact that the demixing phase separation exists with E E the LCST at about xBE = 006 and at 50  C, and that HBE and TSBE both have positive slopes in this composition region are consistent with Mixing Scheme II. The more direct evidence for the existence of clusters came from the findings by light (LS) (Ito et al., 1983) and small angle x-ray scattering (Hayashi & Udagawa, 1992b) experiments. The forward scattering intensities increase to the maximum in this composition region, and the maximum values increase on increasing temperature to that of the LCST. From the measured scattering intensity, the zero angle partial structure factor, Sxx(0), can be extracted (Bhatia & Thornton, 1970), and is shown to be equal to the concentration fluctuation of AL−H2 O N <xAL 2 >; namely, as shown in eqs. (III-22) and (IV-23),   E  AL Sxx0 = N <xAL 2 > = RT1 − xAL  + xAL

RT xAL





(V-13)

The right-hand side of eq. (13) provides the route to calculate N <xAL 2 > from EAL data by vapor pressure measurements (VP). The left-hand side is obtained from LS or SAXS. See Kato & Fujiyama (1976) or Koga (1984) for extracting Sxx(0) from LS or SAXS, respectively. Fig. V-37 shows the results calculated from our vapor pressure measurements and those in the literature from scattering experiments. Fig. 37(f) is for BE−H2 O, discussed already in Section [V-1]. The concentration fluctuation by LS (Ito et al., 1983) reaches the maximum at about xBE = 006, the LCST composition, and it doubles by raising temperature from 21 to 32  C. On increasing further from 32 to 42  C, it almost quadruples, though not shown in Fig. 37(f). The values of N <xAL 2 > for AL = BE by VP are rather scattered with large error bars, as shown in Fig. 37(f). This is inevitable for large values of N <xAL 2 > calculated from the VP route. For N <xAL 2 > to take a large value, the denominator of eq. (13) is small. While the second term of the denominator is always positive and could be accurate, the first term is negative with

Eq. (V-13)

131

1.0

Hu et al., 2003, 25 °C - VP Morita, 2004, 100 °C - SAXS Morita, 2004, 150 °C - SAXS

N<(ΔxME)2>

0.8

(a) ME

0.6

0.4

0.2

Ideal sol.

0.0 0.0

0.2

0.4

0.6

0.8

1.0

xME

Fig. V-37(a). The concentration fluctuation in methanol (ME)−H2 O N <xME 2 >, at 25  C by vapor pressure (VP) method. (Hu et al., 2003). The error of the data obtained from vapor pressure measurements is at most ±002. The data for 100 and 150  C (at 7 MPa) are due to Morita and Nishikawa (2004) by SAXS method. I thank the original authors for providing the values together with the error bar at xME = 05. Reproduced with permission from the Canadian Journal of Chemistry, 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

1.4

Hu et al., 2003, VP Nishikawa & Iijima, 1993, SAXS

1.2

N<(ΔxME)2>

1.0

(b) ET

0.8 0.6 0.4 0.2 0.0

0.0

0.2

0.4

0.6

0.8

1.0

xET

Fig. V-37(b). The concentration fluctuation in ethanol (ET)−H2 O N <xET 2 >, at 25  C. The VP data are by Hu et al. (2003), and the SAXS data by Nishikawa and Iijima (1993). Reproduced with permission from the Canadian Journal of Chemistry, 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

132

Fig. V-37(c) to V-37(f)

Hu et al., 2003, VP Hayashi et al., 1990, SAXS

6

N<(Δx1P)2>

5 4

(c)

3 2 1 0 0.0

0.2

0.4

0.6

0.8

1.0

x1P

Fig. V-37(c). The concentration fluctuation in 1-propanol (1P)−H2 O N <x1P 2 >, at 25  C. The VP data are by Hu et al. (2003), and the SAXS data by Hayashi et al. (1990) and Hayashi & Udagawa (1992a). Reproduced with permission from the Canadian Journal of Chemistry, 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

3.0

Hu et al., 2003, VP 2.5

N<(Δx2P)2>

2.0

(d) 2P

1.5

1.0

0.5

0.0 0.0

0.2

0.4

0.6

0.8

1.0

x2P

Fig. V-37(d). The concentration fluctuation in 2-propanol (2P)−H2 O N <x2P 2 >, at 25  C. The VP data are by Hu et al. (2003). Reproduced with permission from the Canadian Journal of Chemistry, 81, 141, (Hu, J. et al.), Copyright (2003), NRC Research Press, National Research Council of Canada, Ottawa.

133

3

2

N<(ΔxTBA)2>

(e) TBA 1

0

Koga et al., 1990, VP, 25 °C Nishikawa et al., 1989, SAXS, 28 °C Same as above, 40 °C Same as above, 55 °C

-1

0.0

0.2

0.4

0.6

0.8

1.0

xTBA

Fig. V-37(e). The concentration fluctuation in tert-butanol (TBA)−H2 O N <xTBA 2 >, at various temperatures by vapor pressure measurements (VP) and by small angle x-ray scattering measurements (SAXS). The VP data are by Koga et al. (1990a), and the SAXS data by Nishikawa et al. (1989). 8

6

(f) BE

N<(ΔxBE)2>

4

2

0

Ito et al., 1983, Light Scatt., 21 °C Same as above, 32 °C Koga, 1991, VP, 25 °C

-2

0.0

0.2

0.4

0.6

0.8

1.0

xBE

Fig. V-37(f). The concentration fluctuation in 2-butoxyethanol (BE)−H2 O N <xBE 2 >, at various temperatures. Note large error bars for N <xBE 2 > by vapor pressure measurements (VP). This is inevitable when the values become large. For such a case, the light or the small angle x-ray scattering (SAXS) measurements are much more advantageous (see text). Also note that N <xBE 2 > increases as temperature increases reflecting the fact that the de-mixing critical point, LCST, is close. In fact the LCST is known to be at xBE = 006 and at 50  C. The VP data are by Koga (1991), and the LS data by Ito et al. (1983).

134

some uncertainty. A conservative estimate for EAL is ±001 kJ mol−1 . Through the differentiation process the uncertainty of the first term will increase sizably. The sum of these two terms ends up with a small (always positive) value. The absolute error of the first term in the denominator is carried to the sum, resulting in a large relative error in the net result. The error bars in Fig. 37(f) and in all others, 37(a) to 37(e), are calculated based on the error in EAL being ±0005 kJ mol−1 ; a slightly aggressive estimate. The uncertainty in Sxx(0) from LS or SAXS is basically due to the background correction and its absolute value remains constant. As a consequence, the relative error becomes smaller for higher values of Sxx(0). Not all literature data from LS/SAXS routes report the uncertainty, however, In any case, LS/SAXS and VP methods are thus complementary, in that the former result is more reliable when the value of Sxx(0) is large, while the VP method is more accurate when Sxx(0) is small. For ME−H2 O (Fig. 37(a)), the values of N <xAL 2 > from the VP route is much more reliable, being rather small. Its error bar is at most ±002, about the size of the symbol. In comparison, Sxx(0) from SAXS has inevitably large uncertainties. Recently, Morita and Nishikawa (2004) made a careful SAXS study on ME−H2 O at 100 and 150  C under pressure, using a strong x-ray source from a photon factory. These results with error bars are shown in Fig. 37(a). In comparison between N <xAL 2 > by VP at 25  C and those by SAXS at higher temperatures, there seems to be a trend of increase in N <xAL 2 > on increasing temperature, which is consistent with a possibility of phase separation with an LCST at much higher temperature than 150  C. We recall that E E and TSME−ME are both positive in the intermediate composition, which also HME−ME supports a possibility of de-mixing with an LCST, as discussed in Section [V–2]. For TBA−H2 O, Fig. 37(e) indicates that the maximum of N <xAL 2 > from SAXS shows a marginal increase on increasing temperature from 28 to 40  C, while it increases by about 30% from 40 to 50  C. At 20  C the values of the concentration fluctuation (Nishikawa et al., 1987) are practically identical with those at 28  C shown in Fig. 37(e). This trend of accelerating increase of the maximum of N <xAL 2 > is no doubt due to the fact that the slope of chemical potential which may be decreasing linearly (say) is inversely proportional to N <xAL 2 >, as shown in eq. (13). Accordingly, an LCST is expected moderately above the boiling point, as suggested by Euliss and Sorensen (1984). For ET−, 1P−, and 2P−H2 O, while there are no temperature dependence data available at present, there are maxima in the mid-composition range. This and the analogy from the ME−, TBA−, and BE−H2 O cases discussed above lead us to suggest that Mixing Scheme II is operative in the intermediate composition for all AL that mix with H2 O, i.e. the solution is a mixture of two kinds of clusters: One rich in H2 O and the other in AL.

[V-5] MIXING SCHEME III – SECOND AND THIRD DERIVATIVE QUANTITIES IN THE ALCOHOL-RICH REGION In the AL-rich region, we suggested that Mixing Scheme III is operative whereby AL molecules cluster together in a similar manner as in the pure state. The aqueous solutions

135

in this region are therefore mainly alcohol, and the individual nature of the mixture should depend heavily on each alcohol. However, the partial molar quantities of AL in this region are close to zero (which is instrumental in suggesting AL clustering), and show little individuality. The partial molar quantities of H2 O, on the other hand, and their corresponding third derivatives would identify variations among AL via information about the effect of H2 O on each AL. Thus, Liltorp et al. (2005) determined the excess partial molar enthalpy of H2 O HWE accurately in the AL-rich region, from which TSWE were calculated using the data of EW (Hu et al., 2003). The density data were also determined from which the partial molar E volumes, VAL and VWE were calculated. Alcohols studied are 1-butanol (1B), 1-pentanol (1Pe), 1-hexanol (1Hx), 1-octanol (1Oc) and 1-decanol (1De), in addition to ME, ET, 1P, 2P, TBA, and BE. The results are plotted in Figs. V-38 to 41. The values of HWE  TSWE and VWE at the infinite dilution of H2 O xW = 0, or xAL = 1, are listed in E E Table V-4. In comparison with equivalent HAL and TSAL in the infinite dilution of AL, xAL = 0, in Table 1 (p. 110), the thermodynamic effects of H2 O in AL are modest. We note in Table 4 (p. 141) a trend in VWE at xW = 0 that the value increases from about −1 −4 cm3 mol for ME and ET, and levels off at about zero for large AL (1Pe, 1Hx, 1Oc, and 1De). This suggests that H2 O molecules exist in the latter large alcohols as H2 O clusters of about the same volume as in pure H2 O, while for smaller alcohols H2 O takes the value closer to a van der Waals, non-hydrogen bonded H2 O. The molar volume of the −1 −1 latter H2 O is estimated to be about 12 cm3 mol  −6 cm3 mol in the excess volume,  assuming H2 O to be a sphere of radius 15 A by using the van der Waals radii of H and O (Suzuki, 1980). Similar semi-systematic trends of HWE and TSWE are evident as seen in Table 4 (p. 141). A modest enthalpy gain of −31 kJ mol−1 and an equally modest entropy (times T ) loss of −43 kJ mol−1 decreases to an enthalpy loss 29 kJ mol−1 and an entropy loss −16 kJ mol−1 , as the size of the alkyl group increases from ME to 1De. For smaller alcohols (ME and ET), the values of HWE and TSWE may be consistent with the following picture. H2 O molecules break hydrogen bonds with an average of 60% bond probability (Stanley & Teixeira, 1980) in pure H2 O and form a hydrogen bond to −OH of ME clusters as a single molecule, with a net increase in hydrogen bonding from 60% to 100% and with loss of entropy due to localization at −OH site. As the size of hydrophobic moiety increases, such a single H2 O interaction with AL cluster might gradually change to H2 O cluster−AL cluster interaction for 1De. However, the values of energetics of a few kJ mol−1 and their variations from AL to AL are rather too subtle to discuss the molecular processes by. Furthermore, TBA and BE seem to be out of “place” in the above line of argument. Indeed, these energetic effects are the results of: (a) H2 O molecules breaking away from mutual interaction in its pure state, (b) producing a void of the size of H2 O in the AL environment, and (c) letting the system settle to the final state. The competing thermodynamic effect of each of the three steps must be an order of magnitude larger than the net result of a few kJ mol−1 . For example, the enthalpy of vaporization is about 50 kJ mol−1 for H2 O which may be comparable to the value for step (a). Hence, a subtle difference due to specificity of each AL in each step may result in a sizable variation in the net result. Thus comparison among alcohols at the HWE and

136

Fig. V-38(a) to V-38(c)

HwE / kJ mol-1

4

ME ET 1P

2

(a)

1B

0

-2

2P TBA BE

-4 0.2

0.4

0.6

0.8

1.0

xAL

2

-0.55 -0.60 -0.65

HwE / kJ mol-1

1

HwE / kJ mol-1

Fig. V-38(a). The excess partial molar enthalpies of H2 O HWE , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). All data are from Liltorp et al. (2005), except for TBA, which are taken from Koga (1988b). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

0.4 0.6 0.8 1.0

xAL 0

-1

(b) -2

0.4

0.6

0.8

1.0

xAL

Fig. V-38(b). The expanded (in the ordinate) version of Fig. 38 (a). For BE−H2 O, the ordinate is further expanded in the inset. The excess partial molar enthalpies of H2 O HWE , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). All data are from Liltorp et al. (2005), except for TBA, which are taken from Koga (1988b). Symbols are identical to Fig. 38 (a). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

Eq. (V-14) to (V-15)

137

4

1B 1Hx 1Oc 1De

HwE / kJ mol-1

3

(c)

2

1

0 0.80

0.85

0.90

0.95

1.00

xAL

Fig. V-38(c). The excess partial molar enthalpies of H2 O HWE , in aqueous 1-butanol (1B), 1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1-De) at 25  C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

TSWE (second derivative) level must be exercised with caution. In accordance with the basic attitude we take throughout this book, the relative changes, i.e. the xAL (or xW ) dependences of HWE and TSWE , are more useful and powerful in gaining a deeper insight into the mixing schemes; namely we use third derivative quantities, HWE −W and TSWE −W , for the remainder of this section. They are written as,  E  E HW HW E HW −W ≡ N = 1 − xW   (V-14) nW xW  TSWE −W

≡N

T SWE nW





T SWE = 1 − xW  xW

 

(V-15)

First, however, we examine the xW dependences of HWE and TSWE in Figs. 38 and 39. For TBA (Fig. 38(a)), 1Oc and 1De (Fig. 38(c)), HWE decreases monotonously in a concave manner as xW increases. For 2P (Figs. 38(a) and (b)), while HWE also decreases, the decrease is in a convex manner at first and then turns to a concave decrease with an inflection point at about xAL = 09. For all remaining AL, HWE increases in a convex mode at first and changes to a concave decrease with an inflection point. The inflection point will become an anomaly in HWE −W as a negative peak, or a valley anomaly (Figs. V-42(b) E in the H2 O-rich region. The latter marks and (c)), equivalent to the anomalies in HAL−AL the onset of transition from Mixing Scheme I to II. In analogy then, the present anomaly in HWE −W in the AL-rich composition could also serve as an indicator of the transition

138

Fig. V-39(a) to V-40(b)

1

ME ET 1P 1B

TSwE / kJ mol-1

0 -1 -2 -3 -4

2P TBA BE

(a)

-5 0.2

0.4

0.6

0.8

1.0

xAL E Fig. V-39(a). The excess partial molar entropies of H2 O times T TSW , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). All data are from Liltorp et al. (2005). The data for TBA−H2 O were calculated using HWE data from Koga (1988b) and the vapor pressure data.(Koga et al., 1990a) which were re-analyzed by the method discussed in Section [III-7]. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

0

1B 1Hx 1Oc 1De

TSwE / kJ mol-1

-1

-2

-3

(b) -4 0.80

0.85

0.90

0.95

1.00

xAL E Fig. V-39(b). The excess partial molar entropies of H2 O times T TSW , in aqueous 1-butanol (1B),  1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1De), at 25 C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

139

VALE / cm3 mol-1

0.0

(a)

ME ET 1P 2P TBA BE

-0.5

-1.0 0.6

0.8

1.0

xAL E Fig. V-40(a). The excess partial molar volumes of alcohol (AL),. VAL , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). The data are calculated using the density data of Sakurai and Nakagawa (1982a and 1982b) for ME-, Benson & Kiyohara (1980) for ET- and 1P-, Sakurai (1988) for 2P-, Sakurai (1987) for TBA-, and Liltorp et al. (2005) for BE−H2 O. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

0.02

VALE / cm3 mol-1

(b)

0.00

-0.02

1B 1Pe 1Hx 1Oc 1De

-0.04 0.85

0.90

0.95

1.00

xAL E Fig. V-40(b). The excess partial molar volumes of alcohol (AL), VAL , in aqueous 1-butanol (1B),  1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1De), at 25 C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

140

Fig. V-41(a) to V-41(b)

VwE / cm3 mol-1

0

(a)

-1

-2

ME ET 1P 2P TBA BE

-3

-4

0.5

0.6

0.7

0.8

0.9

1.0

xAL

Fig. V-41(a). The excess partial molar volume of H2 O,. VWE , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). The data are calculated using the density data of Sakurai and Nakagawa (1982a and 1982b) for ME-, Benson & Kiyohara (1980) for ET- and 1P-, Sakurai (1988) for 2P-, Sakurai (1987) for TBA-, and Liltorp et al. (2005) for BE−H2 O. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa. 1

(b)

VwE / cm3 mol-1

0

-1

1B 1Pe 1Hx 1Oc 1De

-2

-3 0.85

0.90

0.95

1.00

xAL

Fig. V-41(b). The excess partial molar volume of H2 O VWE , in aqueous 1-butanol (1B), 1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1De), at 25  C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

141

Table V-4. The excess partial molar quantities of H2 O at the infinite dilution in alcohol (AL)−H2 O at 25 C xW = 0. Data are from Liltorp et al. (2005) unless specified in (ref). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa. AL

EW (ref) kJ mol−1

HWE (ref) kJ mol−1

E TSW (ref) kJ mol−1

VWE (ref) −1 cm3 mol

ME

1.1 (a)

−43

ET

2.2 (a)

1P

3.0(a)

2P

2.8(a)

−32 −299(i) −14 −201(i) 0.2 0.20(i) 1.6

1B

4.0(c)

−22

TBA

3.0(d)

1.8 1.68(i) 4.0(d)

BE 1Hx

0.4(e) 4.0(a)

−06 3.0 3.05(i)

−10 −10

1Oc

4.6(h)

3.5 3.31(i)

−11

1De

4.5(h)

2.9 3.24(i)

−16

−36 −36(b) −41 −41(b) −30 −26(b) −36 −26(b) −16 −19(b) −11 −15(b) −18 −06 0.0(f) −07(g) 0.1 0.3(f) −02(g) 0.1 0.4(f)

−36 −32 −12

1.0(c, d)

(Ref.) (a) Hu et al. (2003), (b) Sakurai et al. (1994), (c) Tamura et al. (2000), (d) Koga (1988a), (e) Koga (1991), (f) Segatin & Kloufutar (2004a), (g) Sakurai & Nakagawa (1984), (h) Segatin & Kloufutar (2004b), (i) Nilsson (1986).

between Mixing Schemes II and III. The xW dependence of TSWE (Fig. 39) is similar to that of HWE , resulting in a similar behavior in TSWE −W (Fig. V-43). Table V-5 (p. 146) summarizes the values of HWE −W and TSWE −W at xW = 0. Also listed in the table are the loci of xAL of the valley anomalies and their minimum values. Unlike the peak anomalies E E in HAL−AL and TSAL−AL , the end point of the valley anomaly (an equivalent of point Y) is not clearly defined as is evident in Figs. 42 and 43. Thus, whether there is also a transition region between Mixing Schemes II and III is not clear at this point. However, the loci of the II–III transition determined from the bottom of the valley anomaly in HWE −W and TSWE −W are more reliable than those in Table 3 (p. 114) from the shallow E E breaks in HAL−AL and TSAL−AL .

142

Fig. V-42(a) to V-42(c)

20

(a)

Hw-wE / kJ mol-1

10

0

ME ET 1P 1B

-10

-20

2P TBA BE

-30

-40 0.2

0.4

0.6

0.8

1.0

xAL

Fig. V-42(a). The H2 O−H2 O interaction in terms of enthalpy, HWE −W , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). All data are from Liltorp et al. (2005) except for TBA. The data for TBA−H2 O were calculated using HWE data from Koga (1988b). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

4

(b)

0 -2

xBE 0.4

0.6

0.8

-4

1.0 0.4 0.2 0.0

-6

-0.2

Hw-wE / kJ mol-1

Hw-wE / kJ mol-1

2

-8 0.2

0.4

0.6

0.8

1.0

xAL

Fig. V-42(b). The expanded (in the ordinate) version of Fig. 42 (a). For BE−H2 O, the ordinate is further expanded in the inset. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

143

(c)

Hw-wE / kJ mol-1

0

-10

1B 1Hx 1Oc 1De -20 0.85

0.90

0.95

1.00

xAL

Fig. V-42(c). The H2 O−H2 O interaction in terms of enthalpy, HWE −W , in aqueous 1-butanol (1B), 1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1De), at 25  C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

We now examine a general trend in molecular interactions, i.e. details of Mixing Scheme III for each AL, using the HWE −W and TSWE −W , the third derivatives. For the first group of small alcohols (ME, ET, 1P, and 1B) in Table 5 (p. 146), both HWE −W and TSWE −W are positive at xW = 0, or xAL = 1. At the infinite dilution, therefore, the W−W interaction is more unfavorable in terms of enthalpy, but favorable entropy-wise, than the W−AL interaction. This suggests that for these AL’s, H2 O molecules are dispersed as a single molecule each interacting with the AL-clusters and W−W pairs tend to repel each other enthalpy-wise with more favorable entropic situation. As xW increases, H2 O crowding gradually lessens the effects, and HWE −W and TSWE −W decreases to the valley anomaly, thereupon Mixing Scheme II sets in. We note that the same is true for BE, although the value of HWE −W is zero at xW = 0. For 2P and 1Hx, the second group, both HWE −W and TSWE −W are already negative at xW = 0, but they decrease to the valley anomaly similar to the first group. BE’s behavior may belong to the first with both HWE −W and TSWE −W very close to zero. Thus, we suggested in Section [V-1] that BE forms a rather tight cluster of its own, to which H2 O molecules interact one by one. Hence the W−W interaction is almost non existent. We suggest the BE clusters could be of a micelle type with −OH group pointing outwards to which H2 O molecules interact one by one without feeling each other’s presence. This behavior of BE is out of place in comparison with 1B, 2P, and 1Hx. No doubt the presence of ether −O− makes BE unique in the mono-ol series. From the loci of the valley anomaly listed in Table 5, we may rank the effect of the alkyl group on the II–III transition as we did for the I–II transition: Namely, the larger

144

Fig. V-43(a) to V-43(c)

20

(a)

TS w-wE / kJ mol-1

10

0

ME ET 1P 1B 2P TBA BE

-10

-20

-30

-40 0.2

0.4

0.6

0.8

1.0

xAL E Fig. V-43(a). The H2 O−H2 O interaction in terms of entropy times T TSW −W , in aqueous methanol (ME), ethanol (ET), 1-propanol (1P), 2-propanol (2P), tert-butanol (TBA), and 2-butoxyethanol (BE). All data are from Liltorp et al. (2005) except for TBA. The data for TBA−H2 O were calculated using HWE data from Koga (1988b) and the vapor pressure data.(Koga et al., 1990a) which were re-analyzed by the method discussed in Section [III–7]. Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

8 6

(b)

TSw-wE / kJ mol-1

4 2 0

ME ET 1P 1B 2P TBA BE

-2 -4 -6 -8

0.2

0.4

0.6

0.8

1.0

xAL

Fig. V-43(b). The expanded version of Fig. V-43 (a). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

Eq. (V-16)

145

10

TSw-wE / kJ mol-1

(c)

0

-10

1B 1Hx 1Oc 1De -20 0.85

0.90

0.95

1.00

xAL E Fig. V-43(c). The H2 O−H2 O interaction in terms of entropy times T TSW −W , in aqueous 1-butanol (1B), 1-hexanol (1Hx), 1-octanol (1Oc), and 1-decanol (1De), at 25  C. The data are from Liltorp et al. (2005). Reproduced with permission from the Canadian Journal of Chemistry, 83, 420, (Liltorp, K. et al.), Copyright (2005), NRC Research Press, National Research Council of Canada, Ottawa.

the absolute value of HWE −W and TSWE −W at the bottom of the valley and the smaller the value of xW = 1 − xAL , the stronger is the effect of alkyl group. Thus, the ranking is given as, ME
(V-16)

The oddity of BE in the above ranking could again be due to the presence of −O−, and that of 2P must be due to an extra methyl group. If we call eq. (16) the ranking of hydrophobicity, it must be emphasized that this is for the effect of H2 O on a predominantly AL environment, while the ranking given earlier in Section [V-2] as eq. (3) is for the effect of each AL on liquid H2 O in the H2 O-rich environment. These two situations must clearly be specified in discussing the “hydrophobicity” in general terms. The third group consisting of TBA, 1Oc, and 1De show a different propensity. The enthalpic W−W interaction is strongly favorable, while that in terms of entropy requires equally large costs at the infinite dilution. This seems to suggest that H2 O in these AL-rich environment forms clusters of H2 O more organized than pure H2 O. Furthermore there are no signs of a valley anomaly in HWE −W , nor TSWE −W in these AL, as is evident in Figs. 42 and 43, i.e. these AL−H2 O systems do not have Mixing Scheme III. For TBA, Mixing Scheme II extends from xAL = 1 all the way to point Y of the I–II transition listed in Table 3 (p. 114). For 1Oc and 1De, on the other hand, Mixing Scheme II spans from the phase separation boundary to xAL = 1. As given in Table 2 (p. 111), the phase

146 E  − Table V-5. The summary of the H2 O−H2 O interactions, HWE −W and TSW −W in AL H2 O at 25 C. The loci of the II–III boundary are determined with more confidence than listed in Table V–3.

From HWE −W HWE −W /kJ mol−1 AL

at xW = 0 ±07

ME BE ET 1P 1B

6 0 5 4 2

2P 1Hx

−3 −5

1Oc 1De TBA

−17c −18 −31

E From TSW −W

xAL -loci of valley

−1 E TSW −W /kJ mol

xAL -loci of valley

at valley ±05

II-III boundary ±003

at xW = 0 ±07

03 −021 ± 005a −10 −28 −53

057 065 076 080 081

8 0 7 6 5

15 ---b 09 02 −10

058 ---b 073 079 080

−60 −90

088 090

−2 0

−30 −50

089 090

---c -----

---c -----

-------

-------

−15 at xW = 0 −18 at xW = 0 −31 at xW = 0

at valley II-III boundary ±05 ±003

For BE−H2 O, a more sensitive calorimeter was used. HWE for BE−H2 O was determined with more than usual accuracy with much smaller increment. EW data E for the same system do not match the precision of HWE . Hence, TSW could not be calculated. c For 1Oc, 1De, and TBA, no valley anomaly was observed (see Figs. 42 and 43). Hence, there is no Mixing Scheme III in aqueous solutions of theses alcohols. Instead of the loci of the valley anomaly, the values of E HWE −W and TSW −W at the infinite dilution, xW = 0 are given. a

b

boundary in the AL-rich region is at xAL = 0793 and 0.735 (xW = 0207 and 0.265) for 1Oc and 1De, respectively. Thus in these larger alcohols H2 O-rich and AL-rich clusters form at the outset, xAL = 1. From the values of HWE −W and the almost compensating values of TSWE −W at xW = 0, we may rank the effect of alkyl group (“hydrophobicity”) within the three AL in the third group as, 1Oc<1 De
(V-17)

The surprising strength of TBA in this regard (eq. (17)), is no doubt due to two extra methyl groups. Here again, this ranking is applicable only for how H2 O sees in the AL-rich environment. In closing this section, we stress that the behavior of AL in H2 O is compositionspecific as we have seen in this chapter, and that caution must be exercised in the application of the findings of this and other studies for the nature of biopolymers in H2 O. Certainly a blanket application of our findings of any one of the rankings in “hydrophobicity” (eqs. (3), (16) and (17)), without paying attention to the appropriate composition will not bring about a fruitful discussion.

Eq. (V-17)

147

[V-6] MIXING SCHEMES OF AQUEOUS ALCOHOLS (AL) STUDIED BY OTHER TECHNIQUES As we saw in this chapter, Mixing Scheme I is operative only in the H2 O-rich region, xME <0065, and xBE <0017 (the top and the bottom of Table 3 (p. 114)). xME and xBE are the mole fractions of methanol (ME) and 2-butoxyethanol (BE) in the respective aqueous solutions. For such a dilute range, spectroscopic signals are generally over-swamped by those by the dominant solvent H2 O. Hence, data in the range xAL <01 are scarce, and the detailed spectroscopic studies tend to focus in the range xAL >01. Nonetheless, some attempts are found in the literature to cover the range xAL <01. Shirota and Castner (2000) add coumarin 153 into 1-propanol (1P)−H2 O as a fluorescent probe. They found that the wave numbers of the absorption and the fluorescence maxima showed a sharp break in the mole fraction range of 1P, 005<x1P <01, between the first two data points. This suggests the crossover in the mixing scheme that corresponds to our transition from Mixing Scheme I to II. However, the details of the mixing scheme were not given. Sato and Buchner (2005) reviewed their studies on aqueous methanol (ME), ethanol (ET), 1-propanol (1P) and 2-propanol (2P) by dielectric spectroscopy. They concluded that in the AL-rich region, xAL > about 0.5, corresponding roughly to the range of Mixing Scheme III, AL molecules form a zigzag hydrogen bonding chain similar to pure AL to which H2 O molecules are inserted. This is consistent with our Mixing Scheme III (see Tables 3 (p. 114) and 5 (p. 146) for the boundary between Mixing Schemes II and III). From the temperature dependence of the relaxation time of the co-operative dynamics of the H-bonding system, they obtain the excess activation enthalpy and entropy. They then take derivative with respect to xAL and calculated the excess partial molar activation enthalpy and entropy of AL. The latter results were plotted against xAL . In the H2 O-rich region, two peak anomalies are apparent at x1 and x2 , which are listed in Table V-6, together with our point X and Y taken from

Table V-6. The xAL -loci of anomalies in the excess partial molar activation enthalpy of AL. AL = methanol(ME), ethanol (ET), 1-propanol(1P) and 2-propanol(2P). The results of dielectric relaxation spectroscopy (Sato and Buchner, 2005)

ME ET 1P 2P a

x1

x2 a

X

Yb

0045 004 003 003

014 008 006 007

007 006 0049 005

025 019 009 011

The loci of two anomalous peaks in the excess partial molar activation enthalpy and entropy calculated from the xAL and temperature dependence of the relaxation time. See Sato & Buchner (2005) for detail. b The loci of the transition region from Mixing Scheme I to II from our studies. Taken from Table 3 (p. 114).

148

Table 3 (p. 114). Their interpretation for the mixing scheme in the H2 O-rich region is not the same as our Mixing Scheme I. In this low concentration range, the complex permittivity was too small for detailed analysis. Furthermore, the measurements were made at 20, 25, and 30  C. From these data, Eyring’s theory was applied to calculate the activation enthalpy and entropy. While the appropriateness of the use of Eyring’s theory is yet to be clarified, the uncertainty in the resulting activation quantities could be sizable. It follows then that the uncertainty in the partial molar activation quantities would be quite significant. Nevertheless, the fact that anomalies were observed at x1 and x2 , which are in the same range as the loci of our point X and Y (Table 6) support at least the finding that a crossover of mixing scheme is present in the H2 O-rich region. The so-called self-association parameter can be determined by NMR measurements. Holz et al. applied this to tert-butanol (TBA)−H2 O (Mayele et al., 1999; Mayele & Holz, 2000). The self-association parameter is obtained by measurements of the intermolecular H−H dipole–dipole relaxation rate of methyl protons and of self-diffusion coefficients of TBA. It is directly related to the spin density reduced first coordination number of H around a reference H on TBA. Its xTBA -dependence showed the maximum at xTBA = 002∼003 at 25  C. On increasing temperature, the xTBA locus of the maximum decreased; the same temperature dependence of the Koga line. They interpreted that the region below the maximum is dominated by single hydrated TBA molecules, while the higher concentration region is characterized by microstructures in which TBA molecules are hydrophobically associated. Though the boundary value of xTBA is somewhat smaller than our point X of the transition from Mixing Scheme I to II (Table 3 (p. 114)) the mixing schemes given by them are consistent with our Mixing Schemes I and II. Dixit et al. (2000) used high resolution Raman spectroscopy to study methanol (ME)−H2 O. From the xME -dependence of the wave numbers of C−O and C−H stretching modes, they suggested that there are three composition regions each with a distinct mixing scheme. In the ME-rich region, xME > about 0.7, ME remains in the chain structure of pure ME and hydration takes place at the chain ends, −OH group. In the intermediate region, 02<xME <07 H2 O breaks up ME chains and ME molecules also become hydrated individually. By about xME ∼015 −OH groups of ME are surrounded completely by H2 O and the hydration of −CH3 takes place. At about xME ∼005, the hydration of ME at the −CH3 as well as the −OH parts is complete. This interpretation is almost completely consistent with our Mixing Schemes III, II and I, from the ME-rich end. The boundary values also corroborate approximately our values (see Table 3 (p. 114) and Table 5 (p. 146)). The same authors then went on to study ME−H2 O by neutron diffraction with hydrogen/deuterium substitution (Dixit et al., 2002a and 2002b). They used the empirical potential structure refinement method (Soper, 1996) to analyze the diffraction data and to construct a structure model of the mixture. They concluded that at xME = 07 ME molecules cluster together (Dixit et al., 2002a), which is consistent with our Mixing Scheme III. They repeated the same study on the mixture at xME = 005, and concluded that ME also forms clusters of its own and there is no enhancement of H2 O structure surrounding – CH3 (Dixit et al., 2002b), which is in contrast with our Mixing

149

Scheme I and their earlier interpretation of the Raman results (Dixit et al., 2000). More recently, Dougan et al. (2004) applied the same neutron diffraction methodology to ME−H2 O at xME = 005, 0.27, 0.54 and 0.7. In addition to similar conclusions about the self-aggregation of ME throughout the entire composition range, they suggested that in the region 027 < xME < 054, where Mixing Scheme II is operative according to our studies, both H2 O and ME clusters are bond-percolated simultaneously, making this mixture a bi-percolating liquid mixture, which is at odds with our Mixing Scheme II. They also concluded that the percolating H2 O clusters are two-dimensional in this intermediate region. At xME = 005, however, the H2 O clusters are also percolated with three-dimensional characteristics. The latter conclusion at xME = 005 is consistent with our Mixing Scheme I. Exactly the same neutron diffraction method was applied to study tert-butanol (TBA)−H2 O, at the mole fractions xTBA = 002 (Bowron & Finney, 2003), 0.057, 0.1071, and 0.167 (Finney et al., 2000; Bowron et al., 1998). The solution at xTBA = 002 is in Mixing Scheme I region, that at xTBA = 0057 is in the transition region to Mixing Scheme II, and the remaining two mixtures are in the Mixing Scheme II region (see Table 3 (p. 114)). Their conclusion is that in all the mixtures studied, including that at xTBA = 002, TBA molecules aggregate together and thus TBA−H2 O is de-mixed at the molecular level. That TBA molecules cluster together is consistent with our Mixing Scheme II operative in the last two mixtures, but not in dilute mixtures, particularly for the first solution at xTBA = 002. In all of these neutron diffraction studies, the so-called “empirical potential structure refinement procedure” (Soper, 1996) is used to extract the structural model from the diffraction data. It uses an iteration of simulations, starting with a pair-wise potential, following refinement of the potential using the resulting paircorrelation function. As with any other simulation technique, the number of molecules is usually at most of the order of 103 , which is unrealistically small in comparison with the Avogadro number∗ . The rationale behind this is to sample microstates, as represented by the simulated system, over a sufficient number of configurations, so that their average is equivalent to that obtained over a macroscopic system. While this argument works out quite well for systems that are locally homogeneous, it may not be adequate for aqueous mixtures, which are locally heterogeneous. In particular, Mixing Scheme I may require to simulate a coarse grain of the order of 106 . We discussed this issue already in Section [V-1] regarding the long-range interaction between BE molecules operative in Mixing Scheme I in BE−H2 O. While the size issue in simulation may be an important one, there are other approximations that are often made that may affect the outcome, though to a lesser extent. For example, the interaction potential is necessarily approximated as a pair potential without the real many-body effects. When these difficulties are sorted out, this neutron diffraction methodology would no doubt provide a picture which is a step closer to the reality.∗ Recently, Misawa et al. (2006) studied 1-propanol (1P) – H2 O by quasi-elastic neutron scattering at the mole fraction x1P = 002, 0.04, 0.06, 0.1 and 0.167. The Mixing Scheme boundary from I to II is from x1P = 0049 (point X) to 0.09(point Y), as shown in ∗

I thank Aurelien Perera, Pierre et Marie Curie University, for illuminating discussion on this issue.

150

Table 3 (p. 114). The resulting spectra show a smooth variation on x1P without observing any sign of the change-over of mixing scheme. Thus, within a limited number of studies on aqueous alcohols reviewed here, the mixing schemes learned from various techniques seem to be more or less consistent in the intermediate and the alcohol-rich regions, i.e. Mixing Schemes II and III. The findings by other techniques for the H2 O-rich region are rather mixed. Some are consistent with our Mixing Scheme I, while others are not. We note, at the risk of being simplistic, that spectroscopy/scattering/diffraction techniques are generally for gaining structural information. The structure is dictated, in terms of thermodynamics, by the Gibbs energy, G. G is the result of the integration of H S V (the first derivative of G) and other higher derivatives. Limiting our attention at the H and S level for now, H and S tend to compensate each other in aqueous solutions (Lumry & Rajender, 1970; Lumry, 2003). Hence, anything conspicuous in the behaviors of H and S could very well be smoothed out in the net value of G. Thus, while G contains all the thermodynamic information, it could be insensitive to the details conspicuous at the H and S level. By the same token, the information accessible only at the second and third derivative quantities could be obscured at G level, and hence at spectroscopy/scattering/diffraction data. This point will be stressed again in Section [VIII-7-3]. Finally, we add a brief discussion about the consequence of distinct compositiondependent mixing schemes on chemical reactions taking place in the mixed solvent of aqueous alcohol. Kasza & Koga (1989) studied the kinetics of the trans- to cisisomerization of the trans-bis(oxalato)diaquachromium (III) in tert-butanol (TBA)−H2 O and ethanol (ET)−H2 O. They found that the apparent first order rate constant decreases on increasing the composition of alcohol(AL), and its xAL -dependence made a clear break at the Mixing Scheme boundary from I to II. Li et al. (1996) reported a dramatic increase in the solubility of ethane into TBA−H2 O at xTBA ≈ 0055, which corresponds to about half way from point X to point Y. Namely, when TBA−H2 O is beginning to lose the bond-percolated hydrogen bond network of H2 O and to form TBA-rich clusters, ethane molecules dissolve readily presumably into TBA-rich clusters. Oba et al. (1997) found that chlorophyll a’ suspension in aqueous methanol makes a dramatic change in the colloidal behavior at about xME = 015, in between point X and Y. There must be a large number of observations about the change in chemistry at a certain composition when mixed solvents of alcohol−H2 O are used; they are indeed used widely. A comprehensive treatise is awaited on the effect of composition of mixed aqueous solvents on the chemical properties of a solute.

Chapter VI

MIXING SCHEMES IN AQUEOUS SOLUTIONS OF NON-ELECTROLYTES

[VI-1] INTRODUCTION In this chapter we look into the mixing schemes in aqueous solutions of other nonelectrolytes than the mono-ols studied in Chapter V. We learned the effects of the hydrophobic moiety by changing the size of the alkyl group within the mono-ol series. The non-electrolyte solutes studied here are amphiphiles (as are mono-ols) and contain both a hydrophilic and a hydrophobic group with various relative strengths. As will become evident, the effects of hydrophobicity and hydrophilicity on H2 O seem to be cooperative and/or competitive within the H2 O-rich region, and manifest themselves clearly in third derivative quantities. We use here the same methodology of utilizing the second and third derivative quantities, and extract information out of the temperature and the composition dependence of these higher order derivatives. As will be shown below, we learned that there are also three mixing schemes depending on the composition as in the case of mono-ols−H2 O. In the H2 O-rich region the mixing scheme could also be characterized as protecting the integrity of liquid H2 O, whereby the hydrogen bond network is connected throughout the entire bulk of H2 O with a large degree of fluctuation in intensity (amplitude) as well as in extensity (wavelength). Mixing Scheme III in the solute-rich region is the same as for what was discussed in Chapter V in that the solute molecules cluster together as in their pure liquid to which H2 O molecules interact. Mixing Scheme II operative in the intermediate composition is a mixture of two kinds of clusters: One rich in H2 O and the other in solute molecules. Since the mixing scheme in the H2 O-rich region is to retain the integrity of liquid H2 O, we may still call this Mixing Scheme I. However, the detail of how the solute modifies the molecular organization of H2 O depends on the nature of the E as shown in solute. Its manifestation is apparent in the mole fraction dependence of Hi−i E Fig. VI-1. We note that the xi -dependence of Hi−i in the intermediate and the solute-rich regions is the same as in mono-ol−H2 O, suggesting that Mixing Schemes II and III are also operative. In addition to the mono-ols dealt with in Chapter V, we have studied a total of 11 aqueous non-electrolytes, as shown in Table VI-1. Within the systems we studied, these four types sketched in Fig. 1 cover all cases. As discussed in Chapter V, case (a) is typical for a “hydrophobic” solute which enhances the hydrogen bond network in its immediate vicinity (“iceberg” formation) but reduces the hydrogen bond probability of

152

Fig. VI-1(a) to VI-1(d)

X

H i-iE

(a)

Type (a)

Y 0

I

II

III

0

1

xi E Fig. VI-1(a). The xi -dependence of Hi−i for type (a) solutes; alcohols including isobutoxyethanol and acetonitrile. I, II, and III are the regions where Mixing Schemes I, II and III are operative. Points X and Y mark the transition region from I to II. The data are taken from Koga (2003a). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis.

(b)

X

Hi-iE

Type (b)

I

III Y

0

II 1

0

xi E Fig. VI-1(b). The xi -dependence of Hi−i for type (b) solutes; iso-butyric acid, and 2-butanone. I, II, and III are regions where Mixing Scheme I, II, and III are operative. Points X and Y mark the transition region from I to II. The data are taken from Koga (2003a). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis.

153

(c)

H i-iE

X

Type (c) Y I 0

II

III

0

1

xi E Fig. VI-1(c). The xi -dependence of Hi−i for type (c) solutes: dimethylsolfoxide, and 1,2-propanediol. I, II, and III are regions where Mixing Schemes I, II, and III are operative. Points X and Y mark the transition region from I to II. The data are taken from Koga (2003a). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis.

(d)

X Y Type (d)

H i-iE

I II

III

0 1

0

xi E Fig. VI-1(d). The xi -dependence of Hi−i for type (d) solutes; glycerol, acetone, 1,3-propanediol, and tetramethylurea. I, II, and III are the regions where Mixing Schemes I, II, and III are operative. Points X and Y mark the transition region from I to II. The data are taken from Koga (2003a). Reproduced with permission from the Netsusokutei (Journal of the Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis.

154

Table VI-1. The excess partial molar quantities in aqueous none-electrolytes at the infinite dilution. Reproduced with permission from the Netsusokutei (Journal of Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and thermal Analysis. Solute (i)

Abbr.

Type

temp. C



2-iso-Butoxyethanol Acetonitrile

iBE ACN

iso-Butyric acid

IBA

2-Butanone Dimethylsulfoxide 1,2-Propanediol Glycerol 1,3-Propanediol Acetone Tetramethylurea Ethylene Glycol

BUT DMSO 12P Gly 13P AC TMU EG

a a a a a a b b b c c d d d d d

20 6 20 30 37 45 25 27 25 25 25 25 25 25 25 25

HiE /kJ mol−1 ±01 −07 −33 −19 −10 −04 04 −16 −15 −105 −175 −105 −56 −87 −152 −235 −65

Values at xi = 0 (ref.) Ei /kJ mol−1 (a) (b) (b) (b) (b) (b) (c) (c) (d) (f) (g) (h) (g) (j) (j) (n)

(ref.)

ViE initial decrease? Yes/No (ref.)

±01 60

(a)

Yes

(a)

58

(b)

Yes

(b)

80 −69 −03 −21 −04 50 15

(d) (f) (g) (h) (g) (j) (j)

Yes Yes Yes No Yes Yes Yes Yes

(e) (f) (g) (h) (g) (j) (k, m) (o, p)

Fig. VI-2

References:(a) Cheung et al., (2003). (b) Nikolova et al., (2000). (c) Siu et al., (1991). (d) Wong et al., (1992). (e) Boeje & Hvidt, (1971). (f) Lai et al., (1995). (g) Parsons et al., (2003). (h) To et al., (1999). (j) Chen et al., (2000). (k) Jakli & Van Hook, (1996). (m) Phillips et al., (1974). (n) Koga (2003b). (o) Huot et al. (1988). (p) Sakurai (1991).

155

bulk H2 O away from “icebergs.” As shown below, case (d) is typical of a “hydrophilic” solute. Cases (b) and (c) could be regarded as the net result of the mixture of the effects of a hydrophobic and a hydrophilic moiety within a given non-electrolyte. These findings receive more support in Chapter VII in which we adopt the 1-propanol(1P) probing method developed by us (Koga, 2002). Namely, we study a three-component system, E 1P−A−H2 O in the H2 O-rich region. We then obtain the 1P−1P interaction, H1P−1P , and E look for the changes induced in the x1P -dependence pattern of H1P−1P on addition of the third component A. These induced changes turned out to provide important information regarding the effect of A on the molecular organization of H2 O.

[VI-2] TYPE (A) – AQUEOUS SOLUTIONS OF ISO-BUTOXYETHANOL (iBE) AT 20  C AND ACETONITRILE (ACN) AT 6 TO 45  C The type (a) anomaly in Fig. 1(a) is already familiar to us since we discussed it at some length in Chapter V. While iso-butoxyethanol (iBE) is a mono-ol, we did not include it in Chapter V, since iBE−H2 O has the LCST at 256  C at xiBE = 0052 (Perron et al., E 1993), and the available data for HiE and Hi−i i = iBE are only at 20  C. (Cheung et al., E 2003) (see Figs. VI-2 and VI-3). The excess partial molar volume of iBE, ViBE , calculated using the density data (Davis & Chacon,1991) shows the initial concave decrease

0.5 0

I

0.0

-2

H iBE E / kJ mol-1

II

-0.5

III

-4 -1.0

II

-6 -1.5 -8 -2.0 -10 -2.5 0.00 0.02 0.04 0.06

0.0 0.1 0.2 0.3 0.4 0.5 0.6

xiBE

xiBE

E Fig. VI-2. The excess partial molar enthalpy of iso-butoxyethanol (iBE), HiBE in iBE−H2 O at 20  C. I, II, and III are the regions where Mixing Schemes I, II and III are operative. The data are taken from Cheung et al. (2003). Reproduced from Thermochimica Acta, 397, 75, (Cheung et al.), Copyright (2003), with permission from Elsevier.

156

Fig. VI-3 to VI-5(a)

600

15

X

H iBE-iBE E / kJ mol-1

500 10

400 300

5 200 100

0

0

I

II

Y

III

II

-100

-5 0.00 0.02 0.04 0.06 0.08 0.10

0.0 0.1 0.2 0.3 0.4 0.5 0.6

xiBE

xiBE

E Fig. VI-3. The iBE−iBE enthalpic interaction, HiBE−iBE , in iBE−H2 O at 20  C. I, II, and III are the regions where Mixing Schemes I, II, and III are operative. The data are taken from Cheung et al. (2003). Reproduced from Thermochimica Acta, 397, 75, (Cheung et al.), Copyright (2003), with permission from Elsevier.

as for all the mono-ols studied above (see Fig. VI-4). The values of HiE and TSiE remain zero for xi > 035i = iBE. In the intermediate range, 0026 < xi < 035, both slopes E E of Hi−i and TSi−i are small but positive, consistent with the system having an LCST (Cheung et al., 2003). Thus, iBE−H2 O displays all the hallmarks of mono-ols−H2 O, and the mixing schemes are the same as discussed in Chapter V, which is of course not surprising. Surprisingly, acetonitrile (ACN)−H2 O belongs to this class, from the behavior of the second and third derivative quantities. The mole fraction dependence of the enthalpic E interaction, Hi−i for i = ACN, is shown in Fig. VI-5 (Nikolova et al., 2000). Table VI-2 E (p. 159) lists the behavior of Hi−i . The excess partial molar volume, ViE i = ACN, calculated from the density data (Handa & Benson, 1981) indicates an initial decrease in the range xi < 001, although not as conspicuously as mono-ols (see Fig. 4). Other third derivative quantities calculated using the data of heat capacity and speed of sound (Benson et al, 1981) for ACN−H2 O also suggests anomalies at about the same loci of point X shown in Fig. 1(a) (see Cheung et al. (2003) for detail). We plot these anomalies E in Hi−i , i.e. point X, at temperatures from 6 to 45  C in Fig. VI-6, together with those − for BE H2 O and TBA−H2 O, shown in Figs. V-13 and Fig. V-26. As evident in Fig. 6, the Koga line seems to point towards about 85  C at the infinite dilution independent of solute, though we have only three cases so far. The system ACN−H2 O has a UCST at −1  C at xACN = 038 (Armitage et al., 1968) and the small angle x-ray scattering intensity increases dramatically above and near the UCST (Nishikawa et al., 2002), supporting the existence of Mixing Scheme II in the intermediate composition range.

157

2 0

E

V i / cm3 mol-1

-2 -4 -6

iBE

-8

ACN -10 -12 0.00

0.10

0.05

0.15

0.20

0.25

0.30

x i (i = iBE or ACN)

Fig. VI-4. The excess partial molar volumes of i, for i = isobutoxyethanol (iBE) at 20  C, and acetonitrile (ACN) at 25  C. Calculated by Cheung et al. (2003) using the density data by Davis & Chacon (1991) for iBE. For ACN, calculated by Nikolova et al. (2000) using the density by Handa & Benson (1981). 140

140

X

120

120

37 °C 6 °C

100

(a)

H ACN-ACNE / kJ mol-1

100 80 80

II

60 40

60

I

20 40

0 -20 0.00

20

Y 0.05

0.10

0.15

0.20

0

II

-20 0.0

0.2

III 0.4

0.6

0.8

1.0

xACN E  Fig. VI-5(a). The enthalpic interaction in acetonitrile (ACN)−H2 O, HACN −ACN , at 6 and 37 C. The data are taken from Nikolova et al. (2000). Reproduced with permission from the Canadian Journal of Chemistry, 78, 1553, (Nikolova, P. V. et al.), Copyright (2000), NRC Research Press, National Research Council of Canada, Ottawa.

158

Fig. VI-5(b) to VI-6

140

X Error

H E ACN-ACN / kJ mol-1

120

(b)

100

80

60

40

20 0.00

6 °C 20 °C 30 °C 37 °C 45 °C 0.01

0.02

0.03

0.04

0.05

xACN E Fig. VI-5(b). The enthalpic interaction in acetonitrile (ACN)−H2 O, HACN −ACN , at temperatures 6, 20, 30, 37, and 45  C. The data are taken from Nikolova et al. (2000). Reproduced with permission from the Canadian Journal of Chemistry, 78, 1553, (Nikolova, P. V. et al.), Copyright (2000), NRC Research Press, National Research Council of Canada, Ottawa.

100

BE TBA ACN

80

t / °C

60 40 20 0 -20 -40 0.00

0.02

0.04

0.06

0.08

0.10

x i (i = BE, TBA or ACN)

Fig. VI-6. The Koga line for acetonitrile (ACN)−H2 O (Nikolova et al., 2000), together with that for 2-butoxyethanol (BE)−H2 O, taken from Fig. V-13, for tert-butanol (TBA)−H2 O taken from Fig, V-26. Reproduced with permission from the Netsusokutei (Journal of Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis.

Table VI-2. The solute – solute interactions. Reproduced with permission from the Netsusokutei (Journal of Japan Society of Calorimetry and Thermal Analysis), 30, 54, (Koga, Y.), Copyright (2003), the Japan Society of Calorimetry and Thermal Analysis. Solute

type

iBE ACN

a a a a a a b b b c c d d d d d

IBA BUT DMSO 12P Gly 13P AC TMU EG

temp.  C

20 6 20 30 37 45 25 27 25 25 25 25 25 25 25 25

sign of Ei−i −ve −ve

−ve +ve 0 +ve 0 −ve −ve

E Hi−i /kJ mol−1 at xi = 0

E /kJ mol−1 Hi−i at point X

240±30 40±5 40±5 40±5 40±5 40±5 170±10 170±10 139±8 51±5 78±7 42±7 59±5 120±10 220±15 43±5

550±50 133±7 122±7 110±7 100±7 95±7 170±10 170±10 139±8 34±5 47±7 13±3 25±3 40±5 100±8 19±5

Loci of X

Y

±0005 0012 0026 00125 0085 00106 00088 00080 0060 00065 0022 0043 0022 0044 004 021 035 010 019 010 015 010 020 008 025 008 026 011 017

II–III ±01 035

045

056 037 034 030 050 040 029

ref.

(a) (b) (b) (b) (b) (b) (c) (c) (d) (e) (f) (g) (f) (h) (h) (j)

References:(a) Cheung et al., (2003). (b) Nikolova et al., (2000). (c) Siu et al., (1991). (d) Wong et al., (1992). (e) Lai et al., (1995). (f) Parsons et al., (2003). (g) To et al., (1999). (h) Chen et al., (2000). (j) Koga (2003b).

159

160

Fig. VI-7 to VI-8(b)

[VI-3] TYPE (D) – AQUEOUS SOLUTIONS OF GLYCEROL (GLY), ACETONE (AC), 1,3-PROPANEDIOL (13P), AND TETRAMETHYL UREA (TMU) Fig. 1(d) shows that Hi-i E decreases sharply, and then turns abruptly into the region with a lesser slope, as xi increases. We identify the latter part in the intermediate composition region as Mixing Scheme II by the similarity of the value and the slope of Hi-i E to type (a) solutes and also to others, type (b) and (c). Fig. VI-7 shows the excess partial molar volume, ViE , of i. For acetone (AC) (Chen et al., 2000), 1,3-propanediol (13P) (Parsons et al., 2003; Nakanishi et al., 1967) and tetramethyl urea (TMU) (Jakli & Van Hook, 1996; Phillips et al., 1974), ViE (i = AC, 13P, or TMU) shows an initial decrease, suggesting a hydrophobic contribution. On the other hand, glycerol(Gly) has no such trend in ViE (i = Gly) (see also Table 1 (p. 154)). There are a few examples of this non-hydrophobic behavior of ViE ; urea (Boeje & Hvidt, 1971) which is shown in Fig. 7, formamide (Boeje & Hvidt, 1971), and d-fructose (To et al., 2000) to name but a few. While urea and formamide are clearly hydrophilic solutes, fructose has −OH on five C on the fructose ring with a remaining C connected to ether −O−. Gly also has three −OH groups on every C on the propyl chain. Thus Gly could be regarded also as a typical hydrophilic solute. Figs. VI-8(a) and (b) are the plots of the entropy-volume cross fluctuations, SV  and SV , introduced in Chapter IV for Gly−H2 O. Also shown in the figure are those for 1-propanol (1P)−H2 O, reproduced from Fig. V-31 for the purpose of comparison. Recall

2

-2

-4

Gly

-6

13P AC

E

V i / cm3 mol-1

0

TMU-Jakli TMU-Phillips U

-8

-10

0.0

0.2

0.4

0.6

0.8

1.0

x i (i = Gly, 13P, AC, TMU, or U)

Fig. VI-7. The excess partial molar volume of solute i, ViE , for i = glycerol (Gly), 1, 3-propanediol (13P), acetone (AC), tetramethyl urea (TMU), or urea (U) at 25  C. References are; Gly (To et al., 1999); 13P (Parsons et al., 2003; Nakanishi et al., 1967); AC (Chen et al., 2000); TMU with hollow inverted triangle (Phillips et al., 1974); TMU with filled inverted triangle (Jakli and Van Hook, 1996); U (Boeje & Hvidt, 1971).

161

0.30

0.25

(a) SV

δ

0.20

0.15

Gly

0.10

1P

0.05 0.0

0.1

0.2

0.3

0.4

0.5

x i (i = Gly or 1P)

Fig. VI-8(a). The entropy-volume cross fluctuation density, SV , for glycerol (Gly)−H2 O, or 1-propanol (1P)−H2 O at 25  C. The region between two vertical dotted lines is where Mixing Scheme II is operative. The straight line is the extension from region II to xi = 0 and indicates the presumed “normal liquid” behavior. The difference from this line to the plot is the negative contribution to SV  due to putative ice-like patches. As xi increases therefore, this characteristic negative contribution diminishes. The data for Gly are taken from To et al. (1999). 0.08

SV

Δ / JK-1 cm-3

0.07

0.06

(b)

0.05

0.04

0.03

Gly 1P

0.02 0.0

0.1

0.2

0.3

0.4

0.5

x i (i = Gly or 1P)

Fig. VI-8(b). The normalized entropy-volume cross fluctuation, SV , for glycerol (Gly)−H2 O, or 1-propanol (1P)−H2 O at 25  C. The region between two vertical dotted lines is the region where Mixing Scheme II is operative. The straight line is a smooth extension of the xi -dependence of SV  towards xi = 0, and indicates the presumed normal liquid behavior. The difference from this line to the data point is the negative contribution to SV  due to the putative ice-like patches characteristic to liquid H2 O. For Gly, the data are taken from To et al. (1999).

162

SV

Fig. VI-9(a) to VI-9(b)

 is the intensity or the amplitude of the cross fluctuation between entropy and volume, while SV  represents qualitatively the extensity (wavelength) as well as the intensity (amplitude). They are defined by eqs. (IV-9) and (IV-16) in Section [IV-2] (see also eqs. (V-6) and (V-10) in Section [V-3]). Both cross functions, SV  and SV , show initial increases on increasing xi for both i = Gly and 1P. Then SV  slows down its increase, while SV  starts to decrease, after point X. As discussed in Section [V-3] for 1P−H2 O, these initial increases could be attributed to the decrease of the negative contribution to the SV cross fluctuation due the special characteristics of H2 O. Unlike normal liquids, the fluctuation in a volume increase could be associated with the putative formation of ice-like patches which give rise to an entropy decrease. Hence H2 O contains a negative contribution to the SV cross fluctuation. The initial increases of the SV fluctuation as xi increases, shown in Fig. 8, may then represent the process by which this special characteristic of H2 O diminishes by addition of the respective solute. How then does Gly, type (d), differ from 1P, type (a), in modifying the molecular organization of H2 O? Firstly, we note that the degree of this increase is less for Gly than for 1P, as seen in Fig. 8. The thin straight lines are drawn smoothly extending the region of Mixing Scheme II in Fig. 8 for a presumed behavior of normal mixtures for normal liquids. The difference from this line (there is no particular reason why the line should be straight) represents the decrease of the negative contribution. What Fig. 8 indicates is that (a) the normal positive contribution and (b) the decrease in the negative contribution are both smaller for Gly−H2 O than for 1P−H2 O. This could mean that the lesser decrease in the negative contribution is coming from the lesser hydrophobicity of Gly. The effect of three −OH groups in Gly, on the other hand, is to reduce the positive part of SV cross fluctuation more than for 1P with only one −OH. While both 1P and Gly contain three carbon atoms, 1P has a CH3 - group but Gly has not. As will be shown in Chapter VII, the alkyl groups with CH3 - are stronger hydrophobes than those without (which is generally accepted), and the hydrophilicity depends heavily on the number of −OH and −O− groups. We now turn our attention to the curvature of increase in SV  and SV . For 1P, both fluctuation functions increase in a concave manner at first and turn into a convex increase at point X, while Gly makes the convex increase from the outset but the curvature changes at point X, though this effect is not as conspicuous as for HiE , due to the smaller number of data points for the SV cross fluctuations. These changes in curvature are more conspicuous in the next derivative, SV i and SV i shown in Fig. VI-9. Also shown E in the figure is the i−i enthalpic interaction,Hi−i for i = Gly (To et al., 1999). As is E could bring the case for 1P−H2 O, a single scaling factor applied to the values of Hi−i E SV about the slope of i from xi = 0 to point X matching that of Hi−i (see Fig. 9(a)). The same is true with another single scaling factor resulting in the same slopes (the E xi -dependence) between SV i and Hi−i from the infinite dilution to point X as shown in SV E with the Fig. 9(b). Furthermore, for i the slope from point X to Y also matches Hi−i SV same single scaling factor, but not quite so for i , as is evident in Fig. 9. We recall E and SV i in 1P-H2 O. For comparison, the same good match in both slopes between Hi−i the equivalent plots are reproduced in Fig. VI-10 from Fig. V-34. The match is perfect for 1P in that the values themselves in addition to the slopes could be made the same

163

0.8

H Gly-GlyE (arbitrary scale) SV

SV

δGly and H Gly-GlyE

0.6

δGly (a)

0.4

III

II 0.2

I X Y

0.0

0.0

0.1

0.2

0.3

0.4

0.5

0.6

x Gly

Fig. VI-9(a). The partial molar entropy-volume cross fluctuation density of glycerol (Gly), SV Gly , E and the enthalpic Gly–Gly interaction, HGly−Gly (arbitrary units), in Gly−H2 O at 25  C. The E ordinate HGly−Gly is scaled by a single factor. The data are taken from To et al. (1999).

0.3

ΔGly

SV

0.2

(b) X Y II

0.1

III

0.0

I

SV

ΔGly / J K-1 cm-3 and H Gly-GlyE

H Gly-GlyE (arbitrary unit)

-0.1 0.0

0.1

0.2

0.3

0.4

0.5

0.6

x Gly

Fig. VI-9(b). The partial molar normalized entropy-volume cross fluctuation of glycerol (Gly), E SV Gly , and the Gly–Gly enthalpic interaction, HGly−Gly (arbitrary units), in Gly−H2 O at 25  C. E The ordinate of HGly−Gly is scaled by a single factor. The data are taken from To et al. (1999).

164

Fig. VI-10(a) to VI-10(b)

2.5

(a)

2.0

1.5

δ1P

SV

δ1P

SV

H1P-1PE

1.0

0.5

0.0 0.1

0.0

0.2

0.3

x1P

Fig. VI-10(a). The partial molar entropy volume cross fluctuation density of 1-propanol (1P) in E 1P−H2 O, SV 1P , at 25  C. Also plotted are the 1P−1P enthalpic interaction H1P−1P , with arbitrary units. The same as Fig. V-34 (a).

1.0

I II

0.8

(b)

SV

Δ1P / J K-1 cm-3

X 0.6

Δ1P

0.4

SV

H1P-1PE 0.2

0.0

-0.2 0.00

Y

0.05

0.10

0.15

0.20

0.25

0.30

x1P

Fig. VI-10(b). The partial molar normalized entropy volume cross fluctuation of 1-propanol (1P) E in 1P−H2 O, SV 1P , at 25  C. Also plotted are the 1P−1P enthalpic interactions, H1P−1P , with arbitrary units. The same as Fig. V-34 (b).

165

within the experimental uncertainty from xi = 0 to point X. In addition, SV i matches E perfectly with Hi−i from point X to Y as well. For Gly, while the values themselves are not the same, the slopes (the xi -dependences) for the two regions, from xi = 0 to point X and from point X to Y, were made the same by a single scaling factor. This leads us E to suggest that for Gly−H2 O also, Hi−i and SV i share the same cause and hence the i−i interaction is mediated via the bulk H2 O where the fluctuation nature is modified by the presence of solute, just as for 1P−H2 O. Hence the i−i interaction is of a long range. Furthermore, the fact that the match in slope is better in SV i than in SV i for the region from point X to Y suggests the importance of the extensity (wavelength) of fluctuation in considering the nature of aqueous solutions. As for the i−i interaction being of a long range, we could not find the region of Henry’s law down to the mole fraction of 10−5 for Gly−H2 O together with urea−H2 O, tert-butanol−H2 O, and 2-butoxyethanol−H2 O (Koga, 1995a; Westh et al., 1998). If the i−i interaction starts to operate after a number of i are accumulated within the system, E then there should be a finite Henry’s law region with Hi−i = 0 below this composition E does not extrapolate to zero for (see Section [II-4]). Indeed, as is clear from Fig. 9, Hi−i Gly−H2 O and 1P−H2 O (and of course for other mono-ols) (see Table 2 (p. 159) and E Table V-3 (p. 114) for the values of Hi−i at xi = 0). The difference between Gly−H2 O E and SV i (Figs. 9 and 10). For Gly, and 1P−H2 O is clearly the xi -dependence of Hi−i E SV Hi−i and i start at a finite value and decrease to point X, while those for 1P start at a non-zero value and increase to point X. As discussed in Sections [V-1] and [V-3], the effect of 1P on the hydrogen bond probability profile can be sketched as in Fig. V-14. Namely, 1P enhances the hydrogen bond network of H2 O in its immediate vicinity (“iceberg” formation), which produces depletion in the hydrogen bond probability in the vicinity of the “icebergs”. The local depletion is gradually recovered with a tailing profile as sketched in Fig. V-14. These tails interact with each other as xi increases E to bring about the gradual increase in Hi−i (i = 1P) as discussed in Section [V-1]. In E contrast, the effect of Gly in terms of Hi−i as well as SV i is the strongest at the infinite dilution. This suggests that its effect is directly on the existing hydrogen bond network of H2 O. As xi increases, the effect is weakened since the previous solute molecules have already modified the hydrogen bond network. We tentatively suggest that Gly, as well as other type (d) solutes, form hydrogen bonds directly to the existing network through their hydrophilic moieties. By so doing, Gly breaks the H donor/acceptor symmetry of the original hydrogen bond network, which partially cleaves the network reducing the hydrogen bond probability and also makes the network more rigid, resulting in reduction in the degree of fluctuation. As discussed below in Section [VI-6], a recent simulation study on aqueous urea, a typical hydrophile lacking in C−H bonds, showed that urea leads to stiffening of the short-time dynamics of both urea and H2 O (Idrissi et al., 2000). This conjecture will gain experimental support in Chapter VII, in which we learn how urea together with Gly modifies the molecular organization of H2 O by the method using the second and third derivatives of 1P in 1P−Gly−H2 O. E E E Fig. VI-11 shows TSi−i together with Hi−i for i = Gly. As is evident, the values of Hi−i E are larger than those of TSi−i . Namely, the Gly–Gly interaction is repulsive in terms of

166

Fig. VI-11 to VI-12

35

H Gly-GlyE

H Gly-GlyE or TS Gly-GlyE / kJ mol-1

30

X Y

TS Gly-GlyE

25 20 15

II

I

III

10 5 0 -5 0.0

0.1

0.2

0.3

0.4

0.5

0.6

xGly E E Fig. VI-11. The enthalpic and entropic interactions of glycerol (Gly), HGly−Gly and TSGly−Gly , in E E Gly−H2 O at 25  C. Note that HGly−Gly > TSGly−Gly unlike mono-ols−H2 O. Thus the net Gly–Gly interaction in terms of chemical potential, EGly−Gly , is positive and unfavorable driven by enthalpy, or the net Gly−H2 O interaction is negative and attractive. The data from To et al. (1999).

enthalpy and attractive entropy-wise. This is the same as in all the other non-electrolyte solutes studied so far. The uniqueness of Gly−H2 O is that the enthalpic repulsion is E E < Hi−i , resulting in the net repulsion in terms stronger than the entropic attraction, TSi−i of chemical potential, Ei−i > 0. Or, equivalently, the net interaction between Gly and H2 O is attractive (favorable). This may be one of the signatures of hydrophilicity. All the other non-electrolyte solutes including other (d) type solutes studied by the present E E . The DMSO−H2 O is an > Hi−i methods show the net attraction Ei−i < 0, since TSi−i exception having Ei−i > 0, and will be discussed in some detail below. Other type (d) solutes, 1,3-propanediol (13P), acetone (AC), and tetramethyl urea (TMU) show more hydrophobicity than Gly. Judging from the degree of initial decrease in ViE (Fig. 7), TMU seems to be most hydrophobic. However, such effects are not E . Indeed, even Gly, having the alkyl backbone, may also have some conspicuous in Hi−i hydrophobic nature, though not apparent in the behavior of ViE . This point will also be dealt with in Chapter VII.

[VI-4] TYPE (B) – AQUEOUS SOLUTIONS OF ISO-BUTYRIC ACID (IBA) AND 2-BUTANONE (BUT) E , for i = IBA or BUT behaves as shown in Fig. 1(b). The xi -dependence of Hi−i Both systems have liquid–liquid phase separations, the macroscopic manifestation of Mixing Scheme II. For IBA−H2 O, the UCST is present at xi i = IBA = 01 and

167

at 272  C, though the critical slowing down prevents pin-pointing the UCST (Siu et al., 1991). For BUT−H2 O, the phase separation is present in the mid composition, and no critical de-mixing point is observed hidden by boiling or freezing. However, there is a hint of LCST below solidification temperatures. Indeed, the phase boundary in the H2 O-rich region slants with a negative slope, dT/dxi < 0 (i = BUT) (Wong et al., 1992). This (b)-type behavior of Hi−i E in the H2 O-rich region indicates that the i−i interaction in terms of enthalpy stays constant up to point X, rather than increases as in type (a). E (an increase in the repulsive enthalpic interaction) for type (a) The increase in Hi−i was attributed, as discussed in detail in Chapter V, to a decrease in the hydrogen bond probability of bulk H2 O away from solute, through which the solute–solute interaction E is operative. Thus the present case suggests some additional effect of decreasing Hi−i compensating the increase of type (a). This effect must be due to the hydrophilic moiety −COOH for IBA and >C= =O in BUT, against the hydrophobic effect from CH3 2 C – for IBA and CH3 -CH2 - and CH3 - for BUT. The net result appears to be some combination of the effects of the two moieties. This question gains some positive answer as will be discussed below in Chapter VII. The excess partial molar volumes, ViE , were calculated using the density data (Boeje & Hvidt, 1971) for i = BUT and those by Ouerfelli & Bouanz (2006) for i = IBA, and shown in Fig. VI-12. As is evident, BUT−H2 O has

4

E

V i / cm3 mol-1

2 0 -2 -4

12P -6

DMSO BUT

-8

IBA

-10

0.0

0.1

0.2

0.3

0.4

0.5

x i (i = 12P, DMSO, or BUT)

Fig. VI-12. The excess partial molar volume of solute i, ViE , for i = 1,2-propanediol (12P), dimethylsufoxide (DMSO), 2-butanone (BUT) or iso-butyric acid (IBA) at 25  C, except for IBA, which is at 29  C 2  C higher than its UCST. All possess a hydrophobic signature, i.e. ViE shows an initial decrease, except for IBA (see text). They are calculated using density data by Nakanishi et al. (1967) for 12P, Lai et al. (1995) for DMSO, Boeje & Hvidt (1971) for BUT, and Ouerfelli & Bouanz (2006) for IBA.

168

Fig. VI-13 to VI-14

an initial decrease, while for IBA there is no sign of hydrophobic signature in the xi dependence of ViE . In the intermediate composition for BUT−H2 O, in which the phase boundary slants in the way dT/dxi < 0 in the xi − T field, Hi-i E and TSi-i E have proper positive signs (Wong et al, 1992) For IBA, on the other hand, the sign of Hi−i E past the step anomaly (point X to Y in Fig. 1(b)) is properly negative for a system having a phase separation with the UCST at 272  C at xi = 01 (Siu et al., 1991). Thus, we conclude that Mixing Scheme II is operative past point Y.

[VI-5] TYPE (C) – AQUEOUS SOLUTIONS OF DIMETHYLSULFOXIDE (DMSO) AND 1,2-PROPANEDIOL (12P) As shown in Fig. 1(c), the unfavorable i−i interaction in terms of enthalpy, Hi-i E , decreases gradually as xi increases in contrast to types (a) and (b) and at point X it nosedives to the typical behavior for Mixing Scheme II. Thus, the detail of Mixing Scheme I for this type of solutes could be the net result of the two effects by the hydrophobic and the hydrophilic moieties, with a stronger latter effect than for case (b). The excess partial molar volume, ViE , for both i = DMSO and 12P, show the initial concave decrease as shown in Fig. VI-12, just as type (a) solutes, suggesting the hydrophobic contribution apparent in the xi -dependence of ViE . E Fig. VI-13 shows the excess partial molar enthalpy of DMSO, HDMSO , and that of E E H2 O HW . The subtle curvature changes in HDMSO in the H2 O-rich and the middle ranges 2 0

H EDMSO or H Ew / kJ mol-1

-2 -4 -6 -8 -10 -12

H EDMSO

-14

H Ew

-16 -18 -20 0.0

0.2

0.4

0.6

0.8

1.0

x DMSO E Fig. VI-13. The excess partial molar enthalpy of dimethylsulfoxide (DMSO), HDMSO , and that of H2 O HWE , in DMSO−H2 O at 25  C. The data are taken from Lai et al. (1995). Reproduced from the Journal of Solution Chemistry, 24, 89, (Lai, J. T. W. et al.), Copyright (1995), with kind permission from Springer Science and Business Media.

169

H Ei-i or TS Ei-i / kJ mol-1 (i = DMSO)

60

H Ei-i TS Ei-i i = DMSO

50 40

X

30

Y

20 10

I

0

II

III

-10

0.0

0.2

0.4

0.6

0.8

1.0

x i (i = DMSO) E E Fig. VI-14. The enthalpic and the entropic interactions, Hi−i and TSi−i , for i = dimethylsulfoxide E E (DMSO). Note that Hi−i > TSi−i . The net DMSO–DMSO interaction in terms of chemical potential is positive, Ei−i > 0, and is unfavorable (repulsive), or the net DMSO−H2 O interaction is favorable. The data are taken from Lai et al. (1995). Reproduced from the Journal of Solution Chemistry, 24, 89, (Lai, J. T. W. et al.), Copyright (1995), with kind permission from Springer Science and Business Media.

E E become apparent only in the DMSO–DMSO interactions, Hi−i and TSi−i for i = DMSO E shown in Fig. VI-14. We note in Fig. 13 that the values of HDMSO are zero and those of HWE are almost constant in the range, xDMSO > 06, similar to the BE-rich region xBE > 05 in BE−H2 O discussed in Section [V-1] (see Figs. V-3 and V-4 for comparison). Hence, in the DMSO-rich region DMSO molecules cluster together as in the pure liquid and H2 O molecules interact with the DMSO clusters as a single molecule. The xDMSO -dependence E of Hi−i (i = DMSO) (Fig. 14) provides the boundaries of mixing schemes. The loci of point X, Y and the transition point from Mixing Scheme II to III are listed in Table 2 E (p. 159). We note in Fig. 14 that the values of Hi−i (i = DMSO) are larger than those E for TSi−i . Namely, the enthalpic repulsion surpasses the entropic (times T ) attraction and the net interaction in terms of chemical potential, Ei−i (i = DMSO) is unfavorable (repulsion) between DMSO molecules; or, equivalently, the DMSO−H2 O interaction is favorable This behavior is similar to Gly−H2 O discussed above. Gly is regarded as a typical hydrophilic, and hence the Gly−H2 O interaction should be net favorable in E terms of chemical potential. DMSO, on the other hand, the xDMSO -dependence of Hi−i E shows some contribution from the hydrophobic moiety, and that of Vi clearly indicates the hydrophobic signature. Yet, the DMSO−H2 O interaction is attractive in terms of chemical potential. We recall that the enthalpy–entropy compensation is prevalent in aqueous solutions (Lumry & Rajender, 1970; Lumry, 2003). Hence some conspicuous

170

Fig. VI-15

E E behaviors in Hi−i and Si−i tend to compensate each other and the net result in Ei−i may E E . Furthermore, Ei−i is the second derivative no longer be conspicuous as in Hi−i or Si−i E E quantity (Table I-1 (p. 20)), while Hi−i and Si−i are the third derivative. If indeed the information contained in a higher order derivative is more detailed as argued in Section [I-8], then the detailed functions of hydrophobic and hydrophilic moieties within the same solute towards H2 O may have to wait until higher order derivatives than the third become available. We note in Column 4 of Table 2 (p.159) that Ei−i for 12P− or 13P−H2 O is zero. Hence the entropy–enthalpy compensation is complete in these solutions.

[VI-6] MIXING SCHEMES OF AQUEOUS NON-ELECTROLYTES STUDIED BY OTHER TECHNIQUES We briefly review here findings on the mixing schemes of aqueous non-electrolytes studied by techniques other than the present thermodynamic methodology. This will be in no way a comprehensive review; rather it will be selective. As pointed out in Section [V-6], there are not many studies in the literature which realize that, in aqueous solutions, the detail of the mixing scheme is crucially dependent on the composition. In Chapter V, we dealt mainly with hydrophobic effects on H2 O by way of studying a series of mono-ols, keeping the hydrophilic −OH the same and varying the size of the alkyl group. In this chapter, the solute non-electrolytes dealt with are amphiphiles and the effects of hydrophilic moieties gain progressively more important roles in modifying H2 O in the H2 O-rich region, as we go from types (a), (b), (c) to (d). We first briefly describe a theoretical study on the effect on H2 O of a typical hydrophilic solute: urea. We then take DMSO (type (c))−H2 O as an example and see the findings about mixing schemes by other techniques. Idrissi et al. (2000) conducted a molecular dynamics study on the short-time dynamics of urea−H2 O at a dilute mole fraction of 0.04 and at near saturation, 0.29. They focused on the directional dependent translation and rotational self-diffusion of both urea and H2 O species, as well as on velocity auto-correlation functions and corresponding power spectra. The results indicated that the addition of urea leads to an overall isotropy and stiffening of the short-time dynamics of both species. Urea lacks any hydrophobic moiety and can be regarded as a typical hydrophilic solute. Thus, these findings may be interpreted as urea forming hydrogen bonds to the existing hydrogen bond network of H2 O and retarding the degree of fluctuation inherent in liquid H2 O, by breaking the H donor/acceptor symmetry. This conclusion will receive some experimental support in the next chapter, when we use the 1-propanol (1P) probing method in ternary 1P–urea−H2 O. For DMSO−H2 O, the mixing scheme boundaries were found by our thermodynamic methodology to be xDMSO = 021 for point X, 0.35 for point Y and 0.56 for the II–III transition. Namely, Mixing Scheme I is operative from xDMSO = 0 to 0.21, II in the range 035 < xDMSO < 056, and III in 056 < xDMSO < 1. In the transition region, 021 < xDMSO < 035, the integrity of H2 O is gradually lost, and in xDMSO > 035, the solution consists of two kinds of clusters: One rich in H2 O and the other in DMSO.

171

1 0

4πr 2(ρ - ρ0) / Å–1

Mixing Scheme III 1 0

Mixing Scheme II

1 0

Mixing Scheme I

0

5

10

15

20

r/Å

Fig. VI-15. The radial distribution function for dimethylsulfoxide (DMSO)−H2 O at room temperature. The mole fractions of DMSO xDMSO for each curve from top down at about r = 56 Å are, 1.0, 0.80, 0.60 (Mixing Scheme III group), 0.50, 0.38, 0.28 (the second group), 0.28, 0.15, 0.07, and 0 (the last group). The data are taken from Koga et al. (2001). Reproduced from the Journal of Solution Chemistry, 30, 885, (Koga, Y. et al.), Copyright (2001), with kind permission from Springer Science and Business Media.

These clusters are reminiscent of Mixing Schemes I and III, respectively. Fig. VI-15 is the radial distribution function calculated from the wide angle x-ray scattering data from DMSO−H2 O at room temperature at mole fractions covering the entire composition range (Koga et al., 2001). The results are displayed in three groups of mole fraction, corresponding to Mixing Schemes I, II and III, respectively. The first intermolecular peak at about 5 Å in Fig. 15 could be regarded as basically the S−S and S−O interferences since the x-ray cross section of S is the largest. For pure liquid DMSO xDMSO = 1, the scattering data were consistent with the structure of solid DMSO with about 3% increase in the unit cell (Koga et al., 2001; Itoh & Ohtaki, 1987). Accordingly, the first intermolecular peak at 5.6 Å for xDMSO = 1 can be interpreted as the S−S and =O groups. This could be true the S−O distance of the two adjacent anti-parallel >S= also in the DMSO-rich region (Mixing Scheme III). Fig. VI-16 shows the loci of the first intermolecular peak in the radial distribution function against xDMSO . It is evident

172

Fig. VI-16

6.0

Locus of the First Peak / Ao

X

Y

5.5

5.0

III II 4.5

I 4.0 0.0

0.2

0.4

0.6

0.8

1.0

x DMSO

Fig. VI-16. The loci of the first intermolecular peak in the radial distribution function (Fig. 15), as a function of the mole fraction of dimethylsulfoxide (DMSO) in DMSO−H2 O at room temperature. The data are taken from Koga et al. (2001). Reproduced from the Journal of Solution Chemistry, 30, 885, (Koga, Y. et al.), Copyright (2001), with kind permission from Springer Science and Business Media.

that the 5.6 Å peak remains at the same locus as that of pure DMSO within Mixing Scheme III. This is consistent with our suggestion discussed above in Section [VI-5]: Namely, DMSO molecules form clusters with the same local structure as in pure DMSO with the S−S and the S−O distance being 5.6 Å. H2 O molecules do not form any structure to be conspicuous in the distribution function, i.e. H2 O molecules interact with clusters of DMSO as a single molecule. In the group of results for Mixing Scheme I in Fig. 15, the 2.9 Å peak for xDMSO = 0 is due to the hydrogen bonded O−O interference. Other characteristic O−O interferences of liquid H2 O appear at 4.5 and 7.1 Å. They are due to a strong correlation between the central O to the second and the third neighbor O atoms through the hydrogen bond network. On addition of DMSO, theses peaks characteristic to liquid H2 O diminish gradually. As xDMSO increases to 0.07 and further to 0.15, the 2.9 and 7.1 Å peaks decrease, while the 4.5 Å peak shifts towards a larger value and grows in height. This may be due to the involvement of S. But at xDMSO = 028 in the transition region between point X and Y, the value stops short at 5.3 Å rather than 5.6 Å. This may suggest that in Mixing Scheme I some DMSO may exist as a more tightly bound “smaller” cluster than that in Mixing Scheme III. This “smaller” cluster in the H2 O-rich region as opposed to the clusters in Mixing Scheme III may consist of a number of DMSO molecules held =O dipoles in an anti-parallel fashion. If so, the hydrophobic together mainly by >S= methyl groups point outward of the “smaller” cluster. This effect could be responsible

173

for a hydrophobic showing in ViE (i = DMSO) (Fig. 12). We shall come back to this point below in Section [IX-3]. Safford et al. (1969) studied DMSO−H2 O by means of neutron inelastic and x-ray scatterings. They concluded that for xDMSO < 02, DMSO causes a small cooperative orientation of H2 O around DMSO and also rigidifies the water structure. As discussed above for urea, a hydrophilic solute forms hydrogen bonds with the existing hydrogen bond network of H2 O and in effect rigidifies the fluctuation nature of H2 O. At the same time, the cooperative reorientation of surrounding H2 O, reminding “iceberg”, could be due to the hydrophobic effect of DMSO. Thus, the findings by Safford et al. in the region xDMSO < 02 are consistent with our Mixing Scheme I in the same mole fraction range. They found that at higher mole fractions the water structure is broken due to the formation of hydrogen bonded DMSO−H2 O complexes. The latter picture is consistent with our Mixing Scheme II and/or III. Indeed, Soper and Luzar (1992) did neutron scattering with H/D substitution for two mole fractions, 0.21 and 0.35., exactly at point X and Y (see Table 2 (p. 159)). They concluded that the percentage of H2 O molecules that are hydrogen bonded to each other is reduced, while an increasing fraction of H2 O molecules are hydrogen bonded to the lone pair electrons of DMSO. Later, Luzar and Chandler (1993) showed by molecular dynamics calculation that for DMSO−H2 O at xDMSO = 033, in the middle of the transition region from Mixing Scheme I to II hydrogen bonds between DMSO−H2 O are longer lived than H2 O−H2 O. The hydrogen bond lifetimes determined by “reactive flux correlation function approach” were about 5 and 3 ps for DMSO−H2 O and H2 O−H2 O respectively in the mixture, while that for H2 O−H2 O in pure H2 O is about 1 ps. These findings are again consistent with our suggestions on the hydrophilic effect of an amphiphile, DMSO. Finally, we note the observation by Areas et al. (1996) that the viscosity of 5 mM lysozyme−H2 O showed a dramatic increase when DMSO or tetramethyl urea (TMU) was added to the mole fraction xi = 043 for (i = DMSO), and xi = 021 for (i = TMU), respectively. This dramatic change is no doubt related to the transition of mixing scheme in DMSO−H2 O and TMU−H2 O. Indeed, the critical mole fractions are close to our xi loci of point Y; 0.35 for DMSO and 0.26 for TMU, listed in Table 2 (p.159). The mixed aqueous solvents of non-electrolytes are widely used as reaction or colloidal dispersion media, as those of alcohols discussed in Section [V-6]. There is no doubt that the distinct mixing schemes and their transitions in aqueous solutions must have devastating effects on the reactions or on the colloidal characteristics. A comprehensive review on this subject is awaited.

Chapter VII

EFFECTS OF NON-ELECTROLYTES ON THE MOLECULAR ORGANIZATION OF H2 O: 1-PROPANOL (1P) PROBING METHODOLOGY

[VII-1] INTRODUCTION – 1-PROPANOL (1P) PROBING METHODOLOGY As we saw in Chapters V and VI, the uniqueness of H2 O is retained in the H2 O-rich region of binary aqueous solutions of non-electrolytes (hydrophobes, hydrophiles, and amphiphiles): Namely, the hydrogen bond network of H2 O with a characteristically large degree of fluctuation is maintained against invasion of the solute to the extent that the hydrogen bond connectivity of the hydrogen bond network of H2 O remains connected throughout the entire bulk, up to the threshold composition. This threshold is crucially dependent on the nature of the solute in question. We call the mixing scheme operative in this region Mixing Scheme I. The solute–solute interaction in Mixing Scheme I is mediated by the bulk H2 O (that is modified somewhat by the solute), and of a long range. The effect of a solute on H2 O provides a modified field through which another solute molecule interacts. How H2 O protects its integrity while sacrificing some parts of the liquid H2 O characteristics depends also on the nature of the solute in question. In Chapter V, we studied the series of mono-ols, varying the size of alkyl groups and keeping −OH fixed in order to exemplify the nature of the hydrophobic moiety in aqueous solutions. We learned that in the H2 O-rich region, a hydrophobic solute enhances the hydrogen bond network of H2 O in its immediate vicinity (corresponding to the classical “iceberg formation”) with the concomitant reduction of the hydrogen bond probability of H2 O in the bulk H2 O away from the solute molecules. The solute–solute E interaction, Hi−i , takes type (a) dependence, as shown in Fig. VI-1, where i represents a hydrophobic solute. This dependence, particularly the initial increase up to point X in the figure, is attributed to the loss of the negative contribution to the fluctuation nature of the bulk H2 O due to the reduction in the hydrogen bond probability of bulk H2 O (see the discussion in Section [V-3]). Another signature for a hydrophobic solute is the initial decrease in the excess partial molar volume of i, ViE , or the negative volumetric E interaction, Vi−i , as discussed in Chapter V, Sections [V-1] to [V-3]. In Section [VI-3] we showed that glycerol(Gly) which is clearly more hydrophilic E than mono-ols affects H2 O quite differently. The enthalpic interaction, Hi−i (i = Gly), is of type (d) pattern in Fig. VI-1. We also saw a suggestion that the Gly–Gly enthalpic

176

Fig. VII-1

interaction is mediated by somewhat modified H2 O, from the same mole fraction dependence of the partial molar entropy-volume cross fluctuation, SV i (i = Gly) with E that of Hi−i (see Fig. VI-9). The excess partial molar volume of Gly, ViE (i = Gly) does not show an initial decrease, as shown in Fig. VI-7. While Gly has the propyl backbone that could work as a (weak) hydrophobic moiety, urea (U) lacks in a possible hydrophobic part and could work as a typical hydrophile. Unfortunately, the binary urea–H2 O system has not yet been studied using the second and third derivative quantities. Excess partial molar enthalpy data are available (Egan & Luff, 1966), but the data are not closely E spaced to warrant the graphical differentiation to obtain Hi−i (i = U). We point out, however, that the excess partial molar volume of urea could be calculated by accurate and closely spaced density data (Boeje & Hvidt, 1971) which show no sign of the initial E decrease as shown in Fig. VI-7. For other amphiphiles, the xi -dependence pattern of Hi−i indicates that both the hydrophobic and the hydrophilic parts modify H2 O individually E and the i–i interaction is operative, giving rise to the net effect on the Hi−i pattern. Hence it appears to be intermediate between type (a) and type (d). As we saw in Fig. VI-1, there are indeed types (b) and (c) which are intermediate of types (a) and (d). Namely, the bulk H2 O away from the solute is modified by the hydrophobic and the hydrophilic moieties in their respective characteristic manners. The bulk H2 O which is modified but still keeps its integrity, provides the field through which the enthalpic solute–solute interaction is operative. If this is indeed the case in the H2 O-rich region, then the heterogeneous interaction between the two different species in the multi-component system should also occur via the modified bulk H2 O, as long as the entire system is within Mixing Scheme I. This led us to the idea of the 1-propanol (1P) probing methodology of studying the effect of a sample solute (S) on the molecular organization of H2 O in the ternary 1P−S−H2 O system. We determine the excess partial molar enthalpy of 1P in E 1P−S−H2 O H1P , by perturbing the amount of 1P only. We then graphically differenE E E tiate H1P data to obtain H1P−1P . We know the x1P -dependence of H1P−1P for the binary 1P−H2 O without S as shown in Fig. VII-1. We recall that the partial molar normalized SV cross fluctuation of 1P, SV 1P , displays exactly the same dependence as shown in Fig. 1. This information was instrumental in suggesting that the 1P–1P interaction is operative via the bulk H2 O, which is modified yet still retains its integrity. Now, in the E presence of S, since S already modified H2 O in its own way, the H1P−1P will have a different pattern from that without S. This induced change is a reflection of the way S modifies the molecular organization of H2 O within Mixing Scheme I. Hence this induced change provides the information about the effect of S on H2 O in the H2 O-rich region. As E pattern reserves will become apparent below, in the limited H2 O-rich region, the H1P−1P the type (a) anomaly, which confirms that the system is still within Mixing Scheme I. We apply this idea for a variety of a sample solute S in this and the next chapters. First, we review the variable conversion discussed in Section [II-6] and learn how E H1P−1P can be evaluated in the 1P−S−H2 O ternary systems. For convenience, we use the p T xi i = 1 2     k − 1 N variable system in the k-component mixture, rather than the p T ni i = 1 2    k used in defining thermodynamic quantities.

Eq. (VII-1) to (VII-2)

177

250

X H1P-1PE (left scale)

H1P-1PE / kJ mol-1

200

SVΔ

150

1P

(arbitrary units)

100

50

0

-50 0.0

0.1

0.2

0.3

x 1P E Fig. VII-1. The 1P–1P interaction function, H1P−1P , in binary 1P−H2 O and the partial molar entropy-volume cross normalized fluctuation of 1P, SV 1P , as a function of x1P . The ordinate scale E is for H1P−1P , and in arbitrary units for SV 1P . The same graph as Fig. V-34(b), but the label of the ordinate is altered. Reproduced with permission from the Canadian Journal of Chemistry, 77, 2039, (Koga, Y.), Copyright (1999), NRC Research Press, National Research Council of Canada, Ottawa.

Here, N = defined as,

k 

ni . We recall the excess partial molar enthalpy of the i-th component is

i=1

 HiE ≡

H E

ni

 

(VII-1)

pTnj

where j = i. Thus no matter how many components are present, it is only the target component i that is perturbed in its amount, ni , in order to determine HiE . Having determined HiE by titrating the i-th component, the enthalpic i–i interaction is defined in the p T ni i = 1 2    k variable system as,  E Hi−i ≡N

HiE

ni

 

(VII-2)

pTnj

E where j = i. To calculate Hi−i , however, it is more convenient to do so in the p T xi i = 1 2     k − 1 N variable system. Consider now the ternary 1P−S−H2 O. With

178

p, and T kept constant, any thermodynamic quantity F is written in both variable systems as, Fn1P  nS  nW  = Fx1P  xS  N 

(VII-3)

For each variable system, the perfect differential is written as,  dF =

F

n1P



 dn1P + nS nW

F

nS



 dnS +

n1P nW

F

nW

 dnW 

(VII-4)

n1P nS

and  dF =

F

x1P





F dx1P +

xS xS N





F dxS +

N x1P N

 dN

(VII-5)

x1P xS

and the left-hand sides of eqs. (4) and (5) are equal. We are here interested in the case E F = H1P , which is an intensive quantity, and hence the last term on the right of eq. (5) E we keep nS and nW constant in the right of eq. (4), is identically zero. To obtain H1P−1P E and use the definition for H1P−1P , eq.(2) with i = 1P. Hence,  E H1P−1P ≡N

 =N

E

H1P

n1P E

H1P

x1P

 nS nW

  xS

x1P

n1P



 +N nS nW

E

H1P

xS



 x1P

xS

n1P

  nS n

(VII-6)  Since

x1P

n1P

 nS nW



1 − x1P  = N

and 

E H1P−1P = 1 − x1P 

E

H1P

x1P

xS

n1P





=− nS nW

 − xS xS

E

H1P

xS

xS  it follows that N

 

(VII-7)

x1P

E Thus, if H1P data are available in the x1P  xS  variable grids in small enough increments, E E . Often, however, we obtain the H1P data by titrating eq. (7) is used to determine H1P−1P − successively a small amount of 1P in the mixed solvent of S H2 O. In this case, the ratio rS = nS /nW remains constant as the amount of 1P, n1P , increases from zero. Hence instead of eq. (3) we have,

Fn1P  nS  nW  = Fx1P  rS  N

(VII-8)

Eq. (VII-3) to (VII-10)

179

and instead of eq. (5) for an intensive F ,    

F

F dF = dx1P + dr 

x1P rS

rS x1P S

(VII-9)

Since drS = 0, instead of eq. (7) we have, 

E H1P−1P

E

H1P = 1 − x1P 

x1P

 

(VII-10)

rS

This is the same as for a binary system, which is understandable since the system is a pseudo-binary system, S−H2 O being treated as the (mixed) solvent. [VII-2] EFFECTS OF METHANOL (ME), 2-PROPANOL (2P) AND E TERT-BUTANOL (TBA) ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE The characteristics of a hydrophobic moiety in aqueous solutions were elucidated by studying a series of mono-ols−H2 O in Chapter V. We recall the resulting hydrophobicity ranking of mono-ols (eq. (V-3)) in the H2 O-rich region. Accordingly, ME is definitely a weaker hydrophobe than the probing 1P, while 2P is a slightly weaker hydrophobe, if not almost equal. The hydrophobicity of TBA is stronger than 1P. Fig. VII-2, Fig. VII-3 E and Fig. VII-4 are the measured H1P data for S = ME, 2P, and TBA respectively in the ternary systems 1P−S−H2 O with various initial mole fractions of S, xS0 . As is evident E from the figures, the x1P -dependence of H1P is qualitatively different at high values of 0 0 xS . For example, for xS = 01458 and 0.4007 for S = ME (Fig. 2), the x1P -dependence of E H1P is different from those for xS0 = 005786, and 0.03291. We recall from the discussion in Section [V-2] that at xME = 007 Mixing Scheme I ends and starts to change into Mixing Scheme II for the binary ME−H2 O. Hence for xS0 > 007 (S = ME), the probing E 1P sees the system to be already passed Mixing Scheme I, and H1P shows a different x1P -dependence. Similar observations are made in Figs. 3 and 4 with the onset of the transition at x2P = 005 and xTBA = 0045, respectively (see Table V-3 (p. 114)). The E E details of x1P -dependence of H1P is naturally more conspicuous in H1P−1P . Figs. VII-5, E VII-6 and VII-7 show the H1P−1P plots for S = ME, 2P, and TBA, respectively. As is clear, it is within Mixing Scheme I for each binary S−H2 O that the x1P -dependence E pattern of H1P−1P displays the proper type (a) shown in Fig. VI-1(a). Fig. 6 shows the E effect of the almost equally hydrophobic, though slightly less, 2P on the H1P−1P pattern. 0 E As xS (S = 2P) increases within Mixing Scheme I for 2P−H2 O, the H1P−1P pattern seems to shift almost parallel to the left, to a smaller value of x1P , without changing E the height, the value of H1P−1P . If 1P itself were to be added as S in 1P−S−H2 O, the additional probing 1P starts off at the point where S (1P) had left, and follows exactly E the same route to points X and Y with the lesser amount of 1P. Hence the H1P−1P pattern

180

Fig. VII-2 to VII-5

4 2

H1PE / kJ mol-1

0 -2

xME0

-4

0.0000 0.03291 0.05786 0.1458 0.4007

-6 -8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-2. The excess partial molar enthalpy of 1P, H1P , in 1P−S−H2 O (S = ME). The data from Koga (2003b). Reproduced from the Journal of Solution Chemistry, 32, 803, (Koga, Y.), Copyright (2003), with kind permission from Springer Science and Business Media.

5

H1PE / kJ mol-1

0

-5

x2P0 = 0 0.01389 0.02832 0.05002 0.1260

-10

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VII-3. The excess partial molar enthalpy of 1P, H1P , in 1P−S−H2 O (S = 2P) at 25  C. The data from Hu et al. (2001). Reproduced with permission from the Bulletin of the Chemical Society of Japan, 74, 809, (Hu, J. et al.), Copyright (2001), the Chemical Society of Japan.

181

2

H1PE / kJ mol-1

0

-2

x TBA0 0

-4

0.01475 0.02974

-6

0.06017 -8

-10

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-4. The excess partial molar enthalpy of 1P, H1P , in 1P−S−H2 O (S = TBA). The data from Miki et al. (2005). Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

200

Error

X 150

H 1P-1PE / kJ mol-1

x ME0 = 0 0.03291 0.05786 0.1458

100

0.4007 50

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-5. The enthalpic 1P–1P interaction, H1P−1P , in 1P−S−H2 O (S = ME) at 25  C. The data are from Koga (2003b). Reproduced from the Journal of Solution Chemistry, 32, 803, (Koga, Y.), Copyright (2003), with kind permission from Springer Science and Business Media.

182

Fig. VII-6 to VII-7

250

Error

X

X

H 1P-1PE / kJ mol-1

200

x 2P0

150

0 0.01389

100

0.02832 50

0.05002

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-6. The 1P–1P enthalpic interaction, H1P−1P , in 1P−S−H2 O (S = 2P) at 25  C. The data from Hu et al. (2001). Reproduced with permission from the Bulletin of the Chemical Society of Japan, 74, 809, (Hu, J. et al.), Copyright (2001), the Chemical Society of Japan.

250

X Error

X

H 1P-1PE / kJ mol-1

200

150

xTBA0 0

100

0.01475 0.02974

50

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-7. The enthalpic 1P–1P interaction, H1P−1P , in 1P−S−H2 O (S = TBA) at 25  C. The data from Miki et al. (2005). Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

183

must shift parallel to the left, to the smaller value of x1P by the amount of S(1P) added initially. This is approximately what is observed in Fig. 6 for S = 2P. Indeed, 2P is almost equally hydrophobic as 1P (see Table V-3 (p. 114)). Thus, the change in the E H1P−1P pattern in Fig. 6 is summarized as follows: As xS0 increases, (1) the entire pattern E including point X shifts to the left, (2) the value of H1P−1P at the start, x1P = 0, increases, E and (3) the value of H1P−1P at point X remains almost the same. We note that the value E E of H1P−1P at x1P = 0, is a measure of how S modified H2 O in terms of H1P−1P . This corresponds also to the situation of bulk H2 O away from “icebergs” in terms of SV 1P . E pattern induced by addition of ME, a much Fig. 5 shows the changes in the H1P−1P weaker hydrophobe. It is apparent that (1) the left shift of point X, and (2) the increase E in H1P−1P at the start are qualitatively the same as for S = 2P. However, (3) the value of E H1P−1P generally decreases together with that at point X. The third effect is consistent with the interpretation that the 1P − 1P interaction is mediated via the bulk H2 O modified by the existing ME. Since the manner in which ME modifies is the weakest in the monool series (Table V-3 (p. 114)), the 1P − 1P interaction of the probing 1P experiences the E field of the bulk H2 O that was modified by ME. Hence the value of H1P−1P becomes smaller. By the same token, but in the opposite direction, for a stronger hydrophobe, E TBA, the value of H1P−1P becomes higher as xS0 (S = TBA) increases, of course within Mixing Scheme I as shown in Fig. 7. For S = TBA, the first two effects are qualitatively the same as for ME and 2P. Fig. VII-8 summarizes the quantitative effects of ME, 2P, and TBA on the value E of H1P−1P at point X and at the start, x1P = 0. Fig. 8(a) is an indicator of the relative hydrophobicity of S relative to the probing 1P. While the effect of 2P is small, almost null E within the estimated uncertainty, there is a systematic tendency of suppression of H1P−1P 0 as xS S = 2P increases. If this is real, it means that 2P is slightly less hydrophobic than 1P, consistent with the hydrophobicity ranking (eq. (V-3)), applicable in the H2 Orich environment. Thus, the slope in Fig. 8(a) ranks the relative hydrophobicity. The more positive the slope, the stronger its hydrophobicity than the probing 1P. Hence, we suggest from Fig. 8(a) that the hydrophobicity is ranked as TBA > 2P > ME, which is consistent with eq. (V-3). Another point to note in Fig. 8(a) is that the slope for ME is negative. As we will see in the next section, this behavior is typical of a hydrophilic S. This suggests that while ME is classified as hydrophobic, though the weakest in the mono-ols, it could be regarded as being more hydrophilic than the probing 1P. Hence, E the increase/decrease of H1P−1P at point X could be used as a convenient continuous scale for relative hydrophobicity/hydrophilicity. E The xS0 dependence of the value of H1P−1P at x1P = 0 (effect (2)), Fig. 8(b), is the consequence of the two effects; (1) the parallel shift to the left due to hydrophobicity E of S, and (3) the increase or decrease of H1P−1P at point X depending on the relative E hydrophobicity as compared with that of the probing 1P. For S = 2P, the value of H1P−1P at the start increases smoothly by the overall parallel shift to the left (1). For TBA, in E due to its addition to the effect (1), there is the general increase in the value of H1P−1P E stronger hydrophobicity. Consequently the value of H1P−1P at x1P = 0 increases faster, E as is evident in the figure. For ME, the behavior of the value of H1P−1P at the start is

184

Fig. VII-8(a) to VII-9

240

(a)

H 1P-1PE at point X

220

ME 2P 200

TBA

180

160

0.00

Error 0.02

0.04

0.06

0.08

x S0 (S = ME, 2P, or TBA) E E at point X in the H1P−1P pattern for 1P−S−H2 O at 25  C. The Fig. VII-8(a). The value of H1P−1P original data are from Koga (2003b) for ME, Hu et al. (2001) and Miki et al. (2005) for TBA.

250

(b)

H 1P-1PE at X1P = 0

200

150

100

ME 50

2P TBA

0 0.00

0.02

0.04

0.06

0.08

x S0 (S = ME, 2P, or TBA) E Fig. VII-8(b). The value of H1P−1P at the start, x1P = 0, for 1P−S−H2 O at 25  C. The original data are from Koga (2003b) for ME, Hu et al. (2001) for 2P, and Miki et al. (2005) for TBA.

185

complex because of the competition between the increase due to the left shift and the E general decrease in H1P−1P due to its lesser hydrophobicity (more hydrophilicity). Similar observations will be made for general amphiphiles studied in the following sections. We now examine the x1P -locus of point X, the left shift, more quantitatively. Fig. VII-9 shows the effect of xS0 (S = ME, 2P, or TBA) on the x1P -locus of point X. The vertical dotted line shows the upper limit of xS0 below which the present 1P-probing methodology is valid; namely, the binary S−H2 O system reaches the beginning of the transition from Mixing Scheme I to II. It is clear that the more hydrophobic S shifts the x1P locus of point X further. Here again, the ranking, TBA > 2P > ME, holds for hydrophobicity. Figs. VII-10 and VII-11 are the results of an attempt to find the origin of hydrophobicity within the present 1P-probing methodology. We focus on the x1P -locus of point X, since the shift of point X to the left seems to be purely due to hydrophobicity of S. A simplistic view is that the degree of hydrophobicity depends on the number of carbon atoms in the alkyl chain. Thus, in Fig. 10 the value of x1P at point X is plotted against 0 the carbon atom fraction, xcarbon (= xS0 × the number of carbon atoms). It is apparent in the figure that these plots seem to converge into a single curve within the estimated uncertainty. In turn, the degree of the left shift of point X will serve as a hydrophobicity scale within the present methodology. A recent examination indicated that the standard

0.08

ME Error

2P

x 1P at point X

0.06

TBA

0.04

ME 0.02

2P I TBA

0.00 0.00

0.02

0.04

0.06

0.08

x S0 (S = ME, 2P or TBA) E Fig. VII-9. The x1P - locus of point X of the H1P−1P pattern in 1P−S−H2 O at 25  C. The raw data are from Koga (2003b) for ME, Hu et al. (2001) for 2P, and Miki et al. (2005) for TBA. The vertical dotted line shows the upper limit of xS0 below which Mixing Scheme I is operative in binary S−H2 O for S = ME, 2P or TBA, i.e. below which the present 1P-probing methodology is valid.

186

Fig. VII-10 to VII-12

0.06

ME

x 1P at point X

2P TBA 0.04

0.02

Error

0.00 0.00

0.05

0.10

0.15

0.20

x carbon0 E Fig. VII-10. The x1P locus of point X in the H1P−1P pattern in 1P−S−H2 O (S = ME, 2P or TBA)  at 25 C. The abscissa is the mole fraction of carbon atom. See text.

ME

0.06

2P

x 1P at point X

TBA

0.04

0.02

Error

0.00 0.0

0.1

0.2

0.3

0.4

0.5

x (C-H bond)0 E Fig. VII-11. The x1P locus of point X in the H1P−1P pattern in 1P−S−H2 O (S = ME, 2P, or TBA) at 25  C. The abscissa is the mole fraction of C−H bonds. See text

187

free energy of transfer from H2 O to hexadecane of alkanes, alkenes, alkadienes and arenes scale nicely with the number of C–H bonds (Kyte, 2003). This argument uses the zeroth derivative of G, while ours uses the x1P -dependence of the third derivative E 0 . Nonetheless, Fig. 11 shows the same plots vs. xC−Hbonds . These plots quantity, H1P−1P also seem to converge to a single curve. We discuss this issue further below.

[VII-3] EFFECTS OF UREA (U), TETRAMETHYL UREA (TMU) AND E ACETONE (AC) ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE =O and the −NH2 work as hydrophilic Urea (U) lacks in alkyl group, while the >C= =O as a proton accepcenters. U could readily form hydrogen bonds to H2 O with >C= − tor and NH2 a donor. It is generally known that the O · · · H−O∗ type hydrogen bond is stronger than the N–H · · · O∗ , where O∗ is the oxygen of H2 O. The H· · · O distance is shorter for the former case than that of the latter (Huheey, 1978). Thus, =O is lopsided in the strength of the proton donor −NH2  and the acceptor >C= U. Nonetheless, urea could be regarded as a typical hydrophilic solute. Indeed the excess partial molar volume of U in the binary U−H2 O (Fig. VII-12) shows no sign 2

0

ViE / cm3 mol-1

-2

-4

-6

U TMU

-8

AC -10 0.0

0.1

0.2

0.3

0.4

xi (i = U, TMU, or AC)

Fig. VII-12. The excess partial molar volume of i, ViE (i = U, TMU, or AC) in the binary i−H2 O at 25  C. Calculated using the density data by Boeje & Hvidt (1971) for U, Phillips et al. (1974) and Jakli & Van Hook (1996) for TMU, and Boeje & Hvidt (1971) for AC.

188

Fig. VII-13 to VII-14

of the initial decrease typical of a hydrophobic solute. Also shown in Fig. 12 are the excess partial molar volume of tetramethyl urea (TMU) and that of acetone(AC). We will discuss below how the effect of S changes, when the −NH2 of urea is converted to −NCH3 2 and further to −CH3 on the molecular organization of H2 O. Since U is a solid at room temperature, the excess partial molar enthalpy of U, HUE , has not been determined directly, although it is possible to titrate aqueous urea instead and convert the data to HUE , as discussed in Appendix E. The heats of solution of urea into H2 O have been determined in the literature (Egan & Luff, 1966) but their data are not so closely spaced in composition as to determine HUE with confidence, let alone HUE−U . Hence, we have no information as to the location of point X, the onset of transition from Mixing Scheme I to II, if any, in U−H2 O. The present 1Pprobing methodology, on the other hand, is convenient in that 1P is readily titrated in the binary U−H2 O mixed solvent to learn about the effect of U on H2 O in E in the ternary 1P−U−H2 O (To et al., the H2 O-rich region. Fig. VII-13 shows H1P 0 1998). It is immediately clear that the xS (S = U)-dependence is qualitatively differE data ent from those for S = ME, 2P, and TBA shown in Figs. 2 to 4. The H1P−1P E were re-evaluated using the original H1P data; the results are shown in Fig. VII-14. E pattern reserves type (a) dependence up to As is evident in the figure, the H1P−1P 0 xS S = U = 02, confirming that Mixing Scheme I is still operative. The figure indicates

0

H 1PE / kJ mol-1

-2

xU0 = 0

-4

0.03040 0.04985

-6

0.1020 0.1501

-8

0.2035 -10

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E , in 1P−U−H2 O at 25  C. The data are Fig. VII-13. The excess partial molar enthalpy of 1P, H1P from To et al. (1998). Reproduced with permission from the Journal of Physical Chemistry B, 102, 10958, (To, E. C. H. et al.), Copyright (1998), American Chemical Society.

189

250

xU0 = 0 Error

X

0.03040 0.04985 0.1020 0.1501 0.2035

H 1P-1PE / kJ mol-1

200

150

100

50

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VII-14. The 1P–1P enthalpic interaction, H1P−1P , in 1P−U−H2 O at 25  C. Revaluated using E the original H1P data by To et al. (1998).

E that the x1P -locus of point X remains the same, while the value of H1P−1P is suppressed 0 as xS (S=U) increases up to 0.2. The solubility in the binary U−H2 O at 25  C is 0.266 in the mole fraction of U, beyond which is the macroscopic manifestation of Mixing Scheme III. Unfortunately, we have no information at this stage regarding the transition from Mixing Scheme I to II, if any. From the observation from Fig. 14, however, it is likely that the point X, the onset of the transition in binary U−H2 O, could be in the range 02 < xU < 027. Nevertheless, it is clear that U does not change the x1P locus of point X. Namely, it requires the same amount of 1P to drive the system to point X, the onset of the loss of the percolation nature of the hydrogen bond network of H2 O. We suggest, therefore, that U does not alter the connectivity of the hydrogen bond network of H2 O. This is probably due to urea’s capability of forming hydrogen bonds with the existing network of H2 O. E , or equivalently that of SV 1P , by the presence The suppression of the value of H1P−1P of U shown in Fig. 14, could be attributed to the lesser decrease of the negative contribution to SV 1P , or to the direct loss of the positive part of SV 1P . The former case is, as discussed above, due to the loss of ice-like patches in the bulk H2 O, and was E is due to the hydrophobic effects of the solute in question. If the reduction in H1P−1P due to a weaker hydrophobe as ME discussed above, then the x1P -locus must also be shifted to the left. Since it remains the same, as is evident in Fig. 14, we suggest that the

190

Fig. VII-15 to VII-16

E effect of U on the value of H1P−1P , or that of SV 1P is due purely to its hydrophilicity and to reducing the positive part of SV 1P . Namely, the degree of fluctuation inherent in liquid H2 O is reduced by the presence of U. This may occur by U participating in hydrogen bonding with the existing network of H2 O, and thus keeping the hydrogen bond connectivity intact. The intervening U in the network, however, breaks the existing proton donor/acceptor symmetry of H2 O. As a result, the degree of fluctuation, the positive part of SV-cross fluctuation inherent in H2 O is reduced, while the hydrogen bond percolation is retained. Indeed, Idrissi et al. (2000) concluded that the addition of U leads to an overall isotropy and stiffening of the short-time dynamics of both species, as mentioned above. This finding is consistent with our interpretation based on Fig. 14. The capability of forming hydrogen bonds of U is no doubt due to the −NH2 =O group for H-acceptor. Therefore, the group working as the H donor and the >C= H donor/acceptor symmetry which is the very source of the unique nature of H2 O is severely retarded by the presence of U. Not only may the geometry of the hydrogen bond network, but also the strength of the H-donor/acceptor, be affected. To see how these effects manifest themselves within the present methodology, we use tetramethyl urea (TMU) and acetone (AC) for S in the present 1P−S−H2 O methodology. TMU replaces all H’s in −NH2 with −CH3 ’s and AC changes the two −NH2 ’s to −CH3 ’s. Figs. VII-15

250

xTMU0 = 0 0.01245 0.02502

H 1P1PE / kJ mol-1

200

0.05006

150

0.07507 0.1003

100

50

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P E , for 1P−TMU−H2 O at 25  C. The data are Fig. VII-15. The 1P–1P enthalpic interaction, H1P−1P from Chen et al. (2001). Reproduced from the Fluid Phase Equilibria, 189, 31, (Chen, D. H. C. et al.), Copyright (2001), with permission from Elsevier.

191

250

xAC0 = 0

X

0.03205

200

0.06103 H 1P1PE / kJ mol-1

0.09006 150

0.1604 0.2005

100

50

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P E , in 1P−AC−H2 O at 25  C. The data are Fig. VII-16. The 1P–1P enthalpic interaction, H1P−1P from Chen et al. (2001). Reproduced from the Fluid Phase Equilibria, 189, 31, (Chen, D. H. C. et al.), Copyright (2001), with permission from Elsevier.

E and VII-16 show the H1P−1P data for 1P−TMU−H2 O and for 1P−AC−H2 O, respectively (Chen et al., 2001). As expected, the effect of the hydrophobic moieties is now apparent in that point X is shifted to the left for both TMU and AC. We show the x1P E and the H1P−1P loci of point X in Fig. VII-17, together with those for U. Fig. 17(a) shows the effect of hydrophobic effects on the x1P -locus of point X by TMU and AC. Clearly, the effect is stronger for TMU with four −CH3 groups than that for AC with two. The effect of the number of −CH3 is not twice, however. This may be due to the difference between N and C, to which −CH3 is bonded. Fig. 12 shows the excess partial molar volume signature of hydrophobicity for TMU and AC. The initial decrease in ViE indicates that the former is a substantially stronger hydrophobe than the latter. U, on the other hand, does not show the initial decrease in VUE (Fig. 12) and does not alter the x1P -locus (Fig. 17(a)). E Fig. 17(b) indicates that the effects of U, TMU, and AC on the value of H1P−1P SV and hence those of 1P are surprisingly similar, considering the uncertainty. Perhaps, the hydrophilic effect of TMU is stronger than U and AC. This observation suggests =O, and the contribution from that the hydrophilic effect is predominantly due to >C= −NH2 in U towards hydrophilicity is marginal. As discussed above, it is known that the O · · · H−O∗ type hydrogen bond is stronger than the N−H · · · O∗ type, where O∗ is the oxygen of H2 O (Huheey, 1978).

192

Fig. VII-17(a) to VII-18

0.08

Error

(a)

x 1P at point X

0.06

0.04

Urea TMU

0.02

AC

0.00 0.00

0.05

0.10

0.15

0.20

0.25

0.30

xs0 (S = U, TMU, or AC) E Fig. VII-17(a). The x1P locus of point X in the H1P−1P pattern for 1P−S−H2 O (S = U, TMU, or AC) at 25  C. The data arew from To et al. (1998) for U, and from Chen et al. (2001) for TMU and AC.

250

Error H 1P1PE / kJ mol-1 at point X

U 200

TMU AC

(b) 150

100

50

0.00

0.05

0.10

0.15

0.20

0.25

xs0 (S = U, TMU or AC) E E Fig. VII-17(b). The value of H1P−1P at point X in the H1P−1P pattern of 1P−S−H2 O (S = U, TMU, AC) at 25  C. The data are from To et al. (1998) for U, and Chen et al. (2001) for TMU and AC.

193

[VII-4] EFFECTS OF ETHYLENE GLYCOL (EG), 1,2- AND 1,3-PROPANEDIOL (12P AND 13P), GLYCEROL (GLY) AND POLY (ETHYLENE GLYCOL) (PEG) ON H2 O E PATTERN CHANGE AS PROBED BY THE H1P−1P In this section, we study the effect of poly-ols on H2 O by the same 1P-probing methodology. Poly-ols are clearly more hydrophilic than the probing 1P. We will learn the E effect of −OH on the H1P−1P pattern. We will also learn, as shown below, how the alkyl E backbone manifests its hydrophobicity in the H1P−1P pattern. Poly(ethylene glycols) (PEG) are known to have a strong affinity to H2 O, i.e. hydrophilic. This effect is often attributed to a good match of the inter-oxygen distance in liquid H2 O and that between the neighboring ether oxygen atoms in the PEG helicoidal chain (Giordano et al., 1995). We wish to see the difference, if any, between the end −OH and the ether −O− on the respective hydrophilic effect. E Figs. VII-18, VII-19, VII-20 and VII-21 are the H1P−1P plots for S = EG(2C/2OH), 13P (3C/2OH), 12P(3C/2OH), and Gly(3C/3OH) in the ternary 1P−S−H2 O. The number of the C atoms in the alkyl chain and that of the hydroxyl −OH groups are shown in brackets for convenience. As is evident from the figures, the induced changes of the E H1P−1P pattern by the presence of S are in both ways in terms of point X, one shifting E point X to the left or reducing the x1P -locus, and the other lowering the value of H1P−1P . 200

xEG0 = 0

H 1P-1PE / kJ mol-1

150

0.02033 0.03984

X

0.08089

100

0.1687 0.5238

50

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VII-18. The 1P–1P enthalpic interaction, H1P−1P , for 1P−EG−H2 O at 25  C. The data are from Koga (2003b). Reproduced from the Journal of Solution Chemistry, 32, 803, (Koga, Y.), Copyright (2003), with kind permission from Springer Science and Business Media.

194

Fig. VII-19 to VII-21

H 1P-1PE / kJ mol-1

200

X

x13P0 = 0 0.05008 0.1007

150

100

50

0

0.00

0.05

0.10

0.15

0.20

x1P E Fig. VII-19. The 1P–1P enthalpic interaction, H1P−1P , in 1P–1,3-propanediol (13P)−H2 O at 25  C. The data are from Parsons & Koga (2002). Reproduced with permission from the Journal of Physical Chemistry B, 106, 7090, (Parsons, M. T. and Koga, Y.), Copyright (2002), American Chemical Society.

200

X

x12P0 = 0

H 1P-1PE / kJ mol-1

0.02513 150

0.05030 0.07529 0.09998

100

50

0

0.00

0.05

0.10

0.15

0.20

x1P E Fig. VII-20. The 1P–1P enthalpic interaction, H1P−1P , in 1P–1,2-propanediol (12P)−H2 O at 25  C. The data are from Parsons and Koga (2002). Reproduced with permission from the Journal of Physical Chemistry B, 106, 7090, (Parsons, M. T. and Koga, Y.), Copyright (2002), American Chemical Society.

195

H 1P-1PE / kJ mol-1

200

X

xGly0 = 0

150

0.02493 0.05249 100

0.07541 0.10125

50

0 0.00

0.05

0.10

0.15

0.20

x1P E Fig. VII-21. The 1P–1P enthalpic interaction function, H1P−1P , in 1P–glycerol (Gly)−H2 O at  25 C. The data are from Parsons et al. (2001). Reproduced from the Journal of Solution Chemistry, 30, 1007, (Parsons, M. T. et al.), Copyright (2001), with kind permission from Springer Science and Business Media.

E The increase/decrease of the value of H1P−1P at the start, x1P = 0, is complex due to the similar competition discussed for ME above. Namely, the increase due to the parallel shift to the left and the decrease due to hydrophilicity compete, and the net results seem complex as shown in Figs. 18 to 21. We focus for the remainder of this section on the behavior of point X only. From the discussion in the above two sections, the left shift is the signature of hydrophobicity and the decrease in the height of the peak is that of hydrophilicity of S, relative to the probing 1P. Indeed, the latter effect seems stronger for Gly (3C/3OH) (Fig.21) than EG(2C/2OH), 13P(3C-2OH), or 12P(3C/2OH) (Figs. 18–20). This point will be clearer when we show the locus of point X as a function of xS0 below. E plots for S = PEG’s. Fig. 22 is for PEG with Figs. VII-22 and VII-23 are the H1P−1P the average molecular weight of 203, abbreviated as PEG-2, which contains 8.4 carbon and 5.2 oxygen atoms, including the ether −O−. The PEG used for Fig. 23 is of the average molecular weight of 600 (PEG-6) and the number of carbon atoms is 26.4 and that for oxygen is 14.2. Of course, the two end oxygen atoms in both PEG are in the hydroxyl form −OH. The figures show that xS0 < 003 for S = PEG-2 (Fig. 22) and for E pattern remains as that of type (a). This suggests that xS0 < 001 for PEG-6, the H1P−1P within these H2 O-rich regions the integrity of H2 O is retained and Mixing Scheme I is operative in the respective binary PEG−H2 O system. Similar to poly-ols (Figs. 18–21), PEG’s also cause point X to shift to the left as well as downward.

196

Fig. VII-22 to VII-24

250

error

xPEG-20 = 0 0.008088 0.01588 0.02299 0.04179 0.07099 0.1205

X

H 1PE / kJ mol-1

200

150

X

100

50

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VII-22. The 1P–1P enthalpic interaction, H1P−1P , in 1P–poly(ethylene glycol) 200 (PEG 2)−H2 O at 25 C. The data are from Miki et al. (2005). Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

250

H 1P-1PE / kJ mol-1

200

X

xPEG-60 = 0

error

0.004595 0.009511 0.01649 0.04247

150

X 100

50

0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VII-23. The 1P–1P enthalpic interaction, H1P−1P , in 1P–poly(ethylene glycol) 600 (PEG6)−H2 O at 25  C. The data are from Miki et al. (2005). Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

197

240

PEG-2

Error

220

H 1P1PE at point X / kJ mol-1

PEG-6 200

EG Gly

180

13P ME

160

2P TBA

140

12P 120 100 80 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

xs0 (S=PEG-2,PEG-6,EG,Gly,12P, 13P, ME, 2P, or TBA) E E Fig. VII-24. The value of H1P−1P at point X of the H1P−1P pattern for 1P−S−H2 O at 25  C. S = PEG-2, PEG-6, EG, Gly, 12P, 13P, ME, 2P, or TBA. The abscissa is the initial mole fraction of S, xS0 . See text. Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

E Fig. VII-24 summarizes the relative hydrophilicity, in terms of the value of H1P−1P 0 at point X, of various mono-ols, poly-ols including PEG. The PEG-6 requires xS (S = E PEG-6) ≈ 0.015 to reduce the value of H1P−1P to about one half of the original, while 0 Gly needs xS (S = Gly) ≈ 0.07, about 5-fold, to do the same. The ratio of the oxygen in each molecule is 142/3 ≈ 5. This rather suggests that it is the number of oxygen that is E responsible for suppression of H1P−1P regardless of whether O is the end OH or the ether E −O−. Indeed, the most effective H1P−1P reducer apparent in Fig. 24 is PEG-6, having the largest number of O in the solutes studied. We note that ME, 2P, 12P and TBA contain −CH3 group. For those, hollow symbols are used for convenience in Fig. 24 and the subsequent three figures. Fig. VII-25 shows the plots of the x1P -locus against xS0 . As we showed in Fig. 9, the more hydrophobic the sample solute S, the more negative the slope. Similarly, from Fig. 25 a general trend may be detected. Again. PEG-6 having the largest number of C seems to be most effective for the left shift, i.e. the most hydrophobic. Thus, the number of oxygen and carbon atoms could be an important factor characterizing hydrophobicity and hydrophilicity of the present set of amphiphiles. E Fig. VII-26 shows the plots of the value of H1P−1P against the initial mole fraction 0 0 of oxygen atom, xoxygen (= xS × the number of O in S). Except for 2P and TBA, which together with ME and possibly 12P may be operating by the hydrophobic mechanism,

198

Fig. VII-25 to VII-27

0.06

PEG-2 PEG-6 EG Gly 13P

x1P at point X

Error

0.04

0.02

ME 2P TBA 12P

0.00 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16

xs0 (S=PEG-2,PEG-6,EG,Gly, 12P, 13P, ME, 2P, or TBA) E Fig. VII-25. The x1P locus of point X in the H1P−1P pattern for 1P−S−H2 O at 25  C. S = PEG-2, PEG-6, EG, Gly, 12P, 13P, ME, 2P, or TBA. The abscissa is the initial mole fraction of S, xS0 . See text. Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

PEG-2 PEG-6 EG Gly 13P

H 1P1PE at point X / kJ mol-1

Error 200

ME 2P TBA 12P

150

100

0.0

0.1

0.2

0.3

0.4

xOxygen0 E Fig. VII-26. The value of H1P−1P at point X in 1P−S−H2 O at 25  C. S = PEG-2, PEG-6, EG, Gly, 13P, ME, 2P, TBA, or 12 P. The abscissa is the initial mole fraction of oxygen atom in S. See text. Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

199

all seem to converge into a single curve within the estimated uncertainty. Namely, the hydrophilicity, which is caused by the hydroxyl −OH and the ether −O−, scales with the number of oxygen and the difference in the effect between the end −OH and the ether −O− is not significant within the present methodology. For ME (1C(methyl)/1OH) and 12 P (3C with 1methyl/2OH), the balance of the hydrophobic and the hydrophilic E 0 moieties is such that the H1P−1P vs. xoxygen falls on almost the same curve. Fig. VII-27 is the hydrophobic propensity in terms of the x1P -locus of point X against 0 xcarbon (= xS0 × number of carbon in S). The plots seem to converge into two curves, one group without methyl group and the other with methyl groups. A recent analysis of the standard free energies of transfer from water to hexadecane of alkanes, alkenes, alkadienes, and arenes scales well with the number of C−H bonds (Kyte, 2003). On converting the abscissa of Fig. 27 to the mole fraction of the number of C−H bonds, however, the two curves did not converge. Tanford (1973) used the logarithm of the partition coefficient of a solute in H2 O and in an organic solvent and showed a linear relationship with the number of carbon atoms in the solute molecule. However, for a relative ranking covering from hydrophobicity to hydrophilicity of amphiphiles and ions (which will be dealt with in Chapter VIII), the choice of the organic solvent makes a devastating effect on the results (Marcus et al., 1988; Taylor et al. 1991). Furthermore, the logarithm of the partition function is equal to the excess chemical potential difference 0.06

PEG-2 Error

PEG-6 EG Gly 13P

x1P at point X

0.04

0.02

ME 2P TBA 12P 0.00 0.0

0.1

0.2

0.3

0.4

xCarbon0

Fig. VII-27. The x1P locus of point X in 1P−S−H2 O at 25  C. S = PEG-2, PEG-6, EG. Gly, 13P, ME, 2P, TBA, or 12P. The abscissa is the initial mole fraction of carbon atom in S. See text. Reproduced with permission from the Journal of Physical Chemistry B, 109, 3873, (Miki, K. et al.), Copyright (2005), American Chemical Society.

200

Fig. VII-28 to VII-29

of the solute in H2 O and in the chosen organic solvent. We use here the behavior of the third derivative of G to monitor hydrophobicity and hydrophilicity, as opposed to the first derivative, chemical potential. Hence, the more subtle difference could be detected by the present method. Indeed, the present method could distinguish the methyl group containing solutes from those without.

[VII-5] CONCLUDING REMARKS – SUMMARY The present 1P-probing methodology is a powerful tool for sorting out hydrophobicity and/or hydrophilicity of a non-electrolyte. For a hydrophobic solute, its effect on the E H1P−1P pattern could be summarized in Fig. VII-28. The original pattern for the binary 1P−H2 O [O] in the figure shifts parallel to the left, in the presence of an equally E hydrophobic S. Thus, the value of H1P−1P at point X remains the same, the x1P locus E decreases, and the value of H1P−1P at the start, x1P = 0, increases steadily as xS0 increases. For a weaker (stronger) hydrophobe, in addition to the shift to the left, the value of

[O] Binary [A] Equal hydrophobe

X

[B] Weaker Hydrophobe

1P)

[C] Stronger Hydrphobe

H1P-1PE (or

SVΔ

[O]

[A]

[C]

[B]

0 0

x1P E Fig. VII-28. The induced changes in the H1P−1P pattern by the presence of a hydrophobic S, in E − − 1P S H2 O. [O] is the H1P−1P pattern for the binary 1P−H2 O; [A] shows that with an equally hydrophobic S as the probing 1P; [B] is that with a weaker, and [C] that with a stronger hydrophobe. Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

201

E H1P−1P decreases (increases) in the entire range of x1P within Mixing Scheme I. As a E at point X decreases (increases), and the rate of increase result, the value of H1P−1P E of H1P−1P at x1P = 0 decelerates (accelerates). The mechanism for these effects on the E pattern comes from the way a hydrophobic S affects the molecular organization H1P−1P of H2 O within Mixing Scheme I. Namely, a hydrophobic S enhances the hydrogen bond network of H2 O in its immediate vicinity (“iceberg” formation) and concomitantly reduces the hydrogen bond probability of bulk H2 O away from S. This in turn reduces the chance of putative formation of ice-like patches. As a result, the negative contribution to the SV cross fluctuation decreases. The effect of 1P on the rate of this decrease increases E . At the as x1P increases, resulting in the initial increase in SV 1P and hence in H1P−1P same time, the probing 1P−1P interaction occurs via the bulk H2 O away from S, which is modified by S to its own strength and is influenced by the mixed ways in which S and 1P modify the bulk H2 O. Hence, if S is a weaker(stronger) hydrophobe than 1P, the E is suppressed (enhanced). The effect of a hydrophobe is thus mainly on value of H1P−1P the retardation of “ice-like patches” in the bulk H2 O away from “icebergs”, reflecting the “mixture model” of H2 O. E pattern. The Fig. VII-29 is a sketch of the effect of a hydrophile on the H1P−1P E x1P locus remains unchanged but the value of H1P−1P decreases by the presence of a

[O] Binary [D] Hydrophile

1P)

X

[E] Stronger Hydrphile

H1P-1PE (or

SVΔ

[O]

[E]

[D]

0 0

x1P E pattern by the presence of a hydrophilic S, in Fig. VII-29. The induced changes in the H1P−1P 1P−S−H2 O. Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

202

Fig. VII-30

hydrophile. A hydrophile forms hydrogen bonds to the hydrogen bond network of H2 O, and thus the connectivity is retained, but the H donor/acceptor symmetry of H2 O is broken, not only in the geometry but also in the hydrogen bond strength. As a result, the hydrogen bond network becomes more rigid and loses the degree of fluctuation inherent in liquid H2 O. Namely, the regular positive part of SV cross fluctuation decreases in the presence of a hydrophilic S. Thus, its effect is more readily understood in terms of the “bent hydrogen bond” model of H2 O. E pattern are some combination of the For amphiphiles, their effects on the H1P−1P above two cases, depicted in Fig. 28 and Fig. 29. Fig. VII-30 sketches the resulting E by an amphiphile. A typical case is seen in Fig. 20 for S = 12P. change in H1P−1P For poly-ols including poly(ethylene glycols), the hydrophobic effect on the x1P locus of point X seems to scale with the number of carbon in the alkyl chain. For −CH3 containing solutes, however, their hydrophobic effect is found to be stronger. The E seems to scale with the number of oxygen, hydrophilic effect on the value of H1P−1P regardless of the hydroxyl −OH, or the ether −O−. The effects of other hydrophilic

[O] Binary [F] Amphiphile

SVΔ ) 1P

X

H1P-1PE (or

[O]

[F] 0 0

x1P E Fig. VII-30. The induced changes in the H1P−1P pattern by the presence of an amphiphile S, in E 1P−S−H2 O. The hydrophobic part shifts point X to the left and raises the value of H1P−1P at E the start, x1P = 0. The hydrophilic part, on the other hand, lowers the value of H1P−1P in the entire range including at point X. The net result is a combination of these two effects. Reproduced with permission from the Netsusokutei (Journal of Japan Society of Calorimetry and Thermal Analysis), 34, 3 (Koga, Y.), Copyright (2007), the Japan Society of Calorimetry and Thermal Analysis.

203

moieties than −OH or −O− are yet to be studied. Preliminary indication shows that the =O appears stronger than the proton donor −NH2 within the present proton acceptor >C= methodology. As demonstrated, this 1P-probing methodology is powerful and convenient. It would therefore find wide application in aqueous solutions of biological importance. Biopolymers are large molecules and almost colloidal. Their surfaces are covered with hydrophobic and hydrophilic moieties and each part would interact with H2 O differently. Hence, we suggest that the same methodology using another probing species, one more hydrophilic than the present 1P, is mandatory for further studies.

Chapter VIII

THE EFFECTS OF SALTS ON THE MOLECULAR ORGANIZATION OF H2 O: 1-PROPANOL (1P)-PROBING METHODOLOGY

[VIII-1] INTRODUCTION – HOFMEISTER SERIES In this chapter, we examine the effects of various salts (and the constituent ions) on the molecular organization of H2 O using the same 1P-probing methodology successfully used in Chapter VII. We showed in Chapter VII that the relative hydrophobicity/hydrophilicity of a sample amphiphile (S) could be sorted out and ranked in terms E of the induced changes in the H1P−1P pattern in the ternary 1P−S−H2 O system. By this method, we could distinguish the effect of the hydrophobic part and that of the hydrophilic part within the same molecule of an amphiphile. Here we apply the same methodology for S = salts. What electrolytes do to H2 O is an important and interesting subject in its own right. However, the relative influences of anions and cations on the properties and functions of biopolymers have been compared and ranked as the so-called Hofmeister series initiated in 1887 by the Hofmeister group (Kunz et al., 2004). The original ranking (Hofmeister, 1887) showed the effectiveness of the ion in precipitating the hen egg white from aqueous solution. Since then almost the same ranking applies, particularly for anions, to ion effects on a wide range of processes in aqueous solutions; not only biochemical processes but also colloidal sciences. The generality of this ion ranking has suggested that the effects of ions on H2 O are important. So modified mixed solvents of aqueous electrolyte in turn influence the behaviors of the examined processes. This belief triggered a massive interest in basic studies of binary aqueous solutions of various salts. This approach has provided important advances through analysis of various factors including surface tension increments, lyotropic number, viscosity B-coefficients, zetapotentials, salt activities, etc. However, not all of these physical properties carry the hallmarks of the Hofmeister effects. Thus, the molecular level understanding is only fragmentary (Kunz et al., 2004; Collins & Washabaugh, 1985; Cacace et al., 1997). There is another issue regarding the generality of the effects of ions on H2 O over the specificity of the examined processes. While the effects of anions are generally stronger and the specificity may not obscure the ranking, cations tend to exert weaker effects on the molecular organization of H2 O. Hence, the cation ranking in the literature is mixed. Table VIII-1 lists a number of the Hofmeister ranking in randomly selected reviews and

206

Fig. VIII-1

Table VIII-1. Hofmeister series. #

← Kosmotropic, salting out, stabilizing

Destabilizing, salting in, chaotropic →

Anions a

SO4 2− >HPO4 2− >citrate>tartrate>CrO4 − ≈ Cl− >NO3 − >ClO3 −

b

citrate>tartrate>SO4 2− >acetate>Cl− >NO3 − >Br − >I− >SCN−

c

SO4 2− ≈ HPO4 2− >F− >Cl− >Br − >I− ≈ ClO4 − >SCN−

d

SO4 2− >HPO4 2− >acetate>citrate>tartrate>Cl− >NO3 − >ClO3 − >I− >ClO4 − >SCN−

e

F− >PO4 3− >SO4 2− >acetate>Cl− >Br − >I− >SCN−

f

citrate>SO4 2− >HPO4 2− >F− > acetate>Cl− >Br − >I− >NO3 − >ClO4 − >SCN− Cations

a

Na+ >K+

b

Th4 >Al3+ >H+ >Ba2+ >Sr 2+ >Ca2+ >K+ >Na+ >NH4 +

d

NH4 + >K+ >Na+ >Li+ >Mg2+ >Ca2+ >guadinium

e

+ + 2+ + 2+ 2+ + CH3 4 N+ >CH3 2 NH2 + > NH+ 4 >K >Na >Cs >Li >Mg >Ca >Ba

f

+ Mg2+ >Ca2+ >H+ >Na+ >K+ >Rb+ >Cs+ >NH+ 4 >CH3 4 N



a, a :(Hofmeister, 1887). b, b :(Marshall, 1978). c:(Collins & Washabaugh, 1985). d, d :(Creighton, 1993). e, e :(Cacace et al., 1997). f, f  :(Lopez-Leon et al.,2003).

textbooks. As is evident from the table, the anion ranking is more consistent than that for cations. Even among anions, the relative position between SO4 2− and F− varies. In scheme # c and # f, SO4 2− >F− , while in #e the opposite ranking F− >SO4 2− is shown. For cations, the rankings of # b and of # d are almost completely opposite, so are those of # e and # f . This discrepancy is no doubt due to the fact that the specificity of the examined process is more important when the difference in the ion effect becomes subtle. Indeed, Collins and Washabaugh (1985) did not list the cation ranking although (or because) the reference covered in the review is massive. In spite of these discrepancies, the generality of the Hofmeister ranking suggests that an important factor is how each salt modifies the molecular organization of H2 O. As we pointed out above, and will be doing so throughout the book, the mixing scheme in aqueous solutions is crucially dependent of the composition. It is in the H2 O-rich region only where the integrity of liquid H2 O is maintained. In this chapter we study what the sample salt (S) does to H2 O by the 1P-probing E methodology. Namely, we determine H1P−1P in the ternary 1P−S−H2 O, and monitor E first that the H1P−1P pattern remains of type (a) (Fig. VI-1(a)). This ensures that our attention is limited within Mixing Scheme I, in which the integrity of H2 O is maintained E even in the presence of a given salt. We then examine the induced changes in the H1P−1P pattern (still in type (a)) to learn about the effect of the salt on H2 O.

207

E [VIII-2] EFFECTS OF NaF AND NaCl ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE E Figs. VIII-1 and VIII-2 show the H1P data for 1P−S−H2 O with S = NaF and NaCl, respectively (Westh et al., 2006). The arrows in the figure indicate approximate points where the thermal response after titration slows down considerably due to liquid–liquid phase separation. We do not include for further analysis the series with the phase E separation prematurely before the inflection point in the H1P plots against x1P . The E corresponding H1P−1P plots are shown in Figs. VIII-3 and VIII-4. As is evident in the figures, for xS0 S = NaF<00043 and xS0 S = NaCl<0035, the x1P -dependence of E H1P−1P retains the type (a) form. Thus, within these concentration ranges, the aqueous solutions are in the Mixing Scheme I region. In Fig. 4, our earlier data (Matsuo et al., 1999) are also included. What is striking in Figs. 3 and 4 is that on addition of S (S = NaF or NaCl), the starting point at x1P = 0 remains fixed while the locus of point X E shifts to the left, to smaller values of x1P , without changing the value of H1P−1P , and hence those of SV 1P . This immediately suggests that the availability of H2 O for 1P to interact is reduced as xS0 increases. Hence, it requires a lesser x1P to drive the system to point X. We suggest that this is due to hydration of the salt. From the observation that E the values of H1P−1P remains the same at the start as well as at point X, we conclude that the molecular organization of bulk H2 O away from the hydration shell remains the

2 0

H 1PE / kJ mol-1

-2 -4

xNaF0 = 0 0.002194 0.004347 0.009081 0.01368

-6 -8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-1. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P−NaF−H2 O at 25  C. The arrows in the figure indicate the point at which the liquid–liquid phase separation occurs. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

208

Fig. VIII-2 to VIII-4

2 0

H 1PE / kJ mol-1

-2 -4

xNaCl0 = 0 0.01798 0.02425 0.03808 0.04715

-6 -8 -10

-12 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16

x1P E Fig. VIII-2. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P−NaCl−H2 O at 25  C. The arrows indicate the point at which the liquid–liquid phase separation occurs. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

250

xNaF0 = 0 X

H 1P-1PE / kJ mol-1

200

0.002194 0.004347

X

150

100

Error 50

0 0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P E Fig. VIII-3. The enthalpic 1-propanol (1P)–1P interaction, H1P−1P , in 1P−NaF−H2 O at 25  C. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

209

250

X

X

H 1P-1PE / kJ mol-1

200

xNaCl0 = 0 0.01798 0.02425

150

100

Error

50

0 0.01092 0.01962 0.03503

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P E Fig. VIII-4. The enthalpic 1-propanol (1P)–1P interaction, H1P−1P , in 1P−NaCl−H2 O at 25  C. The data from Westh et al. (2006). The data for hollow symbols are taken from Matsuo et al. (1999). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

same as that in pure H2 O. The recent femto second pump-probe spectroscopic studies could distinguish the dynamics of H2 O in the hydration shell from that of bulk H2 O outside of the hydration shell. They showed that the former in the hydration shell for Na+ salts of Cl−  Br − and I− is much slower than the latter (Kropman & Bakker, 2001; Bakker et al., 2004), providing a direct support for the existence of the hydration shell. These studies also suggested that the dynamics of bulk H2 O away from the hydration shell is not affected at all by the presence of the above halide ions as well as SO4 2− and ClO4 − (Omta et al., 2003). A similar conclusion was drawn by a neutron diffraction study with the empirical potential refinement analysis (Soper & Wechstrom, 2006). It was suggested that the first hydration shells of F−  Cl−  Br − and I− together with the counter cation, K + , are strongly perturbed, but the bulk H2 O outside of the hydration shell is only mildly affected. These findings regarding the Na+ salt of F− and Cl− are completely consistent with our conclusion. However, as will be shown below, for the Na+ salts of all the other anions (i.e. Br −  I−  SO4 2− , and ClO4 − ), our findings regarding their effects on the bulk H2 O are qualitatively different. We discuss the discrepancies between our findings and those by these techniques below in Section [VIII-7]. Figs. VIII-5 and VIII-6 show the loci of point X, which is the onset of the transition from Mixing Scheme I to II, together with the phase boundary of 1P−S−H2 O (S = NaF and NaCl). The detection of the onset of the phase separation could be facilitated by a novel method within the margin of a single drop of 1P (Kato et al., 2004). The figures indicate that the loci of point X seem to form a straight line; namely, the shift

210

Fig. VIII-5 to VIII-6

0.020

S+L

X NaF0

0.015

0.010

0.005

II

I 0.000 0.00

0.02

0.04

0.06

0.08

0.10

x1P

Fig. VIII-5. The mixing scheme and phase diagram of 1-propanol 1P−NaF−H2 O at 25  C. The filled line is the phase boundary, and the dotted line is the mixing scheme boundary, indicating the onset of the transition from Mixing Scheme I to II. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society. 0.12

S+L

S+L+L

0.10

X NaCl0

0.08

L+L 0.06

0.04

0.02

I

0.00 0.00

0.02

II

0.04

0.06

0.08

0.10

0.12

x1P

Fig. VIII-6. The mixing scheme and phase diagram of 1-propanol 1P−NaCl−H2 O at 25  C. The solid lines are the phase boundaries and the dotted line is the mixing scheme boundary indicating the onset of the transition from Mixing Scheme I to II. The data from Westh et al. (2006). The data points with hollow symbols are taken from Matuso et al. (1999). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

211

Table VIII-2. The hydration numbers of salts, nH , determined by the 1P-probing methodology. The 1P-probing methodology

The values of nH in the literature

Salt

Ion

nH (ref)

xS0

%

Na2 SO4

both#

27 ± 3 (a)

<0015

NaF NaCl

both both

19 ± 1.5(c) 7.5 ± 0.6(c)

<0004 <0035

3.9 ± 0.5

NaBr NaI

Na+ Na+

5.2 ± 0.5(c) 5.6 ± 0.5(c)

<0.06 <007

3.9 ± 0.5 3.9 ± 0.5

CaCl2 NH4 Cl CH3 4 NCl

both both Cl−

11 ± 1.6(f) 3.3 ± 0.4(f) 2.1 ± 0.4(f)

<0023 <0062 <008

12 ± 2 1.8 ± 0.5

range

nH (Zavitsas)&

nH (Dielectric relaxation, no kinetic depolarization) (ref.) 27 at xS0 = 0 decreasing to 20 at 0.016 (b) 13 ± 1 (d) 8.6 to xS0 = 009 decr. to 6.7 ± 0.2at 0.035 (e)

xS0 is the initial mole fraction of salt. Zavitsas (2001). # Plus an extra hydrophobe-like effect. See text. (a) Koga et al. (2004b). (b) Buchner et al. (1999a). (c) Westh et al. (2006). (d) Buchner et al. (1994). (e) Buchner et al. (1999b). (f) Koga et al. (2006). % &

of point X to the left is linear to xS0 . If this is due to hydration, as we suggest, then its slope provides an estimate of the hydration number. Alternatively, the xS0 intercept of the straight-line Mixing Scheme boundary gives the mole fraction of S such that all H2 O molecules would be used up for hydration. Thus, the xS0 intercept is equal to 1/1 + nH , where nH is the hydration number within the context of the present 1Pprobing methodology. Thus, the effects of NaF and NaCl on H2 O are that they hydrate 19 ± 15 and 75 ± 06 of H2 O respectively, and that their presence does not affect the properties of the bulk H2 O away from the hydration shells. These values are listed in Table VIII-2.

E [VIII-3] EFFECTS OF NaBr AND NaI ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE E Figs. VIII-7 and VIII-8 show the H1P plots for 1-propanol 1P−S−H2 O at 25  C for S = NaBr and NaI. There is an apparent trend of decrease in the initial slope as xS0 E increases, unlike NaF and NaCl (Figs. 1 and 2). The corresponding H1P−1P plots are 0 shown in Figs. VIII-9 and VIII-10. While point X shifts to the left as xS (S = NaBr or E NaI) increases, the values of H1P−1P at point X and at the start, x1P = 0 decrease. These E induced changes in the H1P−1P pattern, particularly for S = NaI (Fig. 10), resembles the effect of glycerol (Gly) shown in Fig. VII-21. Thus the effects of NaBr and NaI on

212

Fig. VIII-7 to VIII-10

2 0

H 1PE / kJ mol-1

-2

xNaBr0 = 0

-4

0.01017 0.01945 0.02874 0.03972 0.04820 0.05982 0.07021 0.08580

-6 -8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-7. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P−NaBr −H2 O at  25 C. The arrows indicate the points where the liquid–liquid phase separation occurs. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

2 0

H 1PE / kJ mol-1

-2

xNal0 = 0

-4

0.008203 0.01889 0.02912 0.04422 0.05138 0.06335 0.07293 0.09086

-6 -8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-8. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P−NaI−H2 O at 25  C. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

213

250

xNaBr0 = 0 0.01017 0.01945 0.02874 0.03972 0.04820 0.05982

X

200

H 1P1PE / kJ mol-1

X

150

100

50

0

Error -50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-9. The enthalpic interaction, H1P−1P , in 1-propanol(1P)−NaBr −H2 O at 25  C. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

250

xNaI0 = 0 0.008203 0.01889 0.02912 0.04422 0.05138 0.06335 0.07293

X

200

H 1P1PE / kJ mol-1

150

100

X

50

0

-50

Error -100 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-10. The enthalpic interaction, H1P−1P , in 1-propanol (1P)−NaI−H2 O at 25  C. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

214

Fig. VIII-11 to VIII-12

0.16

S+L+? 0.14 0.12

L+L

xNaBr0

0.10 0.08 0.06 0.04

II

I 0.02 0.00 0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P

Fig. VIII-11. The mixing scheme and phase diagram for 1-propanol (1P)−NaBr −H2 O at 25  C. The solid line shows the phase boundary, the broken line a presumed phase boundary, and the dotted line shows the onset of transition from Mixing Scheme I to II. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

H2 O have a hydrophilic propensity. For S = Gly, the shifts of point X to the left was attributed to the hydrophobicity of the alkyl backbone, as discussed in Section [VII-4]. For the present cases, S = NaBr or NaI, is the shift of point X to the left due to the hydration to Na+ ? Fig. VIII-11 and Fig. VIII-12 show the plots of the x1P -loci of point X vs. xS0 , representing the shift to the left by whatever mechanism. They appear to form a straight line, and its intercept provides nH = 52 ± 06 for NaBr, and 56 ± 06 for NaI. This suggests that Na+ hydrates about 5 H2 O, and therefore Br − and I− work strictly as hydrophiles, as urea (Fig. VII-14). White et al. (2000) used first-principles molecular dynamics simulations which avoid empirically determined parameters or a restriction to pair-wise interaction potentials and studied the Na+ −H2 O system. The results showed richer structural properties than the classical simulations, and indicated that a Na+ hydrates 5.2 molecules of H2 O and leaves the orientational dynamics of the bulk H2 O away from the hydration shell unperturbed (White et al., 2000). More recently, Varma & Rempe (2006) also used ab initio molecular dynamics simulations and suggested that the hydration number of Na+ is 5. Experimentally, another recent study (Ansell et al., 2006) by the ‘difference method’ of neutron diffraction with isotopic substitutions and the anomalous x-ray diffraction concluded that the hydration number of Na+ is also 5. Thus it appears to be safe to use the hydration number of Na+ to be

Eq. (VIII-1)

215

0.20

S+?

xNal0

0.15

II 0.10

0.05

I

0.00 0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P

Fig. VIII-12. The mixing and phase diagram for 1-propanol (1P)−NaI−H2 O at 25  C. The broken line shows a presumed phase boundary and the dotted line the onset of the transition from Mixing Scheme I to II. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

5.2 from our NaBr results, and evaluate that for counter anion accordingly. Thus, the hydration number, nH , for F− is 14 ± 2 and 23 ± 06 for Cl− . These values are listed in Table VIII-3 (p. 216). E Fig. VIII-13 shows the values of H1P−1P at point X for NaF, NaCl, NaBr, and NaI. For NaF and NaCl, the value stays constant, having the hallmarks of the hydration center E as discussed in the previous section. For NaBr and NaI, however, the value of H1P−1P 0 decreases as xS (S = NaBr, or NaI) increases, just as a typical hydrophile, urea (see Fig.VII-17(b)). Since the effect of Na+ is to shift point X to the left, the decrease in E H1P−1P is due to a hydrophilicity of counter anions, Br − and I− . Namely, as discussed in Chapter VII for hydrophiles, Br − and I− form hydrogen bond directly to the existing hydrogen bond network of H2 O and retard the degree of fluctuation inherent to liquid H2 O. Fig. 13 suggests that I− is a stronger hydrophile than Br − . It is known that halide ions are capable of forming hydrogen bonds with H2 O as proton acceptors (Schulte & Horning, 1961;Huheey, 1978; Jeffrey;1997), and their strength decreases in the order of F− >Cl− >Br − >I− 

(VIII-1)

216

Fig. VIII-13

Table VIII-3. The hydration number, nH , of ions by the 1P-probing methodology. Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics. Ion

nH

ref.

Na+

Westh et al. (2006)

SO4 2− F− Cl−

5.2 (reference) 17 ± 3 14 ± 2 2.3 ± 0.6

Koga et al. (2004b) Westh et al. (2006) Koga et al. (2004b)

Ca2+ NH4 +

6±2 1±1

Koga et al. (2006) Koga et al. (2006) Parsons & Koga (2005) Parsons & Koga (2005) Parsons & Koga (2005)

−

OOC–CH2 –NH3 + GL+− 

7.0 ± 0.6

−

OOC–CH2 –NH2 GL− 

4.8 ± 1

HOOC–CH2 –NH3 + GL+ #

2.7 ± 1

#

There is an extra hydrophilic effect in addition to hydration. See text.

H 1P-1PE at point X / kJ mol-1

250

200

150

100

NaF NaCl NaBr NaI

Error

50

0.00

0.02

0.04

0.06

0.08

0.10

xsalt0 E Fig. VIII-13. The values of H1P−1P at point X in 1-propanol (1P)−Na-halides−H2 O at 25  C. For E NaF and NaCl, the value of H1P−1P remains constant, both Na+ and F− or Cl− being a hydration center. For NaBr and NaI, each anion works as hydrophiles. The hydrophilicity is stronger for I− than for Br − , the slope being more negative. The data from Westh et al. (2006). Reproduced with permission from the Journal of Physical Chemistry A, 110, 2072, (Westh, P. et al.), Copyright (2006), American Chemical Society.

217

We suggest that F− and Cl− , being stronger proton acceptors, strip a number of H2 O molecules from the existing hydrogen bond network of H2 O and form the hydration shells, which do not extend their influences over to the bulk H2 O. The latter effect (non-effect) contrasts with what hydrophobes do to H2 O. As discussed at some length in Chapters V and VII, hydrophobes form “icebergs” and influence concomitantly the bulk H2 O away from “icebergs” in that the hydrogen bond probability of the bulk H2 O is reduced progressively to the hydrogen bond percolation threshold. Clearly F− is a stronger hydration center than Cl− , and nH is 14 ± 2 for F− and 23 ± 06 for Cl− . Thus, in terms of the original Hofmeister’s “H2 O withdrawing power” (Hofmeister, 1887), the Hofmeister ranking F− >Cl− is consistent. Br − and I− , on the other hand, are weaker proton acceptors than F− and Cl− . As a result, they form hydrogen bonds to the existing network of H2 O, rather than stripping any H2 O molecules out of the network. By so doing, the H2 O network now contains impurity centers that break the H donor/acceptor symmetry inherent in H2 O. Hence the hydrogen E and SV 1P decreases. This effect is bond network of H2 O becomes more rigid and H1P−1P − − stronger for I than Br , as shown in Fig. 13. In terms of the Hofmeister series, (Table 1, p. 206), I− is more chaotropic than Br − . This suggests that chaotropes are hydrophiles in the context of the present 1P-probing methodology. Thus, the effects of a series of halide ions on H2 O are not governed by a single parameter (or mechanism). Kosmotropes work as the hydration centers without influencing the H2 O characteristics of the bulk H2 O away from the hydration shells. By hydration, they limit the availability of H2 O to the examining system to interact. A stronger kosmotrope, F− , hydrates more H2 O molecules than Cl− . A recent molecular dynamics simulation study (Chowdhuri & Chandra, 2006) concluded that the residence time of an H2 O molecule directly hydrogen bonded to a halide ion is longer than that in the bulk H2 O in the decreasing order of F− >Cl− >Br − >I− ≈ H2 O. This trend is not inconsistent with our findings, but our 1P probing methodology found a qualitative difference between the F−  Cl−  and the Br −  I−  pairs. An oxygen K-edge x-ray absorption spectroscopy (Cappa et al., 2005) on aqueous solutions of Na-salts of Cl−  Br − , and I− observed general red shifts of the H2 O transitions as the salt concentration increases. They were attributed to a strong, direct perturbation of the unoccupied molecular orbital of the adjacent H2 O by the presence of salts. The magnitude of this effect followed the trend I− ≈ Br − >Cl− . They concluded that the observed red shifts were not due to the distortion of the hydrogen bond network of H2 O beyond the hydration shell. As mentioned above, the femto second pump-probe spectroscopies (Kropman & Bakker, 2001; Bakker et al., 2004; Omta et al., 2003) and the neutron diffraction study (Soper & Weckstrom, 2006) concluded that all of F−  Cl−  Br −  I−  SO4 2− , and ClO4 − leave the bulk H2 O unaffected. Our findings are that Br − and I− do affect the fluctuation nature of the bulk H2 O. Both SO4 2− and ClO4 − do alter the properties of bulk H2 O in different ways (Koga et al., 2004c), as discussed below. In Section [VIII-7] we comment on the insensitivity of these spectroscopic studies in distinguishing the effects of ions on the bulk H2 O away from solutes

218

Fig. VIII-14 to VIII-16

[VIII-4] EFFECTS OF Na2 SO4  NaOOCCH3 (NaOAc), NaClO4 , AND NaSCN E ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE In this section we examine the effects of the sodium salts of SO4 2−  OOCCH3 − OAc−  ClO4 − , and SCN− by the same 1P-probing methodology. These anions together with halides studied in the previous section span the Hofmeister series from the kosmotropic to the chaotropic ends. These anions are almost unequivocally ranked as, SO4 2− >F− >OAc− >Cl− >Br − >I− >ClO4 − >SCN− 

(VIII-2)

except that the relative order between SO4 − and F− is mixed (see Table 1, # c and # f vs. # e). E Figs. VIII-14, VIII-15, VIII-16 and VIII-17 show the H1P−1P pattern in 1P−S−H2 O  at 25 C with S = Na2 SO4 , NaOAc, NaClO4 , and NaSCN. Na2 SO4 (Fig. 14) shows the E value at the start, x1P = 0 hallmark of a hydration center in two respects: The H1P−1P remains constant and the x1P -locus of point X shifts to the left, to smaller values of x1P , E as xS0 S = Na2 SO4  increases. Unlike NaF or NaCl, however, the value of H1P−1P at point X increases. If the effects of cations and anions are additive, as we are taking as E a premise, then that of Na+ would not increase the value of H1P−1P at point X. Thus, 2− E the observed increase in H1P−1P at point X is due to SO4 . We recall that tert-butanol 300

xNa2SO40 = 0

X 250

0.004840

H 1P1PE / kJ mol-1

0.01032

X

200

0.01488 150

100

50

0

0.00

I Error 0.02

II 0.04

0.06

0.08

0.10

0.12

0.14

x1P E , in 1P−Na2 SO4 −H2 O at Fig. VIII-14. The 1-propanol(1P)−1P enthalpic interaction, H1P−1P 25  C. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

Eq. (VIII-2)

219

250

X

H 1P-1PE / kJ mol-1

xNaOAc0 = 0

X

200

0.01302 0.02715 0.04176

150

100

50

I 0

0.00

Error

II

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-15. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P−NaOAc−H2 O at  25 C. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

250

xNaClO40 = 0 200

Error

0.01077

X

H 1P-1PE / kJ mol-1

0.01989 0.03138

150

0.04771 0.09014

100

X 50

I 0

II 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-16. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P−NaClO4 −H2 O at  25 C. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

220

Fig. VIII-17 to VIII-18

200

X

xNaSCN0 = 0

Error

0.01094

150

H 1P-1PE / kJ mol-1

0.03094 0.04890 100

0.1002 X

50

0

I II

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-17. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P−NaSCN−H2 O at 25  C. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

E (TBA), a stronger hydrophobe than the probing 1P, makes the value of H1P−1P increase E both at the start and at point X (Fig. VII-7). Since the value of H1P−1P remains the same at the start, SO4 2− is not a typical hydrophobe. Thus, the detail of the effect of SO4 2− is not clear at present. It could be that SO4 2− works as a hydration center in the E absence of 1P in the system, resulting in the unchanged value of H1P−1P at x1P = 0. As x1P increases, however, a direct interaction between SO4 2− and 1P starts to operate, in addition to the H2 O-mediated 1P−SO4 2− interaction. If so, then the Hofmeister ranking depends crucially on the test biopolymer, and thus the ranking could become unstable for SO4 2− . In this context also, a series of similar studies using a probe other than 1P is awaited. We plan to do this, using glycerol (Gly) as a probe and follow the induced E changes in the HGly−Gly pattern. Further studies on Na2 SO4 such as a dielectric relaxation spectroscopy on aqueous MgSO4 (Buchner et al, 2004) are awaited. Nonetheless, from E the fact that the H1P−1P values at x1P = 0 remains constant, we may classify SO4 2− as a hydration center. Thus, from the mixing scheme boundary (Fig. VIII-18), the hydration number, nH , is estimated as 27 ± 3 for Na2 SO4 . It follows using the reference nH = 52 for Na+  nH for SO4 2− is 17 ± 3. They are listed in Tables 2 (p. 211) and 3 (p. 216). Fig. 15 indicates that NaOAc works basically as a hydrophobic species with a strength equal to the probing 1P, as shown in Fig. VII-6 for 2-propanol (2P). However, there E is a tendency of slight suppression of the increase in the value of H1P−1P at x1P = 0 than that shown for the effect of a typical hydrophobe as type [A] in Fig. VII-28.

221

0.05

S+L S+L+L

xNa2SO40

0.04

0.03

L+L 0.02

0.01

I II 0.00 0.00

0.02

0.04

0.06

0.08

0.10

0.12

x1P

Fig. VIII-18. The mixing scheme and phase diagram for 1-propanol (1P)−Na2 SO4 −H2 O at 25  C. The solid line is the phase boundary. The broken line is a presumed phase boundary. The dotted line is the mixing scheme boundary, the onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

The latter suppression must be due to the hydration of Na+ , which tends to keep the E H1P−1P value constant at x1P = 0. Namely, the hydrophobic behavior apparent in Fig. 15 is due to OAc− . Thus, the observed left shift of point X is due to a combination of hydration by Na+ and hydrophobicity by OAc− . Clearly the −CH3 group is the main cause of this. Why then an expected hydrophilicity from −COO− is not apparent is yet to be elucidated. We speculate that there could be some ion pairing operative such that the hydrophilic moiety is consumed by this pairing leaving the hydrophobic moieties pointing outward. Fig. VIII-19 shows the loci of point X in 1P−NaOAc−H2 O, the onset of transition E from Mixing Scheme I to II as probed by H1P−1P , together with the phase boundary. The thick dotted line in the figure indicates the loci of point X if the shift to the left of point X is due only to hydration by Na+ with nH Na+  = 52. The actual plots with the thin dotted line lie below this due to the hydrophobicity of OAc− . Fig. 16 S = NaClO4  and Fig. 17 (S = NaSCN) show that both effects are similar to that by glycerol (Gly) shown in Fig. VII-21. Namely, in addition to some tendency E of the left shift, there is a clear decrease in the value of H1P−1P at point X. Thus, both salts give a hydrophilic propensity. The left shift, on the other hand, could be due to the hydration by Na+ or some hydrophobicity inherent in each anion. This point is

222

Fig. VIII-19 to VIII-21

0.12

S+L 0.10

xNaOAc0

0.08

L+L

0.06

0.04

0.02

0.00 0.00

II

I

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P

Fig. VIII-19. The mixing scheme and phase diagram for 1-propanol (1P)−NaOAc−H2 O at 25  C. The solid line is the phase boundary. The broken line is a presumed phase boundary. The thick dotted line is the expected loci of point X if the shift to the left of point X is due only to hydration of Na+ . The thin dotted line is the observed onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

discussed below. Judging from the figure, however, the hydrophilic effects are more − conspicuous for both ClO− 4 and SCN . Fig. VIII-20 shows the mixing scheme boundary for 1P−S−H2 O (S = NaClO4 or NaSCN). The thick dotted line corresponds to the case where the left-shift is due only to hydration by Na+ . The actual loci of point X seem to lie below this, suggesting that both ClO4 − and SCN− contain some hydrophobic propensity, or at least some mechanisms to reduce the hydrogen bond probability of the bulk H2 O away from the solute. However, this point requires further clarification E including repeats of H1P measurements. E Fig. VIII-21 is a summary of the effects of salts on the value of H1P−1P at point X including halide salts. Those showing the signature for a hydration center (including Na2 SO4 ) and a hydrophobe are indicated by hollow symbols in the figure. The hydrophiles are shown by filled symbols. Within these hydrophiles, the relative effect of the anion is represented by the slope in the figure, since the counter cation is fixed at Na+ . Accordingly, the slope becomes more negative in the order of, Br −
(VIII-3)

Eq. (VIII-3)

223

0.30

xS0 (S = NaClO4 or NaSCN)

S+L 0.25

0.20

0.15

II

0.10

NaClO4 NaSCN

0.05

I 0.00 0.00

0.02

0.04

0.06

0.08

0.10

x1P

Fig. VIII-20. The mixing scheme and phase diagram for 1-propanol (1P)−NaClO4 or NaSCN−H2 O at 25  C. The broken line is a presumed phase boundary. The thick dotted line is the expected loci of point X if the shift to the left of point X is due only to hydration by Na+ . The thin dotted line is the observed onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2004b). Reproduced with permission from the Journal of Physical Chemistry A, 108, 8533, (Koga, Y. et al.), Copyright (2004), American Chemical Society.

H 1P-1PE at point X / kJ mol-1

300

Na2SO4 NaF NaOAc NaCl NaBr NaI NaClO4

250

200

150

100

50 0.00

NaSCN

Error

0.02

0.04

0.06

0.08

0.10

xsalt0 E Fig. VIII-21. The values of the 1-propanol (1P)–1P enthalpic interaction, H1P−1P , at point X  on addition of salt (S) in 1P−S−H2 O at 25 C. S = Na2 SO4, NaF, NaOAc, NaCl, NaBr, NaI, NaClO4 , or NaSCN. Na+ is a hydration center, with nH = 52. OAc− (hollow square symbol) is a hydrophobe. The remaining anions with hollow symbols are all hydration centers. the rest of the E anions with filled symbols are all hydrophiles. A stronger hydrophile lowers the values of H1P−1P more. The data from Koga et al. (2004b & 2004c) and Westh et al. (2006).

224

Fig. VIII-22

This ranking resembles the chaotropic half of the Hofmeister series (see Table 1, p. 206). In Section [VIII-7] below, we discuss the effects of anions and cations on the molecular organization of H2 O elucidated by the present 1P-probing method in relation to the Hofmeister ranking E In closing this section, we summarize the signature of a hydration center on the H1P−1P E pattern changes as sketched in Fig. VIII-22. The point to note is that the values of H1P−1P both at the start, x1P = 0, and at point X remain constant, while the x1P -locus of point X shifts to the left, to smaller values of x1P as xS0 increases (S is a hydration center). From these observations, we suggest that the hydration center hydrates a number of H2 O molecules and leaves the bulk H2 O outside of the hydration shell unperturbed, as we discussed above repeatedly. We have still the unsettled case of SO4 2− . As sketched in Fig. 22, SO4 − carries two hallmarks out of three required to be “the hydration center” by the present definition within the 1P-probing methodology. The fact that the value of E H1P−1P at x1P = 0 remains the same indicates that the bulk H2 O away from the hydration shell remains the same as pure H2 O at the start, x1P = 0. As x1P increases, the rate of E increase is more rapid as if the third component is a stronger hydrophobe as H1P−1P tert-butanol (TBA) as shown in Fig. VII-7 or the case [C] in Fig. VII- 28. Since the E value of H1P−1P at x1P = 0 remains the same, SO4 − cannot be a hydrophobe. Thus, we tentatively suggest that in the presence of 1P, SO4 − , while hydrating a large number

[0] Binary X [A] Hydration Center H 1P-1PE or (SVΔ1P)

[B] Na2SO4 [0]

[A]

[B]

0 0

x1P E Fig. VIII-22. The H1P−1P pattern changes for a typical hydration center (HC) in 1P−HC−H2 O. See text for discussion of the Na2 SO4 case. Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

225

17 ± 3 of H2 O molecules, begins to modify the bulk H2 O away from the hydration shell just as a hydrophobe does to H2 O. We recall that the number of H2 O molecules in the “iceberg” surrounding a hydrophobic alcohol is typically of the order of 20 or more, while the hydration number of SO4 − is in the same range. This may suggest that the hydration number may be important for influencing the bulk H2 O away from the hydration shell or “icebergs”. Even so, how the presence of 1P encourages SO4 − to do the same as a hydrophobe does to the bulk H2 O is yet to be elucidated.

[VIII-5] EFFECTS OF CaCl2  NH4 Cl AND TETRAMETHYL AMMONIUM E CHLORIDE (TMACL) ON H2 O AS PROBED BY THE H1P−1P PATTERN CHANGE In this section, we choose, as the sample salt (S), the chloride salt of Ca2+  NH4 + or CH3 4 N+ TMA+  in 1P−S−H2 O. By fixing the counter anion at Cl− , we seek to E learn the effects of these cations on the H1P−1P pattern, and in turn, their effects on the molecular organization of H2 O. As we showed above, NaCl is a typical hydration center, and hydrates 75 ± 06 molecules of H2 O, leaving the bulk H2 O away from the hydration shell unaffected. We then choose the hydration number for Na+ to be 5.2 as discussed in Section [VIII-3]. It follows then that Cl− hydrates 23 ± 06 molecules of H2 O and leaves the bulk H2 O away from the hydration shell in the same state as in pure H2 O. See Table 3 (p. 216). E Figs. VIII-23, VIII-24 and VIII-25 show the H1P−1P results for S = CaCl2  NH4 Cl, E and TMACl, respectively. Fig. 23 indicates that the H1P−1P pattern changes for S = CaCl2 E are the same as a typical hydration center. Namely, the values of H1P−1P at the start and at point X remain the same, while the x1P -loci shift to the left as xS0 S = CaCl2  increases. Thus, Ca2+ together with Cl− is a hydration center without modifying the molecular organization of the bulk H2 O away from the hydration shell. For NH4 Cl E Fig. 24 shows that the values of H1P−1P at x1P = 0 may be decreasing on increasing 0 xS S = NH4 Cl albeit within the estimated uncertainty. If this decrease at x1P = 0 is E due to hydrophilicity, then the value of H1P−1P at point X should equally decrease. Since E the latter values remain constant, we are inclined to conclude that the values of H1P−1P at x1P = 0 remain constant within the estimated uncertainty and NH4 Cl also shows the signature of a hydration center. The effects of TMACl (Fig. 25) show surprisingly the hallmarks of a hydrophile, rather than those of an expected hydrophobe with four methyl groups. As discussed at some length in Chapter VII, this simply means that TMA+ is less hydrophobic (more hydrophilic) than the probing 1P. We note, however, that the values E of H1P−1P at x1P = 0 show the tendency of an initial increase followed by a decrease as 0 xS S = TMACl increases, albeit the variation is within the estimated uncertainty. This trend is similar to typical amphiphiles, acetone and tetramethyl urea (see Figs. VII-15 and VII-16). Thus, we suggest that TMA+ ion is an amphiphile and its hydrophilicity is stronger than its hydrophobicity in comparison with 1P.

226

Fig. VIII-23 to VIII-25

250

200

xCaCl20 = 0

H 1P-1PE / kJ mol-1

0.01198 150

0.02374

100

50

Error

0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E , in 1P−CaCl2 −H2 O at 25  C. Fig. VIII-23. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

250

xNH4Cl0 = 0

200

H 1P-1PE / kJ mol-1

0.02336 0.04579

150

0.06231 100

50

0

0.00

Error

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-24. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P−NH4 Cl−H2 O at 25  C. The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

Eq. (VIII-4)

227

250

xTMACl0 = 0 200

0.02440

H 1P-1PE / kJ mol-1

0.04618 150

0.06894 0.08466

100

50

Error 0

-50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-25. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P–tetramethyl ammonium  − chloride (TMACl) H2 O at 25 C. The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

Figs. VIII-26, VIII-27 and VIII-28 show the mixing scheme and phase diagram for 1P−S−H2 O at 25  C for S = CaCl2  NH4 Cl, and TMACl, respectively. All the mixing scheme boundaries appear to be straight lines. Since CaCl2 and NH4 Cl are hydration centers, the xS0 intercepts of the dotted lines in Figs. 26 and 27 provide the hydration numbers. The results are 110 ± 16 for CaCl2 and 35 ± 04 for NH4 Cl. Since nH = 23 ± 06 for Cl− (Table 3, p. 216), nH = 6 ± 2 for Ca2+ and 1 ± 1 for NH+ 4 . They are listed in Tables 2 (p. 211) and 3. The loci of point X, i.e. the mixing scheme boundary shown in Fig. 28 for TMACl, also show the left shift as xS0 (S = TMACl) increases. This left shift must contain that due to hydration by Cl− . Indeed, the hydration number from the intercept amounts to 21 ± 04! We thus conclude that the shift to the left of E point X is due to hydration by Cl− . Hence, the decrease in H1P−1P at point X must be + + the effect of TMA . We thus conclude that TMA is a more hydrophilic amphiphile than the probing 1P. In relation to the Hofmeister cation ranking, our hydration number decreases in the order of, Ca2+ 6>Na+ 5>NH4 + 1

(VIII-4)

where the number given in brackets is the respective hydration number. From the original idea by Hofmeister (1887) of the “H2 O withdrawing power”, then the above ranking

228

Fig. VIII-26 to VIII-28

0.14

Solid + ? 0.12

xcacl20

0.10 0.08

Liquid - Liquid 0.06 0.04 0.02

I

II

0.00 0.00

0.05

0.10

0.15

0.20

x1P

Fig. VIII-26. The mixing scheme and phase diagram in 1-propanol (1P)−CaCl2 −H2 O at 25  C. The solid line is the phase boundary. The broken line is a presumed phase boundary. The dotted line is the onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan. 0.25

0.20

xNH4Cl0

Solid + ? 0.15

0.10

Liquid - Liquid

0.05

II

I 0.00 0.00

0.05

0.10

0.15

0.20

x1P

Fig. VIII-27. The mixing scheme and phase diagram in 1-propanol (1P)−NH4 Cl−H2 O at 25  C. The solid line is the phase boundary. The broken line is a presumed phase boundary. The dotted line is the onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

Eq. (VIII-5)

229

0.35

0.30

xTMACl0

0.25

0.20

0.15

Liquid - Liquid

0.10

0.05

I 0.00 0.00

II 0.05

0.10

0.15

0.20

x1P

Fig. VIII-28. The mixing scheme and phase diagram in 1-propanol (1P)−TMACl−H2 O at 25  C. The solid line is the phase boundary. The dotted line is the onset of the transition from Mixing Scheme I to II. The data from Koga et al. (2006). Reproduced with permission from Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

coincides in decreasing order of the precipitating power, or the order from the kosmotrope to the chaotrope. Since TMA+ is found to be a hydrophile, analogous to the effect of SCN− or ClO4 − discussed above, we suggest TMA+ is a chaotrope. Thus the overall ranking could be written as, Ca2+ >Na+ >NH4 + >TMA+ 

(VIII-5)

from the kosmotropic to the chaotropic end. This ranking is consistent with some (# b’ and # f’) but not with others (# d’ and # e’) in Table 1 (p. 206). The cations’ effects are weaker than anions’, and the ranking is not unequivocal, influenced by the specificity of the test system: Proteins, temperature, pH, concentrations etc.

[VIII-6] HYDRATION NUMBER OF GLYCINE AND ITS SALTS AS PROBED E PATTERN CHANGE BY THE H1P−1P We applied the same 1P-probing methodology to experimentally estimate the hydration number of glycine and its salts. Glycine in aqueous solution is known to take predominantly the zwitter ion form, − OOC-CH2 -NH3 + . We abbreviate glycine as GL and

230

Fig. VIII-29 to VIII-30

its zwitter ion GL+− . There have been a number of theoretical works in the literature regarding the proton transfer to give rise to the GL+− form from the neutral GL that is stable in gas phase. Some of the studies support a direct intra-molecular transfer from the carboxylic to the amino groups (Rzepa & Yi, 1991; Tortonda et al., 1998; Nagaoka et al., 1998; Tunon et al., 1998). More recent theoretical works, however, favor H2 O-mediated proton transfer mechanisms (Fernandez-Ramos et al., 2000; Chaban & Garber, 2001; Balta & Aviente, 2003; Balta & Aviente, 2004; Ahm et al., 2005; Leung & Rempe, 2005). The latter studies consider a super-molecule of the type GLH2 Om and show that the zwitter ion GL+− H2 Om is the stable form. However, the value of m varies from 1 (Chaban & Garber, 2001) to 6 (Balta & Aviente, 2004) depending on the theoretical method and the chosen potential field. The recent ab initio calculation concludes that the stable m changes from 5 for the neutral GL to 8 for GL+− (Leung & Rempe, 2005). We apply here the 1P-probing methodology with S = GL+−  NaOOC−CH2 −NH2 (NaGL), and HOOC−CH2 −NH3 Cl (GLCl) in the ternary 1P−S−H2 O at 25  C. E data. The arrows in Fig. 30 indicate Figs. VIII-29, VIII-30 and VIII-31 show the H1P E results are shown the onset of liquid–liquid phase separation. The corresponding H1P−1P in Figs. VIII-32, VIII-33 and VIII-34. Figs. 32 and 33 carry the hallmarks of the proper hydration center, as sketched for [A] in Fig. 22. This suggests that both GL+− and Na+ GL− are the proper hydration centers. Thus, the hydration numbers, nH , are estimated from the x1P -loci of point X vs. xS0 (S = GL+− and Na+ GL− ) plots, shown in Fig. VIII-35. The results are nH = 7 ± 06 for GL+− and 48 ± 1 for GL− , taking into account that nH = 52 for Na+ . They are listed in Table 3 (p. 216). Fig. 34 indicates that GL+ Cl− has a hydrophilic propensity, while there is a weak shift to the left on point X. If the latter part is due to hydration of both GL+ and Cl− , then the slope shown in Fig. 35 for GL+ Cl− provides the total hydration number being 5.0 ± 0.5, out of which 2.3 ± 0.6 is due to Cl− . It follows then that nH for GL+ 2.7 ± 1, if the shift to the left of point X is due to hydration to GL+ . Alternatively, this left shift could be due to a hydrophobicity of GL+ , if the latter is an amphiphile. There is no way of unequivocally distinguishing the two possibilities within the present methodology. We point out, however, that such a typical hydrophilic amphiphile as tetramethyl urea or E at the start, x1P = 0, is complex because acetone, the behavior of the value of H1P−1P of the competition between an increase due to hydrophobicity and a decrease due to E hydrophilicity (see Figs. VII-15, VII-16 and VII-30). In contrast, the value of H1P−1P − + 0 at x1P = 0 in Fig. 34 decreases monotonously as xS S = GL Cl  increases. We are thus inclined to favor the hydration scenario for the left shift, and hence to conclude that GL+ works as a hydration center as well as a hydrophile, and its hydration number is 2.7 ± 1. These results are summarized in Table 3. We note in Table 3 (p. 216) that the sum of nH for GL− and nH of GL+ amounts to 7.5 ± 2, the same as that for GL+− , 7.0 ± 0.6. It is therefore tempting to speculate that the −COO− part hydrates 4.8 and the −NH3 + part hydrates 2.7 molecules of H2 O. If so, the un-dissociated −COOH in E as seen in Fig. 34. GL+ works as a hydrophile resulting in the reduction of H1P−1P − On the other hand, −NH2 in GL would be a very weak hydrophile with an almost

231

2 0

H 1PE / kJ mol-1

-2 -4

xGL0 = 0 0.01253

-6

0.02490 -8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-29. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P–glycine(GL)−H2 O  at 25 C (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

2 0

H 1PE / kJ mol-1

-2 -4

xNaGL0 = 0 -6

0.01998 0.04020

-8

0.05918 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-30. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P–sodium glycinate (NaGL)−H2 O at 25  C (Parsons & Koga, 2005). The arrows indicate the sign of phase separation. Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

232

Fig. VIII-31 to VIII-34

2 0

H 1PE / kJ mol-1

-2 -4

xGLCl0 = 0

-6

0.02007 0.04003 0.05986

-8 -10 -12 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-31. The excess partial molar enthalpy of 1-propanol (1P), H1P , in 1P–glycine hydrogen  chloride (GLCl)−H2 O at 25 C (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

250

X

200

H 1P1PE / kJ mol-1

xGL0 = 0 0.01253

150

0.02490 100

50

Error

0

I 0.00

0.02

II 0.04

0.06

0.08

0.10

0.12

0.14

x1P E , in 1P–glycine (GL)−H2 O Fig. VIII-32. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P at 25  C (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

233

300

250

H 1P1PE / kJ mol-1

xNaGL0 = 0

X

200

0.01998 0.04020

150

0.05918 100

Error

50

0

I 0.00

0.02

II 0.04

0.06

0.08

0.10

0.12

0.14

x1P E Fig. VIII-33. The 1-propanol (1P)–1P enthalpic interaction, H1P−1P , in 1P–sodium glycinate  (NaGL)−H2 O at 25 C (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

250

X

200

H 1P1PE / kJ mol-1

Error

xGLCl0 = 0 0.02007

150

0.04003 0.05986 100

50

II I 0

0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P E Fig. VIII-34. The 1-propanol(1P)–1P enthalpic interaction, H1P−1P , in 1P–glycine hydrogen chlo − ride (GLCl) H2 O at 25 C (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

234

Fig. VIII-35

0.07

xS0 (S = GL, NaGL, or GLCI)

0.06

GL NaGL

0.05

GLCl 0.04

0.03

II

0.02

0.01

0.00 0.030

I

0.035

0.040

0.045

0.050

x1P

Fig. VIII-35. The mixing scheme boundary, the loci of the onset of the transition from Mixing Scheme I to II, in 1P−S−H2 O at 25  C, where S = glycine (GL), sodium glycinate (NaGL), or glycine hydrogen chloride (GLCl) (Parsons & Koga, 2005). Reproduced with permission from the Journal of Chemical Physics, 123, 234504, (Parsons, M. T. and Koga, Y.), Copyright (2005), American Institute of Physics.

E negligible hydrophilic effect on the H1P−1P pattern seen for NaGL in Fig. 33, in which E at point X. We recall a similar there is no sign of a decrease in the value of H1P−1P =O and of −NH2 . conclusion in Section [VII-3], comparing the hydrophilic effects of >C= Thus, all our findings discussed above seem to be self-consistent within the 1P-probing methodology.

[VIII-7] CONCLUDING REMARKS IN RELATION TO OTHER STUDIES ON AQUEOUS ELECTROLYTES IN THE LITERATURE [VIII-7-1] Ion pairing Salts (S) in aqueous solutions have been regarded as completely dissociated at least in dilute aqueous solutions. However, there is an increasing number of studies suggesting some form of ion-pairing even at a concentration low enough that the system is still in E . For example, the Mixing Scheme I region as probed by the type (a) pattern of H1P−1P dielectric relaxation spectra were shown to be deconvoluted into a number of Debye relaxation processes. The slower dispersion processes were assigned to the H2 O-shared

235

ion pairs. This is true in the range of xS0 <003 for S = Na2 SO4 (Buchner et al., 1999a). A Raman band shape analysis of aqueous solutions of various sulfate salts suggested that it is the dynamics of the ion pairing that is more important than the static equilibrium between various ion-paired species (Watanabe & Hamaguchi, 2005); see also the exchanges in opinion (Buchner et al., 2006; Watanabe & Hamaguchi, 2006). Our 1P-probing methodology provides the overall effect of a given salt on the molecular organization of H2 O, and cannot give information about the actual ion species (separated, paired, or H2 O-separated). However, as demonstrated above, the information about what the anion and the cation do to H2 O could be identified. This is true whether the ions are completely dissociated or completely paired, just as the case for the amphiphiles studied in Chapter VII. Even within the same molecule of an amphiphile, the effects of the hydrophobic and the hydrophilic parts were identified by the present 1P-probing methodology. This may be thanks to the fact that we have E , and their xS0 dependence (i.e. the fourth the data of third derivative quantity, H1P−1P derivative).

[VIII-7-2] Hydration number, nH E Within our 1P-probing methodology, a salt (S) that causes the change in the H1P−1P pattern depicted as [A] in Fig. 22 is the hydration center, and the bulk H2 O away from the hydration shell remains the same as in pure H2 O. Namely, the x1P locus of point E at the start, X shifts to the left, to a smaller value of x1P , while the values of H1P−1P 0 x1P = 0, and at point X remain unchanged as xS increases. For such a case, hydration occurs for both the cation and the anion and the hydration number, nH , evaluated as shown above is the sum of those by both ions. In Table 2 (p. 211), the second column explains this fact. We then chose the value of nH =52 as the reference as discussed in Section [VIII-3], and hence nH for Cl− becomes 2.3 ± 0.6. Thus, from the results shown in Table 2, we assigned nH for each ion. The results are listed in Table 3 (p. 216). For NaBr, NaI and CH3 4 NCl (TMACl), the shift of point X to the left is attributed to hydration of Na+ in the first two and that of Cl− for the third, respectively. Thus the hydration numbers, nH , as assigned in Table 3, are self-consistent. In turn, we conclude Br − and I− are purely hydrophiles while TMA+ is a more hydrophilic amphiphile than the probing 1P, within the context of the present methodology. In the literature, there are a great many estimates of “hydration number”. Here is a list of randomly selected (Hinton & Amis, 1971; Marcus, 1985; Ohtaki & Radnai, 1993; Kiriukhin & Collins, 2002). The definition of “hydration number” is directly linked to the method of evaluation used under a variety of assumptions. Thus, comparison among different sets of “hydration number” is meaningless. In addition, there is another issue: The experimental estimates are always given to neutral salts, and a reference standard must be assigned for a chosen anion or cation. A number of such assignments choose zero for Cl− , without a convincing argument.

236

Nonetheless, we comment on a recent evaluation of hydration number by Zavitsas (2001). It is shown that the freezing point depression of aqueous salt solutions obey Raoult’s law in a wide concentration range, as long as a specific number of H2 O molecules is regarded as hydrated to each ion, and the solution is assumed to consist of free H2 O and the hydrated ion species. The freezing point depression is written as, (see eq. (D-12) in Appendix D), T ∗ − T  RT ∗ = − ln xW W  T∗ fus HW∗

(VIII-6)

where T and T ∗ are the freezing points of the solution and H2 O fus HW∗ is the enthalpy of fusion of ice, and W is the activity coefficient of H2 O, or RT ln W = EW , the excess chemical potential of H2 O. Hence, the right of eq. (6) is equal to −W /fus HW∗ at T ∗ with pure H2 O is taken as the reference. What is observed in Zavitsas’s work is that the right-hand side of eq. (6) is rewritten as, T ∗ − T  RT ∗ = − ln XW  T∗ fus HW∗

(VIII-7)

Where XW is the readjusted mole fraction of free H2 O, without including the presumably hydrated H2 O. The plots of 1/T against ln XW are impressively straight with marginal scatters. However, this does not necessarily mean that the real solution is an ideal solution of the hydrated ions and free H2 O. Rather, this only indicates that the chemical potential of H2 O is insensitive to the reality. Of course, EW and ES (S = salt) contain information about the reality of the solution, but the minute detail of its variation is often not conspicuous due to enthalpy–entropy compensation prevalent in aqueous solutions (Lumry & Rejender, 1970; Lumry, 2003). Thus some conspicuous variation in the excess partial molar enthalpy and entropy of H2 O HWE and SWE , may be obscured in the resulting EW . We recall that the chemical potential is the first derivative of G. In our 1P-probing E , and its methodology, on the other hand, we use a third derivative quantity, H1P−1P induced changes in the presence of the salt, i.e. equivalent of the fourth derivative. Nevertheless, it is remarkable that both sets of results for salts fall in the reasonable ballpark, as shown in Table 2 (p. 211). Dielectric permittivity measurements on solutions with a wide frequency range provide important information on solvation and complexation as well as on the impact of long-range ion–solvent interactions on the cooperative solvent dynamics. They are particularly useful for studies of aqueous solutions with the hydrogen bond network. The concentration dependence on the static permittivity could also provide the effective “hydration number”. The results show weak concentration dependence, which are summarized in the last column of Table 2 (p. 211). This evaluation comes from the dielectric relaxation point of view, and is qualitatively different from our 1P-probing and that by Zavitsas. Nevertheless, all three sets of “hydration numbers” fall within the reasonable range.

Eq. (VIII-6) to (VIII-7)

237

Table VIII-4. The effects of ions on the molecular organization of H2 O, as probed by the induced E changes in the H1P−1P pattern.

↑

Ion

What the ion does to H2 O

SO4 2−

Hydration Center, nH = 17 ± 3, plus an extra hydrophobe-like effect. See text.

F− CH3 COO− Cl− Br − I− ClO4 −

Hydration Center, nH = 14 ± 2. Hydrphobe with an equal strength as 1P. Hydration Center, nH = 23 ± 06. Hydrophile. Hydrophile, stronger than Br − . Overall hydrophile, stronger than I− , but with an extra hydrophobic effect.

SCN−

Overall hydrophile, almost equal to ClO− 4 , with an extra hydrophobic effect.

Ca2+

Hydration Center, nH = 6 ± 2.

Na+ NH+ 4

Hydration Center, nH = 52 (reference). Hydration Center, nH = 1 ± 1.

CH3 4 N+

Overall hydrophilic, with some hydrophobic effect. More hydrophilic amphiphile.

kosmotropic

chaotropic ↓ ↑ kosmotropic

chaotropic ↓

[VIII-7-3] Hofmeister series Table VIII-4 summarizes our findings from the 1P-probing methodology about what each ion does to H2 O. The top part of the second column lists anions in the order of Hofmeister ranking from the kosmotropic to the chaotropic end following the scheme # c and # f in Table 1 (p. 206). The bottom part of column 2 shows cations following the ranking scheme # b and # f . From Table 4 the following conclusions could be drawn: (1) A kosmotropic ion is either a hydration center or a hydrophobe. For hydration centers, the more kosmotropic the ion, the larger the hydration number. This is in complete accord with the classical idea by Hofmeister (1887), the ranking of “the H2 O withdrawing power.” While we have only a single data point for a hydrophobe, OAc− , we suggest that the more kosmotropic the ion, the more hydrophobic. The relative position in the Hofmeister series between a hydration center and a hydrophobe is yet to be established. Hence, a single parameter (or mechanism) cannot explain the ranking even within kosmotropes. (2) The ions regarded as chaotropes are hydrophiles, since they reduce the value of E within the 1P-probing methodology. This tendency of reduction is stronger H1P−1P

238

for a more chaotropic ion. Thus, while the Hofmeister ranking seems continuous, the mechanism is clearly different between kosmotropes and chaotropes. (3) The boundary between kosmotropes and chaotropes may be between Cl− and Br − . We stress that the results summarized above in Table 4 are the effects of each ion E on the molecular organization of H2 O probed by the x1P dependence pattern of H1P−1P within Mixing Scheme I. While there is no doubt that the properties of the mixed solvent, the aqueous salt solution, play a major role in dictating the structure and the function of biopolymers, there must be a great many more issues to be sorted out before linking the above results to the actual Hofmeister effect for each testing system. Namely, there is another equally important question remaining as to how a given biopolymer behaves in the aqueous salt solution. This is particularly true for the cation ranking. There is yet another issue to be considered regarding the specificity of a biopolymer. Biopolymers are giant molecules whose surfaces are covered with hydrophobic and hydrophilic moieties in varying ratios. We used above 1-propanol (1P) in which the balance of the polar to non-polar surfaces tips towards the more hydrophobic side. It seems therefore mandatory to study the complete replica of the present 1-P probing study by using another probe that is more hydrophilic than 1P. Glycerol (Gly) could be an ideal candidate for this, with 3 −OH’s on the same propyl backbone. In the binary E Gly−H2 O the anomaly in the Gly−Gly enthalpic interaction, HGly−Gly , associated with the transition from Mixing Scheme I to II takes the type (d) pattern as discussed in E Section [VI-3] (see Figs. VI-1(d) and VI-11). It is thus interesting to see how the HGly−Gly pattern changes on addition of typical hydrophobes, hydrophiles and various salts. In closing this chapter, we comment on the recent findings on aqueous solutions of salts by the femto second pump-probe spectroscopy and the dielectric relaxation spectroscopy. The femto second two-color mid-infrared pump-probe spectroscopy can distinguish the dynamics of H2 O in the hydration shell from that in the bulk H2 O away from the hydration shells. By this fine technique, the rotational dynamics of H2 O in the hydration shell was found to be much slower than in the bulk H2 O for Cl−  Br − and I− (Kropman & Bakker, 2001; Bakker et al., 2004) and for SO4 2− and ClO4 − (Omta et al., 2003). Furthermore, the dynamics of the bulk H2 O away from the hydration shell were found to remain the same as that in pure H2 O. It is thus tempting to conclude that the above ions do not enhance or weaken the hydrogen bond network of H2 O away from the hydration shells. More recently, Soper & Weckstrom (2006) drew a similar conclusion by neutron diffraction studies. While the conclusions for F− and Cl− are consistent with our findings (Table 4), those for SO4 2−  Br −  I− , and ClO4 − are not. The dielectric relaxation studies (Wachter et al., 2005) has drawn a similar conclusion. Namely, the impacts of anions spanning from a kosmotropic SO4 2− to a chaotropic ClO4 − on the H2 O dynamics have almost no effect on the relaxation time of the bulk H2 O network, although there are very subtle differences in the concentration dependence of the relaxation time. Why are these very fine, microscopic techniques insensitive in distinguishing the effects of different salts on the molecular organization of bulk H2 O, different from our findings listed in Table 4 and the century-long biochemical experiences since Hofmeister (1987)?

239

When we can provide an answer to this question, we will, no doubt, be a step closer to the reality of aqueous solutions of electrolytes. Spectroscopies/diffractions are generally regarded as a superior technique to thermodynamic investigations in order to advance the molecular level understandings. Or are they? At the risk of being simplistic, we gain structural information by these techniques. After all, the structure is governed, in terms of thermodynamics, by the Gibbs energy, G, which is the net results of all the possible higher derivative quantities. To begin with, G is the difference of the H and TS terms, and the more detailed information given in H and TS separately could be obscured by the H −S compensation prevalent in aqueous solutions (Rajender & Lumry, 1970; Lumry, 2003). H and S are consisting of the partial molar enthalpies and entropies, which in turn compensate with each other obscuring the details in the resulting chemical potentials. And so on. Thus, in order to advance the molecular level understanding, it is mandatory to obtain higher order derivative quantities experimentally for thermodynamic studies. And a further development of a finer non-linear spectroscopy is awaited. There is no doubt that the afore-mentioned techniques have provided the first step towards this end.

Eq. (IX-1) to (IX-2)

Chapter IX

INTERACTIONS IN TERNARY AQUEOUS SOLUTIONS – GENERAL TREATMENT

[IX-1] INTRODUCTION E pattern (the third In Chapters VII and VIII, we saw the effects of sample S on the H1P−1P derivative) in ternary 1P−S−H2 O systems. In effect, therefore, we studied the fourth derivative quantities. By means of this method, we were successful in elucidating the effect of S on the molecular organization of H2 O with unprecedented clarity, almost equal, if not superior, to what modern spectroscopic studies have provided. However, in these two chapters we exclusively used the enthalpy information only. In this chapter, we attempt to gain the entropy analogues. Furthermore, we determine the heterogeneous interaction between different species. For the ternary system abbreviated as B–D–H2 O E E , etc. (Koga, 2004b) in addition to EB−B  HB−B , etc. (W), we determine EB−D  HB−D E As shown in previous chapters, Hi (i = B or D) data are readily determined. It is only the target component i, the amount of which is perturbed and the thermal response of the entire system is measured by calorimetry. To obtain entropy analogues, the excess chemical potential of i, Ei , is necessary as discussed in Section [III-5]. For binary systems, the vapor pressure determination with the Boissonnas analysis was shown to be advantageous in determining Ei (see Section [III-6]). In ternary systems, however, the analysis equivalent to the Boissonnas one is not trivial, while the total vapor pressures are readily determined. We show below two examples, one system by the numerical iteration based on the Gibbs–Duhem relation as for the Boissonnas method, and the other by head-space gas chromatography (Halpern, 1997). We follow the variable conversion from the n1  n2      nm  variable system in which thermodynamic quantities are defined to the more convenient x1  x2      xm−1  N, as discussed in Sections [II-6] or [VII-1]. Here m is the number of component, and N is the total amount of the system. Thus,

 Ei−j

≡N 

E Hi−j

≡N

Ei nj HiE nj







Ei = 1 − xj  xj 



HiE = 1 − xj  xj

−

m−1 

 xk

k=j

 −

m−1  k=j

Ei xk

 xk



HiE xk



(IX-1)

 

(IX-2)

242

 E Si−j ≡N

 E Vi−j

≡N

SiE nj ViE nj



 = 1 − xj 



SiE xj





ViE = 1 − xj  xj

−

m−1 

 xk

k=j

 −

m−1  k=j

SiE xk

 xk



ViE xk



(IX-3)

 

(IX-4)

Note that Ei−j = Ej−i etc. as long as thermodynamic quantities are differentiable. Timasheff’s “preferential interaction parameter” (1998) is equivalent to i−j in our terminology, and it differs from our Ei−j (eq. (1)), by the last two terms in eq. (IX-5) below:    m−1     ln xi  ln xi E  (IX-5) − xk i−j = i−j + RT 1 − xj  xk xj k=j The last two terms on the right of eq. (5) are coming from the ideal mixing entropy. Thus, our Ei−j emphasizes the non-ideality of the system excluding the ideal mixing term. Either way, both are useful in describing the net i−j interaction in terms of chemical potential; namely, the effect of the j-th component on the (excess) chemical potential of the i-th. However, due to enthalpy–entropy compensation prevalent in aqueous solutions (Lumry, 2003; Lumry & Rejender, 1970), the details of the mixing scheme are more E E E conspicuous in Hi−j and Si−j . We note that Ei−j is the second and the latter Hi−j and E Si−j are the third derivatives of G. As argued above, the third derivatives provide more E detailed information than the second ones. Vi−j is also the third derivative quantity.

[IX-2] SOLUTE–SOLUTE INTERACTIONS IN TERT-BUTANOL (TBA) – DIMETHYLSULFOXIDE (DMSO)–H2 O We consider here the ternary system abbreviated as B–D–W, where B, D, and W represents TBA, DMSO and H2 O, respectively. In order to evaluate the i−j interactions (i, j = B, D or W), an accurate determination of HiE , for example, is required. As discussed above in Sections [III-3] and [III-4], HiE and ViE are readily measured experimentally even in a multi-component system. The determination of Ei , however, becomes non-trivial as soon as a third component is added. In this section, we show how the total vapor pressure data for the B–D–W system were treated using the Gibbs–Duhem relation in the range of the mole fraction of B xB < 012 and that of D xD <035. The total pressures, p, were determined by means of the vapor pressure apparatus shown in Section [III-5]. The mixed solvent D–W, with the initial mole fraction of D, xD 0 , was placed in the cell. A small amount of B was added successively and p was determined to ±0001 Torr. As xB increases, xD decreases as xD = xD 0 1 − xB . The values of p were then interpolated graphically to ±001 Torr in the closely spaced grid

Eq. (IX-3) to (IX-7)

243

Table IX-1. The working table to calculate partial pressures, pi (i = B, D, or W) in B-D-W system. Each element in the table is named (m,n) by the row and column numbers. Each element contains the known quantities, xi (i = B, D, or W) and p. The unknown pi ’s are to be solved. To solve the target element(m,n) we use the solved pi values of the “up” and the “left” elements. See text. xD \xB 0 0.01 0.03 0.05

0

0.005

0.01 0.015

(1,1) (2,1) (3,1)

(1,2) (2,2) (3,2)

(1,3) (2,3) (3,3)

(m,1)

(m,2)

(m,3)

0.11

0.12

(1,n)

“left” (m,n − 1)

“up” (m − 1,n) “target” (m,n)

0.33 0.35

of xB and xD (see Trandum et al. (1998) for detail). The resulting interpolated values of p in the grid of xB  xD , a portion of which is sketched in Table IX-1, is the starting point for our numerical iteration based on the Gibbs–Duhem relation. The first element (1, 1) in the table is pertinent for pure H2 O, i.e. xB = xD = 0 pB = pD = 0 and p = pW . pi is the partial pressure of i (= B, D, or W). The first column (m, 1) is for the binary mixtures of D–W, while the first row (1, n) for B–W. Both binary systems have already been solved (Lai et al., 1995; Koga et al., 1990a). Consider the target element (m, n). Suppose that the “up” element (m − 1, n) and the ‘left” element (m, n − 1) are already solved. Thus pB U  pD U  pW U  pB L  pD L , and pW L are known, in addition to the total pressures pU and pL . The superscript U indicates the data for the “up” and L for the “left” elements. Our task is to solve the partial pressures of the target element (m, n), given the values of the partial pressures of the “up” and the “left” elements. We assume or make the table such that the change in mole fraction from the “up” or the “left” element to the target is sufficiently small. Thus the changes in chemical potentials (not excess ones), i (i = B, D, or W) are related by the Gibbs–Duhem relation. Thus, xB B + xD D + xW W = 0

(IX-6)

where represents the small change from the “up” element. The chemical potentials are related to the logarithm of partial pressures (eq. (II-9)) and eq. (6) is rewritten as, 

pB xB ln pBU





x + xDU + D 2





pD ln pDU





U x + xW + W 2





pW ln U pW

 = 0

(IX-7)

244

Similarly, for the change from the “left” element, 

         L xB + xBL xW + xW pW pD pB + xD ln + = 0 ln ln L pBL pDL 2 pW 2 pB + pD + pW = p

(IX-8) (IX-9)

The above three equations (eqs. (7), (8) and (9)), are necessary and sufficient to solve three unknowns, pB  pD and pW of the target element. To start this procedure, however, we need to make some assumptions for the partial pressures in the second column and in the second row elements. Recall the Boissonnas analysis for a binary system in Section [III-6], in which we assumed that the second data point with the most dilute solution is in the Henry’s law region and hence the major component obeys Raoult’s law. This assumption is not rigorously correct, though approximately acceptable. Even so, the Boissonnas method has its own built-in procedure of self-correcting the initial error as the analysis proceeds from point to point. We discussed this at some length in Section [III-6]. For the ternary B–D–W system, we make the following assumptions: (1) In the solution at the lowest mole fraction of a volatile component B, xB = 0005 in the present case (i.e. the elements in column 2), the partial pressure of H2 O pW , is written as, pW = 1 − xB pW ’

(IX-10)

where pW ’ is the partial pressure of the “left” element, i.e. for column 1. Column 1 is for binary D–W and the partial pressures were already solved. Eq. (10) implies that the mixed solvent D–W, the major component, obeys Raoult’s law. (2) In addition to (1), the partial pressures of the less volatile D, pD , in the second column remains the same as that of the “left” element. Namely, the presence of a minute amount of B does not alter the chemical potential of D. This may be approximately acceptable, since the vapor pressure of D turned out to be much less than 1/100th of that of B, as is evident from the results shown in Table IX-2 (p. 246). Thus, the values of pB in column 2 are calculated by eq. (9). We have thus all the partial pressures for column 2. To start iterations based on eqs. (7), (8) and (9), we need the element in row 2 to be assumed. We proceed by guessing a value of pD in row 2-column 3, element (2,3), similar to that in the “left” element (2, 2). With this value of pD , eqs. (8) and (9) provide the solution for pB and pW for (2, 3). Then the “down” element, (3, 3) can be solved uniquely by using eqs. (7), (8) and (9) with the values of the “up”, (2, 3), and the “left”, (3, 2), elements. In this

Eq. (IX-8) to (IX-11)

245

manner, we proceed element by element downwards. Having completed column 3, we start column 4 by guessing the value of pD in element (2, 4). Now, the above two assumptions and a guessed value of pD similar to the “left” in row 2 element for each column may be approximately acceptable, but are most likely not rigorously true. Is there a self-correction operative similar to the Boissonnas analysis for binary systems? It turned out that the slightest wrong guess for pD for the row 2 element results in a wild oscillation in the values of pB and pD resulting in a negative value towards the bottom of the column. This contrasts with the Boissonnas method for binary systems, and may be due to one extra freedom for ternary mixtures. We take this property to our advantage for monitoring the correct initial guess for pD . Since eqs. (7) and (8) use the logarithms of partial pressures, it is a good relative uncertainty, an appropriate number of significant figures, that is required in guessing the appropriate value of pD , although the absolute uncertainty in the smoothed values of p is ±001 Torr. It turns out that the oscillation in the results of pD and pB is extremely sensitive to the choice of pD , and the value of pD must be guessed at least to ±0000001 Torr, to about 4 significant figures, before the oscillation disappears. For the columns for xB > 009, the value of pD even with six significant figures could not stop the oscillation towards the end of the column. The results are shown in Table IX-2. Note that the results are not available for xB = 009 and xD >033, and xB = 012 and xD >023, etc. due to the trouble of oscillation mentioned above. The excess chemical potentials of i (i = B, D, or W) are calculated as a function of xB  xD  using the data in Table 2 (p. 246) by,   pi E  (IX-11) i = RT ln xi pi∗ where i = B, D, or W, and pi∗ is the vapor pressures of pure i. Figs. IX-1, IX-2 and IX-3 show the results. In Fig. 1, the xB dependence of EB shows a subtle change from xD = 015 (up-triangle) to 0.23 (down triangle). This boundary corresponds to the onset of transition from Mixing Scheme I to II in binary DMSO D−H2 O xD = 021 (see Table VI-2 (p. 159)). Similarly in Fig. 2, the xD dependence of EB may show a subtle change at about xB = 005 (down triangle), which corresponds to the boundary region of the I–II transition for binary TBA B−H2 O (see Table V-3 (p. 114)). This is no doubt due to changes in the manner by which each solute modifies H2 O within Mixing Scheme I, and the effect of transition from Mixing Scheme I to II. These subtle differences in the behavior of Ei will be more conspicuous in the second derivative quantities, Ei−j , E E HiE and SiE , and more so in third derivatives, Hi−j and Si−j . The HBE data determined 0 experimentally for a given initial mole fraction of D, xD , of the mixed solvent D−H2 O are shown in Fig. IX-4. The data are obtained as a small amount of B was successively titrated into the mixed solvent D−H2 O. Fig. IX-5 is the values of TSBE , calculated from EB and HBE . Comparing the ordinate scale with that in Fig. 1, the enthalpy–entropy compensation is apparent (Lumry, 2003; Lumry & Rejender, 1970). Similarly, Figs. IX-6 and IX-7 show HDE and TSDE .

246

Table IX-2. Partial pressures, pi / Torr (i = B, D, or W) in tert-butanol (B)–dimethyl sulfoxide (D)–H2 O. xD \xB

0.0

0.005 23.62 0.0 2.448

0.01 23.50 0.0 4.775

0.015 23.39 0.0 6.981

0.02 23.28 0.0 9.060

0.025 23.16 0.0 11.27

0.03 23.05 0.0 13.38

0.035 22.95 0.0 15.32

0.04 22.85 0.0 17.06

0.045

0.0

pW 23.68 pD 0.0 pB 0.0

22.77 0.0 18.56

0.01

23.46 23.34 23.22 23.11 22.99 22.88 22.78 22.68 22.60 pW 23.52 pD .000649 .000649 .000666 .000669 .000672 .000674 .000676 .000679 .000681 .000683 2.421 4.774 7.030 9.154 11.39 13.50 15.44 17.17 18.65 pB 0.0

0.03

23.05 22.90 22.78 22.67 22.55 22.43 22.33 22.24 22.15 pW 23.11 pD .001608 .001608 .001667 .001665 .001670 .001675 .001683 .001684 .001689 .001694 2.400 4.839 7.181 9.380 11.64 13.74 15.68 17.38 18.82 pB 0.0

0.05

22.53 22.41 22.28 22.16 22.03 21.92 21.81 21.72 21.63 pW 22.59 pD .002777 .002777 .002778 .002799 .002820 .002840 .002851 .002872 .002884 .002896 2.411 4.885 7.309 9.570 11.84 13.92 15.84 17.50 18.88 pB 0.0

0.07

21.93 21.78 21.65 21.53 21.41 21.29 21.19 21.10 21.02 pW 21.99 pD .004243 .004243 .004304 .004331 .004360 .004384 .004414 .004423 .004442 .004458 2.441 4.998 7.475 9.773 12.02 14.07 15.96 17.55 18.87 pB 0.0

0.09

21.24 21.09 20.96 20.83 20.71 20.61 20.50 20.42 20.34 pW 21.30 pD .006107 .006107 .006192 .006234 .006274 .006310 .006330 .006368 .006391 .006414 pB 0.0 2.482 5.093 7.607 9.924 12.14 14.14 15.98 17.52 18.77

0.11

20.49 20.33 20.20 20.08 19.97 19.86 19.76 19.68 19.61 pW 20.55 pD .008440 .008440 .008558 .008610 .008659 .008700 .008752 .008767 .008797 .008822 2.522 5.175 7.704 10.02 12.19 14.14 15.92 17.39 18.57 pB 0.0

0.13

pW 19.74 pD .01133 pB 0.0

19.68 .01133 2.556

19.52 .01148 5.229

19.39 .01154 7.753

19.28 .01160 10.05

19.17 .01165 12.17

19.07 .01167 14.05

18.97 .01172 15.77

18.89 .01176 17.17

18.82 .01179 18.29

0.15

pW 18.88 pD .01487 pB 0.0

18.83 .01487 2.577

18.67 .01504 5.248

18.55 .01510 7.746

18.44 .01516 10.00

18.33 .01520 12.06

18.23 .01528 13.89

18.15 .01529 15.53

18.07 .01533 16.86

18.01 .01536 17.92

0.17

pW 17.99 pD .01915 pB 0.0

17.94 .01915 2.580

17.79 .01933 5.226

17.67 .01940 7.680

17.56 .01946 9.891

17.46 .01952 11.87

17.38 .01952 13.63

17.29 .01958 15.19

17.22 .01962 16.46

17.16 .01967 17.48

0.19

pW 17.07 pD .02427 pB 0.0

17.03 .02427 2.562

16.88 .02447 5.161

16.78 .02452 7.551

16.68 .02456 9.696

16.59 .02458 11.59

16.50 .02466 13.30

16.42 .02467 14.78

16.36 .02470 15.97

16.30 .02471 16.95

0.21

pW 16.15 pD .03034 pB 0.0

16.10 .03034 2.523

15.97 .03051 5.050

15.87 .03054 7.360

15.78 .03057 9.429

15.69 .03062 11.24

15.62 .03057 12.88

15.55 .03061 14.28

15.48 .03063 15.41

15.42 .03068 16.37

0.23

pW 15.22 pD .03743 pB 0.0

15.17 .03743 2.462

15.05 .03757 4.896

14.96 .03757 7.115

14.88 .03756 9.095

14.81 .03749 10.81

14.73 .03754 12.40

14.67 .03752 13.72

14.61 .03753 14.80

14.56 .03752 15.72

0.25

pW 14.29 pD .04562 pB 0.0

14.25 .04562 2.380

14.15 .04570 4.704

14.07 .04563 6.816

13.99 .04556 8.700

13.91 .04558 10.35

13.86 .04547 11.85

13.79 .04545 13.10

13.74 .04542 14.12

13.69 .04547 15.05

247

0.05

0.055

0.06

0.07

0.08

0.09

0.10

0.11

0.12

xB \xD

22.69 0.0 19.80

22.63 0.0 20.75

22.59 0.0 21.48

22.52 0.0 22.41

22.47 0.0 22.97

22.44 0.0 23.35

22.42 0.0 23.54

22.41 0.0 23.66

22.40 0.0 23.74

pW pD pB

0.0

22.53 .000683 19.86 22.08 .001701 19.95

22.47 .000686 20.79 22.03 .001700 20.84

22.42 .000687 21.49 21.98 .001708 21.58

22.36 .000690 22.41 21.91 .001719 22.42

22.31 .000693 22.97 21.86 .001731 22.98

22.28 .000698 23.36 21.82 .001738 23.40

22.26 .000700 23.54 21.79 .001763 23.57

22.25 .000703 23.67 21.77 .001784 23.75

22.24 .000709 23.74 21.75 .001799 23.81

pW pD pB pW pD pB

0.01

21.57 .002901 19.95

21.51 .002921 20.82

21.47 .002927 21.45

21.39 .002951 22.35

21.34 .002978 22.92

21.28 .003022 23.38

21.25 .003045 23.56

21.21 .003077 23.76

21.18 .003130 23.85

pW pD pB

0.05

20.95 .004482 19.90

20.90 .004485 20.72

20.85 .004510 21.33

20.77 .004550 22.24

20.71 .004596 22.83

20.66 .004623 23.29

20.61 .004697 23.51

20.55 .004767 23.74

20.52 .004821 23.85

pW pD pB

0.07

20.28 .006424 19.75

20.22 .006465 20.55

20.17 .006478 21.15

20.09 .006530 22.05

20.02 .006587 22.65

19.95 .006688 23.14

19.90 .006747 23.39

19.85 .006816 23.63

19.80 .006926 23.79

pW pD pB

0.09

19.54 .008867 19.52

19.49 .008871 20.29

19.45 .008917 20.88

19.35 .008995 21.80

19.27 .009093 22.43

19.21 .009144 22.90

19.15 .009273 23.20

19.07 .009420 23.47

19.02 .009536 23.66

pW pD pB

0.11

18.76 .01181 19.20

18.70 .01187 19.96

18.66 .01190 20.54

18.57 .01199 21.46

18.50 .01207 22.10

18.40 .01226 22.61

18.34 .01238 22.95

18.28 .01250 23.21

18.21 .01267 23.43

pW pD pB

0.13

17.94 .01543 18.82

17.89 .01543 19.54

17.84 .01549 20.13

17.75 .01562 21.06

17.65 .01580 21.74

17.59 .01588 22.22

17.52 .01605 22.59

17.43 .01630 22.89

17.36 .01651 23.15

pW pD pB

0.15

17.11 .01967 18.34

17.04 .01977 19.07

16.99 .01982 19.65

16.90 .01995 20.58

16.83 .02004 21.26

16.73 .02031 21.78

16.65 .02054 22.19

16.59 .02071 22.48

16.52 .02092 22.74

pW pD pB

0.17

16.23 .02481 17.81

16.19 .02480 18.50

16.14 .02487 19.09

16.05 .02501 20.03

15.94 .02531 20.75

15.88 .02545 21.26

15.80 .02565 21.67

15.71 .02598 22.00

15.62 .02635 22.30

pW pD pB

0.19

15.38 .03065 17.19

15.32 .03078 17.89

15.27 .03085 18.48

15.17 .03106 19.43

15.10 .03112 20.13

15.01 .03142 20.67

14.92 .03178 21.11

14.85 .03203 21.43

14.80 .03217 21.70

pW pD pB

0.21

14.49 .03765 16.55

14.45 .03763 17.22

14.41 .03768 17.81

14.33 .03777 18.75

14.22 .03817 19.50

14.14 .03846 20.05

14.06 .03870 20.47

13.97 .03914 20.83

-------

pW pD pB

0.23

13.65 .04536 15.83

13.60 .04547 16.51

13.55 .04555 17.10

13.44 .04592 18.07

13.37 .04604 18.79

13.29 .04632 19.35

13.19 .04678 19.79

13.12 .04708 20.14

-------

pW pD pB

0.25

0.03

(Continued)

248

Fig. IX-1 to IX-2

Table IX-2. (Continued) xD \xB

0.0

0.005

0.01

0.015

0.02

0.025

0.03

0.035

0.04

0.045

0.27

pW pD pB

13.39 .05495 0.0

13.35 .05495 2.282

13.26 .05496 4.481

13.18 .05486 6.485

13.11 .05474 8.271

13.06 .05457 9.820

12.99 .05453 11.27

12.94 .05445 12.44

12.89 .05442 13.43

12.84 .05434 14.32

0.29

pW pD pB

12.51 .06543 0.0

12.47 .06543 2.171

12.39 .06538 4.241

12.32 .06517 6.126

12.27 .06493 7.802

12.20 .06488 9.288

12.15 .06469 10.66

12.10 .06460 11.77

12.05 .06452 12.71

12.00 .06457 13.61

0.31

pW pD pB

11.67 .07698 0.0

11.63 .07698 2.054

11.56 .07683 3.988

11.50 .07658 5.761

11.44 .07643 7.346

11.39 .07613 8.740

11.34 .07599 10.04

11.29 .07585 11.09

11.24 .07578 12.01

11.21 .07559 12.87

0.33

pW pD pB

10.87 .08946 0.0

10.84 .08946 1.938

10.77 .08934 3.749

10.71 .08908 5.415

10.66 .08872 6.890

10.60 .08863 8.231

10.55 .08844 9.447

10.50 .08836 10.46

10.45 .08834 11.35

10.38 .08859 12.21

0.35

pW pD pB

10.13 0.1026 0.0

10.10 0.1026 1.832

10.03 0.1025 3.531

9.972 0.1022 5.095

9.909 0.1022 6.497

9.860 0.1019 7.763

9.809 0.1018 8.888

9.766 0.1016 9.855

9.726 0.1014 10.71

9.712 0.1008 11.48

7

xD 0.03 0.07 0.15 0.23 0.27 0.31 0.35

µ BE / kJ mol-1

6

5

4

3 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

xB

Fig. IX-1. The excess chemical potential of TBA, EB , against mole fraction of TBA, xB , in TBAB−DMSO D−H2 O (W) at 2495  C. The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

249

0.05

0.055

0.06

12.78 .05451 15.12

12.74 .05452 15.78

11.97 .06438 14.36

xD \xB

0.07

0.08

0.09

0.10

0.11

0.12

12.69 .05461 16.37

12.63 .05455 17.29

12.54 .05486 18.03

12.45 .05517 18.61

12.40 .05531 19.02

-------

-------

pW pD pB

0.27

11.92 .06439 15.01

11.89 .06437 15.58

11.77 .06487 16.55

11.68 .06531 17.30

11.56 .06607 17.93

11.37 .06771 18.47

-------

-------

pW pD pB

0.29

11.14 .07580 13.63

11.09 .07595 14.29

11.02 .07630 14.88

10.93 .07657 15.82

10.84 .07701 16.56

10.79 .07702 17.13

-------

-------

-------

pW pD pB

0.31

10.34 .08851 12.93

10.30 .08851 13.57

10.28 .08825 14.10

10.22 .08806 15.00

10.21 .08747 15.66

-------

-------

-------

-------

pW pD pB

0.33

9.685 0.1005 22.54

9.668 0.1001 12.77

9.633 0.1000 13.29

9.574 .09982 14.18

9.445 0.1009 14.94

-------

-------

-------

-------

pW pD pB

0.35

This table is reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum et al.), Copyright (1998), Americal Chemical Society.

7

µ BE / kJ mol-1

6

5

4

xB 2 0.00

0.05

0.10

0.06 0.08 0.10 0.12

0.01 0.02 0.035 0.05

3

0.15

0.20

0.25

0.30

0.35

0.40

xD

Fig. IX-2. The excess chemical potential of TBA, EB , against mole fraction of DMSO, xD , in TBA B−DMSO D−H2 O (W) at 2495  C. The data from Trandum et al., (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

250

Fig. IX-3 to IX-6

0

µ DE / kJ mol-1

-2

-4

-6

-8

-10 0.00

0.21 0.27 0.31 0.35

xD 0.03 0.09 0.15 0.02

0.04

0.06

0.08

0.10

0.12

0.14

xB

Fig. IX-3. The excess chemical potential of DMSO, ED , against mole fraction of TBA in TBA B−DMSO D−H2 O (W) at 2495  C. The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society. 10

H BE / kJ mol-1

5

0

x D0 = 0 0.0478 0.1031 0.1999 0.2341 0.2735 0.3540

-5

-10

-15

-20 0.00

0.02

0.04

0.06

0.08

0.10

0.12

xB

Fig. IX-4. The excess partial molar enthalpy of TBA, HBE , against mole fraction of TBA, xB , in TBA B−DMSO D−H2 O (W) at 25  C. xD0 is the initial mole fraction of DMSO in the mixed solvent of DMSO−H2 O into which TBA was titrated to obtain HBE . The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

251

5

TS BE / kJ mol-1

0

-5

xD0 = 0

-10

0.0478 0.1031 0.1999 0.2341 0.2735 0.3540

-15

-20

-25 0.00

0.02

0.04

0.06

0.08

0.10

0.12

xB

Fig. IX-5. The excess partial molar entropy of TBA (times T ), TSBE , against mole fraction of TBA, xB , in TBA B−DMSO D−H2 O (W) at 25  C. The data from Trandum et al. (1998). xD0 is the initial mole fraction of the mixed solvent, D–W. Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society. -2 -4

H DE / kJ mol-1

-6 -8

xB0 = 0

-10

0.00954 -12

0.0302

-14

0.0507 0.0724

-16 -18 0.00

0.05

0.10

0.15

0.20

0.25

0.30

0.35

xD

Fig. IX-6. The excess partial molar enthalpy of DMSO, HDE , against mole fraction of DMSO, xD , in TBA B−DMSO D−H2 O (W) at 25  C. The data from Trandum et al. (1998). xB0 is the initial mole fraction of TBA in the mixed solvent of TBA–W, into which DMSO was titrated to obtain HDE . Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

252

Fig. IX-7 to IX-9

0

TS DE / kJ mol-1

-2

-4

xB0 0.00954

-6

0.0302 0.0507

-8

0.0724 -10 0.00

0.05

0.10

0.15

0.20

0.25

0.30

0.35

xD

Fig. IX-7. The excess partial molar entropy of DMSO (times T ), TSDE , against mole fraction of DMSO, xD , in TBA B−DMSO D−H2 O (W) at 25  C. The data from Trandum et al. (1998). xB0 is the initial mole fraction of the mixed solvent B–W before D is mixed. Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

Next, we determine the i−j interaction functions in terms of chemical potential by eq. (1) assuming that the xB and xD increments are small enough. The plots of EB−B against xB are shown in Fig. IX-8 and those of EB−D against xB are shown in Fig. IX-9. The self-consistency of the partial pressure data was checked by noting that the values of EB−D were equal to those of ED−B within ±001 kJ mol−1 (see Trandum et al. (1998) for detail). It is now clearer in Fig. 8 that the xD dependence of EB−B changes from xB = 00425 to 0.065, corresponding to point X (0.045) and point Y (0.065) in binary TBA B−H2 O (see Table V-3 (p. 114)). Similarly, the xB dependence of EB−D shows a change from xD = 024 to 0.34, the Mixing Scheme transition region from I to II, for binary D–W (see Table VI-2 (p. 159)). The B—B interacE E tions in terms of enthalpy and entropy, HB−B are shown in Fig. IX-10. and TSB−B E In calculating them by eqs. (2) and (3), the HB data at a given xD0 were interpolated at fixed values of xD , noting that xD = xD0 (1−xB  (see Westh & Koga (1996) E for detail). The HB−B pattern changes on addition of D are similar to what we saw in E the 1P-probing methodology discussed in Chapter VII. The H1P−1P pattern is replaced E here with HB−B . Accordingly, Fig. 10 indicates that DMSO (D) is an amphiphile in comparison with the probing TBA (B). From the degree of the left shift of point E X of the HB−B pattern, DMSO seems to contain a rather strong hydrophobic effect, in addition to the reduction of the peak height which indicates a hydrophilic contribution. We discuss the relative effects of the hydrophobic and the hydrophilic

253

5

0.0125 0.0225 0.0325 0.0425 0.053 0.065 0.085 0.105

xB

0

µ BBE / kJ mol-1

-5 -10 -15 -20 -25

-30 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40

xD

Fig. IX-8. The TBA–TBA interaction in terms of chemical potential, EB−B , in TBA B−DMSO D−H2 O (W) against mole fraction of DMSO, xD , at 25  C. The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

2

µ BDE / kJ mol-1

0

-2

-4

xD -6

-8 0.00

0.02

0.04

0.06

0.08

0.04 0.08 0.12 0.16 0.24 0.34 0.10

0.12

0.14

xB

Fig. IX-9. The TBA–DMSO interaction in terms of chemical potential, EB−D , in TBA B−DMSO D−H2 O (W) against mole fraction of TBA, xB , at 25  C. The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

254

Fig. IX-10 to IX-12

xD = 0 0.0478 0.1031 0.1999 0.3540

HBBE or TSBBE / kJ mol-1

500

400

X

300

200

100

0

-100 0.00

0.02

0.04

0.06

0.08

0.10

xB E , (filled symbol) and entropy, Fig. IX-10. The TBA–TBA interaction in terms of enthalpy, HB−B E TSB−B , (hollow symbols) in TBA B−DMSO D−H2 O (W) at 25  C. Point X is determined E using HB−B data. The original data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

contribution within DMSO in the next section, where we use 2-butoxyethanol (BE) E instead of TBA for B in B–D–W. We note in Fig. 10 that the HB−B pattern preserves type (a) anomaly (Fig. VI-1(a)) within Mixing Scheme I of D–W, xD < 02, suggesting that the B–B interaction is bulk H2 O-mediated which is modified by the presence of both B and D. We have discussed this repeatedly in the previous chapters. Another point to note in the figure is the strong enthalpy–entropy compensation effect. E The heterogeneous B–D interactions in terms of enthalpy and entropy, HB−D and E TSB−D are also calculated by eqs. (2) and (3) in the same manner as the B–B interactions. The results are shown in Fig. IX-11. The enthalpy–entropy compensation effect is also E apparent in the figure. The xB dependence pattern of HB−D shows type (b) (Fig. VI-1(b)) (there is a hint of type(a) anomaly) at a low xD composition. As xD increases, however, E the pattern changes to type (c), the same as HD−D for binary DMSO−H2 O, as shown in Fig. VI-14. This supports the fact that the B–D interaction is also bulk H2 O-mediated. E In Fig. IX-12, we plot the values of HB−D extrapolated to xB = 0 against xD . While E the data points are scarce, its xD dependence seems to take type (c) as HD−D for the E binary DMSO D−H2 O, as shown in Fig. VI-14. This indicates that HB−D at xB = 0 reflects how DMSO modifies H2 O, and thus the B–D interaction is operative via the bulk H2 O-modified by the presence of both solutes, as long as the composition of solutes is low enough for the system to be within Mixing Scheme I. This point has been stressed throughout Chapters V to VIII.

255

140

xD 0.0239 0.0478 0.1031 0.1999 0.2735

X

HBDE or TSBDE / kJ mol-1

120 100 80 60 40 20 0

0.00

0.02

0.04

0.06

0.08

0.10

xB E Fig. IX-11. The TBA–DMSO interaction in terms of enthalpy, HB−D , (filled symbols) and entropy, E TSB−D , (hollow symbols) against mole fraction of TBA, xB , in TBA B−DMSO D−H2 O (W) at 25  C. The data from Trandum et al. (1998). Reproduced with permission from the Journal of Physical Chemistry, 102, 5182, (Trandum, Ch. et al.), Copyright (1998), American Chemical Society.

140

HB-DE (xB=0) / kJ mol-1

120 100 80 60 40 20 0

0.0

0.1

0.2

0.3

0.4

0.5

xD E Fig. IX-12. The TBA–DMSO enthalpic interaction, HB−D , at the infinite dilution of TBA, xB = 0, − − in TBA B DMSO D H2 O (W) against mole fraction of DMSO, xD , at 25  C. The original data from Trandum et al. (1998). The xD -dependence is type (c) pattern discussed in Section [VI-1] E (see Fig. VI-1(c)). The DMSO–DMSO enthalpic interaction HD−D in binary DMSO−H2 O takes the type (c) pattern also. See text.

256

Fig. IX-13 to IX-14

[IX-3] SOLUTE–SOLUTE INTERACTIONS IN 2-BUTOXYETHANOL (BE)–DIMETHYL SULFOXIDE (DMSO)–H2 O The vapor pressure analysis successful in the previous section is not applicable to this BEB−DMSO D−H2 OW system, since the system has an azeotrope at about xB ≈ 0005. For binary BE−H2 O, the Boissonnas method was applied from the solute-rich end down to xBE = 001 (see the difficulty associated with the presence of an azeotrope in applying the Boissonnas method discussed in Sections [III-6] and [III-7]). For ternary systems, the following methods could be possible solutions: (1) At a given xB  xD , the total pressure, p, is determined by the vapor pressure apparatus described in Section [III-5]. Then the gas phase composition of B, D, and W is determined by gas chromatography. The sampling for gas chromatography could be problematic, but if successful, this is the most reliable method. (2) If component D is much less volatile so that the pD is a few orders of magnitude smaller than pB , the absolute value of pB alone can be determined by head-space gas chromatography (Halpern, 1997). The binary B−H2 O, in which the values of pB are already known, could be used as the standard samples for gas chromatography. Thus, we obtain EB as a function of xB and xD . This was used for the present BE−DMSO−H2 O. (3) If we limit ourselves to the very dilute aqueous solution, we can assume that the partial pressure of B, pB , obeys Henry’s law and hence that of the mixed solvent D–W follows Raoult’s law, being the major component. Thus the partial pressure of D–W is calculated by, pW + pD  = pW + pD 0 1 − xB , where pW + pD 0 is the total pressure of the binary D–W(i.e. xB = 0) at the identical xD . This is exactly what we did in the previous section to evaluate the partial pressures of column 2 in Table 1, in order to start the iteration. If in addition we limit ourselves to the case in which component D is non-volatile, such as a salt, then it follows that pW = pW 0 1 − xB  and pB = p − pW with pD = 0. This route was used recently to calculate E1P−S in 1-propanol (1P)–salt (S)–H2 O (Miki et al., 2006). We used method (2) above to determine the partial pressure of BE, pB , in BE (B)– DMSO (D)–H2 O (W). The resulting excess chemical potential of BE, EB , is shown in Fig. IX-13, as a function of the mole fraction of BE, xB , and the initial mole fraction of D in the initial mixed solvent D–W, xD0 . Fig. IX-14 shows the plots of HBE . From E and EB−B are calculated by eqs. (2) and (1) at the grid of xB  xD  and are these, HB−B E plotted in Fig. IX-15. Similarly, HB−D and EB−D are calculated and plotted in Fig. IX-16. E E . This is the consequence The value of i−j is an order of magnitude smaller than Hi−j of the enthalpy–entropy compensation (Lumry, 2003; Lumry & Rajender, 1970). The E general behaviors of Hi−j are all similar to the equivalent plots for TBA−DMSO−H2 O discussed in the previous section. Some quantitative differences are due to the difference E in the hydrophobicity between TBA and BE. The values of HB−D at xB = 0 are obtained by extrapolating the data in Fig. 16, as we did in the previous section. Alternatively, we

257

9

xD0 = 0 0.0159

8

µ BE / kJ mol-1

0.0517 7

0.0899

6

5

4

3 0.00

0.01

0.02

0.03

0.04

0.05

0.06

xB

Fig. IX-13. The excess chemical potential of BE, EB , against mole fraction of BE, xB , in BEB−DMSOD−H2 O (W) at 25  C. xD0 is the initial mole fraction of DMSO in D–W, into which a small amount of B was added for gas chromatographic analysis. The data from Westh & Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 433, (Westh, P. and Koga, Y.), Copyright (1996), American Chemical Society.

5

HBE / kJ mol-1

0

xD0 = 0 0.0126 0.0250 0.0545 0.0899 0.2006 0.3284 0.6607

-5

-10

-15

xD0 = 1 0.00

0.02

0.04

0.06

0.08

xB

Fig. IX-14. The excess partial molar enthalpy of BE, HBE , against mole fraction of BE, xB , in BEB−DMSOD−H2 O at 25  C. xD0 is the initial mole fraction of DMSO−H2 O into which BE is titrated to measure HBE . The data from Westh & Koga (1996). Reproduced with Permission from the Journal of Physical Chemistry, 100, 433, (Westh, P. and Koga, Y.), Copyright (1996), American Chemical Society.

258

Fig. IX-15 to IX-17

1200

xD = 0 0.02 0.04 0.06 0.08

µ B-BE or H B-BE / kJ mol-1

1000 800

HB-BE

600 400 200 0

-200 -400 0.00

µ B-BE 0.01

0.02

0.03

0.04

0.05

xB E Fig. IX-15. The BE–BE interaction in terms of chemical potential, EB−B , and enthalpy, HB−B , against mole fraction of BE in BEB−DMSOD−H2 O (W). The data from Westh & Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 433, (Westh, P. and Koga, Y.), Copyright (1996), American Chemical Society.

160

X

xD = 0 0.02 0.04 0.06 0.08

µ B-DE or H B-DE / kJ mol-1

140 120 100 80

HB-DE

60 40 20 0 -20 0.00

E

µ B-D 0.01

0.02

0.03

0.04

0.05

xB E Fig. IX-16. The BE–DMSO interaction in terms of chemical potential, EB−D , and enthalpy, HB−D , against mole fraction of BE, xB , in BE−DMSO−H2 O at 25  C. The data from Westh & Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 433, (Westh, P. and Koga, Y.), Copyright (1996), American Chemical Society.

259

10

HBE at xB = 0/kJ mol-1

5

0

Point Y for DMSO - H2O

-5

-10

Point X for DMSO - H2O

-15

-20 0.0

0.2

0.4

0.6

0.8

1.0

xD

Fig. IX-17. The extrapolated value of the excess partial molar enthalpy of BE, HBE , to the infinite dilution of BE, xB = 0, against mole fraction of DMSO, xD , in BE−DMSO−H2 O at 25  C. It is striking that the xD -dependence resembles that for binary DMSO−H2 O. The data from Westh & Koga (1996). Reproduced with permission from the Journal of Physical Chemistry, 100, 433, (Westh, P. and Koga, Y.), Copyright (1996), American Chemical Society.

first extrapolate the HBE data (Fig. 14) to xB = 0 to obtain HBE data at xB = 0. We then plot them against xD as shown in Fig. IX-17. It is striking that the values of HBE at the infinite dilution of B in the presence of D makes a clear change in the xD dependence E slope at about point X and Y in HD−D of purely binary D–W, as is evident in Fig. 17. E at xB = 0 From the smooth curve drawn in Fig. 17, we evaluate graphically HB−D by eq. (2). The results from both routes, from Fig. 16 and from Fig. 17, are plotted in Fig. IX-18. As for B = TBA shown in Fig. 12, Fig. 18 shows a type (c) pattern, E Fig. VI-1(c), similar to HD−D for binary DMSO−H2 O (Fig. VI-14). Namely, the heteroE geneous B–D interaction at xB = 0 HB−D xB = 0, shows the same xD dependence of the E homogeneous D–D interaction, HD−D , for binary D–W. This suggests that upon addition E of the infinitely small amount of B in the mixed solvent D–W, the HB−D xB = 0 merely E reflects the mixing situation of D–W, or the HB−D xB = 0 probes the mixing scheme in D–W. This suggests again that the solute–solute interaction, whether homogeneous or heterogeneous, is bulk H2 O-mediated within Mixing Scheme I. For the heterogeneous interaction, the effects of two species act competitively or cooperatively. Furthermore, E for D–W could also be said to probe the state of mixing scheme modified by the HD−D D at the infinite dilution. E E xB = 0 pattern and that of HD−D for Thus, quantitative differences in the HB−D binary D–W indicate the relative differences in the way in which TBA, BE, and DMSO modify H2 O at the infinite dilution. Fig. IX-19 compares these. Thus the values of

260

Fig. IX-18 to IX-19

160

From data in Fig. IX-17 From data in Fig. IX-16

HB-DE at xB = 0/kJ mol-1

140 120 100

X 80 60 40

Y 20 0 0.00

0.05

0.10

0.15

0.20

0.25

0.30

0.35

xD E Fig. IX-18. The BE–DMSO enthalpic interaction, HB−D , at the infinite dilution of BE, xB = 0, − − against mole fraction of DMSO, xD , in BE DMSO H2 O at 25  C. This shows type (c) pattern, [VI-1]. The data from Westh & Koga (1996).

160

HD-DE for D-W

140

H TBA-DE at x TBA = 0

HB-DE at xB = 0 or HD-DE

HBE-DE at xBE = 0

XBE

120 100 80

XTBA 60 40

X

20

Y

0 -20 0.0

0.1

0.2

0.3

0.4

0.5

xD E E at xB = 0 for B−D−H2 O for B = BE and TBA and HD−D Fig. IX-19. Comparison between HB−D for binary D – W at 25  C. The data from Westh & Koga (1996) for B = BE, and Trandum et al. (1998) for B = TBA. Those for the binary D–W from Lai et al. (1995).

Eq. (IX-12)

261

E HB−D xB = 0 are 150, and 122 kJ mol−1 for B = BE and TBA at xD = 0, while that of E HD−D is 51 kJ mol−1 at the infinite dilution in D–W. (See Table VI-2 (p. 159)) The xD = BE, 0.17 for B = = TBA and 0.21 for binary loci of point X in Fig. 19 are 0.09 for B = D–W. In the previous chapters, we showed that the xi locus of the onset of transition to Mixing Scheme II, i.e. point X, occurs at a smaller value for a stronger hydrophobe. E Furthermore the value of Hi−i at point X is larger for a stronger hydrophobe, and smaller if the hydrophilicity is stronger. Accordingly, the above observations indicate that the net hydrophobicity of the solute in the H2 O-rich region is ranked as,

BE>TBA>DMSO

(IX-12)

Thus, the heterogeneous interaction may be useful in ranking the species belonging to the E different classes; BE and TBA are mono-ols and show type (a) Hi−i pattern, while DMSO E has type (c) Hi−i . This inter-class comparison was not possible in the previous chapters, in which only the homogeneous interaction functions were used. While the heterogeneous interaction is the net result of the competition or cooperation of the two different E species, it may not be linearly additive. If so, the value of HB−D at the infinite dilution, E E xB = xD = 0, would be equal to the average of HD−D at xD = 0 and HB−B at xB = 0 in the respective binary systems. As is evident in Table IX-3, it is not the case, while both values are parallel. The number of examples, however, is limited at this point in time. E Fig. 10 for B = TBA and Fig. 15 for B = BE are equivalent to the H1P−1P pattern changes shown repeatedly in Chapter VII in the presence of an amphiphile. Here we use TBA or BE as a probe to elucidate the effect of DMSO on H2 O instead of 1-propanol (1P) in Chapter VII. As discussed in Chapter VII, the net effect of DMSO can be separated into the hydrophobic and the hydrophilic contributions relative to the probing E B. Namely, the degree of the left shift of point X in HB−B pattern reflects hydrophobicity E and that of the decrease in HB−B at point X shows the strength of hydrophilicity. E Fig. IX-20 shows the xB locus of point X of HB−B for B–D–W, normalized to that for binary B–W xD = 0. The figure indicates that DMSO has a hydrophobic moiety stronger than BE and even more so than TBA. This surprising indication may support the view discussed in Section [VI-5] that in Mixing Scheme I, DMSO molecules exist as = O pointing CH3 -groups outward a small tight cluster bound together by dipoles of >S = E of these clusters. Fig. IX-21 shows the change in the value of HB−B at point X. The E Table IX-3. The values of HB−D at xB = xD = 0 in B−DMSO−H2 O.

B TBA BE a b

E HB−D xB = xD = 0 kJ mol−1

E E

HB−B xB = 0a + HD−D xD = 0b /2 kJ mol−1

122 150

70 220

See Table V-3 (p. 114) See Table VI-2 (p. 159)

262

Fig. IX-20 to IX-21

1.1

BE TBA

xB / xB * at point X

1.0

0.9

0.8

0.7

0.6

0.5 0.00

0.02

0.04

0.06

0.08

0.10

0.12

xD E Fig. IX-20. The xB locus of point X in the HB−B pattern in B–D–W against xD . xB ∗ is that for binary B–W, xD = 0. The data from Westh & Koga (1996) for B = BE, and from Trandum et al. (1998) for B = TBA.

1.1

BE TBA

HB-BE / HB-BE* at point X

1.0

0.9

0.8

0.7

0.6

0.5

0.4 0.00

0.02

0.04

0.06

0.08

0.10

0.12

xD ∗

E E E Fig. IX-21. The value of HB−B at point X in the HB−B pattern in B–D–W at 25  C. HB−B is that for binary B–W. The data from Westh and Koga (1996) for B = BE and Trandum et al. (1998) for B = TBA.

263

figure suggests that DMSO has a hydrophilic moiety stronger than TBA and more so than BE, which is understandable, in view of the hydrophobicity ranking, BE>TBA, as discussed in Chapter V. Thus, it is clear that the heterogeneous interaction functions in many other ternary or multi-component aqueous solutions will be useful in understanding the molecular processes in aqueous solutions, and particularly in those of biological importance. More generally, model-free thermodynamic data of the second, third and higher order derivatives of G are powerful in advancing the molecular level understanding of aqueous or any other solutions. The accumulation of such data without resorting to any fitting function is awaited.

Chapter X

IN CLOSING – EXECUTIVE SUMMARY ON THE EFFECT OF SOLUTE ON H2 O

The nature of an aqueous solution is crucially dependent on its composition; there are three distinct composition regions, in each of which the mixing scheme is qualitatively different. The unique characteristics of liquid H2 O are retained in the H2 O-rich region even in the presence of solute. In the solute-rich composition range, solute molecules tend to cluster together, and thus the individuality of the solute dictates the nature of the mixture. In the intermediate region, the aqueous solution consists of two kinds of clusters, one rich in H2 O and the other in the solute molecules. These clusters are reminiscent of the mixing schemes operative in the H2 O-rich and the solute-rich regions. For convenience, we call these three mixing schemes operative in each region Mixing Scheme I, II and III from the H2 O-rich end. Mixing Scheme I is such that the integrity of liquid H2 O is retained. In particular, the hydrogen bond network is connected throughout the entire bulk (bond-percolated) and the hydrogen bond angle and its strength fluctuates widely. If the solute in question is a hydrophobe, classically called a “structure-maker”, the hydrogen bond network of H2 O is enhanced in the immediate vicinity of the solute corresponding to the classical “iceberg formation”. In addition, however, the hydrogen bond probability of bulk H2 O away from “icebergs” is reduced progressively. If a hydrophile, a “structure-breaker”, is added, it forms hydrogen bonds directly to the existing hydrogen bond network of H2 O and the solute acts as an impurity center within the network. Consequently, the degree of fluctuation inherent in liquid H2 O is suppressed and the hydrogen bonds are partly broken. These effects are due to the loss of H donor/acceptor symmetry enjoyed in pure liquid H2 O. If the solute is an amphiphile, the competition or cooperation of the effects of the hydrophobic and the hydrophilic moieties results in the net effect on H2 O. Either way, as long as the composition is low enough, the global hydrogen bond probability is still high enough that the bond percolation of H2 O is intact. When the solute composition increases to a threshold value, depending on the nature of solute, the global average of the hydrogen bond probability reaches the bond-percolation threshold for the bond connectivity of the ice Ih type. Thus, the integrity of liquid H2 O is lost and Mixing Scheme II sets in. Within Mixing Scheme I, the solute–solute interaction is bulk H2 O-mediated and of an extremely long range. The transition from Mixing Scheme I to II is associated with anomalies in the third derivatives of Gibbs energy, G. This contrasts with “phase” transitions for which the second derivative quantities show anomalies. Thus, the transition in the

266

Fig. A-1

mixing scheme is more subtle than “phase” transitions. The mixing scheme transition occurs at a lower solute composition as temperature increases. Its loci in the mole fraction–temperature field extrapolate to about 83 C at the infinite dilution, close to the hydrogen bond percolation threshold in terms of the average probability of hydrogen bond in pure H2 O estimated by Stanley et al. (Stanley & Teixeira, 1980). For tert-butanol or 2-butoxyethanol, the boundary ends towards low temperatures at the incongruent melting point of an addition compound of BH2 Om type, where B is tert-butanol or 2-butoxyethanol and m is about 7 or 38 respectively. This mixing scheme boundary is sometimes called the “Koga line”. Similar conclusions are drawn for electrolytes. Some ions (e.g. acetate) act as a hydrophobe, while others as a hydrophile (e.g. Br −  I−  ClO4 − , and SCN− ). In addition, some other ions (e.g. Na+  NH4 +  F− , and Cl− ) are the hydration centers and they hydrate a number of H2 O molecules. The bulk H2 O away from the hydration shell, however, is not affected by the presence of this type of ions. We thus learned that there are basically three kinds of effects of solute on the molecular organization of H2 O, as long as the solution is within Mixing Scheme I. In relation to the Hofmeister ranking of anions, kosmotropes (salting out, stabilizing agents) are either hydration centers or hydrophobes. Within hydration centers, a stronger kosmotrope hydrates a larger number of H2 O. Chaotropic ions, on the other hand, are hydrophiles; the stronger its hydrophilicity, the more chaotropic the ion. In ternary (and multi-component) solutions within Mixing Scheme I, the solute– solute interactions are also bulk H2 O-mediated, and reflect how each solute modifies H2 O. Just as for an amphiphile described above, the net interaction, whether it is homogeneous (inter-same-species) or heterogeneous (inter-different-species), is the result of competitive or cooperative effects of the combination of the above three kinds, depending on the solutes in question. The competition or cooperation between two different species does not seem linearly additive, however. All of these findings by the present thermodynamic studies are completely model-free. They are our natural interpretation of the behavior of the second and third derivatives of G that are determined experimentally and/or numerical differentiation without resorting to any fitting function. These findings based on higher order derivatives than hitherto utilized provide the molecular understanding equal, if not superior in some cases, to those elucidated by modern spectroscopic techniques.

Appendix A

GRAPHICAL DIFFERENTIATION BY MEANS OF B-SPLINE

In computer graphics, it is common to draw a smooth curve through data points using the B-spline method. We have been doing manual graphing using a flexible ruler followed by reading the data off the smooth curve drawn at a given interval. These data are used to approximate the slope of the mid point. The reliability of such an approximation has been discussed elsewhere at some length (Parsons et al., 2001;Koga, 2004a). The advantage of manual graphing is that human judgment in the quality of each data point is readily applicable. Fig. A-1 shows the enthalpic interaction between 1-propanol (1P) E molecules, H1P−1P , in binary 1P–H2 O (Koga et al., 2006). This is the best data set to our knowledge. Also shown in Fig. A-1 are the results by B-spline on the same data set. As is evident, the standard cubic B-spline and our manual graphing provide the same

350 300

Manual B-spline

H 1P-1PE / kJ mol-1

250 200 150 100 50 0 -50 0.00

0.02

0.04

0.06

0.08

0.10

0.12

0.14

0.16

x1P

Fig. A-1. Cubic B-spline vs. manual graphing. A satisfactory case. Note, however, the two very bad points by B-spline at x1P ≈ 012. Reproduced with permission from the Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

268

Fig. A-2

160 140

H 1P-1PE / kJ mol-1

120 100 80 60 40

Manual

20

B-spline

0 -20 0.00

0.01

0.02

0.03

0.04

0.05

0.06

0.07

0.08

x1P

Fig. A-2. Cubic B-spline vs. manual graphing. An unsatisfactory case. See text. Reproduced with permission from the Bulletin of the Chemical Society of Japan, 79, 1347, (Koga et al.), Copyright (2006), the Chemical Society of Japan.

E H1P−1P plots within an acceptable uncertainty. It is apparent from the figure, however, that in the region where the slopes are close to zero, i.e. x1P >008, the resulting slopes by B-spline are unstable. Note the two bad points at about x1P ≈ 012. Another data set (Koga et al., 2006), in which a number of data points were known to be taken at a temperature higher than the set point by an unknown degree, were subjected to the cubic B-spline procedure. The results are shown in Fig. A-2 together with the manual graphing method applying human judgment as to how much to correct for these data points. There are more recent advancements in B-spline which allow uncertainties in both abscissa and ordinate and also flexible boundaries (Nunhez et al., 2004). However, for this newest method to work, the curve to be interpolated must not contain an inflection point. This is not suitable for the present purpose, since we seek an inflection point in our data. Further developments in a more robust computer interpolation technique are awaited, in view of the tediousness of manual graphing and subsequent reading of the data off the smooth curve drawn.

Eq. (B-1) to (B-5)

Appendix B

GIBBS-KONOVALOV CORRECTION

Consider the equilibrium in a binary system (B, W) between the gas and the liquid phases (see Fig. II-2). The thermodynamic quantities for the gas phase are with prime  , and those for liquid are without. Since the system is in equilibrium, B  = B

and

W  =  W 

(B-1)

The independent variables in the liquid phase are p T xB . xB  for the gas phase with common p and T . Consider now a small variation in chemical potential due to small changes in the independent variables along the equilibrium line. Recalling eq. (I-36), we   i write the resulting variation for the gas phase, d , instead of dG in eq.(I-36) as, T 



i d T

V H 1 = − 2i dT + i dp + T T T



 i  dxB   xB 

(B-2)

recalling also eq. (I-39) and eq. (I-40). Here i = B or W. Similarly for the liquid phase, d

  i

T

H 1 V = − 2i dT + i dp + T T T



 i dxB  xB

(B-3)

Since we consider the above variation along the equilibrium line,  d

i  T

 =d

  i

T



(B-4)

for i = B and W. Now we wish to keep the composition of the liquid phase constant, i.e. dxB = 0. Since the partial molar volume of gas phase is much larger than that of the liquid mixture, we can ignore the latter, i.e. Vi  Vi ≈ 0. Furthermore, we assume that the gas phase mixture is ideal in comparison with the degree of non-ideality of the liquid mixture. Therefore, following eq. (II-6), Vi =

RT  p

(B-5)

270

and using eq. (II-11), 

B  xB 



RT =  xB

 and

W  xB 

 =−

RT  1 − xB  

(B-6)

By virtue of eqs. (2), (3), (4), (5) and (6), we write two equations for B and W as, 

 hB R R dT = dp +  dxB   T2 p xB   hW R R dT = dp − dx   T2 p 1 − xB   B hi = Hi − Hi = vap Hi∗ − HiE 

for i = B and W

(B-7)

(B-8) (B-9)

where vap Hi∗ is the heat of vaporization of pure i, and HiE is the excess partial molar enthalpy of i. By multiplying eq. (7) by xB  and eq. (8) by 1 − xB   and by adding them in order to eliminate dxB  , we obtain, 

dp dT

 = xB

pB hB + pW hW   RT 2

(B-10)

This equation provides the means to correct for temperature variation in determining the total pressure by the method described in Section [III-5]. While the temperature of the bath is controlled to within ±0001 K during the measurement, the day-to-day variation of ±002 K appears inevitable. To correct for the latter temperature variation, we first calculate pB and pW by the Boissonnas method (Section [III-6]), ignoring the temperature variation. We use these values to correct for the total pressure, p, by eq. (10) at the fixed mole fraction in the liquid phase, xB , using HiE separately determined. Using these corrected values of p, we recalculate pB and pW . We could reiterate the process until the values of pB and pW converge, but our experience suggests that one iteration appears sufficient. We note that a general treatment on the displacements along an equilibrium line, eq. (10) above being an example, is given in Prigogine & Defay (1954), though their approach is unique and non-standard.

Eq. (B-6) to (B-10)

Appendix C

HEAT CAPACITY ANOMALIES ASSOCIATED WITH PHASE TRANSITIONS – TWO LEVEL APPROXIMATION

In the literature, there is a voluminous accumulation of anomalies in heat capacity, Cp , the second derivative of G, associated with phase transitions. In most of these studies, some concurrent information is given such that the transition in question shows the changes in certain structural, spectroscopic, or other macroscopic properties. Some examples are shown in Fig. C-1. As shown in Fig. C-1(a), the antiferromagnetic transition of FeF2 is associated with a -type anomaly in Cp and its value seems to diverge at the transition point. Hence this is clearly the second order transition by Ehrenfest’s classification. While CoF2 may also show the same tendency in the figure, for MnF2 the value of Cp at the transition point appears finite at about 2.5R (R is the gas constant), and the pattern of anomaly seems smooth and round at the top (Stout & Catalano, 1955). Hence, the next derivative with respect to T would not show discontinuity and this transition is not readily classified by the Ehrenfest scheme. It is, however, the antiferromagnetic ordering, i.e. a change in a macroscopic property occurs across the transition point. Fig. C-1(b) is that of Cp associated with the ferromagnetic transition of EuO, which also does not show divergence at the Curie point, although the pattern is skewed resembling the -type (Kornblit & Ahlers, 1975), just as MnF2 above. Fig. C-1(c) shows the Cp anomaly associated with the changes in the twist angle of the benzene ring in p-quaterphenyl, which also has the finite value of about the same size of that for EuO, but the Cp pattern is almost symmetric (Saito et al., 1985). A similar change in the molecular conformation in biphenyl (Saito et al., 1988) results in the Cp anomaly with a small round symmetric hump of about 0.1R at 40 K (Fig. C-1(d)), while at 17 K the lock-in phase transition gives an extremely small symmetric Cp anomaly as shown in Fig. C-1(e). Fig. C-1(f) is a more recent example associated with the superconducting transition of -BEDT-TTF2 CuNCS2 (Yamashita et al., 2005). Thus, it seems reasonable to state that phase transitions are associated with anomalies in Cp , the second derivative of G, no matter how small or smooth they might be. Hence the anomalies in the third derivative quantities observed throughout this book are associated with more subtle changes than “phase” transitions.

272

Fig. C-1(a) to C-1(d)

60

Cp (elect.) / J mol-1 K-1

50

(a)

40

30

FeF2 MnF2

20

CoF2 10

0

0

20

40

60

80

100

120

140

160

T/K

Fig. C-1(a). The electronic part of heat capacity anomaly associated with the antiferromagnetic phase transitions. The data listed in Stout & Catalano (1955) were used to draw the graph. Reproduced with permission from the Journal of Chemical Physics, 23, 2013, (Stout, J. and Catalano, E.), Copyright (1955), American Institute of Physics.

50

Cp (J mole-1 K-1)

SAMPLE I SAMPLE II

40

30

65

67

69

71

T(K)

Fig. C-1(b). Heat capacity anomaly near the Curie point of EuO. Reproduced with permission from the Physical Reviews B, 11, 2678, (Kohnblit, A. & Ahlers, G.), Copyright (1975), American Physical Society.

273

15

ΔCp / J K-1 mol-1

(c)

10

5

0

160

180

200

220

240

260

280

300

T/K

Fig. C-1(c). Anomalous part of heat capacity associated with the molecular conformational change in p-quaterphenyl. The data listed in Saito et al. (1985) were used to draw the graph. Reproduced from the Journal of Chemical Thermodynamics, 17, 539, (Saito et al.), Copyright (1985), with permission from Elsevier.

ΔCp / J K-1 mol-1

0.6

0.4

1% of Cp

0.2

0

30

40

50

T/K

Fig. C-1(d). The excess heat capacity anomaly due to the twist transition of biphenyl. Reproduced with permission from Bulletin of the Chemical Society of Japan, 61, 679, (Saito et al.), Copyright (1988), the Chemical Society of Japan.

ΔCp

Fig. C-1(e) to C-1(f)

J K-1 mol-1

274

+0.1 0

1% of Cp

-0.1

θD / K

134

132

130

14

15

16

17

18

19

T/K

Fig. C-1(e). The excess heat capacity anomaly due to the lock-in transition of biphenyl. The lower part of the graph shows the Debye characteristic temperature, the deviation from the smooth curve of which was converted to the excess heat capacity. Reproduced with permission from Bulletin of the Chemical Society of Japan, 61, 679, (Saito et al.), Copyright (1988), the Chemical Society of Japan.

0.8

Cooling rate slow

ΔCp / J K-1 mol-1

fast 0.6

0.4

0.2

0.0 8

9

10

T/K

Fig. C-1(f). The heat capacity anomaly associated with the superconductive phase transition of a single crystal of -BEDT-TTF2 CuNCS2 determined by Yamashita et al. (2005). I thank Professor Nakazawa for providing the raw data from which the above graph was drawn. Reproduced from the Thermochimica Acta, 431, 123, (Yamashita, et al.), Copyright (2005), with permission of Elsevier.

Eq. (C-1) to (C-3)

275

Here we show that a variety of Cp anomalies associated with phase transitions can be understood even within the mean-field approximation by considering the degeneracy difference between the ground and the excited states in the two-level thermodynamic treatment originated by Strassler & Kittel (1965). They consider the system consisting of a large number, N , of subsystems that could be in the ground state A or the excited state B. As T increases, the population density, , of B increases cooperatively. The cooperativity is taken into account only in the energy of excitation such that the energy gap decreases in proportion to the square of , i.e. in the mean-field approximation. However, they stressed the importance of the difference in the degeneracy between the two states in relation to the order of phase transition. We applied this idea and calculated the Cp anomalies. The results showed surprisingly rich features depending crucially on the degeneracy difference, or the standard entropy difference of the two states (Koga, 1975 & 1978; Takai et al, 1994). Here we reiterate our earlier treatment (Koga, 1978). Consider a system of N subsystems which are either in the ground state A or the excited one B. If the subsystems interact with each other (which of course is necessary for the system to exhibit any phase transition), the partition function, Q, is written as, Q=

N! N N q Aq BQ  NA !NB ! A B c

(C-1)

Here NA and NB are the number of subsystems in state A and B respectively, and N = NA +NB . qA and qB denote the subsystem partition functions of A and B. The system free energy for a mole of subsystems, Fm , is calculated using eq. (1) as, Fm = −RT ln Q, which yields, Fm = 1 −  fA∗ + RT ln1 −  +  fB∗ + RT ln  + f EX

(C-2)

where  = NB /N fA∗ = −RT ln qA fB∗ = −RT ln qB , and f EX = −RT ln Qc . fA∗ and fB∗ are the free energy of pure A and pure B respectively, and f EX is the excess free energy due to inter-subsystem interactions. The configurational partition function is written as in the mean-field approximation,   Qc = exp− ij /RT 

  where ij is the average nearest neighbor interaction energy. If for simplicity, the nearest neighbor subsystems have the attractive interaction, BB < 0, only when both  of them are in the excited state, then ij can be expressed as,   Nm k ij = f EX = 2 = − 2

2 BB 2

(C-3)

276

Fig. C-2 to C-3

where m is the number of the nearest neighbors and k = −Nm BB > 0. We define the energy gap ∗ and the degeneracy of the ground and the excited states gA and gB , hence fB∗ − fA∗ = ∗ − Ts∗

(C-4)

where s∗ = R lngB /gA . Now by minimizing eq. (2) with respect to , using eq. (3) and eq. (4),  0 = ∗ − k − RT

  s∗ − ln  R 1−

(C-5)

Eq. (5) provides the value of  given ∗ , s∗ and k at a required temperature T , and the thermodynamics of the system is fully described. In particular, the total molar energy for the system, Em , and the molar heat capacity, Cm , can be written as, Em = ∗  − k/22

(C-6)

Cm = ∗ − k / T 

(C-7)

The solution of  in eq. (5) is the intersection of the following two curves in the  − z plane. 

1 −    k ∗ s∗ z= − +  RT k R z = ln

(C-8)

(C-9)

Fig. C-2 shows the curve of eq. (8) against 0 <  < 1. Eq. (9) is a straight line with the slope k/RT with the fixed point W ∗ /k s∗ /R. Here ∗ /k > 05 is assumed, otherwise the roles of state A and B are reversed. If the values of ∗ s∗ and k are such that the fixed point W is above the tangent PQR to the curve of eq. (8) at point Q(0.5, 0), the system displays the characteristic first order phase transition. Namely, at T = 0, the stable solution is  = 0, and all the subsystems are at the ground state. As temperature increases, the value of  increases and that of heat capacity, Cm , changes according to eq. (7). At a certain temperature, three solutions for  exist as shown by a broken line in Fig. 2. However, the lowest value gives the smallest free energy, until T = T ∗ at which the straight line of eq. (9) cuts point Q(0.5, 0). At this temperature, the values of Fm become the same at the two end solutions 1 and 2 , and the first order transition associated with the jump from 1 to 2 occurs. Thus T ∗ is equal to the first order phase transition temperature. When point W is on the line segment QR, the second order phase transition is apparent in the Ehrenfest classification. Fig. C-3 shows the three cases with fixed point

Eq. (C-4) to (C-9)

277

10

W

T = infinity

ρ2

5

z

Q

T

=

R

T*

0

T=0

P

ρ1

-5

-10 0.0

0.5

ρ

1.0

1.5

Fig. C-2. Graphical solution of eq. (5). Line PQR is the tangent to eq. (8) at point Q(0.5, 0). See text. Reproduced with permission from the Collective Phenomena, 3, 1, (Koga, Y.), Copyright (1978), Taylor & Francis.

14 12

2(a)

10

2(b)

Cm /R

2(c) 8 6 4 2 0 0.0

0.5

1.0

1.5

T / T*

Fig. C-3. The molar heat capacity anomaly associated with the second order phase transition within the mean-field two state approximation for cases 2(a), 2(b), and 2(c). See Table C-1 for the chosen parameters. Reproduced with permission from the Collective Phenomena, 3, 1, (Koga, Y.), Copyright (1978), Taylor & Francis.

278

Fig. C-4

Table C-1. The chosen values for k/RT∗ , ∗ /k, and s∗ /R for the examples of the second and the higher(?) order phase transitions. Reproduced with permission from the Collective Phenomena, 3, 1, (Koga, Y.), Copyright (1978), Taylor & Francis. k/RT ∗

Case

Second order transitions 2(a) 4 2(b) 4 2(c) 4 Higher order(?) transitions 3(d) 25 3(e) 38 3(f) 2

6

∗ /k

s∗ /R

05 06 14

0 04 36

15 06 06

25 038 02

3(d) 3(e)

Cm / R

3(f) 4

2

0 0.0

0.5

1.0

1.5

T / T*

Fig. C-4. The heat capacity anomaly associated with the higher (?) order phase transition within the mean-field two state approximation for cases 3(d), 3(e) and 3(f). See Table C-1 for the chosen parameters. Reproduced with permission from the Collective Phenomena, 3, 1, (Koga, Y.), Copyright (1978), Taylor & Francis.

279

W at (0.5, 0), (0.6, 0.4) and (1.4, 3.6). All these are on line QR. Case 2(a) with W (0.5, 0), i.e. s∗ = 0 or gA = gB , corresponds to the Bragg–Williams approximation for the order–disorder transition, the Weiss–Heisenberg theory for spin 1/2 ferromagnets, and other equivalent theories with the classical mean-field approximation without taking into account of the degeneracy difference between the ground and the excited states. As is evident from the figure, Cp shows discontinuity but does not diverge. If we relax the condition of the iso-degeneracy, the -type Cp anomaly is recovered (see case 2(b)). If we make s∗ to be large, then the Cp pattern appears almost symmetric (case 2(c)). In these cases also, T ∗ is identified as the transition temperature. Thus, in addition to the above requirement that point W lies on line QR, k/RT∗ is equal to the slope of the tangent PQR, i.e. 4. Table C-1 lists the values for the examples of the second order phase transitions shown in Fig. 3. We note in passing that a required condition for the system to exhibit the first order phase transition is k/RT ∗ > 4 (Strassler & Kittel, 1965). When the fixed point W lies below line QR but above the abscissa ∗ /k > 0, the Cp anomaly does not show discontinuity. The Cp pattern appears as a smooth and round hump and its size and shape depends crucially on the values of k ∗ /k and s∗ as shown in Fig. C-4. See Table 1 for the chosen values for the three cases shown in Fig. 4. In this case T ∗ does not coincide with the peak position and the Ehrenfest classification is not at all applicable. Thus a variety of patterns and sizes in Cp anomalies associated with phase transitions, only a small number of examples of which are shown in Fig. 1, could be understood even within the mean-field approximation by additionally taking into consideration the differences in degeneracy of the ground and the excited states.

Eq. (D-1) to (D-8)

Appendix D

FREEZING POINT DEPRESSION

Consider a dilute solution of B in solvent W at equilibrium with the pure solid W. Imagine the situation in which a minute amount of pure W precipitates on cooling the solution. This occurs at T . The pure solid W is in equilibrium with the solution. Hence, S

L

 W = W  L

(D-1)

S

where W and W are the chemical potentials of W in the liquid and the solid phases L respectively (see eq. (II-31)). Following eq. (II-12), W is written as, ∗L

L

W = W + RT ln xW + EW 

(D-2) ∗L

where xW is the mole fraction of W in the liquid phase, and W is the chemical potential of pure liquid W. EW is the excess chemical potential of W due to the non-ideality of the system, and depends on the nature of inter-molecular interactions in solution B–W. Since we take the symmetric reference system (see Section [II-1]), EW = 0 at xW = 1. For the solid phase, the pure solid W precipitates separately from the mother liquor. Hence ∗S

S

W =  W  ∗S

where W

(D-3)

is the chemical potential of the pure solid W. Eqs. (1), (2) and (3) yield, ∗L

∗S

− ln xW − EW /RT = W − W /RT = fus HW∗ − Tfus SW∗ /RT

(D-4)

with ∗L

W

∗S

∗L

= HW

∗L

− TSW 

∗S

∗S

∗L

− HW 

W = HW − TSW  fus HW∗ = HW

∗L

∗S

∗S

fus SW∗ = SW − SW

(D-5) (D-6) (D-7) (D-8)

282

Generally, both heat and entropy of fusion of pure W are positive, i.e. fus HW∗ > 0 and fus SW∗ > 0. Consider now the equilibrium between the pure liquid W and the pure solid W. This occurs at the freezing (or melting) point of pure W, T ∗ . Since EW = 0 at xW = 1, eq. (4) yields, 0 = fus HW∗ − T ∗ fus SW∗ /RT ∗ 

(D-9)

We assumed here that the values of fus HW∗ and fus SW∗ do not change from T to T ∗ , which is an acceptable assumption since T ≈ T ∗ . Due to eq. (9), eq. (4) is rewritten as, − ln xW − EW /RT = fus HW∗ − Tfus HW∗ /T ∗ /RT or, T ∗ − T = −RT ∗ 2 /fus HW∗ ln xW + EW /RT 

(D-10)

where we used a numerical approximation TT ∗ ≈ T ∗ 2 . We now take an arithmetic approximation that ln xW = ln1−xB  ≈ −xB . For example, if xW = 0 99 and xB = 0 01, then ln0 99 = −0 01005. If xW = 0 999 ln0 999 = −0 0010005. We make another more serious assumption that the solution in question is dilute enough that the solvent obeys Raoult’s law (Section [II-4]), i.e. EW = 0. Thus, eq. (10) is rewritten as, T ∗ − T  = RT ∗ 2 /fus HW∗ xB

(D-11)

Since fus HW∗ > 0 T ∗ − T  > 0, i.e. the freezing point depression by addition of a minute amount of B. Since, the proportionality factor, RT ∗ 2 /fus HW∗ , concerns only solvent W, eq.(11) indicates that it is only the quantity of solute, xB , but not the quality of B, that determines the freezing point depression as long as the solvent is fixed. Thus, the freezing point depression is called one of the “colligative properties”, together with the boiling point elevation, the vapor pressure suppression, the osmotic pressure etc. This is true only when the Raoult’s law assumption is correct, while this assumption is approximately acceptable in dilute enough solutions. As discussed in Sections [III-6] and [V-1], this does not occur even at xB ≈ 10−5 . Thus, strictly speaking, the “colligative property” is not realistic. If the solute dependent EW is non-zero in eq. (10), it should be rewritten with the activity coefficient of W, W  EW /RT = ln W , T ∗ − T  = −RT ∗ 2 /fus HW∗  ln xW W

(D-12)

Eq. (D-9) to (D-14)

283

With the osmotic coefficient, , defined as,  = 1 + ln W / ln xW 

(D-13)

eq. (12) is written as, together with the arithmetic approximation for lnxW , T ∗ − T  = −RT ∗ 2 /fus HW∗  xB  and W are now dependent, though weakly, on what the solute B is.

(D-14)

Eq. (E-1) to (E-10)

Appendix E

TITRATION CALORIMETRY WITH DILUTE TITRANT

Isothermal titration calorimetry is designed for a liquid sample of low viscosity. Thus, the method is not applicable for a solid sample. If the titrating sample is a viscous liquid, the titration becomes problematic in quantitative delivery. In such cases, a dilute solution of the sample, instead of pure liquid can be titrated. In this manner, a solid sample could also be used. The conversion of the raw data to the excess partial molar enthalpy is straightforward (Westh et al., 1998; Koga, 2003b). t Consider a small amount of an aqueous solution of B of the fixed mole fraction xB t t with r = xB /1 − xB  titrated into the titrand mixture consisting of nB of B and nW of W. Thus, the process is described as, nB of B nW of W + nB of B nW of Wt → nB + nB  of B nW + nW  of W

(E-1) t

where super-script (t) stands for the titrant whose composition is fixed at xB . The total enthalpy difference for the above process (eq. (1)), is determined as q. Hence, q = HnB + nB  nW + nW  − HnB  nW  − H t nB  nW 

(E-2)

with t

t

H t nB  nW  = nB HB + nW HW 

(E-3)

It follows then that   q HnB + nB  nW + nW  − HnB  nW + nW  h≡ = nB nB t

+

HnB  nW + nW  − HnB  nW  H t − HB − W  rnW r

(E-4)

286

We used the relationship r = nB /nW in deriving eq. (4). If nB and nW are sufficiently small, the right-hand side of eq. (4) becomes, t

h = HB +

HW H t − HB − W  r r

(E-5)

where HB and HW are the partial molar enthalpy of B and W respectively. The last two terms on the right of eq. (5) are associated with the titrant and constant. A small increment of eq. (5), h, can be written, using the Gibbs–Duhem relation, as,   xB HW  (E-6) = HB 1 − h = HB + rxW r Since we take the symmetric reference states, eq. (6) is rewritten by using the excess partial molar enthalpies, HBE = HB − HB∗ etc. as, h = HBE +

  x HWE = HBE 1 − B  rxW r

(E-7)

where HB∗ is the enthalpy of pure B, which is taken as zero. The same is true for W. Therefore, HBE = 

h  x 1− B rxW

(E-8)

It follows then, HBE =



HBE + constant

(E-9)

The integration constant on the right of eq. (9) can be determined if a value of HBE is available at one point. For ethylene glycol (EG)−H2 O, the values of HmE are available E (Koga, in the literature and they show a minimum. Its minimum value was used as HEG 2003b). Recall eq. (II-75),  E Hm + HmE  HBE = 1 − xB  xB E , is then readily calculated using HBE data obtained The B−B enthalpic interaction, HB−B by eq. (9) as,  E HB E  (E-10) HB−B = 1 − xB  xB

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INDEX

1-propanol probing methodology, 175–179, 180–202, 205–239 Activity, 25 coefficient, 25, 236, 282 Amplitude (intensity) of fluctuation, 70, 71 Anomalies in second derivatives of G, 271–9 in third derivatives of G, 102–103, 116, 143, 155–70, 265 Azeotropy, 48–9, 65, 66 Boissonnas method, 49, 60–5, 66, 241, 244–5, 256, 270 Chemical potential of the i-th component, 13, 17, 24, 25, 26, 58 Coarse grain, 69–78, 84, 98, 118, 149 Colligative property, 282 Concentration fluctuation, 22, 48–9, 67–9, 84–5, 130, 134 Dielectric relaxation spectroscopy, 147, 220, 238 Diffusional stability criterion, 22, 41, 43 Ehrenfest’s classification of phase transition, 271 Entropy-Giles’ derivation, 11–13 Excess chemical potential of the i-th component, 24, 25, 58–60 Excess partial molar enthalpy, 27, 51, 53, 54–7, 93, 106, 135, 168, 176, 188, 236, 285

Excess partial molar entropy, 58–60 Excess partial molar quantities, 19, 21, 27, 37, 69, 110 Excess partial molar volume, 58, 100, 102, 104, 108, 155, 156, 160, 167–8, 175–6, 187–8, 191 Excess quantities, 18–19 Extensity (wavelength) of fluctuation, 21, 70, 71, 72, 73 Extensive quantity, 26, 36, 38 Femto second pump-probe spectroscopy, 238 First law of thermodynamics, 9, 11 Fluctuation density, 71, 117, 119 Fluctuation functions, 69–78, 117–30 Freezing point depression – molecular level interpretation, 104, 105 Freezing point depression, 281–3 Giles’ derivation of entropy, 11–13 Gibbs-Duhem relation, 25–30, 35, 49, 60, 63, 103, 241–3, 286 Gibbs energy, 9, 14, 35, 36, 85, 150, 239, 265 Gibbs-Konovalov correction, 59, 269–70 Gibbs phase rule, 30 Heat capacity, 19, 41, 69, 70–1, 95, 98, 102, 156, 271–8 anomaly, 95, 271–8 Helmholtz energy, 14 Henry’s law, 33–5, 62 Higher order derivatives, 1, 21, 37, 151, 170, 263

296

Hofmeister series (or effects), 4, 205–206, 217, 218, 224, 237–8 Hydration center (or number), 207–11, 224–5, 227, 229–37 Hydrophilicity, 211–7, 222, 261 Hydrophobicity, 109, 113, 145–6, 151, 162, 166, 179, 183, 185, 191, 193, 195, 197, 199, 205, 214, 221, 230, 256, 261 ranking, 179, 183, 263

Mixing Scheme II, 93, 99, 102, 105, 109, 118, 130, 145, 149, 160, 162, 168, 173, 265 Mixing Scheme III, 92, 94, 134, 143, 147, 151, 171, 189

I-j interaction chemical potential, 22, 242 enthalpy, entropy, and volume, 22 Iceberg formation, 51, 93, 100, 103–5, 111, 116, 118, 151, 165, 175, 201, 265 Intensity (amplitude) of fluctuation, 70, 71 Intensive quantity, 26, 38–9, 178 Interaction function, 3, 4, 21–2, 39–41, 69, 95, 110–1, 252, 261–3 Ion pairing, 221, 234–5 Irreversibility of process, 11, 12, 13 Isobaric thermal expansivity, 19 Isothermal compressibility, 19, 41, 69, 85

Partial molar enthalpy, 18, 27, 51–4, 57–8, 93, 103, 135, 168, 176, 188, 236, 285–6 Partial molar entropy, 18, 51, 58–60 Partial molar fluctuation density, 71 Partial molar normalized fluctuation, 72 Partial molar quantity, 17–18, 21, 22 Partial molar volume, 51, 58, 84–5, 100, 102, 104, 108, 135, 155–6, 160, 167, 168, 175–6, 187, 188, 191, 269 Partial structure factor, Sxx(0), 68, 130

Kirkwood-Buff integral, 22, 84, 85–6 Koga-line, 102, 103, 104, 116, 148, 156, 158, 266 Lower critical solution temperature (LCST), 41–8 Many-body problem, 18 Mean-square fluctuation, 69–70 Mechanical stability criterion, 19, 41 Mixing entropy, 18, 35–7, 40–5 Mixing Scheme, 1, 3, 25, 37, 40, 89–179, 220–3, 261, 266 Mixing Scheme I, 91, 93, 94, 98, 102, 103, 104, 105, 109, 116, 117, 118, 130, 147–150, 179, 183, 188, 207, 221, 261, 266

Neutron diffraction, 148–9, 209, 214, 217, 238 Normalized fluctuation, 71, 72, 117, 120, 177 Order-theoretical approach, 3, 11, 13

Raoult’s law, 33–5, 62, 236, 244, 256, 282 Response functions, 1, 3, 19, 21, 69–70 Second law of thermodynamics, 9 Small angle X–ray scattering, SAXS, 67, 68, 92, 130, 156 Spinodal point, 45 Stability criteria, 10, 15, 16–17, 22, 39, 41 Symmetric reference, 25, 281, 286 Third derivatives, 1, 3, 21–2, 23, 69, 71, 80, 86, 95, 98, 100, 102, 135, 143, 165, 242, 245, 265–6 Titration calorimetry with dilute titrant, 285–6 Transition of mixing scheme from I to II, 95–8, 101, 114 from II to III, 101, 114, 146 Upper critical solution temperature (UCST), 41–8, 156, 166–8 Wavelength (extensity) of fluctuation, 3, 72–3, 78, 122, 165