Advances in Catalysis and Related Subjects, Volume 8

ADVANCES IN CATALYSIS AND RELATED SUBJECTS

VOLUME VIII

This Page Intentionally Left Blank

ADVANCES IN CATALYSIS AND RELATED SUBJECTS VOLUME VIII EDITED B Y

V. I. KOMAREWSKY

W.G.FRANKENBURG

Chicago, Zlt.

Lancasber, Pa.

E. K. RIDEAL London, England

ASSOCIATE EDITOR

PAULB. WEISZ Paulsboro, N . J .

ADVISORY BOARD

PETERJ. DEBYE Ithaca, N.Y.

W. E. GARNER Bristol, England

D. D. ELEY Nottingham, England

W. JOST Gottingen, Germany

P. H. EMMETT Pittsburgh, P a .

P. W. SELWOOD Evanston, Ill.

H. S. TAYLOR Princeton, N .J .

1956 ACADEMIC PRESS INC., PUBLISHERS NEW YORK, N.Y.

COPYRIGHT @ 1956, BY ACADEMIC PRESS INC. 125 East 23rd Street, New York 10, N.Y.

All Rights Reserved No part of this book may be reproduced in any form, by photostat, microfilm, or any other means without written permission from the publishers. Library of Congteaa Catalog Card Number: 49-7755

PRINTED IN THE UNITED STATES 06 AMERICA

CONTRIBUTORS TO VOLUME VIII J. R. COLEY,Catalysis Laboratory, Illinois Institute of Technology, Chicago, Illinois* J. H. DE BOER,Staatsmijnen in Limburg, Central Laboratory, Geleen, The Netherlands

E. E. DONATH, Koppers Company, Inc., Pittsburgh, Pennsylvania J. ARVIDHEDVALL,Chalmers University of Technology, Goteborg, Sweden

EDWINK. JONES,Universal Oil Producis Co., Des Plaines, Illinois

v. I. KOMAREWSKY, Catalysis Laboratory, Illinois Institute of Technology, Chicago, Illinois

RUDOLPH M. LAGO,Research and Development Laboratory, Socony Mobil Oil Company, Inc., Paulsboro, New Jersey G. A. MILLS,Houdry Process Corporation, Linwood, Pennsylvania

CHARLESD. PRATER,Research and Development Laboratory, Socon y Mobil Oil Company, Inc., Paulsboro, New Jersey S . W. WELLER,Houdry Process Corporation, Linwood, Pennsylvania

* Present address: Standard Oil Company of Indiana, Whiting, Indiana.

V

This Page Intentionally Left Blank

CONTENTS CONTRIBUTORS TO VOLUME VIII . . . . . . . . . . . . . . . . . . . . .

v

Current Problems of Heterogeneous Catalysis

BY J . ARVIDHEDVALL. Chalmers University of Technology. Goteborg. Sweden Adsorption Phenomena

1 -

. . DE BOER.Staatsmijnen in Limburg. Central Laboratory. Geleen. T h e Netherlands

BYJ H

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Adsorption and Catalysis . . . . . . . . . . . . . . . . . . . . .

I11. Physical Adsorption and Chemisorption . . . . . . . . . . . IV. Difficulties Experienced in Calculating Adsorption Energies . . . V. The Forces that Cause Adsorption . . . . . . . . . . . . . . VI . Cooperation among Various Forces . . . . . . . . . . . . VII . Mobility and Orientation . . . . . . . . . . . . . . . . . . . VIII . Physical Adsorption Phenomena at High Degrees of Occupation . I X . Chemisorption Phenomena a t High Degrees of Occupation . . . X . Simultaneous Adsorption of Different Molecules or Atoms . . . X I . Some Remarks on Catalysis and Chemisorption . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . .

.

. . . . . . . .

.

18 -19 . . . 20 . . . 22 . . . 29 . . . 64 . . 81 . . . 98 . . . 107 . . . 140 . . . 148 . . 150

Activation of Molecular Hydrogen by Homogeneous Catalysts

BY S. W . WELLERAND 0. A. MILLS.Houdry Process Corporation. Linwood. Pennsylvania

I . Introduction . . . . . . . . . . I1. Homogeneous Catalytic Systems . I11. Summary Discussion . . . . . . . References . . . . . . . . . . .

. . . . . . . . . . . . . . . . .

163

. . . . . . . . . . . . . . . . . T65

. . . . . . . . . . . . . . . . .

. . . . . . . . . . . . . . . . .

200 204

Catalytic Syntheses of Ketones

BY V. I . KOMAREWSKY AND J . R . COLEY.Catalysis Laboratory. Illinois Institute of Technology. Chicago. Illinois

I . Introduction I1. Development 111. Summary . . References .

. . . . . . . . . . . . . . . . . . . . . . . . . . . 207 of Ketonization Process . . . . . . . . . . . . . . . . 208 . . . . . . . . . . . . . . . . . . . . . . . . . . . 216 . . . . . . . . . . . . . . . . . . . . . . . . . . 216

Polymerization of Olefins from Cracked Gases

BY EDWINK . JONES. Universal Oil Products Co., Des Plaines. I . Introduction . . . . . . . . . . . . . . . . . . . . . . . I1. Sources of Feed Stock . . . . . . . . . . . . . . . . . . . I11. Basic Factors in Polymerization . . . . . . . . . . . . . . vii

Illinois

. . . . 219 . . . . 220 . . . . . 221

viii

CONTENTS

IV . Polymerization Reactions . . . . . . . . . . . . . . . . . . . . . . 225 V. Polymerization Processes Using Phosphoric Aoid . . . . . . . . . . . 236 VI . Materials of Construction . . . . . . . . . . . . . . . . . . . . . 237 VII . Future Outlook for Polymerization . . . . . . . . . . . . . . . . . 238 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 238 Coal-Hydrogenation Vapor-Phase Catalysts

.

BY E E. DONATH, Koppers Company, Inc., Pitbbutgh. Pennsylvania

.

I Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 239 I1. Brief History of Vapor-Phase-Catalyst Development . . . . . . . . . . 242 I11. Tungsten Disulfide Catalyst . . . . . . . . . . . . . . . . . . . . 245 IV. Nonsplitting Catalysts . . . . . . . . . . . . . . . . . . . . . . . 264 V. Splitting Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . 276 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 291 The Kinetics of the Cracking of Cumene by Silica-Alumha Catalysts

.

BYCHARLES D.PRATERA N D RUDOLPH M LAGO.Research and Development Laboratory, Socony Mobil Oil Company. Ine., Paulsboro. New Jersey

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 294 I1. The Results of Previous Studies . . . . . . . . . . . . . . . . . . . 294 I11. The Kinetics of Cumene Cracking as Determined by Differential-Reactor Studies . . . . . . . . . . . . . . . . . . . . . . . . . . . I V. Determination of Adsorption Constants with an Integral Reactor V. Coke Formation . . . . . . . . . . . . . . . . . . . . . . . VI . Discussion and Summary . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . .

. . . 305 . . . . 324 . . . 329 . . . 337 . . . 338 AUTHOR INDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 341 SUBJBCT INDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

348

Current Problems of Heterogeneous Catalysis J. ARVID HEDVALL Chalmers University of Technology, Goteborg, Sweden

Only a few decades ago the term catalysis in the sense indicated by the title was practically a collective designation for a number of chemical reactions and adsorption processes, the nature and mechanism of which were very little known. Text books dating from the turn of the century, or even a decade or so later than that, usually contain a collection of more or less unsystematic observations and patents. For a long time in the school curriculum a catalyst was defined as a substance which accelerated a reaction without participating in it-a chemical hocus pocus, indeed ! The veils shrouding the mode of operation of the catalyst were rent by the kinetic and thermodynamical investigations by Langmuir, Paneth, Polany, Volmer, Bodenstein, and Taylor. Debye’s studies of dipole effects constituted decided progress in this field, and the same may be said of Born, Franck, and London’s treatment of these problems from the point of view of wave mechanics. The literature on catalysis a t this stage ( 1 ) is entirely different from the earlier works;-among modern works may be mentioned Rideal’s “An Introduction to Surface Chemistry,” Schwab’s ‘‘ Katalyse vom Standpunkt der chemischen Kinetik,” the “Handbuch der Katalyse” in six volumes*edited by Schwab, and the “Papers Presented for a Discussion a t a Joint Meeting of Soci6t6 de Chimie et de Physique and Faraday Society ” (Bordeaux, 1947). The next steps toward a deeper understanding of the catalytic processes were possible through researches which elucidated the nature of crystal structures and the geometrical positions of electrons, ions, or atom groups and their energetic conditions and th a t also determined the factors which influence the mobility of these particles within the lattice or a t phase boundaries. Already in his early work Taylor had pointed out that particles a t surfaces, edges, corners, or points are not energetically equivalent with particles in symmetrical positions, because of their exposed and, with regard t o affinity, unsaturated positions. Taylor held the opinion that those particles had a particularly high catalytic activity (“the active centers”). It was Taylor I think who drew attention to the fact that no phenome1

2

J. ARVID HEDVALL

nological differences exist between adsorption processes and chemical reactions a t the phase boundaries; he called adsorption processes “faint chemical reactions ” ( I ) . The activity of the particles in unsaturated positions have also a purely physical and crystallographic effect. At places with high activity the recrystallization that usually is inadvantageous for catalytic processes occurs a t lower temperatures or a t constant temperatures with higher velocity than a t less active places on the surface of the solidstate catalyst. Goetz showed that there exists a block formation on the surfaces of a crystal and that in this formation the distances between particles are slightly different from those in the interior of the structure; and Cohen observed that transitions a t low temperatures can occur in such structures (2,s). There will be occasion t o return t o these phenomena in connection with the general significance of those crystal states where new phases are formed. It might be pointed out here, however, th a t far too little is . known about the structural and energetic conditions of surfaces, in spite of their fundamental importance not only for catalytic processes, but for chemical reactions in general. After all, those reactions start at surfaces. The necessity of further studies in the mechanism of formation of nuclei, of changes of modification, and of crystal growth should also be underlined, although a solid foundation for further work on these problems has been laid by the researches of several investigators, for instance Stranski and Volmer (4,s). Quite recently there was an important advance toward a deeper understanding of the nature of catalytic processes. The mobility of elect.rons a t surfaces and the exchange of electrons a t phase boundaries between catalyst and substrate were investigated and rules for these processes formulated. The liberation of electrons from the surfaces of solids through irradiation or by other means has attracted the attention of several investigators, working with different objects in view. These researches are mentioned here, because, viewed in perspective, they also emphasize the common traits in the mechanism of adsorption, reaction, photochemical, physical, and crystallographic processes, as well as those pertaining to electronics. Regularly recurring symposia and congresses in several countries, covering the physical and chemical characteristics of the solid state, give further information accentuating these common traits. Among specialists those gatherings are well known and particular references consequently unnecessary.

PROBLEMS OF HETEROGENEOUS CATALYSIS

3

I n 1943 Schwab demonstrated a connection between electron concentration and catalytic activity. Similar problems have been investigated by Hauffe and Wagner, among others (6,Y).In this connection it is necessary to mention the important ideas and investigations by Weyl on the screening effect in crystals (7a). Schwab and Suhrmann and Sachtler have proceeded further along this road and shown that coupling an electron-spending to an electronaccepting substance apparently is a fundamental condition for a process to take place, whether it is a purely chemical reaction or a catalytic process or the preIiminary stage of the cataIytic process, adsorption (8,9). These investigations seem to be a sound basis for further research on catalysis. It should be possible one day to compute the most advantageous combinations of carrier, catalyst, and substrate in order to obtain the desired course of the reaction. Here for instance are common research approaches for flotation and ion exchange processes. The preceding paragraphs were a rough outline of the gradual dispersal of the miasma veiling the catalytic processes; further, it should once more be underlined that catalysis, from the phenomenological standpoint, does not occupy any specific position in the group of processes which take place at phase boundaries. Results from recent research are, of course, extensions of earlier results concerning polarization effects in the innermost adsorption layers and consequent decrease of the energy of activation-in the cases where a reaction is facilitated. Even though the elucidation of the adsorption and reaction mechanism is close a t hand, there are still many problems of theoretical interest and great practical importance. The following paragraphs will touch on these matters. Better methods must be evolved both for determining and for exploiting the surface of the catalyst. Closely allied to these issues is the task of elaborating new methods of preventing recrystallization, which lowers the effect in the purely geometrical and, especially, in the energetic sense. This problem is undergoing intensive study, as the author had opportunity to observe at a recent visit to American research laboratories. I n part, a t least, the solution lies in introducing into the lattice such guest particles as retard the transport of matter across the boundary lines of the crystallites. It constitutes an urgent problem in the caw of both metallic and nonmetallic catalysts (10). Recrystallization is one form of aging in catalysts. There is a further form in the case of the industrially important mixed catalysts. The latter type occurs because there are gradually formed compounds or solid solutions of the components. Theoretically there is nothing t o prevent the

4

J. ARVID HEDVALL

new-formed phases from being even better catalysts than the original one, but it is a remote possibility that an improvement should occur with reference to the relevant reaction. Aging of this type is consequently often combined with a change not only in the activity of the catalyst, but also of its selectivity. One type of aging, frequently leading t o complete inactivity, is caused by poisoning. It may consist of a purely chemical change of the catalyst, but also of a decrease of the active surface, caused by the fact that substances formed a t the reaction are deadsorbed more slowly than the substrate is adsorbed.

La*

1

14’ 0

I

N

2

I

1 0 .

I

I

I

(J1

I

800

900

1000

1100

4 Temperature OC

FIG.1 . Changes in the adsorption activity of ferric oxide after heating in oxygen (curve 0 ) and in nitrogen (curve z).The quantities of ferric oxide required for adsorption of a fixed quantity of the adsorbate are shown on the abscissa. It will be seen that the adsorption capacity of a ferric oxide preparation declines more rapidly if it has been preheated in nitrogen than if it has been heated in oxygen.

These phenomena are closely connected with the mobility of so-called surface molecules or moleculoids, the mobility of which is also influenced by energetic or material disturbances in the interior of the lattice (11,12). It should be pointed out that in many cases it seems uncertain which substance it is that constitutes the veritable catalyst. Particularly for metallic and oxidic catalysts each separate case must be investigated for formation of a monomolecular layer of compounds or adsorbates, for instance, sulfides, carbides, hydrides, etc., which constitutes the real catalyst after an individual activation period of the metal or the oxide. In such cases the electron exchange between the film and the substrate will, of course, be the decisive factor (13). These phenomena underline the necessity of research on the geometric, structural, and energetic problems of the surface. Interferometric, radiographic, and X-ray methods as well as determinations of surface

PROBLEMS OF HETEROGENEOUS CATALYSIS

5

conductivity and electron emission are among the procedures th a t should be exploited. The effect of dissolved gases that are not chemically reactive in the ordinary sense on the activity of surfaces and more deeply situated layers has attracted too little attention although it is industrially important. The gases may remain in solid solution or they may be liberated again,

FIQ.2. The change in the dissolution velocity of AlzOa on the heating of kaolin in various chemically inactive gases. The effect consists of a change in the recrystallization velocity and is of interest with reference to the problem of the carrier (- - - - - - air; - - - CO; ___ cod. depending on the temperature and the substance. I n both cases there is a disturbance of the lattice structure-activating or deactivating according t o the composition of the system (14). Here we approach the group of problems concerning unstable lattice states, brought about in one way or another, problems which are common t o adsorption, catalysis, phosphorescence, luminescence, photochemistry, electronics, etc. The geometric distortions of structure, which are naturally combined with energetic changes, are of several types, e.g., "frozen-in" reversible lattice defects, incomplete crystallographic transitions, hereditary lattice imperfections, guest particles of different atomic volume or charge in the lattice, and stoichiometric deviations from ideal formulas. The significance of the latter defects for semiconductors, photo-

6

J. ARVID HEDVALL

reagents, phosphors, and luminophors is common knowledge. Their influence on the adsorption capacity, demonstrated by the author, is perhaps less well known ( 1 4 ~ ) . The influence of the other factors on the activity has been shown by the author in all conceivable cases. With reference to transformations, 20

Ii I

I

IS

10

3H 3.7

5

0

+ 100

300

c -Temperature

'C

FIG.3 A. Adsorption of fast red at different pH values as dependent on the heating temperature of pholerite.

the results have been summed up in a rule of general validity: every kind of transition process or formation of new phases corresponds to a relative maximum of reactivity (16). Figures 2 t o 6 give some illustrations of this fact. It should be pointed out that though these effects have already demonstrated their usefulness in practice, higher effects and greater selectivity may reasonably be expected through adaptations of the apparatus. The same may be said with reference to the following considerations.

PROBLEMS O F HETEROGENEOUS CATALYSIS

7

When an increased yield per unit of time of a chemical reaction in heterogeneous systems was desired, temperature, grain size, grain shape, and contact surface were altered. For catalytic reactions there was further the method of admixing guest particles to the crystal lattice or creating 4e

4c

3e

3c

25

20

15

10

5

0 -Temperature

"C

FIQ.3 B. Adsorption of methylene blue a t different pH values as dependent on the heating temperature of pholerite.

phase boundaries by purely mechanical mixture with added substances (mixed catalysts). Neither at catalysis nor at ordinary chemical reactions, except for photochemistry in the restricted sense of the term, does there appear to have been a thought that there existed other forms of energy than heat that might be altered. When the author started experiments in alteration of energy forms

8

J. ARVID HEDVALL

some twenty years ago, they were generally regarded with marked skepticism. Nobody doubted, of course, th at magnetic, electric, or irradiation states should indeed exert an influence on the activity. It was assumed, however, that those effects would be extremely small. There were ,.6 I.4

I .2

1.0

0.8

0.9 Fro. 4. The formation of SO3 with quartz as catalyst in the range of the p * a transition of quartz (575"). Ordinate: yield in % of theoretical yield.

even mathematical computations that appeared to indicate such a result. It should be pointed out in this connection th a t there is a considerable danger in trying, by starting from a lattice in a state of equilibrium, to calculate the conditions in a lattice where, from one cause or another, the 3 2 I

' 9 2

93

94

95

96

97 98 99

---.Temperature *C

FIG.5. Maximum of vulcanization velocity of rubber at the transition of sulfur (minimum of free sulfur). Ordinate: free sulfur in %.

equilibrium has been disturbed. Even the calculations of the states involving reversible defects depend more on models and ordinary temperatures or roentgeno(Frenkel, Schottky) than on real-electronographic graphic-pictures of the structure. Regarding the real structure of irreversible defects, which in practice are immensely important, there do not

PROBLEMS O F HETEROGENEOUS CATALYSIS

9

exist such investigations, with the exception of a few simple metal lattices. The difficulties naturally increase with heightened temperature and complicated lattices.

-

Preheating Temperature OC

FIG.6. The influence of the transition from anatase t o rutile on the oxidation of linseed oil when Ti02 is used as pigment.

--b

Temperature

OC

FIG.7. Abrupt changes in the catalytic decomposition of NZ0 on passing the Curie point (360") of pure nickel. The two curves refer to two experiments with different velocities of the gas currents. Ordinate: yield in % of theoretical decomposition.

To be able to regulate the activity and selectivity of the catalyst, however, is one of the most important problems in research on catalysis. It can be done only through a deeper understanding of the displacement of material particles and electrons. Results from the experiments are given here by means of a few fig-

10

J. ARVID HEDVALL

ures; these results indicate the great influence exerted on the activity by changes of nonthermal type, i.e., changes in the magnetic, electric, and irradiation state of a substance and the influence of supersonic vibrations (Figs. 7 to 11). In all cases the systems have been selected to ex-

--+

B4 T

FIQ.8. Similar change observed on kinetic measurement of the catalytic decomposition of formic acid with a Co-Pd alloy as catalyst (Curie interval 135" to 185").

I" 77.0

# I

E

a

785

76.0

640

760

700 *Temperature

OC

FIG.9. Changes in catalytic effect on loss of ferromagnetism a t the reaction HzO with cast iron as catalyst. Reaction mixture: 1 part CO, CO Hz+ CH, 3 parts Hz; gas velocity: 3.17 cm.S/min.; temperature increase: curve 1-1.15, curve 2-0.9; curve 3-0.8"/min. The early increase of the Hz content is caused by the subsidiary reaction CO H t = HzOli, C. The decreasing percentages above the Curie temperature correspond to the formation of CH,.

+

+

+

+

d u d e the possibility of crystallographical variations. Composite effects are consequently excluded (16). When effects from changes in the magnetic state were investigated, ferromagnetic substances were selected, as there is then no change in modification on passing of the Curie interval. The influence of the magnetic state was investigated in a large num-

PROBLEMS O F HETEROGENEOUS CATALYSIS

11

ber of ferromagnetic substances, Nil Fe, Co-Pd and Heusler alloys, ferrites, etc. In all cases the paramagnetic state exhibited higher catalytic effect. Several different substrates were used. This so-called “magnetocatalytic” effect was later confirmed by Forestier (17). 1.4670

1.4680 1.4690

125

135

155

145

165

175

4 Temperoture O C

FIQ.10. The change in the velocity of hydration of an unsaturated oil (cottonseed 02) at the passage of the Curie interval of E Co-Pd alloy (152 to 160”). The effect was determined by measuring the refraction index (ordinate).

The result of the kinetic investigations can be expressed as follows: the paramagnetic state has a higher activity than the ferromagnetic state, in spite of the fact that the gross value of activation is smaller in the latter case. It has been pointed out earlier that the interplay of spending and accepting electrons is an essential factor, and the above-mentioned phe-

-

(Oe)

FIG. 11. Effect on the ortho-+ para transition of hydrogen. [Cf.Justi W.1 Ordinate: magnetic field; abscissa: yield in %.

nomena probably mean that this interplay can take place more easily in the paramagnetic state. However, it must be pointed out that further research work must be carried out here, it being too early to explain these phenomenon in details. Some experiments were made recently in collaboration with Justi in order to ascertain the effect of exterior magnetic fields; an effect on the ortho + para transition of hydrogen was registered (18). Similar effects are obtained also in the case of substances with so-called “electrical Curie

12

J. ARVID HEDVALL

points” (repolarization a t certain temperature), as may be seen from Fig. 12 (16). It must u priori be assumed that every factor which influences the interplay of spending and accepting electrons at a phase boundary must also influence the processes occurring there.

1 ‘8

19

20 21 22 -Temperature

23 24 25 “C

FIG.12. The change in the dissolution velocity of Rochelle salt at the passage of the electrical Curie point (22”) of the salt. Ordinate: relative values of dissolved quantities.

Figure 13 shows the so-called ~ h o t o - u d ~ o r ~ efect ~ i o n found b y the author (16). I n the case of the substance illustrated above, and in several other substances used as adsorbens, the adsorption equilibrium is largely displaced on irradiation with absorbable light.

n L

5:

95 90

0 0 a8

95 280 c

75

V

70

a

t

65 6o0

60

I20 -Minutes

180

240

FIG.13. Regression of the adsorption, activated by irradiation, of lanasol green on the phosphor ZnS(Cu) in darkness.

Experiments have also demonstrated the influence not only on adsorption but on dissolution processes too (19). Considerable differences can be observed if the wave lengths of the irradiated light are absorbable. A colored substance will show the highest effect if the irradiation is done with complementary light. If light the same color as the substance is used, the result will be the same as in darkness. Every substance, then, has its own “chemistry” in darkness and in light. If the substance is

PROBLEMS O F HETEROGENEOUS CATALYSIS

13

white, effects will be obtained on irradiating it with ultraviolet light of suitable wave length. In such an instance the author has also demonstrated the individual sensitiveness of different crystal faces. The prismal faces of CdIz, for instance, are blackened by irradiation with ultraviolet light, but not so the base planes. CdBrz exhibits the same peculiarity, but only in the form isomorph with CdIz (20). Naturally these cases, too, can be viewed from the “electronic” aspect. Different electron concentrations and bonds-obviouslyl we may say nowadays-exert an influence to give the results indicated here.

FIG.14. Tarnishing of copper in iodine vapor. Ordinate: thickness of tarnishing film, measured interferometrically; abscissa : time in minutes. Curve 1: with influence of siipersonic vibrations (coupled to a supersonic drive, 300 kHertz); curve 2: without supersonic vibrations. Evidently these later experiments open up not only virgin fields of research, but also fields which so far have given only a poor crop of technical applications. The concepts magnetochemistry, electrochemistry and photochemistry have also acquired a further content. The results cannot be disregarded at methodical changes in catalytic processes. Finally, there should be some mention of the influence of supersonic vibrations on the tarnishing of metals. As in all previous experiments, no other factor (e.g., temperature) but the intended one was altered. The thickness of the tarnishing layers was determined interferometrically. It will be seen from the figure that the influence is of considerable magnitude, a fact that will be of some interest in the study of the mechanism

14

J. ARVID HEDVALL

regulating the formation of adsorption or tarnishing layers, problems of considerable import (21). The large number of comparatively recent effects surveyed here will surely indicate that there are many roads open for catalytic research and technical applications.

REFERENCES 1 . The work of the authors cited so far will be too well-known t o workers in the field of catalysis to necessitate quotations. A particular reference to the already mentioned “Handbuch der Katalyse” might not be misplaced, however, as there are many more contributors than those named here and many valuable contributions. 1. Goete, A., Proc. Nall. Acad. Sci. (U.S.) 16, 99 (1930); Phys. Rev. 61, 151, 159 (1937). 3. Cohen, E., 2. physik. Chem. 86, 419 (1913); ibid. 87, 409 (1914); ibid. 89, 638 (1915). 4 . Stranski, J. N., see survey by M. Straumann in “Handbuch der Katalyse” (G. M. Schwab, ed.), Vol. I, pp. 269 ff., Springer, Berlin, 1931; and a number of later works, for instance, 2.anorg. Chem. 262, 241 (1944); 2. Krist. AlO6, 287 (1943); Trans. Chalmers Univ. Technol. Gothenburg, No. 114; Ann. Physik [6] 1, 169 (1947). 6. Volmer, M., “Kinetik der Phasenbildung.” Steinkopf, Leipeig, 1939; and survey in “Einfuhrung in die Festkorperchemie” (J. A. Hedvall, ed.), pp. 135 ff. Vieweg, Braunsehweig, 1952. 6. Schwab, G. M., survey in “Dedication Volume to J. Arvid Hedvall,” pp. 533 ff. Chalmers University of Technology, Gothenburg, 1948. 7 . Hauffe, K., and Engell, H. J., 2.Elektrochem. 67, 776 (1953). 7a. Weyl, W. A., A new approach to surface chemistry and to heterogenous catalysis, Miner. Ind. E q . Station Bull. No. 67, 1951. 8. Schwab, G. M., Abstracts of Papers at the 13th International Congress of Pure and Applied Chemistry, Stockholm, 1953. 9. Suhrmann, R., 2. Elektrochem. 44, 478 (1938); Abstracts of Papers at the 13th International Congress of Pure and Applied Chemistry, Stockholm, 1953; Suhrmann, R., and Sachtler, W., 2.Nuturforach 9a, 14 (1954); Dissertation, Rraunschweig, 1952. 10. Milligan, W. O., and Weiser, H. B., J. Phys. & Colloid Chem. 62, 942 (1948). 11. Schwab, G. M., Proc. Intern. Symposium Reactivit,ily of Solids, Giileborg, p. 515 (1953) ; de Boer, J. H., Elektronenemission und Adsorptionserscheinungen.” Barth, Leipeig, 1937. 1% Gomer, R., and Smith, C. S., “Structure and Properties of Solid Surfaces.” Chicago U. P., Chicago, 1952; Read, W. T., Jr., “Dislocations in Crystals.” McGraw-Hill, New York, 1953; Massachusetts Institute of Technology Quarterly Reports on Solid State and Molecular Groups. 1% Schwab, G. M., “Dedication Volume to J. Arvid Hedvall,” p. 539. Chalmers University of Technology, Gothenburg, 1948. 14. Hedvall, J. A., “Einfuhrung in die Festkorperchemie,” pp. 211 ff. Vieweg, Braunschweig, 1952; Forestier, H., Proc. Intern. Symposium Reactivily of Solids, Goteborg, p. 41, 881 (1953). Ida. Hedvall, J. A., Arkiu Kemi 17A, No. 11 (1943).

PROBLEMS O F HETEROGENEOUS CATALYSIS

15

16. Hedvall, J. A., “Reaktionsfahigkeit fester Stoffe,” pp. 127 ff. Edwards, Ann Arbor, 1943. Hedvall, J. A., “Einfuhrung in die Festkorperchemie,” pp. 182 ff. Vieweg, Braunschweig, 1952. 16. Hedvall, J. A., “Reaktionsfiihigkeit fester Stoffe,” pp. 162 ff. Edwards, Ann Arbor, 1943; Hedvall, J. A., “Einfuhrung in die Festkorperchemie,” pp. 196 ff. Vieweg, Braunschweig, 1952. 17. Forestier, H., Proc. Intern. Symposium Reactivity of Solids, Goteborg, p. 181 (1953). 18. Justi, E., and Vieth, G., Z. Naturforsch. 8a (1935). 19. Hedvall, J. A., Z . anorg. Chem. 239, 113 (1938); Nature 143,330 (1939); Kolloid-Z. 94, 57 (1941). 20. Hedvall, J. A., and Wallgren, P., Trans. Faraday SOC.230, 697 (1940). 21. Hedvall, J. A., “Einfuhrung in die Festkorperchemie,” p. 210. Vieweg, Braunschweig, 1952.

This Page Intentionally Left Blank

Adsorption Phenomena J. H. DE BOER Staatsmijnen in Liinburg, Central Laboratory, Geleen, The Netherlands Page I. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Adsorption and Catalysis ..................................... 19 111. Physical Adsorption and isorption . . . . . . . . . . . . . IV. Difficulties Experienced in Calculating Adso 1. General Remarks.. . . . . . . . . . . . . . . . . . . . . . 2. The Structure of a Smooth Surface of an 3. The Distance between an Adsorbed Molecule and the Surface.. . . . . . . 24 4. The Influence of the Repulsion Forces.. . . . . . . . . . . . . . . . . . . . . . . . . . . . 26 V. The Forces that Cause Adsorption .................... 29 1. The Nonpolar van der Waals’ Attraction Forces.. . . . . . . . . . . . . . . . . . . 29 2. Nonpolar van der Waals’ Forces on Conducting Surfaces. . . . . . . . . . . . . 3 1 3. The Adsorption of Ions on Metal Surfaces.. . . . . . . . . . . . . . . . . 32 4. The Adsorption of Ions on Dielectric Sur 5. The Adsorption of Polar Molecules.. . . . . . . . . . . . . . . . . . . . . . 35 ielectric Adsorbent. . 37 6. The Polarization of an Adsorbed Molecule 7. The Polarization of an Adsorbed Molecule by a Conducting Adsorbent. 38 8. Chemical Bonds in Adsorption Phenomena on Metals. . . . . . a. The Formation of Ions.. . . . .................... 39 b. Sharing of Electrons.. . . . . . 9. Heats of Adsorption and Desorption nected with Chemisorption on Metals. . . . . . . . . . . . . . . . . . . . . . . 48 10. The Influence of the Electronic Structu 11. Chemical Bonds in Adsorption Phenomena on Nonmetallic Surfaces.. . 57 57 a. Carbon Monoxide Adsorption.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .......................... 59 b. Hydrogen Adsorption. . . . . . c. Oxygen Adsorption.. . . . . . . ............... d. Hydrocarbons. . . . . . . . . . . . .................. e. Hydrogen Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61 12. Active Spots. . . . . . . . . . . . . . .......................... 61 VI. Cooperation among Various Forces. . . . . . . . . . . . . . . . . . . . . . . . . 64 1. Physical Adsorption on Charcoal (and Metals). . . . . 2. The Adsorption on Ionic Surfaces .......................... 65 3. The Adsorption of Hydrogen.. . . . . . . . . . . . . . . . . . . . . . . . . . . 68 4. The Chemisorption of Oxygen.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 76 5. Optical and Other Physicochemical Changes by Adsorption.. . . . . . . . . . 79 VII. Mobility and Orientation. . . . . . . . . . . . . . . . . . . . . . 1. Mobility on Charcoal. . . . . . . . . ............................ 81 2. Orientation or Rotation., . . . . . . . . . . . . . . . . . . . . . . . 3. Hopping Molecules. . . . . . . . . ............................ 83 4. The Time of Adsorption.. . . . . . . . . . . . . . . . . . . 17

18

J. H . D E BOER

Page 91 92 7. Solution in the Adsorbent (Catalyst). , . . . . . . . . . . . . . . . . . 96 VIII. Physical Adsorption Phenomena a t High ccupation.. . . . . . 98 1. General Remarks.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 98 2. The Heat of Adsorption as a Function of the Degree of Occupation in Physical Adsorption Phenomena on Conducting Adsorbents.. . . . . . . . . 98 3. The Heat of Adsorption as a Function of the Degree of Occupation in Physical Adsorption Phenomena on Ionic Surfaces. . . . . . . . . . . . . . . . . . 100 4. Two-Dimensional Condensation and Multimolecular Adsorption. . . . . . 104 IX. Chemisorption Phenomena at High Degrees of Occupation.. . . . . . . . . . . . . 107 1. The Decrease of the Heat of Chemisorption with Increasing Degrees of Occupation.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 107 2. Factors that Cause a Surface Heterogeneity for Chemisorption. . . . . . . 109 3. Experimental Methods to Study the Hete Character of a Surface for Chemisorption. . 4. The Heats of Adsorption of Cesium Atoms o .......................................... 116 e of the Bonds at High Degrees of Occupation. . 123 6. The Decrease of the Heat of Chemisorption with Increasing Degree of Coverage for Other Adsorptives. . . . . . . . . . . . . . . . . . . . . . . . 125 7. Advantages and Disadvantages of Metal Surfaces Prepared by Different Methods. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 126 8. Other ExpIanations for the Decrease f Chemisorption with Increasing Coverage. . . . . . . . . . ............................ 128 9. Changes in Activation Energy creasing Degree of Occupation. . 131 134 10. Equations for Chemisorption Isotherms. . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Some Effects in Chemisorption Phenomena that are Connected with ...................................... 136 Activation Energies, . 12. Restricted Chemisorption Caused by the Increase of the Activation Energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 138 13. Some Final Remarks with Respect to the Decrease of the Heat of Chemisorption with Increasing Amount of Adsorbed Material. . X. Simultaneous Adsorption of Different Molecules or Atoms.. . . . . . . 1. Simultaneous Adsorption of Different Molecules in Physical Adsorption ..................... 140 2. Simultaneous Adsorption of Different Species in Chemisorption; the 141 Relative Amounts that are Chemisorbed. . . . . . . . . . . . . . . . . . . . . . . . . . 3. The Chemisorption of Different Atoms Giving Dipoles of the Same Sign 143 144 4. Contaminated Surfaces. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5. Mutual Assistance of Chemisorbed Atoms.. . . . . . . . . . . . . . . . . . . . . . . . 147 XI. Some Remarks on Catalysis and Chemisorption.. . . . . . . 1. Heat of Chemisorption and Catalysis.. . . . . . . . . . . . . . 2. Endothermic Chemisorption ................................. 149 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 150 5. Mobility and Reactivity.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6. Induced Mobility of Atoms of the Surface.. . . . . . . . . . . . . . . . . . . . . . . . .

I. INTRODUCTION The field of adsorption is usually divided into two main domains, viz., the domain of physical adsorption and the domain of chemical adsorption,

ADSORPTION PHENOMENA

19

or chemisorption. Both domains have already been treated, separately, in this series, physical adsorption by Hill ( I ) and, more recently, chemisorption by Kwan ( 2 ) . Both authors make definite statements with respect to the scope of their articles. Hill’s article gives recent aspects of the statistical thermodynamical theory of physical adsorption, which has been the field of much of his own work, and Kwan devotes a great part of his article to recent contributions by Japanese authors. We shall, therefore, try to focus our attention somewhat more closely on some other aspects in both domains of the field of adsorption. The author shall not give a detailed survey of his own work on adsorption, which started in 1925; nevertheless examples from his own work may be used t o illustrate points in the argument. The author will also take the opportunity to put forward, in this article, his own views on some aspects. These views have not been published before; some have proved their fruitfulness in the work of several of the author’s present collaborators, others may be considered to be of a tentative character. It will be clear that some “ established” views of 25 to 30 years ago may be less established now, and the author intends to concentrate attention on one or two cases where a completely new interpretation of experimental facts has revolutionized our ideas or is about to do so. 11. ADSORPTION AND CATALYSIS The present article is meant to throw some light on adsorption problems related to catalysis. It is obvious that only the so-called “heterogeneous” catalysis is concerned. Here the reaction takes place a t the interface between the catalyst (solid or liquid) and the phase which contains the reacting molecules (liquid or gas). Such a heterogeneous catalyst acts as an adsorbent for at least one of the reacting elements. Adsorption may be caused by “physical” forces, comparable with those which are responsible for liquefaction of gases, or by “chemical” forces, similar to those acting in the formation of normal chemical compounds. It is customary, consequently, to distinguish between physical adsorption (also called van der Waals’ adsorption, because of the nature of these physical cohesion forces) and chemical adsorption, or chemisorption. In some cases of heterogeneous catalysis it is the mere fact of adsorption which, by increasing the L‘concentration’)of the reacting molecules, accelerates the reaction. The expression for the rate v of a chemical reaction can be given as v =

f

a

e-E/RT.

where f is the so-called frequency factor, E the activation energy, R the gas constant, and T the absolute temperature. TC: is the product of the

20

J. H. DE BOER

concentrations Ci,each t o its appropriate power n. It is the last concentration factor which is influenced in this case. The adsorption may be either of a physical or of a chemical nature. There are examples of both kinds of adsorption leading to catalysis by this mechanism. A homogeneous catalyst cannot act in this way. I n other cases catalysis is caused by the catalyst-homogeneous or heterogeneous-taking away the heat of reaction, thus influencing the frequency factor f. A heterogeneous catalyst is advantageous in this case, because it has far more possibilities of dissipating energy than a homogeneous catalyst. The adsorption, again, may be caused by physical as well as by chemical forces. I n most cases of catalysis, however, the increase in the rate of the reaction is caused by the catalyst lowering the activation energy, El of the reaction. I n order to do this, the catalyst, by entering into a chemical combination with one of the reactants, must alter the properties of its molecules. I n cases of homogeneous catalysis addition-compounds are formed between the catalyst and one of the reacting elements, or there is an exchange of an electron between the catalyst and the molecule which is influenced. I n heterogeneous catalysis the same happens. When a molecule of one of the reactants enters into a chemical combination with the surface of the catalyst (which means with one or more of the constituent atoms of the surface or with the surface as a whole), it is chemisorbed. There is, in chemisorption phenomena, very often an exchange of an electron between the adsorbed molecule-which may or may not break up into smaller units (atoms, ions, or radicals)-and the adsorbent. I n other cases of chemisorption the adsorbed atom or radical shares electrons with one of the constituent atoms of the adsorbing surface. One might be inclined to think that only chemisorption would lead to a n enhanced reactivity in this type of catalysis. Adsorption by physical forces only tends t o lower the reactivity of the adsorbed molecules. It is, however, difficult to give such definitions of physical adsorption and chemisorption that the fields are clearly separated. We shall, therefore, discuss some of the differences and similarities between these two kinds of adsorption.

111. PHYSICAL ADSORPTION AND CHEMISORPTION

It is in many cases difficult to decide whether a certain adsorption phenomenon belongs to the physical adsorption type or whether it is a case of chemisorption. If we define physical adsorption as the phenomenon which occurs when the molecules are bound t o the surface of the adsorbent by van der Waals’ cohesion forces in the widest sense of the word, hence

ADSORPTION PHENOMENA

21

including quadrupole, permanent dipole, and induced dipole attraction, and if we define chemisorption as the phenomenon which occurs when the binding of the molecules to the surface is caused by an exchange or a sharing of electrons, we may have a distinguishing principle, but we certainly have not an easy means of analysis. There is sometimes a tendency t o believe that the heat of adsorption gives a n indication of the kind of adsorption. The forces acting in physical adsorption phenomena are the same as those which cause liquefaction. It may therefore be expected that the heats of adsorption will be of the same order of magnitude as the heats of liquefaction of gases, The forces which are responsible for chemisorption are the same as those leading t o chemical combinations, and one expects the heats of chemisorption t o be of the same magnitude as the heats of formation of these compounds. There is consequently a tendency to believe that the heats of adsorption in physical adsorption phenomena do not include high values and th a t chemisorption is characterized by high heats of adsorption. True, the heat of adsorption in physical adsorption phenomena does, generally speaking, not include very high values. Figures of 20 kcal./mole or higher do, however, occur. I n many cases the heat of adsorption in chemisorption phenomena is high; the heat of adsorption of oxygen on some metals may be of the order of magnitude of a few hundred kilocalories/mole. On the other hand the heat of adsorption in other cases of chemisorption may even be negative, as is the case with the negative heat of formation of endothermic compounds. Chemisorption is often identified with the so-called “activated ” adsorption, where the rate of adsorption is governed by an activation energy. We shall have t o conclude in our following discussions th a t this criterion is not valid. One might suppose that in chemisorption phenomena the adsorbed atoms or molecules occupy fixed places on the surface and that physically adsorbed molecules show some freedom of movement over the surface. We shall see, however, that “mobile” adsorption and adsorption on definite adsorption sites may occur in both cases. The absorption of light by certain molecules is sometimes drastically changed when they are adsorbed. This phenomenon undoubtedly has some relation t o the nature of the binding forces and one might be inclined t o consider it as chemisorption. We shall be compelled in our following discussions t o shatter this conception. It is, of course, in many cases quite easy to decide whether we are dealing with chemisorption forces or with forces of a physical nature only. There are, however, many other cases where the decision will be less obvio w . There is no sharp boundary between the two domains of adsorption.

22

J. H. DE BOER

IV. DIFFICULTIES EXPERIENCED IN CALCULATING ADSORPTION ENERGIES

I. General Remarks If the forces causing adsorption were known with a suficiently high degree of accuracy and if their dependence on the distance of the adsorbed atom or molecule from the adsorbing surface were also exactly known, it would, theoretically, be possible to calculate the change in potential energy accompanying the act of adsorption. Though we know the exact laws of the mutual attraction of two ions and though we have also a fair knowledge of some other forces which may participate in adsorption phenomena, there are other forces, especially the repulsion forces, of which the magnitude and the dependence on distance are hardly known. It is true that all molecular and atomic forces ultimately find their root in the mutual behavior of the constituent parts of the atoms, viz., the nuclei and the electrons. They may theoretically all be derived from the fundamental wave equations. It is, however, convenient, as in other branches of physics and chemistry, t o treat the various forms of mutual interaction of atoms as different forces, acting independently. We shall therefore follow the usual procedure and treat such forces as the nonpolar van der Waals (dispersion) forces, the forces of the electrostatic polarization of atoms or molecules by ions or by dipoles, the mutual attraction or repulsion Coulomb forces of ions and of dipoles, the exchange forces leading t o covalent bonds, the repulsion forces due to interpenetration of electronic clouds, together with the Pauli principle, etc., all as different, independently acting forces. I n all cases a t least two of these forces act simultaneously; a summation of their contributions t o the adsorption energy over all participant atoms has t o be made for each of the participating forces separately. Such a summation can in principle be applied with any desired degree of accuracy t o some of the forces mentioned, provided that the distances between the participating atoms are exactly known. The application of this procedure to calculating or estimating the magnitude of the contribution toward the energy of adsorption due t o these forces cannot lead, however, t o a high degree of accuracy, because we have, first, hardly any knowledge of the real structure of the adsorbing surface (even of a smooth surface) and, second, we know little about the real distance of the adsorbed atom from the surface. We shall deal first with these two serious setbacks (3).

ADSORPTION PHENOMENA

23

2. The Structure of a Smooth Surface of an Adsorbent

Usually in theoretical calculations of adsorption energies the surface of the adsorbent is idealized as if it were obtained by cutting a crystal into two halves with an ideally sharp razor blade. It is then generally assumed t ha t the atoms (or ions) of these freshly created surfaces do not alter their positions. From specular reflection experiments with molecular beams of hydrogen or helium ( 4 ) it may be concluded that cleavage surfaces of LiF or NaCl are very smooth surfaces indeed, the inequalities being caused only by the temperature movement and amounting to the order of magnitude of cm. Diffraction spectra obtained with helium beams ( 5 ) , indicate moreover (6) that the distance between two fluorine ions in the outer layer of LiF is exactly the same as in the crystal. The distances, however, from the ions of the outer layer to those of the second layer may have undergone material alterations. I n a theoretical study Verwey ( 7 ) came to the conclusion th a t the negative surface ions of the free surfaces of alkali halide crystals are generally displaced, so that their distances from the next layers of the lattice are increased, and the positive surface ions are displaced toward the inside of the lattice. The numerical results of his calculations for NaCl are t ha t the distance between the sodium ions of the outer layer and the ions of the second layer is 2.66 8.;whereas the chloride ions of the outer layer are located a t a distance of 2;86 8. from the second layer, the normal distance in the lattice being 2.81 A. The electrical double layer formed by the negative chloride ions, being located in a plane 0.20 A. distant from the plane of the positive sodium ions, is almost compensated for by the effect of the dipoles set up in the negative ions. Displacement of the outer ions and polarization of the negative ions have been used by various authors in attempts to calculate the surface energy of NaC1. The various figures obtained in these calculations vary from 77 erg/cm.2 to 155 erg/cm.2 The latter figure, calculated by Shuttleworth (8), is obtained from a model which included surface distortion, while the influence of the van der Waals forces was also accounted for. Experimental data for the surface energy of NaC1, are, however, mostly higher than the theoretical figures. The most recent experimental figure of 305 erg/cm.2 a t 25°C. b y Benson and Benson (9) is based on the cubic face, as are the theoretical figures. We may, from the rather scarce theoretical and experimental evidence, conclude t ha t the distance between like atoms (ions) in the surface layers will probably be the same as in the lattice itself. The distances between

24

J. H. DE 130ER

different ions will probably be materially different from those in the lattice, and also the distances of all atoms (ions) of the surface layer from those of the second layer are different from the distance in the lattice. We may expect the atoms of the outer layer of a metal surface to be somewhat less close to the second layer than corresponding to the distances in the metal lattice. The edge atoms of the two-dimensional layers constituting the lattice of graphite (charcoal) may be expected to lie a t a somewhat smaller distance from their neighboring atoms than the normal C-C-distance in graphite (I 0). Considering the fact that both single (gaseous) molecules of alkali halides and crystals of alkali halides consist of ions, we may conclude that the outer layers of the crystals also consist virtually of ions. The same may be true for crystals of the alkaline earth halides and of some fluorides of other metals, such as lead fluoride, or of some complex fluorides such as potassium zirconium fluoride (K2ZrF6)(11). Several oxides, such as the alkaline earth oxides, form in their crystalline state lattices consisting of ions. The single molecules of these oxides, however, must be considered as having largely a homopolar character (12). It is, therefore, not certain that the constituent atoms of the outside layers of the crystals of these oxides are in the same ionic state as the ions inside the crystal; a more homopolar character is certainly possible. The same holds for many halides of heavy metals, such as AgCl and AgI. When these halides are suspended in water (colloidal state), the binding may well be of a more ionic character, owing to the surface hydration. 3. The Distance between an Adsorbed Molecule and the Surface

A more serious setback in calculating adsorption energies is our complete lack of knowledge of the real distance of the adsorbed atom or molecule from the adsorbing surface. Unfortunately most of the forces which cause adsorption greatly depend on this distance. As a matter of fact, the equilibrium distance between the adsorbed atom and the surface is determined by the equilibrium of all attraction and repulsion forces acting on the atom. Calculations of lattice energies and of sublimation energies from the various simultaneously acting forces have been very successful, because the distances between the ions or molecules in the crystals are well known from other sources. I n adsorption phenomena the distance between the adsorbed atom or molecule and the surface is completely unknown. Very often this distance is assumed to be the sum of the “radii” of the adsorbed atom and a surface atom. These radii are then derived from experimentally determined figures of distances of the same atoms in other combinations. I n a calculation of the energy of physical adsorption the radius of the adsorbed atom or molecule is often taken from an-

25

ADSORPTION PHENOMENA

other combination of a similar physical nature (a so-called “van der Waals’ radius”). The same should be done for the radius of the surface atoms, but generally far smaller figures are used in t8hiscase. Many calculated figures of adsorption energies given in the literature which, because of their close approximation to experimental values, are said t o prove that a supposed force is responsible for the adsorption phenomenon, are in fact fallacious, because of the improbably low figures for the “distances” used in obtaining these calculated results ( I S ) .

_-_--_--

to B

FIG.1. When the adsorption force depends on the polarizability of the adsorbent, the acting distance between the adsorbed molecule and the surface is far smaller if the region of idcal polarizability starts a t the outer peripheries of thc atoms of the adsorbent (van dcr Waals’ radii) (I A ) than if this region starts at a plane through the centers of the outer atoms (I B ) .

Another difficulty is met with in adsorption on metallic surfaces. Metals, or rather conducting bodies, are considered as adsorbents with an ideal polarizability. Accepting this view as true does not make it clear whether the metallic properties leading to this ideal polarizability should be assumed t o start at the outer peripheries of the surface atoms of the metal or whether we must assume these properties to be found from & plane through the centers of the surface atoms. The choice of the outer boundary of the region of conducting electrons is very important, however, for the assessment of the “distance” of the adsorbed atom to the metal. As we shall see in Secs. V,7,8 and VI,l, the forces th a t cause adsorption on metal surfaces are exercised by the adsorbent body as a whole, rather than by the constituent atoms. It is, therefore, important t o know the

26

J. H. DE BOER

distance between the adsorbed atom and the metal surface. A glance at Fig. 1 will reveal that, depending on the assumption of where the polarizability starts, this distance may vary by a factor of about 2.

4. The InJluence of the Repulsion Forces Even if the exact structure of the adsorbing surface and also the exact distance between the adsorbed molecule and the surface were known t o us, we still should not be able to calculate the energy of adsorption, mainly because of our lack of knowledge of the repulsion forces. These repulsion forces arise from the interpenetration of the electronic clouds of the atoms. This, together with the Pauli exclusion principle, prevents the atoms from approaching each other too closely. It is extremely difficult to formulate the resulting repulsion forces in a mathematical expression. They increase strongly with decreasing distance between the atoms, but certainly not according t o a simple law. For certain cases simplified equations have proved useful. I n calculations of the lattice energies of ionic lattices of the alkali halides and the oxides of the alkaline earths, good results are obtained with a simple expression for the contribution t o the potential energy, owing to the repulsion forces, suggested by Born and Mayer (14):

Ere, = +be-r/p

(1)

where b and p are constants and r is the distance between the atoms (15,16). (The expression for this contribution to the potential energy E is given a positive sign, because it causes an increase in potential energy; the contributions to the potential energy given by attraction forces will, consequently, be given negative signs.) When, however, the .formation of single molecules of the alkali halides and their condensation to crystals is studied, expression (1) fails completely, as may be seen from Fig. 2, where the ratio of the interionic distances in the alkali halide molecules rm and in the solid alkali halide crystals r, is given as a function of r,. An earlier empirical expression

where b and n are constants, gives excellent results in this case (I?'). It transpires t ha t for the larger distances between the ions, as used in the calculations of lattice energies-integration from infinitely large distances t o the equilibrium distance in the lattice-expression (1) can be used. For smaller distances, however, it is better t o use the older expression (2) with a high value of n.

27

ADSORPTION PHENOMENA

As we shall see in the next section, the potential energy of two atoms attracting each other with nonpolar van der Waals’ forces will decrease according t o a Eaut = - (3) r6 where a is a constant and r is the distance between the atoms.

3.5

3.0

-

4.0

rc

FIG.2. Ratio r,,,/rc as a function of r,. The theoretical lines A and B are drawn with the repulsion law (2) for TZ = 12 and n = 10 respectively. Line C is drawn with repulsion law (1) and p = 0.345 A., as found from the lattice energies of alkali halides. Experimental points for several alkali halides of the NaCl type are shown a8

+.

When van der Waals’ forces and repulsion forces act together, the potential energy is given by a b E = Eattr Erep = - p 7 (4)

+

+

If n = 12, we see that the repulsion forces balance the attraction forces a t P = rm, when 6a - 12b ,‘I

-13

or

At the equilibrium distance rm the total decrease in potential energy is, therefore,

28

J. H . DE BOER

FIQ.3. E as afunction of r [Eq. (4)].Curve A gives the energy of attraction, curve R the energy of repulsion; curve W is the resultant curve.

which is only half of the value which would have been found but for the contribution of the repulsion forces (Fig. 3). Using a more general expression yields a

E=-r"+T"

b

(5)

we obtain for the decrease of potential energy at the equilibrium distance rm,

Without having a better knowledge of the laws of the repulsion forces we cannot expect to make exact calculations of the energies of adsorption. The influence of the repulsion forces is often neglected or accounted for by subtracting a fixed percentage, say 40%, from the adsorption energy as calculated with the attraction forces only (18).When adsorption is caused by van der Waals' forces only, the contribution of the repulsion

ADSORPTION PHENOMENA

29

forces is nearly completely counterbalanced by the contribution of correction terms for the attraction forces (19-21)- (See Sec. V,l.)

V. THE FORCES THAT CAUSEADSORPTION 1. T h e Nonpolar van der Waals’ Attraction Forces (22)

The attraction forces which act most frequently in physical adsorption are the nonpolar van der Waals’ forces. Since London (23) described the close connection between their nature and the cause of optical dispersion, they may also be called dispersion forces. The main contribution to the nonpolar van der Waals’ forces arises from the interaction of continually changing inducing dipoles and induced dipoles. The interaction energy of a pair of atoms due to this contribution is inversely proportional to the sixth power of the distance:

E,-

C -T6

(7)

The constant C depends on the properties of the atoms. London (23) derived the following approximation for C:

where a1and a2 are the polarizabilities of the participating atoms, h is Planck’s constant, and v 1 and v 2 are characteristic frequencies of the optical dispersion curve of the atoms. The energies hvl and hv.2 are, therefore, characteristic energies in the dispersion equation; they are often approximately equal to the ionization energies of the atoms. If, therefore, no data on the dispersion are available, the following expression may be used :

where I1 and I , are the ionization energies of the atoms. Other approximations for C have been derived; of these we shall mention only

given by Slater and Kirkwood (24), where E and m are the charge and the mass, respectively, of an electron. ar and a z , as above, are the polarizabilities of the atoms, and nl and n2are the numbers of electrons in the outer shells of the atoms. The latter equation always results in somewhat higher figures than expression (8) or (9).

30

J. H. DE BOER

Apart from the term in Eq. (7) there are other contributions toward the nonpolar van der Waals’ interaction energy arising from the interaction of continually changing quadrupoles with dipoles and quadrupoles with quadrupoles. The total expression should, therefore, be written as

where D and E are, like C, constants depending on the nature of the atoms. The contribution by the second term of this equation may amount t o from 15 t o 30% of the first term, and in some cases this contribution may be even higher than that of the first term. In numerical calculations of adsorption energies, however, expression (7) is mostly used. It is assumed that the two last terms of Eq. (11) are counterbalanced by the contribution of the repulsion forces (see Sec. IV,4). Expression (7) gives the interaction energy between two atoms. I n order t o evaluate the adsorption energy, the interaction energies of the adsorbed atom with all individual atoms of the adsorbent should be calculated and added together. This addition is allowed, as the dispersion forces are, a t a first approximation, additive. If a molecule instead of a n atom is adsorbed, the summation should be made for all atoms of the molecule. I n the latter case we may sometimes expect deviations from the additive law. Many molecules show different polarizabilities in different directions. If the position of an adsorbed molecule is fixed, the angles of its various axes of polariaability with respect t o the surface enter into the calculations (25). If, however, the molecule rotates freely, which is often the case in physical adsorption, this correction is not necessary. Instead of carrying out an elaborate calculation of all individual contributions, Polanyi and London (26) replaced the summation by a n integration :

where Na is the number of atoms of the adsorbent per cubic centimeter and ro is the shortest distance between the adsorbed atom and the surface. According to this expression the decrease in potential energy due t o the nonpolar van der Waals’ forces in the case of adsorption on a surface is inversely proportional to the third power of the distance. According t o London (27), integration is permissible only if T O >> l/(Na)VJ.I n practice the actual equilibrium distance in adsorption is always too small, and consequently expression (12) always gives too low values. A better method is to evaluate the interaction energies of the adsorbed

ADSORPTION PHENOMENA

31

atom with all the individual atoms of the adsorbent within a certain distance T D , to add all these contributions, and then to add the contributions of the rest of the atoms of the adsorbent beyond the chosen distance by integration (28). The values thus obtained turn out t o be about 2.5 times the values calculated by use of Eq. (12). The surface of an adsorbent is not smooth but shows a roughness of molecular or higher dimensions. Many catalysts used in practice are deliberately prepared to contain a great number of capillaries of submicroscopic dimensions. There are many places on the highly developed inner surface areas of such microporous adsorbents where the adsorbed molecules come into direct contact with many more atoms of the adsorbent than would be possible if the surface were an ideally smooth plane. Such places where an increased number of atoms of the adsorbent are in direct contact with the adsorbed molecules form “active places” or “active spots” for van der Waals’ adsorption (28-30). Crevices, cavities, the inside of cracks, recessed parts of the surface, and especially the inside of capillaries are all more “active” for van der Waals’ adsorption than is a plane surface. The heats of adsorption will be higher on these active spots and, therefore, the first molecules t o be adsorbed generally show a higher differential heat of adsorption. The heat of adsorption mostly tends to fall with increasing amount of adsorbed matter (see Sec. V,12). 2. Nonpolar van der Waals’ Forces on Conducting Surfaces

The adsorption by nonpolar van der Waals’ forces on metal surfaces demands a separate treatment. Many attempts have been made t o consider the metal as an ideally polarizable structure. As Margenau and Pollard (31) pointed out, there is a serious objection against such a use of the so-called “image” picture. The inducing fields of the continually changing dipoles in a nonpolar molecule change so rapidly that the conduction electrons in the metal are incapable of following their movements. With respect to van der Waals’ forces a metal behaves as a dielectric body. Margenau and Pollard describe the van der Waals’ interaction energy between an adsorbed atom and the adsorbing metal as a sum of two contributions : E, = Ei Ez (13)

+

where E l results from the polarization of the metal by the continually changing dipoles in the adsorbed atom and Ez arises from the polarization of the adsorbed atom by the electronic movements in the metal. El turns out to be positive instead of negative. This part, therefore, re-

32

J. H. DE BOER

sults in a repulsion instead of an attraction. The final expression is

where e is the charge of a n electron, a0 is the polarizability of the adsorbed atom, ro is the shortest distance of the atom t o the metal, h is Planck’s constant, ne is the number of conduction electrons per cubic centimeter in the metal, m is the mass of an electron, v o is the characteristic frequency of the adsorbed atom, C is a numerical constant approximately equal t o 2.5 and re is the radius of a sphere in the metal containing one conduction electron. We shall use this expression in estimating the nonpolar van der Waals’ interaction energy of physically adsorbed atoms or molecules on metals and on charcoal. We see that also in this expression the potential energy is inversely proportional to the third power of the distance. 3. The Adsorption of Ions on Metal Surfaces

When an ion is adsorbed on a metallic surface, the electric charge of the ion will polarize the metal in such a way that the action may be described as if an electric charge of opposite sign were formed (electric image) a t a distance below the surface equal to the distance between the actual inducing charge and the metal surface. The attraction which the adsorbed ion experiences by this phenomenon may, therefore, be described as the attraction between the ion and its image a t a distance 2 r, the distance between the ion and the surface being r . It is here that we meet the difficulty (Sec. IV,3) of not knowing where we should locate the surface of the metal, or rather the boundary of the region of conducting electrons. The image force being given by

where e is the charge of an electron and ni denotes the number of elementary charges of the ion, the contribution to the adsorption energy is

where r0 is the equilibrium distance between the ion and the metal surface. An ion adsorbed on 8 metal surface is also attracted by the normal van der Waals’ forces [Eq. (14)] and by some other minor forces, which we shall discuss presently in this section. It is, moreover, repelled by the

ADSORPTION PHENOMENA

33

repulsion forces [Eq. (2)], balancing the attraction forces at the equilibrium distance ro. When all these contributions are added together and numerically evaluated for the adsorption of a cesium ion on tungsten, the unknown value of ro can be estimated from the known value of the heat of adsorption of the Cs+-ion, viz., 54.5 kcal./mole. Then ro is found to be 1.74 A. (SZ), which seems to fit the assumption that the boundary of the region of ideal polarisability starts at a plane through the centers of the surface atoms of tungsten. The possibility of this boundary starting at the outer peripheries of the tungsten atoms is, however, not completely excluded (32).

4. The Adsorption of Ions on Dielectric Surfaces When an ion is adsorbed on the surface of a dielectric, itself consisting of ions, we may expect Coulomb forces to act between the ions of the adsorbent and the adsorbed ion. A positive ion, adsorbed on top of a negative ion of an adsorbent, is attracted by this ion, but it is repelled by the ions surrounding the one with which it is in direct contact, attracted again by the then following ions, etc. The result is a rather weak attraction. Huckel (33) derived the following equation for the electrostatic field emanating from a cubic face of the surface of an alkali halide crystal :

where e is the charge of the ions, r, is the shortest interionic distance in the crystal, and r is the distance from the surface. If an ion is adsorbed on top of a surface ion of the crystal having an opposite charge t o its own, and its distance from this ion is exactly the same as the interionic distance, re, in the crystal, the energy contribution of the Coulomb forces, resulting in the force of Eq. (17), is 2 E , = -0.0662 (18) rc

This energy is only 6.6?4, of the energy contribution by Coulomb forces which we would have found if the adsorbing ion of the crystal surface had not been surrounded by all other ions of the crystal. The collaboration of all the ions of the crystal results in relatively small energy contributions and, moreover, in attraction forces of a very short range. At a distance r = 2rc the force is negligibly small. When an ion approaches the adsorbing crystal, foIlowing a line perpendicular to the surface and ending in a surface ion of the same charge as its own, it will be repelled. The electrostatic part of the repulsion

34

J. H. DE BOER

forces is also given by Eq. (17), the force acting in opposite direction. Surface spots exactly in between surface ions, do not exercise electrostatic forces on an adsorbed ion. Therefore, a t the short distances over which the electrostatic surface forces are active they cause a periodical unhomogeneous field. The movement of a single adsorbed ion over the surface may be hampered by these variations. I n all these electrostatic considerations the surface of the ionic crystal was idealized, as described in Sec. IV,2, as if it were cut by means of an ideally sharp razor blade. Our lack of knowledge of the structural deviations of the surface arrangements with respect t o the structure inside the crystal renders it impossible for us t o make any quantitative or semiquantitative statements regarding the actual adsorption energies caused by electrostatic forces. We can only say that in most ionic crystals negative ions i.e., halide ions or oxide ions, tend t o form the outside (adsorbing) surface. We shall have an opportunity (see, for example, Sees. V,5 and VI,5) t o revert t o this phenomenon. The foregoing electrostatic calculations hold, moreover, only for positions in the middle of a cubic face of a crystal of the NaCl type. Any deviation from this situation may result in a stronger electrostatic bond. Corners and edges of crystals, other crystallographic faces, lattice disturbances, etc., may all form “active spots” where the electrostatic adsorption of ions is relatively strong. We shall return t o the problem of active spots in Sec. V,12. I n Sec. V,3 we dealt with the adsorption of ions on metallic surfaces as the problem of the polarization of an ideally polarizable structure by the ion. Dielectrics have a more restricted polarizability, the polarization resulting in the shifting of the electrons in the atoms or in groups of atoms of the dielectric to which they belong or in the mutual shifting of ions as well (34). Instead of Eq. (16), which holds for an adsorbent of ideal polarizability, we obtain for the adsorption energy contribution due t o the electrostatic induction of a dielectric:

where K is the static polarizability of the dielectric. If the polarizability of the dielectric is due only to electronic shifts and not to a displacement of ions, we have

K

=

n:

where nr is the refractive index. Because of the relation between the mean polarizability E and refraction, a = -3 . Ln 2 - 1 4rN3 nf 2

+

ADSORPTION PHENOMENA

35

we obtain

where Na is, as in Eq. (12), the number of atoms (or centers of polarization) per cubic centimeter. I n all cases, however, where K > n:, Eq. (19) must be used instead of (21). This is the case for all metallic halides and oxides. The contribution of this polarization toward the adsorption energy of ions on plane surfaces of ionic dielectrics is far more important than the contribution by the Coulomb electrostatic attraction (34). 5 . The Adsorption of Polar Molecules Molecules that have dipoles, such as organic halides, ethers, ketones, nitrocompounds, etc., will be attracted by the electrostatic field emanating from the surface of an ionic crystal. The contribution toward the adsorption energy is given by E p = -Fp (22) where F is the electrostatic field, as given by Eq. (17), and p is the dipole moment of the molecule. As the field F is not very strong and decreases very strongly with increasing distance, a significantly high contribution to the adsorption energy can be expected only, if a dipole of sufficiently high dipole moment is situated a t such a place in the polar molecule that it can approach to within a very short distance of the surface. Hydrogen bearing compounds, like water, ammonia, organic hydroxy compounds (alcohols, phenols), organic amines and acids, possess dipoles between the hydrogen atoms of the functional groups and the atoms t o which these hydrogen atoms are bound. All these dipoles are situated near the periphery of the molecules. These peripheral dipoles, moreover, all point with their positive end toward the outside of the functional group. As the ion arrangement of most ionic surfaces is such th a t the negative ions form the outer layer (Secs. IV,2, and V,4), these peripheral dipoles show a strong tendency t o take up an oriented position perpendicular to the surface, the H atom of each dipole tending tQmake a close contact with one of the negative surface ions (a halide ion or a n oxide ion) (36).Surfaces of inorganic salts and oxides have, consequentiy, a great tendency to adsorb water molecules tenaciously (36),and organic alcohols, phenols, amines, and acids are also adsorbed remarkably well. Protein surfaces have negative parts pointing outward, such as the oxygen of the CO groups, and peripheral dipoles may be adsorbed in a way similar t o the case of inorganic surfaces. Carbohydrate surfaces and protein surfaces,

36

J. H. DE BOER

moreover, possess peripheral dipoles of their own (OH groups and NH groups) and they can easily attract dipole molecules, such as water. In this case the oxygen of an adsorbed water molecule comes into direct contact with the hydrogen atom of a surface dipole (37). In later years all these types of bonds formed by peripheral dipoles have often been classified as cases of hydrogen bonding (38). Calculations of the contributions toward the adsorption energy by these forces reveal them to be quite appreciable and, indeed, a strong and oriented adsorption will always be found. These calculations result in somewhat too low values when the dipole moment I.( is used, because of the short distances between the dipoles and the negatively charged attracting centers. I n many cases the actual contribution may be 10% higher than the calculated one (39). A dipole that is not situated near the periphery of the adsorbed molecule makes a far smaller contribution toward the adsorption energy than do peripheral dipoles. Such nonperipheral dipoles may, however, cause enough difference in adsorption energy t o effect a preferential adsorption between molecules that otherwise would show equal adsorption energies. They also may lead to a fixed and an oriented position of the adsorbed molecule instead of an adsorption of more or less freely moving and rotating molecules. A similar behavior may be found when the charge distribution in the molecules is more complex. In carbon dioxide the charge distribution is of the character of a quadrupole. Lenel(40) calculated the influence of the interaction of this quadrupole with the surface of an alkali halide crystal and reached the conclusion that a substantial contribution of roughly 3 kcal./mole is to be expected from this polar interaction. Recently Drain (41) succeeded in approaching the remarkable fact that the heat of adsorption of Nz on ionic crystals is often appreciably higher than that of O2and A, which is not the case when these gases are adsorbed on nonionic surfaces. He shows that the quadrupole of N2 may be responsible. We shall return t o this problem in Sec. VI,2. When a dipole molecule is adsorbed on the surface of a metal or on other conducting surfaces (charcoal), the attraction may be described by the image force. The energy of interaction is given by

E, = when the dipole is oriented perpendicularly to the surface and

(23)

ADSORPTION PHENOMENA

37

when the dipole is parallel to the surface. If the direction of the dipole forms an angle /3 with the normal to the surface, we obtain

All these contributions are very low and they may, in most cases, be neglected. The polarization of a dielectric adsorbent by an adsorbed dipole may, similarly, be neglected. 6 . The Polarization of a n Adsorbed Molecule by a Dielectric Adsorbent

The electrostatic field emanating from the surface of an ionic crystal will also polarize a molecule adsorbed on it. The energy of combination due to this effect is F2CX E,= - -

2

where P is the strength of the field (such as, given by Eq. (17) for a cubic face of NaC1) a t the center of polarisability of the adsorbed molecule, and a is the polarizability of that molecule. As we have seen in Sec. V,4, the field over a smooth face of an ionic crystal is small and falls very rapidly with the distance r . At the center of polarizability of the adsorbed molecule, F is usually so small that the contribution according to Eq. (26) is of minor importance. Calculating this contribution for an argon atom on top of a potassium ion of the cubic face of KC1, and taking the distance between the center of the argon atom and the center of the potassium ion as r = 3.14 A. (the same as the smallest interionic distance, re, in the KCl crystal), we obtain E , = 0.25 kcal./mole, if the ~ m . ~ polarizability of the argon atom is a = 1.68 X Lenel (40)and also Teller (42) have pointed out that the electrostatic field of the surface of the ionic lattice is too unhomogeneous for even an argon atom to be treated as a whole with one average value for the polarizability, centered in the center of the atom. More elaborate calculations, taking into account a more likely distribution of the field strength and of the polarizability, lead to a value of E, = 0.45 kcal./mole for the argon atom in this position. It is to be noted that an argon atom, adsorbed just over a center of a square formed by four surface ions of the cubic face of KCl, would, according to Eq. (26), not contribute to the adsorption energy by static polarization as the electric field is zero in this position. According to Lenel’s method, however, a contribution of 0.37 kcal./rnole is obtained. All these contributions, however, are certainly low. It is only at “active spots,” as already mentioned in Sec. V,4, that the electrostatic polar-

38

J. H. DE BOER

ization leads t o important contributions toward the adsorption energy. We shall return to this problem in Sec. V,12. 7. T h e Polarization of an Adsorbed Molecule by a Conducting Adsorbent Only a few years ago it was discovered by Mignolet (43) that a metal also polarizes a molecule which is adsorbed on its surface. Mignolet, by measuring contact potentials, found that nonpolar molecules, adsorbed on metallic surfaces by purely physical adsorption forces, show, nevertheless, rather appreciable dipole moments. T o give an example, Mignolet found a potential change of 0.85 volt by the adsorption of xenon on a nickel surface. Assuming that he had a fully occupied xenon-adsorption layer, this means that each xenon atom shows a n induced dipole moment of ,u = 0.42 X e.s.u. (0.42 Debye). The direction of the dipoles is such that the positive part is pointing away from the adsorbing surface. The same conclusion may be obtained from the study of the behavior of many gases adsorbed on charcoal. We shall discuss the mobility of adsorbed molecules in Sec. VII, but we may mention here one of the results of such studies. Many gases, such as A, Nz, 02,CO, CH4, etc., when adsorbed on charcoal, behave as two-dimensional nonideal gases (44). This behavior can be described by a two-dimensional van der Waals’ equation, from which a two-dimensional van der Waals’ constant a2 (comparable with the normal three-dimensional van der Waals’ a) may be derived. The two-dimensional van der Waals’ constants can also be calculated from the three-dimensional values of a (46). The experimental results show that the actual a2 constants for gases adsorbed on charcoal or on mercury are always far lower than the theoretical ones and are very often even negative (46). The adsorbed molecules tend t o repel each other instead of showing a mutual attraction. This behavior also points t o a polarization of the adsorbed molecules by the field of the charcoal or of the mercury (47). We may assume this polarization t o be caused by an electric double layer formed by an electron distribution over the surface of the conducting adsorbent and corresponding positive charges in the metal. The dipole moment induced in the adsorbed molecules by the field of this double layer may be calculated from the difference between the theoretical value of a2 and the actual value which is found. This difference forms the a2 contribution caused by the dipole and is given by the expression

where p is the dipole moment and d is the diameter of the molecules (48).The actual value of the field may then be derived in each case from p

39

ADSORPTION PHENOMENA

by using the expression

where a is the polarizability of the molecule. E q u a h n (26) could then be used t o calculate the contribution t o the adsorption energy due t o this polarization. It is not necessary to calculate F , however, as E, is immediately given by

Some figures obtained in this way for gases adsorbed on charcoal are given in Table I. TABLE I P,

Gas

Nz

Debye

GO

1.1 1.1

CH,

1.16

C2He

1.7

C3Hs

2.0 1.5

GO 2

’””, ff,

10-24

cm.8

1.76 1.95 2.60 4.47 6.29 2.65

2a

kcal./mole 5.0 4.4 3.6 4.6 4.6 6.1

It is clear that this effect leads t o a rather important contribution t o the adsorption energy, if not t o the most important, in those cases of adsorption of gases on conducting surfaces where no chemical interaction results. 8. Chemical Bonds in Adsorption Phenomena on Metals a. The Formation of Ions. The adsorption of atoms of alkali metals, alkaline earth metals, and some other metals on surfaces of other metals may easily lead to the formation of ions from the adsorbed atoms. An electron is transferred from the adsorbed atom t o the adsorbing metal surface. The formation of ions from alkali atoms by a heated tungsten filament was discovered, nearly simultaneously, by Ives (@) and by Langmuir and Kingdon (50). The ions evaporate from the filament, and the condition for this conversion and the subsequent evaporation of the ions is t ha t the ionizing potential Vi, of the alkali atoms be less than the work function, ‘p, of the metal filament. A metal may absorb electrons as well as emit them, and for cases of chemisorption we may compare a metal as a whole to an atom, possessing an “ionization energy,” eVi = w, as well as a n electron affinity - Ed = --ep (61).If etp of the metal is

40

J. H. DE BOER

larger than eV, of the adsorbed alkali atom, a n electron will be transferred from the alkali atom to the metal and on desorption a t sufficiently high temperatures an alkali ion will evaporate. If the alkali atoms strike a metal surface with a lower temperature, they will still be ionized, and so the electrons will still be transferred to the adsorbing metal. At lower temperatures they will not evaporate but

4

8

12

16

ZO------aD

-c t i n A

FIQ.4. Potential curves relating t o the combination of sodium and chlorine to a molecule.

stay on the surface as adsorbed ions. Initially it was thought that Eq. (30) describes also the condition for this ionic adsorption. The author (66), however, showed that the condition is where Q. ia the adsorption energy of the alkali atom in atomic form and Qi its adsorption energy in ionic form. The adsorption of a sodium atom on a tungsten surface leading to an ionic adsorption may be compared directly to the formation of an ionic molecule of NaCl from the atoms of N a and C1. Figure 4 gives the change in potential energy as a function

ADSORPTION PHENOMENA

41

of the distance between an N a atom and a C1 atom. When the atoms (energy level A ) approach each other and the transfer of a n electron could be avoided, only a weak van der Waals' attraction (minimum B ) would result. The transfer of an electron from a n Na atom t o a remote C1 atom would mean a shift from energy level A to level D,the energy distance being given by (€Tii - Eel), viz. the difference between the ionization energy of the sodium atom and the electron affinity of the chlorine atom. When the ions (level 0 ) approach each other, a strong Coulomb attraction brings us t o the minimum €3, where the repulsion forces

FIQ.5. Potential curves relating to the adsorption of sodium on a tungsten surfme.

balance the attraction forces. The energy difference between level A (atoms) and the minimum E (ionic molecule) gives the heat of formation of the ionic molecule from the atoms. A similar curve may be drawn for the system of a sodium atom and a tungsten surface (Fig. 5). Minimum B (adsorption of N a in atomic form) would be appreciably lower than in the corresponding Fig. 4, and minimum E is higher than in Fig. 4. The difference between the energy levels A and D is here given by (cVi - ep), i.e., by the difference between the ionization energy of the sodium atom and the electron affinity of the tungsten metal. As level E is still appreciably lower than level B [condition (31)], the atom will adsorb in an ionic form; the real change in potential energy connected with the mutual approach of the sodium atom and the tungsten surface being given by the line ASEF. If, by heating, the adsorbed ion is desorbed, line E S A will be followed and the ion desorbs in atomic form while drawing a n electron from the metal

42

J. H. D E BOER

with it (see also See. IX,5). The condition for this atomic evaporation is given by the unequality (30), viz. the difference between the levels A and D; hence level A is lower than level D. When a cesium atom is adsorbed on a tungsten surface, level A is higher than level D (Fig. 6 ) and the desorption of the cesium is in ionic form, provided that no external electric fields are used that will force atoms to evaporate. The potential curves of Fig. 6 are completely comparable to the formation of the ionic molecule of CsF from the atoms of cesium and fluorine (Fig. 7).

ur i n A

FIG.6. Potential curves relating to the adsorption of cesium on a tungsten surface.

The adsorbing metal as such behaves in these cases as an atom of an electronegative element. Some years ago, in preparing an article about adsorption forces, the author (53) wrote: “One of the first statements that has to be made before discussing atomic forces and adsorption is that there are no special adsorption forces.” Generally speaking this statement may be true. The author, however, overlooked the special nature of the forces on conducting surfaces, as discussed in this section and in Sec. V,7. As the conduction electrons, which play an important role in the forces discussed in these two sections, do not belong t o definite atoms of the conducting body, but t o the body as a whole we may speak of special adsorption forces in their own rights. The question may arise whether, in forming adsorbed ions on surfaces, the metal may sometimes also act as the electron donor, resulting

ADSORPTION PHENOMENA

43

in the formation of adsorbed negative ions. From experiments of Rijanoff and Lukirsky (54) on selective photoelectric emission of potassium exposed t o atomic hydrogen, it may be concluded th a t hydrogen atoms when colliding with a surface of potassium will each receive a n electron from the metal and form an adsorbed layer of negatively charged hydrogen ions on the surface (55).This surface compound may be compared t o the ionic compound of lithium hydride or other hydrides. E

in kcal/rnolc

FIQ.7. Potential curves relating to the combination of cesium and fluorine atoms to a molecule.

b. Sharing of Electrons. I n the last example of the previous section we compared the adsorption of atomic hydrogen to the formation of a n alkali hydride. We may now ask whether surface combinations exist that may be compared t o chemical compounds of hydrogen in which this element is the electropositive partner, as is the case in hydrogen chloride. It is known that HC1 cannot be described as a purely ionic compound. Nevertheless some decades ago many attempts were made t o describe i t as a n ionic compound, in which the negative chloride ion was

44

3. H. DE BOER

polarized by the positive hydrogen ion to such a n extent that part of the electrons did belong to both partners. Later it became clear that only one pair of electrons, shared by both partners, was responsible for the bond, the electrons forming a pair when their spins are opposed. The resulting HC1 molecule, however, is a polar molecule, the electron distribution being such that the center of the negative end of the dipole is more on the side of the chlorine atom and the center of the positive end on the side of the hydrogen atom. It is now customary to describe the actual situation as a hybrid between the ionic molecule H+Cl- and the homopolar atomic molecule HCI, a situation which may be expressed by one of the systems of formulation used in these cases of “resonance” phenomena, such as H+C1- H HCl or H+C1 The second formula means merely that the HC1 molecule is a resonance hybrid between the ionic molecule H+Cl- and the molecule with the purely covalent bond, the direction of the arrow giving the direction in which the electrons have, on the average, been displaced (66). As, however, such an arrow is used by others (57), for indicating a coordinate link (semipolar double bond) caused by a lone electron pair of the donor atom, which likewise produces a dipole with its positive end on the donor side and its negative one on the acceptor side, the author suggests t h a t the symbol be used for the normal covalent bond, which, by resonance with an ionic structure, possesses a dipole. The point of this half arrow also indicates the direction of the negative end of the dipole. The full arrow -+ will then be reserved for the coordinate link. Both links play their roles in chemisorption, and it may be useful for the purposes of this article t o introduce relatively simple symbols. According to this principle HC1 should be formulated as H-C1. I n many cases of chemisorption normal covalent bonds are formed, where an electron of the adsorbed atom and one of the metal form a pair. The adsorbed atoms share their electrons either with the atoms of the metal on which they are adsorbed or with the metal as a whole. They form dipoles on the surface of the metal, and the direction of these dipoles are of great importance for chemisorption and catalysis. Just as a metal may play the role of a halogen atom in the adsorption of alkali ions discussed in the previous section, it may in other cases act similarly to the chlorine atom in HCl and form a covalent bond. Hydrogen atoms adsorbed on the surface of platinum may serve as an example. The dipoles point with their positive ends away from the metal and may

-

ADSORPTION PHENOMENA

45

be regarded as resonance hybrids :

which, according to the above-mentioned suggestions, we can formulate as

a symbol t ha t will be simplified t o

I n other cases the adsorption of hydrogen atoms results in surface hydride dipoles pointing with their negative ends away from the metal, as is observed in the adsorption on nickel,

R

which is t o some extent comparable to the existence of alkali hydrides, or rather t o that of free molecules of metallic hydrides. The free molecules of metallic oxides can best be described as having covalent links, possessing dipoles (68). Similarly oxygen atoms adsorbed on metallic surfaces form covalent bonds, sharing two pairs of electrons with the metal (69)or with one specific atom or two atoms of the metal. Their dipoles point with the negative ends away from the metal. We may, t o give an example, express the situation of the adsorption of oxygen on silver by

,+, +, or

Similar covalent bonds may be f rm d between metal surfaces and many other atoms, including atoms forming part of molecules or radicals. I n many cases the dipoles point with their negative ends away from the metal surface. I n other cases, however, as with CzHzand CzH4on nickel, they form dipoles pointing with their positive poles away from the surface (60).

46

J. H. DE BOER

The surface hydrides, surface oxides, and other surface compounds, mentioned above, need not be formed by the action of free atoms with free valencies on metal surfaces, but, just as in normal chemical reactions, these compounds may result from the reaction of the metal surfaces with molecules. The chemisorption of an H z molecule on a metal surface may lead t o the chemisorption of two separate hydrogen atoms and so may the action of a n 0 2 molecule on a metal surface lead to the chemisorption of two oxygen atoms, the action of an NPmolecule t o the chemisorption of two nitrogen atoms, etc. Surface hydrides, oxides, and nitrides are, then, the result of normal chemical reactions of these gases with the surfaces of the metals. or Nz could not be This does not mean that the molecules H z , 02, chemisorbed in their molecular state without being split into atoms. We shall see in Sec. V,11 that molecular chemisorption also comes into the picture. Other molecules may also lead to such “dissociative” chemisorptions. I n some cases of chemisorption of NH3 this molecule may split into a n H atom and an NH2 radical which are separately adsorbed on the metal surface; sometimes the dissociation may go even further. I n this type of surface reaction CHI and other hydrocarbons may also split into H atoms a lid hydrocarbon radicals which are bound by chemisorption. Ethylene, when adsorbed on nickel or similar metals, may be supposed t o open its double bond and be adsorbed as HZC-CH, /

V”/’

/

but it is also possible that a dissociative chemisorption will take place as is given by H HC=CH

75

H

‘a’ VL

The chemisorbed CzHzradical may, in such a case, be supposed to be stabilized by a resonance between structures such as + + HC -CH

HC =CH

HC-CH

*+-*++As already stated above, the dipole connected with chemisorbed ethylene points with its positive end away from the metal surface.

47

ADSORPTION PHENOMENA

Coordinate links between molecules having at least one lone pair of electrons and acting as the electron donor, and metal surfaces acting as electron acceptors may also lead to chemisorption of these molecules. Maxted (61) and co-workers found that many molecules with such lone pairs of electroils are extremely effective poisons for catalytic processes because of their strong adsorption on the metallic surfaces acting as catalysts. Maxted ascribes this strong adsorption to the coordinate links forming in these cases. We may give the following examples t o illustrate the type of chemisorption:

using the full arrow t o indicate the coordinate link. These donor elements,

S,As, N and also Se, P, etc., do not form catalyst-poison molecules when they are in their highest state of valency; there are no lone pairs of electrons then. The same holds for ions of these elements. The sulfite ion acts as a poison; the sulfate ion does not:

The former is bound by a coordinate link, aided by an image force of the ionic charge, which, however, is rather weak here because of the relatively large distances of the negative excess charges from the metal. The sulfate ion exhibits this image force only. In both cases there are the van der Waals’ forces as well, which also will be weak here. Suhrmann (62) explains the strong increase of the normal photoelectric effect of metals caused by the adsorption of water molecules and also by the molecules of ammonia, by accepting similar coordinate links t o function in the chemisorption of these molecules. Dipoles are formed which point with their positive poles away from the surface, thereby decreasing the work function and, consequently, increasing the normal photoelectric effect: H

48

J. H . DE BOER

9. Heats of Adsorption and Desorption and Activation Energies

Connected with Chemisorption on Metals When sodium is adsorbed on a tungsten surface, it is transferred into an ion. The heat of adsorption can be read from Fig. 5 (Sec. V,8,a) as the difference between level A (the atom) and the minimum E of curve DEF. The difference in energy between levels D and E is given mainly by Eq. (16), Sec. V,3, and is modified by contributions from van der Wads’ forces, polarization forces, and repulsive forces. The total amount of this energy difference is about Q; = 77 kcal./mole. I n order to obtain the heat of adsorption, the energy difference between levels A and D , viz., tV;

- q o = 14 kcal./mole

must be subtracted from this value, the resulting value being (QJi = about 63 kcal./mole

symbolizes the heat of adsorption of the atom in the form of where (QJi an ion. The heat of desorption is also given by the same difference between levels E and A and is, consequently, equal to the heat of adsorption; Qsd..

=

This equality is nornial for all cases of physical adsorption and, as we have seen, also holds in the present case of chemisorption. When we now consider the adsorption of cesium on a tungsten surface, we see from Fig. 6 that the heat of adsorption of the atom is given by (Qa); =

- (tV; -

Qi

tq) =

54

+ 14.8 = 68.8 kcal./mole

In this case cesium is desorbed in the ionic form, the heat of desorption being no more than Qde.. = Q = i 54 kcal./mole which shows that the heat of desorption is smaller than the heat of adsorption. As a matter of fact we cannot speak of a purely reversible phenomenon here. We shall not discuss the details of the calculations of the energies involved in Figs. 5 and 6 of Sec. V,8,a; a reference to some earlier calculations by the author may suffice (63). When atoms of hydrogen are adsorbed on a metal surface, the changes in potential energy may be schematically represented by a single curve as shown in Fig. 8. Our present knowledge of the forces responsible for the formation of the covalent bond described in Sec. V,8,b does not enable

ADSORPTION PHENOMENA

49

us to calculate this curve. Once the heat of adsorption is known and some other figures are given, the form of the curve may be constructed as a socalled “Morse” curve (64). The calculation of the heat of adsorption may be approached in a semi-empirical way as was shown by Eley (66),who uses an equation of Pauling (66) for calculating the bond energy of a covalent bond between the atoms A and B:

- B ) = s { D ( A - A ) + D ( B - B ) ) + 2 3 . U 6 ( ~-~ X B ) ~ (32) where D ( A - B ) represents the bond energy between A and B in kcal./ D(A

mole and

ZA

and

ZB

are the “electronegativities” of A and B.

ro I

I

-r

FIQ.8. Schematical representation of the potential curve of the adsorption of a hydrogen atom on a metal surface.

I n the case of the surface bond between an H atom and a tungsten surface, we obtain

D(W

-

H) = l/i(D(W

- W)

+ D(H - H)) + 23.06(zw - za)2

(32a)

D(H - H) is known to be 103.2 kcal./mole. Eley calculates D(W - W) from the sublimation energy of the metal and puts B(W - W) = %-6x (33) where S is the sublimation energy. (xW- zH) may be estimated with Pauling’s approximation, which assumes that this difference is equal t o the dipole moment of the bond, expressed in debyes (1 debye = 10-8 e.s.u.). This dipole moment may be obtained from the change in contact potential caused by the adsorbed layer. This value is mostly known experimentally for a fully occupied ad-

50

J. H. DE BOER

sorbed layer, and so for a complete layer of dipoles on the metal (at 0 0 being the degree of occupation). In that case we obtain

=

1,

where pe-1 is the dipole moment per bond a t 0 = 1; AV is the change in contact potential, and ne,l is the number of dipoles (adsorbed atoms) per square centimeter a t e = 1. As a matter of fact, for calculating the term (zw - zH)with (zw - z), = PO-0 (35) we need pe-0, this being the dipole moment for a surface bond a t 6 = 0. Eley assumes th at the difference between pe-1 and pe-o will not be so large as t o disturb the calculation seriously. The calculation of the terms of Eq. (32a) leads to x { D ( W - W) % { 33.8

+ D ( H - H) 1 + 23.06(~w+ 103.2 1 + 4.9 D(W - H)

hence t o

=

SH)'

D(W - H) 73.4

= =

73.4 kcal./mole

The experimental value is 74.1 kcal./mole. The result is fantastically good in this example and it cannot be expected t o fit so closely in other cases. I n Table I1 we have given some calculated and observed values for D ( M - H ) for different metals M , as calculated by Eley; the observed value for Cu, however, has been obtained from more recent data (67). TABLE I1

__

~

~

Bond Energies for Surface Hydrides in Kcal./Mole Metal

D(M D(M

- H) calc. - H) obs.

Ta 67.6 71.1

W 73.4 74.1

Cr 59.5 74.1

Fe 60.1 67.6

Ni

Rh

cu

60.2 67.1

63.1 65.6

58.4 69.1(67)

It is remarkable that the observed figures differ so little, far less than the calculated ones. The formation of surface hydrides, oxides, and nitrides is usually a result of the dissociative chemisorption of the molecules of these gases. As can be seen from Fig. 9, which gives the potential curves for such a n adsorption, the heat of adsorption is given by the difference in energy level between A and E , this difference being given by (Qm)a

=

2Qa - Dm

(36)

ADSORPTION PHENOMENA

51

where (Q,), stands for the heat of adsorption of the molecule in atomic form, Q. for the heat of adsorption of the atom, and D, denotes the dissociation energy of the molecule. [Compare D(H - H) in Eq. (32a).] Calculated values of (Q,), can be easily derived from calculated values of Q. and the known values of D,. Trapnell (68) gives a survey of calculated and observed values of (&), for various gases on various metals. I n all these cases he gives the so-called “initial” heats of adsorption, i.e., the heat of adsorption on a bare surface. It is these figures which E in kcal/mole

-

rinA

FIG.9. Potential curves relating to the dissociative chemisorption of a molecule M (H2) on a metal Me without an activation energy,

have to be compared with the calculated ones. At higher degrees of occupation there is generally a strong decrease in heat of adsorption (see Sec. IX). According to Trapnell’s tables the experimental values of the initial heats of adsorption of various gases are highest when the gases are adsorbed on tantalum and lowest when they are adsorbed on copper or gold; they follow the order:

Ta

> W, Cr > Fe > Ni > Rh > Cu, Au

which is not the order which may be expected from calculations with the aid of the equations given in this section. We shall return to this problem in the next section. I n Fig. 9 the intersection point S of the curves ARC and DEF lies

52

J. H. DE BOER

lower than level A . This means that there is a constant decrease of potential energy during the mutual approach of the molecule and the metal ; there is, in other words, no activation energy involved in the chemical reaction :

M

f

F]

At At

A t the intersection point S , or rather a t a somewhat lower energy level (stabilization by resonance), we find the so-called “activated complex” of this chemical reaction and so the transition state between the original reaction partners and the reaction product. Metal surfaces seem to act as if they really have free valencies. There are no activation energies in chemisorption reactions between atoms and metal surfaces. I n reactions of molecules with metal surfaces, on the other hand, activation energies may be expected. If it is true, however, that metal surfaces behave as free atoms or as surfaces with free valencies, the activation energies should be small for we know from other chemical reactions that the reaction between molecules and free atoms or free radicals proceed without activation energies or with small values of these energies. From the behavior of Hz with chlorine atoms, for example (no activation energy), we may expect the activation energy for the chemisorption of H2 on metal surfaces to be very small or even negligible. Other molecules, such as O2 and N2, having far larger dissociation energies, may, however, need some activation in order to react with a metal surface. We may approach this problem somewhat more closely with the aid of potential curves. There is no activation energy in Figs. 5 and 9; the slope of the right part of curve DEF is such that the intersection point S does not come above level A . I n Fig. 10 the location of the intersection point S is such th a t a n activation energy E , will govern the speed of the reaction. It may be clear that point S will be higher the larger the difference between levels A and D, or the smaller the difference between levels D and E. The activation energy will also be higher the smaller the equilibrium distance 9-0 a t point E and the steeper the slope of the part DE of curve DEF. I n Fig. 11 the influence of these variations may be seen. The smaller the ionic contribution to the adsorption energy of the atom, the steeper will be part DE and the higher point S and, hence, the larger the activation energy. I n recent years more and more experimental evidence has come to point t o the conclusion that the chemisorption of Hz on really pure metal surfaces does not involve an activation energy (69-72) or at most in-

ADSORPTION PHENOMENA

53

volves such a small one that even a t very low temperatures the adsorption proceeds practically instantaneously. The adsorption of Nz on an iron surface, (73,74) even on a very pure iron surface, involves an activation energy, and the situation may be described by Fig. 10. The chemisorption of Nz on a tungsten surface, on the other hand, seems to proceed without an activation energy or a t most a very small one (76-77). E in kral/mole

FIQ.10. Potential curves relating to the dissociative chemisorption of a molecule M on a metal Me with an activation energy.

Contaminated or incompletely reduced metal surfaces, however, may also in their chemisorption reactions with Hz lead to appreciable values of activation energies governing the speed of the reaction. We shall discuss these phenomena later (Sec. X,4). In case an activation energy is noticed the heat of desorption is higher than the heat of adsorption. A glance a t Fig. 10 shows that the speed of the desorption reaction will be governed by the difference in height between level E and intersection point S , which is the sum of the heat of adsorption ( A to E ) and the activation energy ( A to 8). If the energy level of the minimum E lies above level A (Fig. 12), the heat of adsorption will be negative (increase of potential energy). Such an adsorption would be an endothermic one. I n physical adsorption p h e

54

J. H. DE BOER

nomena endothermic adsorption does not occur; as the entropy always decreases in physical adsorption phenomena, the heat of adsorption has t o be positive in order to enable adsorption to take place. I n chemisorption processes, however, endothermic adsorption cannot be excluded. b

I C E

F'

F

E -r

-r

d

C

0

i E' E

E - 1

-r

FIG.11. Various possibilities for the relative situations of potential curves relating to dissociative chemisorption. (a) A higher value for D, gives a higher act. energy, (b) a higher value for &. gives a lower act. energy, (c) a shorter distance ro gives a higher act. energy, (d) a higher ionic contribution gives a lower act. energy.

Figure 12 indicates a possibility of such a case, which is comparable t o the existence of endothermic chemical compounds. Once formed, such a n endothermic surface molecule may have a certain time of existence before dissociating. This time may be long enough t o enable it t o react with other molecules. I n many catalytic processes the intermediate role

55

ADSORPTION PHENOMENA

of such endothermic surface compounds is not to be excluded. It may be expected that such a n endothermic chemisorption plays a rather important role in catalysis. The time of existence (for time of adsorption, see Sec. VII,4) depends on the heat of desorption (E',S),which is smaller than I

2

4

-r

6

8

10

12

FIG. 12. Schematic representation of potential curves relating t o anlendothermic chemisorption.

the heat of activation (A,X) in this case. We shall discuss possible cases of endothermic chemisorption in Secs. V,9, VI,3,4,5 and X,4. 10. The InJEuenceof the Electronic Structure of the Metals

I n the previous section we saw that the experimental heats of chemisorption of many gases are highest on tantalum and lower on other metals, the decrease following a certain order. The order, mentioned there, is not exactly the same as the order for the heats of sublimation (78).If,however, Eq. (32) is valid and the term of the electronegativities may be neglected or is the same for all metals, the term containing the bond strength between metal atoms would be the only one which differs from metal to metal, and the term would follow the heat of sublimation. It has struck many workers in the field that the transition metals and near-transition metals are the best catalysts for many gas reactions (78), and i t was suggested (79) th at probably the covalent bonds between the chemisorbed atoms and the metal were formed by sharing of an electron of the atom with a d electron of the metal. It is, indeed, not improbable that d electrons of the metals play a n important role in these chemisorption phenomena. It is known from their role in the bond strengths of complex chemical compounds that pairing of other electrons with d electrons leads to strong bonds. An incorporation of d elec-

56

J. H. DE BOER

trons from the metallic adsorbents would, therefore, lead to higher heats of adsorption (80). Beeck (82)drew attention to the fact that the order of decreasing heats of chemisorption on various metals is the same as the order of increasing d character of these metals. Increasing d character, according to Pauling (82),means that more d electrons are used for the mutual cohesion of the metal atoms in the metallic crystal lattices. According to Pauling’s metal theory these d electrons are not available for other chemical bond formations. An increasing d character, therefore, means fewer d electrons for chemisorption. The best indication that d electrons may play an important role in catalytic reactions and in chemisorption phenomena is obtained from those experiments where a reaction is studied, that is catalyzed by the surfaces of alloys of two metals composed in such a way that in the range of alloys the d band is completely filled from a certain composition onward. I n several cases a sudden change in catalytic activity is found a t such a critical composition (83,84). Another, more direct, indication is obtained from the work of Dilke, Eley, and Maxted (85),who found from the change in magnetic susceptibility of palladium on adsorption of dimethylsulfide that an electron of the sulfide had entered the d band of the metal. As we saw in Sec. Vl8,b, a coordinate link is formed in this case. Although the availability of d electrons will certainly influence the ease of formation and the strength of the covalent bonds in chemisorption phenomena on metals, we may not expect a simple relationship between the heat of chemisorption and some more or less simple property which is related to the d electrons. There will certainly be other properties of the metals which play a role. I n expression (32) the last term containing the electronegativities has some relation to the availability of electrons from the metal, and it must be said that the order in which the heats of chemisorption fall (See. V,8,b) is nearly the same as that in which the work function of the metals rises. I n the formation of the dipoles between the adsorbed atoms and the metal, work has to be done against the work function; we may expect that less work will be necessary to form these dipoles and that the dipole moment will be larger, the smaller the work function. I n the forming of these polar bonds, electrons of the metals are withdrawn from the metal. The binding of electrons can be shown by the increase of the secondary electron emission (86,8?), and conductivity measurements (88,89) and measurements of contact resistances (90,91) show that conduction electrons have been occupied by these bonds. The physical adsorption of a gas on a metallic surface, on the other hand, causes a slight increase of the conductivity of the metal (92,93).

ADSORPTION PHENOMENA

57

If. Chemical Bonds in Adsorption Phenomena on Nonmetallic Surfaces Cheniisorption on surfaces other than those of metals have been studied mostly on oxides or salts. Oxides may be either semiconductors or insulating dielectrics. Semiconductors may in some respects act similarly t o metals; we may speak about a work function, more or less free electrons in partially occupied energy bands for electrons, etc., as is customary in the more physical theories about metal structure. It is more convenient for our purposes to consider them from the chemical viewpoint and t o describe them as ionic crystals having ions of the same atom (homonymous ions), but in different valencies, on crystallographically identical places (94).Such ions may be statistically distributed among their sites in the crystal lattice. The surfaces of such crystals contain, therefore, also these ions of different valencies. I n some special cases the lattices of the oxides or salts may contain these ions already in their stoichiometric compositions (Fe3O4)(95), but in most cases homonymous ions of different valencies are linked with deviations of the stoichiometric composition or they are controlled by the addition of some ions of other metals having other valencies which form solid solutions in the crystals, a method devised by Verwey and collaborators (96,97). Homonymous ions of deviating valency (charge) may also be formed in some processes of chemisorption; they will then be formed on or near the surface but may subsequently diffuse into the interior of the cryst,al lattice. There is no boundary between semiconductors and insulators; the distinction has been made only for practical reasons and a arbitrary boundary has been chosen at a certain value of the electrical resistance. We may expect that also metal ions in so-called “insulators” can be transferred into ions of other valency or into atoms. Such reactions may proceed more easily a t the surface than in the interior of the crystal and we have t o be aware of this possibility in cheinisorption processes. a. Carbon Monoxide Adsorpiion. Normal CuzO, containing a slight excess of oxygen with respect to its stoichiometric composition, contains Cuzf ions and consequently is a semiconductor. The electron transfer in the lattice can be pictured as

cu+

+ cuz+ i3 cu2+ + c u +

At room temperature carbon monoxide is adsorbed readiIy while simultaneously the semiconductivity is decreased (98).Probably a covalent bond is formed between CO and the copper ions of the surface, thus immobilizing electrons for the electron transfer in the lattice. Whether the bonding of the CO molecule is caused by donating two electrons from the

58

J. H. DE BOER

CO t o a Cu2+ion or by sharing electrons with a Cu+ ion is not known. I n both cases d electrons (or d levels) would be involved. Carbon monoxide is also readily adsorbed in ZnO (99) and a donation of electrons to the Zn2f ions or a sharing of electrons with Zn+ ions or Zn atoms, which may both be present in ZnO (loo),might be considered. A similar bond may be formed in the chemisorption of CO on C r 203a t low temperatures (liquid-air temperature). When the adsorbed gas is desorbed in this latter case (101),and also from ZnO and C u 2 0 , i t is desorbed as such. If i t is chemisorbed on Cr203 a t room temperature, however, it is desorbed as COz (10.2). Apparently a more complicated surface reaction has taken place in this case. The CO probably combines with two oxygen ions of the surface to form a CO:- ion while simultaneously some metal ions are reduced t o a lower valency. On the surface of Crz03this reaction may perhaps be pictured as CO

+ 202- + 2cr3+--t C0:- + 2Cr2+

When the adsorbed gas is desorbed by heating, the surface carbonate ions decompose: C O i - 4 COZ 0 2 -

+

and COz is desorbed instead of CO. Such a n irreversible chemisorption of CO has also been found with other oxides. I n all such cases the production of metal ions of lower valency results in a tendency to adsorb oxygen, or rather t o enhance the capacity to adsorb oxygen. Sometimes the extra amount of oxygen that may be adsorbed after an irreversible CO adsorption is-stoichiometrically-equal to half the amount of CO t ha t was adsorbed. Carbon monoxide on CuO, for example,

CO

+ 202- + 2CUZ+--t c0;- + 2 c u +

induces a n enhanced O2 adsorption

XOZ

+ 2Cu++

2Cu2+

+

02-

oxidizing practically all the Cu+ t o Cu2+ again (105).In other eases the enhancement of the oxygen chemisorption is less. The reversible adsorption of CO on ZnO mentioned above does not lead t o a n enhancement of oxygen adsorption. The net result of the irreversible adsorption of CO, followed by the adsorption of the right amount of O2is

CO

+

SO2

+

O2-+

c0:-

and i t has been proved (104) that a direct adsorption of C 0 2on the same oxides leads t o identical surface carbonate layers.

59

ADSORPTION PHENOMENA

b. Hydrogen Adsorption. Hydrogen, just as carbon monoxide, may either react with an oxide surface t o cause a reversible chemisorption or may be bound in a stronger way which, on desorption, leads to the evaporation of water molecules. Both ehenisorptions must be considered t o be of the dissociative type. The reversible adsorption of H z and of Dz on Crz03a t liquid-air temperatures, with initial heats of adsorption of 5.1 and 5.4 kcal./mole respectively (105), are active in the exchange of H z and DZand we must, consequently, assume that both gases are adsorbed in the atomic form. We may think of a sharing of electrons between Cr3+ ions and H atoms, while also a n electron transfer from the H atom t o a Cr3+ ion, forming a CrZf ion, may be considered. Both possibilities may contribute t o the real situation, the bond being a resonance between them:

ci

or

Cr*O3

At room temperature H z is chemisorbed by Crz03 in a different way (106); the initial heat of adsorption is far higher, viz. 72 kcal./mole, and

on desorption HzO is given off. We must assume th a t on adsorption OHions are formed from Oz- ions and metal ions are reduced t o a lower valency Cr3t

02-

Cr3+

r l Cr2+ 0 Cr2+ 0

02-

(c.,o,I+n,-

rl-

On desorption HzO evaporates and an ion vacancy is left:

Cr2+

o

Cr2+ O

Cr2f

T

l

... Crz+ +

02-

H

Z

O

After adsorption of Hz in this way an enhancement of the tendency to chemisorb O2 should be expected which, indeed, is found t o be the case. After desorption of water, one may expect chemisorbed oxygen t o fill u p

60

J. H. DE BOER

the vacant site in the surface, I n our schematical way of describing surface phenomena we tend t o place the adsorbed species in a layer above the layer of the surface atoms. In reality there are often open sites in the surface layers and adsorbed atoms or ions may tend to fill such gaps. We shall see in a later section (VI1,G) that adsorbed atoms or ions and original surface atoms or ions may often change places; adsorption is, therefore, not restricted strictly to the outer surface. c. Oxygen Adsorption. It will be obvious from what has been said molecules may also be expected t o chemisorb readily when metal t ha t 0% ions of lower valency are present. This is indeed found t o be the case, and when 0 2 is adsorbed on CuzO the semiconductivity is increased ( l o r ) ,because Cu2+ions are created simultaneously next to the Cuf ions. Zinc oxide is normally nonstoichiometric in composition, there being a n oxygen deficiency. When it is prepared in air, a large amount of chemisorljed oxygen is present on the surface, which is thus more stoichiometric than the bulk (108). Oxygen chemisorption penetrates gradually into deeper layers, as will be discussed in Sec. VI1,B. It may be stated here that the oxidation of metals by oxygen to form metallic oxides proceeds essentially along these lines. The surface reaction is a chemisorption of oxygen, which is dissociated and ionized into two 02-ions, while, simultaneously metal atoms or metal ions of lower valencies are oxidized t o higher valencies. Metal atoms or metal ions of lower valencies are continuously produced by the part of the metal that has not yet been oxidized; they diffuse through the already formed oxide layer and provide the electrons for transferring the 0 2 molecules into 02-ions. As oxygen may, a t a sufficiently high temperature, be desorbed from the oxide, it may be expected th at there will be a continual exchange ions of the oxide surface, provided th a t between O2 of the gas and 02the temperature is high enough to enable the diffusion in the lattice to take place. The diffusion in the lattice seems to be the rate-governing process in exchange reactions between O2 (containing the stable isotope 1 6 0 ) and oxygen ions of Crz03,A1203, MgO, T h o z , etc. The apparent activation energy is 27 kcal./mole for Tho2 between 440' and 540°C. and 29.5 kcal./mole for Cr203 at temperatures below 410°C.;a t higher temperatures the exchange with Crz03requires a n activation energy of about 1 kcal./mole only (109). d . Hydrocarbons. Dissociative chemisorption on the metal ions of transition metal oxides, similar to the low-temperature reversible chemisorption of HI?on these oxides, has been suggested for the chemisorption of saturated hydrocarbons on these oxides a t high temperature. This adsorption is suggested as the first step in the dehydrogenation and aromatization of paraffins (110).

ADSORPTION P H E N O M E N A

61

e. Hydrogen Ions. Many oxides adsorb hydrogen ions from acids. We may in many cases assume them to be bound to the OH groups of the surfaces, where they are bound as in H30+ions in aqueous solutions. An acidified alumina, therefore, may be pictured as

H H

I

H

\ / Of

H

H

0

0

I

I

the anion of the acid being bound in the neighborhood of the positive charge. I n the lattice of A1203we must assume the constituent atoms to be in ionic form, hence Al3+ and 0 2 - . The surface OH groups may be bound mainly by covalent forces (in resonance with the ionic bond). When H+ is bound to such an OH group an oxonium compound is formed. The synthetic cracking catalyst consisting of SiOz and A1203 has no additional anion, the negative charge being part of the surface itself. The bonds in the silica lattice have mainly a covalent character; A1 may replace Si on the surface, provided that it can assume a negative formal charge t o enable it to show four covalencies like Si; the electrons are provided by the oxygen atom of the water molecule assuming a positive formal charge itself.

1.2. Active Spots (111)

I n Sec. V , l we discussed the influence of crevices, cavities, the inside of cracks, recessed parts of the surface, and especially the inside of capillaries. I n all these “active spots” for nonpolar van der Waals’ forces the adsorbed molecules can find far more direct neighbors than on a plane surface, and consequently the heat of adsorption is far higher in these spots than on plane surfaces. Owing to their structure many dielectric adsorbents, adsorbing molecules with nonpolar van der Waals’ forces show a rather heterogeneous distribution of adsorption sites of various strengths. If this were not the case, no smooth adsorption isotherms would be found, but isotherms showing sudden jumps, separated by hori-

62

J. H. DE BOER

zontal parts. A so-called “stepwise adsorption” would occur. The very fact that smooth-running adsorption isotherms are found proves the heterogeneous character of the surfaces for physical adsorption (112). It is worth while to see whether this conclusion may also be drawn for the other types of adsorption discussed in Sec. V. It is, in principle, the same for the nonpolar van der Waals’ forces discussed in Sec. V,2. Charcoal within molecular dimensions, acting principally as a conducting adsorbent, shows by its very nature a rather flat adsorbing surface (113). There is, however, a marked difference between the basal faces and the hexagonal prism faces of the graphitic structure, and normal charcoal is therefore not homogeneous enough in its surface to show stepwise physical adsorption. When charcoal is graphitized at very high temperatures, however, an adsorbent of a homogeneous nature is obtained and recently stepwise adsorption isotherms have been found for krypton or graphitized carbon black (114). Step-wise adsorption results from two-dimensional condensation (Xec. VIII,4) on homogeneous surfaces. If, at the same time, multimolecular adsorption takes place, steps may be found in the building up of every successive layer. These steps, however, do not coincide with the filling up of every successive layer (215). The forces between ions and metal surfaces, discussed in Sec. V,3 are far less influenced by active spots. Those spots that are active for nonpolar van der Waals’ forces are not active here. According to the simplified picture described in Sec. V,3, all crystallographic faces should give the same attraction if the equilibrium distance ro were the same. This distance, however, will not be the same and for this reason as well as because of other minor differences, we may expect the actual surfaces also to be heterogeneous with respect t o this contribution of adsorption forces though quantitatively far less outspoken than for the nonpolar van der Waals’ forces. The Coulomb forces between an ion and a cubic face of an ionic crystal of the NaCl type are small, as we saw in Sec. V,4. The corresponding forces on other crystallographic faces may be appreciably stronger (three to four times). Crystal edges and especially crystal corners also exercise far stronger forces than the plane surface. Those spots on a crystal face where during the forming of the crystal the growth stopped show very high attractive fields for ions of opposite charge. This type of adsorption, therefore, which may be negligibly small on some smooth crystal faces, may be of major importance on active spots of the character just mentioned. Actual surfaces, therefore, will again show a heterogeneous distribution of adsorption strengths for this type of adsorption. It may, however, be noted that the active spots for van der Waals’ forces and

ADSORPTION PHENOMENA

63

those for Coulomb forces do not coincide. It is rather so that active spots for van der Waals’ forces are not active for Coulomb forces and vice versa (116). As discussed in Sec. V,4, ions also polarize dielectrics [Eqs. (19) and (21)], and we saw that the contribution toward adsorption energy arising from this effect may be even more important than that from the Coulomb forces. Just as in the case of the interactions of ions and metal surfaces, these polarization effects are far less influenced by active spots than Coulomb forces are. The adsorption of polar molecules on surfaces of ionic crystals (Sec. V,5) is influenced by active spots of the same kind as influence the action of Coulomb forces. The effect of these active spots is, quantitatively, less €or dipole-containing molecules than for ions. The eff ect of dipoles on metal surfaces is small (Sec. V,5), and active spots are not expected to give appreciably higher contributions. The effect of active spots on the polarization of adsorbed molecules by a dielectric absorbent (Sec. V,6) is very great. The nature of the active spots is the same as of those which affect the attraction of ions or dipoles. Edges or corners of crystals, other crystallographic faces, and especially those places where the growth of individual crystal faces stopped, as well as lattice disturbances in the surface, will be active. The polarization of adsorbed molecules by conducting surfaces (Sec. V,7) cannot be expected to be highly influenced by most active spots, which are effective for van der Waals’ forces, nor can such an influence be expected in the case of the formation of ions by transfer of an electron between the adsorbed atom or molecule and a metal surface (Sec. V,8,a and partially also b). As the forces that govern the transfer of electrons are related to the work function, however, we may expect these forces to differ for different crystallographic faces, as we know t h a t the different crystallographic faces show different work functions. It is for this reason that actual metal surfaces, consisting of various crystallographic faces, ivill show some heterogeneous character with respect t o these forces. The same holds for the covalent forces between adsorbed molecules and metal surfaces discussed in Sec. V,8,b. Other active spots will be of less importance again in this case. Covalent bonds are more individual than bonds caused by van der WaaIs’ forces or by Coulomb or dipole attraction, and cooperation of other surrounding atoms of the adsorbent has, therefore, less influence. This just means that active spots are less important for covalent bonds. Covalent bonds between adsorbed molecules and oxide or salt surfaces (Sec. V,ll) may perhaps be highly dependent on active spots, because the formation of these bonds alters seriously the distribution of

64

J. H. D E BOER

electric charges in their neighborhood. It is, therefore, not only the strength of the bond which is formed that concerns us in this case, but also the change in bond strength around it. These rather complicated effects may probably result in a rather heterogeneous distribution of heats of those chemisorptions on the surfaces of oxides and of salts where the surrounding ions suffer changes of character and charge simultaneously. Other crystallographic faces will also give variations in the strength of these chemisorptions (see Sec. VIII) . VI. COOPERATION AMONG VARIOUSFORCES The various forms of interaction between a molecule and a surface, which are conveniently (see Sec. IV,l) treated as different forces and were discussed in Sec. V, cooperate with the repulsion forces, dealt with in Sec. IV,4, t o create the phenomenon of adsorption. They determine the magnitude of the adsorption energy and the distance between the adsorbed molecule and the surface. Nonpolar van der Waals’ forces and repulsion forces are always present and are, therefore, always among the cooperating forces in every case of adsorption. There are, however, few cases where these two general forces are the only ones that operate. The adsorption of noble gases on nonpolar dielectric surfaces will be governed by these two forces only. I n practically all other cases one or more of the other forces of Sec. V will cooperate. We shall, in this section, discuss a few selected cases of such cooperation. I . Physical Adsorption on Charcoal (and Metals)

The adsorption of many gases on charcoal is mostly taken as an example of cooperation of van der Waals’ forces and repulsion forces only. When London used his equations for the dispersion forces to calculate adsorption energies [Eqs. (8) and (12)l he found (117) a good agreement between his figures and the experimental values of the heats of adsorption of gases such as He, A, Nz, CO, CHI, and CO, on charcoal. H e had unfortunately made a calculation error, and so his figures were ten times too high. The discrepancy could partly, but only partly, be bridged by the application of a summation instead of an integration, as described in Sec. V,1. I n 1934 we gave a solution (118) of the problem, namely that the adsorption of these gases occurs in pockets, tubes, and cavities of the charcoal, hence that the adsorption would mainly take place on active places. This view was generally accepted and refined by Brunauer (119) when he suggested that all adsorbed molecules, in the very narrow capillaries of charcoal, would be in contact with two layers of carbon atoms instead of one. This view is entireIy true, but the calculated figures for

ADSORPTION PHENOMENA

65

the adsorption energy still tended t o be too low. One has to bear in mind that the repulsion forces were neglected in these last calculations. In Sec. IV,4 we saw how big the influence of the repulsion forces may be. Recent investigations showed (120) that all the gases mentioned above, when adsorbed on charcoal, are very mobile and behave as two-dimensional gases. It was discovered (121) in the course of the same investigations that the gases were polarized by the electric field of the charcoal (Sec. V,7) and that this polarization resulted in a n important contribution t o the heat of adsorption, probably the most important one. Physical adsorption of these gases on charcoal, therefore, must be regarded as being caused by the cooperative action of a polarization by the field of the charcoal nonpolar van der Waals forces and repulsion forces. The same picture holds for physical adsorption on metal surfaces. The polarization of the adsorbed molecules causes dipoles pointing with their positive ends away from the metal surface. The work function of the metal will be lowered by this effect, and it seems as if the increase of the normal nonselective photoelectric emission of metals by the adsorption of water molecules (122) or molecules of organic substances such as pyridine, propionic acid, and benzene (123) or alcohol, diethyl ether, and acetone (124) is caused by this effect. The explanation, which, many years ago, was given by the author (125), viz., polarization by positive hydrogen ions which should still be present, may seem to be unnecessary and obsolete. As already mentioned in Sec. V,8,b, water molecules may even cause dipoles by forming a coordinate bond with the metal surface, these dipoles working in the same sense as those formed by polarization in the molecules themselves. When chemical action enters into the picture, in other words when chemisorption can also take place, more forces come into action, as in the above-mentioned adsorption of water molecules. We shall discuss the chemisorption of hydrogen and of oxygen on charcoal and on metal surfaces in subsequent sections. 2. The Adsorption on Ionic Surfaces

There have been numerous theoretical and experimental investigations on the adsorption of argon, oxygen, and nitrogen on potassium chloride (126-128) and in this connection we may refer to a survey in Brunauer’s book on physical adsorption (129). There seems to be a general agreement that the most favorable positions for the adsorbed atoms or molecules will be found just above the center of a lattice cell. The electrostatic polarization is minimum a t such spots, but the nonpolar van der Waals’ forces are at their maximum and dominate (130).Drain

66

J. H. DE BOER

(131) drew attention to the fact that the adsorption energy of nitrogen is generally higher than that of argon or oxygen when adsorbed on ionic surfaces, and th at the energies are practically the same on nonionic adsorbents. He ascribes this effect to the quadrupole moment of nitrogen and calculates the contribution of the mutual attraction by the quadrupole of nitrogen and the field over a cubic face of KC1. According t o these calculations the sites just above the center of a lattice cell still represent the most favorable positions; the calculated contribution by the quadrupole attraction is of the right magnitude to explain the extra heat of adsorption of nitrogen on this surface as compared with that of oxygen and argon [about 500 cal./mole (128)l.Drain assumed the nitrogen molecules to lie flat on the surface. Drain and Morrison (132) studied the thermodynamic properties of O2 and Na on rutile experimentally and derived the conclusion that the adsorbed molecules do not freely move over the surface (see also Sec. VII,2) and that both rotational degrees of freedom of nitrogen are considerably hindered whilst oxygen has more freedom of rotation. We saw in Sec. V,5 that molecules with peripheral dipoles, such as OH, N H 2 and, COOH groups are attracted strongly by the electrostatic field of the surface. These dipole forces form the most important contributions toward the adsorption energy when such molecules are adsorbed on ionic surfaces, nonpolar van der Waals’ forces and electrostatic polarization giving smaller contributions (133). Healy et al. (134) studied experimentally the heats of adsorption of many polar and nonpolar gases on polar and nonpolar surfaces by means of their heats of immersion. It was found that the heat of immersion of rutile on a series of straight-chain compounds was a linear function of the dipole moment of the wetting liquid. In a later article (135)-this work was extended and it is shown that nearly the entire heat effect on immersion of the clean solid surface is due to adsorption of molecules in the first layer. From the slope of the line, giving the values found for the net heat of adsorption as a function of the dipole moments, the average field strength, F , of rutile can be found by means of Eq. (22). The experimental value found by these investigators is

F

=

2.72 X lo6 e.s.u.

which is the value of the field strength at the distance to the center of the dipole. By means of Hiickel’s equation for the dependence of the field strength on distance [Eq. (17)J the average distance between the center of the dipole and the rutile surface was calculated t o be 2.08 A. We may remark that the polarizing field of charcoal (Secs. V,7 and VI,1) has approximately the same strength a t this distance (136). I n the

ADSORPTION PHENOMENA

67

latter case the sign of the field is just opposed to th a t of the field over a rutile surface, and consequently peripheral dipoles having their positive poles (H atoms) pointing outward are strongly absorbed by the surface of rutile, but not attracted a t all, even repelled, by the surface of graphite (charcoal). This result may probably be generalized to the statement tha t the electrostatic field over an ionic crystal always attracts peripheric dipoles strongly (the negative ions practically always form the outside layer of these surfaces) and the polarizing field of metals repels them. A further analysis of the various contributions toward the adsorption energies (135) has revealed th at the adsorption energy of alcohol on rutile consists mainly of the contribution of the attraction of the dipole; the nonpolar van der Waals’ forces contribute less than 40% of this part and electrostatic polarization less than 10 %. The adsorption energies of hydrocarbons on rutile are mostly due to the van der Waals’ forces, and half the amount of the van der Waals’ contribution (one third of the total) originates from the electrostatic polarization. The adsorption of water, by its peripheral dipoles, on the surface of some inorganic salts, such as CaFz, is so strong that it cannot be desorbed as such. On heating, a reaction takes place and HF desorbs (1.37) instead of HzO. The resulting salt surface is left occupied with OH groups instead of F ions a t the outer layer. The physical adsorption of HzO molecules has been transformed into a chemisorption of OH groups. On further heating, HzO desorbs by the reaction of two OH groups t o form one HzO molecule and the surface hydroxide is converted into a surface oxide. The conversion of the surface layers of many oxides into surface hydroxide layers is the result of the chemisorption of water on these oxides. On heating, HzO desorbs and the OH groups are converted into 0 ions again. Some organic molecules with OH dipoles behave similarly to water molecules when adsorbed on salt surfaces such as CaFz, BaC12, or NaCl surfaces. On heating they do not desorb as such, but HF or HC1 evaporates and their ions are left behind as an adsorbed layer on the depleted salt surface. Many di- and polyhydroxy anthraquinones, like alizarin (I%), behave in this way. The chemisorption reaction is accompanied by characteristic color changes, which could be studied by their absorption spectra, using completely transparent adsorbent layers as obtained by vacuum sublimation of these salts. The reaction of these hydroxy anthraquinones is strictly confined to the surface layer of ions. With other substances, like picric acid, the reaction resulting in the formation of HF or HC1 and picric ions proceeds further into the lattice and ultimately the whole salt layer is converted

68

J. H. DE BOER

into a picrate layer (139). Other organic molecules, like monohydroxy anthraquinones or ortho- and paranitrophenol, are strongly adsorbed with their OH dipoles, but on heating they desorb as such. The surface reactions with HzO or with alizarin may be used for estimating the surface areas of such salt layers (140). Estimations of surface areas may in other cases also successfully be performed with the aid of the physical adsorption of other molecules with peripheric dipoles, such a s lauric acid molecules on alumina (141). 3. T h e Adsorption of Hydrogen

Various examples of chemisorption of hydrogen on metals or on oxides have already been mentioned in Sees. V,8,b through 11. I n this section we shall discuss some problems related t o the transformation of physically adsorbed hydrogen to chemisorbed hydrogen and the relations between various forms of chemisorbed hydrogen. We shall restrict ourselves to some selected problems and refer to the articles of Eley (142) and of Beeck (143) in this series for many other questions concerning chemisorption. Unlike other gases hydrogen seems t o be chemisorbed on charcoal a t very low temperatures. It is only a t extremely low temperatures, viz., in the neighborhood of 20"K., that we can speak of a physical adsorption of hydrogen on charcoal (144). The heat of adsorption is very low in this case, about 0.37 kcal./mole, which is comparable with the value for helium (0.36 kcal./mole). I n the temperature range from 60' to 90°K. hydrogen and deuterium are chemisorbed on charcoal and on graphite, the heat of adsorption being about 1.5 kcal./mole. The entropy of adsorption points to localized adsorption but it cannot be established whether one or two adsorption sites are occupied. A dissociative adsorption, however, with a free and random distribution of the hydrogen atoms has t o be excluded in this temperature range. It may be th a t dissociation has taken place but th at the atoms cannot move apart (see Sec. VII,3). At still higher temperatures (50" to 100°C.) the chemisorption of hydrogen has a dissociative character and the atoms have moved apart t o form an at-random distribution over the surface; the heat of adsorption is roughly 2.5 kcal./mole. This chemisorption probably does not lead t o the formation of covalent bonds between C atoms and H atoms; such bonds are formed a t higher temperature, hence by a chemisorption process with a n appreciably high activation energy and a far higher heat of adsorption. Burstein (146) in an extensive study of the exchange between Hz and D2 on charcoal found th at this process is governed by the dissociative adsorption of hydrogen in the range of temperatures between 500°K. and

ADSORPTION PHENOMENA

69

90°K. The charcoal had to be degassed thoroughly (at high temperature and for a long time). Adsorption of hydrogen at higher temperatures, viz. 5OO0C., poisons the surface for the exchange reaction, probably owing t o formation of covalent C-H bonds at the surface, as mentioned above. There is one indication that hydrogen may be adsorbed physically on pure-metal surfaces a t very low temperatures with a heat of adsorption comparable t o the value of 0.37 kcal./mole which was mentioned above for the physical adsorption on charcoal. Eucken and Hunsmann (146) give a heat of adsorption of Hz on Ni of 1.2 kcal./mole in the initial stages of adsorption at 20"K., a value that decreases to 0.4 kcal./mole on further adsorption. If the first value indicates active spots, the latter might give the real value. At liquid-air temperatures hydrogen adsorbs chemically on many metal surfaces, provided th a t they are really pure and not contaminated with oxygen or other gases (147). (See also Secs. V,9 and X,4.) The activation energy on pure-metal surfaces seems t o be very small. On oxides or on metal surfaces which are not completely pure, appreciably high activation energies may be found. We saw that hydrogen when adsorbed on charcoal can give rise t o different types of chemisorption. We may ask whether similar effects are found with metal surfaces. Emmett and Harkness (148) found two different sorts of chemisorption of hydrogen on an iron catalyst for NH3 synthesis. These chemisorptions were found in the neighborhood of - 100" and 100°C. respectively and were called type A and type B adsorption. The experiments were repeated by Gauchman and Royter (149), who obtained identical results. According t o Emmett (150) the chemisorption a t -100°C. can be obtained only when the iron is prepared by reduction of the oxide a t about 500°C. in a rapid stream of hydrogen which has been freed of oxygen and well dried before being passed into the reduction zone. It seems as if surface contaminations have something t o do with the type-A adsorption with the lower energy of activation. It is possible that type-A adsorption occurs on parts of the surface that are well reduced and free of contaminations or contain at, least fewer contaminations than the parts on which type-B adsorption takes place. We shall discuss the influence of these contaminations more fully in Sec. X,4. Another possible explanation was recently offered by Zwietering (151), who discusses the possibility th at type-A adsorption is due t o hydrogen atoms having a dipole with the negative pole pointing away from the surface, and that type-B represents the type where the dipole has a reversed direction. Type-A adsorption is comparable to chemisorption on pure-metal surfaces in the form of wires or of films obtained by sublimation; the sign of the dipoles of the hydrogen chemisorption at very low

+

70

J. H. DE BOER

temperatures is indeed such that the negative end of the dipole points away from the surface. Photoelectric measurements a t higher temperatures have often shown tha t chemisorbed hydrogen forms dipoles pointing with their positive ends away from the surface, also in those cases where in later investigations a t low temperatures (liquid-air temperatures) contact potential measurements revealed dipoles of reversed direction. It is, therefore, not to be excluded th at both types of chemisorption may occur at the same metal surfaces and it is t o be expected then t h a t the latter type, which in the case of iron is called A-type chemisorption:

occurs a t a lower temperature (lower activation energy) than the B type: B

Figure 13 may elucidate the energy relations. With type A the hydrogen is the electron acceptor and with type B it is the electron donor. As the ionization energy of a hydrogen atom has a high value (312 kcal./ mole) and energy is gained when an electron is joined to a hydrogen atom (16.4 kcal./mole) and as the work function of the metal is exactly the same as its electron affinity, the transfer of a n electron from the hydrogen atom t o the metal (type B ) costs more energy than the transfer in the opposite direction (type A). It is for this reason th a t the difference between the energy levels DA for type A and level A (the energy level for the Hz molecule; compare Figs. 9, 10, and 11) is far less than the difference between levels D g and A. As the H atom can approach nearer t o the metal in case B than in case A (a positively charged hydrogen atom has negligibly small dimensions), we may expect a higher activation energy and a larger value for the heat of adsorption in case B than in case A , just as has been observed by Emmett and Harkness and later investigators. We shall return to these possible explanations for type-A and type-B adsorption in Sec. X,5 but may remark here that both explanations offered above need not be considered as alternatives. The influence of the contaminations is probably such th at they facilitate the formation of the type-B dipole. Both type-A and type-B adsorption on iron have a poisoning effect on the Hz and Dz exchange on iron catalysts a t very low temperatures

71

ADSORPTION PHENOMENA

(- 196OC.). We have here the same relationship as we saw with the adsorption of hydrogen on charcoal. There is, apparently, a t those low temperatures a type of chemisorption leading to far smaller heats of adsorption and heats of desorption than are caused by the processes which dominate at higher temperatures (162). Whether this low-temperature chemisorption is of a dissociative type or not cannot be decided yet; because of the H2 and D2 exchange we might be inclined to believe it t o be of the dissociative type. The bond strength between the two H atoms must have been appreciably loosened anyhow. One might suppose that

L

I)

1

3

2

4

5

6

7

-r

FIG.13. Potential curves giving a possible explanation for the existence of two chemisorptions of hydrogen on metals.

a t liquid-nitrogen and liquid-air temperatures other electrons of the metals are involved in the chemisorption processes than at higher temperatures, where probably d electrons participate in the bonding. The Hz and Dz exchange also proceeds easily (153), that is, a t 90°K. on the surfaces of some oxides, such as Cr203.In a recent study Molinari and Parravano (164) showed ZnO t o be also a good catalyst for this reaction. Pure, nonsintered ZnO gives the exchange reaction only very slowly. As we saw in Sec. V , l l , the ZnO surface may be more stoichiometric than the bulk of the oxide. Molinari and Parravano succeeded in increasing the speed of the exchange reaction (and in lowering its activation energy) by various methods, as by sintering in vacuum, by a reducing activation,

72

J. H. DE B O E R

or by incorporating three-valent ions such as A13+ or Ga3+. All these measures result in changing the ratio between cations and anions in the ZnO lattice in favor of the cations. This means that more Zn ions of lower valency are produced. It is also stated th a t no hydrogen chemisorption could be detected on the nonsintered ZnO, in spite of its larger surface area. The more Zn ions of lower valency, the more sites will be available for hydrogen chemisorption and the exchange reaction. As the chemisorption leading to this exchange will be of the dissociative type, we may conclude that hydrogen atoms will form bonds with the metal ions (probably with those of lower valency). As the heat of adsorption is 5.1 kcal./mole in the case of Crz03, we must conclude th a t the bond strength between the adsorbed hydrogen atom and the chromium ion will be about 54 kcal./mole. Atomic hydrogen adsorbs very strongly on many surfaces including glass (155-157). From the desorption curve (167) (in the form of Hz molecules) i t may be concluded that the heat of adsorption of the hydrogen atoms will be about 43 t o 48 kcal./mole (168). Twice this amountfor the two hydrogen atoms that are formed from one molecule-is less than the dissociation energy of a hydrogen molecule. A dissociative chemisorption of Hi on glass will therefore be of an endothermic character. The activation energy, moreover, will be very high, as the activation energy for the desorption is already nearly 25 kcal./mole. Figure 14 gives potential curves which probably represent the relations in this case. The bond strength between the hydrogen atom and the glass surface is, neverbheless, remarkably high, namely of the same order of magnitude as the strengths of hydrogen in potassium hydride (44.5 kcal./mole) and in arsenic hydride (47.3 kcal./mole). The nature of this bond is unknown, but we may suggest that it is formed by a n electron transfer from the hydrogen atom to one of the metal ions of the glass surface (Ca2+, Na+, Pb2+) which the hydrogen atom can approach very closely. Hydrogen atoms are adsorbed far more strongly on CaFz films obtained by sublimation in a high vacuum (157). The CaFz surface when exposed to atomic hydrogen is covered with adsorbed atoms t o saturation, when there is one adsorbed hydrogen atom for every fluorine ion of the surface (159). The hydrogen is not desorbed at room temperature, and from the rate of desorption a t elevated temperatures a n activation energy of desorption of more than 40 kcal./mole may be estimated (160). The heat of adsorption of atomic hydrogen on CaFz may be estimated from the latter figure to be roughly 60 kcal./mole. As shown in Fig. 15, the dissociative chemisorption of molecular hydrogen on CaFz should be an exothermic process. The activation energy for adsorption, however, is extremely high, and this is probably the reason th a t this chemisorption phenomenon has never been found.

73

ADSORPTION PHENOMENA

E in kcrl/mole 125

I

10c

73

so

25

0

B i

1

2

3

4

5

6

8

inn

-r

FIQ.14.Potential curves representing the adsorption of atomic hydrogen on glass. E in Kcal/,,,,lc 10c

SO

0

-25 1

2

-

3

4

5

6

7

rinA

FIQ. 15. Potential curves representing the adsorption of atomic hydrogen on CaFs films.

74

J. H. DE BOER

The high values of the activation energies connected with the interaction of hydrogen with glass surfaces or with a CaF2 surface may be caused partly by the large difference in distances between the hydrogen molecule and “the surface” on the one hand and between the hydrogen

AH* mo‘ecu‘e

f Hatorn ions a-ions

etc

FIG. 16. Cross section through a CaF2lattice perpendicular t o an octahedral face. F ions form the outside surface. An Hzmolecule cannot come so near to a Ca ion as an H atom can.

atom and the metal ion to which it is bound on the other. The metal ions in CaFz-and in the majority of inorganic salts-are not situated on the outside of the surface; it is the negative ions that form the outside. The situation is visualized in Fig. 16 and the nature of the bond may be pictured as

r, C$+

t ,

or

CaF,

Other salt surfaces, as the surface of LiF, will also be covered to saturation when exposed to atomic hydrogen. Hydrogen atoms colliding on such a saturated unimolecular layer of adsorbed hydrogen atoms will not be adsorbed to any appreciable extent; they will even partly reflect specularly. It is in this way that an apparent discrepancy must be solved according t o which hydrogen atoms are extremely strongly adsorbed t o the surface of such inorganic saIts whiie on the other hand cleavage surfaces of LiF give specular reflection and even remarkably sharp diffraction patterns with beams of atomic hydrogen (161). Atomic hydrogen reflects on the first adsorbed layer of hydrogen atoms, which is so firmly bound that even the chemical reaction of a n impinging atom and a n

ADSORPTION PHENOMENA

75

adsorbed one t o form a hydrogen molecule is not a n important reaction (168). Atomic hydrogen is only weakly adsorbed on the surfaces of ice and paraffins. There is, nevertheless some adsorption on these surfaces at very low temperatures, and a combination reaction

H+H-+HZ takes place under these circumstances. This catalytic reaction is probably of great importance for the formation of molecular hydrogen from E in kcal/rnolc

-

rinA

FIG.17. An endotherniic chemisorption with low activation energies for adsorption and desorption.

the atomic hydrogen in the interstellar gas clouds on the surfaces of interstellar dust. On metallic oxides and salt surfaces, on the other hand, atomic hydrogen is adsorbed strongly and, depending on the relative position of the potential curves, this may lead t o an exothermic or endothermic dissociative adsorption of molecular hydrogen on these surfaces. For the exchange between Hz and DZ, it is necessary that a dissociative adsorption take place. But it is also a necessity that the activation energies for the adsorption and for the desorption be low enough t o allow the reaction t o proceed a t measurable speed. An endothermic adsorption as pictured in Fig. 17 may, therefore, also lead to this catalytic exchange rextion. It is possible t ha t such a situation governs the rather quick Hz and Dz ex-

76

J. H. D E BOER

change reaction on alumina (162), even at -8O"C., where no measurable adsorption is found. A measurable (exothermic) chemisorption is, therefore, not a necessity for catalysis.

4. The C h e ~ ~ s o r of ~ ~Oxygen ~on Whilst hydrogen enters into a chemisorptive bond with charcoal at very low temperatures, oxygen remains physically adsorbed unless relatively high temperatures are reached. At liquid-air temperatures the adsorption entropy of oxygens shows that the adsorbed molecules are completely free t o move and rotate over the surface (163). At room temperature and higher temperatures the van der Waals' adsorption changes slowly into a chemisorption (164). This behavior of oxygen is clearly shown by its properties t o catalyze-by its paramagnetic properties-the ortho-paraconversion of hydrogen. When adsorbed on charcoal a t low temperatures oxygen promotes this conversion, but when adsorbed a t higher temperatures it poisons the effect (165,166). The reaction of oxygen with the surface of charcoal, therefore, requires a n activation energy. I n the case of adsorption on metals, the activation energy may be zero or negligibly small. Oxygen spontaneously forms a chemisorbed layer of surface oxide molecules on a cesium surface a t liquid-air temperatures. It is, however, quite possible that this chemisorption is nondissociative (see later in this section). A molybdenum film, obtained by evaporation in a high vacuum, however, requires a higher temperature t o convert the physical adsorption into chemisorption. This can be shown by the decrease of conductivity of the film by chemisorption of oxygen (167). A similar behavior is found with the adsorption of oxygen on nickel or platinum (168). There are some cases where oxygen is chemisorbed without being dissociated into atoms. As a matter of fact one can easily understand that a molecular chemisorption may be the first step of chemisorption. This first step may well be the formation of an 0; ion, a n oxygen molecule that has taken up one electron. Oxygen molecules have a positive electron affinity of 2 to 3 kcal./mole (1691, which means that this energy is gained when electrons are taken u p by oxygen. 0; ions are found in the normal oxidation of the higher alkali metals, like potassium, when lattices of superoxides result from the reaction of the metal with oxygen (KOz). The paramagnetic and other physicochemical properties of these superoxides suggest them t o be ionic compounds, the negative ions being 0; ions. 0; ions are also formed, as a first step, in electrical reduction processes where O2is reduced at a cathode (170). It is quite reasonable to assume, therefore, th a t oxygen can be chemi-

ADSORPTION PHENOMENA

77

sorbed on some metals as 0; ions or as O2 molecules accepting an electron from the metal t o form a covalent bond with the metal surface. As already stated above, a cesium-metal surface, when exposed to oxygen a t liquid-air temperatures, is spontaneously covered with a chemisorbed layer of oxygen. As the fully oxidized product which is formed by oxidation a t higher temperatures proves t o be cesium superoxide (CsOz), we may assume that the chemisorbed layer formed a t - 180°C. also consists of chemisorbed 0; ions. As the work function of cesium has a low value, the difference in energy between levels A and D in Fig. 18 is rather small and, consequently, minimum E of the chemisorption curve will be appreciably lower than level A . E in kcal/mole

c

E I

1

2

3

4

5

----- m

rhA

FIG.18. Potential curves for the forming of a surface Cs+Oa- layer on Cs-metal.

Other metals, like silver, copper, or platinum, have far higher work functions and it is possible, even probable, th a t the corresponding minima E will be higher than level A (Fig. 19). We have again a n example of endothermic chemisorption. It will hardly be possible to obtain a measurable amount of chemisorbed oxygen in this way, but in catalytic processes other molecules may well be oxidized b y the very active 0; ions, which will form in great numbers when the activation energy E , can be easily obtained from thermal energy. Molecular oxygen ions depicted in Fig. 19 will be catalytically active, though their times of adsorption are very small; molecular oxygen ions depicted in Fig. 18 will not be catalytically active, because their heat of desorption is too high. I n constructing Figs. 18 and 19 we assumed th a t the distance be-

78

J. H. DE BOER

t

100-

,--

D

/

/ I

/ /

C

75.

/

/

0

.--0

/ / /

c

/ /

/

,/

A

-------

----

1 1

2

3

4

5

U r i n A

FIG.20. An alternative set of curves for Fig. 19.

03

ADSORPTION PHENOMENA

79

tween the oxygen and the surface in minimum E is not only smaller than the minimum of the van der Waals’ curve ABC, but also th a t i t falls on the left-hand side of curve BC. It is not impossible that the size of the 0; ion is such that the situation would be better represented by the relative positions of the curves of Fig. 20. There is no principal difference between Figs. 20 and 19. The formation of some organic hydroperoxides b y oxidation with molecular oxygen is catalytically promoted by metals like silver or copper (171). A dissociative chemisorption of oxygen cannot be active in these processes; they probably proceed via the chemisorption of 0; ions (or O2 molecules forming a covalent bond resonating with a n ionic bond). 5 . Optical and Other Physicochemical Changes by Adsorption A potential curve of an endothermically chemisorbed atom or molecule represents an excited state with respect to the normal state of the physically adsorbed atom or molecule. When cesium atoms are adsorbed on salt layers or on cesium oxide, they are adsorbed as atoms and not, as they would be on metal surfaces, as ions. Ionization can be brought about by absorption of light (172) or by thermal excitation (173). The potential curves of the adsorption of cesium on a CaFz surface are given in Fig. 21, which shows that the curve for the ion represents a n endothermic chemisorption. By the absorption of light of suitable wave length the system is transferred from minimum B to a point P of the upper curve and an electron is freed and may be drawn off as a photoelectron. The phenomenon of the selective photoelectric effect could be fully explained by this photoionization process (174).By thermal excitation the transfer can be effected a t point S and this mechanism may serve t o explain the electron emission of oxide cathodes. Point S is reached by taking up a n amount of energy, which may be called the work function of the oxide cathode in this case but which is completely comparable with the energy of activation in chemisorption discussed in Sec. V,9 and subsequently. We shall not discuss these phenomena in .this article but refer t o a book of the author where these subjects are dealt with in detail (174). A similar relationship may be found in the adsorption of other atoms or molecules. Many organic substances with peripheric dipoles when adsorbed on salt layers or on the surfaces of metallic oxides show absorption spectra which are shifted appreciably to the red side of the spectrum. Thus p-nitrophenol, having a maximum of light absorption at 316 mp. when adsorbed on CaF2, has its absorption spectrum shifted to the red side and is yellow instead of colorless (175), its absorption maximum being at 365 mp. (176). Adsorbed on BaFz it shows a n absorption maxi-

80

J. H. DE BOER

mum at 413 my. (177). Similarly phenolphthalein when adsorbed on CaFz shows a bright red color and has a maximum of light absorption at 475 mp. but is red violet in color when adsorbed on BaFz with a maximum of absorption a t 536 mp. (17'8). I n all these cases the shift of the absorption spectrum has nothing to do with salt formation or ionization but results from the fact that, apparently, the excited state of the molecule is adsorbed more strongly than the ground state. The act of adsorption, therefore, decreases the energy difference between the ground level and the excited level (178).

i

3 -rinA

4

6

CO

FIG.21. Potential curves for the thermal and photo ionization of an adsorbed cesium atom.

Similar color changes were reported later by Weitz and his collaborators (179), apparently without knowing the older work of the author (180). They describe the bright-red coloration resulting from the physical adsorption of phenolphthalein. I n all these cases we are concerned with physically adsorbed molecules, pointing with their peripheric dipoles to the negative ions of the surface. The excited states of these molecules are far more polar in character; light absorption causes an electron shift in the molecule in a direction away from the surface, resulting in a far stronger bond with the negative ion on which the molecule is adsorbed (181). We may illustrate

ADSDRPTION PHENOMENA

81

this point with just one example. The structure of the ground state of p-hydroxyazobenzene is given mainly by the formula

with the normal resonance structure for the benzene rings and the OH group. The molecule is oriented with the OH group t o the negatively charged oxygen ions when adsorbed on alumina. The excited state will m:i.inly have the structure

light absorption causing a iransition of an electron from the 0 of the OH group t o an N of the azo group. The positively charged OH group of this excited structure is more strongly adsorbed on the alumina than the ground state, hence a shift of the absorption spectrum to longer wave lengths; p-hydroxyazobenzene is deep red in color when adsorbed on dry alumina. Its original yellow color can be shown by admission of water, water molecules replacing the organic molecules (182).

VII. MOEIILITY AND ORIENTATION 1 . Mobility o n Charcoal

The study of the adsorption entropies gives a great deal of information about the mobility of a.dsorbed atoms or molecules along the surface, as Kemball (185) has shown. A systematic study of the entropies of gases adsorbed on charcoal (184) along similar lines showed th a t many gases, including CO, 0 2 , N2, and many hydrocarbons behave as twodimensional gases, moving freely over the surface while rotating freely. At lower temperatures and at, higher degrees of occupation there is some restriction of the free movements, which increases in strength as the temperature is lower. The free translatory movement is the first t o be restricted at lower temperatures. The free rotation is hardly affected a t all. A distinction between the restrictions of the translation and the rotation could be made by comparing various gases, including the noble gases which have no rotations. I n drawing the conclusions we found th a t the area occupied per molecule also proved to be a n important entity. Rotating molecules show a molecular surface area which is approximately equal t o the value of the two-dimensional van der Waals’ b(b2) (185,186). Carbon disulfide molecules, methyl-, ethyl-, and n-propylchloride molecules, and also diethylether molecules, are more or less strongly hindered in their movements a t not too high temperatures. Their rota-

82

J.

n.

DE BOER

tions, however, are not hindered. The same holds for n-pentane and n-heptane molecules; these have lost their translatory movements, but not their rotations, when adsorbed on charcoal a t room temperature. The values for the molecular areas of these molecules as found experimentally on charcoal surfaces are compared with their bz-values in Table 111, which also shows some values of molecular areas occupied by the molecules on other, polar, surfaces. These figures have either been taken from the article by Livingston (185) or have been newly determined (18.4). TABLE 111 Molecular areas in A2 Gas

cs2 CzH,C1 C3H7C1 (C2Hs)zO C4Hio GHi2

b2

On charcoal

23.5 23.7 31.7 36.0 33.2 37.3 47.0

25.2 24.8 32.1 34.1 37.1 46.2

On polar adsorbents 37.9

44.6

> 44

59.6

The approximate equality of the bz value of the two-dimensional van der Waals’ equation and the molecular area is not so obvious as might be thought (186a). The bz value is by definition twice the surface area of a molecule, the diameter of which is d. This diameter is derived from the distance of approach of two colliding molecules. The molecular surface area, however, is derived from the density of liquids, each molecule being assigned its own sphere with a diameter dmin,on the assumption th a t the molecules are closely packed. As it happens t o be th a t for a great number of molecules d,, = about 1.37d a two-dimensional closely packed hexagonal arrangement will allow the molecules a surface area of

Nd&, fi = 0.865&

=

1.62d2

As b z = g r d 2 = 1.57d2 we see that the surface area which the molecules really occupy-provided that they are freely rotating-is reasonably given by the bz value. 2. Orientation or Rotation There is, apparently, hardly any orientation of the molecules when they are adsorbed on charcoal. This is due t o the rather unspecific nature

ADSORPTION PHENOMENA

83

of the polarizing forces emanating from the charcoal surface (Sec. V,7). The same may hold for the physical adsorption on metals. In physical adsorption phenomena on dielectric ionic substances conditions are different. As we have seen already, there are more specific fields over such surfaces, reriulting in alternating fields when ions of opposite signs are being passed over via the centers of surface cells (Secs. V,4 and 5 ) . Many molecules having dipoles (peripheral or not) or quadrupoles may, therefore, tend to be oriented and to lose their rotations. As discussed in Sec. VI,2, Drain and Morrison (187) assume nitrogen molecules to lie flat on the surface of rutile, owing to their quadrupole moments. The molecular areas of other molecules, as experimentally found on polar substances, also suggest flat positions in many cases. Molecules with peripheral dipoles are directed with these dipoles to negative spots of the surfact:, as discussed in Secs. V,5 and VI,2 and 5. Such molecules may erect estch other so as to leave the dipoles directed to the surface and the rest of the molecules, being parallel to each other, pointing away from it (188). We shall not discuss these points in this article but only remark that they may be of great importance for the understanding of some selective catalytic processes. 3. .Flopping Molecules

If the entropy data indioate a localized or site adsorption, the adsorbed molecules need not be considered as immobile in the course of time. The mere existence of a,n adsorption equilibrium between a gas and an adsorbed layer implies a mobility of the adsorbed molecules along the surface. Owing to the regular pattern of crystalline matter, the surface of crystalline adsorbents will show periodical fluctuations. There will, consequently, be a regular alternation of spots where the strength of the adsorption forces is somewhat greater than the average and others where it is lower. The adsorption energy may, therefore, be different if the molecule is situated on the top of a surface ion or if it is just over the center of a surface cell. If, in the case of dynamical equilibrium, a molecule can pick up such an amount of energy from the thermal energy fluctuations that it can desorb, we may expect that by assembling a smaller amount of energy it will be able t o move from one spot to another without losing its contact with the surface altogether. In the case of normal dyi~amicalequilibrium of adsorption, a molecule striking the surface will stay there for an average time, T , which we may call the time of adsorption. During its stay of T sec. a t the surface, it remains for a much shorter time, T' sec., at a given site. After this average halting time, 7 ' ) it jumps to a neighboring site, where it stays again

84

J. H . DE BOER

for a n average time, T I , after which it jumps again, etc. During its adsorption time of 7 sec., it will hop T / T ' times to one of the neighboring sites, thus moving along the surface. The time of adsorption, 7 , may be related to the heat of adsorption by the expression 7 = T0eQ./RT (37) where T O is a constant, which we shall discuss in the next section and is the heat of adsorption. Similarly we may relate the halting time, with the activation energy for the hopping movement, Em, by

Qa

r',

The constant 7; is of the same order of magnitude as 7 0 . Em may be considered t o be the difference between the heat of adsorption when the molecule is adsorbed on a preferential site of the normal regular surface pattern and the heat of adsorption of the same molecule adsorbed on a spot just in between two such preferential sites. Em is, therefore, essentially smaller than Qd. If Em is very small with respect t o R T , the molecules will move freely over the surface; we are then dealing with the two-dimensional gases of Sec. VI1,l. I n many cases of physical adsorption on polar surfaces the value of Em may be of the order of one third t o one half the value of Qa. Let us assume Qa to he 10 kcal./mole and Em t o be 5 kcal./mole. The time of adsorption 7 a t room temperature is then roughly 3 X sec. and the halting time T I is roughly 5 X 10-lo sec. The molecule, therefore, will during its stay a t the surface make a n average of 6,000 hops from one spot to another. Let every hop be equal t o the shortest distance between the surface atoms of the adsorbent, hence about 3 8. The molecule would then cover a distance of roughly 2 X cm. This, of course, is not its rectilinear displacement, because the directions of the successive hops are arbitrarily distributed so that the displacement of the molecule is far smaller than the total length of all the hops together. Migration of adsorbed niolecules over the surface of the adsorbent in the case of physical adsorption or of adsorbed atoms or radicals in the case of chemisorption is a normal phenomenon provided the temperature is high enough to enable them to overcome the energy of activationif there is one. I n the case of adsorption of cesium ions on tungsten the energy of activation for migration is about 14 kcal./mole. A t room temperature the adsorbed ions may be considered t o be localized a t definite adsorption sites, but a t elevated temperatures there is a vivid migration. Oxygen atoms chemisorbed on a metal surface a t room temperature may be considered t o be bound to definite spots; a t higher temperatures, how-

85

ADSORPTION PHENOMENA

ever, they migrate and a t sufficiently high temperatures they may even behave as a free mobile two-dimensional gas. (See end of Sec. VII,4.) Migration of atoms over the surface of their own crystals plays a n important role in crystal growth. It is very likely that pairs of ions of opposite sign move far more freely over the surface of a n ionic crystal than single ions do (189). Such a migration of ionic pairs, probably by hopping movements, will play an important role in sintering phenomena of catalysts, leading t o a considerable reduction of the large surface area of the capillary systems of rnicroporous substances.

4. T h e T i m e of Adsorption I n the previous section we related the time of adsorption, heat of adsorption, Q a : 7 = eQz/RT

r,

t o the

(37)

This equation originates from Frenkel (190), who identified the constant the time of an oscillation of the adsorbed molecule, namely with the reciprocal frequency of tthe vibration perpendicular t o the surface. The identification of r o with the time of an oscillation of adsorbed sec., the molecules leads to the assumption th at r o will be about latter being the time of oscillation of bound atoms (191). I n many cases of adsorption 70, indeed, proves t o be about see., although this has nothing t o do with a time of oscillation of molecules. Frenkcl’s original derivation holds for the special case where the perpendicular vibration of the adsorbed molecules contributes t o the entropy, t ha t is, for the special case which Kemball (192) has termed “supermobile adsorption.” The adsorbed molecules, which in this case move freely over the surface, have not lost the third degree of translatory freedom altogether; this freedom is transformed into a freedom of vibration, perpendicular t o the surface. The reciprocal value of the frequency of the vibratory flight, which these molecules make over the surface, equals 1 0 . I n all other cases r o has the dimensions but not the meaning of a reciprocal frequency (193). The time of adsorption can be calculated by means of statistical mechanics from the partition functions of the gaseous and the adsorbed molecule (193). The equilibrium condition for the adsorption may be written as r o with

where J t r , and ofvihr stand for the partition functions of the translations, rotations and vibrations of the gaseous molecules respectively;

86

J. H. DE BOER

.frat, and are the corresponding functions for the adsorbed molecules. N, and Na stand for the number of molecules in the gaseous and adsorbed phases respectively and Qo is the heat of adsorption at absolute zero. The translational partition function of an ideal gas is equal to

Jtr,

where m is the mass of a molecule, k and h are the constants of Boltzmann and of Planck respectively, and P is the pressure. Similarly we can write for the translational partition function of an ideal two-dimensional gas :

where 0 is the total area of the adsorbent. I n the evaluation of Eq. (38) the gas may be considered as an ideal gas. The two-dimensional gas-we consider mobile adsorption-need not be a n ideal one and we may write aftr

=

at e e t i

X

af t r I _

offme tr

Inserting this in Eq. (38) we obtain with (38), (39), and (40)

X

afrot

X

afvihr

X

eQoIRT

(41)

The internal vibrations of the gaseous and the adsorbed molecules may be taken to be equal. The adsorbed molecules, however, may have a vibration perpendicular to the surface which has taken the place of the lost translation (see above). We may, therefore, write afvzhr

=

ofvibr

X fs

(42)

where fz is the partition function of the vibration of the adsorbed molecule perpendicular to the surface. The amount of molecules adsorbed per unit area, CT = N,/O, may then be written as

According t o the kinetic theory of gases the number of molecules falling on 1 cm.2 in 1 sec. is equal to n=

n r

d 2 x T

(44)

ADSORPTION PHENOMENA 7

87

and n are related to u by (194) (45)

u = n7

Solving 7 from (45), (44), and (43) results in

Comparing (46) with (37) we see that the two equations differ in th a t the exponent of (46) contain,s the heat of adsorption a t absolute zero and not a t the temperature where the adsorption takes place and th a t the expression in brackets in (46) stands for r 0 : TO

=

&fh aftr

afmt

a ireetr gfrot

(47)

We may evaluate T O in a few selected cases. 1. For a freely translating and freely rotating adsorbed molecule, which moreover has retained part of the entropy of the third direction of translation, by converting it into an entropy contribution from the vibration perpendicular to the surface (fz > 1) (supermobile adsorption), we may put aftr

=

a5.e tr

afm

=

sfrot

TO

=

-fs

and

so that h kT

If the frequency of the supermobile adsorption v, is so low th a t hv, is much smaller than the kinlstic energy kT, the partition function

hence To

=

1

VZ

(49)

This is the case which Frenkel had in mind when he derived Eq. (37). We shall illustrate this case with a few examples. Cassel and Neugebauer (195) estimated the adsorption of xenon on mercury. An evaluation of their data (196) leads t o the conclusion that the adsorbed xenon molecules are not in their lowest state of vibrations perpendicular to the surface,

88

J. H. DE B O E R

but that the entropy contribution of this vibration is still 7.2 entropy units (at 283°K.). This means that the frequency of the vibration is v, = 4.3 )( 10" set.-'

and consequently TO =

2.3 X 10-'2 sec.

Argon adsorbed on charcoal a t 215'K. has still a slight supermobile character (197), the entropy contribution of the perpendicular vibration being 2 entropy units. Its frequency is v, = 4.5 X 10l2 set.-'

and consequently T~

=

2.2 X

sec.

From these figures for 7 0 and the heats of adsorption of xenon on mercury (3,255 cal./mole) and of argon on charcoal (3,470 cal./mole) and the temperatures of the experiments (283°K. for xenon and 215°K. for argon) the times of adsorption may be calculated with Eq. (37) [or (46)]. We obtain T = 7.8 X 10-lo sec. for xenon on mercury and T = 8.0 X sec. for argon on charcoal Despite the somewhat larger heat of adsorption and the lower temperature, the time of adsorption of argon on charcoal therefore is practically the same as the corresponding figure for xenon on mercury. The higher the entropy, hence the more mobile the adsorbed molecule is, the longer is its time of adsorption, other quantities, suck as heat of adsorption and temperature, being equal. 2. If the perpendicular vibration is practically in its lowest state, because the vibrational quantum hv, is much greater than kT and, therefore, does not contribute to the entropy, the partition function f, = 1 and we obtain

This holds true for all adsorbed molecules which, on adsorption, have just lost one degree of translatory freedom and which, in the adsorbed state, move with complete freedom over the surface as molecules of a two-dimensional gas. They also rotate freely. It was mentioned in Sec. VII,1 that many gases which are physically adsorbed on charcoal belong t o this category and we may expect this to be also the case when they are physically adsorbed on metal surfaces.

ADSORPTION P H E N O M E N A

89

sec. and we see At 300°K. their value of T~ amounts to 1.60 X that, although the meaning is different, the numerical value of h / k T in the temperature range where the adsorption is usually measured is of the order of sec., which also happens t o be the order of a reciprocal molecular frequency. We shall just give one example. The adsorption of ethyl chloride on charcoal a t 331°K., measured by Pearce and Taylor (198), can be described as a case where the adsorbed molecule moves and rotates freely over the surface. When the surface is covered with adsorbed molecules t o about 27%, the heat of adsorption is 9.1 kcal./mole. The value of h/k?' is 1.45 X 10-13 sec.; hence TO =

1.45 X 10-13 sec.

which leads t o T =

1.45 X

sec.

3. If the two translations along the surface are seriously hindered, uftr

uffrce t r

and we obtain smaller values for T~ [Eg. (47)]. The same holds when rotations are hindered or lost during adsorption, because then UfFUt

< Jrat

Kemball (199) has shown that the adsorption of benzene on mercury may be described by assuniing the molecules to be oriented with their rings parallel to the mercury surface while moving freely over this surface and rotating in the plane of the ring. Numerically, we have (200) fi

,fro,

= 1 =

6.8 X lop4 X

Jrot

and 1.1 X

sec. 4. When, in the act of adsorption, all three translations are lost we are concerned with a case of localized adsorption. Then, T~ has to be calculated in a somewhat different way, because a localization partition function, taking account of the number of possible ways of distributing Nu molecules over N , adsorption sites, comes into the picture. Instead of Eq. (46) we obtain 70

=

where 8 indicates the fraction of the adsorption sites that is occupied by adsorbed molecules, 8 = N,/N, (52)

90

J. H . DE BOER

and f z and f, are the partition functions for the vibrations of the adsorbed molecule along the surface. The effect of the internal vibrations of the molecules does not (or hardly) contribute and has been left out of consideration. A formal application of Eq. (51) leads to the conclusion th a t the time of adsorption appears to be dependent on the amount adsorbed and drops to zero for a fully occupied layer (200). This is caused by the assumption of the “monolayer ”-conception, according to which molecules striking the occupied parts of that layer are supposed to reflect, without being adsorbed even for a moment. Equation (51), therefore, gives 7 as the average value for such an “ideal monolayer.” The real time of adsorption is given by Eq. (48) by omitting the factor (1 - 0) and we see that it can be written again i n the form of Eq. (46) by putting

If 7 0 is evaluated in such a way th at Jz = f, = fi = 1, and SO all vibrations of the adsorbed molecule (atom or radical) are in their ground state, we see that the value is lower than that of Eq. (50)-the case of the two-dimensional gas with free translations and free rotations-by an amount

Taking the adsorption of water on charcoal (201) as an example and accepting for a moment that the adsorbed water molecules have rewe obtain (300°K.) tained their rotations and taking NB/O = 10i5/~n1.2, T~

=

sec.

As the water molecules have also lost some of their rotational degrees of freedom, the actual value is lower, viz., T~

=

10-lB sec. (200)

Comparing the adsorption of ethyl chloride on charcoal (see 2 above) with the adsorption of water on charcoal, mentioned here, we obtain TABLE IV C,€I&l

T Qa ro r

331°K. 9 . 1 kcnl./mole 1.45 X sec. 1.45 X lo-’ sec.

€120

300°K. 10.8 kcal./mole see. 1.8 X sec.

ADE,ORPTION PHENOMENA

91

Despite the somewhat lower temperature and the higher heat of adsorption, the time of adsorption of the water molecules is roughly ten times shorter than th at of ethyl chloride molecules. It is again the influence of the entropy on the free energy t h a t causes this effect. The time of adsorption is higher the higher the mobility of the adsorbed molecule. I n those cases of chemisorption where the adsorbed atoms or radicals are fixed by covalent bonds, each t o its own atom on the adsorbing sursec. At higher temperatures face, T~ will be appreciably lower than there is always an appreciable mobility of the adsorbed atoms. As long as their movements can be described by the concept of hopping molecules, which mesns th at the activation energy for their movements is appreciably higher than RT, this mobility has no effect on the time of adsorption. At still higher temperatures the activation energy for the mobility may be comparable with R T and migration of the adsorbed molecules may be better described as a more or less free mobility. Johnson and Vick (202) in 1935 measured the time of adsorption 7 for oxygen atoms a t a tungsten surface in the neighborhood of 2,200"C. They found the following figures: a t 2,548'K.: 7 = 0.36 sec. a t 2,362"K.: 7 = 3.49 sec. which can be represented by = 8 X 10-14 X e147.000/RT As a t this temperature h /kT = 2.2 X we see that the chemisorbed atoms seem to behave as a supermobile twodimensional gas under these circumstances. 5 . Mobility and Reactivity Two species of molecules reacting on the surface of a catalyst may both be bound by chemisorption forces, or it may be that only one of the reacting species is bound. I n the latter case-which is known as the Rideal-mechanism-both sorts of molecules hit the surface of the catalyst, but only one of the splxies is chemisorbed. The molecules of the other sort hit the chemisorbed molecules and form a n "activated complex" which leads t o reaction. They may, however, also be adsorbed by van der Waals' forces and react with the chemisorbed reaction partner from a van der Waals' layer. It may be stated th a t entropy considerations show that such reactions will proceed more easily the smaller the mobility of the adsorbed molecules is, other quantities, such as the activation energy of reaction, being the same.

92

J. H. DE BOER

This also holds for the other mechanism-known as the LangmuirHinshelwood mechanism-where the two reactant molecules have to be chemisorbed side by side. Such pairs can be formed statistically by the molecules hitting the surface and being bound by chemisorption forces a t the very spots where they hit (localized adsorption) or they may result from collisions of the molecules moving along the surface of the catalyst (mobile adsorption). It may again be stated th a t the formation of pairs in the case of immobile localized adsorption is more favorable for reaction than in the case of mobile adsorption. As a n example we shall take the unison of two chemisorbed hydrogen atoms forming a hydrogen molecule which desorbs. When hydrogen atoms are adsorbed on a glass wall they are very strongly bound; theii rate of desorption a t room temperature is negligible (Sec. V1,3). There is, however, a very slow desorption of molecular hydrogen, resulting from the combination of hydrogen atoms on the surface. The rate of this reaction may be calculated with the aid of Eyring’s (203) theory of absolute reaction rates. The experimental rate of the evolution of molecular hydrogen from a layer of hydrogen atoms on glass, which is only partially occupied, may be understood by accepting the concept of localized adsorption of the hydrogen atoms and an energy of activation for the reaction of 25.1 kcal. /mole. If, however, the hydrogen atoms are assumed t o move freely over the surface, this energy of activation would give a rate of production of hydrogen ten times lower than the observed rate. With mobile atoms the energy of activation is required to be lower; 23.7 kcal./mole would then give the observed rate ($04). 6 . Induced

Mobility of Atoms of the Surface

Potassium metal in bulk is not attacked by dry molecular hydrogen at room temperature. Atomic hydrogen, however, reacts with the surface of metallic potassium. About 20 years ago this reaction was studied thoroughly by Lukirsky and Rijanow (206). By considering the photoelectric behavior when known amounts of atomic hydrogen were taken up by potassium a t various temperatures they found that a t -180°C. only a unimolecular layer of chemisorbed hydrogen atoms is taken up. Thus, the photoelectric activity in visible light is reduced to a low value, which points t o the fact th at the hydrogen atoms form surface hydride molecules (206). The resulting electrical double layer, with its negative side pointing away from the metal, increases the work function of the metal to such a value t ha t visible light cannot release electrons from the metal. When,

ADSORPTION PHENOMENA

93

after the forming of this urlimolecular layer, the rest of the hydrogen is pumped away and the metal is heated up to room temperature, a change will occur in the surface structure which causes a high photoelectric sensitivity. The same high photoelectric activity is found when atomic hydrogen is taken u p by the surface of the potassium a t room temperature. room temperature-increases linearly The photoelectric activity-at with the amount of hydrogen atoms th at are chemisorbed, until a maximum is reached, after which the activity falls again t o a value which is lower than t hat of the pure metal (Fig. 22). The maximum is reached when the amount of hydrogen taken u p is, again, just sufficient t o form

-

time

FIG.22. Amount of hydrogen taken up by potassium (curve a) and photoelectric current of the layer (curve b ) ail a function of the time of exposure to hydrogen atoms (20,5).

one unimolecular layer of cliemisorbed hydride molecules. Apparently, a t room temperature another layer is adsorbed in addition to the first one. The explanation of these phenomena is that a t room temperature the mobility of the potassium atoms of the surface of the metal is high enough to cause a migration of these atoms onto and over every area of surface hydride which is formed by the take up of hydrogen atoms. Consequently the photoelectric current rises in direct proportion t o the number of atoms adsorbed and thus with time (Fig. 22). The photoelectric cathode which is formed by these phenomena may be represented by t,he symbol (206):

[K]-KH-K After the whole surface has been covered with such a composite layer built up of a unimolecular hydride layer and a superposed unimolecular layer of potassium, hydrogen is taken up by the latter film, converting it into

[K]-KH-KH

94

J. H. DE BOER

which shows no photoelectric sensitivity. This is the final state a t room temperature. The final state a t - 18OoC., however, is represented by

[K]-KH which, again, shows no photoelectric sensitivity. If, a s reported above, this system is heated t o room temperature, it converts t o [ K]-KH-K

which shows a maximum of photoelectric sensitivity and can take u p another amount of atomic hydrogen while losing its photoelectric sensitivity a t the same time. Other investigators have prepared similar and more complicated layers with hydrogen on top of metals along the same principles, hence the combined action of chemisorption and surface migration. A similar behavior may be shown by oxygen. When oxygen is adsorbed on cesium a t - 18OoC., the photoelectric current rises sharply, passes a maximum, and falls to zero when oxygen is supplied continuously (207). Apparently the same happens as in the system potassiumatomic hydrogen a t room temperature. At - 180°C. cesium atoms are mobile enough to migrate on top of the first layer of surface oxide, whereupon these atoms are oxidized in their turn. The final state a t - 180°C.on continuous supply of oxygen-seems t o be a bimolecular layer of cesium oxide on top of the cesium metal. The metal is now protected against further attack. When oxygen is acting on cesium a t room temperature the mobility of the cesium atoms is so high th at a polyatomic layer of cesium atoms forms on top of any oxide layer that has been formed. We may also say that, apparently, the cesium oxide is absorbed by the metal, i.e., it is dissolved in the cesium. The appearance of the surface is unchanged, and so is its photoelectric behavior. Only after the cesium has been almost completely oxidized on continuous supply of oxygen, does a very thin adsorbed cesium layer appear on the cesium oxide and does the photoelectric current temporarily rise sharply until, on further supply of oxygen, also these last cesium atoms are converted into oxide. The mobility of potassium is smaller than th a t of cesium. The exposure of potassium metal to oxygen a t room temperature leads to a complete oxidation of the metal, but, unlike cesium, the potassium atoms, during the oxidation, do not form a polyatomic layer on top of the oxide formed, but a layer of only one-atom thickness. Consequently the photoelectric current rises and continues to rise with the amount of oxide formed until a maximum is reached-when 4 X lop4 g. of oxygen has been taken up/cm.2 of potassium surface (208)-after which it decreases

ADElORPTION P H E N O M E N A

95

again. Apparently the thickness of the oxide layer forms such a hindrance for the passage of the electrons t h a t the photoelectric sensitivity has t o fall. ’ The mobilities of the atoms of alkali metals are high. Atonis of other metals, however, may also migrate onto and on surface oxide layers. The oxidation of many metals starts in a way similar to the mechanisms discussed above. An oxygen molecule chemisorbed on the metal surface is converted into an oxygen ion by electrons from the metal. A metal ion may move up to take a position next to the chemisorbed oxygen ion. Thus a unimolecular layer of metal oxide molecules may be formed. More metal ions can move on top of this layer and the oxidation may continue. Also other mechanisms of transport of ions (metal ions from the metal through the oxide layer via vacant places in the lattice of the oxide layer or oxygen ions moving in the opposite direction) may play a role. We touch here the general phenomena of the so-called “tarnishing” reactions, which have been studied from many sides during the last 20 to 25 years; we may refer to a few outstanding articles by Wagner (209) and M ot t (210) or refer to the excellent little monograph of Rees (211). I n special cases the action of oxygen does not lead to a complete oxidation of the metal but to a reaction which is restricted to the surface or t o the surface region. These are the cases which are of importance for catalysis. The chemisorbed, catalytically active species need not necessarily be bound to the very outside of the catalyst; we may also find them in the second or third layer under the surface; for instance, in the surface region, let us say in the region of the first four or five layers. We have already discussed some consequences of this conception in Sec. V,11 and it may suffic~: here t o point to the fact that the rate of diffusion from this surface region to the surface and in the reversed direction is often sufficiently high to consider the surface region of, say, four layers as a potential source for catalytically active species. I n zinc oxide and in a-FesOs such surface regions can play a role in the equilibrium with oxygen a t low pressures (212), giving rise t o an induced heterogeneity of the surface. Frumkin and collaborators (213) investigated the action of oxygen on iron which leads t o passivity. Here again, iron atoms can migrate on top of the first unimolecular layer of oxide, giving rise to an increased capacity for electron emission. Depending on temperature, the oxidation process ends after formation of two to four layers of oxide on the iron metal, which protect the metal against further oxidation in the same way as two layers of cesium oxide protect cesium at - 180°C. A t 100°C. a maximum in potential difference with tungsten (minimum in work function) appears when 22 X l0l4 moles oxygen are adsorbed/cm.2 of real surface

96

J. H . DE BOER

of the iron (Fig. 23); at 270°C. a maximum is found when 72 X l O I 4 moles of oxygen have been taken up. Afte,r the maximum has been passed, the original potential difference with tungsten is reached a t 60 X 1014 and 100 X 1014 moles 02/cm.2 respectively. The higher the temperature, the thicker the layer which is formed by the simultaneous action of chemisorption and migration. Summarizing, we may state th at by the migration of metal atoms (ions) onto and over oxide (or other) layers formed by chemisorption, oxygen (or other atoms or ions) may be incorporated into the metal lat-

FIG.23. Change of potential difference, a t two teniperatures, between W and iron when oxygen is taken up by the latter (213).

tide, or oxides (or other compounds) may be formed. There are many cases where the rate of diffusion in the surface region, hence in a region of three to four layers under the surface, is great enough to enable the adsorbed or absorbed atoms to take part in catalytic reactions. 7. Solution in the Adsorbent (Catalyst)

Many two-atomic gases can dissolve into metals. They are split into atoms and diffuse into the metal in atomic form. I n the dissolved state they behave as having a positive or a negative charge (Zf4). Hydrogen atoms, dissolved in palladium, nickeI or iron are, partly, present in the form of protons (215); oxygen atoms in solution in zirconium are, partly, negatively charged (216). I n many cases the dissolution of the gas into the metal may be an exothermic process; in other cases, however, including the dissolution of hydrogen into nickel, iron, and platinum, the process is of an endothermic nature. I n the latter case the solubility of the hydrogen increases with increasing temperature. It is a n important fact th at in many cases the rate of diffusion of the dissolved atoms in the metals is high as compared with the rate of dissolution of the gases into the metal or the rate of desorption from the

ADSORPTION P H E N O M E N A

97

metal. The temperature coefficient of the diffusion of hydrogen through nickel or platinum is entirely given by the heat of desorption from the surface of those metals (217),where molecular hydrogen has t o be formed on desorption. Molecular hydrogen does not dissolve easily into iron with a smooth surface a t temperatures below 200°C. Atomic hydrogen, however, enters the iron easily even a t room temperature (218). I n the case of molecular hydrogen it is the activation energy at the surface which governs the process. On a smooth or con1,aminated iron surface, it is the rate of chemisorption which governs the iota1 rate. We shall return t o this special case in Sec. X,4. Oxygen may be taken up by platinum. A study of this reaction has revealed that at some stageis at not too high temperatures, e.g., 200"C., there may be a temporary increase of the work function of the platinum, showing that some of the oxygen is still at the surface; after some time the work function returns to the proper value for platinum, the oxygen atoms having diffused into the metaI (219). Nitrogen may in atomic form dissolve in and diffuse through a! iron. Up to 0.4 atomic % may be taken up. Iron and nitrogen may form compounds of variable composition. e iron nitride, for example, has a close-packed hexagonal lattice of iron atoms with nitrogen atoms occupying some of the octahedral interstices in an ordered manner. The composition may be between 35.5 and 99.3 N atoms/100 iron atoms at 400°C. The N atoms diffuse readily through this lattice. Goodeve and Jack (620) studied the evolution of Nz gas from this solution and found that the nitrogen atoms on reaching the surface behave as a two-dimensional gas and combine to form Nz molecules, which evaporate. The rate-determining process is the unison of two ehemisorbed N atoms to form a N?:molecule. Oxygen and nitrogen may also dissolve in the lattice of zirconium or titanium, where again they are situated a t interstices of the metallic lattice. I n these metals the atoms are too strongly bound to be of any value for catalytic purposes. The diffusion of chemisorbed atoms into the metals and the migration phenomena of the previcms section may show that in chemisorption phenomena it is distinctly possible to find the adsorbed atoms below the surface of the adsorbent. Chemisorbed atoms of the surface region, hence including those below the actual surface, may take part in catalytic processes. When the nitrogen atoms of the t iron nitride phase meet hydrogen at the surface of their lattice, ammonia is formed. Hydrogen atoms dissolved in nickel may well react with chemisorbed olefines reacting from underneath. There is also the possibility of hydrogen compounds, for example, hydrocarbons, splitting off a hydrogen atom which immediately

98

J. H. DE BOER

disappears into the lattice of the catalyst, the rest of the molecules remaining chemisorbed at the surface. Such chemisorption phenomena play a role in catalytic isomerization processes of vegetable oils on certain nickel catalysts.

VIII. PHYSICAL ADSORPTIONPHENOMENA AT HIGH DEGREES OF OCCUPATION 1. General Remarks

Most of the preceding sections were concerned with singly adsorbed atoms or molecules. At higher degrees of occupation, when the distance between the adsorbed atoms or molecules decreases, the interaction forces between them may become strong enough t o influence the strength of the adsorption. These forces may be repelling forces or mutual attraction forces. I n theoretical derivations of adsorption isotherms, giving the amount of adsorbed molecules as a function of the equilibrium pressure of the coexistent gas or the equilibrium concentration in the coexistent solution, it is mostly assumed that the heat of adsorption is the same over the entire surface of the adsorbent. As we discussed in Sec. V,12 such a conception can hardly be maintained for practical adsorbents. Chemisorbed atoms or ions alter some properties of the adsorbents rather seriously. It is especially the work function for electron emission (and, therefore, also the electron affinity) which is either decreased or increased. Conversely this change affects again the adsorption energy. As a result of these various effects the strength of the adsorption may depend on the degree of occupation. The present section deals with physical adsorption phenomena a t high degrees of occupation. Chemisorption phenomena a t high degrees of coverage will be dealt with in See. IX. 2. T h e Heat of Adsorption as a Function of the Degree of Occupation in Physical Adsorption Phenomena on Conducting Adsorbents

As we discussed in Sec. VI,1 physical adsorption on charcoal and on metal surfaces is caused by the polarization of the adsorbed molecules in the electronic field over the surface of the conducting adsorbent (Sec. V,7) , together with the nonpolar van der Waals’ forces between the adsorbent and the adsorbed molecules (Sec. V,2). As mentioned in Sec. V,12, the magnitude of the polarization of the adsorbed molecules by the electronic field is not seriously influenced by so-called “active spots” or by surface heterogeneity. The contribution by the nonpolar van der Waals’ forces, however, is more influenced by a heterogeneous character of the surface of the adsorbent. As those forces cooperate and as the surface of a metallic

99

ADS0 RPTION PHENOMENA

adsorbent generally shows heterogeneity, we may expect the heat of adsorption t o decrease with inmcreasing amount of adsorption. This is true for the physical adsorption on normal, polycrystalline metals and metal powders. Usually the heat of adsorption starts a t low coverage with a figure appreciably higher than the heat of liquefaction of the adsorbed gas. With increasing degree of coverage, 0, the heat of adsorption decreases skadily and gradually approaches the heat of liquefaction. Rhodin (921) has succeeded in preparing three different well-defined crystallographic faces of copper and in measuring the adsorption of nitrogen on these faces a t various (low) temperatures. As may be seen from

I

Mo'

1

FIG.24. Heats of adsorption of nitrogen on copper as a function of t h e degree of coverage, 8. The solid lines are for the crystallographic faces, which:are indicated; the dashed line is for polycrystalline copper (221).

Fig. 24, i t is remarkable th at the heat of adsorption on all three crystallographic faces is practically constant a t low degrees of coverage and th a t a t 6 values higher than 0.5 the heats of adsorption increase considerably. At low 0 values we may coiiclude th at the nitrogen molecules, which repel each other by their dipoles induced by the electronic cloud of the surface and attract each other by their mutual van der Waals' forces, do not influence each other; this may be largely due to the counteraction between their forces. At higher 0 values the attraction by their mutual van der Waals' forces, leading to a two-dimensional condensation (Sec. V111,4), raises the heats of adsorption considerably. It may be noted that the three faces give rise to different heats of adsorption. The {IlO] face, which is the least densely packed with copper atoms, gives the smallest value; the most densely packed { 1 1 1 ) face gives the highe6t heat of adsorption. Polycrystalline copper, a t low

100

J. H. DE BOER

8 values, gives heats of adsorption which are higher than those on any of the :three faces. Apparently, active places, e.g. crystal boundaries, are responsible for this behavior (222). Charcoal, activated or graphitized, and graphite itself may be expected t o offer mainly their basic planes for adsorption. We may therefore expect a rather homogeneous character of the surface. The heat of adsorption of physically adsorbed molecules reflects this character. The heat of adsorption of many gases, including argon, nitrogen, oxygen, and many hydrocarbons, is practically constant (223a-e) . Sometimes heats of adsorption are reported which decrease slightly with increasing degree of occupation; the heats of adsorption for ethylchloride on charcoal, reported by Goldmann and Polanyi (ZdSe,%4), decrease from 12.5 kca1.l mole for 0 = 0.09 t o 9.5 kcal./mole for 0 = 0.60. The heats of desorption of n-pentane (bdSc), carbon disulfide (22%) , and diethylether (223e) on the same charcoal show exactly the same dependence on the degree of occupation. We may, therefore, conclude that in such a case the surface heterogeneity causes the decrease. Argon on graphite, a t higher e values, shows the same behavior as nitrogen on sinyle-crystal copper, as has been described above (223). The mutual van der Waals’ attraction forces may, in this case, 1ea.d t o two-dimensional condensation, as is also shown by the entropy data. We shall discuss two-dimensional condensation in Sec. VIII,4. I n all other cases the mutual repulsion of the induced dipoles and the mutual attraction by van der Waals’ forces, both rather weak in themselves, balance each other in the adsorption of gases on charcoal and graphite, causing the heat of adsorption t o be practically independent of the degree of occupation. It is only a t very low e values that the influence of active spots is noticeable; a t low degrees of coverage higher heats of adsorption are reported, which fall rapidly with increasing 0 values to practically constant values (625). 3. The Heat of Adsorption as a Function of the Degree of Occupation in Physical Adsorption Phenomena o n Ionic Surfaces

As we saw in Sec. VI,2 physical adsorption of normal gases on ionic surfaces results from a combined action of van der Waals’ forces and polarization of the molecules by the electric fields of the surface. Active spots (Sec. V, 12) influence both effects. Actual heterogeneous surfaces of ionic adsorbents, therefore, showing various Crystallographic faces, crystal boundaries, edges, vacant ionic sites and many other types of active places, will in all practical cases adsorb the first molecules with a relatively high heat of adsorption. The heat of adsorption will decrease appreciably with increasing degree of coverage (626). Crawford and Tom-

ADSORPTION PHENOMENA

101

kina (227), investigating the adsorption of SO2, COZ, and other gases on CaFz and BaFz, noticed a decrease in the heats of adsorption with increasing amounts of adsorbed molecules. They ascribe the effect t o the nonuniformity of the surfaces, in addition t o the presence of different crystallographic planes. Both SOz and C 0 2 possess quadrupole moments and it might be asked whether the main contiribution toward the heat of adsorption comes from the nonpolar van der Waals' forces or from the electrostatic forces.

FIQ.25. Isosteric heats of adsorption a t 0°K. for nitrogen on rutile (curve A ) and for argon on rutile (curve B ) ;contribution due to the quadrupole attraction of nitrogen (curve C); contributions due to the electrostatic polarization of argon (curve 0 ) and to the dispersion forces (curve E ) [data from Morrison et al. (23O)J.

Active spots for these effects do not coincide (Sec. V,l2). Active spots for van der WaaIs' forces are not active for electrostatic effects and vice versa. A distinction between the two effects has recently been made by Drain and Morrison (628). The heats of adsorption of argon, oxygen, and nitrogen on rutile decrease markedly with an increasing amount of adsorbed molecules (229). In Fig. 25 the isosteric heats of adsorption for nitrogen (curve A ) and argon. (curve B ) , reduced t o O'K., are given according t o the data of Morrison and collaborators (230). The curves are similar in character, and Drain and Morrison (268) succeeded in reducing them by linear transformations t o one single curve. The curve for oxygen could also be reduced t o the same curve. A single form of distri-

102

J. H. D E BOER

bution curve may be used for the three gases. The excess of heat of adsorption of nitrogen over that of argon varies greatly with the degree of occupation, 8, and Drain and Morrison assume that this excess of heat is caused by the attraction of the quadrupole of nitrogen by the electrostatic fields (see also Sec. V1,2) of the various parts of the surface. The difference between curves A and 3 (curve C ) may then be taken as an indication of the heterogeneity of the distribution of these electrostatic active spots. The known quadrupole moment of nitrogen enabled them t o calculate the electric field, acting a t the various surface coverages, whereupon, assuming argon to be adsorbed on the same electrostatic active spots, they calculated the contribution of the electrostatic polarization of argon by means of the known polarizability of this gas. Thus curve D results; curve El being the difference between B and D, gives the contribution of the nonpolar van der Waals’ forces. We see t ha t a t low # values the electrostatic polarization is more important than the attraction by van der Waals’ forces. This would, according t o our interpretation, mean that the electrostatic fields force the molecules t o be adsorbed on their active spots and not on the active spots of the van der Waals’ forces. The antagonism of the activities of both kinds of active spots is in this picture reflected in the increase of the van der Waals’ contributions (curve E ) with increasing degree of occupation. Relations of this kind can, of course, be found only when the molecules are not freely mobile over the surface but are actually localized on their active spots. Morrison and collaborators (228,230) found from their entropy figures that the adsorptions of argon, nitrogen, and oxygen on rutile are of the localized character; this does not mean (see Sec. VII,3) tha t the molecules cannot move over the surface. During the calorimetric experiments they reach their equilibrium positions by hopping movements. These experimeiits were done between 75” and 180°K. At still lower temperatures even the hopping movement may not be sufficiently strong for the molecules to reach their right positions; in such cases too low values are found for the heats of adsorption a t low 0 values. Pace, Dennis, Greene, and Heric ( M I ) found that the temperature should be higher than 73°K. in order t o obtain reversible adsorption, When we calculate the st.rengths of the electrostatic fields corresponding to the points a and b indicated in Fig. 25, we find the values 4.1 X lo6 e.s.u. and 0.9 X 106 e.s.u. respectively. These figures compare very well with the average value of the field over rutile, viz.,

F

=

2.72 X lose.s.u.

derived by Chessick, Zettlemoyer, Healey, and Young (232) from their

ADSORPTION PHENOMENA

103

completely different experiments, which were discussed in Sec. VI,2. As already stated in See. VI,2, this average field is of the same order of magnitude as the field caused by the electronic cloud distribution over charcoal (Secs. V,7 and V1,l). The field over charcoal-and metals-however, is far more homogeneously distributed over the surface than the electrostatic field emanating from the surface of rutile. Moreover, as already stated in Sec. VI,2, the two fields have opposite directions. Here we differ with Drain and Morrison (228), who tentatively ascribe the field t o the positive titanium ions. In th at case the peripheric dipoles of alcohols would not be so strongly adsorbed as they apparently are.

f

- e

FIG.26. Heats of adsorption of argon on cesium iodide (curve A ) and on potassium chloride (curve B ) as a function of 0 (233).

It is remarkable th at the heats of adsorption of argon, oxygen, and nitrogen on rutile do not increase when the e value approaches unity, in other words, when the mutual distances of the molecules approach those of a completely filled unimolecular layer. Apparently the heterogeneity of the surface prevents this effect. I n Fig. 24 it is shown th a t such a n increase, caused by the mutual attraction of the adsorbed molecules, is found with the adsorption of nitrogen on copper. Such a n effect is also found by Orr ( W S ) , who estimated the differential heats of adsorption from adsorption isotherms of argon, oxygen, and nitrogen on potassium chloride and on cesium iodide a t about 80°K. Two of his curves for argon are given in Fig. 26. We see, again, th a t at low coverages the heats of adsorption fall with increasing e values, but th a t if @ approaches unity the heats of adsorption rise t o a maximum. This increase is caused by the mutual van der Waals’ attraction of the adsorbed molecules. Kemball (284) deduced from entropy data th a t the adsorbed molecules are freely mobile in this case, up to a 0 value of about 0.8. Iodine molecules, adsorbed on barium chloride layers obtained by sublimation, also show an increased heat of adsorption with increas-

104

J. H. DE BOER

ing degree of coverage (636),the heat of adsorption being 11.8 kcal./mole a t 6 = 0.57 and 13.7 kcal./mole at 0 = 0.74. The heat of sublimation of solid iodine is 15.5 kcal./mole; the adsorbed molecules moving freely over the surface have a far higher entropy than those of the solid. Again the mutual van der Waals' attraction of the adsorbed molecules causes the increase of heat of adsorption. The heterogeneous character of the surface of CaF2, SrF2, and BaFz layers, obtained by sublimation, is shown by the absorption spectra of iodine molecules which are adsorbed on these surfaces. The very first molecules show a very high absorption coefficient (236); molecules adsorbed at 0 values higher than e = 0.005 show far lower values for the absorption coefficient. The absorption maxima shift continuously t o longer wave lengths when the molecules are adsorbed at higher 0 values (23237)'

4. Two-Dimensional Condensation and Multirnolecular Adsorption The mutual van der Waals' attraction of adsorbed molecules may lead t o two-dimensional condensation phenomena. We shall not discuss these phenomena in detail but shall refer t o some recent reviews and general treatments (238,239). A few remarks will, therefore, be sufficient here, Two-dimensional condensation of a nonideal two-dimensional gas (mobile adsorption) requires a temperature lower than the two-dimensional critical temperature, which, depending slightly on the theory used, may be calculated to be either exactly half the normal three-dimensional temperature (two-dimensional van der Waals' equation) or a figure somewhere in this neighborhood. Experimental evidence corroborates this. Mutual orientation may play a n important role in two-dimensional condensation phenomena. When anisotropic molecules erect each other during their condensation, the two-dimensional critical temperature will be higher. When they are condensed while lying flat on the surface, the temperature will be lower than half the value of the three-dimensional critical temperature. The normal critical temperature of nitrogen being T, = 12G°K., the normal (isotropic) two-dimensional critical temperature would be T,, = 63°K. ; if, however, the nitrogen molecules are condensed while lying flat on the surface, T,,would be 5G.8"K. If the condensation led to an arrangement with all the long axes of the molecules parallel to each other and perpendicular to the surface, T,, would be 92°K. There are indications that the two-dimensional critical temperatures of N z and CO on steel are above 78°K. and a t about 93°K. respec-

ADSORPTION PHENOMENA

105

tively; this may mean that those molecules erect each other when condensing on the surface of steel. Two-dimensional condensation may even be an endothermic process when the difference between the heats of adsorption of the molecules in their flat positions and in their positions perpendicular t o the surface is larger than the contributicln of the mutual attraction of the van der Waals’ forces during two-di rnensional condensation (240). Parallel oriented dipoles lower the value of T,,, as does a decrease in the heat of adsorption caused by a heterogeneous surface. Hill (241) has discussed the two-dimensional condensation phenomena in localized adsorbed layers. Two-dimensional condensation from a dilute localized adsorbed layer to a relatively condensed localized layer may also occur in this case at temperatures lower than a two-dimensional critical temperature. Two-dimensional condensation-on homogeneous surfaces-leads to sudden jumps in the adsorption isotherm. These jumps may already be found at very low pressures of the gas which is in equilibrium with the adsorbed layer (242). Heterogeneous surfaces do not give rise t o sudden jumps but to gradual slopes (Sec. V,12). There is sometimes a tendency to consider such jumps as indications of multimolecular adsorption; this is not correct. It is of course true that stepwise adsorption can also occur together with multimoleculair adsorption. (See also Sec. V,12.) A condensed unimolecular layer may form a suitable new surface for another adsorbed layer t o be formed (243). A second layer, and more, may also be adsorbed on top of a supercritical densely packed unimolecular layer. I n many cases of di- or multimolecular adsorption, the isotherm shows a region where the slope changes from a rather steep course to a more gradual or sometimes nearly horizontal one. The adsorbed amount in this region often coincides approximately with the amount which fills a unimolecular layer. Any method, therefore, which analyzes mathematically or graphically the isotherm to produce a point of that region can be used for the estimation of surface areas. This explains the success of the well-known B.E.T. method for this analysis. After the excellent Idiscussion by Hill (244) of the B.E.T. and the Huttig theories, in which he points out the weaknesses of the first and the fallacy of the latter, and after the analysis by Halsey (.245),who indicates when a B.E.T. isotherm of “satisfactory ” character is obtained on a heterogeneous surface, little need he said here. It may, however, be of solme importance t o indicate once more (246) that all kinds of isotherms, urjually taken as proof of multimolecuIar adsorption, may be obtained with unirnolecular layers and gradual varia-

106

J. H. DE BOER

tion of adsorption forces on heterogeneous surfaces. The B.E.T. method, nevertheless, is very useful as a practical tool. Molecules can adsorb in a second layer only when their heat of adsorption in that layer is higher than the heat of liquefaction (or solidification) or when in the second layer the entropy is higher than in the liquid (or solid). This entropy effect can be of assistance for only the top layer, for if a third layer has to be adsorbed on top of the second one, the entropy in the second one may not be very high. Consequently multimolecular adsorption requires a heat of adsorption in the second layer and higher (up till the last but one) layers which is higher than the heat of liquefaction. Hill (247) and also Halsey (248) assume the van der Waals’ field of the surface t o transmit energy to the second and higher layers. In addition to this possibility it may be pointed out th a t in physical adsorption on charcoal and on metals as well as on ionic surfaces, the adsorbed molecules are polarized, as we have seen in the preceding sections. The field of their dipoles may also influence the molecules in a second layer, etc. This is essentially the basic idea of the oldest conception of multimolecular adsorption (2.49). Unfortunately, many authors in more recent literature have erroneously stated that, according t o our old conception, polarization had to do the entire job. It has, in fact, only to provide the small amount of extra energy above the heat of liquefaction t o enable a next layer to form. The forces responsible for a second or third layer to be adsorbed cannot be very strong. We may, therefore, not expect multimolecular adsorption t o occur at moderately low relative pressures on flat and homogeneous surfaces. By means of absorption spectra a bimolecular adsorption of p-nitrophenol could be established on CaFz or BaFz layers (250). The absorption spectrum of the first layer was quite different from th a t of p-nitrophenol itself, the absorption spectrum of the second layer being practically the same as th at of pure p-nitrophenol. The first layer is apparently sufficiently polarized by the salt surface, to enable a second layer t o adsorb; the second layer, however, cannot accumulate a third layer. It may not be entirely excluded, however, that there is no more place in the capillary space between the salt layers t o accommodate a third layer of p-nitrophenol. Recently Deryagin and Zorin (251), by optical means, investigated the adsorption of alcohols, water, benzene, nitrobenzene, and other gases on a n optically polished glass surface. Multimolecular adsorption did not occur until the relative vapor pressure was higher than 0.95. I n capillary systems of microporous substances more layers can easily be formed by the combined action of two or more walls, giving a gradual transition to capillary condensation (262).

107

ADSORPTION PHENOMENA

IX. CHEMISORPTION PHENOMENA AT HIGHDEGREES OF OCCUPATION 1. The Decrease of the H m t of Chemisorption with Increasing Degrees of Occupation

I n most cases of chemisorption it is experimentally found th a t the differential heat of chemkorption decreases seriously with increasing 0 values. This phenomenon has been the subject of many discussions during the last few years (253-258). In order t o draw the attention to the magnitude of the effect we shall give a few examples in Figs. 27 50 kcal/mole

t

40

30

20

I0

- (12_0.4

0.6

0.8

- 9

FIG.27. Heats of chemisorption of hydrogen on tungsten; A, on tungsten films (253,260), B, on tungsten filament [curve through the points of Roberts (261)],C, on tungsten powder (269).

and 28. Curves A and B of Fig. 27, giving the heat of chemisorption of hydrogen on tungsten films (259,260) and on tungsten filaments (261) respectively, show that the initial values of the heat of chemisorption (at e = 0) are practically the same on both forms of this adsorbent and that the decrease as a function of e also follows practically the same line. Curve C of the same figure shows the heats of chemisorption of hydrogen on tungsten powder according to Frankenburg (26.2).The initial heat is practically the same as on the other forins of the tungsten adsorbent, but the decrease, as a function of el is much steeper. As discussed fully by Beeck (254),it is not improbable that the surface of the tungsten powder used by Frankenburg was seriously contaminated with some impurities (BBS), such as, for example, silica (see Sec. IX,2). Although in the case of Fig. 27 the initial heats of chemisorption are

108

J. H. DE BOER

practically the same on the three forms of adsorbent used, this is not so in the examples of Fig. 28. I n this figure curve A gives the heat of chemisorption of hydrogen on a nickel film (264) and curve B the same on nickel powder obtained by reduction of nickel oxide (666).As the reduction of the oxide in the latter case consisted of a short reduction at the relatively low temperature of 280"C., a contamination of Euckens' nickel sample by oxide, or chemisorbed oxygen, is not impossible (see Sec. X,4). Curves C and D were also obtained with reduced nickel powder, at 100' and 300°C. respectively (266);the initial heats of those curves %re comparable with the heat of curve E, which was found for thoroughly reduced nickel supported on silica (267).

I

FIQ.28. Heats of chemisorption of hydrogen on nickel: A, on a nickel film a t room temperature (864);B, on nickel powder at 0°C. (86'6); C, on nickel powder a t 100°C. (266);D, on nickel powder at 300°C. (866);E, on nickel supported on silica a t 0°C. (67).

Whereas the heats of chemisorption of hydrogen, found on various nickel adsorbents, are higher on films than on reduced powder, copper films do not bind hydrogen by chemisorption (f?68,W69) , although copper powder does (270,271) (see Sec. X,4). We shall not mention more examples here. The same picture as shown by hydrogen is given by other gases when chemisorbed on metals. We may refer t o the literature mentioned in the beginning of this section. I n Secs. VIII,2 and 3 we saw that two phenomena may cause a decrease of the heat of physical adsorption of gases on metals or on ionic adsorbents: first the mutual repulsion of parallel oriented dipoles and, second, a heterogeneous character of the surface, Since we also learned that the mutual repulsion of dipoles does not materially contribute t o a decrease of the heat of adsorption and since i t can easily be proved that, also in the case of chemisorption on metals,

ADfIORPTION PHENOMENA

109

the dipoles of the chemisorbed atoms are not large enough by far to explain the effect (272), we shall not further consider the possibility of this cause. The mutual repulsion of the chemisorbed atoms by these dipoles can, in the most favorable case, account for only a minor part of the decrease of the heat of chemisorption with increasing degree of covering. A heterogeneous character of the adsorbing surface could, of course, also lead t o a decrease in the heat of chemisorption with increasing e if the heterogeneity is conceived as the existence of a wide distribution of strengths of chemical bonda between the chemisorbed atoms and the surface. Ever since Taylor (273) suggested the surface to be heterogeneous, many authors have accepted the view th at the surface would show varying properties from place t o place. Other authors, however, are inclined to consider metal surfaces which are carefully (physically as well as chemically) prepared as essentially homogeneous. Well-prepared filaments t ha t can be outgasseld a t high temperatures (such as the filaments of tungsten, used by Roberts (,%‘GI)), and films of metals, obtained by evaporation and condensation of metals in a high vacuum, would presumably show more homogeneous surfaces than powders, obtained by reduction of oxides, would do. Anyhow, one would not expect the same degree of heterogeneity t o result independently of the method of preparation of the metal. If the surface has t o be considered as a homogeneous one, another explanation must be given for the strong decrease of the heat of chemisorption. We shall discuss the influence of the heterogeneous character of the surface in See. IX,2, which will be followed (Sec. IX,3) by a discussion of experimental methods to distinguish between a heterogeneous and a homogeneous surface. I n subsequent sections we shall examine the decrease of the heats of chemisorption with increasing degrees of occupation on a surface of a homogeneous character. 2. Factors that Cause a Surface Heterogeneity for Chemisorption

The undeniable fact that; the surface may show a dominating heterogeneity for physical adsorption, hence for van der Waals’ attraction forces or for electrostatic polarization by local fields of the surface (Secs. VII1,B and 3), does not mean th at they should be heterogeneous for chemisorption as well. As was stated in Sec. V,12 the forces between ions and metal surfaces and the covalent forces between chemisorbed atoms or molecules and metal surfa,ces are far less influenced by “active places” of the surface than are some of the forces leading to physical adsorption. It is especially the cracks and fissures of the surface, which may give it a pronounced heterogeneoua character for physical adsorption, th a t do not influence the chemisorption bonds very much (274).

110

J. H. D E BOER

It is, therefore, quite possible that a surface, heterogeneous in character for physical adsorption shows a homogeneous nature for chemisorption. As discussed in Sec. V, the work function of a metal measuring the energy which has to be provided to extract an electron from it and a t the same time indicating the electron affinity of the metal plays a n import a nt role in many heats of chemisorption. The actual value of the work function is different for different crystallographic faces of the metal. In a qualitative way this is shown spectacularly by the emission patterns obtained with Muller's field-emission microscope. I n 1937 Muller (275), using a tungsten single-crystal point, observed that the { 110) face shows the weakest emission of electrons; the emission from a 1211) face was stronger; a {loo) face emitted more strongly again, and the strongest emission was obtained from a { 111) face. It is undecided yet whether these crystal faces are really present or not (276a,b,277).It may be th a t surface migration takes place and results in the forming of really small crystal faces, the crystal tending to shape itself into a n equilibrium form (278). Such surface migrations of the tungsten atoms, even taking place in preferred directions, have been observed (276b,278) when tungsten atoms are condensed on the emitting point a t 1000 t o 1500°C. Similar phenomena are observed when nickel atoms are condensed on a n emitting nickel point (279), and we may accept, with Gomer, th a t the singlecrystal point has a spherical form with small crystal faces of low indexes, separated from each other b y bond regions without sharp edges. The influence of the natural roughness of the surface on the emission seems t o be rather small (280). All these observations seem t o indicate th a t the variation of work function with crystallographic direction is not necessarily connected with the actual presence of these faces a t the outside of the crystal. One may, perhaps, conclude tha t the work function depends on the direction in which the electron leaves the metal, the direction of leaving perpendicular to a { 1111 plane, hence in a [lll]direction in tungsten, resulting in a lower work function than in a [110] direction. Nichols (281) measured the work function of tungsten for various directions. His figures are reported in Table V, where also for the various directions the numbers are given of the nearest (Nl) neighbors of a tungsten atom of the surface as well as the numbers of the next following atoms (N,) and the numbers of the atoms that follow as third neighbors (Na);these numbers have been taken from Stranski and Suhrmann (278). There are indications that the work function for the [110] direction is far higher than indicated (280). It seems as if the number of neighbor atoms located a t short distances determines the magnitude of the work function; the larger the number of direct neighbors at the shortest and

111

ADSCiRPTION PHENOMENA

next shortest distances, the higher the work function. It must be stated, however, that Benjamin and Jenkins (282) found the emission of { 110) to be greater than that of { 1111. From a comparison between various modifications of metals it was already known that the work function of a metal (average work function over all directions) is higher the denser the metal (283). Recently Sachtler dkcovered a rough parallelism between the work function and the product of the density of the metal and the ionization potential of the individual metal atoms (684). The heat of chemisorption also depends on the orientation of the surface. I n a qualitative way this is clearly demonstrated by electron emission microscope pictures of metals on which various atoms are adsorbed. TABLE V Crystallographic direction [hkll

N1

N2

N3

6 4 5 4 4

4 5 3 3 3

7 8

7 9 7

'p

(e volt)

4.65 4.53 4.66 4.36 4.36

The various orientations of the crystallites of a metal specimen show up clearly, and i t can be shown that the electron emission depends on the degree of covering with adsorbed atoms, which itself depends on orientation and on temperature (2135-287). We may expect the adsorbability to depend on the work function and also on geometrical faciiors, such as for example the fitting in of the dimensions of the adsorbed rztoin or molecule in the two-dimensional pattern of the surface. Geometrical factors have played an important role in the catalytic literature a t various times (288,289). Both factors depend on the orientation of the crystal faces of the surface. They may either work in opposite directions or may collaborate in the same direction. The chemisorption of cetsium or other alkali metal atoms on tungsten, leading t o the formation of ions (Sec. V , l l ) , will be largely governed by the work function, As positive ions are formed, we may expect the heat of chemisorption to be larger, the higher the work function. Johnson and Shockley (g90) found that a tungsten filament, consisting of a single crystal, having a (110) plane perpendicular t o the axis and emitting electrons a t 2,OOO"K. shows 11111 and { 100) faces, which is in accordance with the work functions mentioned in Table V. When the tungsten is

112

J. H. DE BOER

heated a t lower temperatures in cesium vapor, other faces appear on the fluorescent screen, depending on temperature, viz., { 211) a t 900°C. and 1110) a t 850°C. The (211) faces adsorb cesium most strongly, next follow the (110) faces. At 700°K. the amount of cesium adsorbed on these two faces is already too large for maximal emission; the faces around (211) and (110) show the highest emission. { 100) and { 111))however, do not show up yet; apparently they do not adsorb cesium at this Gemperature and cesium pressure, These results are completely in harmony with the work-function figures of Table V. The higher the work function the better cesium is adsorbed. Similar results were obtained with tungsten single-crystal spheres (291). The adsorptions of sodium, barium and thorium on tungsten and molybdenum have been studied along similar lines. Sodium on tungsten and molybdenum and barium on molybdenum are preferably adsorbed on 12111 faces (292); barium on tungsten on 1111) faces; thorium, however, is preferably adsorbed on { 111) faces but not on { 110) faces (29.2). It seems as if other factors are operative besides the work function. Oxygen atonis, chemisorbed on tungsten or molybdenum, seem to prefer (100) faces (293). In the chemisorption of oxygen the metal has t o provide the binding electrons, and it might be expected that a low value of the work function would favor the process. Steric factors, however, may also play their role here. We may conclude that on polycrystalline material, where various crystallographic orientations will be present at the surface, a certain degree of heterogeneity will result from the different heats of adsorption on the different crystal faces. As we see, however, from the difference in the work functions of tungsten, these differences are not extremely large. We may perhaps expect th at a certain part of the observed decrease of heats of chemisorption with increasing 0 values may be ascribed to this heterogeneity, but it seems doubtful th at the whole effect should be caused by it. I n Secs. V,5 and VI1,6 we discussed the influence of irregularities of atomic dimensions in the surface. We also saw how chemisorption may cause new deviations of normal valencies or may create vacancies in the regular pattern of the surface layer or of the layers in the surface region beneath the surface. Vol’kenshtein (294) treated chemisorption phenomena on solids with surfaces that are provided with such sorts of microdefects (295). These microdefects possess a mobility which is characterized by a n energy of activation. They have the ability of interacting with one another and they give rise to new microdefects when chemisorption takes place on them. Instead of the concept of active centra, therefore, which are bound

ADSORPTION PHENOMENA

113

to definite sites, a picture is developed of nonlocalized active spots. Taylor (296) suggests that such active spots may be operative on the surfaces of actual catalysts and on metal surfaces which are contaminated. Vol'kenshtein considers the microdefects to be partly of a biographical origin, partly of a thermal oyigin. The creation of the latter sort may consume a certain amount of energy. Chemisorption means a reaction between the atom to be chemisorbed and a microdefect. The heat of chemisorption results from the cantributions by the heats of these reactions and the opposite contribution by the heats of creation of new (thermal) microdefects that form new adsorption sites. I n this conception the surface is not essentially of a heterogeneous character; the heterogeneities are created during and by the chemisorption process. As this creation consumes energy, the heat of chemisorption decreases. As already stated above, Taylor assumes the aforementioned microdefects t o be present on contaminated surfaces. Contaminated surfaces, however, may also show heterogeneity for chemisorption in a more classical way. As we have seen in Sec. V, oxides of metals may chemisorb gases, such as hydrogen, as well as metals do. Surface contaminations extending over small, but dej'inite, areas may, therefore, offer possibilities for chemisorption th at are different from those on clean or less contaminated parts of the surface. PL heterogeneity in this sense may, therefore, also be expected on contaminated surfaces. Roginskil (297) showed that in many cases catalytic activity may be given t o the catalyst b y very small amounts of gases acting as promotors. Heterogeneity of the kind #just indicated may, therefore, be expected with actual catalysts. 3. Experimental Methods to Study the Heterogeneous or Homogeneous Charucter of a Surface for Chemisorption

The occurrence of different crystal faces a t the surface of polycrystalline adsorbents may cause a cert,sin degree of heterogeneity. The differences, however, in heats of chemisorption on the various crystal faces will probably not be large enough to explain the general decrease of heats of adsorption with increasing 8 values. It is true that orientation may result in an appreciable diff erenee in catalytic activity or that catalytic activity is particularly shown b y selected crystal faces, but this does not necessarily point t o a difference in heat of adsorption. Beeck, Smith, and Wheeler (298) prepared nonoriented and oriented nickel films, the latter showing { 110] planes parallel to the substrate on which they were condensed. They showed that the catalytic hydrogenation of ethylene a t 0°C. proceeds five times more quickly on these oriented films than on randomly oriented films. The heats of chemisorption of

114

J. H. D E BOER

hydrogen on both types of films, however, are the same, which induced Beeck t o conclude that the difference in catalytic activity is caused by a difference in entropy of activation (299). Sachtler, Dorgelo, and van der Knaap (300) showed that oriented nickel layers, made according to the methods of Beeck, Smith, and Wheeler, are, indeed, oriented with a (110) plane parallel to the surface on which they are condensed but that the crystal faces forming the outside of the layers are not (110) planes. This could mean that the border faces of the oriented layers are randomly distributed, forming a surface just as heterogeneous as th at of nonoriented layers. If catalysis, like electron emission but unlike chemisorption, is governed more by orientation than by the actual presence of oriented bordering planes, differences in catalytic activity cannot be used to study the homogeneous or heterogeneous chemisorption character of a surface. Various studies of catalytic activities on single crystal spheres of copper (301) suggest that orientation may indeed be responsible for differences in the rates of the catalytic reactions. The reaction of hydrogen and oxygen shows the highest rate on those places of the surface of the copper sphere that are parallel to ( I l l ) directions. Those parts of the surface that are parallel to { 100) directions are seriously roughened by the reaction, though the rate of the reaction is lower there than on the ( 111) parts, which are not roughened (305). It is a s if on the parts which are parallel t o the { 100) planes both H and 0 atoms penetrate into the copper to some depth and react underneath the actual surface (see Secs. VII,6 and 7), and the quicker rate on the (111) parts (parallel t o the { 111) planes that are not actually present) prevents the atoms from penetrating. There is no direct parallelism between heats of chemisorption and catalytic activities. Another argument, used to prove the existence of an important heterogeneity of the adsorbing surface, is taken from the inhibition that catalytic reactions may suffer from the addition of quite small quantities of strongly adsorbed poisons (303). A chemisorption of carbon monoxide, covering only 1% of the surface of a copper-powder catalyst, may reduce the activity in ethylene hydrogenation at 0°C. by a factor of 9 (304).Since Wheeler (305) has shown that such drastic reductions of the reaction rates may quite reasonably be expected with porous catalysts when the poison is selectively adsorbed on the pore mouths, these arguments are not to be taken as proof of the existence of heterogeneity in chemisorption properties. The occurrence of slow chemisorptions following fast chemisorptions is in many cases explained by a heterogeneous character of the surface, The rate of these slow chemisorptioiw is governed b y a n activation en-

ADSORPTION PHENOMENA

115

ergy, and it has been observled th at the rate decreases exponentially with the amount that is adsorbed. The activation energy increases linearly with the degree of occupation 8. We shall return t o this relationship in Secs. IX,9 and 11, where we shall see that this phenomenon cannot necessarily be taken as proof of the existence of heterogeneity. Taylor and collaborators (306) have found that in many cases a n increase in temperature during such a slow adsorption causes a rapid desorption, followed by a slow readsorption. This effect has been observed for a number of adsorption3 on oxide and metal-powder surfaces, and also in chemisorption of hydrogen on tungsten films (307).This phenomenon, t o be sure, can be explained from a heterogeneous character of the surface if it is assumed that there are areas with relatively low heats of adsorption and activation energies as well as areas where higher heats of adsorption and activation energies are found. There are, however, also other possibilities to explain such phenomena, and the occurrence of this effect should, therefore, not be taken as an absolute proof of the existence of heterogeneity (Sec. IX,ll). An experimental method of investigation which may indicate in a direct way the existence or nonexistence of heterogeneity for chemisorption has been devised by ltoginskii and collaborators (308). By this method the adsorption is performed in two steps b y use of different isotopic forms. Should the surface be of a homogeneous nature, the adsorbed atoms would all be in the same condition. When, therefore, the gas is subsequently desorbed in two steps, the isotopic composition of the two portions should be the same. If, however, the surface has a heterogeneous character, the fra'ction of the gas added first during the adsorption process should be removed last from the surface during the desorption process and the desorbed fractions should have the same isotopic differences as the adsorbed portions. The method was used for the study of the chemisorption of hydrogen on nickel and on zinc oxide, and Keier and Roginskii could demonstrate the heterogeneous character of the surfaces of these adsorbents ($09). They could also demonstrate the heterogeneous character of the surface of charcoal for the chemisorption of hydrogen (310). Kummer and Emmett (31 1 ) studied the chemisorption of carbon monoxide on a n iron ammonia catalyst by the same method and found a partial mixing of the desorbed gas, as if the surface consisted of a heterogeneous complex of homogeneous parts. Similar results were obtained by Eischens (312). The chemisorption of nitrogen on an iron ammonia catalyst was also studied by Emmett and Kurrtmer (313),who found that the surface behaved as if it were of a homogeneous character.

116

J. H. D E BOER

Recently Schuit used the method for studying the chemisorption of hydrogen deuterium mixtures on thoroughly reduced nickel, supported by silica (314). He concluded that such a nickel surface behaves as a completely homogeneous surface. The various investigations do not, therefore, lead t o a definite conclusion. It is quite possible that some of the surfaces which were investigated were homogeneous for the type of chemisorption that was studied. A slight degree of heterogeneity, as may be expected for a polycrystalline material, may still lead to a homogenization of the isotopic mixture which is adsorbed, provided the activation energy for surface migration is not too high. It is, therefore, also quite understandable that experiments with one sorbate (nitrogen) may lead to the conclusion that the surface shows a homogeneous character, while the use of another sorbate (carbon monoxide) seems to prove the existence of a certain degree of heterogeneity, as we have seen from the results of Emmett and Kummer. The nature of the surface depends, undoubtedly, on the way the adsorbent has been prepared. There are various reasons to agree with the conclusion drawn by Kwan (315) that the surface of a number of metallic catalysts is of a homogeneous nature for chemisorption as long as these catalysts are prepared by a very careful reduction and are kept free from any poisoning materials. We shall see that it is very difficult to fulfill these conditions.

4. The Heats of Adsorption of Cesium Atoms on Tungsten at High Degrees of Occupation I n Sec. V,9 we spoke about the heat of adsorption of a cesium atom chemisorbed in the form of an ion on a tungsten surface. The value (Q& = 68.8 kcal./mole holds for the adsorption on a bare tungsten surface. It has been known for many years that the heat of adsorption decreases with an increase in the amount of cesium atoms adsorbed (316). I n Fig. 29 the full-line curve shows the decrease according to the figures of Taylor and Langmuir (317). These authors could represent Q O (the heat of adsorption as a function of e) by an empirical equation: Qs =

1

64 kcnl./mole + 0.7148

This equation holds between e = 0.06 and e = 0.60; a t 6 values lower than 0.06 the heat of chemisorption increases more strongly with decreasing 0 (full-line curve) than was calculated with the equation (dotted curve). This deviation from the empirical equation was explained by Taylor and Langmuir by the presence of active areas on the tungsten

117

ADSORPTION PHENOMENA4

surface. If 0.5% of the surface consists of such active areas, the adsorption by the rest of the surface can be represented by the equation. Apart from this heterogeneity of 0.5%, they consider the surface to be homogeneous and Langmuir (328) explains the “normal” fall of the heat of adsorption from the mutual repulsion of the dipoles which are formed by the cesium ions and the negative charges in the metal opposite these ions. Applying the two-dimensional van der Waals’ equation, one can express the spreading force of the adsorbed film in terms of the dipole moment p and the degree of occupation 8.

40. 6

FIG.29. Heat of adsorption of cesium on tungsten (317).

By use of Gibbs’ adsorption equation, the spreading force can also he expressed as a function of {;herate of evaporation of the atoms. From the figures for the spreading force, the dipole moments pi can be calculated. In this way the following values were found by Langmuir: pi a t 0 = 0, at 8 = 0.50, pi a t 8 = 0.90, pi

= = =

16.16 X lo-‘* e.s.u. 8.28 X 10-l8 e.s.u. 6.06 X 10-l8 e.s.u.

The marked fall in dipole moment with increasing 8 values indicates a strong mutual depolarization of the dipoles. As the figures for the dipole moments, indicated above, are calculated on the assumption that the decrease of the heat of adsorption is caused by the mutual repulsion of the dipoles, they may not be used,

118

J. H. DE BOER

reversely, for calculating the magnitude of this decrease, as is sometimes done in literature. The presence of the dipoles pointing with their positive poles away from the metal facilitates the emission of electrons from the metal. The work function of the metal is lowered by an amount AV

=

4aeoopi

where uo is the maximum number of cesium atoms that can be adsorbed per square centimeter, 8 is the degree of occupation, and pi the dipole moment. The electrons are influenced by the forces of the dipole layer only after they have left the surface; hence Langmuir suggested that when they have covered half the distance between the positive and negative poles of the dipoles, only half the foregoing amount is responsible for the actual lowering of the work function. However, if we denote the effective part of the dipole protruding from the metal by p instead of p i , we can use A p = 4Tfbop (54)

As was pointed out by the author ( S 1 9 ) ,the lowering of the work function has a pronounced effect on the heat of adsorption. I n See. V,9 we learned that the heat of adsorption of a cesium atom, transformed into a cesium ion by the chemisorption process, is given by (&a);

= Qi

- (evi - W )

(55)

where ( Q a ) i symbolizes the heat of adsorption of the atom in the form of an ion, Qi is the heat of adsorption of a cesium ion, eVi is the ionization energy of the cesium atom, and p is the work function of the metal (see also Sec. V,8a). As soon as some cesium is adsorbed on the surface, the work function of the metal is lowered, which means that less energy is gained during the act of adsorption of the next atoms. Introducing Eq. (54) into (55) and denoting the heat of adsorption by the symbol Qo, t o indicate that it depends on 8, we obtain Qe = Qi - (evi -

or, introducing (54), =

Q; -

-

Q(O

+ CAP)

+

47re~,p)

(56)

(57)

If Qi remained constant and if p did not change during the progress of further adsorption, there would be a linear decrease of the heat of adsorption with 8. The heat of adsorption of the cesium ion, Q;, however, does not remain constant, but increases during further adsorption. This effect is

119

ADSORPTION P H E N O M E N A

caused by the fact th at the dipole layer does not give a homogeneously distributed electric double layer, but an electric double layer where the dipoles are situated a t discrete spots (3dO19d1).An electric double layer with homogeneously distributed charges does not exert any electric force

+

-

2"

FIG.30. Change of potential when an electron passes through a double layer with homogeneously distributed electric charge.

outside the planes of its charges; there is only a potential gradient between those planes (Fig. 30). Jf, however, the dipoles are situated a t discrete spots, such as indicated in Fig. 31, the double layer also exerts forces outside the planes of the double layer. Positive charges are attracted from the side of the positive plane, negative E charges from the side of the negative plane. I n Fig. 32 t he change of potential is given for a n electron passing through such B discrete dipole layer; the distance between the dipoles is 26 in this case, if 6 is the length of the dipole. The solid line in the figure holds for the case when the electron passes through the planes between the charges, the charges being arranged in a regular square pattern. If the electron passes FIG.31. Two planes of along the line DE in Fig. 31, it meets one of the discrete charges; distance charges and is repelled, as is indicated by the 6; distance of the charges dotted lines in Fig. 32. in the planes d = 26. Line ABC passes through The positive charge of a cesium ion will the planes between the therefore be attracted owing to the discrete charges. Line DE passes distribution of dipoles brought about by the through the planes a t cesium ions already present. Consequently the points occupied by heat of adsorption of the cesium ion is increwsed. charges. This was already knownfrom the work of Becker (322) and could be explained by the author (323) by means of these considerations. The increase of the heat of adsorption of the cesium ion, as found by Becker, giving direct proof for the discreteness of the dipole distribution, means a less strong decrease of the heat of adsorption of the cesium atom. When we raise level D of Fig. 6 (Sec. V,8,a) by an amount

I;

120

J. H. DE BOER

4 . r r & ~ ~ level p, E should be consequently raised by a smaller amount. The difference between levels E and A gives the heat of adsorption of the atom. The theory developed in 1935 leads to the view that the effective decrease of the heat of adsorption of the atom depends more or less on the ratio between the length of the dipoles (6) and their distance from

E'\

FIG.32. Charge of potential when an electron passes through a dipole layer along lines ABC or DE of Fig. 31.

each other ( d ) . We may suggest that the effective decrease is approximated by 6 6 AQ,tr = - eAp -- 2e - dUop

d

2d

This leads t o

QB=

Qo

- AQeir

= Qo

-

where Qo is the heat of adsorption a t 0 = 0 [here As we may, again as an approximation, write

(59) see Eq. ( 5 5 ) ] .

$$€6 = p

and

According t o this equation the decrease of the heat of adsorption should be proportional t o 0 3 4 ; hence &o = Qo -

a@."$

(611 The constant a can be calculated by means of the known values for uo = 3.56 X atoms/cm and p = 6.8 X 10-l8 e.s.u. (324) and amounts t o a = 56.4 kcal./mole

ADElORPTION P H E N O M E N A

121

Curve B in Fig. 33 is calculated by Eq. (61), by use of this value for a and Q0 = 68.8 kcal./mole, which is the experimental value for e = 0. Curve B should be compared with curve A , which gives the experimental data (this curve is the same as the solid curve of Fig. 29). We see th a t the two curves have complthely diff eredt characteristics.

FIG.33. Curve A : heat of adhorption of cesium on tungsten, curve B: theoretical curve according to Eq. (60) with H constant dipole moment, curve C: theoretical curve according to Eq. (60) with a variable dipole moment, b u t a constant polarizability, CY = 10 X curve D is curve C fitted on A at 8 = 0.3.

In our calculations the dipole moment p has been considered constant. This is not correct: i t decreases as a result of mutual depolarization of the dipoles and this may seriously affect the curves. The depolarization can be accounted for by an equation given by Roberts (325): p e == p0/[i

+ gLYe(~~i

(62)

where p e is the dipole moment a t the degree of occupation 8, po the dipole moment given by the first cesium ion on a bare tungsten surface, and a the “polarizability.” e and 130 have the same meaning as in the other equations. It is not known which value of LY should be used for the polarizability in a chemisorption case like this, neither whether it would be constant and independent of 8. Using a constant value a = 10 X and using the known value for p = 6.8 X loF2*at e = 0.07 (SW/i), we

122

J. H. DE BOER

have calculated pe as a function of 0. The figures, obtained in this way, have been substituted in Eq. (SO), which is now written in the following form: Qs = Qo - bpifP’ (63) where b =4~(~0)9$ b amounts t o b = 1.22 when p e is given in Debye units. The resulting curve for Qe is given as curve C in Fig. 33. It may be observed that the slope of this curve for higher 0 values is approximately

0.1

0.2

0.3

0.4

0.5

C

- 8

FIG.34. Curve il: heat of adsorption of cesium on tungsten, curve B : theoretical curve according to Eq. (611, curves C, 1),and E : calculated with Eq. (63) and (L = 10 X lopz4,o = 30 X U P 4 , and LY = 40 X 10-24 respectively. All curves are fitted on curve A at 0 = 0.2 and 0.6.

the same as the slope of curve A . If curve C is not started a t Qo = 68.8 kcal./mole but drawn in such a way that it fits on curve A at 0 = 0.3, we obtain curve D. This curve covers curve A well for e values between 0 = 0.2 and e = 0.6. Apart from a slight bend a t lower 6 values, it is practically a straight line. The first part of curve A cannot be represented by these curves. We may try whether Eq. (61) or (63) will give better results when we esti-

ADSORPTION PHENOMENA

123

mate the constants a orb, and also Qo, by fitting the curves at two points of the experimental one. We have chosen the points a t 6 = 0.2 and 0 = 0.6, and some of the results may be seen in Fig. 34, where curve A is the experimental curve again. Curve B is calculated by means of Eq. (61) and curves C, D,and E are calculated by means of Eq. (63). The values have been calculated assuming a = 10 X 30 X and 40 X respectively, by means of a depolarization equation th a t differs slightly from Eq. (62). We have used pe =

+

~ / [ l g ~ ( e d ~ i

(64)

which holds for mobile dipoles; whereas Eq. (62) holds for located dipoles. Curve C approaches curve -4 very well for B values higher than B = 0.2; it turns out t o be practically identical with curve D of Fig. 33. It seems not possible to devise an equation for Q Owhich covers curve A completely, a t least not ;along the lines of thought developed above. We may, probably, conclude th at the tungsten filaments used in these experiments, being of a polycrystalline nature, show different crystallographic faces or directions, which gives rise to slight differences of Q over the surface. The surface seeins to be slightly heterogeneous. 6. A Change in the Nature of the Bonds at High Degrees of Occupation

When, at increasing degree of occupation, level D of Fig. 6 (Sec. V,8,a) is raised by the growing dipole layer, it will, at a certain e value, reach level A . When this happens, condition (30) of Sec. V,8 will not be fulfilled any more. From this degree of occupation onward the adsorbed cesium, though still adsorbed in the ionic form, will not desorb in the ionic, but in the atomic form. It is at 0 = 0.07 th a t this happens (326). At a still higher degree of occupation level D of Fig. 6 (Sec. V,8,a) will be raised t o such an extent that it reaches the same height as the minimum B of the curve giving the adsorption of cesium in atomic form. This happens a t e = 0.134 (327), and a t higher 0 values no ions can be formed on the surface. 131 our earlier considerations (528) we suggested t ha t atoms would be adsorbed next t o the ions and be polarized by them. Though we stated there th at a sharp distinction between ions and atoms could not be made, we had better consider this old picture obsolete and replace it by thte picture of the adsorption of atoms, which are still strongly polarized by the metal surface. Even physically adsorbed atoms are polarized in the same direction when adsorbed on metal surfaces (Sees. V,7 and V1,l). The nature of the adsorption, therefore, changes at higher degrees of occupation. There is experimental evidence of this change. Mayer (329) bombarded with potassium ions filaments of platinum, copper, and aluminum,

124

J. H. DE BOER

on which sodium was adsorbed. He found a n emission of sodium light (sodium D line) during the induced evaporation of sodium. Apparently the adsorbed sodium ions are desorbed by the bombarding potassium ions and while evaporating they are neutralized to atoms via several excited states t hat cause light emission. Light emission was not observed when smaII amounts of sodium were adsorbed ; it became noticeable with increasing amounts of adsorbed sodium passed through a maximum and disappeared again when a larger amount of sodium was present. At low 8 values the sodium ions are liberated as such, a t increasing 0 values the liberated ions are neutralized, and a t still higher 0 values sodium atoms are adsorbed which do not need to be neutralized when they are liberated. Other experimental evidence may be obtained from photoelectric measurements. The normal (nonselective) photoelectric emission of a tungsten filament on which sodium is adsorbed shows a marked increase with raising temperature at relatively low 0 values. The photo effect, however, decreases with rising temperature when the degree of covering is higher. Both effects are reversible (350). Apparently a t relatively low e values, when the adsorbed sodium is in the ionic form, a n increase in temperature means a slight increase of the mean distance of the ion from the surface, hence a slightly increased dipole moment and a slightly decreased work function. When the adsorbed sodium is in the atomic state the dipoles, which are now formed by polarization of the atoms by the field of the metal, decrease when their mean distance from the metal increases a t higher temperatures. The main differences between our present views and those held in 1934-1937 may be briefly stated. Contrary t o our older views we have t o accept a certain degree of heterogeneity of the tungsten surface. We cannot understand the form of curve A of Fig. 33 without accepting the presence of areas where the adsorbed ions are more strongly bound than on the rest of the surface. This heterogeneity is certainly not caused by impurities or foreign atoms; the presence of different crystallographic faces may be responsible. Second, we do not maintain the picture of atoms adsorbed next t o ions a t higher degrees of coverage. At low e values all atoms are adsorbed in the ionic state; at higher degrees of coverage the type of binding changes and from a certain @ value onward atoms are adsorbed as atoms. They are polarized by the field of the metal. At increasing coverage their dipole moments decrease by mutual polarization and a minimum is found for the work function when the decreasing dipole moment per atom is not compensated any more by the increasing amount of dipoles per square centimeter. It is likely th a t the polarizability which is effective in this depolarization is different for the various states of binding; hence it will not be a constant, independent of

ADSNORPTION PHENOMENA

125

coverage. This means that Eqs. (60) and consequently (61) cannot be used in a more quantitative way. 6. T h e Decrease of the Haat of Chemisorption with Increasing Degree of Coverage for Other Adsorptives

Most chemisorbed atoms form dipoles on the surface of their adsorbents. Either the positive or the negative poles of these dipoles may point away from the metal (Sec. V,8b). I n both cases the dipoles influence the work function of the metal, increasing it when the dipoles point with their negative poles away from the metal and decreasing it when the dipoles point in the other direction. As a negative dipole (negative pole pointing away from the surface) is formed by shifting a n electron from the metal t o the adsorbed atom, work is done against the work function. By the increase of the work: function more work will be required for the formations of new dipoles when the degree of coverage increases. T h e heat of chemisorption will therefore decrease. I n case positive dipoles are formed, the electron affinity of the metal facilitates the effect. As the electron affinity decreases with an increase in the amount of adsorbed atoms, the result is again t,hat the heat of chemisorption will decrease with increasing degree of coverage (331). The cause of the decrease of the heat of chemisorption is, therefore, the same a s we discussed in the preceding sections. Boudart (332) elaborated the idea recently, and he proved that it can explain the order of magnitude of the experimental figures for the decrease. I n his equations Boudart uses the conception of an electric double layer with homogeneously distributed charges. As, however, the decrease of the heat of chemisorption AQ is numerically smaller than the change in work function Acp, Boudart introduces the relation AQ =

(65)

and tries t o justify this by assuming that the electron is bound halfway between the planes of the dipole layer. This assumption is a strange one, because it is just these electrons th at have to contribute to the dipole layer. Mignolet (333), therefore, criticizes this assumption and shows that i t is superfluous if more correct data than Boudart's are used. I n his reply t o this criticism EIoudart (334) once more points out that his conception may be taken a3 only an approximate one, just t o give the order of magnitude. As we have seen in Sec. lX,4 the dipole layer may not be treated as a double layer with a continuous distribution of charge. Discrete charge distributions have t o be assumed. Gomer (335)criticizes Boudart on this point. H e evaluates potential curves for a discrete distribution of dipole

126

J. H. DE BOER

charges and, rightly, remarks that a t low 6 values the effect will be smaller than that calculated by Boudart’s method. It is, unfortunately, very difficult to compare the experimental data on the decrease of the heats of adsorption with the observed values for the changes in q o o . It is possible only in the case of the adsorption of alkali metals on tungsten filaments, where sufficient reliable data are available. Contact potentials, measured when the gases are adsorbed on filaments, are less reliable. Most investigations concerning contact potentials on films have yielded values of surface potentials that are known for nearly complete films, but not for adsorbed layers with low e values. There are, moreover, rather serious deviations among the experimental values published by different authors (336).

7. Advantages and Disadvantages of Metal Surfaces Prepared by Diflerent Methods Comparing the various surfaces that have been used for measurements of heats of adsorption and contact potentials, one may say t h a t none of them seems to be ideal for the purpose. It is true that filaments, especially those of tungsten, may be obtained in a pure state, free from contaminations. However, when no special precautions are taken, they do not stay in this condition. It is remarkable how quickly impurities are taken up from the residual gas in the highest vacuum that can be applied. Unless chemical binding agents are used (“getters”), the filaments will be contaminated during the cooling down from the flashing temperature which is used to clean them. A recent publication of Thomas and Schofield (337) on the accommodation coefficients of helium has revealed that even Roberts’s filaments were not clean, though i t was thought that they had unquestionably clean metal surfaces (SSS). It is very fortunate that in the experiments on the adsorption of atoms of alkali and alkaline earth metals on filaments these metals themselves produce and maintain a high vacuum. The only drawback of these filaments is that they are polycrystalline and, therefore, show a slight heterogeneity. Films of metals, produced by evaporation and condensation, will also take up impurities from the “vacuum.” During the preparation of the films the vacuum produced by the evaporation is very high. Afterward traces of gas, liberated from various parts of the apparatus, form a source of contamination. Owing to their very large surface areas, however, the films can be maintained in a clean state for a far longer time than filaments. Many films seem to have the additional advantage that they consist mainly of one crystallographic plane. The method of preparation provides a unique process of crystallization (3.39); it may be that the

ADSORPTION PHENOMENA

127

crystallographic face that develops is not the same as the faces found when the metal is prepared in other ways. There is, however, a disadvantage in the use of films. Because of their extremely large surface area per unit of weight of the metal (340),their surface energy is high. The average thickness of the primary laminae t h a t build up the total film is very small, and it is known th at the films do not behave electrically as normal metals (341). Many of these films show a somewhat expanded lattice of 1-2% (342). It ii3 only after thorough sintering th a t they approach a more normal metallic state. Mignolet (343) observed th a t the work function of films increases during sintering and approaches the value of the normal metal. During such a sintering process impurities may well be taken up. The development of a very large surface area, hence a n extremely open structure, deplends also on the rate of evaporation and condensation, Tungsten films will show a slightly more normal behavior than nonsintered nickel films. Metals produced by a thorough reduction of oxides may be obtained in such a state that their surface is clean. The possibility is not excluded, however, that these metals contain an appreciable amount of dissolved impurities (e.g., atoms of tlhe reducing gas, hydrogen) which may interfere with the chemisorption processes carried out afterward. Boudart (344) explains the peculiar behavior of the chemisorption of hydrogen on tungsten powder, as investigated by Frankenburg, from such an interference by dissolved atoms. I n Sec. IX,l we saw that the initial heat of chemisorption of hydrogen on this powder has the same value as obtained by the chemisorpticm on tungsten films and tungsten filaments (Fig. 27), but th at the decrease of the heat of chemisorption with increasing 8 values proceeds very quickly. It does not seem impro'bable th at thoroughly reduced metals, supported on a carrjer, give the best approach to pure metallic surfaces. The small, crystalline, metal particles cannot sinter together; they therefore cannot include impurities, and, moreover, they are formed under circumstances so nearly ideal that, being so small, they may consist of single crystals, probably showing one main crystallographic face. It is remarkable that with nickel on silica prepared in this way, Schuit and de Boer (345) obtained curve E of our Fig. 28 (Sec. IX,l), which may be compared with curves C or D of Fig. 33. It is not improbable that the decrease of the heat of chemisorption in this case may be fully understood from the decrease in the work function together with the mutual depolarization of the dipoles. Schuit (346) found (Sec. IX,3) that this surface behaved as a homogeneous one. It may be remarked that curve A in Fig. 28, giving the heat of chemisorption of hydrogen on a nickel film, also resembles curves C and D of Fig. 33, but th at the absolute values are higher than

128

J. H. DE BOER

those of curve E . Curve A was obtained by Beeck with a nickel film of large surface area (see above). After sintering of this film the absolute values of the heats of chemisorption were decreased (347) to practically the same level as those of Schuit and de Boer. The higher heats of chemisorption of the negatively charged hydrogen atoms found with the unsintered film are in accordance with the lower work function of such a film. 8. Other Explanations for the Decrease of the Heat of Chemisorption with Increasing Coverage

Suggestions have been made to explain the decrease of the heat of adsorption with increasing amount of chemisorbed atoms in terms of kinetic energy of the electrons used or set free by the bonding of the atoms. Eley (348) remarked that when a chemical bond is formed on the surface, which involves an electron t o enter the metal, this electron will occupg the lowest available energy level. The following electron that enters the metal when the following atom is chemisorbed has to occupy the next higher energy level, thus having a somewhat higher kinetic energy in the electron gas in the metal. Conversely, when on the surface a chemical bond is formed that requires an electron to take part, this electron will come from the highest energy level in the metal; the next>comes from the next lower energy level, etc. In either case the differential heat of chemisorption will decrease with increasing number of atoms adsorbed. Similar remarks were made by Schwab (349) and by Mignolet (560). As the conductibility band belongs t o the collective conduction electrons of the whole metal, it seems rather unlikely that electrons entering its permitted levels or coming from its occupied levels would bring about such drastic effects when they are produced by or used in the production of chemical bonds on the surface of the metal. It is for this reason, th at Temkin (351) introduced the idea of a surface electron gas. He suggested the presence of a two-dimensional electron gas a t the surface of the metal, which apparently behaves in complete independence of the normal three-dimensional electron gas. Accepting the same exclusion principles and statistic distribution for this separate twodimensional electron gas as for the normal three-dimensional one, Temkin derives the following expression for AQ:

where h is Planck's constant and m is the mass of an electron, u is, as usual, the number of adsorbed atoms per square centimeter, ua is, again,

ADSORPTION PHENOMENA

129

the number of chemisorbed atoms in a completely covered unimolecular layer, and e = u/uo. This expression would give a linear decrease with 0. If uo = 3.56 X l O I 4 atoms/cm.2 were used in the case of the adsorption of cesium on tungsten (as was done before; see Sec. IX,4), AQ would amount t o 19.6 kcal./mole between B = 0 and 0 = 1. This value, though of the right order of magnitude, is too low; it would predict a decrease of 7.85 kcal./mole between 0 == 0.2 and 0 = 0.6, although the actual decrease (Fig. 29) is 11.7 kcal./mole. When applied to the chemisorption of hydrogen on nickel supported on silica, where (352) (Secs. IX,1 and 7) a t 0°C. AQ = 14.48 Eq. (66) demands th at uo should be 2.6 X l O I 4 atoms/cm.2 t o account for this figure. As 6 0 for hydrogen atoms on nickel will certainly be far higher than this, we reach the conclusion that in this case Eq. (66) gives too high a figure. Both examples show one of the imperfections of Eq. (66). Temkin, realizing that his AQ, generally speaking, would be too large, suggests that the mass of the electron, m, has to be replaced by a n I( effective mass.” Temkin’s theory, on account of its simplifications, does not contain anything of the specific character of the metals but the various values for uo. This is certainly too strong a simplification. A more serious remark, however, was made by Boudart (353), when he discussed the constancy of the heat of solution of gases in metals. From the fact that the heat of solution of hydrogen in the interior of metals does not fall with increasing concentration, Boudart drew the conclusion t ha t such an effect may not be expected at the surface either. The heat of solution of hydrogen in /3 titanium (354) has the constant value of 27.83 kcal./mole for concentrations lower than 10 atomic % and increases slowly t o only 28.3 kcal./mole for 30 atomic %. This increase is caused by the dilatation of the lattice. I n this instance of solution there will also be an exchange of electrons with the collective conduction electrons of the metal. If then, in the case of solution, this does not involve a change in heat, it may also bme expected that the exchange of electrons between chemisorbed atoms and the surface layer of electrons does not involve such a change. Temkin, on the other hand, refers to a study by Federova and ) ~ found that the heat of solution of hydrogen in the Frumkin ( 3 5 1 ~who P-phase of palladium depends largely on the concentration. There are some indications in literature, however, that may endorse the non-dependence. I n Sec. VII,6 we learned that the chemisorption of oxygen on iron was not restricted to the surface only; iron atoms (ions) moving on top of the chemisorbed oxygen atoms (ions) bring about a n

130

3. H. DE BOER

incorporation of the chemisorbed oxygen ions into the surface layers of the metal. I n a n extensive study of heat of chemisorption on iron films, Bagg and Tompkins (355) observed that the heat of sorption of oxygen by iron films a t room temperature was 71 kcal./mole, independent of coverage, They also observed that oxygen penetrated into the lattice during their measurements. A constancy of the heat of adsorption at increasing coverage may, t o be sure, also be found when the molecules or atoms d o not move freely over the surface before entering into chemisorption. When they stay at the first point of the surface that they reach, the adsorbed layer will be built up gradually from those parts of the microporous system (metal films, like films of inorganic salts, are microporous systems (356)) th a t

',I

I

L

0.2

0.4

0.6

0.11

1.0

- 8

FIG. 35. A : heat of chemisorption of hydrogen on an iron film a t 23°C.; B : the same at -180°C. (857).

are in direct contact with the gas and proceed gradually t o the interior. In this case the average value for the heat of chemisorption is found, independent of the amount adsorbed. It is only when the molecules move freely over the whole surface before being chemisorbed th a t the chemisorbed atoms will be distributed a t random or, by their mutual repulsion, take up positions as far apart from each other as possible, and it is only in these cases that the usual curves, showing a high initial heat of adsorption falling off with illcreased 0 values, are found. This effect is clearly demonstrated by the curves found by Beeck (357) for the heats of chemisorption of hydrogen on iron films a t - 180°C., where hydrogen does not move freely prior to chemisorption or after chemisorption, and a t 23"C., where it does move over the surface (Fig. 35). It is, therefore, understandable th at Bagg and Tompkins explain the constancy of'the heat of sorption th at they found for oxygen on iron films a t room temperature in terms of immobility. I n view of what we learned in Sec. VII,G, however, a solution of oxygen in the surface layers is more probable.

ADSORPTION PHENOMENA

131

Maxted and Hassid (35CI), and later Kwan (359),found the heat of sorption of hydrogen on platinum t o be independent of the amount th a t was taken up. It is quite probable that, a t the temperatures they used in their investigations, hydrogen penetrates the surface layers of the metal. As the atoms that penetrate into the surface layers of the metal do not produce a regular dipole layer on the surface, they do not change the work function in a regular way either. The constancy of the heats of solution and of the heats of sorption in those cases where, apparently, the atoms penetrate the surface layers of the metal does not indicate that the kinetic energy of the electrons in the metal plays the role th at T'emkin needs for his theory. We cannot, of course, entirely exclude the possibility that in all these cases, including the solution of hydrogen in 8-titanium, the constancy of the heats of sorption is caused by immobility of the sorbed atoms. We are, meanwhile, inclined t o think th at Ternkin's picture does not give the solution here. If i t did, the obvious conclusion would be that neither the changes in work function nor in contact potential would be caused by alleged dipole layers, but by the changes in the occupation of the energy levels of the surface electron gas. The consequences of the adoption of this idea would be far reaching. Alternatively, we might assume th a t part of the change in work function and in contact potential-say half the effectwould be due t o the change in the kinetic energy of the electrons in the sense described by Temkin, and the other part would be caused by surface dipoles.

9. Changes in Activation Energy with Increasing Degree of Occupation

It is not only the heat of chemisorption th a t changes during progressing adsorption; the attivation energy, if present, also gradually alters its value. I n general, the activation energy increases with increasing degree of occupation. Figure 36 shows the increase in activation energy of the chemisorption of nitrogen on iron in the temperature range between 200" and 250"C.,as measured by Zwietering and Roukens (360). As we saw in Sec. V,9, the chemisorption of nitrogen on iron is one of the rare cases where the chemisorption of a gas on a pure-metal surface involves a n activation energy. I n most other cases of chemisorption on pure metal surfaces the rate of chemisorption, even a t low temperatures, is so high that, apparently, there is no, or only a very small, activation energy. Even if there is no activation energy when the first atoms are being chemisorbed, there may be cine when the degree of occupation becomes higher. A study of potential curves (Fig. 37) reveals th a t the decrease in the heat of adsorptian with mereasing 0 i s likely t o be accompanied by

132

J. H. DE BOER

the deveIopment of an activation energy from a certain 6 value onward. E will then increase with decreasing Q. I n Fig. 37 curve 4 shows the occurrence of a small activation energy, smaller than the heat of chemisorption (Qk), and the activation energy of curve 5 (EL)is larger than the heat of adsorption ( Q s ) .

.-/

k-l/rdole

FIG,36. Activation energy of the chemisorption of nitrogen on a n iron catalyst (360).

It has recently been realized that in many cases a quickly proceeding chemisorption may be followed by a slower uptake of the same gas. When the pressure is raised, more gas is taken up and the additional uptake de-

.o

I

-

2

4

6

e x

Oistance from s u r f a c e

FIG.37. A gradual decrease in the heat of chemisorption Q may give rise t o the development of an activation energy E, the latter increasing with decreasing Q.

pends on the pressure. I n many cases this phenomenon has been studied with the aid of metal films obtained by evaporation and condensation of metal vapors. As these films have a microporous character and as even the rates of physical adsorption of gases on microporous systems, like charcoal, may show an activation energy (361) (caused by the activation

ADSORPTION PHENOMENA

133

energy of surface migration) , the results obtained with films may perhaps not be taken as conclusive evidence of the occurrence of an activation energy in the later stages of' adsorption. We shall return to this question in Sec. IX,11. A careful study of the ad,sorption of various gases on carefully cleaned mercury surfaces has led Kemball (362) to the conclusion that such gases as carbontetrachloride, hexachlorethane, and chloroform are chemisorbed

'O1

t

a,

0.1

0.3

- 8

FIG.38. Heats of chemisorption of nitrogen on a n iron film as a function of 0 (363).

on the surface of mercury. I n the initial stages there is no activation energy, but after roughly half the surface is covered (e slightly more than 0.5) an activation energy comes to govern the rate. This activation energy for CC14 on mercury is practically zero for e values up to 0 = 0.50 and assumes the values of 4.1 kcal./mole for 0 = 0.62, 9.6 kcal./mole for 0 = 0.69, and 19.2 kcal./mole for e = 0.76. I n all cases where an activation energy is already present at 0 = 0 (Nz on iron, H, on contaminated metal surfaces; see Sec. V,9) it increases with increasing 3 values. The increase of the activation energy is slower than the decrease of the heat of chemisorption. An examination of Fig. 37 shows that this should be so. The maxima in the potential curves are shifted to the left and the minima of the curves are either a t the same distance from the eurface-as we have assumed in our figure-

134

J. H . DE BOER

or are shifted t o the right when the bond with the surface becomes less strong. It is, therefore, t o be expected th at

< AQ

AE

(67)

Figure 38 gives the heats of chemisorption of nitrogen on iron films, recently published by Bagg and Tompkins (363),and we may compare the decrease of this curve with the increase shown by the curve of Fig. 36. 10. Equations for Chemisorption Isotherms

The activation energy for the chemisorption of nitrogen on a n iron catalyst, as measured by Zwietering and Roukens (360) (hence the curve of Fig. 36 just mentioned), can be represented by the expression

E

=

9.2

+ 72.78 kcal./mole

The rate of chemisorption a t a pressure of 20 cm. mercury could be represented by

'' -

dt-

446.5 X

e107.28/R

x e-(9200+7Z,700S)/RT

The entropy term partly counteracts the term containing the activation energy. Restricting ourselves to one temperature, we can write the equation as dB _ -- 446.5e at

0200 72,700 1072 -_ RT x e ( - X F + * ) e

= kle-keS

(68)

The constant kz in this equation contains that part of the equation for the activation energy that depends on 8, as well as the &dependent term of the entropy factor. . Equation (68) has the same form as a n empirical expression found by Zeldowitsch (364) in 1934; viz., vads

=

k,,pe-g8

(69)

where Veda is the rate of adsorption (chemisorption accompanied b y a n activation energy), p is the pressure, and k , and g are constants. This equation is, of course, valid only at constant temperature. Prior to the formulation of Eq. (69) it had been found by Langmuir (385) that a n empirical equation vdes =

kdehe

(70)

where k d and h are constants, offers a good expression for the rate of desorption (vdeJ of alkali atoms chemisorbed on metals. Both expressions (69) and (70) give a good representation of the

ADSORPTION PHENOMENA

135

rates of adsorption and desorption in many cases of chemisorption. At equilibrium, Usds

=

vdss

hence kape-g8

(71)

= kdehe

Taking logarithms and rearranging, we obtain S(g

k, + h) = In --p kd

Introducing g

+h =f

k,/kd

and

= Oo

we obtain 1

6 = -In uop

f

This empirical equation of the adsorption isotherm, giving the relationship between e and the premure, excellently represents many characteristics of chemisorption. Equation (72) is introduced by Frumkin and Slygin (366),who derived it from their electrochemical investigations on hydrogen electrodes. The equation has played a n extensive role in the successful theory of ammonia catalysis of Temkin (367) and it has in literature been termed the Temkin equation (368), although Temkin himself and other Russian investigators call it the logarithmic adsorption isotherm. Equation (72) can, of course, represent e values as a function of the pressure p (or the concentration c) provided th a t these values are not too close t o zero or too close to unity. I n the range of medium values of 0, it adequately represents many cases of chemisorption (369). This isotherm equation will always fit the experimental data when the heat of chemisorption shows an approximately linear decrease with increasing 0 values; i t is not necessary that an activation energy be present. This may be seen from Eqs. (71) and (72); even when h = 0, Eq. (72) results. The validity of Eq. (72:) does, of course, reveal nothing about the cause of the decrease of the heat of chemisorption. Any time th a t the heat of chemisorption may be represented by

Q~ =

- ce

(73)

the logarithmic adsorption equation

e

RT

= - In C

(uop)

(74)

136

J. H. DE BOER

which is identical with Eq. (72), can be derived, whether it be with the aid of the conception of a heterogeneous or of a homogeneous surface. Any time that the heat of chemisorption does not decrease linearly with increasing 8 value, but may be represented by

Qe

=

const. - LY log

e

(75)

an isotherm equation of the form 0 = const. pRTIa

(76)

can be derived. This latter form is the well-known empirical exponential equation, generally called the Freundlich isotherm. We shall not here give the various derivations of these equations but refer to the publications of RoginskiI (370), Temkin ( 3 7 l ) , Brunauer, Love, and Keenan (37W), Zeldowitsch (373), Halsey and Taylor (374), Trapnell (376), and Bokhoven et al. (376). Practically all experimentally obtained data on chemisorption isotherms may, in fact, be expressed in either Eq. (74) or Eq. (76). 11. Some E$ects in Chemisorption Phenomena that are Connected

with Activation Energies I n Sec. IX,9 we discussed the phenomenon that the quick chemisorption of many gases on films of metals is often followed by a slow uptake of the same gas. With nickel films Beeck, Ritchie, and Wheeler (377) found that the amount of gas taken up in this slow sorption process is independent of the degree of sintering of the film. The surface areas of the films that are freely available to the gas-also t o a physically adsorbed gas such as krypton-markedly decrease on sintering. The amount of hydrogen taken up in the fast chemisorption process shows the same decrease. The amount of hydrogen taken up in slow chemisorption, however, is not connected with the surface area but with the weight or the volume of the films. Similar results with tungsten films were obtained by Trapnell (378). I n a more recent investigation Porter and Tompkins (379) found that, on sintering, the amount of hydrogen sorbed as a result of the slow uptake by iron films decreases roughly in proportion to the decrease observed on the fast chemisorption process and hence t o the freely available surface area. They could also prove that the slow process follows the fast chemisorption process continuously. It is, therefore, not a solution in the lattice of the metal films, as was originally suggested by Beeck. Such a dissolution process, moreover, is an endothermic process. Porter and Tompkins (379) suggest that the phenomenon is caused by surface heterogeneity and that fast chemisorption takes place, with-

ADSORPTION PHENOMENA

137

out the occurrence of an activation energy, on the most active parts of the surface and the slow uptake, which also is a chemisorption phenomenon, is accompanied by an activation energy. As we have seen in Sec. IX,9, it is not necessary t o introduce active and nonactive parts. I n the case of iron films, as investigated by Porter and Tompkins, both phenoinena apparently occur on the surface which is freely available. Fast chemisorption may quite normally be followed by SIOWchemisorption associated with a n activation energy when th e potential curves follow the picture laid down in Fig. 37. The activation energy is a normal consequence of the decrease in the heat of chemisorption. With the nickel films of Beeck, Ritchie, and Wheeler and the tungsten films of Trapnell, however, the situation is more complicated. Films of metals, like films of inorganic salts, offer surface areas th a t are proportional t o the weight of the films (380). On sintering, the surface areas decrease strongly (%I), but on adsorption of a strongly bound adsorptive the result of the sintering may be annihilated. CaFz films, previously sintered, showing a strong decrease in the amount of iodine th a t could be taken up by adsorption, did not show a decrease in the cesium-absorption capacity. Apparently cesium desintered the films, for, when after the adsorption of cesium i t was desorbed carefully-without raising the temperature-the original surface area was restored, as could be measured by the amount of iodine that could be adsorbed (382). Such a desintering might also be responsible for the slow uptake of hydrogen in the case of nickel and of tungsten films. The activation energy would then be due at least partly t o the work that is required to reopen the capillary space closed by the previous sintering. A desintering effect may also have played a role in the experiments of Taylor and collaborators, already mentioned in Sec. 1X,3, who found that in many cases a n increase in temperature during slow (activation energy) adsorption causes rapid desorption, followed by slow readsorption (383). As remarked in this section, this behavior has often been taken as a proof of heterogeneity for chemisorption. Areas were assumed with a relatively low heat of adsorption and a relatively low activation energy, together with areas where both the heat of chemisorption and the activation energy have higher values. The occurrence of these types of sites next t o each other on the same surface does not seem very probable. If during the experiments quick chemisorption were accompmied by a slow desintering process, the same phenomena would be observed. On the temperature being raised, desorption would occur, resulting in a slightly smaller degree of occupation corresponding with the temperature and pressure employed. This quick desorption would be followed by a slow uptake.of gas molecules th a t go

138

J. H. DE BOER

on reopening capillaries th at had previously been closed in the sintering (“stabilizing”) process. The slow rate of desintering should then be ascribed t o the activation energy of the desintering and not to a n activation energy of the chemisorption proper. The fact th a t the desorption process is quicker than the slow chemisorption means that the heat of chemisorption must be lower than the activation energy acting in the slow process of uptake. However, there is still another possibility t o explain this behavior. We saw in Sec. VI,3 that hydrogen may be chemisorbed in more than one way. If the general scheme of Fig. 13 could be applied to the phenomena that we are discussing here, the fast chemisorption process would be a chemisorption of hydrogen in the minimum E A of curve ABEAFA; i t may be t hat the activation energy ( E J A is practically negligible. If the temperature is not too low, this chemisorption will be accompanied by a slow process along curve ABEBFB with the activation energy (EJB. Raising the temperature would produce a desorption of atoms adsorbed a t EA, but not of those a t E,. This desorption would be followed by the continuation of the slower process, leading to a n adsorption of atoms a t EB. If this picture should be accepted, we do have sites with a high heat of chemisorption together with a high activation energy as well as sites where both quantities are low. In this picture these sites do not exist next t o each other, however; there is only one type of site, on which the atoms may be bound in two ways. There is, of course, still the possibility of heterogeneity. The experiments, however, do not prove a heterogeneous character of the surface. 12. Restricted Chemisorption Caused by the Increase of the Activation Energy

If, with increasing degree of occupation, the activation energy becomes higher and higher (see Fig. 37) it may rise so high that no further chemisorption will take place, a t least not a t the temperature of the experiment. Higher temperature and higher pressure would result in a n increase of the chemisorption. According t o Fig. 36 the activation energy for the chemisorption of nitrogen on iron a t 0 = 0.2 is about 24 kcal./ mole. This means that even at a nitrogen pressure of 1 atm. a n amount of roughly lo6 molecules of nitrogen only would be chemisorbed/sec. and/cm.Z a t room temperature, leading t o a n increase of e with about 10-6/hr. This means that the chemisorption has practically come t o a standstill. Experimentally it was found b y Beeck (384) th a t a t room temperature the chemisorption of nitrogen on iron films does not proceed further than t o about 8 i=0.2. (See also Fig. 38.) At liquid-air tempera-

ADSORPTION PHENOMENA

139

tures quite a different type of chemisorption of nitrogen takes place on iron films. Chemisorption is then fast and there is no activation energy, and the heat of chemisorption is only 10 kcal./mole, decreasing to 5 kcal./mole with increasing: coverage. Theref ore any time an activation energy plays a role, increasing with increasing degree of occupation, we are not sure whether the whole surface will be occupied fully by a unimolecular chemisorbed layer. 13. Some Final Remarks with Respect to the Decrease of the Heat of Chemisorption with Increasing Amount of Adsorbed Material

I n the preceding pages of this section we have repeatedly been obliged t o draw attention to the disagreements of opinion with respect to the question of homogeneity or heterogeneity of the surface. Summarizing and reviewing this section we may draw the following conclusions. 1. The decrease of QS u ith increasing e values on metal surfaces is due mainly t o the change in the work function of the metals b y the discrete dipole layers formed by the chemisorbed atoms. The author wishes t o state here that he does not recommend the phrase induced heterogeneity, introduced b y Boudart (585). He would rather recommend the term work-function effect or surface potential effect. 2. An a priori heterogeneity may be, and is mostly, superimposed on the surface potential effect This heterogeneity may be caused by the presence of different crystallographic faces or of impurities, by contaminations of the surface, or by dissolved impurities in the lattice or between the grains. 3. The mutual depolarization of the surface dipoles may be responsible for a third contribution. This effect is of importance in the chemisorption of atoms of alkali metals on metal surfaces. It causes, also in the chemisorption of other gases a minimum or a maximum in th e surface potential as a function of 8. 4. T he surface potential effect together with the depolarization may bring about a more or less linear decrease of QSwith 8. This leads to the logarithmic adsorption isotherm (74). A slight heterogeneity together with the surface potential effect may give the same result. 5. A pronounced contribution of surface heterogeneity may result in a more or less exponential decrease of QO with 8, leading t o th e exponential adsorption isotherm (76). 6. On the surface of semiconducting or insulating oxides th e decrease of Q Owith 0 may be caused by microdefects (see Sec. IX,2), a s discussed by Vol’kenshtein (386). If practically all of these microdefects were of biographical origin, th e heat of chemisorption would be constant; if, on

140

J. H. DE B O E R

the other hand, they were all of thermal origin, a n exponential adsorption isotherm would be obtained.

X. SIMULTANEOUS ADSORPTION OF DIFFERENT MOLECULES OR ATOMS I . Simultaneous Adsorption of Diferent Molecules in Physical Adsorption Phenomena

We shall make only a few remarks with respect t o the physical adsorption of different molecules on the same adsorbent. The physical adsorption of solvent molecules in catalytic reactions carried out in solutions is often overlooked. The author cannot resist the impression that in many of Maxted’s (387) experiments on the poisoning of catalysts the adsorption of solvent molecules plays a more important role than Maxted thinks. It may be expected that a critical study of the adsorption of the solvent molecules and their influence on the adsorption of the poison molecules will lead t o even more important conclusions on poisoning than have already been reached by Maxted and collaborators in their systematic and elaborate work. Another remark may be made with respect to those estimations of surface areas where use is made of the adsorption of strongly polar molecules, such as those of fatty acids. Generally a unimolecular layer of mutually erected long molecules with their polar ends on the surface is assumed t o be formed on adsorption from solutions. This holds only for surfaces of ionic or sufficiently strong polar character and even in those cases only if the solvent molecules do not interfere. Since the work of Harkins and Gans (388),benzene is frequently used as a solvent for the fatty acids. Houben (389),in his thesis, showed that even this solvent is not completely indifferent and that a mixed layer of benzene and lauric acid is formed on the surface of alumina. It is only with still less strongly adsorbed, nonpolar, solvents, such as n-pentane, that reliable results are obtained. Fortuin (390) found that lauric acid, from n-pentane solutions, is adsorbed t o form a complete unimolecular layer on alumina surfaces that are depleted of chemisorbed water (OH groups) as well as on those that are completely covered with OH groups. Lauric acid forms a complete unimolecular layer even on alumina surfaces possessing a complete chemisorbed water layer (OH groups), on top of which a second layer of physically bound water is adsorbed with strong polar forces. The lauric acid molecules do not replace any of these water molecules; they are only adsorbed on top of them. If, however, a multimolecular water layer is present (in capillaries wide enough to admit lauric acid molecules), lauric

ADSORPTION PHENOMENA

141

acid replaces all but the two water layers just mentioned and is adsorbed on top of these two.

in Chemisorption; the Relative Amounts that are Chemisorbed Catalysis offers numerous examples of the simultaneous chemisorption of different molecules or atoms. The extremely important problems of poisoning, of promoting and of selective catalysis depend on this phenomenon. I n the past numerical estimations of the relative amounts that will be bound by chemisorption in equilibrium were mostly based on the assumption that Langmuir’s adsorption isotherm would be valid. I n such a case the following equation is easily derived: 2. Simultaneous Adsorption of Di$erent Species

where el and O2 are the degrees of occupation for species 1 and species 2 respectively, p l and p2 the partial pressures of the two gases in equilibrium with the adsorbed layer, and Q 1 and Qz the heats of adsorption of the two species. I n the conceptions underlying Langmuir’s adsorption isotherm the heats of adsorption are constant and independent of the amount that is adsorbed. We learned, in Sec. IX, th at this does not hold for chemisorption and we shall, therefore, derive a n expression based on the logarithmic adsorption isotherm [Eq. (7 1) or (74)l. We shall assume th at the fall of the heats of chemisorption and the rise of the activation energies are caused by the surface potential effect (Secs. IX,4,6,9, and 13) and that these changes vary linearly with 8. If we now also assume that the dipoles of both sorts of chemisorbed atoms are directed in the same sense (hence either both are positive or negative), it is a logical consequence th at the decrease in the heat of chemisorption of species 2 caused by the dipoles of species 1 is the same as the decrease t ha t the dipoles of species 1 cause in the heat of chemisorption of their own species. And we shall, therefore, also assume th a t the activation energies of both species are influenced in the same way by the dipole layers of both species. Both activation energies and heats of chemisorption depend, therefore, on the sum of the degrees of occupation of 8.J. Consequently we can write both species: (el

+

+ +

(El)(e,+e,)= OEI Y l e l (E2)(8,+8,)

=

OE2

(&J(e,+e,)

= =

0‘21

(&2)(e1+e,)

oQ2

?’I&

+ +

yzez y202

- ciei - C Z ~ Z - C1e1 - ~ 2 0 2

(78) (79) (80) (81)

142

J. H . DE

BOER

where the symbols stand for the initial values of the activaand tion energies and heats of chemisorption respectively and y and c are constants. The rate of adsorption of species 1 may now be written as (vadJl = k a l p l ( l -

el - e 2 ) e - ( o E i + y i ~ i + r 2 ~ 2 ) / R T

(82)

and its rate of desorption as (vdea)

=

kdlOle- ( ~ Q I - c01-cz I

O~+aEi+yi Oi+yiOz)/RT

(83)

for, the activation energy for desorption is the sum of the heat of chemisorption and the activation energy for adsorption (See. V,9). Adsorption equilibrium means (uads)l = ( v d d l

On rearranging terms we obtain

A similar equation holds for

02:

Dividing (84) b y (85) gives

The product ( k a l / k d l ) ( k d 2 / k a 2 ) may be p u t equal to unity, just as is done in deriving Eq. (77). Its deviation from unity is negligible with respect t o the e power. We, therefore, obtain

The only difference between Eqs. (87) and (77) is th a t in Eq. (87) we have t o take the initial heats of chemisorption as they are when no surface potential effect is noticeable as yet. It may be remarked that the same result would have been obtained if we had started from the conception of Temkin (Sec. IX,8). Equation (87) can be used only when the dipoles of both species have the same electrical direction. If one of the species forms a positive dipole layer and the other a negative one, the one dipole layer will decrease the activation energy and increase the heat of chemisorption of the other. The two sorts of

ADSORPTION PHENOMENA

143

molecules (atoms) will, therefore, facilitate each other’s chemisorption, in regard t o both rate and drength. We may then write (&)(61+&)

+

- Yzez

= o ~ l Ylel

+ ( Q ~ ) ( ~ , + ~ ~ ~) clel + czez

(E~)(~= , +o ~ ~~~z )ylel and

=

( Q ~ ) ( ~ , +=~ ~o , )~

yZez

+ clel - czez z

Instead of Eq. (87) we obtain

If 81 and 0 2 do not differ too much and the pressures are of the same order of magnitude we may probably write 2(ci81 - czez) =

OQI

-

oQz

(89)

Two sorts of chemisorbed, atoms, giving dipoles of opposite signs, will attract each other and no homogeneous distribution may be expected. The formation of pairs or of agglomerates on the surface th a t will result from this attraction will influence the validity of Eq. (88) or (89). These equations have never been put t o the test. The atoms will in any case, as noted above, facilitate each other’s chemisorption, in regard t o both rate and strength. 3. The Chemisorption of Digereat Atoms Giving Dipoles of the

Same Sign The simultaneous chemisorption of CzHz and C2H4 on nickel may serve as a n example. The initial heats of chemisorption are 67 kcal./mole for CzHz and 58 kcal./mole for CzH4. At equal pressures the proportion of chemisorption a t 50°C. will be [Eq. (87)] -~ eCoHz = 106 eC@~

The surface of the catalyst is, therefore, practically covered with acetylene only, and when a mixture of the two gases is hydrogenated it is only acetylene that is converted to ethylene until practically all acetylene has disappeared. The selective hydrogenation of nonsubstituted and substituted acetylene and ethylene mixtures may be ascribed to the selective chernisorption of the gases (991). A special type of selective catalytic reaction is the poisoning of the catalytic formation of ammonia by oxygen or oxygen-containing gases,

144

J. H. D E BOER

such as CO, COZ, or HzO. All these gases react readily with the surface of the iron catalyst, producing chemisorbed oxygen atoms on th a t surface (392). I n the reaction with hydrogen these oxygen atoms compete with the nitrogen atoms. As oxygen has a n appreciably higher heat of chemisorption than nitrogen, the presence of even a very small amount of oxygen in the gas mixture results in a serious poisoning effect. I n both cases mentioned above, the two gases, chemisorbed simultaneously, do not react with each other. When such reactions can occur, more complicated relationships may be expected. Relatively few measurements have been made by direct chemisorption studies. Beeck (393) studied the simultaneous chemisorption of nitrogen and hydrogen on iron films. As we have seen in Sec. IX,12, the activation energy for nitrogen adsorption, after it has covered 20% of the surface has practically become too high for further chemisorption a t room temperature. Beeck found that when hydrogen is first adsorbed to a degree of coverage O, the surface will take up less nitrogen, namely to a n amount of ON

=

0.2(1 - 6,)

Boudart (394) explains this by assuming th a t the dipole layer formed by hydrogen affects the activation energy of nitrogen in the same sense as does nitrogen itself, but quantitatively t o a smaller degree. If a completely covered hydrogen film affects the activation energy of nitrogen just as much as a film of nitrogen covering only 20% of the surface, and if both effects bear a linear relation to 0, the above-mentioned relationship may be understood. The heat of chemisorption of hydrogen, adsorbed on iron that has previously been covered with nitrogen u p t o 6 = 0.18, is, indeed, lower than the heat of chemisorption of hydrogen adsorbed on a clean surface (395). The heat of chemisorption of CO on a n iron film partly covered with nitrogen is also lower than on a clean film, but Bagg and Tompkins (395) found t hat hydrogen when adsorbed on a film partly covered with CO shows a higher heat of adsorption than when adsorbed on a clean film. Beeck (393) as well as Bagg and Tompkins (395) has included oxygen as one of the gases in these investigations. As, however, oxygen penetrates easily into the surface layers of the adsorbents (Sec. VII,6), complications arise.

4.

Contaminated Surfaces

Contaminations or poisons on surfaces may have a far more import a nt effect than the blocking of a part of the surface. Apart from the quantitatively large influence that they may have when specifically adsorbed in a microporous system (396), they may also influence the giving

145

ADSOIZPTION PHENOMENA

off or taking up of electrons by the surface. Roginskii (3977, who has extensively studied the influence of surface contaminations, classifies them into four groups, the most important of which is constituted by the socalled “modifiers.” It is these modifiers that influence the electric properties of the catalysts (398). These modifying contaminations exercise a great influence also on chemisorption phenomena. We may give just one example. It is well

c-

d i s t a n c e in fhe m e t a l

-

d i s t a n c e f r o m the m e f a l

FIG.39. Potential curves indicating the chemisorption on and the dissolution in iron having a noncontaminated surface.

known t ha t the dissolution of hydrogen into iron is a n endothermic process (Sec. VII,7). Hydrogen atoms, either obtained in a gaseous atmosphere or produced by the action of an acid medium on iron, however, penetrate easily into iron. The presence of sulfide ions on the surface of the iron facilitates this proctm (S99). Sulfide ions may be assumed t o form a dipole layer with the negative poles pointing away from the metal. The sign of this dipole layer is, therefore, the same as that of the dipole layer formed by the chemisorption of hydrogen itself when this is adsorbed a t room temperature. Consequently, the heat of chemisorption of hydrogen on iron will be severely reduced by the presence of the sulfur contamination. At the same time a n energy of activation for the dissociative chemisorption of molecular hydrogen will result from the action of the sulfur layer. The relations are schematically shown in Figs. 39 and 40. I n Fig.

146

3. H. DE BOER

39 curve A B C D gives the dissociative chemisorption of molecular hydrogen; there is no activation energy on a pure iron surface (level C lower than level A ) . The heat of chemisorption is rather high, which involves a low level for the minimum D. Molecular hydrogen can penetrate (in atomic form) into iron, provided that the kinetic energy is high enough t o reachlevel E , from which it can move into the metal (level F ) . The exact place of the surface cannot be pictured in such a potential-curve scheme; it is somewhere in

-

distance in the m e t a l

-

distance from the rnclal

FIG.40. Potential curves indicating the chemisorption on and the dissolution in iron having a contaminated surface.

the region of the vertical lines indicated by dashes. Atomic hydrogen, coming from level G, has sufficient energy t o penetrate into the metal (curve GDEF) but, owing to the fact that level D is far lower than E , most of it stays a t the surface in chemisorbed form. When a contamination is present which produces a dipole layer, as sulfur does, the dissociative chemisorption of molecular hydrogen is given by (Fig. 40) ABC’D’; there is an activation energy (difference between levels C’ and A ) ;the heat of adsorption is severely reduced; it is even negative (endothermic chemisorption) . The dissolution of molecular hydrogen proceeds less easily than in the case of a pure-iron surface. The kinetic energy has t o be sufficient to overcome the difference between C’

ADSORPTION PHENOMENA

147

and A in this case. Atomic hydrogen coming from G will penetrate easily into the metal. Far fewer hydrogen atoms will stay a t the surface, because D’is not much lower than E’. ( I t may even be higher.) Surface contaminations that produce a dipole layer of the same electric sign as that produced by 1,he adsorbed atoms themselves decrease the heat of chemisorption and produce and increase a n activation energy. This is why, activation energies are always found for the chemisorption of hydrogen on metal surfaces which are not sufficiently reduced or which are contaminated with impurities forming negative dipoles or which are partially oxidized (Secs. V,9 and IX,9). The heats of chemisorption will be lower a t the same time. If the nickel powder of Eucken described in Sec. IX,l (Fig. 28, curve B) were contaminated with so many oxygen ions (insufficiently reduced) that these would produce a surface-potential effect of the same magnitude as would be created by a chemisorption of hydrogen itself of $ = 0.3 curve B of Fig. 28 would have t o be shifted t o the right over a distance of 0 = 0.3 and it would practically coincide with curve E of Schuit and de Boer. 6. Mutual Assistance of Chemisorbed Atoms

I n Sec. X,2 we learned that atoms producing dipoles of opposite character will facilitate each other’s chemisorption, in regard t o both rate and strength. There are many examples to be found in the literature on electron emission where this effect is obvious (400). When the work function of tungsten is strongly increased by the chemisorption of oxygen, not only the atoms of alkali metals and alkaline earth metals will be bound as ions, but even metal atoms with far higher ionization energies will be chemisorbed in the ionic state. Phenomena of this kind play a role in the action of oxygen on the surfaces of iron, copper, or nickel when iron, copper, or nickel ions move next, to or on top of chemisorbed oxygen ions (Sec. VII,6), causing the penetration of the chemisorbed oxygen into the surface layers of the metal and causing a reverse of the surface potential a t the same time. Similarly cesium, adsorbed “on top of” chemisorbed oxygen on tungsten, is far more strongly bound than cesium chemisorbed on a bare tungsten surface; as the result of the simultaneous chemisorption of both species the work function is decreased to such a low value as cannot be obtained with cesium alone. I n Sec. IX,1 we remarked th at copper films do not chemisorb hydrogen, whereas copper powder does. We must add here that also pure and thoroughly reduced copper powder does not chemisorb hydrogen (401). However, when copper powder is not completely reduced it shows a lower work function than does pure copper (40Z),because copper ions on

148

J. H . DE BOER

top of oxygen ions produce, together with these latter ions, such a surface potential that, just as in the case of iron (Sec. VII,6), the work function is decreased. This complex layer has, therefore, produced a dipole layer with its positive side pointing away from the surface. Such a dipole layer facilitates the chemisorption of hydrogen, forming dipoles of opposite direction (403). We may also say th a t the reduced work function, caused by the copper ions and the oxygen ions together, facilitates the formation and chemisorption of negatively charged hydrogen. Hydrogen chemisorption a t lower temperatures or even at room temperature leads to dipoles pointing with their negative poles away from the metal. At higher temperatures another type of hydrogen chemisorption ( B adsorption, Sec. VI,3) occurs, which is probably connected with the formation of positively charged hydrogen. It is not improbable th a t an oxygen contamination of the surface promotes the B adsorption or, as already stated in Sec. VI,3, prevents the A adsorption. A dipole layer, formed by oxygen, with the negative pole pointing away from the surface could indeed reduce the heat of chemisorption of negatively charged hydrogen atoms ( A adsorption) to such a low value th a t this type of adsorption would not take place any more, although on the other hand this layer would promote B adsorption. Technical ammonia catalysts, consisting of iron with “promotors” such as aluminum oxide, potassium oxide, etc., working a t high temperatures (400” t o 500°C.) most probably adsorb hydrogen in the positively charged B form. This form of chemisorption will facilitate the chemisorption of nitrogen, which is negatively charged in its chemisorbed state. It is well known that nitrogen can easily be taken up by iron when it is heated in ammonia, which dissociates into hydrogen and nitrogen. A presorption of hydrogen markedly facilitates also the chemisorption of molecular nitrogen and its dissolution into iron (404). Conversely a previous adsorption of nitrogen can give rise to a n increase in the chemisorption of hydrogen (405). X I . SOMEREMARKS ON CATALYSIS AND CHEMISORPTION 1. Heat of Chemisorption and Catalysis

The history of the views on the mechanism of the parahydrogen conversion is a very illustrative one. For a long time it was thought th a t the heat of chemisorption of the chemisorbed hydrogen atoms was too high t o enable the atoms to react together on the surface so as t o produce molecular hydrogen. The hydrogen chemisorption was even thought t o be irreversible at or below room temperature. Subsequently it was found (406) that a t sufficiently high hydrogen pressures a normal reversible

ADSOltPTION PHENOMENA

149

chemisorption could be measured. The decrease of the heat of chemisorption with increasing amount a,dsorbed proved to proceed so far th a t sufficiently low values were obtained a t such e values as are found during the catalytic reaction. The low values of the heat of chemisorption enable the catalytic reaction to take place. Similar effects may be obtained with “promotors ” producing such surfaee-potential effects that the heat of chemisorption is low enough for chemical action. Low heats of chemisorption are especially needed in the case of the Langmuir-Hinshelwood mechanism (Sec. VII,5). A promotor may also cause a sufficient amount of chemisorbed atoms to be present. Boudart (40’7) remarks that the presence of aluminum or potassium on an iron ammonia catalyst may produce a lower work function, resulting in the chemisorption of more nitrogen than would otherwise be possible (0 = 0.2; see Sec. IX,12). The first nitrogen atoms bound by such a “promotion” caused by aluminum or potassium and oxygen together (just as with copper promoted by oxygen; see Sec. IX,5) will be bound more firmly. At the same time, however, a larger amount of nitrogen can be chemisorbed and the heat of chemisorption is sufficiently decreased for all these nitrogen atoms to react. It may be repromoted by marked t ha t the increased hydrogen-B-chemisorption, the “promotors,” as discussed in Sec. X,5, may also facilitate the nitrogen chemisorption in the same sense. Here again it is of major importance t ha t the heat of chemisorption be reduced to a sufficiently low value. As the rate of the ammonia catalysis is governed by the rate of the chemisorption of nitrogen, hence by the activation energy of nitrogen, i t is important that this activation energy be sufficiently lowered. During this catalysis no adsorption equilibria will be reached. The nitrogen is taken away by the hydrogen a,t a quicker rate than by its own desorption. 2. Endothermic Chemisorption

The heat of chemisorption, which must be low in order t o enable catalysis t o take place, may even be negative. I n various sections we have seen that endothermic ohemisorption may play a n important role (Secs. V,9, VI,3,4,5, and X,4:).Figure 40 shows th a t surface contaminations can “promote” endothermic chemisorption. I n nickel, a s in iron, hydrogen atoms can be dissolved endothermically. It is highly probable that dissolved hydrogen atoms react from the metal phase with chemisorbed hydrocarbons. Hydrogen atoms can also be taken up from hydrocarbons by such catalysts, the hydrogen atomn immediately disappearing in the metal.

150

J. H. DE BOER

A measurable amount of chemisorbed atoms need not be present during catalysis. It is essential, however, that the heat of chemisorption be sufficiently low, or even negative.

REFERENCES 1. Hill, T. L., Advances in Catalysis 4, 211 (1952). 2. Kwan, T., Advances in Catalysis 6, 67 (1954). 5’. de Boer, J. H., Advances in Colloid Sci. 3, 1 (1950).

4. Fraser, R. G. J., “Molecular Rays.” Cambridge, New York, 1931. 6. Estermann, J., Frisch, R., and Stern, O., 2. Physik 78, 348 (1931). 6. de Boer, J. H., Advances in Colloid Sci. 8, 5 (1950). ?. Verwey, E. J. W., Rec. trav. chim. 66, 521 (1946). 8. Shuttleworth, R., Proc. Phys. SOC.(London) 62A, 167 (1949). 9. Benson, G. C., and Benson, G. W., Can. J . Chem. 33, 232 (1955). 10. de Boer, J. H., Rec. trav. chim. 69, 826 (1940). 11. de Boer, J. H., Koninkt. Ned. Akad. Wetenschap. Proc. 49, 1103 (1946). 12. de Boer, J. H., and Verwey, E. J. W., Rec. trav. chim. 66, 443 (1936). 13. de Boer, J. H., Advances in Colloid Sci. 3, 35, 38 (1950). 14. Born, M., and Mayer, J. E., 2. Physik 76, 1 (1932). 16. Mayer, J. E., and Maltbie, M., 2. Physik 76, 748 (1932). 16. de Boer, J. H., and Verwey, E. J. W., Rec. trav. chim. 66, 443 (1936). 17. Verwey, E. J. W., and de Boer, J. H., Rec. trav. chim. 69, 633 (1940). 18. Brunauer, S., “The Adsorption of Gases and Vapors,” Chapter IV. Oxford, New York, 1943. 19. de Boer, J. H., Trans. Faraday SOC.82, 10 (1936). 20. Hellmann, H., “Einfuhrung in die Quantenchemie,” Chapter V. Deuticke, Leipzig, 1937. 21. de Boer, J. H., Advances in Colloid Sci. 3, 21 (1950). 2%.For a further discussion, see Margenau, H., Revs. Mod. Phys. 11, 1 (1939); de Boer, J. H., Advances in Colloid Sci. 3, 20 (1950). 23. London, F., 2. Physik 63, 245 (1930). 24. Slater, J. C., and Kirkwood, J. G., Phys. Rev. 37, 682 (1931). 26. de Boer, J. H., and Heller, G., Physica 4, 1045 (1937). 26. Polanyi, M., and London, F., Naturwissenschaften 18, 1099 (1930). 27. London, F., 2. physik. Chem. B11, 246 (1931). 88. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26,225 (1934); de Boer, J. H., Trans. Faraday SOC.32, 10 (1936). 89. Brunauer, S., “Physical Adsorption of Gases and Vapors,” Chapter VII. Oxford,

New York, 1943. 5’0. de Boer, J. H., Advances in Colloid Sci. 3, 27, 46 (1950). 31. Margenau, H., and Pollard, W. G., Phys. Rev. 60, 128 (1941).

32. de Boer, J. H., Advances in Colloid Sci. 3, 44 (1950). 33. Huckel, E., “Adsorption und Kapillarkondensation,” p. 126. Akademieche

Verlagsges., Leipzig, 1928.

$4. For a more detailed treatment see de Boer, J. H., Advances in Coolloid Sci. 3, 13 (1950). 36. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26, 225 (1934). 36. de Boer, J. H., and Dippel, C. J., 2. physik. Chem. B26, 399 (1934). 37. de Boer, J. H., and Dippel, C. J., Rec. trav. chim. 62, 214 (1933).

ADSORPTION PHENOMENA

151

38. Wheland, G. W., “The Theory of Resonance.” Wiley, New York, 1945. 39. de Boer, J. H., Advances in Colloid Sci. 3, 33 (1950). 40. Lenel, F. V., Z. physik. Chem. B23, 379 (1933).

41. Drain, L. E., Trans. Faraday Xoc. 49, 650 (1953).

4.2. Brunauer, S., “The Adsorption of Gases and Vapors,” p. 28. Oxford, New York, 1943. 43. Mignolet, J. C. P., Discussions Faraday SOC.No. 8, 105 (1950); J . Chem. Phys. 21, 1298 (1953). 44. de Boer, J. H., and Kruyer, S., Koninkl Ned. Akad. Wetenschap. Proc. B66, 451 (1952); B66, 67 (1953); B66, 236 (1953); B66, 415 (1953); B67, 92 (1954); B68, 61 (1955). 46. de Boer, J. H., “The Dynamical Character of Adsorption,” p. 146. Oxford, New York, 1953. 46. de Boer, J. H., “The Dynamical Character of Adsorption,” p. 168. Oxford, New York, 1953; Ned. Tijdschr. Natuurk. 19, 283 (1953); Kruyer, S., Thesis, Delft, 1955. 47. cf. also Magnus, A., Z. physik. Chem. A142, 401 (1929); Trans. Faraday SOC.28, 386 (1932). 48. de Boer, J. H., “The Dynarnical Character of Adsorption,” p. 155. Oxford, New York, 1953. 49. Ives, H. E., Boston Meeting American Physical Society, Dec. 1922; see J . Franklin Inst. 201, 47 (1926). 60. Langmuir, I., and Kingdon, K. H., Science 67, 58 (1923). 61. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” 58. Cambridge, New York, 1935. 52. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” 520. Cambridge, New York, 1935; de Eloer, J. I<., and Veenemans, C. F., Physica 1, 753 (1934). 63. de Boer, J. H., Advances in Colloid Sci. 3, 2 (1950). 54. Riyanoff, S., Z. Physik 71,325 (1931); Lukirsky, P. I., and Riyanoff, S., ibid. 76, 249 (1932); Lukirsky, P. I., Physik. 2.Sowjetunion 4, 225 (1933). 65. de Boer, J. II., “Electron Emission and Adsorption Phenomena,” p. 215. Cambridge, New York, 1935. 66. Wheland, G. W., “The Theory of Resonance,” p. 229. Wiley, New York, 1945. 67. Sidgwick, N. V., “The Electron Theory of Valence.” Oxford, New York, 1927; see also Wells, A. F., “Structural Inorganic Chemistry,” p. 49. Oxford, New York; Pauling, L., “The Nature of the Chemical Bond,” p. 7. Cornell U. P., Ithaca, 1948. 58. de Boer, J. H., and Verwey, IS. J. W., Rec. trav. chim. 66, 443 (1936). 69. Verwey, E. J. W., and de Boer, J. H., Rec. trav. chim. 66, 675 (1936). 60. Mignolet, J. C. P., Discussions Faraday Soc. NO.8, 105 (1950). 61. Maxted, E. B., Advances in Catalysis 3, 129 (1951). 62. Suhrmann, R., 2. Elektrochem. 66, 351 (1952). 63. de Boer, J. H., “Elektronenemission und Adsorptionserscheinungen,” p. 59. Leipsig, 1937; see reference 32. 64. Morse, P. H., Phys. Rev. 34, 57 (1929). 66. Eley, D. D., Discussions Faraday SOC.NO. 8, 34 (1950). 66. Pauling, L., “The Nature of ,,he Chemical Bond,” p. 60. Cornell U. P., Ithaca, 1948. 67. Kwan, T., Advances in Catalysis 6, 91 (1954).

152

J. H. DE BOER

68. Trapnell, B. M. W., “Chcmisorption,” pp. 147-155. Butterworths’, London, 1935. 69. Roberts, J. K., Proc. Roy. SOC.A162, 445 (1935). 70. Beeck, O., Advances i n Catalysis 2, 151 (1950). 71. Trapnell, B. M. W., “Chemisorption,” p. 61. Butterworths’, London, 1935. 72. Schuit, G. C. A,, and de Boer, J. H., Nature 168, 1040 (1951). 73. Beeck, O., and Wheeler, A., J. Chem. Phys. 7, 631 (1939). 74. Zwietering, P., and Roukens, J., Trans. Faraday SOC.60, 178 (1954). 76. Roberts, J. K., Nature 137, 659 (1936). 76. Becker, J. A., and Hartman, C. D., J . Phys. Chem. 67,153 (1953). 77. Trapnell, B. M. W., Trans. Faraday SOC.48, 160 (1952). 78. Trapnell, B. M. W., “ChemisorptionJJJChapters VI and VII. Butterworth-< London, 1935. 79. Dowden, D. A., Research 1, 239 (1948). 80. See also a discussion in de Boer, J. H., Chem. Weekblad 47, 423 (1951). 81, Beeck, O., Discussions Faraday SOC.No. 8, 126 (1950). 82. Pauling, L., Proc. Roy. SOC.A196,343 (1949). 83. Couper, A,, and Eley, D. D., Discussions Faraday SOC.NO. 8, 172 (1950). 84. Dowden, D. A , , and Reynolds, P. W., Discussions Faraday Soc. No. 8, 184 (1950) 85. Dilke, M. H., Eley, D. D., and Maxted, E. B., Nature 161,804 (1948). 86. de Boer, J. H., and Bruining, H., Physica 6,941 (1939). 87. Briiining, H., and de Boer, J. H., ibid. 4, 473 (1937); Bruining, H., ibid. 8, 1161 (1941). 88. de Boer, J. H., and Kraak, H. H., Bee. trau. chim. 66, 1103 (1937). 89. Suhrmann, R., Z. Elektrochem. 66,351 (1952). 90. Went, J. J., Physica 8,233 (1941). 91. Twigg, G. II., Trans. Faraday SOC.42, 657 (1946). 92. Braunbek, W., 2. Naturforsch. 36, 216 (1948). 93. Koller, I<., Vcrhandl. deut. physik. Ges. [3] 21, 11 (1940). 94. de Boer, J. H., and Verwey,-E. J. W., Proc. Phys. Soc. (London)49,extra part 59 (1937). 95. Verwey, E. J. W., and de Boer, J. H., Rec. trav. chim. 66,531 (1936). 96. Verwey, E . J. W., Haayman, P. W., and Romeijn, F. R., Chem. Weekblad 44, 705 (1948); Verwey, E. J. W., Bull. SOC. chim. France D93 (1949); Verwey, E. J. W., in “Semiconducting Materials” (H. K. Henisch, ed.), p. 151. Butterworths’, London, 1951. 97. See also Rces, A. L. G., “Chemistry of the Defect Solid State,” Methuen’s Monographs, London, 1954. 98. Garner, W. E., Gray, T. J., and Stone, F. S., Proc. Roy. SOC.Al97, 294 (1949). 99. Garner, W. E., and Veal, F. J., J . Chem. SOC.,p. 1487 (1935). 100. de Boer, J. H., and Verwey, E. J. W., Proc. Phys. SOC.(London) 49, extra part 59 (1937). 101. Beebe, R. A , , and Dowden, D. A,, J. Am. Chem. SOC.60, 2912 (1938). 102. Dowden, D. A., and Garner, W. E., J. Chem. SOC.,p. 893 (1939). 103. Garner, W. E., Stone, F. S., and Tiley, P. F., Proc. 1 2 0 ~SOC. . A211, 472 (1952). 104. Garner, W. E., and Ward, T., J . Chem. SOC.,p. 857 (1939); Garner, W. E., J . Chem. SOC.,p. 1239 (1947). 105. Beebe, R. A,, and Dowden, D. A,, J. A m . Chem. SOC.60, 2912 (1938). 106. Dowden, D. A , , and Garner, W. E., J . Chem. SOC.,p. 893 (1939). 107. Garner, W. E., Gray, T. J., and Stone, F. S., Discussions Faraday SOC.No. 8, 246 (1950).

ADSORPTION PHENOMENA

153

108. Hauffe, K., and Engell, H. J., 2. Elektrochem. 66, 366 (1952).

109. Winter, E. R. S., Discussions Faraday SOC.No. 8, 231 (1950). 110. Herington, E. F. G., and Rideal, E. K., Proc. Roy. SOC.A184, 434, 447 (1945); A190, 289, 309 (1947). 111. For a discussion of various active spots see also de Boer, J. H., Advances in Colloid Sci. 3, 8-38 (1950). 112. de Boer, J. H., Discussions Faraday SOC.No. 8, 93 (1950). 113. de Boer, J. H., Ned. Tijdschr. Natuurk. 19,283 (1953). 114. Gulbransen, E. A., and Andrew, F. F., Znd. Eng. Chem. 44,1039 (1952);Polley, M. H., Schaeffer, W. D., and Smith, W. R., J. Phys. Chem. 67, 469 (1953); Singleton, J. H., and Hakey, G. D., ibid. 68, 1011 (1954); Amberg, C. H., Spencer, W. B., and Beebe, R. A., Can. J. Chern. 33, 305 (1955). 116. de Boer, J. H., “The Dynarnical Character of Adsorption,” Chapters VIII and IX. Oxford, New York, 1953. 116. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26, 238 (1934). 117. London, F., 2. physik. Chem. B11, 246 (1931). 118. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26, 225 (1934). 119. Brunauer, S., “The Adsorption of Gases and Vapors,” p. 206. Oxford, New York, 1943. 120. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 451 (1952); B66, 67, 236, 415 (1953); B67, 92 (1954); B68, 61 (1955). 121. Kruyer, S., Thesis, Delft, 1955; de Boer, J. H., Ned. Tijdschr Natuurk. 19, 283 (1953). 122. Klumb, H., 2. Physik 47, 652 (1928). 123. Abendroth, B., 2. Physik 86,530 (1933). 12.4. Ouellet, C., and Rideal, E. K., J. Chem. Phys. 3, 150 (1935). 126. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” p. 160. Cambridge, New York, 1935. 126. Lennard Jones, J. E., and Dent, B. M., Trans. Faraday SOC.24, 92 (1928). 127. Lenel, F. V., Z. physik. Cherrz. B23, 379 (1933). 128. Orr, W. J. C., Proc. Roy. &rc. A173, 349 (1939); Trans. Faraday SOC.36, 1247 (1939). 169. Brunauer, S., “The Adsorption of Gases and Vapors,” Chapter VII. Oxford, New York, 1943. 130. de Boer, J. H., Advances in CooEEmd Sci. 3, 39 (1950). 131. Drain, L. E., Trans. Faraday SOC.49, 650 (1953). 132. Drain, L. E., and Morrison, J. A., Trans. Faraday SOC.49, 654 (1953). 133. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26, 225 (1934); for a corrected calculation see de Boer, J. H., Advances in Colloid Sci. 3, 41 (1950). 13.4, Healey, F. H., Chessick, J. J., Zettlemoyer, A. C., and Young, G. J., J. Phys. Chem. 68, 887 (1954). 136. Chessick, J. J., Zettlemoyer, A. C., Healey, F. H., and Young, G. J., Can. J . Chem. 33, 251 (1955). 136. de Boer, J. H., Ned. Tijdschr. Natuurk. 19, 283 (1953); Kruyer, S., Thesis, Delft, 1955. 137. de Boer, J. H., and Dippel, .;1 J., 2. physik. Chem. B26, 399 (1934). 138. de Boer, J. H., and Broos, J , 2. physik. Chem. B16, 281 (1932). 139. de Boer, J. H., Z. physik. Ckem. B16, 397 (1932). 1.40. de Boer, J. H., 2. physik. Ckem. B16, 300 (1932). 141. Houben, G. M. M., Thesis, Delft, 1951. 142. Eley, D. D., Advances in Caralysis 1, 157 (1948).

154

J. H. DE BOER

149. Beeck, O., Advances in Catalusis 2, 151 (1950).

144. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 236 (1953); Kruyer, S., Thesis, Delft, 1955. 146. Burshtein, R., Acta Physicochim. U.R.S.S. 8, 857 (1938). 146. Eucken, A , , and Hunsmann, Mi.,2. physik. Chem. B44, 163 (1939). 147. Cf. Trapnell, B. M. W., Cheniisorption,” pp. 56-61. Butterworths’, London,

1955. Emmett, P. H., and Harkness, R. W., J . Am. Chem. SOC.67, 1631 (1935). Gauchman, S. S., and Royter, W. A., J . Phys. Chern. U.S.S.R. 13, 593 (1939). Communicated by Taylor, H. S., Advances in Catalysis 1, 11 (1949). Zwietering, P., in a contribution by Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 316. Reinhold, New York (1955). 162. Kummer, J. T., and Emmett, P. H., J . Phys. Chem. 66, 258 (1952). 169. Beebe, R. A., and Dowden, D. A., J . Am. Chem. Soc. 60, 2912 (1938). 164. Molinari, E., and Parravano, G., J . Am. Chem. SOC.76, 5233 (1953). 156. Langmuir, I., J . Am. Chem. SOC.34, 1310 (1912); 38, 2270 (1916). 166. Johnson, M. C., Proc. Roy. SOC.A123, 603 (1929); A132, 67 (1931); Trans Faraday SOC.28, 162 (1932). 167. de Boer, J. H., and Lehr, J. J., 2. physik. Chem. B22, 423 (1933). 168. de Boer, J. H., and van Steenis, J., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 572, 578, 587 (1952). 169. de Boer, J. H., and Lehr, J. J., 2. physik. Chem. B24, 98 (1934); de Boer, J. H., and Dippel, C. J., ibid. B26, 399 (1934). 160. de Boer, J. II., and van Steenis, J., Koninkl. Ned. Akad. Wetenschap. Proe. B66, 587 (1952). 161. Johnson, T. H., J . Franklin Inst. 206, 301 (1928); 207, 629, 639 (1929); 212, 507 (1931); Phys. Rev. 36, 1432 (1930). 162. Private communication by P. Zwietering. 163. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 236 (1953). 164. Verwey, E. J. W., and de Boer, J. H., Rec. trav. chim. 66, 675 (1936). 165. Bonhoeffer, K. F., Farkas, A., and Rummel, K. W., Z.physik. Chem. B21, 225 (1933). 166. Juaa, R., 2nd Langbein, R., 2. Elektrochem. 46, 689 (1939). 167. de Boer, J. H., and Kraak, H. H., Rec. trav. chzm. 66, 1103 (1937). 168. Suhrmann, R., 2. Elektrochem. 66, 351 (1954). 169. See Massey, H. S. W., “Negative Ions,” p. 64. Cambridge, New York, 1938. 170. Kolthoff, I . M., and Jordan, J., J . Am. Chem. SOC. 74, 570 (1952). 171. Fortuin, J. P., Thesis, Delft, 1952. 172. de Boer, J. H., and Teves, M. C., 2. Physik 66, 489 (1930); 73, 192 (1931); 83, 521 (1933). 173. de Boer, J. H., Chem. Weekblad 29, 34 (1932). 174. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” Chapter IX. Cambridge, New York, 1935. 175. de Boer, J. H., Z.physik. Chem. B16, 397 (1932). 176. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26, 238 (1934). 177. Custers, J. F. H., and de Boer, J. H., Physica 3, 407 (1936). 178. de Boer, J. H., 2. physik. Chem. B18, 49 (1932). 179. Weitz, E., and Schmidt, F., Ber. 72, 2099 (1939); Weitz, E., Schmidt, F., and Singer, J., 2. Elektrochem. 46, 222 (1940); Weitz, E., ibid. 47, 65 (1941). 148. 149. 160. 161.

ADSORPTION PHENOMENA

155

180. de Boer, J. H., 2. Elektrochem. 44, 488 (1938). 181. de Boer, J. H., and Houben G. M. M., Koninkl. Ned. Akad. Wetenschap. Proc.

B64, 421 (1951). 182. Fortuin, J. M. H., Thesis, Delft, 1955. 183. Kemball, C., Advances in Catalysis 2, 233 (1950). 184. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 451

(1952); B66, 67, 236, 415 (1953); B67, 92 (1954); B68, 61 (1955); Kruyer, S., Thesis, Delft, 1955. 186. Livingston, H. K., J. Colloid Sci. 4, 447 (1949). 186. Hill, T. L., Advances i n Catalysis 4, 219 (1952). 186a. de Boer, J. H., “The Dynamical Character of Adsorption,” p. 172. Oxford, New York, 1953. 187. Drain, L. E., Trans. Farada,) SOC.49, 650 (1953); Drain, L. E., and Morrison, J. A., ibid. 49, 654 (1953). 188. For a discussion of the factors governing this, see de Boer, J. H., “The Dynamical Character of Adsorption,” pp. 150-170. Oxford, New York, 1953. 189. de Boer, J. H., Advances i n Colloid Sci. 3, 43, 54 (1950). 190. Frenkel, J., 2. Physik 26, 117 (1924). 191. See de Boer, J. H., “The Daamical Character of Adsorption,” p. 34. Oxford, New York, 1953. 192. Kemball, C., Advances i n Catalysis 2, 233 (1950); Proc. Roy. SOC.A187, 73 (1946). 193. Kruyer, S., Koninkl. Ned. A/cad. Wetenschap. Proc. B68, 73 (1955). i94. See de Boer, J. H., “The Dynamical Character of Adsorption,” p. 3. Oxford, New York, 1953. 196. Cassel, H. M., and Neugebauer, K., J . Phys. Chem. 40, 523 (1936). 196. de Boer, J. H., “The Dynarnical Character of Adsorption,” pp. 107, 112, 115, 117. Oxford, New York, 1953. 197. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 451 (1952). 198. Pearce, J. N., and Taylor, A. L., J . Phys. Chem. 36, 1091 (1931). 199. Kemball, C., Proc. Roy. SOC.A187, 73 (1946). 200. Kruyer, S., Koninkl. Ned. Ahad. Wetenschap. Proc. B68, 73 (1955). 201. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B68, 61 (1955). 602. Johnson, M. C., and Vick, F. A,, Proc. Roy. SOC.A161, 308 (1935). 203. Eyring, H., J . Chem. Phys. 3, 107 (1935); Glasstone, S., Laidler, K. J., and Eyring, H., “The Theory of Rate Processes.” McGraw-Hill, New York, 1941. 604. de Boer, J. H., and van Steenis, J., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 578 and 587 (1952). 206. Riyanov, S., 2. Physik 71, 325 (1931); Lukirsky, P. I., and Riyanov, S., ibid. 76, 249 (1932); Lukirsky, P. I., I’hysik. 2.Sowjetunion 4, 225 (1933). 606. See de Boer, J. H., “Electron Emission and Adsorption Phenomena,” pp. 215 ff. Cambridge, New York, 1935. 207. Koller, L. R., Phys. Rev. 36, 1643 (1930). 208. Timofeew, P. W., and Nalimow, W. W., 2. Physik 81, 687 (1933). ,909. Wagner, C., 2. physik. Chem. B21, 25 (1933); B32, 447 (1936). 610. Mott, N. F., Trans. Faraday SOC.43, 429 (1947). 211. Rees, A. L. G., “Chemistry of the Defect Solid State.” Methuen, London, 1954. 216. Bevan, D. J. M., and Anderson, J. S., Discussions Faraday SOC.8, 238 (1950). 219. Burshteln, R. K., and Surovct, M. D., Doklady Akad. Nauk S.S.S.R.61,75 (1948);

156

J. H. D E BOER

Shumilova, N. A., and Burshtein, R. K., ibid. 61,475 (1948); Burshteln, R. K., Surova, M. D., and Zardenman, I. A,, Zhur. Fiz. Khim. 24, 214 (1950). $14. Fast, J. D., Philips Techn. Rev. 6, 369 (1941); 7, 74 (1942); see also Smithells, C. J., “Gases and Metals,” Chapman & Hall, London, 1937. 216. Coehn, A., and Specht, W., 2. Physik 62, 1 (1930); Coehn, A., and Jurgens, H., ibid. 71, 179 (1931); Francini, T. (1931) cited by Franck, J., Nachr. Ges. Wiss. Gottingen, p. 293 (1933). 216. de Boer, J. I€.,and Fast, J. D., Rec. trav. chim. 69, 161 (1940). 217. Ham, W. R., J . Chem. Phys. 1, 476 (1933). 218. de Boer, J. €I., and Fast, J. D., Rec. truer. chim. 68, 984 (1939). 219. Kalish, T. V., and Burstern, R. K., Doklady Akad. Nauk S.S.S.R. 81,1093 (1951). 220. Goodeve, C. F., and Jack, K. H., Discussions Faraday SOC.NO.4, 82 (1948). 221. Rhodin, T. N., J . Am. Chem. SOC.72, 5692 (1950). 22.9. Halsey, G. D., Advances in Catalysis 4, 259 (1952). 223. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 451 (1952). 2 2 3 ~ de . Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 67 (1953). dd3b. de Boer, J . H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 236 (1953). 223c. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B66, 415 (1953). 223d. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B67, 92 (1954). 223e. de Boer, J. H., and Kruyer, S., Koninkl. Ned. Akad. Wetenschap. Proc. B68, 61 (1955). 224. Goldmann, F., and Polanyi, M., 2. Physik. Chem. 132, 321 (1928). 226. Barrer, R. M., Proc. Roy. SOC.Al61, 476 (1937); see also Brunauer, S., “The Adsorption of Gases and Vapors,” pp. 246 ff. Oxford, New York, 1943. 226. Halsey, G. D., J . Chem. Phys. 16, 931 (1948); Sips, R. J., ibid. 16, 490 (1948); Schreiner, G. D. L., and Kemball, C., Trans. Faraday SOC.49, 292 (1953). 227. Crawford, V. A., and Tompkins, F. C., Trans. Faraday SOC.46, 504 (1950). 228. Drain, L. E., and Morrison, J. A., Trans. Faraday SOC.49, 654 (1953). 229. Honig, J. M., and Reyerson, L. H., J . Phys. Chem. 66, 140 (1952). 230. Morrison, J. A,, Los, J. M., and Drain, L. E., Trans. Faraday SOC.47, 1023 (1951); Drain, L. E., and Morrison, J. A., ibid. 48, 316, 840 (1952). $31. Pace, E. L., Dennis, K. S., Greene, S. A., and Heric, E. L., Can. J . Chem. 38, 245 (1955). 232. Chessick, J. J., Zettlemoyer, A. C., Healey, F. H., and Young, G. J., Can. J. Chem. 33, 251 (1955). 293. Orr, W. J. C., Proc. Roy. SOC.A173, 349 (1939). 234. Kemball, C., Advances in Catalysis 2, 233, 242 (1950). Z35. de Boer, J. H., 2. physik. Chem. B16, 300 (1932). 236. de Boer, J. H., and Custers, J. F. H., Z. physik. Chem. B21, 208 (1933). $37. Custers, J. F. H., and de Boer, J. H., Physica 3, 1021 (1936). 238. Hill, T. L., Advances in Catalysis 4, 212 (1952). 239. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapters V I I and VIII. Oxford, New York, 1953. 840. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapter VII, Sections 108 and 109. Oxford, New York, 1953. 241. Hill, T. L., Advances in Catalysis 4, 222 (1952).

ADSORPTION PHENOMENA

157

242. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapter VIII, Sections 118 and 119. Oxford, New York, 1953. ,845. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapter IX. Oxford, New York, 1953. 2&. Hill, T. L., Advances in Catalysis 4, 227 (1952). 246. Halsey, G. D., Advances in (htalysis 4, 259 (1952). 246. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapter VIII, Section 122. Oxford, New York, 1953. 247. Hill, T. L., J. Chem. Phys. 16, 181 (1948); Advances in Catalysis 4, 236 (1952). 248. Halsey, G. D., Advances in Catalysis 4, 265 (1952). 249. de Boer, J. H., Proc. Roy. Acad. Sci. (Amsterdam) 31, 906 (1928); de Boer, J. H., and Zwikker, C., 2. physik. Chem. B3, 407 (1929). 260. de Boer, J. H., and Custers, J. F. H., 2. physik. Chem. B26,238 (1934); Custers, J. F. H., and de Boer, J. H., Physica 3, 407 (1936). 261. Deryagin, B. V., and Zorin, Z. M., Doklady Akad. N a u k S.S.S.R. 98, 93 (1954). 262. de Boer, J. H., “The Dynamical Character of Adsorption,” Chapter IX, Section 129. Oxford, New York, 1953. 263. Miller, A. R., Discussims Faraday Soe. No. 8, 57 (1950); Beeck, O., ibid. No. 8, 118 (1950). 26.4. Beeck, O., Advances in Catalysis 2, 151 (1950). 266. Iiwan, T., Advances in Catalysis 6, 67 (1954). 866. Laidler, K. J., in “Catalysie” (P. H. Emmett, ed.), Vol. I, p. 75. Reinhold, New York, 1954. 267. Trapnell, B. M. W., “Chemisorption,” Chapter VI. Butterworths’, London, 1955. 258. Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 265. Reinhold, New York, 1955. 269. Beeck, O., and Wheeler, A,, Revs. Mod. Phys. 17, 61 (1945). 260. Rideal, E. K., and Trapnell, B. M. W., J. chim. phys. 47, 126 (1950); Trapnell, B. M. W., Proc. Roy. Sac. A206, 39 (1951). 261. Roberts, J. K., Proc. Roy. SUC.A162, 445 (1935). 262. Frankenburg, W., and Hodler, A., Trans. Faraday Sac. 28, 229 (1932); Frankenburg, W., J. Am. Chem. Sac. 66, 1827, 1838 (1944). 265. See also Frankenburg, W., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 222. Reinhold, New York, 1955. 264. Beeck, O., Advances in Catalysis 2, 174 (1954). 266. Eucken, A., Z. Elektrochem. 63,285 (1949), 64, 108 (1950). 266. Kwan, T., Advances in Catalysis 6, 89 (1954). 267. Schuit, G. C. A., and de Boer, J. H., Rec. trav. chim. 72, 909 (1953). 268. Beeck, O., Smith, A. E., and Wheeler, A., Proc. Roy. Sac. A177, 62 (1940). 869. Kington, G. L., and Holmes, J. L., Trans. Faraday SOC.49, 417-(1953). 270. Taylor, H. S., and Kistiakowsky, G. B., 2. physik. Chem. 126, 341 (1927). 871. Beebe, R. A., and Taylor, H. S., J. Am. Chem. Sac. 46, 43 (1924). 272. Trapnell, B. M. W., “Cheniisorption,” p. 164. Butterworths’, London, 1955;

Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 298. Reinhold, New York, 1955. $75. Taylor, H. S., Proc. Roy. Soc. A108, 105 (1925). 274. de Boer, J . H., Advances in Colloid Sci. 3, 1 (1950). 276. Miiller, E. W., 2. Physik 106, 541 (1937). 276a. Miiller, E. W., Z. Physik 108,668 (1938). b76b. Miiller, E. W.,2.Physik 126, 642 (1949).

158

J. H. DE BOER

877. Becker, J. A., and Hartman, C. D., J . Phys. Chem. 67, 154 (1953). 278. Stranski, J. M., and Suhrmann, R., Ann. Physik. 16) 6,153 (1947); Drechsler, M., Z . Elektrochem. 68,334 (1954). 279. Comer, R., J. Chem. Phys. 21, 293 (1953). 280. Drechsler, M., and Muller, E. W., Z. Physik 134, 208 (1953). 281. Nichols, M. €I., Phys. Rev. 67,297 (1940). 282. Benjamins, M., and Jenkins, R. O., PTOC. Roy. SOC. A176, 262 (1940). 285. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” p. 21. Cambridge, New York, 1935. 284. Sachtler, W. M. H., Z. Elektrochem. 69,119 (1955). 285. Brucke, E., 2. Physik 98, 77 (1936). 286. Schenk, D., 2. Physik 98,753 (1936). 287. Burgers, W. G., and Ploos van Amstel, J . J. A., Physica 4, 5, 15 (1937); 6, 305. 313 (1938). 288. Balandin, A . A., 2.Physik. Chem. 126, 267 (1927); B2, 289 (1928). 289. Twigg, G. H., and Rideal, E. K., Proc. Roy. SOC. A171,55 (1939); Trans. Faraday SOC.36, 533 (1940); Proc. Roy. SOC.A177, 62 (1940). 290. Johnson, R.P., and Shokley, W., Phys. Rev. 49, 436 (1936). 291. Martin, S. T., Phys. Rev. 66,947 (1939). 292. Benjamins, M., and Jenkins, R. O., Proc. Roy. SOC.A176, 262 (1940); A180, 225 (1942). 293. Muller, E. W., 2. Physik 106,541 (1937). 294. Vol’kenshteIn, F. F., Zhur. Fiz. Khirn. 22, 311 (1948), 23, 917 (1949).

S., and Field, E., on the Contributions of Russian Scientists in Advances in Catalysis 6, 249 (1953). 296. Taylor, H. S., Discussions Faraday SOC.No. 8, 9 (1950). 297. ItoginskiT, S. Z., Zhur. Priktad. Khim. 17,3, 97 (1944); Roginskil, S.Z., “Adsorption and Catalysis on Non-Uniform Surfaces.” Acad. Sci. U.S.S.R., 1948. 698. Beeek, O., Smith, A. E., and Wheeler, A., Proc. Roy. SOC.A177, 62 (1940). 299. Beeck, O., Advances in Catalysis 2, 174 (1950). 300. Sachtler, W. H. M., Dorgelo, G., and van der Knaap, W., J. Chim. Phys. 61,491 (1954). 501. Gwathmey, A. T., and Leidheiser, H., J . Am. Chem. SOC.70, 1200 (1948); Leidheiser, H., and Meelheim, R., ibid. 71, 1122 (1949); Leidheiser, H., and Gwathmey, A. T., ibid. 70, 1206 (1948). 302. Wagner, J. B., and Gwathmey, A. T., J. Am. Chem. SOC.76,390 (1954); Cunningham, R. E., and Gwathmey, A. T., ibid. 76, 391 (1954). 303. Trapnell, B. M. W., “Chemisorption,” p. 166. Butterworths’, London, 1955. $04. Pease, R. N., and Stewart, L., J. Am. Chem. SOC.47, 1235 (1925). 305. Wheeler, A;, Advances in Catalysis 3, 308 (1951). 306. Taylor, H. S., and Liang, S. C., J. Am. Chem. SOC.69,1306, 2989 (1947) ; Taylor, H. S., and Sadek, H., ibid. 72, 1168 (1950). 307. Rideal, E. K., and Trapnell, B. M. W., J. chim. phys. 47, 126 (1950). 508. Roginskil, S. Z., and Todes, 0. M., Acla Physicochim. U.S.S.R. 21, 519 (1946). 309. Keier, M. P., and RoginskiI, S. Z.,Zzvest. Akad. Nauk S.S.S.R. (1950); see the article of Tolpin, J. G., John, G. S., and Field, E., in Advances in Catalysis 6, 248 (1953). 510. Kefer, M. P., and Roginskil, S. Z., Doklady Akad. Nauk S.S.S.R. 67, 151 (1047). 311. Kummer, J. T., and Emmett, P. H., J. Am. Chem. SOC.73, 2886 (1951). 31%. Eischcns, R.P., J . Am. Chem. SOC.74,6167 (1952). 313. Emmett, P. H., and Kummer, J. T., J . chim. phys. 47, 67 (1950). 696. See also the article by Tolpin, J. G., John, G.

ADSORPTION PHENOMENA

159

314. Schuit, G. C. A., Proc. Intern. Symposium on the Reacfivity of Solids, Gothenburg, 1952, p. 571 (1954). 315. Kwan, T., Advances in Catai‘ysis6, 95 (1954). 316. Langmuir, I., and Kingdon, K. H., Proc. Roy. SOC.A107, 61, 76 (1925). 317. Taylor, J. B., and Langmuir, I., Phys. Rev. 44, 432 (1933). 318. Langmuir, I., J. Am. Chem. SOC.64, 2798 (1932); Nobel Lecture, Angew. Chem. 46, 728 (1933); Chem. Revs. 13, 147 (1933). 319. de Boer, J . H., and Veenenmns, C. F., Physica 1, 953 (1934); de Boer, J. H., “Electron Emission and Adsorption Phenomena,” pp. 80 ff. Cambridge, New York, 1935. 320. de Boer, J. H., and Veenemans, C. F., Physica 1, 960 (1934). 321. de Boer, J. H., “Electron Elmission and Adsorption Phenomena,” p. 52. Cambridge, New York, 1935; see especially German edition, p. 40 (1937). 383. Becker, J. A., Phys. Rev. 28, 357 (1926); Trans. Am. Electrochem. SOC.66, 169 (1929). 383. de Boer, J. H., “Electron Emission and Adsorption Phenomena,’, p. 84. Cambridge, New York, 1935; see especially German edition, p. 62 (1937). 324. de Boer, J. H., “Electron Emission and Adsorption Phenomena,” p. 85. Cambridge, New York, 1935. 986. Roberts, J. K., “Some Problems in Adsorption.” Cambridge, New York, 1939. 386. Taylor, J. B., and Langmuir, I., Phys. Rev. 44, 423, 428, (1933). 327. de Boer, J . H., and Veenemans, C. F., Physica 2, 521 (1935). 328. de Boer, J. H., “Electron Emission and Adsorption Phenomena,,’ Chapter 111. Cambridge, New York, 1935. 389. Mayer, H., Phil. Mag. [71 16, 594 (1933). 330. de Boer, J. H., and Veenemans, C. F., Physica 2, 529 (1935). 391. de Boer, J. H., Discussions Faraday Soc. No. 8, 94, 208 (1950); Chem. Weekblad 47, 416 (1951). 338. Boudart, M., J. Am. Chem. SOC.72, 3556 (1952). 333. Mignolet, J. C. P., J . Chem. Phys. 23, 753 (1955). 334. Boudart, M., J. Chem. Phys. 23, 753 (1955). 335. Gomer, R., J. Chem. Phys. 21, 1869 (1953). 336. Mignolet, J. C. P., Rec. trav. chim. 74, 685 (1955). 937. Thomas, L. B., and SchofieIcl, E. B., J . Chem. Phys. 23, 861 (1955). 338. Trapnell, B. M. W., “Chemisorption,” p. 54. Butterworths’, London, 1955. $39. de Boer, J. H., Koninkl. Ned. Akad. Wetenschap. Proc. 49, 1103 (1946). 340. Trapnell, B. M. W., Trans. Paraday Soc. 61, 368 (1955). 341. de Boer, J . H., and Kraak, 11. H., Rec. trav. chim. 66, 941 (1936). 34.9. Shishakov, N. A., Exptl. and Theoret. Phys. (U.S.S.R.) 22, 241 (1952). 343. Mignolet, J. C. P., Rec. trav. chim. 74, 685 (1955). 344. Boudart, M., J . Am. Chem. SOC.74, 3556 (1952). 545. Schuit, G. C. A., and de Bow, N. H., Rec. Irav. chim. 72, 909 (1953). 346. Schuit, G. C. A., Proc. Intern. Symposium on the Reactivity of Solids, Gothenburg, p. 571 (1952). 347. Beeck, O., Advances in Catalysis 2, 185 (1950). 348. Eley, D. D., J. Phys. Chem. 66, 1017 (1951). $49. Schwab, G. M., Trans. Faraday SOC.42, 689 (1946). 360. Mignolet, J. C. P., J . Chem. Phys. 23, 753 (1955). 351. Temkin, M. I., “Symposium on Problems of Chemical Kinetics, Catalysis and Reactivity,” Akad. Nauk.,S.S.S.R., 1955. 361a. Federova, A. J., and Frumkin, A. N., J . Phys. Chem. (Russ.) 27, 247 (1953).

160

J. H. DE BOER

362. Schuit, G. C. A., and de Boer, N. H., Rec. trav. chim. 72, 909 (1953). 363. Boudart, M., J. Am. Chem. SOC.74, 3556 (1952). 364. McQuillan, A. D., Proc. Roy. Sac. A204, 309 (1950). 366. Bagg, J., and Tompkins, F. C., Trans. Faraday SOC.61, 1071 (1955). 366. de Boer, J. H., and Dippel, C. J., 2. physik. Chem. B26, 399 (1934). 367. Beeck, O., Advances in Catalysis 2, 177 (1950). 358. Maxted, E. B., and Hassid, N. J., J. Chern. SOC.30, 3313 (1931). 369. Kwan, T., Advances i n Catalysis 6, 92 (1954). 360. Zwietering, P., and Roukens, J., Trans. Faraday SOC.60, 178 (1954). 361. Berl, E., and Weingaertner, E., Z. physik. Chem. A173,35 (1935); Damkbhler, G.,

ibid. A174, 222 (1935). Kemball, C., Proc. Roy. SOC.A201, 377 (1950). Bagg, J., and Tompkins, F. C., Trans. Faraday SOC.61, 1071 (1955). Zeldowitsch, J., Acta Physicochim. U.S.S.R. 1, 449 (1934). Langmuir, I., J. Am. Chem. SOC.64, 2798 (1932). Frumkin, A. N., and Slygin, A. J., Acta Physicochim. U.S.S.R. 3, 791 (1935). Temkin, M. I., and Pyzhev, V., Acta Physicochim. U.S.S.R. 12, 327 (1940); Temkin, M. I., and Kiperman, S. L., J. Phys. Chem. U.S.S.R. 14, 1250 (1940); Temkin, M. I., ibid. 24, 1312 (1950); Kiperman, S. L., and Temkin, M. I., Acta Physicochim. U.S.S.R. 21, 267 (1926); Kiperman, S. L., and Temkin, M. I., J. Phys. Chem.. U.S.S.R. 20, 369 (1946); Kiperman, S. L., ibid. 21, 1435 (1947); Kiperman, S. L., and Granovskaya, V., ibid. 26, 557 (1951); see also Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, pp. 265, 318. Reinhold, New York, 1955. 368. Trapnell, B. M . W., ‘’ Chemisorption,” pp. 124. Butterworths’, London, 1955; Bokhoven, C., van Heerden, C., West.rik, R., and Zwietering, P., in “Catalysis” I11 (P. H. Emmett, ed.), Vol. 3, p. 291. Reinhold, New York, 1955. 369. Trapnell, B. M. W., “Chemisorption,” pp. 126 ff. Butterworths’, London, 1955. 370. RoginskiI, S. Z., Proc. Acad. Sci. U.R.S.R. 46, 66, 206 (1944). 37i. Temkin, M. I., J. Phys. Chem. U.S.S.R. 16, 296 (1941). 3’7% Brunauer, S., Love, K. S., and Keenan, R. G., J. Am. Chem. SOC.64,751 (1942). 373. Zeldowitsch, J., Acta Physicochiin. U.S.S.R. 1, 961 (1934). 3’74. Halsey, G. D., and Taylor, H. S., J. Chem. Phys. 16, 624 (1947); Advances i n Catalysis 4, 259 (1952). 376. Trapnell, B. M . W., “ Chemisorption,” Chapter V. Butterworths’, London, 1955. 3’76. Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 265. Reinhold, New York, 1955. 3’77. Beeck, O., Ritchie, A. W., and Wheeler, A., 3. Colloid Sci. 3,505 (1948); see also Beeck, O., Advances i n Catalysis 2, 151 (1950). 3’78. Trapnell, B. M. W., Proc. Roy. SOC.A206, 39 (1951). 3’79. Porter, A. S., and Tompkins, F. C., Proc. Roy. SOC.A217, 529 (1953). 380. de Boer, J. H., 2. physik. Chem. B14, 149 (1931). 381. de Boer, J . H., and Dippel, C. J., Z. physik. Chem. B21, 198 (1933). 382. Dippel, C. J., and de Boer, J. H., Rec. trav. ckim. 67, 277 (1938), 67, 1087 (1938). 383. Taylor, K. S., and Liang, S. C., J . Am. Chem, SOC.69, 1306, 2989 (1947); Taylor, H. S., and Sadek, H., ibid. 72, 1168 (1950). 384. Beeck, O., Advances i n Catalysis 2, 182 (1950). 386. Boudart, M., J . Am. Chem. SOC.74, 3556 (1952). 386. Vol’kenshteIn, F. F., J. Phys. Chem. U.S.S.R. 22, 311 (1948); 29, 917 (1949); Advances in Catalysis 6, 249 (1953). 366. 363. 364. 366. 366. 367.

ADSORPTION PHENOMENA

lGl

387. Maxted, E. B., Advances in Catalysis 3, 129 (1951). 388. Harkins, W. D., and Gans, R., J. Am. Chem. SOC.63, 2804 (1931). 389. Houben, G. M. M., Thesis, p. 87. Delft, 1951. 390. Fortuin, J. M. H., Thesis, Chapter VI B. Delft, 1955. 391. Bond, G. C., and Sheridan, J., Trans. Faraday SOC.48, 651, 658, 664 (1952). 392. Bokhoven, C., Proc. 2nd Radio Isotope Conf. Oxford, p. 53 (1954). 393. Beeck, O., Advances in Catalysis 2, 181 (1950). 394. Boudart, M., J. Am. Chem. SOC.74, 3556 (1952). 396. Bagg, J., and Tompkins, F. C., Trans. Faraday SOC.61, 1071 (1955). 396. Wheeler, A., Advances in Catalysis 3, 250 (1951). 397. Roginskil, S. Z., Doklady Akad. N a u k S.S.S.R. 87, 1013 (1952). 398. Cf. a survey of the work of RoginskiI and collaborators by Jabrova, G. M., J . chim. phys. 61, 769 (1954). 399. de Boer, J. H., and Fast, J. D., Rec. trav. chim. 68, 984 (1939). 400. de Boer, J. H., “Electron Ihission and Adsorption Phenomena,” Cambridge, New York, 1935. 401. Benton, A. F., Trans. Faraday SOC.28, 209 (1932). 402. Fianda, F., and Lange, E., Z. Elektrochem. 66, 237 (1951). 403. Boudart, M., J . Am. Chem. Soe. 74, 3556 (1952). 404. Fast, J. D., private communicaton. 406. Bokhoven, C., van Heerden, C., Westrik, R., and Zwietering, P., in “Catalysis” (P. H. Emmett, ed.), Vol. 3, p. 312. Reinhold, New York, 1955. 406. Rideal, E. K., and Trapnell, B. M. W., Discussions Faraday Soc. No. 8, 114 (1950). 407. Boudart, M., J . Am. Chem. Soc. 74, 3556 (1952).

This Page Intentionally Left Blank

Activation of Molecular Hydrogen by Homogeneous Catalysts S. W. WELLER AND G. A. MILLS HOUdTg

Process Corporation, Linwood, Pennsylvania Page

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Homogeneous Catalytic Systems.. 1. Cuprous and Silver Salts in Org a. Cuprous Salts in Quinoline., . . . . . . . . . . . . . . . ( 1 ) Reduction of quinone.. . . . . . . . . . . . . . . . . (2) Reduction of cupric acetate.. ................................ (3) Reduction and exchange by use of deuterium.. . . . . . . . . . . . . . . . . (4) Ortho- parahydrogen interconversion. ........................ ( 5 ) Effect of anion, solvent, complexing agents.. . . . . . . . . . . . . . . . . . .

168 171 171 172

. . . . . . . . . . . . 180

. . . . . . . . . . . . . . . . . 181

.........................................

181

. . . . . . . . . . . . . . . . . 199 c. Cobalt Cyanide.. . . . . . . . . . . . . . .

References.. . . . . . . . . . . . . . . . . . . . . . . . . .

......................

204

I. INTRODUCTION A few solutions have been found t o possess the unusual property of activating molecular hydrogen so as to bring about reduction of organic or inorganic compounds or to came H z - D ~ exchange or ortho-para-hydrogen interconversion. Such catalytic systems have been termed homo163

164

S. W. WELLER AND G . A . MILLS

geneous t o emphasize their physical state and to distinguish them from usual solid hydrogenation catalysts. Catalytic hydrogenation in homogeneous liquid solutions is significant at the present time chiefly because of the unique insight provided into the mechanism of catalytic activation of molecular hydrogen. That is to say, it is the unusual opportunity to understand the chemistry of catalysis afforded by such systems, rather than the discovery or performance of new or difficult reductions, which makes such homogeneous catalysts of interest. As mentioned above, hydrogenation catalysts, almost without exception, are solids. The nature of solids has been little understood, however, and the properties of the surfaces of solids have remained even more unknown. By contrast, the nature of molecular species which exist in solution has been relatively well established. There is much hope, therefore, t ha t catalytic hydrogenation in homogeneous systems can be related t o molecules, ions, and complexes of known chemical properties. For the most part, catalysis has developed faster as a n art than a science. However, there has been a growing recognition that groups of seemingly unrelated catalysts could be placed in general classes. Two broad classifications have emerged : one, an acid (base) group-including as acids such materials as HC1-A1C13, H20-Si02-A1203, H2S04,etc.and a second group, typified by metals and transition metal oxides, characterized by “an electron defect structure” and capable of activating hydrogen and oxygen. Fortunately, the fundamental behavior of acid (base)-type catalysts has been rather well elucidated through the carbonium (carbanion) ion mechanism theory and so such catalytic reactions proceed in a generally predictable manner. No such general mechanism, however, has been devised for the catalytic activation of hydrogen, and this has remained an important problem. For this type of catalyst there has been a gradual development of ideas, progressing from the concept of a n active site at the surface due to unusual but vague forces as a consequence of geometric corners and edges, through theories of unfilled 3d electronic levels and Brillouin zones, to consideration of semiconductivity as fundamental to this type of catalysis. It seems unlikely, however, th a t catalysis will be explained by simple physical measurements, for example by conductivity measurements, just as a detailed knowledge of the electronic levels of hydrogen would not provide for the deduction of the chemical properties of hydrogen given on the pages of a freshman chemistry text. On the other hand, it seems logical to insist th a t catalysis is essentially a chemical phenomenon. The high selectivity for acceleration of certain chemical reactions emphasizes the necessity of such a conclusion. The functioning of homogeneous catalysts involves the properties of coordination compounds. Metal porphyrin complexes such as exist in

ACTNATION OF MOLECULAR HYDROGEN

165

hemoglobin and chlorophyll have long been recognized as important in biooxidation and reduction. Such complex structures have exactly balanced oxidation-reduction potentials and, in many instances, exactly specified geometric factors. Homogeneous catalysts capable of activating molecular hydrogen are relatively rare. In the last few years a renewed interest, in such homogeneous catalysts has led to the discovery of several new systems. It appears, therefore, that there may be rriany more homogeneous catalysts found. A t this stage of development of the field it is appropriate t o consider in turn each of the known homogeneous catalyst syst,ems. These are (1) cuprous and silver Salk in organic bases, (2) cupric and mercuric salts in aqueous medium, (3) cobalt carbonyl, (4) platinous ethylene chloride, (5) base catalysts, and ( 6 ) miscellaneous.

11. HOMOGENEOUS CATALYTIC SYSTEMS 1. Cuprous and Silver Salts in Organic Bases

a. Cuprous Salts in Quinoline. I n 1938 Calvin (1) discovered th a t cuprous acetate or salicylaldehyde dissolved in quinoline catalyzed the reduction by molecular hydrogen of (p-benzo-) quinone or of cupric ion. The amount of hydrogen taken up corresponds quite well with that expected stoichiometrically for the reduction of the quinone to semiquinone or cupric t o cuprous ion. The solution remains optically clear over the course of these reductions. The rate of reduction is rapid a t 100°, being complete in about 1 hr., with convenient pressures and solution concentrations. Following the reductions of the quinone or cupric to cuprous ion, hydrogen absorptiori either stops or proceeds a t a very much slower rate depending on the conditions. Under those circumstanres where there is a slow uptake of hydrogen, there is then a deposition of finely divided metallic copper from what was previously a clear solution. The amount of metallic copper corresponds to the amount of hydrogen absorbed a t the slow rate, that, is, after the “break point” in rate of uptake of hydrogen. WiImarth and co-workers (2-4) and Weller, Mills, and Wright (6,6) have investigated this system and related systems intensively. Their work serves t o confirm the original work of Calvin and to define the catalytic mechanism in more detail. Kinetic measurenients have served primarily t o elucidate the mechanism. These investigators are in agreement th a t for the reduction of either quinone or cupric ion, the rate of uptake of hydrogen is given by the expression

- -dH2 - - k.PH,.(conc. of at

cuprous dimer)

166

S. W. WELLER AND G . A. MILLS

Thus the rate of reduction is independent of the substrate concentration.

It is very dependent on the environment surrounding the copper, such as solvent, anions, and complexing compounds. (1) Reduction of quinone: The homogeneous catalytic hydrogenation of quinone was first reported by Calvin in 1938 and later studied in more detail (.2,6,7). The course of a typical hydrogenation of quinone (p-benxoquinone), catalyzed by cuprous acetate in quinoline solution, is shown in Fig. 1. As indicated here, the quantity of hydrogen absorbed increases linearly with time until a point is reached where a change in slope occurs. The 60

0 20

40 60 Time, minutes.

80

100

Fxa. I. Hydrogenation of quinone (6).

solution is optically clear up to the break point, but past this point metallic copper is precipitated. The color of the solution is deep red throughout the hydrogenation. The amount of metallic copper found is usually stoichiometrically equivalent to the amount of hydrogen absorbed past the break point. If the run is terminated prior t o the break point, no copper is found on filtration. The amount of hydrogen absorbed u p to the break point is approximately that necessary for reduction of quinone t o the semiquinone (or quinhydrone) stage. It appears th a t the hydrogen absorption up to the break point is associated with a reduction of quinone catalyzed by dissolved cuprous acetate; hydrogen absorption past the break point corresponds t o reduction of cuprous acetate t o metallic copper. The rate of quinone hydrogenation depends in a nonlinear way on the concentration of cuprous acetate. Calvin carried out a brief kinetic study in which he noted that the rate of reaction was somewhat more than pro-

ACTIVATION OF MOLECULAR HYDROGEN

167

portional to the amount of cuprous acetate in solution. From this Calvin suggested that it may be a dimer of copper that is active. Figure 2 presents the results of a much more detailed kinetic study. I n this figure each individuai point represents a separate experiment. If the rate of

0

0.04 0.08 0.12 Cuprous acetate, moles/l.

0.16

FIG.2. Rate of quinone hydrogenation vs. cuprous acetate concentration (6).

hydrogen uptake were proportional to the concentration of a cuprous acetate dimer, the following equations would obtain a t constant hydrogen pressure : Rate = ~ [ ( C U ' ) ~ ] (2) and I( = /(CUl)Z]/[CUI]* (3) where k is a specific rate constant including the effect of hydrogen pressure, [(CU*)~] represents the concentration of dimer, (Cu') is the concentration of monomer and K is the dimerization constant. If the total concentration of cuprous acetate is expressed as C, where

c = (CU')

+ 2[(CU')Z]

(4)

then the rate of hydrogenation is given by

I n the low concentration range, where dissociation of the dimer is almost complete, the rate expression becomes Rate = kKC2

(6)

I n the high concentration range, where dimerization is almost complete, it becomes Rate = +$kC (7)

168

S. 7.7'. WELLER AND G. A. MILLS

The variation of rate with concentration should thus change from a square dependence a t very low concentrations to a linear dependence a t high concentrations. The continuous curve shown in Fig. 2 was calculated from Eq. (5) by use of the following values for the two parameters: K = 11.2 mole-' liters, k = 41.9 ml. (at 515 mm.) min.-' moie-I. This dimerizatioti constant is in agreement with the approximate value established by separate ebulliometric experiments. It is clear that within experimental error the data can be fit by an expression of the form of Eq. ( 5 ) ;i.e., the data are consistent with the hypobhesis that the active catalyst is a dimer of cuprous acetate. The activation energy for quinone hydrogenation was found t o be 14.3 (7) and 15.3 (6) kcal./mole in the range 78"-118". (2) Reduction of cupric acetate: An unreduced solution of cupric acetate in quinoline is a dark green color; the fully reduced solution is a clear ruby red. When quinone has been reduced, as described above, further treatment with hydrogen produces metallic copper. In contrast, when cupric ion has all been reduced to cuprous ion, no further reduction occurs within convenient experimental times, and the solution remains clear. The reduced cupric solution can be reoxidized rapidly by oxygen at room temperature. The rate of reduction of cupric ion is illustrated in Fig. 3 ( 1 ) . Curve 21 represents the hydrogen pressure drop observed in a typical experiment. Curve 23 is a repetition of this reduction, starting, however, with the reduced solution of evperimeiit 21, which was oxidized by molecular oxygen from a red reduced solution to a green oxidized one. The final slopes (reduction rate) are the same for 21 and 23, the only difference being t ha t 23 has a much longer induction period. If the reaction is stopped just before the fiat portion of the curve (21 and 23) is reached and the solution filtered through a fine filter paper, no precipitate is found. The shape of the reduction curve indicates that the reduction of cupric ion is autocatalytic. This conclusion is confirmed by the almost complete disappearance of the induction period when cuprous acetate is added, strikingly illustrated by a comparison of curves I and I1 in Fig. 4 (6). I n both cases the quantity of hydrogen absorbed is close to the theoretical value for the reduction of Cu++ t o Cu+. There is evidence that the presence of aniline in the quinoline is effective in providing some cuprous ion from cupric to initiate sufficient catalyst to start the reaction. The individual points in Fig. 5 represent instantaneous rates, picked off the curves for nine separate experiments in which cupric acetate or oxyacetate was reduced, as a function of the total amount of cuprous acetate present a t the corresponding time. The total cuprous acetate was th e

169

ACTIVATION OF MOLECULAR HYDROGEN

FIG.3. Curve 21, 0.4 g. (2 mmoles) cupric acetate (H20) in 50 co. quinoline; temperature 105". Curve 23, curve 21 after reoxidation with 02 at room temperature (1).

30

P5

20

v

-ij

f

10

Cd

5 0 100

200 Time, min.

300

400

FIG 4. Hydrogenation of cupric acetate: effect of cuprous acetate (6).

170

S. W. WELLER AND G. A. MILLS

sum of the cuprous acetate initially added (if any) and that produced by reduction of the cupric salt. The solid curve in Fig. 5 is a curve drawn through the individual points of Fig. 2; it represents the observed variation of the rate of quinone hydrogenation with concentration of cuprous

..

w

0

0.04

0.08

0.16

0.12

Total cuprous acetate, moles/l. FIG.5. Reduction of Cu"-quinoline

solution (6).

acetate. It is a striking fact that, within a range of 10-20%, the rates of quinone hydrogenation and Cu" hydrogenation are identical for a given cuprous acetate concentration. This result strongly suggests that the same step is rate determining for both reactions.

' ,

80

h

2

60

v

4 40 ,f

- 4 MlLLlMOLES Cu'Ace ' Hz.0 2 MlLLlMOLES Cu AG I

20

40

I

60 Time, min.

I

I

__

I

I l l 80

100

120

FIG.6. Hydrogenation of cupric acetate: comparison of theory with experiment (6).

On this basis one can deduce the course of cupric acetate hydrogenation during any single experiment,, the rate again being assumed t o be proportional t o the concentration of ouprous acetate dimer. Figure 6 shows the agreement obtained between the theoretical curve (solid curve) and the observed data (circles) for a n experiment in which the hydro-

ACTIVATION OF MOLECULAR HYDROGEN

171

genation of 4 mmoles of cupric acetate was carried out in the presence of 2 mmoles of cuprous acetate. The fit is excellent u p to the point of exhaustion of the cupric acetate (break point). It is clear that, as far as the kinetics are concerned, the catalyzed hydrogenations of both quinone and cupric acetate are consistent with the same mechanism. (3) Reduction and exchange by use of deuterium: Calvin ( I ) made a preliminary experiment in which deuterium was used to reduce quinone in the cuprous acetate-quinoline system. The rate was 40% slower than with Hz. Weller and Mills (5) found that the rate of quinone reduction by D, was within 5-10% of the rate observed with hydrogen. They also investigated in detail the exchange of deuterium. It was found that when D2 was contacted with a quinoline solution of cuprous acetate, HD and H, appeared in the gas phase. Some hydrogen donor obviously was present in the solution. It was established th at (1) when a reducible substrate is present, verylittle exchange occurs prior to the time of complete reduction of the substrate, but extensive exchange occurs after this time; (2) occurrence of exchange does not depend on the initial presence of either quinone or cupric acetate since quinoline solutions of cuprous acetate alone can bring about exchange; (3) neither quinoline, acetic acid in quinoline, dry cuprous acetate, nor dry cuprous acetate-cupric acetate mixture is able t o catalyze the exchange; (4) freshly reduced copper powder does not cause the exchange; and (5) the exchange is homogeneously catalyzed. Furthermore, only a limited amount of hydrogen exchanged from the solution even after an extended time. Aniline, present in small amounts as an impurity in the quinoline, was indicated t o be the hydrogen donor. The rate of deuterium-hydrogen donor exchange is comparable both with the rate of quinone reduction (at the same cuprous acetate concentration) and with the rate of ortho-para-hydrogen interconversion (3). (4) Ortho- parahydrogen interconversion: Calvin ( 1 ) found th a t 75 "/o parahydrogen reduced quinone a t the same rate as a normal hydrogen. He also reported that there had been no conversion of the parahydrogen in the gas examined just after the break in the absorption curve. A similar experiment in which cupric acetate alone was reduced gave no conversion of parahydrogen in the gas examined just a t the point of completion of the Cu" to Cut reaction. If, however, the reduced solution was allowed t o cool to room temperature and the shaking of the parahydrogen continued, the conversion was complete in less than 10 hr. Wilmarth and Barsh (3) studied in detail the conversion of 0-p H, by solutions of cuprous acetate in quinoline with the purpose of relating the conversion process t o the hydrogenation process. Over the tempera-

172

8 . W. WELLER AND G. A. MILLS

ture range of fiOo-lOOo the conversion rate showed a second-order dependence on the total concentration of cuprous acetate. The data for 100' are shown in Fig, 7, which shows that the product of the experimental first-order rate constant k" and the cuprous acetate concentration C varies linearly with the square of the cuprous acetate concentration.

FIG.7. Effect of concentration on the conversion rate constants a t 100" (3).

Over the range of temperature and concentration studied, the rate of conversion by cuprous acetate in quinoline is expressible by the equation

(- 16000/RT)(1. mole-' min.-1)2(CuOAc)2(Hz) (8) Wilmarth and Barsh concluded that the most reasonable mechanism involves the assumptions t h a t (1) the cuprous acetate is involved in a rapid monomer-dimer equilibrium; (2) the dimer is the active species for conversion; and (3) the concentration of the dimer is always small with respect t o that of the monomer. Wilmarth and Barsh also compared the rate of reduction of cupric ion with that of 0-p H2 conversion. They observed a close correspondence between rate of reduction of cupric ion and a-p Hz conversion, catalyzed by cuprous acetate, which was taken as experimental verification that the same activated intermediate is formed in both reactions. (6) E f e c t of anion, solvent, complexing agents: Salts of cuprous ion other than acetate have been tried as catalysts. Cuprous salicylaldehyde is effective in activating hydrogen at, about the same rate as is cuprous acetate. Cuprous salicylaldimine, salicylaldehyde-urea, salicylaldehyde-

ACTIVATION OF MOLECULAR HYDROGEN

173

ethylene diamine and o-phenylene diamine, and aceto acetic ester were found by Calvin ( 1 ) not to be catalysts. Cupric chloride is reduced only at a low rate, even at 135". The effect of twenty different solvents has been investigated by Weller and Mills (6) for the reduction of quinone or cupric acetate. Methyl quinolines, pyridine, and pyridine derivatives, as well as a number of amines and oxygenated compounds, were tried. It was found that, within the range of solvents studied, it seems to be a sufficient condition for the activation of hydrogen that the solvent be a nitrogen base, not necessarily heterocyclic, and free of complicating features, such as unusual chemical reactivity, steric factors, or exceptionally strong chelating tendency. There are some indications (4-methyl quinoline, isoquinoline, 4-methyl pyridine) that an increase in the base strength of the solvent results in increased catalytic activity, but the data are insufficient to permit a definitive conclusion. The effect of additives in altering the catalyst environment has been studied by Wright and Weller (6). Chelation is known to alter the oxidation-reduction properties of ions and therefore could be expected to alter the activation of hydrogen. Ethylene diamine and ethylene diamine tetraacetic acid inhibit both the rate and extent of reduction of quinone and cupric acetate monohydrate in quinoline solution at 100". The production of metallic copper is related to the presence of organic material. Reduction of CuII in dodecylamine and pyridine will be discussed later. (6) Discussion: The mechanism of the catalytic activation of hydrogen depicted in the following equations was first suggested by Calvin (7) in 1939: Cul Q Cu'Q (fast) (9) 2Cu'Q (CU'Q)~ (fast) (10) (Cur&)* Hz (Cu1Q)2.H2 (slow) (11) (CUIQ)~.H~2Su + (CuIQ), 2HSu (fast) (12)

+

+

+

+

where Q = quinone and Su = reducible substrate, quinone or Cu". Cuprous acetate monomer, complexed with the quinoline solvent, is in rapid equilibrium with dimer. The equilibrium constant is such that dimer formation is incomplete. Activation of the hydrogen occurs by a slow reaction between dimer complex and dissolved molecular hydrogen. Following activation of the hydrogen, the substrate quickly reacts with the hydrogen. Reaction (1 1) is believed rate controlling. Weller and Mills (6) attempted to establish whether the reaction went through a two-step oxidation and reduction of the Cu' catalyst; however, the conclusion was that the reaction depicted above best fits the observed facts.

174

8. W. WELLER AND G . A. MILLS

The Hz-D2exchange occurs to an appreciable extent only under conditions when Hg (or D2) is not being absorbed rapidly in a chemical reaction such as the reduction of quinone or cupric acetate. This is consistent with the idea that the rate-determining step in the reduction is the activation of hydrogen by cuprous acetate [reaction (ll)],the activated hydrogen being rapidly removed by reaction with quinone or CuJrbefore it can return to the gas phase. The fact that the exchange rate is comparable with that of the reduction means that the exchange is not a slow, secondary reaction. Here also reaction (11) is presumably the rate-determining step, and the exchange step proper may be schematically written as

+

( C U ' ) ~ . D ~RH (in soln.)

(Cu1)2.HD

+ RD (in soln.)

(13)

The exchange and hydrogenation reactions are competitive, in the case that a reducible substrate is present. It is not possible to determine, from these data, the relative rates of (12) and (13); both are faster than reaction (11). It should be noted that the formation of HD from a mixture of H q and D q does not require a direct interchange. Collisions between activated D, and HZare not likely. Instead, it would appear probable that Dz is activated and exchanges with a hydrogen donor according t o Eq. (13) and that subsequently Hz is activated and exchanges with the deuterated donor. The occurrence of the exchange reaction is of great importance. Since it is to be expected that the difficult step in the over-all reaction is the dissociation of the stable hydrogen molecule, it is probable that the hydrogen is dissociated on activation and reaction (11) is more properly to be written in the form (CU')~ Hq (Cu1)2.2H (14)

+

the symbol on the right signifying that the hydrogen is dissociated on reaction with cuprous acetate. No information is available t o determine whether the dissociation is symmetrical or whether it results in the formation of a positive and negative ion. The former seems t o be the simpler and more reasonable assumption. I n view of the fact that breaking a hydrogen-hydrogen bond requires the expenditure of a Iarge amount of energy, it is probable that two strong copper-hydrogen bonds are formed in this system when the hydrogen molecule is dissociated, as in reaction (14). If this is the case, consideration of the energetics of the system shows that each copper-hydrogen bond must have an energy of at least 45 kcal./mole (6). The detailed structure of the dimer and of the compound with hy-

175

ACTIVATION O F MOLECULAR HYDROGEN

drogen are not known and can only be conjectured. Calvin suggested structure I below for the dimer; I may react with hydrogen t o form II.* (Q reprcsents a quinoline molecule here.)

CI-I%,

A

0

I

c:1 I .i

I

C

0

I

cu-Q I

+(;I1

I

/

/ \

0

+ Hz * I+CU

I

\

0

I

11 11 -c:u--c2 I

I

(15)

I

I

c11;i

CH ;;

I

I1

For steric reasons only threefold coordination may be possible around the copper ion in I. Two possibilities for the outer electronic structure of copper in coniplex I1 are 3d 4s 4 p 0

+-

11: 2 2 2 'L 1 p j21 B: 2 2 2 2

"p"

1 dsp2

Cuprous ion complexes with four ligands are normally tetrahedral, illvolving s p 3 hybrid orbitals (electronic distribution A ) . However, the cuprous hydrogen complex IT, which is of the form (Cu'XSH), is isoelectronic with four coordinate complexes of cupric ion, of the form (CulIXq), which are k1io1vn to be planar and to use dsp2 orbitals (distribution B ) . It seemed possible, therefore, that because of its unusual electronic structure, complex TI was also planar. Construction of scale (Fischer-TaylorHirschfelder) models indicates that this is probably not the case. A planar model of I can be constructed but not of 11; insufficient space exists to accommodate the hydrogen atoms between the copper ions in 11. If, however, tetrahedral coordination is permitted about the copper ions, no * It is of corisiderablc interest that some 14 years after Calvin postulated these structures, van Niekerk arid Schoeriing (8)showed Iiy X-ray analysis t h a t cupric wetate monoliydratc, Cu (OAc) 2.H20, exists in the solid as dimeric molecules, the two copper ions being joined by !our bridging acetate groups. The copper-copper distance in the dimer is only 2.64 A., as compared with 2.56 4. in the metal.

176

S. I\.. V-ELLEIi AND G. A . MILLS

steric dificulties occur in either case. I n fact, the flexibility of the molecule is such bhat quite large variations can occur in the copper-copper distance and in the angle formed by the two copper-hydrogen bonds; this would be an advantage if some optimum geomctric configuration w r c

FIG.8. Model of c:uprous ncetate-cluinaline-hydroRcn dimer structural ioriiiula 11. X = hydrogen atoms ntt,:i.c:licdt,o copper. (A): Hydrogen atoms close t o one anothcr; (H) : hydrogen atoms scparated by rotating copper atoms.

necessary for the activation of hydrogen. This flexibility is illustrated by the atom models show1 in Fig. 8. It is of iuterest, that,, with tetrahedral coordination, both cis and trans f o r m of I aiid I1 exist. Only in the cis form do the two uiifilled orbitals of the c.uprous ions in complex I ap-

ACTIVATION O F MOLECULAR HYDROGEN

177

proach sufficiently closely t o react with a hydrogen molecule, and presumably only the cis form would be catalytically active. The electronic state of copper when cupric disalicylaldehyde was hydrogenated in pyridine solution has been investigated by Tyson and Vivian (9),who observed that the original green solution changed t o ruby red upon hydrogenation and th at no precipitate occurred, nor was the Tyndall effect evidenced in either original or hydrogenated solutions. Magnetic susceptibility measurements were made, the results of which corresponded exactly to one unpaired electron per copper atom on either the original or hydrogenated sample. However, later work by Wilmarth, Barsh, and Dharmatti (4) makes it appear that the unchanged paramagnetism after hydrogenation probably resulted from exposure of the reduced solutions to atmospheric oxygen. Wilmarth et al. found th a t in all instances the hydrogenated cupric solutions were diamagnetic when kept out of contact with air, showing that copper was present in the cuprous form. Exposure of the solutions to air very quickly caused the magnetic susceptibility to become that of one unpaired electron per copper, corresponding t o oxidation back to the cupric ion. The problem of determining in detail which properties of cuprous acetate permit it t o act as a hydrogenation catalyst remains largely unsolved. The most difficult step in the hydrogenation is the activation of hydrogen, and this appears possible in this system through the circumstance th a t (a) the cuprous acetate is present in part as a dimer and ( b ) the electronic and geometric structure of the dimer-solvent complex is such th a t two strong, metal-hydrogen bonds can be simultaneously formed. Even less can be stated about the details of the subsequent fast reaction, which may involve transfer of either a hydrogen atom or an electron from the copper-hydrogen complex to the reducible substrate. Since this reaction is not rate determining, kinetic studies furnish little information concerning its detailed mechanism. b. Cuprous and Silver Salts in Pyridine and Dodeeylamine. ( 1 ) Cuprous acetate: A kinetic study has been reported of the catalyzed hydrogenation of cupric acetate by pyridine and dodecylamine solutions of cuprous acetate (10) in the range 78°-1000, pHaof 515 mm., and cuprous acetate concentration of 0.01-0.18 M . The course of a typical hydrogenation is given in Fig. 9. The reaction is autocatalytic. I n all cases the quantity of hydrogen absorbed is close to the theoretical value for the reduction of cupric t o cuprous ion. After reduction to the cuprous state, further reduction t o copper metal was not observed. The reduction was found t o be first order in hydrogen pressure. Under fixed conditions of hydrogen pressure and cuprous acetate concentration, the initial rate was, to a first approximation, independent of cupric acetate concentration. This re-

178

8. W. WELLER AND G . A . MILLS

sult is qualitatively different from that of Dakers and Halpern ( 1 1 ) . to be discussed later, who found that the rate of reduction of aqueous cupric acetate solutions a t 100" and 190 Ib./sq. in. hydrogen pressure was proportional t o the cupric ion concentration. The dependence of rate of reduction of cupric acetate upon concentration of cuprous ion is shown in Figs. 10 and 11 in pyridine and in dodecylamine solution. These figures were calculated from data such as shown in Fig. 9 and show the instantaneous rate of hydrogenation as a 2.0 MILLIMOLES CU(OAC)~ ' He0 0.98 MiLLlMOLES CuOAc 40 ML. WRlDiNE

FIG.9. Reduction of cupric acetate monohydrate in pyridine a t 100" (10).

function of the instantaneous concentration of cuprous ion during the course of reaction. The conclusion is clear that the rate of reduction is expressed by Rate = kp,,(Cu+) (16) where (Cu+) expresses the total concentrations of cuprous ion. (2) Silver salts: Before the foregoing data are discussed, the results with silver acetate will be given, as they are so similar. The course of a typical hydrogenation is given in Fig. 12 (10). The quantity of hydrogen adsorbed increases, a t a rate decreasing with time, until it is approximately that calculated for reduction of silver acetate t o silver metal; this was checked by weighing the amount of silver a t the conclusion of a run. The solution does not remain homogeneous during the hydrogenation, in contrast to all the examples previously given. However, the silver metal

179

ACTIVATION O F MOLECULAR HYDROGEN

TOTAL CUPROUS ACETATE ,MOLES / L.

FIG.10. Reduction of cupric acetate; instantaneous rate vs. total cuprous acetate concentration in pyridine, 100" (10).

6

TOTAL CUPROUS ACETATE, MOLES /La

FIG.11. Reduction of cupric acetate; instantaneous rate vs. total cuprous acetate concentration in dodecylamine, 100" (10).

produced dpring the hydrogenation is not catalytically active, as shown by the following occurrences: (1) the rate of hydrogenation continuously decreases as the reaction proceeds and (2) in two successive experiments the metal produced during the first does not influence the rate of the second experiment. Like the cuprous reduction in pyridine, silver acetate reduction is pro-

180

S.

W. WELLER AND G . A. MILLS

portional to the first power of the concentration of Ag+. The apparent activation energy for the reaction in solution is about 13 kcal./mole. I n the experiments just described, pyridine was used as a solvent; it was found t ha t dodecylamine is also suitable. However, aqueous silver acetate was not reduced, nor did reduction proceed when glacial acetic acid was the solvent. Also, silver chloride dissolved in pyridine was not reduced. 2.0 MlLLlMOLES AoOAc 40 ML PYRlDlNE

When either pure D2 or an H2-DZ mixture is used in the reduction of silver acetate in pyridine a t 78", no exchange is observed either during or after the reduction. (5)Discussion: The first-order dependence on the total concentration of silver or cuprous acetate for the hydrogenation rate can be interpreted on the basis either that the catalytically active species is a monomer or that i t is a dimer, provided, in the latter case, th a t the dimerization constant is so high that substantially all the catalyst molecules are present as dimers. However, molecular-weight measurements showed th a t over the concentration range in which kinetic studies were made, dimerization of cuprous or silver acetate was smali. This means t,hat in this system it is the monomer of silver or cuprous acetate which is catalytically active, in marked contrast to the cuprous acetate-quinoline system (see above). It seems probable that the difference between the quinoline and the pyridine systems is related to the different size of the solvent molecules. Formation of monomeric CuQsOAc is sterically much less favorable than

ACTIVATION OF MOLECULAR HYDROGEN

181

that of CuPysOAc, for example, and the apparent dimerization constant for cuprous acetate in quinoline (about 11 mole-' at 100°C.) is 10 to 20 times greater than i t is in pyridine. As will be discussed in the next sections, a number of other examples have been established in the last few years which argue conclusively that the homogeneous activation of Hz does not require the action of a dimeric metal catalyst species. It is still necessary, for energetic reasons, that both fragments of the hydrogen molecule be able simultaneously to form bonds of considerable stability. I n the case of cuprous and silver acetates in pyridine, both a heterolytic split of hydrogen (into a proton and a hydride ion) and a homolytic split (into two hydrogen atoms) should be considered as possible ratedetermining steps. If two atoms are formed, both may become attached to the metal ion; although no metal salt of this nature is known t o exist, there is some evidence that this may be the case for iron hydrocarbonyl, H,Fe(CO)a. Alternately, one hydrogen atom may go to the metal ion and one to a pyridine molecule with the formation of a radical. If a split into a hydride ion and a proton occurs, the hydride ion might react with, for example, the silver ion to form a solvated silver hydride molecule, and the proton might react with pyridine to form pyridinium ion. It is of interest that AH" for the reaction is in the neighborhood of 4-30-35 kcal. Since the over-all activation energy must at least equal the (endothermic) heat of the rate-determining step, this high AH" value is consistent with the failure of aqueous silver acetate t o be reduced by hydrogen. The corresponding AH" value for pyridine solutions cannot be computed because of inadequate data, but qualitative considerations of the difference in heats of solvation of silver and hydrogen ion indicate that the value might be 5-10 kcal. lower than in water. I n pyridine solution the cuprous or silver acetate molecules probably exist as the solvated complexes CuPy30Ac and AgPy,OAc. (In the case of silver, coordination with less than four groups may also occur.) Presumably one of the coordinated solvent molecules is displaced when the activated complex with hydrogen is formed. This will not change the electronic configuration about the metal ion, provided that either two hydrogen atoms or a hydride ion become attached t o the metal; in either case, the 3 4 4s, and 4p (4d, 5s, and 5 p for silver) orbitals remain filled. 2. Cupric and Mercuric Salts in Aqueous Solution

a. Cupric Salts. ( 1 ) Introduction: Ipatieff and co-workers long ago reduced aqueous solutions of cupric and other metal salts with molecular

182

S. W. WELLER AND G. A. MILLS

hydrogen (12). Typical conditions, as given by Ipatieff, Corson, and Kurbatov (IS),are a solution of cupric sulfate, 12 hours a t 150°, and a hydrogen partial pressure of 50 atm. Metallic copper is obtained, along with some residual dissolved cuprous salt. The interest of these investigators was to obtain pure metallic copper for catalytic studies. Other than the initial reduction of the solution to produce the first copper, it was not evident whether or not this was a heterogeneous reaction. Recently there has been an increasing interest in inorganic hydrogenation reactions involving the reduction of metallic ions in aqueous solutions. Reactions of this type have found important metallurgical application as methods of recovering metals from leach solutions (14,16).The reduction of ions, such as Ni++, Co++, UOz++, and VOZ-, has been examined kinetically. All these reactions were shown t o proceed heterogeneously and to require the presence of hydrogenation catalysts of the usual type such as finely divided metallic nickel or cobalt. Recently Halpern and Dakers (16) studied the reduction of cupric acetate in aqueous solution. They found, as did Ipatieff and Werchowsky (I,%'),that the reaction takes place at relatively moderate temperatures and hydrogen pr'essures with the formation of cuprous oxide under certain conditions. Halpern and Dakers found that the reduction of cupric ion does not require an added catalyst and that the activation of hydrogen proceeds homogeneously in solution. Their experiments were carried out in a stirred autoclave between 80-140" and hydrogen partial pressures ranging from 6.8 to 34.0 atm. Samples were withdrawn periodically and the unreacted cupric ion determined spectroscopically by use of the blue cupric ammine color developed by addition of ammonia. The solutions of cupric acetate used in the reduction experiments were buffered to a pH between 4 and 5. The stoichiometry of the reaction appears to be represented by the following reaction : 2Cu(OAc)z Hz HzO -+ CuzOJ, 4HOAc (18)

+ +

+

The experiments were generally continued until about 98% of the cupric acetate had reacted. (2) Kinetics: Rate plots depicting the course of some typical experiments, through which the temperature and hydrogen partial pressure were held constant, are shown in Fig. 13. The cupric acetate concentration always fell off according t o a first-order kinetic law, as illustrated by the linearity of the log CU(OAC)~ vs. time relationships shown. Further indication of first-order kinetic behavior was provided by the constancy of the determined rate constant with vayiation in the initial cupric acetate concentration.

ACTIVATION OF MOLECULAR HYDROGEN

183

The effect of hydrogen pressure on the reaction was measured. Figure 14 shows that the first-order rate constant is directly proportional t o the hydrogen partial pressure, PH1.Since the solubility of hydrogen in most aqueous solutions obeys Henry’s law over this range of pressure, this relation also implies that the first-order rate constant is proportional t o the concentration of molecular hydrogen in solution (H2)

ki = kz(H2) where a: is the Henry’s law constant.

= JCZCQH~

TIME, MINUTES

FIQ.13. Typical first-order rate plots showing effect of varying initial cupric acetate concentration (11).

The influence of temperature on the reaction is shown by the plot of data given in Fig. 15. The plot of log k l vs. 1/T was a n excellent straight line whose slope corresponded to an activation energy of 24.2 kcal./mole. Consideration of this reduction of cupric ion led Peters and Halpern (17) t o propose that cupric acetate should also be capable of functioning as a homogeneous catalyst for the reactions with hydrogen of other compounds which are thermodynamically reduced more readily than cupric acetate itself but which do not react with hydrogen in the absence of a catalyst for kinetic reasons. In accordance with this theory, it was found th a t the homogeneous reduction of dichromate salts by molecular hydrogen in aqueous solution, represented by the equation Crz07-

+ 3H.2 + 8Hf -+ 2Cr+++ + 7Hz0

(20)

184

S . W. WELLER AND G . A. MILLS

He PARTIAL PREsSwtE

- ATM.

FIG,14. Dependence of the rate on hydrogen partial pressure a t 130" (11).

0

100

TIME

Eoo

- MINUTES

FIQ.15.First-order plots a t different temperatures: NaOAc, 0.25 M./liter; HOAc, 0.44 M./liter; HI partial pressure, 13.6 atm. (11).

ACTIVATION O F MOLECULAR HYDROGEN

185

would proceed in the presence of dissolved cupric acetate, the latter apparently acting as a true catalyst. I n the absence of cupric acetate no reaction between sodium dichromate solution and hydrogen could be detected up t o temperatures of 160" and hydrogen partial pressures of 30 atm. When cupric acetate was present, the reduction of dichromate proceeded at temperatures as low as 80".

FIQ.16. Typical rate curws for the reactions of dichromate and cupric acetate with hydrogen: temperature, loo", H1 partial pressure, 13.6 atm. ( 1 7 ) .

The course of a typical experiment is shown in Fig. 16. At constant temperature and hydrogen pressure, the concentration of Crz07- always decreases linearly with time. This was further confirmed by finding equal rates of reduction over a wide range of original CrzO?- concentration. Figure 16 also shows that the concentration of cupric acetate remained constant as long as dichromate was undergoing reaction. Only when the reduction of dichromate was complete did the cupric acetate react with hydrogen to form cuprous oxide. Apparently the reduction of cupric acetate is not affected by the previous dichromate reaction or by the presence of small amounts of chromic salts in the solution. The dependence of the rate of the dichromate reaction on the concentration of cupric acetate provides further support for the catalytic role of the latter. The results in Fig. 17 show that the rate increases in a nearly linear manner with increasing cupric acetate concentration; i.e., ko

=

ka(Cu(OAc)2)

(21)

Actually the relation is not quite linear, ka showing a tendency t o fall off slightly with increasing (Cu(0Ac)J. The reduction of dichromate was studied a t hydrogen partial pressures ranging from 0 to 27.2 atm. The results show the rate t o be di-

186

S. W. WELLER AND G. A. MILLS

'

O

f

1'"

l

Cu(OAC)*- MOLES PER LITRE

FIG.17. Dependence of the rate of reaction of dichromate on the concentration of cupric acetate: temperature, looo, Hz partial pressure, 13.6 atm. (17).

rectly proportional to the partial pressure of hydrogen throughout this range. The kinetics of the reaction thus conform to the bimolecular rate equation:

-d(Cra07-) dt

=

k4'(Cu(OAc) 2)Pa,= kr(Cu(OAc)2) (H2)

(22)

and The rate of the catalyzed dichromate reaction was measured at temperatures ranging from 80" to 140". The results were found to give a good Arrhenius plot. The activation energy is 24.6 kcal./mole, in close agreement with the value of 24.2 kcal./mole found for the reaction of cupric acetate itself with hydrogen. (3) Discussion: The kinetic similarity of the dichromate reduction in the presence of cupric acetate and the reduction of cupric acetate itself support the view that both have the same rate-controlling step. Halpern and co-workers suggested that the rate is determined by a bimolecular process involving one molecule of cupric acetate and one molecule of hydrogen. They have therefore, proposed the following reaction scheme.

+

k

CU(OAC)~ Hz ---+ Cu(OAc)?.Ha slow

(25)

ACTIVATION O F MOLECULAR HYDROGEN

187

followed by (a)

+ Cu(0Ac)z + HzO fast

Cu(0Ac)z.H~

-

or

+

(b) ~ C U ( O A C ) ~ . H Z CrzOy-

+ 8H+ fast

+Cu2O

2Cr+++

dt

(27) =

At 100" k2' was found t o be 4.45 X of Cr?O,-d(Crz07-) dt

(26)

+ 7Hz0 + ~ C U ( O A C ) ~

For the reduction of CU(OAC)~

-~ [ C(OAC) U 21

+ 4HOAc

=

kz'ICu(OAc) z]PH, atm.-' min.-l For the reduction ~ ~ [ C U ( O A ~ ) ~ ] P ~ ~ (29)

k4' extrapolated to zero Cu(0Ac)z concentration = 7.45 X 10-6 atm.-' min.-1 The ratio kz' :k4' is thus found experimentally to be equal to 5.95, in agreement with the theoretical value of 6.0 corresponding to the fact that reduction of Crz07- involves six electrons, compared with one for Cu(0Ac)z. I n a consideration of the reduction of cupric acetate, if the solubility of hydrogen is taken to be independent of temperature (known to be nearly so in the range considered), then for a n a estimated t o be 6.4 X moles atm.-' 1.-'

where k is the rate of reaction (25) and kz is the over-all rate of CU(OAC)~ reduction. The factor 2 reflects the fact that each time a molecule qf hydrogen is activated, as shown in reaction (25), two molecules of CU(OAC)~ are reduced in the over-all reaction. With the use of the value of k equally obtained from Cu(0Ac)z or CrzOl- reduction, an activation entropy AS,' of -6.5 e.u. a t 100" was calculated. In the preceding discussion the reaction has been represented as involving undissociated Cu (OAc) 2 molecules rather than simple cupric ions or other cupric complexes. This appears t o be consistent with the kinetic results taken in conjunction with independent measurements of the dissociation constants of cupric acetate. These measurements indicate that in qolutions of the composition used in these experiments, cupric acetate is predominantly undissociated. The degree of dissociation to CuOAc+ ranges from about 10% a t the highest NaOAc (used as a buffer) concen-

188

5. W. WELLER AND G. A. MILLS

tration of 0.75 M./liter t o about 25% a t the lowest concentration of 0.25 M./liter. The dissociation to Cu++ is negligible. Similar considerations discourage the suggestion that cuprous acetate, which may be formed as an intermediate, has a n appreciable catalytic influence on the reaction. Cuprous acetate is unstable in aqueous solution where it is hydrolyzed and precipitated quantitatively as cuprous oxide. It may be concluded that only trace concentrations of cuprous acetate were present in these solutions during reaction. Any reaction mechanism in which an important catalytic influence is attributed to cuprous acetate would thus appear to be inconsistent with the observed kinetics. 6. Mercuric Salts. Inspired by their results on the activation of hydrogen by aqueous solutions of cupric acetate, Halpern, Korinek, and Peters (18) tested the acetate salts of a number of other divalent metals, including Mg, Ca, Mn, Co, Nil Zn, Cd, Hg, and Pb, for similar catalytic activity. An aqueous solution of each salt was heated in a stainless steel agitated autoclave under 13.6 atm. of hydrogen, to temperatures u p t o 150'. Evidence was sought either for direct reaction between hydrogen and the metal acetate or for homogeneous catalysis by the metal acetate of reduction of dichromate. Of these salts only mercuric acetate appeared t o activate hydrogen. At about 100' it readily underwent reduction to mercuric acetate, as represented by the equation 2Hg(OAc)z

+ Hz

-+

Hgz(OAc)zJ

+ 2HOAc

(31)

The mercurous acetate product was only slightly soluble. Typical rate curves a t constant temperature and pressure define a decrease in mercuric acetate concentration according to a first order kinetic law, finally leveling off a t a small constant value. During this last period mercurous acetate was being reduced to metallic mercury. For the initial period concerned with reduction of mercuric acetate, a preliminary kinetic study found the reaction to be represented by -d[Hg(oAc)zl dt

=

~[H~(OAC)~](H~)

where k = 8 X 10l2 exp (-20,70O/RT) 1. mole-' min.-l The proposed reaction mechanism is

+

Hz -+ Hg(OAc)z.Hz slow Hg(OAc)z Hg(0Ac)zHz Hg(0Ac)z Hg,(OAc)z 2HOAc fast

+

-+

+

(33) (34)

3. Ethylene Platinous Chloride

Platinum olefin compounds have long been known, the most familiar being Zeise's salt, K(PtCl2C2H4)C1.The fundamental compound of such

ACTIVATION O F MOLECULAR HYDROGEN

189

complexes can be considered t o be (PtClzC2H4)2, bis-(ethylene)-dichlorop-dichlorodiplatinum (11))or ethylene platinous chloride. This compound forms light orange crystals that decompose when heated above 130'; it dissolves in many organic solvents and decomposes in water. It exists as a monomer in acetone and as a dimer in benzene. At room temperature solid ethylene platinous chloride is quantitatively reduced by hydrogen according t o the equation

This reaction was investigated recently by Flynn and Hulburt (19,20) with the view t o utilize such information to elucidate the role of metal catalysts i n the heterogeneous reduction of olefins. When a toluene solution of ethylene platinum chloride was reduced by hydrogen, they found that (1) in dilute solutions a t room temperature a n induction period of several minutes might occur before platinum began t o form and that once reduction had begun, platinum precipitated quite rapidly; (2) the presence of platinum from a previous run would catalyze the reduction of the complex a t temperatures as low as -70"; (3) a t -30" the platinum formed in the reaction vessel catalyzed hydrogenation of ethylene by hydrogen; (4) reduction of a solution of ethylene platinum chloride with deuterium a t -23.7" resulted in the formation of ethanes ranging from CzH6 t o C2Ds and ethylenes from C2H4 t o CeD4. When solid (PtC12 C2H4)2 was reduced with Dz a t - 2 2 O , the results of a detailed mass spectroscopic study indicated that deuterium did not add unsymmetrically to the carbon double bond. From these results it appears th at although the initial reaction may be homogeneous, the reaction becomes heterogeneous, being autocatalyzed by the platinum formed in the reduction. Flynn and Hulburt, however, went further into a study of the hydrogenation of ethylene by hydrogen when a stream of these gases was passed through a solution of ethylene platinum chloride. They found th a t ethyIene inhibits the formation of platinum metal. This inhibition seems t o indicate t hat the first step of the reduction is a dissociation of the complex with ethylene as one of its products. T h e platinum formed as a result of the initial dissociation and reduction would catalyze the reduction of the complex. Ethylene, however, still inhibits the reaction markedly in the presence of platinum. They found that when the ethylene-to-hydrogen ratio was sufficiently high and the temperature sufficiently low, hydrogenation occurred without a n y deposition of metallic platinum. The minimum value for the ratio of ethylene t o hydrogen needed to inhibit effectively the formation of plat-

190

8. W. WELLER AND G . A. MILLS

inum increases rapidly with temperature. Under the conditions used, it was found that a t 0" in toluene a ratio of 2.6 is not great enough; a t -4.8" in acetone, a ratio of 6.4 is sufficient; and at -33.5" a ratio of 0.67 is sufficient. I n runs in which metallic platinum was not formed, the rate of formation of ethane was quite low or zero at temperatures above - loo, but at temperatures below -14" the rate was higher, falling off slowly with decreasing temperature. This is illustrated in Fig. 18.

-40

-30 -20 Temp.. "C.

-10

0

FIQ.18. Rate of ethane formation vs. temperature, 3.2 g./L (C2H4PtCIz)zin acetone; (C2H4)/(H2) = 6.4; flow rate = 0.11 l./min. (19).

Flynn and Hulburt interpreted these results to indicate that the formation of ethane a t -10" without the deposition of platinum occurred by a mechanism different from the reduction a t higher temperatures. They point out that Chatt (21)found that PtC1z(C2H4)2 can be formed by bubbling ethylene through an acetone solution of (PtC12C2H4) at low temperatures and that this compound was stable in an atmosphere of ethylene below about -6". They suggest that the low-temperature r e duction in the presence of ethylene may consist of the sequence of steps given below :

ACTIVATION OF MOLECULAR HYDROGEN

191

to be repeated in a chain fashion, giving a catalytic reaction for the hydrogenation of ethylene. It is their opinion that the evidence points to the low-temperature reaction being o truly homogeneous reaction. If so, they point out, it appears that this catalytic property of platinum can be accounted for by the properties of the individual atoms rather than some macroscopic property of the metal catalyst. It is obvious that additional work is needed to define the nature of the reaction between the pIatinum compIex and hydrogen in this interesting system. 4. Cobalt Carbonyl One of the most interesting catalytic reactions to be discovered is the so-called ( ( O X O " reaction. The 0x0 reaction consists of the catalytic addition of carbon monoxide and hydrogen to olefins to form, primarily, aldehydes possessing one carbon atom more than the original olefin. This hydroformylation reaction was developed during World War I1 by Roelen and co-workers (22) in Germany. While they utilized solid FischerTropsch cobalt-thoria catalyst, it became apparent to them that the hydroformylation reaction was probably a homogeneous catalytic process with either dicobalt octacarbonyl or cobalt hydrocarbonyl as the catalyst. The use of solutions of dicobalt octacarbonyl under hydrogen and carbon monoxide pressure has been described in detail (%?,%$)as effective in adding hydrogen or hydrogen and carbon monoxide to unsaturated organic compounds. Thus in the 0x0 reaction hydrogen as well as carbon monoxide is activated by what is believed to be homogeneous solutions of dicobalt octacarbonyl. In the normal 0x0 reaction a certain amount of hydrogenation occurs, a minor amount of olefins being converted to paraffins; in the case of certain olefinic compounds hydrogenation indeed occurs to the exclusion of hydroformylation. It is a remarkable fact that this catalytic reaction occurs in the presence of carbon monoxide and also of sulfur compounds, although cobalt metal is notoriously poisoned by traces of these compounds. The significance of this was pointed out by Adkins and Krsek (23) and Wender, Orchin, and Storch (25) in terms of the concept that the hydroformylation catalyst is a homogeneous one, not sensitive to carbon monoxide or sulfur compounds and in this respect different from usual solid cobalt catalysts. Wender, Orchin, and Storch (25) sought to test critically for the homogeneous nature of the hydrogenation reaction. They measured the hydrogenation reaction of butyraldehyde a t 185"under 2000 psi hydrogen preasure. They performed a series of three experiments using variable

192

S. W. WELLER AND G. A. MILLS

pressures of carbon monoxide along with the hydrogen and incorporating into the reaction finely divided cobalt metal. The reduction of the butyraldehyde to butanol-1 is shown below: 185", 2000 psi Hs

Carbon monoxide, psi Reduction of butyraldehyde Catalyst functioning

None Yes Co metal

1000 300 Yes No or [ C O ( C O ) ~ ] ~ None HCO(CO)~

As noted above, the butyraldehyde was reduced to the alcohol in these experiments when no carbon monoxide was added and when 1000 psi was added, but not when 300 psi was added. When no carbon monoxide was present the reduction was catalyzed by the metallic cobalt. When the 1000 psi carbon monoxide was used, it was presumed that the reaction was homogeneous, soluble dicobalt octacarbonyl or cobalt hydrocarbonyl being the catalyst. It is known, however, that a t 150" a carbon monoxide pressure ofiat least 600 psi is needed to keep [Co(CO).& from decomposing t o cobalt metal. When only 300 psi of carbon monoxide was present, therefore, the cobalt would remain as metal and be inactive because it was poisoned by the carbon monoxide. I n another experiment butyraIdehyde was treated with a benzene solution of dicobalt octacarbonyl a t 158" with 2000 psi initial hydrogen pressure in the absence of carbon monoxide. No hydrogenation of the butyraldehyde occurred. The carbonyl was reduced t o cobalt, which did not function as a catalyst because of carbon monoxide poisoning. Additional experiments showed that sulfur compounds did not retard hydrogenation when cobalt metal was used in the presence of hydrogen and carbon monoxide but did when hydrogen alone was employed. The structure of cobalt carbonyl and the mechanism of the 0x0 reaction have been the subject of a number of investigations. Cobalt metal It is believed th a t possesses 27 electrons arranged ls22s22p63p23s63d14s2. dicobalt octacarbonyl has a structure in which each cobalt is attached t o three terminal carbon monoxide groups and t o two bridging carbon monoxide groups and that a metal-metal bond exists in addition. Thus each cobalt acquires, by covalent sharing, 8 electrons from the surrounding carbon monoxide molecules and one from the other cobalt atom. With the original 27 extranuclear electrons there is a total of 36 electrons, the electronic structure of the rare gas krypton. The electronic structure from the viewpoint of the cobalt would be 3d'04s24pe.It has also been pointed out t ha t the reaction of dicobalt octacarbonyl with hydrogen to form

ACTIVATION OF MOLECULAR HYDROGEN

193

cobalt hydrocarbonyl

co

0 C

co

\ / \ / CO-CO ~ C O - C / \ / \

co

C 0

O

+ H2

-+

ZHCo(CO)4

(38)

co

provides a n alternative manner for cobalt to achieve the rare-gas electron structure. I n the case of the hydrocarbonyl each hydrogen atom provides one electron for each cobalt atom t o augment its original 27 extranuclear electrons. I n this regard it is significant that nickel with 28 extranuclear electrons does not form a hydrocarbonyl and does not behave like cobalt in the 0x0 reaction. Since the energy required for splitting the hydrogen molecule is over 100 kcal./mole, the bond formed by each hydrogen atom in the hydrocarbonyl must be quite stable. The fact that, as shown in the previous equation, two cobalt atoms contained in the octacarbonyl participate in activating one molecule of hydrogen is significant,. For one thing, this means that there are two H-Co(C0)4 bonds which can together provide the energy t o split the high-energy H-H bond. After activation of the hydrogen molecule by the octacarbonyl as indicated above, the 0x0 reaction, it has been suggested, proceeds probably by a carbonium-ion mechanism. Numerous studies carried out at the U.S. Bureau of Mines and elsewhere strongly support this postulate. I n water HCo(CO)&behaves as a strong mineral acid. For hydrogenation, addition of a proton t o an olefinic material must be followed by acquisition of a second hydrogen atom. For acid catalysts, such as HCl-AlClS or HOH-A1203-Si02, this has frequently been considered t o be achieved through abstraction of a hydride ion from another hydrocarbon molecule. However, for the cobalt hydrocarbonyl the reaction has been postulated t o occur through reactions such as PhZCH+

+ cO(c0)4- -+

Ph&HCo(C0)4

+

HCO(CO) 4 I

PhzCh

+ [Co(C0>4]~

(39)

Alternative mechanisms have been proposed involving other cobalt carbony1 polymers as well as ion-radical sequences (26). At the present time the complete reaction must be regarded as not known in a.11 its details.

194

S. W. WELLER AND G . A. MILLS

6. Bases

Perhaps the simplest example, and certainly onc of the most interesting cases, of the homogeneous activation of hydrogen is the base-cataIyzed exchange between hydrogen gas and liquid water or ammonia. The phenomenon was discovered in 1936 by Wirtz and Bonhoeffer (dy), who found that when D,O containing KOH was heated with hydrogen in sealed tubes at loo”, exchange occurred between the DzO and Hz.Reaction times of the order of 10 to 60 hours were used, and KOH concentrations in the range 0.1 to 1 N . The extent of exchange increased with increasing reaction time and with increasing KOH concentration. The exchange does not occur in neutral or acid solution. No further work was done on this system for several years. Then Abe reported in 1941 (28) th at the exchange observed by Wirtz and Bonhoeffer (27) was caused not by hydroxyl ion catalysis, but probably by colloidal iron oxides present as impurities in the KOH used. Abe found, for example, that the activity was destroyed by diffusing the KOH through a collodion membrane or by prolonged heating of the solution and that the activity was restored by temporary introduction of a n iron wire into the deactivated solution. Abe’s results, if correct, meant th a t the exchange observed in the system of Wirtz and Bonhoeffer was of relatively trivial interest from the standpoint of representing a novel type of hydrogen activation. After Abe’s work the problem again lay dormant for a number of years until i t was taken up by Wilmarth and his co-workers. Claeys, Baes, and Wilmarth (29) in 1948 reported that a liquid ammonia solution of potassium metal rapidly catalyzed o-p Hz conversion, a half-time in solution of 37 sec. being obtained a t -53”. I n order to establish that this result was due t o dissolved metal and not to amide ion impurity, Claeys, Dayton, and Wilmarth (30)studied the o-p H z conversion in the presence of potassium amide in liquid ammonia. Rates were obtained comparable with those occurring with the metal solution. The mechanism of the conversion was different for the two cases, however, since the amide solution also catalyzed exchange between gaseous deuterium and liquid ammonia, while the metal solution did not. ft was assumed that the metal acted by a “paramagnetic” mechanism and the amide ion by a “chemical” mechanism. In the same note Claeys, Dayton, and Wilmarth (SO) reported confirmation of Wirtz and Bonhoeffer’s results on the aqueous alkali system and questioned the validity of Abe’s objections. Both the amide ion-liquid ammonia and the hydroxide ion-liquid water systems were subsequently studied in greater detail by Wilmarth’s group. Wilmarth, Dayton, and Flournoy (32) made a kinetic study of the

ACTIVATION OF MOLECULAR HYDROGEN

195

0 - p H z conversion and the Dz-H20 exchange in the KOH-HzO system. They first established th at the reaction was homogeneously catalyzed by showing t ha t NaOH gave the same rate constant as KOH and that the reaction was not influenced by the surface:volume ratio of the reactor or by dialysis of the alkaline solution. The rate of conversion of p-Hz in the temperature range 80-110" was then shown to be first order in concentration of hydroxyl ion up to an ion concentration of 1 M. Figure19 is a n Arrhenius plot of their data for p-Hz conversion. The quantity k" is the second-order rate constant (ie., corrected for hydroxyl ion and

I/T x

lo3, *K?

FIG.19. Temperature coefficient of parahydrogen conversion by aqueous potassium hydroxide (31).

hydrogen concentration) for the reaction in solution; its units are 1. moles-' min.-1 The activation energy computed from this line is 23,800 kcal. mole-l. Over the temperature range studied, k" is given b y k" = 4.7 X 1013exp (-23,80O/RT) 1. mole-' min.-l

(40) Comparison of this result with Eq. (30) indicates the striking similarity in rates for this reaction and for the reduction of aqueous cupric acetate, studied by Halpern. The exchange of Dz with aqueous KOH to produce HD and Hz was found by Wilmarth, Dayton, and Flournoy (31) t o proceed a t a slightly slower rate than p-Hz conversion under comparable conditions. Figure 20 shows the decrease in deuterium content and the initial rise and subsequent fall in hydrogen deuteride content of the gas as the exchange proceeds a t 100". (Since water is present in large molar excess, the final gas composition is almost pure hydrogen.) Reaction rate constants, computed from such curves, for both the DZ-HZO and HD-H20 exchanges are plotted in Fig. 21 as a function of hydroxyl ion concentration. The

196

S. W. WELLER AND G. A. MILLS

100

80

rc

60

20

0 0

80

160

240

320

TIME, MIN.

FIG.20. Variation in moles of Dt and HD with time during exchange with aqueous potassium hydroxide a t 100" (31).

CONCN., MOLE I:'

FIG.21. Rate of exchange in solution by aqueous potassium hydroxide at 100" (51).

ACTIVATION O F MOLECULAR H Y D R O G E N

197

uppermost line is that for p-Hz conversion, which is thus seen t o go somewhat faster than the exchange reaction. Values of the bimolecular rate constants k" for these three cases are summarized in the following table. Rate of Activation by Aqueous Potassium Hydroxide a t 100" Molecule activated

k" (1. mole-' rnin.7)

p-Hz Da HD

0.34 0.21

0.56

The assumption was made that all the species of hydrogen have identical solubilities. As the authors point out, only the rates for D2 and H D are directly comparable, as p-H2 conversion may occur in some collisions of insufficient energy to cause exchange. The lower rate constant for HD relative t o D2 is apparent rather than real; as only half the exchange reactions of H D with HzO result in a loss of HD, the true rate constant for H D is twice that shown in the preceding table. Wilmarth and Dayton (32) reported similar results for the activation of hydrogen by potassium amide in liquid ammonia a t -53". T h e rate of p-Hz conversion is again linear in concent,ration of amide ion, and the exchange of deuterium with the solvent occurs a t a rate comparable with, but slightly lower than, the rate of p-H2 conversion. Although the activation energy for the conversion was not measured, the authors computed that it would be about 10 kcal. mole-' if the preexponential factor were the same as for the hydroxyl ion-catalyzed reaction. This lower activation energy was attributed to the greater base strength of the amide ion relative t o the hydroxyl ion. Of particular interest was the magnitude of the rate of hydrogen activation. At -53" the rate constant for amide ion in ammonia was about lo4 times that a t +loo" for hydroxyl ion in water; a t a n amide concentration of 2.7 X mole L-I, the half-time for exchange in ammonia solution a t -53" was less than 1 min. Because of the simplicity of the chemical systems involved, the reaction mechanism of these base-catalyzed activations of hydrogen should be particularly susceptible to determination. Wilmarth and his coworkers established th at the rate of activation is proportional t o concentration of base (OH- or NHz-). Although they did not establish th a t the rate is first order with respect t o hydrogen, it is probable th a t this is so. (Because, in the absence of complications, every exchange reaction is apparently first order during any single experiment, it is necessary t o vary the initial concentration t o establish the true order.)

198

8. W. WELLER AND G. A. MILLS

It is probable, then, that the activated complex for exchange of deuterium with water contains deuterium, hydroxyl ion, and possibly water molecules. Wirtz and Bonhoeffer had suggested the following two-step mechanism : DP OH--+ DDOH (41) DH20 -+ H D OH(42)

+

+

+ +

As Wilmarth, Dayton, and Fluornoy point out, the reaction kinetics would require that reaction (41) be rate determining. It becomes pertinent to examine the thermodynamics for this reaction, for comparison with the observed activation energy. Wilmarth et al. calculated that reaction (41) was endothermic by only +20 kcal. mole-’. (A figure of 22.kcal. mole-’ was obtained by the authors of this review, a different method of calculation being used. This was considered important t o do because the enthalpy of formation of OH- (g) given in the National Bureau of Standards compilation (33) is high by about 30 kcal. mole-’.) This is remarkably close t o the observed activation energy of 24 kcal. mole-’. Although this result does not a t all establish the correctness of the mechanism represented by Eqs. (41) and (42), it is a t least consistent with it. There is a problem associated with this mechanism. I n order that the energetics of Eq. (41) come out as indicated above, it is necessary t o assume that the hydride ion is stabilized by the hydration energy appropriate to a stable anion of the same size. It is questionable, in view of the reactive nature of hydride ion in water, whether this ion could survive long enough t o permit the orientation of water dipoles necessary for production of the full hydration energy. Wilmarth, Dayton, and Flournoy suggested an alternate mechanism in order to avoid this dilemma. They pointed out that the hydrogen molecule could be simultaneously attacked by a hydroxyl ion from one side and by a water molecule from the other. The reaction would be written

HO-

+ D : D + H:OH+

HOD

+ DH + OH-

(43)

Since the products are the same chemical species as the reactants, the over-all reaction is substantially thermoneutral; except for activation energy, the problem of energetics is thus side-stepped. The apparent termolecular reaction required by Eq. (43) is also no problem, as the dissolved molecules are essentially in constant collision with water molecules. Wilmarth, Dayton, and Flournoy questioned the adequacy of this “concerted attack” mechanism, however, as they believed that it would predict acid catalysis of the exchange reactions, as well as base catalysis.

ACTIVATION OF MOLECULAR HYDROGEN

199

The acid-catalyzed reaction, written analogous to Eq. (43),would be

OH3+

+ D: D + OH2 -+ OH2 + HD + DOHzf

(44)

Wilmarth et al. found that no exchange occurred after 5 hr. in 10 M sulfuric acid at 120". Similar experiments using deuterium with either concentrated HzS04 or HF-BF, have likewise shown no exchange (34). As comparison of Eqs. (43) and (44) shows, however, the nature of the bonds (other than D-D) which have to be broken and made is quite different in the two cases. The authors of this review do not see any reason, therefore, why the existence of the base-catalyzed exchange should require the existence of an acid-catalyzed exchange under any given set of conditions, although such exchange might occur under sufficiently severe conditions. 6. Miscellaneous

a. Hydrogenolysis of R M Compounds. Gilman, Jacoby, and Ludeman (36) have reported that organoalkali metal compounds such as phenyl sodium are reduced by molecular hydrogen a t room temperature, the more reactive compounds being reduced at atmospheric pressure. The products are the unsubstituted hydrocarbon and the alkali metal hydride, as indicated by the following typical equation:

The system is not a homogeneous one, as the organometallic compounds and alkali metal hydrides are insoluble in benzene, the solvent used; however, the rate of reaction was found not to be affected by the addition of platinum or palladium catalysts. The reaction is formally similar to the base-catalyzed activations of hydrogen studied by Wilmarth ; this becomes clear if Eq. (45) is written as

In both systems the hydrogen molecule is apparently split into ions, rather than atoms. Weight is Ient to this conclusion by the fact that the order of increasing ease of hydrogenolysis of phenyl derivatives is Ca, Li, Na, K, Rb, and Cs; this is also the order of increasing polarity of the carbon-metal bond. b. Hydrogenase. Stephenson and Stickland (36) a quarter of a century ago discovered that a number of bacteria possess the ability to activate molecular hydrogen at room temperature; they gave the name hydrogenase to the enzyme responsible for the catalytic activity. Among

200

S. W. WELLER AND G. A. MILLS

the compounds which can be hydrogenated by such bacteria are methylene blue, molecular oxygen, and sulfate, nitrate, and fumarate ion. I n 1934 Green and Stickland (37) reported th a t a n inert metal in a suspension of Escherichia coti behaves as a reversible hydrogen electrode, and Cavanagh, Horiuti, and Polanyi (38) found that several strains of E. coli and of Aerobacter aerogenes catalyze exchange between deuterium gas and water. This bacterial activation of hydrogen should presumably be classed as homogeneous catalysis, as individual enzyme molecules are probably responsible for the activation. There is no general agreement as to the nature of the enzyme, though inhibitor studies (39) suggest that hydrogenase is a n iron porphyrin enzyme. It seems likely th a t most progress in understanding the enzymic action will come from work with cell-free extracts of the bacteria, a field which is being actively studied (40-44). c. Cobalt Cyanide. When an aqueous solution of potassium cobaltous cyanide is acidified, hydrogen is liberated readily, owing t o the high redox potential of the cobalt cyanide which permits it t o reduce water according t o the following equation (there is some evidence that five rather than six cyanide ions are in the cobaltous and cobaltic complexes) :

Thus molecular hydrogen is generated in a homogeneous liquid system. Farkas and Farkas (45) rcported that the hydrogen generated was in isotopic equilibrium with the solution in which the cobaltous cyanide was dissolved. Ogg (46) reportled that if this reaction is carried out with deuterium gas present, H D is found. He proposed th a t the reaction depended upon atomic hydrogen and proceeded, as follows, in a manner similar t o that observed with molecular oxygen exchange triggered by ozone : H+DP-+HD+D

(48)

Blank runs demonstrated th at the reactants and final products caused no exchange between H, and Dz.

111. SUMMARY DISCUSSION The catalytic systems discussed in this chapter are summarized in Table I, which describes the catalyst, the solvent, the over-all reduction reaction, and the proposed reaction mechanism. I n the latter column the reactions are presented which have been proposed by the original research workers, references for which are also listed. It appears t o be significant that among the relatively few catalysts

TABLE I Summary of Catalytic Homogeneous Activation of Molecular Hgdrogen . Solvent

Catalyst Cu'OAc

$uinoline

Reduction cu++-r cu+,

Kinetics

Proposed reaction

Reference

k.p~,(CuOAc)'

CH3

I

C

Quinone + Semiquinone, 0-P Hz, HrDz

+

(Q c ~ H . o A ~ )c~u + ++ Q-&-H

H-&-Q

0

0

1

1

\

4, 5, 6, 7

/

C' I hl

s

I CuIOAc

Pyridine, dodecylamine Pyridine,

AgIOAc

10

++

PtCl2CzH4 Toluene [CO(CO)~]ZCyclohexane

OHNHz-

1

Water Ammonia

Butyraldehyde + Butanol-1 or Crotonaldehyde + Butyraldehyde 0-p Hs, Hz-Dz O-p Hz, Hz-Dz

++

AgOAc Hz+ AgH HOAC Slow AgH AgOAc + 2Agl HOAc Fast Cu(OAc)z Hz+ Cu(OAc)Z.Hz Slow Cu(0Ac)tHt Cu(0Ac)z-t CUZO 4HOAc Fast Hg(OAc)z Hz 3 Hg(OAc)z.Hz Slow Hg(OAc)z.Hz Hg(OAc)z+ Hgi(OAc)z +2HOAc Fast (PtCIzC2H4)z 2CzH4 2PtClz(CzH,)z ZPtClz(CzH4)z 2Hz * (PtClzCZH,), 2CzHtj [Co(CO)alz Hz + 2HCo(CO), HCo(CO)( substrate -+ SH CO(CO)~ SH CO(CO)r HCo(CO)r-+ SHz [CO(CO)&

+ + ++

+

++

++ +

Dz D-

++ OH--+ HzO+

+

-+

+

D- +DOH HD OH-

+

10 ii,i6,i r 18

19,20 Zb, 23

24125 26

27, 28, 29 so, 31, 32

202

S. W. WELLER AND G . A. MILLS

known, activity is found for CuI, Agl, and HgII, which form a series of similar electronic arrangements :

CU'

Ag'

HgII

Is

2s 2 p

39 3 p 3d

2

2 6

2 G

4s 4 p 4d 4.f

5s 5 p 5d

10

2 G

10 14

2 6

10

However, CuI1 (as the acetate), PtII (as ethylene chloride), and Coo (as the carbonyl), which do not have this electron distribution, are also effective homogeneous catalysts for activation of hydrogen. It does not appear possible at present, therefore, to set forth a simple generalization of electronic qualifications for the known active catalysts, much less to predict new systems. Of course what should be considered is the electronic structure of the metal in the catalytically active complex, rather than the electron distribution of the ion alone as shown above. If a system is known to be active, it is possible to write metal complexes with possible electronic structures which are in accord with the observations. Such post hoc structures each appear special and without generalization in any detail. It is clear that several of the active metal catalysts do not have unfilled d orbitals, whereas unfilled d orbitals have been considered necessary for this type of catalysis in some theories (47). However, in each of the instances discussed here there have been some unfilled electronic levels. The more general requirement may be considered to be, therefore, the existence of unfilled levels, not necessarily d orbitals, of low energy suitable for formation of a chemical complex with the molecule t o be chemically activated. It is evident that the activation of molecular hydrogen does not require the existence of a solid metal with properties associated with a group of metallic atoms. It seems probable that the present information on molecularly dispersed catalysts will apply to atoms or ions, present on the surface of metals or metal oxides or sulfides, which singly or in small numbers constitute active sites. Of course the atom on the surface must be affected by the bulk solid, which is part of its environment. While geometric factors are undoubtedly of considerable influence in many catalytic systems, it is likely that for homogeneous activation of hydrogen the geometric factor is not prominent, although it may be responsible for lack of activity observed on addition of certain complexing agents t o otherwise active systems. Insofar as the substrate hydrogen is concerned, because of the high

ACTIVATION OF MOLECULAR HYDROGEN

203

bond energy of the hydrogen molecule, the present authors believe that it is generally necessary that the interaction between hydrogen and catalyst system be such that both fragments of the hydrogen molecule be stabilized either by formation of a new bond (with a catalyst or solvent molecule) or by solvation in a medium of high dielectric constant. This applies t o both soluble metal and base catalysis and is illustrated by the following examples:

+

+

+ +

H:H H20 + HOH H(Hz0) [CO(CO),]~ Hz + 2HCo(C0)4 HO

Hg++

+ Hz

[

++

-+

Little is known concerning the reduction steps which follow activation of the hydrogen. This lack of knowledge is due to the kinetic situation that the reduction step is more rapid than the hydrogen activation step. For catalytic reduction to occur of a substrate other than the soluble metal which activates the hydrogen, it is necessary that the rate of reaction of the catalyst-H2 complex with the substrate be faster than the rate of the reaction, catalyst-Hz metal H+. The elucidation of the actual reduction step following hydrogen activation remains a future field of great interest. Only a few years ago it appeared that only one case of homogeneous hydrogenation catalysis was known-the cuprous-acetate-in-quinoline system. The uniqueness of this system appeared to define a very special set of chemical and physical circumstances. However, recent searching for this type of catalyst has disclosed a number of active catalysts. It appears possible that many more will be found in the future and that the chemical reactivity of hydrogen a t low temperatures ha5 not been fully appreciated. As new and old systems become better defined, there is high hope that this scientific approach to hydrogenation catalysis will continue to provide critical information of theoretical and, ultimately, of great practical importance. ---f

+

Note added in proof: Halpern and his associates have recently extended their studies of the activation of hydrogen by aqueous solutions of cupri,c, mercuric, mercurous, and silver salts. The catalytic activities of a series of cupric complexes decrease in the following order: butyrate, propionate > acetate > sulfate > chloride > water (perchlorate solution) > glycine,

204

S. W. WELLER AND 0. A. MILLS

ethylene diamine (Peters and Halpern, 48). At 100" the butyrate and propionate are one hundred and fifty times more active than the perchlorate (i.e., the uncomplexed cupric ion). The activation energy for the reaction involving cupric perchlorate is 26.6 kcal./mole, which is 2 kcal./ mole higher than that for cupric acetate (Peters and Halpern, 49). I n all cases the reactions are first order with respect t o concentration of metal salt. Korinek and Halpern (50)report that solutions of mercuric perchlorate are reduced by hydrogen about ten times more rapidly than is mercuric acetate, and t hat the product mercurous perchlorate is itserf reducible t o mercury by a slower, homogeneous reaction with hydrogen. The total rate of reaction of hydrogen is given by the equation

where k, = 4.2 X 10'" exp

and kn

=

1.2

x

10" exp

Silver perchlorate is reduced by hydrogen in a reaction which is second order with respect to silver ion concentration (Webster and Halpern, 51). The activation energy is 15.2 kcal./mole. This is the first demonstrated example of a reduction in aqueous solution which is not first order in metal salt concentration.

R.EFERENCES 1 . Calvin, M., Trans. Faraday SOC.34, 1181 (1938). 2. Wilmarth, W. K., Thesis, University of California, May, 1942.

3. Wilmarth, W. K., and Barsh, M. I<., J . A m . Chem. SOC.76, 2237 (1953). 4. Wilmarth, W. K., Barsh, M. K., and Dharmatti, S. S., J . Am. Chem. Sac. 74,

5035 (1952). Weller, S., and Mills, G. A., J . Am. Chem. SOC. 76, 769 (1953). Wright, L. W., and Weller, S., J . Am. Chem. SOC.76, 3345 (1954). Calvin, M., J. Am. Chem. Soc. 61, 2230 (1939). van Niekerk, J. N., and Schoening, F. R. L., Acta Cryst. 6, 227 (1953). Tyson, G . N., and Vivian, R. E., J . Am. Chem. SOC.63, 1403 (1941). 10. Wright, L. W., Weller, S., and Mills, G. A., J . Phys. Chem. 69, 1060 (1955). 1 1 . Dakers, R. G., and Halpern, J., Can. J . Chem. 32, 969 (1954). 12. Ipatieff, V. N., and Werchowsky, W., Ber. 42, 2078 (1909). 19. Ipatieff, V. N., Corson, B. B., and Kurbatov, I. D., J . Phys. Chem. 43,589 (1939). 14. Forward, F. A., Bull. Inst. Metals 2, 113 (1954). 16. Forward, F. A., and Halpern, J., Trans. Can. Znst. Mining Mel. 66, 355 (1953). 16. Halpern, J., and Dakers, R. G., J . Chem. Phys. 22, 1272 (1954). 17. Peters, E., and Halpern, J., Can. J . Chern. 33, 356 (1955). 18. Halpern, J., Korinek, G. J., and Peters, E., Research 7, 561 (1954). 19. Flynn, J. H., and Hulburt, H. M., J. Am. Chem. Sac. 76, 3393 (1954). 20. Flynn, J. H., and Hulburt, H. M., J . Am. Chem. SOC.76, 3396 (1954). 21. Chatt, J., J . Chem. Soc. p. 2622 (1952). 6. 6. '7. 8. 9.

ACTIVATION O F MOLECULAR HYDROGEN

205

2.2'. British Intelligence Objectives Sub-committee Misc. Repi. 113 ; FIAT Final Rept.

1000. 23. Adkins, H., and Iirsek, G., J . Am. Chem. SOC.71, 3051 (1949). 24. Wender, I., Levine, R., and Orchin, M., J . Am. Chem. SOC.72, 4375 (1950). 26. Wender, I., Orchin, M., and Storch, H. H., J . Am. Chem. SOC.72, 4842 (1950). 26. Wender, I., private communication. 87. Wirtz, K., and Bonhoeffer, K. F., 2.physik. Chem. A177, 1 (1936). 28. Abe, S., Sci. Papers Znst. Phys. Chem. Research (Tokyo) 38, 287 (1941). 39. Claeys, Y. M., Baes, C. F., and Wilmarth, W. K., J. Chem. Phys. 16, 425 (1948). 30. Claeys, Y. M., Dayton, J. C., and Wilmarth, W. K., J . Chem. Phys. 18,759 (1950). 31. Wilmarth, W. K., Dayton, J. C., and Flournoy, J. M., J . Am. Chem. SOC.76,4549

(1953). 3.2'. Wilmarth, W. K., and Dayton, J. C., J . Am. Chem. Soe. 76, 4553 (1953). 33. Selected Values of Chemical Thermodynamic Properties, Natl. Bur. Standards

Circ. 600 (1952). 34. Mills, G. A., Weller, S., and Wheeler, A., unpublished results. 36, Gilman, H., Jacoby, A. L., and Ludeman, H., J . Am. Chem. SOC.60, 2336 (1938). 36. Stephenson, M., and Stickland, L. H., Biochem. J . 26, 205, 215 (1931). 37. Green, D. E., and Stickland, L. H., Biochem. J . 28, 898 (1934). 38. Cavanagh, B., Horiuti, J., and Polanyi, M., Nature 133, 797 (1934). 39. Hoberman, H. D., and Rittenberg, D., J . Biol. Chem. 147, 211 (1943). 40. Bovarnik, M., Proc. SOC.Exptl. Bid. Me&.47, 191 (1941). 41. Back, K. J. C., Lascelles, J., and Still, J. L., Australian J. Sci. Research 9, 25 (1946). 42. Joklik, W. K., Australian J . Exptl. Biol. Med. Sci. 28, 321, 331 (1950). 43. Gest, H., J . Bacterial. 63, 111 (1952). 44. Krasna, A. I., and Rittenberg, D., J . Am. Chem. SOC.76, 3015 (1954). 46. Farkas, A., and Farkas, L., J. Chem. Phys. 2, 468 (1934). 46. Ogg, R. A., Jr., Am. Chem. SOC.Abstracts 24P (1953). 47. Dowden, D. A., J. Chem. SOC.p. 242 (1950). 48. Peters, E., and Halpern, J., Can. J . Chem., in press. 49. Peters, E., and Halpern, J., J . Phys. Chem. 69, 793 (1955). 60. Korinek, G. J., and Halpern, J., J . Phys. Chem., in press. 61. Webster, A. H., and Halpern, J., J . Phys. Chem., in press.

This Page Intentionally Left Blank

Catalytic Syntheses of Ketones V. I. KOMAREWSKY AND J. R. COLEY Catalysis Laboratory, Illinois Institute of Technology, Chicago, Illinois Page I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Development of Ketonization Process. . . . . . . . . . . . . . 1. Dehydrogenation of Secondary Alcohols. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 208 2. Decarboxylation Condensation of Acids. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 208 3. Tischenko Condensation Synthesis . . . . . . . . . . . . . . . . . . . 209

. . . . . . . . . . . . . 216 References. .......................

............................

216

I. INTRODUCTION Heterogeneous catalysis has played an important role in the synthesis of aliphatic oxygenated compounds. The present article is devoted specifically to the preparation of ketones and particularly to the work carried out by the authors in the catalysis laboratory of Illinois Institute of Technology. The preparation of ketones by dehydrogenation of secondary alcohols over zinc and copper catalysts and the decarboxylation condensation of acids over manganous oxide or thoria have been adequately covered by standard reference books on catalysis. However, the more complete but equally serviceable catalytic syntheses involving either an aldol or a “Tischenko ” ester type of condensation have been virtually ignored. Primary alcohols of n-carbon atoms are readily converted to symmetrical ketones of 2n - 1 carbon atoms by vapor-phase contact with various catalysts. A synthesis involving a Tischenko condensation (1) is a variation of an ester synthesis used commercially in Russia (2). In the conversion of ethyl alcohol to acetone by reaction with steam, ethylacetate is considered an important intermediate in the chain of reactions involved. In the presence of chromia catalyst (3) primary alcohols of n-carbon 207

208

V. I. KOMAREWSKY

atoms are converted t o 2 n - 1 symmetrical and mixed ketones by a different mechanism, as follows: 1. Dehydrogenation of alcohol to aldehyde. 2. Aldol condensation of the aldehyde. 3. Removal of carbon monoxide from the -CHO group of the aldol leaving a secondary alcohol. 4. Dehydrogenation of this alcohol t o ketone. Both of these dehydrogenation-condensation ketone syntheses are equally applicable to numerous members of the primary alcohol series.

11. DEVELOPMENT OF KETONIZATION PROCESS 1. Dehydrogenation of Secondary Alcohols

The dehydrogenation of alcohols was first studied by Ipatieff, who obtained the corresponding aldehydes or ketones by treatment of methyl, ethyl, isopropyl, isobutyl, and isoamyl alcohols with such catalysts as a platinum tube, zinc rods, and brass a t suitable temperatures. The work of Sabatier and Senderens and later Constable and Palmer added t o the understanding of this industrially important reaction.

6. Decarboxylation Condensation of Acids Squibb (4) transformed the well-known method of preparing ketones by calcination of calcium or barium salts of monobasic organic acids into a catalytic method. Acetic acid vapors on vapor-phase contact with barium carbonate a t 500' undergo a continuous decomposition into acetone, water, and carbon dioxide with yield of 90% or better. Investigation of this reaction revealed th at acetic acid and barium carbonate react a t 400" t o yield barium acetate with the liberation of carbon dioxide and water. If the acid flow was discontinued and the temperature raised to 500', the acetate was decomposed and the carbonate and liberating acetone were regenerated. This is a good example of a catalytic reaction involving a n unstable intermediate. This intermediate formation ceased t o be apparent when the acid was passed over the oxide a t a higher temperature, because the formation of the salt was then balanced by its rapid destruction. For certain oxides, like thoria and titania, the intermediate cannot even be perceived, as doubtless the formation does not take place a t a lower temperature than the decomposition, but the analogy is so close that we cannot fail t o assume similar mechanisms with all the oxides. This ketone synthesis was developed further by Ipatieff and by Sabatier, Mailhe, and Senderens. Thoria and manganous oxide proved to be more effective catalysts than the carbonates. The method was extended t o higher fatty acids and to the production of mixed ketones.

209

CATALYTIC S Y N T H E S E S O F K E T O N E S

Vapor-phase ketonization of higher fatty acids is possible a t reduced pressures, but undesirable side reactions make the liquid-phase reaction at the milder conditions more attractive commercially. 3. Tisehenko Condensation Synthesis

The first direct conversion of a primary alcohoI to a ketone was reported by Donath (ii) in 1889. The mechanism of the conversion of ethyl alcohol into acetone by reaction with steam has been examined by Kagan and his co-workers ( I ) , who consider th a t ethyl acetate is an important intermediate in the chain of reactions involved. The mechanism consists of the following steps :

4CzH6OH + 4CH3CHO

+ 2Hz

2H20

4CH3CHO 4 ~ C H ~ C O O C Z H ~ZCH3COOH 2CHsCOOH + CHaCOCH3 COz HzO

+

+

+ 2CzH6OH

The over-all reaction may be expressed by the equation

2CzHsOH

+ HzO + CHsCOCH3 + COz + 4Hz

The oxides of zinc, cadmium, manganese, nickel, cobalt, and chromium and their mixtures are satisfactory catalysts. I n a later work Dolgov and Golodnikov (6) developed a n activated copper catalyst and produced a mixture of esters and ketones from alcohol. The reactions proceed by ester mechanism, and a t lower temperatures (275"-300") the formation of esters predominates. This ketone synthesis is equally applicable to higher members of the primary alcohol series.

4. Aldol Condensation Synthesis a. Development. The conversion of primary alcohols to ketones without any participation of water in the reaction was first reported by the authors of this article (3). This discovery was made during a n extended study of the action of the metallic oxide catalysts on oxygenated compounds. Primary alcohols ranging from n-propyl t o n-octadecyl alcohol of n carbon atoms were converted by vapor-phase contact with chromium oxide at 400" t o symmetrical ketones with 2n - 1 carbon atoms. The molal conversion of reacted alcohol to ketone exceeded 50% in most cases (Table I). Subsequent analysis of the noncondensable product gas indicated th a t the gas consisted largely of hydrogen and carbon monoxide present in a ratio of three volumes of hydrogen to one of carbon monoxide. The over-all reaction may be expressed b y the equation

2RCH2CHzOH + RCHzCOCHzR

+ CO + 3Hz

210

V. I. KOMAREWSKY

As would be expected from the Le Chatelier principle, the use of reduced pressure gave improved yields. n-Octyl alcohol at 125 t o 135 mm. pressure yielded 73.9% di-n-heptyl ketone compared with 56.0% conversion obtained at atmospheric pressure. TABLE I

Reactant charged n-Propyl alcohol n-Butyl alcohol n-Amy1 alcohol

n-Hexyl alcohol n-Heptyl alcohol n-Octyl alcohol n-Decyl alcohol n-Octadecyl alcohol

Temp., "C.

Molal conversion of reacted alcohol to ketone, %

Unreacted alcohol, per cent

425 400 425 375 425 425 425 400 400 400

48.8 27.8 46.9 55.5 47.0 46.5 57.5 56.0 83.2 47.6

5.2 28.3 12.2 8.9 5.5 16.7 11.6 13.3 3.5 19.8

This ketone synthesis method is equally serviceable for the preparation of unsymmetrical ketone and is recommended in particular for the preparation of unsymmetrical methyl ketones, as no acetone is formed. Typical yields of unsymmetrical ketones are shown in Table 11. TABLE 11 Temp., "C.

Reactant charged

Molal conversion of reacted alcohol to ketone, %

Unreacted alcohol, per cent

Ethyl alcohol, 25% by vol. n-Octyl alcohol, 75% by vol.

}

425

CHaCOCHs 0 . 0 CzHrOH 0 . 0 C H S C O C ~ H I ~41.7 CaH17OH 5 . 9 C ~ H ~ ~ C O C26.2 ~HL~

n-Amy1 alcohol, 50% by vol. n-Decyl alcohol, 50% by vol.

I

400

CaHgCOCaHs 13.5 CpHsCOCaHis 2 7 . 2 CgHigCOCaHig 20.9

CsHiiOH 2 . 4 CioHirOH 1 . 5

I n a preliminary consideration of the reaction mechanism, aldehydes and aIdols seem t o be logical intermediates. It was found that aldehydes undergo the same type of condensation, producing ketones in yields considerably superior to those obtained from the corresponding alcohol. When the aldols of n-butylaldehyde and n-heptaldehyde were subjected to the same reaction conditions, excellent yields of ketones were produced

21 1

CATALYTIC SYNTHESES OF KETONES

directly by a catalytic decomposition reaction. The order of conversion of the reactants to ketones was aldol > aldehyde > dcohol. These data are shown in Table 111. TABLE I11

Reactant charged n-Butyl alcohol n-Butyraldehyde Aldol of n-butyraldehyde n-Heptyl alcohol n-Heptaldehyde Aldol of n-heptaldehyde

Temp., "C.

Molal conversion of reacted reactant to ketone, %

Unreacted reactant, per cent

400 400 400 425 400 400

27.8 31.4 61.0 57.5 66.9 82.8

28.3 23.6 0.0 11.6 2.5 0.0

One exception to the genera1 application of these ketone syntheses was failure of compounds having an alpha-substituted carbon atom such as isobutyl alcohol or 2-ethylhexanol to undergo the dehydrogenation (7') condensation reaction. This failure of alpha-subst,ituted reactants to undergo the ketone synthesis was unexpected as the aldol condensation of alpha-substituted aldehydes with one labile hydrogen atom occurs readily. TABLE IV

Reactant Aldol of isobutyraldehyde (CH,),CHCH(OH)C(CHa) &HO Mixed aldol of acetaldehyde and isobutyraldehyde CH&H(OH)C(CHa)zCHO Aldol of n-butyraldehyde C~H&HZCH(OH)CH(CZH~)CHO Aldol of n-heptaldehyde C,H~,CH&H(OH)CH(C,HII)CHO

Ketone

Molal conversion of reacted reactant to ketone, %

[(CHa)zCHIzCO

6.5

CHaCOCH(CH3)z

12.0

(C3H7)zC0

61 .O

(c6 H13)zCO

82.8

An investigation of the catalytic decomposition of aldols of several alphasubstituted aldehydes supplied the answer to this irregularity (8). The results summarized in Table I V indicate that the major decomposition product of aldols containing no labile hydrogen atoms was the original aldehyde; whereas the aldols with labile hydrogen atoms underwent a decarbonylation-dehydrogenationreaction to ketones. The step blocking the conversion of alpha-substituted aldehydes was not the primary

212

V. I. KOMAREWSKY

condensation step, but the secondary decarbonylation-dehydrogenation step. It was noted that sodium hydroxide-precipitated catalysts were quite active in contrast with comparatively inactive ammonia-precipitated catalysts. Investigation of a series of catalysts with varying sodium content proved that the sodium content of a chromium oxide catalyst has a

SODIUM IN CATALYST,

FIG.1. The effect of sodium oxide on the conversion of n-butyraldehyde t o di-npropyl ketone.

definite effect on the catalytic activity for ketone synthesis (9). As shown in Fig. 1, the optimum sodium content was found t o be approximately 1.0%. b. Reaction Mechanism. The aldol mechanism proposed for the dehydrogenation-condensation reaction is a logical choice and is supported by the experimental data. However, in addition t o the aldol mechanism, another equally plausible mechanism must be considered. This mechanism involves a Cannizzaro or Tischenko condensation. A tracer technique was eniployed in an effort t o gain further information about the reaction mechanism (20).One alpha hydrogen atom of n-octyl alcohol was replaced with a n atom of deuterium. The steps of the two reaction mechanisms under discussion are presented in Fig. 2 . If the exchange between hydrogen and deuterium is small, the presence of significant quantities of deuterium in the gaseous product may be explained only by mechanism (1). The exact experimental details of this investigation will now be discussed.

(1) ALDOL MECHANISM

D 2R-

c:I

-CHzOH+

’

H

c:& -

2R-

&I b = O

4R-A-L

B

=

0

R-

&-&-6-‘ A JH

R-

1

- -R

R-LC-

1

0

+ 2Hz

D D -COOCH%- -R

c:I

LI

+ 2R-

A

2R-A-COOH

I H

+ R-

-R

c:I

-CHzOH

-rMm

U

I

II I

-R

+ COZ+HsO

H O H

D

LI

z

4

H

1

II

Tz1

D

+ CO

b* c3

D

fc:-COOH

%

5

H

H

2R-

=O

d1

LI-Ll Ll

D

+ HD

B

=

D

H OH H

H OD H

-

12H20

D H D

+ CO

A(!!

4R-

0 ---* 2R-

1

D H H

H O H

-+

D H R H =

II I

A-CHzOH

D H

H

D H R H

AI -bl -c l:-R

4R-

(2) CANNIZZARO MECHANISM D H

HI

R--

I

1

I

+ 2H2

C:b =O I

+

H

R-

0

D H

R--

ob 1

=

€L

D H

R-(!-(!-(!d

D

D H

+ Hz

H O H

FIG. 2. The effect of sodium oxide on the conversion of n-butyraldehyde to di-n-propyl ketone.

E Fi

214

V.

I. KOMAREWSKY

( 1 ) Preparation of l-Octanol-24, CHa(CHJ&HDCHzOH: A mixture of 460 g. (4.1 moles) of redistilled octane-1 and 20 g. of benzoyl peroxide was placed in a l-liter round-bottomed flask fitted with a motor-driven glass stirrer, an inlet tube extending to the bottom of the flask, and a condenser cooled with dry ice connected to a liquid-air-cooled trap. The flask was immersed in an ice bath. DBr, made in an HBr generator (11) by passing deuterium gas (1 mole) through bromine, was passed through the octene-1 over a period of 8 hr. The unreacted DBr, which collected in the liquid-air trap, was repassed through the octene-1. Undissolved benzoyl peroxide was filtered off, and the product, after being washed with a ferrous sulfate solution and water and then dried over anhydrous sodium sulfate, weighed 514 g. Distillation in a packed column gave 87 g. (0.45 mole) of CH3(CH2)&HDCH2Br,n: 1.4519. The deuterated bromide was then hydrolyzed in the following manner: 78 g. (0.4 mole) of CH3(CH2)&HDCH2Brand 200 ml. of 8% NaOH were placed in a 500-ml. stirring autoclave, which was heated to 165" for 1 hr. The product was washed free of alkali, dried over anhydrous sodium sulfate, and distilled in a packed column to give 37.4 g. of alcohol, b.p. 73"75" at 3.5 mm., nio 1.4305. This yield corresponds to a conversion of 72%. Mass spectrometric and infrared analysis of the deuterated alcohol indicated that the synthesis of the l-octanol-2-d was successful. (2) Preparation of l-Octanol-d, CHB(CHZ)&H*OD: A mixture of 4.87 g. (0.0375 mole) of octanol-1, 15.38 g. (0.69 mole) of D 2 0 (99.8% purity), and 0.1 g. of anhydrous potassium carbonate was placed into a 50-ml. glass-stoppered flask, shaken for 15 min., and allowed to stand for 60 hr. After an additional 1-hr. shaking period, the two phases were separated and the alcohol layer was dried over anhydrous potassium carbonate. The water layer was distilled in a small distilling flask. Density determination of the original and the reacted heavy water indicated that 92 & 4% of the l-octanol had been converted to l-octanol-d. (3) Catalytic conversion of l-Octanol-2-d to ketone: A quantity (29.7 ml.) of l-octanol-2-d was charged (space velocity 0.2) to a 5-mm. reactor tube containing 18 ml. of 8- to 10-mesh chromium oxide catalyst maintained a t 400". The 27.7 ml. of liquid product was fractionated in a concentric-tube column. A 60.6% yield of di-n-heptyl ketone was obtained. Approximately 17% of the alcohol was recovered*unconverted. Mass spectrometric analysis of the gaseous product showed the atomic ratio of deuterium to hydrogen to be 0.106. The molal yields of deuterium, hydrogen, and carbon monoxide produced per mole of ketone were 0.216, 2.040, and 0.815 respectively. The unconverted alcohol recovered from the reaction product was examined closely for evidence of hydrogen-deuterium exchange. Infrared

CATALYTIC SYNTHESES OF KETONES

215

analysis was very well suited for this work as the O-D and C-D stretching vibrations occur in a region of the spectrum that is usually free of other interfering absorptions. I

I

o-n c-n O-D

I

I

6

6

c-o

E

tn 2 W

I-

E

3

4

4

VYVE LENQTH IN MICRONS

Fro. 3. Infrared spectra of octanols.

A Beckman IR-2 infrared spectrometer equipped with a sodium chloride prism was used. The wave-length scale was calibrated against known absorption maxima of liquid toluene and of atmospheric water vapor and carbon dioxide. Wave lengths are accurate t o k-0.02 p. The infrared absorption spectra of four octyl alcohols, using a cell length of 0.1mm., are shown in Fig. 3: (a) I-octanol, (h) 9.2 vol. % sohtion of l-octanol-d in I-octanol, (c) 1 octanol-2-d, (d) 1-octanol-2-d exposed t o ketone-synthesis conditions.

216

V. I. KOMAREWSKY

A comparison of the C-D and O-D band intensities obtained for 1-octanol-2-d exposed to synthesis conditions (curve d ) with reference to curves (c) and ( b ) indicated that the exposed alcohol consisted of 80 vol. % l-octanol-2-d, 7.4 vol. % l-octanol-d, and minor quantities of impurities containing a carbonyl band. Di-n-heptyl ketone and the di-n-heptyl ketone prepared catalytically from 1-octanol-2-d were examined in carbon tetrachloride solutions with a cell having a path length of 1.0 mm. The ketone prepared from l-octanol-2-d showed an absorption band characteristic of C-D a t 4.63 of diheptyl ketone. The number of C-D bonds per ketone molecule was estimated to be in the range of 1.0 to 1.5. A more accurate estimate of the C-D content of the ketone would require a reference sample of deuterated ketone, which was not available. c. Discussion of Results. The extent of O-D bond formation in unconverted l-octanol-24, which was exposed t o ketone synthesis conditions, was a good measure of the magnitude of hydrogen-deuterium exchange since the hydroxyl hydrogen is the most vulnerable of any hydrogen atom in the alcohol molecule (12)to H-D exchange. Since 80% of the exposed l-octanol-2-d was unchanged and only 7.4% l-octanol-d was formed, hydrogen-deuterium exchange occurred t o only a minor extent a t ketonesynthesis conditions. The experimental atomic ratio of deuterium to hydrogen of 0.106 corresponds closely to the ratio calculated for the aldol mechanism (0.091) on the supposition that deuterium and hydrogen atoms in the alpha position participate equally in the condensation. The calculated molal yields of deuterium, hydrogen, and carbon monoxide produced per mole of ketone (0.250D2, 2.75H2, and l.OC0) d o not differ appreciably from the experimental values 0.216, 2.04, and 0.815. The estimated C-D content of the ketone produced from l-octanol-2-d is consistent with the proposed aldol mechanism. I n view of the indicated absence of any marked hydrogen-deuterium exchange under ketone-synthesis conditions, the presence of substantially the theoretical atomic ratio of deuterium to hydrogen in the gaseous product provides additional support for the proposed aldol mechanism. 111. SUMMARY I n the presence of chromium oxide catalyst primary aliphatic alcohols of n-carbon atoms are converted t o symmetrical ketones of 2n - 1 carbon atoms. The reaction proceeds by an intermediate aldol formation.

REFERENCES 1. Kagan, M. Y . , Sobolev, I. A., and Lybarskii, G.D., Ber. 08, 1140 (1935). 8. Dolgov, B. N., Koton, M. M., and Lelchuk, S. L.,OTQ.Chem. Ind. (U.S.S.R.) 1,

CATALYTIC SYNTHESES O F KETONES

217

70 (1936); Ivannikov, P. Y., Tsygankova, M., and Gavrilova, E. Y., ibid. 6, 53 (1938); Veltistova, M. V., and Lelchuk, S. L., ibid. 6,657 (1939); Ivannikov, P. Y., J . Appl. Chem. (U.S.S.R.) 13, 118 (1940). 3. Komarewsky, V. I., and Coley, J. R., J . Am. Chem. SOC.63, 700, 3269 (1941). 4. Squibb, E. R., J . Am. Chem. SOC.17, 187 (1895); ibid. 18, 231 (1896). 6. Donath, E., Chem. Ztg. 12, 1191 (1889). 6. Dolgov, B. N., and Golodnikov, G. V., J . Gen. Chem. (U.S.S.R.) 24, 987, 1167, 1364 (1954). 7. Komarewsky, V. I., and Smith, L. G., J . Am. Chem. SOC.66, 1116 (1944). 8. Coley, J. R., and Komarewsky, V. I., J . Am. Chem. SOC. 68, 716 (1946). 9. Coley, J. R., and Komareweky, V. I., J . Am. Chern. Soc. 74, 4448 (1952). 10. Komarewsky, V. I., Zimmerschied, W. J., and Coley, J. R., “Farkas Memorial Volume,” p. 181. Research Council, Jerusalem, 1952. 11. Ruhoff, J. R., Burnett, R. E., and Reid, E. E., “Organic Synthesis,” Vol. 2, p. 338. Wiley, New York, 1943. 12. Anderson, L. C., and MacNaughton, N. W., J . Am. Chem. SOC.64, 1456 (1942); Farkas, A., and Farkas, L., ibid. 61, 1336 (1936).

This Page Intentionally Left Blank

Polymerization of Olefins from Cracked Gases EDWIN K . JONES Universal Oil Products Co., Des Plaines, Illinois Page

11. Sources of Feed Stock.. . . . . . . . . . . . . . . . . . . . .................... 111. Basic Factors in Polymerization. . . . . . . . . . . . . . . . . . . . 1. Process Variables.. . . . . . . . . . ..........................

220 222

. . . . . . . . . . . . . 223

3. Propylene Motor Polymer., . . . . . . . . . . . . . . . . . . . . 4. Petrochemical Reactions.. . . . .............................

d. Heptene. . . . . e. Dimer.. . . . . . . . . . . . . . . . .

..........................

230

236

1. U.O.P. Solid Phosphoric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 236 2. Solid Copper Pyrophosphate of the Polyco Type.. . . . . . . . . . . . . . . . . . . . 236 Research Type . , . . 237 3. Liquid Phosphoric Acid-on-Q VI. Materials of Construction.. . . . . . . . . . . . . . . . . . . . . .237 VII. Future Outlook for Polymerisat References. . . . . . . . . . . . . . . . . . . .

I. INTRODUCTION Catalytic polymerization is a leading process of the refining industry in the production of high-quality motor fuels and certain petrochemicals. The Solid Phosphoric Acid catalytic process was developed by Vladimir Ipatieff and has since been extensively modified and improved. By this method the catalyst is loaded in a chamber or placed in tubes and the heated charge is passed through it. The olefinic feed stock generally is diluted with low-olefin recycle. The first commercial unit employing Solid Phosphoric Acid was built to polymerize catalytically propylene and butylenes from thermal cracking into motor gasoline. The product has a research method (F-1) octane 219

220

EDWIN K. JONES

rating of 100 after tetraethyl lead is added. This is exceptionally good in comparison with thermal gasoline, and the so-called “poly gas” is today considered one of the better motor gasolines available. The “solid” catalyst is a calcined mixture of phosphoric acid and kieselguhr. The kieselguhr acts as a catalyst support but it also enters into chemical combination with the phosphoric acid. Although several other catalysts were developed earlier, and some subsequently, the acidand-kieselguhr solid catalyst excels all others. The first polymerization units were designed for low-pressure operation, a condition which rather quickly deposited considerable quantities of heavy polymer, or tar, which had to be removed from the product, and also laid carbon, or “coke,” on the catalyst. The latter in particular made i t necessary to install regeneration facilities to burn the coke from the catalyst in order to secure reasonable operating economy. Experience showed that higher pressures and better temperature control retarded the production of coke and extended the catalyst life, so that the use of catalyst-regeneration facilities was no longer economically justified. The catalyst was placed in tubes surrounded by a cooling-water jacket in the first units designed for more precise temperature control. It also was found t ha t temperature control within the required precision limits could be obtained in chamber-type units by separating the catalyst into several beds and employing a quench of effluent hydrocarbon low in olefins between the beds. The polymerization units have also been used t o produce large quantities of high-octane aviation gasoline for the armed forces. Isopropylbenzene, or “ cumene,” was required t o upgrade available lower octane-number stocks. Many motor polymerization units were converted t o the production of cumene, propylene and benzene being used as feed stocks. Other polymerization units were converted to the production of impure iso-octene, or ‘‘ codimer,” the charge stock being mixed butylenes. The codimer was shipped to centrally located hydrogenation units for conversion into a n aviation gasoline containing a high proportion of iso-octanes. The recent use of large quantities of synthetic detergents has created a n increased demand for dodecene, or “tetramer.” This is currently being produced by catalytically polymerizing four or five molecules of propylene t o form a Cl2-CIsolefinic mixture.

11. SOURCES OF FEEDSTOCK Propane propylene, butane butylene and ethane ethylene comprise the main feed stocks t o the polymerization process. Benzene is charged in addition when cumene or ethylbenzene is produced. Thermal- and

221

POLYMERIZATION O F OLEFINS

catalytic-cracking, thermal-reforming, and steam-cracking units are the essential sources of the olefinic stocks. The feed to the polymerization unit should be low in sulfur and basic nitrogen compounds, as the sulfur in the feed forms sulfur compounds in the product and the basic nitrogen poisons the catalyst. When the hydrogen sulfide content of the polymerization-unit feed stock is high, it can be removed by a regenerative type of amine scrubber. However, a well-designed gas-recovery unit on the cracking or reforming unit will normally cause the hydrogen sulfide content t o be low enough to require only a simple caustic scrubber for removal. When the mercaptan content TABLE I C S C ~Charge to Motor Polymer Unit and Hydrocarbons Rejected in AbsoTber Gas ~

_

_

_

_

Absorber gas, mole % Hydrogen Methane Ethylene Ethane Propylene Propane Butadiene Butylene Isobutane Normal butane Pentanes

20.5 38.5 13.5 20.6 3.9 1.0

0.4

-

1.4 0.2 100.0

_

C8-C4poly feed, mole %

Trace 0.1 18.8 17.8 0.1 30.4 19.7 13.0 0.1

100.0

of the feed to the polymerization unit is sufficiently high to give excessive sulfur in the product, the mercaptan can be removed by a simple regenerative caustic scrub. Basic nitrogen is removed by adding a countercurrent water wash on the polymerization feed. Table I shows the products from a well-designed gas-recovery unit in a typical refinery having a catalytic-cracking unit and a thermal-cracking unit. Where only the propane propylene is charged to the polymerization unit a depropanizer is added to separate the Ct and lighter from the Cc and heavier, shown in the last column of the table.

111. BASICFACTORS IN POLYMERIZATION There are a number of basic factors to be considered in polymerization in addition to process variables, types of catalyst, and catalyst poisons. Possibly the most important is that essentially all polymerization

222

EDWIN K. JONES

processes used today employ a catalyst, although noncatalytic polymerization of propylene and butylenes occurs to some extent and units have been built to take advantage of this noncatalytic reaction. The polymerization process is exothermic. The average heat of reaction when motor polymer is made from propylene is 670 B.t.u./lb. I n contrast, the average heat of reaction of butylenes t o motor polymer is only 400 B.t.u./lb. It is easier, therefore, to control the temperature of a butylene polymerization unit than to control propylene polymerization. Propylene polymers are relatively stable, but the butylene polymers can be broken down under many conditions. Such a breakdown occurs when butylene polymers are charged to a n alkylation unit, when the individual butylene molecules react with the isobutane. The stability of the propylene polymers is one of the reasons why they are preferred as starting materials for such reactions as the synthesis of dodecyl benzene in the ma,nufacture of detergents. Propylene normally is more difficult t o polymerize than the average mixture of butylenes. However, when propylene is polymerized in conjunction with butylenes, the speed of reaction is reversed and the propylene reaches the higher conversion rate. 1, Process Variables

The important process variables in the polymerization of olefins from cracked gases are time, catalyst activity, pressure, hydration, and temperature. All these variables are related t o some extent and can be interchanged within certain narrow limits. Insufficient reaction time will give low conversion rates for propylene and butylenes charged t o the catalyst. Conversely, excessive time will produce polymer of a higher boiling range than gasoline has. Motor polymer gasoline made under good operating conditions normally will have an A.S.T.M. Engler distillation end point between 410' and 420'F. The reaction time is normally measured as space velocity and referred to in units of gallons of combined feed per hour per pound of catalyst. The various processes require space velocities all the way from 0.12 t o 0.46 and must be controlled very closely in order t o obtain the proper results. Figure 4 shows the space velocities required for any given olefin conversion when making C3-C4motor polymer. While older units required high recycle rates, increased efficiency of recycle and quench has allowed the recycle rate to be lowered considerably on new units. Lower recycle rates require smaller catalyst volumes, lower utility consumptions, and smaller equipment sizes. Reaction time is interrelated with catalyst activity and, within limits, to operating pressure. The Solid Phosphoric Acid catalyst is the most ac-

POLYMERIZATION OF OLEFINS

223

tive commercially employed for polymerization. Kieselguhr has the greatest active surface of materials used to support the catalyst. Increasing the pressure in the reaction zone increases the residence time by increasing the density of the hydrocarbon phase. This densephase hydrocarbon keeps the heavy polymer or tar washed off the catalyst and prolongs catalyst life. The optimum pressure depends upon the composition of the feed stock. Pressures of between 400 and 1000 p.s.i. are normally considered best for commercial operation. Reaction temperature and time are interchangeable between very narrow limits. Although the limits appear to be quite wide in actual operation, the fact that the higher temperatures tend to produce a tar coating on the catalyst and to reduce its activity makes raising temperature of doubtful benefit. Normally this tar is produced slowly, but in time it will reduce the activity of the catalyst to such an extent that the average polymer production for the full run is sometimes less when the unit is operated a t the high end of the temperature range. If shorter catalyst life can be tolerated, on the other hand, the catalyst can be changed more often to maintain the catalyst activity and thereby to eliminate part of the low-conversion period a t the end of the run usually experienced with operation at the higher temperatures. The best temperatures for good operation are between 350" and 435°F. Both catalyst activity and tar formation are directly affected by the state of hydration of the phosphoric acid-kieselguhr type of catalyst. At the higher temperature it is more difficult to maintain proper hydration. Hydration control is required because the catalyst has an optimum water content which determines the activity and selectivity of the catalyst. The water-vapor pressure varies at different catalyst temperatures and it is important to keep the water content of the hydrocarbon in equilibrium with that of the catalyst. I n those units where water of saturation in the feed is insufficient, additional water must be injected into the feed as catalyst requirements dictate. The solid phosphoric acid type of catalyst contains the proper amount of water when manufactured and the art of catalyst hydration has reached such a point that catalyst in properly operated polymerization units no longer fails from coke formation or loss of activity. 2. Types of Catalyst A widely used commercial catalyst is the extruded pellet type of calcined phosphoric acid and kieselguhr developed by Universal Oil Products Company. It is the most active of the catalysts used in polymerization and has the greatest active surface area. When thermal cracking was the predominant method of making cracked gasolines, a copper pyrophosphate catalyst (I) was used to some

224

EDWIN K. JONES

extent in the polymerization of cracked gases. Although there still are some units using this catalyst, its lower activity for the conversion of propylene and butylene-2, which are made in greater quantities by catalytic-cracking units than by thermal crackers, has caused many of these units t o be converted to the Solid Phosphoric Acid catalyst. Liquid phosphoric acid on various supports has been tried from the beginning of commercial polymerization but none has been so successful a s the solid type of catalyst. However, a commercial-sized unit was built in 1937 t o use liquid phosphoric acid on quartz(2). Several years then were spent in developing the process and some units later were constructed and operated. This process is based on the principle of manufacturing the catalyst in situ by placing the supporting material in the reaction zone, flooding it with liquid phosphoric acid, and then draining off the excess acid prior to the introduction of the olefin feed stock. I n view of the rather small surface area of the catalyst support employed, relatively large reactor sections are required for a given polymer product, Sulfuric acid (3) has been used in the past as a polymerization catalyst both in the Cold Sulfuric Acid Process and the Hot-Acid Polymerization Process. Its main utility today lies in its use for selectively absorbing isobutylene from mixed butane-butylene streams for use in the production of synthetic rubber. The products are a di-isobutylene polymer and a butane-normal butylene mixture. The di-isobutylene is also used t o a small extent by the chemical industry. A silica-alumina catalyst also has been used in several commercial units. It has the disadvantage of polymerizing only a relatively small percentage of the olefins charged. 3. Catalyst Poisons

Essentially all the commercial catalysts for the polymerization of olefins from cracked gases contain phosphoric acid and therefore are poisoned by alkaline materials in the reactor feed. T h e most commonly encountered poisons of this type are ammonia and combined organic nitrogen compounds of a basic nature. Crude oils contain a certain amount of combined nitrogen which sometimes breaks down in thermal crackers to form these harmful nitrogen compounds. California, West Texas, and Venezuelan crudes seem to break down this way much more readily than other crudes. Catalyticcracking units convert the nitrogen compounds in their feed to these polymerization catalyst poisons almost without exception, Other basic materials which have poisoned polymerization catalyst a t times are sodium hydroxide and diethanolamine. Both of these materials are used extensively for the removal of hydrogen sulfide from the feed to polymerization units. Catalyst poisons of a basic nature can be removed from the

225

POLYMERIZATION OF OLEFINS

feed to the polymerization unit by scrubbing the hydrocarbons with water. Although neither oxygen nor butadiene can be technically classified as catalyst poisons, the presence of these compounds in the feed to the polymerization process will have a deleterious effect on the catalyst life. While oxygen does not react with the catalyst, it tends to modify the polymerization reaction to form long-chain polymers boiling much above the gasoline-distillation range. This heavy polymer, or tar, remains on the catalyst, coating the surface and plugging the catalyst voids to such an extent that the pressure-drop buildup requires that the catalyst be Benzene (alternate Water iniection

--

Fresh propylene. propane

Spent C3

Alternate motor polymer to storage

Benzene alternate hght polymer or butane

Cumene bottoms alternate tetramer bottoms

Cumene alternate tetramer

FIG.1. U.O.P. cumene unit-(alternate tetramer or motor polymer).

changed before the end of normal life expectancy. A well-designed catalytic-cracking-unit gas-recovery section should have a stripper for the removal of oxygen from the feed to a polymerization unit. Butadiene overpolymerizes to tar and thus causes pressure drop in the catalyst. However, units have been designed to handle up t o 3% butadiene in the feed with acceptable catalyst life. The copper pyrophosphate catalyst is poisoned by sulfur compounds in addition to those mentioned above.

IV. POLYMERIZATION REACTIONS Polymerization of olefins from cracked gases today covers a broad range of products from motor fuel to petrochemicals. The petrochemical list is expanding rapidly with many of these products being made from propylene. Figure 1 shows a typical chamber type unit for producing the important petrochemicals, tetramer and cumene.

226

EDWIN K. JONES

Although many polymerization units have been built in the past to use gaseous feed, the trend is now toward liquid-feed units. In preparing the liquid feed, both the ethane and hydrogen sulfide are usually reduced considerably, which reduces the polymerization-catalyst and feed-treating requirements.

Catalyst tower

Butane to storage

LPG to storage Polymer gasoline to storage

charge

FIG.2. U.O.P. catalytic polymerization unit-chamber

type.

In charging gas to the polymerization unit, either a one- or two-stage compressor is required to take the gas up to the catalyst-chamber pressure. Units charging the total gas from a refinery usually in the first compression stage take the absorber or receiver gas from about 70 lb. pressure up to the stabilizer pressure, which is about 150 lb. The gas is then ,,Tubular

reactors'

Spent butane Polymer Propane

FIG.3. U.O.P. catalytic polymerization unit-reactor

type.

combined with the stabilizer gas, treated for hydrogen sulfide and basic nitrogen, and compressed in the second stage to 500 lb. Only the second stage of compression is used when only the stabilizer gas is charged. The size of polymerization units varies greatly, the smallest liquidfeed unit having a charge rate of 136 bbl./day and producing 35 bbl./day

227

POLYMERIZATION O F OLEFINS

of 2-lb. R.V.P. polymer. The largest charges 13,900 bbl. of fresh feed/day and produces 5700 bbl. day of polymer. Figure 2 shows a typical C3-C4chamber-type polymerization unit and Fig. 3 shows a typical C&4 reactor-type unit. These two units are described in detail in Sec. V. 1. Propylene-Butylene Motor Polymer This product is the most predominant of those made by polymerization of the olefins from cracked gases. It is a clean-burning fuel having a TABLE I1 Typical CS-CIMotor Polymer Reid vapor pressure, p.s.i. Gravity, "A.P.I. at 60°F. A S.T.M. distillation, "F. I.B.P. 5% 10 30 60

70 90 95

E.P. % Over yo Bottoms % Loss Bromine number Dispersion, A20 Refractive index, n 20/d Peroxide number A.S.T.M. Gum, mg./100 ml. Cu. dish gum, mg./100 ml. Induction period, min. Corrosion number Total sulfur, wt. % Octane rat:ngs F-I, Clear 1.0 ml. TEL/gal. +3.0 ml. TEL/gal. F-2, Clear 1.0 ml. TEL/gal. +3.0 ml. TEL/gal.

+

+

9.2 66.1 96 130 149 201 228 263 348 405 422 96.5 1.0 2.5 134 85.5 1.40959 0.04 0.8 5 720 0 0.01 96.0 98.5 100 82.5 84.8 86.3

leaded research octane number of 100. Properties of a typical Ca-Ca motor polymer are shown in Table 11.Figure 4 shows the space velocities required to make this polymer a t various conversion levels. Figure 5 shows a true-boiling-point curve of a similar polymer.

228

EDWIN K. JONES

Space velocity, gal. feed/hr./lb. catalyst

FIG.4. Olefin conversions for chamber type polymerization, CI-C, feed.

0 Liauid volume, x

FIQ.5. C3-C4polymer.

Most of the hydrogen sulfide and mercaptans in the feed to catalytic polymerization are converted during the reaction to high-boiling mercaptans and will appear in the gasoline. Very little, if any, HsS will remain after passing over the catalyst. Therefore, the sulfur content of the feed should be kept as low as practical in order to avoid a polymer excessively high in sulfur. High-sulfur polymer not only has a low tetraethyl lead

TABLE 111 Blending Value of Polymer in Straight Run at Various Percentages of Sulfur in Polymer and Straight Run

S.R.

Pennsylvania Mid-Continent

California Mid-Continent West Texas

Poly

F-1 Blending Octane

% S.R.

% Poly

%S

F-1 C1

%S

F-1 C1

90 100 90 80 70 60 0 90 80 70 60 0 90 0 90 90 80 70 95 90 85 95 90 85

10 0 10 20 30 40 100 10 20

0.047 0 0 0 0 0 0 0 0 0 0 0 0.04 0.04 0.04 0.02 0.02 0.02 0.05 0.05 0.05 0.03 0.03 0.03

53.3 53.3 60 60 60 60 60 70 70 70 70 70 69.3 69.3 69.3 46.2 46.2 46.2 63.8 63.8 63.8 66.6 66.6 66.6

0.13 0.04 0.04 0.04 . 0.04 0.04 0.04 0.04 0.04 0.04 0.04 0.04 0.55 0.55 0.04 1.0266 1,0266 1,0266 0.02 0.02 0.02 0.02 0.02 0.02

82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5 82.5

30 40

100 10 100 10 10 20 30 5 10 15 5 10 15

Clear

3 CC.

F-2 Blending Octane Clear

112

134

119

106

123

117

124 124 117 108 98 103

114 120 117 126 117 119

3 cc.

125 110 102 96 82.2 99 94 91 89 82.2 98

102 98 92 92 84.9 87 87 86 86 84.9 72

100 134 119 112 106 105 97 78 80 89.5

87 111 107 106 130 117 107 108 103 101

w

0

0

q

N N

W

230

EDWIN K. J O N E S

susceptibility, but also will impair the lead susceptibility of any lowsulfur gasolines with which it might be blended. Table I11 shows the blending octanes of polymers having various sulfur contents when blended with a variety of straight-run gasolines. Hydrogen sulfide may be removed from polymerization feed stock by scrubbing with ethanolamine or sodium hydroxide, When the mercaptan content of the feed is sufficiently high to give a Doctor-sour polymer, a regenerative caustic wash usually is inserted in the feed-preparation train t o remove them. Mercaptan scrubbers ordinarily are not necessary on polymerization feed streams from catalytic-cracking units, but are required on feeds produced from sour crudes by thermal cracking. 2. Butylene Motor Polymer and Iso-Octane

When there was a shortage of aviation gasoline, many cat-poly units were converted to codimer production. Codimer was made by carefully segregating the butane butylenes, or the so-called “B-B cut,” for polymerization and controlling the ratio of n-butylene t o i-butylene in reactor feed. A few units still operate in this manner. The resulting polymer was shipped to central points for hydrogenation and rerunning. The final product averaged 160 F-4 performance number and 106 F-3 octane number with 4.6 cc. TEL/gal. and was blended with other aviation-gasoline base stocks of lower performance number to produce the desired fuel. 3 . Propylene Motor Polymer

Some propylene polymerization units are operating for production of motor polymer only. This now is becoming uncommon because many refiners have changed their operations to make tetramer, and others fractionate some tetramer, trimer, or dimer out of the polymer leaving very little, if any, motor gasoline. Table IV shows the properties of a typical CS motor polymer of fullboiling range. It is somewhat heavier than the Ca-Cd polymer and the octane rating is possibly a one-half number lower on the Cp polymer. The required equipment is the same as shown in Fig. 1. When C3 motor polymer is made, however, the debutanizer and rerun columns are not required. Figure 6 shows a true-boiling-point curve of typical Cs motor polymer.

4. Petrochemical Reactions After the introduction of the alkyl-aryl-sulfonate type of detergent, catalytic polymerization became increasingly important in the production of olefins to be used as raw materials and intermediates in further

23 1

POLYMERIZATION OF OLEFINS

petrochemical reactions. The development of various oxidation processes has further increased the demand for polymers of bot,h propylene and propylene-butylene mixtures. The most common of these polymers are listed in Table V. Although cumene and ethylbenzene are produced by an TABLE IV Typical Ca Motor Polymer Reid vapor pressure, p.s.i. Gravity, "A.P.I. A.S.T.M. distillation, "F. I.B.P. 5% 10 30 60

70 90 E.P. A.S.T.M. gum mg./100 ml. F-1 Octane number clear +1.0 ml. TEL +3.0 ml. TEL Blending octane number clear

4.0 61.4

103 194 228 266 282 302 378 42 1 1.0 95.5 98.0 99.5 106-126

Liquid volume, x

FIG.6. Ca polymer.

alkylation process, the reaction also occurs over Solid Phosphoric Acid in a typical polymerization unit and hence has been included. Di-isopropylbenzene, cymene, and secondary butylbenzene also can be produced in these units.

232

EDWIN K. JONES

a. Tetramer. Tetramer, as known t o the refining industry, is a relatively straight-chain olefin product made by the controlled polymerization of propylene and consists primarily of Clz olefins with a minor amount of C16olefins. The Csand C e olefins, which also are formed during the reaction, are recovered from the product along with unreacted

TABLE V Typical Polymers Used in the Petrochemical Industry Formula

CsHie C7H14 CeHis C12H21 ik CISHSO CiaHm

Trade name Dimer Heptene Trimer Tetramer Pentamer

TABLE VI A.S.T.M. Engler Distillation of Streams i n Tetramer Operation

Gravity, "A.P.I. A.S.T.M. Engler distillation, I.B.P. 10% 50 90 95 E.P. R.V.P. wt. % s.

Light polymer

Tetramer

Polymer bottoms

62.5

51.0

36.4

131 244 275 290 300 329 8.0

353 361 371 397 420 450

503 520 534

OF.

-

-

0.01

propylene and are recycled to the reactor. Properties of typical tetramer streams are shown in Table VI. A true-boiling-point curve for tetramer is shown in Fig. 7. The quantity of butylene in the feed must be kept a t a minimum in order t o ensure a good product, because a polymer made from butylene is

233

POLYMERIZATION OF OLEFINS

less stable and tends to break down in the alkylation process t o which the tetramer is charged. Essentially all tetramer is made by use of Solid Phosphoric Acid because of the high catalyst activity required for high conversions.

Liquid volume, x

FIG.7. Tetramer. TABLE VII Yields from Tetramer Operation

Charge: Fresh propane propylene % Propylene Products: Spent P.P. Motor-fuel polymer Tetramer Polymer bottoms % Propylene conversion

Bbl./day

Vol. %

595

100.0 49.3

325 2 173 8

54.6 0.3 29.1 1.3 92.3

-

-

The essential equipment requirements for a tetramer unit are a combined feed-charge drum, charge pump, catalyst chamber, depropanizer, recycle column, rerun column, and attendant pumps and exchangers. Figure 1 shows such a unit as designed to produce tetramer and cumene alternatively. Table VII shows yields from the tetramer operation when a 350°-485'F. boiling-range tetramer is produced. b. Cumene. Cumene was used widely as a high-performance-number component of aviation gasoline, but most of these units were shut down

234

EDWIN K. JONES

because of little or no demand. After the devdopment of the Hercules process for the oxidation of cumene in the production of phenol and acetone, cumene units again were put into operation. The equipment required for the production of cumene is almost identical to that shown in Fig. 1. The recycle column takes the recycle benzene overhead while the cumene is removed as an overhead product of the rerun column. Table V I I I shows typical yields from cumene operation. TABLE V I I I Typical Yields from Cumene Operation

Charge: Fresh propane propylene Benzene

Bbl./day

Vol. %

525 240

68.7 31.3 100.0

Products :

L.P.G. Cumene Bottoms

260 380 15

34.0 49.7 2.0

The charge to the unit is treated refinery propane-propylene along with recycle benzene from the recycle column overhead. Make-up benzene is added to the recycle. Nitration-grade benzene is usually used so t ha t a drag stream of benzene is not required to remove contaminants from the unit. Table I X shows the component analysis of the various streams in the cumene process. G. Trimer. Trimer is the trade name for nonene, which is made by polymerizing propylene. Dimer (hexene) and tetramer are also produced in the operation. Dimer can be recycled t o the reactor t o increase the production of trimer if desired. Trimer has been used to make nonyl naphthalene and decyl alcohol. The equipment required to make trimer is the same as th a t shown in Fig. 1. The dimer is taken off as an overhead product of the recycle column and the trimer is a n overhead product of the rerun column. An additional fractionation column is required when it is desired t o make a rerun tetramer while trimer is produced. Typical trimer has an A.S.T.M. Engler distillation range of 260" t o 300°F. It is not normally made when high yields of tetramer are required because such severe operating conditions lower the bromine number of the trimer.

235

POLYMERIZATION O F O L E F I N S

d. Heptene. Heptene is used for making iso-octyl alcohol in the 0x0 process. It can be recovered easily from C3-C4 motor polymer by adding two more fractionators to the polymerization unit. Normal motor polymer operation using C3-C4 feed will yield somewhat more than 15 liquid TABLE IX Component Analyses of Streams in the Cumene Process % by vol. Recycle benzene Benzene Toluene Ethylbenzene Cumene Nonaromatics

95.3 0.012 0.014 0.081 4.58 % by wt.

Cumene Olefins Cumene Bottoms

0.5 99.4 0.07 % by wt*

Cumene bottoms Unsaturates (other than a-methylstyrene) Cumene a-Methylstyrene t-Butylbenzene +But ylbenzene Cumene bottoms Other monoalkylbenrencs Di-isopropylbenzene 0-

m-

POther dialkylbenzenes 0-

m-

PUnaccounted for Bromine numbers Benzene recycle Cumene Cumene bottoms -

1.8 6.6 0.04 0.8 0.5 % by wt.

19.4 16.2 23.5 26.5 1.8 0.3 0.8 1.8

1.16 0.46 2.08

~~

volume % of the heptene cut. Selective operation will produce slightly higher yields. The A.S.T.M. Engler boiling range of a typical narrow cut heptene is 188" t o 200°F. and the gravity is 67.8"A.P.I.

236

EDWIN K. JONES

e . Dimer. Although dimer (hexene) is not used by the chemical industry so widely as some of the other propylene polymers, it is certain to be one of the important building blocks as more uses are developed. On some polymerization units the dimer is charged back into the reactor as recycle for the production of either trimer or tetramer. f. Pentamer. Pentamer is finding considerable use for the production of lubricating oil detergents. It is also considered beneficial when combined with tetramer for the production of household detergents. g . Ethylbenzene. Although much ethylbenzene is produced using aluminum chloride catalyst, refiners are looking toward the solid phosphoric acid process as a more simple unit which can charge the relatively impure ethylene streams that are available. This type of unit is capable of producing a lower cost ethylbenzene than other processes.

V. POLYMERIZATION PROCESSES USING PHOSPHORIC ACID 1. U.O.P. Solid Phosphoric Acid The Solid Phosphoric Acid catalyst is produced from a controlled mixture of liquid phosphoric acid and kieselguhr. The catalyst is white or gray cylindrical-shaped pellets. It is hard when dry but picks up water on exposure to moist air for extended periods of time. Two types of polymerization units are designed by Universal Oil Products Company. The U.O.P. Reactor-type unit contains the catalyst in tubes which are surrounded by water in a jacket for the purpose of removing the heat liberated by the exothermic polymerization reaction. The steam generated in the water jacket normally is used to preheat the feed. A feed-to-products heat exchanger furnishes the remaining heat requirements. Conventional depropanizer and debutanizer columns are used to fractionate the product. Figure 3 shows a flow diagram of a reactor type of polymerization unit. The U.O.P. Chamber-type unit contains the catalyst in a reactor in which the catalyst is separated into a number of beds. Temperature control is accomplished by quenching between the beds with cold liquid effluent and by recycling spent propane and butane effluent into the reactor feed. This is not a heat-balance operation and a preheater using an outside source of heat is required in addition to the feed-to-products exchanger. A depropanizer and debutanizer are usually used to fractionate the effluent. Figure 2 shows a flow diagram of a chamber type of polymerization unit. 2. Solid Copper Pyrophosphate of the Polyco T y p e (1)

The Copper Pyrophosphate Catalyst was developed for use in the Kellogg polymerization units. These units are similar t o the two U.O.P. types with the following exceptions:

POLYMERIZATION OF O L E F I N S

237

a. The tubes are considerably larger in diameter in the reactor-type units. b. Forced circulation of the cooling medium is usually used instead of thermal circulation. The coolant may be either gas oil or water. A disadvantage of the catalyst is that it must be mixed with charcoal in order t o minimize plugging. This requires a larger reactor in order to secure the required catalyst activity for good conversion rates. New Copper Pyrophosphate Catalyst requires a n activating period before satisfactory olefin conversion can be obtained. This activating period may range up t o 1 or 2 days, a factor which requires additional plant capacity if high conversions are required during these periods. This type of unit requires very efficient sulfur removal from the feed, as sulfur acts as a catalyst poison. Furthermore this catalyst is poisoned by the same materials which poison the Solid Phosphoric Acid type of catalyst, which contains no copper. 3. Liquid Phosphoric Acid-on-Quark of the California Research T y p e ( 2 )

Units designed to use the Liquid Phosphoric Acid-on-Quartz type catalyst developed by California Research Corporation are of the chamber design. The first units had solid-quartz beds but the temperature control of the eatalyst beds was poor, resulting in excessive coke. The quartz was split into beds with the quench between beds in later units. Early commercial units utilized 10- to 20-mesh quartz but, because of the low surface area and consequent low activity of the catalyst, the quartz size was reduced to 28 to 35 mesh. The quartz is activated by pumping the reactor full of 75% phosphoric acid, allowing the excess acid t o drain out, and then charging hot hydrocarbon t o the unit. Even with the smaller quartz particles the catalyst activity is much lower than th a t of the Solid Phosphoric Acid catalyst. This lower activity has resulted in lower olefin conversion in this type of unit. Increased conversion has been obtained by separating the olefins from the product with very efficient fractionators and returning them t o the reactor. However, this factor requires relatively large units with high utility consumption. These units utilize commercially available liquid phosphoric acid t o coat the quartz bed. However, a srrious disadvantage develops in reactivating spent catalyst, in that when the spent acid is washed from the quartz the tars are not also removed. This requires the periodic removal of the quartz so that the tars may be burned off.

VI. MATERIALS O F CONSTRUCTION The solid catalysts of the U.O.P. and Polyco types are noncorrosive under all ranges of operating conditions and therefore require only carbon steel for all equipment.

238

EDWIN K. JONES

The Liquid Phosphoric Acid catalyst is very corrosive a t elevated temperatures and therefore all hot lines are alloy and all hot vessels are alloy clad, including the catalyst chamber and all catalyst-chamber internals. Stainless steel is used for this corrosion protection and in some severe cases nickel has been used for some internal parts.

VII. FUTURE OUTLOOK FOR POLYMERIZATION Commercial polymerization was once used only for converting the olefins from cracked gases into motor fuel. However, it is rapidly becoming very important in the production of such petrochemicals as heptene, propylene dimer, trimer, tetramer and pentamer; and the alkylated aromatics such as ethylbenzene, isopropylbenzene, cymene, and butylbenzenes. This list may be expected to grow as new uses are found for the heavier olefins. Propylene is becoming more important as one of the building blocks for petrochemical reactions, and the polymerization process is very important in producing the required higher olefins. REFERENCES 1. Steffens, J. H., Zimmerman, M. U., and Laituri, J. M., Chem. Eng. Progr. 45, 4, 269 (1949). 2. Langlois, G. E., and Walkey, J. E., Petroleum Refiner 31, No. 8, 79 (1952). 3. McAllister S. H., Proc. Am. Petroleum Znst. Part 10, Sec. 111, 78 (1937).

Coal-Hydrogenation Vapor-Phase Catalysts E. E. DONATH Koppers Company, Inc., Pittsburgh, Pennsylvania

Page I. Introduction. . . . . . . . . . . . . . . . . 239 11. Brief History of Vapor-Phase.................... 242 111. Tungsten Disulfide Catalyst. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 245 1. Preparation and Properties.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 245 2. Hydrogenation Reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 250 3. Reaction Mechanism 4. Factors Affecting Catalyst Activity.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 257 5. Influence of Catalyst-Pellet Size IV. Nonsplitting Catalysts. . . . . . . . . . . . V. Splitting Catalysts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 291

I. INTRODUCTION The term coal hydrogenation usually means the addition of hydrogen t o coal and t o oils derived from coal or from other sources combined with cracking of the raw material. The first successful attempt to add hydrogen t o coal was undertaken by Berthelot ( 1 ) . The use of nascent hydrogen from hydriodic acid a t 270°C. gave 60% of the weight of the coal as oil. The industrial development of the coal-hydrogenation process started with Bergius's work on the constitution of coal, which led to the conclusion t ha t molecular hydrogen reacts with coal a t eIevated temperature and high pressure. I n 1913, a t a pressure of about 200 atm. and a temperature between 400" and 500"C., Bergius succeeded in converting coal into liquid products ( 2 ) . It was a long way from this first experiment in an autoclave t o a large-scale use of the process; yet the first obstacle t o continuous operation, the injection of coal into the high-pressure reactors, was overcome by Bergius. Coal was made into a paste with oil (3) and this paste was injected by means of pumps. I n 1916 an experimental plant was erected by Bergius near Mannheim; however, little progress was made until 1921. I n this plant coal paste was hydrogenated in horizontal reactors. In these reactors heated nitrogen a t reaction pressure was circulated between a liner and the pressure-retaining wall t o supply heat and to prevent reaction hydrogen from corroding the steel of the pressure vessels. The products obtained in this plant 239

240

E. E. DONATH

contained gasoline, Diesel fuel, and fuel oil. Their properties were similar t o those of low-temperature tar obtained from the same coal. The boiling range and properties of the oil could be changed within narrow limits only, and the fuels obtained were not competitive in quality with petroleum products. The tar acids and bases in coal-hydrogenation oils, which are a disadvantage for the use of these oils as fuels, have now turned into an asset, and the separation of compounds of this type and their use as chemical raw materials is gaining in importance. Only the use of catalysts could improve the quality of the fuels and direct the coal-hydrogenation reactions to better yields of the desired product. At the plants of Badische Anilin- und Soda-Fabrik the Haber highpressure ammonia synthesis was developed into a commercial process by C. Bosch beginning in 1910. The production of methanol from carbon monoxide and hydrogen was introduced as a second commercial highpressure synthesis in 1924 by M. Pier. These developments and the experience with high-pressure techniques and catalysts were prerequisites for the further development of the coal-hydrogenation process. Since the necessity of using catalysts for the coal-hydrogenation reaction was clearly understood by Pier, catalysts were at first studied on a model reaction which permitted rapid screening of catalysts. The reduction and hydrogenation of cresols was used for this purpose. It was found that the iron catalyst used for ammonia synthesis was a satisfactory catalyst for this reaction; however, it was poisoned rapidly by sulfur. It was then found that iron sulfide, prepared from metallic iron and hydrogen sulfide, and especially cobalt, molybdenum, and tungsten sulfides were very active catalysts. I n 1925 brown-coal tar was hydrogenated at 200 atm. with a molybdenum catalyst and converted in one stage into essentially pure gasoline hydrocarbons without coke formation. These experiments were the basis for the interest of Badische Anilin- und SodaFabrik in coal hydrogenation. For these tests reactors of 1-in. I.D. with 50-ml. granular catalyst were used. At a temperature of 450°C. the tar throughput (4) was about 0.2 L.H.S.V. (liquid hourly space velocity based on the volume of catalyst) and the feed partial pressure in 200 atm. of hydrogen was below 1 atm.; thus a huge excess of hydrogen was used. For commercial application, higher throughput and feed partial pressure (lower hydrogen excess) were a necessity. I n such experiments on a larger scale it was found that the catalyst activity decreased rapidly. High-boiling asphaltic hydrocarbons of low hydrogen content were deposited on the catalyst. Later experiments showed that a t 450°C.and 200 atm. hydrogen pressure, aromatic compounds with four or more condensed rings are hydrogenated very slowly and, in addition, form polymerization products of higher molecular weight on the catalyst.

COAL-HY D ROGENATION VAPOR-PHASE CATALYSTS

24 1

A fundamental improvement was achieved by the division of the hydrogenation process into two stages ( 5 ) . 1. Liquid-phase hydrogenation for the conversion of coal, tar, or heavy oil into an intermediate product-the middle oil-with a boiling range up t o about 325°C. (In this stage finely divided catalysts distributed in the liquid or liquefiable feed were used.) 2. Vapor-phase hydrogenation using a fixed-bed catalyst for the conversion of middle oil into gasoline whereby the middle oil was vaporized under pressure in a stream of hydrogen. This two-stage process for the hydrogenation of tar proved useful also for the hydrogenation of coal and petroleum oils; in subsequent operations these two phases of the coal-hydrogenation process developed separately, requiring different catalysts and equipment. At first, in 1924, coal, especially brown coal, was hydrogenated without pasting oil. To obtain comparative data for different coals and catalysts a t well-defined temperature and residence time, a “tilting converter” was used. The coal sample confined in a wire-mesh container was inserted into the cold end of the converter. The other end was heated to reaction temperature and a stream of hydrogen passed through the tube. After reaction temperature had been established a t the heated end, the converter was tilted and the coal entered the heated zone. Coalliquefaction products were carried out of the reaction space with hydrogen and condensed in a trap. After the predetermined reaction time had passed, the converter was tilted back and the coal residue removed from the reaction zone. I n this way pure coal-hydrogenation products which were free of pasting oil were obtained and the suitability of different coals and the activities of catalysts were determined. However, this procedure was not developed into a commercial, continuous process. The commercial development of coal hydrogenation in liquid phase was based on the use of a coal paste as proposed by Bergius. The coal hydrogenation in the liquid phase, aided by catalysts, was developed commercially as the first stage of a n integrated fuels-production process in which middle oil was produced. An excellent summary of coal hydrogenation in liquid phase has been prepared by Storch (6). The industrial status of the coal-hydrogenation process has been reviewed in a series of papers by Sherwood (7). Coal liquefaction in liquid phase is possible without catalyst. However, in commercial synthetic-liquid-fuels plants the use of catalysts is of definite advantage. Their main functions are t o increase the reaction rate and prevent the build-up of high-molecular-weight constituents, which are difficult to split into lower boiling fractions. Vapor-phase hydrogenation without a catalyst is impractical ; the influence of the catalyst on rate and direction of the reaction is by far

242

E. E. DONATH

greater in this phase than in liquid-phase hydrogenation. I n this paper the development of vapor-phase catalysts and their properties will be described.

11. BRIEFHISTORYOF VAPOR-PHASE-CATALYST DEVELOPMENT At first the primary purpose of the vapor-phase hydrogenation step was the conversion of middle oils obtained from liquid-phase coal hydrogenation, or other sources, into fuels, especially gasoline. The catalyst development during the years prior to 1927 was based mainly on the use of molybdenum and tungsten. During these studies many initially active catalysts were eliminated because of insufficient mechanical strength, difficulty of preparation, or unsatisfactory product quality and yield. The first catalyst which was used on a commercial scale in the hydrogenation plant a t Leuna for many years consisted of equal molar amounts of molybdenum, zinc, and magnesium oxides-the composition in weight percentage being Moos 53.5, ZnO 30, and MgO 16.5. The catalyst was prepared by mixing an aqueous paste of the oxides, followed by drying. The activity depended upon the purity of the components and the particle size of the zinc and magnesium oxide. This catalyst was widely used in commercial vapor-phase hydrogenation a t 200 atm. pressure. With brown-coal middle oil, about 30 % conversion was obtained with a n hourly space time yield of 0.1 t o 0.2 volume of gasoline/volume of catalyst a t a temperature of 450°C. The formation of gaseous hydrochrbons was about 20% of the converted feed. T o obtain complete conversion of middle oil to gasoline, recycle operation was required. Pilot plant results indicated th at satisfactory catalyst life could be realized by gradual temperature increase t o off set the decrease in activity with time. This decrease in activity was caused mainly by the formation of more refractory aromatic recycle oils. I n commercial operation the activity loss of the catalyst was more rapid. The decline in catalyst activity could be slowed down by decreasing the end point of the feed middle oil or by withdrawing small amounts of heavy ends formed. Commercial operation indicated further that the use of recycle hydrogen was a cause of the more rapid loss of catalyst activity. Ammonia and volatile ammonium salts formed by the reduction of tar bases in the feedstock might have been a factor in the accelerated-catalyst-activity loes. As an example of a more active catalyst, which was not adopted on a commercial scale, a catalyst consisting of ninety parts Moo3, ten parts Cr03, twenty parts kaolin with the addition of one-half part of powdered aluminum may be mentioned. Although of higher activity than the Mo-Zn-Mg catalyst, it lacked the mechanical strength when heated

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

243

slowly t o reaction temperature and it disintegrated, therefore, in commercial operation. Investigation of used catalysts showed that the molybdenum was always completely converted to sulfide. The sulfide formation, however, was not the cause of the loss of activity, as pretreatment of the catalyst at atmospheric pressure with hydrogen sulfide gave catalysts with unchanged or slightly increased catalyst activity. However, pretreatment of a tungsten-zinc oxide mixture with hydrogen containing 10% HzS at 200 atm. total pressure for an extended period of time a t a temperature of 200°C. and higher gave a catalyst of improved activity. Because of the high hydrogen sulfide partial pressure required, this pretreatment was not practical for commercial catalyst production, and other methods for the preparation of molybdenum and tungsten sulfides were investigated. I n 1930 a very active tungsten sulfide catalyst was prepared from ammonium sulfotungstate. With this catalyst, splitting of middle oil proceeded a t a temperature of about 400°C. with an hourly space time yield of 0.5 volume of gasoline or more per volume of catalyst. The conversion to gasoline was 50% and higher, and the gas yield based on converted feed, around 8 %. While the molybdenum-zinc-magnesia catalyst did not give satisfactory conversion of bituminous-coal middle oil into gasoline, the tungsten sulfide catalyst hydrogenated those oils without difficulty. Presence of ammonia or organic bases did not cause a decrease of the activity with time, although t o some extent it decreased the activity level of the catalyst. One difference between the MoOa-ZnO-MgO and the WSZ catalyst was pronounced. The former, for example, converted brown-coal middle oil into gasolines containing about 30% aromatics, and the aromaticity of the gasoline increased with operating temperature. With the WSz catalyst, even from purely aromatic feedstocks, gasoline with a maximum aromatics content of 10% was obtained. The splitting reaction, as will be shown later, did not start until the hydrogenation of aromatic rings was essentially complete. The high splitting activity of the catalyst and the high content of branched-chain paraffins in the product may be explained by a carbonium ion mechanism. The pattern of products obtained by splitting with the MoOa-Zn0-Mg0 catalyst in the presence of hydrogen, on the other hand, resembles t o some extent the products obtained by thermal cracking of oil, which follows a free radical mechanism. However, one difference in comparison with the cracking process is obvious, the absence of the large amounts of coke, polymerization products, and olefins in the hydrogenation products. Since the tungsten sulfide catalyst could not be used at high temperatures for the production of highly aromatic gasolines, the develop-

244

E. E. DONATH

ment, of vapor-phase hydrogenation catalysts branched out. Specific catalysts were developed for different reactions. The octane number of gasolines obtained with the WSz catalyst was somewhat low. The search for catalysts continued, therefore, with two objectives. The first was the development of a catalyst of the WS2 type with decreased hydrogenating activity and increased ability t o produce branched-chain hydrocarbons. The second was the development of improved catalysts for the production of highly aromatic gasolines, th a t is, catalysts t ha t give higher conversion and lower formation of gaseous hydrocarbons than those of the Mo-Zn-Mg type. I n 1934 a n improved WSz type of catalyst was obtained by use of activated fuller’s earth, which was treated with hydrofluoric acid as support for 10% WS2. This diluted catalyst gave the same conversion with petroleum oils as the WSZ catalyst, but produced gasolines with a 5- to 10-point higher octane number. However, if used with brown- or bituminous-coal middle oils, it lost its activity rapidly. It was found, as will be shown later, that nitrogen compounds and, to some extent, oxygen compounds are the catalyst poisons. Extraction of ta r acids and bases with caustic and acid was possible. However, the loss caused b y such treatment made removal of oxygen and nitrogen by selective hydrogenation t o water and ammonia more attractive. Such middle oils, therefore, had t o be pretreated b y vapor-phase hydrogenation (saturation or prehydrogenation) to remove oxygen and nitrogen compounds. Tungsten sulfide as a saturation catalyst gave satisfactory removal of these compounds; however, it produced some gasoline of comparatively low octane number. It was found th a t the splitting activity of tungsten sulfide can be virtually suppressed by adding 15% nickel sulfide. This catalyst found commercial application, especially for the hydrogenation of diisobutene to isooctane. A catalyst with a higher nickel sulfide content was used for dehydrogenation reactions. A catalyst of similar properties and activity, but lower cost, was obtained with active alumina as carrier. This catalyst contained 70% active alumina, 27 % tungsten sulfide, and 3 % nickel sulfide and is used commercially as a saturation catalyst. Improved fuller’s earth-WSn splitting catalysts were also developed. Hydrofluoric acid-treated fuller’s earth alone is a splitting catalyst that produces high-octane gasoline. It required higher operating temperatures and therefore, to avoid activity loss, had to be used a t pressures above 300 atm. Iron compounds instead of tungsten sulfide on HF-treated fuller’s earth also gave gasolines of improved octane number. Probably catalysts of this type were used commercially by the I.C.I. in Billingham, England (8) for the hydrogenation of creosote oil and by the A.N.I.C. (9)

COAL-HYDROGENATION

VAPOR-PHASE CATALYSTS

245

in Italy for the hydrogenation of petroleum oil. Commercial use of such catalysts in Germany was contemplated during the war, but a lower gasoline yield and production rate compared with the tungsten-containing catalyst and the catalyst sensitivity were the main obstacles. Considerable greater octane-number improvement in comparison with the WSz-fuller's earth catalyst was obtained with synthetic aluminum silicates as carrier. Such a catalyst with 0.5% Moos as activator produced gasolines with an aromatics content of about 50% from Nz-free aromatic feedstocks a t 400°C. (10). Thus a catalyst type was found which permitted splitting of aromatic rings, probably by a carbonium ion mechanism, at temperatures around 400°C. simultaneously with the formation of paraffins with predominantly branched chains. This catalyst type thus is basically different from catalysts used for the manufacture of highly aromatic gasolines a t 500°C. Catalysts of this type are used a t present commercially (11) for the hydrogenation of petroleum oils in Germany. Parallel with this development of low-temperature vapor-phase catalysts went the development of catalysts which a t higher temperature, about 450" to 500"C., permitted conversion of aromatic middle oils into highly aromatic gasoline. Two such commercially used high-temperature catalysts may be mentioned here. For operation a t 300 atm. pressure, a catalyst containing 80% activated carbon, 15% chromium oxide, and 5% vanadium oxide was used. For operation a t 700 atm. pressure a t the Welheim plant of Ruhrol GmbH, a catalyst using fuller's earth as carrier was developed, This catalyst contained 0.6 % molybdenum, 2 % chromium, 3 % zinc, and 5 % sulfur on hydrofluoric acid-treated fuller's earth. It gave higher conversion and catalyst life length than the 300-atm. catalyst. Both catalysts gave gasoline containing about 50% aromatics from bituminous-coal middle oil. A great variety of catalysts has been developed for vapor-phase highpressure hydrogenation reactions. I n the next section the tungsten disulfide catalyst and its properties will be described in detail. The following sections will be devoted t o specialized catalysts. I n such catalysts the hydrogenating activity of the tungsten sulfide catalyst can be greatly enhanced while its splitting activity is suppressed, and vice versa. Catalysts are thereby obtained which permit production of a variety of highquality fuels from different feedstocks.

111. TUNGSTEN DISULFIDECATALYST 1. Preparation and Properties

Tungsten disulfide (12) is a bluish-black substance with a density of 7.5 which crystallizes in the hexagonal system. It is stable in the absence

246

E. E. DONATH

of oxidizing agents at temperatures up t o 1100°C. but at higher temperatures it decomposes with sulfur evolution. It cannot be dissolved without decomposition. For the preparation of a catalytically active tungsten disulfide, a method had t o be found which avoided high temperatures and permitted formation of a material with a large surface. Such a method was the formation of oxygen-free ammonium sulfotungstate and its aubeequent decomposition t o tungsten disulfide. For the production of tungsten disulfide, tungstic acid is dissolved in ammonia a t 70"C., whereby a solution of ammonium tungstate is formed. From this solution near its boiling temperature ammonium thiotungstate is precipitated with hydrogen sulfide. An HZS pressure of about 1.2 atm. abs. is used t o form a n oxygenfree compound according t o the equation

(NH4)zWOd

+ 4H2S = (NH4)2WS4 + 4Hz0

The crystallized sulfosalt is then decomposed a t 400°C. with hydrogen according t o the following equation:

The decomposition temperature of 400°C. was found most suitable for commercial catalysts, although decomposition at lower temperatures produced catalysts of higher activity for some low-temperature reactions. The tungsten disulfide was pressed into pellets of a bulk density of about 2.5 g./ml. and a crushing strength of 300 kg./cm.2 The fresh catalyst usually contains a 10% excess of sulfur as compared with the formula WSz. Excessive contact with air during preparation should be avoided; however, the catalyst usually contains small quantities of moisture and sulfuric acid which are formed by oxidation of absorbed hydrogen and of sulfur during handling. The catalyst-pellet shrinkage during use is negligible; a slight decrease in pellet diameter is almost compensated b y an increase in pellet height. I n normal vapor-phase operation the catalyst maintained its activity for 1 t o 2 years, although some batches have been used for 5 years. The decrease in activity of the catalyst is caused by the irreversible absorption of high-moIecular-weight hydrocarbons. Ash (e.g., iron phenolates) -containing feedstocks cause a more rapid loss of activity. Analyses of used catalysts indicated carbon contents of about 2%. Tungsten disulfide is isomorphous with molybdenum sulfide, and many of the catalytic properties of these two compounds are similar. X-ray analyses of the catalyst showed a hexagonal lattice (13). The primary size of the crystallites is about 3 X lo-' cm. in height and twice th a t in width. Microscopic examinations showed that the crystals are pseudomorphous t o the

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

247

monoclinic ammonia sulfotungstate crystals from which they were formed. They also show evidence of shrinkage and cracks which might be the cause of the large surface of the catalyst. The surface (14) of the tungsten sulfide catalyst was calculated from adsorption isotherms a t pressures u p to 100 mm. Hg obtained a t temperatures between 0" and 150°C. with argon. T h e method used for the TABLE I Surface Area Surface area Activated carbon Silica gel Tungsten sulfide Pumice

m.a/g.

m.l/ml.

400-800 110 20 1

55 50

calculation is described by Huckel ( 1 5 ) . I ts accuracy was checked by measurement of the surface area of metal foils. Because of adsorbed sulfur, sulfur, sulfuric acid, etc., on the catalyst, these measurements are difficult and the results depend on the pretreatment of the catalyst. Nevertheless, the data in Table I show th a t the internal surface of the WSZ catalyst per unit volume is of a similar magnitude t o t h a t of silica gel. From this internal surface area a particle size is calculated which is TABLE I1 Gas Adsorption on Digkrent Catalysts Argon

Ethylene

Ethane

Hydrogen

(Ml. gas/ml. catalyst) Saturation pressure, mm. Hg Activated carbon Silica gel Tungsten sulfide

5 0.02 0.005 0.013

5 2.7 0.25 7.8

5 2.6 0.1

40 0.02

2.1

0.55

0.02

about ten times larger linearly than the primary particle size obtained by X-ray measurements. From the adsorption isotherms about 2500 t o 3500 cal./mole is obtained as heat of argon adsorption, which is normal for nonactivated adsorption. Adsorption measurements with other gases in Table I1 show th a t the gas adsorption on the WSz catalyst is selective toward hydrogen and unsaturated hydrocarbons. The pretreatment of the tungsten sulfide catalyst influences the amount of hydrogen th a t is adsorbed. Figure 1 shows in curve I a n isotherm a t OOC. with a catalyst th a t was degassed for

248

E. E. DONATH

several hours a t 300°C. Curve I1 shows the isotherm for a catalyst which has been pretreated with hydrogen a t 300°C. for several hours and then degassed. Hydrogen pretreatment of the catalyst thus increases the amount of hydrogen adsorbed. The hydrogen adsorption proceeds in two stages. At 0°C. about one third, and a t 160°C. more than one half, of the hydrogen is adsorbed within 1 min., but establishment of the final adsorption equilibrium requires several hours. Pertinent data are shown in Table 111. At higher 0.25-

-4 6

rJ u)

8' 5

%02--0.5 * d P

a

L

m

a

w

h 0

w

-0.4 m K 0

0

u)

a

-*0.32

2

0

4

I.

CATALYST DEGASSED AT 300'C. AND DEGASSED AT 300%.

n. CATALYST H z T R E A T E D

FIG. 1. Adsorption of hydrogen on tungsten disulfide. I. Catalyst degassed a t 300°C. 11. Catalyst Hz treated and degassed a t 300°C.

temperatures the amount of adsorbed hydrogen is increased as well as the adsorption rate. That selective adsorption of the unsaturated hydrocarbons as well as of hydrogen on tungsten sulfide is a n activated adsorption is shown by the data in Table I11 and Fig. 2, which also shows the rate of hydrogen adsorption on tungsten disulfide. Figure 2 also shows that the rate of hydrogen adsorption is increased if the catalyst had adsorbed an equivalent amount of butylene before the admission of hydrogen. The olefin adsorbed on the WSz catalyst is hydrogenated by the added hydrogen, which is adsorbed from the gas

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

249

phase. Under the conditions used, the paraffin formed remains essentially completely adsorbed and the pressure change is therefore a direct measure of the hydrogen reacting with the olefin. The butane formed on the catalyst by the reaction can be completely desorbed by evacuation. TABLE I11 Hydrogen Adsorption on W S Z ~~

Temperature, "C. Time, min.

Pressure, mm. Hg Adsorbed hydrogen, ml./g. WSZ

0

1 30

56.9 44.4

0.05 0.18

50

1 215

58.0 33.6

0.10 0.26

1

180

123.8 60.9

0.31 0.61

1 140

42. 7 1.3

0.25 0.49

loo* 160

* For the experiment at 100°C. twice the amount of hydrogen was used.

MINUTES

FIG.2. Hydrogen adsorption rate with and without adsorbed butylene at 0 and 50°C., WS, catalyst.

Ethylene reacts somewhat more slowly than butylene. This is the opposite of the behavior observed with nickel (16) catalysts. The olefin hydrogenation rate with nickel catalyst decreases with increasing molecular weight of the olefin. The experiments were made with WSz th a t was

250

E. E. DONATH

treated with hydrogen a t 300°C. and degassed before olefin adsorption. WSz catalyst which was not pretreated with hydrogen showed the same olefin adsorption; however, a t 0°C. it failed to adsorb added hydrogen. A tungsten sulfide catalyst which had been inactivated by extended operation had, as expected, a considerably lower adsorption capacity than new catalyst. This is shown in Fig. 3 for the adsorption of ethane a t 100°C. The adsorption rate of gases on used, inactive WSZ catalyst seems t o be decreased to an even greater extent. 0.4

t; *-I a 0.3 t

U

0

*(D 0.2 I (u

0

E,

OJ

0

0

20

40

60

80

100

PRESSURE mm Hq,

FIG.3. Adsorption of ethane on new and used WS, catalyst a t 100°C.

The reaction between butylene and hydrogen a t room temperature is comparatively slow. Another reaction-the combination of oxygen and hydrogen-is rapid at room temperature; however, the catalyst is poisoned during this reaction. Thfs observation might have some bearing upon facts that will be reported later which indicate sensitivity of some hydrogenation catalysts toward oxygen-containing compounds. The results of adsorption measurements made a t room temperature cannot be used to draw direct conclusions about the reaction mechanism on this catalyst a t its usual operating temperature and pressure. Temperatures around 400°C. and oil partial pressures of several atmospheres in the presence of high-pressure hydrogen are used in actual operation. 2 . Hydrogenation fieactions

The adsorption experiments have shown that tungsten sulfide catalyzes the hydrogenation of olefins t o paraffins a t room temperature. They give, however, no indication that WSZ catalyzes other reactions. Two types of reactions are a. Reactions without change of the carbon skeleton of the feed. b. Reactions that involve change of the carbon skeleton of the feed. The first group of reactions includes hydrogenation of olefins and aro-

25 1

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

matic rings and reduction of organic sulfur, oxygen, and nitrogen compounds t o the corresponding hydrocarbons and HzS, HzO, and NHa. Temperatures of 200°C. or more are used. The reactions of the second group, like isomerization and splitting, require higher temperatures, about 400°C. The catalyst activity is enhanced by sulfur and decreased by nitrogen and oxygen compounds, and these elements influence the reaction temperature. TABLE IV Hydrogenation with Tungsten DisulJide as Catalyst Reaction without Change of the Carbon Skeleton Reaction conditions

Reaction type

Feed

Product

Hydrogenation Diisobutylene Isooctane of olefins Hydrogenation Cyclohexane of aromatics Benzene Naphthalene Decalin Lignite M.O. PrehydroReduction of genated oil organic com- containing pounds of 18% Phenols Oxygen Nitrogen 4% Nitrogen bases Sulfur 2.5% Sulfur compounds

Feed par- Feed tial rate, Temp., Pres., pres., kg./ Conver"C. atm. atm. (l.)(hr.) sion, %

2

99

3 3

0.1 0.9

99 90

8

1.0

216

250

20

322 336

200 200

380

200

99.5 99.5 99

Reactions of pure compounds have been studied with tungsten sulfide in continuous systems, that is, while a stream of hydrogen and oil is fed a t reaction temperature and pressure over the catalyst. Reactions without change of the carbon skeleton are shown in Table IV. Hydrogenation of olefinic double bonds becomes rapid at temperatures around 200°C. and for the hydrogenation of aromatic rings to hydroaromatics, temperatures of 300" t o 350°C. are needed. For complete reduction of oxygen, nitrogen, and sulfur compounds as contained in coal-hydrogenation middle oils, a temperature of 380°C. is needed. Tungsten sulfide catalyst should accelerate not only the hydrogenation reactions, but also dehydrogenation of hydroaromatics to aromatics. Thermodynamic calculations show that such reactions are probable a t about 500°C. a t a hydrogen pressure of about 50 atm. Use of th e tung-

252

E. E. DONATH

sten sulfide catalyst under such conditions results in partial dehydrogenation accompanied by extensive splitting of the hydroaromatics with rather rapid loss of catalyst activity. However, as mentioned before, the splitting activity of tungsten sulfide can be suppressed by addition of nickel sulfide. With a nickel sulfidetungsten sulfide catalyst a t 485°C. and 50 atm. total pressure and 6 atm. partial pressure of methylcyclohexane or cyclohexane, the corresponding aromatics can be obtained with a conversion of 90 and SO%, respectively. An L.H.S.V. of 0.6 can TABLE V Hydrogenation with Tungsten Disulfide as Catalyst Reactions with Change of the Carbon Skeleton Reaction conditions

Reaction type Isomerization

Splitting of paraffins

Splitting of naphthenes

Feed

Product

Isobutane Methylcyclopentane MethylcycloBenzene pentane Lower boilIsooctane ing hydrocarbons Gasoline Paraffins 260"-320°C. 180°C. E.P. Lower boiln-Heptane ing hydrocarbons Lower boilDecahydroing hydronaphthalene carbons

n-Butane Cyclohexane

Feed parFeed tial rate, Temp., Pres., pres., kg./ Conver"C. atm. atm. (l.)(hr.) sion, % 400

408

200 250

160 27

0.5 1.0

35 90

408

250

27

1.0

90

400

250

18

1.5

40

408

220

5

1.0

90

435

250

20

1.5

15

408

220

5

1.1

90

be used and side reactions, like splitting into gaseous hydrocarbons, occur t o only a minor extent. Like other hydroforming catalysts, the nickel.tungsten sulfide loses activity with time; however, i t cannot be regenerated by simple burn-off of the deposited hydrocarbons. The loss of catalyst activity appeared more rapid if the content of nonhydroaromatic hydrocarbons in the feedstock was high. To effect changes in the carbon skeleton of the feed with tungsten sulfide catalyst, temperatures of about 400°C.and higher are needed. Examples of such reactions obtained with pure compounds are shown in Table V. Isomerization reactions with tungsten sulfide catalysts are

COAL-HYDROGENATION VAPOR-PHASE

CATALYBTS

253

possible. Normal butane can be converted to isobutane. For this reaction a high partial pressure of the butane can be used. The conversion to isobutane reaches the thermodynamic equilibrium. Practically no splitting to lower molecular-weight hydrocarbons was observed under these reaction conditions with butane. Isomerization of higher paraffins was obtained together with splitting to lower molecular-weight hydrocarbons. I00

HANE

HANE

00 $

P

240 P

20 5 10 15 20 25 TIME HOURS

0

-

FIQ. 4. Splitting hydrogenation with WSI catalyst at 200 atmospheres pressure: Propane. Temp. 460'C.

T I M E - HOURS-

FIG. 5. Splitting hydrogenation with ws, catalyst at 200 atmospheres pressure: n-Heptane. Temp. 425°C.

NE

0

5 10 15 20 TIME-HOURS

25

FIG.6. Splitting hydrogenation with M's2 catalyst at 200 atmospheres pressure: Cetrtne. Temp. 425°C.

TIME - H O U R S

FIG.7. Splitting hydrogenation with WS, catalyst at 200 atmospheres pressure: Isooctane. Temp. 390°C.

The converaion of cyclohexane to methylcyclopentane is nearly complete at 408°C. Under identical conditions benzene is also converted to methylcyclopentane. This shows that the hydrogenation of benzene t o cyclohexane, which can be considered the first step in this reaction, is very rapid. At approximately the same reaction conditions, isooctane can be split, whereby isobutane is the main product. Similarly, a paraffinic oil boiling between 260" and 320°C. and decahydronaphthalene are converted almost completely into lower boiling hydrocarbons. Splitting of normal heDtane, however, is slower. Even at 435°C. the conversion to

254

E. E. DONATH

lower boiling hydrocarbons is only 15 %. Decahydronaphthalene is split almost completely at 408°C. Similar reaction products and rates are obtained with naphthalene in pilot plant experiments. I n commercial plants removal of the heat of reaction is often the controlling factor. Hydrogenation of naphthalene, because of the greater heat of reaction, would then require a lower throughput than hydrogenation of decahydronaphthalene. To obtain further data for the mechanism of hydrocarbon splitting, autoclave experiments (17) were made with several paraffinic hydrocarbons with tungsten sulfide in the presence of hydrogen at about 200 atm. pressure. A shaking type of autoclave, charged with 2.3 kg. tungsten sulfide pellets, was used and the catalyst was confined by screens. The autoclave was charged with 100 g. of the hydrocarbon used and pressurized with hydrogen, It was then heated and after it had reached reaction temperature, gas samples were taken a t suitable intervals. Ethane, propane, heptane, cetane, and isooctane were used as feed materials. It was found that the reaction rate depends greatly upon the length of the hydrocarbon chain and chain branching. Some results of experiments with propane, n-heptane, cetane, and isooctane are shown in Figs. 4 t o 7. T he data for heptane, compared with the rate of splitting shown in Table V, indicate that the reaction rate in these autoclave experiments is slower than in experiments in a converter using a flow of heptane vaporized in a stream of hydrogen. 3. Reaction Mechanism

The autoclave experiments permit speculation about the mechanism of splitting. The content of isobutane in the butane fraction obtained from straight-chain paraffins is about 50% and decreases with temperature in the range investigated. It is above the equilibrium value, which in this temperature range corresponds t o a n isobutane content of the butane fraction of about 33%. The high isobutane content therefore indicates that splitting of the higher hydrocarbons such as n-heptane is preceded by isomerization as the slowest reaction step. Such a reaction mechanism explains the high isocontent of the butane fraction obtained from higher paraffins. This is also in accord with the general experience that higher paraffins cannot be isomerized on tungsten sulfide catalyst without simultaneous splitting. Isoparaffins split directly, and the splitting point is a t the carbon atom with the lowest hydrogen content. Thus a high yield of isobutane is obtained from isooctane. Since the isooctane used in these experiments was pure, it is probable that some of the isobutane formed by primary splitting had been isomerized to normal butane. One possible reaction mechanism for the splitting of straight-chain

COAL-HYDROGENATION

255

VAPOR-PHASE CATALYSTS

paraffins is that a methyl group is introduced by isomerization in the 2 position and, to a lesser extent, in the 3 position. Splitting of isomerized paraffins with more than four carbon atoms, then, is faster than the isomerization, For normal heptane the following scheme has been suggested. The observed products are underscored. a. Main reaction via 2-methyl isomerization :

n-C 7H 16

isomerization

c c-c-c-c-c-

-C

isomeriaation 4

C

c-c-c-c-c

C

b. Side reaction via 3-methyl isomerization :

-/ isomerizattion

n-C,Hle

c

C-C-C-C-C-C

+HZ/ / k !

+

'\

\+Hz \

L

C3Hs n-CqHio % -~ Further conversion as in main reaction

n-CsH12

J.

+

CzHs C3H8 -

The rate of splitting of straight-chain paraffins depends on the length of the hydrocarbon chain, as shown in Fig. 8, which indicates the halflife for the hydrocarbons used in the autoclave experiments. Ethane is practically stable a t 460"C., and propane splits with a half-life of about 30 hr. at this temperature. A t 425°C. propane was found practically stable, and the half-life of butane was 5 to 10 hr. At the same temperature normal heptane splits in 1 hr., on the average, into two fragments, and cetane into four fragments. From the temperature coeffeient of the splitting reactions an apparent activation energy of the order of 20 kcal. can be calculated. The splitting of isooctane is about twenty times more rapid than that of normal heptane at the same temperature. The splitting hydrogenation of decahydronaphthalene a t 300 atm. pressure with tungsten sulfide catalyst at 408°C. gave an insight into the mechanism of the conversion of this condensed ring system. The results suggest that splitting is preceded by isomerization of decalin into methyl-

256

E. E. DONATH

indane followed by the opening of the five-membered ring. The cyclohexane homologs formed are then further isomerized t o cyclopentane homologs. Although individual hydrocarbons were not identified, the hydrocarbon-group analysis and the physical constants of the liquid fractions indicate, as expected, negligible formation of paraffins; however, some gaseous paraffins were obtained. The fractions boiling below 155°C. consisted principally of cyclopentane and cyclohexane homologs with fewer than ten carbon atoms and several side chains. The 155" to 175°C. 0.1 -

020.5 -

ISOOCTANE

I-

CETANE

2510 W

20

-

50100200-

500

f

0

IS0 n +IS0

-

loooL i60

BUTANE

ETHANE

a'80

4bO 420 TEMP. .C

4;O

4b

FIG.8. Splitting hydrogenation 0f:paraffins with WS2 at 200 atmospheres hydrogen pressure.

fraction consisted of C10H20compounds containing one ring, and the residue boiling above 175°C. consisted of decahydronaphthalene isomers in addition t o unchanged feedstock. Cawley (18) described a somewhat different mechanism for the splitting hydrogenation of decahydronaphthalene. According to his data, naphthalene is hydrogenated to decahydronaphthalene, whereupon both rings are isomerized and then split. The experiments in the preceding paragraph, however, indicate that both cyclohexane and cyclopentane homologs were obtained as splitting products. This means that splitting

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

257

with WS2 catalyst may begin a t the five-membered ring, while the sixmembered ring stays intact.

4 . Factors Aflecting Catalyst Activity It has been shown that the splitting rate of paraffins increases with chain length and branching. The hydrogenation and splitting of aromatic rings such as benzene or naphthalene is a reaction of similar velocity. Different results are obtained with condensed] high-molecular-weight aromatics. The thermodynamic equilibrium for the hydrogenation of aromatic hydrocarbons in general is less favorable with increasing molecular weight and number of fused rings. The tendency for condensation reactions also increases, and a change in reaction conditions is necessitated for the hydrogenation of higher condensed aromatics in comparison with oils boiling below 325°C. It is one of the basic reasons for separating the coal-hydrogenation process into liquid- and vapor-phase stages. As an example, data of a somewhat qualitative nature for coronene (CZ4Hl2) will be used (19).The experiments were made in a rotating autoclave containing 300 mi. (800 g.) of pelleted tungsten sulfide catalyst. The coronene was treated in decahydronaphthalene solution. At 287°C. and 250 atm. hydrogen pressure and with 200 g. of eoronene dissolved in 800 g. of decahydronaphthalene, about 97 % of the coronene is converted to perhydrocoronene. With higher coronene concentrations the catalyst activity is impaired and conversion to perhydrocoronene decreases. At 600 to 800 atm. hydrogen pressure of as much as 300 g. of coronene can be used in 700 g. of decahydronaphthalene, and about 99 % of the coronene is perhydrogenated. At higher temperatures hydrogenation is less complete. At 408°C.and 250 atm. hydrogen pressure perhydrocoronene is dehydrogenated to a complex consisting of one molecule each of coronene and perhydrocoronene, and at stiIl higher temperature dehydrogenation proceeds further. Comparison with Tables IV and V shows that under similar reaction conditions naphthalene is rapidly hydrogenated and split. At 600 to 800 atm. hydrogen pressure and 450"C., perhydrocoronene is not dehydrogenated noticeably; however, some splitting occurs. These experiments show that coronene can be hydrogenated in the same temperature range as, e.g., naphthalene, but splitting is by far slower. It seems, however, that this is due not to higher stability of the coronene itself, but rather to a poisoning of the catalyst activity. It can be speculated that at hydrogen pressures above 800 atrn., pressures which would prevent condensation reactions, coronene could be split easily. Support for this assumption is the observation that the presence of coronene decreases the splitting activity of the WS2 catalyst toward other

258

E. E. DONATH

molecules. Results of autoclave experiments with sarious amounts of coronene in decahydronaphthalene solution are shown in Fig. 9. The strong inhibiting effect of coronene on the splitting activity of the catalyst a t 600 t o 700 atm. hydrogen pressure is evident. At 250 atm. hydrogen pressure the inhibiting action is even stronger; about 100 g. of coronene reduces conversion of decahydronaphthalene from 100 t o 10%, but a t 600 atm. the decrease is only 50%.

9. CORONENE /1000ml OECAHYORONAPHTHALENE

FIG.9. Splitting of decahydronaphthalene in the presence of coronene. Autoclave reaction conditions: Temperature, 425°C.; Hydrogen pressure, 600-700 atm.; Catalyst, 300 cc. (800 g.) WS,; Charge, 1000 ml. decahydronaphthalene and 0-600 g. coronene.

Catalysts used in splitting experiments with coronene contain about 5% carbon which cannot be removed by extraction with benzene, a solvent for coronene. They have also lost their activity t o a great extent. To prevent rapid loss of activity of splitting hydrogenation catalysts, conditions must be selected under which the thermodynamic equilibrium does not permit formation of condensation products. This can be achieved by high hydrogen partial pressure. However, for hydrogen partial pressures in the range up t o 700 atm. the concentration of compounds like coronene must, in addition, be kept sufficiently low t o avoid condensation reactions. Probably the same reaction, the dehydrogenation, condensation, and irreversible adsorption of highly condensed aromatic hydrocarbons, causes the loss of activity of the WS2 catalyst when used a t pressures well below 200 atm. As an example (go), in experiments with a paraffin-base petroleum-oil fraction boiling between 200" and 325°C. the results in the accompanying tahulation were obtained :

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

Hydrogen pressure, atm. 200

100 First day Sixth day Return to 200 First day Final value after several days

259

Conversion to gasoline, % 90 (constant value)

45 22 27 42

The catalyst activity, measured b y the conversion t o gasoline, drops immediately after the pressure is decreased from 200 to 100 atm. and drops further very rapidly during the 6-day, 100-atm. operation. After return t o 200 atm. pressure, conversion to gasoline increases slowly and remains well below the original activity. Thus poisoning of the catalyst by operation a t 100 atm. pressure is partly reversible and partly irreversible. The reversible part of the catalyst deterioration changes gradually into irreversible damage. For instance, after short operation at 150 atm., the catalyst activity is recovered completely by operating it again a t 200 or 250 atm. pressure, while prolonged operation (e.g., for 1 week) at 150 atm. causes permanent activity decrease. The lowest pressure a t which high catalyst activity can be maintained for months depends upon operating temperature, throughput, hydrogen-oil ratio, and boiling range and aromaticity of the feedstock. For the prehydrogenation of coal middle oil and for the splitting hydrogenation, this pressure seems t o be just slightly below 200 atm. The reduction of the activity of hydrogenation catalysts by nitrogen and oxygen compounds has been mentioned. Experiments for the increase of the cetane number of Diesel fuels by hydrogenation with the WS2 catalyst indicate that ammonia or nitrogen compounds in this case may be beneficial, as they reduce the splitting rate t o a greater extent than the pure hydrogenation activity. The splitting activity of the WSZ catalyst is not noticeably affected by the addition of nitrogen bases (aniline) up t o a nitrogen content of 0.02% based on the feed. Increase of the organic nitrogen content t o 0.05% on the feed decreased the splitting rate considerably. A temperature increase of around 20°C. was necessary to maintain the same conversion t o gasoline as without aniline addition. The WS2 catalyst is less sensitive to NHB-forniing compounds than strongly splitting catalysts on fuller’s earth as support; the latter have a weaker hydrogenating activity. It is more sensitive, however, than the Mo-Zn-Mg catalyst. It appears that the carbonium-ion-splitting mecha-

260

E. E. DONATH

nism is especially sensitive to ammonia or organic nitrogen compounds. Similarly, catalytic cracking catalysts lose their activity in the presence of nitrogen compounds. On the other hand, the activity of WSz is enhanced by the presence of hydrogen sulfide, and its presence is necessary t o maintain the activity of the catalyst, especially when oxygen-containing oils are treated. Usually, a sulfur content of the feedstock of 0.2% or higher is sufficient t o prevent a decrease of the WS2 catalyst activity. Higher sulfur contents up t o 2% increase the catalyst activity b y 10 to 20%, but a further increase has little influence. Increased oxygen content of the feed seems to increase the required amount of sulfur. I n commercial use of WS2 as a prehydrogenation catalyst, sulfur addition was not necessary with most commercial sulfur-containing oils. Sulfur as such, or as hydrogen sulfide, was added t o prehydrogenated, practically sulfur-free feedstocks of splitting catalysts. Because of the corrosion caused by hydrogen sulfide, its addition was kept as low as permitted by catalyst activity. The mechanism of the catalyst activation by sulfur is not understood. The amount of sulfur compounds necessary t o maintain or increase the catalyst activity depends in some cases on the stability of the heavymetal sulfide component of the catalyst. Thus molybdenum sulfide seems t o require a higher hydrogen sulfide concentration than tungsten sulfide. However, some catalysts that do not contain elements that can form sulfides under reaction conditions also showed an increased activity when sulfur compounds were added t o the feed. Hydrogen sulfide in many cases decreases the catalyst sensitivity t o nitrogen compounds and thus causes a n activity increase. Sufficient data for pure compounds are not available t o permit segregation of these effects. 6. Injluence of Catalyst-Pellet Size

Tungsten sulfide as vapor-phase catalyst is used in the form of pellets which are pressed from the powder in commercial tableting machines under a pressure as high as 5000 kg./cm.2 (75,000p.s.i.). Three-millimeter pellets were used in pilot plants and 10-mm. pellets in commercial plants. The feedstocks used for vapor-phase hydrogenation are mixtures containing a range of compounds of varying molecular weight (boiling range). Therefore, differences in the course of the vapor-phase reaction are t o be expected with different sized catalyst pellets if one considers th a t the bulk of the reaction occurs in the interior of the catalyst pellets, which can be reached by the feedstock molecules only by diffusion or surface migration. Experiments ( 2 1 ) were made with tungsten sulfide catalyst of different sizes, as shown in Table VI, to evaluate this variable. The 3- and

26 1

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

10-mm. pellets were made from the same batch of WSZ powder and, t o obtain a smaIler size, the 10-mm. pellets were crushed and the 2- to 4-mm. screen size used for the experiments. The average height of these particles was considerably smaller than their width of 2 to 4 mm. and the average weight, therefore, less than one half of the 3-mm. pellets. TABLE VI W S z Catalyst Properties ~

~~

10 mm. Pellets crushed to 2 to 4mm. size

Catalyst in

10-mm. Pellet

3-mm. Pellet

Pellets Height, mm. Diameter, mm. Volume, ml. Weight, g. Apparent density, g./ml. Bulk density, g./ml. Voids between pellets, % Pore volume of pellets assuming WS2 specific gravity 7.5, %

9.8 10.2 0.8 3.0 3.8 2.3 40

3.9 3.1 0.03 0.14

Irregular shape

4.6 2.8 40

3.8 2.0 47

49

38

49

With these three catalysts, experiments were made for the (a) prehydrogenation of bituminous-coal middle oil and ( b ) splitting hydrogenation of petroleum gas oil a t 250 atm. pressure. Because of the differences in bulk density of the catalyst in this case, the throughput per weight of catalyst was kept constant, although usually catalysts are compared with equal throughputs per volume of catalyst. The properties of the feedstocks are shown in the accompanying tabulation. The operating condi-

Specific gravity, 20°C. Initial B.P., "C. Final B.P., "C. Aniline point, "C. Tar acids, % Nitrogen, %

Bituminous-coal middle oil

Petroleum gas oil

0.970 180 325 -20 15 0.6

0.820 180 330 70 0 0

tions and results are shown in Table VII. I n the prehydrogenation experiments it can be seen th a t the hydrogen

262

E. E. DONATH

content* of the residue above 180°C. increases with decreasing catalystpellet size. The hydrogen content of the gasoline, as well as conversion to gasoline and its low boilers (- 100°C.) content, shows the opposite trend. Hydrocarbon gas formation as percentage of feed is small and essentially constant. The increased degree of hydrogenation of the higher fractions with decreasing catalyst size, shown by the aniline points of gasoline and residue, can be more clearly seen in Fig. 10 for 20°C. fractions. TABLE VII Prehydrogenation and Splitting Experiments with WS2 Catalysts of Different Sizes at 250 Atm. Pressure and 0.35 Kg. of Feed/(Kg. Catalyst)(Hr.) ~~

Prehydrogenation of bituminous-coal middle oil Temperature, "C. Hydrogen, m.s/kg. oil Catalyst size, mm. Gasoline, % in product Gasoline production rate kg./(kg. catalyst) (hr.) kg./(l. catalyst)(hr.) Hydrocarbon gas formation % of feed % of converted feed Gasoline 180°C. E.P. Specific gravity, 20°C. Aniline point, "C. % Boiling u p to 100°C. % Aromatics Residue above 180°C. Specific gravity Aniline point, "C. Final B.P., "C.

10 38

400 4 3 38

2-4 29

0.12 0.13 (0.29) (0.36)

0.10 (0.20)

3

4

-

4

-

-

~

~~~

~

~~~~

Splitting hydrogenation of petroleum gas oil

10 59

0.19 (0.43)

382 3 3 62

2-4 68

0.20 0.21 (0.56) (0.43)

14

12.5

11

0.766 37 18.5 15

0.777 34 14.0 16

0.784 29.5 12.5 19

0.717 62.5 29 1

0.726 62.5 21 1

0.725 62.5 21 1

0.858 48.5 306

0.851 54.5 301

0.851 55.0 305

0.785 76.0 295

0.783 76.3 284

0.776 77.0 265

For the splitting hydrogenation of petroleum gas oil, a lower temperature than that required for the prehydrogenation of bituminous-coal middle oil with smaller conversion t o gasoline was sufficient. This is caused mainly by the higher nitrogen content of the bituminous-coal middle oil, which decreases catalyst activity. As shown in Table IV, aromatic rings like naphthalene are saturated with WSZ catalyst a t a temperature of 335°C.

* The aniline point (critical miscibility temperature with aniline) increases with increasing hydrogen content of a hydrocarbon fraction and thus can be used as an index of paraffinicity. It increases also a t the same hydrogen content with a n increase in boiling range. An approximate figure for the aniline point of hydrocarbons in the gasoline boiling range is paraffins, 70°C.; o l e h s , 40°C.; aromatics, below 0°C.

COAL-HYDROGENATION VAPOR-PHASE

CATALYSTS

263

With decreasing catalyst size the following changes can be observed. Formation of gaseous hydrocarbons (methane to butane) and low boiler (- l0OOC.) content of the gasoline decrease, and conversion t o gasoline increases. The hydrogen content of the residue above 180°C. indicated by specific gravity and aniline point increases, and its final boiling point decreases.

AVERAGE BOILING POINT OF FRACTIONS OC

FIG. 10. Influence of WSZ catalyst size on aniline points of prehydrogenation product fractions.

From these experiments, in which the range of catalyst-particle diameter (from 10-mm. pellets to crushed pellets of 2 t o 4 mm. width and about 1 mm. height) has been changed by a factor approximating ten, the following can be concluded: a. The total number of individual hydrogenolysis steps on the catalyst for which, for example, the hydrogen consumption could be used as an index, is, within experimental accuracy, apparently independent of the catalyst size in the observed range. I n the prehydrogenation experiments a lower hydrogen content of the gasoline is accompanied by a higher hydrogen content of the residue. The corollary in the splitting experiments is that a lower conversion to gasoline is accompanied by a higher formation of gaseous hydrocarbons and low boilers in the gasoline, which requires several splitting steps. b. With increasing catalyst size, conversion of the higher boiling, higher molecular-weight fractions is retarded. Thus it can be concluded th at diffusion of the larger molecules into

264

E. E. DONATH

the center of the larger pellets is slowed down. However, these inner parts of the catalyst are available for reactions of the smaller molecules of the feedstock or of splitting products. It is not believed that the results are influenced to any noticeable extent by higher interior temperature of the larger pellets. If this had been the case, greatly increased hydrocarbon gas formation in prehydrogenation as well as splitting experiments would have been expected. I n addition, this effect should have been more pronounced in prehydrogenation because of the greater heat of reaction. I n the paper by Reitz (21) an attempt is made to estimate the diffusion coefficient of feed molecules in the catalyst pores. It is probable that diffusion in the vapor phase is not rapid enough t o explain the observed reaction rates and that surface diffusion must be an important factor. Other experiments indicated th at extrapolation of these results t o other catalysts and especially supported catalysts which might have a different structure and porosity is not permissible. With larger (15 mm.) WSz pellets, also, a definite decrease in catalyst activity was observed.

Iv. NONSPLITTING

CATALYSTS

The tungsten sulfide catalyst permitted conversion of middle oils into gasolines of any desired boiling range a t high conversion rate and with high yield. The disadvantage of the catalyst was the comparatively low octane number of the gasoline. The branching of the paraffinic hydrocarbons obtained with WS2 catalyst was insufficient t o compensate for the loss in knock resistance caused by the essentially complete saturation of aromatics. As mentioned in Sec. 11, catalysts were developed th a t produced gasolines of higher octane number than were made when WS2 was used. These catalysts were poisoned by nitrogen bases and t o some extent by oxygen compounds. It was therefore necessary to remove nitrogen and oxygen from coal middle oils. A separate hydrogenation step, the prehydrogenation, was used for this purpose. It preceded the second vaporphase stage, in which middle oil underwent splitting into gasoline. The prehydrogenation thus is a combined refining and hydrogen addition step for middle oil and is similar t o the hydrogenation refining of gasoline or coke-oven light oil; the upgrading of kerosene, Diesel fuel, or lubricating oils; or the hydrogenation of recycle oils from catalytic cracking. The purpose of hydrofining reactions may vary with the feedstock. Gasolines or coke-oven light oil are treated to remove gum formers, th a t is, readily polymerizing olefins and sulfur compounds, while hydrogenation of aromatics must be avoided. Hydrogenation of kerosenes and Diesel

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

265

fuels has as its main purpose the saturation of aromatics to improve burning properties and the removal of sulfur. Upgrading of catalytic cracking recycle oils is achieved by saturation of aromatics and reduction of nitrogen compounds, which reduce the activity of cracking catalysts. This is analogous t o the sensitivity of the splitting hydrogenation catalysts t o nitrogen compounds. These different requirements can be fulfilled by different catalysts, and different operating conditions will be required. I n general, high pressure increases the rate of the saturation of aromatic rings. Therefore, for the hydrofining of gasolines or coke-oven light oil, pressures below 50 atm. are used. TABLE VIII Temperature of Prehydrogenation Reactions with W S ZCatalyst Reaction Tar-acid reduction Tar-base reduction Hydrogenation Splitting *At 380%.

Begin

Complete

310 310 310 350

380 380 max. 410’-460°C.

*

12 % gasoline boiling below 150%. formed by splitting hydrogenation.

Similar reactions are the hydrogenation of lubricating oils, whereby oils with a flatter viscosity-temperature curve are produced, as well as low-temperature hydrogenation of brown-coal tar, which became known as the “TTH” process. Although in these two processes the bulk of the feedstock remains liquid and reacts as liquid with hydrogen on the fixedbed catalyst, the path of the reaction is the same as in the vapor-phase operation. As the experiments described with WS2 catalyst have shown, the splitting rate increases with the chain length of the paraffins. Therefore these high-boiling feedstocks, even under mild conditions, undergo considerable splitting. Saturation-addition of hydrogen-is accompanied by splitting reactions and partial reduction in molecular weight. With the introduction of the two-stage vapor-phase operation in 1935, the WSz catalyst was used a t first as a prehydrogenation catalyst. The cost of this catalyst and the properties shown in Table V I I I were the reasons for the development of other prehydrogenation catalysts. The data in Table VIII were derived from experiments for the prehydrogenation of bituminous-coal middle oil a t 250 atm. hydrogen pressure and a throughput of 0.8 kg. of feed/liter of catalyst. Practically complete removal of nitrogen compounds is needed for the splitting catalyst. Therefore, a prehydrogenation temperature above th a t

266

E. E. DONATH

at which splitting begins must be used. Since the gasoline formed by the prehydrogenation catalyst has a lower octane number than th a t obtained from the splitting stage, a prehydrogenation catalyst with minimum splitting activity is desirable. Such catalysts were found in combinations of tungsten sulfide with nickel sulfide and a less active WS2-FeS combination; however cheaper catalysts were also developed. An evaluation of some selected supports ( 2 2 ) as t o splitting of gas oil at atmospheric pressure gave results shown in a qualitative way in Table IX. The supports are arranged in order of increasing splitting activity. TABLE IX Splitting Activity of Catalyst Supports

Support MgO A1208

Fez08 SiOa-FetOa SiOl gel (large pores, precipitated at high pH) SiOz gel (small pores, low pH) AlzOa H F SiOa A1203

+ +

Splitting activity

Low Low Idm Low

(1%) (3%) (3%) (3%)

Remarks

Marked gas formation

Some (10%) Medium (20%) Strong (35%) Strong (45%)

Vapor-phase hydrogenation results and experimental evidence of this type lead t o the conclusion that catalysts on “basic” supports are suitable for nonsplitting prehydrogenation-type reactions and that “ acidic” supports are best used for splitting catalysts. Activated alumina was found t o be the best support because of rapid reduction of tar acids. Especially, alumina precipitated from aluminum salts at constant p H was satisfactory and produced catalysts th a t could be formed into pellets of high mechanical strength. The first suitable activated-alumina catalyst contained 10% MoOa activated by addition of 3% nickel oxide. This catalyst was superseded by a more active tungsten catalyst containing 70% activated alumina, 27 % tungsten sulfide, and 3% nickel sulfide. This catalyst is used in commercial plants for the prehydrogenation of middle oils and also for the direct hydrogenation of shale oil and lignite tar (TTH process). The influence of varying amounts of the active component in Mo-Ni catalysts on alumina is shown in Fig. 11. The prehydrogenation tests 17% were made with a bituminous-coal middle oil (aniline point-l2’C., t a r acids, 0.75% Nz) with the addition of 1% CS2. Carbon bisulfide is

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

267

rapidly converted into HzS and CHd on the catalyst and provides a convenient means of supplying hydrogen sulfide for catalyst activation. Addition of sulfur in pilot-plant experiments was required to duplicate more closely the conditions in commercial plants. I n commercial plants, hydrogen is recycled. With this hydrogen the hydrogen sulfide obtained from sulfur in the feed is also recycled, and thus a higher H2Spartial pressure

g COMPONENT/LITER ALUMINA

FIQ. 11. Mo-Ni prehydrogenation catalysts supported on activated alumina. Components: I. Ni203. 11. Mool. 111. MooI with constant 20 g./liter Ni203.

is obtained than in pilot plant operation without hydrogen recycle. These tests were made a t 250 atm. pressure, 430°C., an hourly, throughput of 0.8 kg. of feed/liter of catalyst, and 3 cbm (NTP) of hydrogen/kg. of feed. The aniline point of the middle oil obtained was used as a measure of the catalyst activity. I n those tests molybdenum and nickel were added assoxides and converted into sulfides during the reaction. The activity of these catalysts increased t o a constant value during the first 3 days because of conversion into sulfide. This activity was the same as observed for a few points with reference catalysts prepared from sulfides. As shown in Fig. 11, alumina alone does not hydrogenate the middle oil. Addition of nickel alone (curve I) gives insufficient hydrogenation. Ad-

268

E. E. DONATH

dition of increasing amounts of Mooa (curve 11) produces catalysts with increasing hydrogenation as well as refining activity. However, a limiting value is being approached. The most active molybdenum-aluminum catalyst produces middle oil of insufficient purity with an aniline point below that obtained with the WS2 catalyst (44°C.). The combination of molybdenum and nickel (curve 111) shows great improvement and pro-

0

* 10

0

FIG.12. Influence of Ni content on alumina-supported Mo and W prehydrogenation catalysts. I. 150 g. MoOa/liter constant. 11. 250 g. Wo03/liter constant.

duces a satisfactory prehydrogenation catalyst. However, a catalyst containing 27% WS, and 3% NiS is more active. The activity of such molybdenum and tungsten catalysts with varying amounts of nickel is shown in Fig. 12. The molybdenum catalyst shows maximum hydrogenation with a nickel content of 20 g. of Ni203/liter,and the activity of the tungsten catalyst increases slightly with nickel contents above that value. To investigate the splitting activity, experiments with bituminouscoal middle oil were made under the conditions given in the preceding paragraph, except for varying temperatures. The differences in the splitting activity of four catalysts are shown in Fig. 13. As a measure of the splitting activity, the percentage boiling t o 150°C. and the product end

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

269

point were selected. The lowest temperature a t which a product was obtained t ha t could be satisfactorily split in the second vapor-phase stage is indicated. I n comparison with th at of the WS2 catalyst, the splitting activity is greatly reduced by addition of nickel and by the aluminum c

I

a 0 9 320

{Al- W- Ni

TEMP.%

FIQ.13. Temperature and splitting activity of prehydrogenation catalysts.

support. A singular behavior was observed with the tungsten nickel catalyst. It refines satisfactorily a t a lower temperature than does WS2 and has about the same splitting activity a t temperatures up t o 425"C., but it loses its splitting activity within a week a t this temperature. The cause of this conversion point is not known and X-ray analyses gave no indication of changes; however, electron diffraction indicated the

270

E. E. DONATH

presence of only WS2 in fresh WNi catalyst, but in the used catalyst the lines of NiS were also observed. I n the same experiments the degree of hydrogenation of the middle oil has been determined. Figure 14 is a graph of the aniline point of the prehydrogenation middle oil (180' t o 325°C.) for the four catalysts a t varying temperatures. The aniline point increases with reaction temperature

"\\

FIG.14. Temperature and hydrogenating activity of prehydrogenation catalysts.

t o temperatures between 400" and 440°C. The temperature a t which hydrogenation begins is different for the four catalysts, as is the rate of aniline-point increase with temperature. The steep rise for the WS2 catalyst between 360" and 380°C. might be connected with the splitting action which becomes noticeable for this catalyst above 350°C. The curves for all the catalysts show that satisfactory refining, that is, reduction of tar acids and bases, is reached a t the beginning of the plateau of maximum hydrogenation. I n this temperature range aromatics are essentially saturated. With further increased temperatures the extent of hydrogenation of aromatics decreases. This decrease is due t o increased stability of the aromatics in the aromatic-hydroaromatic equilibrium in spite of the high hydrogen partial pressure. Differences in the middle-oil aniline point

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

271

a t temperatures above 450°C. are probably due to the differences in splitting, which produce middle oils of a different boiling range. The alumina-tungsten-nickel catalyst thus shows a much smaller splitting activity than the WS2 catalyst a t the same level of refining and hydrogenation, in spite of somewhat higher reaction temperature. These characteristics of the catalyst also appear in the hydrogenation of lubricating-oil fractions and the direct hydrogenation of brown-coal tar. The commercial hydrogenation of lubricating-oil fractions was introduced by the Standard Oil Company of New Jersey (23). I t s main purpose is obtaining an oil with a flatter viscosity-temperature curve, a higher viscosity index. This is achieved by the saturation of aromatic rings of the feed. Splitting in this process is undesirable as it reduces the yield of high-boiling lubricating-oil fractions. TABLE X Comparison of A1203-W-Niand W S ZCatalysts Catalyst Temperature, "C. Vacuum residue, yield % Viscosity 'Engler at 100°F. 'Engler at 210°F. Viscosity index

wsz 374 51.2 47.5 3.12 83

AlzOa-W-Ni 391 70.0 47.4 3.12 83

A comparison of the two catalysts (24) operated under such conditions t ha t oil of the same viscosity index is obtained is shown in Table X. These experiments, made a t the same throughput with a propane raffinate from German crude oil, show that, in spite of higher operating temperatures, a higher lubricating-oil yield of the same viscosity index and viscosity is obtained with the A1203-W-Ni catalyst, in comparison with WSZ. The two catalysts may also be compared under conditions that produce the same yield of oil of a given flash point. Under such conditions the A1203-W-Ni catalyst again requires higher temperature but produces oils of a higher viscosity index. As expected, increase of the hydrogen pressure, for example, from 250 t o GOO atm., improves the viscosity indexyield relationship with both catalysts, Increase of the hydrogen pressure increases the rate of nonsplitting hydrogenation reactions and has little influence on the splitting reactions. Mild hydrogenation of lubricating-oil fractions can also be used t o improve the color, odor, acid value, carbon residue, and sulfur content and to lessen the tendency to emulsify. A commercial application of such a

272

E. E. DONATH

process developed by Imperial Oil, Ltd., has been recently described (226). An alumina molybdena catalyst is being used. The hydrogenation replaces acid treatment; it provides high yield and eliminates acid-waste disposal problems. Similarly, a lower splitting activity of the A1203-W-Ni catalyst in comparison with WSZ has been observed in the TTH process, the lowtemperature hydrogenation of lignite tar for the preparation of Diesel TABLE XI Influence of the Hz-Feed Ratio in the Hydrogenation of Brown-Coal Tar with W& and AhOX-W-Ni

+

Pressure: 280 atm. (80% H2, 19% CH, Nz, 1% HzS) Throughput: 0.375 kg. tar/(liter catalyst)(hr.) Feed : Specific gravity, about 0.92; tar acids 10%;middle-oil aniline point 32°C. Product

Cbm. gas/kg. feed

Spec. Middle41 grav. % Tar acids aniline point, "C.

% -250'

WSZcatalyst, 331'C. 2.7 10.0 WSz catalyst, 314°C. 2.7 10.0 AlzOs-W-Ni catalyst, 327°C. 0.7 1.3 2.7 3.3 5.0 10.0

0.837 0.840

0.05 0.16

66 64

32 30

0.889 0.888

3.4 2.5

43 41

22 22

0.861 0.864 0.860 0.866 0.866 0.861

0.10 0.05 0.10 0.05 0.08 0.09

51 52 53 51 52 53

27 26 27 26 27 24

fuel, lubricating oil, and paraffin wax. A recent (26) publication contains confirming data obtained over a range of temperatures, throughputs, and hydrogen-feed ratios. The data obtained with different hydrogen-feed ratios are shown in Table XI. The hydrogen rate used in commercial operation is about 2.7 cbm/kg. of feed. Change of the hydrogen rate over a wide range has no great effect on the result. Essentially the same result has been obtained for the prehydrogenation of middle oils. However, in splitting hydrogenation reactions, especially a t temperatures above 450°C. the hydrogen-feed ratio has a noticeable effect with a flat optimum a t 2.5 cbm. of Hz/kg. of feed.

COAL-HYDROGENATION VAPOR-PHASE

273

CATALYSTS

The activity of alumina-supported catalysts for the refining of highboiling feedstocks such as petroleum residues can be increased by adding silica t o the alumina ( 2 7 ) .Such an alumina catalyst containing 20% SiOz, 10% MOOS,and 3% Niz0 3has been described as suitable for the treating of a petroleum residue at 250 atm. pressure and 420°C. Dewaxing of the treated residue permits separation of a pure paraffin wax. The activating effect of nickel in these catalysts is highly specific. Replacing nickel by iron produced catalysts of decreased hydrogenating and refining activity (28). Results obtained with bituminous-coal middle oil are shown in Table XII. The same experimental conditions were used as indicated for Fig. 11. TABLE XI1 Iron and Cobalt Replacing Nickel in the AL2O3-WS2-NiS(70-27-3) Prehydrogenation Catalyst Middle oil above 180°C. Iron-group metal Nickel Cobalt Iron

Product, % 180°C.

Aniline point, "C.

23 23 20

47 11.5 -5.5

Tar bases as Phenols, % mg.NHt/liter 0.01

0.03 0.05

3 40 700

Cobalt takes an intermediate position between nickel and iron. Iron does not increase the hydrogenating activity of the tungsten alumina combination (see Fig. 12, curve II), but cobalt increases it t o some extent. Tar-acid reduction, while decreasing from nickel to cobalt to iron, is still almost complete. However, reduction of tar bases shows greater differences. I n rough figures the remaining tar bases amount to 0.1, 1, and 10% of those in the feedstock for the nickel, cobalt, and iron catalyst, respectively. Thus the cobalt-containing catalyst has good refining activity with only a fraction of the hydrogenating activity of the nickel catalyst. These great differences between the metals of the iron group lead t o further study of partial replacement. Results are shown in Fig. 15 for mixtures of nickel cobalt and nickel iron in varying proportion as activators for WSZ alumina. These mixtures show additivity as far as taracid and -base reduction is concerned. The hydrogenating activity for the nickel cobalt combination is also additive. However, use of u p t o 20% iron in place of nickel does not decrease the activity of the nickel catalyst. With higher percentages of iron and decreasing nickel content, the hydrogenating activity drops t o the value of the iron catalyst. The de-

274

E. E. DONATH

scribed prehydrogenation catalysts reduce oxygen, nitrogen, and sulfur compounds to a great extent without splitting of C-C-bonds and have varying hydrogenating activity, as shown in the example of the Alz03-WSzcatalyst with nickel or cobalt as activator. At pressures above 200 atm., aromatics are hydrogenated. For some reactions, for instance, 0

A

0 100

-

Co - Ni COMBINATION COMBINATION

----- F c - Ni

20 80

40

60

60

40

80 20

100%, N; 0 % Co or Fe

PIG.15. Iron and cobalt 8s a substitute €or nickel in the AlzOa-~~~2-NiS(70-27-3) prehydrogenation catalysts. 0 Fe-Ni comCo-Ni combination. bination.

the refining of gasoline or coke-oven light oil, removal of sulfur compounds and gum-forming olefins is essential, but hydrogenation of aromatics must be avoided. This is achieved by the use of pressures below 200 atm. The cobalt-molybdenum combination was also investigated later by the Union Oil Company for the desulfurization of petroleum fractions (29). It was found that, for a bentonite-supported catalyst, cobalt oxide

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

275

and molybdic oxide together were more active than either of the oxides, and an even more active catalyst was cobalt molybdate on bentonite. For different feedstock temperatures between 385" and 450"C., pressures between 7 and 17 atm., feed rates from 1 to 2 vol./(vol.)(hr.), and hydrogen rates between 500 and 850 vol./vol. of feed were used. Under these conditions the sulfur, olefinic, diolefinic, and phenolic compounds were largely hydrogenated, but the aromatic hydrocarbons remained substantially unaffected. The cobalt molybdate catalyst is also suitable for the reduction of organic nitrogen bases (30). From Colorado-shale-oil fractions containing 2% Nz and 0.7% sulfur, jet and Diesel fuels have been obtained with 0.01 to 0.1% Nz and 0.03 to 0.04% sulfur. The hydrogenation of coke-oven light oil was originally developed by the BASF in 1931 and is now used on a commercial scale (31). It produces benzene, toluene, etc., with a very low sulfur content of below 0.005%. The product is free of gum formers. Hydrogen at 60 atm. pressure is employed in one plant, and in the other coke-oven gas at 35 atm. pressure is used. An operating temperature of 350°C. is given for an alumina-supported catalyst. A German patent application (32) describes a catalyst containing 10% molybdic acid on alumina precipitated from aluminum nitrate at 95°C. and a pH of 7. The composition of the hydrogenated light oil is described in detail by Grothe (33).The analysis indicates that the olefins in the crude light oil were hydrogenated t o paraffins and naphthenes, but aromatic rings such as benzene and toluene were virtually not attacked, as indicated by the cyclohexane and methylcyclohexane content of 0.11 and 0.1496, respectively. Similar alumina-supported catalysts (34) are used a t pressures up to 5 atm. for the refining of gasolines for the reduction of nonhydrocarbons and the removal of gum-forming substances. The olefins are not attacked. Catalysts for the refining of gasoline (36),which improve the octane number by splitting of higher boiling fractions without hydrogenation of aromatics, consist of activated alumina and an iron-group metal, which are treated with hydrofluoric acid. For such catalysts a hydrogen pressure of 50 atm. a t a temperature of 400°C. is proposed. It is probable that catalysts of this type have also been used for the desulfurization of Diesel fuels and crude oils (36).Operating pressures of 50 atm. with 0.2 cbm Hz/kg. of feed and throughput of 2 kg./(liter catalyst) (hr.) are reported with catalyst-regeneration cycles of up to 150 hr. Table XI11 shows the reported results; temperature and yield were not given. Comparison of these results with those reported for prehydrogenation shows that the selectivity of these catalysts for sulfur removal has been

276

E. E. DONATH

increased. The hydrogenating activity has been decreased by catalyst selection and use of lower hydrogen pressure. The commercial use of nonsplitting hydrogenation processes for the desulfurization of middle and light distillates in the petroleum industry in the United States is rapidly increasing. I n 1954 forty petroleum hydrogenation units (37) were in operation, under construction, or planned. Comparatively low pressure will be used in these plants in contrast with the high-pressure process used in the thirties (2.3) for the hydrogenation TABLE XI11 Desulfurization of Diesel Fuel and Crude Oil Aramco crude oil

Diesel fuel % Desulfurization Specific gravity Aniline point, "C. Initial B.P. 200°C. 225°C. 300°C. 350°C. % Sulfur

Feed

70

94

Feed

65

75

0.833 69.3 202 11 69 98 0.75

0.827 70.5 199

0.823 70.8 187

0.849 72.5

0.841 74

0.835 73

-

18.5

18.9

19.5

12 71 98 0.22

15 72 98 0.04

62 1.4

64 0.48

67.5 0.35

-

of lubricating oils, kerosenes, etc. One impetus for the present development appears to be the availability of by-product hydrogen from hydroforming, platforming, and similar plants in petroleum refineries.

V. SPLITTING CATALYSTS With a catalyst containing 10% WS2 on activated fuller's earth (terrana) treated with hydrofluoric acid (38), the gasolines obtained had higher octane numbers than those made when the WS2 catalyst was employed. A comparison of this catalyst with WSz is given in Table XIV. The experiments were made with a petroleum distillate boiling between 209' and 327"C., of 0.85 sp. gr., a t 200 atm. pressure, 407'C., and hourly throughput of 2 kg. of feed/liter of catalyst. The data show th a t conversion and gasoline yield are alike. The gasoline obtained with the fuller's earth catalyst has a higher octane number. Only about one half of this octane-number increase can be ascribed t o the somewhat higher content of aromatics. This indicates a higher content of more knock-resistant branched-chain paraffins. The influence of feed fractions of various boiling ranges on the overall conversion was investigated with a bituminous-coal middle oil. The

277

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

oil was prehydrogenated with the A1203-W-Ni catalyst and the middle oil obtained was distilled into six equal fractions. The properties of the fractions and the results of splitting hydrogenation with the fuller's earth-WSz catalyst at 250 atm. pressure and an hourly throughput of TABLE XIV Comparison of Catalysts

WSa

Catalyst

HF-Terrana

Conversion 65 % Hydrocarbon gas as % of converted feed 8 Gasoline 0.716 Specific gravity 40 %, 100°C. E.P., "C. 188 Octane number 57 78 % Paraffins 20 % Naphthenes 1.5 % Aromatics 0.5 % Olefins

+ 10% WSZ

65%' 8 0.722 41 202 64 72 19.5

8 0.5

TABLE XV Influence of Feed Fractions on Conversion Feed fraction Specific gravity 50% Point, "C. Aniline point Nitrogen bases (mg. NH3/1.)* Splitting hydrogenation Temperature, "C. Gasoline Space time yield, kg./(l.)(hr.) Specific gravity 7% 100°C. Paraffins Naphthenes Aromatics Octane number (M.M.)

1

2

3

4

5

6

0.825 0.841 0.856 0.875 0.887 0.894 263 290 164 192 214 243 43.5 45.5 48.5 51.5 54.0 58.5 3.4 4.8 4.8 1.0 1.0 1.7 374

374

374

374

374

391

0.88 0.64 0.82 0.81 0.76 0.89 0.735 0.735 0.725 0.730 0.736 0.730 48 45 38 62 58 51 42 42.3 43 39 48 44 53 48.5 49 56 59 54 5 3 2 1 3 2 71 75 76 75 74 71.5

* The content of nitrogen bases was determined by anhydrous titration with perchloric acid in glacial acetic acid and expressed as mg. NH8/liter oil (39). 1.5 kg. of feed/liter of catalyst are shown in Table XV. The gasolines were cut to 150°C. end point. The lowest and highest fractions show deviations from the four middle fractions. The high conversion and low gasoline volatility obtained with the first-fraction are probably con-

278

E. E. DONATR

nected. This fraction has a high content of one-ring naphthenes. A single opening of a naphthenic ring of this low-boiling fraction would decrease its boiling point sufficiently to convert it into gasoline without formation of fractions boiling below 100°C. Fractions 2-5 produce gasolines of similar nature; the gasoline volatility (%, 100°C.) decreases, as expected, with increasing molecular weight. The highest fraction requires somewhat higher temperatures for a splitting rate comparable with th a t obtained with the other fractions. The content of ta r acids and bases of this fraction was not higher than in the lower fractions and cannot be the cause. The need for a greater number of splitting steps to convert this fraction, which consists of molecules with a n average of seventeen carbon atoms, into gasoline or the presence of hydrocarbons of a condensed nature might be the cause of the refractory behavior. The fuller’s-earth catalyst thus produces more branched-chain paraffins and does not saturate aromatics t o the same extent as does WSz. It even permits splitting of part of the aromatics of the feed without saturation. Further development of catalysts which permit preservation of a higher percentage of the aromatics in the feed and production of more branched-chain paraffins appeared possible t o obtain gasolines of still higher octane number. The activity of the WSz catalyst is decreased by nitrogen bases. The sensitivity of the fuller’s earth-supported catalyst, however, is much greater. Figure 16 shows the change in conversion of a n acid-washed, nitrogen-free gas oil with the addition of nitrogen bases. Addition of the equivalent of 50 mg. NHa/liter decreases conversion b y about 50%. Figure 16 indicates that increasing amounts of bases become proportionally less effective; however, the catalyst is practically inactive a t 400°C. when bituminous-coal middle oil containing tar bases equivalent to 5 g. NHS/liter is used as feed. However, fuller’s earth-supported catalysts show remarkable splitting activity with nitrogen-containing feedstocks at temperatures higher than 450°C., as will be described later. The catalyst is not permanently damaged by ta r bases. Figure 17 shows the decrease in conversion observed with petroleum gas oil when 0.1 ?6 of aniline is added. Operating conditions similar t o those shown in Table XV were used. After aniline addition, conversion drops to one half of the original value. When the aniline addition is discontinued, the conversion returns t o its original value rapidly. The great differences in the influence of aniline on the splitting activity of several catalysts are shown in Table XVI. The high sensitivity of the fuller’s earth catalysts is evident. The choice of aniline as nitrogen base is arbitrary. Comparative tests were therefore made with an addition of 0.015% nitrogen in the form of

COAL-HYDROGENATION VAPOR-PHASE

CATALYSTS

279

FIG.16. Influence of increasing amounts of nitrogen bases on the activity of the fuller's earth-WSz catalyst. 'L ANILINE IN FEED

OPERATING DAYS

Fro. 17. Nitrogen bases and activity of fuller's earth-WSz catalyst.

different nitrogen compounds. Tests similar to those shown in Fig. 17 indicated that saturated cyclic amines, such as piperidine or hexahydroaniline, reduced conversion by about 20%. With pyrrole, dibutylamine, and aromatic amines (pyridine, quinoline, aniline, and naphthylamine), the reduction was 40 to 70%.

280

E. E. DONATH

The properties of gasolines (40) from various liquid-phase-hydrogenation middle oils obtained by prehydrogenation and splitting hydrogenation with the fuller's earth-WSz catalyst are shown in Table XVII for gasolines of about 190°C. end point and with 40% boiling up t o 100°C. TABLE XVI Influence of Aniline on Catalyst-Splitting Activity Conversion in % of conversion without aniline Nitrogen added as aniline, % Mo-Zn-Mg WSa WS, on HF-Terrana at 200 atm. at 600 atm. HF-Terrana, 600 atm.

0.01

0.017

0.04

65

44 58 34

35 25

0.22

0.5

The data show that the aromaticity of the feedstock determines the proportion of cyclic to acyclic compounds in the gasoline. Analyses of fractions of a similar bituminous-coal gasoline of 150°C. end point by Raman spectrography by J. Goubeau gave the composition shown in TABLE XVII Properties of Gaaolines

Origin Gasoline, sp. gr. Aniline point, "C. % Paraffins % Naphthenes % Aromatics % Olefins Octane number (R.M.)

Cracked petroleum residue 0.736 53 58 37 4 1 70

Brown coal

Bituminous coal

Coke-oven tar residue

0.736 51 53

0.740 40 30 59 10 1 73

0.75 34 20 64 15 1 75

40

6 1 66

Table XVIII. The results appear to be in agreement with the reaction mechanisms given in Sec. III,3 for the WSz catalyst. The higher octane number of the gasoline obtained with the fuller's earth catalyst indicates that a greater percentage of highly branched paraffins is present. The high percentage of normal butane in the butane fraction is due to the removal of most of the butanes from the gasoline

28 1

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

TABLE XVIII Raman Analysis of Gasoline from Bituminous Coal Paraffins Normal Branched Naphthenes Cyclopentanes Cyclopentane Methylcyclopentane Dimethylcyclopentane Methylethylcyclopentane Higher

39.3 12.5 26.8 49 3 20.3 5.3 6.0 6.I 0.8 2.1 20.3

Cyclohexanes Cyclohexane Methylcyclohexane Dimethylcyclohexano Higher

29.0 6.2 4.3 3.5 15.0 __

29.0

Aromatics Benzene Toluene Xylenes Higher

Branching

ca C4 C6 C6

c7 Ce

11.4 2.5 3.9 3.6 1.4

__

11.4 Detailed Analysis of Parafins None Single Double Triple

-

Total

0.7 4.9 6.5 5.7 1.0

2.0

__

-

1.2 2.8 2.0

0.1 3.4 7.6 11.8 10.7 5.7

12.5

18.8

6.0

2.0

39.3

0.1 2.7 2.7 4.1 2.2 0.7

-

-

by fractionation. Thereby the higher boiling normal butane remains preferentially in the gasoline, but the isobutane is stripped off. Catalysts that produce gasolines of higher octane number than the fuller’s earth-WSs catalyst might be obtained by using less of the strongly hydrogenating WSz component and adding other activators for the splitting reaction. It was attempted to develop such catalysts that have high-splitting activity at temperatures around 400°C. Only a t such low temperatures

282

E. E. DONATH

can high gasoline yields be obtained, as the conversion to hydrocarbon gases increases with the reaction temperature. The properties of the HF-treated fuller's earth without activators will be discussed first. This material a t 200 to 300 atm. pressure showed small splitting activity and high sensitivity to nitrogen bases. At 600 atm. pressure, considerably increased splitting activity was observed. It seems that the activation of the HF-treated fuller's earth obtained by the WS2, which converts it into a dual-function catalyst (41) of increased activity, can also be achieved-at least to some extent-by higher hydrogen pressure. The results obtained with the petroleum distillate used for the experiments in Table XV are shown in Table XIX. The splitting activity TABLE XIX Fuller's Earth Splitting Catalyst

Catalyst Pressure, atm. Temperature, "C. Gasoline Space time yield Specific gravity %, 100°C. End point "C. % Paraffins % Naphthenes % Aromatics % Olefins Octane number

HF-Terrana 10% ws2

HF-Terrana

Activated carbon

200 407

600 425

600 475

1.05 0.722 41 202 72 19.5 8 0.5 64

0.6 0.728 39 180 66 21 12.5 1.5 72

0.24 0.738 28 178 65.5 21 13 0.5 51

+

of the H F Terrana a t 600 atm. is smaller than that of the WS2-containing catalyst, and a somewhat higher temperature had t o be used. However, the splitting activity is remarkable in comparison with that of activated carbon, which requires considerably higher temperature t o begin. The active-carbon-catalyst gasoline is of much lower octane number, indicating different splitting mechanism. The HF-Terrana catalyst produces gasoline of higher aromatics content and octane number than the catalyst containing WS2. The increase in aromatics content only partly explains the increased knock resistance, as is the case in the comparison of the concentrated WSZ and the WSz-Terrana catalyst. Various compounds were found to increase the splitting activity of the fuller's earth. A catalyst of activity simiIar t o th a t of the catalyst containing WSZ was Terrana containing 20% FeF3 (42). This catalyst gave the same conversion as the WSz-Terrana catalyst at a slightly lower

283

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

temperature and produced gasolines of about two points higher octane number. The composition of the gasolines produced by these two catalysts in respect t o hydrocarbon classes is very similar; one difference is the higher olefin content, as much as 2% obtained with the F e F 3 catalyst. Similar catalysts were developed under the direction of G. Kaftal (9) by the A.N.I.C. in Italy. Results obtained at the I.C.I. plant at Billingham with a fuller's earth-iron catalyst were given by G. Gordon (8). Table XX shows results obtained with a prehydrogenated creosote oil, which is a distillate fraction from the highly aromatic coke-oven tar. The experiments were made a t 250 atm. pressure, and results obtained in single-stage operation with WSz catalyst are also given. I n this case, replacTABLE XX Creosote Oil Hydrogenation ~~

Single-stage Process catalyst

Gasoline % Aromatics % Naphtlienes % Branched-chain paraffins % Straight-chain paraffins Octane number (CFR M.M.)

ws2

3 50 31 18 68

~~~

Prehydrogenation with WSZand splitting with fuller's earth

7 54 33 6 75

15 48 32 5 77

ing the WS2 in the fuller's earth-supported catalyst by iron increased the aromaties content a t the cost of naphthenes only with little change of the ratio of branched- to straight-chain paraffins. Remarkable increases in the aromatics content of gasolines and the octane number were obtained with synthetic-aluminum-silicate-supported catalysts ( 4 3 ) , as shown in Table XXI. The experiments were made a t 250 atm. pressure with petroleum gas oil of 0.846 sp. gr. and a boiling range of 187" t o 330°C. Temperatures around 400°C. were used and were adjusted to obtain approximately the same conversion for both catalysts. The silicate catalyst is less active. The high aromatics content obtained with the aluminum silicate catalyst is especially evident in the higher fractions. The amount of branched-chain paraffin hydrocarbons is not greatly different, as shown by the octane number of the aromatic free fractions, used as a n index. These tests indicate that, with the aluminum-silicatesupported catalyst, aromatics can be split with less ring hydrogenation than observed

284

E. E. DONATH

with fuller's earth catalysts. This was demonstrated by the splitting of aromatics like naphthalene or of hydroforming residues with an aluminum-silicate-molybdenum catalyst a t about 400°C. With catalysts containing 0.5 % molybdenum, gasolines with about 50% aromatics were obtained. With higher molybdenum contents, gasoline aromaticity decreased, but the conversion rate increased. Similar catalysts were used on a semicommercial scale (44). TABLE XXI Fuller's Earth and Alumina Silicate Catalysts Catalyst Temperature, "C. Gasoline, sp. gr. %, 100°C. % Aromatics % Naphthenes % Paraffins Octane number (M.M.) Fraction 75"-100°C. % Aromatics Octane number (M.M.) Octane number (M.M.) 0 % aromatics* Fraction 14Oo-16O0C. % Aromatics Octane number (M.M.) Octane number (M.M.) 0% Aromatics*

HF-Terrana 10% WS2

Alumina-silicate 20% ZnS, 1% WS2

382 0.718 49 10 27 63 71.4

416 0.726 53 18 27 55 75.8

+

72 70.8

12 73.6 72.0

16 58.4 49

38 69.4 46.5

8

* Aromatics were separated by low-temperature extraction with a liquid mixture of SO2 and CaHa. Specially prepared synthetic aluminum-silicate catalysts containing heavy metals are being used a t present for the splitting hydrogenation of petroleum oils in German hydrogenation plants. It is stated th a t these silicate catalysts can be used for the splitting hydrogenation of low-boiling gas oils as well as for higher boiling oils with a boiling range above 325°C. Precipitation of the silicate a t a low pH or coprecipitation with heavy metals has been described in patents (45). A comparison (46) of gasolines produced from petroleum gas oil with the fuller's earth-WSz and the silicate catalyst is shown in Table XX II. With the new catalyst, gasolines of higher aromatics and isoparaffin content are obtained. The yield is somewhat lower. The additional aromatics content appears mostly in the higher boiling part of the gasoline. This is shown in Fig. 18, and the aniline points of 20°C. gasoline fractions are used to indicate aromaticity. The operating temperature of the splitting catalysts described is about 400"C., but the first commercial catalyst (Mo-Zn-Mg) required tempera-

285

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

BOILING POINT,

C.

FIG.18. Aniline point of gasoline fractions. Splitting hydrogenation of gas oil. TABLE XXII Fuller's Earth and Silicate Catalysts Catalyst Gasoline Specific gravity % Boiling to 100°C. End point, "C. $& Aromatics Octane number Motor method clear Motor method 1.5 ml. TEL/gal. Research method clear Research method 1.5 ml. TEL/gal.

+

+

HF-Terrana

+ WSS

0.715 47 190 8 63 76 64 77

Silicate

0.723 45 190 17.5 72 83 74 85

286

E. E. DONATH

tures of 450°C. and more for splitting hydrogenation, and a t temperatures of about 500°C. gasolines of high aromatics content were obtained. Because of the high knock resistance of aromatics, improved high-temperature splitting (aromatizing) catalysts were developed. Most of the tests with such catalysts were made with middle oils obtained from liquid-phase hydrogenation of bituminous coal. Such oils are highly aromatic and therefore are suitable for the conversion into highly aromatic gasolines by splitting hydrogenation at pressures between 200 and 600 atm. A number of supports and activating compounds were investigated and some rules were found. Basic supports like alumina or activated carbon require activators of only weak hydrogenating activity, such as vanadium or chromium oxides, t o produce stable, tar-acid-free gasolines. The life of such catalysts was satisfactory a t 250 atm. operating pressure. The disadvantage of catalysts of this type is th a t the gasolines are somewhat deficient in low-boiling components and that the octane number of the nonaromatic portion of the gasoline is low. Acidic supports, such as fuller's earth or synthetic aluminum silicate, require strongly hydrogenating activators like molybdenum oxide. I n this case satisfactory tar-acid reduction and stable gasolines are obtained from bituminous-coal middle oil. Although these catalysts normally begin to show splitting activity a t 400"C., it is necessary t o use temperatures in excess of 450°C. with bituminous-coal middle oil of 0.5% nitrogen content. At a pressure of 250 atm., satisfactory catalyst life was found with several catalysts producing high-aromatic-content gasolines. However, at 600 atm. long catalyst life was achieved with a wider range of catalysts. The properties of basic supports can be varied b y addition of acidic components like silica or silicates or of fluorides to alumina. Inversely, acidic supports can be modified by addition of magnesia, zinc oxide, etc. The high operating temperature causes high formation of hydrocarbon gases, and the gas loss increases with increasing aromaticity of the gasoline. With the HF-Terrana WSz catalyst a t 400°C. hydrocarbon-gas yield for motor-gasoline production is only 8% of the converted feed. I n Fig. 19 the percentage of hydrocarbon gas formation is plotted for different catalysts and the production of gasolines of different aromatics content. I n Fig. 19 only the catalyst supports are named. The different aromatic content was obtained either by different additives or different amounts of these or by changes in operating conditions, temperature, and pressure. The individual experimental points are not shown and, although the scatter is wide, the positions given for the different supports appear a t least qualitatively correct. The composition of the nonaromatic portion of the aromatization

287

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

gasolines varies also with the catalyst support. As an over-all index of its quality, the octane number of the gasoline after extraction of the aromatics with liquid sulfur dioxide and propane can be used.

I

CATALYST SUPPORT:-

10 20 30 % HYDROCARBON GAS ON CONVERTED FEED

FIG.19. Aromatication of bituminous coal middle oil. TABLE XXIII Catalyet Support and Octane Number

Octane number of aromatic-free gasoline Boiling to 100°C. Catalyst support Activated carbon Alumina Synthetic aluminum silicate} Fuller’s earth

30 %

40 %

50%

61

64

65 -

68

70

70

72

Table XXIII shows that increasing octane numbers were obtained in the following order: activated carbon, alumina and synthetic aluminum silicate, fuller’s earth. Branching of paraffin hydrocarbons will be parallel to changes in octane number. It should be remarked that these experiments were made in 1941 and that newer aluminum silicates might produce more branched paraffin hydrocarbons.

288

E. E. DONATH

One example will be given to show the influence of molybdic acid addition to fuller's earth, which was modified with minor amounts of MgO, ZnO, and Cr2O3.The experiments were made a t 600 atm. pressure with bituminous-coal middle oil as feed. Figure 20 shows the splitting activ-

TEMPERATURE F.

FIG.20. Splitting rate with fuller's earth catalyst containing different amounts of molybdic acid. (Bituminous coal middle oil hydrogenated a t 600 atm. pressure.)

ity of three catalysts containing 0, 0.5, and 5 % molybdic acid. The splitting rate after addition of only 0.5% molybdic acid to the support is doubled. With increased molybdenum addition the aromatics content of the gasoline and the recycle middle oil decreases as shown in Table XXIV. TABLE XXIV Splitting Activity and Molybdenum Content

"C.,

Gasoline

% MOO*in catalyst Temperature, "C. % Aromatics 0

0.5 5.0

500 493 459

57 44 37

Aniline-point % 100°C. recycle mid. oil 20 26 29

-33 -13

+7

The increased splitting rate with the molybdenum-containing catalysts is expressed also in the higher content of low-boiling fractions in the gasoline. The experiments were made a t different temperatures in order t o compare the catalysts a t similar conversions. However, comparison at the same temperature would show the same sequence in aromatics content.

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

289

A sidelight to the reaction mechanism of splitting catalysts a t different reaction temperatures is furnished by the composition of the gaseous hydrocarbons obtained as by-products. The butane content of the total hydrocarbon off-gas and the isobutane content of the butane fraction are

FIG.21. Splitting hydrogenation with fuller’s earth catalysts. Butane content in CI-C, fraction and isobutane content of butane fraction.

especially significant ( 4 7 ) . With increasing operating temperature, the butane content of the off-gas, as well as the isobutane content in the total butane, decreases. I n Fig. 21, curves are shown for fuller’s earthsupported catalysts. The activator used on fuller’s earth does not influence this relationship greatly; neither are the operating pressure and the nature of the middle oil used as feed of great importance. The d a ta shown

290

E. E. DONATH

were obtained in l-liter converters a t 250 t o 600 atm., with petroleum and prehydrogenated coal and lignite middle oils a t temperatures up t o 425”C., and at higher temperatures with liquid-phase-coal middle oil. Indicated are points obtained in a 1000-atm. experiment with fuller’s earth, without activator, as the catalyst, with decalin at 250 atm.; one point was obtained with an active carbon-Cr-V catalyst. The butane content decreases with increasing temperature and the isobutane content shows a similar trend. The isobutane content of the butane fractions is higher than the equilibrium figure at temperatures below 5OO0C., an indication that isomerization precedes splitting on fuller’s earth catalysts in a manner similar t o that described for the tungsten sulfide catalyst. However, the isomerizing action of the fuller’s earth catalysts is stronger. The isobutane content at 4OO0C., according to Fig. 21, is 65%, and an isobutane content of 50% was observed with tungsten sulfide. The active-carbon-supported catalyst shows much lower butane content in the off-gas and the isobutane content is lower than the equilibrium curve. This indicates a different splitting mechanism. Similar behavior of these catalyst types has been shown for the higher boiling fractions of the gasoline. The aromatics content of gasolines from bituminous coal can be varied within wide limits with corresponding changes in gasoline yield and quality. Feedstocks other than liquid-phase’bituminous-coal middle oil werealso tested. The aromatics content depended upon the aromaticity of the feedstock. An active-carbon-supported catalyst containing C rz 0 3and V z 0 6 produced gasoline with 50% aromatics from bituminous-coal middle oil. From coke-oven tar middle oil, with the same catalyst, gasolines containing 55 t o 60% aromatics were obtained. Lignite middle oils and cracked petroleum oils or extracts produced gasolines with a n aromatics content of 20 t o 40%. For the commercial production of gasolines with high aromatics content, two catalysts were used. For operation a t 300 atm. pressure at the Scholven Plant, an activated-carbon-supported catalyst containing 15% C r z 0 3and 5% Vz06 was used. Bituminous-coal-hydrogenation middle oil was converted into gasoline containing 50% aromatics. At 700 atm. pressure, a fuller’s earth-supported catalyst developed by the Ruhrol GmbH was used. At the Welheim Plant middle oil obtained from liquid-phase hydrogenation of coke-oven tar pitch was converted into gasoline with an aromatics content of 45 %. The catalyst contained, in addition to fuller’s earth, hydrofluoric acid, Crz03,ZnO, sulfur and 0.5% molybdenum. The same catalyst was used a t the United States Bureau of Mines Coal t o Oil Demonstration Plant, Louisiana, Missouri, for the process-

291

COAL-HYDROGENATION VAPOR-PHASE CATALYSTS

ing of various coal and lignite middle oils (48). The catalyst was reproduced and tested in experimental equipment at the Bruceton Station of the United States Bureau of Mines (49).

ACKNOWLEDGMENTS This article is affectionately dedicated to Professor Matthias Pier. Much of the work covered herein was conducted under the inspiring direction of Dr. Pier in the highpressure department of the Badische Anilin- und Soda-Fabrik at Ludwigshafen/Rhein, Germany. Review of the manuscript and many helpful suggestions by L. L. Hirst, Maria Horing, G. V. McGurl, and H. H. Storch are gratefully acknowledged.

REFERENCES 1. Berthelot, P. E. M., Bull.

chim. [2] 11, 278 (1869); Ann. chim. phys. 141 20, 526 (1870). Translation see Abhandl. Kohle 1, 156 (1915). 2. Bergius, F., and Bilwiller, J., German Patent 301231 (November 26, 1919). 9. Bergius, F., German Patent 303272 (August 22, 1922). 4. Horing, M., in “Die Chemie der Braunkohle” (Lissner-Thau, ed.), Vol. 11. W. Knapp, Halle, 1953. 6. I. G. Farbenindustrie, French Patent 612,504 (March 14, 1925). 6. Storch, H. H., in “Hydrogenation of Coal and Tar, Chemistry of Coal Utilization” (Lowry, ed.), Vol. 11, p. 1750. Wiley, NewYork, 1945. 7. Sherwood, P. W., Petroleum Refiner 28, No. 12, 97 (1949); 29, No. I, 119 (1949); 29, No. 2, 106 (1949); 29, No. 3, 150 (1949); 29, No. 5, 123 (1949); 29, No. 6, 113 (1949); 29, No. 7, 124 (1949); 29, No. 8, 107 (1949). 8. Gordon, K., J . Inst. Fuel 20, 42 (1946). 9. Technical Oil Mission, Reel 173, Frames 418, 529; Reel 254, Frames 418-423. 10. Technical Oil Mission, Reel 253, Frames 958.984. 11. Oettinger, W., Erdol u. Kohle 6, 693 (1953). 12. Pier, M., 2. Elektrochem. 63, 291 (1949); see also Technical Oil Mission, Reel 199. IS. Technical Oil Mission, Reel 205, Frame 330. 1.6. Schuster, C., Technical Oil Mission, Reel 205, Frames 825-859. 16. Hiickel, E., “Adsorption and Kapillarkondensation.” Leipzig, 1928. 16. Schuster, C., 2.Elektrochem. 38, 614 (1932). 17. Technical Oil Mission, Reel 205, Frames 736-754 or Reel 201, Frames 292-310. 18. Cawley, G.M., Research (London)1, 553-561 (1948). 19. Technical Oil Mission, Reel 205, Frames 770-774. 20. Technical Oil Mission, Reel 205, Frames 792-799. 21. Reitz, O., Chem. Ing. Tech. No. 21-22, 413-417 (1949). 22. Technical Oil Mission, Reel 254, Frame 48. 23. Haslam, R. T., Russell, R. P., and Asbury, W. C., Proc. World Petroleum Congr. (1933) 2, 302 (1933); Petroleum Times 30, 297 (1933). 24. Technical Oil Mission, Reel 166, Frames 579-617. 26. Jones, W. A., Oil and Gas J . 63, No. 26, 81 (1954). 26. Guenther, G., Freiberger Forschungs. Zssue A17, 38 (1953). 27. German Patent Application 12g, 4/01, p. 43150, assigned t o Badische Anilin und Soda-Fabrik (May 17, 1949). 28. Technical Oil Mission, Reel 202, Frames 442-459. SOC.

’

292

E. E. DONATH

29. Byrns, A. C., et al., Ind. Eng. Chem. 36, 1160 (1943); Berg, C., et al., Chem. Eng. Progr. 43, No. 1, T I , TI34 (1947). 30. Berg, C., Petroleum Processing 7, 186 (1952). 31. Urban, W., Erdol u. Kohle 4, 279 (1951); Novak, H., and Liebich, M. G., Brennstof-Chemie 36, 308 (1954). 32. German Patent 869,198, to BASF (May 19, 1950). 38. Grothe, W., Erdol u. Kohle 6, 450 (1953). 34. German Patents 802,398, 825,868, and 843,279, to Badische Anilin und SodaFabrik (1951 and 1952). 36. German Patents 886,897, 898,443, and 915,330, to Badische Anilin und SodaFabrik (1953 and 1954). 36. Pier, M., Erdol u. Kohle 6, 690 (1953); Oettinger, W., ibid. 6, 693. 37. Anonymous, Oil and Gas J. 63, 62 (1954). 38. U . S. Patents 2,154,527 (1939) and 2,194,186 (1940). 39. Wittmann, G., 2.Angew. Chem. 00, 330 (1948). 40. Bureau of Mines Translation T 388 (1947). 41. Heinemann, H., Mills, G. A., Shalit, H., and Briggs, W. S., Brennstof-Chemie 36, 368 (1954). 42. Technical Oil Mission, Reel 201, Frames 368-381. 43. Technical Oil Mission, Reel 181, Frames 6453-6459. 44. Technical Oil Mission, Reels 138 and 197. 46. German Patents 821,684 (1951), 854,346 (1952), 869,200 (1953), and 902,845 (1954) to BASF; British Patent 504,614 (1939). 46. Oettinger, W., Erdol u. Kohle 6, 693 (1953). 47. Pier, M., Brennstof-Chemie 32, 129 (1951). 48. Hirst, L. L., Clarke, E. A., and Chaffee, C. C., U . S. Bur. Mines Rept. Ineest. No. 4676 (1950); No. 4944 (1953); No. 6043 (1954). 49. Wolfson, M. L., Pelipetz, M. G., Dornick, A. D., and Clark, E. L., Znd. Eng. Chem. 43, 536 (1951); Pelipetz, M. G., Frank, L. V., Ginsberg, H. H., Wolfsrn, M. L., and Clark, E. L., Chem. Eng. Progr. 60, 626 (1954).

The Kinetics of the Cracking of Cumene by Silica-Alumina Catalysts CHARLES D. PRATER AND RUDOLPH M. LAG0 Research and Development Laboratory, Socony Mobil Oil Company, Ine., Paulsboro, New Jersey Page I. Introduction.. ..... 11. The Results of ...................................... 294 1. Kinetics. . . . . . . . ........................................ 294 ts of Previous Studies.. . . . . . . . . . . . . . . . . . . .295 2. Uncertainties in a. The Use of a Method Insensitive to the Kinetics.. . . . . . . . . . . . . . . . . . 295 b. The Diffusion-Transport Effects.. ............................... 301 c. The Presence of Inhibitors in Cumene.. .......................... 304 305 3. Discussion ....................................................... 111. The Kinetics of Cumene Cracking as Determined by Differential-Reactor Studies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ........................ 305 1. Experimental Procedure. ..... ........................ 305 2. Kinetic Scheme.. ........... ............................. 309 3. Studies with Pure Cumene.. ....................................... 310 a. Dependence of the Rate of Cracking on Pressure and the Determina310 tion of the Values of Constants.. ................................ b. The Temperature Dependence of the Rate of Cracking.. . . . . . . . . . . . 314 c. The Number of Active Sites, Bo................................. 315 316 4. Effect of Inhibitors.. ............................................. a. The Pressure Dependence of the Rate of Cracking., . . . . . . . . . . . . . . . 316 b. The Dependence of the Rate on the Amount of Cumene Hydroperoxide and Other Inhibitors.. . . . c. Competition of Benze

..........

. . . . . . . . . . . 319

5. Adsorption and Desorption of Chemisorbed Cumene Hydroperoxide.. .. a. Reversibility of the Chemisorption of Cumene Hydroperoxide.. . . . . . b. Rate Constants of Desorption and Adsorption of Cumene Hydro........................ peroxide. .......................... 6. Determination of Adsorption Constants for Products from Studies of Diffusion Transport Effects. ....................................... IV. Determination of Adsorption Constants with an Integral Reactor. . . . . . . . . 1. Integration of the Kinetics.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Calculation of Adsorption Constants.. .... 3. Influence of the Back Reaction on the V with an Integral Reactor.. .. ....... ......... V. Coke Formation. ................................................... 293

320 320 321 322 324 324 328 329

294

CHARLES D. PRATER AND RUDOLPH M. LAGO

Page 329 a. Types of Coke.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 329 b. Kinetics of Coke Formation.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 330 (1) Coke Formation as a Function of Mole Fraction of Inhibitor. . . 330 (2) Coke Formation as a Function of Time., . . . . . . . . . . . . . . . . . . . . . 331 . . . . . . . . . . . . . . 331 (3) Coke Formation as a Function of Pressure (4) Empirical Kinetic Equation for Coke For . . . . . . . . . . . . . . 331 (5) Relationship between Coke Formation and Cracking Sit 2. Coke Formation from Gas Oil.. .............................. a. Coke as a Function of Mole Fraction of Inhibitor b. Coke as a Function of Time., ........................... c. Regeneration of Catalytic Activity by Desorption of Inhibitors.. . . . . 334 VI. Discussion and Summary.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 337 .................................................... 338 1. Coke Formation from Cumene Hydroperoxide

....................

INTRODUCTION A knowledge of the kinetics of the reaction at the active sites is of primary importance in determining the nature of catalytic action in heterogeneous catalysis. Information about the nature of the catalystsubstrate* interaction can be obtained from the way in which the “rate constants” in the kinetics change on variation of such parameters as temperature, catalyst treatment, and catalyst composition. I n addition, these constants are the quantities which should correlate with structural information such as that obtained by the methods of solid state physics. However, the true kinetics at the active sites is not always obtained unless certain precautions are taken, as has been pointed out in a recent volume of Advances in Catalysis ( I ) . I n this paper we shall review the information found in the literature concerning the kinetics of cracking of cumene on silica-alumina. We shall also present some new experimental studies which take into account the precautions outlined in the preceding reference. I.

11. THE RESULTSOF PREVIOUS STUDIES I . Kinetics The cracking of cumene has received considerable attention in recent years as a reaction typical of one class of cracking reactions, namely dealkylation of aromatics. Among the studies of cumene cracking found in the literature there are several attempts to determine the kinetics of

* The term substrate designates the substance acted upon by the catalyst. This is the definition used by the enzyme chemist. We use it rather than the term adsorbate to focus attention on the interaction of the reactant molecules with the chemisorbtion sites which are active for the reaction studied in contrast to the interactions which may occur with any other adsorption sites of the catalyst.

295

KINETICS OF THE CRACKING OF CUMENE

the reaction. The reason for this attention to the kinetics is th a t this reaction is particularly simple, yielding essentially only benzene and propylene. This results in a simplification of kinetic analysis. Typical analyses of the gaseous products of this reaction are shown in Table I. TABLE I Composition of the Gas Formed f r o m Cumene Cracking by Silica-Alumina Catalyst

Reaction temperature, "C. Conversion, vol. % Gas Hz CH4 Composition Volume % ccs z c4

I

+ +

Plank and Nace (8)

Plank and Nace (8)

Differential reactor experiment, Fig. 4

427 30-50 1 95-97 4-2

482 50-65 1-2 93-96 5-3

420 <1 0.8 96.4 2.8

The results from the published studies of the kinetics are summarized in Table 11. All except one were made by means of a flow-type integral reactor. An integral reactor is one which yields data already containing an integration of the kinetics over a changing concentration of reactant in the catalyst bed, th at is, a reactor operated so that the conversion of reactant t o products is relatively high. The integral-reactor studies were, in each case, one of three types, conversion being measured either (1) as a function of residence time of reactant in reactor (space velocity), (2) as a function of pressure, or (3) as a function of concentration of inhibitors of the cracking reaction added to the cumene charged to the reactor. 2. Uncertainties in the Results of Previous Studies

On examinination of the results of the integral-reactor studies, one or more sources of uncertainty were found in each case, which makes the kinetic conclusion drawn from them doubtful. The three major sources of uncertainty are (1) the use of a method which is insensitive to the precise functional forms of the kinetics, (2) the presence of diffusiontransport effects which modify the kinetics, and (3) the presence in the cumene used of strong inhibitors of the cracking reaction. a. The Use of a Method Insensitive to the Kinetics. The method which is insensitive t o the functional form of the kinetics is the study, at constant pressure, of the variation of conversion as a function of residence time of the charge in the reactor (space velocity = amount of reactant charged per unit time per unit amount of catalyst). I n this method the assumed kinetics is used to calculate a functional relationship between conversion and velocity of flow of reactants (or residence time of reactant in catalyst bed). Agreement between the functional forms of the calcu-

TABLE I1 Summary of Cumene-Cracking Kinetics Reported i n the Literature Investigators

Proposed kinetics Ref.

Kinetics

Type of reactor

Integrated form

Integral

Obolentsev 9,4 and Gryaeev

Ballod et al.

6,6

Topchieva 7 and Ponchenkov

_ -- B~ + B F

~ Z+ I n

_ -- cl[z+ I n F

(I

- z)*

(I - z)l*

Integral

Integral

Integral

Type of study

Tempera- Catalyst ture composirange, "C. tion

Conversion 300-450 Silica alumina as a function of space velocity a t constant pressure Conversion 400-500 Silica alumina, as a func10-30 % tion of alumina space veby weight locity at constant pressure Silica Conversion 400 as a funcalumina ranging in tion of space vecomposition from locity a t constant pure silica to pure pressure alumina Conversion 455-488 Silica alumina, as a function of 10% alumina pressure by wt.

TABLE 11. (Continued) Summary of Cumene-Cracking Kinetics Reported in the Literature

Investigators Weiss and Prater

Proposed kinetics Ref.

1

%Plankand 2 Nace

*

Kinetics P ,*

dn

z-

'lBo

dn -=

P.

+ G K I P I+ G K 2 , + G

kP,* K P , +--'PI K,

Integrated form

W = Elz + EzIn (1 - z)*

Type of reactor Differential

F

Rate equation not integrated

Integral

Type of study dn

zas a

Tempera- Catalyst ture composirange, 'C. tion 360-420

function of temp., pressure, and concentration of inhibitors 426-482 Effect of nitrogen compounds on conversion

Silica alumina, 10 % alumina by wt.

Silica alumina, 10% alumina by wt.

See pages 298. 310, and 325 of text for meaning of symbols. In addition, z is mole fraction of reactant converted; ze is mole fraction of reactant converted a t equilibrium; A's, B's, C's, D's, E's, and b are constants.

298

CHARLES D. PRATER AND RUDOLPH M. LAG0

lated and observed relationship of conversion to space velocity is considered to mean that the assumed kinetics is the actual kinetics. However, this procedure is not very sensitive t o the form of the kinetics assumed except at high conversions (observed conversion/conversion at thermodynamic equilibrium greater than one half). At high conversions other difficulties arise, such as, for example, the necessity of including the back reaction in the kinetics. I n addition, the probability increases that complicating side reactions will become significant. Furthermore, the possibility increases that adsorption competition between the reac'tion products and the reactant for catalyst sites will make important contributions t o the kinetics. Those complicating factors make the interpretation of data, except in special cases, difficult and uncertain. To show the insensitivity of this procedure to kinetics, a cylindrical reactor filled with catalyst will be considered. We have for each element of cross section an at

Fdx=-dW

(1)

where F is the rate of charge (moles/sec.) W is the weight of catalyst dn - is the rate of reaction [moles/(sec.)(wt. of catalyst)] at z is the mole fraction of reactant converted.

Rearranging and integrating, we obtain

The well-known values of this integral [see, for example, (9)] for several kinetics are given in Table 111. In all cases we have neglected the back reaction. In three cases it is assumed that there is no volume expansion during the reaction. In one case a volume expansion of 1 mole of reactant producing 2 moles of product is assumed. To test an assumed kinetics with experimental conversion data, we calculate, using experimental values of 2, values for the function defined by the integral [Eq. (2)) for the kinetics under test. These values are then plotted against the experimental values of W / F . If the points lie on a straight line, the assumed kinetics is considered to be the true kinetics. For example, to test a first-order kinetics with no volume expansion, we would plot W /F against In (1 - 2). Let us calculate conversion data which would be observed for zeroand second-order kinetics by using the equations in Table I11 to see how

KINETICS O F T H E CRACKING O F CUMENE

299

well we can distinguish these from first-order kinetics, To do this we plot the calculated zero- and second-order values of W/F for each conversion against In (1 - x). These plots are shown in Fig. 1. Below 50% conversion these data lie on an approximate straight line and would be conTABLE I11 Differential and Integral Forms of Various Kinetics

dn

Order 0

Conditions

a = f(x)

No back reaction

36 No back reaction, no volume 1

expansion Same as above

2

Same as above

1

No back reaction, 2/1 volume expansion

sidered t o he first order if we allowed for the existence of the usual and reasonable experimental errors in observed values. A comparison of half order with first order with a 2/1 volume expansion is also shown in Fig. 1 and leads t o essentially the same conclusion. Thus, with the usual experimental procedures, a comparison of conversion data as a function of space velocity (residence time) does not tell us whether the kinetics differ from first order in the region of conversion below 50% for the range of order tested (zeroth t o second). On the contrary, first-order kinetics can be used to represent the conversion as a function of residence time for a wide range of situations. Some investigators have been aware of this approximate first-order behavior of integral reactors, as shown by the statement th at “even complex catalytic systems approximate a pseudo- first order relationship when only space velocity is varied . . . ” (10). A variation on the aforementioned test procedure used bjr several investigators is to plot one part of the function against another. For example, consider the first-order kinetic equation with volume expansion, as used by Topchieva and Panchenkov (7) :

Fx is plotted against F In (1 - z). If a straight line is obtained, the

300

CHARLES D. PRATER AND RUDOLPH M. LAG0

kinetics under test is considered to be the true kinetics. T h a t this procedure is also insensitive to kinetics is shown by calculating a set of values for F and 2 from the half-order kinetic equation given in Table 111 and plotting i t according t o the foregoing procedure. The results are I0 20

% CONVERSION 30 40

I1

3llL 0.4 02

ZERO ORDER v 5. FIRST ORDER

c a p

0.2

0.4

0.6 -In [I-X]

0.8

1.0

1.2

96 CONVERSION

SECOND ORDER

vs.

FIRST ORDER

0.2 0.0

02

0.4 0.6 -In [I-X]

0.8

1.0

1.2

i r/i % CONVERSION

n

-F, X

1.2

HALFSROER

0.6

FIRC,TOHRDER

0.2 0.0

2/1 VOLUME EXPANSION 02

0.4 0.6 0.8 - ~ - 2In [I-X]

1.0

1.2

FIQ.1. Conversion data from zero-, half-, and second-order kinetics tested against first-order kinetics.

shown in Fig. 2. As can be seen, with a reasonable experimental error in the observed values allowed, the half-order kinetic values behave as if they were the kinetics of Eq. (3). I n kinetic studies of heterogeneous catalysis, many schemes involving adsorption of reactants, reaction on the surface, and desorption of products have to be considered which differ less from first-order kinetics than the extremes given in Table 111. However, it is just these schemes which we need t o distinguish one from the other.

KINETICS O F THE CRACKING O F CUMENE

301

An examination of the results of Table I11 suggests a method for determining the dependence of the kinetics on the partial pressure of the reactant. We see that the exponent of co exhibits the order of the reaction. Thus, if the value of co is varied at constant conversion by changing the pressure in the reactor, we can determine the dependence of the kinetics on the partial pressure of the reactant. This is the procedure used by Corrigan et al. (8). % CONVERSION

F

Ln ( I - x )

FIG. 2. Conversion data from half-order kinetics tested against the kinetics of Topchieva and Panchenkov (7).

b. T h e Difusion-Transport Efects. When the rate of reaction becomes too large, the insufficient rate of diffusion transport within the porous catalyst particle will modify the concentration of reactants and products seen by the active sites. Under these conditions incorrect reaction kinetics, activation energies, and rate constants will be obtained (1,12). TABLE IV Rates of Reaction Above Which Diffusion Effects Infiuence the Kinetics for Various Particle Radii

Particle radius, cm.

Rate, , moles (CU. cm.1(set.) from Eq. (4)

5.6 x 10-3 2.9 x 6.3 X 1.75 X 10-l

6.65 x 10-4 3.05 x 6.3 X 8.4 x 10-7

Temperature, "C. 690 380

345 307

How easily experimental investigations of cumene cracking by porous catalysts can suffer from diffusion effects can be demonstrated by use of the data in Table IV. These data were calculated from the criterion (1)

302

CHARLES D. PRATER AND RUDOLPH M. LAO0

that diffusion effects are negligible if dn 1R 2 51 - . dt c Dc

(4)

dn where - is the rate of cracking per unit volume of catalyst [moles/(cm.a) dt (sec.11 . ._ c is the concentration of reactant (m~les/cm.~) R is the radius of the catalyst particles (cm.) Dc is the effective diffusivity of cumene in the porous catalyst particle at the reaction temperature (cm.2/sec.). Equation (4) gives the values of the various parameters for which diffusion-transport phenomena begin to have a significant effect on the kinetics. This condition has been shown to be, for all practical purposes, independent of kinetics (11). The value of DCused to calculate the data in Table I V was obtained from direct measurement of the diffusion of hydrogen in the catalyst by the porous plug method (1, p. 189, method b). The value used was for the spherical beads of cogelled silica-alumina cracking catalyst used in the experiments to be reported here. The catalyst contained 10% A1203 by weight and had a surface area of 350 m.2/g. The value of the effective diffusivity of the catalyst particle for Hz at 27°C. (DHJ was found t o be 7 x 10-3 cm.2/sec. The value of the effective diffusivity of the catalyst particle for cumene Dc, a t reaction temperature was calculated from this measured hydrogen cliff usivity by the equation

where ME, = molecular weight of hydrogen Mc = molecular weight of cumene T = reaction temperature (OK.) T o = temperature of measurement of DH2 (= 300'K.). Equation ( 5 ) holds for the catalyst used, as diffusion in the porous catalyst particle is of the Knudsen type (pore radius << mean free path of the gas molecules). For this catalyst the average pore radius, 7, calculated from the equation, for cylindrical pores

is 30 A. I n this equation V , is the pore volume per gram of catalyst and S

is the surface area per gram of catalyst. Assuming that cumene vapor at cracking temperature behaves as a

KINETICS O F THE CRACKINQ O F CUMENE

303

perfect gas and using values of known parameters a t 1 atm., we obtain by combining Eqs. (4)and ( 5 )

dn = 1.82 x

1

-

1 0 - 9 ~ &

(7)

dt

Column I1 of Table IV gives, for each particle size, the rate of cracking of cumene above which diffusion phenomena will influence the observed kinetics for the silica-alumina catalyst described. Column I11 of Table IV gives the temperature at which these rates are observed for very pure cumene. This temperature was determined from the experiExtent of Digusion Limitation

TABLE V as a Function of Particle Radii for the Experiments of Corrigan et al. (8)

1)

Radius of catalyst particle, cm.

0.0225 0.165 0.215 0.265

W / F for 28% conversion 1.3 5.7 7.6 10.6

1)

0.72 0.16 0.12 0,091

mentally observed d n / d t vs. 1/T data of Fig. 8. Since this cracking reaction is usually studied in the temperature range of 300”-500”C. with catalysts similar to those mentioned above (Table 11), the possibility of diffusion effects being present is clearly shown. Corrigan et al. present sufficient data to enable us to show directly that their data were strongly affected by diffusion-transport phenomena. They made a study of the variation of W/F at constant conversion (28%) as a function of particle radius. The data, which are reproduced in Table V, can be analyzed for each particle radius by the “triangle” method (1) to obtain the ratio of the observed reaction rate to the reaction rate when no diffusion effects are present (9).This analysis is possible for these data as the ratio of values of W/F at constant conversion should be proportional to the ratio of the rates. The values of r ] obtained are given in Table V. Since the catalyst particles used in Corrigan’s studies had a radius of 0.215 cm., the existence of large diffusion effects is clear. As will be shown below, although Corrigan et al. chose the correct kinetic scheme, the experimental results cannot be considered to support this scheme because of the presence of such large diffusion effects. How their results can be interpreted as a combination of diffusion and inhibitor effects has already been discussed elsewhere (1).

304

CHARLES D. PRATER AND RUDOLPH M. LAG0

Attention should be called to the danger of using activation-energy arguments to determine whether a given study is free of diffusion effects. It is often considered that an observed high activation energy (say, -10 kcal./mole) is a guarantee for the absence of diffusion effects; i.e., the presence of strong diffusion effects is often presumed to lead to an apparent activation energy in the range of 1 to 2 kcal./mole characteristic of the temperature dependence of the diffusivity constant Dc. That this is not necessarily the case has been pointed out in previous volumes of the Advances in Catalysis (1,1,2).The most that diffusion effects within porous particles can do for catalytic reactions is to reduce the activation energy to one half the true activation energy. c. The Presence of Inhibitors in Cumene. Another source of uncertainty in the studies of cumene-cracking kinetics as reported in the literature TABLE VI Eflect of Distillation and Chromatographing Procedures on Rate oj Cracking of Cumene by a Silica-Alumina Catalyst

Cumene sample Eastman cumene P-1481 P-1481 After atmospheric distillation, 30-plate column P-1481 After chromatographing through silica-gel-clay column

Rate of cracking, moles/(m.t)(sec.) x 10% a t 420°C. and 1 atm. of cumene 40 130

260

arises from the presence of inhibitors in the cumene charge. All samples of the cumene of commerce tested by us contained considerable amounts of cumene hydroperoxide. As will be shown later, and as already discussed by Weisz and Prater (l), this substance is a strong inhibitor of the cracking reaction. The kinetics observed in the presence of the inhibitor can be entirely different from that of pure cumene. Conventional distillation is not a good method for removal of the cumene hydroperoxide, as can be seen by the results given in Table VI. The results show that a single distillation, even with a 30-plate column, does not remove all the inhibition. This is probably due to the decomposition of cumene hydroperoxide a t the boiling point of cumene to give other inhibitors with a lower boiling point, which then pass through the column. We find that the inhibitors will be removed, however, if a vacuum distillation is made a t room temperature. It has been pointed out that polar inhibitors can be removed from

KINETICS O F THE CRACKING OF CUMENE

305

cumene by chromatographing through silica gel (I,%'). This seems to be the most satisfactory method of inhibitor removal. 3. Discussion The original integrated forms of the kinetics as reported in the various studies have been rearranged in Table 11 t o show their similarities. All Bs C In (1 - s). The differexcept one are of the form W/F = A ences among the kinetics, such as order of reaction, adsorption properties for reactant and products, inhibitor effects, are contained in the constants A , B, and C. For the particular forms of these constants the original references should be consulted. The exception to the preceding form is the kinetics of Obolentsev and Gryazev (3,4), which is of the form W/Fb = A B In (ze - 2). Because of the three sources of uncertainty outlined in the preceding section, the experimental results for the integral reactor studies cannot be considered to yield conclusive data concerning the respective kinetic schemes. The kinetic conclusions of Obolentsev and Gryazev (3,4), of Biallod et al. (5,6), and of Topchieva and Panchenkov (?') are in doubt because of the method of study used, i.e., variation of space velocity at constant pressure. In addition, as will be shown below, the probability is high that either the rates of reaction observed in their studies were diffusion limited or the cumene used by them contained cumene hydroperoxide. I n the study of Corrigan et al. (8) the presence of large diffusion effects has already been demonstrated. Any study of cumene cracking made above 400°C. with commercial cracking catalyst of conventional activity levels, with particle size of approximately 1 mm. or greater (as is the usual practice in integral reactor studies) , and with pure cumene will almost certainly be diffusion limited. If this is not the case, it is by virtue of the fact that the cumene contains inhibitors which reduce the reaction rate and, of course, alter the kinetics of the reaction. This can be seen by considering the diffusion results in Table I V and the discussion of effect of inhibitors. In order to resolve the question as to the precise kinetics of this reaction, differential-reactor experiments were made under conditions that eliminated the three sources of uncertainty given above.

+ +

+

AS DETERMINED BY 111. THE KINETICSOF CUMENECRACKING DIFFERENTIAL-REACTOR STUDIES

1 . Experimental Procedure The advantages of the differential reactor in kinetic studies in heterogeneous catalysts have been discussed in a previous volume of Advances

306

CHARLES D. PRATER AND RUDOLPH M. LAG0

in Catalysis ( 1 ) . A differential reactor is one which yields directly the reaction rate a t a known concentration of reactant, that is, a reactor in which the conversion of reactants to products is small ( < 1%). The apparatus is a modification of the Schwab type of reactor described by Weisz and Prater (1) and is illustrated in Fig. 3. In view of the

FIG.3. Differential reactor for kinetic study. The parts of the system in contact with the reactant are of Pyrex. 1. Filling reservoir 13. Cumene condenser 2. Drainage stopcock 14. Stirrer 15. Magnetic valve 3. BoiIer tube sled with crapillaries 4. Boiler furnace 16. Benzene condenser 5. Preheater furnace 17. Drainage stopcock 6. Preheater coils 18, Stirrer 7. Loading arm 19. Condenser 8. Loading-arm furnace 20. Known volume V O 9. Reactor furnace 21. Vent stopcock 10. Reactor 22. Water manometer 23. Manometer vent 11. Catalyst tray 12. Thermocouple wells

simplicity of the cumene-cracking reaction as illustrated by the data of Table I, the rate of gas production can be considered to be essentially identical with the rate of cumene cracking. The rate of production of

KINETICS OF THE CRACKING OF CUMENE

307

gas is measured by noting the pressure rise on a water manometer in a given interval of time, say 30 sec., when the system is closed to the atmosphere by stopcock (part 21, Fig. 3). The calibration of the systein necessary to convert the pressure reading to moles of propylene produced is accomplished by means of duplicate readings made when a known volume V O(part 20, Fig. 3) is added to the system. An application of the gas law gives

dn at

where - is the rate of gas production (moles/sec.)

?!!At

is the rate of pressure rise a t the time t

V Ois the known volume To is the temperature of the known volume R is the gas constant AP8 is the pressure rise in the time interval At when the known volume is attached to the system A P is the pressure rise in the time interval At without the known

volume attached to the system

A is the cross-sectional area of the manometer The last term contributes less than 1 % to d n / d t and can be neglected. Since propylene is extremely soluble in cumene and benzene, some provision must be made for removing it quantitatively from the liquid condensate. This is accomplished by the two condensers (13 and 16 in Fig. 3). Condenser (13) is operated at 148°C. and serves to condense the cumene (b.p. 151OC.). The temperature of the gases is dropped from 148" to 75°C. in the short section of air condenser between condensers (13) and (16). Condenser (16) is operated a t 75°C. and serves t o remove the benzene (b.p. 80") which is not condensed by condenser (13). Condenser (19) serves to reduce the temperature of the gases to room temperature. Most of the condensation occurs in condensers (13) and (16). The catalyst used in these studies was a commercial type of bead silica-alumina cracking catalyst (Socony Mobil Oil Company, Inc.) containing 10% alumina by weight and having originally a bead diameter between 2 and 5 mm. This catalyst had a surface area of 350 m.*/g. and a CAT-A activity of 42 A.I. (IS). The catalyst was crushed to -100 200 mesh particle size (0.14 - 0.07 mm. diam.) for use in the apparatus. It was spread in a thin layer on a 2- by 5-cm. glass tray (Fig. 3, part 11). For conditions of operation which lead to high catalytic activity, 20 mg. of catalyst was used and for

+

308

CHARLES D. PRATER AND RUDOLPH M. L A G 0

low-activity conditions as much as 100 mg. The liquid circulation rate was -60 cc./hr. Because of the small particle size used and the small depth of catalyst bed, diffusion effects are negligible for temperatures up to 650°C. and P, = 1 atm., as seen by the application of Eq. (4) and the observed rates (dnldt). (See Table IV.) I

0

lo00

I

2000

3000

4000

T I M E , SEC.

FIG.4. Time course of the rate of cracking of cumene by silica-alumina catalyst at 420°C. and atmospheric pressure.

The apparatus is opened for catalyst loading by breaking the glass loading arm and then resealing. This is done because of the difficulty of finding a convenient cumene-insoluble, high-temperature, reversible cement which could be used to seal a ground-glass joint. The cumene used was Eastman pure grade or Phillips pure grade, which was further purified to remove polar inhibitors, mostly cumene hydroperoxide. This was accomplished a t room temperature by passing the cumene at a rate of -5 cc./min. through a 5- by 120-cm. cylindrical column containing alternate layers of silica gel and freshly burned clay. The alternate layers were about 5 cm. deep. The time course of the observed rate, dn/dt, for a typical run is shown in Fig. 4.Experiments with longer runs showed that after about 3000 sec. dn/dt became essentially constant. Consequently, the standard length of run was chosen t o be between 3,000 and 4,000 see. The transient in the first part of the run is caused in part, if not entirely, by the necessity of establishing thermal equilibrium in all parts of the apparatus. Because of the small rate of gas production, even small

KINETICS OF THE CRACKING OF CUMENE

309

temperature fluctuations or slowly moving condensing fronts will cause changes in the apparent rate. 6. Kinetic Scheme We have found the data obtained from extensive diff erential-reactor studies to be consistent with a kinetic reaction scheme which involves the following: 1. The adsorption of cumene and inhibitors on active cracking sites follows a Langmuir type of isotherm. This means that there is little or no interaction among the chemisorbed molecules on the surface. This might be expected to be the case as studies of the chemisorption of the inhibitor quinoline by similar catalyst (14) show that the surface is sparsely covered with active sites ( <5% of internal surface area covered with chemisorbed quinoline a t 315°C.). In addition, the active sites are homogeneous with respect to adsorption energies. 2. The rate of cracking is proportional to the number of cumene molecules chemisorbed on active cracking sites.

Since the investigation to be reported was made under differentialreactor conditions, the conversion of reactants to products was less than 1 mole % in all cases. As will be shown below, the chemisorption of the products on active sites is much weaker than the chemisorption of cumene. This means that the products do not compete with cumene very strongly for active sites and that their presence can be neglected when their concentration is small. Thus, the back reaction of products t o reactants can be neglected. Expressed in a more formal manner, the foregoing considerations lead to the following scheme: ki

k3

S + A eSA -+ products kz

310

CHARLES D. PRATER AND RUDOLPH M. LAG0

where S represents the substrate, cumene, in the gas phase A represents the free active-catalyst sites SA represents cumene chemisorbed on active sites Ii represents the ithinhibitor present in the gas phase IiA represents the i t h inhibitor chemisorbed on active sites. Inhibitors have been included in the scheme since studies will be made of their effect. This scheme leads to the set of equations: ki(S X A) h,(Ii X A)

. .

Bo = (SA)

= =

+

(kz kd(SA) k8,(11A)

. .

+ A + 2 (LA) i

These equations can be solved for the rate of cracking d n / d t and yield

d n -_ dt p8

+

kaBoP, G KriPIc-k G

2

(11)

i

where B Ois the total number of active cracking sites P, is the partial pressure of cumene Pri is the partial pressure of the ithinhibitor Kri is the adsorption equilibrium constant for the ithinhibitor and is equal to 1c7;/kgi Gis a constant related to the chemisorption of cumene and is equal to (kz k3)/k1.

+

3. Studies with Pure Cumene

With no inhibitors present, Eq. (11) reduces to

This scheme will be shown to be consistent with experimental data on the variation of the rate of cracking with pressure and temperature. a. Dependence of the Rate of Cracking o n Pressure and the Determination of the Values of Constants. An examination of Eq. (12) shows that, when P, is large compared with G, d n / d t becomes independent of pressure and equal to ksB0. On the other hand, when P, is small compared with

311

KINETICS OF THE CRACKING OF C U M E N E

G, d n l d t becomes essentially first order in P, with a rate constant equal to k3BO/Q. In the range where P, is the same order of magnitude as G , no simple-order relationship holds. The value of the constants k3Bo and G cannot both be accurately determined in arbitrary pressure ranges. When P, is large compared with G, d n / d t measures directly the value of k3B0. In the range where P,is of the same order of magnitude as G , studies of d n l d t as a function of pressure yield values of both kaBoand G . However, when P, is small compared with G , only the ratio k3BolG can be determined with any degree of accuracy.

300 -

200

-

,

0

,

l

l

l

,

l

I

.8 PRESSURE, ATMOSPHERES .4

C

,

l

1.2

FIQ.5. The rate of cracking of pure cumene as a function of absolute pressure for various temperatures.

The rate of cracking as a function of pressure was studied in the pressure range of 0.4 to 1 atm. in order to test the applicability of the kinetic scheme and to determine the values of the constants k3Bo and G . Operation of the differential reactor at the required pressure was achieved by operating the gas outlet (21, Fig. 3) and the manometer outlet (23, Fig. 3) attached t o a large reservoir kept at the required pressure. This made it possible to use the manometer system as before. The condensers were operated a t the appropriate temperatures relative to the boiling points of cumene and benzene at the pressure used. It was found that for the pressure range of 0.4 to 1 atm. and a temperature range of 430" to 380°C': the condition that P, is the same order of magnitude as G holds sufficiently well for a determination of botth k3Bo and G to be made as shown by the data in Fig. 5. Below 350°C.

312

CHARLES D. PRATER AND RUDOLPH M. LAG0

G becomes negligible compared with P, and k3B0 is obtained directly. This is illustrated by the small pressure dependence which is observed at 360°C. (Fig. 5). Above 430'C. G will become too large compared with P, for accurate determinations of k3Bo and G at atmospheric pressure. Going to higher pressures and temperatures causes the rate of cracking to increase so rapidly that diffusion limitations are encountered in the region of good sensitivity to both k3Boand G.

,o-4,

,

5:O

TEMPERATURE OC 410 3qO

3y

4y

m"

3

10-7

10'8.

LO-^^ 1.1

I

1.2

0

,

,

I

L3

1.4

1.5

10'

f

1.6

1.7

3

T'K

FIG.6. Plot of In kaB, vs. 1/T.

The values of kaBo and G at 420" and 390°C. were calculated for each pair of experimental points in Fig. 5. The average values of k3Boand G for these two temperatures were used to calculate the solid curves in the figure from Eq. (12). The agreement is good between the calculated curves and the experimental points. Measurements of d n f d t at 1-atm. partial pressure of cumene were made for the temperatures 300" and 330°C. The value of dn/dt under these conditions is equal to k&. The values of kaBo a t 420" and 390°C.

KINETICS OF T H E CRACKING OF C U M E N E

313

obtained from the pressure studies and the values obtained at 300" and 330°C. are given in the In k3Bo vs. 1 f T plot (Fig. 6). A linear relationship is obtained as required if ksBo is to obey an Arrhenius relationship

k3B0 = Ae-Q/RT

If the linear relationship between In k3Bo and 1/T is assumed to hold up to 500"C., we can use measurements of dn/dt at I-atm. pressure to 100

W W

0:5,

TEMPERATURE *C 4:O 3:O

4y

3?0

10

a I

a

-m 0 I I4

I1

u 0.I

.o I

FIG.7. Plot of In G vs. 1/T.

determine the value of G . Using extrapolated values of koBo and the measured rates of cracking at 1-atm. partial pressure of cumene, we obtain values of G at 500", 480°,and 460°C. The values of G obtained a t 420" and 390°C. along with the foregoing values are given in the In G vs. 1 / T plot (Fig. 7). Again a linear relationship is obtained. Thus, if kSBo obeys an Arrhenius relationship in the temperature range of 420" to 500"C., G also obeys an Arrhenius relationship in this range. A linear relationship will be obtained for G = (k2 k 3 ) / k l only if either (1)

+

314

CHARLES D. PRATER AND RUDOLPH M. LAGO

kz < k3, ( 2 ) k ) < kz, or ( 3 ) the activation energy of kz is equal t o the activation energy of k8. Condition (2) seems most likely. This condition makes the cracking of the chemisorbed cumene, characterized by the rate constant k3, the rate-determining step. The data plotted in Figs. 6 and 7 yield

32.6 X 10%

G

=

3.1 X 1 0 l 0 e - 7

for the temperature dependence of krBoand G . 500

TEMPERATURE ' C 450 390 330 I

4AO

I

1.1

1.2

I

1.3

I

0

420

360

300

I

I

I

I

1.4

1.5

1.6

1.7

10'

3

T OK

FIQ.8. The rate of cracking, dnldt, of cumene by silica-alumina catalyst as a function of I I T . The solid line is the theoretical curve dn/dt = kJio(l/l G ) . The value of ksBo and G were obtained from Figs. 6 and 7.

+

b. The Temperature Dependence of the Rate of Cracking. The temperature dependence of dn/dt for 1-atm. partial pressure of cumene is shown by the In (dn/dt)p,,, vs. 1 / T plot in Fig. 8. The theoretical curve (solid line) was calculated by use of Eqs. (13) and (14) in Eq. (12). The experi-

KINETICS O F T H E CRACKING O F CUMENE

315

mental values are shown as circles. This plot illustrates the large changes in apparent activation energies of the rate of reaction, d n l d t , which can occur even in a relatively short temperature range in heterogeneous catalytic investigations. The activation energy of d n / d t changes on passing from 380" to 480°C. from 40 kcal./mole characteristic of the activation energy of IcsBo to 7.4 kcal./mole characteristic of the diff ereace in the activation energies of ksBoand G. c. The Number of Active Sites, Bo. An interesting use can be made of the frequency factor of Eq. (13). When P,>>G, the pure cumene rate becomes dn S keBa (15) The reaction is then essentially zero order. Absolute-reaction-rate theory (16)gives the rate of reaction ( d n / d t ) of zero-order reactions on heterogeneous catalysts to be*

where k is Boltzmann's constant and h is Planck's constant. Comparing Eqs. (13), (15), and (16), we see that

The number of active sites BO can be calculated, as the values of T (42OoC.), k, and h are known. Bo = 0.87 X 10x7 sites/m2. We may compare this finding with the value for the number of chemisorption sites per square meter for quinoline obtained from experiments by Mills, Boedeker, and Oblad (14) for catalysts similar to the one used by us but containing 12.5% alumina. Quinoline is a strong inhibitor for the cumene-cracking reaction; i.e., it is strongly chemisorbed on the active cracking sites. Furthermore, the cracking activity of the catalyst for cumene approaches zero asymptotically as the amount of quinoline is increased. Therefore, the number of quinoline-adsorption sites will be equal to or greater than the number of cumene-cracking sites. The data for the number of quinoline-adsorption sites are summarized in Table VII and yield a value of 1.27 X 10'7 sites/mZ. The agreement between this and the value for Bo, above, is remarkable in view of the completely different methods used to obtain the values.

* The

usual assumption ie made that

f* - equals unity, where f*

and fa are the fa partition functions for the activated complex and adsorbed molecule respectively.

3 16

CHARLES D. PRATER AND RUDOLPH M. LAG0

TABLE VII Adsorption of Quinoline by Silica-Alumina Catalysts from the Studies of Boedeker, and Oblad ( 1 4 ) ~

~~

~

Mills,

~

~~

Catalyst No.

Ability to chemisorb quinoline at 315"C., ml./g.

sq. m./g.

IA I1 A IV A VA

0.06 0.044 0.021 0.009

273 193 111 44

Area,

Chemisorption sites/m.a 1.33 x 1.36 x 1.15 X 1.24 x

1017 1017 101' 1017

Average 1.27 X 1017

4. Efect of Inhibitors a. The Pressure Dependence of the Rate of Cracking. The way in which inhibitors influence the dependence of d n / d t at 420°C. on total pressure is illustrated by studies made on the effect of cumene hydroperoxide

-

4.6 X to-' MOLE FRACTION CUM E N E HYDROPEROXIDE -0

lo-' MOLE FRACTION CUMENE HYDROPEROXIDE

X <

0 1

0

-

0 I

.4 .8 PRESSURE, ATMOSPHERES

I

I . .

1.2

FIG.9. Rate of cracking of cumene containing cumene hydroperoxide at 420°C. as a function of absolute pressure. The solid lines are theoretical curves derived from

the kinetics scheme [Eq. (9)].

addition to cumene. The results are presented in Fig. 9. The presence of the inhibitor has an effect similar to lowering the reaction temperature, as seen by comparing Fig. 5. One experimentd value of dn/dt (Pa = I atm.; PI = 4.6 X atm.)

KINETICS O F THE CRACKING O F CUMENE

317

was used to determine the value of K I. Its value is 4.96 X lo3atm.-'. This value is used in Eq. (11) along with values of k3BO and G from the pure cumene studies to calculate the theoretical curves in Fig. 9. The agreement of the theoretical curves with the remaining three experimental points is good. 300

0

1.0 x lom3 2.0 x loJ 3.0 x 10-3 4.0 x I O - ~ MOLE FRACTION OF CUMENE HYDROPEROXIDE

FIG.10. Rate of cracking of cumene at atmospheric pressure as a function of mole fraction of cumene hydroperoxide present. The solid curve is a plot of Eq. (ll), where the values of k& and G were determined from pure-cumene pressure studies and the value of K I was determined from pressure studies with cumene containing cumene hydroperoxide.

Cumene hydroperoxide is probably decomposed under the conditions of the experiment. Consequently, the inhibition action observed is probably due to one or more decomposition products. b. T h e Dependence of the Rate on the A m o u n t of Cumene Hydroperoxide and Other Inhibitors. Experiments were made at 1-atm. total pressure and 420°C. to determine d n l d t as a function of the partial pressure of cumene hydroperoxide added to the cumene. The results are given by the circles in Fig. 10. The theoretical curve (solid line) was calculated from Eq. (11) with values of k3Bo and G from the pure cumene studies and the value of K I = 4.96 X lo3 atm.-1 obtained from the above-mentioned pressure studies. Good agreement between experimental data and the prediction from the kinetic scheme is again obtained. This type of experiment was also made at 1 atm. with acetophenone, xylene, and cyclohexane as additives. The results for the inhibitor acetophenone a t 420°C. are given as the circles in Fig. 11. I n this figure the ratio of the rate in the presence of the inhibitor to the rate for pure

318

CHARLES D. PRATER AND RUDOLPH M. L A G 0

cumene is shown. The relative rate is given by

R -

Z

-

+

P, G P , 4- GKIPI 4- G

where Rois the rate of cracking of pure cumene R is the rate of cracking of inhibited cumene. I.o

0.8

R -

0.6

-

R" 0.2 0-

0.0

I

I

t

I

,

MOLE FRACTION OF ACETOPHENONE

FIQ.11. Relative cracking activity as a function of the amount of acetophenone present. The value of K I for acetophenone was determined from the experimental data given in this plot. Solid line is theoretical curve calculated from Eq. (18).

A value of K I was calculated for the point at 2 X 10-3 mole fraction acetophenone given in Fig. 11 by use of Eq. (18) and the value of G from the pure cumene data (Fig. 7). The value obtained was 1.90 x 108 U

50

*

-

CYCLOHEXANE DILUTION

0

\-

20 -

l--O -'t 01.0

.9

.8

.7

.6

.5

MOLE FRACTION OF CUMENE

FIG. 12. The rate of cracking a t 360°C. of cumene diluted with cyclohexane or xylene.

atm.-'. The theoretical curve (solid line) was calculated by use of this value of KI. Again, reasonable agreement is obtained between the predictions of the kinetic scheme and the experimental data. The data obtained with added cyclohexane and xylene at 360°C. are given in Fig. 12. At 360°C. G is small compared to unity, and so d n l d t

KINETICS OF THE CRACKING O F C U M E N E

319

exhibits an approximate zero order above P , = 0.6 atm. Thus, when a substance is added to cumene which does not compete for chemisorption on active sites, the rate of cracking will be independent of the amount added down to 0.6-atm. partial pressure of cumene. This is the behavior exhibited by cyclohexane.

PRESSURE, ATMOSPHERES

FIG.13. Competition of benzene with cumene for chemisorption on active sites. The rate of cracking as a function of partial pressure of pure cumene is given by the solid curve. The rate of cracking of eumene diluted with 51 mole per cent benzene is given by the circle.

Since G describes the adsorption characteristics of cumene, the value of GKr gives the adsorption strength of an inhibitor relative t o cumene. The value of GKI for xylene a t 360°C. is unity, as seen from the data in Fig. 12. This means that xylene and cumene compete on a n equal basis for actual sites. c. Competition of Benzene with Cumene for Chemisorption o n Active Xites. The competition of the products for active sites is a necessary part of the investigation of the kinetics. Figure 13 shows the effects on the rate of reaction of diluting pure cumene with benzene to give a charge stock containing 49% cumene on a mole basis. The solid curve is a reproduction of the 420°C. curve from Fig. 5 and gives the effect expected from the reduction of the partial pressure of cumene alone. Benzene is seen to have an additional competition effect. The value of K I for benzene can be determined from Eq. (11) by using the value of G and k,Bo obtained from the pressure studies. This gives the value GKr = 0.4 E K I / K# where K , = k l / k 2 g l / G (since ka appears to be <
320

CHARLES D. PRATER AND RUDOLPH M. L A G 0

d. Adsorption Constants for Various Inhibitors. Maatman, Lago, and Prater (16) have measured the value of Kr for a variety of substances by using a differential reactor and the kinetics given by Eq. (11). Some typical values at 420°C. taken from their study are reproduced in Table VIII. TABLE VIII Adsorption Equilibrium Constants for Various Inhibitors on Silica-Alumina Cracking Catalysts

Compound Thiophene a-Methyl styrene Naphthalene Acetophenone Cumene hydroperoxide Pyridine Quinoline

Adsorption equilibrium constant a t 420°C. 9.15 X 101 1.39 x 102 1.03 X 102 1.90 x 10’ 4.97 x 10’ 8.02 x 104 3.42 X 106

The large value of K I for naphthalene shows that the presence of one hydrocarbon may greatly influence the rate of reaction of another hydrocarbon on the catalyst. Thus, the cracking behavior of a mixture of hydrocarbons need not be given by a simple sum of their pure cracking rates weighted according to their individual partial pressures. 5. Adsorption and Desorption of Chemisorbed Cumene Hydroperoxide

a. Reversibility of the Chemisorption of Cumene Hydroperoxide. One of the requirements of the kinetic scheme Eq. (9) is that adsorption of the inhibitor be reversible. I n the discussions in the preceding sections we have assumed this to be true. Plank and Nace (2) have shown that the inhibition action of quinoline is reversible. The inhibition action of cumene hydroperoxide or its decomposition products is reversible as TABLE IX Reversibility of Inhibition of the Catalyst by Cumene Hydroperoxide Reactant Pure cumene Cumene +O.OS% (mole) cumene hydroperoxide Pure cumene

Time, sec.

moles Rate

lo*(m.”(sec.)

Coke, 5% by wt.

0-4000

260


4000-8000 8000-12,000

73 252

1.1 1.2

KINETICS O F THE CRACKING O F CUMENE

321

shown by the data presented in Table IX. The rate of cracking for pure cumene was established by passing pure cumene over the catalyst in the differential reactor for 4000 sec. The pure cumene was then replaced by cumene containing 0.08 mole % cumene hydroperoxide. This cumene was passed over the catalyst for 4000 sec. to establish a steady state cracking rate. The rate dropped to 28% of the pure cumene value. The cumene containing the inhibitor was then replaced by pure cumene, which was passed over the catalyst for 4000 sec. to desorb the inhibitor on the catalyst. The rate of cracking at the end of 4000 sec. was 97% of the original pure cumene rate. b. Rate Constants of Desorption and Adsorption of Curnene Hydroperoxide. The transient in the rate of cracking which occurs when pure cumene is replaced by inhibited cumene or when inhibited cumene is replaced by pure cumene can be used to determine the values of the desorption and adsorption rate constants ks and k7. If the assumption is made that the rate of adsorption and desorption of cumene is much faster than the rate of desorption of inhibitor, the growth process can be considered a succession of steady states. The transient rate Rt for the growth of activity after the inhibitor has been removed is then given by

Rt = RF - ( R F- RI)e-k8t

(19)

where Rt is the rate at time t RF is the final rate (t = a ) Rr is the initial rate ( t = 0). This reduces to

Rt - RI RF - RI. Again, by means of the succession of steady states approximation, the decay of activity when pure cumene is replaced by cumene containing inhibitors is given by where ARt

=

ARo

=

Rt

=

RF

+ (RJ - RF)e-kg(RI/RF’f

(21)

This can be written

where ARt = Rt - RF ARo = Rr - R F . The values for In ( a R l / a R o )for the decay and In 11 - ARt/ARo)] for the growth of activity is plotted against time in Figs. 14 and 15, respectively,

322

CHARLES D. PRATER AND RUDOLPH M. LAGO

The data used were obtained from the reversibility experiment given above. The data give linear plots as required. The value of ks can now be calculated independently from either the growth or the decay curves. The results give ks = 1.05 X lomsl/sec. and 1.08 X I/sec., respectively. We can now also obtain kl by use of the value of

K I = 4.96 X 103atrn.-' = We obtain kl

=

5.30 atm.-' set.-' (at 420°C.). 1.0

0.1

I

-

-

.o I 0

400 TIME,

aoo

I200

SEC.

FIG.14. Decay of activity when pure cumene is replaced by cumene containing cumene hydroperoxide.

0.I

0

lo00 2000 TIME, SEC.

3000

Fro. 15. Growth of activity when cumene containing cumene hydropcroxide is replaced by pure cumene.

6 . Determination of Adsorption Constants for Products f r o m Studies of Diffusion Transport E$ects

There is an interesting method of obtaining the approximate value of GK, for the products from diffusion-transport effects. It is convenient to express the effect of diffusion by means of a diffusion factor q , which is defined as the ratio of the observed reaction rate t o the reaction rate when no diffusion effects are present. The diffusion factor q is a function of a parameter cp which contains all the independent variables of the system. For example, for a first-order reaction

KINETICS O F T H E CRACKING OF CUMENE

323

4 = r d v ewhere r is the particle radius, k, is the rate constant per unit volume, and D,is the effective diffusivity of the catalyst particle. Thiele (I?’), Wheeler ( l a ) ,and Weisz and Prater (1) have given q vs. 4 curves for some integral-order reaction kinetics. However, the kinetics of cumene cracking does not exhibit a simple constant “order.” Instead, the order is a function of partial pressure of reactants and products. To determine the q vs. 4 curve for this kinetics, the diffusion equation

must be solved with dn/dt given by Eq. (27). Equation (27), given in section IVl, includes the effects of the products on the reaction rate, 1.0

’I

0.1

-

0.I

1.0

10

0 FIG.16. 7 vs. +curves for the kinetics given by Eq. (27) for various values of GK, (no back-reaction term).

which must be considered in this discussion. In this equation C , is the concentration of cumene and D, is the effective diffusivity of cumene in the catalyst structure. It was found desirable to simplify this differential equation by using Eq. (27) without the back-reaction term (1/Ke)P,,P,.This is permissible in view of the rather smalI effect of the back reaction, which will be discussed later. Equation (23) has been solved numerically for spherical catalyst particles with an IBM electronic computer for values of GK, = 10, 1.0, 0.3, 0.10. In this case q5 = r 1/(k3BO)v/Dc at 1 atm. where ( ~ J B ois ) .the rate constant per unit volume. The curves obtained are plotted in Fig. 16. The various curves have different shapes. This becomes more apparent

324

CHARLES D, PRATER AND RUDOLPH M. LAG0

if all curves are replotted with a common point. This change of shape can be used to determine OK,. Experimental data on v as a function of particle radius for the catalyst used throughout this study were determined for the cracking of cumene at 420°C. We do not have t o know the value of (kaBo), and I), as their ratio can be determined by fitting one point to the curve and making use of the change of shape to choose the correct value of GK,. This procedure has been followed in the presentation of the data in Fig. 17. The 7 = 0.86 point is fitted to the GK, = 1.0 1.0

7

0.1

I

0.1

I

I .o

10 li

Fro. 17. Determination of GK, from diffusioneffects.

and the GK, = 0.1 curve. A better agreement for the remaining three points is obtained with the G K , = 0.1 curve than for the GK, = 1.0 curve. Thus, the products are adsorbed less strongly than cumene. This agrees with the previous results that GK, = 0.4 (benzene competition experiment given in section III4c). This example shows how quantitative analyses of diffusion-transport effects can yield kinetic information.

IV. DETERMINATION OF ADSORPTION CONSTANTS WITH INTEGRAL REACTOR

AN

1. Integration of the Kinetics

The adsorption equilibrium constant K I can be determined by integration of the complete kinetic scheme, including the back reaction, over the conversion in the catalyst bed. The complete scheme can then be written

KINETICS OF T H E CRACKING OF C U M E N E

325

where m and n are the products. The desorption of one of the products (n) is included in the step characterized by the rate constants k3 and k4, For the steady state, this scheme leads to the set of equations d-n -- ki(S X A) - kzSA dt dn - = k3SA - ka(mA X n) dt d-n -- k m A - ks(m X A) dt Bo = SA f IA mA A

+

(25)

-+

These equations can be solved for the rate of cracking d n / d f in terms of the partial pressure of reaction products, inhibitors, and the k’s, t o obtain

P,K,

-+ 1 + k6 KsPm -+ KIP,

1

K , is the equilibrium constant for the reaction cumene $ propylene

+ benzene

If the bond-breaking step is slow compared with the rate of desorption products, that is ka << ke and k4 << ks, Eq. (26) simplifies to

fEn dt - kaBQpa -+ GK,,,P,,,

+ GK,P, + G

This equation contains only constants already known or which can be easily determined. Substituting Eq. (27) for d n / d t in Eq. ( 1 ) and integrating, we obtain

W F

+ y - PSI In (1 + 2) + y + P7]In (1 - s) - 27s

2 - k3Bos3= [as2

[as2

(28)

326

CHARLES D. PRATER AND RUDOLPH M. LAG0

where a

1

=

+ GKIPI + G/P + GKIPI + GK,

/3 = 2G/P

r=G/P+GKm-I K, 1 62 = K, P = total pressure fraction of cumene conversion x = fraction of cumene converted a t thermodynamic equilibrium

+

~

When PI

=

0 a = p = 7 =

+

1 G/P 2G/P GK, G/P GK,,, - 1

+ +

Eliminating 2WIF k3Bob3by using in Eq. (28) the two conditions th a t

PI

> 0;2

PI

=

=

x2

and

0 ;x

= 21

and then solving for K I , we get

where

rC.

= X

1

+

21

In -- 0 In (1 - ~ 1-21

+

1

-~ p )x ~

+

G/P) p/2 6(2G/P GKm) 2(G/P G K , - 1) total pressure. This gives K I in terms of the conversion with pure cumene (xl) and inhibited cumene ( 2 2 ) when W/F is held constant. If the “pure” cumene used contains cumene hydroperoxide, the value of KI obtained from Eq. (29) must be corrected for the presence of this inhibitor. This correction is shown by Maatman, Lago, and Prater (16) to be given b y X = b2(1

e

= p = P =

+

+

$0

KI

-h-

= $1

($0

- *I

-:)1$ KI’

where $J~, $1, $I are the values of $ [see Eq. (29)]for conversion with pure cumene, impure cumene, and impure cumene containing added inhibitors, respectively, and u1 and UI are the values of 6 In (1 z/1 - 2) -

+

KINETICS O F T H E CRACKING O F CUMENE

327

In(1 - $2) for conversion with impure cumene and impure cumene containing added inhibitor, respectively. Since the value of knBo is known, the relationship =

$0

is used to determine

$0

2

W

63k3Bo

in Eq. (30).

2. Calculation of Adsorption Constants

The data of Plank and Nace on the inhibition of cumene cracking by nitrogen compounds can be used with Eq. (29) to determine the values of Kr for these substances. Since the “pure” cumene used by them con1.0

o.8I1 @/

0.8

-

-

,%

0.6

0.2 0.4 -

P

0.0 0.0

d

I -

,

I

0.4 0.6 0.8 1.0 FRACTION OF ORIGINAL SURFACE ARE A/G M R E M A I N I N G

0.2

FIG.18. Correlation between cracking activity and surface area of Si/Al catalyst after various steam treatments.

tained considerable amounts of cumene hydroperoxide, Eq. (30) must also be used. A complete analysis of these data will be found elsewhere [Maatman, Lago, and Prater (IS)]. It is of interest to compare the results obtained with the two types of reactors, differential and integral. This can be done since the data obtained by Plank and Nace were for the same type of catalyst used in our study but with a lower surface area. This reduction in surface area, which results in a lower activity per gram of catalyst, was obtained by treating the catalyst with steam. The activity per unit surface area, however, remains essentially unchanged, as shown by the data presented in Fig. 18. I n this figure the fraction of original activity per gram is plotted against fraction of original surface area per gram remaining after steam treatment. If the activity per unit surface area remains unchanged, the points will lie on a straight line drawn from 1.0 to the origin. This is approximately the case. The independence of d n l d t per unit area makes it possible to compare Kr obtained from these data with that obtained with the

328

CHARLES D. P R A T E R AND RUDOLPH M. L A G 0

350 m.2/g. catalyst. This comparison is shown for quinoline in Fig. 19, which is a plot of In KI against 1/T. The agreement is considered fair, as the value of K , obtained is very sensitive to small effects. The agreement would probably have been better if it had not been necessary for Plank and Nace to overheat the

FIG.19. Arrhenius plot of adsorption-equilibriumconstant for quinoline: 0 differential reactor values, A integral reactor values.

first part of the catalyst bed by 5" to 10°C. to compensate for the endothermic heat of reaction. 3. Influence of the Back Reaction on the Values of Parameters Obtained

with an Integral Reactor The data of Plank and Nace offer a convenient opportunity to investigate influences of the back reaction on the values of parameters, such as the value of the initial rate (dn/dt)p,l and K I obtained from Eqs. (28) and (29). Values of KI for pyridine and styrene and values of initial rate were calculated from these equation with and without the back-reaction term (l/K,)P,,J',. The results are summarized in Table X. Contrary to what might be anticipated, the integral reactor determination of d n / d t and KI is surprisingly insensitive to the back reaction even a t 85% of equilibrium conversion. This is in spite of the fact that this method of determination must become highly dependent on the back-reaction term when equilibrium conversion is approached sufficiently closely. An error

329

KINETICS O F T H E CRACKING OF CUMENE

of only 10.4% is made in the value of (dn/dt)pFr and a n error of only 7.5% is made in KI a t a final conversion, which is 84.4% of the equilibrium conversion. TABLE X Eflects of Including the Back-Reaction Term in Eq. (27) on Calculated Values of Znitial Rate (dnldt) and the Adsorption Equilibrium Constant K I Temperature, "C. Conversion % of equilibrium value

Calc. using integration of Eq. (27) with back-reaction term Calc. using integration of Eq. (27) with no back-reaction term % Difference

426

482 84.4

56.2

K I , for styrene

moles/(g.) (see.)

pyridine

moIes/(g.) (sec.)

3.71 X 10-8

4.72 X 10'

13.5 X 10-8

5 . 9 X lo2

3.39 X 10-6 8.6

4.54 X 104

1 2 . 1 X 10-6 10.4

5.46 X lo2 7.5

4

V. COKEFORMATION 1. Coke Formation from Cumene Hydroperoxide

a. Types of Coke. I n all the discussion so far we have ignored the accumulation on the catalyst of carbonaceous materials usually referred t o as coke. This carbonaceous material can be shown t o be composed of two parts: (1) strongly chemisorbed inhibitors not desorbed at the time of the shutdown operation for removal of the catalyst from the reactor and ( 2 ) harmless carbonaceous materials which do not affect the catalyst activity ( 2 ) . The materials in class 2, to which the name coke will be restricted in the remainder of the discussion, are present in by far the greater amount in most cases. This coke is not desorbable under any conditions that we have found; whereas the adsorbed inhibitors reponsible for the type 1 carbonaceous material can be desorbed if sufficiently high temperatures are used. For experiments with cumene containing only cumene hydroperoxide, the latter is easily desorbed a t 420°C., as shown by the reversibility experiment. The fact th a t coke in itself is not an inhibitor can be seen from the coke data given in Table IX. These data were obtained by repeating the reversibility experiments up to the end of the exposure (1) to pure cumene and (2) t o cumene containing the

330

CHARLES D. PRATER AND RUDOLPH M. LAGO

inhibitor in order to obtain values of coke a t these times. It shows th a t the coke formed from cumene hydroperoxide does not desorb and does not interfere with the return of activity after desorption of cumene hydroperoxide inhibitors. As pointed out by Plank and Nace (Z), pure cumene produces less than 0.1% of coke. However, the inhibitors themselves are heavy coke producers.

MOLE FRACTION OF CUMENE HYDROPEROXIDE

FIG.20.Percentage of coke by weight on catalyst at 4000 sec. as a function of mole fraction of cumene hydroperoxide in cumene.

b. Kineiics of Coke Formation. The kinetics of the formation of coke and the relationship of this kinetics to the active catalyst site can be obtained from the following experimental data on coke formation as a function (a) of mole fraction of inhibitor present, (b) of time, and (c) of pressure. (1) Coke Formation as a Function of Mole Fraction of Inhibitor: Cumene charges containing various mole fractions of cumene hydroperoxide were prepared and passed over the catalyst for 4000 see. T h e percentage of coke by weight on catalyst was determined after a 15-min. oxygen-free nitrogen flush a t the temperature of cracking reaction, 420°C. The values obtained are plotted in Fig. 20. A linear relationship exists

KINETICS O F THE CRACKING OF CUMENE

33 1

between mole fraction of cumene hydroperoxide and percentage of coke by weight. This linear relationship was also observed in experiments a t 2000 and 6000 sec. ( 2 ) Coke Formation as a Function of Time: The amount of coke on the catalyst as a function of time was determined by a series of cracking runs for various lengths of time a t 420°C. The data obtained with cumene containing 0.28% (mole) cumene hydroperoxide are plotted in Fig. 21. The 't function (solid curve). experimental points are approximated by a vThe same functional relation was obtained with 0.14 and 0.56 mole % cumene hydroperoxide.

TIME, MIN.

FIG.21.Percentage of coke by weight on catalyst as a function of time for cumene containing 0.28 mole % cumene hydroperoxide.

(3) Coke Formation as a Function of Pressure: The amount of coke formed is independent of total pressure (Pa P I ) in the range tested (76 t o 44 cm. Hg). This is shown by the data for cumene hydroperoxide at 420°C. given in Table XI. Th e length of the runs was 4000 sec. and the mole fraction of cumene hydroperoxide used is indicated. T h e fact that the rate of coke formation does not depend on pressure means th a t the rate of coke deposition must be independent of the absolute concentration of cumene hydroperoxide and depend only on the mole fraction present. (4) Empirical Kinetic Equation for Coke Formation: The experimental observations given in section V l b ( l ) , (2), and (3) lead to the empirical equation for coke formation c = (2k,t)4521 (31)

+

332

CHARLES D. PRATER AND RUDOLPH M. LAG0

TABLE X I Effect of Pressure on Coke Formation from Cumene Hydroperoxide

Pressure, cm. Hg

Coke (wt. %)

Ratio coke

Mole % cumene hydroperoxide

44 76 44 76 44 76

2.81

0. 93

0.30

3.02 1.63 1.54 9.5 9.5

1.03

0.14

1.0

1.00

Avg. 0.99

where k , is the rate constant for coke formation. This shows the rate of coke formation to be

( 5 ) Relationship between Coke Formation and Cracking Sites: An obvious extension of the kinetic scheme t o include coke formation [see Plank and Nace (a)] is given by

+ A$kiSA+pkaro d u ct + A kr kQ I + A IA coke + A ks

S

k2

$

(33)

4

Coke combined with active sites is not included, as the type of coke under consideration does not deactivate the catalyst. The coke steps are symmetrical with the cracking steps. This leads to an equation for the rate of coke formation analogous with Eq. (11).

Equation (1 1) is transformed into this equation by replacing P I, P,, G, Kr by P,, PI , ~ / K I ,1/G, respectively. The values of G and K I are already known from inhibitor experiments with cumene hydroperoxide and from pure-cumene experiments. Their values are 1.47 atm. and 4.96 X lo3 atm.-', respectively. For P , = 1, Eq. 34 becomes dc _

dt - lrgBo PI

PI

+ 3.39 X

(35)

KINETICS OF THE CRACKINQ OF CUMENE

333

This is a steady state solution and was obtained because the transient has essentially decayed by 400 sec. (See Fig. 14.) Thus, the right side of Eq. (35) is independent of time. On integrating (35) we obtain c = keBo

PI

pr

+ 3.39 x 10-4

Thus, we get a linear dependence on time instead of a square-root dependence. In addition, G will be, according to Eq. (36), nonlinear in Pr at much too high a partial pressure of inhibitor to agree with the data in Fig. 20, as a simple calculation will verify. Thus, we conclude that the proposed mechanism is not the actual mechanism of coke formation, I n fact, since IA becomes approximately constant (essentially all sites occupied) a t about the middle of the range of concentrations of inhibitor mole fraction of inhibitor, Fig. lo), it is difficult to studied (2.5 X see how a reasonable kinetics operating from IA can account for the observed range of linear behavior as a function of concentration. Below approximately 2.5 x mole fraction the kinetics has to give a linear dependence of coke on mole fraction of inhibitor when the amount of chemisorbed inhibitor on cumene-cracking site is rapidly varying. On the other hand, above approximately 2.5 X 10-3 mole fraction the linear dependence has to be obtained with an essentially constant amount of inhibitor chemisorbed on cracking sites. This suggests that the coke is not formed by the inhibitor adsorbed on active cracking sites but by some other mechanism possibly involving other sites which chemisorb inhibitors. 2. Coke Formation from Gas Oil I n order to determine whether the relatianships discovered for coke formation from cumene hydroperoxide might reflect a more general behavior concerning coke formation in cracking reactions, experiments were undertaken with a light East Texas gas oil (LETGO). Such gas oils contain a large variety of inhibitors. a. Coke as a Function of Mole Fraction of Inhibitor. The kinetics of coke formation from LETGO exhibits the same time and mole-fraction dependence as that observed for the coke formation from cumene hydroperoxide. The mole-fraction dependence can be demonstrated by preparing mixtures of cumene and LETGO containing various volume fractions of LETGO and passing the mixture over catalysts in the differential reactor. The results are shown in Fig. 22. The relation between volume fraction and the amount of coke at 4000 sec. is again linear. b. Coke as a Function of Time. The relationship % coke a l/= holds when LETGO is passed over silica-alumina catalyst in the diff eren-

334

CHARLES D. PRATER AND RUDOLPH M. L A G 0 3.a

2.0 W

i

0 V

8 1.0

0

I

0

1

I

1

.20 .40 .60 .80 VOLUME FRACTION OF LETGO

Fro. 22. Percentage of coke by weight as a function of volume fraction of light East Texas gas oil present in cumene.

TIME, M I N .

FIG.23. Effect of amount of coke deposited on catalyst as a function of time for cracking of light East Texas gas oil. The solid line represents the relationship % coke = 0.47 (min.).

tial reactor, as shown by the data presented in Fig. 23. The solid curve is % coke = 0.47 d t ( m i n . ) . A square-root-of-time relationship has been shown to hold for the cracking of gas oil in integral reactors by Voorhies (18) and by Crawford and Cunningham (19). c. Regeneration of Catalytic Activity by Desorption of Inhibitors. The coke formed from LETGO contains the two types of carbonaceous

KINETICS OF T H E CRACKING O F CUMENE

335

materials mentioned above: (1) desorbable materials, which are responsible for a loss of catalytic activity, and (2) nondeactivating nondesorbable coke. However, the desorbable inhibitors require a higher temperature for their removal than for the removal of chemisorbed cuniene hydroperoxide (or its decomposition products). This is shown by experiments in which cumene containing 10% by volume of LETGO is passed over the catalyst in the differential reactor a t 420°C. The inhibitors in LETGO

;

TIME, SEC.

FIG. 24. Catalytic reactivation by desorption of inhibitors deposited on the catalyst by light East Texas gas oil.

act as inhibitors for the cumene cracking, leading to a decrease in activity moles/(m.*)(sec.) to 20 X moles/(m.2)(sec.) in from 260 X 4000 sec. The results are shown in Fig. 24. This is a plot of cracking activity vs. time. Curve 1 shows, for the first 4000 sec., the time course of the rate of cracking of cumene containing 10% by volume of LETGO. This charge was removed and the catalyst was flushed for 1 hr. with oxygen-free nitrogen a t 420°C. A second 4000-sec. cracking run was then made using pure cumene. The resulting activity is shown by curve 2 of the figure. N o recovery of activity was observed. However, when a 1-hr. nitrogen flush a t 510°C. was used, partial recovery was observed, as shown by curve 3. Desorption studies were made at high temperatures by use of catalyst exposed t o LETGO in an integral reactor similar t o th a t described by Hansford (23). Light East Texas gas oil was passed over 100 cc. of catalyst at the rate of 1.2 cc./min. to give a space velocity of 0.72 [cc. charge/ (cc. catalyst) (hr)]. The temperature of the cracking runs was 426°C. The sequence of events was as follows: 1. A 20-min. cracking run was made. The 410°F. gasoline yield was measured as described by Hansford (13). 2. Without being removed from the cracking apparatus, the catalyst was subjected to one of three Orfree nitrogen flushing procedures:

336

CHARLES D. PRATER AND RUDOLPH M. LAG0

a. 5-min. flush at 426°C.; flow rate 1 liter/min. b. 1-hr. flush a t 509°C.; flow rate 1 liter/min. c. 1-hr. flush at 650°C.; flow rate 1 liter/min. The 6rst flush is designed to give negligible desorption of inhibitors, and the second to give some desorption of inhibitors. The third flush procedure is used to give almost, complete removal of the adsorbed inhibitors. 3. After the flushing period, another 20-min. cracking run a t 426°C. was made. The gasoline yield was again determined.

The results are summarized in Table XII. The high-temperature flush restored 90 % ' of the loss in conversion measured after the low-temperature flush, in spite of the fact that approximately the same amount of coke is on the catalyst at the end of each flush. Thus, regeneration of the catalyst TABLE XI1 Regeneration of Silica-Alumina Catalysts by Desotption of Chemisorbed Inhibitors

First 20-min. cracking run to produce spent catalysts

Experiment No.

410°F. Gasoline yield vol. % of charge

(1) (2) (3)

42.0 42.8 42.4

Coke at end of second run

Avg. 42.4 Second 20-min. cracking run following a 5-min. nitrogen flush a t 800°F. Second 20-min. cracking run following a 1-hr. nitrogen flush a t 950°F. Second 20-min. cracking run following a 1-hr. nitrogen flush at 1200°F.

(1)

34.6

1.17

(2)

37.9

1.29

(3)

41.6

1.15

can be accomplished without the necessity of burning the coke. That such flushing procedure restores the catalytic activity has been previously reported (60)but was interpreted as due t o dehydrogenation of coke. That the process actually involves the desorption of inhibitors can be demonstrated by collecting the material desorbed. For a sample of catalyst exposed to LETGO for 20 min. a t 426°C. and removed from the apparatus after a 20-min. oxygen-free nitrogen purge a t 426"C., 0.28% coke by weight was collected by desorption with 02-free nitrogen at 650°C. for 1 hr. The permanent coke left on the catalyst after the desorption procedure was 1.0% by weight. This desorbed material had a high

KINETICS O F T H E CRACKING O F CUMENE

337

inhibition action for cumene cracking when added to cumene, as shown by diff erential-reactor studies.

VI. DISCUSSION AND SUMMARY The kinetic scheme given by Eq. (9) and its more complete form, Eq. (24), have been found to be consistent with the results of studies of several variables: 1. Partial pressure of cumene, P,. 2. Temperature. 3. Mole fraction of inhibitor added to cumene. 4. Total pressure with cumene containing cumene hydroperoxide. The studies reported were made over a wide temperature range, 300" t o 500"C., and under conditions such th a t diffusion transport phenomena do not affect the kinetic conclusions. The best methods known to us were used t o free the cumene from inhibitors of the cracking reaction. The presence of appreciable amounts of inhibitors in the cumene would affect the value of the constant G obtained. Let G' be such a measured constant and G be the true constant. Then G' = G ( l K I X ~ )

+

where K I and xI are the adsorption constants and mole fraction, respectively, of the unknown inhibitor. The true value of the constant G will have an exponential temperature dependence according to the equation Thus, Since K I is a n exponential function of temperature, the experimental constant G' will give a straight line on plotting In G' vs. 1/T only if K I x r is negligible compared t o unity over the entire temperature range studied. Such a linear relationship is observed between In G' and l / T . We conclude, therefore, that the cumene used by us was essentially free of inhibitors. The results are consistent with the following descriptions of the catalyst : 1. The surfaces of the pore walls are sparsely covered with active cracking sites (at most 10'7 sites/m.P). 2. Little or no interaction occurs between molecules chemisorbed on cracking sites. 3. The cracking sites are homogeneous with regard to adsorption properties.

The cracking of cumene can be used as a probe to study the chemisorption of substances on precisely defined active sites. The studies above

338

CHARLES D. PRATER AND RUDOLPH M. L A G 0

show how the adsorption equilibrium constant can be determined for substances on the cracking sites and how the rate constants for adsorption and desorption can be determined if the rates are not too rapid. From the temperature dependence of the values of the adsorption equilibrium constant, free energy, heat, and entropy of adsorption can be determined. This type of study can be reversed in that, after the adsorption characteristics of a substance are known, the information can be used to design experiments to give better direct measurements of the number of active cracking sites, Bo, than have been available to date. The method used by Mills, Boedeker, and Oblad (14) has an important limitation in t ha t sites which may also adsorb the compound used t o count sites but may not participate as cracking sites are not distinguished from active sites. If such noncracking sites have adsorption properties different from those of cracking sites, the material chemisorbed on them can be distinguished from that adsorbed on the cracking sites if the chemisorption properties of the cracking sites are known in advance. Good experimental values of Bo would give a check for the absolute rate-theory equation (16). The type of study outlined above yields meaningful constants, G and Ic3Bo, which can be used to distinguish two possible sources of differences between various cracking catalysts. The constant G can be considered to measure the chemisorption capability of the catalyst for the reactant, and k3Ro measures its capability to crack the chemisorbed reactant. The coke studies lead to the interesting result that the nondesorbable coke, which does not deactivate the catalyst and is produced by inhibitors, is probably not produced by a mechanism involving the cracking sites. Coke formed from light East Texas gas oil behaves in a manner similar t o that of the coke formed by cumene hydroperoxide. Two varieties of carbonaceous material were found for light East Texas gas oil coke, corresponding t o the two varieties found with cumene hydroperoxide coke. One is desorbable and is responsible for catalyst deactivation. T h e other is not desorbable and appears harmless except for possible poreblockage effects. However, there is still a possibility th a t some deactivation of active sites might be due to inhibitors which decompose on heating t o give carbon and free active sites instead of desorbing.

REFERENCES 1. Weisz, P. B., and Prater, C. D., Advances in Catalysis 6, 143 (1954). 2. Plank, C. J., and Nace, D. M., Ind. Eng. Chem. 47, 2374 (1955). 3. Obolentsev, R. D., and Gryazev, N. N., Dolclady Akad. Nauk. S.S.S.R. 73, 121 (1950). 4. Obolentsev, R. D., and Gryazev, N. N., J . Gen. Chem. (U.S.S.R.) 21, 860 (1951).

KINETICS O F THE CRACKING O F CUMENE

339

6. Ballod, A. P., and Gurbich, L. V., Korbov, V. V., and Frost, A. V., Vestnik Moskov. Univ. 6, No. 2, Ser. Fiz. Mat. i Estestven N a u k No. 1, 57 (1951). 6. Ballod, A. P., and Patsevich, I. V., Fel’dman, A. S., and Frost, A. V., Doklady Akad. Nauk S.S.S.R. 78, 509 (1951). 7. Topchieva, K. V., and Panchenkov, G. M., Doklady Akad. N a u k S.S.S.R. 74, 1109 (1950). 8. Corrigan, T. E., Garver, J. C., Rose, H. F., and Kirk, R. S., Chem. Eng. Progr. 49, 603 (1953). 9. Laidler, K. J., in “Catalysis” (P. H. Emmett, ed.), Vol. I, p. 123. Reinhold, New York, 1954. 10. Hougen, 0. A., and Watson, K. M., “Chemical Process Principles,” Part 111, p. 962. John Wiley and Sons, New York, 1943. 11. Weisz, P. B., Science 123, 887 (1956). i2. Wheeler, A., Advances in Catalysis 3, 250 (1951). 13. Hansford, R. C., Advances in Catalysis 4, 1 (1952). 14. Mills, G. A,, Boedeker, E. R., and Oblad, A. G., J . Am. Chem. SOC.72,1554 (1950). 15. Glasstone, S., Laidler, K. J., and Eyring, H., “The Theory of Rate Processes,” Chapter VII. McGraw-Hill, New York, 1941. 16. Maatman, R. W., Lago, R. M., and Prater, C. D., to be published. 17. Thiele, E. W., Ind. Eng. Chem. 31, 916 (1939). 18. Voorhies, A,, Ind. Eng. Chem. 37, 319 (1945). 19. Crawford, P. B., and Cunningham, W. A., Petroleum Refiner 36, 169 (1956). 80. Mills, G. A,, U S . Patent 2,647,042 (1953).

This Page Intentionally Left Blank

Author Index Numbers in parentheses are reference numbers and are included to assist in locating a reference where the author’s name is not mentioned in the text. Numbers in italics indicate the page on which the reference is listed. Boedeker, E. R., 309(14), 315, 338, 559 Bokhoven, C., 69(151), 107(258), 109

A

Abe, S., 194, 201(28), 206 Abendroth, B., 65(123), 155 Adkins, H., 191, 201(23), 206 Amberg, C. H., 62(114), 165 Anderson, J. S., 95(212), 165 Anderson, L. C., 216(12), 217 Andrew, F. F., 62(114), 155 Asbury, W. C., 271(23), 276(23), 291

B Back, I(.J. C., 200(41), 206 Baes, C. F., 194, 201(29), 206 Bagg, J., 130, 133, 134, 144, 160, 161 Balandin, A. A., 111(288), 168 Ballod, A. P., 296, 305, 559 Barer, R. M., 100(225), 166 Barsh, M. K., 165(3, 4), 171, 172, 177, 201(3, 4), 204 Becker, J. A., 53(76), 110(277), 119(323),

(272), 135(367, 368), 136, 144(392), 148(405), 164, 167, 160, 161 Bond, G. C., 143(391), 161 Bonhoeffer, K. F., 76(165), 164, 194, 201 (27), 206 Born, M., 26, 160 Boudart, M., 125, 127, 129, 139, 144, 148 (403), 149, 169, 160, 161 Bovarnik, M., 200(40), 206 Braunbek, W., 56(92), 152 Briggs, W. S., 282(41), 292 Broos, J., 67(138), 163 Briicke, E., 111(285), 168 Bruining, H., 56(86, 87), 162 Brunauer, S., 28(18), 31(29), 37(42), 64, 65(129), 100(225), 136,160,161,153, 156, 160

Burgers, W. G., 111(287), 168 Burnett, R. E., 214(11), 217 Burshtein, R. K., 68, 95(213), 96(213), 97

162, 158, 169

(219), 164, 166, 166

Beebe, R. A., 58(101), 59(105), 62(114), 71(152), 108(271), 152, 163, 164, 167 Beeck, O., 52(70), 53(73), 56,68, 107, 108 (264, 268), 113, 114, 128, 130, 136, 138, 144,162, 164, 167, 168, 169, 160,

Byrns, A. C., 274(29), 292

C Calvin, M., 165, 166, 168(1, 7), 169(1),

161

Benjamin, M., 111, 112, 168 Benson, G. C., 23, 150 Benson, G. W., 23, 160 Benton, A. F., 147(401), 161 Berg, C., 274(29), 275(30), 292 Bergius, F., 239, 291 Bed, E., 132(361), 160 Berthelot, P. E. M., 239, 291 Bevan, D. J. M., 95(212), 156 Bilwiller, J., 239(2), 291

171, 173, 201(1, 7), 204

Cassel, H. M., 87(195), 166 Cavanagh, B., 200, 206 Cawley, G. M., 256, 291 Chaffee, C. C., 291(48), 292 Chatt, J., 190, 204 Chessick, J. J., 66(134,135), 67(135), 102, 163, 166

Claeys, Y. M., 194, 201(29, 30), 206 Clark, E. L., 291(49), 291 Clarke, E. A., 291(48), 292

341

342

AUTHOR INDEX

Coehn, A., 96(215), 156 Cohen, E., 2, 14 Coley, J. R., 207(3), 209(3), 211(8), 212 (9, lo), 217 Corrigan, T. E., 296,301,303(8), 305,539 Corson, B. B., 182, 205 Couper, A., 56(83), 152 Crawford, P. B., 334, 339 Crawford, V. A., 100, 101(227), 156 Cunningham, R. E., 114(302), 168 Cunningham, U ' . A., 334, 359 Custers, J. F. H., 31(28), 35(35), 63(116), 64(118), 66(133), 79(176), 80(177), 104(236, 237), 106(250), 150, 165, 154, 156, 157

Deryagin, B. V., 106, 167 Dharmatti, S.S., 165(4), 177, 201(4), 204 Dilke, M. H., 56, 162 Dippel, C. J., 35(36), 36(37), 67(137), 72 (159), 130(356), 137(381, 382), 150, 155, 154, 160 Dolgov, B. N., 207(2), 209, 216, 217 Domick, A. D., 291(49), 292 Donath, E., 209, 21 7 Dorgelo, G., 114, 158 Dowden, D. A., 55(79), 56(84), 58(101, 102), 59(105, 106), 71(153), 152, 164, 202(47), 206 Drain, L. E., 36, 66, 83, 101, 102, 103, 151, 153, 165, 156 Drechsler, M., 110(280), 158

D E Dakers, R. G., 178, 182, 183(11), 184(11), 201(11, 16), 204 Dayton, J. C., 194, 195, 196(31), 197, 201 (30, 31, 32), 205 de Boer, J. H., 4(11), 14, 22(3), 23(6), 24 (10, 11, 12), 25(13), 26(16, 17), 29 (19, 21, 22), 30(25), 31(28, 30), 33 (32), 34(34), 35(34, 35, 361, 36(37, 39), 38(44,45,46, 48), 39(51), 40, 42, 43(55), 45(58, 591, 48, 52(72), 56(80, 86, 87, 88), 57(94, 95), 58(100), 61 ( l l l ) , 62(112, 113, 115), 63(116), 64 (118), 65(120, 121, 125, 130), 66(133, 136), 67(137, 138), 68(139, 140, 144), 72(157, 158, 159, l60), 75(158), 76 (183, 167), 79(173, 174, 175, 176), 80 (177, 178, 180, 181), 81(184), 82 (184), 83(188), 85(189, 191), 87(194, 196), 88(197), 90(201), 92(204, 206), 93 (206), 96 (216), 97 (2181, 100(223a-e), 104(235, 236, 237, 239), 105 (240, 242, 243, 246), 106(249, 250, 252), 108(267), 109(274), 111(283), 118(319), 119, 120(324), 121(324), 123(327, 328), 124(330), 125[331), 126(339), 127(341), 130(356), 137 (380, 381, 382), 145(399), 147(400), 160, 151, 156, 155, 164, 165, 156, 157, 168, 159, 160, 161 de Boer, N. H., 127, 129(352), 159, 160 Dennis, K. S.,102, 156 Dent, B. M., 65(126), 153

Eischens, R. P., 115, 158 Eley, D. D., 49, 56(83), 68, 128, 151, 166, 153, 159 Emmett,P. H.,69,71(152), 115, 154,158 Engell, H. J., 3(7), 14, 60(108), 153 Estermann, J., 23(5), 150 Eueken, A., 69, 108(265), 154, 157 Eyring, H., 92, 156, 315(15), 539

F Farkas, A., 76(165), 154, 200, 205 Farkas, L., 200, 205 Fast, J. D., 96(214, 216), 97(218), 145 (399), 148(404), 166, 161 Federova, A. J., 129, 159 Fel'dman, A. S., 296(6), 305(6), 539 Fianda, F., 147(402), 161 Field, E., 112(295), 115(309), 158 Flournoy, J. M., 195, 196(31), 201(31), 205 Flynn, J. H., 189, 190, 201(19, 20), 804 Forestier, H., 5(14), 11, 1 4 , 15 Fortuin, J. M. H., 81(182), 140, 156, 161 Fortuin, J. P., 79(171), 154 Forward, F. A., 182(14), 204 Francini, T., 96(215), 156 Franck, J., 96(215), 156 Frank, L. V., 291(49), 292 Frankenburg, W., 107, 167 Fraser, R. G. J., 23(4), 150

343

AUTHOR INDEX

Frenkel, J., 85, 155 Frisch, R., 23(5), 150 Frost, A. V., 296(5, 6), 305(5, 6), 339 Frumkin, A. N., 129, 135, 159, 160

G Gans, R., 140, 161 Garner, W. E., 57(98), 58(99, 102, 103, 104), 59(106), 60(107), 152 Garver, J. C., 296(8), 301(8), 303(8), 305(8), 339 Gauchman, S. S., 69, 154 Gavrilova, E. Y., 207(2), 217 Gest, H., 200(43), 205 Gilman, H., 199, 205 Ginsberg, H. H., 291(49), 292 Glasstone, S., 92(203), 155, 315(15), 339 Goetz, A., 2(2), 14 Goldmann, F., 100, 156 Golodnikov, G. V., 209, 21 7 Gomer, R., 4(12), 14, 110(279), 125, 158, 169 Goodeve, C. F., 97, 156 Gordon, K., 244(8), 283, 291 Granovskaya, V., 135(367), 160 Gray, T. J., 57(98), 60(107), 15.2 Green, D. E., 200, 605 Greene, S. A,, 102, 156 Grothe, W., 275, 892 Gryazev, N. N., 296, 305, S38 Guenther, G., 272(26), 291 Guinier, A., 4(11), l 4 Gulbransen, E. A., 62(114), 153 Gurbich, L. V., 296(5), 305(5), 339 Gwathmey, A. T., 114(301, 302), 158

H Haayman, P. W., 57(96), 152 Halpern, J., 178, 182, 183, 184(11), 185 (17), 186(17), 188, 201(11, 16, 17, 18), 204, 204 Halsey, G. D., 62(114), lOO(222, 2261, 105, 106, 136, 153, 156, 157, 160 Ham, W. R., 97(217), 156 Hansford, R. C., 307(13), 335, 339 Harkins, W. D., 140, 161 Harkness, R. W., 69, 154 Hartman, C. D., 53(76), 110(277), 152, 158

Haslam, R. T., 271(23), 276(23), 891 Hassid, N. J., 131, 160 Hauffe, K., 3, 14, 60(108), 153 Healey, F. H., 66, 67(135), 102, 163, 156 Hedvall, J. A., 5(14), 6(14a, 15), 10(16), 12(16, 19), 13(20), 14(21), 14, 16 Heinemann, H., 282(41), 292 Heller, G., 30(25), 160 Hellmann, H., 29(20), 150 Heric, E. L., 102, 156 Herington, E. F. G., 60(110), 163 Hill, T. L., 19, 150, 81(186), 104(238), 105, 106, 155, 156, 157 Hirst, L. L., 291(48), 292 Hoberman, H. D., 200(39), 205 Hodler, A., 107(262), 157 Horing, M., 240(4), 291 Holmes, J. L., 108(269), 157 Honig, J. M., 101,229), 156 Horiuti, J., 200, 205 Houben, G. M. M., 68(141), 80(181), 140, 153, 155, 161 Hougen, 0. A., 299(10), 339 Hiickel, E., 33, 150, 247, 291 Hulburt, H. M., 189,190,201(19,20),604 Hunsmann, W., 69, 154

I Jpatieff, V. N., 182, 204 Ivannikov, P. Y., 207(2), 217 Ives, H. E., 39, 151 J

Jabrova, G. M., 145(398), 161 Jack, K. H., 97, 156 Jacoby, A. L., 199, 205 Jenkins, R. O., 111, 112, 158 John, G. S., 112(295), 115(309), 158 Johnson, M. C., 72(156), 91, 154, 155 Johnson, R. P., 111, 158 Johnson, T. H., 74(161), 154 Joklik, W. K., 205(42), 205 Jones, W. A., 272(25), 291 Jordan, J., 76(170), 154 Jurgens, H., 96(215), 156 Justi, E., 11, 16 Juza, R., 76(166), 154

344

AUTHOR INDEX

K Kaftal, G., 244, 283, 291 Kagan, M. Y., 207(1), 209, 216 Kalish, T. V., 97(219), 156 Keenan, R. G., 136, 160 KeIer, M. P., 115, 168 Kemball, C., 81, 85, 89, 100(226), 103, 133, 155, 156, 160 Kingdon, K. H., 39, ll6(316), 151, 159 Kington, G. L., 108(269), 167 Kiperman, S. L., 135(367), 160 Kirk, R. S., 296(8), 301(8), 303(8), 305

’”

(8)J

Kirkwood, J. G. , 29, 160 Kistiakowsky, G. B., 108(270), 167 Klumb, H., 65(122), 163 Koller, K., 56(93), 162 Koller, L. R., 94(207), 155 Kolthoff, I. M., 76(170), 164 Komarewsky, V. I., 207(3), 209(3), 211 (7, S), 212(9, lo), 21’7 Korbov, V. V., 296(5), 305(5), 339 Korinek, G. J., 188,201(18),204, ,804, 205 Koton, M. M., 207(2), 216 Kraak, H. H., 56(88), 76(167), 127(341), 152, 164, 159 Krasna, A. I., 200(44), 206 Krsek, G., 191, 201(23), 205 Kruyer, S., 38(44, 46), 68(144), 76(163), 81(184), 82(184, 186a), 85(193), 88 (197), 89(200), 90(200, ZOl), 100 (223a-e), 161, 164, 166, 166 Kummer, J. T., 71(152), 115, 164, 168 Kurbatov, I. D., 182, 204 Kwan, T., 19, 50(67), 107(255), 108(266), 116, 131, 160, 161, 157, 169, 160

L Lago, R. M., 320, 326, 327, 339 Laidler, K. J., 92(203), 107(256), 165, 167, 298(9), 315(15), 339 Laituri, J. M., 236(1), 338 Langbein, R., 76(166),164 Lange, E., 147(402), 161 Langlois, G. E., 237(2), 238 Langmuir, I., 39, 72(155), 116, 117, 123 (326), 134, 161, 154, 159, 160 Lascelles, J., 200(41), 206 Lehr, J. J., 72(157, 1591, 154

Leidheiser, H., 114(301), 168 Lelchuk, S. L., 207(2), 216, 217 Lenel, F. V., 36, 37, 65(127), 151, 163 Lennard Jones, J. E., 65(126), 163 Levine, R., 191(24), 201(24), 206 Liang, S.C., 115(306), 137(383), 168, 160 Liebich, M. G., 275(31), 292 Livingston, H. K., 81(185), 82, 155 London, F., 29, 30, 64(117), 160, 163 Los, J . M., 101(230), 102(230), 156 Love, K. S.,136, 160 Ludeman, H., 199, 206 Lukirsky, P. I,, 43, 92, 93(205), 161, 166 Lybarskii, G. D., 207(1), 209(1), 216

M Maatman, R. W., 320, 326, 327, 339 McAllister, S. H., 224, 238 McNaughton, N. W., 216(12), 217 McQuillan, A. D., 129(354), 160 Magnus, A., 38(47), 151 Maltbie, M., 26(15), 160 Margenau, H., 29(22), 31, 150 Martin, S. T., 112(291), 168 Massey, H. S. W., 76(169), 164 Mayer, H., 123, 169 Mayer, J. E., 26, 160 Maxted, E. B., 47, 56, 131, 140, 161, 159, 160, 161 Meelheim, R., 114(301), 158 Mignolet, J. C . P., 38, 45(60), 125, 126 (336), 127, 128, 161, 159 Miller, A. R., 107(253), 157 Milligan, W. O., 3(10), 14 Mills, G. A., 165, 166(5), 167(5), 168(5), 169(5), 170(5), 171, 173, 174(5), 177, 178(10), 179(10), 180(10), 199(34), 201(5, lo), 20.4, 206, 282(41), 292, 309(14), 316, 336, 338, 539 Molinari, E., 71, 164 Morrison, J. A., 66, 83, 101, 102, 103, 163, 166, 166

Morse, P. H., 49, 161 Mott, N. F., 95, 156 Muller, E. W., 110, 112(293), 167, 168

N Nace, D. M., 295, 297, 305(2), 320, 329 (2), 330, 332, 338 Nalimow, W. W., 94(208), 165

345

AUTHOR INDEX

Neugebauer, K., 87(195), 166 Nichols, M. H., 110, 168 Novak, H., 275(31), 292

0 Oblad, A. G., 309(14), 315, 316, 338, 339 Obolentsev, R. D., 296, 305, 338 Oettinger, W., 245(11), 284(46), 291, 292 Ogg, R. A., Jr., 200, 106 Orchin, M., 191, 201(24, 25), 206 Orr, W. J. C., 65(128), 66(128), 103, 163, 166 Ouellet, C.,65(124), 163

P Pace, E. L., 102, 166 Panchenkov, G. M., 296, 299, 301, 305,

ss9 Paravano, G., 71, 164 Patsevich, I. V., 296(6), 305(6), 339 Pauling, L., 44(57), 49, 56, 161, 162 Pearce, J. N., 89, 166 Pease, R. N., 114(304), 168 Pelipetz, M. G., 291(49), 196 Peters, E., 183, 185(17), 186(17), 188,201 (17, 18), 204, 804 Pier, M., 245(12), 275(36), 289(47), 291, 292 Plank, C. J., 295, 297, 305(2), 320, 329 (2), 330, 332, 338 Ploos van Amstel, J. J. A., 111(287), 168 Polanyi, M., 30, 100, 160, 166, 200, 806 Pollard, W. G., 31, 160 Polley, M. H., 62(114), 165 Porter, A. S., 136, 160 Prater, C. D., 294(1), 297, 301(1), 302(1), 303(1), 304, 305(1), 306, 320, 323, 326, 327, 338, 339 Pyzhev, V., 135(367), 160

R Read, W. T., Jr., 4(12), 14 Rees, A. L. G., 57(97), 95, 162, 166 Reid, E. E., 214(11), 217 Reitz, O., 260(21), 264, 291 Reyerson, L. H., 101(229), 166 Reynolds, P. W., 56(84), 162

Rhodin, T. N., 99, 166 Rideal, E. K., 60(110), 65(124), 107(260), 111(289), 115(307), 148(406), 163, 167, 168, 161 Ritchie, A. W., 136, 160 Rittenberg, D., 200(39, 44), 206 Riyanov, S., 43, 92, 93(205), 161, 166 Roberts, J. K., 52(69), 53(75), 107, 109, 121,162, 167, 169 RoginskiK, S. Z., 113, 115, 136, 145, 168, 160, 161 Romeijn, F. R., 57(96), 162 Rose, H. F., 296(8), 301(8), 303(8), 305 (8),339 Roukens, J., 53(74), 131, 132(360), 134, 162, 160 Royter, W. A., 69, 164 Ruhoff, J. R., 214(11), 217 Rummel, K. W., 76(165), 164 Russell, R. P., 271(23), 276(23), 991

S Sachtler, W. M. H., 3, 14, 111(284), 114, 168 Sadek, H., 115(306), 137(383), 168, 160 Schaeffer, W. D., 62(114), 163 Schenk, D., 111(286), 168 Schmidt, F., 80(179), 164 Schoening, F. R. L., 175, $04 Schofield, E. B., 126, 169 Schreiner, G. D. L., 100(226), 166 Schuit, G. C. A,, 52(72), 108(267), 116, 127, 129(352), 162, 167, 169, 160 Schuster, C., 247(14), 249(16), 291 Schwab, G. M., 3(6), 4(11, 13), 14, 128, 169 Shalit, H., 282(41), 292 Sheridan, J., 143(391), 161 Sherwood, P. W., 241,191 Shishakov, N. A., 127(342), 169 Shokley, W., 111, 168 Shumilova, N. A., 95(213), 96(213), 166 Shuttleworth, R., 23, 160 Sidgwick, N. V., 44(57), 161 Singer, J., 80(179), 164 Singleton, J. H., 62(114), 163 Sips, R. J., 100(226), 166 Slater, J. C., 29, 160 Slygin, A. J., 135, 160

346

AUTHOR INDEX

Tsygankova, M., 207(2), 21 7 Smith, A. E., 108(268), 133, 167, 168 Twigg, 0. H., 56(91), 111(289), 152, 158 Smith, C. S.,4(12), 14 Tyson, 0. N., 177, 604 Smith, L. G., 211(7), 617 Smith, W. R., 62(114), 169 U Smithells, C. J., 96(214), 166 Sobolev, I, A., 207(1), 209(1), 616 Urban, W., 275(31), 191 Specht, W., 96(215), 166 Spencer, W. B., 62(114), 163 v Squibb, E. R., 208, 61 7 Steffens, J. H., 236(1), 138 van der Knaap, W., 114, 168 Stephenson, M., 199, 606 van Heerden, C., 69(151), 107(258), 109 Stern, O., 23(5), 160 (272), 135(367), 368), 136(376), 148 Stewart, L., 114(304), 168 (4051, 164, 167, 160, 161 Stickland, L. H., 199, 200, 606 van Niekerk, J. N., 175, 604 Still, J. L., 200(41), 606 Stone, F. S., 57(98), 58(103), 60(107), van SteeIiis, J., 72(158, 160), 751158), 92 (204), 164, 166 161 Storch,’H. H., 191, 201(25), 206,241,291 Veal, F. J., 58(99), 166 Veenemans, C. F., 40(52), 118(319), 119 Stranski, J. N., 2, 14, 110, 168 (320), 123(327), 161, 169 Straumann, M., 2(4), 14 Suhrmann, R., 3, 14, 47, 56(89), 76(168), Veltistova, M. V., 207(2), 217 Verwey, E. J. W., 23, 24(12), 26(16, 17), 110, 161, 161, 164, 168 45(58, 59), 57, 58(100), 76(164), 160, Surova, M. D., 95(213), 96(213), 166,166 161, 162, 164

T Taylor, A. L., 89, 166 Taylor, H. S., 69, 108(270, 271), 109, 113, 115, 136, 137(383), 164, 167, 168, 169, 160

Taylor, J. B., 116, 117(317), 123(326), 169 Teller, E., 37 Temkin, & I., I. 128, 135, 136, 169, 160 Teves, M. C., 79(172), 164 Thiele, E. W., 323, 339 Thomas, L. B., 126, 169 Tiley, P. F., 58(103), 162 Timofeew, P. W., 94(208), 166 Todes, 0. M., 115(308), 168 Tompkins, F. C., 100, 101(227), 130, 133, 134, 136, 144, 166, 160, 161 Tolpin, J. G., 112(295), 115(309), 168 Topchieva, K. V., 296,299, 301, 305,339 Trapnell, B. M. W., 51, 52(71), 53(77), 55 (78), 69(147), 107(257, 260), 109 (272), 114(303), 115(307), 126(338), 127(340), 135(368, 369), 136, 148 (406), 161, 164, 167, 168, 169, 160, 161

Vick, F. P.,91, 166 Vieth, G., 11(18), 16 Vivian, R. E., 177, SO4 Vol’kenshtein, F. F., 112, 139, 168, 160 Volmer, M., 2, 14 Voorhies, A., 334, 339

W Wagner, C., 95, 166 Wagner, J. B., 114(302), 168 Walkey, J. E., 237(2), 638 Wallgren, P., 13(20), 16 Ward, T., 58(104), I52 Watson, K. M., 299(10), 339 Webster, A. H., 204, 606 Weingaertner, E., 132(361), 160 Weiser, H. B., 3(10), 1.4 Weisz, P. B., 294(1), 297, 301(1), 302(1, 11), 303(1), 304, 305(1), 306, 323, 338, 339 Weitz, E., 80(179), 164 Weller, S., 165, 166(5), 167(5), 168(5), 169(5), 170(5), 171, 173, 174(5), 177, 178(10), 179(10), 180(10), 199(34), 201(5, 6, lo), 604, 606

347

AUTHOR INDEX

Wells, A. F., 44(57), 161 Wender, I., 191, 193(26), 201(24, 25, 26),

Wright, L. W., 165, 173, 177, 178(10), 201(6, lo), 204

806

Went, J. J., 56(90), 162 Werchowsky, W., 182, 204 Westrik, R., 69(151), 107(258), 109(272, 135(367, 368), 136(376), 148(405), 164, 167, 160, 161

Weyl, w. A., 3, 14 Wheeler, A., 53(73), 107(259), 108(268), 113, 114, 136, 144(396), 162, 167, 168, 160, 161, 199(34), 206, 301(12), 323, 339 Wheland, G. W., 36(38), 44(56), 161 Wilmarth, W. K., 165, 166(2), 171, 172, 177, 194, 195, 196(31), 197, 199, 201 (2, 3, 4, 29, 30, 31, 32), 804, 206 Winter, E. R. S., 60(109), 163 Wirta, K., 194, 201(27), 206 Wittmann, G., 277(39), 892 Wolfson, M. L., 294(49), 292

Y Young, G. J., 66(134, 135), 67(135), 102, 163, 166

Z

Zrtfdenrnan, I. A., 95(213), 96(213), 166 Zeldowitsch, J., 134, 136, 160 Zettlemoyer, A. C., 66(134, 1351, 67(135), 102, 163, 166 Zimmermrtn, M. U., 236(1), 238 Zimmerschied, W. J., 212(10), 217 Zorin, Z. M., 106, 167 Zwietering, P., 53(74), 69, 76(162), 107 (258), 109(272), 131, 132(360), 134, 135(367, 368), 136(376), 148(405), 168, 164, 167, 160, 161

Zwikker, C., 106(249), 167

Subject Index A

mobility and reactivity, 91-92 physical, 20-21 on charcoal, 64-65 at high degrees of occupation, 98106 on ionic surfaces, heat of adsorption as function of degree of occupation, 100 simultaneous adsorption of different molecules in, 140-141 physicochemical changes by, 79-81 optical, 79 ff. of polar molecules, 35-37 relation between mobility and reactivity of adsorbed molecules, 91-92 time of, 85-91 Alcohols, secondary, dehydrogenation of, 208 Aldol condensation synthesis of ketones, 209 development, 209-2 12 reaction mechanism, 212-216 Alumina-silicate catalyst, see Silicatealumina catalysts Alumina-tungsten-nickel catalyst, activity, effect of replacing nickel with other metals on, 273-276 comparison with tungsten disulfide catalyst, 271 hydrogenation of brown coal tar with, 272 Argon, adsorption on ionic surfaces, heat of, 101, 103 Atoms, chemisorbed, mutual assistance of, 147- 148

Acids, decarboxylation condensation of, 208209 Activation, of molecular hydrogen by homogenous catalysts, 163-206 Adsorbent (s), conducting, physical adsorption phenomena on, 100 dielectric, polarization of adsorbed molecule by, 37-38 solution of two-atomic gases into, 9698 Adsorption, see also Chemisorption catalysis and, 19-20 on dielectric surfaces, 33-35 effect of electronic structure of metals on, 55-56 energies, calculation of, 23-29 effect of distance between adsorbed molecule and surface on, 24-26 of repulsion forces on, 26-29 of structure of a smooth adsorbent surface on, 23-24 forces causing, 29-64 heat of, associated with chemisorption on metals, 48 ff. of cesium atoms on tungsten, 116123 on ionic surfaces, 65-68 of ions, 32-35 on metal surfaces, 32-33, 65, 126-128 chemical bond in, 39-47 of ions on, 32-33 multimolecular, two-dimensional condensation and, 104-106 on nonmetallic surfaces, active spots, B 61-64 chemical bonds in, 57-64 Bases, phenomena of, 17-161 activation of hydrogen by, 194-199, mobirity and orientation, 81-91 202 348

SUBJECT INDEX

Benzene, competition with cumene for chemisorption, 319 Bituminous coal middle oil, aromatization of, 287 hydrogenation of, 288 Butylene motor polymer, 230 L;

Carbon, activated, gas adsorption on, 247 Carbon monoxide, adsorption of, 57-58 Catalysis, adsorption and, 19-20 heat of, 148.149 heterogenous, problems of, 1-15 Catalysts, aee also individual catalysts cracking, regeneration b y desorption of chemisorbed inhibitor, 334-337 homogenous, 165-200 activation of molecular hydrogen by, 163-206 for polymerization of olefins, 223-224 poisons for, 224 Catalytic syntheses, of ketones, 207-217 development of ketonization process, 208 Charcoal, mobility of adsorbed atoms or molecules, 81-83 physical adsorption on, 64-65 Chemical bonds, in adsorption on metals, 39-47 formation of ions, 39-43 sharing of electrons, 43-47 in adsorption on nonmetallic surfaces, 57-64 change in, a t high degrees of occupation, 123-125 Chemisorption, 21-22, see also Adsorption activation energies, changes during progressing adsorption, 131-134 effects connected with, 136-138 restricted chemisorption due to increase in, 138-139 contmninated surfaces and, 144-147

349

of different atoms giving dipoles of the same sign, 143-145 endothermic, 75, 77, 149-150 equations for isotherms, 134-136 factors causing a surface heterogeneity, 109-113 heat of, 148-149 decrease with increase in chemisorbed material, 107-109, 125126, 128-131, 139-140 on metals, activation energies associated with, 48 ff. heats of adsorption and desorption associated with, 48 ff. phenomena, a t high degrees of occupation, 107-140 restricted, caused by increase of activation energy, 138-139 simultaneous adsorption of different species in, 141-143 surface for, factors causing heterogeneity of, 109-113 experimental methods for study of, 113-1 16 Coal, hydrogenation of, vapor-phase catalysts for, 239-292 Cobalt carbonyl, catalytic activity, 191-193, 201 Cobalt cyanide, activation of hydrogen by, 200 Condensation, decarboxylation, of acids 209 two-dimensional, muitirnolecular adsorption and, 104-106 Copper pyrophosphate, as catalyst i n polymerization processes, 236-237 Creosote oil, hydrogenation, 283 Crude oil, desulfurization of, 276 Cumene, 232, 233-234 component analyses of streams of, 235 cracking by silica-alumina catalysts, 293-339 coke formation, 329-333 from cumene hydroperoxide, 329333 as a function of mole fraction of inhibitor, 330-331 of pressure, 331

350

SUBJECT INDEX

of time, 331 reduction in organic bases, 168-171, kinetics of, 330, 331-332, 337-338 177-178 relation between cracking sites Cuprous salts, and 332-333 activation of hydrogen by, 165-181 types of coke, 329-330 factors affecting, 172-173 composition of gas formed, 295 mechanism of, 174-177 determination of adsorption constants with an integral reactor, D 324-329 Decahydronaphthalene, effect of back reaction on paramsplitting in presence of coronene, 257 ff. eter values, 328-329 Dehydrogenation, integration of kinetics, 324-327 of secondary alcohols, 208 determination of values of constants, Desorption, 310-316 activation energies, associated with diffusion-transport effects, 301-304 chemisorption on metals, 48 ff. adsorption constants of products heats of, associated with chemisorption from, 323-324 on metals, 48 ff. kinetics of, 293 ff. comparative conversion data, 300, Deuterium, reduction of quinone by, 171 301 Diesel fuel, as determined by differential-reacdesulfurization of, 276 tor studies, 305-324 Differential reactor, differentia1 and integral forms of, for kinetic studies, 306 299 Dimer, 232, 236 previous studies, sources of uncertainties, in, 295, 298-305 E scheme of, 309-316 Ethane, use of methods-insensitive to, 295adsorption on tungsten disulfide cata301 lyst, 247, 250 rate of, effect of inhibitors on, 304Ethylbenzene, 232, 236 305, 316-319, 320-322 Ethylene platirious chloride, of pressure and temperature, reduction of, 189-191 308, 310-315 structure, 232 F U.O.P. cumene unity, 225 yields, 234 Forces, Cumene hydroperoxide, causing adsorption, 29-64 chemisorption, adsorption arid desorpcooperation among, 64-81 tion of chemisorbed, 320-322 repulsive, effect on calculation of adrate constants of, 321-322 sorption energies, 26-29 reversibility of, 320-321 van der Waals’, see van der Waals’ coke formation from, 329-333 forces effect on cracking of cumene, 316-319, Fuller’s earth catalysts, see also individ320-322 ual catalysts Cupric salts, activity, effect of molybdic acid on, 288 activation of hydrogen by, 168-178, of splitting, 282 201, 203-204 Fuller’s earth-tungsten disulfide catalyst reduction in aqueous solutions, 181-182 (Terrana), kinetics, 182-188 activity, nitrogen bases and, 279

351

SUBJECT INDEX

comparison with alumina silicate catalyst, 284, 285 with tungsten disulfide catalyst, 277

G Gas oil, coke formation from, 333-337, 338 as a function of mole fraction of inhibitor, 333, 334 of time, 333-334 splitting of, activity of catalyst supports for, 266 Gases, adsorption on catalysts, 247 polymerization of olefins from cracked, 219-238 two-atomic, solution into adsorbent, 96-98 Gasoline(s), from bituminous coal, Raman analysis of, 281 fractions, aniline point of, 285 properties of, 280

H Heptene, 232, 235 Hydrides, bond energies for surface of, 50 Hydrocarbons, chemisorption of, 60 Hydrogen, activation by cupric salts in organic bases, 168-178, 201 of molecular, by homogenous catalysts, 163-206 adsorption of, 59-60, 61, 68-76 amount taken by potassium, 93 on metals, 107 heats of, 108, 130 on potassium, 92 ff. on tungsten disulfide, 247-248, 250 Hydrogenase, 199-200 Hy drogenolysis, of R M compounds, 199

I Ions, adsorption on dielectric surfaces, 33-35 on metal surfaces, 32-33 Iso-octane, 230

K Ketones, catalytic syntheses of, 207-217 development of ketonization process, 208-216

L Langmuir-Hinshelwood mechanism, 92

M Mercuric salts, activation of hydrogen by, 188, 201, 203-204 Metals, adsorption on, chemical bonds in, 3947 effectof electronic structureon, 55-56 physical, 65 Molecule(s), adsorbed, polarization by conducting adsorbent, 38-39 by dielectric adsorbent, 37-38 hopping, 83-85 polar, adsorption of, 35-37 simultaneous adsorption of different, 140-1 41 Molybdenum-nickel prehydrogenation catalyst, 267 ff. Motor fuels, see also individual compounds derived from polymerized olefins, 227230 polymerization reactions, 225-230 polymerization units, 226 properties of, 229, 231

N Nitrogen, adsorption on metals, activation energy of, 132, 134 heat of, 99, 101, 133 0

1-Octanol-d, preparation of, 214

352

BUBJECT INDEX

1-Octanol-2-d, conversion to ketone, 214-216 preparation of, 214 Octanols, infrared spectra, 215 Olefins, from cracked gases, polymerization of, 219-238 catalysts used in, 223-224 polymerization reactions, 225-236 polymerization units using phosphoric acid, 236-238 Orthohydrogen, interconversion with p-hydrogen, 171172 “ 0 x 0 ” reaction, 191-193 Oxygen, adsorption of, 60 on metals, 95-96 chemisorption of, 76-79

P Parahydrogen, interconversion with o-hydrogen, 171172 Pentamer, 232, 236 Petrochemicals, 230-236 see also individual compounds polymers used as, 232 Phosphoric acid, as catalyst in polymerization processes, 236-238 liquid, on-quartz, 237, 238 polyco type, 236-237 solid U.O.P., 236, 237 Polarization, of an absorbed molecule by a conducting adsorbent, 38 by a dielectric adsorbent, 37-38 basic factors in, 221-225 of olefins from cracked gases, 219-238 catalyst poisons, 224 types of catalysts used, 223-224 Potential curves, relating to adsorption of cesium on a tungsten surface, 42 of hydrogen on metal, 49 on nonmetallic surfaces, 73 of sodium on a tungsten surface, 41

relating to chemisorption, 51 ff. dissociative, 51, 53, 54 endothermic, 55 of oxygen, 77, 78 to the combination of cesium and fluorine, 43 to combination of sodium and chlorine, 40 to photo and thermal ionization of adsorbed cesium, 80 Prehydrogenation catalyats, 267 ff. activity, temperature and, 269-271 Propylene motor polymer, 230

Q Quinoline, adsorption by silica-alumina catalysts, 316 Quinone, reduction of, 166-168 by deuterium, 171

R R M compounds, hydrogenolysis of, 199 Rideal mechanism, 91

s Silica-alumina catalysts, adsorption of quinoline by, 316 comparison with Fuller’s earth-WSz catalyst, 284, 285 cracking of cumene by, 293-339 kinetics of, 295 ff. number of active sites, 315-316 inhibitors of, 304-305, 316-322 adsorption constants for, 320 regeneration b y desorption of chemisorbed inhibitors, 336 Silica gel, gas adsorption on, 247 Silver acetate, reduction in organic bases, 178-181, 20 1 Silver salts, activation of hydrogen by, 203-204

353

SUBJECT INDEX

T Terrana, see Fuller’s earth-tungsten disulfide catalyst Tetramer, 232-233 boiling point curve, 233 properties, 233 yields, 233 Tischenko condensation synthesis of ketones, 209 Trimer, 232, 234 Tungsten , adsorption of cesium on, heats of, 116123 Tungsten disulfide, adsorption of gases on, 247-250 Tungsten disulfide catalysts, 245-264 activity, factors affecting, 257-260 influence of pellet size on, 260-264 comparison with alumina-tungstennickel catalyst, 271 with Fuller’s-earth-WSz catalyst, 277 hydrogenation of brown coal tar with, 272 hydrogenation reactions, 250-254 mechanism of, 254 ff. prehydrogenation reactions, temperature of, 265

preparation, 245-246 properties, 246-250, 261

U U.O.P. catalytic polymerization units, 226

V Vapor phase Catalysts see also individual catalysts activity, effect of aniline on, 280 for coal hydrogenation, 239-292 development of 242-245 nonsplitting, 264-276 splitting, 276-291 catalyst support and octane number, 287 effect of feed fraction on conversion, 277 reaction mechanism, 289

W van der Waals’ forces, nonpolar, attraction forces, 29-31 on conducting surfaces, 31-32

This Page Intentionally Left Blank