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Nuclear Energy
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VEZIROGLU Energy Conversion — A National Forum
DE WINTER Sun: Mankind's Future Source of Energy
ZALESKI Nuclear Energy Maturity
RELATED JOURNALS PUBLISHED BY PERGAMON PRESS
International Journal of Hydrogen Energy Annals of Nuclear Energy Progress in Nuclear Energy Solar Energy Sun World Progress in Energy and Combustion Science Energy Conversion Energy
HYDRIDES FOR ENERGY STORAGE Proceedings of an International Symposium held
in Geilo, Norway, 14 - 19 August 1977
Edited by
and
ORGANIZED BY The Netherlands Norwegian Reactor School,
Institutt for Atomenergi, Kjeller, Norway
Published on behalf of the
INTERNATIONAL ASSOCIATION FOR HYDROGEN ENERGY
by'
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Copyright © 1978 International Association for Hydrogen Energy All Rights Reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means: electronic, electro­ static, magnetic tape, mechanical, photocopying, recording or otherwise, without permission in writing from the copyright holders. First edition 1978
British Library Cataloguing in Publication Data
Hydrides for energy storage. 1. Hydrogen as Fuel - Congresses 2. Hydrogen - Storage - Congresses 3. Metal hydrides - Industrial applications - Congresses I. Andresen, A F II. Maeland, A J III. Netherlands - Norwegian Reactor School 665'.81 TP359.H8 78-40501 ISBN 0-08-022715-5
In order to make this volume available as economically and as rapidly as possible the authors ' typescripts have been reproduced in their original forms. This method unfortunately has its typographical limitations but it is hoped that they in no way distract the reader.
Printed in Great Britain by William Clowes & Sons Limited London, Beccles and Colchester
FOREWORD
Hydrogen is considered as one of the most promising fuels for the future. It is non-polluting, fully recycleable and has an almost unlimited supply potential. It can be distributed through pipe­ lines or stored in containers for automotive use. However, the con­ ventional means of hydrogen storage has serious short-comings. Pressure cylinders are heavy, expensive and volume demanding. Lique­ faction is energy consuming and require complicated cryogenic equip­ ment. In both cases the safety aspects pose serious problems. It has long been known that many hydrides contain more hydrogen per unit volume than liquid or even solid hydrogen. Some alloy systems absorb and release hydrogen at room temperature at pressures close to atmospheric pressure. Indeed, metal hydrides offer a reversible chemical means for storing and supplying hydrogen which can con­ veniently be used for both mobile and stationary purposes. For several years extensive research has been carried out in many labo­ ratories to find suitable alloy systems. However, a fully satis­ factory system has not yet been found. All the reported hydrides suffer at least one of the following drawbacks: Too low a ratio between hydrogen and metal weight, too costly metals involved, the absorption or release of hydrogen is difficult and slow or sensitive to poisoning phenomena. The intention with this symposium was to bring together research workers active, either from a practical or fundamental point of view, in the field of hydrides. By discussing the fundamental properties of hydrides we intended to stress the possibilities and limitations which exist and possibly bring out new ideas for future research. With the wide range of activities now being carried out in this field, we felt that there was a need for a survey of the activities and a review of the present state of the art.
IX
SYMPOSIUM COMMITTEES Programme Committee; A.F. Andresen Institutt for Atomenergi P.O.B. 40, 2007 Kjeller Norway T.B. Flanagan University of Vermont Burlington, Vermont 05401 USA G.G. Libowitz Allied Chemical Corp. P.O.B. 1021R Morristown, N.J. 07960 USA A.J. Maeland Allied Chemical Corp. P.O.B. 1021R Morristown, N.J. 07960 USA H.H. van Mai N.V. Philips1 Gloeilampenfabrieken Eindhoven The Netherlands K. Videm Institutt for Atomenergi P.O.B. 40, 2007 Kjeller Norway
Organizing Committee: E. Andersen The Netherlands-Norwegian Reactor School Institutt for Atomenergi P.O.B. 40, 2007 Kjeller Norway A.F. Andresen Institutt for Atomenergi P.O.B. 40, 2007 Kjeller Norway G. Jarrett The Netherlands-Norwegian Reactor School Institutt for Atomenergi P.O.B. 40, 2007 Kjeller Norway
xi
ACKNOWLEDGEMENTS The symposium commitees gratefully acknowledge the financial support and services rendered by Institutt for atomenergi, Kjeller, Norway They are also grateful to Allied Chemical Corporation, New Jersey, U.S.A., for financial support. These proceedings were published under a grant from the United States Department of Energy, Washing­ ton, D.C., U.S.A. Thanks are also due to all the invited speakers and the other lec­ turers for their cooperation in preparing the manuscripts.
XI11
Achard, J.C., Equipe de Chimie Métallurgique, Bellevue-Meudon, France Angus, H.C., INCO Europe Ltd., Birmingham, U.K. Bergsma, J., Netherlands Energy Research Foundation, Petten (NH)
The Netherlands Bowman, R.C.,Jr., Monsanto Research Corp., Miamisburg, Ohio, U.S.A. Bronger, W., Inst. für Anorganische Chemie, Achen, West-Germany Buchner, H., Daimler-Benz AG, Stuttgart, West-Germany Busch, G.A., Lab. of Solid State Physics ETH, Zürich, Switzerland Buschow, K.H.J., Philips Research Labs., Eindhoven, The Netherlands Cannon, J.G., Molycorp Inc., White Plains, N.Y., USA Darriet, B., Lab. de Chimie du Solide, Université de Bordeaux, France Davidov, D., Racah Inst. for Physics, Hebrew University, Jerusalem,
Israel Didisheim, J.J., Lab. de Cristallographie, Université de Geneve,
Switzerland Douglass, D.L., Boelter Hall, University of Cal., Los Angeles, U.S.A. van Essen, R.M., Philips Research Labs., Eindhoven, The Netherlands Flanagan, T.B., Dept. of Chemistry, University of Vermont, Burlington,
Vermont, U.S.A. Furrer, A., Inst. für Reaktortechnik ETHZ, Würenlingen, Switzerland Gelatt, CD., Pierce Hall, Harward University, Cambridge, Mass., U.S.A Halstead, T.K., Dept. of Chemistry, University of York, U.K. Harris, I.R., Dept. of Phys. Metallurgy and Science of Materials,
University of Birmingham, U.K. Hempelmann, R., Westfälische Wilhelms-Universität, Inst, für Phys.
Chemie, Münster, West-Germany Kleppa, O.J., The James Franck Institute, Chicago, 111., U.S.A. Korn, C , Dept. of Physics, Ben Gurion University, Beer Sheva, Israel Libowitz, G.G., Allied Chemical Corp., Morristown, N.J., U.S.A. Lewis, D., AB Atomenergi, Nyköping, Sweden Lewis, F.A., Chemistry Dept., Queens University, Belfast, Northern
Ireland, U.K. Lundin, C , Denver Research Institute, University of Denver, Col., U.S Maeland, A.J., Allied Chemical Corp., Morristown, N.J., U.S.A. van Mal, H.H., N.V. Philips1 Gloeilampenfabrieken, Eindhoven, The
Netherlands Meier, M., Inst. für Anorganische Chemie der TH, Achen, West-Germany Mintz, M.H., Dept. of Nuclear Engineering, Ben Gurion University,
Beer Sheva, Israel Müller, P., Inst. für Anorganische Chemie der TH, Achen, West-Germany Northrup,0.J.M., Jr., Chemical Technology Div., Sandia Labs.,
Albuquerque, N.M., U.S.A. Otnes, K., Institutt for atomenergi, Kjeller, Norway Pedersen, B., Dept. of Chemistry, University of Oslo, Norway Pernestâl, K., Inst. of Physics, University of Uppsala, Sweden de Pous, 0., Batelle Institute, Carouge, Switzerland Radelaar, S., Inst. of Physics, University of Utrecht, The Netherlands Rebiere, J., C.E.N.G., Lab. A.S.P., Grenoble, France Reilly, J.J., Dept. of Applied Science, Brookhaven National Lab.,
Upton, L.I., N.Y., U.S.A. XV
xvi List of Participants
van Rijswick, M. , Philips1 Research Labs., Eindhoven, The Netherlands Ron., M., Dept. of Materials Engineering, TECHNION, Haifa, Israel Sandrock, G., The International Nickel Co., Inc., Sterling Forest,
Suffern, N.Y.,U.S.A. Stohrer, H., Daimler-Benz AG, Stuttgart, West-Germany Schlapbach, L., Lab. für Festkörperforschung, ETH, Zurich, Switzerland Sheft, I., Chemistry Div., Argonne Nat. Lab., Argonne 111., U.S.A. Slotfeldt-Ellingsen, D., Central Inst. for Industrial Research, Oslo,
Norway Suda, J., Kogakuin University, Tokyo, Japan S0rensen, 0. Toft, Research Establishment Ris0, Roskilde, Denmark Venema, W., Natuurkundig Lab., Vrije Universiteit, Amsterdam, The
Netherlands Videm, K., Institutt for atomenergi, Kjeller, Norway Vigeholm, B., Research Establishment Ris0, Roskilde, Denmark von Waldkirch, Th., Eid. Technische Hochschule,Zürich, Switzerland Wallace, W.E., Dept. of Chemistry, University of Pittsburgh, Pa., U.S,A. Weaver, H.T., Org. 2354, Sandia Labs., Albuquerque, N.M., U.S.A. Wenzl, H., Inst. für Festkörperforschung, KFA, Julien, West-Germany Yamadaya, T., Matsushita Research Institute Tokyo Inc., Kawasaki, Japan
THE PROSPECTS OF HYDROGEN AS AN ENERGY CARRIER FOR THE FUTURE
George G. Libowitz Materials Research Center, Allied Chemical Corporation
Morristown, New Jersey, U.S.A. 07960
ABSTRACT
In order to evaluate the possibilities of achieving a "Hydrogen Economy", scientific problems involved in the production, storage, transmission, and utilization of hydrogen are discussed. This in­ cludes such topics as catalysis, solid state electrolysis, photo- electrolysis, thermochemical generation of hydrogen, and metal-hydro­ gen interactions. The importance of the last topic is emphasized.
INTRODUCTION The term "Hydrogen Economy" has been adopted to describe the use of hydrogen as an energy carrier. In recent years, there have been many articles published on a possible Hydrogen Economy, both in the tech­ nical literature [1] and in the popular press [2]. Therefore, a de­ tailed description of a Hydrogen Economy will not be given in this paper. However, one point, which is not always clearly presented in some of the more popular articles, should be emphasized. Namely, that hydrogen is not a primary source of energy, but rather it is a convenient and environmentally desirable way of storing, transporting, and using energy. Consequently, hydrogen must be generated from other sources of energy such as nuclear power, solar energy, etc. In order to determine the prospects of a Hydrogen Economy in the future, it is necessary to become familiar with some of the problems which must be overcome before hydrogen can be used efficiently as an energy carrier. An indication of some of the scientific problems and possible solutions are given in this paper. Since this is a sym­ posium of physical scientists, economic or political considerations related to a Hydrogen Economy are not discussed. The emphasis is on materials problems which may be associated with (1) the generation of hydrogen, (2) its utilization and (3) transmission and storage.
GENERATION OF HYDROGEN Catalysts for Production from Coal Although coal itself can be easily shipped and stored, the advantage of converting coal to hydrogen would be to obtain cleaner burning fuel. Also, hydrogen is a more convenient form of energy for some applications such as automobile fuel.
1
2 G. G. Libowitz One method of producing hydrogen from coal is by reaction with steam as shown:
Coal + H20(g) -* CO, C02, H2 (1) CO + H20(g) C02 + H2 (2)
The relative amounts of the components of the synthesis gas formed in the first reaction depend upon the type of coal used, the temperature, and other conditions of reaction. The water shift reaction (2) re­ quires catalysts in order to proceed at a sufficiently rapid rate. However, one problem is that most heterogeneous catalysts which could be used for this reaction tend to become poisoned by the sulfur in the coal. With the increased use of high sulfur coals, it will be neces­ sary to find new catalysts which, in addition to having good catalytic properties, must not be poisoned by sulfur or sulfur oxides. Various possible sulfide catalysts are being investigated including sulfo- spinels, layered transition metal sulfides, and rare earth sulfides. Water Electrolysis An established method of generating hydrogen, which should become more important with the increased availability of nuclear energy, is the electrolysis of water. This method will also be significant in the development of newer sources of energy such as solar, wind, and ocean thermal gradients. Because of problems associated with corrosion and variation in con­ centration of aqueous electrolytes, the use of solid state electro­ lytes are being explored. For the electrolysis of water, the mi­ grating species must be either hydrogen or oxygen. An example of a solid electrolyte, in which ionic transport is via hydrogen, is a perfluorinated sulfonic acid polymer developed at General Electric [3]. The behavior of this electrolyte is shown schematically in Fig. 1. Water is introduced at the anode and is de­ composed to form oxygen which is evolved, electrons which move through the external circuit, and H ions which migrate through the electrolyte as hydrated ions passing from one sulfonic acid group to the next, and finally evolving as H2 gas at the cathode. Since the sulfonic acid groups are fixed in the electrolyte, the concentration of electrolyte remains constant. Other advantages of this electro­ lyte include its ability to operate at higher pressures, the fact that it is non-corrosive, and reduced power requirements. Inorganic defect solids capable of ionic conduction such as yttria, zirconia, and thoria are also being investigated as possible solid state electrolytes. One such electrolyte system [4] (also developed at G.E.) using calcia stabilized zirconia is illustrated in Fig. 2. Some of the Zr4+ ions in the Zr02 lattice are substituted by Ca2+, and in order to maintain electroneutrality, oxygen vacancies VQ, are formed in the lattice. Water, that has been vaporized by the neat of coal oxidation (which also may be used to generate the electrical power), is introduced at the cathode and reduced to form hydrogen gas, while oxygen fills the lattice vacancies to form oxygen on normal lattice sites, 0Q. At the anode, CO reacts with the lattice oxygen to re-form the vacancies, as shown. The oxygen migrates through the electrolyte as lattice vacancies. In addition to some of the advan­ tages mentioned above for the polymer electrolyte, such cells may
The Prospects of Hydrogen as an Energy Carrier for the Future 3 operate at temperatures as high as 800-1000°C, leading to increased efficiencies. Thermochemiea1 Production It is possible to thermally decompose water by direct application of heat; however, temperatures in excess of 2500°C would be required. Although one such scheme has recently been proposed [5] using solar energy, high temperatures are difficult to obtain by the usual methods of energy production. However, water can be thermally decomposed at lower temperatures by using a series of reactions in which all the re- actants (except water) are re-generated, such that the overall result is the decomposition of water to hydrogen and oxygen. One such system, suggested by Wentorf and Hanneman [6], is illustra­ ted by the set of Eqs. (3) to (7):
3FeCl2 + 4H20 + Fe304 + 6HC1 + H2 450-750°C (3) Fe304 + 8HC1 -> FeCl2 + 2FeCl3 + 4H20 100-110°C (4) 2FeCl3 -> 2FeCl2 + Cl2 300°C (5) Cl2 + Mg(OH)2 -* MgCl2 + ^0 2 + H20 50- 90°C (6) MgCl2 + 2H20 -> Mg (OH) 2 + 2HC1 350°C (7)
It can be seen that the sum of these equations is merely: H2° "** H2 + I°2 (8)
Note that the maximum temperature required for any of these reactions is 750°C. Therefore, lower grade heat, such as that available from nuclear reactors, may be used to thermally decompose water. This method is referred to as thermochemical water splitting. Many such sets of reactions have been proposed and investigated [6]. However, there are also many problems to be solved. The relative kinetics of the reactions are important and side reactions must be avoided. These require appropriate catalysts. Methods of separating the intermediate products must be developed. Also, materials com­ patibility is important when corrosive intermediates such as HC1 and CI2 are present. Photoelectrolysis A relatively new concept for producing hydrogen from solar energy has received a considerable amount of interest recently; the electro­ chemical photolysis of water, or photoelectrolysis. Although the energy required to decompose water is 2.46eV (corresponding to light of wavelength of about 500nm), direct solar photolysis does not occur because water does not absorb light until well into the UV portion of the spectrum where the solar irradiance is weak. However, by using light absorbing semiconductor electrodes immersed in an aqueous solu­ tion, as shown schematically in Fig. 3, the normal electrochemical potential of 1.23eV is required to dissociate water. This corres­ ponds to about lOOOnm; therefore, the visible range of the spectrum can be used.
4 G. G. Libowitz The cell in Fig. 3 may be viewed as a semiconductor p-n junction, separated by an electrolyte, so that band bending occurs near the semiconductor-electrolyte interface as shown. If the semiconductors are irradiated with light whose wavelength is such that hv> band gap, electron-hole pairs will be formed in each semiconductor electrode. Excess electrons will flow from the p-type semiconductor (cathode)in­ to the semiconductor-electrolyte interface to reduce the H+ ions in the electrolyte according to the reaction (in acidic electrolyte):
2H+ + 2e" > H2 (9) Similarly, holes, h , from the n-type semiconductor electrode (anode) will oxidize the water as follows:
H20 + 2h+ + 2H+ + ^0 2 (10) The two electrodes are, of course, connected through an external cir­ cuit to permit current flow. If the electrolyte is alkaline, then the reactions corresponding to Eqs. (9) and (10) are
2H20 + 2e~ -> H2 + 20H~ (11) and
20H~ + 2h+ + io2 + H20 (12)
The sum of Eqs. (9) and (10) or of Eqs. (11) and (12) correspond to the decomposition of water [Eq. (8)]. The concept of photoelectrolysis was first proposed and partially demonstrated by Fujishima and Honda [7] using T1O2 as the n-type semiconductor anode and platinum metal as the cathode. This type of cell, with only one semiconductor electrode, has been referred to as a Schottky barrier analogue cell [8], and is illustrated in Fig. 4. In the Schottky-type cell, the band gap of the semiconductor must be greater than 1.23eV in order that the excited electrons have suffi­ cient energy to decompose water. However, as can be seen in Fig. 5, semiconductors with band gaps greater than 2.5eV will absorb only a relatively small portion of the solar spectrum. For example, T1O2 which has a band gap of 3eV absorbs only about 8 to 10% of the solar spectrum. Therefore, for maximum efficiency, the band gap of the semiconductor electrode in a Schottky-type cell should be higher than 1.3eV (additional energy is needed to overcome irreversible losses in the cell) and less than 2.5eV in order to absorb a sufficient porticn of the solar spectrum. In a p-n cell, the total energy ideally available for photoelectroly­ sis is the sum of the band gaps of the n-type and p-type semiconduc­ tors (if two different materials are used) [8]. Therefore, the band gap of each semiconductor may be less than leV, which means a larger percentage of the solar spectrum could be absorbed with correspond­ ingly greater efficiencies of operation. However, there are other requirements of a semiconductor electrode. First, the semiconductors must be electrochemically stable. This is particularly important for n-type semiconductors which tend to become
The Prospects of Hydrogen as an Energy Carrier for the Future 5 oxidized when acting as an anode. For example, CdS will oxidize to Cd2+ ions in solution and free sulfur [9] and GaP will oxidize to Ga+3 ions and phosphoric acid [10]. Secondly, the positions of the energy levels in the semiconductors relative to the redox levels in the electrolyte are also important. For example, the bottom of the conduction band, Ec, in the p-type semiconductor must be at a higher energy than the H+/H2 redox level (see Fig. 3) so that the photo-excited electrons do not have to over­ come an energy barrier in order to reduce the H+ ions [Eq. (9)]. Similarly, the top of the valence band, Ev, in the n-type semiconduc­tor should be below the OH~/C>2 redox level (since holes flow up) . In a Schottky-type cell, a bias voltage can be used to overcome the mis­ match of energy levels [11]. However, the energy difference between Ec and the H+/H2 redox level in the p-type semiconductor (or between Ev and the OH"/02 level for the n-type semiconductor) must not be too large because this energy difference is not available for dissoci­ ation of water, and therefore the efficiency of the cell is decreased [12].
Finally, the relative positions of the flat band potentials (positions of the original Fermi levels before the semiconductor equilibrates with the electrolyte) in the two semiconductor electrodes should not differ too much because this would lead to a large degree of band bending and a corresponding loss of energy [12]; i.e. the energy of the electron at the electrolyte interface would be much less than its energy in the bulk of the semiconductor. Thus, it can be seen that the requirements of semiconductors for this application are rather stringent and there is need for much further research in order to find appropriate materials [13]. UTILIZATION One major advantage of a Hydrogen Economy is the ability to conve­ niently store electricity. Excess electricity may be used to elec- trolyze water and the hydrogen thus formed is stored. The hydrogen may then be transformed back to electricity via fuel cells. In order that this concept be economically feasible, the efficiency of pres­ ently available fuel cells must be improved. To some degree a hydrogen fuel cell may be viewed as the opposite of an electrolytic cell; instead of electrolyzing water, H2 and O2 are re-combined to generate electricity. Possible new electrolytes for such a cell were discussed under the section "Water Electrolysis" above. However, in developing new fuel cells it is also necessary to find new electrode materials and electrocatalysts. An electro- catalyst is a substance which activates the reacting molecules (H2 and O2 in this case) such that electron transfer will occur rapidly at the electrode-electrolyte interface. The catalyst can be incor­ porated into the electrode, or in some cases, the electrode material itself may act as a catalyst. Other requirements of electrode materials are that they have high electronic conductivities and yet be corrosion resistant. These re­ quirements are frequently mutually exclusive. Oxide layers will usually protect a metal from corrosion, but it will also decrease the conductivity of the material. Types of materials under investigation
6 G. G. Libowitz are carbides such as WC [14], conducting spinels such as NÍC02O4 [15] and heavily doped oxide semiconductors such as Li-doped nickel oxide [16]. Some of the new metallic conducting polymers [e.g. polythiazyl, (SN) ]are also being considered as possible electrode materials [17]. A significant advantage in using hydrogen as a fuel is its versatility; besides direct combustion, and conversion to electricity via fuel cells, hydrogen can be catalytically oxidized at relatively low tem­ peratures. The advantages of this method of utilizing hydrogen in­ clude safety, since there is no open flame, and no formation of ox­ ides of nitrogen. Therefore, catalytic oxidation would be desirable for home heating and in appliances such as space heaters and camp food warmers. One problem in using this method however, is the limi­ ted life of available catalysts. Therefore, new catalysts for this application also must be developed.
TRANSMISSION AND STORAGE Hydrogen Embrittlement Proponents of a Hydrogen Economy have suggested that existing natural gas pipelines may be used to transport hydrogen gas. It has been estimated [18] that, over long distances, the cost of transmitting hydrogen by pipeline will be almost an order of magnitude less costly than transmitting the same amount of energy by electricity. However, in using this method of transporting hydrogen, the problem of hydro­ gen embrittlement must be considered. There are three general types of hydrogen embrittlement of metals [19], (1) hydrogen reaction, (2) internal, and (3) hydrogen environ­ ment embrittlement. Hydrogen reaction embrittlement is due to the reaction of hydrogen to form internal phases. For example, in hy­ dride forming metals, the formation of hydrides which have volumes 15 to 25% greater than the corresponding metal, will cause stresses and tend to crack the metal. In carbon steels, the hydrogen may react with the carbon to form methane gas which can cause cracking or blistering. In internal hydrogen embrittlement, hydrogen, which is formed from water during melting, casting, pickling, welding, plating or by corrosion, becomes dissolved in the metal. The hydrogen then concentrates at the tips of existing cracks in the metal and tends to propagate the crack through the metal. The first two types of hydrogen embrittlement may be avoided by elim­ inating the conditions which cause the embrittlement. For example, in the case of carbon steels, the thermodynamic activity of carbon may be reduced by adding molybdenium so that the carbon no longer re­ acts with hydrogen. The third type, hydrogen environment embrittlement, is more difficult to control because its nature is not yet fully understood. In this case, the metal degrades only when in the presence of hydrogen. It is a temperature dependent process with maximum embrittlement usually occurring at room temperature. Small amounts of oxygen impurity in the hydrogen gas will usually inhibit embrittlement, and this is also frequently true for SO- and 002 impurities. One possible mechanism for hydrogen-environment embrittlement is
The Prospects of Hydrogen as an Energy Carrier for the Future 7 based upon the strong interaction between hydrogen and transition metals. Gilman [20] has suggested that the strong surface adsorption of hydrogen, particularly near crack tips in the metal, will suppress plastic deformation by increasing the energy necessary to create the surface shear step. Thus the tendency towards cleavage will be en­ hanced, with resulting embrittlement. However, other mechanisms have been proposed and there is need for a great deal of further research on the nature of hydrogen embrittlement [21].
Hydrogen Storage
Hydrogen may be stored as a gas, as a liquid, or in easily dissoci­ ated compounds such as metal hydrides, which is the major topic of this symposium. Storing hydrogen as a gas requires large volumes. Even under compression the volume storage efficiency of gaseous stor­ age is not as high as liquid hydrogen, and the weight of the storage cylinder becomes a major disadvantage. Although the volume effi­ ciency is improved when hydrogen is liquefied, the energy required for liquefaction and the need for well insulated containers are dis­ advantages. Also, when storing for long periods of time there is still considerable loss of hydrogen due to evaporation.
Storing hydrogen as a metal hydride has several advantages. First, with respect to volume, hydrogen can be stored more efficiently than in liquid, or even solid, hydrogen as illustrated in Table 1, which
TABLE 1 Hydrogen Densities in Some Hydrogen-Containing Compounds
3 -22 Compound Number of H atoms/cm xlO
Liquid Solid Water LiH TiH2
ZrH2
YH2 UH3
hydrogen hydrogen
(20< (4.2"
°K) °K)
4.2 5.3 6.7 5.9 9.2 7.3 5.1 5.4 6.4 5.7 8.2
shows the number density of hydrogen atoms in some representative hy­ drides. In every case, the number of hydrogen atoms per cm^ is greater than that of liquid, or even solid, hydrogen; and in T1H2, the number density is more than double that in liquid hydrogen. How­ ever, it can be seen that water also has a relatively high hydrogen density. This points up the second major advantage of metal hydrides, the ease of reversibility of the formation reaction:
8 G. G. Libowitz
M + | H 2 X ΜΗχ (13) The formation of the hydride is an exothermic and usually spontaneous reaction, but the hydrogen can be easily recovered by heating the hy­ dride. The use of metal hydrides is an unusually safe method of storing hy­ drogen because hydrides are generally quite stable below their disso­ ciation temperatures. Also, since the reverse of Eq. (13) is an endo- thermic reaction, the self-cooling effect will suppress any loss of hydrogen if a leak develops in the storage system. This method of storing hydrogen requires no thick-walled containers or heavy insu­ lation, and the possibility of explosion due to high pressures is lessened. The properties required of an efficient metal hydride storage medium are summarized in Table 2. High hydrogen retentive capacity corre-
TABLE 2 Desired Properties of a Hydrogen Storage Material
High hydrogen retentive capacity Low temperature of dissociation (!100°C) High rates of hydrogen uptake and discharge Low heats of formation Low cost of alloy Light weight Stable towards oxygen and moisture
sponds to hydrides with high hydrogen-to-metal (H/M) ratios. Low dis­ sociation temperatures are necessary so that the hydrogen will be easily recoverable when needed. Low heats of formation are desir­ able to minimize energy requirements when recovering the hydrogen, and also because there will be less heat to dissipate during formation of the hydride. Light weights are desirable for applications in which the fuel is portable, such as hydrogen-powered vehicles. None of the known binary hydrides meet all, or even most, of these requirements. Therefore, it is necessary to develop new alloy hy­ drides which will have the desired properties listed in Table 2. A knowledge of the fundamental properties of metal hydrides, in general, would be of value in designing new hydride system. Such properties have been reviewed in the past [22], and updated reviews of the fundamental properties are presented in following papers by Maeland, Wallace, Flanagan, and Andresen, among others. There are two general approaches which can be taken in the develop­ ment of new alloy hydrides. One is modification of the properties of known hydrides by appropriate alloying or variation of the compo­ sitions of intermetallic compounds. This approach is described in the papers by Douglass in the case of magnesium hydride, and Machida et al and Davidov £t al for intermetallic compound hydrides. The Rule of Reversed StabTTity, which states that for a given series of intermetallic compounds, the thermodynamic stabilities of the
The Prospects of Hydrogen as an Energy Carrier for the Future 9 corresponding hydrides will decrease with increasing stability of the intermetallic compound, can be of value in this latter approach. The rule was proposed by VanMal et al [23] and is discussed in following papers by Buschow and Miedema, Gelatt, and Davidov et al. The second approach to developing new hydrides for hydrogen storage is to synthesize new intermetallic compounds capable of forming hy­ drides with appropriate properties. This approach has led to several promising systems such as FeTi hydride developed at Brookhaven [24] and the rare earth-transition metal compounds discovered at Philips- Eindhoven [25] . In general, the properties of intermetallic compound hydrides appear to have little, or no, resemblance to those of the constituent metal hydrides. For example, Table 3 shows some typical intermetallic com-
TABLE 3 Intermetallic-Compound Hydrides Intermetallic Compound
Hydride LaNi5H6.7 DyCo3H5 ZrNiH3 ThCoH4
Constituent Hydride LaH3 PrH3 ZrH2 Th4H15
Ref. [25] [26] [27] [28]
pound hydrides which take up more hydrogen than would be expected on the basis of the constituent metal hydrides. In Table 4, the proper-
TABLE 4 Comparison of ZrNiH3 With ZrH2
ZrH2 ZrNiH3
Structure Tetrag. Orthorhombic (distorted fluorite) -9 Dissoc. Press, at 250°C 3x10 Torr 200 Torr
o o Zr-H distance 2.09A 1.96A
o o Closest H-H distance 2.22A 2.04A
ties of ZrNiH3 are compared to those of ZrH2 in more detail. It can be seen that the crystal structures are different and that although the intermetallic compound hydride is less stable (dissociation pres­ sure is higher by a factor of 10 1 1), the Zr-H and H-H distances are smaller in that compound [29] . It may be convenient to consider in­ termetallic compound hydrides as pseudo-binary hydrides. Since there is a very large number of possible intermetallic com­ pounds and an infinite number of compositional variations, it would be desirable to have some way of predicting which intermetallic com­ pounds will react with hydrogen to form hydrides having the proper­ ties required of a good hydrogen storage medium. The Rule of Re­ versed Stability could have some degree of success in this respect [30], but at present, it appears to be of greater value in the first approach; i.e. in predicting the effect of alloying elements on the
10 G. G. Libowitz thermodynamic stability of known hydrides [23].
Certainly, the relationship between the electronic band structure of an intermetallic compound and its behavior with hydrogen is important. Therefore, a better understanding of the electronic structures of in­ termetallic compounds and how they are modified by interaction with hydrogen would be of value in predicting new intermetallic compound hydrides. There are many papers at the symposium which cover that aspect, including those by Wallace, Pedersen, Korn, Griessen et al, and Gelatt.
The importance of electronic structure relative to crystal structure can be studied by investigating the hydrogen uptake of a metallic glass (sometimes called amorphous) alloy, whose composition is iden­ tical to that of a known intermetallic compound. Such studies on Ti-Cu alloys are reported by Maeland in a following paper.
CONCLUSION
The scientific problems discussed in this paper are an indication of the technical difficulties which must be overcome before hydrogen may be efficiently utilized as an energy carrier. Nevertheless, I be­ lieve that there will be a Hydrogen Economy in the future. However, it will be attained gradually over a period of time, and probably not all aspects of the Hydrogen Economy will be achieved. Fleet vehicu­ lar systems (such as busses) look promising, but the use of hydrogen in private autos appears unlikely in the near future. Off-peak power storage is another promising possibility. Also, as the newer inter­ mittent sources of energy such as solar and wind are developed, the use of hydrogen for energy storage will become more attractive.
However, it is obvious that there is need for a great deal of further research before the Hydrogen Economy becomes a reality.
REFERENCES
(1) For example: D. P. Gregory, Sei. Am. 228, 13 (January 1973); W. E. Winsche, K. C. Hoffman, F. J. Salzano, Science 180, 1325 (1973); C. E. Bamberger and J. Braunstein, Am. Sei. 63, 438 (1975) / In addition the International Journal of Hydrogen Energy pro­ vides articles concerned with various aspects of the Hydrogen Economy in more detail.
(2) For example: "The Coming Hydrogen Economy" Fortune, November 1972 and "Here Comes the Hydrogen Era" Readers Digest, December 1973.
(3) L. J. Nuttall, A. P. Fickett, and W. A. Titterington, Proc. Hydrogen Economy Miami Energy Conf., T. N. Veziroglu, Ed. pp. S9-33 to S9-37, Univ. of Miami, Coral Gables, Fla. (1974).
(4) W. W. Aker, D. H. Broun, H. S. Spacil, and D. W. White, U.S. Patent No. 3,616,334, Oct. 26, 1971.
(5) E. A. Fletcher and R. L. Moen, Science 197, 1050 (1977).
The Prospects of Hydrogen as an Energy Carrier for the Future 11
(6) R. H. Wentorf and R. E. Hanneman, S c i e n c e 1 8 5 , 311 ( 1 9 7 4 ) .
(7) A. F u j i s h i m a and K. Honda, N a t u r e 2 3 8 , 37 ( 1 9 7 2 ) .
(8) A. J . N o z i k , A p p l . P h y s . L e t t . 2 9 , 150 ( 1 9 7 6 ) .
(9) R. W i l l i a m s , J . Chem. P h y s . 3 2 , 1505 ( 1 9 6 0 ) .
(10) A. J . N o z i k , P r o c . 1 s t World Hydrogen Energy C o n f e r e n c e , V o l . I I , U n i v . o f Miami, C o r a l G a b l e s , F l a . p p . 5 B - 3 1 t o 5B-34 ( 1 9 7 6 ) .
(11) T. O h n i s h i , Y. N a k a t o , and H. Tsubomura, B e r . B u n s e n g e s . P h y s i k . Chem. 7 9 , 523 ( 1 9 7 5 ) .
(12) A. J. Nozik, Proc. Conf. on the Electrochemistry and Physics of Semiconductor Liquid Interfaces Under Illumination, A. Heller, Ed., The Electrochemical Soc. Inc., Proceedings Vol. 77-3, Princeton, N.J., pp. 272-289 (1977).
(13) A. J. Nozik, J. Cryst. Growth 39, 200 (1977). (14) H. Bonn, Electrochim. Acta 15, 1273 (1970). (15) W. J. King and A. C. C. Tseung, Electrochim Acta 19, 485 (1974). (16) H. L. Bevan and A. C. C. Tseung, Electrochim. Acta 19, 201
(1974) . (17) R. J. Nowak, H. B. Mark, A. G. MacDiarmid, and D. Weber, J.
Chem. Soc, Chem. Commun. (1977) 9. (18) W. E. Winsche, K. C. Hoffman, and F. J. Salzano, Science 180,
1325 (1973). (19) W. T. Chandler and R. J. Walter, Proc. Hydrogen Economy Miami
Energy Conf., T. N. Veziroglu, Ed., Univ. of Miami, Coral Gables, Fla., (1974) pp. S6-15 to S6-31.
(20) J. J. Gilman, Phil. Mag. 26, 801 (1972). (21) Effect of Hydrogen on Behavior of Materials, A. W. Thompson and
I. M. Bernstein, Eds., Metallurgical Soc. of AIME, (1976). (22) G. G. Libowitz, The Solid State Chemistry of Binary Metal
Hydrides, W. A. Benjamin Inc., New York (1965); W. M. Mueller, J. P. Blackledge, and G. G. Libowitz, Metal Hydrides, Academic Press, New York (1968); G. G. Libowitz, MTP Internatl. Rev. Sei., Inorg. Chem. Ser. 1, Vol. 10, Solid State Chemistry, L. E. J. Roberts, Ed., Butterworths Ltd., London (1972) pp.79- 116.
(23) H. H. Van Mai, K. H. J. Buschow, and A. R. Miedema, J. Less Common Metals 35, 65 (1974).
(24) J. J. Reilly and R. H. Wiswall, Inorg. Chem. 13, 218 (1974). (25) J. H. N. van Vucht, F. A. Kuijpers, and H. C. A. M. Bruning,
Philips Res. Repts. 25, 133 (1970).
12 G. G. Libowitz
(26) T. T a k e s h i t a , W. E . W a l l a c e , and R. S . C r a i g , I n o r g . Chem. 1 3 , 2283 ( 1 9 7 4 ) .
(27) G. G. Libowitz, H. F. Hayes, and T. R. P. Gibb, J. Phys. Chem. 62, 76 (1958).
(28) W. L. Korst, U.S.A.E.C. Report No. NAA-SR-6881 (1962). (29) S. W. Peterson, V. N. Sodana, and W. L. Korst, J. Phys. (Paris)
25, 451 (1964). (30) K. H. J. Buschow, H. H. Van Mal and A. R. Miedema, J. Less
Common Metals 42, 163 (1975).
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Arnulf J. Maeland Materials Research Center, Allied Chemical Corporation
Morristown, New Jersey, U.S.A. 07960
ABSTRACT
Binary hydrides, conveniently classified according to bonding as saline, metallic and covalent are reviewed and surveyed with re­ spect to structure and physical properties. Hydrides of inter- metallic compounds, which are of major interest at this meeting, may be considered to be pseudo-binary hydrides and are included in the survey.
INTRODUCTION Hydrogen with its unique electronic structure of one electron in a Is orbital forms compounds with most of the elements in the periodic table. Compounds in which there is a metal-hydrogen or metalloid-hydrogen bond are collectively referred to as hydrides. Based on the nature of the metal-or metalloid-hydrogen bond and the resulting physical properties, the hydrides may be classified in three major categories: (1) Saline or ionic hydrides, (2) metallic hydrides, and (3) covalent hydrides. Saline hydrides have typically, high enthalpies of formation, high melting points, and are electrically conducting in the mol­ ten state. The saline hydrides include the binary hydrides of the alkali and alkaline earth (except beryllium) metals. The physical properties of the alkali and alkaline earth hydrides are in many respects similar to the corresponding halides. The similarity extends to the crystal structure as well, particularly in the alkali series. The alkali hydrides have the sodium chlor­ ide structure, while the alkaline earth hydrides (except MgH~) have an orthorhombic structure which is related to the structure of the barium halides. The crystal lattices of the saline hy­ drides consist basically of hydrogen anions and metal cations. This description is not to be construed as exclusive. In lithium hydride, for example, theoretical calculations [1] and diffrac­ tion experiments [2] suggest that the electron transfer from lithium to hydrogen is between 0.8 and 1 electron. This implies a strong ionic bond, but with some covalent character. Magnesium hydride occupies a special position. Although classified here as a saline hydride, its physical properties are intermediate be­ tween the ionic hydrides and covalent beryllium hydride. MgH2 may thus be regarded as a transition hydride between the saline
19
20 A. J. Maeland
and covalent hydrides. The dihydrides of europium and ytterbium are isostructural with the alkaline earth hydrides, and may also be regarded as saline hydrides. Ternary hydrides such as LiBaH^, LiSrH^, and LiEuHU are basically saline. Metallic hydrides have, as the name implies, metallic properties such as luster, hardness, metallic conductivity (except the higher hydrides of the rare earths), but unlike metals they are quite brittle. Another characteristic of metallic hydrides is their deviation from stoichiometry which in many cases is unusu­ ally large. Hydrides of those transition metals which form bi­ nary compounds with hydrogen (Groups IIIA through VIIIA) are classified as metallic hydrides. This includes the rare earth hydrides (except Eu and Yb) and the actinide hydrides. Many of the intermetallic compound hydrides which are discussed at this meeting, e.g. TiFel^/ LaNißHg and related compounds, have proper­ ties which suggest that they be classified as metallic hydrides. For convenience these hydrides may be regarded as pseudo-binary hydrides.
The nature of the chemical bonding in the metallic hydrides has been the subject of much controversy[3-7]. Two opposing models have been proposed: the protonic and the anionic. In the pro- tonic model[8] hydrogen is assumed to donate its electron to the d-band of the transition metal forming essentially an alloy with the metal. Hydrogen may thus be considered to exist as protons, partially screened by the conduction electrons, in the metal sub- lattice. The opposing view[9] asserts that hydrogen accept elec­ trons from the metal to form hydride anions and metal cations, i.e. a saline hydride. Major support for the protonic model has come from the fact that most metallic hydrides are metallic con­ ductors. Libowitz has pointed out, however, that the trihydrides of the rare earths become semiconductors and their electronic properties are more readily explained by the ionic model[10]. The relatively large enthalpies of formation of most metallic hydrides (they are comparable to the enthalpies of formation of the saline hydrides in many cases) appear to favor the anion model. Experi­ mental results from Mossbauer spectroscopy, positron annihilation magnetic susceptibility, and nuclear magnetic resonance as well as theoretical calculations have not been conclusive, but have been interpreted in favor of one or the other of these two models[7]. Recent energy band calculations by Switendick[ll] offers a solu­ tion to the dilemna. Switendick proposes a model which quali­ tatively may be described as follows: In the metallic hydride structure mixing and hybridization takes place between the Is orbitals of hydrogen and the metal band states. In monohydrides with octahedral hydrogen, e.g. palladium hydride, the sp metal bands mix with the hydrogen bonding orbitals to form a modified band, lowered in energy. The energy states of this band are to a large extent already filled in the metal (below the Fermi energy, Ep) and the added electrons from hydrogen will go into the empty metallic states, above EF, and thus appear to donate its electron to the metal, i.e. the proton model. The stability of the hydride is determined by the extent to which the states
Survey of the Different Types of Hydrides 21 in the modified band are empty, and by the amount the energy of the modified band has been lowered. In dihydrides and trihy- drides new low-lying states, associated with the hydrogen atoms, are formed and the additional electrons occupy these states, i.e. the ionic model. The energies of the new bands are dependent upon the hydrogen-hydrogen distance which in turn is determined by the type of sites occupied by hydrogen in the metal sub-lat­ tice and the size of the metal atom. The relative positions of the various energy bands determine the formation and stability of these hydrides. The problem of bonding in metal hydrides is discussed more fully by Professor Wallace in the next paper. Most covalent hydrides have low melting and boiling points, and are, in fact, liquid or gaseous at room temperature; those that are solid are thermally unstable. This is of course a reflection of the weak van der Waals forces existing between covalent mole­ cules. Those metals and metalloids of Groups IB through VB which form binary hydrides belong to this category. Also included is beryllium hydride. The structure of many covalent hydrides is believed to be polymeric. Since the covalent hydrides are out­ side the area of interest at this meeting, they will not be dis­ cussed any further.
STRUCTURES OF THE BINARY METAL HYDRIDES The room temperature metal sub-lattice structures of the binary saline hydrides are listed in Table 1; those of the metallic hy­ drides are listed in Table 2. The structures of the metals are also indicated. It may be noted that in most cases the metal sub-lattice of the hydride is different from that of the metal from which it is formed. This supports the idea that the metal hydrides are chemical compounds rather than interstitial solid solutions, since the latter implies essentially no change in metal structure on forming the hydride. In the few cases where the metal sub-lattice remains unchanged, (palladium, cerium and
TABLE 1 Structures of the Binary Saline Hydrides Metal Hydride
Metal Structure Hydride Structure Alkali Li-Cs Mg Ca,Sr Ba Eu Yb
metals, b.c.c. h.c.p. f.c.c. b.c.c. b.c.c. f.c.c.
LiH-CsH MgH2 CaH<2, SrHo BaH2 EuH2 YbH2
YbH3-x
Source: Ref.[5]
Metal Ti,Zr V
f . c e (a=5
.89A)
b.c.t. orthorhomb:
Hydride TiH 2,ZrH 2,HfH 2
VH V H 2
LaH2,PrH2,NdH2 YH2,Gd-TmH2,LuH2 YH3,GdH3-TmH3,LuH3
CeH2 SmH2 SmH3 ACH2 T h H2 Th4H15 PaH3 a-UH3 3-UH3 NPH2 NPH3
CmH2,BkH2 PUH2 PuH3 AmH2 AmH^
Hydride Structure
f . c c , f .c t. b.c.t. f . c c orthorhombic f . c c orthorhombic f . c c (a=4.03A) f . c c f . c c f . c c hexagonal f . c c (a=5.58A) f . c c hexagonal f . c c (a=5.67A) f.c.t. b . c c cubic (8-W) b.c.c. cubic (3-W f . c c hexagonal f . c c f . c c hexagonal f . c c (a=5.35A) hexagonal
Source: CmH. ref.[12], BkH2 ref.[13], AmH2 and AmH3 ref.[14] All others, ref.[5]
Survey of the Different Types of Hydrides 23 actinum), there is a large discontinuous increase in the lattice parameter as the hydride forms from the metal. The metal sub- lattice in the rare earth trihydrides YH^, GdHL· through TmH^ and LuH3 is also the same as the metal phase (hexagonal), but in these systems there is an intermediate dihydride phase which has a different metal sub-lattice (f.c.c). In addition, the volume of the hexagonal metal hydride phase is larger (14-25%) than that of the corresponding metal.
PRESSURE-COMPOSITION ISOTHERMS The saline and the metallic hydrides are generally prepared by reacting the metal or alloy with hydrogen gas at elevated tem­ peratures. The overall, reversible, exothermic reaction may be written:
M + | H2 Z MHs (1)
where s=l for a monohydride, 2 for a dihydride, etc. The progress of reaction 1 can be followed by measuring the equili­ brium hydrogen pressures as a function of hydrogen content in the metal and thus pressure-composition isotherms, such as the one shown in Fig. 1, are determined. Hydrogen first dissolves in the metal according to equation 2
M + Σ H2 î MH (solution phase) (2)
to form a solid solution whose composition depends on the hydro­ gen pressure. The solid solution region is represented by the steeply, rising portion on the left hand side of the isotherm in Fig. 1. When the solid solution becomes saturated with hydrogen, the nonstoichiometric hydride, ΜΗχ, begins to form. With further addition of hydrogen more of the saturated solid solution is con­ verted to hydride while the pressure remains invariant, as re­ quired by the Phase Rule, across the two-phase region indicated by the horizontal portion of the isotherm in Fig. 1. This in­ variant plateau pressure is the equilibrium dissociation pressure of the hydride at the temperature of the isotherm. Equation 2 represents the reaction taking place in the plateau region:
MHy + *ZX H2 Î ΜΗχ (3) After complete conversion to the hydride phase, the hydrogen pressure increases again (right hand side of Fig. 1) as the non­ stoichiometric hydride absorbs hydrogen according to equation 4.
ΜΗχ + s- x- H2 ?MHg (solution phase) (4)
If a second hydride phase forms, another plateau region follows as seen in Fig. 2. The extension to systems in which there are more than two hydride phases, is obvious.
24 A. J. Maeland THERMODYNAMIC PROPERTIES
If the solubility of hydrogen in the metal phase is negligible (y-0) and the deviation from stoichiometry is small (x^s), then the standard enthalpy of formation, AHf, of a hydride, MH , can be calculated from the van't Hoff equation,
9£n K ΔΗ^ -£ « - ! (5) 9T RT
where K p is the equilibrium constant for reaction 1, T is the ab­ solute temperature and R is the gas constant. K p is given by equation 6 :
which becomes
K = P„ ~s / 2 (atm.) (7) P Ho
when the standard states of hydride and of metal are taken as the pure solids in each case (a..„ = a M = 1) and a„ = P„ (atm) . If
MH M no nn — s/2 P„ is substituted for K in equation 5 and the equation is H2 P
integrated (assuming AHf is constant over a reasonably large tem­ perature range) equation 8 is obtained
£nP„ (atm) = 2 AHf ,R. H2 "FRT + C (8)
where P represents the plateau pressure at a particular temper-H 2 ature, T, (Fig. 1) and C is the constant of integration. The en­ thalpy of formation of the hydride is calculated from the slope of the straight line obtained by plotting £nP„ versus 1/T (Fig. 3) . H 2 The equilibrium dissociation pressure at a given temperature, allows one to calculate the standard free energy of formation, AGf, v , of the hydride at that temperature,
AGf(T) = "RT£nK p = f RT £nPH (atm) <9>
and the standard entropy of formation can then be evaluated from
= A H f " A G f (10) T
Comparison of equations 8, 9 and 10 shows that the constant in­ tegration in equation 8 may be identified with 2 ASf
Survey of the Different Types of Hydrides 25 Most metal-hydrogen systems exhibit both appreciable solubility of hydrogen in the metal phase and significant deviation from stoichiometry in the hydride phase. The value of ΔΗ, obtained from a plot of £nP__ versus 1/T, is, therefore, not the standard H2 enthalpy of formation of the stoichiometric hydride, MHg, but is generally assumed to be the enthalpy of reaction 3. The situ­ ation is further complicated, however, by the fact that the solu­ bility of hydrogen in the metal phase and the deviation from stoichiometry in the hydride phase change with temperature. Since the relative partial enthalpies of solution of hydrogen in the metal phase and in the hydride phase both vary with composition, it is indeed remarkable that £nP„ versus 1/T is linear (dis-H2 cussed more fully in a later paper by Professor Flanagan). It is nevertheless observed experimentally, that most metal hydrides obey a relation of the form
£nPR (atm.) = - | + B (11)
over fairly large temperature ranges, and it is customary to as­ sociate the enthalpy determined from such relationships with re­ action 3. The enthalpy of formation of the hydride, MH (formed according to equations 2-4 and illustrated in Fig. 1), is accu­ rately given by summing the heat of solution of hydrogen in the metal (equation 2), the heat of reaction 3, and the heat of solu­ tion of hydrogen in the nonstoichiometric hydride (equation 4). Dissociation pressure data have been used extensively in the past to evaluate and tabulate the thermodynamic properties of hy­ drides [4,5,15]. Calorimetric measurements, however, have not been made on most metal-hydrogen systems and relatively few com­ parisons can, therefore, be made. For the alkali hydrides the enthalpies of formation have been determined by measuring the heats of reaction of the hydride and the corresponding metal with water or dilute acids:
M + H20 + MOH + ~H2; ΔΚ^ (12)
MH + H20 -> MOH + H2; ΔΗ2 (13) The enthalpy of formation of the hydride, ΔΗ^, is given by ΔΗ..-ΔΗ2. Calorimetric enthalpies of formation have also been ob­ tained by measuring the heat of combustion of MH :
MHs(s) + (2n4"5)Q2(g) + MOn(s) + |H20(g) (14)
The thermodynamic parameters for oxygen, water and the metal ox­ ides are known and the values for the hydrides can, therefore, be calculated. The enthalpies of formation of TiH« and MgH2, for example, have been obtained by this procedure. Table 3 lists the enthalpies of formation of a number of saline and metallic hy­ drides for which data from both dissociation pressure measure­ ments and calorimetric measurements are available. The agreement between the values obtained by the two methods is quite good, es­ pecially for the alkali hydrides and uranium hydride. The solu-
26 A. J. Maeland TABLE 3 Enthalpies of Formation Determined From Dis­
sociation Pressure Measurements Compared With Calorimetric Values
T 25° -AHf (diss.Pressure) -AHf (calorimetric) Ref.
Hydride KJ/mole H2
LiH 183.8(500°-650°) NaH 116.6(250°-415°) KH 118.1(288°-415°) RbH 108.6(246°-350°) CsH 112.8(245°-378°) MgH2 74.43(440°-560°) CaH2 184.0(600°-780°) SrH2 199.l(to 1000°) BaH2 175.2(470°-550°) TiH2 133 (450°) Zrl^ 5 188(500°-553°) UH3 85.4 av. (260°-650°)
Ref. 16 18 18 18 18 19 21 23 25 26 28 4
KJ/mole H2
181.28 112.88 115.65 104.60 108.07 90.8
188 177 171.4 125 174 84.8
17 17 17 17 17 20 22 24 24 27 29 30
bility of hydrogen in the metal phase and the deviation from stoichiometry are both small in these systems in the temperature range of the measurements (see Table 3) and equation 8 is there­ fore applicable. The poor agreement in the case of MgH2 is prob­ably due to impurities in the samples used in these determina­ tions. The effect of oxygen, for example, on the hydrogen equili­ brium pressures are unknown, but could very well lead to large errors in the enthalpy determination. For SrH2 the equilibrium-dissociation data are incomplete and the value given in Table 3 is not well established. The agreement for the other hydrides, CaHj, BaH^, TiH2, and ZrH1 5, is satisfactory. It thus appears that despite the difficulties discussed above; dissociation pres­ sure data are useful in estimating thermodynamic properties and will in many cases give excellent agreement with calorimetric measurements.
The enthalpies of formation of the alkali hydrides range from -52 to -91 kJ/mole hydride (-104 to -182 kJ/mole H2, Table 3) while the alkaline earth hydrides CaH2, SrH2, and BaH2 have en­thalpies of formation between -171 and -188 kJ/mole. The low value for MgH2, -90.8 kJ/mole reflects the partially coyalent character of this hydride. The enthalpy of formation of YbH2 is -181 kJ/mole[31] which is considerably less tabsolute value) than that of the other rare earth dihydrides, but very close to the value for CaH2. The bonding (ionic) in YbH2 is basically differ­ ent from that of the other rare earth dihydrides. There is no data available for EuH2, but for the same reason the enthalpy of formation is expected to be more in line with the alkaline earth
Survey of the Different Types of Hydrides 27 hydrides than with the rare earth dihydrides. The enthalpies of for­ mation of the other rare earth dihydrides (including Sc and Y) as de­ termined from dissociation pressure data, range from -196 to -227 kJ/mole[15]. The enthalpies of formation of the rare earth trihy- drides points out the fact that the excess hydrogen, beyond MH2, is much more weakly bound. For LaH3, CeH3, PrH3, and NdHß, for example, AHf is approximately -243 kJ/mole hydride[15] which is equivalent to -162 kJ/mole hydrogen. The corresponding values for the dihydrides are -208, -206, -208, and -213 kJ/mole hydrogen[15].
REFERENCES (1) H. Shull, J. Appl. Phys. 33, 292 (1962).
(2) R. S. Calder, W. Cochran, D. Griffiths, and R. D. Lowde, J. Phys. Chem. Solids, 23, 621 (1962).
(3) T. R. P. Gibb, Jr. (1962) Progress in Inorganic Chemistry, Vol. 3, 315, Interscience, New York.
(4) G. G. Libowitz, (1965) The Solid-state Chemistry of Binary Metal Hydrides, W. A. Benjamin, Inc., New York.
(5) W. M. Mueller, J. P. Blackledge, and G. G. Libowitz (1968) ,Metal Hydrides, Academic Pressf New York.
(6) K. M. Mackay (1966) Hydrogen Compounds of the Metallic Elements, E. & F. N. Spon Ltd., London.
(7) G. G. Libowitz (1972) MTP Internat. Rev. Sei., Inorg. Chem. Ser. 1, Vol. 10, Solid State Chemistry, L. E. J. Roberts, Ed., Butterworths Ltd., London, pp. 79-116.
(8) N. F. Mott and H. Jones (1936) The Theory of the Properties of Metals and Alloys, Dover Publications, Inc., New York, Chapter VII.
(9) G. G. Libowitz and T. R. P. Gibb, Jr. J. Phys. Chem., 60, 510 (1956).
(10) G. G. Libowitz, Ber. Bunsenges. Phys. Chem. 76, 837 (1972).
(11) A. C. Switendick, Solid State Commun. 8, 1463 (1970); Int. J. Quant. Chem. 5, 459 (1971); Ber. Bunsenges. Phys. Chem., 76, 535 (1972); Proc. Hydrogen Econ. Miami Energy Conf., T. N. Veziroglu, Ed., Univ. of Miami, Coral Gables, Fia., p. S6-1 (1974); J. Less-Common Metals, 49, 283 (1976).
(12) B. M. Bansal and D. Damien, Inorg. Nucl. Chem. Letters, 6, 603 (1970).
(13) J. A. Fahey, J. R. Peterson, and R. D. Baybarz, Inorg. Nucl. Chem. Letters, 8, 101 (1972).
(14) W. M. Olson and R. N. R. Mulford, J. Phys. Chem., 70, 2935, (1966).
28 A. J. Maeland (15) G. G. Libowitz and A. J. Maeland (1978) Handbook on the Physics
and Chemistry of Rare Earths, K. A. Gschneider and L. Eyring, eds., North-Holland Publishing Co., Amsterdam, Chapter 26.
(16) C. B. Hurd and G. A. Moore, J.A.C.S., 57, 332 (1935).
(17) S. R. Gunn, J. Phys. Chem., 71, 1386 (1967).
(18) A. Herold, Ann. Chim. (Paris), 6, 536 (1951).
(19) J. F. Stampfer, Jr., C. E. Holley, Jr., and T. F. Suttle, J.A.C.S., 82, 3504 (1960).
(20) V. I. Pepekin, T. N. Dymova, Yu. A. Lebedev, and A. Ya. Apim, Zh. Fiz. Khim., 38, 1024 -(1964).
(21) R. W. Curtis and P. Chiotti, J. Phys. Chem., 67, 1061 (1963).
(22) J. N. Brönsted, Z. Electrochem., 20, 81 (1914).
(23) M. D. Banus and R. W. Bragdon, A Survey of Hydrides, USAEC Re­
port CF-52-2-212, Metal Hydrides, Inc., Feb. 1, 1952.
(24) A. Guntz and F. Benoit, Ann. Chim., 20, 5 (1923).
(25) W. C. Schumb, E. F. Sewell, and A. S. Eisenstein, J.A.C.S., 69,
2029 (1947).
(26) R. M. Hagg and F. J. Shipko, J.A.C.S., 78, 5155 (1956).
(27) B. Stalinski and Z. Bieganski, Bull. Acad. Polon. Sei., Ser.
Sei. Chim., 10, 247 (1962).
(28) 0. M. Katz and E. A. Gulbransen, J. Chem. Ed., 37, 533 (1960).
(29) A. G. Turnbull, Austral. J. Chem., 17, 1063 (1964).
(30) B. M. Abraham and H. E. Flotow, J.A.C.S., 77, 1446 (1955).
(31) C. E. Messer, T. Y. Cho, and T. R. P. Gibb, Jr., J. Less-Common Met. 12, 411 (1967).
(32) A. D. McQuillan, J. Less-Common Met., 49, 431 (1976).
Equilibrium Hydrogen Pressure
30 A. J. Maeland
z LU O CC Û > X
o
1 1
1 ± 2
system in which two hydride phases form.
Survey of the Different Types of Hydrides 31
!2Z T(K)
Fig. 3. Logarithm of dissociation pressure vs. 1/T(K) for lutetium dihydride[32].
STRUCTURE AND BONDING IN METAL HYDRIDES*
W. E. Wallace and S. K. Malik Department of Chemistry, University of Pittsburgh
Pittsburgh, Pennsylvania 15260, U.S.A.
ABSTRACT
Bonding in hydrides of the alkali metals, of the alkaline earth and of Eu and Yb is essentially ionic in nature and hydrogen in these materials is anionic. This is indicated by a variety of properties of these hydrides - structures, stoichio- metries and lattice energies. In the early studies of the magnetism of the Pd-H system it was concluded that hydrogen in this hydride was protonic, its Is elec­ tron populating states in the d-band of the host metal. More recent work on the band structure of transition metal hydrides indicates that the simple protonic model is incorrect. Hydrogen participates, along with the ion cores of the host metal, in establishing the potential within which the delocalized electrons move. Hydrogen contributes states as well as electrons, in contrast with its behavior according to the protonic model in which it contributes only electrons.
The special complexities which can arise for sub-stoichiometric hydrides - order- disorder phenomena on the hydrogen sub lattice, hydrogen structure polymorphism, etc. - are illustrated by reference to the well studied system Ta2H.
INTRODUCTION
This paper is intended to serve as a short review of the structures of and bonding in metallic hydrides. Only the salient features are described and these for only a few hydrides carefully selected to typify the several classes of known hydrides. Reference is made to the anionic and protonic models employed earlier to charac­ terize hydrogen in metallic hydrides. A brief account is also given of the alloy model which is currently regarded as providing the best description of transition metal hydrides.
A large fraction of the metals in the Periodic Table form hydrides. These hydrides exhibit a variety of structural types (1). The rock salt and fluorite structures are frequently observed. Since these are structures characteristic of ionic materials, superficial considerations might lead one to believe that metal hydrides are ionic in nature. While this is true of some hydrides, e.g., hydrides of the alkali, the alkaline earth and the rare earth metals, it is certainly not true in general.
Bonding in metallic hydrides is intimately related to the issue of the electronic make-up of hydrogen in these materials. Prior to the present decade two models have been employed (2) to describe the electronic configuration of hydrogen - the anionic model for hydrides such as LiH, Ca^, etc., and the protonic model for
*This work was assisted by a grant from the Petroleum Research Fund.
33
34 W. E. Wallace and S. K. Malik transition metal hydrides. In the simplest version of the protonic model hydrogen is presumed to exist as a bare (i.e., unscreened) proton. Clearly this model is conceptually implausible and it is virtually certain that there are no materials for which the simple protonic model is applicable. The original protonic model has been superseded by the alloy model in which the protons contribute, along with the metal ion cores, to the electrostatic potential in which the delocalized electrons move. This point of view has been given expression in the works and publications of Switendick referred to later on in this paper.
In regard to the structures of metal hydrides hydrogen is usually, but not always, found in the tetrahedral or octahedral interstices (Table 1). Unique hydrogen sublattice structure is found for many hydrides, e.g., NaH, CeH2, H0D3, etc. However, in some metallic hosts the energetics are unfavorable for total occupancy of a particular set of interstitial sites and only partial occupancy occurs. This gives rise to arrangements that are not structurally unique and the possibility of order-disorder transformations involving the occupied and unoccupied sites, features which confer structural complexity on these kinds of metallic hydrides. Some of these complexities have been revealed in detailed studies of Ta H and are elaborated upon in a later section of this paper.
HYDRIDES EXHIBITING IONIC BONDING
Alkali and Alkaline Earth Hydrides
NaH and NaD were examined by Shull et al. (3) using neutron diffraction in the early days of this technique. They established that these hydrides* occurred in the NaCl structure which is typical of ionic materials. In these hydrides hydrogen is situated in octahedrally coordinated sites. The easily ionized alkali metal loses its valence electron to the hydrogen atom to form the hydride anion and the system is stabilized by its Madelung energy. Thus the bonding in these materials is essentially ionic
Similar considerations hold for the alkaline earth hydrides. CaH2, an example of these hydrides, has been extensively studied, most recently by Andresen, Maeland and Slotfeldt-Ellingsen (4). It occurs in the C29 structure. The metal ions in this structure are arranged in a slightly distorted cph structure. In the normal cph structure treated using the orthohexagonal cell b/a = /T* = 1.73. In the orthorhombic CaH2 structure the metal ions are displaced so that this ratio is increased by about 10% to 1.89. Half of the hydride ions are situated in tetra­ hedral interstices, but because of peculiarities in the C29 structure, not all metal atoms are equidistant from the central H~ ion. Ca-H distances range from 2.24 to 2.28 A. The remaining hydride ions have 5 Ca2+ ions as near neighbors of distances ranging from 2.38 to 2.63 A.
Lattice energies calculated by Gibb (5) are listed in Table 2 along with experi­ mental values established through the Bom-Haber cycle. The close agreement of the calculated and experimental lattice energies strongly supports the notion that bonding in the hydrides of the alkali and alkaline earth metals is essentially ionic in nature.
*Both hydrides and deuterides will be referred to in this paper simply as hydrides, since there is no indication that bonding and/or structure in hydrides or deuterides differ with the possible exception of the V-H and V-D system. (See D. G. Westlake, M. H. Mueller, International Conference on Hydrogen in Metals, Jlilich, Germany, 1972.)
Structure and Bonding in Metal Hydrides 35
Table 1 Location of Hydrogen In Metallic Hydrides a b c Octahedral Interstices Alkali metals ; Pd ; Ni
Tetrahedral Interstices Alkaline earths (half of the hydrogen) ; rare earth dihydridese; Tif, Zrg, Hff; Uh; Thg; Ta1; rare earth trihydrides.
Other Alkaline earths (half of the hydrogen) ; rare earth trihydrides (one-third of the hydrogen)J
a. Ref. 3 b. J. E. Worsham, Jr., M. K. Wilkinson and C. G. Shull, J. Phys. Chem.
Solids 3_, 303 (1957). c. J. W. Cable, E. 0. Wollan and W. C. Koehler, Int. Colloquium on
Diffraction and Diffusion of Neutrons, CHRS Publ. No. 12t (1964), p. 36.
d. Ref. 4 e. D. E. Cox, G. Shirane, W. J. Takei and W. E. Wallace, J. Appl. Phys.
34., 1342 (1963) and C. E. Holley, R. N. R. Mulford, F. H. Ellinger, W. C. Koehler and W. H. Zachariasen, J. Phys. Chem. 59.» 1 2 2 6 (1955).
f. S. S. Sidhu, L. Heaton and D. D. Zauberis, Acta Cryst. 9^ 607 (1956). g. R. E. Rundle, C. G. Shull and E. 0. Wollan, Acta Cryst. .5, 2 2 (1952). h. R. E. Rundle, J. Am. Chem. Soc. 73» 4 i 7 2 (1951). i. Ref. 19 j. M. Mansmann and W. E. Wallace, J. Phys. (Paris) 25, 454 (1964).
Table 2 Lattice Energies of Saline Hydridesa
LiH NaH KH RbH CaH CaH2
SrH2
BaH0
Value le)
a. Ref. 5
Hydrides of Europium and Ytterbium
While most of the rare earth elements (and chemically similar ytterbium) react with hydrogen to form trihydrides or hydrides of composition approaching the trihydride,* Yb and Eu under modest hydrogen pressures (<100 atm.) form only the dihydride. Crystallographic work by Korst and Warf (6) showed that the metal sublattice has the same structure as that in CaH2 and presumably the three dihydrides are isostructural.
The limiting composition of the hydrides of Eu and Yb is consistent with the known valence of these elements and is also consistent with the concept of anionic hydro­ gen. These elements form only the dihydride because Eu and Yb are divalent. (This is in contrast with the trivalency exhibited by the other rare earths and the con­ sequent formation of trihydrides in these cases [7,8].)
The dipositive character of Eu and Yb is a consequence of the exceptional stability of the half-filled and filled 4f shell. Eu3+ would exist in a 4f66s25d configura­ tion. Because of the special stability of the half-filled 4f shell Eu3+ captures an electron from the conduction band to achieve the 4f7 configuration and hence becomes dipositive. Similar considerations hold for Yb3+; it captures a conduction electron to achieve a filled 4f shell. The limiting stoichiometries of Eu and Yb hydrides and their structures strongly support the view that these are saline hydrides, stabilized primarily by ionic bonding, and hence they should be represent­ ed as R2 +(H~)2 where R = Eu or Yb.
TRANSITION METAL HYDRIDES
Early Pd-H Studies - The Birth of the Protonic Model
Extensive hydride formation is exhibited by metals in the titanium and vanadium groups. Pd forms hydrides as well as Ni, under appropriate conditions. The issue of the nature of the bonding in these materials is a matter of great complexity, rivalling the problems presented by bonding in the elemental transition elements and their alloys.
Pd was the first and has been the most extensively studied transition metal for the effects of hydrogénation on the host metal (9-12). It was found that the strong paramagnetism of Pd is gradually diminished as it is progressively hydrogenated. Mott and Jones ascribed (13) this to the filling of the d band by the Is electron supplied by the incoming hydrogen, and in so-doing spawned the protonic model. In effect this implied that the solute (H) contributes electrons but not states to the host metal (Pd), or more correctly the states being contributed by hydrogen are above the Fermi limit and are hence unused.
This simplistic picture of the Pd-H system gave way in time to a more satisfying concept (14). It is well known that hydrogénation of Pd leads initially to the formation of a primary solid solution (the a phase) which is rather H-poor. As hydrogénation is continued a second phase (the 8 phase) somewhat richer in hydrogen is formed. The 8 phase has a filled d-band and is diamagnetic. The decrease in susceptibility of Pd as it is hydrogenated is a consequence of the decrease in amount of a phase material, which is paramagnetic, and a simultaneous increase in
*For a discussion of the structure and bonding in rare earth trihydrides see paper by W. E. Wallace, Electric and Magnetic Properties and Rare Earth Intermetallic Hydrides, proceedings of this conference.
Structure and Bonding in Metal Hydrides 37 the amount of the diamagnetic 3-phase. Thus the magnetic effects accompanying the hydrogénation of Pd are dominated by the details of the phase diagram rather than by progressive band filling as postulated by Mott and Jones (13).
Photoemission studies by Eastman, Cashion and Switendick (15) clearly revealed that hydrogen in the Pd-H system contributes states as well as electrons and these states lie well below the Fermi energy. To appreciate the electronic configuration of the hydrides of Pd and other transition elements one must resort to more sophisticated methods of analysis than that involved in considerations based upon the simple protonic metal. To put it differently, it is quite incorrect to assume that all the states supplied by hydrogen lie well above the Fermi energy and are therefore not involved in the bonding in the hydrides.
As noted in the Introduction, the concept of a bare proton is implausible. The field generated by the proton interacts with the sea of delocalized electrons and this field, along with the potential supplied by the ion cores, determines the detailed band structure of the hydride. This problem has been attacked by Switendick (16-18) and some of the results obtained are briefly summarized in the next section.
Band Structures of Transition Metal Hydrides
Switendick has used the APW (augmented plane wave) method to establish the band structure of several transition metal hydrides. Among the objectives of such work is the assessment of the anionic and protonic models and (2) the elucidation of the structures and stoichiometries of transition metal hydrides in terms of the deep fundamentals of these systems. In the furtherance of these objectives Switendick has made calculations on several mono-, di- and trihydrides for a few prototype cubic structures. To elucidate the nature of the states which are filled when hydrogen is introduced, he has performed a detailed analysis of the charge density associated with various states in the energy band and has made a spherical decomposition of the charge density about the metal and hydrogen posi­ tions to establish the symmetry of the states.
The approach employed by Switendick is conveniently illustrated by citing a few of his results on compounds with x = 0, 1, 2 and 3. The YHX compounds are assumed to exist in the fee, NaCl, CaF2 and BÍF3 structures for x = 0, 1, 2 and 3, respec­ tively. This structural assumption is made for YH notwithstanding the fact that elemental yttrium is hexagonal. In YHQ, the hydrogen potential is replaced by regions of zero potential. With YH3, the two kinds of hydrogen are assumed to have the same potential, and to have the same potential as in YH2.
Monohydrides. Comparison of the band structure obtained for Pd and PdH (or for YH and YH) reveals the following:
1. The d- and f-like states are only slightly perturbed and are very similar in the two.
2. The metal s and p states hybridize with the hydrogen ls-orbital; they are strongly perturbed, lowered substantially in energy and become very much hydrogen-like. In the case of PdH these low-lying states are obtained by modification of already existing (and filled) states. Therefore, the additional electrons from hydrogen must fill other states, the lowest-lying unoccupied states. For PdH these are of two kinds. The first of these are states (0.36 per Pd atom) primarily of d-character, lying just above the Fermi limit. For small hydrogen concentrations only these states are occupied and it appears superficially as if hydrogen is contributing electrons
38 W. E. Wallace and S. K. Malik
to the band of the host metal. However, detailed charge density calculations show that there is 'UKó electron of Is character inside the hydrogen APW sphere, which is slightly larger than the 0.5 electron inside the same size sphere for the hydrogen atom. Thus it is incorrect to state that hydrogen donates electrons to the host metal.
The other states which are available to accept electrons, apart from the 0.36 d-states, are those originally unoccupied and associated with the s-p band. These states interact with the hydrogen Is orbital and have their energy lowered below the top of the d-band.*
Dihydrides. As noted earlier, these hydrides form in the fluorite structure; they contain 8 hydrogen atoms per unit cell. Band structure calculations indicate the following:
1. The Fermi energy falls in a band derived from the metal d-states.
2. A new band in the dihydride appears below the d-band. This band arises from the antibonding combination of the Is orbitals of two hydrogen atoms in the unit cell. (This does not occur for mono- hydrides, which are only half as concentrated in hydrogen.)
The stability of the dihydrides is attributable to this new band which can accommo­ date electrons at a rather low energy. The position of this band is determined by the hydrogen-hydrogen separation. If the hydrogen-hydrogen separation is small, the new band has high energy and is unlikely to be filled by electrons originally associated with hydrogen. For large H-H separations the band is low-lying and will be filled. For hypothetical PdH2 the "antibonding band" lies well above the metal d-band and filling it is energetically unfavorable. In contrast, in PdH the electrons brought in by hydrogen can be easily accommodated. This accounts for the existence of PdH (slightly sub-stoichiometric in hydrogen) and the non- existence of PdH2.
For Y and Pr the new band falls below the metal d band (and below the Fermi energy), and for these the formation of a dihydride is more favorable than the formation of a monohydride. (In fact, the monohydrides do not exist.)
Trihydrides. When one fills both the octahedral and tetrahedral sites, one gets the BiF3 structure of the cubic trihydrides. Calculations show that another band now appears below the d-band, and this arises because of the interaction between octahedral hydrogen-tetrahedral hydrogen antibonding states. The interstitial separations determine the energy of this band. In turn, the interstitial separation is determined by the metal-metal distance. In going from cubic rare earth dihy­ drides to cubic trihydrides the metal-metal distance decreases, while it increases or remains the same in going from cubic dihydrides to hexagonal trihydrides. This influences the position of energy levels and hence the stability of structures. Thus the cubic structure is preferred for light rare earth trihydrides and hexag­ onal structure by heavy rare earth trihydrides. For the case of Ti the additional band falls above the Fermi level and therefore TÍH3 does not form. Switendick has given plausible reasoning for the smooth change of structure from dihydride to
*It is these extra states which led to the earlier erroneous conclusions that Pd contains 0.6 vacant d-states.
Structure and Bonding in Metal Hydrides 39
trihydride for light rare earths and for the occurrence of two phases and a structure change in the case of the heavy rare earths.
Ta2H - AN EXAMPLE OF STRUCTURAL FEATURES OF A SUBSTOICHIOMETRIC HYDRIDE
Body centered cubic Ta readily takes up hydrogen at elevated temperatures (>300°C) to form a hydride which retains cubic symmetry. Neutron diffraction work has shown that hydrogen in Ta resides (19,20) in the tetrahedral interstices of which there are six per metal atom. The theoretical composition is therefore TaHg were all sites filled. In Ta2H, which has been extensively studied, only one-twelfth of the available sites are filled.
Some years ago the thermodynamics of the Ta-H system were exhaustively investigated by Wallace and his associates (21-23). The entropy of formation of Ta2H was determined at 300 C and from this the entropy of Ta2H at this temperature was readily established since entropies of elements are known. From the known heat capacity data for Ta2H it was established additionally that this hydride has vanishing entropy at 0 K.
The heat capacity of Ta2H measured by Saba et al. (Figs. 1 and 2) exhibited (23) three λ-type heat capacity anomalies, peaking at 306, 322 and 333.5 K.* This implied that there are at least three polymorphic varieties of Ta2H - one form (3i) existing below 306 K, another (32) between 306 and 333.5 K and the third (a) above 333.5 K. The thermal anomalies are attributed to the energy involved in rearranging hydrogen in the twelvefold more abundant sites. These can therefore be viewed as a consequence of a type of order-disorder transformation of the hydrogen on the sublattice of tetrahedral sites.
The configurâtional entropy information acquired by Saba et al. (23) is summarized in Table 3. The 3χ form is a structurally unique material. This was documented
Table 3 Configurational Entropy of the Polymorphic Varieties of Ta2H
Random Distribution of Hydrogen Entropy (cal/deg.g.atom of H) over 12 sites/cell 6.84 a form at 340 K 5.1 32 form at 317 K 0.9 3χ form at 290 K 0.2 3χ form at 0 K 0b
a. This is the entropy associated with the disorder resulting from the various arrangements of hydrogen in the twelvefold more abundant sites.
b. Saba et al. give -0.39 ± 0.30 cal/deg.g.atom of H for this quantity. This is regarded as zero within the limit of error.
by the study of Wallace (19), which showed superlattice lines in the neutron diffraction pattern of 3i Ta2D. The enlarged unit cell was not observed at 326 K, indicating that the hydrogen superlattice is destroyed in 32 and aTa2D.
*The existence of two closely spaced points of 332 and 333.5 K may be an artifact introduced by compositional variation in the sample employed.
40 W. E. Wallace and S. K. Malik This study is instructive in that it indicates that a variety of hydrogen sub- lattice structures may be observed in sub-stoichiometric metallic hydrides. In the example cited the arrangements range from the 3χ form which is structurally unique at low temperatures to the a form whose configuration has begun to approach that of a random distribution of hydrogen on the lattice composed of the tetra- hedral sites.
REFERENCES
1. See, for example, G. C. Libowitz, The Solid-State Chemistry of Binary Metal Hydrides, W. A. Benjamin, Inc., New York (1965), Chap. 3.
2. Ref. 1, pp. 5 and 6. 3. C. G. Shull, E. 0. Wollan, G. A. Morton and W. L. Davidson, Phys. Rev. 72,
842 (1948). 4. A. F. Andresen, A. J. Maeland and D. Slotfeldt-Ellingsen, J. Solid State
Chem. 2£, 93 (1977). 5. T. R. P. Gibb, "Primary Solid Hydrides," in Progress in Inorganic Chemistry,
F. A. Cotton, ed., Interscience Publishers, Inc., New York (1962), p. 397. This contains an excellent summary of the various facets of metallic hydrides. It covers inferences about bonding drawn from metal-hydrogen distances. This has had to be omitted from the present short review.
L. Korst and J. C. Warf, Acta. Cryst. 9_, 452 (1956). E. Sturdy and R. N. R. Mulford, J. Am. Chem. Soc. 78.» 1Q83 (1956). Pebler and W. E. Wallace, J. Phys. Chem. 66 , 148 (1962). Biggs, Phil. Mag. 32 , 131 (1916). Aharoni and F. Simon, Z. Physik. Chem. B4_, 175 (1929). Svensson, Ann. Phys. Leipzig JU, 699 (1932) and 18, 299 (1933). C. Jamieson and F. D. Manchester, J. Phys. F. 2_, 323 (1972). F. Mott and H. Jones, The Theory of the Properties of Metals and Alloys,
Dover Publications, New York (1958), p. 200. 14. For additional details see W. E. Wallace in Hydrogen in Metals, G. Alefeld,
ed., Springer-Verlag, New York, to appear in 1978. 15. D. E. Eastman, J. K. Cashion and A. C. Switendick, Phys. Rev. Lett. 2_7, 35,
(1971). 16. A. C. Switendick, Solid State Commun. _8, 1463 (1970). 17. A. C. Switendick, Int. J. of Quantum Chem. 5., 459 (1971). 18. A. C. Switendick, Ber. der Bunsen-Gesellschaft T6_, 535 (1972). 19. W. E. Wallace, J. Chem. Phys. 35^, 2156 (1961). 20. V. A. Somenkov, A. V. Gurskay, M. G. Zempyvanov, M. E. Kost, N. A.
Chernoplekov and A. A. Chertkov, Solid State Phys. 1£, 2797 (1968). 21. P. Kofstad, W. E. Wallace and L. J. Hyvonen, J. Am. Chem. Soc. 1» 5015
(1959). 22. W. E. Wallace, P. Kofstad and L. J. Hyvonen, Pure and Applied Chem. _2> 281
(1961). 23. W. G. Saba, W. E. Wallace, H. Sandmo and R. S. Craig, J. Chem. Phys. 35.,
2148 (1961).
2 4
O .I8
li ne
) an d
Ta H
( fu ll
l in
250h
200l·
I50F
«S
50h
2 9 0 3 0 0 310 3 2 0 3 3 0 3 4 0 3c TEMPERATURE, ( K )
Fig. 2 Heat capacity versus temperature for Ta (dashed line) and Τβ2Η (full line). The long range order is lost at the thermal anomaly at 305 K.
THERMODYNAMICS OF METAL, ALLOY AND INTERMETALLIC/HYDROGEN
SYSTEMS
Ted B. Flanagan Department of Chemistry, University of Vermont, Burlington, Vermont
ABSTRACT
The thermodynamics of solution of hydrogen in metals (alloys or intermetallic compounds) is reviewed. The conversion of experi­ mental data from conditions of essentially constant pressure to conditions of constant volume is discussed. The fundamental significance of thermodynamic parameters obtained from the tem­ perature dependence of the two-phase coexistence pressures is considered and the necessary criteria for which these values correspond to the thermodynamics of the reaction: J^2 (QÏ + M(H"sa t u r a t : ed) "* metal hydride are developed. Some experimental methods are reviewed.
SINGLE PHASE REGIONS OF SOLUBILITY The solution of hydrogen in a metal can be represented schemat­ ically by
where [H] indicates hydrogen atoms interstitially dissolved in the metal. The term metal will be used for convenience with the understanding that the discussion refers equally well to alloys and intermetallics. The equilibrium condition for hydrogen dis­ solved in a metal can be given as
UH(g) = K(g) = 'H (2)
where y„ - H„ - TS„, etc. The chemical potential of gaseous hydrogen is related to its pressure (fugacity) by equation 3
2 H2(g) 2 H2 H2
and using equation 2, RT in P;/2 = ΔμΗ = μΗ - Ι μ ^ = ΔΗΗ - ΤΔ8Η (4)
43
44
where
T. B. Flanagan
ASH = SR - isR (6) H H 2 H2 ( g J
ΔΗ„ and AS„ are the relative partial molar enthalpy and entropy of solution of hydrogen, respectively. Partial molar quantities are needed since the thermodynamic changes upon solution of hydrogen are dependent upon the hydrogen content of the metal within single phase regions. (In keeping with recommendations of I.U.P.A.C. [1] the subscripts on the thermodynamic functions are sufficient to designate partial molar quantities.) Thus measurements of the hydrogen pressure in equilibrium with a solid phase containing dissolved hydrogen give (equation 4) a measurement of Δμ„. The temperature dependence of Δμ„ at a given hydrogen content yields values of ΔΗ„ and AS„ provided that the same solution process obtains over the temperature range where data are analyzed. Changes of AS„ with hydrogen content are dominated by the large,
c H' and rapidly varying partial configurational entropy, Sc