Hydrides for Energy Storage. Proceedings of an International Symposium Held in Geilo, Norway, 14–19 August 1977
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BLAIR et ai
Hydrogen in Metals
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VEZIROGLU & SEIFRITZ Hydrogen Energy System
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VEZIROGLU Energy Conversion — A National Forum
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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'
PERGAMON PRESS OXFORD NEW YORK · TORONTO · SYDNEY · PARIS ·
FRANKFURT
U.K. Pergamon Press Ltd., Headington Hill Hall, Oxford OX3 OBW,
England
U.S.A. Pergamon Press Inc., Maxwell House, Fairview Park, Elmsford,
New York 10523, U.S.A.
CANADA Pergamon of Canada Ltd., 75 The East Mall, Toronto, Ontario,
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FEDERAL REPUBLIC Pergamon Press GmbH, 6242 Kronberg-Taunus, OF
GERMANY Pferdstrasse 1, Federal Republic of Germany
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,
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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
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