Syntheses and Physical Studies of Inorganic Compounds

231 Pages·1972·9.92 MB·English

by  C. F. Bell (Auth.)

SYNTHESES AND PHYSICAL
STUDIES OF INORGANIC COMPOUNDS
BY
C. F.BELL, M.A., D.PHIL., F.R.I.C.
Lecturer in Inorganic Chemistry
Brunei University, London
PERGAMON PRESS
OXFORD · NEW YORK
TORONTO · SYDNEY · BRAUNSCHWEIG
Pergamon Press Ltd., Headington Hill Hall, Oxford
Pergamon Press Inc., Maxwell House, Fairview Park, Elmsford,
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Pergamon Press (Aust.) Pty. Ltd., 19a Boundary Street,
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Vieweg & Sohn GmbH, Burgplatz 1, Braunschweig
Copyright © 1972 C. F.Bell
All Rights Reserved. No part of this publication may be
reproduced, stored in a retrieval system, or transmitted, in any
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Pergamon Press Ltd.
First edition 1972
Library of Congress Catalog Card No. 79-178772
Printed in Germany
08 016651 2
TO
SHEILA
PREFACE
THE renaissance of inorganic chemistry during the past three decades owes much of its
origin and continuing impetus to advances in three directions. Firstly, existing preparative
techniques have been improved and new ones developed, resulting in the discovery of many
new kinds of inorganic compounds. Secondly, established methods for studying the physico-
chemical properties and structures of chemical compounds have been supplemented by a
number of important, recently developed techniques, notably nuclear magnetic resonance,
electron spin resonance and Mössbauer absorption spectroscopy. Thirdly, progress in
theoretical chemistry, particularly in the evolution of the molecular orbital and ligand field
theories, has facilitated the rationalization of much experimental data and deepened our
understanding of the forces which bind atoms together.
The present state of inorganic chemistry is reflected in the research publications appearing
in the literature. Many current papers are devoted to establishing the detailed physical
and chemical properties of specific compounds with the emphasis very much on the ap­
plication of physical principles and investigational techniques and the theoretical inter­
pretation of experimental data.
Concurrently with the great expansion of research activity in inorganic chemistry, the
teaching of this subject has veered sharply towards comparative treatments of the chem­
istry of the elements based on the principles underlying the Periodic Table. In such treat­
ments, the student gains an overall appreciation of inorganic chemistry without being
burdened with excessive detail. There still remains, however, a significant gap between
this approach and many of the topics which excite the interest of research chemists
today.
There are various ways in which recent advances can be presented to the undergraduate
or postgraduate student. One is to adopt the approach exemplified in two recent books
(Drago, R., Physical Methods in Inorganic Chemistry, Reinhold, New York, 1965, and
Hill, H. A. O., and Day, P., Physical Methods in Advanced Inorganic Chemistry, Interscience,
London, 1968) on physical methods as applied to inorganic chemistry. The emphasis in
this approach is on the principles and experimentation of the methods discussed and
undoubtedly it provides a sound foundation for the advanced study of the subject.
In this book I Jiave chosen an alternative approach. This is based on the consideration,
in some depth, of the synthesis, properties and reactions and structures of a number of
compounds, selected on the criterion that the study of each has resulted in important
contributions to our present practice and understanding of inorganic chemistry. Details
of experimental procedures are, in general, not included, for these may be found else­
where. For example, the syntheses of almost all the compounds discussed in the following
la SPS IX
PREFACE
pages are described in full in the appropriate volume of the series Inorganic Syntheses
(McGraw-Hill).
My aim throughout has been to combine the results obtained in many different physico-
chemical investigations of each compound in order to show the current state of our know­
ledge. This approach also serves to illustrate how closely interrelated and interdependent
inorganic and physical chemistry have now become and the artificiality of trying to maintain
sharplydefined boundaries between the two disciplines. I trust that this book will introduce
the reader to many of the significant research papers of recent years and will encourage
him to pursue his own studies of the chemical literature.
In the text, the S.I. units are used for the expression of most physical data. There are
the following exceptions. In the cases of magnetic and dipole moments, traditional units
are preferred because of the awkward magnitudes of the S.I. units. The S.I. equivalent
of each is given in the text on the occasion of the first reference to it and thereafter the
traditional unit is used. Temperatures are quoted in degrees Celsius, °C, and wave-num­
bers in reciprocal centimetres, cm-1. Althoug neither is strictly in accordance with S.I.
practice, these are still the units in common usage and it is convenient to retain them here
for ease of crossreference with the literature.
I wish to record my grateful thanks to the authors and publishers who have given
permission to reproduce figures as indicated in the text and to all those research chemists
without whose original inspiration and creative work this book would never have been
possible.
I would like especially to thank Professor H. M. N. H. Irving for his original idea
regarding the writing of this book and for his continued interest and help throughout its
preparation.
Finally, I must record my gratitude to my wife for her considerable help in the typing
and editing of this work and my deep appreciation of her patience and encouragement in
bringing it to completion.
x
CHAPTER 1
DIBORANE
Introduction
The first observation of compound formation between boron and hydrogen appears to
have been made by Sir Humphrey Davy1 in 1809. He noted that a gas was obtained when
the product of the reaction of boron trioxide with potassium was treated with water or
dilute hydrochloric acid. This was mainly hydrogen, but also contained another compo­
nent which burnt with a blue flame tinged with green. He identified this as a volatile com­
pound of boron and hydrogen.
From time to time during the next 100 years, hydrides of boron were made by other
research workers. These compounds were not correctly identified, probably because of
the difficulty in purifying and analysing such reactive substances.
Alfred Stock and his co-workers were the first to characterize the hydrides of boron
(or boranes as they are now commonly called). Their work2 is remarkable for the develop­
ment of experimental techniques for the handling of highly reactive, volatile compounds.
They were able to estabUsh the existence of as many as ten boranes. Stock's work provides
the basis for modern synthetic practice and it remains one of the great, classical contribu­
tions to the development of inorganic chemistry in this century.
For many years, interest in borane chemistry was stimulated by the recognition that,
in their stoichiometry, these compounds did not conform with accepted theories of valence.
They, and some other similar inorganic molecules, are 'electron-deficient' because there
are insufficient valence electrons available to make conventional two-electron bonds between
each pair of atoms bonded together. It has become necessary to introduce the concept
of a multicentre bond into valence theory to account for the bonding in such com­
pounds.
More recently, boranes have been intensively studied in view of their potentialities as
chemical fuels. From a consideration of energy content per unit weight, boranes are
theoretically much superior to hydrocarbons as fuels, although the difficulty of handling
them in large quantities has impeded their exploitation on an industrial scale. Some borane
derivatives, notably the tetrahydroborates, are applied to an increasing extent as reducing
agents for both inorganic and organic compounds.
As the simplest known borane, diborane (B H ) has attracted special interest and it has
2 6
been the subject of very many physico-chemical studies.
la* 1
SYNTHESES AND PHYSICAL STUDIES OF INORGANIC COMPOUNDS
Preparation
Stock prepared a mixture of tetraborane, B H , pentaborane, B H , hexaborane,
4 10 5 9
B H , and decaborane, B H , by the reaction between magnesium boride and hydro­
6 10 10 14
chloric acid. Separation from impurities such as carbon dioxide, hydrogen sulphide and
silanes proved difficult and yields were low. Better yields were obtained by the use of
phosphoric instead of hydrochloric acid. Diborane itself is rapidly decomposed in aqueous
solution and so is not formed directly in the above reaction and it was made by Stock by
the pyrolysis of B H .
4 10
Schlesinger and Burg3 greatly improved the synthesis of B H by the action of an electric
2 6
discharge on a mixture of BC1 and excess hydrogen at low pressure. Monochlorodiborane,
3
B H C1, is formed. This disproportionates at 0°C and atmospheric pressure into B H
2 5 2 6
andBCl .
3
Neither of these reactions is convenient in practice and they have been superseded4 by
the reduction of BC1 or BF with lithium aluminium hydride, LiAlH .
3 3 4
3LiAlH + 4BC1 -► 3LÌC1 + 3A1C1 + 2B H
4 3 3 2 6
LiAlH is dissolved in diethyl ether and allowed to react with the diethyl ether complex
4
of BC1 , (C H ) 0 · BC1 , at a low temperature. B H is released when the reaction
3 2 5 2 3 2 6
mixture is warmed to room temperature.
A method recommended5,6 for the preparation of small quantities of diborane involves
the reaction of potassium tetrahydroborate with 85% ortho-phosphoric acid.
2KBH + 3H P0 -> B H + 2H + 2KH P0
4 3 4 2 6 2 2 4
The reaction is carried out in vacuo at room temperature and the volatile products led
into a cold trap at -196°C. High-purity diborane, m.p. -165°, b.p. -90°C, in 40-50%
yield, is produced.
Better yields of B H result from the reaction of a metal tetrahydroborate with halides
2 6
of bismuth, tin, mercury or antimony.7 For example, up to 98% yield has been achieved
by the reaction between sodium tetrahydroborate and mercury(I) chloride, both in solution
in 'diglyme' (the dimethylether of diethyleneglycol), at room temperature.
2NaBH + Hg Cl -> 2Hg + 2NaCl + B H + H
4 2 2 2 6 2
Halides of other metals are less efficient. When SbCl is used, some stibine contaminates
3
the product and with PbCl only 20% yield is achieved.
2
The synthesis of diborane directly from readily available materials would have clear
advantages over earlier syntheses, such as those requiring the use of highly reactive boron
halides. B H has been made8 from boron trioxide, B 0 , by reaction with hydrogen under
2 6 2 3
high pressure at 175°C in the presence of aluminium and aluminium chloride. Conversions
of up to 50% have been achieved. Another synthetic route is the low-pressure hydrogénation
of boron monoxide, B 0 , at 1200°C.9 B 0 is made in the pure state by the reaction
2 2 2 2
between Ti0 or Zr0 and boron carbide, B C, in vacuo. The major product of this
2 2 4
hydrogénation reaction is B H in yields of up to 20%.
2 6
Reactions
Diborane, like the other boranes, is a highly reactive compound. It is thermally unstable
and decomposes rapidly at 100°C to give hydrogen and higher boranes. It is readily oxidized
by air but does not, unlike some of the other boranes, inflame spontaneously in air.
2
DIBORANE
When B H comes into contact with water, immediate hydrolysis occurs.
2 6
B H + 6H 0 -> 2H3BO3 + 6H
2 6 2 2
One or two hydrogen atoms can be directly replaced by halogen.
B H + Br -> B H Br + HBr
2 6 2 2 5
B H C1 can be made in a similar way, but it is difficult to isolate because of some dispro-
2 5
portionation into BC1 and B H at room temperature. The reaction between B H and
3 2 6 2 6
BCI3 gives,10 at room temperature, a mixture of BHC1 and B H C1, but at 100°C the
2 2 5
product is almost exclusively BHC1 . This indicates that the replacement of a second hy­
2
drogen atom by halogen causes a cleavage of the B H molecule into two parts.
2 6
There is no evidence for the independent existence of the borane radical, BH , although
3
in many of its reactions diborane behaves as two BH fragments. For example, B H
3 2 6
reacts with CO under high pressure at 100°C to give borane carbonyl, BH · CO. This
3
adduct dissociates into CO and B H at ordinary temperature and pressure. Reaction of
2 6
B H with excess trimethylamine gives trimethylamine borane, (CH ) N · BH . In this and
2 6 3 3 3
the carbonyl, the BH fragment acts as a Lewis acid in combination with an electron donor
3
(Lewis base) molecule. In both reactions, the B H molecule undergoes a symmetrical
2 6
cleavage. In the case of the reaction with trimethylamine, an intermediate of composition
(CH ) N · B H has been isolated,11 indicating that the diborane molecule is coordinated
3 3 2 6
by one donor molecule before cleavage occurs.
μ-Dimethylaminodiborane, (CH ) NB H , has been made12 by the interaction of
3 2 2 5
sodium dimethylamidotrihydroborate, Na(CH ) NBH , with diborane in diglyme solution.
3 2 3
(CH ) NBH - + B H -> //-(CH ) NB H + BH "
3 2 3 2 6 3 2 2 5 4
The structure of this derivative has been established by electron diffraction. Like B H ,
2 6
it has a bridge structure (see below) in which a dimethylamino group has replaced one of
the bridge hydrogen atoms of diborane.
Different products arise from the reaction between B H and ammonia, depending on
2 6
the experimental conditions. At — 120°C, reaction with excess ammonia forms a salt-like
solid of stoichiometry, B H · 2NH . This contains BH _ ions and is formulated as an
2 6 3 4
ionic compound, [H B(NH ) ]+[BH ]_. It results from the unsymmetrical cleavage of the
2 3 2 4
B H molecule.
2 6
When diborane is heated with excess ammonia, boron imide and, finally, boron nitride
are produced. If B H and NH in a mole ration of 1:2 are heated together, or if
2 6 3
[H B(NH ) ]+[BH ]- is heated above 200°C, the cyclic compound borazine, B N H ,
2 3 2 4 3 3 6
is formed.
The B H molecule is also split unsymmetrically by dimethylsulphoxide.13 Their reaction
2 6
in a mole ratio of B H :(CH ) SO = 1:2 in dichloromethane solution at -78°C gives
2 6 3 2
BH [OS(CH ) ] +[BH ]-.
2 3 2 2 4
Reactions between diborane and metal alkyls produces metal tetrahydroborates. Thus
B H and ethyllithium react at room temperature to form lithium tetrahydroborate,
2 6
LiBH .
4
2LiC H + 2B H -> 2LiBH + (C H ) B H
2 5 2 6 4 2 5 2 2 4
The BH ~ ion can be regarded as formed by combination between the Lewis acid, BH ,
4 3
and the base, H~. Sodium and lithium tetrahydroborates are used as selective reducing
agents for organic compounds. They will reduce carbonyl groups in aldehydes and ketones
to alcohols, but not those in acids, esters or acid anhydrides.
3
SYNTHESES AND PHYSICAL STUDIES OF INORGANIC COMPOUNDS
The great reactivity of B H and, in many cases, the products of its reactions are consis­
2 6
tent with a cleavage of the molecule by a process which requires comparatively little energy.
The chemistry of diborane can be most readily understood from a knowledge of its struc­
ture, which has now been firmly established in a series of important investigations using a
variety of physical techniques.
Structure
Structures which have been proposed from time to time for molecular diborane are
essentially of three kinds:14 (a) an analogue of ethane, H B · BH ; (b) a bridged molecule,
3 3
H B(H )BH , in which two of the hydrogens act as bridge atoms between the borons,
2 2 2
the remaining four being equivalent to one another but not to the bridge hydrogens; (c) an
ionic structure, [H B = BH ][H+] , in which the two different kinds of hydrogen are bound
2 2 2
respectively by covalent and ionic bonds.
Early X-ray15 and electron diffraction16 results were interpreted in terms of an ethane­
like structure for diborane. Interatomic distances found by electron diffraction were:
B, B = 0-186 nm (cf. C, C in ethane = 0-154 nm) and B, H = 0-127 nm.
The experimental data are equally well in agreement with a bridged structure17,18 so no
conclusive structural assignment can be made from these results.
The vibrational spectrum of diborane (see below) clearly favoured a bridged structure,
so it was necessary for the electron diffraction investigation to be repeated.19 In this study,
the most satisfactory agreement between theoretical radial distribution curves and those
calculated from a visual examination of the electron diffraction photographs was found
for the bridged model (b) (Fig. 1.1). Table I summarizes the data of Hedberg and Scho-
FIG. 1.1. The bridged structure of diborane.
maker and the more recent results of Bartell and Carroll20 on B H and B D . In the latter,
2 6 2 6
electron diffraction photographs were examined microphotometrically to obtain objective
measurements of intensity. There is close agreement between the molecular parameters
found from these two studies.
In the solid state between its melting-point, —165 °C, and liquid nitrogen temperature,
diborane is believed to be capable of existing in three phases, oc, ß and ω. The oc- and ß-forms
were first characterized by their X-ray powder patterns21 and were found to co-exist over
a considerable range of temperature. It is therefore likely that the two phases are closely
related structurally.
ίχ-Diborane is deposited from diborane vapour at —268-8 °Cand transforms slowly into
ß-diborane at c. — 213 °C. β-Diborane has also been made by deposition from the vapour state
at -196 °C followed by annealing above -183 °C.
4
DIBORANE
TABLE I
Molecular Parameters for B H and B D from Electron Diffraction Data
2 6 2 6
B H B D
2 6 2 6
Hedberg and Bartell and Bartell and
Schomaker19 Carroll20 Carroll20
B, B (nm) 0177 ±00013 01775 ± 00003 01771 ± 00003
B, H(nm) 01187 ±0003 01196 ±00007 01198 ±00006
t
B, H(nm) 01334 ± 00027
H-B6-H υ0 1u3n 39_+ 0000000026 0\) 1I3ÔÔ3Ô3 +_ Q0.0Q000Q2 4
t f
angle (deg.) 121-5 ± 7
H -B-B
b
angle (deg.) 48.5 48-5 ± 0-15 48-4 ± 0-15
B-B-H,
angle (deg.) 120-6 ± 0-9 119-4 ±09
The X-ray diffraction analysis of β-diborane has been carried out22 and this shows the
crystal is monoclinic with the molecular parameters :
B, B = 01776 ± 0001 nm; B, H; = 0-108 ± 0-002nm;
B, H/ = 0110 ± 0002nm; B, W = 0-123 ± 0-002 nm;
b
B, H£' = 0125 ± 0-002 nm; Hi-fi-Pi;' = 124 ± 1°;
H^-B-H;' =90+1°.
The B, B distance agrees well with the values determined by electron diffraction. The
B, H bond lengths are systematically about 0-01 nm shorter than the electron diffraction
values. This shortening appears to be a general phenomenon23 associated with parameters
determined by X-ray crystallography and so the electron diffraction data may still be
regarded as reliable.
The oc- and ß-phases have been differentiated by differences in their polarized infrared
spectra.24 jS-Diborane shows a very intense component of absorption on the low-frequency
side of the band at 1600 cm-1, but there is no corresponding feature in the spectrum of
the ix-form. From a detailed study of the intensity of the infrared absorption bands of the
<%-form as a function of polarizer setting, it has been concluded that #-diborane is monoclinic.
Its detailed structure has not yet been published.
The co-form has been obtained as single crystals, but within hours or a few days these
transform spontaneously into polycrystalline masses. Its structure is unknown.
Vibrational Spectra
Raman data25 on liquid diborane and the infrared spectrum26 of the gas were at one
time regarded as consistent with an ethane-like molecular structure. It was, however,
recognized that the infrared spectrum of diborane was much more complex than that of
ethane, but this was erroneously attributed to the existence of some low-lying excited
electronic levels which could participate in some of the vibrational transitions.
Ethane has D or D symmetry: the bridge structure of diborane, assuming the boron
3d 3h
valencies are tetrahedrally disposed, belongs (like ethylene) to the symmetry group D .
2h
Bell and Longuet-Higgins27 have calculated, by normal coordinate analysis, the frequen-
5