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THE GROUP 13 METALS
ALUMINIUM, GALLIUM,
INDIUM AND THALLIUM
Editors
SIMON ALDRIDGE, Department of Chemistry, University of Oxford, UK
ANTHONY J. DOWNS, Department of Chemistry, University of Oxford, UK
The last two decades have seen a renaissance of interest in the chemistry of the Main Group elements. In
particular, research on the metals of Group 13 (aluminium, gallium, indium and thallium) has led to the synthesis
and isolation of some novel and unusual compounds with implications, for example, in organometallic synthesis,
in new materials development, and in biological, medical, and environmental systems.
The Group 13 Metals Aluminium, Gallium, Indium and Thallium aims to cover new facts, developments, and
applications in the context of more general patterns of physical and chemical behaviour, with an eye to both the
homogeneity and heterogeneity displayed by the elements. Particular attention is paid to the main growth areas,
including the chemistry of lower formal oxidation states, cluster chemistry, solid oxides and hydroxides, methods
of formation of III-V and related compounds, the biological significance of Group 13 metal compounds, and the
growing importance of the metals and their compounds in the mediation of organic reactions. Chapters cover:
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general features of the Group 13 elements
Group 13 metals in the +3 oxidation state: simple inorganic compounds
formal oxidation state +3: organometallic chemistry
formal oxidation state +2: metal-metal bonded vs. mononuclear derivatives
Group 13 metals in the +1 oxidation state
mixed or intermediate valence Group 13 metal compounds
aluminium and gallium clusters: metalloid clusters and their relation to the bulk phases, to naked
clusters, and to nanoscaled materials
simple and mixed metal oxides and hydroxides: solids with extended structures of different
dimensionalities and porosities
coordination and solution chemistry of the metals: biological, medical, and environmental relevance
III-V and related semiconductor materials
Group 13 metal-mediated organic reactions
The Group 13 Metals Aluminium, Gallium, Indium and Thallium provides a detailed, wide-ranging, and
up-to-date review of the chemistry of this important group of metals. It will find a place on the bookshelves of
researchers and other experienced scientists as well as students working in inorganic, organic, organometallic,
and materials chemistry.
THE GROUP 13 METALS ALUMINIUM,
GALLIUM, INDIUM AND THALLIUM
Chemical Patterns and Peculiarities
Editors
ALDRIDGE
DOWNS
THE GROUP 13 METALS
ALUMINIUM, GALLIUM,
INDIUM AND THALLIUM
Chemical Patterns and Peculiarities
Editors | SIMON ALDRIDGE | ANTHONY J. DOWNS
The Group 13 Metals Aluminium, Gallium, Indium
and Thallium: Chemical Patterns and Peculiarities
The Group 13 Metals Aluminium, Gallium,
Indium and Thallium: Chemical Patterns
and Peculiarities
Editors
SIMON ALDRIDGE
Inorganic Chemistry Laboratory, University of Oxford, Oxford, UK
ANTHONY J. DOWNS
Inorganic Chemistry Laboratory, University of Oxford, Oxford, UK
This edition first published 2011
Ó 2011 John Wiley & Sons, Ltd.
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Library of Congress Cataloging-in-Publication Data
The Group 13 Metals Aluminium, Gallium, Indium and Thallium: Chemical Patterns and Peculiarities / editors Simon Aldridge,
Anthony J. Downs.
p. cm.
Includes bibliographical references and index.
ISBN 978-0-470-68191-6 (cloth)
1. Group 13 elements. I. Aldridge, Simon II. Downs, Anthony J., 1936- III.
Title: The Group 13 Metals Aluminium, Gallium, Indium and Thallium: Chemical Patterns and Peculiarities.
QD466.C495 2011
546’.67–dc22
2010034290
A catalogue record for this book is available from the British Library.
ISBN 9780470681916
e-book - 9780470976555
o-book - 9780470976562
e-pub - 9780470976685
Set in 9.5/11.5pt, Times by Thomson Digital, Noida
Printed in Singapore by Fabulous Printers Pte Ltd.
Contents
Preface
ix
List of Contributors
xi
1 New Light on the Chemistry of the Group 13 Metals
Anthony J. Downs and Hans-Jörg Himmel
1.1 Reprise of the General Features of Group 13 Elements
1.2 Developments in Methodology
1.3 Redox Chemistry of the Group 13 Metals: Access to Oxidation States Lower than +3
1.4 Bonding Aspects
1.5 Solid Compounds with Specific Electronic, Structural or Other Properties
1.6 Coordination Chemistry of M(III) Compounds
1.7 Mediation of Organic Transformations by Group 13 Metal Compounds
References
2 The Chemistry of the Group 13 Metals in the +3 Oxidation State: Simple
Inorganic Compounds
Simon Aldridge
2.1 Introduction
2.2 Hydrides
2.3 Halides and Pseudo-Halides
2.4 Oxides and Oxo- Derivatives
2.5 Chalcogenides and Chalco-Derivatives
2.6 Compounds with Bonds to Group 15 Atoms
References
3 Formal Oxidation State +3: Organometallic Chemistry
Simon Aldridge, Anthony J. Downs and Deborah L. Kays
3.1 Introduction
3.2 Organo Derivatives with a Metal–Carbon Primary Framework
3.3 Derivatives with Bonds to Group 15 Elements
3.4 Derivatives with Bonds to Group 16 Elements
3.5 Organometal Halides
3.6 Organometal Hydrides
3.7 d-Block and f-Block Compounds with Organo-Group 13 Metal(III) Fragments
References
1
1
6
25
28
38
43
57
60
75
75
76
99
107
118
122
132
148
148
151
168
192
210
215
220
227
vi
4
5
6
7
8
Contents
Formal Oxidation State +2: Metal–Metal Bonded Versus Mononuclear Derivatives
Werner Uhl and Marcus Layh
4.1 Introduction
4.2 Subhalides Containing M–M Bonds
4.3 Homoleptic Chalcogen Compounds
4.4 Homoleptic Dielement Compounds with Pnicogen Atoms Coordinated
to the M–M Bonds
4.5 Heteroleptic Compounds Containing Donor Atoms of Groups 15 to 17
4.6 Homoleptic Dinuclear Organoelement(II) Compounds
4.7 Heteroleptic Organoelement(II) Compounds
4.8 Mononuclear Element(II) Compounds
References
246
The Chemistry of the Group 13 Metals in the +1 Oxidation State
Cameron Jones and Andreas Stasch
5.1 Introduction
5.2 Aluminium
5.3 Gallium
5.4 Indium
5.5 Thallium
References
285
Mixed or Intermediate Valence Group 13 Metal Compounds
Benjamin F. T. Cooper and Charles L. B. Macdonald
6.1 Mixed Valency
6.2 Halides
6.3 Arene-stabilised Mixed Valent Species
6.4 Chalcogenide and Other Non-Halide Salts
6.5 Discretely Bonded Systems
6.6 Donor–Acceptor Compounds
6.7 Conclusions
References
342
Aluminium and Gallium Clusters: Metalloid Clusters and their Relationship to the
Bulk Phases, to Naked Clusters and to Nanoscaled Materials
Hansgeorg Schnöckel and Andreas Schnepf
7.1 Introduction
7.2 Explanations of Special Terms
7.3 The Naked Al13 Cluster
7.4 Metalloid Al/Ga Clusters
7.5 Interactions between Cluster Species within the Crystal
7.6 Summary and Outlook
References
Simple and Mixed Metal Oxides and Hydroxides: Solids with Extended Structures
of Different Dimensionalities and Porosities
Andrew M. Fogg
8.1 Introduction
8.2 The Parent Oxides and Hydroxides
246
247
250
251
254
257
265
274
277
285
286
298
310
317
329
342
346
353
363
368
390
396
397
402
402
404
407
416
471
480
482
488
488
489
Contents
8.3 Layered Materials
8.4 Framework Materials
References
9 Coordination and Solution Chemistry of the Metals: Biological, Medical
and Environmental Relevance
Penelope J. Brothers and Christy E. Ruggiero
9.1 Introduction
9.2 Hydrides
9.3 Halides
9.4 Group 13 Complexes of N-donor Ligands
9.5 Complexes of the Monovalent M(I) Group 13 Metals
9.6 Divalent M(II) Complexes
9.7 Chemistry Relevant to Environmental and Biological Systems
9.8 Environmental Abundances, Uses and Sources
9.9 Biological Systems
References
10
11
vii
494
502
509
519
519
521
542
548
557
568
569
572
576
593
III–V and Related Semiconductor Materials
Mohammad Azad Malik and Paul O’Brien
10.1 Introduction
10.2 III–Nitrides
10.3 III–Phosphides
10.4 III–Arsenides
10.5 III–Antimonides
References
612
Group 13 Metal-Mediated Organic Reactions
Samuel Dagorne and Stéphane Bellemin-Laponnaz
11.1 Aluminium
11.2 Gallium
11.3 Indium
11.4 Thallium
References
654
Index
612
619
627
631
636
645
654
675
681
688
692
701
Preface
“Evolution . . . is – a change from an indefinite, incoherent
homogeneity, to a definite, coherent heterogeneity.”
Herbert Spencer, First Principles, 1862, Chapter 16.
It was homogeneity, modulated by predictable variations, that enabled Mendeleev in 1870 to anticipate with
celebrated fidelity the properties of gallium, then the missing link in what we now call Group 13. While the kinship of
the elements has never been in doubt with the evolution of our knowledge of their chemistry, it is the peculiarities
which have more often left their mark – for example, the discovery of a wide variety of compounds in which the
Group 13 element M assumes a formal oxidation state other than þ 3; the identification of diverse compounds with
MM-bonded frameworks; the finding of catalytic activity in compounds that varies radically according to the
nature of M; the mediation of organic reactions in ways that differ widely from one Group 13 element to another; and
the development of solids with extended structures and absorption properties more or less specific to a particular
member of the Group.
So eccentric is boron, the non-metal with its propensity for forming strong localised or delocalised covalent
bonds, that it is most aptly separated from the other members of Group 13. Its chemistry has been comprehensively
reviewed, for example in volumes of the Gmelin Handbook up to the later years of the 20th century and in
numerous other books. This contrasts with the generally meagre and piecemeal treatment of the heavier members
of the Group, all of them metals forming a more closely knit family, but each with its own distinctive personality.
For none of these does the Gmelin Handbook offer more than a specific volume or two dating beyond the first half
of the 20th century.
This book seeks to remedy the imbalance with a definitive, wide-ranging and up-to-date review of major aspects
of the chemistry of these elements. It has two obvious reference points. The first is the book entitled The Chemistry
of Aluminium, Gallium, Indium and Thallium, written by Wade and Banister, published first in 1973 as part of
Comprehensive Inorganic Chemistry, and appearing as a separate volume in 1975. The second is a book bearing
the same title and edited by one of us and that first saw the light of day roughly two decades later (1993). With the
passage of nearly two more decades that have seen a wealth of activity, it seemed to us timely once again to take
stock. This we have sought to do not as a mere catalogue, but within a framework designed to present a wider
picture that places new facts, developments and applications in the context of more general patterns of physical
and chemical behaviour that is, with an eye to both the homogeneity and heterogeneity displayed by the
elements. The various chapters have been written by members of an international team of authors selected as
experts with practising research experience in the particular field under review.
Chapter 1 sets the scene with an outline of the areas of Group 13 metal chemistry that have seen most progress in
the past two decades. Chapters 2–7 are organised according to the formal oxidation state of the metal, a concept
which, for all its imperfections, is likely to be most widely appreciated. After treatments of first the inorganic and
then the organic derivatives of the metals in the dominant þ 3 state in Chapters 2 and 3, respectively, Chapter 4
x Preface
addresses the þ 2 state with its prevailing theme of MM bonding, while Chapter 5 is concerned with the þ 1
state, which has gained hugely in significance in recent years. Chapter 6 is devoted to compounds in which M
occurs in more than one oxidation state, as exemplified by the classical case of GaIGaIIICl4. Mixed oxidation states
are also a feature of many of the remarkable cluster compounds, including so called metalloid clusters, that have
lately caused such a storm, particularly through the pioneering research of the Karlsruhe group led by Schn€
ockel.
An authentic and challenging account of this area is presented in Chapter 7. There follows in Chapter 8 a review of
simple and mixed Group 13 metal oxides and hydroxides including zeolites, detailing the extended structures of
different dimensionalities and porosities that they form in the solid state. If this is preoccupied with the solid state,
the coordination chemistry of the metals, as described in Chapter 9, is intimately related to their behaviour in
solution, with its relevance in biology, medicine, and the environment. The solid state is again to the fore in
Chapter 10 which deals with III-V and related semiconductor materials. Last but far from least, the role of the
Group 13 metals and their compounds as reagents or mediators in organic synthesis is taken up in Chapter 11.
A book on this scale cannot possibly emulate Gmelin. Even with references to some 5000 original papers, books,
and review articles, some published as recently as 2010, it makes no pretence of being comprehensive. All the
authors have been given licence to treat their subjects as they see fit. We are well aware that some compounds and
some topics have as a result received little or no attention. Such is the case, for example, with Zintl and related phases
containing more or less negatively charged clusters and networks of Group 13 metal atoms. Nor is the bonding in
Group 13 metal compounds made the exclusive preserve of any one chapter. We are aware too of the overlap existing
between some of the chapters, all having been written as self-sufficient accounts. While it may mean that our
coverage is not everywhere as efficient as it might be, we dare to hope that there are compensations from the different
perspectives, as well as the cross-linking between chapters, that will actually help to broaden any appeal the book
may have.
In aiming for a clear and structured treatment with the bare minimum of specialist jargon and annoying acronyms,
we have tried also to achieve an accessible style in a text that is generally readable by non-specialist no less than
specialist readers. We see the book therefore not just as a contemporary source-book on Group 13 metal chemistry,
but as a monograph that can be read with some profit by scientists in different walks of life. It is of course directed
mainly at chemists, but includes sections likely to be of interest to physicists, biochemists, and materials,
environmental and industrial scientists.
S.A.
A.J.D.
List of Contributors
Simon Aldridge, Inorganic Chemistry Laboratory, University of Oxford, Oxford, UK
Mohammad Azad Malik, School of Chemistry, University of Manchester, Manchester, UK
Stéphane Bellemin-Laponnaz, IPCMS, CNRS, Strasbourg, France
Penelope J. Brothers, Department of Chemistry, The University of Auckland, Auckland, New Zealand
Benjamin F. T. Cooper, Department of Chemistry and Biochemistry, University of Windsor, Windsor,
Ontario, Canada
Anthony J. Downs, Inorganic Chemistry Laboratory, University of Oxford, Oxford, UK
Samuel Dagorne, Institut de Chimie, Université de Strasbourg, Strasbourg, France
Andrew M. Fogg, Department of Chemistry, University of Liverpool, Liverpool, UK
Hans-Jörg Himmel, Anorganisch-Chemisches Institut, Ruprecht-Karls-Universität Heidelberg,
Heidelberg, Germany
Cameron Jones, School of Chemistry, Monash University, Clayton, Victoria, Australia
Deborah L. Kays, School of Chemistry, University of Nottingham, Nottingham, UK
Marcus Layh, Institut für Anorganische und Analytische Chemie, Westfalische Wilhelms-Universität
Münster, Münster, Germany
Charles L. B. Macdonald, Department of Chemistry and Biochemistry, University of Windsor, Windsor,
Ontario, Canada
Paul O’Brien, School of Chemistry, University of Manchester, Manchester, UK
Christy E. Ruggiero, Los Alamos National Laboratory, Los Alamos, NM USA
Andreas Schnepf, Institut für Anorganische Chemie, Universität Duisburg-Essen, Essen, Germany
Hansgeorg Schnöckel, Institut für Anorganische Chemie, Karlsruher Institut für Technologie (KIT),
Karlsruhe, Germany
Andreas Stasch, School of Chemistry, Monash University, Clayton, Victoria, Australia
Werner Uhl, Institut für Anorganische und Analytische Chemie, Westfalische Wilhelms-Universität
Münster, Münster, Germany
1
New Light on the Chemistry
of the Group 13 Metals
Anthony J. Downs1 and Hans-J€org Himmel2
1
Inorganic Chemistry Laboratory, University of Oxford, Oxford, UK
Anorganisch-Chemisches Institut, Ruprecht-Karls-Universit€
at Heidelberg, Heidelberg, Germany
2
A little learning is a dangerous thing;
Drink deep, or taste not the Pierian spring:
There shallow draughts intoxicate the brain,
And drinking largely sobers us again.
Alexander Pope, An Essay on Criticism, 1711
1.1 Reprise of the General Features of Group 13 Elements
First impressions may seize upon the commonality of the Group 13 elements – boron, aluminium, gallium, indium
and thallium – arising out of the common configuration ns2np1 shared by the valence electrons in the ground state of
each of the atoms. Witness, for example, the dominance of the formal oxidation state þ 3 and the acceptor properties
that characterise the resulting derivatives, arising partly from the positive charge, partly from the inability of the
Group 13 atom effectively to engage all its valence orbitals in bonding. There is harmony in the variation of
properties dictated by the generally increasing atomic size and decreasing hold of the nucleus on the valence
electrons as the atomic number increases from boron to thallium. But there is also counterpoint, reflecting the
discontinuous build-up of the Periodic Table. Hence, each member of the Group has its own individual personality,
‘with quirks of character and not always evident dispositions’,1 manifesting the infinite variety of the Periodic
Kingdom that is perhaps the most remarkable phenomenon in the universe.
Quirkiest of the Group 13 elements is undoubtedly boron.2–11 As revealed in the numerical properties summarised
in Table 1.1, the boron atom is disproportionately smaller and its valence electrons are more tightly held in relation to
The Group 13 Metals Aluminium, Gallium, Indium and Thallium: Chemical Patterns and Peculiarities. Edited by Simon Aldridge and Anthony J. Downs
2011 John Wiley & Sons, Ltd
Property
B
(i) Properties of the isolated atom
Atomic number
Naturally occurring isotopes
5
Relative atomic mass
(12 C ¼ 12.0000)
Ground-state electron
configuration (term)
Ga
13
Al (100%)
69
B (19.9%)
11
B (80.1%)
10.811(7)
26.9815386(8)
2
P1=
2
[Ne]3s23p1
2
In
31
Ga (60.108%)
71
Ga (39.892%)
69.723(1)
27
10
[He]2s22p1
Al
P1=
2
[Ar]3d104s24p1
2
P1=
2
Tl
49
81
In (4.29%)
115
In (95.71%)
114.818(3)
203
[Kr]4d105s25p1
2
P1=
[Xe]
10 2
1
4f145d
6s 6p
2
P1=
113
Tl (29.524%)
Tl (70.476%)
204.3833(2)
205
2
1
Ionisation energies (kJ mol )
M ! Mþ
Mþ ! M2þ
M2þ ! M3þ
M3þ ! M4þ
Thermal neutron capture crosssection (barns)
Nuclei accessible to NMR
measurements (nuclear spin I)a
(ii) Properties of the bound atom
Electronegativity, x
Pauling scale
Allred scale
Sanderson scales
Pearson scale (eV) Atomic (metallic) radius (A)
Singlebond
covalent radius for
MIII (A)
van der Waals radius (A)
2
800.637
2427.07
3659.75
25 025.9
10
B 3840, 11 B
0.005
½10 Bð3Þ; 11 B (3/2)
577.539
1816.68
2744.78
11 577.46
0.230
578.844
1979.41
2963
6175
69
Ga 1.68, 71 Ga
4.7
[69 Ga (3/2)], 71 Ga
(3/2)
558.299
1820.71
2704
5210
113
In 12, 115 In 205
589.351
1971.03
2878
(4900)
203
Tl 11, 205 Tl 0.11
[113 In (9/2)],
(9/2)
[203 Tl (1/2)].
(1/2)
2.04
1.61
1.81
1.78
2.01
1.88
1.53 [BI], 2.28
[BIII]
4.29
0.80–0.90
0.88
1.47
1.54
0.84 [AlI], 1.71
[AlIII]
3.23
1.431
1.25
1.82
2.10
0.86 [GaI], 2.42
[GaIII]
3.2
1.22–1.40
1.25
1.49
1.88
0.71 [InI], 2.14
[InIII]
3.1
1.62–1.68
1.50
1.62[TlI], 2.04
[TlIII]
1.44
1.96
0.99 [TlI], 2.25
[TlIII]
3.2
1.704 (a-form)
1.55
2.08
2.05
1.90
1.90
2.00
27
Al (5/2)
115
In
205
Tl
2 New Light on the Chemistry of the Group 13 Metals
Table 1.1 Some properties of the Group 13 elements boron, aluminium, gallium, indium and thallium12–23
Covalent bonds in MIII
compounds:
length (A) [mean bond
enthalpy (kJ mol1)]
{coordination number}
M–H
M–Cl
1.74b [444] {3}
M–O
1.37b [559] {3}
1.48(4)d {4}
1.94(5)d {4}
1.58b [376] {3}
1.63(4)d {4}
1.77(6)d [585] {4}
1.89(6)d {6}
2.25(7)d {4}
1.96b [280] {3}
1.97(3)d {4}
1.590 [290]
1.232 [342]
1.263 [732]
1.715 [427]
1.888 [391]
2.131 [361]
a-B, tetragonal
b-B, tetragonal
a-B,
rhombohedral
b-B,
rhombohedral
M–S
M–Cg
Covalent bonds in diatomic
molecules: length (A) [D298
(kJ mol1)]
M–M
M–H
M–F
M–Cl
M–Br
M–I
(iii) Properties of the elements
Crystal structure
1.58c {3}
1.56(11)d [282] {4}
1.63b [591] {3}
1.815(46)d {6}
2.07b [426] {3}
2.14(6)d {4}
1.57c {3}
1.51(13)d [260] {4}
1.71b [602] {3}
1.95(7)d {6}
2.10b [363] {3}
2.20(8)d {4}
1.74c {3}
1.79(17)d [225] {4}
2.10(4)d [ca.525]{6}
1.76c [181] {3}
2.28(13)e [460] {8}
2.43(5)d [368] {4}
2.52(6)d {6}
1.91(8)d [ca. 430]
{4} 1.96(5)d {6}
2.26(5)d {4}
1.97b [245] {3}
1.985(38)d {4}
2.26b [327] {3}
2.40(5)d {4}
2.48(10)d {6}
2.15(10)d [ca. 360]
{4} 2.18(10)d {6}
2.48(7)d {4}
2.16b [162] {3}
2.18(4)d {4}
2.466 [133]
1.645 [288]
1.654 [675]
2.130 [502]
2.295 [429]
2.537 [370]
2.70 [115]
1.662 [265]
1.774 [584]
2.202 [463]
2.352 [416]
2.575 [334]
2.97 [78]
1.836 [243]
1.985 [507]
2.401 [430]
2.543 [388]
2.754 [331]
3.41 [59]
1.873 [195]
2.084 [439]
2.485 [373]
2.618 [331]
2.814 [285]
a-Al, f.c.c.
High pressure form
b-Al, hexagonal
a-Ga,
orthorhombic
b-Ga, monoclinic
g-Ga, rhombic
d-Ga,
rhombohedral
Face-centred
tetragonal
a-Tl, hexagonal
b-Tl, cubic
g-Tl, f.c.c.
T(a ! b) 503 K
2.38(11)d [375] {6}
2.48(2)f {4}
2.21b [125] {3}
2.17(4)d {4}
(continued)
Reprise of the General Features of Group 13 Elements 3
M–F
1.19b [377] {3}
1.06(9)d {4}
1.31b [646] {3}
(Continued )
Property
Melting point (K)
Normal boiling point (K)
DHo atomisation at 298.15 K
(kJ mol1)
Density (kg m3)
Electrical resistivity at 273 K (Wm)
(iv) Cationic and redox behaviour
Ionic radius for six-fold
coordination
M3þ (A
)i
þ M (A)
DfHo[M3þ (g)] (kJ mol1)
DfHo[Mþ (g)] (kJ mol1)
Thermodynamic properties for
aqueous species (kJ mol1)
DfGo[M3þ (aq)] std state,
m¼1
DfHo[M3þ (aq)] std state, m ¼ 1
DhydrationGo[M3þ (g)], single
ion
DhydrationHo[M3þ (g)], single
ion
B
Other polymorphs
reported and
partially characterised, e.g.
cubic, rhombic,
monoclinic and
hexagonal forms
2348
4273
565
2340 (b-rhomb),
293 K
1.8 104
(0.27)
–
7468.8
1369.6
Al
Ga
In
Tl
933.47
2792
330.9
T(g ! a) 238 K
High pressure
forms:
Ga(II), b.c.c.
Ga(III), b.c.
tetragonal
Ga(IV), f.c.c.h
302.92
2477
271.96
429.75
2345
243
577
1746
182.2
2698, 293 K
5907, 293 K
7310, 298 K
11 850, 293 K
2.417 108
13.6 108
8.0 108
15 108
0.535
(ca. 1.00)
5484.0
910.09
0.620
1.13
5816
861.9
0.800
1.32
5345.3
807.8
0.885
1.50
5639.2
777.73
967.7 [B(OH)3]
485.0
159.0
98.0
þ 214.6
1072.8 [B(OH)3]
n.a.
531.0
4540
211.7
4550
105.0
4020
þ 196.6
4000
n.a.
4680
4690
4110
4110
4 New Light on the Chemistry of the Group 13 Metals
Table 1.1
Standard reduction potentials,
Eo (V):
MIII(aq) þ 3e ! M(s),
aHþ ¼ 1
MIII(aq) þ 3e ! M(s),
a
OH ¼ 1
I
M (aq) þ e ! M(s), aHþ ¼ 1
MIII(aq) þ 2e ! MI(aq),
aHþ ¼ 1
(v) Environmental properties
Abundances
Continental crust (ppb)
Ocean (ppb)
Chief ores and sources
a
1.662
0.549
0.3382
1.79
{[OB(OH)2]}
—
—
2.328
{[Al(OH)4]}
( þ 0.3)
(2.7)
1.219
{[OGa(OH)2]}
0.2
0.72
1.007
{[In(OH)4]}
0.129
0.443
104
4.4 103
Borax and kernite,
Na2[B4O5
(OH)4] xH2O;
colemanite,
Ca2[B3O4
(OH)3]2 2H2O
8.2 107
2
Bauxite, found as
boehmite and
diaspore, AlO
(OH), and gibbsite and hydrargillite, Al(OH)3
1.8 104
3 102
Occurs up to 1% in
other minerals;
recovered as a
by-product of
zinc and aluminium refining
ca. 2 106 (B2O3)
270 106 (as
B2O3)
ca. 4 107 (Al)
6 109
ca. 100 (Ga)
—
2.5 102
3 104
Occurs up to 1% in
zinc and lead
sulphide ores;
obtained as a byproduct of zinc
and lead
smelting
ca. 450 (In)
>1500
þ 0.741
0.163 {Tl2O3}
0.334
þ 1.279
8.5 102
1.3 102
Rare; dispersed in
potash, feldspar
and pollucite; byproduct of zinc
and lead smelting
and H2SO4
manufacture
ca. 30 (Tl)
—
Nuclei enclosed in square brackets are seldom used in NMR studies.
Refers to a specific gaseous molecule, that is MH3, MF3, MCl3, M(OH)3 and M(CH3)3 (M ¼ B, Al, Ga, In or Tl).
Calculated value (Ref. 21, p. 3313).
d
Average length of a terminal bond determined for numerous compounds by single crystal X-ray diffraction studies;20 numbers in parentheses are the standard deviations of the last digits.
e
Refers to a single crystal structure, that of TlF3 (Ref. 12, p. 3).
f
Refers to a single crystal structure containing 3 terminal Tl–S bonds to non-chelating ligands.24
g
C is an alkyl carbon atom.
h
Reference 25, and references cited therein.
i
Reference 26.
b
c
Reprise of the General Features of Group 13 Elements 5
World production (tonnes year1)
Reserves (tonnes)
0.870
6 New Light on the Chemistry of the Group 13 Metals
the atoms of its vertical neighbours. Accordingly, the element itself is not a metal but a semiconductor with several
hard and refractory allotropic forms characterised by unique and elaborate structures based on the B12 icosahedron.
Here and in metal borides and boron hydrides, too, evidence is found of boron’s propensity to form polyboron
branched and unbranched chains, cages, planar networks and three-dimensional arrays. These typify an extensive
and unusual type of covalent (molecular or macromolecular) chemistry in which multicentre bonding is one of the
most distinctive features. Moreover, the relatively compact 2p orbitals of the boron atom share with those of the
carbon atom the ability to engage in relatively efficient p-type interactions, providing a mechanism for supplementing substantially the bonding to electron-rich centres. Such interactions subscribe to the relative abundance and
stability of three-coordinate environments for the boron atom, as in the trihalides, boric acid, amidoboranes and
borazine, H3B3N3H3, formally analogous to benzene. Another feature peculiar to boron is the relative closeness in
energy of the valence 2s and 2p orbitals, which favours a major contribution from the 2s orbital to the bonding of
boron(III) compounds. This, combined with the inherent strengths of the bonds that boron forms, acts against the
univalent state; accordingly, boron(I) is rarely encountered outside the realm of ‘high temperature’ molecules such as
BF and BCl.
In these and other respects, boron must be seen as a special case, with many idiosyncracies that separate it from the
other Group 13 elements. Although it serves as a vital reference point for understanding and evaluating the
chemistries of these elements, reasons of space and balance defy its inclusion, except by allusion, in this volume.
Otherwise, there would be the risk of having something of ‘an elephant in the living room’. This view finds support in
the precedents of not only two earlier books treating exclusively the metallic members of Group 13,12,27 but also an
extensive literature devoted specifically to boron2–9 (including the only reasonably up-to-date coverage in Gmelin2
for any member of the Group).
All the other members of the Group are then metals. If they show a closer kinship to one another than they do to
boron, their properties are, however, far from uniform. Symptomatic of the irregularities are the ionisation potentials
of the atoms which, unlike those of the corresponding metals of Groups 1 and 2, vary in a discontinuous way as
a function of atomic number (Table 1.1),12–22 so that I3, for example, follows the order B Al < Ga > In < Tl. This
sawtooth variation is a consequence of changes in the makeup and shielding of the electron core, that is [He], [Ne],
[Ar]3d10, [Kr]4d10 and [Xe]4f145d10. That it is more marked for the valence ns than for the np electrons reflects the
superior penetration of the core by the ns electrons. Relativistic effects23 make a significant contribution to the
binding energies for the later elements but do not change the overall pattern. The energies of the valence electrons of
the free atom are a major, but not exclusive, influence on the strengths of the bonds and on the type of compounds it
forms. They account, at least in part, for the chemical reactivity of the elements under normal conditions and the
emergence in each case of a relatively well defined cationic chemistry.8,12,27,28
The Group 13 metals are now acknowledged to have rich and distinctive chemical lives of their own, no longer
overshadowed by that of boron. Salient features of these lives are: (i) their varied redox chemistry with what is now
seen to be a wide range of formal oxidation states (including not only non-integral but also negative ones); (ii) the
natural acidity and associated coordination chemistry of their MIII compounds; and (iii) the great variety of structures
and other properties displayed by their compounds. The drive for new discoveries and a fuller understanding has
been urged not so much by the familiar commercial importance of aluminium and alumina-based solids,12 but by
a number of other developments of more recent origin.
1.2 Developments in Methodology
‘I keep six honest serving-men
(They taught me all I knew);
Their names are What and Why and When
And How and Where and Who.’
Rudyard Kipling, Just So Stories, ‘The Elephant’s Child’, 1902.
Developments in Methodology
7
1.2.1 Introduction
Developments in Group 13 metal chemistry in the past two decades have often been driven, at least ostensibly, by
the promise of practical applications, for example in realms as diverse as the sourcing and storage of energy; the
production and elaboration of materials with specific electronic, structural, thermal or chemical properties; and
medical diagnosis and therapy. While some important discoveries have indeed been made in pursuit of such
pragmatic causes, disinterested scientific curiosity continues to be the main life force of progress. Advances
have come in three principal areas.
(i) Synthesis. New compounds of a variety of sorts have been prepared. Most notable, perhaps, have been
those in which the Group 13 metal assumes a low formal oxidation state (i.e. < þ3),12,21,22,29 and may engage in
metal–metal bonding (as in so-called ‘metalloid’ derivatives).11,22,30–32 Other compounds are made notable by
the presence of weak, reactive M--H bonds;12,21,22 by the coordination environment of the metal, which may
display an unusual geometry or an uncharacteristically low coordination number (e.g. 133 or 234,35); by the
nature of the bonding, which sometimes challenges conventional wisdom regarding primary or secondary
interatomic interactions;36–38 or by their lability under normal conditions.21,29 At the same time, much effort and
ingenuity have been expended on the synthesis of solid materials with extended frameworks in which homo- or
hetero-nuclear assemblies, including the metal atoms, are bridged by non-metal atoms, typically from Groups
15 and 16.8,12,39–50 Specific objectives have included the achievement of particular topologies and/or
morphologies. The coordination chemistry of the metals has also expanded materially, with the preparation
of new complexes.8,10,12,51–53 Here the metal centre, as MIII, usually plays its conventional acceptor role, but
there is now ample evidence, usually involving MI compounds, to show that the metal can also function
predominantly as a donor.22,35,54–60 How different synthetic approaches have sought to achieve their respective
targets is reviewed briefly in the Section 1.2.2.
i
i
Pr
Pr
+
i
In i
Pr
[B(C6H3-3,5-(CF3)2)4 ]−
CO
Fe
Pr
....... .......
M
Fe
CO
OC
i
OC
(M = Ga or In)
Pr
i
1
Pr
2
(ii) Experimental characterisation. Closer, more extensive physical scrutiny, using new or established
experimental techniques (e.g. electronic, vibrational, microwave, mass and photoelectron spectroscopies, and
X-ray, neutron and electron diffraction) has extended knowledge of the structures, bonding and other properties
of both new and known compounds.61 Some of the major techniques in question are identified in Section 1.2.3.
(iii) Theoretical studies. With improvements in reliability, sophistication, scope and accessibility, quantum
chemical methods have played an increasingly important part in the advancement of the chemistry of this group
of metals.62 In some cases, they have pointed the way to the stable existence of a hitherto unknown compound
8 New Light on the Chemistry of the Group 13 Metals
(e.g. Ga2H2),63 or to the intermediacy of species in vapour transport (e.g. MP5 and MAs5, for M ¼ Ga or In);64a in
other cases, they have offered a rationale for the instability of a compound (e.g. MH3, where M ¼ In or Tl).64b–66
More often, they have been invoked to assist in the identification and characterisation of a new compound that
experiments have brought to light. To an increasing extent, they are now being exploited in the effort to gain a better
understanding of the reactivity of Group 13 metal species, for example Ga263,67 and AlCl.68 It is theoretical methods
that must also be deferred to in any questions of intramolecular or intermolecular bonding.36–38,69 This particular
theme is taken up in Section 1.4.
1.2.2 Synthetic methods
‘All progress is based upon a universal innate desire on the part of every organism to live beyond its income.’
Samuel Butler, Notebooks, Chapter 1, 1912
In an attempt to encompass a topic as vast and diffuse as this, it seems appropriate to treat it phase by phase, that is
according to whether the reaction takes place in the gas phase, involves the solid phase exclusively (as in a matrix, for
example) or in part, or is hosted by the liquid phase (as is most commonly the case).
1.2.2.1 Gas Phase Synthesis
The gas phase is home to the limited number of Group 13 metal compounds that are substantially volatile at ambient
temperatures, for example Al(BH4)3 and GaMe3,10,12 and otherwise to high-energy species such as the metal atoms,
M, metal clusters and their ions, and the diatomic molecules MIX [M ¼ Al, Ga, In or Tl; X ¼ H, F, Cl, Br or I].12,21,22
Some MI compounds, such as InCl or (h5-C5R5)M [M ¼ Al, Ga, In or Tl; R ¼ H or Me], vaporise on heating without
decomposition or disproportionation. At sufficiently high temperatures, the entropy advantage drives even robust
MIII molecules such as GaCl3 and In2O3 to decompose, at least partially, to the corresponding MI compound and
elemental non-metal. Other MI compounds require high-energy reactions for their formation in the gas phase. For
example, GaCl is generated by the reaction of the metal with Cl2, HCl or GaCl3 at 800–1000 C,70 and InF is formed
by heating together the metal and InF3.71 The reaction of the laser-ablated metal vapour with cyanogen or acetonitrile
in an argon carrier gas affords the two isomers MCN and MNC [M ¼ Al, Ga or In].72 Similarly, with ethyne in
a helium or neon carrier gas, laser-ablated aluminium vapour forms the linear AlCCH molecule.73 This, together
with the molecules AlNC and AlCH3 (formed in a discharge reaction between aluminium vapour and HgMe2),74 is
of interest as a potential carrier of the metal in the interstellar medium. The remarkable dialuminium compound
Al2(h5-C5H5), with a half-sandwich structure and an Al2 dimer unit located on the fivefold axis of the C5H5 ring, has
been prepared in a pulsed supersonic molecular beam by the reaction of laser-ablated aluminium vapour with
cyclopentadiene.75
The vapour formed by normal heating of a Group 13 metal consists mainly of atoms with only a low concentration
of dimers M2 and still lower concentrations of larger clusters. Spectroscopic and theoretical studies of the M2 species
indicate weakly bound molecules having a triplet ground state 3 Pu .76,77 Laser evaporation or ion bombardment
(sputtering) of the metal or a compound of the metal can deliver to the gas phase not only atoms and dimers, but
also larger clusters in either neutral or charged states.11,38,78–80 These gaseous species have been investigated
experimentally by different types of mass spectrometry (e.g. secondary ion mass spectrometry (SIMS) and Fourier
transform (FT) ion cyclotron resonance (ICR) mass spectrometry), photoionisation spectroscopy, or even calorimetric measurements, and theoretically at levels ranging from simple shell models to more sophisticated DFT
methods. Thus, aluminium clusters78,79 and, to a lesser degree, gallium80 and indium81,82 clusters have excited
considerable interest in the general search for a better understanding of how the transition is made from the atomic/
molecular state to the bulk metal. How the physical and chemical properties of the clusters vary as a function of size
Developments in Methodology
9
has, therefore, been a primary focus of enquiry, with the constantly teasing issue of just when a cluster can be
justifiably described as a ‘metal’. Intriguingly, density functional theory (DFT) calculations for Aln clusters suggest
that icosahedral packing is favoured only for n ¼ 13, whereas decahedral packing is most stable with n near 55, and
fcc packing is energetically preferred for n > 80.79 Neutral and charged Inn clusters with n up to 200 have been
produced by bombarding a pure indium surface with 15 keV Xeþ ions,82 but with yields that depend only partially,
and to an indeterminate extent, on their intrinsic thermodynamic properties.
Low concentrations and long mean-free-paths make the gas phase generally ill suited to useful synthetic
reactions, the predominant reactions being either simple addition or bond breaking. For example, spectroscopic
studies involving laser fluorescence excitation, one-photon and resonance-enhanced multiphoton ionisation
(REMPI), or zero electron kinetic energy (ZEKE) spectroscopy have revealed, typically in a supersonic beam,
the formation of such weakly bound adducts of aluminium atoms as AlH2,83 AlN2,84 AlOH2,85 AlNH3,86
AlCH487 and Alether [ether ¼ Me2O, Et2O or thf].88 Some of these are potential precursors to further change,
for example as represented by Equation 1.1, but, irrespective of thermodynamic considerations, such a change is
usually opposed by a substantial activation barrier. While the provision of additional energy – for example, to
promote the metal atom to an excited electronic state – may overcome this barrier,29,89 the surplus energy carried
by the product is likely, in the absence of an efficient means of relaxation, to result in its rapid disintegration.
Cation complexes, such as AlþH290 Alþ (OH2)n (n ¼ 1 or 2)91 and Alþ (CH3OH)n (n ¼ 1–4),92 which have also
been characterised, are rather more strongly bound and may be more prone to undergo spontaneous metalinsertion reactions analogous to that in Equation 1.1.22,29 Particularly fascinating, though, are the recent
experimental studies carried out to explore the primary reaction steps of the unusually stable cluster [Al13] and
involving severally Cl2, HCl and O2 (Chapter 7);93 these provide what may be regarded as ‘snapshots’ of the
reactions likely to take place on the base metal surface. Thermal decomposition of gaseous complexes such as
Me3AlNH3 and quinuclidineGaH3 or of gaseous mixtures of Me3Ga and AsH3 is important for the chemical
vapour deposition (CVD) of epitaxial films of III–V compounds or composites on an appropriate substrate.46,94
It is difficult, however, to distinguish between the reactions that occur strictly in the gas phase (and are likely
mainly to feature bond rupture) and those occurring on the surface of the substrate.
.
Al.HE
Al
hν
H
E
(E = H, OH, NH2 or CH3)
ð1:1Þ
The gas phase alone offers, then, few opportunities for the rational synthesis of compounds, and certainly on
a scale exceeding a few milligrammes. Instead, condensation is needed in order to take advantage of the greater
control that is afforded by the denser condensed phases. For example, the vapour of the metal or of a metal
compound may be condensed with a potential reagent and/or solvent to exploit any reactions occurring during
condensation or on subsequent warming of the reaction mixture. Such methods, pioneered in particular by
ockel et al. to capitalise on the inherent
Timms and Skell,95 have been deployed with great success by Schn€
reactivity of aluminium(I) and gallium(I) halides.70 Little is likely to happen in the gas phase, and while some
interaction between the reagents may occur during condensation, most, if not all, of the action occurs once the
reagents have entered the liquid phase; the course of events is accordingly subject to the usual considerations of
choice of solvent, involvement of protective functions, temperature, and so on (Section 1.2.2.3). Alternatively,
the vapour may be co-condensed with an excess of an inert gas doped with a potential reagent, as in the
technique of matrix isolation.29,96 Spectroscopic analysis of the resulting solid matrix then provides the means
of monitoring the reactions of the trapped species that may be activated either thermally or photolytically.
Matrix isolation, which is therefore primarily concerned with the solid phase (Section 1.2.2.2), is not a method
of synthesis in the conventional sense, but it has led to the first sighting and characterisation of numerous Group
13 metal compounds that are too labile to be isolated under ambient conditions, for example MH2, MH3
[M ¼ Al, Ga, In, or Tl],65,66,97 M2H2 [M ¼ Ga or In]63 and H2MNH2 [M ¼ Al, Ga, or In].98
10 New Light on the Chemistry of the Group 13 Metals
1.2.2.2 Synthesis with Solids
Group 13 metals and their solid compounds with extended, strongly bound frameworks are seldom wholly
compatible with homogeneous solution chemistry, at least under ambient conditions.45,99 Synthetic reactions in
which they feature, whether as reagents or as products, are most likely to be heterogeneous and to depend on the
presence of one or more solids. The simplest and most common way of preparing Group 13 metal compounds when
all the components are solids is the so-called ‘ceramic’ method.99 Stoichiometric amounts of the solid reagents are
ground together to give a uniform small particle size and heated to whatever temperature is needed to initiate
reaction. Used widely both industrially and in the laboratory, this method gives access to a whole range of materials,
such as mixed metal oxides, sulfides, selenides, nitrides and aluminosilicates. Representative examples of
compounds recently made in this way are: LaGaO3 and related compounds (from La2O3 and Ga2O3),100 LiGa5O8
and LiGaO2 (from Li2CO3 and Ga2O3),101 In4 þ xSn32xSbxO12 (a new transparent conductor, prepared from In2O3,
SnO2 and Sb2O3 in air)102 and Pt2In14Ga3O8F15 (containing [PtIn6]10þ moieties, prepared from InF3, platinum
powder and Ga2O3).103 The case for such solid state synthesis has also been advanced for the formation of molecular
organic and inorganic products and of materials with microporous metal–organic frameworks.104 It has been urged,
in part, by the desire to avoid the use of organic solvents, whether for reasons of reactivity and contamination or out
of concern for the environment. For example, the labile gallanes Ga2H6,105 H2GaBH4106 and H2GaB3H8107 have all
been prepared by mixing together the powdered solids [H2GaCl]2 and MX, where M ¼ Li or Nn Bu4 and X ¼ GaH4,
BH4 and B3H8, respectively.108 In the same vein, the tetrahydoborates Al(BH4)3,109 Me2GaBH4110 and HGa
(BH4)2111 are most easily synthesised from appropriate solid reactants without the intervention of a solvent.
The ceramic method suffers from several disadvantages. As the entire reaction occurs between solid components
and later by diffusion of the constituents through the product phase, diffusion paths necessarily become longer and
longer with the progress of the reaction, and the reaction rate correspondingly slower and slower. Separation of
a desired solid product from the solid mixture is liable to be difficult, if not impossible. Furthermore, the securing of
a compositionally homogeneous product can be problematic, even when the reaction proceeds almost to completion.
Various modifications of the technique have been devised to overcome some of its limitations. One of the main aims
has been to decrease the diffusion path lengths by reducing the particle size, thus effecting more intimate mixing of
the reactants, and this has been the purpose of introducing freeze drying, spray drying, coprecipitation and sol–gel
techniques.99 Microwave irradiation has also been deployed with some success in place of conventional
heating techniques,112 as exemplified by the formation of the diamond-like semiconductor AgInSe2 from the
powdered elements.113
Reactions involving solids are likely to be accelerated when heating results in melting of one or more components.
For example, heating together indium and molybdenum metals with MoO2 at 1150 C must result in melting of the
indium en route to the mixed metal oxide In5Mo18O28.114 The same applies to the synthesis of intermetallic phases,
including so-called Zintl phases,115 from the appropriate elements, for example K10Tl7,116 K39In80,117 YbGaGe
(a material with zero thermal expansion)118 and the thermoelectric material Ba8Ga16Ge30.119 With microwave
activation, gas–solid reactions may also be turned to advantage, as with the conversion of Al2O3 admixed with
carbon to AlCl3 by the action of HCl.120 Similarly, AlN and GaN have been prepared by the direct action of nitrogen
on a mixture of aluminium and charcoal powder in the first case,121 and by that of ammonia on a Ga2O3/amorphous
carbon mixture in the second.122
Water commonly plays a central role as both medium and reagent in the synthesis of solids with frameworks
composed of oxygen atoms bridging either metal atoms or metal and non-metal atoms such as silicon or phosporus.
‘Hydrothermal’ synthesis is the generic term for various techniques which involve the crystallisation of solids from
aqueous mixtures raised to high temperatures (typically 350–600 K) and under high vapour pressure.99,123 The
formation of a crystalline product rather than a powder stands in marked contrast to the normal outcome of ceramic
methods. The process, which is usually a heterogeneous one, occurs in nature, and numerous minerals, including
naturally occurring zeolites, owe their origin to it. Zeolites are also generally prepared in the laboratory by
hydrothermal methods,41,124 and such methods also provide the main means of entry to Group 13 metal
phosphates43,44 and a variety of other microporous and mesoporous materials.41,45 A typical synthetic mixture
Developments in Methodology
11
for making a specific aluminium phosphate consists of alumina, phosphoric acid, water and an organic material such
as a quaternary ammonium salt or an amine, which are heated together in an autoclave at 373–573 K.43,99 It may be
necessary to cater for large differences in solubility of the reactants, as exemplified by the synthesis of yttrium
aluminium garnet, YAG, Y3Al5O12, for which the more soluble Y2O3 needs to be placed in a cooler part and the less
soluble Al2O3 (as sapphire) in a hotter part of the autoclave; YAG crystals form where the two zones meet.
Hydrothermal methods, in common with ceramic methods, have benefited from the introduction of microwave, in
place of conventional thermal, activation;112 sonochemical methods have also been turned to advantage, notably for
the synthesis of porous metal oxides.125
Solids may act not only as reagents but also as mediators of chemical reactions. So the interstices of an open
framework material, such as a zeolite, are potential reaction chambers with the physical and chemical capacity to
catalyse specific modes of reaction.41,42 The same principle of confinement of reagents applies to a solid matrix
composed of noble gas atoms or simple molecules such as hydrogen, nitrogen, or methane.29,61,96 Necessarily
maintained at a low temperature (typically 2–20 K), such a matrix is rigid enough to immobilise any foreign atoms or
molecules that may be entrained at high dilution within the interstices, inert enough to keep perturbation of theseguests
to a minimum, and transparent enough to broad regions of electromagnetic radiation to admit interrogation by various
spectroscopic techniques. In practice, infrared measurements have been the principal agent of detection, identification
and characterisation, with crucial support often coming from the response to changes of isotopic composition. Final
decisions on the identity and properties of a new molecule are then likely to depend on the synergy between experiment
and quantum chemical calculations at an appropriate level of theory. Given the low thermal energy available under the
conditions of the matrix experiment, only simple addition reactions of the guest species opposed by little or no
activationbarrier are likely tooccur spontaneously.Hence, for example, it has been possibleto observe the formationof
adducts of the metal atoms, such as MNH329,98,126 and M(CO)n (n ¼ 1 or 2),29,127 where M ¼ Al, Ga or In. These
adducts range in their estimated total binding energies from a mere 7 kJ mol1 for InN2 to a highly respectable
176 kJ mol1 for Al(CO)2. The latter is believed to have the intriguing structure 3 with a tight C–Al–C angle (near
70 ) and Al–C:O arms that are bent outwards in such a way as to suggest that the two carbon atoms are drawn towards
Al
C
C
O
O
3
each other.127,128 By contrast, a much higher price in activation has usually to be paid for changes that involve bond
dissociation, insertion into a bond or isomerisation. Photons are then the only currency usually available to matrixisolated species, and practical provision must therefore be made to enable the matrix deposit to be irradiated with light
spanning an appropriate range of wavelengths. With an imagery suggestive of Coleridge’s Kubla Khan,
‘It was a miracle of rare device,
A sunny pleasure dome with caves of ice.’
Group 13 metal atoms, once promoted to their 2 S or 2 D excited electronic states (by ultraviolet radiation in the
wavelength range 290–340 nm), are capable of spontaneous insertion into an H--H or H--C bond to give authentic,
paramagnetic M(II) molecules M (H)X (e.g. X ¼ H 29,97 or CH329,129), which have been identified by their IR and
EPR spectra. The pervasive role of photoactivation is revealed by the results of matrix experiments involving
thermally evaporated gallium atoms, summarised schematically in Figure 1.1. By contrast, laser ablation of the metal
gives rise to atoms of high energy, trapping of which, together with potential reagents, in an appropriate solid matrix
is likely to result in spontaneous changes beyond the reach of thermally evaporated metal atoms. Co-condensing
laser ablated metal atoms with hydrogen (H2) has led, for example, to the first detailed characterisation of the Al2H6
*
12 New Light on the Chemistry of the Group 13 Metals
Figure 1.1 Some reactions of matrix-isolated gallium atoms in their ground or, more often, excited electronic states.
Reprinted with permission from [29]. Copyright 2002 American Chemical Society
molecule130 and of various indium hydrides, including the polymeric solid [InH3]n, which decomposes to the
elements at 160–180 K.65 Whereas gallium atoms require a substantial stimulus before they will react with
hydrogen, the Ga2 dimer reacts spontaneously with hydrogen at about 15 K to form the dimeric gallium(I) hydride,
Ga(m-H)2Ga.63 In2 does not, however, follow suit, requiring UV photoactivation before it will form an analogous
product. Both Ga(m-H)2Ga and In(m-H)2In prove to be photolabile under visible light (l > 450 nm), undergoing the
changes outlined in Figure 1.2. These and other Group 13 metal species characterised in recent matrix studies are
listed in Table 1.2.
1.2.2.3 Liquid Phase Synthesis
Whatever the merits of the gas or solid phases in one form or another, homogeneous processes in the liquid phase
have maintained their supremacy as the means of useful chemical synthesis. Nor can custom stale the infinite variety
that a suitable solvent is able to bring to a reaction mixture by way of support, intimate and efficient mixing, control
and potential thermodynamic and/or kinetic influence. Reagents and target compounds often being susceptible to
attack by moisture, organic solvents such as ethers have been the mainstay of much of the synthetic effort. In that
derivatives of the Group 13 metals in oxidation states lower than þ3 are invariably weaker Lewis acids than are the
corresponding MIII derivatives, they are vulnerable to disproportionation under basic conditions, so that the role of
the solvent may be far from innocent. For example, tetrahydrofuran (thf), used quite frequently to support an
indium monohalide in the presence of various reagents, tends to induce disproportionation rather than the
expected metathesis or addition.22 Thus, indium metal is a common product of all the following reactions:
t
Bu3 SiNa þ InBr giving the indium(II) product (t Bu3 Si)2InIn(Sit Bu3 )2;149 Na[MeGa(Pz)3] (Pz ¼ pyrazolyl) þ InI
giving [{MeGa(Pz)3}2In][InI4];150 RLithf þ InBr giving Br(R)InInBr(R), where R is the 1-aza-allyl
ligand (Me3Si)2C(Ph)C(Me3Si)N;151 LiC(SiMe3)32thf þ InBr giving [Li(thf)3][In3Br3{C(SiMe3)3}3];152 and
ArN ¼ CHPy [Ar ¼ C6H3-2,6-i Pr2 ; Py ¼ 2-pyridyl] þ InCl giving InCl3(thf){h2-ArN ¼ CHPy}.153 Nevertheless,
basic organic solvents have played an important part in opening up the chemistry of the univalent metal halides
Developments in Methodology
13
Figure 1.2 Pathways for the matrix reactions of Ga2 and In2 molecules with H2. Reprinted from [128], with
permission from Elsevier
MX [M ¼ Al, Ga or In]. Toluene/ether or similar mixtures give metastable solutions of aluminium(I) and gallium(I)
halides, which survive at low temperatures (190–250 K) and have proved to be invaluable synthons for new
compounds of these metals in low formal oxidation states.70 In the same vein, the less tractable indium(I) halides can
be made to dissolve to a limited extent in a toluene/tmeda mixture (tmeda ¼ Me2NCH2CH2NMe2), and the
metastable solution formed by InI disproportionates, for example, to form indium metal and the indium sub-halide
cluster complex In6I8(tmeda)4 (4).154
InI2(tmeda)
I2In
In
(tmeda)I 2In
InI(tmeda)
InI(tmeda)
4
14 New Light on the Chemistry of the Group 13 Metals
Table 1.2 Group 13 metal atoms, dimers and molecular compounds featuring in recent matrix-isolation studies
Species
Method of formation
Method of
characterisation
Reference
Al, Ga, In or Tl atoms
co-deposition of metal vapour with
noble gas
mass selection of sputtered Al
cluster ions, neutralisation and
co-deposition with Ar
co-deposition of metal vapour with Ar
UV-Vis
29
resonance Raman
131
resonance Raman;
UV-Vis
Raman; UV-Vis
IR; UV-Vis
132
133
29
IR; UV-Vis; EPR
29, 98, 126
IR; UV-Vis
29, 134
IR; EPR
29, 127
IR; UV-Vis
135
IR
97
IR
130
IR
136
IR; Raman
105, 137
IR; UV-Vis
emission
65
IR
66
IR; Raman
63
IR
98
IR
29, 134
IR; EPR
29, 129
IR
135
Al2
Ga2
GaN2
MOH2 (M ¼ Al, Ga, In or Tl)
MNH3 (M ¼ Al, Ga or In)
MPH3 (M ¼ Al, Ga or In)
M(CO)n, M2(CO)n (M ¼ Al, Ga
or In; n ¼ 1 or 2)
MSiH4 (M ¼ Al or Ga)
MHn (M ¼ Al, Ga or In; n ¼ 1–3)
AlHn (n ¼ 1–3), H2AlH3,
Al2H4, Al2H5, Al2H6, Al3H9
GaHn (n ¼ 1–3), Ga2Hn (n ¼ 2,
4 or 6), GaHn (n ¼ 2 or 4)
Ga2H6
InHn (n ¼ 1–3), H2InH3, In2Hn
(n ¼ 2, 4 or 6), [InH3]x
TlHn (n ¼ 1–3), Tl2H2
Ga2H2, In2H2
HMNH2, MNH2, H2MNH2
(M ¼ Al, Ga or In)
HMPH2, H2MPH (M ¼ Al, Ga
or In)
HMCH3, MCH3 (M ¼ Al, Ga or
In)
HMSiH3, MSiH3 (M ¼ Al or Ga)
co-deposition of metal vapour with N2
co-deposition of metal vapour with
noble gas doped with H2O
co-deposition of metal vapour with Ar
doped with NH3
co-deposition of metal vapour with Ar
doped with PH3
co-deposition of metal vapour with
CO or noble gas doped with CO
co-deposition of metal vapor with Ar
doped with SiH4
co-deposition of metal vapour with H2
or H atoms in an Ar matrix þ UV
photolysis
co-deposition of laser ablated Al
vapour with H2 þ UV photolysis
co-deposition of laser ablated Ga
vapour with H2 þ UV photolysis
co-deposition of gallane vapour with
noble gas
co-deposition of laser ablated In
vapour with H2 or Ne doped with
H2 þ UV photolysis
co-deposition of laser-ablated Tl
vapour with H2 or noble gas
doped with H2 þ UV photolysis
co-deposition of metal vapour with Ar
doped with H2 þ UV-Vis
photolysis
UV photolysis of Ar matrix doped with
M atoms and NH3
UV photolysis of Ar matrix doped with
M atoms and PH3
UV photolysis of M atoms trapped in
CH4 or noble gas matrix doped with
CH4
UV photolysis of an Ar matrix doped
with M atoms and SiH4
Developments in Methodology
15
Table 1.2 (Continued)
Species
Method of formation
Method of
characterisation
Reference
HMSnH3, H2M(m-H)2Sn (M ¼
Al or Ga)
HnMX3n (M ¼ Al, Ga or In;
X ¼ Cl or Br; n ¼ 1 or 2)
CH3GaH2, (CH3)2GaH
co-deposition of metal vapour with Ar
doped with SnH4 þ UV photolysis
UV photolysis of a noble gas matrix
doped with MX and either H2 or HX
co-deposition of volatile thermolysis
products of (CH3)2N(CH2)3Ga
(CH3)2 with Ar
UV photolysis of Ar matrix doped with
(h5-C5Me5)Al and H2
UV photolysis of Ar matrix doped with
AlCl and RH
co-deposition of metal vapour with
noble gas doped with H2, HCl, H2O
or CH4 þ UV photolysis
co-deposition of AlCl vapour with Ar
IR
138
IR
21, 139
IR; mass
spectrometry
21
IR
140
IR
141
EPR
29
IR
142
co-deposition of laser ablated Al vapour with Ar doped with X2
co-deposition of metal vapour with
noble gas doped with O2
IR
29
IR
29, 143
co-deposition of Al vapour with Ar
doped with O2
co-deposition of metal vapour with Ar
doped with O2
co-deposition of AlX vapour with Ar
doped with O2
co-deposition of laser-ablated metal
vapour with Ar doped with H2O2 or
H2 þ O2
co-deposition of laser ablated metal
vapour with N2 or Ar doped with N2
co-deposition of laser ablated metal
vapour with Ar doped with NO
co-deposition of metal vapour with Ar
doped with C2H2
co-deposition of (CH3)3AlNH3
vapour with Ar þ UV photolysis
IR
144
IR
145
IR
146
IR
147
IR
29
IR
29
IR
29
IR
148
(h5-C5Me5)AlH2
RAl(H)Cl (R ¼ CH3 or C:CH)
MH2 (M ¼ Al or Ga), HAlCl,
HAlOH, CH3GaH
AlCl, Al(m-Cl)2Al, ClAl(m-Al)
(m-Cl)AlCl
AlX2 (X ¼ F, Cl, Br or I)
MOM, (h2-O2)M, OMO,
MOMO, M(m-O)2M, M(O2)
(M ¼ Al, Ga, In or Tl)
(h2-O2)nAl (n ¼ 1, 2 or 3)
O2M(h1-O2) (M ¼ Ga or In)
XAl(O2)n (X ¼ F, Cl or Br; n ¼ 1
or 2)
M(OH)n (M ¼ Al, Ga, In or Tl;
n ¼ 1, 2 or 3)
NMN (M ¼ Al, Ga, In or Tl)
MNO (M ¼ Ga, In or Tl)
(h2-C2H2)M (M ¼ Al, Ga or In)
(CH3)3Al NH3, (CH3)2AlNH2
The motivation for synthesis has commonly centred on the support of the Group 13 metal atom in an unusual
oxidation state (< þ 3) or coordination environment (featuring a coordination number <4), and been influenced, too,
by the need to control the partial charge carried by the metal. A metal fragment of this sort is usually reactive through
its exposure to associative attack that results in aggregation, coordination, oxidation or disproportionation. With
matrix isolation, protection from such attack is provided physically by trapping the fragment in an inert, solid matrix
at low temperature. For conventional synthetic operations, however, protection must be achieved by chemical means,
namely by the adoption of ligands with appropriate steric and/or electronic properties. Thus, labile species such as
16 New Light on the Chemistry of the Group 13 Metals
AlH3, GaH3 or InH3 may be intercepted by a suitable donor and preserved in the form of thermally more robust
adducts, such as HC(CH2CH2)3NMH3 [M ¼ Al or Ga]94,155 and Mes(NCH¼CHN(Mes)C)InH3 (Mes ¼ mesityl)156 (Section 1.6.1). One of the main guiding principles in recent years has been the recognition of bulk and
specific design as properties of supporting ligands that can be crucial to manipulating the reactivity of the metal
centre and to stabilising previously unknown or unfamiliar bonding types, geometries or electron configurations. The
ligands in question include substituted h5-cyclopentadienyl groups (e.g. C5Me5);157 bulky alkyl and supersilyl
groups (e.g. CH(SiMe3)2, C(SiMe3)3, and Sit Bu3 );158 substituted aryl groups (e.g. C6H2-2,4,6-t Bu3 ), terphenyl and
related substituents;159 b-diketiminate or amidinate derivatives (e.g. {ArNC(Me)}2CH and Cy2NC(NAr)2, with
Cy ¼ cyclohexyl and Ar ¼ C6H3-2,6-i Pr2 );160–163 substituted diazabutadiene derivatives (e.g. RNCH¼CHNR with
R ¼ t Bu or C6H3-2,6-i Pr2 );160,164 and poly(pyrazolyl)borate groups,165 as represented, for example, in Equation 1.2
showing the oxidation of a gallium(I) compound to a discrete molecular gallium(III) compound containing what may
reasonably be regarded as a ‘semipolar metal–chalcogen double bond’.166 A notable example is provided by the
unusually encumbering o-terphenyl ligand –C6H3-2,6-(C6H2-2,4,6-i Pr3 )2; hence, the indium(I) compound
t
t
t
Bu
N
N
t
Bu
H
B
t
N
N
N
N
t
t
Bu
Ga
Bu
t
Bu
E
N
B
t
Bu
N
N
N
N
N
Bu
t
Bu
H
E = Se or Te
t
Bu
Bu
E
Ga
t
Bu
ð1:2Þ
Bu
InC6H3-2.6-(C6H2-2,4,6-i Pr3 )2 has been prepared and shown to form crystals composed of well separated monomers
(1), which are unique in the one-coordination of the metal.33 Even relaxing the bulk of the ligand seemingly quite
slightly in the change from –C6H3-2,6-(C6H2-2,4,6-i Pr3 )2 to –C6H3-2,6-(C6H3-2,6-i Pr2 )2 gives not a monomer but
a ‘dimetallene’ dimer [InC6H3-2,6-Dipp2]2 with bi-coordinated metal atoms linked by a metal–metal bond.167 A
similar structure is adopted by the so-called ‘gallyne’ compound Na2[GaC6H3-2,6-(C6H2-2,4,6-i Pr3 )2]2, superficially analogous to an alkyne and remarkable for displaying a short Ga–Ga bond, the multiple nature of which has
been heatedly debated.36,168,169 Twofold coordination of the MI centre is also found in the neutral six-membered
heterocycles M[h2-(NArCMe)2CH] [M ¼ Al, Ga or In; Ar ¼ C6H3-2,6-i Pr2 ] formed by either reduction or
metathesis, as in Equation 1.3.160,161 The role of ligand bulk is particularly felt in the disproportionation of MI
compounds [M ¼ Al, Ga or In], which normally leads to the elemental metal M0 and the corresponding MIII
K, − KI
“GaI” or InI
N
N
Ar
− M1I
M1
M1 = Li or K
Ar
N
Ar
N
N
..
M
Ar
M = Al, Ga or In
N
Al
Ar
I
Ar
ð1:3Þ
I
compound. With a bulky ligand such as Sit Bu3 , however, the products may be M0 and a metal–metal-bonded MII
derivative, for example (t Bu3 Si)2InIn(Sit Bu3 )2.149 More spectacularly still, similar ligands may frustrate the
formation of the metal in such a disproportionation, yielding instead metalloid cluster species,29–32 such as
Al50(h5-C5Me5)12170 and Ga22(Pt Bu2 )12.171
Much of the ‘variety’ of solution chemistry arises from the diversity of reaction types that can be accommodated
through the greater freedom enjoyed both in the approach of reagents and the separation of products from the
Developments in Methodology
17
immediate reaction sphere, as well as the greater scope open to thermal or other forms of activation. The reactions
range from oxidation/reduction, through metathesis, to disproportionation/synproportionation, as illustrated by the
following representative examples (Equations 1.4–1.16 )22,168,170,172–180 reflecting some of the results of recent
research having to do with the chemistry of the Group 13 metals in low oxidation states. Products of particular note
include metalloid cluster derivatives,170,180 compounds with the potential for multiple metal–metal bonding,168,174,175 and transition metal complexes in which the bare Group 13 metal cation Gaþ or Inþ acts as
a ligand.179 Oxidation in an aprotic medium, such as thf or ether, may also be brought about electrochemically, as
with the recently reported synthesis of AlH3, a potentially attractive material for hydrogen storage.180
(i) Oxidation
R
E
R
M R
M
E
E or Et3PE
R
toluene
M
R
M
M
E
E
M
R
M = Al, Ga or In; E = O, S, Se
or Te; R = C(SiMe3)3 or η5-C5Me5
ð1:4Þ172
R
M
R
M
I2/AlI 3
R
M
n-hexane
I
I
M
R
M
R
M = In; R = C(SiMe3)3
E
R
R
R
Et 3PE
M
M
− Et3P
R
M
R
R
R
M
ð1:5Þ173
R
M = Al, Ga or In; E = S, Se or Te;
R = CH(SiMe3)2
(ii) Reduction
ArGaCl2
Na, Et2O
NaCl
Na+2
Ga
..........
Ar
2–
Ga
Ar
Ar = C6H3-2,6-(C6H2-2,4,6-iPr 3) 2
ð1:6Þ168
18 New Light on the Chemistry of the Group 13 Metals
2−
Ar
Na, Et2O
ArAlI2
Na+2
− NaI
Al
Al
ð1:7Þ174a
Al
Ar
Ar
Ar=C6H3-2,6-(C6H2-2,4,6-Me3)2
Ar
Al
Ar
Ar
Ar
Ar
Li/tmeda, Et2O
Al
[Li(tmeda)2]+
Al
.−
Al
ð1:8Þ175
Ar
Ar
Ar = C6H2-2,4,6-i Pr3
Ar
(iii) Metathesis/Ligand Exchange
R
MX
M
LiR or MgR2, toluene or ether
R
−LiX or RMgX
M
M
M
R
R
ð1:9Þ172
M = Al, Ga, In or Tl; R=
C(SiMe3)3 or η5-C5Me5
Ar
LiAr, Et2O
TlCl
Tl
−LiCl
Tl
ð1:10Þ174b
Tl
Ar
Ar
Ar = C6H3-2,6-(C6H3-2,6-Me2) 2
[(η5-C5Me5)In]n
CF3SO3H, toluene
In+[CF3SO3] −
− C5Me5H
ð1:11Þ176
MR
[MR]n
M'(cod)2, alkane
− cod (= 1,5-cyclooctadiene)
M'
MR
MR
MR
M = Al, Ga or In; M' = Ni or Pt;
R = C(SiMe3) 3 or η5-C5Me5
ð1:12Þ177
Developments in Methodology
L
19
L
Pt
[Ga2Cp*][BAr F4],
fluorobenzene
L
L
In[BAr F4],
fluorobenzene
−L = PPh3
−L = GaCp*
+
Ga
In
*CpGa
Pt
Ph 3P
[BAr F4] -
GaCp*
[BAr F4] -
Pt
*CpGa
ð1:13Þ178a
+
PPh 3
Ph3 P
GaCp*
Ar F = C6H3-3,5-(CF3) 2
Cp* = η -C5Me 5
5
(iv) Disproportionation/Synproportionation
MgCp*2, toluene/THF
AlBr
Al50Cp*12
− [AlCp*]4 − Cp*MgBr − AlBr3
ð1:14Þ170
Cp* = η5-C5Me5
R
R
GaBr
− LiBr − [GaR]4 − GaBr3
Ga
R
Ga Ga
Ga
LiR, toluene/THF
R
Ga Ga
Ga
ð1:15Þ179
Ga
R
R
R = C(SiMe3)3
L
InX + InY3 + 2L
toluene or CH 2Cl2
Y
X
In
Y
In
Y
ð1:16Þ22
L
X,Y = Cl, Br or I; L =
donor, e.g. tmeda
1.2.3 Experimental methods of identification and characterisation61
‘The motto of all the mongoose family is, ‘Run and find out.’
Rudyard Kipling, The Jungle Book, ‘Rikki-Tikki-Tavi’, 1894
Some of the principal methods used in recent years to identify and monitor Group 13 metal compounds and to
characterise their physical and chemical properties are listed in Table 1.3 (see earlier references as cited and
references 181–225). The structural scene in the solid state continues to be dominated by X-ray diffraction studies
Technique
Parameters
Uses
Examples
X-ray diffraction of single
crystals or crystalline
powders
Unit cell symmetry and
dimensions; crystal and
molecular structures;
molecular dimensions and
intermolecular contacts.
Principal method of structural
characterisation in the solid
state; light atoms not always
easily located, and problems
of differentiating between
atoms of similar atomic
number; problems of
disorder and twinning.
Neutron diffraction of single
crystals or, more often,
crystalline powders
Crystal and molecular
structures; molecular
dimensions and intermolecular contacts.
Characterisation of micro- or
nano-structures of solids and
solid surfaces.
Particularly important for the
location of H(D) atoms.
[GaBH6]n;106 [(h5-C5Me5)M]6
(M ¼ Ga or In);181 [MMe3]n
(M ¼ B, Al, Ga, In or Tl);182
Na2[ArGaGaAr] {Ar ¼ C6H32,6-(C6H2-2,4,6-i Pr3 )2};168
Na2[(AlAr)3] {Ar ¼ C6H3-2,6(C6H2-2,4,6-Me3)2};174
Al50(h5-C5Me5)12;170 Ga8{C
(SiMe3)3}6;180 [{(h5-C5Me5)
Ga}4PtGa][BArF4] {ArF ¼
C6H3-3,5-(CF3)2};179 M0
(MR)4 (M0 ¼ Ni, Pd or Pt; M
¼ Al, Ga or In; R ¼ C(SiMe3)3
or h5-C5Me5);178 corundumtype In2O3 (powder);183 In
(O)OH (powder);183
ZnGa2O4 (powder);184
[H2tmdp]2[Ga10(OH)4
(C2O4)2(HPO4)(PO4)8]3H2O
(tmdp ¼ 4,40 -trimethylenedipyridine, powder).185
[D2GaND2]3;187 Al2(CD3)6;182a
KAlD4;188 Sr2AlD7;189a
BaAlD5;189b LiAlD4/
Li3AlD6.190
Decomposition of LiAlD4;191
In2O3, In(O)OH and In(OH)3
nanocrystals;192 SrAl2O4
nanotubes.193
Electron microscopy
Imaging and mensuration of
micro- and nano-structures;
phase detection and
monitoring.
20 New Light on the Chemistry of the Group 13 Metals
Table 1.3 Major experimental techniques in current use for identification, monitoring and characterisation of Group 13 metal compounds
X-ray absorption spectra at the
absorption edge of a particular element; gives distribution of interatomic distances.
Studies of amorphous solids,
liquids, solutions and surfaces for information about
the coordination environment of the metal atom.
Electron diffraction of vapour
samples
Gives vibrationally averaged
molecular structure; molecular dimensions and vibrational amplitudes; geometric
and vibrational problem
commonly
underdetermined.
Most reliable structure usually
determined by combined
analysis of electron diffraction, vibrational and other
spectroscopic data, and the
results of quantum chemical
calculations.
Vibrational spectroscopy
IR and Raman spectra used to
determine vibrational wavenumbers and intensities of
absorption/scattering; symmetry; isotopic effects.
Identification of new molecules
and deduction of likely
structures; identification of
functional groups, e.g. M–H
vs. M–H–M; monitoring of
chemical changes; of particular use in support of matrixisolation studies; quantum
chemical simulation often
important.
Inelastic neutron scattering
(INS)
Vibrational transitions of solids
or adsorbed species.
Access to vibrational properties
not easily determined by
other means; computational
modelling of the INS spectra.
Amorphous AlBrnF3n;194
aqueous Al(III) and Ga(III)
citrate complexes;195
(t Bu4 pc)InIn(t Bu4 pc)2tmeda
(t Bu4 pc ¼ tetra-t-butylphthalocyanine; tmeda ¼
tetramethylethylenediamine.196
Ga2H6;105,229 GaBH6;106
H2GaB3H8;107
[H2GaCl]2;197
HGa(BH4)2;111 [Me2AlH]n
(n ¼ 2 or 3);198
Me3PGaH3;199 (h5-C5Me5)M
(M ¼ Al, Ga or In);181b,200
[t BuGaS]4;201 Mt Bu3
(M ¼ Al, Ga or In).202
M2 (M ¼ Al or Ga);131,132 MH2
and MH3 (M ¼ Al, Ga, In or
Tl);65,66,97,130,136 Al2H6;130
M2H2 (M ¼ Ga or In);63
H2MNH2 (M ¼ Al, Ga or
In);98 HMEH3 (M ¼ Al, Ga or
In; E ¼ C, Si or
Sn);29,129,135,138 MCH3 (M ¼
Al, Ga or In);29,129 (h2-O2)nAl
(n ¼ 1–3);144 [AlH3]n formed
on the Al(111) surface.203a
Identification of molecular aluminium hydrides, [AlH3]n,
during H2 regeneration from
catalyst-doped NaAlH4.203b
(continued)
Developments in Methodology
Extended X-ray absorption fine
structure (EXAFS)
21
(Continued )
Technique
Parameters
Uses
Examples
High-resolution electronic
emission or absorption, IR,
and microwave/millimetre
wave spectroscopies
Vibrational and rotational
constants of molecules in
their ground and excited
electronic states.
MX (M ¼ Al, Ga, In or Tl; X ¼ H,
F, Cl, Br or I);12,13,21,22
CH3Al;74 MCN and MNC
(M ¼ Al, Ga or In);72 (h5C5H5)M (M ¼ In or Tl);204
AlCCH;73 In2O.205
NMR spectroscopy
Chemical shifts, multiplet
patterns, J values, line profiles
and intensities; temperaturedependence.
Detailed characterisation of
structures and other properties of diatomic and other
simple molecules in the gas
phase. Transient species
commonly interrogated in a
supersonic jet.
Routine characterisation and
monitoring of new compounds, mostly in solution
and involving 1 H, 13 C, 11 B,
31
P or 19 F nuclei in the ligand
(s). Studies of 69 Ga, 71 Ga,
115
In, 203 Tl and 205 Tl NMR
spectra more restricted, but
27
Al NMRexperimentswidely
used, particularly for the study
of extended solid structures.206 Analysis of environment of Group 13 metal atom.
Determination of thermodynamic and kinetic properties.
EPR spectroscopy
g- and A-values; multiplet
patterns; unpaired spin
densities.
Identification and characterisation of M(0), M(II), and
other paramagnetic species,
e.g. in cryogenic
experiments.
Barriers to rotation about M–X
bonds (M ¼ Group 13 metal;
X ¼ Group 15 or Group 16
element) with the potential
for multiple bonding;36
zeolites, clays and related
layer silicates, cements, and
aluminophosphates;207,208
fluorination products of
Al2O3;209 organoaluminium
ketoximates;210 in vitro
aluminated brain tissue;211 a
and b isomers of the Keggin
anion [AlIIIW12O40]5;212
defects and local structure in
natural and synthetic alunite
(K,Na,H3O)Al3
(SO4)2(OH)6;213 solid
oxogallium compounds, e.g.
Ga2O3, MGaO2 (M ¼ alkali
metal) and LaGaO3.214
MH2 (M ¼ Al or Ga);29 HAlCl,
HAlOH and CH3GaH;29
AlNH3 and HAlNH2;126
M(CO)n (M ¼ Al, Ga or In;
n ¼ 1 or 2);29 Al(PF3)2;215
[R2MMR2] {M ¼ Al or Ga;
R ¼ CH(SiMe3)2 or C6H22,4,6-i Pr3 }.36,175
*
22 New Light on the Chemistry of the Group 13 Metals
Table 1.3
Mass spectrometry (MS),
including secondary ion (SI),
ion cyclotron resonance
(ICR) and other forms of mass
spectrometry
Molecular masses, isotopic
patterns; peak intensities.
Routine characterisation of new
molecular compounds; fragmentation patterns may help
in structure (connectivity)
elucidation; potential source
of information about thermodynamic properties;
monitoring of ion–molecule
reactions.
UV photoelectron and
photoionisation
spectroscopy, including
anion photoelectron,
multiphoton ionisation, and
ZEKE spectroscopy
Ionisation energies; profiles and
intensities of bands as a
function of the energy of the
exciting radiation; ion
masses; electronic and
structural details.
Energies of occupied electronic
levels of molecules; identification of ground and excited
states of neutral and ionised
species; high resolution details may give vibrational and
structural information.
Ga2, In2, GaIn;77b AlCH2,
Al(CH3)n (n ¼ 1 or 2) and
AlC2H4 studied by neutralisation-reionisation mass
spectrometry;216 photofragmentation products of MMe3
(M ¼ Al, Ga or In);217
oxidation of alumina formers
studied by SIMS;218Al4 Cp*4
and Alx Cp*y (Cp ¼ h5C5Me5; x up to 8 and x >
y);219 studies of reactions of
Al13 and Al13H ions with
Cl2, HCl or O2.93
Ga2H6 and GaBH6;220
(h5-P2C3t Bu3 )In,
(h5-P3C2t Bu2 )In;221
Al2(h5-C5H5);75 Ga2N;222
ZEKE studies of
(h5-C5H5)Al2,75 Al2,223
AlAr224 and InN2.225
Developments in Methodology
23
24 New Light on the Chemistry of the Group 13 Metals
of single crystals or crystalline powders. For all their huge influence, X-ray methods do have their limitations, to
some of which Table 1.3 alludes. There are problems, for example, with the location of light atoms such as hydrogen,
particularly in the vicinity of much heavier metal atoms. For this and other reasons, metal–hydrogen distances
are often rather poorly defined; being measures of the separation
between maxima of electron density, they are in
any case systematically shorter (typically by about 0.1 A) than the internuclear distances determined by neutron or
electron diffraction or by spectroscopic methods. Nevertheless, technical advances, including the routine study of
crystals at low temperature and improved methods of treating issues of twinning and disorder, have gone a long way
towards countering some of the inherent problems. As a consequence of these improvements, it has been possible to
pay more attention to intermolecular, as opposed to intramolecular, features of crystal structures. Not only has
conventional hydrogen bonding been characterised, but non-classical ‘dihydrogen’ bonding has been identified as
a significant intermolecular interaction. This is the case, for example, in solid cyclotrigallazane, [H2NGaH2]3,187
where hydridic hydrogen atoms bound to gallium interact with protonic hydrogen atoms bound to nitrogen to form
four intermolecular Ga--H H--N dihydrogen bonds per molecule, each with an energy estimated to be in the order
of 12 kJ mol1. Similar intermolecular interactions are evident in adducts of AlH3 and GaH3 with primary and
secondary amines.226 Such interactions may even prefigure the elimination of hydrogen molecules that characterises
the thermolysis of such compounds. The importance of intermolecular forces in determining the molecular packing
and crystal structure is well illustrated by the polymorphism of the trimethyl derivatives of the Group 13 elements.182
Regarding the precise location of hydrogen atoms, neutron diffraction of deuterated derivatives, usually in the form
of crystalline powders, offers the best solution, but only quite rarely, as with [D2NGaD2]3,187 Al2(CD3)6182a and
KAlD4,188 has this option been taken.
To determine the properties of diatomic and other simple Group 13 metal-containing molecules in the gas phase,
and hence free from the potentially perturbing effects of intermolecular contacts, one normally turns to highresolution spectroscopy involving electronic or vibrational transitions, either in emission or absorption, or pure
rotational transitions (microwave or millimetre wave spectroscopy). Hence, virtually all the gaseous molecules of
the type MX [M ¼ Al, Ga, In or Tl; X ¼ H, F, Cl, Br or I] have been characterised in detail.12,13,21,22 Other simple
molecules that have lent themselves to such interrogation include CH3Al,74 AlCCH73 and MCN and MNC [M ¼ Al,
Ga or In].72 For more complicated molecules, it is necessary to appeal to electron diffraction, but in this case the
problem of extracting good estimates of all the structural and vibrational parameters that determine the observed
molecular scattering is usually underdetermined by a significant margin. The best way of dealing with this problem is
to carry out a combined analysis that incorporates the geometric and vibrational information carried not only by the
electron diffraction pattern, but also by the rotational constants and an appropriate vibrational force field. A further
improvement has been made with the development of the so-called SARACEN method, whereby parameters that
cannot be refined freely are made subject to restraints derived not only from other experimental sources but also from
an array of quantum chemical calculations.227
Still the determination of molecular and crystal structures is far from being the defining purpose of experimental
studies. Routine characterisation of compounds still depends typically on vibrational, NMR, and mass spectroscopies for information about elemental composition and connectivity, substituents and their mode of coordination,
the environment of the Group 13 metal and nuclearity.61 Not all solids oblige with the formation of crystalline phases
suitable for X-ray and neutron diffraction, and so methods such as those involving extended X-ray absorption fine
structure (EXAFS) may need to be brought into play. Hence, for example, the environment of the Al(III) and Ga(III)
centres in aqueous citrate complexes has been analysed,195 and the solid phthalocyanine indium(II)
(t Bu4 pc ¼ tetra-t-butylphthalocyanine) has been shown to sport an uncomplex (t Bu4 pc)InIn(t Bu4 pc)2tmeda
usually long In–In bond (3.24 A).196 Characterisation of the micro- or nano-structure of solids and solid surfaces
relies heavily on the use of transmission and/or scanning electron microscopy, as exemplified by the cases of various
hierarchical structures of In2O3, In(O)OH and In(OH)3,192 and SrAl2O4 nanotubes.193 NMR and vibrational
spectroscopies are important for the ease with which they can be applied to the monitoring of chemical changes,
usually occurring in solution, with NMR measurements being paramount in the exploration of the kinetics of such
changes. So it is with dynamic processes that may involve rotation about a bond to a Group 13 metal with potential
Redox Chemistry of the Group 13 Metals: Access to Oxidation States Lower than þ3 25
multiple character,169 and in the cases of Mes*2 GaSMe (Mes ¼ C6H2-2,4,6-t Bu3 ),228a Ar2MN(H)Ar0 [M ¼ Al or Ga;
Ar ¼ C6H2-2,4,6-i Pr3 ; Ar0 ¼ C6H3-2,6-i Pr2 ],228b t Bu2 AlN(R)SiPh3 [R ¼ C6H3-2,6-i Pr2 or 1-adamantyl]228c and
MesAl{N(SiMe3)2}2 (Mes ¼ mesityl) and Mes Ga(NHPh)2,228d rotational barriers about the M--N or M--S bonds
in the order of 40 kJ mol1 have been estimated. The heavier nuclei 69 Ga, 71 Ga, 115 In, 203 Tl and 205 Tl have found
relatively limited application in recent NMR studies, although 69 Ga and 71 Ga measurements have been turned to
account, for example, to analyse the environment of Ga(III) in a variety of crystalline oxo compounds.214 By
contrast, the 27 Al nucleus with I ¼ 5/2 has been the focus of numerous studies whose coverage ranges, for example,
from solids such as zeolites with extended framework structures,206–209 through the two different isomers of the
Keggin anion [AlIIIW12O40]5,212 to aluminated brain tissue as analysed in vitro.211 Knowledge of thermodynamic
properties can be gained by using NMR or vibrational spectroscopic measurements to determine how an equilibrium
responds to changes of temperature. For example, IR measurements suggest an enthalpy change of 46(3) and 59(16)
kJ mol1 for the dissociation reactions in Equation 1.17198 and Equation 1.18,229 respectively.
2 [Me2AlH]3 (g)
Ga2H6 (g)
3 [Me 2 AlH]2 (g)
ð1:17Þ
2 GaH3 (g)
ð1:18Þ
Mass spectrometric and photoionisation studies are also important sources of thermodynamic data, some of the
best estimates of dissociation energy for the diatomic molecules Ga2, In2 and GaIn having been secured in
this way.77b
As regards the electronic properties of molecular Group 13 metal (M) compounds, EPR measurements give
a unique instrument for identifying and characterising genuine M(0) and M(II) derivatives, such as M(CO)n [M ¼ Al,
Ga or In; n ¼ 1 or 2],29 AlNH3126 and MH2 [M ¼ Al or Ga],29 as opposed to compounds which owe their
paramagnetism primarily to the ligands. Photoelectron and photoionisation spectroscopies give access to information about the occupied electronic levels in compounds, as well as shedding light on the structural, electronic and
vibrational properties of the states both preceding and following ionisation. Good examples are provided by the UV
photoelectron measurements carried out on the ‘half-sandwich’ molecules (h5-C5H5)M [M ¼ In or Tl]230 and related
indium(I) phospholyls (h5-P2C3t Bu3 )In and (h5-P3C2t Bu2 )In,221 while pulsed-field ionisation zero electron kinetic
energy (ZEKE) photoelectron experiments have been the basis of characterising the dimetal half-sandwich molecule
(h5-C5H5)Al2.75
1.3 Redox Chemistry of the Group 13 Metals: Access to Oxidation States Lower than þ3
1.3.1 Introduction
The pattern displayed by the energies of the valence electrons finds expression in the redox chemistry of the Group 13
metals. Although þ3 persists as the characteristic formal oxidation state for all the members of the Group, the þ1
state gains in importance as the atomic number increases. The situation, evinced for aqueous conditions by the
standard reduction potentials listed in Table 1.1, culminates in thallium with the emergence of a distinct preference
for Tl(I) over Tl(III), which is a powerful oxidising agent under normal conditions. The accessibility and stability of
Tl(I) compounds under most conditions have long been known. Except within the past two decades and apart from
the indium(I) halides, InCl, InBr and InI, and a handful of organoindium(I) compounds,12,22 derivatives of the lighter
Group 13 metals in the þ1 oxidation state have been mainly the stuff of vapours and high-energy conditions. Such
compounds are typically powerful reducing agents prone to disproportionation under normal (aqueous) conditions.
26 New Light on the Chemistry of the Group 13 Metals
Figure 1.3 (a) Structures of the more or less discrete anionic clusters [Tl7]7, [Ga11]7 and [Tl13]11 as found in
K10Tl7, Cs8Ga11, and Na3K8Tl13, respectively. Reprinted with permission from [115e]. Copyright 2000 American
Chemical Society. (b) [100] View of the unit cell of orthorhombic KNa3In9 showing the layered In12 icosahedra
and zigzag chains. Reprinted with permission from [238]. Copyright 2002 American Chemical Society. (c) Building
units of the K39In88 structure: (i) the 20-atom pentagonal dodecahedron [512] about icosahedra A, B, C (turquoise);
(ii) the 28-atom hexakaidecahedra [51264] about clusters C and D (red and grey, respectively); (iii) the K136
clathrate-II network in K39In80. The connections among clusters are not shown. Reprinted with permission
from [117]. Copyright 2003 American Chemical Society. (d) The fullerene-like In74 cluster (blue) and endohedral
Na39 shell (red) connected to equivalent exohedral Na (green ellipsoids) of Na96In97Ni2; the In10Ni core is not
shown. Based on [115d, Figure 18, p. 686] (see colour version of this figure in Colour Plate section)
Redox Chemistry of the Group 13 Metals: Access to Oxidation States Lower than þ3 27
As noted in the preceding section, detection and characterisation of the free molecules then depend on spectroscopic
interrogation of either the vapour itself12,21,22 or of the vapour species trapped in a solid inert matrix at low
temperatures (q.v.).29,96 However, there has also been a dawning recognition that vulnerable M(I) centres can be
preserved under ambient conditions by an appropriate choice of substituent and/or medium. Indeed, the tailoring of
the bulk, specific geometry, electronic properties and charge of ligands has been one of the defining features in the
spectacular development of M(I) and M(II) chemistry for M ¼ Al, Ga and In that has occurred over the past two
decades (Section 1.2.2.3).
1.3.2 Compounds of the metals in integral low oxidation states
The past two decades have witnessed a fever of research activity leading to the synthesis and characterisation of new
compounds in which the metal has a formal oxidation state lower than þ3. These include many M(I) compounds,
ranging from monomeric molecular species, for example GaCl,12,70 or InC6H3-2,6-(C6H2-2,4,6-i Pr3 )2 (1),33 through
weakly bound oligomers or polymers, for example [RM]n [R ¼ h5-C5H5, h5-C5Me5, or C(SiMe3)3],10–12,22,172,231 to
solids reasonably formulated as containing Mþ cations, for example InCF3SO3 or TlCl.12,22,176 The status of M(II)
compounds has likewise changed out of all recognition. These are represented, on the one hand, by paramagnetic,
mononuclear molecules such as MH2 or M(H)CH3 that are short-lived under normal conditions,22,29 and, on the
other, by more robust metal–metal-bonded dimers such as R2MMR2 [M ¼ Al, Ga or In; R ¼ a bulky organic or other
substituent]10,12,22,173,232 and [X3MMX3]2 [M ¼ Ga or In; X ¼ Cl, Br or I].12,22 Nor is the zero oxidation state any
longer confined to the metal itself. While yet to be sustained more than transiently under ambient conditions, species
such as MnL and M2 [e.g. M ¼ Al, Ga or In; L ¼ NH3 or CO] have been established by spectroscopic studies of
vapours at elevated temperatures or of solid inert matrices at low temperatures (Section 1.2). Through all these
developments new chemistry has come to light. For example, the carbene-like character of a molecular M(I)
compound, for example: MCp (Cp ¼ h5-C5Me5), has been revealed by its functioning as a base that coordinates to
conventional Lewis acids, as in Cp MAlt Bu3 ,233 but, more strikingly still, to unsaturated transition metal centres, as
in Cp GaFe(CO)4 and Ni[MC(SiMe3)3]4 [M ¼ Ga or In].10,22,177,234,235
*
*
1.3.3 Metal clusters and the metals in non-integral oxidation states
Molecular MII compounds with a metal–metal bond and discrete M2 diatomic molecules are but preludes to the
spectacular advances that have been made in the realm of Group 13 metal clusters. These fall into two classes: naked
clusters and metalloid clusters.
1.3.3.1 Naked Clusters
Homo- and hetero-atomic anionic clusters (commonly interconnected) feature in Zintl and related intermetallic
phases, which now constitute a burgeoning area of research.11,12,22,115 These are formed by the liaison of the
Group 13 metal with an alkali, alkaline earth or other, more electropositive metal. Figure 1.3a illustrates the
more or less discrete clusters [Tl7]7 (as in K10Tl7),116 [Ga11]7 (as in Cs8Ga11)236 and [Tl13]11 (as in
Na4A6Tl13, where A ¼ K, Rb or Cs),237 Figure 1.3b the phase built of layered indium icosahedra and zigzag
chains in KNa3In9,238Figure 1.3c the three building units which are interlinked in K39In88,117 and Figure 1.3d the
multiply-endohedral Ni@In10@Na37@In70 cluster in Na172In197Ni2.239 Although solids of this sort with their
beguiling, often unprecedented, structures are primarily of academic interest, some of them, nevertheless, hold
promise as thermoelectric materials240 or, in the case of YbGaGe,118 as materials with zero thermal expansion.241 Certain solids have also been shown to contain distinct units that approximate to cationic clusters;22 such
28 New Light on the Chemistry of the Group 13 Metals
is the case, for example, with slightly kinked, chain-like [In5]7þ and octahedral [PtIn6]10þ , as found in the solids
In5Mo18O28114 and Pt2In14Ga3O8F15,103 respectively.
Among the metal clusters formed by laser evaporation or sputtering of the metal or a compound of the metal
(Section 1.2.2.1) it may be noted, in addition to the unusually stable Al13 anion, pyramidal clusters of the type
[AM4], involving an alkali metal A and M ¼ Al, Ga or In, which have been created and studied typically through the
media of negative ion photoelectron spectroscopy and ab initio calculations.38 With an Aþ cation at the apex, such
clusters are noteworthy for the aromatic character imputed to the square M42 basal unit.
1.3.3.2 Metalloid Clusters11,22,30–32
Through the intercession of appropriate bulky ligands, the disproportionation of certain M(I) compounds can be
controlled to give not the metal itself but ligand-sheathed clusters of which the following, illustrated in Figure 1.4, are
typical: [Al50(h5-Cp )12] (5),170 [Al77{N(SiMe3)2}20]2 (6),242 [Ga22R8] [R ¼ Si(SiMe3)3, Ge(SiMe3)3 or Sit Bu3 ]
(7a),243 [Ga22(Pt Bu2 )12] (7b),171 [Ga22{N(SiMe3)2}10]2 (7c),244 [Ga22{N(SiMe3)2}10Br10 þ x]n (x ¼ 1, n ¼ 3;
x ¼ 2, n ¼ 2) (7d),245 [Ga51(Pt Bu2 )14Br6]3 (8)246 and [Ga84{N(SiMe3)2}20]n (n ¼ 3 or 4) (9).31a,247 These
remarkable species, conjured in the past decade by the sophisticated research of Schn€
ockel and his group at
Karlsruhe, are termed ‘metalloid’ to describe their metal-rich condition in which the number of direct metal--metal
bonds exceeds the number of metal–ligand bonds (Chapter 7).
1.4 Bonding Aspects
1.4.1 Introduction
With their valence ns and np orbitals to call upon, Group 13 metal atoms, M, find much common ground with the
heavier Group 14 atoms – silicon, germanium, tin and lead – in their bonding opportunities.8,12 The metals of both
Groups give rise to compounds with discrete metal–metal bonds, although compounds stable under ambient
conditions have come to light for the Group 13 metals only in the past four decades.10–12,22,173,232 As regards their
valence orbitals, species of the type [X3MMX3]2, [X2MMX2]2 and [XMMX]2 [M ¼ Al, Ga, In or Tl;
X ¼ univalent ligand] are formally isoelectronic with the neutral molecules ethane, ethene and ethyne, respectively,
and so the opportunity for multiple M--M bonding arises. Simple monomeric molecules exemplified by X2MNH2 or
XMO must also feature heteronuclear M--N or M--O bonds with some degree of p-character, while being generally
prone to facile aggregation with the formation of M--N--M or M--O--M bridges.
Clusters such as the metalloid species [Al77{N(SiMe3)2}20]2 (6, Figure 1.7)22,30–32,242 or [Ga11]7 (Figure 1.3)
found in intermetallic phases11,12,22,115,236 call for a delocalised treatment of the metal–metal bonding. These invite
comparison with the boron clusters that feature among the hydrides, halides, metal borides and the solid element.2–11
Indeed, Wade’s or related rules,248 first devised for boron clusters, have commonly been invoked to rationalise the
chemical bonding and electron count of some of the clusters formed by boron’s heavier homologues. The
Zintl–Klemm concept115 may not add greatly to the detailed understanding, but has proved to be a useful
rule of thumb for anticipating and interpreting the structures of Zintl-type phases.
The Group 13 metal atoms differ from their counterparts in Group 14 in being significantly less electronegative,
and the bonds they form with most substituents are correspondingly more polar. As a result, their compounds
feature typically substituents that are more nucleophilic and metal centres more electrophilic than do their Group 14
analogues. That Coulombic terms assume greater importance in the bonding of a Group 13 metal atom to a nonmetallic substituent simply reflects the intermediate position occupied by Group 13 between the more electropositive alkali and alkaline earth metals of Groups 1 and 2 and the less electropositive metals of Group 14.
Bonding Aspects
29
Figure 1.4 Examples of some metalloid clusters: (5) arrangement of the 50 Al atoms and 12 Cp (h5-C5Me5)
moieties of Al50Cp 12 showing successive shells of 12 ligand-bearing Al atoms (blue), 30 Al atoms (orange) and the
central core of 8 Al atoms (blue); (6) arrangement of the 77 Al atoms in the anion [Al77{N(SiMe3)2}20]2 showing the
outer shell of 20 ligand-bearing Al atoms (blue), the second shell of 44 Al atoms (metallic grey) and central core of a
single Al atom surrounded by a distorted cuboctahedral/icosahedral array of 12 Al atoms; (7) four different
arrangements of 22 Ga atoms in the cluster species (a) Ga22R8 [R ¼ Si(SiMe3)3, Ge(SiMe3)3 or Sit Bu3 ], (b)
[Ga22R10]2 [R ¼ N(SiMe3)2] (N atoms pink), (c) [Ga22{N(SiMe3)2}10Br10]2 (Br atoms green, N atoms pink) and
(d) Ga22R12 (R ¼ P t Bu2 , P atoms purple]; (8) [Ga51(Pt Bu2 )14Br6]3 (Br atoms green, terminal and bridging P atoms
blue); (9) [Ga84R20]4 [R ¼ N(SiMe3)2, N atoms pink] with its 20 ligand-bearing Ga atoms (blue).32 Reproduced by
permission of The Royal Society of Chemistry (see colour version of this figure in Colour Plate section)
Simple trivalent Group 13 metal molecules of the type MX3, where X is a univalent ligand (e.g. H, a halogen or an
organic group) owe their Lewis acidity not only to the innate electrophilicity of the metal centre, but also to the
energetic accessibility of the Lowest Unoccupied Molecular Orbital (LUMO) concentrated on the metal.
Accordingly, the molecules are susceptible either to coordination by a donor species or to aggregation via M–X–M
bridges, which may be regarded as self-complexation. Again, Coulombic factors are prominent in the coordinate
links or delocalised bridging moieties formed in this way.8,12,51,52
30 New Light on the Chemistry of the Group 13 Metals
In contrast with MIII compounds, MI compounds at large are characterised by the relatively weak Lewis acidity but
significant Lewis basicity of the metal centre and by their unsaturation.22 The lone pair of electrons originating in the
ns2 valence shell of the free Mþ ion and concentrated largely on the metal is the prime source of basicity in
a monomeric MI molecule such as (h5-C5Me5)M: which, as a s-donor and p-acceptor, resembles CO in its ability to
bind as a terminal or bridging ligand to electron-rich transition metal centres, as in M(InR)4 [M ¼ Ni, Pd or Pt;
R ¼ h5-C5Me5 or C(SiMe3)3]178,249 or Fe3(CO)10(m-InR)2 [R ¼ h5-C5Me5].178,250
Where the MI centre approximates to an Mþ ion, as in the indium monohalides, for example, it is usually found in
an environment of high coordination number, where it is seldom easy to differentiate between primary and secondary
metal–ligand contacts. How diffuse the environment can be is well illustrated in Figure 1.5 for InX7 [X ¼ Cl, Br or I],
InBr9 and InF15 polyhedra met in solid InX with the TlI structure type,251 orthorhombic In2Br3251 and
Pt2In14Ga3O8F15,103 respectively. A second, highly distinctive feature is the non-uniform nature of the environment.
With its implication of limited stereochemical activity on the part of the lone pair on the metal centre, this is
a characteristic shared in varying degrees with other metal ions such as Sn2þ and Pb2þ formally possessing an ns2
valence shell.22,252 Theoretical treatments have been ventured at varying levels of sophistication and appealing, for
example, to the second-order Jahn-Teller effect to rationalise this behaviour.22,252,253
The distinction between primary and secondary bonding in the condensed phases can be relatively arbitrary.
Intermolecular interactions range from the general, as represented by Coulombic dispersion forces, to the specific, of
Figure 1.5 Representative coordination geometries
of the InI centres in more or less ionic indium(I) compounds:
þ
(a)
coordination of In by Br (up to 4.50 A) within InBr; (b) coordination of one of the Inþ ions by Br (up to 4.50
A) within In2[In2Br6]; (c) InF15 polyhedron within Pt2In14Ga3O8F15. (a) and (b) Reprinted with permission
from [251]. Copyright 1994 American Chemical Society. (c) Based on [103, Figure 3, p. 2280]
Bonding Aspects
31
which hydrogen bonding is best known.69 In the following sections, some of these themes are developed and
illustrated in more detail.
1.4.2 Metal–metal bonding
Simplest of binary Group 13 compounds expected to have as their core a metal–metal single bond are the
tetrahydrides, H2MMH2. Although molecules with this stoichiometry have been identified for M ¼ Al, Ga or In in
matrix-isolation experiments,65,130,136 they do not appear to survive under normal conditions. Nevertheless,
quantum chemical calculations21,22,254–256 are noteworthy for what they reveal about the conditions required to
sustain an M--M bond. They identify potential energy minima corresponding to the classical D2h and D2d structures
10 and 11. Exploration of the M2H4 hypersurface shows 11 to be the global minimum for M ¼ B, in keeping with the
stable existence of B2X4 molecules [X ¼ F, Cl, Br or I],257 but for the metals M ¼ Al or Ga (and presumably M ¼ In or
Tl), the global minimum adopts not this form, but the tri-hydrogen-bridged structure 12 with C3v symmetry. This can
be described in terms of more or less Coulombic interaction between a cationic MI centre and a tetrahedral, anionic
MIIIH4 one, thereby implying disproportionation. Structure 12 is then about 45 and 35 kJ mol1 more stable than
structure 10 for M ¼ Al and Ga, respectively,
according to MP2 calculations.254 The distance between the two metal
A (Ga)] is actually calculated to be similar to, or even somewhat shorter than, that
atoms in 12 [2.482
A (Al) and 2.563
in 10 [2.608 A (Al) and 2.554 A (Ga)], despite the former’s exhibiting no direct M--M bond.254–256
H
H
M
H
H
M
M
H
H
10
M
M
H
H
11
H
H
M
H
H
12
The hydrides are merely paradigms for all M2X4 systems, and there is generally a fine energy balance between
a M--M-bonded structure akin to 10 or 11 and a charge-separated mixed valence one akin to 12.12,22,258 Which form is
favoured depends not only on the bulk and electronic properties of the substituent X, but also on the phase and choice
of solvent. The M--M bond is weakened by Coulombic repulsion between the partial positive charges residing on
each of the metal atoms, a factor likely to be accentuated as the electronegativity of X increases. Conversely, any
form of ligation that increases the electron density on the metal serves to strengthen the M--M bond. On the other
hand, the charge separation implicit in the mixed valence structure is favoured by the increased permittivity of the
solid state or of a polar solvent. There is, in addition, a significant activation barrier to the interconversion of the M--M
bonded to the mixed valence structure, reflecting the ability of X to act as a bridging ligand between the two metal
centres, and which decreases in the order X ¼ halogen > hydrogen > organic group. In keeping with these general
considerations, it is found that the dihalides MX2 [M ¼ Ga or In; X ¼ Cl, Br or I] adopt mixed valence structures in
the solid state,12,22 whereas M2X4 compounds with an M--M-bonded structure and tri-coordinated metal atoms are
confined mainly to systems incorporating bulky organic or pseudo-organic substituents,10,12,22,173,232 for example
X ¼ CH(SiMe3)2, Si(SiMe3)3 or C6H2-2,4,6-i Pr3 . Preservation of the M--M bond can otherwise be achieved only by
expanding the coordination shell about, and increasing the electron density at, the metal atoms through the addition
of neutral or anionic donor species, as in [M2X6]2 and M2X4L2 [M ¼ Ga or In; X ¼ Cl, Br or I; L ¼ dioxane,
pyridine or PR3].12,22
All the evidence available argues that these M--M bonds are quite weak. One theoretical estimate of the M--M
dissociation energy of X2MMX2 molecules gives values in the range 69–258 kJ mol1 for M ¼ Al > Ga > In > Tl
and X ¼ H > F > Cl > I, in marked contrast to the corresponding B–B dissociation energy (389–454 kJ mol1).259
The M–M distances measured for each metal in crystalline compounds span quite
a wide range. Compounds of the
Ga–Ga
distances
measuring
2.541(1)
A
[R
¼ CH(SiMe3)2],260 2.599(4) A
type R2GaGaR2 feature, for example,
[R ¼ Si(SiMe3)3]261 and 2.515(3) A (R ¼ C6H2-2,4,6-i Pr3 ).262 Intriguingly, though, the same distances in
32 New Light on the Chemistry of the Group 13 Metals
Li2[Ga2Cl6],263 Ga2[Ga2I6]264 and Ga2Br42thf265 are significantly shorter, that is 2.392, 2.387(5) and 2.412 A,
respectively. Hence it appears that the nature of the substituents takes precedence over both the coordination number
at the metal centres and the net charge carried by the assembly in dictating the length, and presumably the strength, of
the M–M bond. Steric effects undoubtedly play a part, but a number of other factors, such as charge separation,
electrostatic repulsion and rehybridisation, need also to be taken into account according to one analysis.232b Density
functional calculations carried out on digallanes of the type R2GaGaR2 indicate a Ga–Ga bond distance that follows
the order R ¼ SiH3 > CH3 > H > NH2 > HNCH¼CHNH, shortening overall by about 10%.266 According to
population analysis, electronegative substituents such as amide exert a –I effect on the metal atoms, leading to
a marked increase in the positive charge they carry. Far from leading to attenuation of the Ga–Ga bond, however, this
has the effect of increasing the relative s-content of the Ga–Ga s-bonding orbital, with a concomitant contraction of
the bond.
An even simpler species with a direct metal–metal bond is the dimer M2, which has been studied in the gas phase
or in a solid noble gas matrix (q.v.).63,76,77,131,132 Here, the M–M bond is even weaker, the best estimates of the
dissociation energies (D298 kJ mol1 ) being13,22,76,77,132 Al2 133 5.8, Ga2 114.5 4.9, In2 78.1 5.7 and Tl2 59.
3 Unlike B2 with its Sg ground electronic state, dissociation energy of 290 kJ mol1 and appreciable
p-bonding,13,267 therefore, the heavier homologues present little evidence of what is usually understood as
‘multiple’ bonding. Perhaps the most noteworthy feature of the molecule Ga2 is the accessibility of excited
electronic states with energies close to the 3 Pu ground state having the valence shell configuration s2g s2u s1g p1u and
1 þ
a formal bond order of one. For example, there are 3 S
g and Sg electronic states which lie, according to CASSCF/
SVP calculations, at energies no more than 7.1 and 19.0 kJ mol1 above the 3 Pu ground state.67 It is this property
which sets Ga2 apart from the bare gallium atom, and makes it so much more reactive with respect, for instance, to
the hydrogen molecule.63,67
The metal–metal bonding in metalloid clusters cannot be understood in terms of any single electron counting
rule.30–32,93,248 Here it is necessary to treat each species individually by appropriate quantum chemical calculations,
although a qualitative description of the topology can often be developed through comparisons with modifications of
the bulk metals themselves. For ligand-free clusters such as Al13, a model commonly applied regards the nuclei and
innermost electrons as forming a positively charged core with an essentially uniform potential.268 All of the valence
electrons from the individual atoms in the cluster are then subjected to this potential, and the so-called ‘jellium’
electronic shell structure emerges, with stable configurations of electrons (2, 8, 18, 20. . .) that differ from those of the
atomic series (2, 10, 18. . .). Such clusters have been described as ‘super atoms’. On this basis, Al13 acts like
a superhalogen, so that Al13I resembles a dihalogen and Al
13 assumes particular stability in resembling a
halide anion.
Weak metal metal interactions between molecules of the metals in the formal oxidation state þ 1 are a common
feature in the solid state, with M M distances close to, or more often significantly longer than, those in the bulk
metal. The interactions do not survive evaporation or dissolution, even in a weakly coordinating solvent. As noted
previously, then, it is more appropriate to regard them as secondary rather than primary interactions, and to defer
further discussion to Section 1.4.5.
1.4.3 Multiple bonding36
‘Oh, let us never, never doubt
What nobody is sure about!’
Hilaire Belloc, More Beasts for Worse Children, ‘The Microbe’, 1897
In principle, the neutral binuclear dihydrides, M2H2, afford the opportunity for the formation of a multiple bond
between the two Group 13 atoms. Like the tetrahydrides, M2H4, these do not survive under ambient conditions, but
have been identified by their vibrational spectra in matrix-isolation experiments.63,65,66,130,136 Diborene, B2H2,
provides a model. Experimental and theoretical studies indicate a linear structure 13 for the triplet ground
Bonding Aspects
33
1 þ
electronic state 3 S
BH fragments) is about 450 kJ mol1
g ; the dissociation energy De (with respect to two S
associated with the presence, in effect, of a s- and two one-electron p-bonds.269 On the other hand, a singlet bis
(m-hydrido) structure 15 with D2h symmetry is found to be the most stable form of M2H2 for all the heavier Group
13 elements;21,108,270 IR bands attributable to this species have been observed in matrix-isolation experiments with
M ¼ Al,130 Ga63 or In.63 Instead of the linear form 13, which is now only a transition state, a singlet trans-C2h
isomer 14 appears as a second minimum on the potential energy surface for M2H2 with M ¼ Al, Ga, In or Tl.
Matrix experiments have shown that the cyclic isomer 15 with M ¼ Ga or In can be photo-reversibly converted to
the trans-bent form 14.63 There are other possible isomers, including M–MH2 (16) with C2v symmetry, which lie at
energies not far above the global minimum; indeed, appropriate photoexcitation of Ga(m-H)2Ga or HGaGaH is
H
M
M
H
M
M
M
H
13
H
H
H
14
M
M
M
H
15
H
16
is calculated to lie at an energy 48 kJ mol1 above
found experimentally to yield GaGaH2.63 Trans-bent HGaGaH
Ga(m-H)2Ga, to feature a relatively long Ga–Ga bond (2.620 A) and to dissociate into two singlet GaH fragments at
an energy cost of only 57 kJ mol1. The difference in the bonding properties of HBBH and trans-bent HGaGaH
can be rationalised on the basis of the energy needed to promote the MH fragments from the singlet ground state
into the triplet excited state. A prerequisite for the linear structure with appreciable p-bonding is then that the
energy gain achieved by p-bond formation exceed this excitation energy, a condition plainly met for HBBH but not
for the heavier homologues. The bonding situation in trans-bent HMMH species is best described as a double
donor–acceptor interaction between the two MH molecules, as represented in 17.
Derivatives of trans-bent M2H2 that survive and can be characterised under normal conditions have been
synthesised by replacing the hydrogen atoms by sterically encumbered organic ligands R.36 One such example is
ArGaGaAr, where Ar ¼ C6H3-2,6-(C6H3-2,6-i Pr2 )2.36d,271 As predicted for HGaGaH, this molecule adopts a transbent geometry with a relatively long Ga–Ga bond (2.6268(7) A) and Ga–Ga–C angles of 123.16(7) . The weakness
of the metal–metal bond is revealed by the behaviour of the compound in a hydrocarbon solution where the redbrown dimer is in equilibrium with the green monomer. The corresponding indium and thallium compounds follow
the same pattern,167,176 although the aluminium compound is known only as an intermediate which undergoes a fast
[2 þ 4] cycloaddition reaction with the toluene solvent.272 If the steric demand of the aryl groups is advanced still
further, the resulting compound is exclusively monomeric, as exemplified by the cases of M[C6H3-2,6-(C6H2-2,4,6i
Pr3 )2], with M ¼ Ga36d or In (1).33
It is with the dianions [HMMH]2 and their derivatives, however, that the noise of the multiple bonding battle
has truly rolled.36 As regards valence electrons, these are after all isoelectronic with HC:CH. The
compound Na2[ArGaGaAr], with R ¼ C6H3-2,6-(C6H2-2,4,6-i Pr3 )2, initiated the debate.168 The structure of the
deeply coloured crystals reveals for the dianion not a linear but a planar, trans-bent C–Ga–Ga–C skeleton with
34 New Light on the Chemistry of the Group 13 Metals
<Ga–Ga–C ¼ 133.5 and an unusually short Ga–Ga bond (2.319(3) A). In reality, though, the crystalline compound
should be seen as an ion triple in which the Naþ counter-ions reside on either side of the Ga–Ga bond. A second
compound of this sort is Na2[Ar0 GaGaAr0 ] [Ar0 ¼ C6H3-2,6-(C6H3-2,6-i Pr2 )2],271 and the aluminium analogue
Na2[Ar0 AlAlAr0 ] has also been characterised.174 The trans-bent structure of the anionic unit in all of these
compounds, as illustrated in Figure 1.6, is consistent with the development of lone pair electron density at the metal
Figure 1.6 (a) Thermal ellipsoid plot (30% probability) of crystalline Na2[Ar0 AlAlAr0 ] where Ar0 ¼ C6H3-2,6(C6H3-2,6-i Pr2 )2 omitting the hydrogen atoms. (b) Representations of selected Kohn–Sham orbitals for the model
species Na2[Ar00 AlAlAr00 ] where Ar00 ¼ C6H3-2,6-Ph2.174 Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission (see colour version of this figure in Colour Plate section)
Bonding Aspects
35
centres, a view confirmed by quantum chemical calculations.36,273 Kohn–Sham frontier-orbital surfaces for the
model compound Na2[Ar00 AlAlAr00 ] [Ar00 ¼ C6H3-2,6-Ph2] generated from DFT calculations are depicted in
Figure 1.6. The Highest Occupied Molecular Orbital (HOMO) is clearly an out-of-plane p-bond; HOMO–1 has
been called a ‘slipped p-bond’ although essentially non-bonding in its effect; and HOMO–2 features s-bonding
between the metal centres and from the metal centres to the ligand carbon atoms.174 The Al–Al Wiberg bond
order estimated for Na2[Ar0 AlAlAr0 ] is 1.13, a value much at odds with the triple bond first claimed for
Na2[ArGaGaAr]. The bonding picture is further clouded by the likely role of the Naþ counter-ions. For example,
the calculated hydrogenation enthalpies for [HGaGaH]2 and Naþ 2[HGaGaH]2 give grounds for believing that
the latter should be described as a Ga2Na2 cluster.274 Such a view gains support from the successful synthesis and
characterisation of other cluster compounds of Group 13 metal and alkali metal atoms, for example
Na2Ga4(Sit Bu3 )42thf 275a and K2Ga4[C6H3-2,6-(C6H2-2,4,6-i Pr3 )2]2.275b It has also been suggested that the
shortness of the M–M distance in compounds of the type Na2[ArMMAr] is due in part to interaction between
the Naþ ions and aromatic rings.276
No doubly reduced derivatives containing anions of the type [R2MMR2]2 [R ¼ H or an organic group] are known
at present.36 Attempted reduction of tetra-alkyls and -aryls of aluminium and gallium has led to the isolation of the
singly reduced radical species [R2MMR2] [M ¼ Al or Ga; R ¼ CH(SiMe3)2 or C6H2-2,4,6-i Pr3 ], which have been
crystallised as solvent-separated ion pairs with a variety of counter-cations.36a The crystal structures show anionic
moieties with more or less planar C2MMC2 skeletons and M–M bond distances up to 0.18 A (about 7%) shorter than
in the corresponding neutral precursors.
Other possible candidates for multiple bonding are the molecules H2MNH2 [M ¼ Al, Ga or In] formally
valence-isoelectronic with ethene. These have been characterised in matrix-isolation experiments as planar,
C2v-symmetric molecules by their IR spectra.98 A criterion frequently used to analyse double bond character in
molecules of this sort is the barrier to rotation about the bond, and quantum chemical calculations for H2MNH2
indicate barriers of 51, 66 and 51.5 kJ mol1 for M ¼ Al, Ga and In, as compared with values of 272 and
162 kJ mol1 for H2CCH2 and H2BNH2, respectively.98 Neither the hydrides nor their derivatives persist as
monomeric species, but rather aggregate as oligomers under normal conditions unless very bulky ligands are
introduced in place of some or all of the hydrogen atoms. For example, dimerisation of H2MNH2 to give
cyclic products with a D2h-symmetric M(m-N)2M core is associated with an enthalpy change of 237 and
245 kJ mol1 for M ¼ Al and Ga, respectively, according to coupled-cluster CCSD(T) calculations.277
Monomeric compounds of the type R2 MNR0 2 have been synthesised and characterised under normal conditions
for M ¼ Al or Ga in the cases where the substituents are such bulky ligands as mesityl, C6H2-2,4,6-t Bu3 or
C6H3-2,6-(C6H2-2,4,6-i Pr3 )2.36a The molecules are distinguished by more or less planar skeletal geometries and
quite short M–N bonds, but no shorter than the difference in electronegativity might lead one to expect.
Experimental and theoretical estimates of the rotational barrier about the M–N bond find values typically in the
range 40–45 kJ mol1.36a,169,228,278
The question naturally arises: why are the p-interactions in all these homonuclear and heteronuclear bonds formed
by the Group 13 metals so weak? As with analogous compounds of the heavier Group 14 elements,36a,279 part of the
answer lies in the M–M distance associated with the minimum energy state for pure s-bonding, and which lengthens
inexorably as the atomic number of M increases. A significant p-interaction can be established only if the two metal
atoms are relatively close together. Yet, closer approach of the two metal atoms is countered by a substantial energy
penalty from the repulsive branch of the s-bond. The trans-bent geometry displayed, for example, by [RMMR]2
anions can be interpreted on the basis of a second-order Jahn-Teller mixing of the s and bonding p-orbitals (and
possibly also the p and bonding s-orbitals) of the hypothetical linear geometry.36e,280 The result of such mixing is
a build-up of lone pair character at the metal centres. Thus, electron density is shifted from the p-symmetric orbitals
to non-bonding orbitals. As the atomic number of M increases, so the M–M–R angle decreases, a change reflecting
the enhanced mixing of s- and p-character as the energies of the relevant orbitals crowd closer together. In the case of
heteronuclear bonds, such as those found in molecules of the type R2 MNR0 2 , there is an additional factor
contributing to the weakness of the p-bond, namely the reduced orbital overlap arising from the distinctly polar
character of the M–N bond.281
*
36 New Light on the Chemistry of the Group 13 Metals
1.4.4 Coordinate links to MIII compounds
Attempts in the past to assess relative Lewis acidities have served only to emphasise the variety of factors
contributing to the effectiveness of coordination and the strength of the coordinate link.8,12,52 So many are the factors
at work – the electronic properties, geometry and bulk of the substituents at the acid and base centres – that it is
doubtful whether any meaningful quantitative scale of acidity can be drawn up for the different metal centres M ¼ Al,
Ga, In or Tl. The issue is, nevertheless, still being actively pursued through both experimental and theoretical
studies.282–284 Such studies have commonly focussed on a 1:1 complex of the type X3MEY3, where M and E are
Group 13 and Group 15 elements, respectively, and X and Y are substituents such as hydrogen, halogen or methyl.
The formation of the complex is then accompanied by reorganisation of the free acid and base fragments, with MX3
changing from a planar to a pyramidal geometry with a lengthening of the M–X bonds and the pyramidal EY3
typically assuming a steeper pitch (Equation 1.19). Recent quantum chemical studies of the complexes X3MEX3
X
+
M
X
X
Y
X
Y
M
E
Y
Y
Y
ð1:19Þ
E
X
X
Y
[M ¼ B or Al; E ¼ N or P; X ¼ H, methyl or Cl] and Ar3MD [M ¼ B, Al or Ga; Ar ¼ C6H5, C6H4F or C6F5; D ¼ NH3,
H2O, PH3, H, CH3 or F] indicate, inter alia, that the bond dissociation energy is not necessarily a good measure
of intrinsic Lewis acid or base strength since the reorganisation energies of the fragments contribute significantly to
this energy.285 A more reliable measure is provided by the interaction energy between the fragments when frozen in
the geometries they adopt in the complex. These interaction energies, unlike the conventional dissociation energies,
are then found to be generally larger for M ¼ B than for M ¼ Al, and to vary with the bases in the order
EMe3 > EH3 > ECl3. The relatively small reorganisation energy of the cage-like amine quinuclidine, HC
(CH2CH2)3N, accounts for the comparative stability of its complexes with MH3, where M ¼ B, Al or Ga.155 A
more extensive theoretical study of complexes of the type X3MD [M ¼ Al, Ga or In; X ¼ F, Cl, Br or I; and D ¼ EH3,
PX3 or X with E ¼ N, P or As] concludes that the conventional dissociation energies decrease in the orders
X ¼ F > Cl > Br > I; Al > Ga < In; and N P As.286 Strikingly, there is no correlation between the dissociation
energy and the degree of charge transfer, although correlations are found between the energy, the distortion of the
MX3 fragment wrought by complexation, and the length of the coordinate link.
1.4.5 Modes of bonding peculiar to monomeric MI compounds
Simple monohalide molecules, such as AlCl or GaI, are characterised by relatively strong bonds.12,13,18 Taking the
chlorides MCl, for example, the dissociation energies, D298 , are found to be 502, 463, 436 and 373 kJ mol1 for
M ¼ Al, Ga, In and Tl, respectively, the values exceeding in each case those for the chloride molecules of the
corresponding Group 1 and Group 11 metals.22 The bonds must draw some of this strength from (p ! p)p-bonding.
In simple valence terms, such molecules are isoelectronic and isolobal with CO. Ab initio MP2 and CCSD(T)
calculations confirm this analogy for a monochloride such as GaCl,68 revealing the frontier orbitals to be as
portrayed in Figure 1.7. Thus, the HOMO housing the ns2 electron pair, confined largely to the metal and having
s-symmetry, lies at higher energy than the corresponding orbital of CO, and clearly invites the prospect of significant
s-donor action. Interactions between the valence p-orbitals on the metal and halogen atoms then give relatively
strongly bound s- and p-orbitals, localised mainly on the halogen and high-energy orbitals of mainly metal np
character that include the p-symmetry LUMO. This, in turn, suggests the possibility of the molecule’s acting as a pacid. The expectation that such a molecule could bind effectively to an electron-rich transition metal centre is amply
Bonding Aspects
37
Figure 1.7 (a) Frontier orbitals of GaCl (1 S). (b) Potential energy curves calculated for GaCl in its 1 S ground and
3
P excited electronic states.68 Reproduced by permission of The Royal Society of Chemistry
realised in the compound [(h5-C5Me5)Fe(dppe)(GaI)]þ [BArF4 ] (dppe ¼ Ph2PCH2CH2PPh2; ArF ¼ C6H3-3,5(CF3)2), being the first authenticated example of a transition metal complex featuring an MIX molecule as
a terminal ligand.35 In one form or another, the valence orbital properties described for GaCl are characteristic
of all MIX-type molecules formed by the Group 13 metals (M) with either a halogen or an organic group (X),22,35,54–
60
although the relative energies of the HOMO and LUMO vary from metal to metal and substituent to substituent.
Organo-Group 13 metal compounds of the type MIR are now known to function as either terminal or bridging
ligands in a variety of organotransition metal10,22,177 and even organoactinide metal compounds, of which
[(Cp Ga)4Rh(GaCH3)]þ [BArF4] [Cp ¼ h5-C5Me5; ArF ¼ C6H3-3,5-(CF3)2]287a and (h5-C5H4SiMe3)3UAl(h5C5Me5)287b are recently reported representative examples. That this should take up pyridine to form
a complex in which the nitrogen base is bound to the gallium atom of the GaCH3 fragment demonstrates the
residual electrophilicity of the metal centre. The anionic gallium(I) heterocycle [:Ga{N(Ar)CH}2] [Ar ¼ C6H3-2,
6-i Pr2 ] acts in the same way as organogallium(I) compounds forming with Group 10 transition metals complexes such
as cis- or trans-M[Ga{N(Ar)CH}2]2(PEt3)2 [M ¼ Ni, Pd or Pt].160b Compounds of the univalent Group 13 metals
are also prone to undergo oxidative addition. Such is the case, for example, with the neutral gallium(I) heterocycle
[:GaN(Ar)C(Me)CHC(Me)]NAr [Ar ¼ C6H3-2,6-i Pr2 ], Ga(DDP), two molecules of which add to GaX3 [X ¼ Me or
Cl] to form the trigallane (DDP)(X)Ga–Ga(X)–Ga(X)(DDP).161f
Matrix-isolation experiments have identified not only monomeric AlCl but also the dimer Al(m-Cl)2Al, and the
trimer, which appears to have the unusual structure 18 in which two aluminium atoms are bridged by a chlorine and
Cl
Cl
Al
Cl
Al
Al
18
an aluminium atom.142 The experiments also show that AlCl and related MIX species have a rich photochemistry.288
UV irradiation results in a metal-centred transition of an electron from the lone pair in the HOMO into a molecular
38 New Light on the Chemistry of the Group 13 Metals
orbital of mainly metal np-character.88 The latter orbital has the right symmetry to interact with s -orbitals in
molecules such as hydrogen or methane. Transfer of electron density from the metal into these anti-bonding orbitals
results in activation and cleavage of the bond with the insertion of the metal atom. Thus, photoactivation of AlCl in
the presence of H2 results in the formation of the planar, C2v-symmetric molecule H2AlCl.139 Although the
corresponding reaction with InCl is endothermic, it can be engineered, nevertheless, by UV photolysis under matrix
conditions.289 Photoactivated AlCl also inserts into a C–H bond of methane or acetylene (HC:CH) to form the
corresponding R(H)AlIIICl species [R ¼ CH3 or HC:C],290 and adds to O2 to form the peroxide ClAl(h2-O2) and
bis-superoxide ClAl(h2-O2)2.291 That organometal(I) compounds can respond in a similar fashion to photoactivation
is demonstrated by the case of (h5-C5Me5)Al, which inserts under matrix conditions into the H–H bond of H2.140
1.4.6 Secondary bonding
Secondary bonding is a common feature in solid derivatives of the metal hydrides. For example, whereas the
complex Me3NGaH3 is a monomer in the solid as well as the vapour phase,292 the analogous aluminium
compound is a monomer only in the vapour phase, forming a crystalline solid containing dimeric units in which the
Me3NAlH3 molecules are bound together via two highly unsymmetrical Al–H Al bridges.293 As noted earlier
(Section 1.2.3), significant secondary interactions may also arise through ‘dihydrogen’ bonding between the protic
and hydridic hydrogen atoms that coexist in compounds such as H3NBH3294,295 and [H2NGaH2]3.187 Recent
studies suggest that these N–H H–M bonds [M ¼ B, Al or Ga] may compare in strength with conventional
hydrogen bonds as found, say, in solid phenol.296 The similarity appears to extend also to the makeup of the
dihydrogen bonds, with the subtle difference that they involve rather weaker Coulombic interactions than do
conventional hydrogen bonds.
Weak secondary MI MI bonding is a feature of compounds of the metals in the univalent state,22,297 as revealed
by the crystal structures of compounds as varied as [(h5-C5Me5)Al]4,298 [(h5-C5Me5)M]6 [M ¼ Ga or In],181
In6La10O6S17299 and InGaSe2.300 In the tetramer [(h5-C5Me5)Al]4, the metal atoms form a regular tetrahedron with
Al Al edges measuring 2.769 A, somewhat longer than the Al–Al bond in [(Me3Si)2HC]
2AlAl[CH(SiMe3)2]2
(2.660 A),301 but rather shorter than the shortest Al Al contacts in the bulk metal (2.86 A).12 To understand the
bonding in the tetramer, the splitting of the s-type orbitals hosting the lone pair on each (h5-C5Me5)Al fragment into
a t2 (anti-bonding) and a1 (bonding) set has to be considered. The t2 set is stabilised by an interaction with the empty
p-orbitals forming the LUMO in each of the monomers, and this is the only interaction serving to stabilise the
tetrameric assembly. Accordingly, the unit really depends on no more than secondary interactions between the metal
atoms. In the cases of the hexamers [(h5-C5Me5)Ga]6 and [(h5-C5Me5)In]6,181 the binding of the roughly octahedral
M6 cores [M ¼ Ga or In] is even weaker, with shortest metal metal distances of 4.073–4.173 A (Ga) and
3.942–3.963 A (In) much longer than the shortest metal metal contacts in the respective metals. Indeed, the
finer details of the structures are determined as much by ligand ligand repulsion as by metal metal binding.
1.5 Solid Compounds with Specific Electronic, Structural or Other Properties
1.5.1 III–V Compounds
Much of our greatly increased knowledge of Group 13 metal chemistry has stemmed, directly or indirectly, from
the applications the metals find in III–V semiconductor compounds.12,46–50,302 These range from the wide band
gap nitrides, of which GaN is the archetype and the current mainstay of so many light-emitting diodes (LEDs);48
through InP and related compounds widely used for the fabrication of optoelectronic devices operating at longer
wavelengths and high-frequency devices;49 to compounds such as InAsxSb1x which may offer the limit in
III–V device performance in terms of speed, long wavelength for optoelectronics, and quantum effects related to
low effective mass.302 Thus, GaAs is well established as a key component of certain devices, such as monolithic
microwave integrated circuits (MMICs), IR light-emitting diodes (LEDs), laser diodes and solar cells and
Solid Compounds with Specific Electronic, Structural or Other Properties
39
detectors.50 These exploit in varying degrees the electronic advantages that GaAs enjoys over silicon, that is
its possession of a direct, as opposed to an indirect, band gap, the higher saturated electron velocity and
electron mobility (allowing transistors made from it to function at frequencies in excess of 250 GHz) and higher
breakdown voltage.
Although some of the most efficient solar cells constructed up to now have depended on GaAs as the absorber
layer, there is a problem, not shared by silicon, with the electrical passivation of GaAs surfaces. One way of
countering this requires the presence of an epitaxial capping layer, such as AlxGa1xAs; another, described more
recently,303 involves chemical functionalisation of the surfaces through Cl-termination followed by reaction with
hydrazine or sodium hydrosulfide. Only with the availability of high-quality InP was the rapid development of optical
fibre telecommunications made possible during the 1980s and early 1990s.49 Modern telephone trunk networks and
integrated IT systems (including telephony, interactive services and broadcast distribution) rely on InP technology.
The last 20 years have also witnessed particular interest in the production of semiconductor nanocrystals (or
quantum dots),304 which may, for example, have the form of one-dimensional nanostructures (i.e. nanotubes). By
virtue of the reduced dimensions, such materials display novel optical, electronic and other properties that may be
turned to advantage, as in tunable fluorescent markers in biological imaging, display devices, photovoltaics and
lasers. For example, colloidal InP quantum dots, rods and wires of tailored dimensions have been successfully
grown.305 InP quantum dots offer important advantages over the CdSe ones that are the current workhorses; with
a ZnS shell for chemical passivation, these have been shown to be efficient, stable emitters of radiation spanning the
blue to near-infrared regions. InAs quantum dots, by contrast, have been optimised for biological imaging in the nearinfrared.305d Nor is activity by any means restricted to III–V compounds. Certain III–VI phases have also attracted
much attention, with InSe nanowires showing especial promise as high-performance photodetectors.305e Signs of the
huge research investment that has been, and continues to be, made in the production, processing and application of
III–V compounds are to be found in a wealth of books, review articles, research reports and original papers.46–50,306
1.5.2 Oxygen-bridged framework solids
The strength of the bonds formed by oxygen with M(III) centres makes them highly resistant to thermal or other
forms of activation.12,16,39 As a result, binary, ternary, quaternary or more elaborate systems containing these
bonds, for example oxo-, hydroxo-, silicato- and phosphato- derivatives of M(III), are typically solids composed
of more or less rigid oxygen-bridged networks that are resistant to both dissolution and thermal rupture. The
structures are varied and often complex. In part, this reflects the incompatibility of the M3þ cation and the O2
or OH anion in each establishing regular sub-lattices that can be successfully intermeshed so as to achieve an
isotropic 3D array with the necessary charge balance. The complexity also reflects the multiplicity of ways in
which MIIIO4, MIIIO6 and even MIIIO5 fragments can be linked together via oxygen bridges to produce nets with
different dimensionalities; the possibility of replacing some of these fragments, say, by SiO4 or PO4 units merely
adds to the embarrassment of options, ‘life’s business being just the terrible choice’. Yet the complexity yields
some compounds, of aluminium in particular, that are of exceptional scientific interest as well as being of
immense technological significance, for example as refractories, abrasives and ceramics, as catalysts and
catalyst supports, as sorbents and hosts for gas storage, ion exchangers and chromatographic media, as
insulators on the one hand and fast-ion solid state conductors on the other, and as cements.8,12,39–45
One of the closest approaches to regularity is presented by a-Al2O3 (and a-Ga2O3) which has the corundum
structure composed of hcp close-packed layers of O2 ions with two-thirds of the octahedral interstices occupied
statistically by the cations.8,12,39 More often, defect structures are adopted and/or the structures are anisotropic with
one- or two-dimensional M–O frameworks exhibiting varying degress of openness. Thus, g-Al2O3 typically assumes
either a cubic or tetragonal version of a defect spinel-type structure which, with a favourable combination of surface
area, pore volume, pore size distribution and acid/base characteristics, leads to its importance as a catalyst and
catalyst support in the automotive and petroleum industries.307 Layer structures are exemplified by diaspore, a-AlO
(OH), and its gallium analogue with structures in which the oxygen atoms are arranged in hcp; continuous chains of
40 New Light on the Chemistry of the Group 13 Metals
edge-shared MO6 octahedra are then stacked in layers, which are further interconnected by hydrogen bonds.8,12,39
By contrast with a-Al2O3, the normal form of In2O3 has the C-type M2O3 structure, which is related to fluorite but
with one-quarter of the anions removed, thereby reducing the coordination number of the metal from eight to six.
This compound, an n-type semiconductor, is of interest for numerous device applications, mainly as a result of its
combining high electrical conductivity with transparency to visible light.308 InO(OH) differs from AlO(OH) in
assuming a deformed rutile structure rather than a layer lattice;8,12 nanocrystals have recently been grown with
a variety of different structures and morphologies, such as nanocubes, nanorods, multipods and nanoparticles.192 Socalled ‘beta-alumina’ with the idealised formula Na2O11Al2O3 has a layer structure where the sodium ions are
present in discrete layers separated by spinel-like Al–O layers, an arrangement that permits facile two-dimensional
diffusion of the Naþ ions.8,12,39,309
It is not just the diversity and complexity of the structures to which the preceding examples allude. A particularly
distinctive and significant feature is the prevalence of defects, voids or channels that pervade the solid structures of
this class of compound. A case in point is provided by porous anodic alumina, which has been exploited for the
templated synthesis of metal nanowires.310 The refractory, porous material – often called ‘alumina membrane’
because of its transparent, thin, sheet-like character – consists of a self-assembled honeycomb network of uniformly
sized parallel channels. Arrays of aligned nanowires of uniform diameter can then be created reproducibly by the
filling of these channels with metal atoms, an objective successfully achieved by electrochemical deposition, by
degradation of an inorganic or organic derivative of the metal (after immersing the template in a suitable solution), or
by chemical vapour transport.
With a commercial history predating that of beta-alumina by the better part of two centuries, calcium aluminates
are critical ingredients of Portland cement, as well as high-alumina cement (ciment fondu).8,40 Thus, tricalcium
aluminate, Ca3Al2O6, which has been shown to contain the cyclic [Al6O18]18 unit made up of six fused AlO4
tetrahedra (19),8 is an essential constituent of Portland cement, accounting for about 11% of the total by mass. The
presence of the [Al6O18]18 anions imparts to tricalcium aluminate a very open structure which facilitates the entry
of water, and, unlike the less negatively charged polysilicate anions, the polyaluminate anion reacts rapidly and
exothermally with water to form hydrates, a property that controls the setting and hardening of the cement. The main
phase of high-alumina cements, made by fusing limestone and bauxite with small amounts of silicon dioxide and
titanium dioxide, is the calcium aluminate CaAl2O4. Once again, setting and hardening depend on the hydration of
the aluminate and the very rapid hardening and high strength developed at an early stage have been linked to the
formation of a phase approximating in composition to the decahydrate CaAl2O4 10H2O. This has recently been
shown to have a quite unusual framework structure, with water molecules hydrogen bonded in nano-sized channels
(Figure 1.8).311 Under ambient conditions, the phase slowly converts into the stable Ca3Al2(OH)12 with a garnet
structure, and gibbsite, Al(OH)3, a process accompanied by substantial volume changes that may have catastrophic
consequences for the strength and porosity of the concrete. This potentially fatal flaw apart, high-alumina cements
have numerous highly desirable properties, for example resistance to corrosion by sea water and sulfate brines or by
weak mineral acids, and refractory properties.
Solid Compounds with Specific Electronic, Structural or Other Properties
41
Figure 1.8 The crystal structure of CaAl2O410D2O determined from neutron diffraction data. The upper part of
the figure shows the ab plane projection of one unit cell. The lower part of the figure shows a projection on the ac
plane of two unit cells along the c direction and one unit cell in the ab plane.311b Reproduced by permission of The
International Union of Crystallography (see colour version of this figure in Colour Plate section)
Other significant structures feature not only M–O chains or layers, but also cyclic groups, cages or extended 3D
networks containing voids or channels.39–45 Very often it is the cavities in these structures that are of especial interest,
there being great concern with the reproducible synthesis of structures having cavities, tunnels or pores of precisely
defined dimensions on the atomic scale.45 To the range of microporous and layered solids afforded by the familiar
zeolites41 and clays42 there have been added both aluminophosphates43 and gallophosphates,44a as well as related
materials, for example aluminoborates.44b Here mention should be made of the case of microporous aluminophosphates, which have developed hugely since they were first announced in 1982 in a seminal report by Wilson
et al.312 By the use of hydrothermal synthetic techniques, and with the mediation of organic amines, quaternary
ammonium cations or macrocycles bearing oxygen and/or nitrogen donor sites313a as templates or structuredirecting agents,313b it has been possible to prepare aluminophosphate and gallophosphate molecular sieves with
a wide range of pore sizes. With an appropriate choice of template it is possible even to produce porous chiral
materials, which hold the prospect of useful applications in the field of enantioselective separation of organic
materials. Both alumino- and gallo-phosphates have been grown in enantiomerically pure forms from resolved
templates.314 AlPO4 itself, in common with SiO2 with which it is isoelectronic, and closely related materials are
typically composed of neutral open frameworks. The majority possess a (4,2)-connected framework, that is with
primary building units of vertex-linked AlO4 and PO4 tetrahedra arranged in strict alternation. There is variety,
however, in that some structures feature five- or six-coordinated aluminium atoms with one or two extra-framework
oxygen substitutents (OH or OH2), as well as four-coordinated ones. This may lead, for example, to unique structures
with no zeolite counterpart. Such is the case with VPI-5, a molecular sieve exhibiting unusually large micropores
42 New Light on the Chemistry of the Group 13 Metals
Figure 1.9 (a) Framework topology of the microporous aluminophosphate VPI-5 (bridging oxygen atoms being
omitted). (b) Phase transformation of the framework from VPI-5 to AlPO4-8 (water molecules being omitted) that
occurs on heating. Reprinted with permission from [43c]. Copyright 2003 American Chemical Society (see colour
version of this figure in Colour Plate section)
generated from rings of 18 T atoms [T ¼Al or P] linked by oxygen bridges.315 The structural building blocks are
AlO4, AlO4(OH2)2 and PO4 units, producing the topology illustrated in Figure 1.9. More radical tailoring of the
topology and other properties of the framework can be achieved by partial replacement of the lattice aluminium and/
or phosphorus atoms with other atoms (e.g. of transition metals, silicon etc.). Objectives may include the
introduction of specific Brønsted acid sites and/or catalytically active metal centres.
Still greater diversity of structure arises in the many aluminophosphates that have anionic open frameworks.
These may extend in three dimensions, two dimensions (layers), one dimension (chains) or no dimensions (clusters).
The structures are formed by the alternation of aluminium-centred polyhedra (AlO4, AlO5 or AlO6) and phophoruscentred tetrahedra P(Ob)n(Ot)4n (n ¼ 1, 2, 3 or 4) featuring both bridging and terminal oxygen atoms (Ob and Ot,
respectively, with Ot representing either P¼O or P–OH units). Examples of the fascinating structural architectures
that can arise are provided by the compounds [(CH2)6N4H3]3þ [Al12P13O52]3 (AlPO-CJB1)316 and JDF-20,317
salient structural features of which are illustrated in Figures 1.10 and 1.11, respectively. AlPO-CJB1 is notable for
affording the first aluminophosphate molecular sieve to possess Brønsted acidity upon removal of the organic
templates by calcination, protons then remaining to balance the negatively charged framework. JDF-20 stands out
for the size of the pores permeating the framework with channel rings of 20 T atoms.
Coordination Chemistry of M(III) Compounds
43
Figure 1.10 (a) Anionic framework of the aluminophosphate [(CH2)6N4H3]3þ [Al12P13O52]3, AlPO-CJB1,
viewed along c (bridging oxygen atoms being omitted). (b) Side view showing the connections of the three types
of cage. Reprinted with permission from [43c]. Copyright 2003 American Chemical Society
Figure 1.11 Anionic framework of the aluminophosphate JDF-20 showing the large diameter channels that
permeate the structure. [43e] Reproduced by permission of The Royal Society of Chemistry
1.6 Coordination Chemistry of M(III) Compounds8,10,12,51,52,318
1.6.1 Acid-base properties
The most significant chemical property the Group 13 metal M brings to neutral MIII compounds is the innate
electrophilicity or acidity of the electropositive metal centre. This finds obvious expression in the more limited
44 New Light on the Chemistry of the Group 13 Metals
Brønsted sense, for example, in the properties of the hydrated cations [M(OH2)6]3þ , which all act in some degree as
proton sources in aqueous solution, as revealed by pK1 values of 4.99, 2.60, 3.9 and 0.6 for M ¼ Al, Ga, In and
Tl,12,52,53 respectively. But it is not just the M3þ cations that behave in this way. All compounds based on the þ3
oxidation state of the metal are typically Lewis acids, and the addition products they form with neutral or anionic
electron-rich species make for an unusually rich and variegated coordination chemistry, as will become still clearer
in Chapter 9. If the complexes are less amenable to ‘fine tuning’ of the metal centre properties by ligand design than
are comparable complexes of the d-block metals, they are the agents of most of the solution chemistry of the Group
13 metals, and have many important applications, for example in the extraction of the metals, in synthesis at large, in
catalysis, in the formation of III–V semiconductor materials and in medical diagnosis and therapy.
Dissolution of a Group 13 metal compound MX3 in a coordinating, monodentate solvent L may result in the
formation of one or more neutral complexes LnMX3 (n ¼ 1, 2 or 3); it may lead to the formation of X anions and
solvated cations of the general type [XnMLm](3n)þ (where n ¼ 1 or 2 and m ¼ 6 n or 4 n); or it may afford
complex cations and anions, as exemplified by the case of [Cl2Ga(NC5H5)2]þ [GaCl4].8,10,12,51–53 What the
outcome will be depends to some extent on M and X; more critically, it depends on the solvent L, with the added
complication that the products are prone to solvolysis, particularly when L is a protic solvent. Some neutral
complexes, usually with 1:1 stoichiometry, are sufficiently strongly bound that they survive in the gas phase, as is
the case, for example, with H3NMMe3 [M ¼ Ga or In]319 and Me3EGaH3 [E ¼ N292 or P199], molecules with real or
potential importance in the chemical vapour deposition of films of the relevant III–V compounds.46 Much of the
chemistry of Group 13 metals in the solid state is also concerned with compounds in which the coordination sphere of
some or all of the metal atoms has been expanded by coordination of neutral or anionic ligands. Examples of this have
already been met in the preceding section with the anionic frameworks displayed in certain aluminophosphates.
There is broad conformity of the observed and calculated dissociation energies of complexes of Group 13 metal
compounds with Pearson’s principle of ‘hard’ and ‘soft’ acids and bases.320 In this context, AlIII stands apart as
a typical hard metal centre that prefers to coordinate with oxygen donors, especially negatively charged ones, such as
carboxylates, phenolates, alkoxides and phosphates.321 Indeed, the aqueous chemistry of AlIII is dominated by its
oxophilicity. By contrast, the heavier MIII congeners are distinctly softer with the result, for example, that the
aqueous chemistry of GaIII and InIII is made much richer than that of AlIII by their greater affinity for softer donor
types, such as neutral nitrogen or anionic sulfur bases. This has an important bearing on the use of the radioactive
isotopes 67 Ga, 68 Ga and 111 In in nuclear medicine (Section 1.6.4). As indicated elsewhere,12,322 the change to softer
acceptor action is a consequence of the greater susceptibility of the heavier, more polarizable MIII centres to shortrange orbital perturbation, and hence to the electron transfer implicit in covalency.
Complexation of an MIII compound MX3 reduces the positive charge on, and at the same time limits access to, M.
Accordingly, it acts to curb the tendency of the parent compound to oligomerise or polymerise and, more generally,
to shield the MX3 fragment from other reactions that require associative activation. This can be of practical
importance in many areas. For example, multidentate oxygen donors, such as (poly)hydroxyl(poly)carboxylates
(including citric acid, tartaric acid, gluconic acid and sugar carboxylates), are capable of binding the AlIII ion so
strongly as to inhibit its hydrolysis under physiological conditions.323 In a rather different vein, binary aluminium
hydride (AlH3) is a relatively intractable, polymeric solid that is difficult to prepare in a pure state.8,12 By contrast,
pure complexes such as Me3NAlH3 and LiAlH4 are much easier of access, as well as being soluble in suitable
organic solvents.12,21 For this reason, it is the complexes rather than the parent compound that are normally favoured
as reducing agents or for the vapour transport of aluminium. Complex aluminium hydrides such as NaAlH4 and
Na3AlH6 are capable, with appropriate catalytic activation, of entering into reaction cycles that involve uptake
and release of dihydrogen.324 As a result, they have attracted keen attention in the quest for hydrogen storage
materials,325 even if recent studies of NaAlH4 cast doubt on the reversibility of the storage process.203 Gallium
hydride (GaH3) differs from its aluminium counterpart in being thermally unstable to decomposition to the elements
at ambient temperatures,105,108 whereas the neutral complex Me3N GaH3 is long lived under these conditions.292
Unsurprisingly, therefore, it is Me3NGaH3 or even more stable complexes, such as HC(CH2CH2)3NGaH3,155 that
are preferred for synthesis, for example selective reduction of organic compounds,326 or chemical vapour deposition.
The case of indium hydride is even more stark. Indium hydride (InH3) survives as a solid only at temperatures up to
Coordination Chemistry of M(III) Compounds
45
180 K,65 yet the carbene complex Mes[NCH¼CHN(Mes)C]InH3 (Mes ¼ mesityl) remains intact in the solid state at
temperatures up to 388 K.156
Where the ligand X in the compound MX3 has the ability to coordinate to the metal through more than one donor
site or to coordinate to more than one metal atom, there is the possibility of intramolecular coordination, as
exemplified by the diverse cases of H2Ga(m-H)2GaH2,8,12,21,105,108 solid MF3,8,12 and aluminium tris(quinolin-8olate), 20. The last is of especial interest in being much the most studied compound with regard to the fabrication of
organic light-emitting diodes (OLEDs).327 The unique electron transport and emissive properties of the compound
find application in so-called molecular OLEDs based on molecules of relatively low molecular weight with weak
intermolecular interactions in the solid phase. With the recent synthesis of heteroleptic aluminium(III) bis(2-methyl8-quinolinate)phenolate complexes such as 21,328 even the previous dearth of blue-emitting materials appears to
have been successfully countered. Intramolecular chelation of the sort found in 20 and 21 may be an important factor
in resisting solvolysis of the metal centre, and in stabilising metal–ligand bonds that are labile with respect to
associative attack. This last feature is exemplified by the neutral compound HIn(2-Me2NCH2C6H4)2 (22), one of the
first indium hydrides to be isolated.329
Me
H2
C
Me
N
H
In
Me
N
Me
C
H2
22
There is a marked diminution in Lewis acidity with the transition from MIII to MII and then to MI. Little is known
about the coordination chemistry of paramagnetic mononuclear MII compounds, which are short-lived under normal
conditions, but M--M-bonded MII centres are known to have the power to take up donor molecules or anions, as in
LX2GaGaX2L [X ¼ Cl or Br; L ¼ tetrahydrofuran, dioxane, diethylamine, triethylamine or pyridine],12,265
[X3MMX3]2 [M ¼ Ga or In; X ¼ Cl, Br or I],12 and [R2InInR2(CCPh)]330 and (R0 NC)R2InInR2(CNR0 )331
[R ¼ CH(SiMe3)2 and R0 ¼ tert-butyl or phenyl]. By contrast, the metal centres of MI compounds are weak, soft
Lewis acids whose coordination by conventional neutral or anionic donors has only rarely been demonstrated by
structural investigations. One example is afforded by the solid 1:1 In2Cl4 complex of dibenzo[18]crown-6, which
takes the form of a mixed valence compound Cl3InIII InI(dibenzo[18]crown-6)Cl (23).332 Another example is
provided by the anionic gallium(I) heterocycles [:Ga{N(R)CH}2], where R ¼ tert-butyl or C6H3-2,6-i Pr2 , for
46 New Light on the Chemistry of the Group 13 Metals
example 24; these prove to be very nucleophilic s-donors exhibiting parallels with the important N-heterocyclic
carbene (NHC) class of ligand.160,164,333,334 Intriguingly, though, MI centres prove to be susceptible to coordination
by arene molecules such as benzene, mesitylene, and [2,2]-paracyclophane, the stability of the resulting complexes
varying in the order GaI > InI > TlI.335 The structure of the benzene complex [InI(C6D6)][(Me3Si)3CInIIII3] (25)
illustrates the h6-mode of coordination of the C6 ring characteristic of this type of complex.336 Much more significant
than the acidic properties, though, are the basic properties of molecular MI compounds stemming ultimately from the
ns2 electron pair localised mainly on the metal (Section 1.4.5).
The much reduced acceptor power of MI and MII with respect to MIII compounds makes them susceptible to
disproportionation in the presence of bases, reactions such as those shown in Equations 1.20 and 1.21 being
driven by the superior strength of the coordinate links that can be established to the MIII centre.
3In þ ðsolnÞ > 2InðsÞ þ In3 þ ðsolnÞ
ð1:20Þ
3In42 þ ðsolnÞ > 2InðsÞ þ 4In3 þ ðsolnÞ
ð1:21Þ
In the following sections, outlined briefly are three areas in which the coordination chemistry of the Group 13
metals finds important expression, namely in catalysis, in ionic liquids and in aqueous media.
Catalysis
47
1.6.2 Catalysis
It is the Lewis acidity of the metal centres that accounts for the primary role of MIII compounds as catalysts,
cocatalysts or activators of both organic and inorganic reactions.337–339 One of the most familiar examples is
provided by the role of the trihalides, and most notably AlCl3, as catalysts for the Friedel–Crafts alkylation and
acylation of aromatic hydrocarbons. The Group 13 metal halide activates the alkyl or acyl halide via complexation
which, in the limit, may proceed by halide anion transfer to the metal, and in the process induces a partial or,
ultimately, full positive charge on the b-substituent (Equation 1.22). The electrophilic character developed by
the b-substituent is then played out in the subsequent attack of an aromatic substrate. Likewise, coordination
dþ
d
RCðOÞCl þ AlCl3 > RCðOÞCl AlCl3 > ½RCO þ ½AlCl4 ð1:22Þ
by the metal trihalide enhances the electrophilic character of the carbon atom of both organic and inorganic carbonyl
compounds, leading, for example, to the activation of transition metal carbonyls with respect to alkyl migration.340
While the same principles apply in these and other Lewis acid-mediated reactions, recent years have seen
a significant move away from traditional single-use, air- and moisture-sensitive catalysts such as AlCl3 to reusable,
more tolerant agents. In this vein, InCl3 is being successfully exploited as a water-tolerant catalyst, notably in
carbon–carbon bond-forming reactions.341
What holds for the metal trihalides holds more or less strongly for any other homoleptic or heteroleptic MIII
compound, in which the metal may be bound to an alkyl, aryl, alkoxide, aryloxide, methanesulfonate (triflate) or
other ligand. For example, AlIII compounds such as RAl(BHT)2 [R ¼ methyl or ethyl] containing the sterically
challenging ligand BHT ¼–OC6H2-2,6-t Bu-4-Me (26), have been shown to activate ketones to alkylation and
reduction.342 With respect to a rather different substrate, some idea of the effects of coordination by an aluminium
Lewis acid may be gained from the finding that the pKa of an alcohol rises by at least seven units.343 Still more subtly,
enantiomeric cyclic amido-aluminium compounds such as 27 render a variety of structurally diverse enolisable
aldehydes effective electrophiles for asymmetric cyclocondensation reactions involving an acyl halide to afford the
appropriate b-lactone with a high degree of enantioselection (Equation 1.23).344 Aluminium triflate, Al(OSO2CF3)3
CH 2Ph
Me
t
Bu
N
t
Bu
i
Al
Me
Me
O
O
i
Pr
Pr
N
F3CO2S
t
t
Bu
Bu
Al
N
R
R = Me or Cl
SO 2CF3
27
26
O
O
O
10 mol % 27
+
Br
Me
i
H
R
O
ð1:23Þ
Pr2EtN, CH2Cl2
R
48 New Light on the Chemistry of the Group 13 Metals
or Al(OTf)3, has also been shown to be an excellent Lewis acid catalyst for the alcoholysis345a and aminolysis345b of
epoxides, as well as being a highly active cocatalyst in the palladium-catalysed methoxycarbonylation reaction
(Equation 1.24).345c
CO2CH3
CO2CH3
Pd catalyst
R
Al(OTf) 3
MeOH, CO, heat
R
ð1:24Þ
+
R
Alkyls of the Group 13 metals, and especially those of aluminium, have evoked intense interest in the past 50 years
through their role as co-catalysts for olefin polymerisation.339 Aluminium alkyls of the type RnAlCl3n [R ¼ methyl,
ethyl etc.; n ¼ 1–3] are important components in Ziegler–Natta catalysis, whether in its classical heterogeneous form
or in homogeneous systems. The reaction between Cp2TiCl2 (Cp ¼ h5-C5H5) and R2AlCl or RAlCl2 involves first
exchange of R and Cl ligands and, subsequently, complexation, as in Equation 1.25, for example.
R Cl
Cp
Cp 2Ti(R)Cl
+
RAlCl 2
R
Al
Ti
Cp
Cl
Cl
ð1:25Þ
[Cp 2TiR] +[RAlCl3]−
The alkylaluminium halide thus serves the dual function of acting as an alkyl source and as a Lewis acid that
activates the Cp2Ti(R)Cl thus formed to produce either a surrogate or real [Cp2TiR]þ fragment, which is the
active centre for ethene insertion. Indeed, there is strong circumstantial evidence from studies involving the
intercession of the silylalkyne Me3SiC:CPh that the active component of the reaction between Cp2TiCl2 and
MeAlCl2 is the cationic species [Cp2TiMe]þ paired with the anion [AlCl4]. Such titanium- or zirconium-based
metallocene/alkylaluminium catalysts usually exhibit low-to-medium activity for the polymerisation of ethene,
but only ethene. Addition of water to the halogen-free, but polymerisation-inactive system Cp2ZrMe2/Me3Al
results, however, in the development of surprisingly high activity for ethene polymerisation, a finding associated
with the formation of a highly efficient activator, namely an oligomeric methylaluminoxane (MAO). This
discovery has served to rejuvenate Ziegler–Natta catalysis; along with major advances in controlling polymer
and stereochemistry architecture, it heralded the era of metallocene and single-site polymerisation catalysis.339
MAO, [–Al(Me)–O]n, typically with n in the range 5–20, affords in combination with Group 4 metallocenes
highly active catalysts for polymerising not only ethene, but also propene and higher a-olefins. It has thus
become a very important co-catalyst for metal-catalysed olefin polymerisation. Despite extensive research,
however, its exact composition and structure are still not entirely clear or well understood.346 Structures that
have been variously proposed are set out in Figure 1.12.339 About its overall function, though, there can be little
doubt. It acts in the same way as does MeAlCl2 with Cp2TiCl2, by accepting an anion, in effect or in ‘plain plump fact’,
and so creating a highly electrophilic, positively charged transition metal centre bearing the metal–carbon bond that is
to be the active site for olefin insertion.
Withdrawal of an anion from a substrate to form a complex Group 13 metal anion of low basicity and so create and
preserve a highly electrophilic cationic fragment is the guiding principle of many other interventions by Group 13
Lewis acids. For example, unusual cations of the type [(arene)2Hg]2þ [arene ¼ C6H5Me, C6H5Et, o-C6H4Me2 or
Catalysis
49
Figure 1.12 Possible structures for methylaluminoxane [–Al(Me)–O–]n, MAO. Reprinted with permission from
[339b]. Copyright 2000 American Chemical Society
C6H3-1,2,3-Me3] are formed in association with [MCl4] anions when mercury(II) chloride reacts with MCl3
[M ¼ Al or Ga] in an arene solution.347 Likewise, polycations such as [E8]2þ [E ¼ Sb or Bi],348a
[Te4]2þ ,348a ½SbHg2 nn þ 348b and ½Sb2 Hg3 nn þ 348b have been successfully isolated as salts of the [GaX4] anion
[X ¼ Cl or Br]. Of particular interest are complex anions with minimal basic, or coordinating properties, the parents of
which must be correspondingly powerful Lewis acids, or even so-called ‘super acids’. Such is the case with fluorinated
alkoxyaluminate anions, for example [Al{OC(CF3)3}4]349 and [(C6F5)3Al–imid–Al(C6F5)3], where ‘imid’ is an
imidazole fragment (28).350 Hence, for example, the powerful cationic Brønsted acids [H(thf)2]þ and
[H(OEt2)2]þ have been isolated and characterised as salts of the [Al{OC(CF3)3}4] anion,349 emerging as potentially
versatile reagents in organometallic synthesis and catalysis. The salt [HNMe(C18H37)2]þ [(C6F5)3Al–imid–Al
(C6F5)3], along with its boron analogue, is reported350 to be among the best known activators for olefin
polymerisation reactions.
(C6F5)3Al
Al(C6F5)3
N
−
N
28
Nor is there any change of general principle in catalytic action with the move from more or less simple molecular
systems to extended porous solids, such as zeolites,338 and several techniques have been devised to probe the surface
acidity and basicity of such materials at the gas–solid or liquid–solid interface.338,351 The efficiency of zeolites as
size-selective catalysts in the cracking reactions of hydrocarbons in petrochemical processes stems from their high
acidity, which is mainly due to the presence of Brønsted sites formed after calcination.352 Betwixt and between the
simple molecular and extended solid systems is found the world of synthetic supramolecular compounds, which
have been developed with well-defined internal environments, for example, to mimic enzyme catalysts.353 Rather
than using many different inner-sphere ligands to effect changes in reactivity, the outer-sphere environment of a selfassembled host cavity may be used to introduce new types of size and shape selectivity. An intriguing example of this
is provided by the supramolecular tetrahedral assembly [Ga4L6]12, where L ¼ bis(bidentate) catecholamide
(Figure 1.13), which can encapsulate organometallic cations such as [(h5-C5Me5)(Me3P)Ir(Me)(h2-C2H4)]þ , and
so provide a means of activating the C–H bonds of aldehyde reagents in a highly selective way.353
50 New Light on the Chemistry of the Group 13 Metals
Figure 1.13 (a) View down the twofold axis of the crystal structure of the tetrahedral supramolecular [M4L6]12
host [M ¼ Al or Ga; L ¼ bis(bidentate) catecholamide] which can encapsulate monocationic guest molecules.
Ligands are coloured for differentiation. (b) Schematic representation showing the structure of one of the six identical
ligands that span the edges of the tetrahedral host.353 Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced
with permission (see colour version of this figure in Colour Plate section)
1.6.3 Ionic liquids
The ability to take up a halide anion and form a strongly bound complex anion of low basicity finds further
expression in a variety of ionic liquids featuring a Group 13 metal trihalide or other MIII compound. A familiar
and commercially vital example of an ionic liquid is provided by ‘cryolite’, Na3AlF6, which melts at
temperatures above 900 C to produce a liquid composed of Naþ and [AlF6]3 ions, with the latter suffering
partial dissociation to [AlF5]2 and [AlF4] ions.12,354 The addition to the melt of not only 2–8% alumina, but
also other ingredients, for example AlF3 and CaF2, then affords the normal electrolyte for the electrodeposition
of aluminium metal.
Over the past two decades, however, the term ‘ionic liquid’ has been largely usurped by salts composed solely
of ions with melting points not near 1000 C, like that of cryolite, but below 100 C.12,355,356 Ideally, such
a liquid is colourless and involatile, has a low viscosity and long liquid range, and is easily handled in air. In
practice, a good approach to some or all of these ideals is often achieved with a salt of an organic cation, for
example tetraalkyl-ammonium or -phosphonium, N-alkylpyridinium, 1,3-dialkylimidazolium or trialkylsulfonium, and a polyatomic anion, for example BF4, AlCl4, Al2Cl7, PF6 or CF3SO3. In order to be liquid at
room temperature, the cation should preferably be unsymmetrical, with different alkyl substituents, for example,
attached to the central atom or imidazolium ring, but the melting point also depends on the anion. In this and
other respects, it is evident that the properties of the liquid can be tuned. Ionic liquids of this sort turn out to be
good solvents for a wide range of both inorganic and organic materials. Since they are typically composed of
weakly coordinating ions, reaction products are seldom isolated as solvates; they are also immiscible with
a number of organic solvents and so provide a non-aqueous, polar alternative for two-phase systems (some are
even hydrophobic and can also be used as immiscible polar phases with water). Of particular importance is their
involatility, leading to the elimination of many of the containment and pollution problems that beset the use of
conventional organic solvents. Accordingly, their development and exploitation offer one of the most promising
strategies in the drive for cleaner, greener technologies. Their applications, established or projected, reach from
organic and inorganic synthesis and catalysis, over separation processes, to electrochemical operations. The
point has certainly not been lost on either industry or academia, as more than 8000 papers (including over 900
patents or applications) bear witness in the past decade. Only the more cynical might remark: ‘Whom the gods
wish to destroy they first call “promising”.’
Ionic Liquids
51
Chloroaluminate-based ionic liquids were among the first to be taken seriously,355,356 being proposed as early
as 1948 as possible bath solutions for electroplating aluminium. The preparation of chloroaluminates that are
liquid at ambient temperatures was delayed, however, until the late 1970s, interest then focussing mainly on
possible electrochemical applications. It was not until the 1980s that an appreciation of the potential of
chloroaluminate melts as non-aqueous, polar reaction media for transition metal complexes and for organic
synthesis and catalysis started to dawn. Fundamental studies carried out on ionic liquids, such as that formed
from AlCl3 and 1-ethyl-3-methylimidazolium chloride, ImCl, reveal Lewis acid–base chemistry governed by
the equilibrium in Equation 1.26.357 The melts are termed acidic or basic if the mole ratio of AlCl3 to the organic
chloride is greater than or less than 1:1.
17
*
2AlCl
4 ðlÞ ) Cl ðlÞ þ Al2 Cl7 ðlÞ K ¼ 10
ð1:26Þ
In order to produce a buffered neutral melt, such as might be useful as a battery electrolyte, an alkali metal chloride
can be added to an acidic melt. By contrast, HCl behaves as a Brønsted super acid when dissolved in an acidic melt
and will protonate a variety of arenes. This acidity is associated with reactions of the general form shown in
Equation 1.27, where X ¼ Cl, HCl2, H2Cl3, AlCl4 or Al2Cl7. Acidic chloroaluminate melts prove to be
HClðgÞ þ X ðlÞ > XHCl ðlÞ
ð1:27Þ
effective media for Friedel–Crafts reactions of not only arenes, but also ferrocene,358a and for stereoselective
hydrogenation of arenes;358b both types of reaction can be performed with excellent yields and selectivities. The use
of the melts as solvents for transition metal catalysts of olefin polymerisation has also been described.356,359 That
they can be good solvents for the synthesis of inorganic materials has also been demonstrated, as exemplified by the
excision, isolation and electrochemical studies of the centred hexanuclear zirconium halide clusters [(Zr6ZCl12)
Cl6]n [Z ¼ Be, B, C, Mn or Fe].360
For all these virtues, chloroaluminate melts, representing the first generation of ionic liquids, have their
limitations. Most notably, they are moisture-sensitive and prone to hydroysis. Through the 1990s, therefore, the
emphasis shifted to a second generation of ionic liquids that are more benign in regard to their stability to hydrolysis.
The cations have remained largely the same, but the anions are more likely to be BF4, PF6, CF3SO3,
[N(SO2CF3)2] or tosylate. Unlike the chloroaluminates, such systems tend to offer greater tolerance to functional
groups, a feature which extends greatly the range of applications, especially in organic synthesis,356a catalysis,356b,d
biotransformations,356d gas storage,361 inorganic synthesis362a and electrodeposition (including that of aluminium).362b,c Nevertheless, the Group 13 metals still feature as either constituents or beneficiaries. Thus, ionic liquids
with the weakly coordinating perfluoroalkyloxyaluminate anions [Al(ORf)4] [Rf ¼ CH(CF3)2 or C(CF3)3] are not
only more hydrolytically robust than chloroaluminate melts, but also show surprisingly low viscosities at elevated
temperatures (60 C);363 at the same time, they offer the potentially attractive advantage of an exceptionally wide
electrochemical window. Chloroindate ionic liquids offer another alternative, also having superior hydrolytic
stability and reduced oxophilicity compared with chloroaluminate melts. They prove to be versatile reaction media
for Friedel–Crafts acylation reactions;364 importantly, the system is said to be catalytic and totally recyclable using
an aqueous workup, with no leaching of indium into the product phase.
Ionic liquids, in general, are good microwave absorbers, a property that has been turned neatly to advantage for
the synthesis of zeolites.365 Many ionic liquids, and especially those based on imidazolium and quaternary
ammonium salts, are very similar chemically as regards the type of organic cation that is commonly used as
a template or structure-directing agent in the preparation of zeolites by the conventional hydrothermal route.
Replacing the solvent and the organic template with a single ionic liquid is the basis of the so-called ‘ionothermal’
method of zeolite synthesis, which has been developed in recent years.365 With microwave activation, this synthesis
can now be accomplished on the bench top in open vessels, rather than in the sealed, teflon-lined autoclaves
normally required.
52 New Light on the Chemistry of the Group 13 Metals
1.6.4 Aqueous media: environmental and medicinal considerations
It is the response to aqueous conditions that shapes the environmental and biological chemistry of the Group 13
metals and how, in particular, they interact with living systems.12,51,52,366 Only in distinctly acidic solutions free from
neutral or anionic bases of significant coordinating ability are the Group 13 metals found as the simple mononuclear,
hexahydrated cations [M(OH2)6]3þ , where M ¼ Al, Ga, In or Tl. Studies involving, variously, X-ray and neutron
diffraction and extended X-ray absorption fine structure (EXAFS) and molecular dynamics simulation of the
solutions indicate that each of the six water molecules in the first coordination shell of the cation formed by
aluminium, gallium and indium is hydrogen-bonded to two water molecules in the second hydration sphere.367
Trading of water molecules between the first coordination shell and the solvent is relatively slow for Al3þ but
becomes progressively faster for the heavier cations, the first-order rate constants estimated by NMR measurements
being about 1, 4 102 and 4 104 s1 for Al3þ , Ga3þ and In3þ , respectively.368 With ligands that bind to Al3þ more
strongly than water, and particularly multidentate ligands such as citrate, the rate of exchange is even slower; this
sluggishness appears to be a factor critical to the general rejection of Al3þ , and other multiply charged metal ions, by
biological systems.366
In the absence of other ligands with significant coordinating ability, an increase of pH causes the M(III) aquo
ions to undergo extensive hydrolysis and polymerisation, leading to an array of hydroxo and oxo solute species
and solids. Most attention has been paid to aluminium, for which hydrolysis sets in above pH 3,12,51,52,369,370 rather
less to gallium12,371 and comparatively little to indium and thallium.12 Taking first the case of aluminium, it is noted
that, in contrast to the sequentially formed mononuclear hydrolysis products [Al(OH)n(OH2)6n](3n)þ (n ¼ 1–6),
polynuclear species are formed in a manner that is at once less facile and highly concentration-dependent.
Soluble polynuclear species include dimers [Al2(OH)2(OH2)4]4þ , trimers [Al3(OH)4(OH2)9]5þ , tridecamers
[Al13O4(OH)24(OH2)12]7þ (Al13-mers) and the recently characterised [Al30O8(OH)56(OH2)26)]18þ cation
(Al30-mers).372
Knowledge of the speciation in aqueous solutions and the dynamics of interconversion at medium to high
aluminium concentrations is of major consequence in relation to materials science, catalysis, geochemistry,369c soil
science and many other fields. For example, the polyoxocations are present in certain catalysts, clay-pillaring
agents and additives used for water treatment.369 Most familiar of the larger polycations are the Al13-mers and their
derivatives. These exist as the Baker–Figgis–Keggin isomers illustrated in Figure 1.14, being characterised by
a central MO4 tetrahedron, in which M may be not only aluminium but other Main Group atoms, for example
gallium or germanium.
Largest of the polyoxocations of aluminium yet to be identified is the remarkable Al30 species, Al30-mers. As
illustrated in Figure 1.15, the structure can be understood as two d-Al13 molecules that face each other at the rotated
trimers and are bonded via a belt of additional AlO6 linkages.369,372 The resulting unit is about 2 nm in length, and
exposes to the aqueous solution oxygen atoms in many different coordination environments. Heteroatom variants of
this cluster, for example [W2Al28O18(OH)48(OH2)24]12þ , have also been characterised.369 All such clusters feature
a central MO4 unit readily identifiable by its sharp 27 Al NMR signal. In a second class of cluster is found a
characteristic core of edge-shared AlO6 octahedra organised into cubane-like moieties that are linked together in
a structure similar to that of brucite [Mg(OH)2]. These molecules are most commonly formed in conjunction with an
aminocarboxylate ligand that reduces the overall positive charge, but also occur in purely inorganic solutions. An
example is the so-called ‘flat-Al13’ species, [Al13(m3-OH)6(m2-OH)18(OH2)6]15þ with the structure illustrated in
Figure 1.16.369,370c Despite the difficulties generally experienced with trying to prepare stable single crystal forms
of polyoxocations, the gallium analogue of this flat-Al13 cation has recently been characterised, in addition to
a notable Ga32 cluster [Ga32(m4-O)12(m3-O)8(m2-O)7(m2-OH)39(OH2)20]3þ isolated as a supramolecular compound
of the macrocyclic cavitand cucurbit[6]uril.371b
Of the four Group 13 metals the three heaviest, gallium, indium and thalium, occur naturally only in trace amounts
(Table 1.1) and are not of general environmental concern.51,52,366,373 The case of aluminium is decidedly different,
for it is the most abundant metal in the earth’s crust. Despite this abundance, low solubility dictates that little of the
metal finds its way into the planet’s waters. The ocean concentrations of about 2 ppb are even lower than that
Environmental and Medicinal Considerations
53
Figure 1.14 The Baker–Figgis–Keggin isomers for the tridecamer [Al13O4(OH)24(OH2)12]7þ , ‘Al13-mers’, shown
in polyhedral representation. The relationship between the isomers can be understood as the stepwise rotation of
trimeric groups of AlO6 octahedra that share corners (light grey) about the m4-O centre so that they share edges with
one another (darker grey). Reprinted with permission from [369b]. Copyright 2006 American Chemical Society
allowed by clay sediments, possibly because of harvesting of aluminium, as well as silicon, by diatoms.373 Most
natural waters contain insignificant amounts of aluminium, except for those in some volcanic regions or emanating
as alum springs. Nevertheless, biological systems could take up the metal; its concentration may be low (1011 M at
pH ¼ 7.0), but it is still six orders of magnitude greater than that of iron (1017 at pH ¼ 7.0), which all biological cells
contain. With increasing acidity, or through the agency of complexing ligands, such as fluoride, citrate374 or organic
phosphates,375 aluminium becomes even more freely available. Yet its role in biological systems is to this day
a subject of uncertainty and debate. In numerous clinical and biological studies over recent years, aluminium’s
accumulation and biotoxicity have been variously linked with (a) its detection in neuritic brain deposits of patients
with Alzheimer’s disease, (b) other human pathological conditions, such as renal dialysis-related encephalopathy,
osteomalacia and endemic amyotrophic lateral sclerosis, and (c) toxic processes arising from the increased
acidification of the environment.373,376 Whether the aluminium exerts a specific or a general effect, and whether
its accumulation is causative or consequential, however, remain generally open to doubt.
It is an intriguing fact of life that the Mg2þ ion is similar in size to, but a much weaker Lewis acid catalyst than,
the Al3þ ion, yet it, and not Al3þ , is essential to living systems.366 It seems likely that the stronger binding of AlIII to
the phosphate groups of nucleotide di- and tri-phosphates, NDA and NTA, and the slowness of association and
dissociation prevent it from being active in kinases. More seriously, though, entry of AlIII would tend, by virtue of
these very properties, to block the access of essential ions such as Mg2þ . In view of its affinity for phosphates, AlIII
needs therefore to be excluded, and how this may be achieved, whether by natural or artificial means, through
appropriate sequestration, for example by citrate, has been the focus of intensive study.366 There is much less doubt,
though, that aluminium, once rendered more soluble by acid conditions, is toxic to many plants because of its effects
on the surface of roots, where AlIII antagonises the CaII binding that is essential to the surface. It has been estimated
that soil acidification, accelerated by certain farming practices and by acid rain, affects about 40% of the arable
land worldwide. The development of aluminium-tolerant plant varieties has, therefore, become imperative. A good
54 New Light on the Chemistry of the Group 13 Metals
Figure 1.15 The cation [Al2O8Al28(OH)56(OH2)26]18þ (aq), ‘Al30’, shown (a) in polyhedral representation, (b) as
a ball-and-stick model and (c) in a polyhedral exploded view. Hydrogen atoms are omitted from the structure in the
interests of clarity. Reprinted with permission from [369b]. Copyright 2006 American Chemical Society (see colour
version of this figure in Colour Plate section)
Environmental and Medicinal Considerations
55
Figure 1.16 The cation [Al13(m3-OH)6(m2-OH)18(OH2)6]15þ (aq), ‘flat-Al13’, is based (a) on a repeated set of
cubane-like moieties arrayed in a lattice of edge-shared Al(OH)6 octahedra, and is shown (b) in polyhedral
representation. Reprinted with permission from [369b]. Copyright 2006 American Chemical Society
example is provided by experiments in Mexico with tobacco and papaya plants that have been genetically engineered
to overproduce citrate as the means of sequestrating the aluminium.377
Thallium, standing necessarily apart as being most likely to occur under aqueous conditions as TlI, is well known
for its toxic properties;12 the aqueous chemistry of gallium and indium has much in common with that of
aluminium.12,366 The difference lies mainly in the weaker binding and faster exchange of ligands. Neither gallium
nor indium has a metabolic role. Importantly, though, they do possess radioisotopes that are signally useful in
medicine for diagnosis or therapy (both radiotherapy and chemotherapy).12,366,378 Of these isotopes 67 Ga, 68 Ga and
111
In are the most significant, with the following decay modes and half-lives: 67 Ga EC/g, 78.1 h; 68 Ga EC/bþ ,g,
1.13 h; 111 In EC/g, 67.4 h. 67 Ga and 111 In have been widely used for many years, but 68 Ga has gained prominence
lately for Positron Emission Tomography (PET).379 67 Ga finds application in the related g-ray imaging technique of
Single Photon Emission Computed Tomography (SPECT). These imaging techniques have been successfully
employed to diagnose cancer, infection, thrombosis and kidney and liver abnormalities, as well as many other
cardiac and neurological disorders.378
In fact, the pre-eminence of the isotopes in nuclear medicine has been a primary driving force in the recent
development of the coordination chemistry of GaIII and InIII. Complexation of the ions is needed for their
administration in vivo.366,378,380 Major requirements of complexation are then thermodynamic and/or kinetic
stability at physiological pH, with metal–ligand binding robust enough to resist hydrolysis of the metal ion and
competition from endogenous molecules that can behave as ligands, notably the plasma protein transferrin.381
56 New Light on the Chemistry of the Group 13 Metals
Moreover, the presence in the ligand of a second functional group that is not involved in complexation of the metal
allows covalent linkage to targeting biomolecules, for example antibodies, small proteins or synthetic macromolecules. A good example is provided by the cell membrane folate receptor commonly overexpressed in malignant
cells, and which may be targeted with the 111 In(III) complex of the ligand 29 in Figure 1.17.382
(a)
R
(b)
R
N
CO2H
HO2C
N
N
N
N
OH HO
R
Me
R=
-CO2-
or
O
Me
S
(d)
(c)
OH
NH
NH
HN
O
O
HO
R = Me, X = O, Htma
R = Et, X = O, Hetma
R = Me, X = NH, Hmppt
R = Me, X = NMe, Hdppt
OH
R
X
HO
O
(e)
O
O
OH
N
N
H2N
N
H
N
O
H
N
N
CO2H
O
N
H
N
N
HO2C
N
29
CO2H
HO2C
(f)
−
O2 P
H+
N
N
+
HN Bz
Bz
+ −
HN O
2P
Bz
−
O2P
(g)
R1
N
H
N
OH
R2
R1 = H, R2 = OMe
R1 = NO2, R2 = H
R1 = R2 = Cl, Br or I
Bz = benzyl
30
Figure 1.17 Representative ligands or ligand parents which have been proposed for the delivery of radioactive
gallium or indium isotopes in medical diagnosis or therapy
Mediation of Organic Transformations by Group 13 Metal Compounds
57
Figure 1.17 also displays other representatives of the wide variety of multidentate ligands which have been
variously ventured as real or potential vehicles for the administration of these isotopes. Some of the complexes hold
major promise in chemotherapy. For example, certain GaIII complexes of the type [Ga(LX)2]ClO4, where [LX] is the
deprotonated form of the asymmetric ligand 30, show activity superior to that of cisplatin in the treatment of the
neuroblastma that accounts for more than 10% of childhood cancer.383 In quite a different area of medicine that
does not involve radioisotopes, GaIII complexes of suitably tailored amine phenol ligands are claimed to be
potent antimalarials, able even to cope with chloroquine-resistant organisms.384 At a fundamental level, the
similarity of size and binding properties has enabled GaIII to be used as a proxy for FeIII, the labile, paramagnetic FeIII
centre being replaced by a kinetically inert, diamagnetic GaIII one more amenable to NMR and other studies. Hence,
for example, it has been possible to shed some light on the mechanism for siderophore-mediated iron uptake in
microorganisms.385
1.7 Mediation of Organic Transformations by Group 13 Metal Compounds
As Chapter 11, together with earlier accounts,386–391 makes clear, Group 13 metal compounds have a long
history as reagents, catalysts and activators of organic transformations. In common with boron, aluminium386,387
and thallium386,388 are well established in this respect, their compounds finding extensive use in organic
synthesis. Gallium and indium were less familiar in this role until the 1980s, but the last two decades have seen
a marked rise in profile as advantages of handling, reactivity and selectivity of gallium386,389,390 and
indium341,386,389,391 have gained recognition. Reference has already been made in Section 1.6 to some aspects
of catalysis and the application of chlorometalate ionic liquids in the general context of the coordination
chemistry of the metals.
The catalytic function of Group 13 metal compounds depends primarily on the Lewis acidity of the metal
centre and its ability to induce an electrophilic site in an appropriate organic substrate. In practice, the active
agent or intermediate in an organic reaction is most likely to be an organo- derivative of the metal. Such
a compound is bifunctional, being at once a Lewis acid by virtue of the metal, and a nucleophile by virtue of the
more or less carbanionic organic substituent. In their intrinsic nucleophilicity, organo- derivatives of the Group 13
metals fall generally some way short of organolithium or Grignard reagents; in Group 13 at large, this
nucleophilicity varies roughly in the order B < Al > Ga In > Tl. Similar considerations apply to hydrides of
the metals, for example i Bu2 AlH or HGaCl2, which act as hydride sources, and hence reducing agents, less
powerful but more tractable than hydride derivatives of the alkali or alkaline earth metals. It is the nucleophilicity
of the substituents that is uppermost when the primary role of the Group 13 metal compound is that of a reagent
rather than a catalyst or activator, although the two functions are commonly intertwined. The nucleophilicity of
the organic or hydride substituent is significantly modified by complexation of the compound, being particularly
enhanced when the base is anionic, as in Li[AlH4], Li[Me3AlR] [R ¼ butyl or phenyl] or M0 [InR4] [M0 ¼ Li,
MgCl or MgBr; R ¼ Me, i Pr, i Bu, s Bu, Ph or C:CPh].392 The ability to tailor the reactivity of organo and
hydrido compounds of the Group 13 metals aluminium, gallium and indium is open to further refinement through
the choice of solvent and/or supplementary catalyst. Such versatility is a primary attraction in the mediation of
organic reactions.
One of the main applications of organoaluminium compounds is the engineering of carbon–carbon bond
formation. Thus, they feature in a variety of reactions of this sort involving mainly, but not exclusively, unsaturated
organic moieties, and leading, for example, to alkylation, arylation, carbometalation, cyanation, addition to carbonyl
compounds, conjugate addition to enones, and epoxide ring-opening reactions.386 Where metalation occurs to give
Al–C bonds, these undergo facile protonation, oxidation or halogenation to produce C–H, C–OH or C–halogen
bonds, respectively. Recent research has tended to be preoccupied with the achievement of a high degree of not just
selectivity, but of enantioselectivity. Two examples serve to illustrate recent trends. The first, represented by
58 New Light on the Chemistry of the Group 13 Metals
Equation 1.28, involves enantioselective conjugate addition of a trialkylaluminium reagent R3Al [R ¼ Me, Et or
vinyl] to an enone in the presence of a catalytic amount of a copper(I) salt and a phosphoramidite ligand
O
O
+
R3Al
CuX, L*
ð1:28Þ
−30 oC, 18 h
R
R′
R'
O
R1
HO
i
+
Ar 3Al THF
Ti(O Pr)4/(S)-binol
Ar
ð1:29Þ
toluene, 0 oC
R1
R2
R2
(L ).393 The second, illustrated by Equation 1.29, involves aryl addition of Ar3Althf [Ar ¼ phenyl, 2-naphthyl or
4-RC6H4, where R ¼ Me, SiMe3 or OMe] to a ketone in a reaction catalysed by a titanium(IV)-(S)-binol complex.394
Hydridoalanes, often in the form of adducts with ethers or amines or as anionic complexes (particularly Li[AlH4] and
Na[(MeOC2H4O)2AlH2]) are well known as powerful and more or less selective reducing agents.12,21,386,395 Of the
various groups open to reduction, carbonyl functions have been the principal targets, but other potential reaction
centres include C–halogen, C¼C, C:C, epoxide, –CN, –NO2 and sulfur-containing moieties.
What organo- and hydrido-gallium and -indium compounds lack in nucleophilicity and availability, compared
with their aluminium counterparts, may well be more than matched by greater selectivity, as well as ease and
cleanliness of manipulation.341,386,389–391 In connection with this last point, the active agents are typically formed in
situ and the necessary starting materials are often simple inorganic compounds, such as the trihalides, or even the
metal itself. Importantly, too, the reduced oxophilicity and sensitivity to hydrolysis makes gallium and indium
reagents much more water-tolerant than are analogous aluminium reagents.
Indium mediation is of particular note in its capacity to effect a variety of reactions in aqueous, rather than
organic, media, with the economic and environmental benefits this implies.396 For example, the Barbier reaction
of an allylic halide with a carbonyl compound to produce a homoallylic alcohol can be brought about efficiently
by the action of indium metal in water (Equation 1.30). Selective C–C bond formation at a carbonyl function is
also evident (Equation 1.31) in the Reformatsky reaction between an aldehyde or ketone and an a-haloester,
which can be efficiently engineered through either gallium- or indium-mediation.389 The first ionisation potential
of the indium atom (558 kJ mol1) is not only the lowest in Group 13, but even comes quite close to that of an
alkali metal (for example Na 496 kJ mol1), while being much lower than that of metals such as zinc
(906 kJ mol1) or tin (709 kJ mol1), which are commonly used as reducing agents. Combined with its stability
to air and water, this makes indium metal an attractive reducing agent for organic substrates, and it has indeed
found use, for example, in the reduction of (i) C¼N bonds in imines, (ii) the heterocyclic ring in benzo-fused
nitrogen heterocycles, (iii) oximes, (iv) nitro compounds, (v) conjugated alkenes and (vi) non-enolisable bicyclic
OH
O
+
R
H
Br
In
H 2 O, 80-97%
ð1:30Þ
R
Mediation of Organic Transformations by Group 13 Metal Compounds
CO 2 R
CO 2 R
O
+
59
ð1:31Þ
X
OH
a-diketones.397 Another role it has found is as a stereoselective dehalogenating agent, for example, of
aryl-substituted vic-dibromides.398 Hydrido derivatives of both gallium and indium,21,399 for example Me3NGaH3,
LInH3 [L ¼ P(cyclo-C6H11)3 or :CN(Mes)CH¼CHN(Mes), where Mes ¼ mesityl], LiMH4 [M ¼ Ga or In], HMX2
[M ¼ Ga or In; X ¼ Cl or Br], are typically milder, more selective reducing agents than corresponding alanes. For
example, Me3NGaH3 selectively reduces the carbonyl group in 4-BrC6H4COBr, whereas quinuclidineAlH3
reduces not only this group but also the adjacent C–Br bond.400
HInCl2, conveniently formed in situ from InCl3 and a hydride source, such as NaBH4, Bu3SnH, or Et3SiH, has
found considerable favour as a reducing agent. For example, it selectively reduces C¼C bonds in conjugated alkenes,
as in a,a-dicyano olefins, a,b-unsaturated nitriles, cyano esters, cyanophosphonates, diesters and ketones.
Importantly, it can work as an effective radical mediator able, for instance, to bring about highly regioselective
radical addition to alkynes, to reduce alkyl halides and to promote C–C bond formation.389 HGaCl2, formed from
GaCl3 and Na[(MeOC2H4O)2AlH2], can act in a similar manner, not only reducing alkyl halides, but also effecting
radical cyclisation of halo acetals (Equation 1.32).401 These monohydrido derivatives of gallium and indium have
been advocated as alternatives in synthetic radical chemistry to the widely used organotin hydrides, which are
usually toxic and difficult to remove completely from the desired reaction products.
RO
O
RO
R′
HGaCl2, Et3B
THF, 0 oC
O
R′
ð1:32Þ
Br
As gallium’s and particularly indium’s stars have waxed with recent synthetic advances, so thallium’s star has
waned, reflecting no doubt the increasing demands now being made in the causes of cleanliness and safety.
Nonetheless, thallium continues to play a significant and distinctive role. It differs from the other three metals in the
oxidising potential of organothallium(III) derivatives and the lack of any hydrido derivatives able to survive under
normal conditions. Although organothallium compounds, R3Tl, can emulate the corresponding compounds of the
lighter metals by acting as a source of R,388 the dominant characteristic of organic transformations mediated by
thallium is the highly favourable reduction of TlIII to TlI. Electrophilic thallation of an arene, ArH, by a thallium(III)
salt, TlX3, is highly regiospecific, affording a monoorganothallium(III) derivative ArTlX2, where X is commonly
trifluoroacetate, CF3CO2. This derivative is then a highly efficient precursor for the synthesis of a variety of
functionalised arenes, for example ArX [X ¼ halogen], ArCN, ArSH, ArSCN, ArNO2, ArAr and so on, with
simultaneous reduction of TlIII to TlI.386,388 Similar activation of certain other unsaturated organic compounds can
60 New Light on the Chemistry of the Group 13 Metals
also be effected by thallium(III) salts en route to overall reactions involving, variously, oxidation, ring-contraction
or ring-expansion and heterocyclisation.386
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