Chemistry 3810 Lecture Notes
3.
Dr. R. T. Boeré
Page 42
Hydrogen and a Survey of the Molecular Hydrides
In Chemistry 2810, we introduced the chemistry of the unique element, hydrogen. Please review those notes, which
correlated roughly to sections 8.1 to 8.3 of Shriver-Atkins, 3rd edition. Our coverage corresponds to all the remaining chapter
sections of the text.
3.1
Classification of Hydrogen Compounds
There are three main types of binary hydrogen compounds, organized by the relative electronegativities of hydrogen and
the other element. The known binary EH n compounds are organized on the following periodic chart:
3.3.1 Saline or Ionic Hydrides
The saline, or Ionic, hydrides occur when hydrogen is combined with the very electro-positive elements. They are
compounds of Mn+ and n H– , and are when very pure colourless ionic compounds with structures similar to those of the ionic
halides. However, the materials that are commercially available are always found to be contaminated by colloidal-size
particles of the metallic element, and typically appear from light to very dark gray.
They exist in the following crystalline forms:
Compound
Crystal Structure
LiH, NaH, KH, RbH, CsH
Rock salt (NaCl)
MgH2
Rutile
CaH2 , SrH2 , BaH2
Distorted PbCl2
Source: A.F. Wells, Structural Inorganic Chemistry, Oxford University Press (1984)
Within these crystal lattices, the radius of H– varies from a 1.26 Å in LiH to 1.54 Å in CsH, a variability that reflects the very
loose hold that the single proton has on the two-electron charge cloud on these ions. Not surprisingly, the dominant reaction
of these hydrides are ones that deliver electrons, i.e. reductions.
The most important from a practical standpoint are LiH, NaH and CaH2 . All of these are extremely reactive towards
water, and calcium hydride granules are routinely used in the laboratory as a powerful drying agent for aprotic, nonhalogenated solvents. The water-scavenging reaction is:
CaH2
+ 2 H 2O → Ca(OH ) 2
+ 2 H2
Oil-dispersions of finely divided NaH are used as powerful reducing agents. The saline hydrides are also used as the source
for the synthesis of the more easily handled hydrides of other elements, such as LiAlH4 or the commercial “Superhydride”:
LiH
+ B(C2 H 5 ) 3
→ Li [ HB(C2 H 5 ) 3 ]
Which is sold as a solution in tetrahydrofuran, and is an very powerful, quantifiable reducing agent in both organic and
inorganic chemistry. Each equivalent of superhydride delivers one H– or one electron to a substrate.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 43
3.3.2 Metallic or Interstitial Hydrides
The more electronegative metals do not form saline hydrides. Instead they form a range of, typically, non-stoichiometric
hydrogen derivatives that are of great theoretical and practical interest, but of little utility in chemistry in general. The know
hydrides and their approximate stoichiometries are shown in the following table:
They find applications in hydrogen storage, and as membranes in
hydrogen gas purification devices, such as the one shown at right, in
which a palladium-silver alloy tube is used which is saturated with
hydrogen, and through which hydrogen but not impuritiy gases can
permeate. The purified hydrogen, at a lower pressure, passes through
the walls of the tube and the impurities are expelled.
We will not say much more about this topic, which is more
closely tied to the chemistry of the transition metals. It is very
important to recognize, however, that many transition metals are used
as catalysts for reactions involving hydrogen. Thus the standard
hydrogen electrode uses colloidal platinum on a platinum surface to
bring the equilibrium between hydronium oi n in acid solution and
hydrogen gas into a reasonably kinetic range. Palladium dispersed on
charcoal is a very common hydrogenation catalysts in organic
synthesis, and porous forms of nickel (e.g. Raney nickel) are used to
hydrogenate fats and oils in the commercial production of margarine, soaps and cosmetics.
3.3.3 Molecular Hydrogen Compounds
By far the most important hydrogen compounds are the molecular compounds, and they will be the focus of our
discussion in this chapter. Molecular hydrogen compounds form for most of the p-block elements. Unlike the halogen
derivatives of the elements, there is only one oxidation state of each of these elements that forms a stable hydrogen
compound. However, there are a few unstable species. Thus carbon, which has CH4 as its stable hydrogen compound, also
has the unstable form CH2 , the carbene molecule. Compared to carbon, for which a vast array of hydrocarbon derivatives
exist, based on chains and rings of the element with hydrogen satisfying the remaining valences, the remaining elements have
far less diversity. Thus nitrogen has ammonia, NH3 and hydrazine, H2 NNH2 . Oxygen has water and hydrogen peroxide.
The heavier group 14 elements do have a few recognized extended structures, but nothing even close to the diversity of
carbon.
We can classify the molecular compounds of hydrogen into three classes, as shown in the following table that also gives
the names of these compounds.
Chemistry 3810 Lecture Notes
Class
Electron-Deficient
Electron-Precise
Electron-Rich
Dr. R. T. Boeré
Group
13
13
13
14
14
14
14
15
15
15
15
16
16
16
16
17
17
17
17
Formula
B2 H6
AlH3 (poylymeric)
Ga 2 H6 (< –30°C)
CH4 and hydrocarbons
SiH4 and silanes
GeH4 and germanes
SnH4 and stananes
NH3
PH3
AsH3
SbH3
H2 O
H2 S
H2 Se
H2 Te
HF
HCl
HBr
HI
Trivial name
Diborane
Alane
Gallane
Methane
Silane
Germane
Stanane
Ammonia
Phosphine
Arsine
Stibine
Water
Hydrogen sulfide
Hydrogen selenide
Hydrogen telluride
Hydrogen fluoride
Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
Page 44
IUPAC name
Diborane(6)
Alane
Gallane
Methane
Silane
Germane
Stanane
Azane
Phosphane
Arsane
Stibane
Oxidane
Sulfane
Sellane
Tellane
Hydrogen fluoride
Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
The electron-deficient hydrides are those that cannot complete an octet of electrons around the central atom. They are
chiefly the Group 13 elements, although the gas-phase only species BeH2 also fits this description. The electron precise
compounds are those that have an octet of electrons, while the electron-rich elements have additional electrons belonging to
the central atom that function as lone pairs. Note that although they are all electron rich, this group of compounds vary
greatly in the availability of these extra electrons.
Note that while the structures of these hydrides follow similar patterns, i.e. all the EH 3 in group 15 are pyramidal, the
detailed structures differ significantly. Consider the following table of bond lengths for both groups 15 and 16:
Group 15 hydrogen compounds
Bond angle Group 15 hydrogen compounds
NH3
106.6
H2 O
PH3
93.8
H2 S
AsH3
91.8
H2 Se
SbH3
91.3
H2 Te
Source: A.F. Wells, Structural Inorganic Chemistry, Oxford University Press (1984)
Bond angle
104.5
92.1
91
89
In the last problem set, you have explored the origins of some of these changes using molecular orbital methods. The
bending of H–E–H systems is due to a Jahn-Teller effect that lowers the energy of what would be a degenerate set of E π
orbitals in the linear molecule. The greater bending in the third and subsequent periods is due to a second-order Jahn-Teller
effect, and is a consequence of the greater similarity in the atomic orbital energies of hydrogen and elements such as sulfur
and phosphorus, compared to the large difference that exists in water and ammonia. Consult the problem set and its answer
guide to review this important point.
An important property of the electron rich hydrogen
compounds is that of hydrogen bonding. In the figure at right,
the boiling points at one atmosphere of the hydrogen compounds
of groups 14 through 17 are graphed in a comparative manner.
We can take the electron precise series from group 14 as a
reference point in that no hydrogen bonding can occur for these
elements. Indeed, the boiling point curve for this class of
compound is a smooth curve, with increasing bp as the mass
increases. Compared to these, all the electron rich compounds
from groups 15, 16 and 17 have higher boiling points, indicative
of stronger dipole forces operating in these polar molecules.
However, the unusual displacements to higher values of H2 O,
HF, NH3 and HCl are indicative of additional inter-molecular
forces, and these have been identified as hydrogen bonding.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 45
Hydrogen bonds have energies considerably smaller than those of the equivalent covalent bonds, but are still
significantly strong. The table below presents some comparisons between H-bonds and E–H covalent bond energies, and
also provides metrical data for the short contacts that are typical for hydrogen bonding.
Hydrogen bond
HS–H⋅⋅⋅SH2
H2 N–H⋅⋅⋅NH3
HO–H⋅⋅⋅OH2
F–H⋅⋅⋅ F–H
HO–H⋅⋅⋅Cl–
[F⋅⋅⋅H⋅⋅⋅F]–
E––E distance, Å
2.94 – 3.15
2.48 – 2.90
2.45 – 2.49
2.95 – 3.10
2.27
∑v.d.Waals radii, Å
3.70
3.00
2.80
2.70
3.20
2.70
Energy (kJ/mol)
7
17
22
29
55
165
Covalent Bond
S–H
N–H
O–H
F–H
Cl–H
H–F
Energy (kJ/mol)
363
386
464
565
428
565
Because of the imprecision of locating hydrogen in X-ray
diffraction studies, the existence of H-bonding is inferred
from the close approach of the elements themselves. Hbonds vary greatly in distance and energy, but some
examples are shown in the figure. The E⋅⋅⋅H–H bond angles
can vary widely between 150° and 180°. The figure at right
presents the potential energy curve for typical H-bonds,
which have double minima. The low barrier between the
minima allows for facile H+ shuffling. The extremely stable
H-bond in the symmetrical [F⋅⋅⋅H⋅⋅⋅F]– ion is unusual in
having a single potential energy minimum. This ion is
always found to be linear in its salts, which include NaHF2
and KHF2 .
The bonding in this ion is explored on the next
assignment.
3.2
Stability of Hydrogen Compounds
The stability of the main group hydrogen compounds is an important consideration. The table below depicts the
Standard Gibbs energies of formation of the hydrogen compounds. Although the saline hydrides are all exoergic, they are
still very reactive species. But among the p-block hydrogen compounds, only the second period derivatives are exoergic, and
in all cases hydrogen compounds become progressively less stable down each group. The weak E–H bonds formed with
these heavy elements is usually attributed to poor overlap between the large valence atomic orbitals of these elements and the
small, compact hydrogen 1s orbital. Note that better energy matching exists for most of these elements with hydrogen than is
the case for C, N, O and F! Bond enthalpies for the E–H bonds are shown in the bar graph.
Chemistry 3810 Lecture Notes
3.3
Dr. R. T. Boeré
Page 46
Electron Deficient Hydrides of the Boron Group
By far the best-defined electron deficient hydrides are those of boron. Aluminum and gallium have few binary hydrides.
However, for the whole group the anionic, electron-precise, hydrides of the type M[EH 4 ] are extremely important. Some
common salts are NaBH4 , a very mild reducing agent that can be used both in organic chemistry where it is a selective
reagent and in inorganic chemistry where less-vigorous reactions are desired. LiAlH4 is a very powerful reducing agent that
is used in finely powdered form, typically using ether type solvents. It is used throughout organic and inorganic chemistry as
a source of hydride in both hydrogenation reactions and in general reduction reactions. We will seem many examples of
these reagents in use. Gallium also forms such hydridic salts. For each, they can be considered at the result of a reaction
between Lewis acid EH 3 and Lewis base H– . The hydride source is typically the corresponding saline hydride MH.
EH 3
+
H−
→
[ EH 4 ]
−
All the Group 13 hydrogen compounds are hydridic, because hydrogen is more electronegative than these elements.
However, for boron this difference is very small. Thus diborane or adducts of borane are common reducing agents in organic
chemistry, but the react by addition to unsaturated bonds, i.e. they react in concerted reactions typical of covalent species
rather than in nucleophilic reactions typical of ionic species. We start our discussion by focusing on the two most important
boron hydrides, BH3 and B2 H6 .
3.3.1 Borane and its adducts
Borane, BH3 is a trigonal planar molecule. The bonding was also developed on the previous problem set, so that you are
quite familiar with it. Here it is again for consideration. Note the good energy match between boron and hydrogen, so that
high covalency is expected for the B–H bonds (as is indeed found to be the case for this molecule).
Noteworthy in this bonding depiction is that the bonding MO’s are quite stable, and hence not expect to be reactive. The
dominant chemistry of borane involves the high-lying 1a” orbital, which corresponds to a pure p orbital unhybridized on the
boron atom. This is a powerful acceptor orbital, and explains the strong Lewis acid character of borane. Indeed, BH3 does
not exists in condensed phases except when it forms such adducts. It is sold commercially as ether adducts in solution.
Although BH3 can exist under special conditions in the gas phase, it reacts with itself to form dimeric B2 H6 , which is itself a
gas at room temperature.
All adducts of borane are the conceptually the result of Lewis acid-base reactions. For example, the trimethylamine
adduct:
BH3
+
NMe3
→ H3 BNMe3
or the THF solution that is a common commercially available reagent for hydroboration
BH3
+ O [ CH2CH 2 ]2
→ H 3BO [ CH 2CH 2 ]2
But in practice, such species are actually prepared from diborane.
Borane adducts (or in some selected cases, diborane itself) are extremely important in organic reaction for
hydroboration, which is initiated by a concerted addition of a B–H bond to an alkene or an alkyne, producing
organoboranes. In most organic methods, these adducts are subsequently destroyed to produce a range of useful derivatives.
For example, peroxide workup produces alcohols, hydroxylamine sulfonic acids produce amines, and carboxylic acids lead to
a hydrogen atom on the carbon where the boron was attached. But the organoboron compounds are of interest in themselves.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 47
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 48
3.3.2 Diborane, or diborane(6), B 2H6
Diborane, or diborane(6) in its systematic name, has the following unusual structure:
H
1.19Å
H
H
H
97° B
B
H
1.37 Å
122°
H
The two planes are exactly orthogonal, and the two middle hydrogen atoms form symmetric bridges between the boron
atoms. Note also that the H–B–H bond is greater than 90°, so that the boron atoms are in fact quite close to one another. One
way to understand the formation of diborane in the reaction:
2 BH 3
→ B2 H 6
is as the approach of two borane molecules, with the electrons in an H–B bond acting as a Lewis base:
H
H
H
B
H
H
H
B
H
H
H
H
B
B
H
H
Diborane is prepared by the following reaction that is carried out in a higher-boiling ether solvent to make the separation
of the product easier:
3 LiEH4
+ 4 BF3
→ 2 B2 H6
+ 3 LiEF4
( E = B, Al )
The
reaction
must
be
undertaken in the complete
absence or air, because
diborane burns spontaneously
in air, usually with a bright
green flash (similar to the
green produced in fireworks
through the addition of borax
to the gunpowder.) Handling
diborane requires the use of a
vacuum line.
A schematic
diagram of such a (glass)
vacuum line is shown next.
Diborane
was
first
prepared by the great German
synthetic
chemist,
Alfred
Stock. Stock first prepared
most of the hydrides of the
heavier p-block elements. In
order to carry out this work
Stock developed the vacuum
line. His first designs were
based on simple extensions of the use of glass tubes to collect stable gases under water. Stock replaced the water with liquid
mercury. Later most of the mercury was replaced by more glass tubing and innovative mercury float valves. The trend in
modern designs has been to move away from mercury altogether by the use of electronic or mechanical pressure gauges. The
stopcocks used for valves on modern vacuum lines can be made either from ground-glass parts, or from a variety of designs
in which a Teflon plunger fits into a carefully machined glass barrel. The advantage of the Teflon stopcocks is that no grease
needs to be used, generally resulting in a cleaner and longer-lasting vacuum system.
Diborane has long posed a challenge to chemists to describe its bonding since it breaks the rules of the two-electron bond
that developed from the Lewis theory. The first explanation for bonding was the introduction of a so-called 3-centre/2electron bond, which was based on a partial MO approach, but used VB theory much the way that organic chemists explain
the π-bond in ethene as a delocalized orbital while using hybrid orbitals to describe the σ-bonds. To do this, we use two sp 3
hybrid orbitals on each boron and the bridging hydrogen 1s atomic orbitals. The terminal B–H bonds are “normal”.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 49
This model uses “regular” two-electron bonds for the terminal B–H bonds, which are correspondingly shorter. They are
produced using standard sp 3 orbitals on boron overlapping with a single H 1s. The bridge bonds are then produced by the
indicated delocalized bonding schemes with a bonding MO Ψ b from all in-phase orbitals, a totally non-bonding MO Ψ n for
which there can be no H-contribution due to orthogonality. Finally there is an out-of-phase combination Ψ a. Note that (A)
this set of MO’s is remarkably similar to that for a bent array of trihydrogen, but that (B) there are two equivalent and
degenerate sets of these for the top and bottom bridges. Filling this MO scheme with four electrons fills Ψ b and leaves
everything else empty, thus is the ideal filling pattern. Although this molecule is clearly begging for a complete delocalized
MO treatment, the concept of a 3c,2e bond is extremely useful, and there are now many recognized examples of such
situations.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 50
A delocalized MO approach to the electron-deficient diborane molecule.
The following description is based on AM1 semi-empirical
calculations, but the results have been checked for accuracy against highlevel ab initio calculations. B2 H6 is a large molecule for determining a full
MO diagram, and in any case, it is not so obvious what fragments one
should consider in the construction of an orbital interaction diagram. Here
we choose to build up the MO diagram for diborane in analogy to the
chemical synthesis, i.e. a 2 + 2 symmetry-allowed concerted dimerization of two BH3 units, as depicted at right. The
dominant interactions in the interaction diagram are shown for this double HOMO-LUMO interaction. The other major rearrangements are in the in-phase and out of phase combinations of the 1a 1 ’ orbitals, leading to the strongly bonding 1a g MO.
Most of the other orbitals do not greatly change their energy compared to the original BH3 orbitals.
BH3
B2H6
BH3
Fragment orbitals
2b 1g
3b 3u
3b 2u
4a g
2b 1u
3a g
1b 3g
2b2u
↑↓ ↑↓
1b 1g
↑↓↑↓2a g
↑↓
↑↓
1b 3u
1b 1u
↑↓
↑↓
↑↓
1b 2u
↑↓↑↓
↑↓
1a g
Molecular orbital sketches
1ag
1b 2u
1b 1u
1b 3u
2ag
2b 2u
1b3g
3ag
4ag
2b 1u
3b3u
2b 1g
1b 1g
3b2u
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 51
There are several remarkable differences between the qualitative VB-MO approach shown first, and this full MO
treatment. First, there are not two equivalent bridging orbitals as the VB treatments suggests. The single 1b u MO indeed
looks like these bridging orbitals, but that’s only one of them. The rest of the bonding in the dimer is from the 1a g and 2a g
orbitals, both of which are fully delocalized over the central B2 H2 unit. Note also that there is substantial direct B––B
bonding in both of these orbitals, consistent with the large H–B–H bond angles in the bridging part of the molecule.
The terminal B–H bonds are comprised of the 1b 2u , 1b 3u , and 1b 1g orbitals. 1a g and 2a g also contribute to terminal
bonding.
Note also that the bridging orbitals are not the HOMO of the molecule, as you might expect from the VB description. In
fact, one of the terminal B–H bonding MO’s is the HOMO, and the fully bridging orbital is buried quite deeply.
In order to verify this bonding description, we should compare it to the photoelectron spectrum. Below are two UV-PES
recorded for B2 H6 , the first using the He I source, which gives better resolution at low energy, the other with HeII, which can
also ionize more deeply held energy levels.
These spectra are taken from John M. Dyke, e.a. J. Phys. Chem. 1996, 100, 2998-3004. These workers have analyzed the
energies of the bands, and find the following values: 11.88, 13.35, 13.93, 14.76, 16.08, 21.42 eV. This corresponds very well to the
AM1 energies for the occupied MO’s: –11.39, –12.12, –13.05, –13.14, –18.69, –27.89, and detailed calculations confirm the
identity of the MO’s with those in our calculation. The band from ionization of the 1b 2u orbital has identifiable vibrational fine
structure in the HeI spectrum. This is consistent with ionization from the 1b 2u orbital which involves terminal B–H bonding:
removing one of these electrons is expected to strongly affect a H–B–H bending vibrational mode. The other bands show broad
envelopes, also suggesting changes in vibrational energy, but these are not resolved.
There are many ways to analyze the bonding in this molecule. If we compare it to the expected pseudo-Lewis structure
bond orders, we would want the equivalent of 4 terminal B–H bonds, and a total of four electrons, or two “bonds” holding the
bridge together. In the MO scheme, we see that 1b 2u , 1b 3u and 1b 1g are purely B–H terminal bonding. However, both 1ag and
1b 1u are split fairly evenly between B–H terminal and bridge bonding, so we could say that half of each of these MO’s
contribute to terminal bonding. That is an almost exact match to the expected bond order. By the same reasoning, 1a g and
1b 1u are worth one bond order for bridging bonding, and the rest of the bridge bond is the 1b 1u MO, for the total expected
bridging bond order of 2.
This is fine insofar as it goes. But the energy level diagram tells a very different story. The fragment orbital energies in
the diagram were computed at a fragment separation of 5 Å, a distance at which there was no change in the fragment orbital
energies compared to that of a free BH3 . Net bonding is therefore only observed for energy levels that change on dimer
formation, and this is seen in the scheme to be primarily the change in 1a g and 1b 2u . These correspond to the in-phase and
out-of-phase combination of the 1a 1 ’ fragment orbitals. Now if these interactions were equal and opposite, there would be
not net dimer bonding! But they are not. Mixing in of one of the 1e’ fragment orbitals causes the 1b 2u MO to lose
antibonding character, and become terminal B–H bonding, but fragment non-bonding. The difference in the dropping in
energy of 1a g and the raising in energy of 1b 2u corresponds to the main bond energy in the dimer.
For the remaining filled orbitals, three of them are essentially unchanged from their fragment roots (1b 3u , 2a g and 1b 1g ).
For the remaining MO, substantial re-organization takes place, with one of the 1e’ set being exchanged with an orbital
derived primarily from 1a 2 ”. However, despite this massive change, the energies of these four MO’s are hardly changed at
all, so these do not contribute to the binding energy of the dimer .
Important to the chemistry of diborane is that the LUMO+1 orbital 1b 3g acts much as the acceptor orbital of BH3 , and the
dominant chemistry is that of a potent Lewis acid, despite its internal dimerization. The HOMO is not a very reactive orbital,
just as in BH3 .
On the next page of these notes is a chart that correlates the orbitals of the fragments using the labels that they will have
in the D2h point group. These labels were determined manually, because the 5 Å-separated fragments do not belong to this
point group.
Symmetry Correlation of the BH3 fragment orbitals with those of B 2H6
Definition of Fragment Orientation:
2b3u
2b1g
4ag
3b3u
3ag
2b1u
1b3g
1b1u
1b3u
1b1g
2ag
2b2u
1ag
1b2u
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 53
3.3.3 Digallane Ga 2H6 and gallaborane GaBH6
Digallane was first reported in 1989, and more recently the mixed molecule gallaborane was prepared. A recent
summary of the chemistry of these species has been published: A.J. Downs, e.a. Inorg. Chem. 2001, 40, 3484-3497. Both are
prepared by similar reactions, and both are unstable at room temperature. They can be handled in the gas phase in
scrupulously clean glass vacuum lines (complete absence of trace water is essential.)
1
[ H 2GaCl ]n
n
+
LiEH 4 →
[ H 2GaH 2 EH 2 ]
+ LiCl
( E = Ga, B)
In the solid state, both species apparently polymerize, usually irreversibly. But in dilute solution and in vacuum, they have
the same bridged structure as diborane. The point groups are D2h for the two symmetrical species, and C2v for gallaborane.
Gallaborane is a fascinating molecule, and the authors of this article include pertinent NMR data in tolune-d 8 solution.
These spectra are depicted below. The small peaks at 3 and 8 ppm in the 1 H NMR spectra are due to residual solvent peaks.
Thus at –15°C, there is a 1:1:1:1 quartet at 2.0 ppm for the two terminal hydrogens on boron, while there is a broad signal at
6.0 ppm. Comparison to other species shows the 6.0 ppm signal to be due to the GaH atoms, while the quartet at 2 ppm is
typical for a BH unit. Only the boron signal at –24 ppm is from gallaborane. The other signal is from some form of
molecular aggregate that forms in solution. The author’s speculate at to the identity of this species, and attribute it to an
eight-membered ring formed by head-to-tail dimerization of monomeric gallaborane.
The origin of the temperature dependence of the 1 H NMR spectrum has been attributed to varying degrees of thermal
decoupling of the protons from the quadrupolar Ga and B nuclei to which they are bound. Be sure to try to interpret these
NMR spectra!
There is intense interest in the chemistry of Ga 2 H6 and related species, because the heavier p-block hydrides are among
the best sources of highly purified elements used in the construction of bulk semiconductors, and most importantly, in the
fabrication of logic chips. Thus, there is a great deal of interest in the GaAs mixed-element semiconductor, which can
support intrinsically faster computing speeds than any silicon-based chip. Besides this, elements such as P, As, Al and Ga are
routinely used to produce nand p zones in the semiconductor, which are the basis for the transistor effect in semiconductors.
The general term to these processes in which a reactive gaseous source of an element is combined with atoms on a solid
surface is chemical vapour deposition, usually going under the name CVD. This is most certainly the diving force behind
all attempts to develop the chemistry of the gallium hydrides.
Chemistry 3810 Lecture Notes
3.4
Dr. R. T. Boeré
Page 54
Electron Precise Group 14 hydrides (other than the hydrocarbons)
Most important are the silanes. These were also first discovered by Alfred Stock. It is likely that both linear and
branched chain polysilanes are known up to Si9 H20 . However, many have not been isolated, and the evidence for these
species is from gas chromatography studies of crude mixtures of the silanes.
Stock identified SiH4 , Si2 H6 , Si3 H8 and two isomers of Si4 H10 . The silanes are thermally less stable (in the absence of
oxygen) than corresponding hydrocarbons. For example, SiH4 is converted to Si and H2 above 500°C. Although all the
hydrocarbons burn easily, silanes ignite spontaneously in air, indicating a greater rate of reaction with oxygen.
There is considerable similarity in the synthesis and reactivity of the silanes and the boranes. This is one example of a
general property of the main group elements, the so-called “diagonal principle”. This rule states that there is often unusual
similarity between two elements when one is located one period lower and one group to the right in the table. The origin of
this rule can be traced to the trend in electronegativity, which increases diagonally across the periodic table up to fluorine and
helium. Consequently, diagonally related elements often have greater similarity in electronegativity than elements in the
same group. Note that the diagonal relationship is not one of exact similarity, so that for example the boron hydrides have
the formulas BH3 , while silane is SiH4 .
3.4.1 Synthesis of Group 14 hydrides
The synthetic routes to such compounds start from either the element halides or oxides, almost invariably using Group
13 reducing agents. Thus for example, in the Chem3810 laboratory we prepare germane, GeH4 by the reaction:
HGeO3−
BH −4
+
+
2H +
→ GeH 4
+
H 3 BO3
This reaction actually produces a mixture, with some polygermanes being formed at the same time. The HGeO3 – is produced
by dissolving GeO2 in a strong alkali solution (OH– ). These higher hydrides are largely H3 GeGeH3 , with some of the threegermanium chains as well. Notice that chain formation still requires reducing agent, the difference being as to whether
reduction occurs with H nucleus attachment, or whether only an electron is transferred allowing for the formation of Ge –Ge
bonds.
Silane is commonly made by the reduction of silicon(IV) fluoride with lithium aluminum hydride:
SiF4
+ LiAlH4
→ SiH 4 + LiAlF4
Such a reaction is readily understood as a form of metathesis reaction, as studied in Chemistry 2810, driven by HSAB
affinities. Under certain conditions, higher hydrides can als o be made.
3.4.2 Properties of these hydrides
Straight and branched chain hydrides are know for Sin H2n + 2 up
to n = 10, and for Ge n H2n + 2 up to n = 9. In principle, they act as
heavier forms of the hydrocarbons, as can be seen from the graph at
right that compares the boiling points of some straight-chain
analogues of C and Si. But there are far fewer known derivatives of
the heavier Group 14 hydrides, and they have significant differences
in properties. Consider these differences:
• The bond energy of C–H bond is about 416, and of a C–Cl
327 and C–O 359 kJ mol–1 , while for Si–H it is 326 and Si–
Cl 391 and Si–O 466 kJ mol–1 . By comparing these values,
we see that bonds from carbon to hydrogen are stronger
than corresponding bonds to Cl or O, while bonds from
silicon to hydrogen are weaker than those to Cl or O. The
same situation occurs for Ge and Sn.
• Methane is chlorinated with some difficulty (radical
chlorination), while silane reacts violently with Cl2 .
• CH4 is stable with respect to hydrolysis, whereas SiH4 is readily attacked by H2 O.
• SiH4 is spontaneously flammable in air and, although it is the kinetic stability of CH4 with respect to reaction with
O2 at room temperature that is crucial, values of the enthalpy of combustion for methane and silane are significantly
more exothermic in the case of silane.
• Catenation (chain-forming) is more common for C than the later group 14 elements, and hydrocarbon families are
much more diverse than for any of their Si, Ge, Sn or Pb analogues.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 55
3.4.3 Bonding in dislane.
The silanes are also more readily cracked when heated in the absence of air. This is similar to the “cracking” reaction of
the petroleum refining, which requires higher temperature and catalysis to occur at a reasonable rate. For example, disilane
undergoes the reaction:
Si2 H6
o
400 C

→ SiH 4
+
H2
higher silanes
We might ask ourselves, Why does the Si–Si bond in dislane, Si2 H6 break upon heating more easily than the carboncarbon bond in ethane? The origin is found in the bonding of these two molecules. At the AM1 level of theory we can
compare the energies of the bonds in ethane, disilane and ethene. Now ethene has a π-bond, and in organic chemistry this is
considered to be a functional group, i.e. a center with special reactivity compared to the saturated hydrocarbons. The
functionality is primarily due to the high-lying HOMO (the π-orbital) and somewhat from the low-lying LUMO (the π*orbital). The small separation between these two orbitals allows a double bond to act as a chromophore, i.e. it absorbs light
at wavelengths longer than 200 nm, which is the cutoff for the transmission of UV light in air. Consider the following energy
level diagrams as calculated in HyperChem for these molecules:
Energy levels of Ethane
Energy levels of disilane
Energy levels in ethene
Here we see that the HOMO-LUMO gap in disilane is considerably smaller than that in ethane (using AM1, we calculate
9.7 eV, compared to 15.8 eV in ethane). In fact, the gap in disilane at the AM1 level is somewhat smaller than the gap in
ethene. This means that a Si–Si single bond should be considered as a functional group in an organosilicon compound. In
addition, silanes with their single bonds are UV chromophores, and this leads to the observation that UV light is capable of
cleaving Si–Si single bonds, something that is not true of the saturated hydrocarbons! We note that whereas excitation of one
π-electron in ethene from the HOMO to the LUMO reduces the bond order by one, the molecule is still expected to remain
bonded (although now free to rotate, and hence UV-light can isomerize alkenes from cis to trans, or vice versa – this the
process involved in the detection of light by the retina.) But if photons of about the same energy are absorbed by disilane, its
bond order is reduced to zero, and the Si–Si bond is expected to cleave.
The topology of the HOMO and the LUMO for ethane and disilane are very similar, as shown in the orbital wire-frame
graphs below. The difference in bond strength then is primarily in the degree of stabilization and destabilization that occurs.
HOMO (σ), 2a 1g
LUMO (σ*), 2a 2u
At this stage you should stop and ask yourself what the reason for the smaller HOMO-LUMO gap in silane is fundamentally
attributable to.
Chemistry 3810 Lecture Notes
3.5
Dr. R. T. Boeré
Page 56
Electron-Rich Hydrides
The common hydrides are as previously mentioned NH3 , PH3 , AsH3 and SbH3 for group 15, H2 O, H2 S, H2 Se and H2 Te in
Group 16, and the hydrohalides in group 17. There are no hydride derivatives of the group 18 elements. In addition, there
are some polyhydrides, such as hydrazine H2 NNH2 and hydrogen peroxide HOOH. The most extensive series of catenated
hydrides occur for sulfur, and indeed sulfur shows the greatest tendency to form homocatenates after carbon. Thus H2 Sn
chains for n = 2 to 6 have been prepared, and these are known as polysulfanes. We have already discussed water and
ammonia extensively in Chemistry 2810, and will not say any more on these compounds at this time. The Group 17
derivatives were also discussed in the previous course.
Hydrazine is an important example of an endothermic compound. Thus ∆H°f is +50.6 kJ mol–1 . However, it is
kinetically stable towards the formation of the elements at room temperature. Hydrazine is used as a rocket fuel in some
rocket systems. For example, some of the rockets used in the Apollo missions used the reaction between N2 H2 and N2 O4
(products H2 O and N2 ) as propellant. It dissolves in water, and forms a stable monohydrate which is the most common
industrially used form of hydrazine. Direct dehydration of this hydrate is difficult (decomposition of the hydrazine sets in),
and anhydrous hydrazine is usually produced directly, e.g. by the reaction:
2NH 3
+
[ N2 H 5 ][ HSO4 ]
→ N2H4
+
[ NH 4 ] 2 [ SO4 ]
The structure of hydrazine in the gas phase has been shown to be the gauche form:
The N–N bond distance is 1.45 Å, and the N–H distance is 1.02 Å. Hydrazine is a powerful reducing agent, and that is its
main application. It finds uses in the agricultural and plastics industries, and is added to boiler water to remove O2 as an
impediment to corrosion.
3.5.2 Heavy group 15 hydrides
The important species are phosphine, arsine and stibine. PH3 is an extremely toxic, colourless gas wich is much less
soluble in water than is ammonia. It is neither Brønsted acidic nor basic, but is an excellent Lewis base. Strong acids do
protonate PH3 to PH4 +, but this ion is decomposed by water. Alkali metals can reduce PH3 to PH2 – in non-aqueous solvents
(e.g. ethers or liquid ammonia), but again this ion is destroyed by water, leading to POH-type species.
There is also a phosphorus analogue to hydrazine, diphosphane, P2 H4 , which has a similar structure. It is also a
colourless liquid.
The heavier hydrides AsH3 and SbH3 are similar to phosphine, but less thermally stable. The bond angles decrease on
going down the group, as was previously mentioned. The bonding issues were dealt with in Problem Set #3.
3.5.3 Heavy Group 16 hydrides.
The parent compounds are hydrogen sulfide (sulfane), hydrogen selenide (selane) and hydrogen telluride (tellane). All
are toxic, foul-smelling gases at room temperature (bp.’s are 214, 232 and 271 K, respectively). H2 S is an important
byproduct of the natural gas industry in Alberta, where it is used as the primary source of elemental sulfur (see Chem2810
notes). The polysulfanes are produced by protolysis of the polyanions. First sulfur is dissolved in aqueous solutions of Na 2 S
to lead the polysulfide anions, Na 2 Sn . Then acid is added to the basic solutions:
2 H 3O +
+
Na2 S n
→
H2 S n
+ 2 Na +
+ H 2O
As mentioned previously, the bond angles in the series of EH 2 decrease steadily from O to Te, and this can be
rationalized by MO theory (Problem set #3). The effect of decreasing the bond angles is to make the E atom a better Lewis
base, and this is indeed the case in practice. At the same time, these hydrides become better Brønsted acids on going down
the group, which can be attributed to weaker E–H bonding (just as HI is a stronger acid than HCl or HF).