Download 4. Main-Group Metals and Organometallic Compounds

Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 57
4. Main-Group Metals and Organometallic Compounds
The information we will be considering is somewhat spread-out in Shriver-Atkins. The s and p-block metals are
discussed in Chapter 9 (sections 9.1 – 9.5; 9.12 – 9.18) and Chapter 15 (all of this chapter). Be sure to read these sections of
the text and study the example problems that are liberally spread through these sections.
This presentation of the periodic table serves to remind us that 80% of all elements are metals. We dedicate a separate
course (Chem3810) to consider primarily the chemistry of the transition elements, i.e. those that have partly filled d orbitals
in at least one common oxidation state. This excludes the elements of Group 12. Also, silver in group 11 is dominated by
chemistry of the +1 oxidation state, so it too does not behave as a typical transition metal. The chemistry of copper in the +1
oxidation state, of Ag(I), Zn(II), Cd(II) and Hg(II) will therefore be briefly considered here by lumping them with the p-block
metals.
The chemistry of the lanthanides (the first f-series) is closely related to that of the Group 3 elements, and like these
elements the dominant oxidation state is always 3+. Thus these elements are pseudo-noble gas ions, and their chemistry is
predominantly that of ionic and weakly coordinated species analogous to the Group 2 elements. On the other hand, the
actinides behave more as the d-block elements do, and there is an extensive chemistry particularly of uranium, since the 238 U
isotope is long-lived and the element is not too radioactive to preclude conventional chemical studies.
In this course we will focus on the Group 1 and 2 metals and the p-block elements, with an emphasis on their covalent
derivatives.
4.1
Some basic properties of the metallic elements
4.1.1 Enthalpies of Vaporization of the metallic elements (kJ/mol)
The graphic at right is a bar graph of the enthalpies of
vaporization of the metallic elements in kJ/mol. For
metallic lattices, this data represents the strength of the
metal-metal interactions in the solid-state metallic lattice.
We recognize that this distribution has the general trend of
being low at the left and the right-hand side. It peaks in
the middle, which corresponds to d-block elements with a
half-filled or close to half-filled d-valence orbital. Thus
the resistance of tungsten towards evaporation has been
attributed to covalent d-d bonding. Consistent with this
interpretation is that tungsten is one of the most brittle of
the metallic elements. It is a conductor, and it is correct to
call it a metal, but it is a metal with strong covalent W–W
bonding in addition to the metallic bond that allows for its
metallic properties. This aspect of tungsten chemistry
accounts for its ubiquitous role as the filament material of
incandescent light bulbs.
The s and p block metals, on the other hand, tend to be the most volatile. The larger enthalpies of vaporization for the pblock metals reflects the presence of partly filled p-orbitals which can lead to stable E–E bonding.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 58
4.1.2 Electron-Deficient Metal Compounds: Clusters
Metals are electron deficient species, and the metallic bond that leads to the classic properties of metals (high thermal
and electrical conductivity, malleability, metallic sheen, typical silver colours) are a consequence of the delocalized bonding
that occurs as they attempt to form the best bonding patterns that are available with the limited number of electrons that are
present.
It should therefore not be a surprise that when these elements are reacted with fewer ligands than required by the octet
rule or the 18-electron rules they tend to form cluster compounds. These can be thought of naively as small chunks of the
metallic element, with the structure usually adopted at room temperature for bulk metal, surrounded on the outside by a
suitable set of ligands. The distribution of such cluster compounds over the periodic table is shown in the following graphic:
There are several element regions where cluster formation is common. We have already mentioned simple boron derivatives
that are electron deficient (e.g. the as seen in self-dimerization of BH3 ). There are also cage structures for the hydrocarbons,
such as tetrahedrane, C4 H4 , and cubane, C8 H8 , but these are exceptions rather than the rule. Also, they are not truly electron
deficient. Clusters are also common for organometallic derivatives of lithium and sodium (see below). With Rb and Cs the
formation of suboxides is common, and there is an extensive fragment population among the transition elements. The heavy
p-block metals tend to form aggregated structures. We will likely not have time to discuss this important class of elements.
4.2
s-Block Metals
The crustal abundances of the s-block elements are shown in the
graphic at right (note that the numbers in this bar graph are the logarithms
of the abundances.) Sodium and potassium are high abundance minerals,
and both are mined in large quantities. Magnesium and calcium are largely
present as a part of rocks and minerals (limestone, dolomite, etc.) Cesium
and beryllium are expensive both because of their lower abundance and
because of difficulties in handling. Cesium must be stored in an inert
atmosphere or under vacuum. Beryllium is highly toxic. The latter
element finds an extremely important application as the most X-ray
transparent element with reasonable structural strength. It is universally
employed as the “window” on X-ray tubes (for both medical and scientific
applications) that allows the X-rays flux to emanate at high intensity from
the vacuum chambers in which they are generated.
In Chemistry 2810 we focused extensively on the reactivity of the
metallic elements of Group 1. In aqueous solution, lithium has the most
negative standard reduction potential, but cesium has the lowest ionization
potential. How can this apparent contradiction be squared? It is of course
due to the difference in solvation between these elements in aqueous
solution. The thermodynamic cycles shown below compare lithium and
cesium in the elemental form with both the gas phase and aqueous phase ions. Although lithium has almost 150% the
ionization energy of cesium, it also has a correspondingly higher enthalpy of hydration. The net balance between these two
terms is that the solution ionization energy of these elements are remarkably similar, and this is true for the intervening
Group 1 elements as well. However, on balance the solution ionization energy (which correlates closely to the redox
potentials) is greater for lithium. The large value of the lithium hydration energy is attributed to its very small size (r(Li+) =
0.90 Å for CN6). This is just one more example of the general principle that the second-row elements have distinctly
different properties from those of the subsequent elements within any chemical group.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 59
Thermochemical cycles for oxidation of Li and Cs
4.2.1 Complex Formation: Crown ethers and cryptands
Modern developments in the chemistry of the s-block elements have largely dealt with non-aqueous environments. In
order to dissolve these reactive elements in suitable solvents, it is essential to supply appropriate ligands that stabilize their
charges and facilitate their dissolution. Typical solvents for this work have included ethers, amines and aromatic
hydrocarbons (benzene, toluene).
By far the most important complexes are those formed between the metal ions and crown ethers and cryptand ligands.
The crowns, developed by Peterson and later by Cram, are cyclic polyethers that have the maximum chelate effect, and
whose size can be adjusted by the number of donor atoms and the spacing of the oxygen atoms by ajustment of the number of
CH2 groups in the rings. They can thus be tuned to select out different ions based on their sizes.
The crown effect can be optimized even further by the cryptand ligands,
which combine crown ethers with trialkylamine functionality. The nitrogen
centers are three coordinate, turning rings into cages. The cryptand ligands have
been extensively developed into selective chelating agents for the s-block
elements. The structures of come examples are provided below, and the graph at
right maps the affinity for 2.2.1 and 2.2.2 crypts for the Group 1 metal cations in
terms of the formation constants for complex formation (again expressed as
logarithms). The former is optimally sized to coordinate sodium ions, and the
latter is best for potassium ions.
4.2.2 Ammoniacal solutions of alkali metals: Electrides
All the Group 1 metals, as well as the more reactive Group 2 metals Ca, Sr, Ba dissolve in anhydrous liquid NH3 .
(NOTE: not the common aqueous ammonia that is widely used in the chemistry laboratory.) Ammonia is a gas at room
temperature, but it is liquefied relatively easily by compression, and the liquid can be dispensed if kept at or below the
boiling point of –33.4 °C. Ammonia, like water, has a high enthalpy of vaporization, and open flasks of the liquid can be
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 60
handled for reasonable periods of time. For work with
electrides, it is essential to keep the solutions dry and a typical
approach to handling them is to use a dry-ice condenser in an
apparatus protected by nitrogen or with a drying tube. Note
that metallic sodium is the standard reagent for drying
ammonia, the main purification method that is required.
Indeed, this is a self-indicating drying agent, for so long as the
blue colour of the electride persists, the ammonia must be dry.
Typical apparatus for purifying and handling liquid ammonia
are shown in the figure at right. The cooling agent most
commonly used is solid dry ice suspended in a suitable heattransfer fluid (methanol is effective).
The solutions of metals in liquid ammonia develop an
intense blue colour. From such solutions, the Group 1 metals
can be recovered unchanged (contrast this with their behaviour
in water!) The Group 2 metals are recovered as [M(NH3 )6 ]
complexes, a form of expanded metal. From ammoniacal
solutions of lithium and sodium, the complexes [Li(NH3 )6 ]
(yellow) and [Na(NH3 )6 ] (blue) can be isolated at low
temperature, but on warming these release ammonia and revert
to finely divided metal (CAUTION: highly reactive state!)
The blue colour of such solutions is attributed to the
formation of electrides. That is, a solvated free electron:
liquid NH 3 , bp
M 
→ M +( solvate ) + e−( solvate)
The solvated electron exist in cavities of 3-4 Å radius as
determined by the volume expansion on additon of the metals
to the liquid. The origin of the colour apparently is due to
quantization of the solvated electron.
The most popular model has been to
solve Schrödinger’s equation for an
electron in a spherical box (related to the
3-D cubic box developed in Chemistry
3730.)
The results are a set of
wavefunctions closely related to those of
the hydrogen atom, but with different
quantum number rules. The solution can
be graphed as shown in the picture at
right. The wavefunctions are labeled the
same as the H-atom orbitals for
convenience, but note that the pattern of energies is different. It is assumed that the blue colour is due to a 1s→2p transition
(absorption) leading to the observed λmax at ~1500 nm in the near infra-red. The blue colour is that of the reflected
wavelengths. This wavelength is not a property of the alkali element used to generate the electride, but is a property of the
medium, as shown by the following spectra:
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 61
As expected for solutions containing free electrons, they have distinctive
EPR spectra. The spectra from liquid ammonia solutions are hard to measure,
and somewhat uninformative. However, using organic amines it is possible to
obtain more meaningful spectra, and the spectrum shown at right is an
example of this. Note that the observed spectrum seems to be a result of
several overlapping sets of signals, implying that in this solvent the electron
has more than just one environment. The observed hyper-fine coupling is to
the 14 N nuclei of the amine solvent.
Electride solutions are powerful reducing agents. In the Chemistry 3810
laboratory they are used to reductively cleave a phosphorus carbon bond.
There are many practical applications for electrides in synthetic chemistry.
The solutions are metastable, and they can be decomposed, especially when
catalyzed with iron salts. The presence of rust in most chemical laboratories
requires special care that vessels for handling electrides are kept extremely
clean. The decomposition reaction is analogous to the reaction of the
elements with water, producing the corresponding alkali or alkaline earth
amides. The most commonly used amide is NaNH2 . It is a colourless ionic
compound, but is very hygroscopic and reacts instantly with water to produce
sodium hydroxide (solvent leveling). In non-aqueous solvents it acts as a
stronger base than sodium hydroxide.
2 NH 3
+ 2e−
→ 2NH 2−
+
H2
4.2.3 Ether, alkylamine and cryptand solutions of alkali
metals: alkalides
Another way to generate electrides is to add an excess of crown ether or
cryptand in a suitable solvent to which is added an alkali metal. This produces
a solution of the type [M(crown)n ]+e– . An example of such an adduct of
cesium, [Cs(18-crown-6)2 ]+e– , is shown in a figure below. However, when
only half an equivalent of a crown or cryptand is added, a different reaction
takes place with elemental alkali metal. For example:
2 Na
solvent
+ crypt − [222] 
→
[ Na (crypt − [222]]
+
Na −
In this case, the free electron is trapped by another sodium atom, and the result is the formation of the sodide ion. Such
anionic alkali elements are called alkalides. Recall from Chem2810 that the alkali elements have exothermic electron gain
enthalpies:
H
–72.77
Li
–59.63
Electron-Gain Enthalpy Values for Some Elements (kJ/mol)*
Be
+48†
B
–26.7
C
–121.85
N
+7
Na
–52.87
Mg
+39
Al
–42.6
Si
–133.6
P
–72.07
K
–48.39
Rb
–46.89
Cs
–45.51
Ca
+30
Sr
+30
Ba
0
Ga
–30
In
–30
Tl
–20
Ge
–120
Sn
–120
Pb
–35.1
As
–78
Sb
–103
Bi
–91.3
(the first entry refers to the formation of X – from X; the first entry refers to the formation of X 2– from X – )
O
–140.98
+844
S
–200.41
+532
Se
–194.97
Te
–190.16
Po
–180
F
–328.0
Cl
–349.0
Br
–324.7
I
–295.16
At
–270
*Data taken from H. Hotop and W. C. Lineberger: Journal of Physical Chemistry, Reference Data, Vol. 14, p. 731, 1985. (This paper also
includes data for the transition metals.) Some values are known to more than two decimal places.
† Elements with a positive electron-gain enthalpy indicate that a stable anion A of the element does not exist in the gas phase.
Chemistry 3810 Lecture Notes
This means that the reaction:
Dr. R. T. Boeré
Na( g )
Page 62
+ e − → Na(−g ) is favoured. Usually this is not a favourable reaction in
solution, but the cryptand ligand reverses normal solvation trends. Adducts of this type have even been crystallized, as
shown in the following figures, which depict structures from X-ray diffraction studies of [Na(crypt-[222])]+Na – . Note the
large size of the Na – ion, as expected based on shielding consideration. Note that alkalides are the formal analogues to the
hydride ion in Group 1.
+
–
[Na(crypt-[222])] Na
4.3
+ –
[Cs(18-crown-6)2] e
p-Block Metals
The Group 11 and Group 12 metals are formally analogous to the Group 1 and 2 elements, respectively, but differ due to
the presence of a poorly shielding d 10 filled shell of electrons. Thus, they have uniformly higher ionization energy, and while
cesium is one of the most reactive metals, gold is among the least reactive! We have already considered many consequences
of this phenomenon in Chem2810, so we will not revisit this material now.
For the p-block elements, the important thing to recognize is that they can exist in more than one oxidation state. In
Chem2810 we discussed this important distinction using Frost diagrams. Such diagrams are reproduced below for the Group
13 and Group 14 elements.
The lower oxidation states, e.g. Tl+ and Sn 2+, are a consequence of not oxidizing the ns2 electron pair. The difficulty of
oxidizing these electrons increases down the periodic table, so that Tl3+ and Pb 4+ are both strong oxidizing agents. But of
course for boron or aluminum, the 3+ oxidation state are not only very stable, but the only accessible state. There are many
practical consequences of these changes in the chemistry relevant to the chemical properties of the compounds of these
elements. Finally we note that the heavier p-block elements are only very mild reducing agents.
4.4
Handling alkali element and organometallic compounds
Most cluster compounds of Groups 1 and 2, and of virtually all organometallic compounds that have been prepared, have
very high reactivity. Specialized methods have had to be developed in order to handle them safely and without
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 63
compromising their integrity. Typically, reactive species are susceptible to water, and indeed to any protic source. Thus they
are typically handled only in dried, aprotic solvents, and their manipulations are carried out under an atmosphere of inert
nitrogen gas (or if available, argon, which is superior as an inert gas medium to nitrogen for several reasons, but is quite a bit
more expensive.) Most organometallics are also quite sensitive to oxygen, with some like methyl lithium spontaneously
burning in air.
4.4.1 The Glove Box
There is a variety of isolation chambers that are marketed (or locally built) depending on their application. Thus, nuclear
chemistry tends to use “hot boxes” where highly radioactive samples can be handled, often using remote control of
mechanical arms and levers.
Biohazards are also
contained in special chambers. However, the concept that
is usually understood by “glove box” is an inertatmosphere box with an automated catalytic purification
system, and an interlock chamber for loading and
emptying the box without contaminating the purified air.
The goal of such glove boxes is to eliminate all water and
oxygen from an atmosphere where otherwise fairly
standard chemistry can be undertaken, the operator
reaching in to manipulate the reagents, flasks, beakers etc.
through full-length rubber gloves.
A typical glove box is shown in the photo at right.
Several modifications are now commercially available,
such as double boxes allowing two workers to work on a procedure either facing each other or side-by-side in otherwise
conventional boxes that have been joined together. Some boxes have integrated freezers, and others have microscopes for
examining products and/or mounting crystals in a protected environment.
4.4.2 The Double Manifold Vacuum Line and
Schlenk Apparatus
Many labs involved in this area of research do the majority of
their work in so-called Schlenk apparatus. Some pieces of Schlenk
apparatus are shown below. They must be used in conjunction with a
specialized gas distribution system known as a double manifold
vacuum line. One possible design for such a line is shown in the
figure at right. The taps can alternately remove all gases from the
apparatus which is attached to the bottom connectors, either directly
or, more commonly, via a vacuum hose. Or it can introduce a purified
inert gas such as dinitrogen or argon into the evacuated vessel.
Schlenk apparatus was
specifically designed to allow
for all the usual operations
that chemists are used to performing in air within a controlled atmosphere. The
arrangment shown in the figure at left shows how a vacuum filtration can be
performed under nitrogen or argon. Once the mother liquor has been removed, the
pieces are separated and the glass joints sealed with caps and stoppers. Then the
filtrate and/or precipitate can be used further. For example, precipitates need to be
thoroughly dried under vacuum. The filtrates are often evaporated under vacuum to
recover the solute.
Many modifications
to the basic Schlenk
concept
have
been
proposed. The following
illustration shows how
side-arm flasks can be
used to peform a
filtration under inert gas.
This procedure would be
used
to
remove
unwanted solids from the
reaction mixture.
Chemistry 3810 Lecture Notes
4.5
Dr. R. T. Boeré
Page 64
Introduction to Main Group Organometallic Compounds
There is a strong conceptual link between the element hydrides and organometallic derivatives. In the language
introduced by Nobel winner Roald Hoffman, a singly-bonded hydrocarbon anion group is “isolobal” with a hydride anion:
C
CH3
H
R
"isolobal" symbol = same electron count and same orbital type (σ)
The table at right places the methyl derivatives of the
main group elements into the periodic system (compare
to the similar table in section 3.1 for the hydrides.) They
fall into the same classification schemes of ionic,
electron poor, electron precise and electron rich. The
series is most extensive for methyl; some elements do not
form stable derivatives of primary longer chain alkyl
groups because of instability towards β-elimination. For
most methyl derivatives, the phenyl derivative is known,
however, and these are often quite stable and
considerably less reactive. There is an extensive modern
chemistry with exotic hydrocarbon derivatives that we
will see examples of later in this section.
β-elimination is the reaction:
MCH 2CH 2 R → MH
+ CH 2 CHR
In which the H atom transferred to the metal comes from
the second, or β-carbon.
4.5.1 Examples of E–CH3 Compounds
Consider these representative examples:
Compound
Category
Bonding
Structure
KMe
CsMe
ionic
ionic
ionic
ionic
K+ CH3 – ionic lattice
Cs + CH3 – ionic lattice
[Be(CH3 )2 ]x
[Mg(CH3 )2 ]x*
poylmeric
poylmeric
(3c,2e)
(3c,2e)
doubly-bridged polymeric chain
doubly-bridged polymeric chain
Li4 Me4
Zn(CH3 )3
HgMe2
Al2 (CH3 )2
electron poor
electron poor
electron precise
electron poor
(4c,2e)
(2c,2e)
2c,2e normal bond
(3c,2e and 2c,2e)
(tetrahedron structure)
linear
linear
SiMe4
GeMe4
electron precise
electron precise
2c,2e normal bond
2c,2e normal bond
tetrahedral
tetrahedral
As(CH3 )3
electron rich
2c,2e bond
pyramidal with lone pair
*Note: the more common magnesium methyl compound is CH3 Mg +Br– , i.e. a Grignard reagent! Dimethyl
magnesium is made from the Grigard by a redistribution reaction, see below.
The thermal stability of element methyl compounds also mirrors that of the hydrides. The graph below presents the M –
C bond enthalpies (for methyl derivatives , in kJ/mol) at 298 K. (Data from M.E. O’Neill and K. Wade, Comprehensive
organometallic chemistry, ed. G. Wilkinson, F.G.A. Stone and E.W.Abel, Vol. 1, Pergamon Press, Oxford, 1982.) The table
presents the standard enthalpies of formation. Note that some are exothermic and some endothermic. Since the C–H bonds
in the methyl portion are strongly exothermic w.r.t. their elements, even mildly exothermic compounds tend to be easily
cleaved at the M–C bonds. Especially the heaviest members of each group have easily cleaved bonds, and for most of the
last century tetraethyl lead (cheaper to manufacture than tetramethyl lead, but very similar in properties) was added to inferior
gasoline grades to improve the smoothness of combustion in gasoline engines (a so called “anti-knock” agent.) This effect is
based on the thermal cracking of PbEt 4 into lead and ethyl radicals which act as initiators for the radical combustion of the
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 65
gasoline hydrocarbons. This agent is now banned in North America only because of feared health hazards associated with
the production of lead oxides released from the exhaust gases of the engines, especially from vehicles.
4.5.2 Synthesis of main group organometallics
There are many synthetic routes for the synthesis of main group organometallics. Some are only used in a laboratory
setting because they are too costly for large-scale industrial use. On the other hand, many of the industrial methods are
difficult to implement in the laboratory because of the need for highly specialized equipment and working conditions. The
following reactions will serve as model equations:
• Direct preparation from a halocarbon (usually chloride or bromide) and the metallic element
Preparation of organolithium compounds:
2 Li + RX → LiR + LiX
Preparation of Grignard reagents:
Mg
+
RX
ether

→
RMgX
Rochow method for the production of methyl silyl chlorides
Cu
Si + CH3Cl 
→ (CH3 ) 2CCl 2 + (CH3) 3SiCl + CH3SiCl3
Fluidizedbedreactor
This last reaction, uses a copper catalyst, and is an essential industrial process; without it the production of silicones,
an important class of synthetic hybrid inorganic/organic polyemr not economical.
•
Transmetallation
2 Ga + 3 Hg(CH 3 ) 2
→ Ga(CH3 )3 + 3 Hg
This is the displacement of one metal from a hydrocarbon group by another. The displacing metal must be less
electronegative than the displaced metal. Hence Hg, with the very high value of χ = 2.0 is often used. This
particular reaction is done at about 60°C in a sealed, shaken, heavy wall glass reactor, sometimes called a Carius
tube).
•
Metathesis (which can, however, best be thought of as nucleophilic substitution.)
Li4 Me4 + SiCl4
Al 2Me 6 + 2 BF3
•
→
→
4LiCl + SiMe 4
2 AlF3
+
2 BMe 3
Redistribution reaction type (this is another form of metathesis, and here the label is highly appropriate
SiCl 4 + SiMe 4
→
SiClMe 3 + SiCl 2Me 2 + SiCl 3Me etc
This reaction has considerable importance for the silicone industry, which is economically very important.
•
Addition to a multiple bond
E-H + CH 2 =CH 2
→
E-CH 2 -CH3
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 66
Here E can be, e.g. BH2 , SiH3 , etc. This is used extensively in organic chemistry in hydroboration for antiMarkovnikov addition to alkenes. [Organikers oxidize off the BH2 with H2 O2 to make alcohols, and with
peroxybenzoic acid to make carboxylic acids.] A variation on this reaction is used industrially to produce ethyl
compounds by direct reaction of the metal and ethene in the presence of hydrogen gas. This is a very important
process for cheap large scale production of main group organometallic compounds.
4.5.3. Properties of organometallics
•
•
•
•
4.6
They tend to oxidize easily. Many are flammable in air! Consequently they are reducing agents.
They have nucleophilic character. The more electropositive
the metal, the more carbanionic the organic group is. Thus
alkyl lithium reagents and Grignard reagents are usually
reacting as nucleophiles. In the Chemistry 3810 laboratory a
methyl grignard (CH3 MgBr) is used with S=PCl3 in a reaction
where it acts both nucleophile and reducing agent!
Electron deficient organometallics are powerful Lewis acids;
they form complexes with Lewis bases, and for this reason,
basic solvents such as ethers are extensively used to stabilize
reactive organometallic reagents: case in point, Grignard
reagents are always made in ethers. An exception is n-butyl
lithium, which is prepared and reacted in hexane solution.
Amines, pyridine and Me2 S are all used as stabilizing agents
for reactive organometallics.
The extremely varied reactions of methyl lithium are depicted
in the radial diagram at right.
Structure and Bonding in Main Group Organometallic Compounds
The structure of, and bonding in, main group organometallics can best be understood using a delocalized molecular
orbital approach. We will consider several representative examples; it is impossible to cover all the many varied examples
that have been identified by chemists to date.
4.6.1 Structure and bonding of the tetrameric cluster methyl lithium
Methyl lithium, readily available from chemical suppliers as a solution in ether, has been shown to exist in the solid state
using X-ray diffraction and in less-coordinating solvents as the tetrahedral Me4 Li4 cage shown below. There is a central Li4
tetrahedron, and an outer C4 tetrahedron. These two clusers interpenetrate one another, but importantly they both have the
same set of point group symmetries. Another way to think of this is as that of each methyl group pointing out from the center
of one face of the Li4 tetrahedral cluster.
There have been several attempts to describe the bonding in this highly unusual carbon compound. First, we recognize
the classification of methyl lithium as an extremely electron deficient compound. Just as diborane forms from the electron
deficient BH3 , so H3 C–Li is not stable. Rather than dimerized, it aggregates into a tetramer in order to obtain sufficient
bonding character. We note first that the carbon-lithium bond in H3 C–Li is expected to be highly polar. The MO’s of this
fragment are very similar to those of the heteronuclear diatomics we treated earlier. At the AM1 level of theory, we get for
this fragment:
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 67
1a1
1e
2a1
The occupied orbitals shown at the right are 1a 1 , 1e and 2a 1 , all of which are quite polarized towards the C with the
exception of 2a 1 . The empty orbitals are localized strongly on the lithium – thus we have a very polar carbon lithium bond,
as we might expect from the electronegativity differences. What happens in the tetramer? There have been numerous
discussion in the literature about the bonding in this very important cluster compound. A definitive answer to the problem is
provided, however, by Bickelhaupt FM, Hommes NJRV, Gu erra CF, Baerends EJ Organometallics 15 2923-2931 (1996).
Their calculations, using density functional theory lead to the energy-level diagram and fragment orbital topologies shown in
the next figure.
The calculations were done using a
neutral metal cluster and a tetrahedral
array of ligand sp hybrid basis
functions. These results from a highlevel calculation can be mimicked by
using two fragments in HyperChem at
the AM1 level of theory. Be sure to set
each as having a spin multiplicity of 5,
or the energy levels will not look right.
The picture developed at right is
essentially displayed in the four highest
lying MO’s in AM1. The remaning
electrons are involved in fairly
standard C–H bonding orbitals.
As a result of the extensive sp
mixing that occurs for the element
lithium, the energy of the 1a1 SOMO
fragment orbital is remarkably similar
to that of the 2a 1 SOMO of the lithium
cluster. These therefore undergo a
remarkably covalent, low polarity
8c,2e
bond.
The
remaining
interactions between the t2 sets of
orbitals are distinctly polarized, and
this accounts for the high negative
charge on the C atoms and high
positive charge on the lithium atoms of the tetramer. Using AM1, set up as described above, the in-phase overlaps of these
eight fragment orbitals correspond the four highest-lying filled orbitals. The lower orbitals are more extensively involved in
C–H σ bonding. Note that these MO’s correspond to linear combinations of the 2a 1 orbitals of the monomer.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 68
This description of the bonding
in methyl lithium explains the
driving force for tetramer formation
to be the creation of the lo wer a1
MeLi monomer
bonding MO. The bonds are not
activation barrier for reaction of MeLi
strong, and methyl lithium remains
a highly reactive compound. But in
Me4Li 4
the absence of other reagents, they E
effectively make do as best as n
possible by forming this cluster e
structure. This idea can be graphed r
g
using a typical reaction profile for a
y
generic reaction of methyl lithium
as shown at right.
Thus the
products of reactions
monomer
corresponds
to
a
with electrophiles
maximum, and Me4 Li4 to a local
minimum. If a suitable electrophile
is added (see diagram above of the choices available) then the reaction will exceed very exothermically to the final products,
which tend to be electron precise, or at the very least less electron deficient than methyl lithium is itself.
4.6.2 Structure and bonding in alkyl and aryl aluminums
The next example that we will treat in detail is the hydrocarbon derivatives of aluminum. So long as the hydrocarbon
groups are not too bulky, these tend to exist in a diborane-like dimerized structure. This is true, for example, in Al2 Me6 ,
Al2 Ph 2 R4 and in the mixed compounds Al2 Cl2 Me4 and Al2 (OR)2 Me4 . The structures of examples of such compounds are
shown below:
Al2Me6
Al2Cl2Me4
Al2Ph2R4
Note that just as in diborane, the angles at the bridging groups is unusually small (e.g. 75°at the methyl groups and similar
values at the bridging aryl. However the chlorine bridged dimer has a more “normal” bond angle of 91°. This again suggests
substantial Al–Al bonding in these dimerized clusters. We note that the preference for occupation of the bridging position is
given by the series: (Cl, OR > aryl > alkyl). The dimers are weak and are readily
broken if stronger Lewis bases are present.
At the PM3 level of theory, the orbitals of Al2 Me6 are highly reminiscent of
those found in diborane. The energy level diagram at the right corresponds to the
highest 6 filled levels in the methyl bridged dimer. The orbital at –11.81 eV is the
sixth down from the top, but is the major contributor to fragment bonding in the
dimer. This MO uses a sp σ type orbital to form the bonds in the bridge, where H
had to use a 1s orbital. However, otherwise this
MO strongly resembles the 2a1g MO of B2 H6 .
Also note that this molecule has effective D2h
symmetry, so long as the methyl groups are
freely rotating in the molecule. The text tries to
approach bonding in closely related Al2 Ph 2 R4
using a VB 3c,2e bond model. The delocalized
MO method is obviously superior here.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 69
Of course, there are now many additional orbitals
present, due to the need to have bonding between the
C and H atoms of the methyl groups.
This
considerably complicates the MO diagram, and
necessitates a very thorough consideration of the
bonding theory and a willingness to reject those
orbitals that make only a small contribution to the
inter-dimer bonds. For this reason theoreticians will
often replace hydrocarbon groups in calculations for
hydrogen atoms (isolobal analogy). However, it is
fairly easy to divide those orbitals that deal with the
cluster from those that deal with terminal atom
bonding.
The inter-dimer bond is weak and can be broken
by typical Lewis bases (ethers and amines) Another
indication for this weak dimerization is what happens
with large R-groups. For example, the structure of the
tris(2,6-dimethylphenyl) analogue is shown at right. The large 2,6-substituted aromatic rings keep the molecule from
dimerizing, and hence it forms a discrete AlR3 fragment, and is planar at the central Al atom.
4.6.3 Structure and bonding in alkyldislanes and polysilanes
Consider the molecule Me3 SiSiMe3 which is expected to have an element-element single bond.
Using the PM3 method, we calculate the indicated energy levels for the frontier orbitals of this compound. The HOMO and
LUMO are also shown. Noteworthy here is the small size of the HOMO-LUMO gap, less than 8 eV. This is much smaller
that the value calculated by the same method for the all-carbon analog 2,2,3,3-tetramethylbutane at 14.1 eV). A consequence
of this as well as the fact that there is only a single bond between these atoms is that the Si–Si bond is readily cleaved by UV
light.
It is possible to prepare catenated silicon organometallics, e.g.:
Si
Si
Si
Si
Si
Si
Si
Si
Si
Si
Si
These are made by the reductive elimination of halides e.g. from R2 SiCl2 . A suitable reagent is for example Li[C10 H8 ].
How do these polymers compare with polyethylene? First, it is possible to make them with R side groups rather than
just H, because they are not as sterically crowded as the smaller organic polymer. What is much more fundamentally
interesting is that they show chemical and spectroscopic behaviour more like polyalkynes! For example:
• low-lying empty orbitals (UV absorptions, which for C-C bonds do not occur at all, are found to be similar to those
for C=C bonds, with the band maxima at ca. 200 nm
• evidence of electron delocalization
• electrochemically easier to reduce, etc.)
The origin of this effect is related to the basic postulate of MO theory: σ stabilization and σ* destabilization is a function
of the extent of overlap. Si-Si overlap is poorer than C–C, thus there is a smaller HOMO-LUMO gap for the Si-Si bond, as
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 70
shown in the comparative energy-level diagram at the right. Moreover, for longer polysilane chains, an effect is seen on the
orbitals because of cross-overlap resulting in delocalized MO's over the whole chain, as shown in the schematic figure at the
right.
E
N
E
R
G
Y
σ*
σ*
2sp3
2sp3
3
3sp3
3sp
σ
σ
C
Si
The net result is that polysilanes approach semi-conductor properties. This is exactly analogous to the difference in the
properties of diamond and elemental silicon, which has the diamond structure. Some data:
Band gap, kJ mol-1
Resistivity, ;
C
580
106
Si
105
6 104
Ge
58
50
Sn
7
1
This property is exploited in the use of poly(alkylsilanes) as photoresists in the making of circuit paterns on solid silica. A
representation of a straight-chain poly(dimethylsilane) is shown in the line diagram at right. The rubbery material is applied,
and then a mask is placed over it. Intense UV light then cleaves the bonds between the silicon atoms, so that these are easily
rinsed off with solvent, and the remaining space can be sputtered with elemental metal (Al or Cu!) After this the residual
photoresist is rinsed off as well, and additional silicon is laid down from SiH4 .
This tendency towards weaker E–E bonds
O
O
O
O
O
increases down the periodic table. The most extensive O
Si
Si
Si
series of heavier analogues to the hydrocarbons is in
Si
Si
this very family of poly(dimethylsilanes). In order to
obtain thermally stable and non-light-sensitive silicon
polymers, it is necessary to replace half the silicon units with a suitable heteroatoms . The best known examples are the
poly(dimethylsiloxanes), commonly known by their trade name of silicones. Such heterochain polymers are common for the
3rd and heavier period elements. They are typically extremely temperature resistant, and remain flexible over a wide range of
temperature.
4.7
Stereochemistry at silicon
The redistribution reaction of SiCl4 and SiR4 was mentioned above. This is a quite general process, and mechanistic
studies show these to be associative in nature:
rate = k [ SiR3 X ][ X ] ]
This is what is called an SN 2 reaction for
carbon (substitution with inversion):
This reaction is drawn to show inversion Y
of configuration, which is normally observed
for the carbon analog. With silicon, either
retention or inversion is seen. Retention
occurs when Y– is a poor leaving group, e.g.
H– or OH– . This is rationalized by a rearrangement of the intermediate. When the Y
elimination of Y– is a slow process, this has
time to occur, otherwise it does not occur and
inversion occurs.
The alternate mechanism is drawn at right.
The increased importance for this over C is
explained by the greater stability of fivecoordinate Si over C. This is primarily an
effect of larger size.
+
Si
B
-
A
A
X
-
Y
C
Si
A
X
Y
-
+
C
B
B C
B
X
A
A
Si
Si
C
+
-
X
Y
Si
X
B
A
Y
Si
B C
A
X
Y
C
Si
B
X
+
C-
Chemistry 3810 Lecture Notes
4.8
Dr. R. T. Boeré
Page 71
p(π )-p(π ) Multiple Bonds Between Elements from Period 3 and Beyond
Text books of inorganic chemistry not so long ago contained statements such as the following:
"Si, Ge, Sn and Pb do not form p multiple bonds under any circumstances. Thus numerous types of carbon
compounds, such as alkenes, alkynes, ketones, nitriles, etc. have no analogs."
Cotton and Wilkinson, "Advanced Inorganic Chemistry, 1972
Statements such as these have now been completely disproved! The major breakthroughs came in the early years of the
1980's. Continuous new discoveries are being made in this whole area of the chemistry of multiple bonds to "heavy" (3rd
period and beyond) main-group elements. Most of these advances are based on organometallic derivatives of these elements
- this is where the big progress has been made. We will cover aspects of multiple bonds for the Group 14 and 15 elements at
this time. This material belongs just as much to the chapters on those elements, but the textbook has chosen to place them in
the organometallic chapter. So we will deal with them now. This is just one example on how difficult it is to categorize
chemical compounds!
4.8.1. Historical survey
There had been many unsuccessful attempts to prepare analogues of 2nd period double-bonded compounds, many being
made in the early years of the previous century. These claims are demonstrated in the names of these supposed compounds,
e.g.:
H
R
Ph
H
N
O
Known compounds:
H
alkenes
H
Si
Supposed new analogues:
H
R
H
R
ketones
Si
N
azobenzene
Ph
R
Si
P
O
P
H
R
R
H
disilenes
silicones
diphosphenes
But by the mid-thirties, all these claims had been refuted. Instead, those compounds that had been prepared were all shown to
be cyclic oligomers or high polymers, in which there were only single bonds.
R
H
H
R
Si
disilenes
H
R
H
R
R
Si R
Si
R
Si
Si
R
Si
Si
R
silicones
P
diphosphenes
O
R
P
R
Si
R
R
Si
O
R
R
P
R
R
n
n
P
n
n
Si
R
P R
P
O
R
R
P
R
R
nn
Si
O
R
R
Formation of small rings (above), and high polymers (below)
R
R
R
Si
O
R
Si
R
Si
nn
Si
R
O
*
R
n
n
Si
R
R
R
*
Presumed polymers
not well characterized
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 72
4.8.2. The double bond rule
Persistent failure over many decades led chemists of the 40's and 50's to formulate the "classical double bond rule",
which stated that: Elements having a principal quantum number greater than 2 are not able to form (p-π) π bonds with
themselves or with other elements. This truism was rationalized by various arguments, such as:
• Overlap
It was argued that there was very poor sideways overlap between p-p orbitals due to the large size or the 3rd period
and heavier elements
• Diffuseness of orbitals
It was suggested that the higher-quantum-number p orbitals were more diffuse than the 2p orbitals. Hence the
electron density in an overlapping orbital would be lower, leading to weak π bonds.
•
•
Energy of orbitals
It was suggested that 3p, 4p and 5p orbital had too high energies to participate readily in bonding.
Thermochemistry
It was suggested that p-bond energies in heavy double bonds would be too small to prevent polymerization. This
was largely an argument from silence, but went something like this:
energy(kJ mol–1 )
347
619
bond
C-C
C=C
2n
R
H
H
Si
R
H
H
H
R
n
Si
R
R
R
n*
H
H
R
R
619
R
n
n
R
R
2
∆ H = large -ve
R
*
2n
R
Si
Si
H
R
Si
*
Si
energy(kJ mol–1 )
226
???
bond
Si-Si
Si=Si
R
∆ H = -75 kJmol -1
*
R
× 347
Notice that even C=C double bonds are thermodynamically less stable than two corresponding single bonds,
although the difference is small. In fact, C=C double bonds are kinetically stabilized, and the preparation of
polymers, e.g. polyethylene (plastic bags, etc) involves using a catalyst to overcome the kinetic barrier and induce an
exothermic reaction. Nevertheless, if the Si=Si bond were much weaker, the driving force for polymerization could
be so large that it would be hard to prevent, or so the reasoning went...
4.8.3 Transient species indicating double bonds might exist
During the 1960’s and 1970’s evidence for the transient existence of species with heavy mutiple bonds was accumulated.
The general philosophy was to produce the species by a suitable reaction in solution or in the gas phase, and then to look
either at trapped adducts, or at the decomposition, and argue for the existence of the unsaturated compound. Thus
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 73
tetramethyldisilene was generated by a retro-Diels Alder reaction, and then trapped again by a different diene with a higher
stability constant. Such results left open the question, however, as to whether a genuine π-bond had been formed, or whether
perhaps a transient diradical species had been formed.
In the case of phosphorus both double and triple bonded species were postulated, but again the evidence was indirect,
since only the decomposition products were actually isolated. All of this work served to whet the appetite for the actual
isolation of multiply bonded species, but none of it was definite proof.
4.8.4 Silicon – the breakthrough
The key synthetic principle that was applied in the first successful isolation of Si=Si double-bonded compounds was to
provide sufficient steric bulk that the spontaneous polymerization was inhibited, but not so bulky that the reactivity of
the new functional group was completely hindered. The early reports were for almost identical compounds produced by two
US labs (Wisconsin and MIT). Robert West’s group at Wisconsin prepared starting materials with Si–Si single bonds, and
then used a photochemical means to cleave the single bond, as we discussed above. The resulting silenes apparently
dimerized to provide the beautiful yellow disilene. His choice of very bulky group was the mesityl group (2,4,6trimethylphenyl) – we saw this group previously as stabilizing planar AlR3 species. The competing group at MIT used an
active metal reduction of Si–Cl bonds. This resulted in trimers (i.e. trisilacylopropanes). However, these trimers convered to
disilenes under the action of either heat or UV-light.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 74
Are the properties of these compounds consistent with a genuine π-type double bond? The evidence in favour was obtained
first of all from a structure determined single-crystal X-ray diffraction study. The evidence was:
• The co-planarity of the central C2 SiSiC2 core is clearly visible.
• The substituents are eclipsed, despite their very large steric bulk.
• Careful measurement of the Si–Si bond distances gave 2.160 Å for the Wisconsin compound, versus a typical Si–Si
bond of 2.32 Å. This represents an 8% shortening. For comparison, tetraphenylethene, a suitable C=C model
system shows 12% shortening from the hydrogenated analogue. Thus the Si=Si is significantly shortened for a nonpolar bond.
• The final evidence was obtained from a measurement of the Z → E isomerization energy, using a closely-related
unsymmetrical disilene, (Mes)(t Bu)Si=Si(t Bu)(Mes), for which Eactivation = 131 kJ mol–1 . For comparative purposes,
the barrier to inversion of Z to E isomerism of trans-stilbene is only 179 kJ mol–1 . Thus the barrier in the disilene is
significantly higher than expected for such Si–Si bonds.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 75
Note that the absence of the π-component to a double bond will allow for free-rotation about the E–E bond. This was
elegantly proven several year’s ago by the following crystal structures. The first one is of the dianion of Ph 2 C=CPh 2 , in
which occupation of the π* orbital reduced the C–C bond order to 1. It is found to be staggered in the solid state.
Thus the dianion [Ph 2 C–CPh 2 ]2– is found in a crystal structure
of one of its salts to be almost completely twisted to the “eclipsed”
form, as might be expected based on steric hindrance between the
large phenyl groups (note that the hydrogen’s are not shown in this
diagram because they are thought to obscure the salient details of
the structure.
The second example is that of a dication of
a different alkene, {Me2 N}2 C=C{Me2 N}2 . Here
the electrons have been removed from the π
orbital, again allowing for essentially free
rotation about the now single C–C bond in
[{Me2 N}2 C=C{Me2 N}2 ]2+.
Other properties of the disilenes that are
consistent with the claim that they have a
genuine σ + π double bond are as follows:
• Their colour is typically yellow or orange,
consistent with a λmax ~ 400 nm.
• Their electrochemical reduction and
oxidation are more facile than similarly
substituted alkenes. Thus Ered = –2.12 V,
while Eox = +0.38 V. Both of these values
are considerably smaller than for the
alkenes.
• Detailed calculations suggest that both the
HOMO and the LUMO of this compound
should be considerably more reactive than
that of alkenes. This suggests that the
compounds will be reactive to both
electrophiles and nucleophiles.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 76
A direct output from the Wisconsin group’s structure is
shown in the graphic at right. The circles represent the volume
of space within which most of the electron density around the
nucleus of these atoms vibrates in the crystal lattice. Thus,
although Si is the largest atom in the structure, it has some of the
smallest atom circles, because it is a heavy element and these
often vibrate the least.
The distinctive yellow colour of these disilenes imply that
they absorb visible light in the blue region of the spectrum.
Molecular orbital studies of this type of molecule has shown the
following results:
The HOMO-LUMO gap in a disilene is indeed much smaller than that in a normal alkene. The band gap is half as high as in
the C=C bond case! But with this kinetic stabilization the product lasted for some time. The following orbitals are generated
from a PM3 calculation. They are the HOMO and the LUMO of a model disilene.
←HOMO
-
LUMO→
4.8.5 Stable disilenes are still reactive molecules
Alkenes are among the most important hydrocarbons, because the double bond is classified as a functional group in
organic chemistry, i.e. a region of heightened reactivity compared to the very low reactivity of saturated hydrocarbon
molecules. This reactivity is due to the high-lying HOMO and low-lying empty LUMO, resulting in both addition and
donation reactions occurring from these two orbitals.
An alkene becomes less reactive with increased subsitution of the sp 2 carbon atoms – eventually the double bond gets
too greatly protected – shielded – so that it cannot donate or receive electrons (especially if these come attached to an atom!)
What then for diselenes? Have we “stericaly protected” them to the point where they cannot react at all? Fortunately, this is
not necessarily the case. Indeed, dimerization and polymerization is a reaction with oneself. If instead these disilenes react
cleanly with reagents more reactive than itself, then a proper reactivity may be expected for this compound. Indeed, the latter
is the case!
The following diagram includes a reaction pattern for West’s original tetra(mesityl)disilene. As can be seen at a gla nce,
most of the common reactions of alkenes, such as addition of halogen or HX apply to the disilene. In addition, there are
some reactions that do not occur with alkenes, but do with the more reactive disilenes.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 77
CH3
Mes Cl
H3C
CH3
Si
Mes
Mes H
Mes
Si
Si
Cl Mes
Mes
Mes
Si
Cl Mes
R = Mes (for mesityl)
Mes
Mes
Si
Si
O
O
HCl
Cl2
Mes
Mes
O
O
Ph
Mes
Ph
Mes
Si
Si
Mes
Ph
Ph
Mes
Mes
Si
O
Mes
Mes
Mes
CMe2
Mes
Si
Si
Ph
ROH
Mes
O
Si
Si
Ph
Mes
Mes
Mes
O2
Me2CO
Ph
Mes
Mes
R = H, alkyl
H Mes
Mes
Si
O
Si
Mes
Ph
Mes OR
Si
Mes
4.8.6 Extensions of kinetically stabilized double bonds to germanium and tin
Since the first reports by West and Masumune, similar ideas have been used to try and make even heavier elements with
stable double bonds.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 78
The following data has been obtained from solid-state crystallographic data. The terms are defined in the sketch above.
Note that all these compounds have considerable bond shortening. MO calculations have been peformed to try and
understand this unusual bond pyramidalizaton.
There have been two main attempts to try and rationalize these large deviations from co-planarity. The first is a delocalized
MO approach (ie. What we are familiar with). The second is that of a VB model put forward by some British scientists as kl
a possible explanation both for the pyramidalization of, especially, the tin compounds. This model considers two “carbenelike” fragments, each undergoing internal Lewis Acid-Base interactions between the filled “sp 2 ” and the empty p orbitals.
This model is most useful in that it explains the existence of free tin-based R2 Sn “carbene-like” fragments.
The MO calculations show a much more systematic approach to these systems. Note how the data is presented: the energies
of bending about the θ-angle are ploted for all the Group 14 elements. Silicon is a lot more flexible than the carbon analogue.
This trend continues until for both Ge and Sn, the flat form is no longer the global minimum, and the structure will be bent.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 79
The origin of this greater “floppiness” of the heavier Group 14 element compounds can be found from molecular orbital
calculations. Consider the following energy level diagram, which compares planar and pyramidalized forms of both
compounds.
At root, it is the much higher-lying nature of the weak π-bond in the Ge =Ge compound that accounts for the much greater
ease of second-order mixing upon symmetry.
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 80
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
4.8.7 Extension of the principle to the Group 15 elements
Stable Group 15 p(π)–p(π) bonds have been known since the early 1980’s.
Page 81
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Some crystal structures of diphosphenes and diarsenes
Page 82
Chemistry 3810 Lecture Notes
The Electronic Structure of R–P=P–R
Dr. R. T. Boeré
Page 83
Chemistry 3810 Lecture Notes
Dr. R. T. Boeré
Page 84
4.8.8 Reactivity of diphosphenes
Again the reactivity of these compounds has not been blocked – just their propensity to oligomerize and polymerize.
The following chart allows us to see some typical reactions, almost all of which are in some ways unique!
4.8.9 Conclusions
1.
2.
3.
4.
5.
Thermodynaically stable p(π) - p(π) bonds exist both between the heavy elements and between heavy and light elements.
Such bonds must be kinetically stabilized against polymerization with bulky ligands.
Despite the bulky groups, the bonds are still reactive, and hence interesting. There are now many derivatives of such
systems due to a wide variety of reactions at the multiple-bond centers.
The homologous compounds of elements down a group show a continuum in behaviour and properties.
Molecular orbital theory is well suited to showing such continuous relationships and why observed changes in properties
within groups occur.