A. Properties of the Elements
Objectives:
By the end of this set of lectures you will be able to use simple electronic and stereoelectronic
arguments for rationalising reactions involving main group elements such as B, Si, S and P.
Specifically, you will:
1)
understand how electron configuration and molecular structure can be used to predict the
mechanisms and outcomes of reactions of organoboranes.
2)
understand the stereoelectronic arguments used to explain how silicon stabilises a beta
positive charge and an alpha negative charge.
3)
have an appreciation of how low lying vacant atomic orbitals can affect the reaction
mechanisms of compounds reacting at or near to period 3 elements.
4)
appreciate how a knowledge of bond strengths and simple concepts such as
electronegativity can be used to rationalise reactions involving main group elements.
A.1 Boron
Position in Periodic Table
Period 2, Group 13 (old group III)
Electron Configuration
2
2
1
1s , 2s , 2p .
Three electrons in the valence shell. Can form up to three covalent bonds (after hybridisation).
The majority of organoboron compounds are tricoordinate in their monomeric state.
An electron count around the central boron atom reveals it has only 6 electrons; two more
electrons would enable it to achieve a "stable" octet of electrons. Borane (BH3) goes some way
to attaining this and full coordination (tetrahedral) by dimerisation; however even in diborane
(B2H6) each boron atom is still electron deficient.
Boranes are Lewis Acidic
A better method by which the boron atom in tricoordinate boron compounds can satisfy the Octet
Rule is for it to behave as a Lewis Acid and accept a pair of electrons from a Lewis base.
BH3 is commercially available as a Lewis Acid-Lewis Base complex in which a THF or DMS
(Me2S) molecule behaves as the Lewis Base component.
Donating a pair of electrons into the unoccupied p AO of the boron atom generates an "ate"
complex. Although the B atom is now coordinatively saturated and has attained its full octet of
valence electrons (electron configuration of Ne), it is now negatively charged. Charged organic
species are generally less stable than neutral compounds.
Thus "ate" complexes, whilst
coordinatively saturated and containing a full complement of valence electrons, are prone to
further reaction.
This usually results in the re-formation of a neutral electron deficient
tricoordinate organoborane. A good example of this is the Matteson rearrangement.
Note that in the above rearrangement it is the Me group - rather than an alkoxide - that migrates
in the 'ate' complex. This is because B−C bonds are weaker than B−O bonds. The migration is
also a stereospecific SN2-like process, proceeding with inversion of configuration.
The Lewis acidity of organoboranes is strongly dependent on the substituents about the boron
atom as these dictate whether the 'empty' p AO is able to accept a pair of electrons.
The order of Lewis acidity for the following boron compounds is:
BF3 > B(OR)3 > B(NR2)3
A number of factors determine the Lewis acidity of the central boron atom:
i) degree of positive charge on the boron. The more electronegative the substituent the greater
the polarisation of the B-X bond and the greater is the positive charge on the B atom. On these
grounds we would expect BF3 to be the most Lewis acidic as the order of electronegativity is:
F > O > N.
ii) Fluorine, oxygen and nitrogen all have lone pairs of electrons which can participate to varying
degrees in pπ-pπ bonding. This will have the effect of filling the boron p AO and rendering the
boron atom less Lewis acidic. Lone pair donating capacity correlates well with electronegativity;
thus the most electronegative fluorine is the poorest electron donor and the least electronegative
nitrogen is the best electron donor. Again we would expect BF3 to be the strongest Lewis acid.
iii) The efficiency of pπ-pπ orbital overlap will also depend on the relative energies and sizes of
the overlapping orbitals. This accounts for the increased Lewis acidity of BX3 on descending the
group 17 (old group VII) elements. Fluorine is in the same period as boron and therefore
provides the most efficient overlap between the p AOs (2p with 2p).
electronegativity effects (order of electronegativity F > Cl > Br).
order of Lewis acidity: BBr3 > BCl3 > BF3
This overrides
Electronegativity of Boron
2.01 (Pauling scale) - more electropositive than carbon; displays metallic character. This, at least
in part, can be used to account for the regioselectivity in hydroboration of unsymmetrical olefins
(steric effects are clearly also very important):
The electropositive nature of the boron atom renders attached hydrogen substituents hydridic in
character. Any build-up of positive charge on the olefin will occur at the more substituted end;
thus the hydridic hydrogen will be transferred to this carbon and the boron will end up on the less
substituted end. Although we discussed the mechanism last semester in CHM3A2-I, we will
examine hydroboration again in more detail later in the course.
Bond Strengths
-1
Bond
Bond Strength (kJ mol )
B−H
389
B−C
372
B−N
446 (in B3N3H3Cl3)
B−O
536
B−F
613
An important point to note is that the B−C bond is weaker than a B−X bond where X is a more
electronegative element. This is because N, O and F can all participate in partial pπ-pπ bonding
i.e. they are not really single bonds.
Summary
1.
Boron is more electropositive than both carbon and hydrogen.
-
This polarises the C−B bond putting a δ charge on the carbon atom and making a
hydrogen substituent hydridic in character.
2.
Tricoordinate boron compounds are Lewis acidic although their Lewis acidity is
modulated by the substituents.
3.
Boron "ate" complexes are coordinatively saturated and the boron atom has a full octet of
electrons. Nevertheless these complexes possess a formal negative charge on the
boron atom and therefore are prone to further reaction.
4.
Partial pπ-pπ bonding makes B−X (X = F, O, N) bonds much stronger than B−C and B−H
bonds.