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The structure of 3-methyl-2-hexenoic acid.

Boron trihalides: a special case

The three lighter boron trihalides, BX 3 (X = F, Cl, Br) form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicity of the adduct-forming reaction:

This trend is opposite to that expected based upon the electronegativity of the halogens. The best explanation of this trend takes into account the extent of π-donation that occurs between the filled lone pair orbital on the halogens and the empty p-orbital on the planar boron ( [link] ). As such, the greater the π-bonding the more stable the planar BX 3 configuration as opposed to the pyramidalization of the BX 3 moiety upon formation of a Lewis acid-base complex, [link] .

Schematic representation of π-donation from filled halogen p-orbitals into empty p-orbital in the halides BX 3 .

The criteria for evaluating the relative strength of π-bonding are not clear, however, one suggestion is that the F atom is small compared to the Cl atom, and the lone pair electron in p z of F is readily and easily donated and overlapped to empty p z orbital of boron ( [link] a). In contrast, the overlap for the large (diffuse) p-orbitals on the chlorine is poorer ( [link] b). As a result, the π-donation of F is greater than that of Cl. Interestingly, as may be seen from [link] , any difference in B-X bond length does not follow the expected trend associated with shortening of the B-X bond with stronger π-bonding. In fact the B-Br distance is the most shortened as compared to that expected from the covalent radii ( [link] ).

The B-X bond distances in the boron trihalides, BX 3 , as compared to the sum of the covalent radii. a Covalent radius of B = 0.84(3) Å.
Compound B-X (Å) X covalent radius (Å) Sum of covalent radii (Å) a Δ (Å)
BF 3 1.313 0.57(3) 1.41 0.097
BCl 3 1.75 1.02(4) 1.86 0.11
BBr 3 1.898 1.20(3) 2.04 0.142
BI 3 2.125 1.39(3) 2.23 0.105

At the simplest level the requirements for bonding to occur based upon the molecular or atomic orbital are:

  • Directional relationship of the orbitals.
  • Relative symmetry of the orbitals.
  • Relative energy of the orbitals.
  • Extent of orbital overlap

In the case of the boron trihalides, the direction (parallel) and symmetry (p-orbitals) are the same, and the only significant difference will be the relative energy of the donor orbitals (i.e., the lone pair on the halogen) and the extent of the overlap. The latter will be dependant on the B-X bond length (the shorter the bond the greater the potential overlap) and the diffusion of the orbitals (the less diffuse the orbitals the better the overlap). Both of these factors will benefit B-F over B-Cl and B-Br. Thus, the extent of potential overlap would follow the order: [link] . Despite these considerations, it is still unclear of the exact details of the rationalization of the low Lewis basicity of BF 3 as compared to BCl 3 and BBr 3 .

Anionic halides

The trihalides all form Lewis acid-base complexes with halide anions, [link] , and as such salts of BF 4 - , AlCl 4 - , GaCl 4 - , and InCl 4 - are common.

In the case of gallium the Ga 2 Cl 7 - anion ( [link] ) is formed from the equilibrium:

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Source:  OpenStax, Chemistry of the main group elements. OpenStax CNX. Aug 20, 2010 Download for free at http://cnx.org/content/col11124/1.25
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