0.6 Molecular geometry and electron domain theory  (Page 4/6)

As an example of a molecule with an atom with less than an octet of valence shell electrons, we consider borontrichloride, $B{\mathrm{Cl}}_3$ . The geometry of $B{\mathrm{Cl}}_3$ is also given in : it is trigonal planar , with all four atoms lying in the same plane, and all Cl-B-Cl bond angles equal to 120°. The three Clatoms form an equilateral triangle. The Boron atom has only three pairs of valence shell electrons in $B{\mathrm{Cl}}_3$ . In applying Electron Domain theory to understand this geometry, wemust place three points on the surface of a sphere with maximum distance between points. We find that the three points form anequilateral triangle in a plane with the center of the sphere, so Electron Domain is again in accord with the observedgeometry.

We conclude from these predictions and observations that the Electron Domain model is a reasonablyaccurate way to understand molecular geometries, even in molecules which violate the octet rule.

Observation 2: molecules with double or triple bonds

In each of the molecules considered up to this point, the electron pairs are either in single bonds or in lonepairs. In current form, the Electron Domain model does not account for the observed geometry of $C_2{H}_4$ , in which each H-C-H bond angle is 116.6° and each H-C-C bondangle is 121.7° and all six atoms lie in the same plane. Each carbon atom in this molecule is surrounded by four pairs ofelectrons, all of which are involved in bonding, i.e. there are no lone pairs. However, the arrangement of these electron pairs, and thus the bonded atoms,about each carbon is not even approximately tetrahedral. Rather, the H-C-H and H-C-C bond angles are much closer to 120°, theangle which would be expected if three electron pairs were separated in the optimal arrangement, as just discussed for $B{\mathrm{Cl}}_3$ .

This observed geometry can be understood by re-examining the Lewis structure. Recall that, although there arefour electron pairs about each carbon atom, two of these pairs form a double bond between the carbon atoms. It is tempting to assumethat these four electron pairs are forced apart to form a tetrahedron as in previous molecules. However, if this were thiscase, the two pairs involved in the double bond would be separated by an angle of 109.5° which would make it impossible forboth pairs to be localized between the carbon atoms. To preserve the double bond, we must assume that the two electron pairs in thedouble bond remain in the same vicinity. Given this assumption, separating the three independent groups of electron pairs about a carbon atom produces an expectation that all three pairs should liein the same plane as the carbon atom, separated by 120° angles. This agrees very closely with the observed bond angles. Weconclude that the our model can be extended to understanding the geometries of molecules with double (or triple) bonds by treatingthe multiple bond as two electron pairs confined to a single domain . It is for this reason that we refer to the model as Electron Domain theory.

Applied in this form, Electron Domain theory can help us understand the linear geometry of $C{O}_2$ . Again, there are four electron pairs in the valence shell of thecarbon atom, but these are grouped into only two domains of two electron pairs each, corresponding to the two C=O double bonds.Minimizing the repulsion between these two domains forces the oxygen atoms to directly opposite sides of the carbon, producing alinear molecule. Similar reasoning using Electron Domain theory as applied to triple bonds correctly predicts that acetylene, $HCCH$ , is a linear molecule. If the electron pairs in the triple bond aretreated as a single domain, then each carbon atom has only two domains each. Forcing these domains to opposite sides from oneanother accurately predicts 180° H-C-C bond angles.