Molecular shape and the VSEPR model Valence-shell electron-pair repulsion theory

The shapes of molecules containing a central p-block atom tend to be controlled by the number of electrons in the valence shell of the central atom. The valence-shell electron-pair repulsion (VSEPR) theory provides a simple model for predicting the shapes of such species. The model combines original ideas of Sidgwick and Powell with extensions developed by Nyholm and Gillespie, and may be summarized as follows:

•  Each valence shell electron pair of the central atom E in a molecule EX_n containing E–X single bonds is stereochemically significant, and repulsions between them determine the molecular shape.
•  Electron–electron repulsions decrease in the sequence: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
•  Where the central atom E is involved in multiple bond formation to atoms X, electron–electron repulsions decrease in the order: triple bond–single bond > double bond – single bond > single bond–single bond.
•  Repulsions between the bonding pairs in EX_n depend on the difference between the electronegativities of E and X; electron – electron repulsions are less the more the E^_X bonding electron density is drawn away from the central atom E.

The VSEPR theory works best for simple halides of the pblock elements, but may also be applied to species with other substituents. However, the model does not take steric factors (i.e. the relative sizes of substituents) into account. In a molecule EX_n, there is a minimum energy arrangement for a given number of electron pairs. In BeCl₂ (Be, group 2), repulsions between the two pairs of electrons in the valence shell of Be are minimized if the Cl^_Be^_Cl unit is linear. In BCl₃ (B, group 13), electron–electron repulsions are minimized if a trigonal planar arrangement of electron pairs (and thus Cl atoms) is adopted. The structures in the lefthand column of Figure 1.1 represent the minimum energy structures for EXn molecules for n = 2–8 and in which there are no lone pairs of electrons associated with E. Table 1.1 gives further representations of these structures, along with their ideal bond angles. Ideal bond angles may be expected

Fig. 1.1 (a) A simplified orbital interaction diagram for CO which allows for the effects of some orbital mixing. The labels 1σ, 2σ . . . rather than σ(2s) . . . are used because some orbitals contain both s and p character. (b) A more rigorous (but still  qualitative) orbital interaction diagram for CO.

Table 1.1 ‘Parent’ shapes for EXn molecules (n = 2–8).

when all the X substituents are identical, but in, for example, BF₂Cl (1.18) some distortion occurs because Cl is larger than F, and the shape is only approximately trigonal planar.

The presence of lone pairs is taken into account using the guidelines above and the ‘parent structures’ in Figure 1.29.

In H₂O (1.19), repulsions between the two bonding pairs and  two lone pairs of electrons lead to a tetrahedral arrangement but owing to the inequalities between the lone pair–lone pair, lone pair–bonding pair and bonding pair– bonding pair interactions, distortion from an ideal arrangement arises and this is consistent with the observed H^_O^_H bond angle of 104.58.e

مواضيع ذات صلة

Ligand field theory

Crystal field theory

Limitations of VSEPR theory

MO theory: heteronuclear diatomic molecules

The octet rule

The bonding in He2, Li2 and Be2

Molecular orbital theory applied to the bonding in H2

 Homonuclear diatomic molecules: molecular orbital (MO) theory

The valence bond (VB) model applied to F2, O2 and N2

The valence bond (VB) model of bonding in H2

Covalent bond distance, covalent radius and van der Waals radius

Bonding models: an introduction