Valence Shell Electron Pair Repulsion (VSEPR) Theory

Valence Shell Electron Pair Repulsion (Vsepr) Theory

      In 1957 Gillespie and Nyhom gave this theory to predict and explain molecular shapes and bond angles more exactly. The theory was developed extensively by Gillespie as the Valence Shell Electron Pair Repulsion (VSEPR) theory. This may be summarized as:

1.    The shape of the molecule is determined by repulsions between all of the electron pairs present in the valence shell.

2.   A lone pair of electrons takes up more space round the central atom than a bond pair, since the lone pair is attracted to one nucleus whilst the bond pair is shared by two nuclei.

3.    It follows that repulsion between two lone pairs is greater than repulsion between a lone pair and a bond pair, which in turn is greater than the repulsion between two bond pairs.

4.   Thus the presence of lone pairs on the central atom causes slight distortion of the bond angles from the ideal shape.

5.   If the angle between a lone pair, the central atom and a bond pair is increased, it follows that the actual bond angles between the atoms must be decreased.

6.   The order of repulsion between lone pairs and bond pairs of electrons follows the order as:

   Lone pair - lone pair repulsion > lone pair – bond pair repulsion > bond pair – bond pair repulsion.

7.   The magnitude of repulsions between bonding pairs of electrons depends on the electronegativity difference between the central atom and the other atoms.

   8. Double bonds cause more repulsion than single bonds, and triple bonds cause more repulsion than a double bond.

     Effect of Lone Pairs: Molecules with four electron pairs in their outer shell are based on a tetrahedron. In CH4 there are four bonding pairs of electrons in the outer shell of the C atom, and the structure is a regular tetrahedron with bond angle H – C – H of 109°28’. In NH3 and N atom has four electron pairs in the outer shell, made up of three bond pairs and one lone pair.

Because of the lone pair, the bond angle H – N – H is reduced from the theoretical tetrahedral angle of 109°28’ to 107°28’.

 In H2O the O atom has four electron pairs in the outer shell.

The shape of the H2O molecule is based on a tetrahedron with two corners occupied by bond pairs and the other two corners occupied by lone pairs.

The presence of two lone pairs reduces the bond angle further to 104°27’.

       In a similar way, SF6 has six bond pairs in the outer shell and is a regular octahedron with bond angles of exactly 90°. 

In BrF5, the Br also has six outer pairs of electrons, made up of five bond pairs and one lone pair. 

The lone pair reduces the bond angles to 84°30’. 

Whilst it might be expected that two lone pairs would distort the bond angles in an octahedral as in XeF4 but it isnot so.

 Actual bond angle is 90°, reason being that the lone pairs are trans to each other in the octahedron, and hence the atoms have a regular square planar arrangement.

       Molecules with five pairs of electrons are all based on a trigonal bipyramid. 

Lone pairs distort the structures as before.

The lone pairs always occupy the equatorial positions (in an triangle), rather than the axial positions (up and down).

Thus in I3– ion, the central I atom has five electron pairs in the outer shell, made of two bond pairs and three lone pairs. 

The lone pairs occupy all three equatorial positions and the three atoms occupy the top, middle, and bottom positions in the trigonal bipyramid, thus giving a linear arrangement with a bond angle of exactly 180°.

       Effect of Electronegativity: NF3 and NH3 both have structures based on a tetrahedron with one corner occupied by a lone pair. 

The high electronegativity of F push the bonding electrons further away from N than in NH3. 

Hence the lone pair in NF3 causes a greater distortion from tetrahedral and gives a F – N – F bond angle of 102°30’, compared with 107°48’ in NH3. 

The same effect is found in H2O (bond angle 104°27’) and F2O (bond angle 102°).