Lone Pair
Valence electrons of atoms in covalent compounds may be shared between atoms to form covalent bonds or they may be unshared. Pairs of electrons that are not shared in covalent bonds are lone pairs. Lone pairs of electrons play a critical role in determining the three-dimensional shape of molecules. A molecule's shape, in turn, determines many of its physical properties and its chemical reactivity.
A very useful method for understanding and predicting molecular shapes that relies upon the role of the lone pair is called the Valence Shell Electron Pair Repulsion (VSEPR) theory. This method has been developed extensively by Professor R. J. Gillespie of MacMaster University. It correctly accounts for the structures of most covalent compounds of elements other than transition metal complexes.
The basic tenet of the VSEPR theory is that atoms in a covalent molecule are arranged to minimize the repulsions of valence shell electrons of the bonded atoms. All valence electrons must be considered. As a result the shape of the ammonia molecule, NH3 is not flat triangular as it would be from considering only the three hydrogen atom bonds. Instead, it is pyramidal, with the three hydrogen atom bonds pushed away from a plane, and towards each other, by the requirement for space of the lone pair of electrons on the nitrogen atom.
The role of the lone pair in VSEPR theory is even more important than a bonding pair of electrons. The repulsions of lone pair of electrons for other electrons are considered to be stronger than the repulsions of bonded electrons. Because bonded electrons are strongly attracted to the nuclei of the two bonded atoms, they are localized between the two atoms. Lone pair electrons are more diffuse and extend into a larger volume. Therefore, they can come closer to other electrons and repel them more strongly. The order of repulsive interactions is: lone pair-lone pair, lone pair-bonding pair, bonding pair-bonding pair. Not only do the lone pair electrons lead to a certain type of shape, pyramidal in the case of ammonia, but they also affect bond angles. If the repulsions of lone pair electrons and bonding electrons were the same the hydrogen-nitrogen-hydrogen bond angles would be those of a perfect tetrahedron, 109°. Instead, the lone pair electrons push the bonding electrons closer together, resulting in a bond angle of approximately 107° in ammonia. In water, Where there are two lone pairs of electrons on the oxygen atom to repel the bonding electrons of each O-H bond, the H-O-H angle is squeezed to 105°.
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