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Stereochemical activity
In compounds with a central atom VSEPR theory can be applied and the geometry about the central atom is determined by the repulsion between bonding and lone pairs. By implication in VSEPR all lone pairs are stereochemically active, with the recognised exception of the transition metals where the lone pairs occupy d orbitals, are not in the valence shell but in the outer sub-shell of the underlying core. . A simple example of the sterochemical activity of lone pairs is in the binary hydrogen compounds of the elements of groups 15 and 16. The group 15 compounds ( NH3, PH3, AsH3, SbH3, BiH3) are all pyramidal, and the group 16 compounds (H2O, H2S, H2Se, H2Te are all bent.

In many solids the effect of stereochemically active lone pairs is apparent, for example the structure of black phosphorus, arsenic, antimony and bismuth which have puckered layer structures clearly shows their infuence where the P, As, Sb and Bi atoms in the layers have trigonal coordination. Other examples include the black α-SnO, tin(II) oxide and red tetragonal PbO, lead (II) oxide. Not all lone pairs are stereochemically active. Transition metals have already been mentioned. There also are many instances of lone pairs that do not affect the geometry, for example TeCl62- is a regular octahedron rather than distorted, solid lead(II) sulfide, PbS, adopts the regular rock salt structure contrasting with tin(II) sulfide, SnS, that has a layer structure similar to black phosphorus. . Explanations when a lone pair is not sterochemically active include ligand crowding which suppresses lone pair inflence in  BrF6- cbut in If6-, as I is larger the crowding is less so the lone pair can exert an influence. A cautionary not e ius that in the solid state the couter ion can infleunce an anions shape.

Hydrogen bonding
Lone pairs are often quoted as being the reason that hydrogen bonds occur, in particlar the reason why ammonia forms accepts only one H wheras water accepts two.

Lewis basicity
Lewis’s original definition was that a base had a lone pair of electrons which may be used to complete the stable group of another atom. Many molecules in groups 14, 15, 16 and 17 can act as Lewis bases.

Group 14
Most of the divalent Si, Ge, Sn compounds have a lone pair and can act as donors One interesting example is F3BSnCl2NMe3 where the Sn(II) atom is acting both as a Lewis acid (acceptor) and Lewis base (donor) The cyanide ion when it is unidentate uses the lone pair on the carbon atom.

Group 15
Well known examples incude ammonia, pyridine, phosphinesand  amines

Group 16
Well known examples are the hydroxyl ion, S2-, SO2 which can bond via both the S and O atoms.

Group 17
The halide ions all can use a lone pairs as can covalently bonded halogens for example AlCl3 dimerises with bridging chlorine atoms.

Shape and orientation of lone pairs
In VSEPR theory the lone pairs are described as domains which are larger and more spread out than those occupied by the bonding pairs. In valence bond theory the lone pairs occupy either hybrid orbitals or atomic orbitals. A typical representation of the lone pair on nitrogen in ammonia is an sp3 orbital lobe containing two dots and in water two sp3 orbital lobes each containing two dots, the so called “rabbit’s ears”. This view of water has been criticised. The photoelectron spectrum of water shows four bands whereas the simple sp3 model would predict just two. The four bands are explained by molecular orbital theory which predicts just one lone pair.

Lone pairs and dipole moments
Lone pairs can make a contribution to a molecules dipole moment. NH3 has a dipole moment of 1.47 D. As the electronegativity of nitrogen (3.04) is greater than that of hydrogen (2.2) the result is that the N-H bonds are polar with a net negative charge on the nitrogen atom and a smaller net positive charge on the hydrogen atoms. There is also a dipole associated with the lone pair and this reinforces the contribution made by the polar covalent N-H bonds to ammonia's dipole moment. In contrast to NH3, NF3 has a much lower dipole moment of 0.24 D. Fluorine is more electronegative than nitrogen and the polarity of the N-F bonds is opposite to that of the  N-H bonds in ammonia, so that the dipole due to the lone pair opposes the N-F bond dipoles, resulting in a low molecular dipole moment.