In chemistry, a lone pair is a valence electron pair without bonding or sharing with other atoms. They are found in the outermost electron shell of an atom, so lone pairs are a subset of a molecule's valence electrons. They can be identified by examining the outermost energy level of an atom—lone electron pairs consist of paired electrons as opposed to single electrons, which may appear if the atomic orbital is not full. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone electrons plus the number of bonding electrons equal the total number of valence electrons in a compound.
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A single lone pair can be found with atoms in the nitrogen group such as nitrogen in ammonia, two lone pairs can be found with atoms in the chalcogen group such as oxygen in water and the halogens can carry three lone pairs such as in hydrochloric acid.
In simple valence bond theory the oxygen atom in water is sp3 hybridized with the lone pairs in equal sp3 orbitals. The H–O–H bond angle is with 104.5° less than the 109° predicted for a tetrahedral angle and this can be explained by a repulsive interaction between the lone pairs [1][2][3]
The pairs often exhibit a negative polar character with their high charge density and are located closer to the atomic nucleus on average compared to the bonding pair of electrons. The presence of a lone pair decreases the bond angle between the bonding pair of electrons, due to their high electric charge which causes great repulsion between the electrons. They are also used in the formation of a dative bond. For example, the creation of the hydronium (H3O+) ion occurs when acids are dissolved in water and is due to the oxygen atom donating a lone pair to the hydrogen ion.
This can be seen more clearly when looked at in two more common molecules. For example, in carbon dioxide (CO2), the oxygen atoms are on opposite sides of the carbon, whereas in water (H2O) there is an angle between the hydrogen atoms of 104.5o. Due to the repulsive force of the oxygen atom's lone pairs, the hydrogens are pushed further away, to a point where the forces of all electrons on the hydrogen atom are in equilibrium. This is an illustration of the VSEPR theory.
A stereochemically active lone pair is also expected for divalent lead and tin ions due to their formal electronic configuration of ns2. In the solid state this results in the distorted metal coordination observed in the litharge structure adopted by both PbO and SnO. The formation of these heavy metal ns2 lone pairs which was previously attributed to intra-atomic hybridization of the metal s and p states[4] has recently been shown to have a strong anion dependence.[5] This dependence on the electronic states of the anion can explain why some divalent lead and tin materials such as PbS and SnTe show no stereochemical evidence of the lone pair and adopt the symmetric rocksalt crystal structure[6],.[7]
In molecular systems the lone pair can also result in a distortion in the coordination of ligands around the metal ion. The lead lone pair effect can be observed in supramolecular complexes of Lead(II) nitrate and in 2007 a study [8] linked the lone pair to lead poisoning. Lead ions in the human metabolism replace native metallic ions in several key proteins, for example: zinc cations in the ALAD protein, which is also known as Porphobilinogen synthase. This seems to be the molecular basis of "lead poisoning", or "saturnism" ("plumbism"). Computational experiments reveal that although the coordination number does not change upon substitution in calcium-binding proteins, the introduction of lead distorts the way the ligands organize themselves to accommodate such an emerging lone pair: consequently, these proteins are perturbed. This lone-pair effect becomes dramatic for zinc-binding proteins, such as the above-mentioned porphobilinogen synthase, as the natural substrate cannot bind anymore: in those cases the protein is inhibited.
Group 14 lone pairs manifest themselves in triple bonds as well. The familiar alkynes have bond order 3 with 180° bond angles (A) but going down the row germanium to germanium formal triple bonds have an effective bond order 2 with one lone pair (B) and trans-bent geometries. In lead the bond order is even 1 with lone pairs for each lead atom (C). In the organogermanium compound D, the bond order is also 1 with complexation of the acidic isonitrile groups based on interaction with germaniums empty 4p orbital [9]