James Richard Fromm
The Lewis structure of a molecule consists of the Lewis structures of its component atoms arranged in an appropriate configuration, with shared electron pairs located between the atoms which share them. In most cases a two-dimensional representation is used so that the structure can be written on paper, although the actual structure can only be properly represented in three dimensions. Shared electron pairs can equally well be shown as paired dots (:) or as single dashes (-). The dashes used in Lewis structures, however, carry the more precise meaning of electron pair used in bonding rather than the older and more elastic meaning of valence.
All of the valence electrons must be shown in the Lewis structure of a molecule even though not all of them are used in forming bonds in that particular molecule. A single unpaired electron is shown as a single dot. A pair of electrons not used in bonding (a lone pair) is shown as a pair of dots. When a single pair of electrons are shared between two atoms in a molecule, they form a single bond between the atoms. It is also possible for two atoms in a molecule to share more than one pair of electrons between them. When two pairs of electrons are shared between two atoms in a molecule, they form a double bond between the atoms. Double bonds are much stronger than single bonds, although they are not usually exactly twice as strong as a single bond. When three pairs of electrons are shared between two atoms in a molecule, they form a triple bond between the atoms. Triple bonds are even stronger than double bonds, although they are not exactly equal in strength to three single bonds. The triple bond in molecular nitrogen is the reason molecular nitrogen is chemically unreactive; few reactions can supply enough energy to break it. Several examples of double and triple bonding are shown in the Figure above. The enthalpies associated with single and multiple bonds were taken up in a different section.