Covalent Bond and Lewis Structures

A covalent bond forms when two atoms — typically two non-metals of comparable electronegativity — share one or more pairs of electrons, rather than one atom transferring electrons to the other outright. Each atom contributes electrons to the shared pair(s), and both nuclei are simultaneously attracted to the shared electron density lying between them, which is what holds the atoms together.

Lewis (electron-dot) structures are the standard way of representing covalent bonding on paper. Constructing one systematically follows a short sequence of steps:

1. Count the total number of valence electrons available — sum the valence electrons of every atom present, and for a polyatomic ion, add one electron for each unit of negative charge or subtract one for each unit of positive charge.
2. Identify a reasonable central atom (usually the least electronegative atom present, other than hydrogen, which is never central) and arrange the remaining atoms around it.
3. Place a single bonding pair between the central atom and each surrounding atom, then distribute the remaining electrons as lone pairs, starting with the outer atoms, so that every atom (other than recognised exceptions) attains an octet.
4. If the central atom still lacks an octet after this, convert one or more lone pairs on the outer atoms into additional shared pairs, creating double or triple bonds as needed.

Applying this to carbon dioxide, $\text{CO}_2$ ($4 + 6 + 6 = 16$ valence electrons in total): carbon sits in the centre, and each oxygen ends up doubly bonded to carbon, each oxygen also carrying two lone pairs. This gives the familiar structure $\text{O}=\text{C}=\text{O}$, with carbon surrounded by two double bonds (an octet, since each double bond counts as four electrons around carbon) and each oxygen surrounded by two lone pairs plus one double bond (also a complete octet).

Real molecules and ions often admit more than one Lewis structure that satisfies the octet rule; formal charge is the bookkeeping tool used to judge which structure is the most realistic one, or how charge is distributed across an ion. The formal charge on any atom is calculated as

$FC = (\text{valence electrons of the free atom}) - (\text{non-bonding electrons on that atom}) - \tfrac{1}{2}(\text{bonding electrons on that atom})$ …