formal charge

When you split a shared bond perfectly down the middle and ask 'is this atom now holding more or fewer electrons than it brought to the party?', the answer is the formal charge. It is an accounting trick, not a real measured charge: a way of keeping track of electrons so you can tell where the pluses and minuses sit on a Lewis structure and decide which structure is most sensible. Think of it as checking each atom's electron 'bank balance' against what it started with.

The recipe is simple. For each atom, take its number of valence electrons (from the periodic table group), then subtract the electrons it 'owns' in the molecule: count all of its lone-pair electrons as its own, plus half of every bonding electron pair it shares. Formal charge equals valence electrons minus lone-pair electrons minus half the bonding electrons. A neutral atom that owns exactly what it brought has formal charge zero; if it owns one fewer, the charge is +1; one more, -1. The formal charges in a molecule must add up to its overall charge.

Formal charge matters for two everyday reasons. First, it tells you where to draw the + or - on ions and reactive intermediates — the positive carbon of a carbocation, the negative oxygen of an alkoxide. Second, when several Lewis structures or resonance forms are possible, the best is usually the one with formal charges closest to zero and any negative charge on the most electronegative atom. The crucial caveat: formal charge assumes electrons are shared perfectly evenly, which is a fiction — it is a label for bookkeeping, not the true partial charge, which electronegativity governs.