Formal charge — Foundations | Organic Chemistry

Formal charge is bookkeeping, and like all good bookkeeping it is boring until it catches an error. Its job is to tell you which atom in a structure is carrying a charge, whether a structure you have drawn is plausible at all, and which of several competing structures is the one that actually describes the molecule. Get fluent with it now; you will be doing it under time pressure for the rest of the course.

The formula

An atom's formal charge compares the electrons it "owns" in a structure with the electrons it would have as a free neutral atom. The convention is that an atom owns all of its lone-pair electrons and exactly half of each bonding pair — one electron per bond.

FC = (valence electrons) − (lone-pair electrons) − (number of bonds)

Two details govern the whole calculation. Valence electrons comes from the group number: C = 4, N = 5, O = 6, halogens = 7, H = 1. And number of bonds counts connections, not electrons — a double bond counts as 2 and a triple as 3, because you are counting the atom's half-share of each shared pair.

An equivalent way to write the last two terms is "lone-pair electrons plus half of bonding electrons," which some books prefer. It gives the same answer; use whichever you can execute faster without a mistake.

Worked examples

Worked examples — three you should be able to do cold

NH₄⁺. Nitrogen has 4 bonds and 0 lone pairs. FC = 5 − 0 − 4 = +1. The hydrogens are each 1 − 0 − 1 = 0. Total = +1, matching the ion's charge. ✓

OH⁻. Oxygen has 1 bond and 3 lone pairs (6 nonbonding electrons). FC = 6 − 6 − 1 = −1. Hydrogen is 0. Total = −1. ✓

H₃O⁺. Oxygen has 3 bonds and 1 lone pair (2 electrons). FC = 6 − 2 − 3 = +1. Total = +1. ✓

The formula run on the three ions worth knowing cold. Read each drawing left to right: count the lines coming off the atom, count the dots sitting on it, subtract both from the group number. The last term is the one people get wrong — it counts bonds, not bonding electrons, because you own only half of each shared pair, which is one electron per bond.FC = valence electrons − lone-pair electrons − number of BONDS count bonds, not bonding electrons — a double bond is 2, not 4

A pattern emerges quickly, and memorizing it is faster than recomputing. For oxygen: 3 bonds means +1, 2 bonds means 0, 1 bond means −1. For nitrogen: 4 bonds means +1, 3 bonds means 0, 2 bonds means −1. For carbon: 3 bonds with a lone pair means −1 (a carbanion), 3 bonds with an empty orbital means +1 (a carbocation), 4 bonds means 0. Once you internalize these, you can read charges off a structure at a glance, which is what you will actually be doing while drawing mechanisms.

The whole of formal charge, memorised as nine pictures rather than recomputed each time. Read a row: give oxygen an extra bond and it goes +1, take one away and it goes −1. Nitrogen behaves the same way one column over, which is why a nitrogen with four bonds always carries a positive charge. Carbon is the odd one out, because its two three-bond forms differ only in whether the leftover pair is there — present makes a carbanion, absent leaves an empty orbital and a carbocation.

The built-in check

The sum of every atom's formal charge must equal the overall charge of the molecule or ion. For a neutral molecule, the sum must be exactly zero. This is not a guideline — it is arithmetic, and it cannot fail for a correctly drawn structure. If your total does not match, you have either miscounted lone pairs or drawn the wrong number of bonds somewhere, and the check has just told you to go back and look.

Use it especially when pushing arrows. Every mechanism step must conserve charge: if the reactants sum to −1, the products must sum to −1. Arrow-pushing errors show up here before they show up anywhere else.

Why it matters: picking the right structure

When a molecule or ion can be drawn in several ways that all satisfy the octet rule, formal charge is the tiebreaker. The best structure, in order of priority:

Has the fewest formal charges overall. A structure with no charges beats one with a +1 and a −1.
Puts negative charge on the most electronegative atom and positive charge on the least. Negative charge on oxygen is comfortable; negative charge on carbon is not.
Avoids like charges on adjacent atoms. Two adjacent positives are electrostatically expensive.

These same three rules, in the same order, reappear in Module 2 as the rules for ranking resonance structures — which is not a coincidence, since ranking resonance contributors is exactly this problem applied to a set of structures that all describe the same real molecule.

Perchlorate, ClO₄⁻, drawn with four single bonds: chlorine has 7 valence electrons, 4 bonds and 0 lone pairs, so FC = 7 − 0 − 4 = +3, sitting next to four oxygens at −1 each. The charges sum correctly to −1, so the structure is legal — but a +3 next to four −1s is exactly the kind of charge separation rule 1 tells you to avoid, and a chemist reading this immediately looks for a structure that spreads the charge better. Formal charge did its job: it flagged the problem.

Formal charge is not real charge. It is an accounting convention that splits every bonding pair exactly down the middle, which ignores electronegativity entirely. The nitrogen in NH₄⁺ carries a formal +1, but the actual electron density there is not especially depleted — the hydrogens bear most of the real positive character, because nitrogen is more electronegative than hydrogen. Formal charge tells you where to write the sign in a structure. It does not tell you where the electron density actually is; for that you want the electronegativity reasoning from the previous section.

The most common arithmetic slip is counting electrons in the bond term instead of bonds. In H₃O⁺, oxygen has three bonds — the term is 3, not 6. If your answers come out consistently too negative, this is almost certainly why.

Reading charges to find reactive sites

Beyond bookkeeping, formal charge is a fast way to spot where a molecule will react. A positively charged atom is short of electrons, so it attracts anything carrying spare ones; a negatively charged atom has electrons to spare, so it goes looking for somewhere to put them. Module 2 gives those two roles their names. A carbon with a formal +1 and only six valence electrons — a carbocation — is both charged and octet-deficient, which is why it is one of the most reactive species in the course.

When you meet a new molecule, assigning formal charges is a reasonable first move. It takes fifteen seconds and it usually tells you what the molecule is going to do.

What carries forward

You will compute formal charges constantly: to validate Lewis structures in the next section, to rank resonance contributors in Module 2, to check that each arrow-pushing step conserves charge, and to identify the charged intermediates — carbocations, carbanions, enolates, tetrahedral intermediates — that every mechanism from Module 6 onward passes through.