AP Chemistry Formal Charge

The formula, and three full worked examples.

Formal charge is how you decide which Lewis structure actually represents a molecule best, when more than one non-equivalent structure seems possible. The calculation is short and mechanical; the skill worth practicing is applying it consistently across every atom in a structure.

The Formal Charge Formula

FC = (valence electrons) − (nonbonding electrons) − ½(bonding electrons), calculated separately for each atom in a drawn Lewis structure. Valence electrons is the atom's normal count; nonbonding electrons are its own lone-pair electrons in that structure; bonding electrons are every electron in every bond attached to that atom (count both electrons of each bond, then halve the total).

Worked Example 1: A Polyatomic Ion (NH₄⁺)

The ammonium ion, a common polyatomic ion, NH₄⁺, has nitrogen singly bonded to four hydrogen atoms, with no lone pairs on nitrogen. Find the formal charge on each atom.

Nitrogen: 5 valence electrons − 0 nonbonding − ½(8 bonding electrons, from its 4 bonds) = 5 − 0 − 4 = +1
Each hydrogen: 1 valence electron − 0 nonbonding − ½(2 bonding electrons) = 1 − 0 − 1 = 0

Sum: (+1) + 0 + 0 + 0 + 0 = +1, matching NH₄⁺'s actual overall charge, confirming the structure is drawn correctly.

Worked Example 2: A Simple Molecule (CO₂)

CO₂'s standard structure is O=C=O, carbon with no lone pairs and a double bond to each oxygen; each oxygen with two lone pairs and one double bond to carbon. Find the formal charge on each atom.

Carbon: 4 − 0 − ½(8) = 4 − 0 − 4 = 0
Each oxygen: 6 − 4 − ½(4) = 6 − 4 − 2 = 0

Every atom has a formal charge of zero, the sign that this structure is already the best possible model, nothing to improve by redrawing it.

Worked Example 3: Comparing Two CO₂ Structures

Compare CO₂'s standard structure (Example 2, all formal charges zero) against an alternative, valid-but-worse structure: O≡C–O⁻, where one oxygen triple-bonds to carbon (one lone pair) and the other single-bonds to carbon (three lone pairs).

Triple-bonded oxygen: 6 − 2 − ½(6) = +1
Carbon: 4 − 0 − ½(8) = 0
Single-bonded oxygen: 6 − 6 − ½(2) = −1

This alternative structure is a valid Lewis structure (every atom has a full octet, and the formal charges still sum to zero overall), but it carries nonzero formal charges (+1 and −1) where Example 2's structure carries none. Since minimizing formal charges is the deciding rule, Example 2's symmetric O=C=O structure is the better model, and it's the one that actually represents real CO₂.

Common Formal Charge Mistakes

Forgetting to halve the bonding electrons. Count both electrons in every bond attached to the atom, then divide that total by 2, don't use the number of bonds directly.
Mixing up formal charge with oxidation number. Formal charge splits bonding electrons evenly; oxidation number assigns them entirely to the more electronegative atom. They're different calculations and often give different answers for the same atom.
Forgetting that formal charges must sum to the actual overall charge. If they don't, the Lewis structure or the formal charge arithmetic has an error somewhere.
Picking a structure by minimizing formal charges alone, ignoring electronegativity. When two structures tie on total formal charge, the better one places any negative formal charge on the more electronegative atom.

This same method, calculate formal charge for every atom, then compare structures, applies across every unit where Lewis structures come up. For every other free AP Chemistry tool and guide on this site, start from the AP Chem Score Calculator.

Frequently Asked Questions

What is formal charge?

The hypothetical charge an atom would have within a Lewis structure if every bonding pair of electrons were split exactly evenly between the two bonded atoms, regardless of real electronegativity differences. It's a bookkeeping tool for evaluating a Lewis structure, not a real, measurable charge.

What is the formal charge formula?

FC = (valence electrons) − (nonbonding electrons) − ½(bonding electrons), where valence electrons is the atom's normal count on the periodic table, nonbonding electrons are its lone-pair electrons in that structure, and bonding electrons are all the electrons in bonds attached to that atom.

How does formal charge help choose between Lewis structures?

When more than one non-equivalent Lewis structure is possible for a molecule or ion, the best structure is the one that minimizes formal charges overall, and places any negative formal charge on the more electronegative atom. A structure with large or poorly-placed formal charges is a worse model of the real molecule.

Do formal charges have to sum to the molecule or ion's actual charge?

Yes, always. The sum of every atom's formal charge in a correct Lewis structure must equal the overall charge on the species, zero for a neutral molecule, or the actual ionic charge for an ion. If they don't match, there's an error in the structure or the formal charge calculation.

Is formal charge the same as oxidation number?

No, even though both are calculated ways of assigning charge to an atom. Formal charge splits every bond's electrons evenly between the two atoms, regardless of electronegativity. Oxidation number assigns both of a bond's electrons entirely to the more electronegative atom. The two methods can give different values for the same atom in the same structure.

This page is not affiliated with or endorsed by College Board. AP® is a trademark registered by the College Board.