Ionic Lewis Structures: Electron Counts and Formal Charge

Charge changes the electron budget

Start by adding the valence electrons of the neutral atoms. For a negative ion, add electrons equal to the magnitude of its charge; for a positive ion, subtract them. Then distribute that total among bonds and nonbonding electrons. Enclose the completed ion in brackets and put its total charge outside.

For ammonium, NH₄⁺, the count is 5 + 4(1) − 1 = 8. Four N–H single bonds use all eight electrons. Nitrogen has an octet, each hydrogen has a duet, and there is no nitrogen lone pair.

Formal charge is a second check

For a proposed structure, assign an atom all of its nonbonding electrons and half the electrons in each bond. Subtract this assignment from the neutral atom's valence count:

$$\mathrm{FC}=V-N_{\mathrm{nonbonding}}-\tfrac12N_{\mathrm{bonding}}.$$

In ammonium, nitrogen's formal charge is 5 − 0 − 4 = +1. Each hydrogen has 1 − 0 − 1 = 0. Their sum matches the ion's +1 charge. Writing a positive charge outside the brackets does not remove the need for this internal audit.

The same procedure works for anions

For hypochlorite, ClO⁻, 7 + 6 + 1 = 14 electrons are available. One Cl–O single bond and three lone pairs on each atom account for all fourteen. Chlorine has formal charge 7 − 6 − 1 = 0; oxygen has 6 − 6 − 1 = −1.

This example shows why the charge outside brackets and the formal charge on an atom serve different purposes: one labels the whole species, while the other evaluates an electron allocation within a particular Lewis structure.

Diagnose common mistakes

If formal charges do not sum to the stated ion charge, check the electron count and missing lone pairs. Do not add a charge symbol merely to make an otherwise incorrect drawing look complete. Formal charge also differs from oxidation state and from measured or calculated partial charge; equal division of bonding electrons is its defining bookkeeping convention.