Formal charge is a bookkeeping tool that helps chemists track how electrons are assigned in a Lewis structure. It does not always equal the real electrical charge on an atom, but it helps compare possible structures for the same molecule or ion. Learning formal charge is important because many molecules can be drawn in more than one way, and not all drawings are equally reasonable.
A good Lewis structure usually keeps formal charges small and places negative charge on more electronegative atoms.
To calculate formal charge, compare an atom's valence electrons to the electrons it owns in the Lewis structure. Lone pair electrons count fully for the atom, while bonding electrons are split equally between bonded atoms. Chemists use these values to choose the best structure, check resonance forms, and predict reactive sites.
Formal charge is especially useful for ions, expanded octets, resonance structures, and molecules with multiple possible bonding patterns.
Key Facts
- Formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons.
- For a single bond, each bonded atom is assigned 1 electron from the bond.
- For a double bond, each bonded atom is assigned 2 electrons from the bond.
- The sum of all formal charges must equal the overall charge of the molecule or ion.
- The best Lewis structure usually has the smallest possible formal charge magnitudes.
- When formal charges cannot be avoided, negative formal charge is usually best placed on the more electronegative atom.
Vocabulary
- Formal charge
- The charge assigned to an atom in a Lewis structure by assuming bonding electrons are shared equally.
- Valence electrons
- Electrons in the outer shell of an atom that are available for bonding and lone pairs.
- Lone pair
- A pair of valence electrons localized on one atom and not shared in a bond.
- Lewis structure
- A diagram that shows atoms, bonds, lone pairs, and sometimes formal charges in a molecule or ion.
- Resonance structure
- One of two or more valid Lewis structures that differ only in electron placement, not atom positions.
Common Mistakes to Avoid
- Counting all bonding electrons for one atom is wrong because formal charge splits each bond equally between the bonded atoms.
- Forgetting lone pairs is wrong because nonbonding electrons are fully assigned to the atom when calculating formal charge.
- Choosing a structure only because every atom has an octet can be wrong because formal charges and electronegativity also affect which structure is best.
- Ignoring the total charge is wrong because the sum of formal charges must match the molecule or ion charge exactly.
Practice Questions
- 1 Calculate the formal charge on nitrogen in NH4+ if nitrogen has four single bonds and no lone pairs.
- 2 In CO2 drawn as O=C=O, calculate the formal charge on each oxygen and on carbon.
- 3 Two Lewis structures for NCO- place the negative formal charge either on nitrogen or oxygen. Which placement is usually more favorable, and why?