Lewis structures, formal charges and resonance

Count the available electrons first

A Lewis structure organizes valence-electron pairs as bonds and nonbonding lone pairs. Start with the relevant atom valence counts and correct for the net charge: add electrons for a negative ion and subtract them for a positive ion. For CO₃²⁻, carbon contributes 4 and the three oxygens 18, then the charge adds 2: 24 valence electrons.

Choose a plausible atom connectivity, distribute bonding and lone-pair electrons, then check the total. Familiar octet and hydrogen-duet patterns guide many small main-group examples, but they have exceptions. Odd-electron, electron-deficient and hypercoordinate species need their appropriate scope rather than forced compliance with one drawing rule.

Formal charge uses equal bond sharing

Formal charge assigns each atom its nonbonding electrons and half the bonding electrons. In electron-count form: formal charge = free-atom valence count − nonbonding count − half the bonding count. Sum the atom formal charges to check that they reproduce the whole species’ net charge.

For a CO triple-bond contributor with one lone pair on each atom, the total valence count is 10. Carbon has formal charge 4 − 2 − 3 = −1; oxygen has 6 − 2 − 3 = +1. Their sum is zero. These are equal-sharing bookkeeping assignments; they are not a measurement showing whole integer charges localized on the actual atoms.

Do not merge three charge concepts

Oxidation-state bookkeeping assigns heteronuclear bonding electrons using a different convention based on electronegativity. For the stated CO case it gives C +2 and O −2, rather than the Lewis formal charges C −1 and O +1. Both sum to the neutral species’ total because they answer different assignment questions.

Partial charges describe a nonuniform electronic distribution under a stated analysis or model and need not be integers. Different schemes can assign different numerical partial charges. Do not infer an actual measured molecular dipole directly from one formal-charge diagram or a naive electronegativity arrow, especially for CO. The convention and the evidence both matter.

Several drawings can represent one state

A carbonate Lewis contributor has one C=O bond and two C–O single bonds. Carbon and the double-linked oxygen have formal charge zero; each single-linked oxygen has formal charge −1. There are three equivalent placements of the double bond, keeping the atom connectivity fixed.

These resonance contributors describe aspects of one delocalized electronic state. The actual ion is not repeatedly switching among three separate Lewis molecules. In the isolated symmetric carbonate description, the three C–O links are equivalent. The bookkeeping average (2 + 1 + 1)/3 = 4/3 can summarize the bond-order picture, but is not a movie of bond orders changing over time.

Worked examples

Try it yourself

1. For a −2 ion, do you add or subtract two electrons in the Lewis total?

   Show the answer and reasoning

   Add two.

   Negative charge means more electrons relative to the neutral atom-count sum.

2. In the stated CO triple-bond contributor, what are the C and O formal charges?

   Show the answer and reasoning

   C −1 and O +1.

   Use equal sharing: 4 − 2 − 3 and 6 − 2 − 3 respectively.

3. Must an atom’s formal charge equal its oxidation state?

   Show the answer and reasoning

   No.

   The assignments distribute bonding electrons using different conventions.

4. Does carbonate resonance mean three structures switch back and forth in time?

   Show the answer and reasoning

   No.

   The contributors are representations of one electronic state with unchanged atom connectivity.

What to carry forward

• Count valence electrons and the net charge first.
• Formal charges use equal sharing of bond electrons.
• Formal charge, oxidation state and partial charge are distinct.
• Resonance contributors represent one delocalized state rather than time switching.