General Chemistry I · Chemical Bonding & Molecular Geometry
Formal Charge
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In 30 seconds
Formal charge is a bookkeeping tool that assigns each atom a hypothetical charge, as if all bonding electrons were shared perfectly equally. It is computed as valence electrons minus lone-pair electrons minus half the bonding electrons. Although formal charge is not the "real" charge on an atom, comparing formal charges across candidate structures tells you which Lewis structure is most reasonable — the one that minimizes formal charges and places any negative formal charge on the more electronegative atom.
Why this matters
Formal charge lets chemists decide among competing Lewis structures and rationalize reactivity. It explains why CO₂ is drawn with two double bonds (all FC = 0) rather than a triple + single bond (which would force nonzero formal charges), and why in a molecule like CO the negative formal charge sits on oxygen (the more electronegative atom). It is also the first step toward understanding resonance hybrids and acid–base behavior.
The college version
Key Ideas
- Definition: FC = (valence electrons) − (lone-pair electrons) − ½(bonding electrons).
- Interpretation: a formal charge of 0 means the atom "keeps" exactly its valence count; +1 means it is one electron short; −1 means one electron extra, relative to the free atom.
- Best-structure rules:
- Structures with formal charges closest to zero are preferred.
- If formal charges are unavoidable, negative charges should sit on the more electronegative atom.
- Sum rule: the formal charges of all atoms add up to the total charge on the molecule or ion.
- Formal charge ≠ oxidation number ≠ real partial charge: it is a model-based assignment, not a measured quantity.
Equations and Variables
- FC = V − N − ½B
- FC = formal charge
- V = number of valence electrons in the free (neutral) atom
- N = number of nonbonding (lone-pair) electrons on the atom in the structure
- B = number of bonding (shared) electrons around the atom (each bond = 2 electrons)
How It Works
- Draw a Lewis structure.
- For each atom, count its lone-pair electrons (N) and its bonding electrons (B = 2 per bond, counting both electrons of every bond it participates in).
- Subtract N and ½B from the atom's valence count V to get FC.
- Repeat for every atom; sum the FCs and confirm they equal the overall charge.
- When several structures are possible, favor the one with the smallest FCs and negative FC on the most electronegative atom.
Worked Example
1. Compute formal charges in CO₂ (O=C=O).
- Carbon: V = 4, N = 0 (no lone pairs), B = 8 (two double bonds). FC = 4 − 0 − ½(8) = 0.
- Each oxygen: V = 6, N = 4 (two lone pairs), B = 4 (one double bond). FC = 6 − 4 − ½(4) = 0.
- Sum = 0, matching the neutral molecule. ✓ All atoms have FC = 0, confirming this is an excellent structure.
2. Compute formal charges in NH₄⁺.
- Nitrogen: V = 5, N = 0, B = 8 (four single bonds). FC = 5 − 0 − ½(8) = +1.
- Each hydrogen: V = 1, N = 0, B = 2. FC = 1 − 0 − ½(2) = 0.
- Sum = +1, which matches the ion's +1 charge. ✓
Common Confusions
- "Formal charge is the atom's real charge." — Wrong: it's a hypothetical equal-sharing assignment; the real distribution is governed by electronegativity.
- "You should always maximize formal charges." — Wrong: you minimize them; structures with FC near zero are favored.
- "Negative formal charge goes on the least electronegative atom." — Wrong: it goes on the most electronegative atom.
- "The sum of formal charges can differ from the ion's charge." — Wrong: they must sum exactly to the overall charge.
Eli explains
The same idea, in plain words
Explain it like I’m 10
Imagine you and a friend are divvying up a shared pizza, but by a rule that ignores who actually paid: each slice is split 50/50. Formal charge is like counting who ended up with more or fewer slices than they "should" have under that artificial rule. If everyone ends up exactly even, that's the "fairest" (most stable) arrangement — the structure with all zeros. If someone must be short a slice, you'd rather it be the hungrier, more electron-loving atom. (The analogy captures the equal-splitting bookkeeping; real electrons are not shared equally — that's what electronegativity and polarity describe.)
Key takeaways
- FC = V − N − ½B (valence minus lone-pair minus half the bonding electrons).
- FC = 0 is ideal; minimize formal charges when choosing structures.
- Negative formal charge should go on the more electronegative atom.
- Sum of formal charges = overall charge of the species.
- Each bond contributes 2 electrons to B for each bonded atom.
- Formal charge is a model assignment, not a real measured charge.
- A neutral molecule must have all FCs summing to zero.
- FC = V − N − ½B.
- Prefer structures with FC closest to zero.
- Negative FC on the most electronegative atom.
- Sum of FC = total charge.
- FC guides the choice among resonance/alternative structures.
- FC ≠ oxidation number ≠ real charge.
Study tools & related lessonsYou’ll learn to · Related
You’ll learn to
- Define formal charge and compute it for any atom in a Lewis structure.
- Use formal charge to choose the most plausible resonance structure.
- Explain why the best structure keeps formal charges close to zero and places negative charge on the most electronegative atom.
- Relate the sum of formal charges to the overall charge on a species.
Sources & references
- OpenStax, *Chemistry 2e*, "7.4 Formal Charges and Resonance."
- OpenStax, *Chemistry 2e*, "7.3 Lewis Symbols and Structures."
- OpenStax, *Chemistry 2e*, "Ch. 7 Introduction."
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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