Chemistry revision sheets

Formal Charge vs Oxidation Number: The Difference

Both numbers sit on the same atom in the same molecule, and both look like charges. Yet carbon in CO has a formal charge of −1 and an oxidation number of +2 — opposite signs on the same atom. That isn't a contradiction. They're answers to two different questions, and each one is deliberately, usefully wrong in its own way.

The short answer: formal charge splits every bonding pair straight down the middle, giving one electron to each atom, and asks what charge each atom would then carry. Oxidation number does the opposite extreme — it gives both electrons of every bond to the more electronegative atom. Formal charge is for picking the best Lewis structure; oxidation number is for tracking redox.

Quick comparison at a glance

FeatureFormal chargeOxidation number
How bonds are splitEvenly — one electron eachEntirely to the more electronegative atom
Assumption madeBonds are perfectly covalentBonds are perfectly ionic
Electronegativity used?NoYes — it decides who takes the pair
Typical valuesSmall: −1, 0, +1Wide: −4 up to +7
What it's used forChoosing between Lewis structuresIdentifying oxidation and reduction
Sum over a speciesEquals the overall chargeEquals the overall charge
Carbon in CO−1+2
Carbon in CH₄0−4
RealityNeither is the real charge — the truth lies between them

What formal charge is

Formal charge asks: if this atom shared every bond perfectly fairly, how many electrons would it end up with compared with the free atom?

Formal charge = (valence electrons) − (lone-pair electrons) − ½(bonding electrons)

Or, equivalently and often faster: valence electrons − lone-pair electrons − number of bonds.

Take the oxygen in water. Oxygen has 6 valence electrons. In H₂O it has 2 lone pairs (4 electrons) and 2 bonds. Formal charge = 6 − 4 − 2 = 0. Each hydrogen: 1 − 0 − 1 = 0. Everything is neutral, which is exactly what you'd hope for such an ordinary molecule.

Formal charge earns its keep when a molecule can be drawn more than one way. The rules for choosing:

  1. Prefer the structure where formal charges are closest to zero.
  2. If charges are unavoidable, put the negative one on the more electronegative atom.
  3. Avoid structures with like charges on adjacent atoms.

For example, carbon dioxide can be drawn as O=C=O or as ⁻O–C≡O⁺. In O=C=O every atom has a formal charge of zero; the alternative manufactures a −1 and a +1 for no reason, and puts the positive charge on oxygen, the most electronegative atom present. Formal charge picks O=C=O — and experiment agrees: both C–O bonds are identical and sit at double-bond length.

What an oxidation number is

An oxidation number (or oxidation state) asks a bluntly different question: if every bond were fully ionic, what charge would this atom carry? The more electronegative atom takes both electrons from every bond it's in.

You rarely have to reason it out from scratch — the standard rules do the work:

  • Element in its standard state → 0 (O₂, Na, S₈, Cl₂ are all 0)
  • Monatomic ion → equal to its charge (Na⁺ is +1, Cl⁻ is −1)
  • Group 1 → +1; Group 2 → +2; fluorine → −1 in every compound
  • Hydrogen → +1 (but −1 in metal hydrides such as NaH)
  • Oxygen → −2 (but −1 in peroxides such as H₂O₂, and +2 in OF₂)
  • Everything must sum to the overall charge on the species

Water: two hydrogens at +1 each, so oxygen must be −2. Sulfate SO₄²⁻: four oxygens at −2 give −8, and the total must be −2, so sulfur is +6.

The example that makes the difference obvious

Carbon monoxide, C≡O, with one lone pair on each atom.

Formal charge — split the triple bond evenly, three electrons to each side:

  • Carbon: 4 − 2 − 3 = −1
  • Oxygen: 6 − 2 − 3 = +1
  • Sum: 0 ✓

Oxidation number — oxygen is more electronegative, so it takes all six bonding electrons:

  • Carbon: +2
  • Oxygen: −2
  • Sum: 0 ✓

Carbon comes out negative one way and positive +2 the other. Both are correct; neither is the actual charge on the atom, which is small and somewhat negative on carbon. The two numbers bracket reality from opposite sides, because one pretends the bond is perfectly covalent and the other pretends it's perfectly ionic.

How to tell which one a question wants

  • "Which Lewis structure is best?" / "Where does the charge go?" → formal charge.
  • "What has been oxidised?" / "Balance this redox equation." → oxidation number.
  • Roman numerals in a name — iron(III) chloride, manganese(IV) oxide — are always oxidation numbers.
  • A charge written on a Lewis structure — the ⊕ on nitrogen in the nitrate ion — is a formal charge.

Notation helps too: oxidation numbers are conventionally written sign-first (+2, −2) and formal charges as a superscript on the atom.

Worked examples

Work out both numbers for the bold atom before reading on.

  • C in CH₄ → formal 4 − 0 − 4 = 0; oxidation −4 (carbon beats hydrogen for electronegativity, so it takes all four pairs).
  • C in CO₂ → formal 4 − 0 − 4 = 0; oxidation +4.
  • N in NH₄⁺ → formal 5 − 0 − 4 = +1; oxidation −3.
  • O in H₂O₂ → formal 6 − 4 − 2 = 0; oxidation −1 (the O–O bond is split evenly even by the oxidation-number rules, because the two atoms are identical).
  • S in SO₄²⁻ (all four bonds drawn as single, four negative oxygens) → formal 6 − 0 − 4 = +2; oxidation +6.
  • N in NO₃⁻ → formal +1 on nitrogen; oxidation +5.

Notice how much larger the oxidation numbers run. That's the tell: formal charges cluster near zero, oxidation numbers spread wide.

Common mistakes to avoid

  • Treating either as the real charge on the atom. Both are deliberate fictions built on opposite extreme assumptions. The genuine partial charge is somewhere in between and is usually a fraction.
  • Using electronegativity when calculating formal charge. Formal charge never looks at electronegativity — it splits every bond exactly in half, even H–F.
  • Forgetting the exceptions in the oxidation-number rules. Oxygen is −1 in peroxides and +2 in OF₂; hydrogen is −1 in metal hydrides. Exam questions are drawn to these cases precisely because they catch out rule-followers.

FAQ

What is the difference between formal charge and oxidation number?
Formal charge splits each bonding pair evenly between the two atoms; oxidation number assigns both electrons to the more electronegative atom. Formal charge is used to choose the best Lewis structure, oxidation number to track redox changes.

How do you calculate formal charge?
Formal charge = valence electrons − lone-pair electrons − number of bonds. For oxygen in water: 6 − 4 − 2 = 0.

Can formal charge and oxidation number be different for the same atom?
Yes, and usually they are. Carbon in CO has a formal charge of −1 and an oxidation number of +2. They coincide only when every bond the atom makes is to an atom of the same element, because a homonuclear bond gets split down the middle by both methods. As soon as the partners differ, oxidation number hands the whole pair to one side — carbon in methane has a formal charge of 0 but an oxidation number of −4, even though carbon and hydrogen are close in electronegativity.

Which one is the real charge on the atom?
Neither. One assumes the bonds are perfectly covalent and the other that they're perfectly ionic, so the true partial charge lies somewhere between them.

The takeaway

Same molecule, two ways of splitting a bond. Formal charge shares evenly and helps you pick the best Lewis structure; oxidation number hands everything to the greedier atom and helps you spot who was oxidised. Learn which question is being asked and the two numbers stop competing.

Uses this → Oxidation vs Reduction: What's the Difference? (Redox). Depends on → What Is a Lewis Structure? Dots, Bonds, and Octets and What Is Electronegativity? Trends and Examples. Next → [What Is Resonance?] (sibling), where formal charge does its best work.

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