Worked example 1
What is the oxidation number of sulfur in H₂SO₄?
Try it first: Assign hydrogen and oxygen first, then let sulfur absorb the remainder.
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Chemical ReactionsOxidation NumbersContent level: Core 20 min
Assigning Oxidation Numbers
What you'll be able to do: Assign oxidation numbers to every atom in a compound or ion using a fixed order of rules.
Introduction
Oxidation numbers are a bookkeeping device, a way of pretending that every bond is fully ionic so electrons can be counted. They are not real charges, but they are the only reliable way to spot a redox reaction.
These are recommended, not required. You can start this lesson at any time.
Learning objectives
- Apply the oxidation number rules in the correct priority order
- Determine the oxidation state of an unknown atom by algebra
- Recognise the standard exceptions for oxygen and hydrogen
- Use oxidation numbers to name compounds of variable-charge metals
Lesson
A deliberate fiction
An oxidation number is the charge an atom would carry if every shared pair of electrons were handed entirely to the more electronegative partner. In CH₄ the carbon is not really at -4, but treating it that way lets us track electron flow consistently across any reaction.
Oxidation numbers are per atom, written with the sign first, as in +2. Ionic charges are written with the sign last, as in 2+.
The priority list
Apply the rules in order and stop when the atom is assigned. First, any atom in its free element form is 0. Second, a monatomic ion equals its charge. Third, fluorine is always -1. Fourth, group 1 metals are +1 and group 2 metals are +2. Fifth, hydrogen is +1 with nonmetals. Sixth, oxygen is -2. Whatever is left is found by making the total come out right.
Order matters. In H₂O₂ oxygen cannot be -2, because hydrogen is assigned first and the total must be zero.
The exceptions worth knowing
Oxygen is -1 in peroxides such as H₂O₂ and Na₂O₂, and +2 when bonded to fluorine in OF₂. Hydrogen is -1 in metal hydrides such as NaH and CaH₂, because the metal is less electronegative. These three exceptions cover almost every trick question you will see.
H₂O₂: H = +1, O = -1
Solving for the unknown
Set up a small equation. Multiply each known oxidation number by how many of that atom appear, add an unknown for the remaining atom, and set the sum equal to the overall charge. This works identically for neutral compounds and for polyatomic ions.
sum of (oxidation number x count) = overall charge
Why chemists care
Variable-charge metals are named by their oxidation state, so iron(III) chloride is FeCl₃. More importantly, comparing oxidation numbers before and after a reaction is the definitive test for whether electrons moved at all.
Identifying redox reactionsAn average oxidation state can hide distinct sites
The charge-sum rule sometimes gives a fractional average oxidation number in a mixed-valence compound. That does not require any individual atom to carry a fractional ionic charge. Magnetite Fe₃O₄ is a familiar example: the formal total from oxygen is -8, so three iron atoms average +8/3, consistent with one Fe²⁺ and two Fe³⁺ in a simple ionic accounting.
Distinguish a formula-average value from oxidation states at individual sites.
Key ideas
Definition
Oxidation number
The hypothetical charge on an atom if all bonds were treated as fully ionic.
Rule
Free elements are zero
Any element in its uncombined form, including O₂, P₄ and Na(s), has an oxidation number of 0.
Rule
Sums are fixed
Oxidation numbers total zero in a neutral compound and equal the charge in a polyatomic ion.
Assumption
Oxygen defaults to -2
Oxygen is -2 except in peroxides, where it is -1, and with fluorine, where it is positive.
Assumption
Hydrogen defaults to +1
Hydrogen is +1 with nonmetals but -1 in metal hydrides.
Equation
Oxidation number sum rule
sum(n x oxidation number) = total charge
- n = number of atoms of that element in the formula
- total charge = zero for a compound, the ion charge for a polyatomic ion
Worked examples
Worked example 2
What is the oxidation number of chromium in Cr₂O₇²⁻?
Try it first: Remember that the total here is not zero.
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Worked example 3
Assign oxidation numbers to every atom in NaH and in H₂O₂.
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Worked example 4
Find the average Fe oxidation number in Fe₃O₄ and give one integer-state breakdown.
Try it first: Assign oxygen its usual -2.
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Common mistakes
Assuming oxygen is always -2.
Why it's wrong: In peroxides oxygen is -1 and with fluorine it is positive.
Check instead: Assign higher-priority atoms first and let the sum rule reveal the exception.
Multiplying the oxidation number by the subscript when reporting it.
Why it's wrong: Oxidation number is a per-atom quantity, the subscript is only used inside the sum.
Check instead: Report the value for one atom, such as O = -2 in H₂O, not -4.
Confusing oxidation number with formal charge or actual charge.
Why it's wrong: Oxidation numbers assume bonds are fully ionic, which real bonds are not.
Check instead: Use oxidation numbers only for tracking electron transfer.
Setting the sum to zero for a polyatomic ion.
Why it's wrong: The sum must equal the ion charge, not zero.
Check instead: Read the superscript charge before writing the equation.
After the lesson · Question bank
Independent practice
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What you should now know
Oxidation numbers are assigned by working through a fixed priority list: free elements are zero, monatomic ions equal their charge, fluorine is always -1, oxygen is usually -2 and hydrogen is usually +1. The sum must equal zero for a neutral compound or the charge for a polyatomic ion, which lets you solve for any remaining unknown atom algebraically.
- Oxidation numbers are bookkeeping, not real charges
- Work through the priority rules in order, then solve for the leftover atom
- Free elements are always zero
- Oxygen is -1 in peroxides and hydrogen is -1 in metal hydrides
- The sum equals zero for compounds and the charge for polyatomic ions
Sources and further reading
This lesson is original Chem Help content. No external sources were adapted.