Assign oxidation numbers using the standard rules (O is -2, H is +1, and so on); an unknown one is found because the oxidation numbers add up to zero in a compound or to the ion's charge.
This page walks through assigning oxidation numbers for SPM Chemistry, step by step: the formula you need, the units to watch, and a worked example.
When you use this
Oxidation number is the tool that decides, in SPM Chemistry redox, what has been oxidised and what has been reduced. You assign a number to each atom using a short list of rules, then find any unknown one from the fact that they must add up correctly.
The rules and units
Learn these standard values:
- An uncombined element has oxidation number 0 (for example Cu, O2).
- A monatomic ion equals its charge (Na+ is +1, Cl− is -1).
- Oxygen is usually -2 (but -1 in a peroxide).
- Hydrogen is usually +1 (but -1 in a metal hydride).
- Group 1 metals are +1, Group 2 metals are +2, fluorine is -1.
- In a neutral compound the oxidation numbers add up to 0; in a polyatomic ion they add up to the ion’s charge.
Units. Oxidation number is a signed whole number (write the sign, as in +6 or -2); it has no unit.
Worked example 1 (easy)
Find the oxidation number of chlorine in NaCl.
- Sodium (Group 1) is +1.
- The compound is neutral, so (+1) + (Cl) = 0, giving Cl = -1.
Worked example 2 (medium)
Find the oxidation number of sulfur in H2SO4.
- Hydrogen is +1, and there are two: 2 x (+1) = +2.
- Oxygen is -2, and there are four: 4 x (-2) = -8.
- Neutral compound: (+2) + S + (-8) = 0, so S = +6.
Worked example 3 (SPM level)
Find the oxidation number of manganese in the permanganate ion, MnO4−, and use it to say whether manganese is reduced when MnO4− becomes Mn2+.
- Oxygen is -2, and there are four: 4 x (-2) = -8.
- The ion’s charge is -1, so Mn + (-8) = -1, giving Mn = +7.
- Going from +7 in MnO4− to +2 in Mn2+ is a decrease of 5, so manganese is reduced (it gains electrons); MnO4− is acting as an oxidising agent.
Using oxidation number to spot redox
Once you can assign the numbers, redox becomes easy to read. An increase in oxidation number is oxidation; a decrease is reduction. The species that is reduced is the oxidising agent, and the species that is oxidised is the reducing agent. Tracking one element’s number from left to right across an equation is the quickest way to justify, in the exam, which change is which, a routine Paper 2 requirement.
Naming from a half-reaction
Many redox questions in Form 5 come as two half-equations. Assign the oxidation number of the key element on each side and read the change: the half-equation where the number rises is oxidation and releases electrons, and the one where it falls is reduction and accepts them. This is how you decide, for example, that zinc displacing copper is oxidised while the copper ion is reduced, and it turns a wordy redox question into two quick sums. The same idea helps you balance a redox equation, because the total rise in oxidation number must equal the total fall.
Common traps
- Dropping the sign. Always write + or - ; an oxidation number without a sign is incomplete.
- Forgetting to multiply. Multiply each element’s value by the number of its atoms before adding.
- Peroxides and hydrides. Oxygen is -1 in H2O2; hydrogen is -1 in NaH. Watch for these special cases.
- Ignoring the ion charge. In a polyatomic ion the numbers add up to the charge, not to zero.
Our teachers have students write each element’s oxidation number above its symbol, so oxidation and reduction can be read straight off the equation and the oxidising and reducing agents named with confidence.
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