Redox, Oxidation Numbers and Oxidising Agents

Redox means reduction and oxidation happening together. The concepts are straightforward; the errors come from the direction of electron transfer and from arithmetic slips in oxidation numbers.

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1. The definitions

There are three ways to define oxidation and reduction. Electron transfer is the one that always works.

	Oxidation	Reduction
Electrons	LOSS of electrons	GAIN of electrons
Oxygen	gain of oxygen	loss of oxygen
Hydrogen	loss of hydrogen	gain of hydrogen
Oxidation number	INCREASES	DECREASES

OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).

Oxidation is LOSS of electrons, not gain. This was recorded repeatedly in both directions — “oxidation is gaining electrons” and “accepting electrons is oxidation” were both stated by students. OIL RIG is the fix, and it is worth writing at the top of your answer sheet.

The oxygen definition is limited. Many redox reactions involve no oxygen at all — for example Mg + Cl₂ → MgCl₂. Uncertainty about identifying redox without oxygen or hydrogen present was recorded; use electrons or oxidation numbers instead.

Reduction sounds like it should mean losing something — and it does: the oxidation number is REDUCED. That is the memory hook that makes the direction stick.

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2. Oxidation numbers — the rules

Apply these in order:

1. An element on its own has oxidation number 0 — including O₂, Cl₂, H₂, Na, Fe. 2. A simple ion’s oxidation number = its charge — Na⁺ is +1, Cl⁻ is −1, Mg²⁺ is +2. 3. Oxygen is usually −2 (except in peroxides). 4. Hydrogen is usually +1 (except in metal hydrides, where it is −1). 5. Group I is +1, Group II is +2, and halogens are usually −1. 6. The oxidation numbers in a NEUTRAL COMPOUND add up to 0. 7. In a polyatomic ION they add up to the ION’S CHARGE.

An uncombined element is always 0. A recorded error gave chlorine in HCl as 0 — but that is a compound, so chlorine there is −1. It is only 0 as Cl₂.

Oxygen is −2, not +2. Recorded directly.

Write the SIGN. “+5”, not “5”. Tutors flagged this — an unsigned oxidation number can lose the mark.

An oxidation number is NOT the same as valency, and NOT the same as the compound’s charge. Both confusions were recorded. Valency has no sign; oxidation number does. And a neutral compound has an overall charge of zero while its individual atoms have non-zero oxidation numbers.

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3. Working out an unknown oxidation number

Method: let the unknown be x, apply the known values, and set the total to 0 (compound) or the ion’s charge (polyatomic ion).

Example — sulfur in H₂SO₄:

• 2(+1) + x + 4(−2) = 0
• 2 + x − 8 = 0
• x = +6

Example — sulfur in H₂SO₃:

• 2(+1) + x + 3(−2) = 0
• 2 + x − 6 = 0
• x = +4

Example — nitrogen in NO₃⁻:

• x + 3(−2) = −1 (the ion’s charge, not 0)
• x − 6 = −1
• x = +5

Multiply by the number of atoms. Forgetting to account for all the atoms — using −2 instead of 3(−2) — was the most frequent arithmetic error recorded. A student who computed nitrogen in NO₃⁻ as −3 had ignored the multiplication entirely.

For a polyatomic ion, the total equals the ION’S CHARGE, not zero. NO₃⁻ totals −1; SO₄²⁻ totals −2.

Example — chlorine in HClO₄:

• (+1) + x + 4(−2) = 0 → x = +7

Halogens are −1 only in simple binary compounds. In compounds with oxygen — HClO₄, ClO⁻ — chlorine takes a positive oxidation number, because oxygen is more electronegative. Forgetting the halogen rule’s exception was recorded.

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4. Identifying oxidation and reduction

Method — the reliable routine:

1. Assign oxidation numbers to every element on both sides
2. Find which ones change
3. Increase → oxidised. Decrease → reduced.

Example: Fe₂O₃ + 3CO → 2Fe + 3CO₂

Species	Before	After	Change
Fe	+3 (in Fe₂O₃)	0 (element)	decrease → REDUCED
C	+2 (in CO)	+4 (in CO₂)	increase → OXIDISED
O	−2	−2	unchanged

Set it out as a chart. Tutors recommended exactly this — write the oxidation numbers above each element, before and after, and the changes become obvious. “Make a chart of oxidation number changes” was direct advice.

Look at INDIVIDUAL atoms, not totals. Tutors flagged this twice. Compare the oxidation number of one iron atom before and after, not the sum for Fe₂O₃.

If nothing changes, it is not a redox reaction. Neutralisation and precipitation are typically not redox.

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5. Oxidising and reducing agents

This is where the most confusion arises, because the naming feels backwards.

An OXIDISING agent OXIDISES something else — so it is itself REDUCED. It GAINS electrons. A REDUCING agent REDUCES something else — so it is itself OXIDISED. It LOSES electrons.

The agent does the opposite to itself. Confusing oxidising with reducing agents was recorded several times. The trick: an oxidising agent gives oxidation to something else, which means it must take electrons, which means it is reduced.

In the iron oxide example above:

• CO is oxidised, so CO is the reducing agent — it reduced the iron
• Fe₂O₃ is reduced, so Fe₂O₃ is the oxidising agent

Common oxidising agents: oxygen, chlorine, acidified potassium manganate(VII) (purple → colourless), potassium dichromate(VI) (orange → green).

Common reducing agents: hydrogen, carbon and carbon monoxide, sulfur dioxide, reactive metals.

Sulfur dioxide is a REDUCING agent. Uncertainty about its role was recorded. It reduces other species while being oxidised itself (S goes from +4 to +6).

Colour changes worth knowing:

Reagent	Change	Indicates
Acidified KMnO₄	purple → colourless	a reducing agent present
Acidified K₂Cr₂O₇	orange → green	a reducing agent present
Aqueous KI	colourless → brown	an oxidising agent present

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6. Half-equations

A half-equation shows one half of the redox process, with the electrons written in.

Oxidation half-equation: electrons on the RIGHT (they are lost) Reduction half-equation: electrons on the LEFT (they are gained)

Examples:

• Oxidation: Mg → Mg²⁺ + 2e⁻
• Reduction: Cl₂ + 2e⁻ → 2Cl⁻

To combine them: multiply so the electrons cancel, then add.

• Mg → Mg²⁺ + 2e⁻
• Cl₂ + 2e⁻ → 2Cl⁻
• Overall: Mg + Cl₂ → MgCl₂

Which side the electrons go on is the whole point. Confusion about electron placement in half-equations was recorded — go back to OIL RIG: if electrons are lost, they are a product, so they go on the right.

Check both the atoms AND the charges balance on each side.

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7. Mistakes that cost marks

Saying oxidation is gain of electrons.

Giving an element in a compound an oxidation number of 0.

Using −2 for oxygen but forgetting to multiply by the number of atoms.

Setting a polyatomic ion’s total to 0 instead of its charge.

Omitting the sign on an oxidation number.

Confusing oxidation number with valency or with the compound’s charge.

Swapping oxidising and reducing agents.

Using totals rather than individual atoms.

Putting electrons on the wrong side of a half-equation.

Assuming no oxygen means no redox.

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Frequently asked questions

What does OIL RIG stand for? Oxidation Is Loss, Reduction Is Gain — of electrons.

What is the oxidation number of an element? Zero — for any uncombined element, including O₂ and Cl₂.

What is oxygen’s usual oxidation number? −2.

What do oxidation numbers add up to? Zero in a neutral compound; the ion’s charge in a polyatomic ion.

How do I find sulfur’s oxidation number in H₂SO₄? Let it be x: 2(+1) + x + 4(−2) = 0, giving +6.

How do I tell what is oxidised? Its oxidation number increases.

What is an oxidising agent? A substance that oxidises something else and is therefore itself reduced — it gains electrons.

Is sulfur dioxide an oxidising or reducing agent? A reducing agent.

What colour change shows a reducing agent with potassium manganate(VII)? Purple to colourless.

Which side do electrons go in an oxidation half-equation? The right — they are lost, so they are products.

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Quick revision checklist

• I know OIL RIG and can apply it
• I know all four pairs of definitions
• I can identify redox without oxygen present
• I know all seven oxidation number rules
• I know uncombined elements are 0
• I always multiply by the number of atoms
• I set polyatomic ions to the ion’s charge
• I write the sign on oxidation numbers
• I don’t confuse oxidation number with valency or charge
• I can build an oxidation number chart for a reaction
• I compare individual atoms, not totals
• I know an oxidising agent is itself reduced
• I know the common agents and their colour changes
• I can write and combine half-equations, with electrons on the right side

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These notes cover redox, oxidation numbers and oxidising agents in the Cambridge IGCSE Chemistry (0620) syllabus and are written for Grade 9–11 / Year 10–11 students. They are based on teaching patterns observed across a large set of one-to-one IGCSE Chemistry lessons, with particular attention to the errors students make most often and the wording examiners reward. Always check the current syllabus and data booklet for your own exam series.