SCH4U Oxidation Numbers and Redox Reactions: Notes and Practice

A steel bike left out in the rain rusts. A phone battery powers a screen. Natural gas burns on a stove. These look like very different events, but chemically they are the same kind of reaction: electrons move from one substance to another. Chemists call this an oxidation-reduction reaction, or redox reaction for short. In Grade 11 you met single displacement reactions, like zinc metal pushing copper out of a copper(II) sulfate solution. This unit explains what is really happening there, and gives you tools to handle much harder cases.

Oxidation and Reduction as Electron Transfer

Drop a strip of zinc into blue copper(II) sulfate solution and a reddish-brown coating of copper forms on the zinc while the blue colour fades. The sulfate ions just watch; they are spectator ions. The real change is between zinc atoms and copper(II) ions. Each zinc atom loses two electrons and becomes a $Zn^{2 +}$ ion. Each $Cu^{2 +}$ ion gains two electrons and becomes a copper atom.

Zn(s) + Cu^{2 +} (aq) \to Zn^{2 +} (aq) + Cu(s)

Net ionic equation: zinc gives two electrons to each copper(II) ion.

Oxidation is the loss of electrons. Zinc is oxidized: $Zn \to Zn^{2 +} + 2 e^{-}$ .
Reduction is the gain of electrons. Copper(II) is reduced: $Cu^{2 +} + 2 e^{-} \to Cu$ .
The two always happen together. Electrons cannot just vanish, so every electron lost by one substance must be gained by another. That is why the reaction is called redox, never just "ox".

Memory aid: LEO says GER

Loss of Electrons is Oxidation; Gain of Electrons is Reduction. Some people prefer OIL RIG: Oxidation Is Loss, Reduction Is Gain. Reduction is named for the number that goes down: the charge is reduced when negative electrons are added.

Oxidation Numbers: Bookkeeping for Electrons

With ions, it is easy to see who lost electrons. But in a reaction like methane burning, $CH_{4} + 2 O_{2} \to CO_{2} + 2 H_{2} O$ , every substance is molecular and no ions form. Electrons still shift, though: in $CH_{4}$ carbon holds the larger share of its bonding electrons (it is more electronegative than hydrogen), but in $CO_{2}$ the more electronegative oxygen atoms pull the shared electrons away from carbon. An oxidation number is a bookkeeping charge that captures this. It is the charge an atom would have if every bond in the substance were fully ionic, with the shared electrons given to the more electronegative atom. It is not a real charge, but tracking how it changes tells you exactly where electrons went.

Rule	Examples
An atom in an uncombined element is 0.	Na, $O_{2}$ , $P_{4}$ , $S_{8}$ all 0
A monatomic ion equals its charge.	$Na^{+}$ is +1, $S^{2 -}$ is −2
Group 1 metals are +1 and group 2 metals are +2 in compounds.	K in KCl is +1, Ca in $CaO$ is +2
Fluorine is always −1 in compounds.	F in $HF$ and $OF_{2}$ is −1
Hydrogen is usually +1, but −1 when bonded to a metal (a metal hydride).	H in $H_{2} O$ is +1, H in NaH is −1
Oxygen is usually −2, but −1 in peroxides and +2 in $OF_{2}$ .	O in $H_{2} O_{2}$ is −1
The oxidation numbers add up to 0 in a neutral compound, or to the ion's charge in a polyatomic ion.	In $SO_{4}^{2 -}$ : S + 4(−2) = −2

Apply the rules from top to bottom; a rule higher in the list wins when two conflict.

To find an unknown oxidation number, assign the atoms you know from the rules, then solve for the one that is left so the total matches the overall charge. For permanganate, $MnO_{4}^{-}$ , four oxygen atoms contribute $4 (- 2) = - 8$ , and the whole ion has a charge of $- 1$ :

x + 4 (- 2) = - 1 \Rightarrow x = + 7 (Mn in MnO_{4}^{-})

Solving for manganese's oxidation number in the permanganate ion.

Spotting Redox and Naming the Agents

If an atom's oxidation number increases from reactant to product, that atom was oxidized (it lost electrons).
If an atom's oxidation number decreases, that atom was reduced (it gained electrons).
The oxidizing agent is the reactant that causes oxidation by taking electrons, so the oxidizing agent is itself reduced.
The reducing agent is the reactant that causes reduction by giving electrons away, so the reducing agent is itself oxidized.
Name the whole reactant as the agent: in methane combustion the reducing agent is $CH_{4}$ , not just "carbon".
If no oxidation number changes, the reaction is not redox. Precipitation and acid-base neutralization reactions are not redox.

Fig 1The activity series. A metal higher on the list gives up electrons more easily, so it is a stronger reducing agent and can reduce the ions of any metal below it.

Balancing by the Oxidation-Number Method

Assign oxidation numbers to every atom and find the ones that change.
Work out the electrons lost per formula of the substance oxidized and gained per formula of the substance reduced. Multiply by the number of those atoms in the formula.
Choose coefficients so that total electrons lost equals total electrons gained (use the lowest common multiple). Once that ratio is fixed, the hard part is done.
Balance the remaining atoms by inspection, usually leaving H and O for last.
Check that every atom balances. For an ionic equation, check that the total charge balances too.

Oxidation Numbers and Redox Reactions

Key things to remember

Oxidation and Reduction as Electron Transfer

Oxidation Numbers: Bookkeeping for Electrons

Spotting Redox and Naming the Agents

Balancing by the Oxidation-Number Method

Worked examples

Where marks get lost

Flashcards 8

Practice