Chemistry • Year 11 • Module 3 • Lesson 8

Redox Reactions & Oxidation States

Apply oxidation state assignment, half-equation writing and OIL RIG reasoning to real data, new reactions and an Australian corrosion context.

Apply • Data & Reasoning

1. Interpret graph — chlorine oxidation state vs corrosion rate in bleach formulations

The graph below shows the average percentage of iron surface corroded after 24 hours when iron coupons were immersed in bleach solutions at different hypochlorite ion (OCl⁻) concentrations. OCl⁻ is the active oxidising species in bleach; it accepts electrons from iron and is itself reduced to Cl⁻. 7 marks

Illustrative experimental data. After Levine & Rhee, Corrosion Science (adapted).

1.1 Describe the overall trend shown by the data. 2 marks

1.2 The oxidation state of chlorine in OCl⁻ is +1 and in Cl⁻ is −1. Using OIL RIG, explain why OCl⁻ is classified as the oxidant (oxidising agent) in this reaction with iron. 3 marks

1.3 Estimate the percentage of iron corroded at 2.5% OCl⁻ concentration. Explain whether this is interpolation or extrapolation. 2 marks

Stuck? Revisit lesson § Card 5 (bleach chemistry) and Card 3 (oxidation number changes). Remember: a species is an oxidant if it accepts electrons — its oxidation number decreases.

2. Compare and contrast — oxidation and reduction

Complete the two-column table comparing oxidation and reduction across five features. Use precise lesson terminology in every cell. 8 marks

Feature	Oxidation	Reduction
Electron movement
Change in oxidation number
Name of the species undergoing this process
Where electrons appear in the half-equation
Real-world example species (from bleach or Fe/Cu reactions)

Stuck? Use OIL RIG as your spine, then add detail from Cards 3–5.

3. Cause-and-effect chain — zinc displacing copper

When a piece of zinc metal is placed in a blue copper(II) sulfate solution, the zinc slowly disappears and a reddish copper solid coats the zinc. The overall ionic equation is: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). Trace the cause-and-effect chain below using redox reasoning. Fill each empty effect box. 5 marks

Cause 1:

Zinc metal (Zn, oxidation state 0) comes into contact with copper(II) ions (Cu²⁺, oxidation state +2) in aqueous solution.

↓

Effect 1 / Cause 2:

Describe what happens to zinc’s oxidation state and identify whether zinc is oxidised or reduced. Name the role zinc plays (oxidant or reductant).

↓

Effect 2 / Cause 3:

Describe what happens to copper’s oxidation state and identify whether Cu²⁺ is oxidised or reduced. Name the role Cu²⁺ plays (oxidant or reductant).

↓

Overall outcome (so…):

State the products that form and explain why this reaction demonstrates the electron-transfer definition of a redox reaction (OIL RIG), rather than merely a definition involving oxygen.

Stuck? Assign oxidation numbers to Zn and Cu on both sides of the equation, then use OIL RIG to classify each species. This reaction is the primary example in Card 2 of the lesson.

4. Apply to a new scenario — sacrificial anodes on Sydney Harbour vessels

Ships docking at the Port of Sydney fit zinc sacrificial anodes to their steel hulls. When a zinc anode contacts the iron hull in seawater, a spontaneous redox reaction occurs: Zn is oxidised preferentially, protecting the iron from corrosion. The relevant half-equations are:
Oxidation: Zn(s) → Zn²⁺(aq) + 2e⁻
Reduction: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
6 marks

4.1 Identify the oxidant and the reductant in this corrosion-protection system. Justify each identification using oxidation number changes. 3 marks

4.2 Combine the two half-equations above to give the balanced overall ionic equation. Show how you balance the electron count. 3 marks

Stuck? For 4.2: multiply the oxidation half-equation by 2 so both half-equations involve 4e⁻, then add and cancel electrons.

Answers — Do not peek before attempting

Q1.1 — Trend description

As OCl⁻ concentration increases from 0% to 5%, the percentage of iron surface corroded after 24 hours increases from approximately 3% to 78%. The relationship is non-linear (the rate of corrosion accelerates at higher concentrations), suggesting that higher oxidant concentration increases the rate of electron transfer from iron to OCl⁻.

Q1.2 — OCl⁻ as oxidant (3 marks)

The oxidation state of chlorine changes from +1 (in OCl⁻) to −1 (in Cl⁻) — a decrease of 2 [1]. According to OIL RIG, a decrease in oxidation number indicates reduction (Reduction Is Gain of electrons) [1]. The species that is reduced is by definition the oxidant (oxidising agent) because it gains electrons from iron, causing iron to be oxidised [1].

Q1.3 — Estimation at 2.5%

Reading between the 2% (27%) and 3% (50%) data points, a reasonable estimate is approximately 38–40% iron corroded at 2.5% OCl⁻ [1]. This is interpolation because the value is estimated within the range of measured data points (0–5%), not beyond it [1].

Q2 — Compare and contrast table

Feature	Oxidation	Reduction
Electron movement	Electrons are lost (leave the species)	Electrons are gained (enter the species)
Change in oxidation number	Increases (becomes more positive)	Decreases (becomes more negative)
Name of the species	Reductant (it causes the other to be reduced)	Oxidant (it causes the other to be oxidised)
Electron position in half-equation	Products (right side): e.g. Zn → Zn²♠ + 2e²	Reactants (left side): e.g. Cu²♠ + 2e² → Cu
Real-world example	Fe(s) or Zn(s) (both lose electrons to an oxidant)	OCl² in bleach / O&sub2; in rusting (both gain electrons)

Q3 — Cause-and-effect: zinc displacing copper

Effect 1 / Cause 2: Zinc’s oxidation state increases from 0 (Zn metal) to +2 (Zn²⁺ in solution). An increase in oxidation number indicates oxidation (OIL RIG: Oxidation Is Loss). Zinc is therefore oxidised; it is the reductant — it donates 2 electrons to the copper(II) ions.

Effect 2 / Cause 3: Copper’s oxidation state decreases from +2 (Cu²⁺) to 0 (Cu metal). A decrease in oxidation number indicates reduction (RIG: Reduction Is Gain). Cu²⁺ is therefore reduced; it is the oxidant — it accepts the 2 electrons from zinc.

Overall outcome: The products are Zn²⁺(aq) (zinc ions dissolving into solution) and Cu(s) (copper metal depositing on the zinc surface). This demonstrates the electron-transfer definition of a redox reaction: electrons are transferred from Zn to Cu²⁺. No oxygen gas is involved at all — proving that redox is defined by electron transfer (OIL RIG), not by the presence of oxygen.

Q4.1 — Oxidant and reductant (3 marks)

Zn: oxidation state changes from 0 (Zn metal) to +2 (Zn²⁺) — an increase → oxidised → reductant [1].
O₂: oxidation state changes from 0 (O₂) to −2 (in OH⁻) — a decrease → reduced → oxidant [1].
Justification: OIL RIG; increasing oxidation number = loss of electrons = oxidation = reductant; decreasing = gain of electrons = reduction = oxidant [1].

Q4.2 — Overall ionic equation (3 marks)

Oxidation half (involves 2e⁻): Zn → Zn²⁺ + 2e⁻. Multiply ×2 → 2Zn → 2Zn²⁺ + 4e⁻ [1].
Reduction half (involves 4e⁻): O₂ + 2H₂O + 4e⁻ → 4OH⁻ — already has 4e⁻ [1].
Cancel 4e⁻ and add: 2Zn(s) + O₂(g) + 2H₂O(l) → 2Zn²⁺(aq) + 4OH⁻(aq).
Charge check: left = 0; right = 2(+2) + 4(−1) = 4 − 4 = 0. ✓ [1]