Oxford AQA International A-level · Chemistry (9620)

Oxidation, reduction and redox equations: Practice Questions

5 multiple-choice questions marked as you go, and 5 written questions with worked solutions. All on Oxidation, reduction and redox equations.

10 questions30 marksFree, no account

Question 1

1 mark

What is the oxidation state of chromium in the compound potassium dichromate, \(K_2Cr_2O_7\)?

Question 2

1 mark

The oxidation of ethanedioate ions (oxalate, \(C_2O_4^{2-}\)) and the reduction of manganate(VII) ions (permanganate, \(MnO_4^-\) in acidic conditions are represented by the following half-equations:

\(C_2O_4^{2-} \rightarrow 2CO_2 + 2e^-\)

\(MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O\)

Combine these two half-equations to form the overall balanced redox equation using the smallest whole number coefficients. What is the coefficient for \(H_2O\)?

Question 3

1 mark

Which of the following is the correctly balanced half-equation for the oxidation of sulfite ions (\(SO_3^{2-}\)) to sulfate ions (\(SO_4^{2-}\)) in acidic solution?

Question 4

1 mark

Calculate the average oxidation state of sulfur in the thiosulfate ion, \(S_2O_3^{2-}\).

Question 5

1 mark

In the reaction below, which species acts as the reducing agent?

\(SO_2(g) + 2H_2S(g) \rightarrow 3S(s) + 2H_2O(l)\)

Question 6

4 marks

Consider the redox reaction: \(\text{Zn}(\text{s}) + \text{Cu}^{2+}(\text{aq}) \rightarrow \text{Zn}^{2+}(\text{aq}) + \text{Cu}(\text{s})\). Identify the oxidising agent and the reducing agent, and state the half-equation for the process occurring at the zinc electrode.

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Question 7

5 marks

The half-equation for the reduction of dichromate(VI) ions to chromium(III) ions in acid is \(\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ + 6\text{e}^- \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}\). Combine this with the oxidation of \(\text{Fe}^{2+}\) to \(\text{Fe}^{3+}\) to derive the overall balanced ionic equation for the reaction.

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Question 8

3 marks

Write a balanced half-equation for the reduction of chlorine gas, \(\text{Cl}_2\), in aqueous solution.

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Question 9

8 marks

Vanadium exhibits several stable oxidation states in aqueous solution, visible as different colours: \( \text{V}(\text{V}) \) (yellow), \( \text{V}(\text{IV}) \) (blue), \( \text{V}(\text{III}) \) (green), and \( \text{V}(\text{II}) \) (violet). The reduction is typically carried out using zinc metal in acidic solution.

(a) Determine the oxidation state of Vanadium in the following ions:
(i) Vanadate(V) ion, \( \text{VO}_2^+ \)
(ii) Vanadium(IV) ion, \( \text{VO}^{2+} \)
(b) Write the balanced half-equation for the reduction of \( \text{VO}_2^+ \) to \( \text{VO}^{2+} \) in acidic solution.
(c) Write the balanced half-equation for the subsequent reduction of \( \text{VO}^{2+} \) to \( \text{V}^{3+} \) in acidic solution.
(d) Write the overall ionic equation for the complete reduction of \( \text{VO}_2^+ \) (V) to \( \text{V}^{2+} \) (II) using zinc (\[ \text{Zn} \]) as the reducing agent.

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Question 10

5 marks

Chlorine gas reacts with cold, dilute aqueous sodium hydroxide solution. This reaction involves the oxidation and reduction of the same element, Chlorine.

(a) Write the half-equation for the oxidation of chlorine to chlorate(I) ions (\[ \text{ClO}^- \]).
(b) Write the half-equation for the reduction of chlorine to chloride ions (\[ \text{Cl}^- \]).
(c) Combine the two half-equations to give the overall equation for the reaction.
(d) What term is used to describe a reaction in which an element is simultaneously oxidised and reduced? Explain how the oxidation states of chlorine confirm that this process occurs.

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