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O Level Chemistry Redox Electrochemistry Quiz

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O Level Chemistry AI Generated Generated by Qwen3.6 Plus Updated 2026-08-17

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Questions

O-Level Chemistry Quiz - Redox Electrochemistry

Name: __________________________
Class: __________________________
Date: __________________________
Score: ________ / 40

Duration: 45 minutes
Total Marks: 40

Instructions:

Answer all questions.
Write your answers in the spaces provided.
The number of marks is given in brackets [ ] at the end of each question or part question.
You may use a calculator.

Section A: Multiple Choice & Short Concepts (Questions 1–5)

Marks: 1 mark each

1. Which statement correctly defines oxidation in terms of electron transfer? A. Gain of electrons B. Loss of electrons C. Gain of hydrogen D. Loss of oxygen

Answer: __________________________ [1]

2. In the reaction below, which species acts as the reducing agent? $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$

A. $Zn(s)$ B. $Cu^{2+}(aq)$ C. $Zn^{2+}(aq)$ D. $Cu(s)$

Answer: __________________________ [1]

3. What is the oxidation state of chromium in the dichromate ion, $Cr_2O_7^{2-}$ ?

Answer: __________________________ [1]

4. During the electrolysis of molten lead(II) bromide ( $PbBr_2$ ) using inert electrodes, what is observed at the anode?

Answer: __________________________ [1]

5. Which of the following ions is preferentially discharged at the cathode during the electrolysis of dilute aqueous sodium sulfate ( $Na_2SO_4$ )?

Answer: __________________________ [1]

Section B: Electrolysis of Aqueous Solutions (Questions 6–10)

Marks: 2–3 marks each

6. Aqueous copper(II) sulfate is electrolysed using carbon (inert) electrodes. (a) Write the ionic half-equation for the reaction at the cathode.

_________________________________________________________________________ [1]

(b) Explain why the blue colour of the solution fades during this electrolysis.

_________________________________________________________________________ [2]

7. The same aqueous copper(II) sulfate solution is now electrolysed using copper electrodes. (a) State the observation at the anode.

_________________________________________________________________________ [1]

(b) Write the half-equation for the reaction at the anode.

_________________________________________________________________________ [1]

(c) Explain why the concentration of copper(II) ions in the solution remains constant.

_________________________________________________________________________ [2]

8. Concentrated aqueous sodium chloride (brine) is electrolysed using inert electrodes. (a) Name the product formed at the anode.

_________________________________________________________________________ [1]

(b) Explain why this product is formed instead of oxygen, referring to the concentration of ions.

_________________________________________________________________________ [2]

9. Dilute sulfuric acid is electrolysed using platinum electrodes. This is effectively the electrolysis of water. (a) Give the volume ratio of the gas produced at the cathode to the gas produced at the anode.

Cathode : Anode = _______ : _______ [1]

(b) Write the half-equation for the formation of the gas at the anode.

_________________________________________________________________________ [2]

10. A student electrolyses aqueous silver nitrate using inert electrodes. (a) Predict the product at the cathode.

_________________________________________________________________________ [1]

(b) Explain your answer using the reactivity series.

_________________________________________________________________________ [2]

Section C: Redox Reactions & Oxidation States (Questions 11–15)

Marks: 2–3 marks each

11. Chlorine gas is bubbled through aqueous potassium iodide. (a) State the observation.

_________________________________________________________________________ [1]

(b) Write the ionic equation for this reaction.

_________________________________________________________________________ [2]

(c) Identify the oxidising agent in this reaction.

_________________________________________________________________________ [1]

12. Iron(II) sulfate solution is added to acidified potassium manganate(VII) solution. (a) State the colour change observed.

From ____________________ to ____________________ [1]

(b) Explain this change in terms of electron transfer involving the iron ions.

_________________________________________________________________________ [2]

13. Determine the oxidation state of the underlined element in each compound: (a) $\underline{N}H_4^+$ : ____________________ [1] (b) $K\underline{Mn}O_4$ : ____________________ [1] (c) $\underline{S}O_3$ : ____________________ [1]

14. Hydrogen peroxide ( $H_2O_2$ ) can act as both an oxidising and a reducing agent. Reaction: $2Fe^{2+} + H_2O_2 + 2H^+ \rightarrow 2Fe^{3+} + 2H_2O$ (a) Show, using oxidation states, that iron is oxidised in this reaction.

_________________________________________________________________________ [1]

(b) Show, using oxidation states, that oxygen in hydrogen peroxide is reduced.

_________________________________________________________________________ [2]

15. Magnesium ribbon is added to dilute hydrochloric acid. $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$ (a) Identify the substance that is oxidised.

_________________________________________________________________________ [1]

(b) Explain why this is a redox reaction by identifying the change in oxidation state for hydrogen.

_________________________________________________________________________ [2]

Section D: Cells and Applications (Questions 16–20)

Marks: 2–3 marks each

16. A simple cell is set up using a magnesium strip and a copper strip dipped in dilute sulfuric acid, connected by a wire and a voltmeter. (a) Which metal acts as the negative electrode (anode)?

_________________________________________________________________________ [1]

(b) Explain your answer in terms of the reactivity series.

_________________________________________________________________________ [2]

17. In the cell described in Question 16: (a) Write the half-equation for the reaction occurring at the copper electrode.

_________________________________________________________________________ [2]

(b) State the direction of electron flow in the external circuit.

From ____________________ to ____________________ [1]

18. Hydrogen fuel cells are used in some vehicles. (a) Write the overall chemical equation for the reaction in a hydrogen fuel cell.

_________________________________________________________________________ [1]

(b) State one advantage of using hydrogen fuel cells compared to petrol engines, other than cost.

_________________________________________________________________________ [1]

(c) State one disadvantage of using hydrogen fuel cells.

_________________________________________________________________________ [1]

19. Iron objects can be protected from rusting by sacrificial protection. (a) Name a metal that can be attached to iron to provide sacrificial protection.

_________________________________________________________________________ [1]

(b) Explain how this metal protects the iron, referring to electron transfer.

_________________________________________________________________________ [3]

20. Electroplating is used to coat a steel spoon with silver. (a) Should the steel spoon be the anode or the cathode?

_________________________________________________________________________ [1]

(b) Suggest a suitable electrolyte for this process.

_________________________________________________________________________ [1]

(c) Write the half-equation for the reaction that coats the spoon.

_________________________________________________________________________ [2]

*** End of Quiz ***

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Answers

O-Level Chemistry Quiz - Redox Electrochemistry (Answer Key)

1. B
Loss of electrons is oxidation (OIL RIG). [1]

2. A
Zn loses electrons to form $Zn^{2+}$ . The species that loses electrons is the reducing agent. [1]

3. +6
Let Cr be $x$ . $2x + 7(-2) = -2 \Rightarrow 2x - 14 = -2 \Rightarrow 2x = 12 \Rightarrow x = +6$ . [1]

4. Reddish-brown vapour / Brown gas
Bromide ions ( $Br^-$ ) are oxidised to bromine ( $Br_2$ ) at the anode. [1]

5. Hydrogen ion ( $H^+$ )
In dilute aqueous solutions, $H^+$ is preferentially discharged over $Na^+$ because hydrogen is lower in the reactivity series. [1]

6.
(a) $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$ [1]
(b) Copper(II) ions ( $Cu^{2+}$ ) are removed from the solution to form copper metal at the cathode. As the concentration of blue $Cu^{2+}$ ions decreases, the colour fades. [2]

7.
(a) The copper anode decreases in size / dissolves. [1]
(b) $Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$ [1]
(c) For every $Cu^{2+}$ ion discharged at the cathode, one Cu atom from the anode dissolves to form a $Cu^{2+}$ ion. The rate of removal equals the rate of formation, so concentration stays constant. [2]

8.
(a) Chlorine gas ( $Cl_2$ ). [1]
(b) Although $OH^-$ is lower in the electrochemical series than $Cl^-$ , the concentration of $Cl^-$ ions is much higher in concentrated solution. Therefore, $Cl^-$ is preferentially discharged. [2]

9.
(a) 2 : 1 [1]
(b) $4OH^-(aq) \rightarrow O_2(g) + 2H_2O(l) + 4e^-$
OR $2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-$ [2]

10.
(a) Silver ( $Ag$ ). [1]
(b) Silver is lower in the reactivity series than hydrogen. Therefore, $Ag^+$ ions are preferentially discharged over $H^+$ ions at the cathode. [2]

11.
(a) The colourless solution turns brown / dark brown. [1]
(b) $Cl_2 + 2I^- \rightarrow 2Cl^- + I_2$ [2]
(c) Chlorine ( $Cl_2$ ). [1]

12.
(a) From purple to colourless. [1]
(b) Iron(II) ions ( $Fe^{2+}$ ) lose electrons to form Iron(III) ions ( $Fe^{3+}$ ). Loss of electrons is oxidation. The $MnO_4^-$ accepts these electrons and is reduced. [2]

13.
(a) -3 [ $x + 4(+1) = +1 \Rightarrow x = -3$ ] [1]
(b) +7 [ $1 + x + 4(-2) = 0 \Rightarrow x = +7$ ] [1]
(c) +6 [ $x + 3(-2) = 0 \Rightarrow x = +6$ ] [1]

14.
(a) Fe changes from +2 in $Fe^{2+}$ to +3 in $Fe^{3+}$ . Increase in oxidation state is oxidation. [1]
(b) O changes from -1 in $H_2O_2$ to -2 in $H_2O$ . Decrease in oxidation state is reduction. [2]

15.
(a) Magnesium ( $Mg$ ). [1]
(b) Hydrogen changes from +1 in $HCl$ to 0 in $H_2$ . The decrease in oxidation state indicates reduction. Since reduction occurs, it is a redox reaction. [2]

16.
(a) Magnesium. [1]
(b) Magnesium is more reactive than copper. It has a greater tendency to lose electrons and form ions, making it the negative electrode (source of electrons). [2]

17.
(a) $2H^+(aq) + 2e^- \rightarrow H_2(g)$ [2]
(b) From Magnesium to Copper. [1]

18.
(a) $2H_2 + O_2 \rightarrow 2H_2O$ [1]
(b) The only product is water (non-polluting) / No carbon dioxide produced / High efficiency. [1]
(c) Hydrogen is difficult/expensive to store / Lack of refuelling infrastructure / Hydrogen is flammable/explosive. [1]

19.
(a) Zinc (or Magnesium). [1]
(b) Zinc is more reactive than iron. Zinc loses electrons more easily than iron. The zinc oxidises (corrodes) in preference to the iron. Electrons flow from zinc to iron, preventing the iron from losing electrons and rusting. [3]

20.
(a) Cathode. [1]
(b) Silver nitrate solution ( $AgNO_3(aq)$ ) / Any soluble silver salt. [1]
(c) $Ag^+(aq) + e^- \rightarrow Ag(s)$ [2]