Secondary 3 Chemistry Redox Electrochemistry Quiz

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Secondary 3 Chemistry Quiz - Redox Electrochemistry

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

Duration: 45 minutes
Total Marks: 40

Instructions:

1. Answer all questions.
2. Write your answers in the spaces provided.
3. For calculations, show all working clearly.
4. The number of marks is given in brackets [ ] at the end of each question or part question.

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

Answer all questions in this section.

1. Which of the following statements correctly defines oxidation in terms of electron transfer? [1] A. Gain of electrons B. Loss of electrons C. Gain of hydrogen D. Loss of oxygen

2. In the reaction below, which species acts as the oxidising agent? [1] $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$ A. $Zn(s)$ B. $Cu^{2+}(aq)$ C. $Zn^{2+}(aq)$ D. $Cu(s)$

3. What is the oxidation state of manganese (Mn) in potassium manganate(VII), $KMnO_4$? [1] A. +2 B. +5 C. +6 D. +7

4. Which of the following changes represents a reduction process? [1] A. $Fe^{2+} \rightarrow Fe^{3+}$ B. $Cl_2 \rightarrow 2Cl^-$ C. $Mg \rightarrow Mg^{2+}$ D. $SO_2 \rightarrow SO_3$

5. In a simple chemical cell consisting of magnesium and copper strips dipped in dilute sulfuric acid, which electrode is the negative terminal? [1] A. Magnesium, because it loses electrons. B. Magnesium, because it gains electrons. C. Copper, because it loses electrons. D. Copper, because it gains electrons.

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Section B: Redox Concepts & Oxidation States (Questions 6–10)

Answer all questions in this section.

6. Consider the reaction between hydrogen sulfide and chlorine gas: $H_2S(g) + Cl_2(g) \rightarrow 2HCl(g) + S(s)$

(a) Identify the substance that is oxidised. [1]

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(b) Explain your answer to (a) in terms of hydrogen transfer. [1]

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(c) Identify the reducing agent in this reaction. [1]

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7. Determine the oxidation state of the specified element in each of the following compounds. [3]

(a) Chromium in $Cr_2O_3$: _______________

(b) Sulfur in $SO_4^{2-}$: _______________

(c) Nitrogen in $NH_4^+$: _______________

8. The following equation represents the reaction of iron(III) oxide with carbon monoxide in a blast furnace: $Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(s) + 3CO_2(g)$

(a) State the change in oxidation state of iron in this reaction. [1] From _______________ to _______________

(b) Explain why carbon monoxide is considered a reducing agent in this reaction. [1]

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9. A student adds aqueous potassium iodide to aqueous chlorine. The solution turns brown.

(a) Write the ionic equation for this reaction. [2]

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(b) State the colour change observed if starch solution is added to the final mixture. [1]

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10. Define the term disproportionation. [1]

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Section C: Electrolysis (Questions 11–15)

Answer all questions in this section.

11. Molten lead(II) bromide, $PbBr_2$, is electrolysed using inert graphite electrodes.

(a) Name the product formed at the cathode. [1]

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(b) Write the half-equation for the reaction occurring at the anode. [2]

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(c) State one observation at the anode. [1]

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12. Concentrated aqueous sodium chloride (brine) is electrolysed using inert electrodes.

(a) List the ions present in the solution. [2] Cations: _________________________ Anions: _________________________

(b) Predict the product formed at the anode and explain why it is preferentially discharged. [2] Product: _________________________ Explanation: _________________________________________________________________

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13. Dilute sulfuric acid is electrolysed using platinum electrodes. This is effectively the electrolysis of water.

(a) Write the half-equation for the formation of oxygen gas at the anode. [2]

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(b) What is the ratio of the volume of gas collected at the cathode to the volume of gas collected at the anode? [1]

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14. Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does. [2]

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15. In the electrolysis of aqueous copper(II) sulfate using carbon electrodes, the blue colour of the solution fades over time. Explain why. [2]

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Section D: Electrochemical Cells & Applications (Questions 16–20)

Answer all questions in this section.

16. A simple cell is set up using a zinc strip and an iron strip dipped in their respective sulfate solutions, connected by a salt bridge and a voltmeter.

(a) Which metal acts as the negative terminal (anode)? [1]

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(b) Write the half-equation for the reaction occurring at the negative terminal. [1]

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(c) If the zinc strip is replaced by a magnesium strip, will the voltmeter reading increase or decrease? Explain. [2]

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17. Describe the process of electroplating a steel spoon with silver.

(a) What material should be used for the anode? [1]

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(b) What electrolyte should be used? [1]

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(c) Write the half-equation for the reaction at the cathode (the spoon). [1]

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18. Copper purification by electrolysis uses impure copper as the anode and pure copper as the cathode.

(a) What happens to the mass of the anode during the process? [1]

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(b) What happens to the impurities (such as gold and silver) that are less reactive than copper? [1]

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19. Hydrogen fuel cells are an alternative to internal combustion engines.

(a) Write the overall chemical equation for the reaction in a hydrogen fuel cell. [1]

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(b) State one advantage of using hydrogen fuel cells over petrol engines. [1]

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20. A student sets up an electrolysis experiment to copper-plate a key. However, they connect the key to the positive terminal and the copper strip to the negative terminal.

(a) Describe what will happen to the key. [1]

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(b) Describe what will happen to the copper strip. [1]

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*** End of Quiz ***

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Secondary 3 Chemistry Quiz - Redox Electrochemistry (Answer Key)

Total Marks: 40

Section A: Multiple Choice & Definitions

1. B [1]

• Explanation: Oxidation is the loss of electrons (OIL RIG).

2. B [1]

• Explanation: The oxidising agent accepts electrons and is itself reduced. $Cu^{2+}$ gains electrons to become $Cu$.

3. D [1]

• Explanation: $K (+1) + Mn (x) + 4 \times O (-2) = 0 \Rightarrow 1 + x - 8 = 0 \Rightarrow x = +7$.

4. B [1]

• Explanation: Reduction is gain of electrons. $Cl_2$ (oxidation state 0) gains electrons to form $Cl^-$ (oxidation state -1).

5. A [1]

• Explanation: Magnesium is more reactive than copper. It loses electrons more readily ($Mg \rightarrow Mg^{2+} + 2e^-$), making it the negative terminal (source of electrons).

Section B: Redox Concepts & Oxidation States

6. (a) $H_2S$ [1] (b) $H_2S$ loses hydrogen to form $S$. [1] (c) $H_2S$ [1]

• Note: The reducing agent is the substance that is oxidised.

7. (a) +3 [1] ($2x + 3(-2) = 0 \Rightarrow 2x = 6$) (b) +6 [1] ($x + 4(-2) = -2 \Rightarrow x - 8 = -2$) (c) -3 [1] ($x + 4(+1) = +1 \Rightarrow x + 4 = 1$)

8. (a) From +3 to 0 [1] (b) Carbon monoxide removes oxygen from iron(III) oxide (or CO is oxidised to $CO_2$). [1]

9. (a) $Cl_2 + 2I^- \rightarrow 2Cl^- + I_2$ [2]

• 1 mark for correct reactants/products, 1 mark for balancing. (b) Blue-black / Dark blue [1]

10. Disproportionation is a redox reaction in which the same element is simultaneously oxidised and reduced. [1]

Section C: Electrolysis

11. (a) Lead [1] (b) $2Br^- \rightarrow Br_2 + 2e^-$ [2]

• 1 mark for correct species, 1 mark for balancing/charges. (c) Brown vapour / Brown gas produced [1]

12. (a) Cations: $Na^+, H^+$ [1] Anions: $Cl^-, OH^-$ [1] (b) Product: Chlorine ($Cl_2$) [1] Explanation: In concentrated solution, chloride ions ($Cl^-$) are preferentially discharged over hydroxide ions ($OH^-$) due to their higher concentration. [1]

13. (a) $4OH^- \rightarrow O_2 + 2H_2O + 4e^-$ [2] * Alternative accepted: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$ (b) 2 : 1 [1] * Hydrogen (cathode) : Oxygen (anode)

14. In solid NaCl, the ions are held in fixed positions in a lattice and cannot move to carry charge. [1] In molten NaCl, the lattice breaks down and the ions are free to move and carry electrical charge. [1]

15. Copper(II) ions ($Cu^{2+}$) are discharged at the cathode to form copper metal. [1] They are not replaced at the anode because inert carbon electrodes are used (oxygen is produced instead), so the concentration of $Cu^{2+}$ decreases. [1]

Section D: Electrochemical Cells & Applications

16. (a) Zinc [1] (b) $Zn \rightarrow Zn^{2+} + 2e^-$ [1] (c) Increase [1] Explanation: Magnesium is more reactive than zinc (further from copper in the reactivity series), creating a larger difference in potential/voltage. [1]

17. (a) Silver [1] (b) Silver nitrate solution / $AgNO_3(aq)$ [1] * Must contain silver ions. (c) $Ag^+ + e^- \rightarrow Ag$ [1]

18. (a) Decreases [1] (b) They fall to the bottom of the cell as anode sludge / sediment. [1]

19. (a) $2H_2 + O_2 \rightarrow 2H_2O$ [1] (b) Only water is produced (no pollution / no greenhouse gases) OR Higher efficiency. [1]

20. (a) The key will dissolve / oxidise / lose mass (if it is reactive) or bubbles of oxygen may form if inert, but typically plating fails. Accept: No copper plating occurs; key may corrode. [1] * Specifically: If key is inert, $OH^-$ discharges. If key is reactive metal, it oxidises. (b) Copper ions from the solution will plate onto the copper strip (cathode). The copper strip gains mass. [1]