Secondary 3 Chemistry Redox Electrochemistry Quiz

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Secondary 3 Chemistry Quiz - Redox Electrochemistry

Name: ____________________ Class: __________ Date: __________ Score: ________ / 50

Duration: 60 Minutes
Total Marks: 50
Instructions: Answer all questions are compulsory. Use a pencil for diagrams and a pen for writing. Show all working for calculations.

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Section A: Fundamentals of Redox (Questions 1–7)

1. Define "oxidation" in terms of the gain or loss of oxygen. [1]
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2. In the reaction $\text{CuO(s)} + \text{H}_2\text{(g)} \rightarrow \text{Cu(s)} + \text{H}_2\text{O(l)}$, identify the substance that is being reduced. [1]
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3. Define a "reducing agent". [1]
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4. For the following reaction, identify the species that is oxidized and the species that is reduced: $\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}$ [2]
   Oxidized: ____________________ Reduced: ____________________

5. Assign the oxidation states of the underlined elements in the following compounds: [2] (a) $\underline{\text{K}}\text{MnO}_4$: __________ (b) $\text{H}_2\underline{\text{S}}\text{O}_4$: __________

6. Explain why the reaction between magnesium and dilute hydrochloric acid is a redox reaction. [2]
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7. A student adds potassium manganate(VII) solution to a solution of iron(II) ions. The purple color of the solution disappears. (a) Is the iron(II) ion acting as an oxidizing agent or a reducing agent? [1] \

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   (b) State the change in oxidation state of the manganese atom. [1] \

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Section B: Electrolysis of Molten and Aqueous Compounds (Questions 8–14)

8. State the name of the electrode where oxidation always occurs. [1]
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9. Describe what happens to the ions at the cathode during the electrolysis of molten lead(II) bromide. [2]
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10. In the electrolysis of aqueous sodium chloride using inert electrodes: (a) Name the gas produced at the anode. [1] \

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    (b) Explain why hydrogen gas is produced at the cathode instead of sodium metal. [2]
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11. Predict the products at the anode and cathode during the electrolysis of aqueous copper(II) sulfate using carbon electrodes. [2]
    Anode: ____________________ Cathode: ____________________

12. Write the ionic half-equation for the reaction occurring at the cathode during the electrolysis of molten aluminium oxide. [2]
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13. During the electrolysis of aqueous $\text{KI}$, a brown solution forms around the anode. (a) Identify the substance responsible for the brown color. [1] \

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    (b) Write the half-equation for the reaction at the anode. [2]
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14. Explain why the electrolysis of molten compounds is generally simpler to predict than the electrolysis of aqueous compounds. [2]
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Section C: Applications of Electrochemistry (Questions 15–20)

15. In the industrial purification of copper: (a) Which electrode is the impure copper? (Anode or Cathode) [1] \

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    (b) Describe the change in mass of the anode over time. [1] \

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16. A student wants to electroplate a steel spoon with silver. (a) What material should be used for the anode? [1] \

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

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17. Describe the function of the electrolyte in a simple chemical cell. [2]
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18. In a simple cell made of zinc and copper electrodes in $\text{H}_2\text{SO}_4$ solution: (a) Which metal acts as the negative terminal? [1] \

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    (b) Explain your answer in terms of the reactivity series. [2]
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19. State one advantage of a hydrogen fuel cell over a traditional combustion engine. [1]
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20. Write the overall cell equation for a hydrogen fuel cell. [2]
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Answer Key - Secondary 3 Chemistry Quiz: Redox Electrochemistry

Marking Scheme & Explanations

1. Oxidation: The gain of oxygen. [1]
2. Substance reduced: $\text{CuO}$ (Copper(II) oxide). [1]
3. Reducing agent: A substance that donates electrons / removes oxygen / increases the oxidation state of another substance and is itself oxidized. [1]
4. Oxidized: $\text{Zn}$ (Zinc) [1]; Reduced: $\text{Cu}^{2+}$ (Copper(II) ions/Copper sulfate) [1].
5. (a) $\text{K}$: +1 [1]; (b) $\text{S}$: +6 [1].
6. Magnesium loses electrons to become $\text{Mg}^{2+}$ (oxidation) [1], while hydrogen ions in the acid gain electrons to become $\text{H}_2$ gas (reduction) [1].
7. (a) Reducing agent [1]. (b) Decreases from +7 to +2 [1].
8. Anode [1].
9. $\text{Pb}^{2+}$ ions migrate to the cathode [1] and gain electrons to be discharged as lead metal [1].
10. (a) Chlorine gas [1]. (b) Hydrogen ions ($\text{H}^+$) are lower in the reactivity series than sodium ions ($\text{Na}^+$) [1], so $\text{H}^+$ is selectively discharged [1].
11. Anode: Oxygen [1]; Cathode: Copper [1].
12. $\text{Al}^{3+}(\text{l}) + 3\text{e}^- \rightarrow \text{Al}(\text{l})$ [2].
13. (a) Iodine ($\text{I}_2$) [1]. (b) $2\text{I}^-(\text{aq}) \rightarrow \text{I}_2(\text{aq}) + 2\text{e}^-$ [2].
14. Molten compounds contain only one type of cation and one type of anion [1], whereas aqueous compounds contain $\text{H}^+$ and $\text{OH}^-$ ions from water which compete for discharge [1].
15. (a) Anode [1]. (b) The mass decreases as copper atoms are oxidized to $\text{Cu}^{2+}$ ions and enter the solution [1].
16. (a) Pure silver [1]. (b) A soluble silver salt (e.g., silver nitrate) [1].
17. It allows the flow of ions [1] to complete the electrical circuit [1].
18. (a) Zinc [1]. (b) Zinc is more reactive than copper [1], so it loses electrons more readily to become the anode/negative terminal [1].
19. Zero emissions / only produces water as a byproduct / higher efficiency [1].
20. $2\text{H}_2(\text{g}) + \text{O}_2(\text{g}) \rightarrow 2\text{H}_2\text{O}(\text{l})$ [2].