O Level Chemistry Redox Electrochemistry Quiz

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O-Level Chemistry Quiz - Redox Electrochemistry

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

Duration: 60 Minutes
Total Marks: 50

Instructions:

• Answer all questions in the spaces provided.
• Show all working for calculations.
• State symbols should be included in chemical equations where appropriate.

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Section A: Oxidation and Reduction (1-7)

1. Define oxidation in terms of electron transfer. [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 and explain your answer in terms of oxygen. [2]
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3. Determine the oxidation state of the underlined element in the following compounds: [3] (a) $\underline{\text{S}}\text{O}_2$: _________ (b) $\text{K}\underline{\text{Mn}}\text{O}_4$: _________ (c) $\underline{\text{Cr}}_2\text{O}_7^{2-}$: _________

4. A student adds acidified potassium manganate(VII) to a solution of iron(II) sulfate. (a) State the color change observed. [1] \

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   (b) Explain why the color change occurs in terms of redox. [2]
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5. Identify the oxidizing agent and the reducing agent in the following reaction: [2] $\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}$ Oxidizing Agent: ____________________ Reducing Agent: ____________________

6. Write an ionic equation for the reaction between concentrated hydrochloric acid and potassium iodide solution. [2]
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7. Explain why aqueous potassium iodide is used to test for the presence of an oxidizing agent. [2]
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Section B: Electrolysis (8-15)

8. Define electrolysis. [1]
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9. During the electrolysis of molten lead(II) bromide ($\text{PbBr}_2$): (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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10. Compare the products formed at the cathode during the electrolysis of molten sodium chloride and aqueous sodium chloride. Explain the difference. [3]
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11. A student electrolyzes a concentrated solution of copper(II) sulfate using inert carbon electrodes. (a) Describe the observation at the anode. [1] \

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    (b) Write the half-equation for the reaction at the anode. [2]
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12. Explain why the concentration of $\text{Cu}^{2+}$ ions in the electrolyte decreases during the purification of copper. [2]
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13. In the electroplating of a steel spoon with silver, the spoon is made the cathode. Explain why the spoon cannot be the anode. [2]
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14. Predict the product formed at the cathode during the electrolysis of aqueous magnesium chloride. Justify your answer using the reactivity series. [2]
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15. State the purpose of using a diaphragm or a selective membrane in some industrial electrolysis cells. [2]
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Section C: Simple Cells and Fuel Cells (16-20)

16. A simple cell is constructed using a zinc electrode and a copper electrode in a salt bridge system. (a) Which electrode acts as the anode? [1] \

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    (b) Explain the flow of electrons in the external circuit. [2]
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17. Describe how the voltage of a simple cell can be increased. [2]
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18. In a hydrogen fuel cell, the fuel is hydrogen and the oxidant is oxygen. (a) State the product formed at the cathode. [1] \

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    (b) Write the overall equation for the reaction in the fuel cell. [2]
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19. Compare a hydrogen fuel cell to a traditional internal combustion engine in terms of environmental sustainability. [2]
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20. A student uses a magnesium strip and a lead strip to create a simple cell. Predict which metal will be the cathode and explain why. [2]
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Answer Key - Redox Electrochemistry Quiz

1. Oxidation is the loss of electrons. (1)

2. $\text{CuO}$ is reduced. (1) It loses oxygen to become copper. (1)

3. (a) +4 (1) (b) +7 (1) (c) +6 (1)

4. (a) Purple to colorless. (1) (b) $\text{MnO}_4^-$ is reduced to $\text{Mn}^{2+}$ (colorless), while $\text{Fe}^{2+}$ is oxidized to $\text{Fe}^{3+}$. (2)

5. Oxidizing Agent: $\text{CuSO}_4$ (or $\text{Cu}^{2+}$) (1) Reducing Agent: $\text{Zn}$ (1)

6. $2\text{I}^- \text{(aq)} \rightarrow \text{I}_2\text{(s/aq)} + 2\text{e}^-$ (2) [Accept balanced molecular equation if correct]

7. $\text{I}^-$ ions are easily oxidized to $\text{I}_2$ (brown color), which indicates the presence of an oxidizing agent that can remove electrons from iodide. (2)

8. The breakdown of an ionic compound using electricity. (1)

9. (a) Lead (1) (b) $2\text{Br}^- \rightarrow \text{Br}_2\text{(g)} + 2\text{e}^-$ (2)

10. Molten: Sodium is formed. (1) Aqueous: Hydrogen is formed. (1) Hydrogen is lower in the reactivity series than sodium, so $\text{H}^+$ ions are preferentially discharged over $\text{Na}^+$ ions. (1)

11. (a) Bubbles of colorless gas (oxygen) are evolved. (1) (b) $2\text{H}_2\text{O(l)} \rightarrow \text{O}_2\text{(g)} + 4\text{H}^+\text{(aq)} + 4\text{e}^-$ (2) [Or $4\text{OH}^- \rightarrow 2\text{H}_2\text{O} + \text{O}_2 + 4\text{e}^-$]

12. $\text{Cu}^{2+}$ ions from the electrolyte are reduced to copper atoms and deposit at the cathode. (2)

13. The anode is where oxidation occurs (metal dissolves). If the spoon were the anode, it would dissolve into the solution instead of being coated. (2)

14. Hydrogen gas. (1) Magnesium is more reactive than hydrogen, so $\text{H}^+$ ions are discharged at the cathode. (1)

15. To prevent the products formed at the anode and cathode from reacting with each other (e.g., preventing $\text{Cl}_2$ and $\text{NaOH}$ from reacting). (2)

16. (a) Zinc (1) (b) Electrons flow from the more reactive metal (Zinc) to the less reactive metal (Copper) through the external wire. (2)

17. Use two metals that are further apart in the reactivity series. (2)

18. (a) Water ($\text{H}_2\text{O}$) (1) (b) $2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}$ (2)

19. Fuel cells only produce water as a byproduct, whereas combustion engines produce $\text{CO}_2$ and $\text{NO}_x$ (greenhouse/pollutant gases). (2)

20. Lead. (1) Lead is lower in the reactivity series than magnesium, meaning it is less likely to lose electrons and will therefore be the site of reduction. (1)