Secondary 4 Pure Chemistry Redox Electrochemistry Quiz

Secondary 4 Pure Chemistry Quiz - Redox Electrochemistry

ANSWER KEY AND MARKING SCHEME

Total Marks: 40

Section A: Multiple Choice (10 marks)

Marking notes:

• Q1: Oxidation is loss of electrons (OIL RIG mnemonic).
• Q2: In molten NaCl, Cl⁻ ions are discharged at the anode: 2Cl⁻ → Cl₂ + 2e⁻.
• Q3: Zinc is more reactive than copper; the more reactive metal acts as the negative terminal (anode in a cell).
• Q4: Potassium dichromate(VI) is a common oxidising agent; it accepts electrons and is reduced.
• Q5: With copper electrodes in CuSO₄(aq), the anode dissolves: Cu(s) → Cu²⁺(aq) + 2e⁻, so its mass decreases.

Section B: Multiple Choice (10 marks)

Marking notes:

• Q6: Neutralisation (HCl + NaOH) is not a redox reaction; no change in oxidation states occurs.
• Q7: In dilute H₂SO₄, H⁺ ions are discharged at the cathode: 2H⁺ + 2e⁻ → H₂.
• Q8: Zinc provides both barrier protection and sacrificial protection because it is more reactive than iron.
• Q9: KMnO₄: K = +1, O₄ = -8, so Mn = +7 (1 + x + 4(-2) = 0, x = +7).
• Q10: The hydrogen fuel cell overall reaction is 2H₂ + O₂ → 2H₂O.

Section C: Structured Questions (10 marks)

Question 11 (6 marks)

(a) Half-equation at the cathode: Pb²⁺(l) + 2e⁻ → Pb(l) [2 marks]

• Award 1 mark for correct species (Pb²⁺ + 2e⁻ → Pb).
• Award 1 mark for correct state symbols (l).

(b) Observation at the anode: Brown fumes / reddish-brown gas evolved [1 mark] Explanation: Bromide ions are oxidised to bromine gas: 2Br⁻(l) → Br₂(g) + 2e⁻ [1 mark]

• Accept: "Brown gas with pungent smell" or "Red-brown vapour".
• Must link observation to the discharge of bromide ions.

(c) Lead(II) bromide must be molten because: In the solid state, the ions are held in fixed positions in the ionic lattice and cannot move / are not mobile. When molten, the ions are free to move and can carry electric current / conduct electricity. [2 marks]

• Award 1 mark for stating ions cannot move in solid.
• Award 1 mark for stating ions are mobile in molten state.

Question 12 (5 marks)

(a) Anode: Magnesium electrode [1 mark] Explanation: Magnesium is more reactive than copper / Magnesium has a greater tendency to lose electrons / Magnesium is oxidised more readily than copper. [1 mark]

• In a simple cell, the more reactive metal acts as the anode (negative terminal) where oxidation occurs.

(b) Half-equation at copper electrode: Cu²⁺(aq) + 2e⁻ → Cu(s) [1 mark]

• Reduction occurs at the cathode (copper electrode).
• State symbols must be correct.

(c) Function of salt bridge: The salt bridge completes the circuit by allowing ions to flow between the two half-cells. Anions move towards the anode and cations move towards the cathode, maintaining electrical neutrality in each half-cell. [2 marks]

• Award 1 mark for "completes the circuit / allows ion flow".
• Award 1 mark for "maintains electrical neutrality" or description of ion movement.

Question 13 (5 marks)

(a) Anode: Impure copper [1 mark] Cathode: Pure copper [1 mark]

(b) Half-equation at cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) [1 mark]

• Copper(II) ions from the electrolyte are reduced and deposited as pure copper on the cathode.

(c) Impurities in the impure copper anode: More reactive metals (e.g., Zn, Fe) dissolve into the solution as ions but are not discharged at the cathode. Less reactive metals (e.g., Ag, Au) do not dissolve and fall to the bottom of the cell as anode sludge / anode mud. [2 marks]

• Award 1 mark for describing the fate of more reactive impurities.
• Award 1 mark for describing the fate of less reactive impurities (anode sludge).

Question 14 (2 marks)

(a) The colour change from orange to green indicates: The unknown solution contains a reducing agent / The unknown solution can reduce potassium dichromate(VI). [1 mark]

• Accept: "The solution is a reducing agent" or "The dichromate(VI) ions are reduced to chromium(III) ions."

(b) Type of reaction: Redox reaction / Oxidation-reduction reaction [1 mark]

• The dichromate(VI) is reduced (orange → green) while the unknown substance is oxidised.

Question 15 (5 marks)

(a) Order of increasing reactivity: Copper < Iron < Zinc < Magnesium [1 mark]

• Must be in correct order from least to most reactive.
• Accept: Cu, Fe, Zn, Mg.

(b) Ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) [2 marks]

• Award 1 mark for correct formulae.
• Award 1 mark for correct state symbols.
• Spectator ions (SO₄²⁻) must be omitted.

(c) Explanation: Copper is less reactive than iron / Copper is below iron in the reactivity series. Copper atoms cannot lose electrons to reduce iron(II) ions / Copper has a lower tendency to form ions than iron. For a reaction to occur, the added metal must be more reactive than the metal ion in solution so that electron transfer can take place. [2 marks]

• Award 1 mark for stating copper is less reactive than iron.
• Award 1 mark for explaining in terms of electron transfer (Cu cannot reduce Fe²⁺ / Cu cannot donate electrons to Fe²⁺).

Section D: Data-Based and Extended Response Questions (10 marks)

Question 16 (5 marks)

(a) Gas produced at the anode: Chlorine gas [1 mark] Explanation: In concentrated aqueous sodium chloride, chloride ions are discharged in preference to hydroxide ions due to the concentration effect / Chloride ions are present in higher concentration than hydroxide ions, so they are selectively discharged despite oxygen being higher in the discharge series. [1 mark]

• Accept reference to "selective discharge" or "concentration effect".
• The relevant half-equation: 2Cl⁻(aq) → Cl₂(g) + 2e⁻.

(b) Explanation of alkaline solution at cathode: At the cathode, hydrogen ions (from water) are discharged in preference to sodium ions: 2H⁺(aq) + 2e⁻ → H₂(g) OR 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq). This leaves behind an excess of hydroxide ions (OH⁻) in the solution around the cathode, making it alkaline. [3 marks]

• Award 1 mark for identifying that H⁺ ions (or water) are discharged at the cathode.
• Award 1 mark for correct half-equation.
• Award 1 mark for explaining that OH⁻ ions accumulate, causing alkalinity.
• Accept: "Sodium ions remain in solution with hydroxide ions, forming sodium hydroxide which is alkaline."

Question 17 (3 marks)

An oxidising agent is a substance that accepts electrons / causes another substance to be oxidised, and is itself reduced. [1 mark] Example: Potassium dichromate(VI) / acidified potassium manganate(VII) / chlorine. [1 mark] Half-equation for reduction: Cr₂O₇²⁻(aq) + 14H⁺(aq) + 6e⁻ → 2Cr³⁺(aq) + 7H₂O(l) (accept other valid examples with correct half-equation). [1 mark]

• Award marks for correct definition, valid example, and correct corresponding half-equation.

Question 18 (2 marks)

In aqueous aluminium chloride, water is present. At the cathode, hydrogen ions (from water) are discharged in preference to aluminium ions because aluminium is more reactive than hydrogen. Therefore, hydrogen gas is produced instead of aluminium metal. [1 mark] In the electrolysis of molten aluminium oxide, there is no water present. The only ions are Al³⁺ and O²⁻, so aluminium ions are discharged at the cathode to form aluminium metal. [1 mark]

• Accept reference to the reactivity series and selective discharge.

Question 19 (3 marks)

(a) Anode material: Silver metal / Pure silver [1 mark]

(b) Suitable electrolyte: Silver nitrate solution / any soluble silver salt solution (e.g., silver ethanoate, silver sulfate) [1 mark]

(c) Half-equation at iron spoon (cathode): Ag⁺(aq) + e⁻ → Ag(s) [1 mark]

• The iron spoon is the cathode where reduction occurs.

Question 20 (2 marks)

Similarity: Both cells involve redox reactions / involve electron transfer / involve oxidation at the anode and reduction at the cathode. [1 mark] Difference: A simple chemical cell converts chemical energy into electrical energy (spontaneous reaction), while an electrolytic cell converts electrical energy into chemical energy (non-spontaneous reaction). [1 mark]

• Accept other valid similarities and differences related to energy conversion and spontaneity.

END OF ANSWER KEY