Secondary 3 Chemistry Redox Electrochemistry Quiz

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Secondary 3 Chemistry Quiz - Redox Electrochemistry

Name: ___________________________
Class: ___________
Date: ____________
Score: _______ / 40

Duration: 60 minutes
Total Marks: 40
Instructions: Answer all 20 questions. Section A is short answer, Section B is structured response, Section C is data and calculation. Show your working where required. Use proper chemical notation.

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Section A: Short Answer (Questions 1–5)

1. State the term used to describe a reaction where both reduction and oxidation occur at the same time. [1]

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2. In the reaction $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$, identify the species that is oxidised. [1]

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3. Write the oxidation state of manganese in $\text{KMnO}_4$. [1]

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4. Name the type of cell that converts chemical energy into electrical energy using a redox reaction. [1]

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5. State the direction of electron flow in a simple galvanic cell: from the ___________ electrode to the ___________ electrode. [1]

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Section B: Structured Response (Questions 6–15)

6. (a) Define oxidation in terms of electrons. [1]
(b) Define reduction in terms of electrons. [1]
(c) State the acronym used to remember the loss and gain of electrons. [1]

7. Consider the half-equations:
$\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^-$
$\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-$
(a) Which species is reduced? [1]
(b) Write the overall redox equation. [2]

8. A student places a strip of magnesium in copper(II) sulphate solution.
(a) State what is observed on the magnesium strip. [1]
(b) Write the ionic equation for the reaction. [2]
(c) Identify the oxidising agent. [1]

9. (a) State two requirements for a standard hydrogen electrode. [2]
(b) Explain why the hydrogen electrode is used as a reference. [1]

10. The reactivity series places metals in order of their tendency to lose electrons.
(a) Which of zinc or silver is the better reducing agent? [1]
(b) Give a reason based on electron loss. [1]
(c) Write a half-equation for the better reducing agent. [1]

11.

Generated diagram for Q11.

Using the diagram above:
(a) Label the anode and cathode. [2]
(b) State the direction of electron flow. [1]
(c) Write the half-equation at the anode. [1]

12. (a) What is meant by an oxidising agent? [1]
(b) Give one example of a common oxidising agent in the lab. [1]
(c) Explain how it causes another species to be oxidised. [1]

13. Electrolysis of molten lead(II) bromide is carried out.
(a) Write the half-equation at the cathode. [1]
(b) Write the half-equation at the anode. [1]
(c) State the product at the anode and a test for it. [2]

14. (a) State the colour change of acidified potassium manganate(VII) when it acts as an oxidising agent. [1]
(b) Write the reduced form of $\text{MnO}_4^-$ in acidic condition. [1]
(c) State the oxidation state change of Mn. [1]

15. A redox reaction occurs: $\text{2I}^- + \text{Br}_2 \rightarrow \text{I}_2 + \text{2Br}^-$
(a) Identify the reducing agent. [1]
(b) State the oxidation state change of iodine. [1]
(c) Explain why bromine is the oxidising agent. [1]

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Section C: Data and Calculation (Questions 16–20)

16. A galvanic cell is set up with $\text{Mg}^{2+}/\text{Mg}$ and $\text{Fe}^{2+}/\text{Fe}$.
Standard electrode potentials:
$\text{Mg}^{2+} + 2e^- \rightleftharpoons \text{Mg}$ $E^\circ = -2.37\text{ V}$
$\text{Fe}^{2+} + 2e^- \rightleftharpoons \text{Fe}$ $E^\circ = -0.44\text{ V}$
(a) Calculate the cell potential. [2]
(b) Write the overall cell reaction. [2]

17. In a titration, $25.0\text{ cm}^3$ of $0.020\text{ mol dm}^{-3}$ $\text{KMnO}_4$ reacts with $\text{Fe}^{2+}$ solution.
Reaction: $\text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}$
(a) Calculate moles of $\text{MnO}_4^-$ used. [2]
(b) Calculate moles of $\text{Fe}^{2+}$ oxidised. [2]

18.

Generated table for Q18.

Using the table:
(a) Which species is the strongest reducing agent? [1]
(b) Write the cell notation for the cell with highest voltage using two of these. [2]
(c) Calculate that maximum cell potential. [1]

19. Electrolysis of aqueous copper(II) sulphate using inert electrodes.
(a) State the product at the cathode and write its half-equation. [2]
(b) State the product at the anode and explain. [2]

20. A sample of $1.00\text{ g}$ of zinc is placed in excess copper(II) sulphate.
$\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$
Relative atomic masses: Zn = 65.4, Cu = 63.5
(a) Calculate moles of Zn reacted. [2]
(b) Calculate mass of Cu deposited. [2]

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Secondary 3 Chemistry Quiz - Redox Electrochemistry: Answer Key

Total Marks: 40
Teaching notes are provided for each question.

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Section A

1. [1] Redox (or reduction-oxidation)
Teaching: A redox reaction is one where reduction and oxidation happen simultaneously. "Redox" is the combined term.

2. [1] Zn (zinc atom)
Teaching: Zn loses electrons to become $\text{Zn}^{2+}$ (oxidation). The species oxidised is the one that loses electrons.

3. [1] +7
Teaching: K is +1, O is -2 (×4 = -8). For neutral $\text{KMnO}_4$: +1 + x - 8 = 0 → x = +7.

4. [1] Galvanic cell (or electrochemical cell / voltaic cell)
Teaching: A galvanic cell uses spontaneous redox to produce electricity.

5. [1] from the anode (negative) to the cathode (positive)
Teaching: Electrons flow from the electrode where oxidation occurs (anode, negative) to where reduction occurs (cathode, positive).

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Section B

6. [3]
(a) [1] Oxidation is loss of electrons.
(b) [1] Reduction is gain of electrons.
(c) [1] OIL RIG (Oxidation Is Loss, Reduction Is Gain).
Teaching: Electron transfer defines redox at Sec 3 level.

7. [3]
(a) [1] $\text{Cl}_2$ (it gains electrons to become $\text{Cl}^-$).
(b) [2] $2\text{Fe}^{2+} + \text{Cl}_2 \rightarrow 2\text{Fe}^{3+} + 2\text{Cl}^-$
Teaching: Multiply Fe half-equation by 2 to balance electrons; add with Cl half-equation.

8. [4]
(a) [1] Reddish-brown copper metal deposits on Mg strip.
(b) [2] $\text{Mg} + \text{Cu}^{2+} \rightarrow \text{Mg}^{2+} + \text{Cu}$
(c) [1] $\text{Cu}^{2+}$ (it gains electrons, gets reduced).
Teaching: Mg is more reactive, displaces Cu²⁺.

9. [3]
(a) [2] Hydrogen gas at 1 atm, $\text{H}^+$ at $1\text{ mol dm}^{-3}$, $25^\circ\text{C}$, Pt electrode. (any two)
(b) [1] It is assigned a potential of 0.00 V as standard reference.
Teaching: SHE defines the zero point for measuring other potentials.

10. [3]
(a) [1] Zinc
(b) [1] Zn loses electrons more readily than Ag.
(c) [1] $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$
Teaching: Better reducing agent = easier to oxidise = more negative $E^\circ$.

11. [4]
(a) [2] Anode = Zn electrode (left), Cathode = Cu electrode (right).
(b) [1] From Zn to Cu (left to right).
(c) [1] $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$
Teaching: Diagram must show Zn oxidised; electrons leave Zn through wire to Cu.

12. [3]
(a) [1] Species that accepts electrons / causes oxidation of another.
(b) [1] Acidified $\text{KMnO}_4$ (or $\text{Cl}_2$, etc.)
(c) [1] It gains electrons from the other species, oxidising it.
Teaching: Oxidising agent itself is reduced.

13. [4]
(a) [1] $\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}$
(b) [1] $2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-$
(c) [2] Bromine; test with starch-iodide paper turns blue-black, or vapour is red-brown.
Teaching: Molten salt electrolysis: cations to cathode, anions to anode.

14. [3]
(a) [1] Purple to colourless.
(b) [1] $\text{Mn}^{2+}$
(c) [1] +7 to +2.
Teaching: In acid, manganate(VII) reduced from Mn(VII) to Mn(II).

15. [3]
(a) [1] $\text{I}^-$ (iodide ion)
(b) [1] -1 to 0.
(c) [1] $\text{Br}_2$ gains electrons to become $\text{Br}^-$, oxidising $\text{I}^-$.
Teaching: Reducing agent is oxidised; oxidising agent is reduced.

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Section C

16. [4]
(a) [2] $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = (-0.44) - (-2.37) = +1.93\text{ V}$
(b) [2] $\text{Mg} + \text{Fe}^{2+} \rightarrow \text{Mg}^{2+} + \text{Fe}$
Teaching: More negative electrode is anode (Mg). Cell potential positive for spontaneous.

17. [4]
(a) [2] $n = c \times V = 0.020 \times (25.0/1000) = 5.00 \times 10^{-4}\text{ mol}$
(b) [2] From equation 1:5 ratio → $5 \times 5.00\times10^{-4} = 2.50\times10^{-3}\text{ mol Fe}^{2+}$
Teaching: Convert cm³ to dm³ by ÷1000 before using $n=cV$.

18. [4]
(a) [1] Li (most negative $E^\circ$ = strongest reducer)
(b) [2] $\text{Li}|\text{Li}^+ || \text{F}_2,\text{F}^-|\text{Pt}$ (or $\text{Li}|\text{Li}^+ || \text{F}_2|\text{F}^-$ with Pt)
(c) [1] $2.87 - (-3.04) = 5.91\text{ V}$
Teaching: Max voltage from most negative and most positive couples.

19. [4]
(a) [2] Cu at cathode: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$
(b) [2] $\text{O}_2$ at anode: $\text{4OH}^- \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4e^-$ (water oxidised, not sulphate).
Teaching: With inert electrodes in aqueous CuSO₄, Cu²⁺ reduced preferentially; OH⁻ oxidised at anode.

20. [4]
(a) [2] $n_{\text{Zn}} = 1.00 / 65.4 = 0.0153\text{ mol}$
(b) [2] 1:1 ratio → $m_{\text{Cu}} = 0.0153 \times 63.5 = 0.971\text{ g}$
Teaching: Mole ratio from equation is 1:1; use $m = n \times M_r$.