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Chemistry · Practice book · Senior 6 · Unit 11 of 14

Electrochemical cells and applications

Senior 6 Chemistry · 14 questions with answers. Try each question on paper, then tap it to see the answer.

Unit 11 Electrochemical cells and applications

Key unit competence: Explain the working principles and industrial applications of electrochemical cells, and predict the feasibility of redox reactions.

Key ideas

  • An electrochemical (galvanic, voltaic) cell turns chemical energy into electrical energy using a spontaneous redox reaction. Two half-cells are joined by a wire and a salt bridge (e.g. filter paper soaked in KNO3), which completes the circuit and keeps the solutions neutral. Oxidation at the anode (negative), reduction at the cathode (positive).
  • The standard hydrogen electrode (SHE): platinum (coated with Pt black) in 1.00 mol dm−3 H+, H2 at 1 atm (100 kPa), 298 K; E° = 0.00 V by definition. Standard electrode potentials E° are measured against it: e.g. Zn2+/Zn −0.76 V, Fe2+/Fe −0.44 V, Cu2+/Cu +0.34 V, Ag+/Ag +0.80 V, Cl2/Cl− +1.36 V.
  • Cell notation: Zn(s)
    Zn2+(aq) || Cu2+(aq)
    Cu(s); E°cell = E°right − E°left = E°(cathode) − E°(anode) = 0.34 − (−0.76) = +1.10 V. A redox reaction is feasible when E°cell is positive; ΔG° = −nFE°cell.
  • Concentration affects E (Nernst equation): E = E° + (0.059/n) log[Mn+] for a metal/ion half-cell at 298 K. Lowering [Cu2+] makes E(Cu) less positive.
  • Corrosion of iron (rusting) needs oxygen and water: anode Fe → Fe2+ + 2e−; cathode O2 + 2H2O + 4e− → 4OH−; then Fe(OH)2 is oxidised to rust, Fe2O3·xH2O. Salt and acids speed it up. Prevention: paint, oil, plastic coating, galvanising (zinc: sacrificial and barrier), sacrificial magnesium blocks, tin plating (barrier only), alloying (stainless steel).
  • Batteries: dry (Leclanché) cell: zinc case anode, carbon rod in MnO2 cathode, NH4Cl paste, 1.5 V; lead–acid accumulator: Pb and PbO2 in H2SO4, 2.0 V per cell, rechargeable; lithium-ion (phones); hydrogen–oxygen fuel cell (2H2 + O2 → 2H2O, 1.23 V, only water as product). Galvanic cells: anode negative, spontaneous; electrolytic cells: anode positive, need an external supply.

Common misconceptions

  • "The anode is always positive." No: In a galvanic cell the anode (where oxidation occurs) is NEGATIVE; in an electrolytic cell it is positive.
  • "Electrons flow through the salt bridge." No: Ions move in the salt bridge; electrons flow only in the external wire.
  • "Tin plating protects iron even when scratched, like zinc." No: Tin is less reactive than iron: once scratched, iron corrodes faster. Zinc protects sacrificially.

Worked examples

Example 1. Describe the Daniell cell: its half-cells, electrode reactions, electron flow and E°.

Solution

Left half-cell: zinc rod in 1.0 mol dm−3 ZnSO4; right half-cell: copper rod in 1.0 mol dm−3 CuSO4; salt bridge of KNO3; voltmeter between the rods.

Anode (−): Zn → Zn2+ + 2e− (zinc dissolves). Cathode (+): Cu2+ + 2e− → Cu (copper deposits; blue fades).

Electrons flow through the wire from Zn to Cu; in the salt bridge NO3− moves to the zinc side and K+ to the copper side.

E°cell = 0.34 − (−0.76) = +1.10 V; cell notation Zn

Zn2+ || Cu2+

Cu.

Example 2. Explain how rusting happens and how galvanising prevents it.

Solution

Water containing dissolved oxygen covers iron. Where oxygen is low, iron is oxidised: Fe → Fe2+ + 2e− (anode).

Electrons flow to oxygen-rich areas: O2 + 2H2O + 4e− → 4OH− (cathode); Fe2+ and OH− give Fe(OH)2, oxidised to rust Fe2O3·xH2O.

Galvanising coats iron with zinc: a barrier to water and air.

If scratched, zinc (E° −0.76 V, more reactive) is oxidised instead of iron: sacrificial protection.

Practice questions

  1. Define standard electrode potential.

    Show answer

    The e.m.f. of a half-cell connected to a standard hydrogen electrode under standard conditions (298 K, 1 mol dm−3, 1 atm).

    Common mistake: Leaving out the reference electrode.

  2. State the standard conditions for the SHE.

    Show answer

    Pt electrode, 1.00 mol dm−3 H+, H2 at 1 atm, 298 K.

    Common mistake: Saying 0 °C.

  3. Why is potassium nitrate used in a salt bridge?

    Show answer

    Its ions do not react with the solutions or electrodes and do not form precipitates.

    Common mistake: Using KCl with a silver half-cell (AgCl precipitates).

  4. Write the overall equation for the Daniell cell.

    Show answer

    Zn + Cu2+ → Zn2+ + Cu

    Common mistake: Writing it the wrong way round.

  5. Give two differences between galvanic and electrolytic cells.

    Show answer

    Galvanic: spontaneous, makes electricity, anode negative; electrolytic: non-spontaneous, uses electricity, anode positive (any two).

    Common mistake: Saying both have negative anodes.

  6. State two ways of preventing rusting.

    Show answer

    Painting, oiling, galvanising, sacrificial protection, alloying (any two).

    Common mistake: Saying 'keep it wet'.

  7. Why is a car battery called a secondary cell?

    Show answer

    It can be recharged by passing a current to reverse the reaction.

    Common mistake: Saying it is the second battery in the car.

  8. Give one advantage of hydrogen fuel cells.

    Show answer

    Only water is produced; high efficiency; quiet (any one).

    Common mistake: Saying hydrogen is easy to store.

  9. Which is the stronger reducing agent: Zn or Cu? Explain.

    Show answer

    Zn: its E° (−0.76 V) is more negative, so it loses electrons more easily.

    Common mistake: Choosing Cu because its E° is larger.

  10. Calculate E°cell for Cl2/Cl− (+1.36) with Br2/Br− (+1.07).

    Show answer

    1.36 − 1.07 = +0.29 V; chlorine oxidises bromide.

    Common mistake: Adding the two potentials.

Matching and drag-and-drop

  1. Matching

    Match each item in column A with the correct answer in column B. Write the letter next to each number.

    Column AColumn B
    1. In a dry cell the negative electrode is the:a. positive
    2. A redox reaction is feasible if E°cell is:b. is oxidised (loses electrons)
    3. The more negative the E° value, the more easily the metal:c. 1.00 mol dm−3
    4. The concentration of H+ in the SHE is:d. zinc case
    Show answer

    1 → d (zinc case) · 2 → a (positive) · 3 → b (is oxidised (loses electrons)) · 4 → c (1.00 mol dm−3)

    • Zinc is oxidised.
    • ΔG° = −nFE° is negative.
    • It is a stronger reducing agent.
    • Standard conditions.

  2. Drag and drop

    Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

    copper  ·  sulfuric acid  ·  0.00 V  ·  −0.76 V
    1. In the Daniell cell, the electrode that gets heavier is .
    2. The standard electrode potential of the SHE is .
    3. The electrolyte in a car battery is .
    Show answer

    1: copper · 2: 0.00 V · 3: sulfuric acid

    • Cu deposits at the cathode.
    • Reference electrode.
    • Dilute H2SO4.

  3. Matching

    Match each item in column A with the correct answer in column B. Write the letter next to each number.

    Column AColumn B
    1. Galvanising protects iron by coating it with:a. zinc
    2. A galvanic cell converts:b. 2.0 V
    3. The only product of a hydrogen–oxygen fuel cell is:c. water
    4. The voltage of one lead–acid cell is about:d. chemical energy into electrical energy
    Show answer

    1 → a (zinc) · 2 → d (chemical energy into electrical energy) · 3 → c (water) · 4 → b (2.0 V)

    • Sacrificial protection.
    • Spontaneous redox reaction.
    • 2H2 + O2 → 2H2O.
    • Six cells make a 12 V car battery.

  4. Drag and drop

    Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

    +1.10 V  ·  anode  ·  positive  ·  voltmeter
    1. In an electrolytic cell, the anode is .
    2. E°cell for the Daniell cell (Zn/Cu) is .
    3. In a galvanic cell, oxidation takes place at the .
    Show answer

    1: positive · 2: +1.10 V · 3: anode

    • Connected to the + terminal.
    • 0.34 − (−0.76).
    • Anode = oxidation.

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