Electrochemistry: Unit Review

Zinc is oxidized; Cu²⁺ is the oxidizing agent.

Oxidation: Na → Na⁺ + e⁻ (twice). Reduction: Cl₂ + 2e⁻ → 2Cl⁻.

Loss of Electrons is Oxidation; Gain of Electrons is Reduction.

Mn: +7. Cr: +6.

S: +6. N: −3.

No: every element keeps the same oxidation number (Ca +2, C +4, O −2).

5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

H₂O for oxygen, then H⁺ for hydrogen, then electrons for charge.

Add OH⁻ to both sides equal to the number of H⁺, combine H⁺ + OH⁻ into H₂O, and cancel water.

MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻

Cu rises by 2; N falls by 3 (+5 to +2). Ratio 3 Cu : 2 NO₃⁻.

Yes: Ag⁺ (+0.80 V) is a stronger oxidizing agent than Cu²⁺ (+0.34 V), so it oxidizes copper; silver crystals form.

No: Mg²⁺ (−2.37 V) is a weaker oxidizing agent than Zn²⁺, so no reaction.

Wear gloves and goggles: it stains and irritates skin and eyes; collect waste for proper disposal.

Two half-cells (electrodes in electrolytes), an external wire and a salt bridge. The salt bridge lets ions flow to keep each half-cell electrically neutral.

Electrons go from zinc to copper through the wire; anions move toward the zinc (anode) half-cell.

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Only differences can be measured, so a reference is chosen: the standard hydrogen electrode is defined as 0.00 V.

1.0 mol/L solutions, 100 kPa for gases and 25 °C.

Strongest oxidizing agent: F₂. Strongest reducing agent: Li.

0.80 − (−2.37) = 3.17 V

E°cell = 0.00 − 0.34 = −0.34 V: not spontaneous.

E° is an intensive property (energy per unit charge); it does not depend on the amount.

Non-standard concentrations or temperature, resistance in the salt bridge and wires, oxide on electrodes, meter draws current. (Any two.)

(1.10 − 1.04) ÷ 1.10 × 100 % = 5.5 %

To remove oxide coatings that add resistance and change the half-reaction.

The anode is positive; oxidation happens there.

Cathode: water is reduced to H₂ and OH⁻ (K⁺ is too weak). Anode: I⁻ is oxidized to I₂.

E°cell = 0.34 − 1.36 = −1.02 V, so more than 1.02 V.

Q = 3600 C; n(e⁻) = 0.03731 mol; n(Cu) = 0.01866 mol; m = 1.19 g.

n(Ag) = 0.04635 mol = n(e⁻); Q = 4472 C; t = 2981 s, about 49.7 min.

n(e⁻) = 1.000 mol; Q = 96 485 C; I = 96 485 ÷ 7200 = 13.4 A.

Impure copper is the anode. Silver and gold are not oxidized; they fall to the bottom as anode sludge.

In water, H₂O would be reduced instead of Al³⁺; cryolite lowers the melting point of Al₂O₃.

Chlorine, hydrogen and sodium hydroxide.

Anode: Fe → Fe²⁺ + 2e⁻. Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻.

Magnesium is a stronger reducing agent, so it is oxidized instead of iron (sacrificial or cathodic protection).

Dissolved ions make the water a better electrolyte, so ions move faster between anode and cathode regions.

Primary cells are single-use; secondary cells can be recharged by reversing the reaction with electrolysis.

1.69 − (−0.36) = 2.05 V

Damaged cells can go into thermal runaway and start fires; their metals should be recovered.

2H₂ + O₂ → 2H₂O; E°cell = 1.23 V.

Grey: made from natural gas, releasing CO₂. Green: made by electrolysis of water with clean electricity.

Life-cycle emissions and the electricity grid, mining of critical minerals and effects on communities, cost, cold-weather range, recycling. (Any two.)