Voltaic (Galvanic) Cell
Voltaic (Galvanic) Cells
A voltaic cell converts chemical energy → electrical energy via a spontaneous redox reaction. The two half-cells are separated so electrons must travel through an external wire.
Key Components
- Anode (−): Oxidation occurs here. The more reactive metal dissolves, releasing electrons into the wire.
- Cathode (+): Reduction occurs here. Metal ions in solution gain electrons and deposit as solid metal.
- External wire: Electrons flow from anode → cathode through the wire, producing a current.
- Salt bridge: Contains unreactive ions (e.g. KNO₃). Allows ion migration to maintain electrical neutrality. Cations move toward the cathode, anions toward the anode.
- Voltmeter: Measures the potential difference (EMF / voltage) of the cell.
Zn–Cu Voltaic Cell: Half-Equations
Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Zinc is more reactive than copper, so it is oxidised. The zinc electrode loses mass while the copper electrode gains mass as Cu is deposited.
Electrolytic Cells
An electrolytic cell uses electrical energy → chemical energy to force a non-spontaneous reaction to occur. An external power supply drives electrons in the opposite direction.
Key Points
- The electrolyte can be a molten ionic compound or an aqueous solution of an ionic compound.
- Anode (+): Oxidation occurs. Connected to the positive terminal of the power supply.
- Cathode (−): Reduction occurs. Connected to the negative terminal of the power supply.
- Cations migrate towards the cathode; anions migrate towards the anode.
Example: Electrolysis of Molten NaCl
Cathode (−): Na⁺ + e⁻ → Na (reduction. Sodium metal deposited)
Anode (+): 2Cl⁻ → Cl₂ + 2e⁻ (oxidation. Chlorine gas evolved)
Overall: 2NaCl(l) → 2Na(l) + Cl₂(g)
Voltaic vs Electrolytic. Comparison
| Voltaic | Electrolytic | |
|---|---|---|
| Energy | Chemical → Electrical | Electrical → Chemical |
| Spontaneous? | Yes (ΔG < 0) | No (ΔG > 0) |
| Anode | Negative (−) | Positive (+) |
| Cathode | Positive (+) | Negative (−) |
| Power source | None (self-generating) | External battery/supply |
In BOTH types: oxidation at anode, reduction at cathode. An Ox, Red Cat.
Standard Electrode Potential E° (HL)
The standard electrode potential (E°) measures the tendency of a half-cell to undergo reduction under standard conditions (298 K, 1 mol dm⁻³, 100 kPa). All values are measured relative to the Standard Hydrogen Electrode (SHE), which is assigned E° = 0.00 V. All E° values in data booklets are written as reduction potentials.
The Standard Hydrogen Electrode (SHE)
- A platinum electrode coated in platinum black (to increase surface area)
- Immersed in 1 mol dm⁻³ H⁺(aq) solution (e.g. 1 mol dm⁻³ HCl)
- H₂ gas bubbled over the electrode at 100 kPa
- Temperature maintained at 298 K (25 °C)
- The half-equation: 2H⁺(aq) + 2e⁻ ⇌ H₂(g), E° = 0.00 V
The SHE acts as a reference. Other half-cells are connected to it via a salt bridge and external wire, and the voltmeter reading gives the E° of the other half-cell.
Interpreting E° Values
- More positive E° → stronger tendency to be reduced → stronger oxidising agent
- More negative E° → stronger tendency to be oxidised → stronger reducing agent
Calculating E°cell
E°cell = E°cathode − E°anode
If E°cell > 0 → reaction is spontaneous (voltaic cell works).
If E°cell < 0 → reaction is non-spontaneous (electrolysis required).
Worked Example: Zn–Cu Cell E°cell
Given: E°(Cu²⁺/Cu) = +0.34 V, E°(Zn²⁺/Zn) = −0.76 V
Cu²⁺ has the more positive E° → it is reduced (cathode).
Zn has the more negative E° → it is oxidised (anode).
E°cell = (+0.34) − (−0.76) = +1.10 V
Positive → spontaneous. This cell will produce electricity.
Gibbs energy & Cell Potential (HL)
ΔG° = −nFE°cell
where n = moles of electrons transferred, F = Faraday constant (96 485 C mol⁻¹).
Positive E°cell → negative ΔG° → spontaneous.
Electroplating
Electroplating uses electrolysis to coat an object with a thin layer of metal. The object to be plated is the cathode; the plating metal is the anode. The electrolyte contains ions of the plating metal.
Electrolysis of Aqueous Solutions (HL)
When electrolyzing an aqueous solution, there is competition between the solute ions and water molecules at each electrode.
Predicting Products at the Cathode (Reduction)
- The species with the most positive (least negative) E° is preferentially reduced.
- Water can be reduced: 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq), E° = −0.83 V
- If the metal cation has a more negative E° than water (e.g. Na⁺, K⁺, Ca²⁺), water is reduced instead, producing H₂ gas.
- If the metal cation has a less negative E° than water (e.g. Cu²⁺, Ag⁺), the metal is deposited.
Predicting Products at the Anode (Oxidation)
- The species with the most negative (least positive) E° is preferentially oxidised.
- Water can be oxidised: 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻, E° = +1.23 V
- For dilute solutions, water is usually oxidised, producing O₂.
- For concentrated halide solutions (e.g. concentrated NaCl), the halide ion may be preferentially discharged despite thermodynamic predictions. This is due to kinetic factors (overpotential).
Worked Example: Electrolysis of Concentrated NaCl(aq)
Cathode (−): Na⁺ has E° = −2.71 V, water has E° = −0.83 V
Water has the more positive E°, so water is reduced: 2H₂O + 2e⁻ → H₂ + 2OH⁻
Anode (+): Cl⁻ has E° = +1.36 V, water has E° = +1.23 V
Thermodynamically, water should be oxidised (lower E°). But with concentrated Cl⁻, kinetic factors (overpotential) mean Cl₂ is produced instead.
Products: H₂(g) at cathode, Cl₂(g) at anode, NaOH(aq) left in solution.
Primary and Secondary Cells
The IB syllabus requires you to distinguish between primary cells, secondary cells, and fuel cells.
Primary Cells (R3.2.6)
- Use a spontaneous, irreversible redox reaction to produce electricity.
- Non-rechargeable: once the reactants are used up, the cell is discarded.
- Examples: zinc-carbon dry cells, alkaline batteries.
- Advantage: cheap, lightweight, no maintenance required.
- Disadvantage: limited lifespan, produces chemical waste.
Secondary (Rechargeable) Cells (R3.2.7)
- Use reversible redox reactions: the discharge reaction can be reversed by applying electrical energy.
- Rechargeable: when charging, the cell acts as an electrolytic cell (electrical → chemical energy).
- When discharging, it acts as a voltaic cell (chemical → electrical energy).
- Examples: lithium-ion batteries, lead-acid batteries (cars).
- Advantage: reusable over many cycles, less waste.
- Disadvantage: higher initial cost, eventual capacity loss.
Fuel Cells
- Convert chemical energy of a fuel (e.g. H₂) directly into electrical energy via a continuous redox reaction.
- Fuel and oxidant are supplied continuously from outside the cell.
- Example: hydrogen fuel cell. H₂ is oxidised at the anode, O₂ is reduced at the cathode. Product: water.
- Advantage: high efficiency, clean product (water), continuous operation.
- Disadvantage: expensive, requires hydrogen storage/infrastructure.
| Primary | Secondary | Fuel Cell | |
|---|---|---|---|
| Rechargeable? | No | Yes | N/A (continuous fuel) |
| Reaction type | Irreversible | Reversible | Continuous |
| Example | Alkaline battery | Li-ion battery | H₂/O₂ fuel cell |
Think About It
Why does the salt bridge contain ions like KNO₃ rather than reactive metal ions?
K⁺ and NO₃⁻ are spectator ions. They don't participate in redox reactions. The salt bridge completes the circuit by allowing ion flow to maintain electrical neutrality in each beaker, without interfering with the cell reaction.
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