Electrochemical cell: device converting chemical energy to electrical energy or the reverse. Electrolytic cell: electrical energy drives a non-spontaneous reaction. Galvanic (voltaic) cell: a spontaneous redox reaction produces electricity.
Electrolyte: substance that conducts when molten or in solution and is chemically changed (NaCl, HCl, NaOH). Non-electrolyte: urea, glucose, sucrose, benzene.
Electrolysis: decomposition of an ionic compound, molten or in solution, by passing an electric current.
Oxidation: loss of electrons (at the anode). Reduction: gain of electrons (at the cathode). OIL RIG; oxidation always at the anode, reduction always at the cathode in both cell types.
Electroplating: depositing a thin layer of one metal on another by electrolysis. Object = cathode, plating metal = anode, electrolyte contains ions of the plating metal.
Electrode potential (E°): tendency of a metal to be oxidised or reduced in its salt solution, measured in volts against the standard hydrogen electrode. Electrochemical series: metals arranged by increasing reduction potential.
Fuel cell: a galvanic cell using hydrogen and oxygen, giving water as the only chemical product. Battery: one galvanic cell or several joined in series.
Electrolytic cell facts
Parts: vessel with electrolyte, two inert electrodes, battery. Electrons flow in the external wire from anode to cathode. Cations move to the cathode, anions to the anode.
Rule for halide solutions: dilute solution gives oxygen at the anode; concentrated solution gives the halogen at the anode. Hydrogen is formed at the cathode in both cases. (In dilute NaCl, OH⁻ is discharged instead of Cl⁻ because of the concentration effect and the overpotential of chlorine.)
Cathode: the less reactive ion (lower in the series) is reduced, so H⁺ is reduced, not Na⁺. Anode: halides are discharged before OH⁻ when the solution is concentrated.
Electrolysis reactions (half-equations)
Molten NaCl (Down's cell, carbon anode, iron cathode): anode 2Cl⁻ → Cl₂ + 2e⁻. Cathode 2Na⁺ + 2e⁻ → 2Na. Gives sodium metal and chlorine.
Molten lead(II) chloride (above 327 °C, Pt or graphite electrodes): anode 2Cl⁻ → Cl₂ + 2e⁻. Cathode Pb²⁺ + 2e⁻ → Pb. Lead collects at the bottom, chlorine at the anode. Chlorine forms at the anode because Cl⁻ is the only anion and is oxidised there.
Concentrated aqueous NaCl (brine, inert electrodes): anode 2Cl⁻ → Cl₂ + 2e⁻. Cathode 2H⁺ + 2e⁻ → H₂. Solution left contains Na⁺ and OH⁻, which is NaOH. This is the industrial route to sodium hydroxide (Nelson cell).
Copper(II) sulphate with platinum or graphite electrodes: cathode Cu²⁺ + 2e⁻ → Cu. Anode 4OH⁻ → 2H₂O + O₂ + 4e⁻. Blue colour fades.
Copper(II) sulphate with copper electrodes: cathode Cu²⁺ + 2e⁻ → Cu. Anode Cu → Cu²⁺ + 2e⁻. The anode dissolves and the cathode gains mass; the blue colour stays.
Refining of copper: impure copper = anode, pure copper = cathode, CuSO₄ with a little H₂SO₄. Gold and silver fall to the bottom as anode mud. Copper obtained is 99.5% pure.
Dilute KI: cathode H₂; anode O₂ (iodine only when concentrated). Molten NaBr: reddish-brown bromine at the anode, sodium at the cathode. Concentrated CuCl₂ with carbon: copper at the cathode, chlorine at the anode.
Electroplating
Zinc plating of steel: Zn anode, steel cathode, electrolyte potassium zinc cyanide with a little sodium cyanide (cyanide is toxic, must not be dumped in rivers). Anode Zn → Zn²⁺ + 2e⁻, cathode Zn²⁺ + 2e⁻ → Zn.
Tin plating: tin anode, SnSO₄ electrolyte with a few drops of dilute H₂SO₄ (prevents hydrolysis). Tin is non-poisonous and stops corrosion, so food cans are tin plated (food acids would react with iron and form toxic substances).
Chromium plating: steel is first plated with copper or nickel (chromium does not stick to steel), then chromium. Chromium anode, chromium sulphate electrolyte with dil. H₂SO₄. Anode Cr → Cr³⁺ + 3e⁻, cathode Cr³⁺ + 3e⁻ → Cr. Used for car and bike parts.
Applications to quote for "three uses": protection from corrosion, decoration and shine, hardness and wear resistance, recovering and refining metals.
Daniell (galvanic) cell
Parts: zinc rod in 1 M ZnSO₄ (anode), copper rod in 1 M CuSO₄ (cathode), salt bridge (inverted U tube with KCl, completes the circuit and keeps neutrality), voltmeter and wires.
Electrons flow from Zn through the wire to Cu. Anode (oxidation): Zn → Zn²⁺ + 2e⁻. Cathode (reduction): Cu²⁺ + 2e⁻ → Cu. Overall: Zn + Cu²⁺ → Zn²⁺ + Cu.
For a diagram: label anode, cathode, salt bridge, voltmeter, electron flow arrow (Zn to Cu in the outer circuit), ions moving through the salt bridge.
Electrochemical series (use for "which metal is more reactive")
The metal with the lower (more negative) reduction potential is more reactive: it loses electrons more easily and acts as the anode. Values: Na⁺ −2.71, Al³⁺ −1.67, Zn²⁺ −0.76, Fe²⁺ −0.44, H⁺ 0.00, Cu²⁺ +0.34, Ag⁺ +0.80.
Hydrogen-oxygen fuel cell
Porous carbon electrodes with platinum catalyst, hot aqueous KOH in the middle, H₂ at the anode, O₂ at the cathode. Anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻. Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻. Net: 2H₂ + O₂ → 2H₂O. Gives about 0.9 V. Used in the American space programme (astronauts drink the water).
Advantages over petrol engines: very efficient (chemical energy straight to electricity), only water vapour as product, no CO₂, NOx or SO₂ so no smog, no moving parts, long life with little maintenance, hydrogen can come from renewable sources, petrol is non-renewable.
Disadvantages: hydrogen is flammable and explosive and hard to store and transport; making hydrogen by electrolysis of water is costly.
Past-paper questions with model answers
Which of Zn (−0.76 V) and Cu (+0.34 V) is more reactive? (B). Zinc, because it has the lower reduction potential, so it loses electrons more easily. Same for Al (−1.67 V) against Cu (D).
Three advantages of the hydrogen-oxygen fuel cell (B, D). Use the list above, pick three.
Reactions at the electrodes of a hydrogen-oxygen fuel cell (M): write the anode, cathode and net reaction above.
Which metal is commonly used for chrome plating, with anode and cathode reactions (M): chromium; anode Cr → Cr³⁺ + 3e⁻, cathode Cr³⁺ + 3e⁻ → Cr.
What is electroplating, two or three uses (B). Definition and uses above.
Sketch a Daniell cell (M). Diagram points above.
Sec C: electrolysis of dilute copper sulphate with copper electrodes, with equations (B). Anode Cu → Cu²⁺ + 2e⁻, cathode Cu²⁺ + 2e⁻ → Cu, copper anode shrinks, cathode grows.
Sec C: compare electrolysis of a halide in dilute and concentrated solutions (D). Dilute: oxygen at the anode (OH⁻ discharged), hydrogen at the cathode. Concentrated: halogen (e.g. Cl₂) at the anode, hydrogen at the cathode. Give the half-equations for NaCl.
Sec C: compare the fuel cell with petrol engines in four points (M). Pollution, efficiency, renewable fuel, moving parts, plus the disadvantages (storage, cost).