Electrolytic cells

The idea

• An electrolytic cell does the opposite of a galvanic cell: it uses an external power supply to force a non-spontaneous redox reaction to happen.

• The reaction has a negative $E°_{cell}$, so it will not go on its own. The power supply pushes electrons the wrong way round, and the applied voltage must exceed the magnitude of that negative $E°_{cell}$.

• Electricity is consumed, not produced. That energy requirement is the defining feature, and the criteria name it: at every grade you must refer to the energy input an electrolytic cell requires.

The parts

• Oxidation is still at the anode and reduction still at the cathode. Only the signs are reversed, because in this cell the electrodes are driven by an external supply rather than driving one.

Electrolysis of a molten salt

The simplest case, because only the salt's own ions are present.

Molten sodium chloride:

• Cathode (−): $\text{Na}^+(l) + e^- \rightarrow \text{Na}(l)$

• Anode (+): $2\text{Cl}^-(l) \rightarrow \text{Cl}_2(g) + 2e^-$

• Overall: $2\text{NaCl}(l) \rightarrow 2\text{Na}(l) + \text{Cl}_2(g)$

• $E°_{cell} = (-2.71) - (+1.36) = -4.07$ V, strongly negative — which is why this needs a substantial applied voltage and a furnace at over 800 °C.

• Observations: molten sodium metal collects at the cathode; pale yellow-green chlorine gas bubbles off at the anode.

Electrolysis of an aqueous solution

• Water is present, and water can be both oxidised and reduced. So at each electrode there is a competition, and $E°$ decides the winner.

At the cathode, the competitors are the metal cation and water:

$$2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^- \qquad E° = -0.83 \text{ V}$$

• The species with the more positive $E°$ is reduced.
  • Cu²⁺ (+0.34 V) beats water → copper metal is deposited.
  • Na⁺ (−2.71 V) loses to water → hydrogen is produced, not sodium.

At the anode, the competitors are the anion and water:

$$\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O} \qquad E° = +1.23 \text{ V}$$

• The species with the less positive $E°$ is oxidised (it gives up electrons more readily).

  • SO₄²⁻ is very hard to oxidise → oxygen from water is produced.
  • Br⁻ (+1.07 V) beats water → bromine is produced.

• The general result: reactive metals cannot be extracted from aqueous solution, because water is reduced in preference. Sodium, magnesium and aluminium must be electrolysed from molten compounds.

Electrolysis of brine

• Concentrated sodium chloride solution is industrially important, producing three useful products at once.

• Cathode (−): $2\text{H}_2\text{O}(l) + 2e^- \rightarrow \text{H}_2(g) + 2\text{OH}^-(aq)$

  • Water beats Na⁺, because −0.83 V is far more positive than −2.71 V.

• Anode (+): $2\text{Cl}^-(aq) \rightarrow \text{Cl}_2(g) + 2e^-$

  • On $E°$ alone, water at +1.23 V should be oxidised in preference to chloride at +1.36 V — but in concentrated brine, chlorine is produced. Two factors override the standard prediction: the very high chloride concentration, which is far from the standard 1 mol L⁻¹, and the fact that oxygen evolution at the electrode surface is kinetically hindered.

• In solution: Na⁺ and OH⁻ remain, so sodium hydroxide is the third product.

• Observations: colourless gas bubbles at the cathode (H₂, which burns with a squeaky pop); pale yellow-green gas with a choking smell at the anode (Cl₂, which bleaches damp litmus); the solution around the cathode turns the indicator purple, showing OH⁻.

• This is a valuable case for the standard because it is a documented example of $E°$ predicting one product and observation showing another — exactly the kind of link between calculation and observation the criteria ask for.

Electroplating

• Electroplating deposits a thin layer of metal onto an object, for appearance or corrosion protection.
• The setup:
  • Cathode: the object to be plated — reduction deposits metal onto it.
  • Anode: the plating metal — it dissolves, replenishing the solution.
  • Electrolyte: a solution of the plating metal's ions.

For silver-plating a spoon:

• Cathode (spoon): $\text{Ag}^+(aq) + e^- \rightarrow \text{Ag}(s)$

• Anode (silver bar): $\text{Ag}(s) \rightarrow \text{Ag}^+(aq) + e^-$

• The overall chemical change is nothing — silver is simply transported from the anode to the cathode. $E°_{cell} = 0$ for the pure transfer, so only a small voltage is needed to overcome resistance.

• The object must be the cathode. Reduction deposits metal, and reduction happens at the cathode.

Electrolysis in New Zealand

• Tiwai Point near Bluff smelts aluminium by electrolysis of Al₂O₃ dissolved in molten cryolite. Aluminium's $E°$ of −1.66 V makes it impossible to obtain from aqueous solution or by carbon reduction, so electrolysis is the only route.
• The smelter consumes around 13% of New Zealand's total electricity, supplied by the Manapōuri power station. The chemistry decides the geography: an electrolytic process must be sited where electricity is cheap and abundant.