Redox · Lesson 3

Forcing reactions to run backwards

Electrolysis uses electrical energy to drive redox reactions that would never happen on their own. Predict the products, plate an object and refine copper.

In 1807 the young English chemist Humphry Davy connected a huge battery to a lump of slightly damp potash (potassium hydroxide). At one electrode, tiny silvery globules appeared, some of which burst into flame. According to his cousin, Davy danced around the laboratory in delight. He had made potassium, an element no one had ever seen, and within days he had made sodium the same way.

Neither metal can be extracted by heating its ore with carbon, because both hold on to their electrons far too tightly. Davy’s battery forced the electrons back onto the ions. Eighty years later the same idea, electrolysis, made aluminium cheap: before 1886 it was a precious metal, and Napoleon III is said to have served his most honoured guests with aluminium cutlery.

Today electrolysis refines the copper in our wiring, chrome-plates our taps and produces the chlorine for our water treatment. This lesson shows how to predict what it will make.

Electrolytic cells

An electrolytic cell uses an external power supply to force a non-spontaneous redox reaction. The power supply pumps electrons into the cathode (connected to the negative terminal), where reduction occurs, and pulls electrons out of the anode (positive), where oxidation occurs. Oxidation is still at the anode and reduction at the cathode; only the signs are reversed compared with a galvanic cell.

	Galvanic cell	Electrolytic cell
Energy change	chemical → electrical	electrical → chemical
Reaction	spontaneous (E°_cell positive)	non-spontaneous, driven by a power supply
Anode	oxidation, negative	oxidation, positive (connected to + terminal)
Cathode	reduction, positive	reduction, negative (connected to − terminal)
Half-cells	usually separate, with a salt bridge	usually one container

Predicting the products

Molten salts contain only the salt’s ions. Molten sodium chloride gives sodium at the cathode (Na⁺ + e⁻ → Na) and chlorine at the anode (2Cl⁻ → Cl₂ + 2e⁻).

Aqueous solutions also contain water, which can be oxidised or reduced itself. Use the E° table:

At the cathode, the strongest oxidant present (highest E°) is reduced. Ions such as Cu²⁺ and Ag⁺ are reduced to metal; ions such as Na⁺, K⁺ and Mg²⁺ are much weaker oxidants than water, so water is reduced instead: 2H₂O + 2e⁻ → H₂ + 2OH⁻.
At the anode, the strongest reductant present (lowest E°) is oxidised. Water is oxidised to oxygen (2H₂O → O₂ + 4H⁺ + 4e⁻) unless a stronger reductant is present, such as I⁻, Br⁻, or the anode metal itself (a reactive anode such as copper dissolves).
An important exception: in concentrated chloride solutions, chlorine is produced at the anode instead of oxygen, even though E° predicts oxygen, because the oxidation of water is slow at the electrode surface. In dilute chloride solutions, mainly oxygen forms.

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Electrolysis cell

Choose an electrolyte and electrodes, then watch where bubbles form and where metal deposits or dissolves. Compare dilute and concentrated sodium chloride.

Electroplating

To coat an object with a thin layer of metal, make the object the cathode, use a piece of the plating metal as the anode, and use a solution of the plating metal’s ions as the electrolyte. To silver-plate a spoon: spoon at the cathode (Ag⁺ + e⁻ → Ag), silver anode (Ag → Ag⁺ + e⁻), silver nitrate solution. The anode replaces the silver ions deposited, so the concentration of the electrolyte stays constant.

Electrorefining copper

Copper for electrical wiring must be extremely pure. Blocks of impure copper are made the anodes and thin sheets of pure copper the cathodes, in acidified copper(II) sulfate.

At the anode, copper is oxidised, along with more reactive impurities such as iron and zinc, which enter the solution as ions.
Less reactive impurities such as silver and gold are not oxidised; they fall to the bottom as anode sludge, which is valuable.
At the cathode, only copper is deposited, because Cu²⁺ is a much stronger oxidant than Fe²⁺ or Zn²⁺. The cathode grows as nearly pure copper.

Predicting electrolysis products for full marks

List every species present, including water (and the electrode metal if it is not inert).
Cathode: choose the strongest oxidant; anode: choose the strongest reductant (remember the concentrated chloride exception).
Write each half-equation and give the observation (bubbles of a gas, a metal coating, a colour change, a change in pH).

Common traps

In electrolysis the anode is positive. Oxidation is still at the anode.
Sodium metal cannot be produced from aqueous sodium chloride; water is reduced instead.
In electroplating, the object to be plated is the cathode, never the anode.

Now put it to workFive worked examples, twenty practice questions and ten harder problems on this lesson.

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