Electrochemistry

Electrochemistry is the chemistry of electricity — using an electric current to break compounds apart, and the industrial processes that depend on it. The central process is electrolysis: the breakdown of an ionic compound (molten or in solution) by passing electricity through it. CSEC Chemistry asks you to explain electrolysis in terms of ions, predict the products, and describe some industrial uses.

Conducting electricity

Ionic compounds conduct electricity only when molten or dissolved in water, because then the ions are free to move and carry charge. In the solid state the ions are locked in the lattice and cannot move, so a solid ionic compound does not conduct. Covalent (molecular) substances generally do not conduct at all, because they have no charged particles that are free to move (graphite is the notable exception, because of its delocalised electrons).

The key terms

• Electrolyte — the molten or dissolved ionic compound that is broken down.
• Electrodes — the rods (often carbon/graphite or platinum) that carry the current into and out of the electrolyte.
• Cathode — the negative electrode. Positive ions (cations, e.g. metals and H⁺) move to it.
• Anode — the positive electrode. Negative ions (anions, e.g. non-metals and OH⁻) move to it.

A helpful memory aid: PANIC — Positive is Anode, Negative Is Cathode.

What happens at the electrodes

• At the cathode, positive ions gain electrons (they are reduced) and a metal or hydrogen is deposited.
• At the anode, negative ions lose electrons (they are oxidised) and a non-metal is released.

Electrolysis of molten lead(II) bromide:

• cathode: lead ions gain electrons → lead metal forms.
• anode: bromide ions lose electrons → bromine gas forms.

Electrolysis of solutions — the rules

In a solution, water also provides H⁺ and OH⁻ ions, so there is a competition. Use these rules:

At the cathode: if the metal is more reactive than hydrogen (e.g. sodium, potassium, calcium), hydrogen gas is given off instead of the metal. If the metal is less reactive than hydrogen (e.g. copper, silver), the metal is deposited.

At the anode: if a halide (Cl⁻, Br⁻, I⁻) is present in reasonable concentration, the halogen is released. Otherwise (e.g. with sulfates or nitrates) oxygen is given off from the OH⁻ ions.

For example, electrolysis of copper(II) chloride solution gives copper at the cathode (copper is below hydrogen) and chlorine at the anode (a halide is present).

Industrial uses

• Extraction of reactive metals. Aluminium is extracted by electrolysis of molten aluminium oxide (dissolved in cryolite to lower the melting point). Metals above carbon in the reactivity series must be extracted this way because they are too reactive to be reduced by carbon.
• Purification of copper. Impure copper is the anode and pure copper the cathode; copper dissolves from the anode and deposits, pure, on the cathode. Impurities fall as anode sludge.
• Electroplating. A thin layer of one metal is coated onto another for protection or appearance — the object to be plated is the cathode and the plating metal is the anode, in a solution of the plating metal's salt (e.g. silver-plating cutlery).
• The chlor-alkali industry. Electrolysis of concentrated sodium chloride solution (brine) produces chlorine (anode), hydrogen (cathode) and sodium hydroxide — all valuable industrial chemicals.

Writing electrode half-equations

You are often asked for the half-equations at each electrode. The cathode is always a reduction (gain of electrons) and the anode an oxidation (loss of electrons). For molten lead(II) bromide:

cathode: Pb²⁺ + 2e⁻ → Pb anode: 2Br⁻ → Br₂ + 2e⁻

For the electrolysis of water-based solutions where hydrogen and oxygen are released:

cathode: 2H⁺ + 2e⁻ → H₂ anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻

Check that the electrons balance between the two electrodes — the total lost at the anode must equal the total gained at the cathode. Writing these correctly, with charges and electrons shown, is a reliable source of marks.

The effect of using active electrodes

The electrode material can change the products. With inert electrodes (carbon or platinum) the electrodes take no part in the reaction. But in copper purification the electrodes are active: at the anode, copper atoms lose electrons and dissolve into the solution (Cu → Cu²⁺ + 2e⁻), while at the cathode copper ions are deposited (Cu²⁺ + 2e⁻ → Cu). So the copper is effectively transferred from the impure anode to the pure cathode, and the impurities drop off as anode sludge (which may itself contain valuable metals such as silver and gold). Recognising the difference between inert and active electrodes explains why copper purification works the way it does.

A note on quantity of product

The more charge passed (a larger current, or the same current for longer), the more product is formed at each electrode. So doubling the time of electrolysis roughly doubles the mass of metal deposited or the volume of gas given off — a simple proportionality that practical questions sometimes test.

Common exam mistakes

• Mixing up the electrodes — the cathode is negative and attracts positive ions (PANIC).
• Forgetting the solution rules — predicting a reactive metal at the cathode when hydrogen is actually released.
• Saying a solid ionic compound conducts — it does not; the ions must be free to move.
• Not stating that ions gain/lose electrons (reduction/oxidation) at the electrodes.

Why electrolysis is a redox process

It helps to see electrolysis as redox driven by electricity. At the cathode, positive ions gain electrons — that is reduction. At the anode, negative ions lose electrons — that is oxidation. So every electrolysis is a redox reaction in which the electrical energy forces electrons to move in a direction they would not go on their own. This is the opposite of a simple cell or battery, where a spontaneous redox reaction produces electricity. Recognising this link ties electrochemistry to the redox topic and explains why the half-equations at the electrodes always show electrons being gained or lost.

Key terms to remember

• Electrolysis — breaking down a molten or dissolved ionic compound using electricity.
• Electrolyte — the molten or dissolved compound that conducts and is broken down.
• Cathode — the negative electrode; positive ions (cations) go there and are reduced.
• Anode — the positive electrode; negative ions (anions) go there and are oxidised.
• PANIC — Positive is Anode, Negative Is Cathode.
• Inert electrode — an electrode (carbon/platinum) that does not react.
• Active electrode — an electrode that takes part, e.g. copper dissolving in copper purification.
• Electroplating — coating an object (the cathode) with a thin layer of metal.

Quick recap

• Electrolysis breaks down a molten or dissolved ionic compound; the ions must be free to move.
• Cathode (−): positive ions gain electrons (reduced) → metal or hydrogen. Anode (+): negative ions lose electrons (oxidised) → non-metal.
• In solutions, a metal more reactive than hydrogen leaves hydrogen at the cathode; a halide gives its halogen at the anode, otherwise oxygen.
• Industrial uses: extracting aluminium, purifying copper, electroplating, and the chlor-alkali process.