Electrolysis

Chapter in a nutshell: Electrolysis is the chemical decomposition of an electrolyte (in molten or aqueous state) by the passage of direct electric current, in which positive ions (cations) migrate to the cathode and are reduced while negative ions (anions) migrate to the anode and are oxidised. The chapter covers electrolytes vs non-electrolytes, the theory of ionic dissociation, the electrochemical series, selective (preferential) discharge of ions, the classic examples (molten lead bromide, acidified water, copper sulphate), and the industrial applications of electroplating and the electro-refining of copper.

1. Conductors, Non-conductors and the Idea of Electrolysis

The word electrolysis comes from electro (electricity / flow of electrons) and lysis (splitting). It is the process of decomposition of a chemical compound in the molten or aqueous state, accompanied by a chemical change, by the passage of direct electric current.

Substances are classified by how they conduct:

2. Electrolytes and Non-electrolytes

An electrolyte is a compound which, in the molten state or in aqueous solution, allows electric current to pass through it and is decomposed by it. Electrolytes are electrovalent (ionic) compounds (metal + non-metal) or polar covalent compounds that form ions in water.

A non-electrolyte is a compound which, neither in solution nor in the molten state, allows current to pass, and is not decomposed at the electrodes. Non-electrolytes contain only molecules in solution.

3. Electrodes, Anode, Cathode, Cations, Anions

Electrodes are the metal plates, graphite rods or gas-carbon rods immersed in the electrolyte, through which the current enters and leaves the cell.

• Anode: the electrode connected to the positive terminal of the battery. It is the oxidising electrode — oxidation occurs here. Anions migrate to it. (Memory aid: A→ADD→positive; Anion to Anode, Cation to Cathode — A to A and C to C.)
• Cathode: the electrode connected to the negative terminal of the battery. It is the reducing electrode — reduction occurs here. Cations migrate to it.

Ions are atoms or groups of atoms carrying a charge equal to their valency.

4. Oxidation, Reduction and the Redox Nature of Electrolysis

Electronic concept: Oxidation = loss of electron(s) (Zn − 2e⁻ → Zn²⁺ ; Na − e⁻ → Na⁺ ; Fe²⁺ − e⁻ → Fe³⁺); Reduction = gain of electron(s) (Cu²⁺ + 2e⁻ → Cu ; Fe³⁺ + e⁻ → Fe²⁺ ; S + 2e⁻ → S²⁻).

Classical concept: Oxidation = addition of oxygen / removal of hydrogen (2Mg + O₂ → 2MgO ; H₂S + Cl₂ → 2HCl + S); Reduction = removal of oxygen / addition of hydrogen (CuO + H₂ → Cu + H₂O ; 3Cl₂ + 2NH₃ → N₂ + 6HCl).

An oxidising agent accepts electrons (MnO₂, conc. HNO₃, O₂, chlorine); a reducing agent provides electrons (C, H₂, CO, H₂S, Na, Zn). Electrolysis is a redox reaction: reduction occurs at the cathode and oxidation at the anode simultaneously, the electrons gained at the cathode exactly equalling those lost at the anode.

5. Theory of Electrolytic Dissociation (Arrhenius, 1887)

Postulates: (1) An electrolyte in water dissociates into free mobile cations and anions, allowing current to flow. (2) All ions carry a charge and conduct. (3) Conductivity depends on the concentration of ions. (4) Total + charge = total − charge, so the solution is neutral (electrolytic equilibrium). (5) Non-electrolytes do not ionise — only molecules. (6) Degree of dissociation = extent of breakup into ions.

Arrhenius vs Modern concept: Arrhenius thought water ionises electrolytes. The modern view: electrovalent compounds are ionic even in the solid state, but their ions are held immobile by strong electrostatic forces; water (or melting) only frees the existing ions.

Dissociation (ions already exist, e.g. NaCl, KCl, PbBr₂, lead bromide) is distinguished from ionisation (ions are produced from polar covalent molecules, e.g. HCl, H₂CO₃, NH₄OH):

6. The Electrochemical (Activity) Series

Metals lose valence electrons to form positive ions, but the ease of doing so varies. Arranging metals in decreasing order of their tendency to lose electrons (form cations) gives the electrochemical series:

K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, (H), Cu, Hg, Ag, Au, Pt

with cations K⁺, Ca²⁺, Na⁺, Mg²⁺, Al³⁺, Zn²⁺, Fe³⁺(/Fe²⁺), Ni²⁺, Sn⁴⁺, Pb²⁺, H⁺, Cu²⁺, Hg²⁺, Ag⁺, Au³⁺, Pt⁴⁺.

Key points: (1) A metal higher in the series displaces one lower from its salt solution. (2) Metals below hydrogen (Cu, Hg, Ag, Au) cannot displace hydrogen from acids; those above can. (3) Ions of metals lower in the series are discharged more easily at the cathode. (4) Hydrogen is placed among the metals as it forms a positive ion.

Series of anions (increasing ease of discharge): SO₄²⁻ < NO₃⁻ < Cl⁻ < Br⁻ < I⁻ < OH⁻ — the lower an anion, the more easily it is discharged at the anode.

7. Selective (Preferential) Discharge of Ions

When two or more ions of the same sign compete at one electrode, only one is preferentially discharged. Three factors decide which: (1) position in the electrochemical series — if other factors are equal, the ion lower in the series wins; (2) relative concentration — a much higher concentration can make an ion discharge even if it lies higher; (3) nature of the electrodes — inert electrodes (Pt, graphite) take no part, active electrodes (Cu, Ni, Ag) themselves enter the reaction.

8. Example I — Electrolysis of Molten Lead Bromide (inert graphite electrodes)

• Cathode (reduction): Pb²⁺ + 2e⁻ → Pb
• Anode (oxidation): Br⁻ − e⁻ → Br ; Br + Br → Br₂
• Overall: PbBr₂(l) → Pb(s) + Br₂(g)

Observations: dark reddish-brown bromine fumes at the anode; greyish-white lead at the cathode.

9. Example II — Electrolysis of Acidified Water (platinum electrodes, Hoffmann voltameter)

• Cathode (reduction): H⁺ + e⁻ → H ; H + H → H₂↑
• Anode (oxidation): 4OH⁻ − 4e⁻ → 2H₂O + O₂ (via OH + OH → H₂O + O ; O + O → O₂)
• Overall: 2H₂O —electrical energy→ 2H₂↑ + O₂↑

Result: H₂ at cathode, O₂ at anode, volume ratio H₂ : O₂ = 2 : 1. Acidity decreases at the cathode and increases at the anode; the net process is simply decomposition of water.

10. Example III — Electrolysis of Copper(II) Sulphate Solution

CuSO₄ solution (acidified with a little conc. H₂SO₄ to raise conductivity) in a glass/porcelain jar. Dissociation: CuSO₄ → Cu²⁺ + SO₄²⁻ ; H₂O ⇌ H⁺ + OH⁻. Ions present: Cu²⁺, H⁺, SO₄²⁻, OH⁻. The product at the anode depends on the electrode:

(a) Platinum (inert) anode, Pt/Cu cathode: Cathode — Cu²⁺ + 2e⁻ → Cu (Cu²⁺ below H⁺, discharged in preference), copper deposits. Anode — 4OH⁻ − 4e⁻ → 2H₂O + O₂, oxygen evolves. Observation: blue colour fades to colourless, as discharged Cu²⁺ is not replaced by the inert anode.

(b) Copper (active) anode and copper cathode: Cathode — Cu²⁺ + 2e⁻ → Cu (cathode grows). Anode — Cu − 2e⁻ → Cu²⁺ (anode dissolves, shrinks); neither SO₄²⁻ nor OH⁻ is discharged. Observation: blue colour does not fade — the dissolving anode replenishes Cu²⁺. This is the principle of copper electro-refining.

11. Application — Electroplating

Electroplating is coating a superior metal over an inferior metal by electric current, to prevent corrosion and improve appearance. Conditions: (a) article at the cathode; (b) electrolyte must contain ions of the plating metal; (c) anode = pure plating metal; (d) article clean; (e) low current for a long time; (f) direct current (a.c. gives an uneven deposit).

Silver-plating (e.g. a copper bell): electrolyte = sodium argentocyanide Na[Ag(CN)₂] (silver nitrate is avoided — it decomposes fast and gives an uneven layer); cathode = article; anode = pure silver. Cathode: Ag⁺ + e⁻ → Ag ; Anode: Ag − e⁻ → Ag⁺.

Nickel-plating: electrolyte = nickel sulphate; cathode = article; anode = pure nickel. Cathode: Ni²⁺ + 2e⁻ → Ni ; Anode: Ni − 2e⁻ → Ni²⁺.

12. Application — Electro-refining (Purification) of Copper

Impurities are removed electrolytically to give very pure metal. Impure copper = anode; thin pure copper = cathode; electrolyte = acidified CuSO₄. Cathode: Cu²⁺ + 2e⁻ → Cu (pure copper deposits); Anode: Cu − 2e⁻ → Cu²⁺ (impure copper dissolves). Pure copper transfers anode → cathode; soluble impurities stay in solution, insoluble ones form anode mud. In all electro-refining, pure metal deposits at the cathode.

13. Acids, Bases and Salts as Electrolytes

Acids ionise to give H⁺ (H₃O⁺), bases give OH⁻, salts give metal cations + acid-radical anions. The activity series shows the tendency to form ions: Na > Mg > Fe > Cu. Zinc, above hydrogen, displaces it from dilute acids: Zn + H₂SO₄(dil.) → ZnSO₄ + H₂↑. Copper, below hydrogen, cannot: Cu + H₂SO₄(dil.) → no reaction. Iron displaces copper: Fe + CuSO₄ → FeSO₄ + Cu, so CuSO₄ cannot be stored in an iron vessel.

Worked Examples

1. Why is molten (not solid) PbBr₂ used for electrolysis? In the solid, Pb²⁺ and Br⁻ are locked in the lattice by electrostatic forces and cannot move. On melting (>380 °C) the lattice breaks, ions become mobile and conduction occurs: Pb²⁺ + 2e⁻ → Pb (cathode); 2Br⁻ − 2e⁻ → Br₂ (anode).

2. On electrolysing dilute CuSO₄ with Cu electrodes, why does the blue colour stay constant, but fades with Pt electrodes? With Cu anode: Cu − 2e⁻ → Cu²⁺ replaces the Cu²⁺ removed at the cathode, so [Cu²⁺] is constant → colour stays. With Pt (inert) anode: OH⁻ is discharged (O₂ evolves), no Cu²⁺ is added, so Cu²⁺ is steadily depleted → blue fades to colourless.

3. Why is H⁺ discharged but not Na⁺ during electrolysis of dilute NaCl? Both Na⁺ and H⁺ reach the cathode, but hydrogen lies well below sodium in the electrochemical series, so H⁺ is discharged preferentially: H⁺ + e⁻ → H; H + H → H₂↑. Sodium is not deposited.

4. Volume of O₂ when 40 cm³ of H₂ is collected (acidified water, same conditions)? Ratio H₂ : O₂ = 2 : 1, so O₂ = 40 ÷ 2 = 20 cm³.

5. Mass of copper deposited by 0.2 mol electrons through CuSO₄? Cu²⁺ + 2e⁻ → Cu, so Cu = 0.2 ÷ 2 = 0.1 mol = 0.1 × 64 = 6.4 g.

6. Moles of electrons to liberate 1 mol O₂ at the anode in water? 4OH⁻ − 4e⁻ → 2H₂O + O₂, so 4 moles of electrons.

7. In copper refining 0.5 mol copper dissolves at the anode — what at the cathode (100% efficiency)? Anode Cu − 2e⁻ → Cu²⁺ gives 1 mol e⁻; cathode deposits 0.5 mol = 32 g pure copper.

8. Why is acidified-water electrolysis an example of catalysis? Dilute H₂SO₄ only speeds the feeble ionisation of water (H₂O ⇌ H⁺ + OH⁻) without being consumed — so it acts as a catalyst.

Key Terms — Quick Glossary

Common Mistakes to Avoid

• Writing that anode is negative — the anode is positive (connected to + terminal); the cathode is negative. Anions go to anode, cations to cathode.
• Saying oxidation occurs at the cathode — oxidation is at the anode, reduction at the cathode.
• Claiming solid NaCl conducts because it contains ions — ions exist but are immobile; only molten/aqueous NaCl conducts.
• Forgetting that in concentrated NaCl, Cl⁻ (not OH⁻) is discharged due to high concentration.
• Saying sodium metal forms at the cathode in electrolysis of aqueous NaCl — it is H₂, because H⁺ lies below Na⁺ in the series.
• Using a.c. or silver nitrate for electroplating — use d.c., and for silver use sodium argentocyanide.
• Confusing electrolytic cell (electrical → chemical) with electrochemical/voltaic cell (chemical → electrical).
• Forgetting that with a Cu anode the CuSO₄ blue colour does not fade (anode replenishes Cu²⁺).

Likely Exam Questions (with crisp answers)

1. Define electrolysis. Decomposition of an electrolyte in the molten/aqueous state by passing direct current, accompanied by chemical change.
2. Why does solid sodium chloride not conduct electricity? Its ions are held immobile by electrostatic forces; there are no free ions.
3. Distinguish ionisation from dissociation. Ionisation forms ions from polar covalent molecules (HCl); dissociation separates ions already present in ionic compounds (KCl).
4. Name a strong and a weak electrolyte. Strong: HCl / NaOH / NaCl. Weak: acetic acid / NH₄OH / carbonic acid.
5. Name a salt that is a weak electrolyte. Ammonium acetate (CH₃COONH₄).
6. Name a base that is not an alkali. Zinc hydroxide, Zn(OH)₂.
7. Name an inert and an active electrode. Inert: platinum (or graphite). Active: copper.
8. What is observed during electrolysis of molten PbBr₂? Reddish-brown bromine fumes at the anode; greyish-white lead at the cathode.
9. Why is a silica crucible used for molten lead bromide? Silica is non-reactive, withstands high temperature and is almost a non-conductor.
10. In electrolysis of acidified water, what is the H₂:O₂ ratio and why dilute H₂SO₄? Ratio 2:1 by volume; dilute H₂SO₄ is non-volatile (unlike HNO₃/HCl).
11. Why is electrolysis of acidified water an example of catalysis? The acid only speeds water's feeble ionisation and is not consumed.
12. In electrolysis of CuSO₄ with Pt electrodes, what gas is at the anode and what happens to the colour? Oxygen at anode; blue colour fades to colourless.
13. Why does CuSO₄'s blue colour stay constant with copper electrodes? Cu anode dissolves (Cu − 2e⁻ → Cu²⁺), replacing Cu²⁺ deposited at the cathode.
14. Why is direct current (not a.c.) used in electroplating? A.c. reverses direction, preventing a smooth, even, adherent deposit.
15. In electroplating with silver, name the electrolyte, anode and cathode. Electrolyte: sodium argentocyanide; anode: pure silver; cathode: article to be plated.
16. In electro-refining of copper, what is the anode, cathode and electrolyte? Anode: impure copper; cathode: thin pure copper; electrolyte: acidified CuSO₄.
17. Why can zinc displace hydrogen from acids but copper cannot? Zn is above hydrogen (more electropositive) in the series; Cu is below it.
18. Electrolysis is a redox reaction — explain. Reduction (gain of e⁻) at cathode and oxidation (loss of e⁻) at anode, simultaneously.

Chemical Equations & Formulas (quick reference)

Common electrolytes: Acids — H₂SO₄, HNO₃, HCl, H₃PO₄ ; Bases — NaOH, KOH, Ca(OH)₂, NH₄OH ; Salts — NaCl, CuSO₄, PbBr₂.

Oxidation (loss of e⁻): Zn − 2e⁻ → Zn²⁺ ; Na − e⁻ → Na⁺ ; Fe²⁺ − e⁻ → Fe³⁺ ; 2Mg + O₂ → 2MgO ; H₂S + Cl₂ → 2HCl + S.

Reduction (gain of e⁻): Cu²⁺ + 2e⁻ → Cu ; Fe³⁺ + e⁻ → Fe²⁺ ; S + 2e⁻ → S²⁻ ; CuO + H₂ → Cu + H₂O ; 3Cl₂ + 2NH₃ → N₂ + 6HCl.

Dissociation / ionisation: NaCl ⇌ Na⁺ + Cl⁻ ; PbBr₂ ⇌ Pb²⁺ + 2Br⁻ ; HCl —H₂O→ H⁺ + Cl⁻ ; H⁺ + H₂O → H₃O⁺ ; H₂O ⇌ H⁺ + OH⁻ ; H₂SO₄ ⇌ 2H⁺ + SO₄²⁻ ; CuSO₄ → Cu²⁺ + SO₄²⁻ ; NiSO₄ → Ni²⁺ + SO₄²⁻ ; Na[Ag(CN)₂] → Na⁺ + Ag⁺ + 2CN⁻.

Molten lead bromide: Cathode Pb²⁺ + 2e⁻ → Pb ; Anode 2Br⁻ − 2e⁻ → Br₂ ; Overall PbBr₂(l) → Pb(s) + Br₂(g).

Acidified water: Cathode 2H⁺ + 2e⁻ → H₂↑ ; Anode 4OH⁻ − 4e⁻ → 2H₂O + O₂↑ ; Overall 2H₂O → 2H₂↑ + O₂↑ (H₂ : O₂ = 2 : 1).

Copper sulphate — inert (Pt) anode: Cathode: Cu²⁺ + 2e⁻ → Cu ; Anode: 4OH⁻ − 4e⁻ → 2H₂O + O₂. Copper sulphate — active (Cu) anode: Cathode: Cu²⁺ + 2e⁻ → Cu ; Anode: Cu − 2e⁻ → Cu²⁺.

Electroplating / refining:

• Silver: Ag⁺ + e⁻ → Ag (cathode) ; Ag − e⁻ → Ag⁺ (anode)
• Nickel: Ni²⁺ + 2e⁻ → Ni (cathode) ; Ni − 2e⁻ → Ni²⁺ (anode)
• Copper refining: Cu²⁺ + 2e⁻ → Cu (cathode) ; Cu − 2e⁻ → Cu²⁺ (anode)

Displacement / activity series:

• Zn + H₂SO₄(dil.) → ZnSO₄ + H₂↑ ; Cu + H₂SO₄(dil.) → no reaction
• Fe + CuSO₄ → FeSO₄ + Cu

Electrochemical series (cathode discharge ease increases down): K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, H, Cu, Hg, Ag, Au, Pt. Anion discharge ease (increases →): SO₄²⁻ < NO₃⁻ < Cl⁻ < Br⁻ < I⁻ < OH⁻.

Faraday: 1 Faraday = 96,500 coulombs of electricity.