Comprehensive E-Note on Electrolysis
Edwin Ogie Library
Electrolysis — Comprehensive E-Note (Detailed)
Comprehensive E-Note on Electrolysis
What you’ll learn — objectives & overview
- Introduction to electrolysis — definitions and concepts
- Electrolytes vs non-electrolytes; ionic conduction
- Faraday’s laws & quantitative electrolysis (worked calculations)
- Electrode reactions for key systems (CuSO₄, H₂SO₄, NaCl, molten salts)
- Factors affecting discharge and electrode potentials
- Electrochemical cells, electrode potentials and the electrochemical series
- Applications: electroplating, extraction, purification, chlor-alkali industry
- Corrosion and protection (cathodic protection, sacrificial anodes)
- Worked examples + 30-question CBT quiz (with timer)
Introduction (short)
Electrolysis uses electrical energy to force a non-spontaneous redox reaction. Passing current through an electrolyte causes ions to move and react at electrodes: oxidation at the anode and reduction at the cathode.
Use the navigation buttons below to move through pages. The quiz is at the end.
Electrolytes vs Non-electrolytes & Ionic Conduction
Electrolytes — substances that dissociate into ions in solution (or molten) and conduct electricity: strong electrolytes (e.g., NaCl, HCl, KOH) dissociate completely; weak electrolytes (e.g., acetic acid) partially dissociate.
Non-electrolytes — molecular compounds that do not form ions in solution (e.g., sugar, ethanol) and do not conduct current.
Ionic conduction (mechanism)
- Under an applied potential, cations migrate to cathode and anions to anode.
- At the cathode, reduction occurs (gain of electrons); at the anode, oxidation occurs (loss of electrons).
- Net current is carried by ion movement in the solution and electron flow in the external circuit.
Example: NaCl dissolved in water → Na⁺ and Cl⁻ ions; apply current: Na⁺ moves toward cathode, Cl⁻ to anode.
Faraday’s Laws & Quantitative Electrolysis
First law: mass ∝ total charge (Q) passed. Second law: masses produced by the same charge are proportional to equivalent weights.
Useful relation: Q = I × t (Coulombs). Faraday constant: F ≈ 96485 C·mol⁻¹ (per mole of electrons).
Worked example 1 — Copper deposition
Given: Q = 1.00×10⁵ C passed through Cu²⁺ solution. Equivalent weight (Cu) = 31.75 g per eq.
Mass = (Q / F) × equivalent weight = (1.00×10⁵ / 96485) × 31.75 ≈ 32.9 g.
Worked example 2 — Determine current
If 0.5 g of Ag (molar mass 107.87 g/mol, Ag⁺ + e⁻ → Ag) is deposited in 10 minutes, what is the average current?
moles Ag = 0.5 / 107.87 = 0.004636 mol. moles e⁻ = 0.004636 (one e⁻ per Ag). Q = n·F = 0.004636×96485 ≈ 447.7 C. Time = 600 s. I = Q/t ≈ 0.746 A.
Electrode Reactions — Common Systems
- Dilute H₂SO₄ (aq): cathode → H₂ (2H₂O + 2e⁻ → H₂ + 2OH⁻), anode → O₂ (2H₂O → O₂ + 4H⁺ + 4e⁻).
- Aqueous CuSO₄ (inert anode): cathode → Cu²⁺ + 2e⁻ → Cu(s); anode → O₂ if anode inert.
- CuCl₂ (aq): cathode → Cu; anode → Cl₂ (2Cl⁻ → Cl₂ + 2e⁻).
- Molten NaCl: cathode → Na(s); anode → Cl₂(g).
Worked example — Why water is discharged instead of Na⁺ (aqueous NaCl)
Reduction potentials: Na⁺/Na is very negative (~−2.71 V) while water reduction is easier under aqueous conditions (→ H₂). Therefore water is reduced and H₂ evolves at cathode.
Factors Affecting Discharge of Ions
Which ion is discharged depends on:
- Standard reduction potential (E°)
- Concentration (Le Chatelier-like effect)
- Nature of electrode (inert vs reactive)
- Overpotential and applied voltage
- Temperature
Example: in brine electrolysis, concentration and electrode type are chosen to favour Cl₂ production and NaOH formation.
Electrochemical Cells & Standard Potentials
The electrochemical series ranks half-reactions by E°. Cell EMF: E°cell = E°(cathode) − E°(anode). Positive E°cell → spontaneous galvanic cell.
Example — Zn/Cu cell
E°(Cu²⁺/Cu)=+0.34 V, E°(Zn²⁺/Zn)=−0.76 V → E°cell = 0.34 − (−0.76) = +1.10 V (galvanic).
Applications — Industry & Laboratory
Major applications:
- Electrolytic refining: copper, silver.
- Electroplating: chromium, nickel plating for corrosion resistance or aesthetics.
- Chlor-alkali industry: production of NaOH, Cl₂ and H₂ from brine.
- Extraction of reactive metals: aluminium by Hall-Heroult (molten cryolite + Al₂O₃).
Worked example — Copper refining (overview)
Impure copper anode dissolves to Cu²⁺; copper is deposited as pure metal at cathode; impurities either fall as anode slime or remain in solution.
Corrosion as an Electrolytic Process & Protection
Corrosion is a redox process where metal (anode area) oxidizes and electrons flow to cathodic areas. Methods of protection include:
- Cathodic protection (sacrificial anode such as Zn on iron)
- Impressed current cathodic protection
- Coatings, paints and electroplating
More Worked Examples
Example — calculating mass from current
Find mass of Cu deposited by a current of 2.0 A after 2 hours from Cu²⁺ solution. (Cu²⁺ + 2e⁻ → Cu)
Q = I t = 2 × (2×3600) = 14400 C. Moles e⁻ = Q/F = 14400/96485 ≈ 0.1493 mol e⁻. Moles Cu = 0.1493/2 = 0.07465 mol. Mass Cu = 0.07465 × 63.55 ≈ 4.74 g.
Summary & Exam Tips
- Always write half-reactions and balance electrons before Faraday calculations.
- Remember which species discharge in aqueous solutions (water often discharges if metal cation reduction potential is very negative).
- Practice industrial examples (chlor-alkali, aluminum, copper refining) and be able to describe the cell set-up and products.
CBT Quiz — Start
When you press Start Quiz you will see 30 questions and a 15:00 timer. Each question shows explanation after submission.
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