Comprehensive E-Note on Electrolysis

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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)

1. Under an applied potential, cations migrate to cathode and anions to anode.
2. At the cathode, reduction occurs (gain of electrons); at the anode, oxidation occurs (loss of electrons).
3. 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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Electrolysis — CBT Quiz (30)

Time: 15:00

Created by Edwin Ogie Library — revisions include expanded theory, worked examples and full quiz. Images included from the original note.