Electrolysis of Aqueous NaCl & H₂SO₄
Section: 4. Electrochemistry | Syllabus: Cambridge IGCSE Chemistry (0620)
What's Different About Aqueous Electrolysis? Aqueous electrolysis is the electrolysis of an ionic compound dissolved in water, rather than molten. The key difference from molten electrolysis: Water molecules are also present in the solution Water can break down: H₂O → H⁺ + OH⁻ We now have four types of ions present (not just two) This creates competition at each electrode Products are less predictable than with molten compounds Ions Present in Aqueous Solutions Example: Aqueous sodium chloride (NaCl solution) From compound From water Na⁺ (sodium ions) H⁺ (hydrogen ions) Cl⁻ (chloride ions) OH⁻ (hydroxide ions) At the cathode: Na⁺ and H⁺ compete At the anode: Cl⁻ and OH⁻ compete Which Ion Wins at the Cathode?
At the cathode (negative electrode), REDUCTION occurs Rule: If the metal is MORE reactive than hydrogen → H⁺ is discharged (hydrogen gas forms) Rule: If the metal is LESS reactive than hydrogen → Metal ion is discharged (metal forms) Reactivity series (partial): K > Na > Ca > Mg > Al > (Carbon) > Zn > Fe > HYDROGEN > Cu > Ag > Au Metals above hydrogen: H⁺ wins (hydrogen gas produced) Metals below hydrogen: Metal ion wins (metal produced) Which Ion Wins at the Anode?
At the anode (positive electrode), OXIDATION occurs Simple rule: If a halide ion (Cl⁻, Br⁻, I⁻) is present → halogen gas forms If NO halide present (e.g., sulfate SO₄²⁻, nitrate NO₃⁻) → oxygen gas forms from OH⁻ Electrolysis of Aqueous Sodium Chloride (NaCl solution) Setup: Electrolyte: Sodium chloride solution (brine) Electrodes: Inert (graphite or platinum) Ions present: Na⁺, Cl⁻, H⁺, OH⁻ At the cathode: Competition: Na⁺ vs H⁺ Sodium is MORE reactive than hydrogen H⁺ wins → hydrogen gas forms 2H⁺ + 2e⁻ → H₂ Observation: Bubbles of colourless gas (hydrogen) At the anode: Competition: Cl⁻ vs OH⁻ Chloride (halide) is present Cl⁻ wins → chlorine gas forms 2Cl⁻ → Cl₂ + 2e⁻ Observation: Bubbles of greenish-yellow gas (chlorine) What's left in solution: Na⁺ and OH⁻ remain This is sodium hydroxide (NaOH) solution!
Industrial Production from Brine The Chlor-Alkali Industry Electrolysis of concentrated sodium chloride solution (brine) produces three important industrial chemicals: Chlorine gas (Cl₂) at the anode - for water purification, plastics (PVC), bleach Hydrogen gas (H₂) at the cathode - for margarine production, ammonia manufacture Sodium hydroxide solution (NaOH) remains - for soap, paper, detergents This is one of the most important industrial processes!
Electrolysis of Dilute Sulfuric Acid (H₂SO₄) Setup: Electrolyte: Dilute sulfuric acid Electrodes: Inert (graphite or platinum) Ions present: H⁺, SO₄²⁻, H⁺, OH⁻ (from water) Note: Extra H⁺ from the acid, plus H⁺ and OH⁻ from water At the cathode: H⁺ ions attracted Hydrogen gas forms 2H⁺ + 2e⁻ → H₂ Observation: Bubbles of colourless gas At the anode: Competition: SO₄²⁻ vs OH⁻ Sulfate is NOT a halide OH⁻ wins → oxygen gas forms 4OH⁻ → O₂ + 2H₂O + 4e⁻ Observation: Bubbles of colourless gas (oxygen) Overall effect: Water is being decomposed: 2H₂O → 2H₂ + O₂ Sulfuric acid concentration increases (water is removed) Volume Ratio in Dilute Sulfuric Acid When electrolyzing dilute sulfuric acid: Volume ratio: Hydrogen : Oxygen = 2 : 1 Why?
Overall equation: 2H₂O → 2H₂ + O₂ 2 moles of hydrogen produced for every 1 mole of oxygen Same volume ratio (at same temperature and pressure) Test: Collect gases in separate test tubes - you should get twice as much hydrogen as oxygen!
Testing the Gases Produced Gas Test Result Hydrogen (H₂) Lit splint Burns with squeaky pop Oxygen (O₂) Glowing splint Relights the splint Chlorine (Cl₂) Damp litmus paper Bleaches it white Comparison: NaCl vs H₂SO₄ Electrolysis Feature NaCl solution Dilute H₂SO₄ Cathode product H₂ (hydrogen) H₂ (hydrogen) Anode product Cl₂ (chlorine) O₂ (oxygen) Solution after NaOH (alkaline) More concentrated H₂SO₄ Industrial use Chlor-alkali industry Producing pure water samples Why Hydrogen Always Forms at Cathode For common salt solutions (Na⁺, K⁺, Ca²⁺, Mg²⁺, etc.): These metals are all MORE reactive than hydrogen H⁺ ions are preferentially discharged The metal ions stay in solution Hydrogen gas always produced at cathode Exception: Copper sulfate solution (Cu²⁺) Copper is LESS reactive than hydrogen, so copper metal forms at the cathode instead!
Summary of Discharge Rules At the CATHODE (negative): If metal is MORE reactive than hydrogen → H₂ forms If metal is LESS reactive than hydrogen → Metal forms At the ANODE (positive): If HALIDE present (Cl⁻, Br⁻, I⁻) → Halogen gas forms (Cl₂, Br₂, I₂) If NO halide (SO₄²⁻, NO₃⁻, etc.) → O₂ forms from OH⁻ Common Mistakes to Avoid Forgetting water ions: Always remember H⁺ and OH⁻ from water are present Wrong reactivity comparison: Compare metal with hydrogen, not with water Ignoring halides: Halides always discharge at anode (not OH⁻) Wrong gas tests: Squeaky pop for H₂, relights for O₂, bleaches for Cl₂ Forgetting NaOH remains: In NaCl electrolysis, sodium hydroxide solution is left Wrong volume ratio…
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