Electrolysis of Aqueous Solutions
Section: 4. Electrochemistry | Syllabus: Cambridge AS Level Physics 9702
Overview of Aqueous Solution Electrolysis When electrolyzing aqueous solutions, we must consider competition between ions from the dissolved compound and ions from water (H⁺ and OH⁻). This topic brings together all the rules for predicting products from different aqueous solutions.
The Complete Discharge Rules At the CATHODE (negative electrode) - REDUCTION: Rule: The ion that is LESS REACTIVE is discharged If metal ion is from a metal BELOW hydrogen → Metal deposits If metal ion is from a metal ABOVE hydrogen → Hydrogen gas forms At the ANODE (positive electrode) - OXIDATION: Rule: In order of preference: Halide ions (Cl⁻, Br⁻, I⁻) → Halogen gas Hydroxide ions (OH⁻) → Oxygen gas + water Other ions (SO₄²⁻, NO₃⁻) stay in solution → OH⁻ discharged instead Reactivity Series Reference Key section for electrolysis: Potassium (K) Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) ═══════════════ Zinc (Zn) Iron (Fe) HYDROGEN (H) ═══════════════ Copper (Cu) Silver (Ag) Gold (Au) Above hydrogen: H₂ gas at cathode Below hydrogen: Metal at cathode Example 1: Copper Sulfate Solution (CuSO₄) Ions present: Cu²⁺, SO₄²⁻, H⁺, OH⁻ At the cathode: Competition: Cu²⁺ vs H⁺ Copper is BELOW hydrogen (less reactive) Cu²⁺ wins → Copper metal deposits Cu²⁺ + 2e⁻ → Cu Observation: Pink/brown copper coats the cathode 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 Solution change: Blue solution fades (Cu²⁺ removed), becomes more acidic (SO₄²⁻ and H⁺ remain) Example 2: Magnesium Sulfate Solution (MgSO₄) Ions present: Mg²⁺, SO₄²⁻, H⁺, OH⁻ At the cathode: Competition: Mg²⁺ vs H⁺ Magnesium is ABOVE hydrogen (more reactive) H⁺ wins → Hydrogen gas forms 2H⁺ + 2e⁻ → H₂ Observation: Bubbles of colourless gas (squeaky pop test) At the anode: Competition: SO₄²⁻ vs OH⁻ OH⁻ wins → Oxygen gas forms 4OH⁻ → O₂ + 2H₂O + 4e⁻ Observation: Bubbles of colourless gas (relights glowing splint) Overall effect: Water is decomposed, MgSO₄ concentration increases Example 3: Potassium Iodide Solution (KI) Ions present: K⁺, I⁻, H⁺, OH⁻ At the cathode: Competition: K⁺ vs H⁺ Potassium is ABOVE hydrogen (very reactive) H⁺ wins → Hydrogen gas forms 2H⁺ + 2e⁻ → H₂ Observation: Bubbles of colourless gas At the anode: Competition: I⁻ vs OH⁻ Iodide is a halide I⁻ wins → Iodine forms 2I⁻ → I₂ + 2e⁻ Observation: Brown solution forms (iodine dissolves) Solution change: Solution turns brown, becomes alkaline (K⁺ and OH⁻ remain = KOH) Example 4: Silver Nitrate Solution (AgNO₃) Ions present: Ag⁺, NO₃⁻, H⁺, OH⁻ At the cathode: Competition: Ag⁺ vs H⁺ Silver is BELOW hydrogen (less reactive) Ag⁺ wins → Silver metal deposits Ag⁺ + e⁻ → Ag Observation: Shiny silver coating on cathode At the anode: Competition: NO₃⁻ vs OH⁻ Nitrate is not a halide OH⁻ wins → Oxygen gas forms 4OH⁻ → O₂ + 2H₂O + 4e⁻ Observation: Bubbles of colourless gas Example 5: Sodium Bromide Solution (NaBr) Ions present: Na⁺, Br⁻, H⁺, OH⁻ At the cathode: Na⁺ is above hydrogen H⁺ wins → Hydrogen gas 2H⁺ + 2e⁻ → H₂ At the anode: Br⁻ is a halide Br⁻ wins → Bromine forms 2Br⁻ → Br₂ + 2e⁻ Observation: Orange/brown solution (bromine dissolves) Solution change: Turns orange, becomes alkaline (NaOH forms) Summary Table: Common Solutions Solution Cathode product Anode product Solution after CuSO₄ Cu (copper) O₂ (oxygen) H₂SO₄ (acidic) AgNO₃ Ag (silver) O₂ (oxygen) HNO₃ (acidic) NaCl H₂ (hydrogen) Cl₂ (chlorine) NaOH (alkaline) KBr H₂ (hydrogen) Br₂ (bromine) KOH (alkaline) MgSO₄ H₂ (hydrogen) O₂ (oxygen) MgSO₄ (concentrated) H₂SO₄ (dilute) H₂ (hydrogen) O₂ (oxygen) H₂SO₄ (concentrated) Concentration Effects Important factor: Concentration can affect which ion is discharged!
Concentrated solutions: More likely to discharge the compound's ions Dilute solutions: More likely to discharge water's ions (H⁺ or OH⁻) Example: Dilute NaCl → mainly O₂ at anode (from OH⁻) Concentrated NaCl (brine) → mainly Cl₂ at anode (from Cl⁻) Using Inert Electrodes The examples above all use inert electrodes (graphite or platinum): Inert electrodes don't react They just allow electrons to flow Products come only from ions in the electrolyte Active electrodes (like copper) can participate in reactions - this is used in electroplating and purification (covered in other topics).
Flow Chart for Predicting Products CATHODE (negative): Is the metal below hydrogen in reactivity series? YES → Metal deposits (Cu, Ag, Au) NO → Hydrogen gas forms (Na, K, Mg, Ca, Al, Zn, Fe) ANODE (positive): Is there a halide ion present (Cl⁻, Br⁻, I⁻)?
YES → Halogen forms (Cl₂, Br₂, I₂) NO → Oxygen forms (from OH⁻) Why These Rules? At the cathode: Reduction occurs (gain electrons) The ion that gains electrons most easily is discharged Less reactive metals gain electrons more easily than H⁺ More reactive metals don't gain electrons easily (H⁺ preferred) At the anode: Oxidation occurs (lose electrons) Halides lose electrons more easily than OH⁻ OH⁻ loses electrons…
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