Electrolysis of aqueous solutions and practicals
Chemical changes • Electrolysis
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Key concepts
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Definition and basic setup
Electrolysis uses an external direct current source connected to two electrodes immersed in an electrolyte. The cathode is the negative electrode that attracts positive ions (cations). The anode is the positive electrode that attracts negative ions (anions). The electrolyte provides mobile ions that carry charge through the solution. Electrode material choice affects observed chemistry. Inert electrodes (graphite or platinum) remain chemically unchanged and allow observation of the electrolyte's reactions only. Reactive electrodes supply or remove metal atoms and change the overall products.
Ions present in aqueous solutions
Aqueous solutions contain ions from the dissolved ionic compound and H+ and OH− from water dissociation. H+ and OH− compete with dissolved ions during electrolysis. The relative ease of reduction or oxidation determines which species reacts at each electrode. Concentration affects which ions reach electrodes in greater number. Higher concentration of a particular ion increases its chance of being discharged, but standard electrode potentials and reactivity also determine which reaction occurs.
Rules for predicting products at the cathode
The cathode attracts cations. Metal cations undergo reduction to the element when the metal is less reactive than hydrogen (lower in the reactivity series). When the metal is more reactive than hydrogen or H+ is present in higher effective reactivity, hydrogen ions reduce to hydrogen gas. Practical rule: if the cation is from a less reactive metal such as Cu2+, the metal plates out at the cathode. If the cation is from a more reactive metal such as Na+, K+ or Ca2+, hydrogen gas forms instead because these metals are harder to reduce in aqueous solution.
Rules for predicting products at the anode
The anode attracts anions. Halide ions (Cl−, Br−, I−) oxidize to form halogen molecules at inert anodes. If no halide ions are present, hydroxide ions (from water) oxidize to oxygen gas. Practical rule: in solutions containing chloride ions, chlorine gas evolves at the anode. In solutions of sulfate or nitrate salts without halides, oxygen gas forms from water oxidation.
Half-equations and charge balance
Half-equations show electron transfer at each electrode. Examples: reduction of hydrogen: 2H+ + 2e− → H2. Oxidation of chloride: 2Cl− → Cl2 + 2e−. Half-equations combine to give the overall cell reaction while conserving mass and charge. Balanced half-equations clarify stoichiometry of gas volumes and mass changes during electrolysis. Electron flow in the external circuit balances the ionic electron transfers at electrodes.
Examples with inert electrodes
Aqueous sodium chloride (NaCl(aq)) produces hydrogen gas at the cathode and chlorine gas at the anode. Sodium ions remain in solution and combine with hydroxide produced at the cathode to form sodium hydroxide. Aqueous copper(II) sulfate (CuSO4) with inert electrodes produces copper metal at the cathode (Cu2+ + 2e− → Cu) and oxygen at the anode (4OH− → O2 + 2H2O + 4e−). Dilute sulfuric acid (H2SO4) gives hydrogen at the cathode and oxygen at the anode because H+ and OH− from water are the reactive species.
Practical method and qualitative tests
Standard practical setup uses a DC power supply, inert electrodes, the aqueous solution, and gas collection methods (delivery tubes into gas jars or test-tube collection). Electrodes must be labelled and current controlled. Observations include gas evolution, metal deposition, pH changes, and solution colour changes. Qualitative tests identify gases: a lit splint that reignites or makes a squeaky 'pop' indicates hydrogen; a glowing splint that relights indicates oxygen; damp blue litmus that turns white or starch-iodide paper that bleaches indicates chlorine. Gas solubility and small volumes can limit detection; repeated tests confirm identity.
Limiting factors and common errors
Factors that alter predicted products include electrode material, ion concentration, current density and overpotential. Reactive (non-inert) electrodes introduce additional reactions such as electrode dissolution. High current density can favour less positive (more easily oxidized) reactions and produce different products. Practical errors include leaking collection tubes, incomplete sealing causing gas loss, impure reagents changing ion availability, and incorrect electrode polarity. Careful control of concentration, electrode inertness and current produces reproducible observations.
Key notes
Important points to keep in mind