Aqueous sodium chloride is electrolyzed with inert electrodes to produce hydrogen gas at the cathode, chlorine gas at the anode, and sodium hydroxide remaining in solution.

1. Prepare a brine electrolyte (e.g., ~3–10% w/v NaCl in water). Pour into an electrolysis cell or beaker to near the top.
2. Insert two inert electrodes (carbon/graphite or platinum) through a stopper or clamp so they do not touch. Connect to a low-voltage DC supply (≈6–12 V) with the ammeter in series.
3. Fill two small test tubes with brine, cover their mouths, invert them over each electrode (still filled) to collect gases by displacement.
4. Switch on the power and adjust current to a steady few hundred mA to a few A for a class demo. Observe bubbles: more at the cathode (H₂) and at the anode (Cl₂).
5. When enough gas is collected, close the taps (or keep tubes inverted) and test:
   1. Cathode tube: bring a lit splint to the mouth—expect a squeaky pop (hydrogen).
   2. Anode tube: hold damp blue litmus near the mouth—expect it to redden then bleach (chlorine). Do not sniff directly.
6. Sample the bulk solution and add a drop of universal indicator—expect green → blue/purple, showing sodium hydroxide formation.
7. (Optional) Continue longer with very dilute brine to observe some oxygen at the anode (less common initially).

Electrolysis of brine using carbon fibre electrodes part 1 MVI 6014 - Nigel Baldwin:

📄 ELECTROLYSIS of aqueous SODIUM CHLORIDE SOLUTION and sodium bromide and potassium iodide solutions NaCl(aq), NaBr(aq) and KI(aq) - Doc Brown: https://www.docbrown.info/page01/ExIndChem/electrochemistry03.htm

• Use a Hofmann voltammeter to measure gas volumes and compare H₂:Cl₂ (theoretical 1:1; observed Cl₂ may be lower due to solubility/side reactions).
• Compare concentrated vs. very dilute brine to discuss when O₂ competes with Cl₂ at the anode.
• Swap carbon for platinum electrodes and compare stability/observations.
• Collect chlorine into a trap leading to dilute NaOH to form bleach (NaOCl), then test with iodide–starch paper (teacher demo only).

• Wear splash goggles, gloves, and a lab coat; work with good ventilation or a fume hood—chlorine is toxic and irritating.
• Keep current modest; do not scale up gas volumes. Use only small tubes for collection.
• Never inhale gases; test chlorine indirectly (indicator paper) rather than smelling.
• Keep electrodes from touching to avoid short circuits; dry hands before handling power leads.
• Neutralize/chlorine-scrub anode off-gas if running for extended times (e.g., through dilute NaOH). Dispose solutions according to local rules.
• Do not ignite gases directly at the anode outlet; perform ignition tests only on small, collected samples away from the cell.

• What half-reactions occur at each electrode?
• What is the overall reaction in brine electrolysis?
• Why is hydrogen produced instead of sodium metal in aqueous solution? (Na⁺ reduction is unfavorable in water; water is reduced to H₂.)
• Why might the measured volume of chlorine be less than hydrogen? (Cl₂ dissolves/reacts with formed NaOH; some hydrolysis.)
• How does dilution affect anode products? (At low [Cl⁻], OH⁻/H₂O oxidation can produce O₂, increasing O₂:Cl₂ over time.)
• How can you confirm NaOH formation in the electrolyte? (Indicator turns blue/purple; pH measurement shows alkaline.)