Water Electrolysis
Summary
Electrolysis1 uses an electric current to force a reaction that won’t happen on its own, splitting water into hydrogen and oxygen gas. You’ll be able to explain why the water needs an electrolyte2 to conduct, what the anode and cathode3 each do, and why the two gases come off in a 2:1 ratio that reveals water’s formula. You’ll also uncover a hidden twist: the current quietly makes acid at one electrode and base at the other, showing that water isn’t really “splitting” in half at all.
History
Water electrolysis was first performed in 1800, just weeks after Alessandro Volta4 announced his electric pile (battery). William Nicholson and Anthony Carlisle connected a Voltaic pile to water and observed gas bubbles - hydrogen at one electrode, oxygen at the other.
This experiment proved that water was a compound of hydrogen and oxygen (H₂O), not an element as Aristotle had taught. It helped establish the law of definite proportions and supported John Dalton’s atomic theory. The 2:1 volume ratio of hydrogen to oxygen provided direct evidence for the molecular formula.
Today, “green hydrogen” produced by electrolyzing water with renewable electricity is seen as a key technology for decarbonizing industry. The same reaction demonstrated in a classroom may help power the future.
Hazards & preparation
PPE: safety glasses.
- Hydrogen is flammable. Keep sparks and open flames away from the collected gas. Only test it deliberately (see Going further), in tiny amounts.
- Electricity in water: use a low-voltage DC source (a 9 V battery or a 6 V supply). Never use mains power.
- Do not use table salt as the electrolyte. With salt, the current oxidizes chloride to chlorine gas at the anode — a toxic, choking gas — instead of oxygen. Use baking soda or sodium sulfate (explained below).
Disposal: a baking-soda or sodium-sulfate solution is harmless — pour it down the drain with plenty of water. See the Safety page.
Materials
- Water - 500mL
- Baking soda - 2 tablespoons (30g) OR sodium sulfate (Glauber5’s salt) - 1 tablespoon (15g)
- 9V battery (or 6V DC power supply for faster results)
- Two graphite electrodes from #2 pencils (expose ~3cm of graphite)
- Two test tubes
- Beaker (600mL) or a tall glass
- Alligator clip wires
Do not substitute table salt for the electrolyte — see Hazards above. Salt seems like the obvious choice, but chloride ions change the anode reaction to produce chlorine gas.
Procedure
- Dissolve the electrolyte in 500mL water in the beaker and stir until clear.
- Fill two test tubes with the solution and invert one over each electrode without trapping air.
- Connect the battery leads to the electrodes with the alligator clips.
- Watch for gas bubbles rising into the tubes — they should appear within seconds.
- Note which side makes gas faster: the cathode (negative) produces hydrogen at about twice the volume of the oxygen at the anode (positive).
What you should see
Fine bubbles stream off both electrodes almost immediately. The negative electrode (cathode) bubbles noticeably faster, and as gas collects in the inverted tubes the hydrogen side fills to roughly twice the volume of the oxygen side. The solution itself stays clear and colorless.
| Symptom | Likely cause | Fix |
|---|---|---|
| No bubbles anywhere | Weak connection, dead battery, or too little electrolyte | Check the clips, use a fresh battery, add more baking soda |
| Bubbles at only one electrode | One clip loose or not submerged | Reseat that connection and make sure the graphite is in the solution |
| Yellow-green tint or a “swimming-pool” smell | Table salt was used as the electrolyte (chlorine!) | Stop, ventilate, and switch to baking soda or sodium sulfate |
| Solution goes cloudy or an electrode crumbles | Metal electrode corroding | Use graphite (pencil leads), not metal |
The reactions
\[\ce{2 H2O -> 2 H2 + O2}\]
water → hydrogen gas + oxygen gas
At the cathode (negative), water is reduced to hydrogen:
\[\ce{4 H2O + 4 e^- -> 2 H2 + 4 OH^-}\]
At the anode (positive), water is oxidized to oxygen:
\[\ce{2 H2O -> O2 + 4 H^+ + 4 e^-}\]
The Science
Electrical energy splits water molecules:
- Cathode (negative): electrons reduce water, producing hydrogen gas.
- Anode (positive): water is oxidized, producing oxygen gas.
- The 2:1 ratio of H₂:O₂ demonstrates the stoichiometry of water.
It isn’t really “splitting”
Look again at the by-products of those two half-reactions. The cathode leaves behind OH⁻ ions, making its surroundings alkaline; the anode leaves behind H⁺ ions, making its surroundings acidic — exactly the blue and red you saw with the indicator. So electrolysis isn’t one water molecule cleanly breaking in two: it’s two independent reactions at two electrodes, each leaving a different ion behind. The hydrogen and oxygen never came from the same molecule, which is precisely why they emerge on opposite sides of the cell. Stir the solution and the H⁺ and OH⁻ recombine into water, cancelling out — so the overall bookkeeping collapses back to the neat \(\ce{2 H2O -> 2 H2 + O2}\) and the acid and base vanish from the summary. That hidden acid/base split is the foundation of a great deal of practical electrochemistry: keep the two sides apart with an ion-exchange membrane and you can harvest the acid and alkali separately, refine metals, or build a flow battery — all from salt, water, and electricity.
The electrolyte makes the water conductive by supplying ions that carry the current between the electrodes. As a simplification we say it isn’t consumed — and for baking soda or sodium sulfate that’s essentially true, because their ions are harder to oxidize or reduce than water itself. This is exactly why table salt is a poor choice: chloride ions are easier to oxidize than water, so at the anode they’re consumed to make chlorine gas instead of oxygen.
Questions to Explore
Why does pure water conduct almost no electricity? Tap water with dissolved minerals conducts electricity, but pure distilled water barely does. Why does electrical conduction in water require dissolved ions — and what are ions actually doing in the solution that makes current flow possible?
Hint / answer
Current in a liquid is carried by moving ions, not by loose electrons. Pure water self-ionizes only very slightly (about 10⁻⁷ mol/L of H⁺ and OH⁻), so there are almost no charge carriers. Dissolving an electrolyte floods the water with ions that drift toward the electrodes and complete the circuit.
Why does the cathode produce twice the volume of gas as the anode? The hydrogen-to-oxygen ratio is exactly 2:1. What does this tell you about the molecular formula of water — and how did observing this ratio in 1800 help prove that water was H₂O and not HO or H₂O₂?
Hint / answer
Splitting 2 H₂O gives 2 H₂ but only 1 O₂ — twice as many hydrogen molecules, hence twice the gas volume (equal volumes of gas hold equal numbers of molecules). That fixed 2:1 ratio pointed directly to two hydrogens per one oxygen in the water molecule.
Why shouldn’t you use table salt as the electrolyte? With salt in the water, the anode produces chlorine gas instead of oxygen — and this happens even with graphite electrodes. Why are chloride ions oxidized in preference to water, and how does the choice of electrolyte steer which reaction happens at the anode?
Hint / answer
At the anode the current oxidizes whatever gives up electrons most readily. Although water is “supposed” to oxidize to O₂ at a lower voltage, that reaction is sluggish (it has a high overpotential), so in practice chloride wins and leaves as Cl₂. Changing the electrolyte to sulfate or bicarbonate — ions that resist oxidation — lets oxygen form instead.
What is the minimum voltage needed for electrolysis? Water is thermodynamically stable — it takes energy to split it. The theoretical minimum voltage for water electrolysis is about 1.23 V. A 9 V battery works but most of that voltage is “wasted.” Where does the extra voltage go, and why is real electrolysis always less efficient than the theoretical minimum?
Hint / answer
The 1.23 V is the ideal thermodynamic minimum. Real cells need more to overcome the overpotential at each electrode (the sluggishness of gas formation) and the resistance of the solution and wires. That surplus energy shows up as heat, which is why practical electrolysis never reaches the theoretical efficiency.
Could you scale this up to produce fuel? Green hydrogen production by electrolysis of water using renewable electricity is a real technology. What challenges would you face scaling from a 9 V battery and a beaker to an industrial hydrogen facility — and what makes electrolysis economically viable or not compared to producing hydrogen from natural gas?
Hint / answer
Scaling up means fighting the same inefficiencies (overpotential, resistance) plus new ones: durable catalysts and electrodes, membranes to keep H₂ and O₂ apart safely, gas purification, and huge amounts of cheap, clean electricity. Today most hydrogen is still made from natural gas because it’s cheaper; electrolysis becomes competitive only when renewable power is abundant and inexpensive.
Why does one electrode turn the water acidic and the other alkaline? With a pH indicator, the anode goes red and the cathode goes blue. What does each half-reaction leave behind in the water, and why does mixing the solution make the colours cancel?
Hint / answer
The anode reaction releases H⁺ ions (acidic, red); the cathode reaction releases OH⁻ ions (alkaline, blue). They form at different electrodes because they come from different reactions — the water isn’t splitting in one place. Stir them together and H⁺ + OH⁻ → H₂O, so the acid and base neutralise and the colour returns to neutral.
Going further
- Confirm the 2:1 ratio by collecting gas in the two graduated test tubes and comparing the levels.
- Make the acid and base visible. Repeat with sodium sulfate plus a pH indicator (or red-cabbage extract) and watch red form at the anode and blue at the cathode — the reveal that water electrolysis is two separate reactions, not a clean split.
- Identify the gases: a lit splint near a little collected hydrogen gives a sharp “pop”; a glowing splint relights in the oxygen. Do this with tiny amounts, away from the rest of the setup.
- Related electrochemistry: put a current to work depositing color instead of gas in Titanium Anodizing.
Footnotes
Electrolysis — Driving a non-spontaneous reaction by passing an electric current through an electrolyte.↩︎
Electrolyte — A substance that conducts electricity when dissolved or molten because it splits into mobile ions.↩︎
Anode / cathode — The two electrodes of an electrochemical cell — oxidation happens at the anode, reduction at the cathode.↩︎
Alessandro Volta — Italian physicist (1745–1827) who invented the first battery, the voltaic pile, in 1800.↩︎
Johann Rudolf Glauber — German-Dutch chemist (1604–1670) who first made sodium sulfate (‘Glauber’s salt’) and potassium permanganate.↩︎