"Water is made of the elements hydrogen and oxygen" — students can recite it from an early age, but what gives that sentence the right to be true? This experiment does. Pass electricity through water and two different gases bubble up on the two electrodes: one you can ignite, and one that relights a glowing splint — and their volumes settle into a steady 2 : 1 ratio. This guide gets that ratio on a kitchen table, with pencil leads and a single battery.

Safety first (read this section before you start)

  • Never, ever use salt water as the electrolyte. Passing current through salt water produces chlorine gas at the positive electrode — irritating and toxic; one sniff is enough to leave you choking. Use baking soda solution (sodium bicarbonate) or sodium sulfate solution instead — their only job is to carry the current. This is the one non-negotiable rule of the experiment.
  • The power source must be a battery: one 9V battery, or three to four AA cells in series. Never plug the wires into a household wall outlet — that can kill.
  • The gas collected at the negative electrode is hydrogen, which is flammable. Collect only a little for testing (one small test tube is enough), point the tube's mouth away from everyone before igniting, keep away from any flame other than your test flame, and have an adult do this step.
  • The solution near the electrodes warms slightly during electrolysis. Pencil leads (graphite) are brittle — handle them gently, and if one snaps, don't fish for it in the water with your fingers.
  • Don't let the two electrodes touch each other — that short-circuits the battery, which will heat up fast.
  • When you finish, disconnect the battery first, then dismantle the setup. The waste liquid is dilute baking soda water and can go straight down the sink.

Materials and equipment

  • One 9V battery (or 3–4 AA cells with a battery holder) and two wires with alligator clips;
  • Two pencil leads as electrodes (mechanical-pencil leads are too thin — better to whittle the wood off ordinary pencils, or buy 2B drafting leads);
  • A clear glass or plastic container holding about 300 mL of water;
  • A small spoonful of baking soda (about 3 g) as the electrolyte;
  • Two small test tubes (or syringe barrels with the needles removed, or clear pen caps) to collect the gases;
  • A long-reach lighter or matches, and a thin wooden splint (a toothpick or the wooden stick of an incense stick works);
  • Paper and pen for records, and a phone (to photograph the bubble volumes at different times).

Procedure

Step 1: mix the electrolyte. Pour about 300 mL of warm water into the container, add a small spoonful of baking soda, and stir until fully dissolved. Pure water barely conducts electricity, so this step cannot be skipped — it is the difference between waiting thirty seconds and waiting thirty minutes.

Step 2: mount the electrodes. Stand the two pencil leads upright in the water, 3–4 cm apart, and clip an alligator clip onto the dry end of each. Tape them to the rim of the glass so they stay vertical and never touch each other.

Step 3: a first test run. Connect the two wires to the battery's terminals. Watch the leads closely: both fizz with bubbles, but one clearly bubbles more, and faster, than the other. Note which battery terminal the busier lead is connected to.

Step 4: collect the gases. Disconnect the battery. Fill both test tubes with water inside the container, seal each mouth with a finger, and invert one over each pencil lead (mouth down, always below the water line), then let go. Reconnect the battery: the bubbles will rise into the tubes and push the water out bit by bit.

Step 5: read the volumes, test the gases. Run the current for 5–10 minutes, until the tube at the negative electrode is more than half full, then disconnect. First compare the gas heights in the two tubes. Then test each: lift out the negative-electrode tube (mouth down) and bring it quickly to a flame — you'll hear a soft pop. Push a glowing splint into the positive-electrode tube — the splint bursts back into flame.

Interactive simulation: this experiment has a drag-and-play simulator in our Interactive Lab. The lab interface is currently Chinese-only — the controls are simple sliders and switches, so it is still easy to explore.

What to record

Check the observation notes against these points:

  • Within a second or two of connecting the battery, both pencil leads are coated in fine bubbles — the one on the negative terminal clearly busier and faster.
  • After collecting for a while, the negative-electrode tube holds roughly twice as much gas as the positive-electrode tube. Measure the gas columns with a ruler and the ratio usually lands between 1.8 and 2.1.
  • Both gases are colorless and odorless — you cannot tell them apart by looking; only the tests can.
  • The negative electrode's gas ignites near a flame with a pale blue flame and a soft pop. That is hydrogen.
  • A glowing splint pushed into the positive-electrode tube relights. That is oxygen — it doesn't burn itself, but it supports burning.
  • The water level slowly drops, but the baking soda is not consumed: evaporate the water afterwards and the white powder is still there.
  • The positive electrode's gas usually comes up a little short of theory, so the ratio doesn't quite reach a full 2 : 1. That is not a mistake — see the explanation below.

How it works

What happens when the current flows is a chemical change, written as:

2H₂O ——electricity——> 2H₂↑ + O₂↑

One memory line covers the products: oxygen at positive, hydrogen at negative — two of hydrogen for one of oxygen. The positive electrode yields oxygen, the negative yields hydrogen, and the hydrogen's volume is double the oxygen's. That 2 : 1 is no coincidence — it comes straight from the coefficients in the equation: every 2 water molecules split into 2 hydrogen molecules and 1 oxygen molecule. Under the same conditions, gas volumes are in the same ratio as molecule counts, so V(H₂) : V(O₂) = 2 : 1.

What deserves even more thought is where the atoms go. Before the reaction: 2 water molecules, containing 4 hydrogen atoms and 2 oxygen atoms in all. After: 2 hydrogen molecules and 1 oxygen molecule — still 4 hydrogen atoms and 2 oxygen atoms. Not one atom appeared or vanished; they only regrouped into new molecules. That is the single most central sentence of middle school chemistry: in a chemical change, molecules can be divided, atoms cannot. The whole logic of the law of conservation of mass lives right here.

Because electrolyzing water yields hydrogen and oxygen and nothing else, we earn the right to the conclusion: water is composed of the elements hydrogen and oxygen. Note the word "composed," not "contains" — there are no hydrogen molecules or oxygen molecules hiding in water, only water molecules; the hydrogen and oxygen molecules are newly made in the reaction.

Also keep the volume ratio separate from the mass ratio. Hydrogen's relative atomic mass is 1 and oxygen's is 16, so 2H₂ weighs 4 while O₂ weighs 32: m(H₂) : m(O₂) = 1 : 8. Volume ratio 2 : 1, mass ratio 1 : 8 — exams love to put the two side by side, and mixing them up loses everything.

So what is the baking soda doing? It doesn't "turn into" anything — it just lets the water conduct. Pure water holds almost no free-moving ions, so the current is nearly unmeasurable and a bubble takes forever to appear. Add a pinch of baking soda (or sodium sulfate, or sodium hydroxide) and the solution fills with mobile ions; the current jumps, and the reaction becomes fast enough to watch. The additive takes no part in the reaction and its mass is unchanged afterwards, which is why it is called an electrolyte rather than a reactant. By contrast, table salt (sodium chloride) also conducts — but at the positive electrode it turns chloride ions into chlorine gas: no longer water electrolysis, and toxic besides. That is why this experiment must use baking soda.

Finally, that "a little short": oxygen is more soluble in water than hydrogen (roughly twice as soluble at 20 °C), so some freshly made oxygen dissolves back into the water; and the carbon (graphite) electrode erodes slightly in the strongly oxidizing environment at the positive electrode, eating a bit more oxygen. So a home-built setup often reads 2.1 : 1 or even 2.3 : 1, with the first gas collected carrying the biggest error. Only with platinum electrodes and a purpose-built water electrolysis apparatus, as in a school lab, does the ratio come very close to 2 : 1.

Tips for teachers and parents

  • Timing and grouping: 30 minutes, in groups of four; the teacher hands out the 9V batteries.
  • Before hands go on: salt water is absolutely forbidden (it makes chlorine gas); the hydrogen test is an adult's job.

The step most often skipped. Let students run pure water first and hit the wall, then add the baking soda — the electrolyte's role explains itself. State the "never salt water" rule separately before anyone starts, and explain why.

How to know it has really landed. A student meets the bar when they can point at the negative-electrode tube, explain why it holds exactly twice as much as the other, and write out the word equation for electrolyzing water.

Going further, and common misconceptions

  • Misconception 1: "Electrolysis proves water contains hydrogen molecules and oxygen molecules" — wrong. Water holds only water molecules; the hydrogen and oxygen molecules are new substances made by the reaction. The correct statement is that water is composed of the elements hydrogen and oxygen.
  • Misconception 2: "Volume ratio 2 : 1, so the mass ratio is 2 : 1 too" — wrong. The mass ratio is 1 : 8, because one oxygen atom is far heavier than one hydrogen atom.
  • Misconception 3: "Electrolysis is a physical change, since it just takes water apart" — wrong. New substances (hydrogen and oxygen) are produced, so it is a chemical change. A physical change only alters state, like water freezing into ice.
  • Misconception 4: "Salt water can stand in for baking soda water" — dangerous. The positive electrode would produce toxic chlorine gas. Never do it.
  • Misconception 5: "The ratio isn't exactly 2 : 1, so the experiment failed" — no. Oxygen's higher solubility and the wearing carbon electrode make some deviation normal; seeing "the negative side clearly doubled" already achieves the goal.
  • Misconception 6: "Electrolyzing water needs a very high voltage" — it doesn't. In theory it starts at 1.23 V; in practice one 9V battery is more than enough.
  • Think about it: what happens if you swap the battery's terminals? (The gases swap sides along with the electrodes — the busier lead is still whichever one is connected to the negative terminal.)
  • Work it out: electrolyzing 36 g of water yields at most how many grams of hydrogen? (Hydrogen's mass fraction in water is 2 ÷ 18 ≈ 11.1%, so 4 g.)
  • Read on with Build an Atom from Modeling Clay to make "atoms cannot be divided" concrete, or with Why Does a Candle Go Out? — which covers exactly the other condition hydrogen needs in order to burn.

When a student can point at the negative-electrode tube and explain why it holds exactly twice as much as the other one, this experiment has earned its place.