In short: passing electric current through water splits it into hydrogen and oxygen gas in a fixed 2:1 ratio by volume. At 5 amps for 1 hour (25 °C, 1 atm, 100% efficient), you get about 2.28 L of hydrogen and 1.14 L of oxygen. Use the calculator above for your exact current, time and efficiency.
These figures are planning aids based on Faraday's laws and the ideal gas law. For hydrogen-handling, ventilation or pressure-vessel work, follow your equipment manufacturer's guidance and applicable safety codes.
How to calculate electrolysis gas yield
Water splits under a DC current as 2H₂O → 2H₂ + O₂. Faraday's law connects the charge you pass through the cell to the moles of gas produced: hydrogen needs 2 moles of electrons per mole of gas, oxygen needs 4, because oxygen's reaction transfers twice as many electrons per molecule.
Worked example. Run 5 A for 1 hour (3,600 s) at 100% efficiency. Charge = 5 × 3,600 = 18,000 coulombs. Moles of H₂ = 18,000 / (2 × 96,485) = 0.0933 mol; moles of O₂ = 0.0466 mol. At 25 °C (298.15 K) and 1 atm (101,325 Pa), the molar gas volume is R×T/P = 8.314 × 298.15 / 101,325 = 24.46 L/mol. That gives 0.0933 × 24.46 = 2.28 L of hydrogen and 0.0466 × 24.46 = 1.14 L of oxygen — exactly what the calculator shows for those inputs.
Why the ratio is always 2:1. Every water molecule holds two hydrogen atoms and one oxygen atom. Splitting it apart can never change that ratio, so hydrogen gas always comes out at exactly twice the moles — and, at equal temperature and pressure, twice the volume — of oxygen gas.
Why efficiency changes the answer
The calculation above is the theoretical maximum — every electron doing useful work splitting water. Real cells lose some current to side reactions, gas that dissolves back into the electrolyte or crosses the membrane, and resistive heating. The fraction of charge that actually produces gas is the Faradaic efficiency. Well-built modern PEM electrolyzers run at 95–99%; alkaline electrolyzers typically run 90–95%; a rough demonstration setup with bare electrodes can be lower still. Lowering the efficiency slider scales both gas volumes down proportionally — the 2:1 ratio itself never changes, since it comes from the reaction stoichiometry, not from how efficiently the cell runs.
Temperature and pressure work differently: they don't change how many moles of gas you get for a given charge, only how much space those moles occupy, through the ideal gas law. Warmer gas and lower pressure both mean a larger volume for the same amount of hydrogen and oxygen.
Gas yield per amp-hour of charge
At 25 °C, 1 atm and 100% Faradaic efficiency — the values this calculator scales for your own current, time and efficiency.
| Charge | Hydrogen (L) | Oxygen (L) | Hydrogen mass (g) |
|---|---|---|---|
| 1 Ah | 0.456 | 0.228 | 0.038 |
| 5 Ah | 2.282 | 1.141 | 0.188 |
| 10 Ah | 4.565 | 2.282 | 0.376 |
| 20 Ah | 9.129 | 4.565 | 0.752 |
| 50 Ah (50 A for 1 hr) | 22.823 | 11.412 | 1.881 |
| 100 Ah | 45.645 | 22.824 | 3.762 |
Worked examples
Car battery charger, 5 A for 1 hour
2.28 L H₂
At 25 °C and 1 atm, ideal efficiency, that's about 2.28 L of hydrogen and 1.14 L of oxygen — enough to fill roughly two large soda bottles with hydrogen.
Bench power supply, 1 A for 10 minutes
76 mL H₂
A small classroom or lab demonstration at 1 A for 10 minutes (600 coulombs) yields about 76 mL of hydrogen and 38 mL of oxygen at room conditions.
Small electrolyzer stack, 100 A for 8 hours at 95%
347 L H₂
A 100 A stack run for 8 hours at 95% Faradaic efficiency produces about 347 L (0.347 m³) of hydrogen — roughly 28.6 g, or 0.0286 kg — plus about 173 L of oxygen.
What else affects the result
Electrolyte and water purity. Pure water conducts poorly, so practical cells add an electrolyte (an acid, base or salt) to carry current. The electrolyte itself is not consumed and doesn't change the Faraday-law math — it only affects how much voltage is needed to push a given current through, which feeds into the optional energy estimate.
Real-gas behavior. This calculator uses the ideal gas law, which is accurate to within a fraction of a percent for hydrogen and oxygen at everyday temperatures and pressures. It becomes less accurate only at very high pressures (tens of atmospheres) or near-cryogenic temperatures, well outside typical bench or industrial electrolysis conditions.
Safety. Hydrogen is flammable and can form explosive mixtures with air across a wide concentration range. Any real electrolysis setup needs adequate ventilation and a source of ignition kept well away — this tool is a planning calculator, not a substitute for equipment safety guidance.
Frequently asked questions
How much hydrogen does electrolysis produce from 1 amp-hour of charge?
At typical conditions (25 °C, 1 atmosphere, 100% Faradaic efficiency), 1 amp-hour of charge produces about 0.456 liters of hydrogen gas and 0.228 liters of oxygen gas — always a 2:1 ratio by volume, per Faraday's laws of electrolysis.
What is Faraday's law of electrolysis?
Faraday's law states that the amount of a substance produced at an electrode is proportional to the electric charge passed through the cell. Splitting water requires 2 moles of electrons per mole of hydrogen produced and 4 moles per mole of oxygen, because water splits as 2H₂O → 2H₂ + O₂.
Why is the hydrogen-to-oxygen ratio always 2:1?
The reaction 2H₂O → 2H₂ + O₂ always produces exactly twice as many moles — and, at the same temperature and pressure, twice the volume — of hydrogen as oxygen, because each water molecule contains two hydrogen atoms for every one oxygen atom.
Why isn't real-world electrolysis 100% efficient?
Some current is lost to side reactions, gas crossover through the membrane, and resistive heating rather than splitting water. This is captured by the Faradaic efficiency, typically about 90–95% for alkaline electrolyzers and 95–99% for modern PEM electrolyzers, versus the 100% theoretical maximum.
How much energy does it take to split water into hydrogen and oxygen?
The thermodynamic minimum is about 1.23 volts per cell, equivalent to roughly 33 kWh of electrical energy per kilogram of hydrogen. Real cells run at 1.8–2.4 volts because of resistive and activation losses, raising practical energy use to roughly 50–55 kWh per kilogram.
Sources & references
- Faraday, M. (1834). Experimental Researches in Electricity, Seventh Series. Philosophical Transactions of the Royal Society.
- CODATA (2018). Recommended values of the Faraday constant and the molar gas constant.
- National Institute of Standards and Technology (NIST). Ideal Gas Law reference data.
John Mac, PE
Water Systems Specialist · Austin, Texas
Texas-licensed Professional Engineer with 13+ years in residential water systems and 180+ completed projects. Writes and reviews the calculators on watercalculator.us. Read full bio
Texas PE (2016) · BS Civil Engineering, Hydraulics (2013)