Water Electrolysis
Part of Electrochemistry
Passing direct current through water splits it into hydrogen and oxygen — two gases with profound applications in energy, metallurgy, and medicine.
Why This Matters
Water electrolysis is one of the most impactful things you can do with electricity. Pure water alone won’t conduct current well, but add a small amount of salt, acid, or alkali, and you have a system that reliably produces hydrogen and oxygen gas on demand, scaled directly to the power you put in. The process requires no rare materials, no complex machinery, and no specialized knowledge beyond what this article covers.
For a rebuilding civilization, the immediate applications are immense. Hydrogen and oxygen together form an oxyhydrogen torch — the hottest manually operated flame achievable, sufficient to cut, weld, and melt nearly any metal. Hydrogen alone is a powerful reducing agent in metallurgy, preventing oxidation during annealing and heat treatment. Oxygen enriches combustion in forges and kilns, raising temperatures beyond what air alone supports.
Every kilowatt-hour of electricity you put into an electrolyzer comes back as chemical energy stored in hydrogen — making it a form of energy storage as well as a gas generator. This matters when your power source (wind, water, solar) produces more electricity than you can use immediately.
Electrolyte Selection
Pure water has extremely high electrical resistance and electrolyzes very slowly. Adding an electrolyte — a dissolved compound that provides mobile ions — increases conductivity dramatically.
Sulfuric acid (H₂SO₄): The traditional electrolyte for water electrolysis. A 20–30% solution by weight gives excellent conductivity. The acid is not consumed — it acts as a catalyst, facilitating ion transport while water is split. Downside: corrosive, requires careful handling, and degrades some electrode materials.
Sodium hydroxide (NaOH / lye): Works equally well as an alkaline electrolyte at 20–30% concentration. Produces the same hydrogen and oxygen outputs. More compatible with nickel and iron electrodes. Can be made from wood ash leaching (potassium hydroxide substitute) if pure sodium hydroxide is unavailable.
Potassium hydroxide (KOH): Highest conductivity of the common electrolytes, preferred in professional alkaline electrolyzers. Made by leaching hardwood ash with hot water and concentrating the filtrate.
Sodium chloride (table salt): Produces hydrogen and chlorine gas (not pure oxygen), plus sodium hydroxide. Useful when you want these products, but not for pure hydrogen/oxygen production.
Practical recommendation: Start with sodium hydroxide at 20% concentration. It is relatively safe to handle compared to sulfuric acid, and potassium hydroxide can be produced locally from ash if lye is unavailable.
Electrode Materials and Cell Construction
Cathode (negative, where hydrogen forms): Less demanding — hydrogen evolution is relatively benign. Stainless steel, iron, zinc, or nickel all work well. Stainless steel is preferred for durability and resistance to the alkaline electrolyte.
Anode (positive, where oxygen forms): Under oxidizing conditions, many metals dissolve or corrode. Suitable materials:
- Nickel — excellent in alkaline electrolytes, durable, minimal contamination
- Platinum — ideal but scarce
- Carbon/graphite — available from battery salvage, slowly consumed
- Stainless steel 316 — adequate in alkaline solutions, not suitable in acid
Cell body: Any acid-resistant or alkali-resistant container. Glass jars, ceramic vessels, polyethylene containers, or polypropylene work well. Avoid metal containers (they short-circuit).
Simple open cell construction:
- Use a wide-mouth glass jar as the cell body
- Cut two electrode plates — stainless steel sheet works well (3–5 cm × 10 cm each)
- Bend the top of each plate to hang over the jar rim without touching each other
- Fill jar with electrolyte, leaving 3–4 cm of headspace for gas collection
- Connect positive terminal to anode plate, negative to cathode plate
- Apply 3–12 V DC; bubbles should appear immediately
Electrode spacing: Closer electrodes mean lower resistance and higher current for a given voltage. Typical spacing: 5–15 mm. Too close risks gas bubbles bridging across and mixing.
Operating Parameters
Voltage: Each electrolytic cell theoretically needs 1.23 V to split water. In practice, 1.8–2.2 V per cell is required to overcome overpotential. If running multiple cells in series, multiply accordingly.
Current density: 100–500 mA/cm² of electrode area is typical for practical operation. Lower current density produces purer gas and is gentler on electrodes; higher density gives faster output.
Temperature: Warmer electrolyte (40–70°C) increases conductivity and efficiency. The cell naturally warms during operation. Avoid exceeding 80°C as evaporation rates increase and electrode corrosion accelerates.
Gas output calculation (Faraday’s law): At 100% efficiency, 1 ampere for 1 hour (1 ampere-hour) produces approximately 0.42 liters of hydrogen and 0.21 liters of oxygen at atmospheric pressure. Real-world efficiency is typically 70–85%.
Divided Cell Design
For separate hydrogen and oxygen collection, a membrane or porous divider separates the cathode and anode compartments while allowing ion flow.
Simple divider: Stretch multiple layers of tightly woven cotton fabric or unglazed ceramic across the cell interior, dividing it into two chambers. Each chamber has one electrode and its own gas outlet.
Collection by water displacement: Fill a bottle completely with electrolyte, invert it over the electrode chamber outlet tube. Gas displaces liquid downward as it accumulates in the inverted bottle. Mark volume graduated lines on the bottle to measure gas production rate.
Pressure equalization: Ensure both gas collection vessels have equal back-pressure, otherwise the divider bulges and eventually leaks gas from one side to the other.
Scaling Up Production
A single 10 cm × 10 cm electrode pair at 500 mA/cm² draws 50 A and produces roughly 21 liters of hydrogen per hour. For a useful oxyhydrogen torch (consuming 2–3 liters per minute), you need 120–180 liters per hour, requiring 6–9 such cells in a bipolar stack.
Bipolar plate stacks: Arrange multiple cells in series, where the back of each anode becomes the front of the next cathode. This allows high-voltage, lower-current operation (easier to supply from generators) while maintaining total gas output. Each cell needs only 2 V, so a 24 V source runs 12 cells in series.
Water electrolysis scales linearly — double the current, double the gas. This predictability makes it straightforward to match production to your power source and consumption needs.