Electrolysis Fundamentals
Part of Electrochemistry
Electrolysis uses electrical current to drive chemical reactions that would not occur spontaneously. It is how you split water into hydrogen and oxygen, extract metals from ores, and produce essential chemicals from salt.
Why Electrolysis Matters for Rebuilding
Electrolysis is the reverse of a battery: instead of chemistry producing electricity, electricity drives chemistry. This gives you the power to:
- Produce hydrogen and oxygen from water for fuel, welding, and medical use
- Extract metals from their ores or solutions (aluminum, sodium, magnesium)
- Purify copper to electrical grade (99.99% pure)
- Produce chlorine and sodium hydroxide from salt — two of the most important industrial chemicals
- Electroplate metals for corrosion protection and decoration
With a working generator and basic electrolysis equipment, a small community gains access to chemical processes that previously required large-scale industrial plants.
How Electrolysis Works
The Basic Setup
Every electrolysis cell has four components:
- Electrolyte: A liquid that conducts electricity (water with dissolved salt, acid, or alkali)
- Anode (positive electrode): Connected to the positive terminal of the power supply
- Cathode (negative electrode): Connected to the negative terminal
- Power supply: DC source providing enough voltage to drive the desired reaction
When current flows through the electrolyte:
- At the cathode: Positive ions gain electrons (reduction) — metals deposit, hydrogen evolves
- At the anode: Negative ions lose electrons (oxidation) — oxygen evolves, metals dissolve, chlorine forms
DC Only
Electrolysis requires direct current. Alternating current reverses the reaction every half-cycle, producing nothing useful. If your power source is AC, you must rectify it first using a diode bridge.
Minimum Voltage
Each electrolysis reaction requires a minimum voltage (decomposition voltage) to proceed:
| Reaction | Minimum Voltage | Practical Voltage |
|---|---|---|
| Water splitting (H2 + O2) | 1.23V | 1.8-2.5V |
| Copper refining | 0.34V | 0.5-1.0V |
| Salt (NaCl) electrolysis | 2.19V | 3.0-4.0V |
| Aluminum extraction | 1.67V | 4.0-5.0V |
The practical voltage is always higher than the theoretical minimum due to overvoltage (extra energy needed to overcome electrode surface effects and resistance).
Faraday’s Laws of Electrolysis
These quantitative laws tell you exactly how much material you will produce for a given amount of electricity:
First Law
The mass of substance deposited or dissolved is directly proportional to the quantity of electricity (current x time).
Mass = (Current x Time x Atomic Weight) / (n x 96,485)
Where n = number of electrons transferred per atom and 96,485 is Faraday’s constant (coulombs per mole of electrons).
Second Law
For the same quantity of electricity, the masses of different substances deposited are proportional to their equivalent weights.
Practical Calculations
| To Produce | Current x Time Needed | Example |
|---|---|---|
| 1 gram copper | 3,040 coulombs (0.85 Ah) | 1A for 50 minutes |
| 1 gram zinc | 2,960 coulombs (0.82 Ah) | 1A for 49 minutes |
| 1 liter hydrogen | ~3,600 coulombs (1 Ah) | 1A for 60 minutes |
| 1 gram chlorine | 2,720 coulombs (0.76 Ah) | 1A for 45 minutes |
Efficiency Rule of Thumb
In practice, electrolysis is 60-90% efficient. Some current goes to side reactions, heating, and gas dissolution. When planning production, assume 70% efficiency and adjust based on actual results.
Electrode Selection
Choosing the right electrode material determines whether your electrolysis works correctly or creates unwanted side reactions.
Cathode Materials
The cathode is where useful products typically deposit or evolve:
| Material | Best For | Notes |
|---|---|---|
| Stainless steel | General purpose, hydrogen production | Durable, inert |
| Copper | Copper refining/plating | Becomes the product itself |
| Iron/steel | Alkaline electrolysis | Cheap, effective |
| Lead | Acid electrolysis | Resists acid corrosion |
| Carbon/graphite | Universal fallback | Cheap, moderately durable |
Anode Materials
The anode is the more challenging electrode because it experiences oxidation (corrosion):
| Material | Durability | Best For | Notes |
|---|---|---|---|
| Graphite | Moderate | Salt electrolysis, general | Slowly consumed |
| Platinum | Excellent | All applications | Extremely rare/expensive |
| Lead dioxide | Good | Sulfuric acid solutions | Forms protective coating |
| Stainless steel | Poor-moderate | Alkaline solutions only | Corrodes in acid |
| Titanium | Excellent | All applications | Hard to find |
| Sacrificial metal | N/A | Plating (dissolves by design) | Copper, zinc, nickel |
Anode Degradation
In acid or salt solutions, iron and steel anodes dissolve rapidly, contaminating your electrolyte. Use carbon or lead anodes for these applications. Inspect anodes regularly and replace before they fail.
Water Electrolysis — Step by Step
The most universally useful electrolysis application:
Setup
- Fill a glass or plastic container with water
- Add electrolyte: 1-2 tablespoons of sodium hydroxide (lye) OR baking soda per liter (do NOT use salt — it produces chlorine gas)
- Insert two stainless steel or nickel electrodes, spaced 2-5cm apart
- Connect to a DC power source at 3-6V
- Place inverted glass jars or bottles over each electrode to collect gas
Operation
- Hydrogen bubbles at the cathode (negative) — collected in one jar
- Oxygen bubbles at the anode (positive) — collected in the other jar
- Hydrogen volume is exactly twice the oxygen volume (H2O → H2 + 1/2 O2)
Gas Safety
Hydrogen is highly explosive when mixed with air (4-75% concentration). Oxygen accelerates combustion of any flammable material. Never collect these gases near open flames. Store separately. Ventilate the workspace. A hydrogen explosion from electrolysis is a real and documented hazard.
Applications of Produced Gases
- Hydrogen: Fuel for heating and cooking, lifting gas for balloons, reducing agent for metal smelting, fuel cells
- Oxygen: Medical support, welding torch (oxy-hydrogen flame reaches 2,800 degrees C), metal cutting
Salt Electrolysis (Chlor-Alkali Process)
Electrolyzing brine (concentrated salt water) produces three essential chemicals simultaneously:
- Chlorine gas at the anode — disinfectant, water purification, chemical feedstock
- Sodium hydroxide (lye) at the cathode — soap making, chemical processing, cleaning
- Hydrogen gas at the cathode — fuel
Simple Membrane Cell
- Dissolve as much salt as possible in warm water (saturated brine)
- Divide a container with a porous membrane (unglazed ceramic, canvas, or multiple layers of cloth)
- Place a graphite anode in the brine side
- Place a steel cathode in the other side (filled with plain water or dilute NaOH)
- Apply 4-6V DC
- Chlorine evolves at the anode (greenish-yellow gas with sharp smell — toxic, work outdoors)
- Sodium hydroxide accumulates in the cathode compartment
- Hydrogen evolves at the cathode
Chlorine Gas Is Toxic
Even small concentrations of chlorine gas cause respiratory distress. Always perform brine electrolysis outdoors or under a fume hood. If you smell a sharp, swimming-pool-like odor, move upwind immediately.
Common Mistakes
- Using AC instead of DC: AC reverses the reaction every cycle, producing nothing useful and wasting electricity. Always verify DC polarity before starting.
- Choosing the wrong electrolyte: Using salt when you want pure hydrogen and oxygen produces chlorine gas instead. Use sodium hydroxide or baking soda for water splitting.
- Electrodes too far apart: Resistance increases with distance, requiring more voltage and wasting energy as heat. Keep electrodes 1-5cm apart.
- Ignoring gas collection: Hydrogen and oxygen mix explosively. If you are not collecting gases separately, ensure vigorous ventilation to keep concentrations below explosive limits.
- Undersized power supply: Electrolysis requires sustained current, not just voltage. A 6V source that can only deliver 100mA will produce negligible output. Match your power supply to your production goals.
Summary
Electrolysis Fundamentals -- At a Glance
- Electrolysis uses DC electricity to drive chemical reactions: splitting water, extracting metals, producing chlorine and lye
- Each reaction has a minimum voltage; practical cells need 50-100% more to overcome losses
- Faraday’s Laws let you calculate exact production rates: mass deposited is proportional to current times time
- Electrode choice determines success: use inert materials (graphite, stainless steel) for anodes in acid/salt solutions
- Water electrolysis produces hydrogen (cathode) and oxygen (anode) — both useful but dangerously flammable/combustion-supporting
- Salt electrolysis produces chlorine, sodium hydroxide, and hydrogen — three essential industrial chemicals from common salt