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⚗️ Electrochemistry

Electrolysis of Water
Virtual Experiment

Pass current through water, collect hydrogen at the cathode and oxygen at the anode, and verify the 2:1 H₂:O₂ volume ratio. Adjust voltage, electrolyte, and electrode material — all governed by Faraday's Laws.

2H₂O → 2H₂ + O₂ Faraday's Laws Virtual Lab FYUGP Chemistry
✓ Reading recorded.
🧪
Live Apparatus
Cathode (−) · Anode (+) · Gas collection tubes
⚡
Electrode Half-Reactions
Real-time reaction equations
Active Reactions
Cathode − 2H₂O + 2e⁻ → H₂↑ + 2OH⁻
Anode + 4OH⁻ → O₂↑ + 2H₂O + 4e⁻
Overall 2H₂O(l) → 2H₂(g) + O₂(g)
Voltage
0.0
Volts
Current
0.00
Amperes
Time
0:00
elapsed
H₂ Volume
0.000
mL
O₂ Volume
0.000
mL
Energy
0.00
Joules
💧→ H₂ (Cathode) 0%
Moles: 0.000000 mol
💧→ O₂ (Anode) 0%
Moles: 0.000000 mol
📊
Observation Table
Recorded readings with all parameters
Sl.ElectrolyteElectrodesV (V)I (A) t (s)H₂ (mL)O₂ (mL)RatioEnergy (J)Del
No readings yet. Run the experiment and click "Take Reading".
🎛️
Controls
Adjust all parameters
Mode
STEP 1 / 5
Set the voltage to at least 1.5 V (minimum for water electrolysis).
Voltage
Voltage (V) 6.0 V
Electrolyte
Electrolyte Concentration
Conc. (%) 5.0 %
Electrode Material
⚗ Setup 2 (shown in purple on graph)
Voltage 9.0 V
Not running. Main sim must be running for comparison.
Calculated Values
Charge (Q)
0.00 C
n(H₂) mol
0.000000
n(O₂) mol
0.000000
Efficiency
—
📈
Live Graph
Gas volume vs time
H₂ gas
O₂ gas
Theory, Notes & Precautions
📖 What is Electrolysis?

Electrolysis is the process of driving a non-spontaneous chemical reaction using electrical energy. An external DC voltage source forces current through an electrolyte (ionic solution or molten salt), causing chemical changes at the electrodes.

In water electrolysis, water is split into hydrogen and oxygen:

2H₂O(l) → 2H₂(g) + O₂(g)

Minimum theoretical voltage required: 1.23 V. In practice, ~1.5–2 V is needed due to overpotential.

⚗️ Cathode & Anode Reactions

Cathode (−) — Reduction:

2H₂O + 2e⁻ → H₂ + 2OH⁻

Anode (+) — Oxidation:

4OH⁻ → O₂ + 2H₂O + 4e⁻

The cathode produces hydrogen (colourless); the anode produces oxygen (colourless). Twice as many moles of H₂ are produced as O₂, giving a 2:1 volume ratio.

📐 Faraday's Laws

First Law: Mass deposited ∝ charge passed. Q = I × t

Charge Q = I × t  (Coulombs)

Second Law: Moles produced depend on stoichiometry.

n(H₂) = Q / (2F)
n(O₂) = Q / (4F)

Where F = 96485 C/mol (Faraday constant).

Volume at STP: V = n × 22400 mL/mol

🔬 Why Add Electrolyte?

Pure water is a poor conductor — it has very few ions. Adding an electrolyte greatly increases conductivity:

  • H₂SO₄ — very effective, widely used in Hoffmann voltameter
  • NaOH — strongly alkaline; efficient
  • Na₂SO₄ — neutral; clean products at both electrodes
  • KOH — highly conductive; used in industrial electrolysers

The electrolyte ions carry current through solution but are NOT consumed (for water electrolysis with inert electrodes).

⚠️ Precautions
  • Use inert electrodes (Pt or C) to avoid contamination
  • Minimum 1.5 V needed for electrolysis to begin
  • Add electrolyte — pure water won't work efficiently
  • Ensure gas tubes are completely filled before starting
  • Do not allow gas mixing — H₂ and O₂ together are explosive
  • Avoid reactive electrodes (they dissolve or react)
  • Collect gas at room temperature for accurate volume measurement
  • Keep current stable throughout the experiment
🧮 Key Calculations
Charge passed:
Q = I × t
Moles of H₂:
n(H₂) = Q / (2 × 96485)
Moles of O₂:
n(O₂) = Q / (4 × 96485)
Volume at STP:
V(mL) = n × 22400
Energy used:
E = V × I × t  (Joules)