Calculate electrolytic cell conditions including charge passed, moles produced, and mass deposited using Faraday's laws of electrolysis
Total Charge Passed
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Moles of Product
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Mass Deposited
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Gas Volume (if applicable)
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Electrolysis is a chemical process that uses electric current to drive non-spontaneous chemical reactions. It involves passing current through an electrolytic cell to decompose compounds or deposit materials at electrodes.
Faraday's law states that the amount of substance produced at an electrode is directly proportional to the quantity of electric charge passed through the electrolyte. The relationship is given by: moles = Q / (n × F), where Q is charge, n is electrons per mole, and F is Faraday's constant (96,485 C/mol).
The value of n depends on the reaction. Common examples: Ag⁺ → Ag requires 1 electron, Cu²⁺ → Cu requires 2 electrons, Al³⁺ → Al requires 3 electrons, and H₂O → H₂ + ½O₂ requires 2 electrons per H₂ molecule.
Current efficiency is the percentage of total current that actually contributes to the desired reaction. It's often less than 100% due to side reactions, resistance losses, or other factors. Higher efficiency means more product per unit of electricity consumed.
Mass is calculated using: mass = moles × molar mass. First determine moles using Faraday's law (moles = Q / (n × F)), then multiply by the molar mass of your product. Remember to account for current efficiency if it's less than 100%.
Faraday's constant (F = 96,485 C/mol) represents the charge of one mole of electrons. It's fundamental to all electrochemical calculations and links the electrical quantity (charge) to the chemical quantity (moles).
For gaseous products, multiply the moles produced by the molar volume of gas. At STP (0°C, 1 atm), the molar volume is 22.414 L/mol. For other conditions, use the ideal gas law or appropriate correction factors.