Auto-fills molar mass and electron count
100% = ideal (no side reactions)
Enter molar mass.
Enter electrons transferred (1–8).
Constant current in amperes
Enter current in amperes.
Convert: 1 hr = 3600 s, 1 min = 60 s
Enter time in seconds.
Mass to deposit or dissolve
Enter target mass in grams.
Result
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Sources & Methodology
Faraday constant F = 96485.33212 C/mol (NIST CODATA 2018, exact). Molar masses from IUPAC 2021 atomic weights.
NIST CODATA 2018 — Faraday Constant
F = 96485.33212 C/mol (exact since 2019 SI redefinition). physics.nist.gov
Faraday, M. (1833) — Experimental Researches in Electricity
Original publication establishing both laws of electrolysis. Philosophical Transactions of the Royal Society, 123.
IUPAC 2021 Atomic Weights of the Elements
Standard atomic weights used for molar mass presets. iupac.org
Faraday’s First Law: m = (Q × M) / (n × F)
Charge: Q = I × t
Combined: m = (I × t × M) / (n × F)
F = 96485 C/mol | m in grams | I in amperes | t in seconds | M in g/mol
Charge: Q = I × t
Combined: m = (I × t × M) / (n × F)
F = 96485 C/mol | m in grams | I in amperes | t in seconds | M in g/mol
m = (I × t × M) / (n × F) | Q = I × t
Copper (n=2, M=63.546) at 2 A for 1 hour:
m = (2 × 3600 × 63.546) / (2 × 96485) = 457,531 / 192,970 = 2.372 g Cu deposited
m = (2 × 3600 × 63.546) / (2 × 96485) = 457,531 / 192,970 = 2.372 g Cu deposited
Faraday's Laws of Electrolysis Explained
Michael Faraday established the quantitative laws of electrolysis in 1833. Faraday’s first law states that the mass of substance deposited or dissolved at an electrode is directly proportional to the total electric charge passed through the solution. Faraday’s second law states that for the same charge, the masses of different substances deposited are proportional to their equivalent weights (molar mass divided by number of electrons).
The key equation combining both laws is: m = (I × t × M) / (n × F), where m is mass in grams, I is current in amperes, t is time in seconds, M is molar mass in g/mol, n is the number of electrons transferred per ion, and F is the Faraday constant (96485 C/mol).
Common Electrolysis Applications
| Application | Metal/Gas | n | Industrial Use |
|---|---|---|---|
| Electroplating | Cu, Ni, Cr, Au, Ag | 1–3 | Corrosion protection, decorative |
| Electrorefining | Cu, Al, Zn | 2–3 | Purification of metals |
| Water electrolysis | H₂ + O₂ | 2 / 4 | Hydrogen fuel production |
| Chlor-alkali process | Cl₂ + NaOH | 2 | Chemical industry |
| Aluminum smelting | Al (Hall-Héroult) | 3 | Primary aluminum production |
💡 Current Efficiency: Real electrolytic processes are never 100% efficient. Side reactions — hydrogen evolution at the cathode, oxygen reduction — consume current without depositing the target metal. Industrial copper plating achieves 95–99% efficiency. Chrome plating is notoriously low at 10–25% (most current goes to hydrogen evolution). Always account for current efficiency in real process design.
Frequently Asked Questions
Mass deposited is proportional to charge passed: m = (Q × M) / (n × F). Q = I × t. F = 96485 C/mol. More charge = more mass deposited.
For the same charge, masses of different substances deposited are proportional to their equivalent weights (M/n). If charge 96485 C is passed: 1 mol Ag (n=1) = 107.87 g; 0.5 mol Cu (n=2) = 31.77 g.
F = 96485 C/mol (exact, CODATA 2018). The charge of one mole of electrons. Equals elementary charge (1.602×10⁻¹⁹ C) times Avogadro's number (6.022×10²³ /mol).
m = (I × t × M) / (n × F). Example: Cu (n=2, M=63.546) at 2 A for 3600 s: m = (2×3600×63.546)/(2×96485) = 2.372 g.
Electrodeposition of a metal coating using electrolysis. Substrate = cathode, plating metal = anode, in electrolyte containing metal ions. Faraday's law governs deposition rate. Used for corrosion protection, decoration, and electronics.
One Faraday = 96485 C/mol. Depositing 1 mol Ag (n=1) requires 96485 C. Depositing 1 mol Cu (n=2) requires 2 × 96485 = 192970 C.
Current efficiency = (actual mass / theoretical mass) × 100%. Side reactions (H₂ evolution, O₂ reduction) consume current without depositing metal. Copper plating: 95–99%. Chrome plating: 10–25%. Always include efficiency in real process calculations.
Rearrange: t = (m × n × F) / (I × M). Example: 5 g Ag (n=1, M=107.87) at 3 A: t = (5×1×96485)/(3×107.87) = 482425/323.61 = 1491 s = 24.9 min.
Always use SI: current in amperes (A), time in seconds (s), mass in grams (g), molar mass in g/mol. Q = I(A) × t(s) = coulombs (C). Convert hours to seconds (×3600) before calculating.
Electrolysis uses external electrical energy to drive a non-spontaneous reaction. Galvanic cells produce electrical energy from a spontaneous reaction. Electrolysis deposits metal at cathode by forcing electrons in; galvanic cells release electrons from spontaneous oxidation.
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