Faraday’s Law - Quantitative Electrochem
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Core Concept
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Practice Tips
Always convert minutes or hours to seconds before multiplying by current.
Double-check the half-reaction — the number of electrons transferred (nₑ) is critical.
Use dimensional analysis to track units: coulombs → mol e⁻ → mol substance → grams.
Memorize Faraday’s constant (96,485 C/mol e⁻) or keep it on your formula sheet.
Think stoichiometrically — the logic is the same as any mass-to-mass problem, just with electricity.
Core Concept
Electric Charge (Q)
Definition: The total electric charge transferred during an electrochemical reaction.
Formula: Q = I ⋅ t where:
Q: Charge (Coulombs (C)).
I: Current (Amperes (A)).
t: Time (seconds (s)).
Faraday's Constant (F)
Represents the charge of one mole of electrons: F=96,485 C/mol
Moles of Electrons (n)
The amount of charge relates to the moles of electrons transferred: $n_{\text{e}^-} = \frac{Q}{F}$
Faraday’s Laws of Electrolysis
First Law:
The mass (m) of a substance deposited or liberated at an electrode is proportional to the charge passed through the electrolyte. m=Z⋅Q Where:
Z: Electrochemical equivalent (g/C).
Q=I⋅t.
Second Law:
For the same amount of charge, the mass of different substances produced is proportional to their molar masses (M) divided by the number of electrons (n) transferred. $m \propto \frac{M}{n}$
Quantitative Relationships
Mass of Product: m = $\frac{M \cdot Q}{n \cdot F}$
Where:
m: Mass of the product (g).
M: Molar mass (g/mol).
n: Number of electrons in the half-reaction.
Volume of Gas:
For gases produced at electrodes, use the ideal gas law: V = $\frac{nRT}{P}$ Where:
V: Volume (L).
R: Ideal gas constant (0.0821 L·atm/mol·K).
T: Temperature (K).
P: Pressure (atm).