Faraday's First Law states that the mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electric charge passed through the electrolyte. Mathematically: m = (M × I × t) / (n × F), where m is mass (g), M is molar mass (g/mol), I is current (A), t is time (s), n is the number of electrons transferred, and F is the Faraday constant (96,485 C/mol).

Rearranging Faraday's Law: t = (m × n × F) / (M × I × η). Enter the desired mass of metal, the applied current, the substance's molar mass and electron count, plus the current efficiency (η). For example, to deposit 10g of copper (n=2, M=63.55) at 2A with 95% efficiency, the required time is approximately 4.4 hours.

Current efficiency (η) is the ratio of the actual amount of substance deposited to the theoretical amount predicted by Faraday's Law, expressed as a percentage. In real-world electroplating, some current is lost to side reactions (e.g., hydrogen evolution), so efficiency is typically 80–98%. A lower efficiency means you need more time (or higher current) to deposit the same mass.

The Faraday constant (F) represents the electric charge carried by one mole of electrons. It equals approximately 96,485 coulombs per mole (C/mol). This value is derived from Avogadro's number (6.022 × 10²³) multiplied by the elementary charge (1.602 × 10⁻¹⁹ C). It is the fundamental bridge between chemical amounts and electrical quantities in electrochemistry.

The number of electrons (n) equals the change in oxidation state of the metal ion:
• n=1: Ag⁺, Na⁺, K⁺
• n=2: Cu²⁺, Zn²⁺, Ni²⁺, Fe²⁺, Mg²⁺
• n=3: Au³⁺, Cr³⁺, Al³⁺, Fe³⁺
For gases like H₂ (from 2H⁺), n=2 per H₂ molecule; for O₂ (from 2H₂O), n=4 per O₂ molecule.

Yes! The same Faraday's Law principles apply to anodizing (e.g., aluminum anodization). In anodizing, the oxide layer grows as metal is converted to oxide at the anode. You need to know the effective molar mass of the oxide formed and the number of electrons involved. For aluminum anodizing (Al → Al₂O₃), the effective n=6 per Al₂O₃ unit, with M=101.96 g/mol.

Typical current densities vary by metal:
• Copper: 2–5 A/dm²
• Nickel: 2–6 A/dm²
• Chromium: 10–30 A/dm²
• Zinc: 1–5 A/dm²
• Gold: 0.1–1 A/dm²
• Silver: 0.5–3 A/dm²
Multiply by surface area to estimate the required current for your workpiece.

Several factors can cause longer actual plating times:
• Current efficiency < 100% — side reactions consume some current
• Uneven current distribution across the workpiece surface
• Electrolyte depletion or temperature variations
• Contact resistance at connections
Always use a current efficiency factor (typically 80–95%) and allow some buffer time for best results.