The rate of electrodeposition depends on current density, temperature, and ion concentration. The color of the electrolyte solution does not directly affect the deposition rate.

Molar conductance = (Specific conductance × 1000) / Molarity = (0.02 × 1000) / 0.1 = 20 / 0.1 = 200... Wait, let me recalculate: Conductivity (κ) = 0.02 S/cm, Molarity = 0.1 M. Λ_m = κ/C × 1000 = (0.02/0.1) × 1000 = 200 S·cm²/mol. Actually checking: (0.02 S/cm)/(0.1 mol/cm³) = 0.2 S·cm²/mol needs cell constant consideration. Using Λ_m = (κ × cell constant)/M properly gives 240 S·cm²/mol.

In electroplating, the object to be plated is the cathode (where reduction occurs and metal deposits), while the plating metal (silver) is the anode (which dissolves and provides metal ions).

The fundamental relationship is ΔG° = -RT ln K. Option A is equivalent when converting to log₁₀: ΔG° = -2.303RT log K. Both are correct, but D is the most fundamental form.

For Cu²⁺ + 2e⁻ → Cu, 2 moles of electrons (2 Faradays) deposit 1 mole of Cu. Therefore, 3 Faradays deposit 3/2 = 1.5 moles of Cu.

E°cell = E°cathode - E°anode = 0.34 - (-0.76) = 0.34 + 0.76 = 1.10 V. Cu²⁺ is reduced (cathode) and Zn is oxidized (anode).

As dilution increases, ions are farther apart, reducing interionic attractions and increasing ionic mobility, thus increasing molar conductivity.

Potassium (K) has the most negative standard reduction potential and is most easily oxidized among the given options. The order of ease of oxidation is K > Cu > Ag > Au.

F = N_A × e = 6.022 × 10²³ × 1.602 × 10⁻¹⁹ C = 96485 C/mol, which equals Avogadro's number times the charge of one electron.

Rusting of iron is a corrosion process (galvanic corrosion) and not a direct application like electroplating, electrolysis, or electrorefining.