Electrolytic Cells and Driven Reactions

If E°_cell for a reaction is negative, the reverse process is spontaneous, but the forward one can still be forced by supplying electrical energy. That is electrolysis. Charging a rechargeable battery is the same idea: the cell operates galvanically while discharging and electrolytically while charging.

Electrolysis drives a reaction against its actual thermodynamic direction under the operating conditions. A negative E for the intended reaction means ΔrG = −nFE > 0 and external work is required. Negative E° is the standard-state version of that test; E° alone does not settle nonstandard conditions. Charging a rechargeable cell reverses its discharge chemistry through supplied electrical work, provided the chemistry and materials permit a reversible cycle.

Oxidation is at the anode and reduction is at the cathode in both cell types. What changes is polarity. In a galvanic cell the anode is negative because it produces electrons; in an electrolytic cell the anode is positive because the power supply is pulling electrons out of it.

The power supply withdraws electrons from the electrolytic anode and delivers them to the cathode. Cations migrate toward the cathode and anions toward the anode in the ordinary electrolyte picture, while the actual electrode products depend on which species react. Polarity identifies the power connection; oxidation and reduction define the electrode names.

Thermodynamics gives the minimum voltage; reality demands more. Extra voltage, called overpotential, is needed to overcome the slow kinetics of forming gas bubbles at an electrode surface. Oxygen evolution has a particularly large overpotential, which is a major reason chlorine, not oxygen, is produced when concentrated brine is electrolyzed.

The reversible minimum opposing voltage follows the actual cell potential at the current activities. To pass useful current, additional applied voltage covers kinetic overpotentials at both electrodes and resistive iR losses. Gas evolution is often kinetically demanding, but overpotential is not simply the energy to make a visible bubble. Electrode material, current density, pH and concentration affect the competition, which is why thermodynamic tables alone may not select the observed brine anode product.

In aqueous electrolysis, water can compete with dissolved ions. For CuSO₄ with inert electrodes under ordinary conditions, Cu²⁺ is reduced to copper at the cathode, while water is oxidized to O₂ at the anode rather than sulfate being oxidized. The solution becomes more acidic as H⁺ is generated. Electrode definitions still hold: reduction at cathode, oxidation at anode.

With inert electrodes in aqueous CuSO₄, a common ordinary-condition prediction is Cu²⁺ + 2 e⁻ → Cu at the cathode and 2 H₂O → O₂ + 4 H⁺ + 4 e⁻ at the anode. Balance electrons to obtain 2 Cu²⁺ + 2 H₂O → 2 Cu + O₂ + 4 H⁺. With a copper anode instead, Cu → Cu²⁺ + 2 e⁻ can replenish the ions plated at the cathode, giving metal transfer rather than the same net acid production. Specify whether electrodes are inert or reactive. In molten NaCl there is no water competitor, so Na⁺ can instead produce sodium metal.