Reduction potentials, measured relative to the SHE, indicate the tendency of a species to be reduced. A more positive reduction potential signifies a stronger oxidizing agent, while a more negative value indicates a stronger reducing agent. For instance, magnesium, with a reduction potential of -2.37 V, is a potent reducing agent, often acting as the anode when paired with metals like copper or iron.
Gibbs free energy (ΔG) is linked to cell potential by the equation ΔG = -nFE, where n is the number of moles of electrons transferred, F is Faraday's constant (approximately 96,485 C/mol), and E is the cell potential. For spontaneous reactions, ΔG is negative, aligning with a positive E°cell in galvanic cells. This relationship underscores the thermodynamic feasibility of electrochemical processes.
Laws Governing Electrolysis
Faraday’s laws of electrolysis provide a quantitative framework for electrochemical reactions. 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, mass = (Q × M) / (n × F), where Q is the charge, M is the molar mass, n is the number of electrons involved, and F is Faraday's constant. This law is fundamental for applications like electroplating and metal refining.
Practical Examples and Applications
The electrolysis of molten NaCl illustrates key principles. At the cathode, Na⁺ + e⁻ → Na, producing sodium metal, while at the anode, 2Cl⁻ → Cl₂ + 2e⁻, generating chlorine gas. This process highlights the role of electrode reactions in industrial chemistry, such as in the production of sodium for various applications.
Another example is determining the cathode in a galvanic cell with magnesium. Given magnesium’s reduction potential of -2.37 V, any metal with a less negative (or positive) reduction potential, such as copper (+0.34 V) or iron (-0.44 V), will act as the cathode, as it is more readily reduced. This comparison is based on standard reduction potential tables, which are essential for predicting cell behavior.

Comparative Analysis of Cell Types

The distinction between galvanic and electrolytic cells is crucial. In galvanic cells, the anode is negative, and electrons flow from anode to cathode through the external circuit, driven by the spontaneous nature of the reaction. In electrolytic cells, the external power source reverses this flow, with the anode positive and the cathode negative, forcing the reaction to proceed. This reversal is evident in the polarity and the need for external energy, contrasting with the self-sustaining nature of galvanic cells.

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