Electrochemistry

Electrochemistry studies the relationship between electrical energy and chemical change. The chapter is built around two types of electrochemical cells. Galvanic (or Voltaic) cells generate electricity from spontaneous chemical reactions, like in a battery. Electrolytic cells use an external electrical current to drive non-spontaneous reactions, a process called electrolysis, which is used in metal plating and refining.

Students will learn to calculate cell potentials using standard electrode potentials and the Nernst equation, which adjusts for non-standard concentrations and temperatures. The chapter also covers the conductivity of electrolytic solutions, introducing concepts like molar conductivity and Kohlrausch's law to understand ion behaviour. Finally, you will apply Faraday's laws of electrolysis to calculate the amount of substance produced or consumed during an electrolytic process. Practical applications like batteries, fuel cells, and corrosion are also discussed.

The most challenging parts are often the application of the Nernst equation to different cell configurations and solving quantitative problems related to Faraday's laws. A solid understanding of the 'Redox Reactions' chapter from earlier studies is essential, as balancing half-reactions is a constant requirement. Concepts from 'Thermodynamics', particularly Gibbs free energy, are also directly linked to cell potential.

Formulas for Electrochemistry

Relation between Gibbs free energy and cell potential: $\Delta_r G = -nFE_{cell}$ n is moles of electrons transferred, F is Faraday's constant (≈ 96500 C/mol).
Standard cell potential from standard electrode potentials: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ Both potentials are standard reduction potentials.
Nernst equation for cell potential: $E_{cell} = E^\circ_{cell} - \frac{RT}{nF} \ln Q$ Q is the reaction quotient, T is temperature in Kelvin.
Nernst equation at 298 K (25°C): $E_{cell} = E^\circ_{cell} - \frac{0.0591}{n} \log_{10} Q$ A practical form for calculations at room temperature.
Relation between standard cell potential and equilibrium constant: $E^\circ_{cell} = \frac{0.0591}{n} \log_{10} K_c$ Holds at equilibrium (E_cell = 0) and T = 298 K.
Conductivity (kappa): $\kappa = \frac{1}{R} \frac{l}{A}$ R is resistance, l/A is the cell constant G*.
Molar conductivity: $\Lambda_m = \frac{\kappa \times 1000}{c}$ κ is in S cm⁻¹, c is molar concentration in mol L⁻¹.
Debye-Hückel-Onsager equation for strong electrolytes: $\Lambda_m = \Lambda_m^\circ - A\sqrt{c}$ Λ°m is limiting molar conductivity, A is a constant.
Kohlrausch's law of independent migration of ions: $\Lambda_m^\circ = \nu_+ \lambda_+^\circ + \nu_- \lambda_-^\circ$ ν is the number of ions per formula unit, λ° is the limiting molar conductivity of the ion.
Faraday's first law of electrolysis: $w = ZIt = \frac{E}{F} It$ w is mass deposited, I is current, t is time, E is equivalent weight.
Charge passed during electrolysis: $Q = It$ Q is total charge in Coulombs, I is current in Amperes, t is time in seconds.