Electrochemistry is the branch of physical chemistry that studies the interconversion of electrical and chemical energy. Electrochemical cells harness redox reactions to generate electricity (galvanic cells) or use electricity to drive non-spontaneous reactions (electrolytic cells).
Redox Reactions in Electrochemistry
Oxidation is the loss of electrons while reduction is the gain of electrons; the oxidizing agent is reduced and the reducing agent is oxidized. Redox reactions can be separated into two half-reactions: oxidation at the anode and reduction at the cathode. The overall cell reaction is the sum of the balanced half-reactions with electrons canceled.
Galvanic (Voltaic) Cells
A galvanic cell converts chemical energy into electrical energy through spontaneous redox reactions, with the Daniel cell (Zn/Cu) as the classic example. The Zn electrode (anode, -) is oxidized: Zn(s) → Zn2+(aq) + 2e-, while the Cu electrode (cathode, +) is reduced: Cu2+(aq) + 2e- → Cu(s). A salt bridge containing KCl or KNO3 in agar maintains electrical neutrality by allowing ion migration between half-cells. Cell notation is written as Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s), where | represents a phase boundary and || represents the salt bridge.
Standard Electrode Potentials
The standard reduction potential (E°) is measured relative to the standard hydrogen electrode (SHE): 2H+(aq) + 2e- → H2(g), E° = 0.00 V. More positive E° values indicate stronger oxidizing agents, while more negative values indicate stronger reducing agents. The standard cell potential is E°cell = E°cathode - E°anode (or E°reduction + E°oxidation), and a positive E°cell indicates a spontaneous reaction.
The Nernst Equation
The cell potential varies with concentration and temperature according to the Nernst equation: E = E° - (RT/nF) ln Q, where n is the number of electrons transferred, F is Faraday’s constant (96,485 C/mol), and Q is the reaction quotient. At 25°C, the equation simplifies to E = E° - (0.0592/n) log Q. At equilibrium, E = 0 and Q = K, giving the relationship log K = nE°/0.0592 at 25°C.
Electrolytic Cells
An electrolytic cell uses an external power source to drive a non-spontaneous redox reaction, where the anode is positive and the cathode is negative — opposite of galvanic cells. Faraday’s Laws of Electrolysis state that the mass of substance produced at an electrode is proportional to the quantity of charge passed: m = (Q/M)/(nF), where Q = I × t and M is molar mass. Applications include electroplating (chrome plating, gold plating), metal refining (copper purification), and electrolysis of water (H2 and O2 production).
Batteries
Primary (non-rechargeable) batteries include the Leclanché dry cell using Zn and MnO2, and the alkaline battery using Zn and MnO2 in alkaline electrolyte. Secondary (rechargeable) batteries include lead-acid (Pb and PbO2 in H2SO4), lithium-ion (LiCoO2 cathode, graphite anode, Li+ electrolyte), and Ni-Cd (NiOOH and Cd in KOH). Fuel cells convert chemical energy directly to electricity — the H2/O2 fuel cell produces water and electrical energy with high efficiency.
Corrosion
Corrosion is the spontaneous oxidation of metals, most notably iron rusting: Fe(s) → Fe2+(aq) + 2e-. Rust formation requires oxygen and water, with the cathodic reaction being O2 + 2H2O + 4e- → 4OH-. Corrosion prevention methods include painting, galvanization (Zn coating), cathodic protection using sacrificial anodes, and alloying to produce stainless steel.
