Electrochemistry studies the production of electricity from spontaneous chemical reactions and the use of electrical energy to drive non-spontaneous reactions. An electrochemical cell has two metallic electrodes dipping in electrolyte solution(s), the electrolyte being the ionic conductor.
There are two types of cells. In a galvanic (voltaic) cell, the chemical energy of a spontaneous redox reaction is converted into electrical work; in an electrolytic cell, external electrical energy drives a non-spontaneous reaction. In the Daniell cell, Zn(s) + Cu2+ gives Zn2+ + Cu(s), with a potential of 1.1 V at unit concentrations; applying an external opposing potential greater than 1.1 V reverses it into an electrolytic cell. Oxidation occurs at the anode (negative) and reduction at the cathode (positive); by convention the anode is written on the left and the cathode on the right, with a single line for a metal-electrolyte boundary and a double line for the salt bridge.
The standard electrode potential of any electrode is defined relative to the standard hydrogen electrode (SHE), Pt|H2(1 bar)|H+(1 M), which is assigned zero at all temperatures. E(standard cell) = E(cathode) - E(anode). The electrochemical series (Table 2.1) ranks couples by E; F2/F- (+2.87 V) is highest (strongest oxidant) and Li+/Li (-3.05 V) lowest (Li is the strongest reductant).
The Nernst equation gives the concentration dependence: E(cell) = E(standard cell) - (RT/nF) ln Q, which at 298 K becomes E(cell) = E(standard cell) - (0.059/n) log Q. The cell is linked to thermodynamics by dG(standard) = -nF E(standard cell) and dG(standard) = -RT ln K, so E(standard cell) = (2.303RT/nF) log K.
Conductivity (kappa) is the inverse of resistivity; conductance G = 1/R = kappa times A/l. Cell constant = l/A = R times kappa. Molar conductivity Lambda m = kappa/c. On dilution, conductivity always decreases (fewer ions per unit volume) while molar conductivity increases; it reaches a maximum, the limiting molar conductivity (Lambda zero m), at infinite dilution. For strong electrolytes Lambda m = Lambda zero m - A times square root of c; for weak electrolytes Lambda m rises steeply on dilution and Lambda zero m is found via Kohlrausch's law of independent migration of ions, Lambda zero m = n+ times lambda zero+ + n- times lambda zero-.
Faraday's laws of electrolysis: the amount of substance reacting at an electrode is proportional to the charge passed (Q = It), and for the same charge the amounts of different substances liberated are proportional to their chemical equivalent weights. One Faraday = 96487 C per mol (about 96500 C). Products of electrolysis depend on the electrode (inert or reactive), the species present and their potentials, with overpotential sometimes changing the product (e.g. Cl2 over O2 in aqueous NaCl).
Batteries and fuel cells are useful galvanic cells. Primary cells (dry/Leclanche, mercury cell) cannot be recharged; secondary cells (lead storage battery, nickel-cadmium cell) can. Fuel cells (such as the H2-O2 cell, overall 2H2 + O2 gives 2H2O) convert combustion energy directly into electricity at about 70 percent efficiency and are pollution free. Corrosion (e.g. rusting of iron) is an electrochemical process in which a metal is oxidised at anodic spots while oxygen is reduced at cathodic spots; it is prevented by barriers, protective metal coatings, or sacrificial electrodes of more reactive metals.