Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes, and with the conversion of chemical energy into electrical energy and vice versa. It forms the scientific basis for batteries, electroplating, corrosion prevention, and many industrial processes. The chapter begins by distinguishing between electrolytic cells, where electrical energy drives a non-spontaneous reaction, and galvanic cells, where a spontaneous chemical reaction produces electrical energy.
Central to the chapter is the concept of electrode potential, the tendency of an electrode to gain or lose electrons. The standard hydrogen electrode serves as the reference for measuring electrode potentials, and the Nernst equation allows the calculation of electrode potentials under non-standard conditions. These concepts enable the calculation of cell potential, free energy change, and the equilibrium constant of cell reactions.
The later sections of the chapter cover electrolysis, the quantitative laws of Faraday, and the applications such as electroplating and extraction of metals. Finally, the chapter examines the phenomenon of corrosion, particularly rusting of iron, and the methods used to prevent it. Conductance of electrolytic solutions, including molar conductivity and its variation with concentration, and Kohlrausch's law of independent migration of ions, complete the theoretical framework.
A galvanic cell converts the chemical energy of a spontaneous redox reaction into electrical energy. In a Daniel cell, zinc is the anode (oxidation) and copper is the cathode (reduction): $$\text{Anode: } Zn \rightarrow Zn^{2+} + 2e^-$$ $$\text{Cathode: } Cu^{2+} + 2e^- \rightarrow Cu$$ $$\text{Overall: } Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$$
The salt bridge completes the electrical circuit and maintains electrical neutrality by allowing the movement of ions. It contains a concentrated solution of an inert electrolyte such as KCl or KNO3 in agar-agar gel.
The cell notation is written as: $$Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)$$
In an electrolytic cell, electrical energy is used to bring about a non-spontaneous chemical reaction. Examples include the electrolysis of molten NaCl and aqueous solutions.
The electrode potential is the potential difference developed between the electrode and its electrolyte. When the concentration of all species is 1 M and gases are at 1 bar pressure, the potential is called the standard electrode potential, E°.
The standard hydrogen electrode (SHE) is assigned a potential of zero and serves as the reference: $$Pt(s) | H_2(g, 1\ \text{bar}) | H^+(aq, 1\ \text{M}) \quad E^\circ = 0.00\ \text{V}$$
The standard cell potential is: $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$
Higher the oxidation potential (or lower the reduction potential), the stronger is the reducing agent. For a spontaneous reaction, the cell potential must be positive.
The Nernst equation relates the electrode potential to the concentration of the species involved. For a general reaction: $$aA + bB \rightarrow cC + dD$$ $$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0591}{n} \log \frac{[C]^c[D]^d}{[A]^a[B]^b} \quad \text{(at 298 K)}$$
For a metal-metal ion electrode: $$E = E^\circ - \frac{0.0591}{n} \log \frac{1}{[M^{n+}]}$$
For the Daniel cell: $$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0591}{2} \log \frac{[Zn^{2+}]}{[Cu^{2+}]}$$
At equilibrium, Ecell = 0, and the Nernst equation reduces to a relation between the standard cell potential and the equilibrium constant: $$\log K_c = \frac{n E^\circ_{\text{cell}}}{0.0591} \quad \text{(at 298 K)}$$
The electrical work done by a cell is related to the free energy change: $$\Delta G = -n F E_{\text{cell}}$$ $$\Delta G^\circ = -n F E^\circ_{\text{cell}} = -R T \ln K_c$$
where $n$ is the number of electrons transferred, $F$ is Faraday's constant (96487 C/mol), and $K_c$ is the equilibrium constant. A negative ΔG indicates a spontaneous reaction.
The conductance of an electrolytic solution depends on the number of ions and their mobility. Key quantities are:
$$\Lambda_m = \frac{\kappa}{c}$$ where $c$ is the concentration in mol/m3. Molar conductivity increases with dilution because the number of ions per unit volume decreases but the degree of dissociation increases, so the total conductivity per mole of electrolyte rises.
The limiting molar conductivity of an electrolyte is the sum of the contributions of its individual ions: $$\Lambda_m^\circ = \lambda_+^\circ + \lambda_-^\circ$$
This law is used to determine the limiting molar conductivity of weak electrolytes and the degree of dissociation: $$\alpha = \frac{\Lambda_m}{\Lambda_m^\circ}$$
The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed: $$m = Z \times Q = Z \times I \times t$$ where $Z$ is the electrochemical equivalent and $Q$ is the charge in coulombs.
When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their equivalent masses.
For electrolysis of water, hydrogen is liberated at the cathode and oxygen at the anode in the volume ratio 2:1. In molten NaCl, sodium is deposited at the cathode and chlorine is liberated at the anode.
Corrosion is the deterioration of metals by oxidation. Rusting of iron is an electrochemical process: $$\text{Anode: } Fe \rightarrow Fe^{2+} + 2e^-$$ $$\text{Cathode: } O_2 + 4H^+ + 4e^- \rightarrow 2H_2O$$
Corrosion can be prevented by galvanising (coating with zinc), tin plating, alloying (e.g., stainless steel), and cathodic protection using a sacrificial anode such as magnesium or zinc.
| Quantity | Equation | Conditions |
|---|---|---|
| Cell potential | Ecell = Ecathode - Eanode | Standard or actual |
| Nernst equation | E = E° - (0.0591/n) log Q | 298 K |
| Equilibrium constant | log Kc = nE°cell/0.0591 | 298 K |
| Gibbs free energy | ΔG = -nFEcell | Always |
| Faraday's first law | m = ZIt | Electrolysis |
| Kohlrausch's law | Λm° = λ+° + λ-° | Infinite dilution |
| Type | Anode | Cathode | Electrolyte | Rechargeable | Example |
|---|---|---|---|---|---|
| Primary | Zn | Carbon + MnO2 | NH4Cl, ZnCl2 paste | No | Dry cell |
| Secondary | Pb | PbO2 | H2SO4 | Yes | Lead storage battery |
| Fuel cell | H2 | O2 | KOH | Continuous feed | H2-O2 fuel cell |
Electrochemistry unifies thermodynamics and electricity through the concept of electrode potentials and the Nernst equation. The chapter explains how spontaneous redox reactions can generate useful electrical energy in batteries and how electrical energy can drive non-spontaneous reactions in electrolysis. Faraday's laws provide the quantitative basis for industrial electroplating and metal extraction, while conductance studies and Kohlrausch's law help determine dissociation constants and molar conductivities. The study of corrosion and its prevention has enormous practical importance in engineering and construction. Mastery of the relationships between E°, ΔG, and Kc is particularly valuable because these concepts reappear in thermodynamics and equilibrium and form the backbone of many numerical problems in board and competitive examinations.