Introduction
Electrochemistry studies the interconversion of chemical and electrical energy — reactions that either produce electricity spontaneously (batteries) or require electricity to proceed (electrolysis). Both directions rest on the same core idea: electron transfer between chemical species.
1. Oxidation, reduction, electrodes, and electrochemical cells
Oxidation is loss of electrons (increase in oxidation state); reduction is gain of electrons (decrease in oxidation state) — the two always occur together (a redox reaction). In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode, physically separated so electrons must flow through an external circuit, generating usable electric current. A galvanic (voltaic) cell generates electricity from a spontaneous redox reaction; an electrolytic cell consumes electricity to drive a non-spontaneous one.
2. Cell potential and electrode potential
Each electrode has a characteristic tendency to gain or lose electrons, quantified as its standard electrode potential (measured relative to the hydrogen electrode, arbitrarily set at 0 V). The overall cell potential, E°_cell = E°_cathode − E°_anode, predicts both the direction of spontaneous electron flow and, via ΔG° = −nFE°, the reaction’s spontaneity and equilibrium constant — directly linking electrochemistry to the thermodynamics of Chapter 3.
3. The Nernst equation and concentration effects
Electrode and cell potentials depend on concentration, not just standard conditions. The Nernst equation, E = E° − (RT/nF)lnQ, corrects the standard potential for actual reactant/product concentrations (Q, the reaction quotient). This explains why a battery’s voltage drops as it discharges (reactant concentrations fall, product concentrations rise, shifting Q) and underlies concentration cells and pH-sensitive electrodes (like the glass electrode in a pH meter).
4. Electrolysis and Faraday’s laws
Electrolysis forces a non-spontaneous redox reaction using an external power source, used industrially to extract reactive metals (aluminium, sodium) and to electroplate surfaces. Faraday’s laws of electrolysis quantify the relationship: the mass of substance deposited or liberated at an electrode is proportional to the total charge passed (Q = It) and to the substance’s molar mass divided by the number of electrons transferred per ion — allowing precise calculation of yields from current and time.
5. Conductance, ionic movement, and applications
Conductance in electrolyte solutions arises from the movement of ions rather than free electrons (as in metals); molar conductivity typically increases with dilution because ion–ion interactions decrease. Key applications built on these principles include batteries (portable galvanic cells, from simple dry cells to rechargeable lithium-ion cells), corrosion (an unwanted spontaneous electrochemical process, typically the oxidation of a metal such as iron in the presence of moisture and oxygen, which can be slowed by sacrificial anodes or protective coatings), and electroplating (a deliberate electrolytic application to deposit a thin protective or decorative metal layer).
Key Takeaways
- Oxidation (electron loss, at the anode) and reduction (electron gain, at the cathode) always occur together; galvanic cells generate electricity spontaneously, electrolytic cells consume it to force non-spontaneous reactions.
- Standard electrode and cell potentials predict reaction direction and spontaneity (via ΔG° = −nFE°); the Nernst equation extends this to real, non-standard concentrations.
- Faraday’s laws give a precise quantitative link between charge passed during electrolysis and the mass of substance produced or consumed.
- Ionic conductance underlies both useful applications (batteries, electroplating) and problems (corrosion), all governed by the same electron-transfer principles.