Chapter 2 — Electrochemistry
Redox reactions move electrons. Make them move through a wire and you have a battery; push them backwards with a power supply and you have electrolysis. This chapter measures, predicts and puts to work both directions of that flow.
Your phone, a pacemaker, an electric scooter and the Mars rovers all run on the same trick: a chemical reaction that would happen anyway is split into two halves, so that its electrons have to travel through a circuit to get from one half to the other. Run the trick backwards and electricity drives chemistry instead. That is how aluminium, chlorine, sodium hydroxide and pure copper are made in millions of tonnes, how a car battery is recharged, and how a thin layer of gold ends up on a connector. Even the rust on a gate is electrochemistry, a tiny galvanic cell that nobody built on purpose.
The chapter starts with the galvanic cell and the standard electrode potentials that tell you which way electrons will flow and with what voltage. The Nernst equation then adds concentration, and the link $\Delta_{r}G^{\circ}=-nFE^{\circ}_{cell}$ turns a voltmeter reading into a Gibbs energy and an equilibrium constant. Next we look at how ions carry current through solutions: conductivity, molar conductivity and Kohlrausch's law, which together reveal how far a weak acid is ionised. Finally, electrolysis and Faraday's laws, the batteries and fuel cells of everyday life, and the electrochemistry of corrosion and its prevention.
| Question | Key relation | Section |
|---|---|---|
| What voltage does a cell give at standard conditions? | $E^{\circ}_{cell}=E^{\circ}_{R}-E^{\circ}_{L}$ | 2.1 |
| How does the voltage depend on concentration? | $E_{cell}=E^{\circ}_{cell}-\tfrac{0.059}{n}\log Q$ | 2.2 |
| How is EMF linked to energy and equilibrium? | $\Delta_{r}G^{\circ}=-nFE^{\circ}_{cell}$, $\ \log K=\tfrac{nE^{\circ}_{cell}}{0.059}$ | 2.2 |
| How well does a solution conduct, per mole? | $\Lambda_{m}=\kappa/c$, $\ \Lambda_{m}^{\circ}=\nu_{+}\lambda_{+}^{\circ}+\nu_{-}\lambda_{-}^{\circ}$ | 2.3 |
| How much is made by electrolysis? | $m=\dfrac{M\,I\,t}{nF}$ | 2.4 |
- describe a galvanic cell (anode, cathode, salt bridge, electron and ion flow), write and read cell notation, and use the standard hydrogen electrode and a table of $E^{\circ}$ values to find $E^{\circ}_{cell}$ and predict redox reactions;
- use the Nernst equation for electrodes and cells, find pH or unknown concentrations from an EMF, and calculate $\Delta_{r}G^{\circ}$ and $K$ from $E^{\circ}_{cell}$;
- measure and calculate resistance, conductivity, cell constant and molar conductivity, explain how they change with concentration for strong and weak electrolytes, and use Kohlrausch's law to find $\Lambda_{m}^{\circ}$, $\alpha$ and $K_{a}$;
- apply Faraday's laws, predict the products of electrolysis of molten and aqueous electrolytes, and explain the working of dry, mercury, lead storage and nickel–cadmium cells and the hydrogen–oxygen fuel cell;
- explain rusting as an electrochemical process and choose a method of preventing corrosion.
Sections in this chapter
Galvanic Cells and Electrode Potentials
The Daniell cell, cell notation, the standard hydrogen electrode and what the electrochemical series predicts.
The Nernst Equation, Gibbs Energy and Equilibrium Constant
EMF at any concentration, pH from a cell, and the links between $E^{\circ}$, $\Delta_{r}G^{\circ}$ and $K$.
Conductance of Electrolytic Solutions
Conductivity, cell constant, molar conductivity, strong and weak electrolytes, and Kohlrausch's law.
Electrolysis, Batteries, Fuel Cells and Corrosion
Faraday's laws, products of electrolysis, primary and secondary batteries, fuel cells, rusting and its prevention.