Introduction

Objectives

After studying this chapter, you should be able to:

• describe what an electrochemical cell is, and tell a galvanic cell apart from an electrolytic cell;
• apply the Nernst equation to find the emf of a galvanic cell away from standard concentrations, and define the standard potential of a cell;
• relate a cell's standard potential to the Gibbs energy of its cell reaction and to the reaction's equilibrium constant;
• define resistivity, conductivity, and molar conductivity of an ionic solution;
• distinguish ionic (electrolytic) conduction from electronic (metallic) conduction;
• describe how the conductivity of an electrolytic solution is measured, and how molar conductivity is calculated from it;
• explain how conductivity and molar conductivity vary with concentration, and define limiting molar conductivity ($\Lambda_m^{o}$), the value at infinite dilution;
• state Kohlrausch's law of independent migration of ions and use it;
• work through the quantitative laws governing electrolysis;
• describe how common primary and secondary batteries, and fuel cells, are built and how they function; and
• explain corrosion as an electrochemical process.

Electrochemistry sits exactly where two familiar ideas meet: a chemical reaction can

release or absorb energy, and an electric current is simply a directed flow of charge.

This chapter is about what happens when those two facts are deliberately linked — when a

spontaneous chemical reaction is harnessed to generate electrical energy, or when

electrical energy is deliberately spent to force a reaction that would never happen on

its own. Both directions matter enormously in practice. The first is exactly how every

battery, mobile-phone cell, and fuel cell works. The second is how reactive metals such

as sodium and aluminium are extracted from their ores, how objects are electroplated, and

how industrial chemicals such as chlorine and caustic soda (sodium hydroxide) are

manufactured at scale. Even biology depends on it — the nerve impulses that carry sensory

signals through your body and back to the brain are, at their core, electrochemical

events.

The chapter develops this subject in three connected strands. It begins with

electrochemical cells — how pairing two half-reactions in separate compartments lets

you tap their combined electron transfer as a current, how to compare different cells

using standard electrode potentials, and how the Nernst equation extends this beyond the

tidy standard-concentration case to predict a cell's emf under any real condition,

tying cell potential back to familiar thermodynamic quantities like Gibbs energy and the

equilibrium constant. It then turns to conductance of electrolytic solutions — how

well an ionic solution carries current, how that ability is measured experimentally, and

why it changes in very different ways for strong and weak electrolytes as a solution is

diluted, a puzzle Kohlrausch's law helps resolve. Finally, the chapter turns to

practical electrochemistry — the quantitative rules of electrolysis, how real

batteries and fuel cells are actually built and what limits them, and how corrosion (the

slow, unwanted electrochemical destruction of a metal) can be understood, and slowed,

using the very same ideas.

A theme worth holding onto throughout: nearly every idea in this chapter is really a

statement about where electrons go and why. Whether the question is which electrode

loses electrons in a battery, why a dilute solution conducts less current than a

concentrated one, or why an iron nail rusts faster near a scratch, the answer traces back

to the same underlying tendency of electrons to move from where they are less stable to

where they are more stable — exactly the same logic you already used for oxidation and

reduction, now put to work generating and consuming real electrical energy.