Chapter 2 - Electrochemistry

Electrochemistry studies the relationship between chemical reactions and electrical energy. It explains how a spontaneous redox reaction can generate electricity in a galvanic cell and how an external current can drive a non-spontaneous reaction in an electrolytic cell. The same principles help us understand batteries, electrolysis, fuel cells, corrosion and the electrical conductance of electrolyte solutions.

The chapter combines redox chemistry with thermodynamics and numerical problem-solving. Students learn to represent electrochemical cells, calculate electrode and cell potentials, apply the Nernst equation, connect cell potential with Gibbs energy and equilibrium, and analyse conductivity data.

Electrochemical cells

A galvanic or voltaic cell converts chemical energy into electrical energy. Oxidation takes place at the anode and reduction takes place at the cathode. Electrons flow through the external circuit from anode to cathode, while ions move through the electrolyte and salt bridge to maintain electrical neutrality.

Cell notation provides a compact representation of the cell. A phase boundary is shown by a single vertical line, a salt bridge by a double line, and the anode is conventionally written on the left. Correctly identifying the two half-reactions is the safest way to determine electrodes and electron flow.

An electrolytic cell uses external electrical energy to carry out a non-spontaneous reaction. The definitions of anode and cathode do not change—oxidation still occurs at the anode and reduction at the cathode—but the signs of the electrodes differ from those in a galvanic cell.

Electrode potential and the standard hydrogen electrode

An electrode develops a potential when it is in contact with a solution containing its ions. Individual electrode potential cannot be measured directly; it is measured relative to a reference electrode. The standard hydrogen electrode is assigned a standard potential of zero volts.

Standard reduction potentials help compare the tendency of species to undergo reduction. A more positive reduction potential indicates a stronger tendency to be reduced under standard conditions. The standard cell potential is obtained from the cathode and anode reduction potentials:

E°cell = E°cathode − E°anode

A positive standard cell potential indicates that the cell reaction is spontaneous under standard conditions.

The Nernst equation

Cell potential changes when concentration, pressure or temperature differs from standard conditions. The Nernst equation relates electrode or cell potential to the reaction quotient. At 298 K, it is frequently used in logarithmic form for numerical problems.

Students must write the balanced cell reaction before constructing the reaction quotient. Pure solids and pure liquids are omitted, and the exponent of each species follows its stoichiometric coefficient. The number of electrons transferred, n, must come from the balanced redox reaction.

Cell potential, Gibbs energy and equilibrium

Electrical work and thermodynamic feasibility are connected through:

ΔG = −nFE

Under standard conditions, ΔG° = −nFE°. At equilibrium, the cell potential becomes zero and the reaction quotient equals the equilibrium constant. This leads to a relation between E°cell and K. These equations explain why a positive cell potential corresponds to a negative Gibbs energy change and a spontaneous cell reaction.

Conductance of electrolyte solutions

Metallic conduction occurs through electrons, whereas electrolytic conduction occurs through the movement of ions. Resistance depends on the dimensions of the conductor; conductivity is an intrinsic property of the material at a specified temperature.

Molar conductivity represents the conductance associated with the ions produced by one mole of electrolyte in solution. On dilution, conductivity generally decreases because the number of ions per unit volume falls, while molar conductivity increases because ions move more freely and ionisation may increase.

Strong and weak electrolytes show different dilution behaviour. Kohlrausch's law of independent migration of ions is used to calculate limiting molar conductivity and can help determine the dissociation of weak electrolytes or the solubility of sparingly soluble salts.

Electrolysis and Faraday's laws

Electrolysis uses electrical energy to produce chemical change. Faraday's laws relate the amount of substance deposited or liberated at an electrode to the quantity of electricity passed. Numerical questions commonly involve current, time, moles of electrons and mass deposited.

Before solving, write the electrode half-reaction. This reveals the number of electrons needed per mole of product and prevents the common mistake of using ionic charge without considering the balanced reaction.

Batteries, fuel cells and corrosion

Primary cells are designed for single use, while secondary cells can be recharged by reversing their cell reaction. The chapter discusses familiar systems such as dry cells, mercury cells, lead storage batteries and nickel-cadmium cells. Fuel cells continuously convert the chemical energy of supplied reactants into electrical energy; the hydrogen-oxygen fuel cell is an important example.

Corrosion is an electrochemical process in which a metal is oxidised by its environment. Rusting of iron involves anodic and cathodic regions on the metal surface. Painting, coating, alloying and sacrificial protection reduce corrosion by interrupting one or more parts of the electrochemical process.

What can students study on Teachoo?

• galvanic and electrolytic cells;

• cell notation, salt bridges and electron flow;

• standard reduction potentials and cell emf;

• Nernst-equation numericals;

• relations among cell potential, Gibbs energy and equilibrium constant;

• resistance, conductance, conductivity and molar conductivity;

• Kohlrausch's law;

• electrolysis and Faraday-law calculations;

• batteries and fuel cells; and

• the mechanism and prevention of corrosion.

Common mistakes to avoid

• Oxidation is always at the anode and reduction is always at the cathode.

• Do not reverse the subtraction in Ecell = Ecathode − Eanode.

• Balance the redox reaction before deciding the value of n.

• Do not include pure solids and liquids in the reaction quotient.

• Distinguish conductivity from molar conductivity and use the correct units.

• Remember that dilution can decrease conductivity while increasing molar conductivity.

Best way to study Electrochemistry

First practise identifying oxidation and reduction half-cells. Next learn cell notation and standard potential calculations, followed by the Nernst equation and thermodynamic relations. Study conductance as a separate block with a unit table. Finish with electrolysis, batteries and corrosion. Drawing the cell and writing both half-reactions before every numerical makes the signs and electron count much easier to control.

Frequently asked questions

What is the difference between a galvanic and an electrolytic cell?

A galvanic cell produces electricity from a spontaneous reaction. An electrolytic cell consumes electricity to drive a non-spontaneous reaction.

What is the purpose of a salt bridge?

It completes the internal circuit, permits ion movement and prevents charge from building up in the half-cells while limiting direct mixing of the solutions.

In which direction do electrons flow?

Electrons flow through the external circuit from the anode, where oxidation occurs, to the cathode, where reduction occurs.

When is an electrochemical reaction spontaneous?

Under the stated conditions, it is spontaneous in the forward direction when Ecell is positive and ΔG is negative.

Why does molar conductivity increase on dilution?

Dilution reduces interionic interactions and, for weak electrolytes, increases ionisation, allowing the ions associated with one mole of electrolyte to conduct more effectively.

What does the Nernst equation calculate?

It calculates an electrode or cell potential under non-standard conditions using temperature, electron transfer and the reaction quotient.

How is rusting an electrochemical process?

Different areas of the iron surface act as anodic and cathodic regions. Iron is oxidised at anodic sites while oxygen is reduced at cathodic sites, ultimately producing hydrated iron oxide.

Use Teachoo's Electrochemistry explanations to connect cell diagrams, equations and units instead of treating the chapter as a collection of unrelated formulas.