Electrochemistry - Lecture 6: Half-Cell Reactions and Constructing Full Electrochemical Cells

A modular guide to redox partitioning, electrode potentials, and cell assembly

Electrochemical systems are built upon half-reactions, the individual oxidation and reduction processes that occur at separate electrodes. When these half-cells are joined by a conductive path and an electrolyte bridge, they form a complete electrochemical cell capable of converting chemical energy into electrical work, or vice versa.

This lecture explains how to identify, write, and balance half-cell reactions, how to assemble galvanic and electrolytic cells, and how potentials combine to produce measurable voltages. You’ll also learn how standard electrodes are referenced, how salt bridges maintain charge neutrality, and how to calculate cell electromotive force (EMF) from tabulated data.

By the end of this session, you should understand how individual redox processes unite to form a functioning electrochemical device, from laboratory cells to commercial batteries.

1. The Concept of Half-Cells

A half-cell is a compartment in which a single redox half-reaction occurs.
Each half-cell contains:

An electrode (solid conductor or gas electrode)
An electrolyte containing ions involved in the redox process

Electrons move through the external circuit, while ions move through the internal electrolyte to maintain electrical neutrality.

Half-cells may be:

Oxidation (anodic) half-cells
Reduction (cathodic) half-cells

Together, they form the complete redox equation:

Reference: LibreTexts – Half-Cell Reactions

2. Writing Half-Reactions

Every redox process can be split into two balanced half-equations.
For example, in the copper–zinc cell:

Oxidation (anode):

Reduction (cathode):

Adding the two gives the overall cell reaction:

When balancing redox equations:

Balance all atoms except hydrogen and oxygen.
Balance oxygen using H₂O.
Balance hydrogen using H⁺ (or OH⁻ in basic media).
Balance charge using electrons.
Combine the two half-equations, ensuring equal electrons are transferred.

Reference: Khan Academy – Balancing Redox Equations

3. Anatomy of a Full Electrochemical Cell

A full cell is formed by joining two half-cells through:

A salt bridge or porous barrier
An external circuit (for electron flow)

Components:

Anode: site of oxidation (electron source)
Cathode: site of reduction (electron sink)
Electrolytes: supply ions and complete the circuit
Salt bridge: maintains charge balance by ion migration

Example – Daniell Cell:

Electrons flow: Zn → Cu
Conventional current: Cu → Zn (through voltmeter)
The cell produces E°₍cell₎ = +1.10 V.

Reference: Royal Society of Chemistry – Daniell Cell Experiment

4. The Role of the Salt Bridge

The salt bridge (or porous barrier) permits ionic conduction while preventing bulk mixing of solutions. It is usually filled with an inert electrolyte such as KNO₃ or KCl.

Functions:

Completes the internal circuit by allowing ion flow
Maintains electrical neutrality in both half-cells
Prevents solution depletion or charge build-up

At the molecular level:

Negative ions migrate toward the anode (where positive ions accumulate)
Positive ions migrate toward the cathode (where cations are consumed)

If the bridge fails, the charge imbalance stops the reaction even though a potential difference remains.

Reference: Chemguide – Salt Bridge and Cell Function

5. Representing Electrochemical Cells

Electrochemical cells are written using cell notation (introduced more fully in Lecture 8), showing phase boundaries and interfaces:

Conventions:

Single vertical line “|” → phase boundary
Double vertical line “||” → salt bridge
Left side = anode; right side = cathode
Electrons flow from left → right

The notation allows rapid interpretation of cell structure and direction of reaction.

6. Measuring Half-Cell Potentials

Individual half-cell potentials cannot be measured directly because there is no absolute potential scale. Potentials are measured relative to a reference electrode, the Standard Hydrogen Electrode (SHE), defined as 0 V under standard conditions:

When connected to another half-cell, the measured potential difference corresponds to the electrode potential of that half-cell relative to the SHE.

Experimental setup:

Platinum electrode in 1 M H⁺
H₂ gas bubbled at 1 atm
Temperature = 298 K (25 °C)

Reference: IUPAC Gold Book – Standard Hydrogen Electrode

7. Standard Conditions and the Sign Convention

Standard electrode potentials (E°) are measured under:

1 atm gas pressure
1 mol dm⁻³ solute concentration
298 K temperature

By convention:

E°₍cell₎ = E°₍cathode₎ − E°₍anode₎
A positive E°₍cell₎ indicates a spontaneous reaction (galvanic cell).
A negative E°₍cell₎ indicates a non-spontaneous process requiring energy input (electrolytic cell).

Example:

Reference: LibreTexts – Standard Electrode Potentials

8. Constructing a Full Cell: Step-by-Step

To construct a simple galvanic cell in the laboratory:

Prepare two beakers – one containing ZnSO₄ (1 M) and one containing CuSO₄ (1 M).
Insert metal electrodes – zinc in the zinc solution, copper in the copper solution.
Connect the electrodes via an external wire to a voltmeter.
Add a salt bridge of KNO₃ in agar between the two solutions.
Record the potential difference (expected ≈ 1.10 V).
Observe the reactions:
- Zinc electrode dissolves (oxidation).
- Copper electrode gains mass (reduction).

Anode (oxidation):

Cathode (reduction):

Reference: RSC Practical – Simple Galvanic Cells

9. Half-Cells in Electrolysis

While galvanic cells generate electricity from spontaneous reactions, electrolytic cells use electricity to drive non-spontaneous ones.

Example: Electrolysis of molten sodium chloride.

Anode (oxidation):

Cathode (reduction):

Overall:

Here, the external power supply enforces electron flow opposite to spontaneity, converting electrical energy into chemical energy.

Reference: BBC Bitesize – Electrolysis Explained

10. Quantitative Aspects of Half-Cell Reactions

Electrode reactions obey Faraday’s laws (see Lecture 5).
The charge transferred relates directly to the number of moles of electrons:

Where N = moles of substance reacted, n = electrons per molecule, and F = 96,485 C/mol.

The relationship between the current and the rate of reaction:

This allows conversion between electrical current and chemical change, crucial for analytical electrochemistry and industrial scaling.

11. Factors Affecting Half-Cell Potentials

Electrode potential depends on:

Ion concentration (activity)
Temperature
Pressure (for gas electrodes)
Nature and surface condition of the electrode
Presence of complexing agents

The influence of concentration and activity is captured by the Nernst Equation (Lecture 17):

At 25 °C:

Changes in concentration shift the potential away from standard values, forming the basis for concentration cells and sensors.

Reference: Chemguide – Nernst Equation Derivation

12. Mixed and Inert Electrodes

Not all half-cells involve solid metals.
Sometimes inert conductors such as platinum or graphite are used to transfer electrons between solution species.

Example:
The Fe³⁺/Fe²⁺ couple:

Electrode: Pt(s) | Fe³⁺(aq), Fe²⁺(aq)

The electrode itself doesn’t react; it simply provides a surface for electron exchange.

13. Practical Cell Design Considerations

When constructing laboratory or industrial cells, engineers must balance:

Electrode area (affects current density)
Separation distance (affects resistance)
Electrolyte composition
Temperature control
Material compatibility

For example, industrial electrolysis cells use:

Graphite or titanium anodes
Stainless steel cathodes
Membranes instead of salt bridges for ion selectivity

Design optimisation ensures high efficiency, low energy consumption, and long electrode life.

14. Experimental Measurement of EMF

To measure EMF accurately:

Connect the cell to a high-impedance voltmeter (minimises current draw).
Use freshly prepared electrolytes at a known concentration.
Maintain temperature at 25 °C.
Avoid contamination of the salt bridge.

The measured EMF should equal:

If the EMF is positive, the reaction proceeds spontaneously from left to right as written.

Reference: Nuffield Foundation Practical – Measuring Cell EMF

15. From Half-Cells to Real-World Systems

Half-cell concepts extend far beyond laboratory glassware:

Batteries: contain multiple half-cells connected in series or parallel.
Fuel cells: operate continuously with reactant feed streams.
Corrosion cells: form spontaneously on metal surfaces exposed to electrolytes.
Sensors: exploit half-cell potentials (e.g. pH electrode, ion-selective probes).

Example:
The pH glass electrode operates as a half-cell sensitive to H⁺ activity, producing a potential governed by the Nernst equation.

Reference: Metrohm – pH Measurement Principles

16. Summary and Learning Outcomes

By the end of this lecture, you should be able to:

Write and balance oxidation and reduction half-equations for redox systems.
Explain the structure and function of a half-cell and how two combine into a full electrochemical cell.
Describe the role of the salt bridge, electrodes, and electrolytes in maintaining circuit continuity.
Use the cell convention to represent electrochemical cells and predict reaction direction.
Calculate cell EMF using tabulated standard electrode potentials.
Recognise how half-cell principles underpin batteries, electrolysis, corrosion, and sensors.

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