Electrochemistry - Lecture 7: Galvanic Cells and Spontaneous Redox Systems

A modular guide to electrochemical energy conversion, electron flow, and spontaneous reaction design

Electrochemical cells can broadly be divided into galvanic (voltaic) and electrolytic systems. While electrolytic cells require an external power source to drive non-spontaneous reactions, galvanic cells operate spontaneously, harnessing the natural tendency of redox reactions to move electrons from higher to lower chemical potential.

These devices convert chemical energy into electrical energy, powering everything from portable batteries to corrosion sensors and fuel cells.

This lecture explores how galvanic cells function, how spontaneity is determined, and how their voltage (electromotive force, EMF) is related to thermodynamics. You’ll also learn about the direction of electron flow, cell notation conventions, and examples such as the Daniell cell and lead–acid accumulator.

1. What Is a Galvanic Cell?

A galvanic cell (also known as a voltaic cell) is an electrochemical system that generates electricity from a spontaneous redox reaction.

It consists of two half-cells:

An anode where oxidation occurs.
A cathode where reduction occurs.

Electrons flow through an external circuit from anode → cathode, while ions migrate through a salt bridge or membrane to complete the circuit internally.

The fundamental requirement for a galvanic cell is that the two half-reactions have different tendencies to gain or lose electrons, creating a potential difference.

Reference: Royal Society of Chemistry – Voltaic Cells

2. The Nature of Spontaneity in Redox Reactions

A spontaneous redox reaction occurs naturally under given conditions without external electrical input.

Thermodynamically, spontaneity is linked to:

A negative Gibbs free energy change (ΔG < 0)
A positive cell potential (Ecell > 0)

The relationship between them is given by:

where
n = number of moles of electrons transferred
F = Faraday constant (96,485 C mol⁻¹)

Thus, a positive EMF corresponds to a negative ΔG, meaning energy is released as electrons flow spontaneously.

Reference: LibreTexts – Gibbs Free Energy and Cell Potential

3. The Daniell Cell: A Classical Example

The Daniell cell is the archetype of a galvanic system and forms the basis for understanding modern batteries.

Construction:

Zinc electrode immersed in 1 M ZnSO₄ solution
Copper electrode immersed in 1 M CuSO₄ solution
Connected via salt bridge (KNO₃ or KCl) and external circuit

Half-reactions:

Overall reaction:

Standard electrode potentials:

The positive EMF indicates that the reaction is spontaneous, with electrons flowing from zinc (anode) to copper (cathode).

Reference: RSC – The Daniell Cell Experiment

4. Electron Flow and Conventional Current

It’s vital to distinguish between electron flow and conventional current:

Quantity	Direction	Description
Electron flow	Anode → Cathode	Actual movement of electrons through a wire
Conventional current	Cathode → Anode	The direction positive charge would flow (used in circuit diagrams)

Inside the cell, ionic movement maintains charge balance:

Zn²⁺ ions enter the solution at the anode.
Cu²⁺ ions are reduced and removed from the cathode solution.
Negative ions from the salt bridge migrate to the anode side, and positive ions migrate to the cathode side.

5. Cell Notation and Representation

The structure of a galvanic cell can be expressed compactly using cell notation:

Notation conventions:

Single line “|” → phase boundary
Double line “||” → salt bridge
Left-hand side → anode (oxidation)
Right-hand side → cathode (reduction)

Tip: Always write the cell so that the spontaneous reaction gives a positive E₍cell₎.

Reference: Chemguide – Cell Diagrams

6. Measuring EMF of a Galvanic Cell

The electromotive force (EMF) or cell potential measures the maximum potential difference between the two electrodes when no current flows (open circuit).

Conditions for standard EMF:

Solute concentrations = 1 mol dm⁻³
Gas pressures = 1 atm
Temperature = 298 K (25 °C)

Instrumentation:

High-impedance voltmeter (to minimise current flow)
Freshly prepared salt bridge
Clean electrodes

Reference: Nuffield Foundation – Measuring EMF

7. Thermodynamic Interpretation

The EMF of a galvanic cell relates directly to the maximum non-expansion work a chemical system can perform:

Since E₍cell₎ > 0, ΔG is negative, indicating spontaneous energy release.

At equilibrium, E₍cell₎ = 0 and ΔG = 0, showing that the cell can no longer perform work.

The Nernst equation (developed later in Lecture 17) extends this relation to non-standard conditions, revealing how EMF depends on reactant and product activities.

8. The Lead–Acid Cell: A Practical Galvanic System

The lead–acid accumulator, used in car batteries, is a real-world galvanic cell that can be recharged (reversible redox system).

Half-reactions (discharging):

Overall reaction:

Cell potential:

When charging, an external power source reverses these reactions, restoring Pb and PbO₂ electrodes.

Reference: BBC Bitesize – Lead–Acid Battery Chemistry

9. Hydrogen–Oxygen (Fuel) Galvanic Cell

Fuel cells are galvanic systems where chemical fuels react continuously with oxidants to produce electricity.

Anode (hydrogen oxidation):

Cathode (oxygen reduction):

Overall:

This reaction produces water as the only product, offering clean, continuous power.

Reference: Fuel Cell & Hydrogen Energy Association

10. Factors Affecting Cell Potential

The actual EMF of a galvanic cell can deviate from standard values due to:

Concentration effects – captured by the Nernst equation
Temperature – affects equilibrium constants and reaction kinetics
Ion activity – non-ideal solutions alter effective concentrations
Electrode surface area and cleanliness
Internal resistance – due to electrolyte conductivity or membrane fouling

Reference: Chemguide – Factors Affecting EMF

11. Polarisation and Overpotential

As current flows, polarisation may occur, a deviation from the equilibrium potential caused by:

Slow electron transfer (activation polarisation)
Mass-transport limitations (concentration polarisation)
Formation of gas films or surface fouling (resistive overpotential)

Polarisation reduces the effective voltage and is a key consideration in battery efficiency and electrode design.

12. Reversible and Irreversible Galvanic Systems

A reversible cell achieves equilibrium potential without net current when the circuit is open and can be reversed by applying an infinitesimal voltage in the opposite direction.

An irreversible cell exhibits significant energy loss, often due to side reactions, diffusion limitations, or polarisation.

Practical galvanic systems (e.g. batteries) aim for quasi-reversibility: high efficiency and recoverable charge over many cycles.

13. Calculating EMF from Standard Potentials

Example problem:

Given:

Calculate E₍cell₎ and the overall reaction.

Solution:

Overall reaction:

→ Spontaneous, since E₍cell₎ > 0.

14. Practical Galvanic Systems in Everyday Life

Application	Description	Cell Type
Dry cell (Leclanché)	MnO₂ cathode, Zn anode, NH₄Cl electrolyte	Primary (non-rechargeable)
Alkaline battery	Zn/KOH/MnO₂ system	Higher capacity, low polarisation
Lead–acid battery	Pb/PbO₂/H₂SO₄	Secondary (rechargeable)
Nickel–cadmium (NiCd)	NiOOH/Cd, alkaline electrolyte	Rechargeable
Lithium-ion cell	Graphite anode, LiCoO₂ cathode	High energy density

Each device is based on the same galvanic principle: electrons flow spontaneously from the material with a lower reduction potential to that with a higher reduction potential.

Reference: ScienceDirect – Battery Fundamentals

15. Energy Efficiency and Work Output

The total electrical work obtainable from a galvanic cell:

The efficiency depends on:

In practical systems, efficiency < 100 % due to internal resistance, heat loss, and polarisation.

Example: A lead–acid battery with ΔH ≈ 210 kJ mol⁻¹ and Ecell = 2.04 V has η ≈ 90 %.

Reference: Nature Energy – Battery Thermodynamics

16. Environmental and Technological Significance

Galvanic systems are pivotal in:

Renewable energy storage (batteries, flow cells)
Transport electrification
Portable electronics
Corrosion prevention (sacrificial anodes)
Analytical instrumentation (reference electrodes, sensors)

In environmental electrochemistry, galvanic coupling explains corrosion processes, as in steel structures, where anodic and cathodic regions form naturally in moist environments.

Reference: UK Corrosion Science – Galvanic Corrosion Overview

17. Linking Galvanic Cells to the Electrochemical Series

The electrochemical series (Lecture 12) lists standard electrode potentials in ascending order.

Metals higher in the series (more negative E°) act as anodes when coupled with those lower in the series (more positive E°).

Rule:
Electrons flow from the species with a lower reduction potential → higher reduction potential.

This series thus predicts the direction of spontaneous electron flow and the feasibility of galvanic combinations.

Reference: Chemguide – The Electrochemical Series

18. Design Considerations for Industrial Galvanic Cells

Key design factors:

Electrode materials – corrosion resistance and conductivity
Electrolyte composition – ionic strength and stability
Temperature control – affects reaction rates
Cell geometry – optimises diffusion paths
Maintenance – electrolyte replenishment, gas removal

Large-scale galvanic systems, such as metal–air batteries or redox flow cells, rely on modular scaling and high surface-area electrodes.

19. Experimental Demonstration: Constructing a Galvanic Cell

Materials:

1 M ZnSO₄ solution
1 M CuSO₄ solution
Zinc and copper strips
Salt bridge (KNO₃ in agar)
Voltmeter and connecting wires

Steps:

Fill two beakers with the respective solutions.
Insert zinc and copper electrodes.
Connect via voltmeter and salt bridge.
Observe the voltage reading (≈ 1.10 V).
Note that zinc dissolves (anode) and copper deposits (cathode).

Safety: Wear goggles and gloves. Dispose of metal solutions responsibly.

Reference: RSC Practical – Simple Galvanic Cells

20. Summary and Learning Outcomes

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

Define a galvanic cell and distinguish it from an electrolytic cell.
Identify anode and cathode reactions in a given redox system.
Write full and half-cell equations and use cell notation correctly.
Predict spontaneity using electrode potentials and Gibbs free energy.
Calculate Ecell from standard electrode potentials.
Explain the operation of real systems such as Daniell, lead–acid, and hydrogen fuel cells.
Understand the role of galvanic principles in batteries, corrosion, and sensors.

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