Advanced Chemical Principles and Molecular Dynamics · 6th Year · Oxidation and Reduction · Summer Term

Electrochemical Cells

Explore how spontaneous redox reactions can be harnessed to generate electrical energy in a galvanic (voltaic) cell.

TL;DR:Ever wondered how a simple battery powers your remote control? We're about to uncover the chemistry that converts chemical energy directly into the electrical energy that powers our world.

• NCCA Curriculum Specifications
• Leaving Certificate Chemistry Syllabus: Physical Chemistry - Electrochemistry

About This Topic

This topic on Electrochemical Cells is a cornerstone of the Leaving Certificate Chemistry syllabus, falling under the broader area of electrochemistry. It builds directly upon students' prior understanding of redox reactions, moving from theoretical electron transfer to the practical application of harnessing this energy. The focus is on galvanic, or voltaic, cells, where spontaneous redox reactions are used to generate a potential difference, or voltage. A deep dive into the Daniell cell (Zn/Cu) serves as the classic exemplar, allowing for a clear illustration of all key components: the anode, cathode, half-cells, external circuit, and the crucial role of the salt bridge in maintaining charge neutrality.

Understanding this topic is essential not only for the Leaving Cert examination, where it is frequently assessed through both theoretical questions and calculations of cell potential (E°), but also for its real-world relevance. It provides the fundamental principles behind all batteries, from simple AA cells to the complex lithium-ion batteries powering our mobile phones and electric vehicles. The curriculum requires students to not only draw and label these cells but also to explain the processes at each electrode, predict the direction of electron flow using the electrochemical series, and understand the function of each component in a complete, working circuit.

Key Questions

1. Explain the function of the salt bridge in an electrochemical cell.
2. Compare the processes occurring at the anode and the cathode.
3. Identify the direction of electron flow in a cell constructed from zinc and copper half-cells.

Learning Objectives

• Define the terms anode, cathode, oxidation, reduction, and salt bridge in the context of a galvanic cell.
• Construct and draw a labelled diagram of a Daniell (Zn/Cu) cell, indicating all key components.
• Explain the function of each component of a galvanic cell, including the role of the salt bridge.
• Use the electrochemical series to predict the direction of electron flow and calculate the standard cell potential (E°cell) for a given pair of half-cells.
• Describe the half-reactions occurring at the anode and cathode for a specified electrochemical cell.

Key Vocabulary

Galvanic Cell	An electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell. Also known as a voltaic cell.
Anode	The electrode where oxidation occurs. In a galvanic cell, it is the negative electrode.
Cathode	The electrode where reduction occurs. In a galvanic cell, it is the positive electrode.
Salt Bridge	A connection containing an inert electrolyte that joins the two half-cells of a galvanic cell, allowing ions to flow to maintain charge neutrality.
Half-Cell	A single electrode immersed in a solution of its own ions, where either oxidation or reduction takes place.
Electrode Potential	The potential difference developed between a metal electrode and the solution of its ions at equilibrium.

Watch Out for These Misconceptions

Common Misconception

Electrons travel through the salt bridge to complete the circuit.

What to Teach Instead

The salt bridge allows for the movement of ions (anions and cations) between the two half-cells, not electrons. This ion flow neutralises the charge build-up that occurs as the reaction proceeds, thus completing the circuit and allowing electrons to continue flowing through the external wire.

Common Misconception

The anode is always the negative electrode and the cathode is always the positive one.

What to Teach Instead

This is only true for galvanic (voltaic) cells. In electrolytic cells, the polarity is reversed. It is more accurate to define the electrodes by the reaction type: oxidation always occurs at the anode (An Ox) and reduction always occurs at thecathode (Red Cat), regardless of the cell type.

Common Misconception

Water can be used in the salt bridge.

What to Teach Instead

Pure water is a very poor conductor of electricity because it has a very low concentration of ions. A salt bridge must contain a solution of a soluble, inert ionic salt, like potassium nitrate (KNO₃) or potassium chloride (KCl), to provide mobile ions to balance the charges in the half-cells.

Active Learning Ideas

Simulation Game

Construct a Daniell Cell

Students work in pairs to build a classic galvanic cell using zinc and copper strips as electrodes, solutions of zinc sulfate and copper(II) sulfate as electrolytes, and a filter paper strip soaked in potassium nitrate solution as a salt bridge. They then use a voltmeter to measure the potential difference generated.

40 min·Pairs

Simulation Game

Fruit Battery Investigation

In small groups, students investigate which combination of metals (e.g., zinc, copper, magnesium, iron) produces the highest voltage when inserted into a lemon or potato. This activity provides a tangible, low-tech demonstration of electrochemical principles.

30 min·Small Groups

Simulation Game

Electrochemical Cell Simulation

Using an online interactive simulation (like those from PhET), students can build virtual cells with various metal combinations without the need for physical materials. They can instantly see the resulting voltage, direction of electron flow, and ion movement in the salt bridge.

25 min·Individual

Real-World Connections

• All disposable and rechargeable batteries, such as those in remote controls (zinc-carbon), mobile phones (lithium-ion), and cars (lead-acid).
• The process of corrosion and rusting, which is essentially an unwanted electrochemical cell forming on a metal's surface.
• Electroplating, where a layer of one metal is deposited onto another using electrochemical principles (though this is an electrolytic process, it shares the core concepts).
• Fuel cells, which generate electricity from the reaction of a fuel (like hydrogen) and an oxidant, used in some modern vehicles and for power generation.
• Biological nerve impulses, which involve the movement of ions across cell membranes, creating an electrical potential.

Assessment Ideas

Exit Ticket

Exit ticket: Provide students with a blank diagram of a galvanic cell and ask them to label the anode, cathode, salt bridge, and direction of electron and ion flow.

Quick Check

A Leaving Cert style problem where students are given two half-equations and their E° values. They must identify the anode and cathode, write the overall cell reaction, and calculate the E°cell.

Quick Check

Students use a traffic light system (red, amber, green) to rate their confidence in explaining the function of the salt bridge, identifying the site of oxidation, and predicting electron flow.

Frequently Asked Questions

What happens if you remove the salt bridge while the cell is running?

If the salt bridge is removed, the circuit is broken. Positive charge would build up in the anode half-cell (as metal atoms become ions) and negative charge would build up in the cathode half-cell (as ions gain electrons). This charge imbalance would quickly oppose the flow of electrons, and the current would stop almost instantly.

Why do we use the electrochemical series?

The electrochemical series (a list of standard electrode potentials) allows us to predict the spontaneity of a redox reaction. By comparing the E° values of two half-cells, we can determine which one will be oxidised (the one with the more negative E° value) and which will be reduced, and we can calculate the overall cell potential (E°cell).

Does the size of the electrodes affect the voltage of the cell?

No, the voltage (potential difference) of a galvanic cell is determined by the nature of the metals and the concentration of the ions in the half-cells, not the physical size of the electrodes. However, a larger electrode surface area can allow for a greater rate of reaction, which means a higher current (amperage) can be drawn from the cell.