21 – Investigating Galvanic Cells (LC 2027) Revision Notes for Leaving Cert Chemistry

21 – Investigating Galvanic Cells

Introduction

A galvanic cell (also called a voltaic cell) is an electrochemical device that converts chemical energy into electrical energy through spontaneous redox reactions. This experiment explores how different metals interact in displacement reactions and how we can harness these reactions to generate electricity.

infoNote

The terms "galvanic cell" and "voltaic cell" are interchangeable, both named after scientists Luigi Galvani and Alessandro Volta who pioneered electrochemical research.

Aims of the experiment

This practical investigation has three main objectives:

Build and test a galvanic cell - constructing an electrochemical cell to understand how it generates electrical energy
Measure cell potential - using a voltmeter to determine the voltage produced by different metal combinations
Investigate displacement reactions - testing how metals of different reactivities interact with metal salt solutions

Key theory and predictions

Understanding galvanic cells

In a galvanic cell, two different metals act as electrodes. The more reactive metal becomes the anode (negative electrode) where oxidation occurs - it loses electrons. The less reactive metal becomes the cathode (positive electrode) where reduction occurs - metal ions gain electrons and deposit as solid metal.

chatImportant

Memory aid: Remember "An Ox, Red Cat" - Anode Oxidation, Reduction Cathode

Predicting cell behaviour

The voltage generated depends on the metals chosen and their positions in the reactivity series. A greater difference in reactivity produces a higher voltage. For example, zinc is more reactive than copper, so in a Zn/Cu cell:

lightbulbExample

Worked Example: Zinc-Copper Cell Reactions

Zinc electrode: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$ (oxidation at anode)
Copper electrode: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$ (reduction at cathode)

The zinc loses electrons (is oxidised) while copper ions gain electrons (are reduced).

Displacement reactions

A more reactive metal will displace a less reactive metal from its salt solution. For instance, zinc metal will displace copper ions from copper sulphate solution, causing copper metal to deposit on the zinc surface.

Materials and apparatus needed

Essential equipment:

Clean metal strips (zinc, copper, iron) - polished and labelled
Metal salt solutions (0.1 M or 1.0 M concentrations) - $\text{ZnSO}_4$ , $\text{CuSO}_4$ , $\text{FeSO}_4$
Small beakers or cells for solutions
Digital voltmeter with connecting wires and crocodile clips
Salt bridge - philtre paper soaked in potassium nitrate ( $\text{KNO}_3$ ) or potassium chloride ( $\text{KCl}$ )

Additional materials:

Sandpaper or steel wool for cleaning electrodes
Measuring cylinders and pipettes
Distilled water
Labels and stirrers
Safety equipment: goggles, gloves, lab coat

infoNote

The salt bridge is crucial for the cell's operation as it maintains electrical neutrality by allowing ions to flow between the half-cells while preventing the solutions from mixing completely.

Safety precautions

chatImportant

Essential safety measures:

Wear safety goggles and gloves throughout the experiment
Handle metal salt solutions carefully to avoid spills
Rinse skin immediately if contact with solutions occurs
Dispose of metal salt solutions according to laboratory waste policy
Never short-circuit the cell - avoid direct connection without a load as this can cause heating and damage

Experimental procedure

Part A: Displacement reactions

This preliminary investigation tests which metals can displace others from solution.

Prepare the metals - clean metal strips with sandpaper and label clearly
Set up solutions - place different metal salt solutions ( $\text{CuSO}_4$ , $\text{FeSO}_4$ , etc.) in separate beakers
Test displacement - immerse one type of metal (e.g. zinc) into a different metal ion solution (e.g. $\text{Cu}^{2+}$ )
Observe changes - look for evidence such as colour changes, solid formation on the metal strip, or solution colour changes
Record results - note which combinations show displacement and write balanced equations for reactions that occur

Part B: Galvanic cell construction

Here you'll build a functioning electrochemical cell.

Prepare electrodes - clean and label metal strips that will serve as electrodes
Set up half-cells - place each metal strip in a beaker containing its corresponding metal ion solution
Connect with salt bridge - link the two solutions using philtre paper soaked in $\text{KNO}_3$ or $\text{KCl}$ . This allows ions to flow while preventing the solutions from mixing completely
Connect to voltmeter - attach crocodile clips to the metal electrodes and connect to a digital voltmeter
Measure EMF - record the electromotive force (voltage) of the cell under open circuit conditions (no additional load)
Compare combinations - try different metal pairs and concentrations to see how voltage changes

lightbulbExample

Worked Example: Setting Up a Zn/Cu Cell

Step 1: Place zinc strip in $\text{ZnSO}_4$ solution (left half-cell) Step 2: Place copper strip in $\text{CuSO}_4$ solution (right half-cell) Step 3: Connect half-cells with salt bridge Step 4: Connect voltmeter (zinc = negative terminal, copper = positive terminal) Expected voltage: approximately 1.1 V

Part C: Voltaic pile construction

This demonstrates how multiple cells can be combined to increase voltage.

Stack alternating metals - create layers of different metal discs (e.g. zinc and copper) with separators soaked in salt solution between each pair
Connect electrodes - attach the voltmeter to the bottom and top electrodes of the stack
Measure total voltage - compare this with the single-cell voltage; multiple cells in series should produce higher total voltage

Data collection and recording

Create organised data tables including:

Metal pair combinations used
Solutions and their concentrations
Measured voltages for each cell
Visual observations (bubble formation, colour changes, metal deposition)
At least 2-3 replicate readings for accuracy

infoNote

Note any anomalies such as contamination, poor electrical contact, or unexpected solution mixing that might affect results. These observations are crucial for evaluating the reliability of your data.

Analysis and calculations

Comparing with theory

Use your measured voltages to compare with known standard electrode potentials if available. This helps verify whether your cells behaved as expected.

Understanding trends

Analyse patterns in your data. For example, you might find that a Zn/Cu cell produces around 1.0 V, while a Zn/Fe cell gives a lower voltage, reflecting their positions in the reactivity series.

Advanced analysis

If studying concentration effects, you may encounter Nernst equation behaviour, where voltage depends on the logarithm of ion concentration.

For displacement reactions, write balanced redox equations and identify which species are oxidised and reduced in each case.

lightbulbExample

Worked Example: Writing Redox Equations

For zinc displacing copper:

Oxidation: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$
Reduction: $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$
Overall: $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

Evaluation and improvements

Common sources of error

infoNote

Contact resistance - poor electrical connections between electrodes and wires can significantly reduce measured voltage. Ensure all connections are clean and tight.

Electrode contamination - impurities on metal surfaces affect reactions and can lead to inconsistent results.

Voltmeter drift - instrument errors over time may cause readings to change even with a stable cell.

Solution mixing - if the salt bridge allows too much solution mixing, the cell may lose effectiveness.

Salt bridge resistance - high resistance limiting current flow can reduce the measured EMF.

Suggested improvements

Clean electrodes thoroughly before each measurement
Use fresh solutions and ensure good electrical contact
Take multiple readings and calculate averages
Try stronger salt bridges or different electrolytes
Control temperature as this affects cell voltage
Test different metal combinations and concentration ranges

Conclusion points

After completing the experiment, you should be able to:

Confirm which metals act as anodes and cathodes in different combinations
Identify successful displacement reactions based on the reactivity series
Explain the relationship between metal reactivity, redox potential, and voltage generation
Evaluate how well your experimental design worked and suggest improvements

bookmarkSummary

Key Points to Remember:

Galvanic cells convert chemical energy to electrical energy through spontaneous redox reactions between different metals
The more reactive metal always becomes the anode (loses electrons), while the less reactive metal becomes the cathode (gains electrons)
Displacement reactions follow the reactivity series - more reactive metals displace less reactive ones from solution
Salt bridges are essential for allowing ion flow while preventing solution mixing that would stop the cell working
Cell voltage depends on the metal combination and increases with greater differences in reactivity

21 – Investigating Galvanic Cells (LC 2027) (Leaving Cert Chemistry): Revision Notes