Electrochemistry: Galvanic Cells and Cell Potential Calculations
This laboratory report examines the fundamental principles of electrochemistry through the construction and analysis of a magnesium-iron galvanic cell. The paper covers core concepts including oxidation and reduction, anode and cathode roles, salt bridge function, and electron flow. Using a magnesium anode rod in aqueous magnesium chloride connected via a potassium nitrate salt bridge to an inert platinum cathode in an iron chloride solution, the report identifies and explains each component's electrochemical role. Half-cell reactions, cell notation, mass changes at each electrode, and cell potential calculations are all addressed. The report concludes that the salt bridge maintains electrical neutrality in both half-cells, enabling sustained current flow as measured by the voltmeter.
- Introduction to Redox Reactions and Galvanic Cells: Redox concepts, galvanic cell theory, and objectives
- Materials and Methods: Lab preparation steps and analysis tasks
- Magnesium-Iron Galvanic Cell: Data and Observations: Cell components, diagram description, and reactions
- Half-Cell Components and Cell Notation: Half-cell structure, notation, and electrode identity
- Discussion of Electrode Behavior and Salt Bridge Function: Electron flow, mass changes, and salt bridge role
- Conclusion: Summary of galvanic cell findings and principles
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What makes this paper effective
- The introduction builds concepts systematically, moving from a simple copper-silver observation to the more complex galvanic cell setup, giving readers a clear conceptual scaffold.
- Chemical equations are presented at each stage — full redox, oxidation half-reaction, and reduction half-reaction — reinforcing the relationship between symbolic notation and physical observation.
- The discussion explicitly connects each observation (mass loss at the anode, unchanged platinum cathode, salt bridge ion movement) to the underlying electrochemical principles, demonstrating strong cause-and-effect reasoning.
Key academic technique demonstrated
The paper exemplifies the use of half-reaction analysis to break a complex redox process into manageable components. By isolating oxidation and reduction into separate equations and then reconciling electron counts, the author shows how systematic stoichiometric reasoning supports both qualitative descriptions and quantitative cell potential calculations — a core skill in undergraduate chemistry laboratory writing.
Structure breakdown
The report follows a standard laboratory format: an introduction establishing theory and objectives, a brief methods section, a data section presenting cell components and reactions, a discussion interpreting observations mechanistically, and a conclusion summarizing findings against the stated objectives. This structure mirrors the scientific method and is typical of undergraduate natural science laboratory reports.
Introduction to Redox Reactions and Galvanic Cells
Consider the result of immersing a clean copper wire into an aqueous silver nitrate solution. A spontaneous change occurs in which the solution turns from colorless to blue and the copper wire that was initially clean becomes covered with a greyish, fuzzy material (Stubbs et al., 2022). These changes result from the processes of oxidation and reduction. Copper undergoes oxidation to produce copper (II) ions, Cu2+ (aq), which cause the clear solution to turn blue, while the silver (I) ions in the silver nitrate solution undergo reduction, producing elemental silver that forms a greyish deposit on the copper wire (Stubbs et al., 2022). The copper metal undergoes oxidation and is thus referred to as the anode, while the silver undergoes reduction and is referred to as the cathode (Stubbs et al., 2022).
Oxidation is the process by which an ion, atom, or molecule loses electrons during a chemical reaction, while reduction is the process by which atoms or ions gain electrons during a chemical reaction (Stubbs et al., 2022). The cathode is the electrode with lower electron potential, and hence electrons flow toward it. Conversely, the anode has higher potential and electrons flow away from it (Stubbs et al., 2022). Electrons are transferred spontaneously from the elemental copper wire to the silver ions in the solution (Stubbs et al., 2022). This is a redox reaction, characterized by changes in the reactants' oxidation states (Stubbs et al., 2022).
The overall redox reaction for the reaction between copper and silver can be represented as follows:
2Ag+ (aq) + Cu (s) → 2Ag (s) + Cu2+ (aq) ………………………………. (i)
This equation can be split into two half-reactions that separate the reduction and oxidation reactions:
Oxidation: Cu (s) → Cu2+ (aq) + 2e− …………………………………… (ii)
Reduction: 2Ag+ (aq) + 2e− → 2Ag (s) …………………………………… (iii)
The reduction reaction (iii) is doubled so that the number of electrons gained in reduction equals the number lost in oxidation.
Now consider a case where the two electrodes (cathode and anode) are placed in separate containers, with no physical contact between the reactants. In such a case, electrons are transferred indirectly with the help of an external circuit connecting the two reactants (Stubbs et al., 2022). This kind of arrangement is an electrochemical cell, characterized by the transfer of electrons from the anode to the cathode via an external circuit (Stubbs et al., 2022). An electrochemical cell in which the redox reaction is spontaneous is referred to as a galvanic or voltaic cell (Stubbs et al., 2022).
Consider a zinc-copper galvanic cell as shown in Figure 1. The two reactants — 1 M Cu(NO3)2 and 1 M Zn(NO3)2 — are placed in separate containers, and a clean copper rod and a clean zinc rod are placed in the two solutions respectively. The ends of both rods are connected by a wire to a voltmeter to create an external circuit joining the two containers; both rods thus serve as electrodes. At this point, however, there is negligible electron flow through the wire because the circuit is open. To close the circuit, a non-reactive, concentrated electrolyte solution such as sodium chloride is used as a salt bridge (Stubbs et al., 2022). The salt bridge is an inert solution used to provide electrical contact between the two containers. This cell can also be represented by the following cell notation:
Zn (s) | 1 M Zn2+ (aq) ‖ 1 M Cu2+ (aq) | Cu (s)
Based on the above context, this laboratory seeks to realize the following objectives:
(i) To enhance understanding of the construction of, and principles behind, the galvanic cell.
(ii) To enhance understanding of spontaneous electrochemical reactions.
(iii) To improve the ability to use standard reduction potentials to calculate cell potential.
(iv) To enhance familiarity with cell diagrams.
The general hypothesis is that the salt bridge contributes to current formation in the galvanic cell through the release of anions and cations that transport current between the two half-cells.
Magnesium-Iron Galvanic Cell: Data and Observations
Magnesium-iron galvanic cells make use of solid magnesium and aqueous iron ions, with an inert electrode such as platinum or gold serving as the cathode. In one such cell, a magnesium rod immersed in an aqueous magnesium chloride solution is connected using a salt bridge to a platinum rod immersed in an aqueous mixture of iron (II) and iron (III) chloride solutions. The inert electrode — platinum in this case — is neither a product of the reaction nor a reactant. This is permitted in cases where the redox couple in a half-cell cannot function as an electrode. Iron is not used as the cathode electrode because it is highly corrosive and is also a solute species (Lumen Learning, 2024). Platinum (Pt) is therefore used as the inert cathode electrode; since it is chemically unreactive, it does not participate in the redox reaction but merely accepts electrons, allowing current to flow (Lumen Learning, 2024). A filter paper saturated with potassium nitrate (KNO3) serves as the cell's salt bridge.
The galvanic cell diagram for the magnesium-iron galvanic cell is presented in Figure 2. The key elements of the cell are as follows:
Blue rod (left half-cell): Magnesium anode (Mg anode).
Red rod (right half-cell): Platinum cathode (Pt cathode).
Green inverted-U tube: Salt bridge — filter paper saturated with KNO3.
Arrow direction: Depicts electron flow from the anode to the cathode (left to right).
Purple aqueous solution (left half-cell): Magnesium chloride solution, MgCl2 (aq).
Grey aqueous solution (right half-cell): A mixture of iron (III) chloride, FeCl3 (aq), and iron (II) chloride, FeCl2 (aq).
Half-Cell Components and Cell Notation
A half-cell is made up of a metal rod immersed in an aqueous solution. The left half-cell comprises the Mg (s)/Mg (II) couple, which consists of an aqueous magnesium chloride solution — 0.1 M MgCl2 (aq) — and a magnesium anode rod. The right half-cell comprises a mixture of 0.3 M iron (II) and 0.2 M iron (III) chloride solution, in which a platinum cathode rod (Pt) is immersed. The cell notation is presented as:
Mg (s) | 0.1 M MgCl2 (aq) ‖ 0.2 M FeCl3 (aq), 0.3 M FeCl2 (aq) | Pt (s) ……………… (iii)
The chemical reaction for the galvanic cell's overall redox process is:
Mg (s) + 2Fe3+ (aq) → Mg2+ (aq) + 2Fe2+ (g) ……………………… (iv)
The chemical reactions for the reduction and oxidation processes in the galvanic cell's half-cells are:
Oxidation: Mg (s) → Mg2+ (aq) + 2e− ……………………………………… (v)
Reduction: 2Fe3+ (aq) + 2e− → 2Fe2+ (g) …………………………………… (vi)
The end result of the galvanic cell reaction is a reduction in mass at the magnesium anode electrode, while the mass of the platinum cathode remains unchanged.
Conclusion
This laboratory sought to enhance understanding of the construction of, and principles behind, galvanic cells, as well as spontaneous electrochemical reactions. The galvanic cell is based on the concepts of oxidation and reduction occurring at the anode and cathode electrodes respectively. It is made up of two half-cells connected via an external circuit. Oxidation is the process by which an anode molecule, atom, or ion loses electrons, while reduction is the process by which the cathode gains electrons.
This laboratory investigated the magnesium-iron galvanic cell. The left half-cell comprised a magnesium anode rod immersed in an aqueous solution of magnesium chloride, while the right half-cell comprised an inert platinum cathode immersed in an aqueous mixture of iron (II) and iron (III) chloride solutions. The salt bridge consisted of a filter paper saturated with aqueous potassium nitrate solution. Since magnesium is more reactive than platinum, it donates electrons in the oxidation process, causing the solid magnesium to dissolve and form aqueous magnesium ions. As a result, the mass of the anode rod decreases. The electrons flow to the cathode, where reduction of iron (III) ions takes place, leading to the formation of gaseous iron (II) ions. There is no change in mass of the inert platinum cathode. The movement of anions and cations from the salt bridge ensures current flows in the cell, as measured by the voltmeter reading.
References
Lumen Learning. (2024). Module 17: Electrochemistry, galvanic cells. https://courses.lumenlearning.com/chemistryformajors/chapter/galvanic-cells/
Stubbs, A., Taylor, N., Stokes, R., Murray, B., Burrows, K., Derbogosian, M., Raphael, S., Bowen, S., & Sharma, L. (2022). Jacaranda chemistry 1. John Wiley & Sons.
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