GALVANIC CELLS AND THE CALCULATIONS OF CELL POTENTIAL LABORATORY 6 ELECTROCHEMISTRY

GEORGIA MILITARY COLLEGE

NATURAL SCIENCE DEPARTMENT

ONLINE CAMPUS

LABORATORY 6 – ELECTROCHEMISTRY: GALVANIC CELLS AND THE CALCULATIONS OF CELL POTENTIAL

NAME

STUDENT NUMBER

CLASS

PROFESSOR’S TITLE AND NAME

Introduction

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 is covered with a greyish fuzzy material (Stubbs et al., 2022). These changes are a result of 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, which 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 towards it. Conversely, the anode has high potential and electrons flow away from it (Stubbs et al., 2022). Electrons are transferred spontaneously from the elementary copper wire to the silver ions in the solution (Stubbs et al., 2022). This is a redox reaction, which is a 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 further 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 that lost in oxidation.

Now consider a case where the two electrodes (cathode and anode) are placed in separate containers and hence, there is 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 cathode to the anode 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 below:

Figure 1: Zinc-copper galvanic cell

The two reactants (1M Cu(No3) and 1M Zn(No3) are placed in separate containers, and a clean copper and zinc rod placed in the two solutions respectively. The ends of both rods are connected using a wire to a voltmeter to create an external circuit joining the two containers. Thus, both rods serve as electrodes. At this point, however, there is negligible electron flow through the wire since 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 reaction can also be represented by a cell notation as shown below:

Zn(s)?1M Zn2+ (aq) ? 1M 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 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.

Materials and Methods

Preparing the Lab

i) Research ‘Magnesium-Iron ‘galvanic cells from available resources.

ii) Develop a hypothesis explaining the process behind the electric current formation in the galvanic cell.

Performing the Lab

iii) Write a summary of your findings on the magnesium-iron galvanic cell (half-page long)

Data Analysis

iv) Using the summary in (iii) above, name the anode, cathode, salt bridge, and show the direction of electron flow in the magnesium-iron galvanic cell diagram.

v) Describe the parts that make a half cell (3-5 sentences)

vi) Write a chemical reaction for the galvanic cell’s redox process.

vii) Write chemical reactions for the reduction and oxidation processes in the galvanic cell’s half-cells.

viii) Describe what happens to each electrode in the galvanic cell in terms of mass

ix) Develop the cell notation for the galvanic cell.

x) Discuss the direction in which electrons flow in the cell.

xi) Use the second method to calculate cell potential for the reaction.

Data and Discussion

Data

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 allowed in cases where the redox couple in a half-cell may not function as an electrode. Iron is not used as the cathode electrode because it is highly corrosive and also a solute species (Lumen Learning, 2024). Thus, platinum (Pt) is used as an inert cathode electrode – since it is chemically unreactive, it does not participate in the redox reaction, but merely accepts electrons, thus allowing the flow of current (Lumen Learning, 2024). A filter paper saturated with potassium nitrate (KN03) is used as the cell’s salt bridge.

The galvanic cell diagram for the magnesium-iron galvanic cell is presented in figure 2 below:

Figure 2: Magnesium-iron galvanic cell

Elements in the Galvanic Cell:

Blue Rod immersed in the left half-cell: Magnesium anode (Mg anode)

Red Rod immersed in right half-cell: Platinum cathode (Pt cathode)

Green tube with inverted-U shape with ends beneath the surface of the solutions in the half-cells: Salt Bridge: filter paper saturated with KN03.

The arrow depicts the flow of electrons from the anode to the cathode (left to right in the standard galvanic cell diagram).

Purple aqueous solution in left half-cell: magnesium chloride solution, MgCl2 (aq)

Grey aqueous solution in right half-cell: iron (III) chloride, FeCl3 (aq), and iron (II) chloride, FeCl2 (aq), mixture.

The half-cell is made up of a metal rod immersed in an aqueous solution. The left half-cell is made up of the Mg(s)/Mg(II) couple, which comprises an aqueous magnesium chloride solution, 0.1M MgCl2 (aq) and a magnesium anode rod. The left half-cell comprises a mixture of aqueous 0.3M iron (II) and 0.2M iron (III) chloride solution, in which a platinum cathode rod (Pt) is immersed. Thus, the cell notation presented as:

Mg(s)?0.1M MgCl2 (aq) ? 0.2M FeCl3(aq), 0.3M FeCl2(aq)? Pt(s)……………(iii)

The chemical reaction for the galvanic cell’s 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 for the magnesium anode electrode, while the mass of the platinum cathode remains unchanged.

Discussion

From the reactivity series, magnesium is more reactive than platinum (Stubbs et al., 2022). Thus, the magnesium rod will serve as the reducing agent, with a higher electron potential, and will lose electrons to the platinum (Stubbs et al., 2022). This makes the magnesium rod the anode, and platinum the cathode. As shown by the right-pointing arrow in the galvanic cell diagram (figure 2), electrons will, therefore, flow from the magnesium anode electrode to the platinum cathode electrode. The process of losing electrons at the anode is referred to as oxidation (as shown by equation iv above). As the solid magnesium rod is oxidized, it dissolves and forms aqueous Mg2+ ions that increase the positive charge in the solution in the left half-cell. The dissolution of the solid magnesium into aqueous ions causes the anode rod to lose mass and become smaller (Lumen Learning, 2024).

The electrons lost in the anode flow to the cathode. However, since platinum is inert, it does not take part in the reduction-oxidation reaction and only allows the flow of current. Therefore, as it gains electrons from the anode, the mass and structure of the platinum cathode does not change. However, as shown by the reduction equation (vi), the electrons gained reduce the positive charge of the aqueous Fe3+ ions in the right half-cell, leading to the formation of gaseous Fe2+ ions, which form yellow-green fumes.

As the electrons move from the anode to the cathode (left to right), the nitrate ions (anions) in the salt bridge move to the left into the MgCl2 and neutralize the Mg2+ produced by the oxidization of magnesium metal, thus ensuring the right electrical neutrality in the left half-reaction. Similarly, the potassium ions (cations) in the salt bridge move into the right half-cell solution and replace the Fe3+ ions that undergo reduction to form gaseous Fe2+ ions (Lumen Learning, 2024). This ensures the right electrical neutrality in the right half-cell. The electrical neutrality in both half-cells due to the salt bridge ensures that current flows between the two compartments. The salt bridge provides the moving ions that transfer current in the galvanic cell. The reading on the voltmeter gives the cell potential, which is a measure of the voltage that exists between two half-cells in a galvanic cell