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Chemistry

Variation of Cell Potential in Zn Cu Galvanic Cell

Variation of Cell Potential in Zn Cu Galvanic Cell

Q: 1. What is the cell potential of a Zn–Cu cell under standard conditions?

The standard cell potential (E°cell) of a Zn–Cu cell is +1.10 V under standard conditions. Standard reduction potentials: Zn2+ + 2e- → Zn(s), E° = −0.76 VCu2+ + 2e- → Cu(s), E° = +0.34 VE°cell = E°cathode − E°anode = 0.34 − (−0.76) = 1.10 VStandard conditions: 1 M solutions, 1 atm pressure, 25°C (298 K).This positive value shows the Zn–Cu galvanic cell is spontaneous.

Q: 3. What is the Nernst equation for the Zn–Cu cell?

The Nernst equation for the Zn–Cu cell at 25°C is E = E° − (0.0591/2) log([Zn2+]/[Cu2+]). Overall reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)Number of electrons transferred (n) = 2Reaction quotient Q = [Zn2+]/[Cu2+]This equation explains how cell potential varies with concentration in a Daniell cell.

Q: 6. What happens to the cell potential when the Zn–Cu cell reaches equilibrium?

When the Zn–Cu cell reaches equilibrium, the cell potential becomes zero (E = 0 V). At equilibrium, the reaction quotient Q equals the equilibrium constant K.From thermodynamics: ΔG = −nFEAt equilibrium, ΔG = 0, so E = 0.This means no net electron flow occurs once equilibrium is established.

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Ritika Singla

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How Concentration and Nernst Equation Affect Cell Potential in Zn Cu Cell

The mechanism of the generation of electricity in a Daniell Cell is remarkable. Students learn how the concentration of the electrolyte can also alter cell potential. It is to be measured by doing an experiment. In this experiment, the variation of cell potential in Zn - Cu cell will be determined by changing the concentration of the electrolytes. Every part of this experiment will be elaborated so that you can easily understand when the standard electrode potential for Daniell cell is 1.1 V and when it varies. Study every section of this article and understand this experiment properly.

Variation in the Potential Difference in a Galvanic Cell: How to Conduct the Experiment?

The voltaic or galvanic cell used in this experiment contains a zinc and copper electrode connected with a voltmeter. The electrodes will carry electrons from the positive end to the negative end and the voltmeter will measure the cell potential. The cell potential is actually the difference we find between the potential of both the electrodes dipped in an electrolytic solution. The potential difference is measured in volts. It is also called the electromotive force of a cell when no circuit is connected to draw a current.

Aim of the Experiment

The aim of this experiment is to measure the variation of cell potential in Zn - Cu cell when the concentration of the electrolyte is changed at room temperature.

Theory of the Experiment

The inter-conversion of different forms of energies into each other has been studied before. In this experiment, you will study how chemical energy is converted into electrical energy in a Zn - Cu cell and how its potential is affected when the electrolyte concentration is altered.

For this experiment, we will use a Daniell Cell where zinc present in the electrode will react with copper sulphate in the electrolyte solution. The same reaction gives the precipitation of copper in a test tube but when carried in a Daniell Cell, it generates electricity. It becomes an electrochemical reaction conducted in a cell where chemical energy is converted into electrical energy. A flow of current will be noticed in the conductor connecting the electrodes. The potential difference between the electrodes will be measured by the voltmeter connected.

The reactions taking place in the electrodes in this galvanic cell are:

Zn Electrode: Zn (s) → Zn2+ (aq) + 2e–

Cu Electrode: Cu2+ (aq) + 2e– → Cu (s)

As you can see, zinc from the anode goes into the solution in the form of cations due to oxidation whereas copper goes into the solution in the form of copper atoms and precipitates. If we summarize both the electrode reactions, we will get:

Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)

The standard electrode potential for Daniell cell is 1.1 V. When the EMF is less than 1.1 V, you will observe that the electrons flow from the zinc anode to the copper cathode. You will also witness the deposition of copper in the cathode and the dissolution of zinc in the anode.

When the standard electric potential is more than 1.1 V, the current flows from the copper to the zinc electrode. The resultant reaction will also be the opposite.

Things Required For the Experiment

Cu Zn anode cathode
1 M solution of Zinc Sulphate (ZnSO4) and Copper Sulphate (CuSO4)
Beakers
Voltmeter
Salt bridge
Conducting wires

How to Set the Apparatus and Conduct the Experiment?

Preparing the salt bridge
Take a U-shaped glass tube. Heat agar-agar gel (20 gm) with potassium chloride (5 gm of KCl) in a clean beaker. Introduce the solution into the U tube by sucking the solution and let it cool.
Prepare 0.1 M solution of Zinc Sulphate and Copper Sulphate and put them in two separate beakers.
Put the salt bridge in such a way that it connects the electrolytes in both beakers.
Dip the zinc electrode with the negative end of the voltmeter and the copper electrode with the positive end with the help of conducting wires
Dip the zinc electrode in the 1 M Zinc Sulphate solution and the copper electrode in the 1 M Copper Sulphate solution.
Check the reading shown in the voltmeter.
Draw 10 ml of 1 M ZnSO4 solution and dilute it to form a 0.1 M solution.
Dip the Zn electrode in this solution and note the reading in the voltmeter.
Dilute 1 M ZnSO4 solution to form a 0.01 M solution and do the same. Take the reading.
Do the same with the CuSO4 solution and take the readings.
Use the data of potential values of both the electrodes for different concentrations to plot a graph. You will find the trend of variation of cell potential in Zn Cu cell due to the changes in the concentration of electrolytes.

Follow these steps for calculating the data you will use in the graph paper.

ECell = E0Cell – log

E0Cell = E0 (cathode) – E0 (anode)

F = 96500C

T = 298K

R = 8.314

n = 2 (where n = electrons gained or lost)

By substituting the values, we get

ECell = E0Cell – log

Results of the Experiment

You will find that the ECell will decrease with the increase in the molar concentration of Zn+2 in the electrolyte. The ECell will increase with the increase in the concentration of Cu+2 in the electrolyte.

FAQs on Variation of Cell Potential in Zn Cu Galvanic Cell

1. What is the cell potential of a Zn–Cu cell under standard conditions?

The standard cell potential (E°_cell) of a Zn–Cu cell is +1.10 V under standard conditions.

Standard reduction potentials:
- Zn²⁺ + 2e^- → Zn(s), E° = −0.76 V
- Cu²⁺ + 2e^- → Cu(s), E° = +0.34 V
E°_cell = E°_cathode − E°_anode = 0.34 − (−0.76) = 1.10 V
Standard conditions: 1 M solutions, 1 atm pressure, 25°C (298 K).

This positive value shows the Zn–Cu galvanic cell is spontaneous.

2. How does concentration affect the cell potential of a Zn–Cu cell?

The cell potential of a Zn–Cu cell changes with ion concentration according to the Nernst equation.

For the reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Nernst equation at 25°C: E = E° − (0.0591/2) log([Zn²⁺]/[Cu²⁺])
If [Cu²⁺] increases, E increases.
If [Zn²⁺] increases, E decreases.

Thus, variation in concentration directly alters the cell potential.

3. What is the Nernst equation for the Zn–Cu cell?

The Nernst equation for the Zn–Cu cell at 25°C is E = E° − (0.0591/2) log([Zn²⁺]/[Cu²⁺]).

Overall reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Number of electrons transferred (n) = 2
Reaction quotient Q = [Zn²⁺]/[Cu²⁺]

This equation explains how cell potential varies with concentration in a Daniell cell.

4. Why does the cell potential of a Zn–Cu cell decrease over time?

The cell potential decreases over time because ion concentrations change as the reaction proceeds toward equilibrium.

[Zn²⁺] increases in the anode compartment.
[Cu²⁺] decreases in the cathode compartment.
According to the Nernst equation, this increases Q and lowers E.
When equilibrium is reached, E = 0.

This explains the gradual drop in voltage in a working Zn–Cu galvanic cell.

5. How do you calculate the cell potential of a Zn–Cu cell at non-standard conditions?

The cell potential at non-standard conditions is calculated using the Nernst equation.

Step 1: Write the balanced reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Step 2: Calculate Q = [Zn²⁺]/[Cu²⁺]
Step 3: Substitute into E = 1.10 − (0.0591/2) log Q

Example: If [Zn²⁺] = 0.1 M and [Cu²⁺] = 1.0 M, then E > 1.10 V.

6. What happens to the cell potential when the Zn–Cu cell reaches equilibrium?

When the Zn–Cu cell reaches equilibrium, the cell potential becomes zero (E = 0 V).

At equilibrium, the reaction quotient Q equals the equilibrium constant K.
From thermodynamics: ΔG = −nFE
At equilibrium, ΔG = 0, so E = 0.

This means no net electron flow occurs once equilibrium is established.

7. How does temperature affect the cell potential of a Zn–Cu cell?

Temperature affects cell potential through the temperature-dependent form of the Nernst equation.

General form: E = E° − (RT/nF) ln Q
R = 8.314 J mol^-1 K^-1, F = 96500 C mol^-1
As temperature increases, the value of (RT/nF) changes.

For the Zn–Cu cell, small temperature increases usually cause slight changes in voltage depending on Q.

8. What is the balanced cell reaction in a Zn–Cu galvanic cell?

The balanced overall reaction in a Zn–Cu galvanic (Daniell) cell is Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s).

Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e^-
Cathode (reduction): Cu²⁺(aq) + 2e^- → Cu(s)

Two electrons are transferred, producing electrical energy.

9. What factors influence the variation of cell potential in a Zn–Cu cell?

The variation of cell potential in a Zn–Cu cell depends mainly on concentration, temperature, and reaction progress.

Ion concentration (via the Nernst equation)
Temperature (affects RT/nF term)
Reaction quotient (Q) during operation
Approach to equilibrium

Changes in these factors directly modify the measured voltage of the Daniell cell.

10. Why is the Zn–Cu cell potential positive?

The Zn–Cu cell potential is positive because copper has a higher reduction potential than zinc.

E°(Cu²⁺/Cu) = +0.34 V
E°(Zn²⁺/Zn) = −0.76 V
E°_cell = 0.34 − (−0.76) = +1.10 V

A positive cell potential indicates a spontaneous redox reaction in the galvanic cell.