Cell Potential Variation in Zn - Cu Daniell 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 Cell
1. What is the standard cell potential (E⁰cell) of a Zn-Cu cell and what does it represent?
The standard cell potential (E⁰cell) of a Zn-Cu Daniell cell is +1.1 V. This value represents the maximum potential difference between the zinc and copper electrodes under standard conditions, which are:
- Concentration of both Zn²⁺(aq) and Cu²⁺(aq) is 1 Molar (1 mol/L).
- The temperature is 298 K (25°C).
- The pressure is 1 atm for any gaseous reactants or products.
A positive E⁰cell indicates that the reaction is spontaneous under these specific conditions.
2. How does the cell potential of a Zn-Cu cell change with the concentration of its electrolytes?
The cell potential (Ecell) of a Zn-Cu cell varies according to the principles described by the Nernst equation. The general trend is:
- Increasing the concentration of Cu²⁺ ions (the reactant ion at the cathode) will increase the cell potential.
- Increasing the concentration of Zn²⁺ ions (the product ion from the anode) will decrease the cell potential.
This is because the reaction equilibrium shifts to counteract the change in concentration, affecting the overall voltage produced.
3. What is the overall cell reaction and the half-cell reactions in a Daniell cell?
In a Daniell cell, the following reactions occur at the electrodes:
- Anode (Oxidation): At the zinc electrode, zinc metal loses electrons.
Zn(s) → Zn²⁺(aq) + 2e⁻ - Cathode (Reduction): At the copper electrode, copper ions gain electrons.
Cu²⁺(aq) + 2e⁻ → Cu(s)
The overall cell reaction is the sum of these two half-reactions:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
4. What is the function of the salt bridge in an electrochemical cell like the Zn-Cu cell?
The salt bridge serves two critical functions in a Zn-Cu cell:
- It completes the electrical circuit by allowing the migration of ions between the two half-cells, which prevents the build-up of charge.
- It maintains electrical neutrality in each half-cell. Anions from the salt bridge flow to the anode beaker to balance the positive charge of newly formed Zn²⁺ ions, while cations flow to the cathode beaker to replace the positive charge of Cu²⁺ ions being consumed.
5. Why does the cell potential decrease if the concentration of Zn²⁺ ions is increased?
According to Le Chatelier's principle, increasing the concentration of a product—in this case, Zn²⁺ ions—shifts the reaction equilibrium to the left (the reactant side). This opposes the natural forward direction of the cell's spontaneous reaction (Zn + Cu²⁺ → Zn²⁺ + Cu). This opposition reduces the overall driving force for electron flow, resulting in a lower cell potential (Ecell). The Nernst equation mathematically confirms this relationship.
6. How does the Nernst equation mathematically explain the variation of cell potential in a Zn-Cu cell?
The Nernst equation provides a precise way to calculate the cell potential (Ecell) under non-standard conditions. For the Zn-Cu cell, the equation is:
Ecell = E⁰cell - (RT/nF) ln([Zn²⁺]/[Cu²⁺])
Here:
- E⁰cell is the constant standard potential (1.1 V).
- The term [Zn²⁺]/[Cu²⁺] is the reaction quotient (Q).
The equation shows that if the ratio Q increases (by increasing [Zn²⁺] or decreasing [Cu²⁺]), the value being subtracted from E⁰cell becomes larger, thus lowering the Ecell. Conversely, if Q decreases, the Ecell increases.
7. What would happen to the cell potential if a student accidentally used a 0.5 M CuSO₄ solution but a 1.0 M ZnSO₄ solution?
In this scenario, the concentration of the reactant ion (Cu²⁺) is lower than standard, and the concentration of the product ion (Zn²⁺) is at the standard level. According to the Nernst equation, decreasing the reactant concentration and/or increasing the product concentration will make the forward reaction less favourable. Therefore, the resulting cell potential (Ecell) would be less than the standard 1.1 V.
8. Can a Zn-Cu cell ever have a cell potential of zero? If so, under what conditions?
Yes, a Zn-Cu cell's potential becomes zero when the cell reaches electrochemical equilibrium. At this point, the forward and reverse reaction rates are equal, and there is no longer a net flow of electrons to generate a voltage. This occurs when the ratio of ion concentrations ([Zn²⁺]/[Cu²⁺]) becomes so large that it equals the equilibrium constant (Kc) for the reaction. At this stage, the cell is considered 'dead' or fully discharged.
