OBJECTIVES

1. To experimentally determine the equilibrium constant of a redox reaction.

2. To compare this experimental equilibrium constant with the theoretical value obtained using Nernst equation

INTRODUCTION

The equilibrium constant of an electrochemical cell’s redox reaction can be calculated using the Nernst equation and the relationship between standard cell potential and free energy. When a reaction is at equilibrium, the change in free energy is equal to zero and the driving force behind the reaction decreases. The reaction progresses forward and backwards at the same rate meaning there is no net electron flow. With no electron flow, there is no current and the potential is equal to zero.

The change in free energy of an electrochemical cell is related to the cell potential by the equation:

ΔG = -nFEcell

whereΔG is the free energy of the reactionn is the number of moles of electrons exchanged in the reactionF is Faraday's constant (96484.56 C/mol)E is the cell potential.

To find the equilibrium by use Nernst equation where is used to calculate the cell potential of a non-standard cell:

Ecell = E°cell - (RT/nF) x log10Q

whereEcell is the cell potentialE°cell refers to standard cell potentialR is the gas constant (8.3145 J/mol·K)T is the absolute temperaturen is the number of moles of electrons transferred by the cell's reactionF is Faraday's constant (96484.56 C/mol)Q is the reaction quotient (reaction quotient, Q is equilibrium constant, K)

Chemicals: 1 M CuSO4, 1 M ZnSO4, 1 M KNO3

Apparatus: voltmeter, small beakers / U-tube, crocodile clips with wires, filter paper, Cu strips,

Zn strip

PROCEDURE

In this experiment, a Zn-Cu cell was selected for study. Concentration of Zn half-cell was maintained at 1.0 M and the Cu half-cell concentration was varying, a series of emf readings was taken.

1. The metal strips was cleaned by rubbing with sandpaper, then was rinsed in distilled water and dried them.

2. One beaker was filled with 1.0 M Zn2+ solution until it covers the Zn electrode.

3. Proper dilution was done for the second beaker to get the following Cu 2+ solutions: 0.00010 M, 0.0010 M, 0.010 M, 0.10 M.

4. 0.00010 M Cu2+ was placed into second beaker until it covers the Cu electrode. The circuit was connected and the voltage was measured.

5. The Cu2+ solution was replaced with the next higher concentration (0.0010 M) and the voltage was measured.

6. This process was continued until concentration of Cu2+ was 1.0 M. the table below was filled.

cell Zinc half-cell (M)

Copper half-cell (M)

1/[Cu2+] Ln 1/[Cu2+] Emf of the cell

1 1.0 0.0001

2 1.0 0.001

3 1.0 0.01

4 1.0 0.1

5 1.0 1.0

With the above data, plot a graph of emf of the cell (Ecell) against ln 1/[Cu2+] for the series of cells. Extrapolate until Ecell = 0

RESULTS & CALCULATIONS

A. From the graph, obtain the following.

1. Concentration of Cu2+ at equilibrium. Using the value, determine experimental equilibrium constant for this redox reaction.

2. Obtain the standard electrode potential (Eº) for this redox reaction.

B. Using Nernst equation, E = Eº - RT/nF ln K, calculate the theoretical value of equilibrium constant for this redox reaction. Compare this value with experimental result.

cell Zinc half-cell (M)

Copper half-cell (M)

1/[Cu2+] Ln 1/[Cu2+] Emf of the cell Emf calculated

1 1.0 0.0001

2 1.0 0.001

3 1.0 0.01

4 1.0 0.1

5 1.0 1.0