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An experiment measuring the potential difference generated by various simple electrochemical cell

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Introduction

An experiment measuring the potential difference generated by various simple electrochemical cells. Aim: The purpose of this experiment was to correctly construct three simple electrochemical cells and then measure the potential difference, including the polarity, between the two metal electrodes in each cell. Apparatus: ==> Chemicals/ substances: * Copper sulphate solution (1M) * Strip of copper foil * Zinc sulphate solution (1M) * Strip of zinc foil * Silver nitrate solution (1M) * Silver wire * Saturated potassium nitrate solution ==> Additional apparatus: * Safety goggles * Emery paper * 4 beakers * 3 connecting leads with crocodile clips * Filter paper * High resistance voltmeter Method: ==> Clean each metal strip with separate pieces of emery paper (if necessary). ==> Construct the three listed electrochemical cells by the following method: Cell Half- cells (1) Zn2+ (aq) + 2e- ? Zn (s) Cu2+ (aq) + 2e- ? Cu(s) (2) Ag+ (aq) +e- ? Ag(s) Cu2+ (aq) + 2e- ? Cu(s) (3) Ag+ (aq) +e- ? Ag(s) Zn2+ (aq) + 2e- ? Zn (s) 1. Place each of the two metal strips required for the cell in a separate beaker. Hold each strip against the inside of the beaker so that it comes over around 2 cm above the rim and fold this projection down over the rim of the beaker, clamping it into position with a crocodile clip attached to a lead.

Middle

Due to the decreased concentration of the Cu2+ ions in the right hand half cell, and so due to electrons arriving from the zinc half cell, positive K+ ions flow into the copper half cell, to replace "lost" positive charge/ to compensate for the increased negative charge. Discussion: We first measured the potential difference of the typical cell, the Daniel cell: a cell consisting of zinc in a solution of its ions and copper also in a solution of its ions and found the zinc electrode to be negative and the copper electrode to be positive so that the measured potential would read a positive polarity (and so indicating that the reactions were feasible). This can easily be explained in terms of the standard electrode potentials of the two metals and their relative positions in the electrochemical series. When the standard electrode potential of zinc metal is measured by coupling a zinc half-cell with the standard hydrogen half-cell (which is given a value of 0) under standard conditions, the measured electrode potential is negative. This negative value indicates that the equilibrium of the zinc electrode, Zn2+ (aq) + 2e- ? Zn (s), lies more to the left compared to the hydrogen equilibrium (2H+ + 2e- ?

Conclusion

If I were to do this experiment again I would perhaps try a different voltmeter with a higher resistance or use a different chemical for the salt bridge, such as potassium chloride, or construct the salt bridge differently by, for example, using a glass tube and then I would see if any of these changes made my results more accurate. Conclusion: From measuring the potential differences of the three simple electrochemical cells, I can conclude various main points which I will continue to discuss in my evaluation to follow: At the negative electrode an oxidation reaction is occurring and electrons are being released to reduce the metal at the positive electrode. At the positive electrode a reduction reaction is occurring as electrons from the negative electrode arrive to reduce the metal. All the cells had a positive measured potential difference, indicating their reactions were spontaneous/ feasible. The largest measured potential difference was the one between silver and zinc and the smallest was the one between silver and copper. A salt bridge is used to complete the circuit and keep the half cells neutral. This consists of filter paper soaked in a substance that doesn't react with the solutions in the half cells, such as, in this case potassium nitrate solution.

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