With the five concentrations of AgNO3(aq), five electrochemical cells were constructed using 1.0 molar CuSO4(aq) as the other electrolyte. For the first electrochemical cell, approximately 80cm³ of the 0.1 molar solution of AgNO3(aq) was poured into a 100ml Pyrex beaker to be used as the silver half-cell. The same volume of CuSO4(aq) was used in another beaker to construct the other half-cell.
The copper and silver metal electrodes were glass papered. The silver electrode was crocodile clipped so it was partly submerged in the silver solution and the copper electrode in the copper solution. The electrodes were connected to a high resistance voltmeter. A strip of filter paper was soaked in saturated KNO3(aq) and it was used to connect the two solutions. The Ecell reading on the voltmeter was recorded.
Four more electrochemical cells were constructed using the same concentration of the copper electrolyte but the remaining four concentrations of the silver electrolyte.
Results Table:
NOTE: The EAg was calculated by rearranging the equation:
Ecell = Erhs – Elhs
- where the Elhs = ECu (+0.34v)
Conclusion:
The Nurst equation:
Where:
Z = number of electrons transferred.
F = faraday constant
R = gas constant
T = temperature
The Nurst equation is used because it is in the form of the equation:
y = mx + c
Where:
Y (axis) = EAg
M (gradient constant) = RT/ZF
X (axis) = Ln [Ag⁺]
C (intercept) = EAg
When plotted on the graph, the results give a straight line according to this equation. The Ln[Ag⁺] is proportional to the electrode potential of the silver electrode. This is because in the Nurst equation, RT/ZF and EAg are constant. Therefore, EAg must be proportional to EAg standard.
The intercept, (c) is the EAg according to the Nurst equation. It should therefore intercept at 0.8 volts. However, the intercept is at 0.845 volts. This is an inaccuracy of 5.625% (0.045/0.8 x 100). The inaccuracies responsible are discussed in the section following. It is worth noting here that the intercept is higher than expected because the EAg readings are too high. A possible explanation for this is that the concentration of the 0.1 molar AgNO3(aq) may be higher than assumed. This is possible since the solution is only accurate to one decimal place.
Also of importance is the fact that the anomalous result of 0.00001 molar AgNO3(aq) reading has been ignored since the other points line up perfectly. Had it been included, a more accurate intercept would have occurred.
Evaluation:
Limitations:
The method was limited by a number of minor inaccuracies. Firstly, the 10cm³ pipette and the volumetric flask were accurate to + or – 0.1cm³. This is a 1% error margin. The voltmeter was accurate to + or – 0.001v which is negligible on its own. The 1.0 molar solution of CuSO4(aq) and the original 0.1M solution of AgNO3(aq) were only accurate to 1 decimal place. Also, the temperature and air pressure may not have been standard.
When the metal electrodes were glass papered, the same piece of glass paper was used for both types of metal. This could have transferred particles of metals onto their opposing electrode. This would have affected their conductivity. Also, the length of time over which the electrodes were connected in the cell may have had an affect on the concentrations. The longer this time was, the more likely it was for a change to have occurred. This length of time was not kept constant.
Finally, the most dilute concentrations of the AgNO3(aq) were so dilute that they were inaccurate. This is because the conductivities of the half-cells in question were approaching that of water. This may well be an explanation for the anomalous result of 0.00001M.
Improvements:
Overall, the method was not very inaccurate, (only 5.625%). Most of the inaccuracies caused by the equipment were practically negligible. When the inaccuracy percentage is of these are added, they are around 3%, so only a small percentage was due to human error.
A possible improvement would be to use a standard hydrogen half-cell or a Calomel half-cell instead of the copper half-cell, (ensuring that the conditions are standard). This would avoid using the equation:
Ecell = Erhs – Elhs
Another advantage is that the platinum electrode is inert. It is far less likely to be chemically corrupted than the copper electrode.
The inaccuracy caused by having solutions, which are too diluted, could be overcome simply by using increasing concentrations of AgNO3(aq), rather than decreasing amounts . For example, the concentration could start at 1.1M and increase from there. The graph could then be extrapolated back to intercept the y-axis. This was not possible in this experiment. However, a control could have been used to ensure that the conductivity of water had not been reached. To do this, a silver half-cell could be set up, but using pure water as the electrolyte. When connected to the copper half-cell, the Ecell reading of this could be compared to those of the lower concentrations of AgNO3(aq), to see if there was still a formidable difference.
