ChemLabSpring07_2_Electrochemistry_Electrolytic_Cells

ChemLabSpring07_2_Electrochemistry_Electrolytic_Cells -...

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2. Experimental Faraday Values Trial Anode/Cathode Faraday Value (C/mol) 1 Anode 88,900 1 Cathode 90,100 2 Anode 95,100 2 Cathode 93,300 3. Mean Experimental Faraday Values and Deviation from Literature Value Mean Faraday (C/mol) Uncertainty (C/mol) Deviation from Literature Value Anode 9.20 x 10^4 +/- 4,000 4.65 % Cathode 91,700 +/- 2,000 4.96% 4. Error Analysis- A systematic error made in the experiment is that we assumed that all the mass that the cathode gained was copper ions plating onto it, which is not true because of the inclusion of water molecules or impurities in the solution which may have attached themselves to the cathode. This means that the cathode would have gained less mass in copper than we measured causing our experimental Faraday values to be lower than actual. Another systematic error was that we took our last current reading and then
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Unformatted text preview: disconnected the circuit leaving a little time for more current to flow through. This would have resulted in a lower Faraday value than actual because the current would not be accounted for while the mass that moved would be. If the stirring bar was not on high enough, the solution could have become supersaturated leading to a layer of insulation could have formed on the electrodes causing the Faraday value to be higher that actual. There was some instrumental errors that are normal as far as equipment, such as ammeters only being able to be read to a certain level of accuracy and the analytical balance only being able to weigh masses accurately to +/- .0001grams....
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This note was uploaded on 04/09/2008 for the course CHEMISTRY 030.105 taught by Professor Pasternick during the Fall '06 term at Johns Hopkins.

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