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This is a labreport on chemical kinetics... From this labreport students of chemical engineering and chemistry students can get a huge advantage in working in their lab..Full description
Elizabeth Gardner Mrs. Shafer AP Chemistry Pd. 3-4 28 March 2011 Electrochemical Cells Objective: The purpose of this lab is to Data: Part 1 Voltage of each half-cell versus the zinc electrode
Zn versus Ag Zn versus Cu Zn versus Fe Zn versus Mg Zn versus Pb
Voltage 1.31V .89V .53V .42V .42V
Anode Zn Zn Zn Mg Zn
Cathode Ag Cu Fe Zn Pb
Predicted and Measured Cell Potentials Anode
Cathode Equation for the Cell Reaction
Mg Fe Fe Mg Pb Cu
Cu Cu Ag Pb Cu Ag
Mg + Cu2+ Mg2+ + Cu 2Fe + 3Cu2+ 2Fe3+ + 3Cu Fe + 3Ag+ Fe3+ + 3Ag Mg + Pb2+ Mg2+ +Pb Pb + Cu2+ Pb2+ + Cu Cu + 2Ag+ Cu2+ + 2Ag
Predicted Potential from Experimental Data (V) 1.31 .36 .78 .84 .47 .42
Measured Potential (V) 1.36 .42 .61 .45 .52 .32
The predicted potentials for this chart were derived from the voltages of the half cells. For example, when the redox reaction between Mg and Cu was balanced, the reduction potential and the oxidation potential of the ions were added:
Part 2
Zn(s)|Zn2+(1.0M)||Cu2+(.0010 M)|Cu(s)
Equation for Cell Reaction Zn + Cu2+ Zn2+ + Cu
Voltage .80V
Predicted Potential .80V
Anode Zn
Cathode Cu
Measured Potential .80V
The potential for the zinc and copper reaction was predicted by using the Nernst equation:
Part 3
Zn(s)|Zn2+(1.0M)||Ag+(unknown M)|Ag(s)
Equation for Cell Reaction Zn + 2Ag+ Zn2+ + 2Ag
Calculated [Ag+] 1.12 10-9
Voltage .78V
Anode Zn
Cathode Ag
Calculated KspAgCl
Reported KspAgCl
1.12 10-9
1.8
10-10
The concentration of Ag+ ions was also found through the Nernst equation:
The Ksp is the solubility constant of a precipitate. Because the products are solid, the denominator is 1 and thus excluded.
Calculations: Part 1 Reduction Equation
Electrode Potential using Zinc as the Standard, E Zn
Ag++ e- Ag Cu2+ + 2e- Cu Fe3+ + 3e- Fe Pb2+ + 2e- Pb Mg2+ + 2e- Mg
1.31V .89V .53V .42V -.42V
Accepted Electrode Potential using Hydrogen as Standard, E .799V .377 -.04 -.126 -2.37
E
Zn-
E
.511V .51V .57V .55V 1.95V
Post-Lab Questions: 1. An electrode potential is created by comparing the potential of a metal, such as zinc in this lab, and potentials under standard conditions, as in published tables. 2. Yes, the ranking was consistent between the reduction equations and the published chart of E values. 3. The values found by using the zinc electrode should be higher than the values from the hydrogen electrode because the standard reduction potential of zinc is lower than that of hydrogen. Thus, when zinc becomes the standard, the voltages shift upward in response to the change. 4. Answer 5. A negative value for a standard potential indicates that oxidation occurs rather than reduction. For example, the electron potential using zinc as the standard of Mg was negative because when it reacted with Zn, Mg was the anode and Zn was the cathode. 6. The change in concentration of the copper ions in part 2 affected the cell potential by making it lower. If the copper ions had not been diluted, the molarity would be 1.0. Log(1)=0, so the value of E would be .89V instead of .80V. Le Chatlier’s principle would predict that because the concentration of Cu2+ ions would be greater, the reaction would shift to the reactants side, yielding more zinc. 7. The solubility product of AgCl was determined by the equation:
