Electrochemical Cells Laboratory #15 Henry Ko AP Chemistry Dulaney High School March 12th, 2009
Abstract: In this experiment, a standard table of reduction potentials of a series of metal ions is constructed using copper, copper, iron, lead, magnesium, magnesium, silver, silver, and zinc. zinc. These half cells are are connected connected by a salt bridge bridge and all potential potentialss are measured measured with respect to the zinc electrode. Also, the measured voltage voltage of a nonstandar nonstandard d copper copper cell cell is calcul calculate ated d throug through h the Nernst Nernst equati equation. on. The solubi solubili lity ty product product constan constantt of AgCl AgCl is also also determine determined d through the Nernst equation. equation. The K sp 33 × × 10 10 11 , sp value for AgCl was determined to be 7 .33 yielding a percent error of 59.3%. The voltage for the cell reaction was experimentally determined to be 0.81 V. −
Theory: electrochemical cell is produced when a redox An electrochemical redox reaction occurs. The resulting resulting electron electron transfer between between the reaction reaction runs through an external external wire. Because Because the oxidati oxidation on and reduction reduction reactions reactions are half-cell reactions. A half cell physically separated from each other, these are called half-cell cell is prepare prepared d from contact with the metal with its solution of ions. Each element’s unique electron configuration allows each to develop a different electrical potential. The standard reduction potential is the voltage that a half-cell, under standard conditions (1 M, 1 atm, and 25 C), develops when combined with the standard hydrogen electrode, that is arbitrarily assigned to a potential of zero volts. A positive E cell value indicates that the redox reaction in that particular cell is spontaneous. Calculatio Calculations ns of nonstandar nonstandard d potential potentialss can be made using the Nernst Equation: ◦
◦
E =
E − ◦
RT ln(Q) nF
(1)
where E is the measured cell potential, E is the standard cell potential, R is the gas constant (8.314 J/mol · K), T , is the temperature in K, n is the number of moles of electrons transferred as shown by the redox reaction, and F is the Faraday constant (9 .65 65 × × 10 104 C/mol). ◦
At STP, the Nernst equation can be simplified to
E =
E − ◦
0 .0592 log(Q) n
(2)
Procedure: A wellplate is set up such that the first row contains approximately 2 mL of 1.0 M Zn(NO 3 )2 solution in each well. well. In the second row, two two mL of Cu(NO3 )2 , AgNO3 , Fe(NO3 )3 , Mg(NO3 )2 , and Pb(NO3 )2 are added in their respective wells. A salt bridge is used to connect to the two adjacent wells (made from filter paper soaked in KNO 3 solution. solution. A voltmet voltmeter er is used to measure the potential potential difference difference for each of the 5 half cells. After, After, measure measure the potential potential difference difference between at least six com combinat binations ions of various arious electrodes electrodes.. Again, use the voltmeter to measure the potential difference. In part 2, Cu(NO 3 )2 is diluted to 0.0010 M. It is added onto a wellplate and measured against the standard zinc half-cell. In part 3, 10 mL of 1.0 M NaCl solution is mixed with one drop of 1.0 M AgNO 3 . After After precipita precipitation tion occurs, some of the solution is poured into the well plate and measured against the standard zinc half-cell.
1
Data Analysis: Part 1 Voltage of Each Half-Cell versus the Zinc Electrode
Volta oltage ge (V) (V) 1.41 0.99 0.553 0.58 0.454
Zn versus Ag Zn versus Cu Zn versus Fe Zn versus Mg Zn versus Pb
Anode Anode Zn Zn Zn Mg Zn
Cath Cathode ode Ag Cu Fe Zn Pb
Predicted and Measured Cell Potentials
Anod Anodee Mg Fe Fe Mg Pb Cu
Cath Cathod odee Cu Cu Ag Pb Cu Ag
Equa Equati tion on for for Cell Cell Reac Reacti tion on 2+ Cu + Mg − Mg −−→ Cu + Mg 2+ 3 Cu 2+ + 2 Fe −−→ 3 Cu + 2 Fe 3+ 3 Ag+ + Fe Fe −−→ 3 Ag + Fe Fe 3+ Pb 2+ + Mg − Mg −−→ Pb + Mg 2+ Cu 2+ + Pb − Pb −−→ Cu + Pb 2+ 2 Ag+ + Cu − Cu −−→ 2 Ag + Cu 2+
Pred Predic icte ted d Poten otenti tial al 0.99 + 0.58 = 1.57 0.99 + (-0.553) = 0.437 1.41 + (-0.553) = 0.857 0.58 + 0.454 = 1.034 0.99 + (-0.454) = 0.536 1.41 + (-0.99) = 0.42
Meas Me asur ured ed Poten otenti tial al 1.55 0.47 0.86 1.0 0.5 0.41
Part 2
Zn| Zn|Zn
2+
Cu
2+
Volta oltage ge 1V
|Cu
Equation Equation for Cell Reaction Reaction 2+ Zn + Cu −−→ Zn 2+ + Cu
Anod Anodee Zinc
Predicted Predicted Potent Potential ial 0.9012 V
Cath Cathod odee Copper Measured Measured Potenti Potential al 1V
Part 3
Zn| Zn|Zn Equati Equation on for Cell Cell Reac Reactio tion n + 2+ Zn + 2Ag −−→ Zn + 2 Ag Ag
2+
+
Ag |Ag
Volta oltage ge 0.81 V
Calcul Calculate ated d [Ag [Ag+ ] 7.33 ×10 × 10 11 −
2
Anod Anodee Zinc
Cath Cathod odee Silver
Calc Calcul ulat ated ed K sp AgCl Cl sp Ag 11 7.33 × 7.33 ×10 10 −
Repor Reporte ted d K sp sp AgCl 1.8 × 1.8 ×10 10 10 −
Calculations: Part 1
1. Write reduction equations for each metal ion, arranging arranging the equations in decreasing order of measured potential potential in the table below. Include Include zinc in the table, using 0.00 volts volts as the potential potential of the Zn — Zn 2+ half-cell. Record the accepted standard potentials using the hydrogen electrode as standard, and calculate the difference between the two standard values. Reduct Reduction ion Equati Equation on Standa Standard rd Zinc, Zinc, E Zn Standard Hydrogen, E E Zn − E + – Ag + e −−→ Ag 1.41 0.80 0.61 2+ – Cu + 2 e −−→ Cu 0.99 0.34 0.65 2+ – Mg + 2 e −−→ Mg 0.58 -1.18 1.76 2+ – Fe + 2 e −−→ Fe 0.55 -0.44 0.99 2+ – Pb + 2 e −−→ Pb 0.45 -0.13 0.58 2+ – Zn + 2 e −−→ Zn 0 -0.76 0.76 ◦
◦
◦
◦
2. Use the electrode potentials from the above table to predict the voltages voltages of the six-half cell combinations combinations selected selected in Part 1, step 10. Record Record thie value and which metal is the cathode and which which is the anode in the Data Table above. Compare the predicted and measured potentials. The predicted and measured potentials were very similar to one another. Part 2
Write a balanced net ionic equation for the reaction occurring in the cell in Part 2. Record this equation in the Part 2 Data Table. Use the Nernst equation to calculate what the expected voltage should be. Record this value in the Part 2 Data Table. Compare this value to the measured voltage.
0 .0592 99 − − log E = 0.99 2 = 0.9017
1 0.001
Part 3
1. Write the balanced balanced net ionic equation equation for the reaction reaction occurring the cell. Use the Nernst equation equation to calculate the concentration of the Ag + ion. Record this value in the Part 3 Data Table. 0 .0592 0.81 = 1.41 41 − − log 2 =
7.33 33 × × 10 10
1 [Ag+ ]2
11
−
2. Calculate Calculate the value value of the solubility solubility product AgCl. Compare Compare the calculated calculated value value to a reported reported value. value. Record this value in the Part 3 Data Table.
= (7.33 33 × × 10 10
K sp sp
= 7.33 33 × × 10 10 11 − 1.8 × 10 × 10 10 1.8 × 10 × 10 −
Percent Error :
−
7.33 33 × × 10 10
10
−
= 59.3%
3
11
−
11
−
)(1.0)
Conclusion: In this experiment, a standard table of reduction potentials of a series of metal ions is constructed using copper, copper, iron, lead, magnesium, magnesium, silver, silver, and zinc. zinc. These half cells are are connected connected by a salt bridge bridge and all potential potentialss are measured measured with respect to the zinc electrode. Also, the measured voltage voltage of a nonstandar nonstandard d copper copper cell cell is calcul calculate ated d throug through h the Nernst Nernst equati equation. on. The solubi solubili lity ty product product constan constantt of AgCl AgCl is also also determine determined d through the Nernst equation. equation. The K sp 33 × × 10 10 11 , sp value for AgCl was determined to be 7 .33 yielding a percent error of 59.3%. The voltage for the cell reaction was experimentally determined to be 0.81 V. Contamination was a huge source of error in this experiment. Because the cells were based in wellplates, any spillage or mixture of two cells could result in a mixing of the concentration of different solutions. This would have produced a different redox reaction than what was intended, and which would have yielded a different different voltage voltage with respect respect to the zinc electrode. electrode. This would have have severely severely impacted impacted calculations calculations concerning cerning the predicted predicted and measured measured cell potential potentials, s, which are used in other parts of the lab. Concentr Concentratio ation n contaminations could impact the cell reaction because a different amount of electrons would have been transferred ferred via the salt bridge, bridge, creaing creaing a different different voltage. voltage. Another Another error that was a factor factor was the voltmeter. voltmeter. Because the voltmeter has a reading that is constantly fluxuating, there is no clear moment in which the measurement should have been taken. Therefore, the measurement for each cell varied with respect to their individual times. This could cause some errors in the experimentally determined voltage, which could have impacted the future calculations. Overall Overall this experiment experiment is neither neither accurate accurate nor precise. precise. Only one set of trials was used, meaning that there is no basis for comparison comparison and the result could have varied varied widely. widely. Also the precision and accuracy depended depended heavily on the voltmet voltmeter er used. Human judgemen judgementt in the reading, reading, as well well as the accuracy accuracy of the voltmet voltmeter, er, makes the reading reading variabl variable. e. This experiment experiment could have have been improv improved ed in equipmen equipments. ts. A more accurate/precise voltmeter could have been used, as well as cleaner and more accurate pipets (in transferring solutions) could have been used to for better results. −
Pre-Lab Questions 1. Which Which ion is most easily reduced? reduced? Copper 2. Which Which metal is most easily easily oxidized? oxidized? Aluminum 3. The copper and aluminu aluminum m electrodes are connected connected to form a battery battery.. a. Which is the anode?
Aluminum b. Which is oxidized?
Aluminum c. What will be the battery voltage?
0.62 0.62 + 1.3 1.388 = 2 V d. Write a balanced net ionic equation for the reaction that takes place.
3 Cu 2+ + 2Al − 2Al −−→ 2 Al 3+ + 3Cu 4. A soltuion is prepared prepared in which which trace or small amounts amounts of Fe 2+ is added to a much larger amount of – solution in which the [OH ] is 1 .0 × 10 2 M. Some Fe(OH) 2 precipitates. The value of K sp sp for Fe(OH) 2 is 8.0 × 10 × 10 10 . −
−
a. Assuming that the hydroxide ion concentration is 1 .0 × 10
ions in the solution. 8.0 × 10 × 10 10 [Fe [Fe 2+ ] = = 8.0 × 10 × 10 1.0 × 10 × 10 2 −
−
8
−
M 4
2
−
M, calculate the concentration of Fe 2+
b. A battery is prepared using the above solution with an iron wire dipping into it as one half-cell.
The other half-cell is the standard nickel electrode. Write the balanced net ionic eqaution for the cell reaction. Ni 2+ + Fe Fe −−→ Ni + Fe 2+ c. Use the Nernst equation to calculate the potential of the above cell.
0 .0592 15 − − log E = 0.15 2 =
8
8 × 10 × 10 −
1 × 10 × 10
2
−
0.3 V
Post-Lab Questions 1. What is an electrode potential potential?? An electrode potential is the difference between the charge on an electrode and the charge in the solution. solution. The electrode potential potential depends on the concentrat concentration ion of substances substances and the temperatur temperature, e, according to the Nernst equation. 2. Did the ranking of reduction reduction equations equations agree with that in the published published chart of E values? ◦
The ranking of the reduction equations should agree with that in a published chart of E values. However, the obtained “standard” table of reduction potentials did not exactly compare with the accepted table. The order of the metals in decreasing reduction potentials was silver, copper, magnesium, iron, lead, and zinc. For the expected values, the order was copper, lead, iron, zinc and magnesium. ◦
3. How should the values values found using the zinc electrode electrode as a standard compare with those in the E table that are based on the standard hydrogen electrode? Did they? ◦
The values found using the zinc electrode as a standard should be greater than the values in the E table table by approxima approximately tely 0.76 V, which which are based on the standard standard hydrogen hydrogen electrode. electrode. This is because because the zinc electrode was used as the standard, with the voltage preset to 0. However, when the hydrogen electrode is used as the standard, the voltage for the zinc electrode is not zero; in fact it is -0.76 V.
◦
4. What factors factors can cause a difference difference between between experimental experimental and reported reported values? values? See Conclusion . 5. What does a negative negative value value for a standard standard potential potential indicate? indicate? A negative value for standard potential indicates that the cell is not galvanic, meaning that the oxidation reaction is more likely to occur than the reduction. 6. Ho How w did the change change in concen concentra tratio tion n of the copper copper ions in Part Part 2 affe affect ct the cell cell potent potential ial?? Is this change in agreement (qualitatively) with that which would be predicted by LeChatelier’s Principle? Did the calculated and measured values agree? The electrode potential depends on the concentration of substances and the temperature (according to the Nernst equation). equation). In this case, the concentrati concentration on of copper was reduced reduced from 1.0 M to 0.001 M. For one molar solution, the reduction potential was 0.99 V, whereas the 0.001 M solution had a potential of 0.81 V. This is in agreement with LeChatelier’s Principle because when the concentration of copper was reduced, the amount of volts generated by the reaction was also reduced. The calculated and measured potential values differed by approximately 0.1 volts. 7. Explain how the AgCl solubility solubility product was determined. The AgCl solubility product was determined using the Nernst equation in conjunction with the equation for K sp unknown in this case, was the concentrati concentration on of Ag + . Therefore Therefore,, with the standard and sp . The unknown nonstandard measures of potential known through the experiment, the concentration of of Ag + was found. found. Because Because the amount amount of Cl – was known to be approximately 1, the K sp sp was also equal to the + concentration of Ag .
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