Science / Electrochemistry: Electrode Potential
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Autor: anton 09 March 2011
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Chemical reactions can be used to produce electricity and electricity can be used to cause chemical reactions through oxidation-reduction reactions. The first part of the experiment measures the standard electrode potentials of five various half cells against the Cu2+(1M)|Cu half cell. The last three half cells are prepared through electrolysis. The electrode potentials of all reactions are positive which means that they are spontaneous. There is a significant percent difference from the theoretical and the experimental standard electrode potentials. The sources of error are incorrect solution preparation and contaminated materials. The second part of the experiment uses standard electrode potentials to predict whether or not a reaction will occur and eventually test that prediction. All the standard electrode potentials are positive which means that they are spontaneous, as directly observed. There is also a percent difference from the theoretical and experimental cell potentials. The sources of error are also incorrect solution preparations, contaminated glasswares and materials, and the resistance of the electrodes.
Keywords: cell potential, half-cell, electrode potential, oxidation-reduction reactions, standard electrode potential
Chemical reactions can be used to produce electricity and electricity can be used to cause chemical reactions. The practical applications of electrochemistry are countless, ranging from batteries and fuel cells as electric power sources, to the manufacture of key chemicals, to the refining of metals, and to the methods of controlling corrosion.
The objective of this experiment is to measure the standard electrode potentials of five various half cells against the Cu2+(1M)|Cu half cell and to use the standard electrode potentials to predict whether or not a reaction will occur. The standard electrode potential, EÐ’Ñ”cell, is the electric potential that develops on an electrode when the oxidized and reduced forms of some substance are in their standard states. When used in electrochemical studies, a strip of metal is called the electrode. An electrode immersed in a solution containing ions of the same metal is called a half-cell. A salt bridge is used to join two half-cells in an electrochemical cell. It salt bridge permits the flow of ions between two half-cells.
The first part of experiment measures the standard electrode potentials of five various half-cells against the Cu2+(1M)|Cu half cell. The last three half cells are prepared through electrolysis. Electrolysis is the decomposition of a substance, either in the molten state or in an electrolyte solution, by means of an electric current.
The second part of the experiment uses the standard electrode potential to predict whether or not a reaction will occur. A reaction is said to occur if the standard electrode potential is greater than 0. That reaction is called a spontaneous reaction. If the standard electrode potential is less than 0, the reaction is said not to occur. That reaction is a called non-spontaneous reaction.
The Cu2+ (0.1 M)|Cu half cell is set-up by immersing a copper electrode in 0.10 M CuSO4. The 0.10 M CuSO4 was prepared by diluting 10 mL of 1.0 M CuSO4 in a 100 mL volumetric flask. The Zn2+ (0.1 M)|Zn half cell is set-up by immersing a zinc electrode in 0.10 M ZnSO4. The 0.10 M ZnSO4 is prepared by the laboratory. The Fe2+ (0.5 M), Fe3+ (0.5 M)|C half-cell is prepared by immersing a graphite electrode in a solution prepared by mixing equal volumes of 0.1 M FeSO4 and 0.1 M FeCl3. The graphite electrode is obtained from a pencil.
The Cu2+|Cu half cell is connected to the other two half-cells in separate set-ups using a salt bridge between the two half cells and a voltmeter to measure the emf. The salt bridge is prepared by adding a salt solution in a fabricated glass U-tube. The two ends of the U-tube are covered with cotton balls. The voltmeter reading is recorded.
The Cl (1M), Cl2|C, Br (1M), Br2|C and the I (1M), I2|C are prepared through electrolysis. The halogen X2 is generated by electrolyzing for about a minute a 1 M solution of potassium halide KX between graphite electrodes using a current from two 1.5 V dry cells connected in series.
The X-, X2|C half cells is connected to the Cu2+|Cu half-cell. The meter is in place before generating the halogen in order not to disturb the X, X2|C half-cell after its preparation. The dry cells are disconnected before connecting the two half-cells.
The second part of the experiment is the applications of the electrochemical cells. Any voltaic cell in which the net cell reaction involves only a change in the concentration of some species is called a concentration cell. A concentration cell consists of two half-cells with identical electrodes but different ion concentrations. The concentration cell is set up by this cell notation:
Cu(s)|Cu2+(aq)(0.01 M)||Cu2+(aq)(0.1 M)|Cu(s).
The second set-up is a redox reaction involving complexes. The electrochemical cell is set up by this cell notation:
Zn(s)|Zn2+(aq)(0.1 M)||[Cu(NH3)4(0.33M), NH3(aq)(1.0 M)|Cu(s).
The [Cu(NH3)4] solution is prepared by mixing 10 mL of 0.1 M Cu(SO4)2 with 20.0 mL 1.0 M NH3.
The final set-up is reaction of partially soluble solid in the electrochemical cell. The electrochemical cell is set up by this cell notation:
The Cu(OH)2 is prepared by mixing 10 mL of 0.10 M Cu(SO4)2 with 30.0 mL of 1.0 M NaOH.
Results and Discussion
As shown in Figure 1 is a simple voltaic cell. A voltaic cell is an electrochemical cell in which a spontaneous chemical reaction produces electricity. The anode is the electrode at which the oxidation occurs while the cathode is the electrode at which reduction occurs. The electrons are joined by a metal wire to permit the flow of electrons from the anode to the cathode. The flow of electric current between the solutions is in the form of migration of ions. The ends of the salt bridge are plugged with a porous material that allows ions to migrate but prevents the bulk flow of fluid.
Figure 1. A simple voltaic (galvanic) cell
As shown in Table 1 are the voltmeter readings of the electrochemical reactions. All the electrochemical reactions are voltaic cells which are shown in Figure 1. The first part of experiment is from set-up 1 to set-up 5 while the second part of the experiment is from set-up 6 to set-up 8. All the reactions have a positive voltmeter which means that they are all spontaneous. The 1st reaction has the greatest voltmeter reading while the 6th reaction has the least voltmeter reading.
Table 1. Voltmeter readings of the electrochemical reactions
Balanced Chemical Reaction Voltmeter Reading (V)
Cu2+ + Zn Ð¿Ñ“Â Cu + Zn2+
Cu + 2Fe3+ Ð¿Ñ“Â 2Fe2+ + Cu2+
Cu + Cl2 Ð¿Ñ“Â 2Cl- + Cu2+
Cu + Br2 Ð¿Ñ“Â 2Br- + Cu2+
Cu + I2 Ð¿Ñ“Â 2I- + Cu2+
Cu2+ + Cu Ð¿Ñ“Â Cu + Cu2+
[Cu(NH3)4]2+ + Zn Ð¿Ñ“Â Cu + Zn2+ + 4NH3
Cu(OH)2 + Zn Ð¿Ñ“Â Zn2+ + Cu + 2OH- 0.852
The voltmeter reading is the cell potential of the electrochemical reactions. The cell potential, EcellÐ’Â¬, is the potential difference or the voltage between two electrodes of an electrochemical cell. This experiment is done in nonstandard conditions. The experimental measurement of cell potential has great significance. The cell potential, Ecell and the standard cell potential can be related by the Nernst equation,
Ð³Ð‚â€“" E" Ð³Ð‚â€”_"cell" "=" "E" _"cell" ^"Ð’Ñ”" "-" "0.0592" /"n" "log" Ð²ÐƒÐŽ"Q" (1)
where Ecell is the cell potential, EÐ’Ñ”cell is the standard cell potential, n is the number of moles of electrons transferred between electrodes, and Q is the reaction quotient. From the Nernst equation, the standard cell potential can be calculated. EÐ’Ñ”cellÐ’Â¬ is given by,
"E" _"cell" ^"Ð’Ñ”" "=" "E" _"cell" "+" "0.0592" /"n" "log" Ð²ÐƒÐŽ"Q" (2)
As shown in Table 2 is the EÐ’Ñ”cell of the five reactions. It is calculated from equation 2 which is derived from the Nernst Equation.
Table 2. EÐ’Â°cell
The standard cell potential is used to calculate the experimental standard electrode potential of the five half cells. By international agreement, the standard electrode potential, EÐ’Ñ”cell measures the tendency for a reduction process to occur at an electrode. Another equation for EÐ’Ñ”cell is given by,
Ð³Ð‚â€“" E" Ð³Ð‚â€”_"cell" ^"Ð’Ñ”" "=" "E" _"red" ^"Ð’Ñ”" +"E" _"oxid" ^"Ð’Ñ”" (3)
where EÐ’Ñ”red is the standard electrode potential of the cathodic half-cell and EÐ’Ñ”oxid is the standard electrode potential of the anodic half-cell. Thus, the experimental standard reduction potential, EÐ’Ñ”red of the five different electrochemical reactions can now be calculated with the Cu|Cu2+ as the reference cell with a standard reduction potential of +0.34 V.
As shown in Table 3 are the half-cell reactions with respect to the Cu2+|Cu electrode. Only reaction 1 has an oxidation reaction which indicates that the reference electrode undergoes reduction. It is because the its standard reduction potential is less than that of the reference electrode. The greater the standard reduction potential, the greater the tendency of the reduction reaction to occur. Half-reactions 2-5 have reduction reactions because their standard reduction potential is greater than the reference electrode.
Table 3. Half-cell reactions with reference to Cu2+|Cu electrode
Zn Ð¿Ñ“Â Zn2+ + 2e-
2Fe3+ + 2e-Ð¿Ñ“Â 2Fe2+
Cl2 + 2e-Ð’Â¬ Ð¿Ñ“Â 2Cl-
Br2 + 2e- Ð¿Ñ“Â 2Br-
I2 + 2e- Ð¿Ñ“Â 2I- Oxidation
As shown in Table 4 is the percent error of the theoretical and experimental EÐ’Ñ”red with the Cu|Cu2+ as the reference electrode. The half-cells measured are labeled at Table 3. Reaction 2 has the least percent error while reaction 5 has the greatest percent error. The average percent error of the five reactions is 33%.
Table 4. Percent error of theoretical and experimental EÐ’Ñ”cell with Cu2+|Cu reference electrode
Half-Reaction Experimental EÐ’Ñ”red (V) Theoretical EÐ’Ñ”red (V) % Error
The percent error of the experiment is significant. The sources of error include incorrect solution preparations, and contaminated glasswares and materials and resistance of the electrodes. The concentration of the prepared solutions may not have been the desired concentrations. They could have been greater than or less than what was desired. Another source of error is contaminated glasswares and materials. The beakers may have been contaminated with other reagents which could have interfered with the electrochemical reactions. The electrodes were completely immersed in the solution, and the alligator clip together with its rubber coating was also immersed in the solution. This alligator clip and rubber coating could have added some considerable resistance which would decrease the actual cell potential of the electrochemical cell.
To generate the halogens, X2 in reactions 3-5, a 1 M solution of the potassium halide, KX should be electrolyzed for about a minute with 3.0 V of electrical voltage. The 3.0 V is more than enough to cause the reaction to occur. There are three complicating factors for the correct calculations.
First, a voltage significantly in excess of the calculated value, an overpotential, may be necessary to cause a particular electrode reaction to occur. Overpotentials are needed to overcome interactions at the electrode surface and are common when gases are involved. During the electrolysis of the halogens, H2 gas is generated.
A second factor is that competing electrode reactions may occur. In the electrolysis of the aqueous potassium halide, there are two oxidation and two reduction half-reactions that need to be considered. As shown in Table 5 are the competing reactions. During electrolysis, the halides undergo oxidation to generate the halogens.
The oxidation of water, H2O and the reduction of potassium, K and water are present in all the reactions.
Table 5 Competing half-cell reactions
Half-Cell Reaction EÐ’Ñ”red (V)
2Cl- Ð¿Ñ“Â Cl2 + 2e-
2Br- Ð¿Ñ“Â Br2 + 2e-
2I- Ð¿Ñ“Â I2 + 2e-
2H2O Ð¿Ñ“Â O2 + 4H+ + 4e-
2K+ + 2e- Ð¿Ñ“Â 2K
2H2O + 2e- Ð¿Ñ“Â H2 + 2OH- -1.36
Note: The Ð²â€”Ð denotes that the reaction occurs in all of the half-cell reactions.
Because the standard reduction potential for the half-cell reaction 1 is somewhat less negative than that of the half-cell reaction of H2O, it must be expected that O2 is the product of the anode. Actually Cl2 is the predominant product because of the high overpotential of O2 compared with that of Cl2. The standard reduction potential of half-cells 2 and 3 are greater than that of H2O, thus it is expected that the halogens are the product. As expected, the reduction half-cell in the electrolysis of KX is the reduction of H2O. The predominant reactions in all of the reactions are the oxidation of the halides and the reduction of water. Therefore, the EÐ’Ñ”cell is just the sum of the EÐ’Ñ”red of the predominant reactions.
As shown in Table 6 is the GibbÐ²Ð‚â„¢s Free Energy, Ð²?â€ G of reactions 6 to 8. When a reaction occurs in a voltaic cell, the cell does work in the form of electricity. The GibbÐ²Ð‚â„¢s Free Energy, Ð²?â€ G is given by,
"Ð²?â€ G=nF" "E" _"cell" (4)
where n is the number of moles of electrons transferred between the electrodes, F the electric charge per mole of electrons, called the FaradayÐ²Ð‚â„¢s Constant equal to 96,485 coulombs per mole of electrons (96,845 C/mol e-) and Ecell is the cell potential.
Table 6. Gibb's Free Energy of selected reactions
Reaction GibbÐ²Ð‚â„¢s Free Energy (Ð²?â€ G)
8 -2.45 kJ
For a process occurring at constant T and P, if Ð²?â€ G<0, the process is spontaneous. If Ð²?â€ G>0, the process is nonspontaneous. If Ð²?â€ G=0, the process is at equilibrium. Since the GibbÐ²Ð‚â„¢s Free Energy of all the reactions are less than 0, the reactions are all spontaneous.
As shown in Table 7 is the percent error of the theoretical and experimental Ecell. The average percent error of reactions 6-8 is 41%.
Table 7. Percent error of theoretical and experimental Ecell
Reaction Experimental Ecell (V) Theoretical Ecell (V) % Error
The Ð²?â€ G for a reaction under any set of conditions can be related to its value for standard conditions, that is Ð²?â€ GÐ’Ñ”. The key term in relating the two is the reaction quotient, Q, formulated for actual, nonstandard conditions. The derivation is,
"Ð²?â€ G=Ð²?â€ GÐ’Ñ”+RTlnQ" (5)
The Ð²?â€ G from equation 5 can now be related to equation 4 to calculate the Ecell, which is the theoretical value for the cell potential, Ecell.
The percent error of the experiment is also significant. The sources of error are similar to the other reactions because the voltmeter used is the voltmeter used for all the reactions.
As shown in Table 8 is the percent error of the experimental and the theoretical Kf of [Cu(NH3)4]2+ and the experimental and theoretical Ksp of Cu(OH)2. The experimental error of Kf and Ksp are 100 and 1020 respectively. The percent error is very significant. The sources of error are similar to the other reactions because the voltmeter used is the voltmeter used for all the reactions.
Table 8. Percent error of experimental and theoretical K
Reaction Experimental Kf or Ksp Theoretical Kf or Ksp % Error
8 Ksp .0138
The EÐ’Ñ”cell and Keq can be related through this equation,
Ð³Ð‚â€“"EÐ’Ñ”" Ð³Ð‚â€”_"cell" "=" "0.0257" /"n" "ln" Ð²ÐƒÐŽÐ³Ð‚â€“"K" _"eq" Ð³Ð‚â€” (6)
The EÐ’Ñ”red can be calculated from equation 3. Equation 6 now becomes,
Ð³Ð‚â€“"EÐ’Ñ”" Ð³Ð‚â€”_"red" "=" "0.0257" /"n" "ln" Ð²ÐƒÐŽÐ³Ð‚â€“"K" _"f/sp" Ð³Ð‚â€” (7)
Thus, the experimental Kf or Ksp can be calculated from equation 7. The value of Kf is very high which means that the reaction will go to completion to form the [Cu(NH3)4]2+ and the value of Ksp is very small which means that the Cu(OH)2 is insoluble in water.
Based on the findings in the experiment, it is concluded that there is a significant percent error between the experimental standard reduction potentials of the various half-cells against the Cu2+|Cu half cells, and the theoretical value of those half cells. The average percent error is 33%. The sources of error are incorrect solution preparation and contaminated glasswares and materials and the resistance of the electrode.
It is also concluded that all the reactions in the second part of the experiment are spontaneous because their GibbÐ²Ð‚â„¢s Free Energy, Ð²?â€ G is less than 0.
 Masterton, William and Cecile Hurley. Chemistry: Principles and Reactions. United States: Brooks/Cole, 2005
 Petrucci, Ralph, and William Harwood. General Chemistry: Principles and Modern Applications. United States: Prentice-Hall, Inc., 1997
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