# Electrochemistry and Concentration Effects on Electrode Potentials Prelab

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1 Electrochemistry and Concentration Effects on Electrode Potentials Prelab 1. What is the purpose of this experiment? 2. a. Calculate the standard cell potential of a cell constructed from Mg 2+ /Mg and Ni 2+ /Ni (Table I). Which is the anode and which is the cathode? b. Using the Nernst Equation, what would be the potential of a cell with [Ni 2+ ] = [Mg 2+ ] = 0.10 M? 1

2 Electrochemistry and Concentration Effects on Electrode Potentials Special mention and thanks goes to Dr. Timothy Friebe for his help in developing this laboratory write-up. Introduction You have already studied oxidation-reduction reactions in the previous experiment. You have an understanding of half-reactions and how they contribute to the overall redox reaction. In those experiments, the oxidizing and reducing agents were in the same container. In this experiment, you will see that oxidation-reduction reactions can occur between reactants in separate containers. This happens when electrons are transferred through a wire that connects the two containers. This is the basis of an electrochemical cell. Before the topic of the electrochemical cell can be explored, you need to understand a few terms. One of the most important is standard reduction potential (E o red). Table I located at the end of this experiment has some of the standard reduction potentials needed in this experiment. The standard reduction potential is a quantitative measure of a substance s tendency to accept electrons under standard conditions (1 atm of gases and 1 M concentration) and is measured in volts. The more positive the reduction potential, the greater the tendency for reduction to occur. For the oxidation half-reaction, the standard oxidation potential (E o oxid) for the half-reaction has the same numerical value as the standard reduction potential but the opposite sign. For example, aluminum ion has a standard reduction potential value of 1.65 V (in going to aluminum metal). The standard oxidation potential for aluminum metal is then V. The standard cell potential (E o cell) is the sum of the standard oxidation and reduction potentials. To construct a nickel-copper electrochemical cell, for example, you would start by placing a strip of nickel metal into a nickel(ii) salt solution. A strip of copper metal is placed into a copper(ii) salt solution in a different container. A wire is connected between the nickel strip and the copper strip. Electrons will pass through this wire during the redox reaction. The E o red value for nickel ion is 0.26 V and V for copper ion. This means that the Cu 2+ ion has a greater tendency to accept electrons than the Ni 2+ ion. Therefore, electrons will pass through the wire from the nickel strip to the copper strip. However, this can only happen if a second connection is provided to complete the electrical circuit. This second connection is called a salt bridge and is placed in both solutions of the two containers. You now have an electrochemical cell. The electrochemical cell is a system that utilizes a spontaneous oxidation-reduction reaction to pump electrons through an electrical circuit. In this example, the nickel half-cell involves nickel metal being oxidized and the copper half-cell involves copper(ii) ion being reduced. The metal strips are called electrodes. The oxidation electrode is called the anode and the reduction electrode is the cathode. The cell potential can be measured by connecting a voltmeter between the halfcells. This is what you will do and you can use Figure 1 as a reference to check your set up (as well as asking your instructor). To predict the observed voltage for this oxidation-reduction reaction, simply add the two half cell potentials together. Since the copper(ii) ion is being reduced in the reaction, its half-reaction 2

3 is the reduction, and the nickel half cell will be an oxidation. The copper electrode voltage will remain a reduction and the sign of E o red of Ni 2+ will be changed to give E o ox Cu e Cu E o red = V Ni Ni e E o ox = V cell reaction Ni + Cu 2+ Ni 2+ + Cu E o cell = V The reading on your voltmeter should be close to V. You are not working under standard conditions, but the ratio of the concentrations is still 1:1. The Nernst equation shows that you will still measure the standard cell potential. Any differences in the readings would come from other sources such as resistance in the salt bridge and the concentrations of the two cells not being exactly the same, for whatever reason. The second half of this experiment deals with the quantitative effect concentration changes have on cell voltage as shown by the Nernst equation. In a redox equation aa(aq) + bb(aq) cc(aq) + dd(aq) (1) the cell voltage is described by the Nernst equation c d o [C] [D] cell cell a b E = E log n [A] [B] (2) where E o cell is the standard cell voltage, E cell is the non-standard voltage and n represents the total number of electrons transferred in the balanced cell reaction. Looking at this equation, you see that increasing the concentration of the products lowers the voltage of the cell since the log term will be larger making the value of E cell smaller. Equation (2) represents the Nernst equation for the entire cell. It can also be written for a halfcell. For example, to determine how the concentration of the silver ion affects the total cell potential, you can first determine how the concentration of the silver ion affects the half-cell reduction potential. The reaction for the half-cell reduction occurring in today s experiment is: The Nernst equation for reaction (3) is: Ag + + e Ag (3) E = E o log (4) + 1 [Ag ] The electrochemical cell you are working with will be made up of two half-cells but in this case one is a saturated calomel electrode while the other is a silver metal electrode. The standard 3

6 Part III should increase each time, but not go above 0.8 V. If they do not, see the instructor. Enter Ag 0.001, Ag 0.01 and Ag 0.1 for the three entries. 1. Prepare a saturated AgCl solution by adding one drop of 0.10 M Ag + to 25 ml of 0.10 M HCl. Insert the electrodes into this solution and the potential as before. Record as AgCl. Click. Data Treatment For each electrochemical cell, calculate the standard cell potential in the same manner shown in the introduction section of this laboratory experiment. Include the half-reactions and electrode potential values. Compare the calculated cell potentials to the measured voltages of your electrochemical cells. Describe the possible sources of error. Using your data from the first three Ag/Ag + cell potential readings, determine the average standard reduction potential (E o ) for silver ion. Use your average value of E o to determine [Ag + ] in the fourth solution. Knowing that [Cl ] is 0.10 M in the fourth cell, calculate K sp of AgCl. Compare with a published value of K sp (give the reference). Describe any source of error. Conclusion List the cell voltages, average E o of Ag/Ag +, the [Ag + ] in saturated AgCl, and your K sp of AgCl. 6

7 Figure 1. Diagram of an Electrochemical Cell Table I. Standard Reduction Potentials for Selected Metals Half-Reaction E o red(v) Mg e Mg 2.37 Al e Al 1.65 Zn e Zn 0.76 Fe e Fe 0.44 Ni e Ni 0.26 Cu e Cu Ag + + e Ag

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