Exp 7 electrochemistry 2022W

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Dec 6, 2023

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EXPERIMENT 7A Electrochemistry: Galvanic Cells and the Nernst Equation PRE-LAB WEEK REQUIREMENTS DURING THE PRE LAB WEEK …... 1. Read the entire experiment. 2. Read the Techniques and Resources: • Techniques: Under glossary – read topics related to Electrochemistry These materials are accessed by clicking on Experiment 7 in the Chemistry 154 Online Laboratory Canvas website . 3. Complete the Experimental Design Form: Think carefully about interpretation of your results as you design some steps of your procedure and answers the questions in the EDF quiz on Canvas. Consult the list of Reagents and Apparatus available in the lab. Be sure to submit the form prior to the lab week (Experiment 7B) on canvas. Do not attempt to leave all of this Pre-Lab Week material until the Lab Week as you will be overloaded if you do!   

EXPERIMENT 7 Electrochemistry: Galvanic Cells and the Nernst Equation Electrochemistry: Galvanic Cells and the Nernst Equation Aim: To investigate redox reactions of several metals and ions. OBJECTIVES The objectives of this experiment are to:  Construct galvanic cells.  Design an experiment to arrange metals in an electrochemical series.  Understand the Nernst Equation and apply it to determine an unknown concentration. This Experiment is presented in two halves: Part I and Part II, to be done together Introduction to Electrochemistry and Part I of the Experiment Any chemical reaction involving the transfer of electrons from one substance to another is an oxidation-reduction (redox) reaction. The substance losing electrons is oxidized and the substance gaining electrons is reduced . Redox reactions are quite common and have a large impact on the world around us. The rusting of iron, for example, is a redox reaction where iron is oxidized and oxygen is reduced. We could make use of redox reactions if we found a way to put to work the energy involved in the electron transfer. Consider the redox reaction between metallic zinc and an aqueous solution of lead nitrate: Zn(s) + Pb 2+ (aq) → Zn 2+ (aq) + Pb(s) (1) The nitrate ions are spectator ions and are omitted in the above net ionic equation because they play no part in the redox reaction. Two electrons are transferred from each reacting zinc atom to a lead ion. The metallic zinc is thus oxidized to produce Zn 2+ ions and the lead ions (Pb 2+ ) are reduced to produce metallic lead. This redox reaction can be carried out by simply putting a piece of zinc in an aqueous solution of lead nitrate. Metallic lead immediately begins to plate out on the zinc, the concentration of aqueous Pb 2+ decreases, and the concentration of Zn 2+ in solution increases as time passes. Since reaction is clearly spontaneous, ∆ G for the process is negative. The reverse reaction does not take place spontaneously; nothing happens when a piece of lead is placed in a zinc nitrate solution. Lead ions, therefore, must have a greater tendency to undergo reduction (accept electrons) than do zinc ions . As shown below in Figure 1, the same spontaneous reaction between zinc and lead nitrate can be carried out without ever putting the two reactants in direct physical contact. A strip of zinc metal is dipped into a solution of zinc sulfate (ZnSO 4 ) and a piece of lead into a solution of lead nitrate, Pb(NO 3 ) 2 . These two metal strips are now the two electrodes of an electrochemical cell. Next, the two solutions are connected by a salt bridge , a porous substance (for example, a piece of a filter paper) saturated with a solution of salt, such as KNO 3 . The porous substance allows ions to pass between the solutions, but prevents the solutions themselves from mixing. If the two

electrodes are connected by a wire, electrons will flow from one electrode to the other, allowing the redox reaction to proceed in the electrochemical cell we have created. By convention, the short form for this electrochemical cell is denoted as: Zn(s) | Zn 2+ (aq) || Pb 2+ (aq) | Pb(s) where | represents the boundary between the electrode and the solution; || represents the boundary between the two half-cells. This convention assumes that reduction always takes place at the right-hand electrode and that oxidation always takes place at the left-hand electrode . In our galvanic cell shown above, as zinc is oxidized on the left-side, Zn 2+ ions enter the solution. The following oxidation half-reaction can thus be written: Zn(s) → Zn 2+ (aq) + 2e - oxidation half-reaction (Left-Hand Side) The electrons that are released pass through the external circuit wire from left to right . As the electrons flow into the lead strip, they are picked up by the lead ions at the metal-solution interface, forming lead atoms, and thus causing metallic lead atoms to plate out on the surface of the electrode. The following reduction half-reaction can thus be written: Pb 2+ (aq) + 2e - → Pb(s) reduction half-reaction (Right-Hand Side) • The electrode at which oxidation occurs is called an anode. (Figure 1, Left -hand side) • The electrode at which reduction occurs is called a cathode. (Figure 1, Right -hand side) Since the electrons flow from the anode to the cathode, this process would lead to an increase of positive charge in the left-hand beaker (the anode), and an increase of negative charge in the right- hand beaker (the cathode), and therefore the process could not proceed. The salt bridge, however, permits a compensating flow of ions to complete the circuit. What is the charge on the ions flowing into each beaker? How does this affect the overall charge in each beaker? We have effectively separated reaction (1) into two components (two half-reactions) and placed them into two separate compartments (half-cells). We have seen that metallic zinc is oxidized in the left beaker (anode) while lead ion is reduced in the right beaker (cathode). The zinc plate is the negative electrical pole as electrons are leaving the metallic Zn at this point, while the lead plate is the positive pole as electrons are being accepted by the Pb 2+ ions to form metallic Pb at this plate. Figure 1: a Zinc/Lead galvanic cell Zn → Zn 2+ + 2e - Pb 2+ + 2e - → Pb oxidation reduction anode cathode 1 M Pb(NO 3 ) 2 1 M ZnSO 4

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The electric potential difference between the two plates in Figure 1 is the motivating force that causes the electrons to flow from the left beaker to the right beaker. This potential difference , ΔE, is called a cell potential ( E cell = ΔE = E right - E left ) or electromotive force (emf) and can be measured by a voltmeter. The electrochemical cell described above operates spontaneously (negative ∆ G) and is called a galvanic cell . Such a cell allows a net conversion of chemical energy into electrical energy, which can then be used to perform work. Quantifying and Calculating Cell Potentials: Different metals, such as zinc and lead, have different tendencies to oxidize; similarly, their ions have different tendencies to undergo reduction. The cell potential of a galvanic cell is due to the difference of the two metals' tendencies to be oxidized (lose electrons), or their ions' tendencies to be reduced (gain electrons). Commonly, a reduction potential , a tendency to gain electrons, is used to represent the relative tendency for a given metal ion to undergo reduction. The potential measured in the cell is the result of the two half-reactions and the magnitude of the potential depends on the ions' concentrations, the temperature, and gas pressures; standard conditions exist when ionic species have a concentration of 1 M and gases have a pressure of 1 atm. When all the ionic concentrations in the zinc/lead system are 1.0 M and the temperature is 25 ° C, the cell voltage is 0.63 volts. It would be a monumental task to assemble a list of all possible cells and report a cell voltage so instead we use the potential of the half-reactions. We cannot measure the half-cell potential directly so we pick another half reaction, call it the standard, construct a cell, measure the cell voltage and report the potential relative to the standard. The standard that has been chosen by convention is: 2 H + + 2e - → H 2 E ° = 0.00 V Here the notation E ° is called the standard electrode potential and the pressure of the hydrogen gas is 1 atmosphere. The measured cell voltage using the standard hydrogen electrode as one of the half cells is, therefore, the potential of the other half reaction. In a teaching laboratory, it is difficult to work with hydrogen gas. Hence, for the purposes of this experiment, we will use Sn/ Sn 2+ as the reference electrode because it has an electrode potential close to zero. 2 Sn 2+ + 2e - → Sn E ° = -0.14 V

Tables of standard half-reaction potentials have been made. The reactions by convention are written as reductions and hence the tables are called tables of standard reduction potentials . A brief excerpt is shown below in Table 1. The greater the tendency of the ion to gain electrons and undergo reduction, the more positive the reduction potential of the ion. Since lead has a greater reduction potential than zinc, the metallic lead plates out on the cathode. Standard reduction potentials can be used to calculate the standard cell potential of any combination of half-cells. Recall that the reaction: Zn(s) + Pb 2+ (aq) → Zn 2+ (aq) + Pb(s) (1) is for the galvanic cell that is denoted as: Zn(s) | Zn 2+ (aq) || Pb 2+ (aq) | Pb(s) (2) By convention, standard cell potentials are calculated from the standard reduction potentials of the constituent half-cells according to eqn (3), where RH means Right-Hand side and LH means Left- Hand side of cell (2) above. E ° cell = E ° RH - E ° LH (3) As by convention the cathode is placed on the right while the anode is placed on the left, equation (3) can also be written as: E ° cell = E ° cathode - E ° anode (4) Using the data in Table 1, we can calculate the standard cell potential for the zinc/lead cell: E ° cell = E ° RH - E ° LH = - 0.13 – (- 0.76) = 0.63 V Table 1: Standard Reduction Potentials at 25 ° C Reduction Half-reaction Potential (V) Cu 2+ + 2e - → Cu(s) +0.34 2H + + 2e - → H 2 (g) 0.00 Pb 2+ + 2e - → Pb(s) -0.13 Zn 2+ + 2e - → Zn(s) -0.76 Mg 2+ + 2e - → Mg(s) -2.37 Li + + e - → Li(s) -3.05

Lab Project Overview: Part I Electrochemical Series: As we have seen, a list that ranks metal ions in order of their reduction potentials is useful. This list can be used to make various predictions, such as whether certain electrochemical reactions are possible; which metal would be the anode and which would be the cathode in a galvanic cell; or even what the actual voltages produced by galvanic cells might be. Accordingly, in Part I of this experiment, you will both create and use such a list (an “electrochemical series” or “table of reduction potentials”) for Cu, Mg, Zn, and Sn as follows: • set up three galvanic cells, using Sn/Sn 2+ as the reference half-cell (0 volts) for each one. The other half-cells examined will be Cu/Cu 2+ , Mg/Mg 2+ , and Zn/Zn 2+ . • measure the reduction potentials produced in the Cu/Sn, Mg/Sn, and Zn/Sn galvanic cells. Some will be positive (+) voltages while others will be negative (-). • use the data obtained and Chart 1 below to create an electrochemical series table. Note that your table is referenced to tin and not the standard hydrogen electrode. • use your table to predict the voltages for Cu/Zn, Cu/Mg, and Mg/Zn galvanic cells. Then, set up these three galvanic cells and measure their voltages. Compare your predicted and measured voltages. Experimental Procedure: Part I Reagents and Apparatus provided in the lab for Part I: • Mult-e-cell • 0.1M KNO 3 , Cu(NO 3 ) 2 , MgCl 2 , ZnSO 4 , and SnCl 2 • small beaker(s) • strips of filter paper • strips of Cu, Mg, Zn, and Sn metal • red and black wires to connect to the interface • sand paper • ceramic tile • Tweezers Sand your metal strips on the ceramic tile and not directly on the bench-top surface. NOTE: The following steps walk you through the process of setting up an electrochemical cell and reading the cell potential. In the lab simulation program, you will simply be reading the voltage from the voltmeter. Also, the student data you will be given for the post-lab assessment was collected in the lab using this set-up. 1. Set-up a galvanic cell consisting of a Cu/Cu 2+ half-cell and a Sn/Sn 2+ half-cell. Connect the half-cells with two salt bridges and fill the center well with KNO 3 . Attach the red and black alligator clips (the leads) to the Cu and Sn metal strips, but do not dip the metals in the cell solutions yet as the electrical circuit must not yet be completed.   

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Note: the picture shows Cu and Zn but it is the same idea 2. Start the Micro Lab computer programme. Complete instructions on how to use the software are on the laminated sheet in your kit. Prepare the computer to read the voltage of your Cu/Sn galvanic cell and type in a sample number but do not press the ‘Enter’ key yet. 3. Dip the metals in the appropriate cell solutions to complete the circuit. Wait 7 seconds and then press the ‘Enter’ key. The computer will take a voltage reading with units of V. Make sure your leads are connected to the metal strips so as to get a positive (+) voltage reading. The black lead connects to the negative terminal of the voltmeter; the red lead connects to the positive terminal. If the electrons flow into the voltmeter through the black(-) lead, and leave through the red (+) lead, the voltmeter displays a positive voltage. When the electrons flow into the voltmeter through the red (+) lead, and leave through the black (-) lead, the voltmeter displays a negative voltage. VOLTMETER Positive Reading Black ( - ) Direction of electrons ( + ) Red VOLTMETER Negative Reading Black ( - ) Direction of electrons ( + ) Red The measured voltage will have a positive sign if the black lead is on the anode (electron provider or electron source) and the red lead is on the cathode (electron receiver or electron sink). Identify and record which metal serves as the anode and which as the cathode. Record the voltage and remove the metals from solutions. Repeat the measurement. You should get approximately the same value; if not, repeat once more. Use the Data Tables provided on the Lab Report Sheet to record all observations. 4. Measure the voltage from the Sn/Mg system and the Sn/Zn system, always having the same lead on the Sn as you did for the Sn/Cu combination (since you continue to use Sn as your reference). 5. A copy of Chart 1 is included on your Post Laboratory Assessment (PLA) form. Using this chart, and tin as your reference (“0” voltage), write in the reduction potentials for Cu, Mg, and Zn above or below tin in order of their relative tendency to gain electrons. Write half- cell reaction equations for the metals you tested as per the example shown for Sn. a Mult-e-cell

Increasing tendency to gain electrons Voltage +1.0 0 Sn 2+ + 2e → Sn 0 E ~ 0 V -1.0 Chart 1: The Electrochemical Series Table. 6. Based on the data in your electrochemical series table and using equation (4), calculate the predicted voltage for Cu/Zn, Cu/Mg, and Mg/Zn galvanic cells. Enter the calculated values on your provided PLA form. 7. Set up these three galvanic cells and measure their voltages. In all cases the red lead should be attached to the cathode to yield positive values. Compare the observed voltages with your calculated values and comment in your PLA on any differences. 8. You will be provided with data/ pictures of a strip of metal and a salt solution of another metal. Using your completed electrochemical series table predict whether a redox reaction will occur between them. Dispose of all used chemicals and paper strips into the large plastic waste tub. Introduction to the Nernst Equation and Part II of the Experiment We have seen how the measured standard reduction potentials of half-cells can be used to theoretically predict the standard potential difference between two electrodes in a galvanic cell. For our predictions to be accurate, standard state conditions must be used. Unfortunately, in the lab, standard-state conditions do not always exist. The Nernst Equation (developed by Walter Nernst during the late 1800s while studying the thermodynamics of electrolyte solutions) allows cell potentials to be predicted when the conditions are not standard: E cell = E ° cell - 2 303 . log RT nF Q or E cell = E ° cell - Q ln nF RT (5) In equation (5), R is the gas constant (8.314 J mole -1 K -1 ), T is the temperature (Kelvin), F is Faraday’s constant (96,485 coulombs/mole), n is the number of electrons transferred in the balanced oxidation/reduction reaction, and Q is the reaction quotient, or ([products]/[reactants]). If the reactions are carried out at room temperature (25 ° C), the Nernst equation becomes: E cell = E ° cell - (6) 0 0591 . log n Q

Note in equations (5) and (6) that if Q = 1, then E cell = E ° cell . In Part II of this experiment, potentials will be measured at various solution concentrations for a copper/zinc galvanic cell and compared with those calculated using the Nernst equation. Note that the measured cell potentials may be slightly below the values predicted by the Nernst equation. This difference is due to the interactions between the ions in solutions. In this experiment we are using concentrations to approximate the activities of the ions in solution. In this part of the experiment, you are given the materials to construct the following electrochemical (galvanic) cell. Cu 2+ (aq) + Zn (s) → Cu (s) + Zn 2+ (aq) (7) Before looking at the next page for some hints, stop here and think about the following, keeping in mind the information given to you above. What do you think will happen to the cell potential, E cell , of (7) as concentrations of the copper solution are varied? Will it change? How does increasing the concentration of Cu 2+ (aq) affect the reaction quotient, Q? How, if at all, should that affect your measured cell potential? Use Le Chatelier’s principle to answer. The Nernst equation allows you to calculate E cell as a function of the reactant and product concentrations. For the above reaction at 25 ° C, the Nernst equation becomes E cell = E ° cell - 0 0591 2 . log {[Zn 2+ ]/[Cu 2+ ]} (8) Why is equation (8) lacking metallic zinc and copper? If you need a hint, what does the ratio of [Zn 2+ ] to [Cu 2+ ] represent? Theoretically, E ° cell for the above reaction is 1.10 V. What conditions must be satisfied if E cell is to equal E ° cell ? What must the ion concentrations be? Lab Project Overview: Part II The Nernst Equation: By measuring E cell for a series of Zn/Cu galvanic cells where the known concentration of Cu 2+ ions is varied, a calibration graph of E cell versus log[Cu 2+ ] can be constructed. Using this graph and the Nernst equation, one could then determine [Cu 2+ ] in a solution of unknown copper concentration by simply measuring E cell for a galvanic cell made from a Zn/Zn 2+ half-cell and a Cu/unknown Cu 2+ half-cell. Accordingly, in Part II you will: • set up five Zn/Cu galvanic cells and measure E cell for each one. • set up a Zn/unknown Cu galvanic cell and measure E cell . • plot a graph of E cell versus log[Cu 2+ ]. • use the graph to calculate the concentration of the unknown copper solution. Experimental Procedure: Part II   

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Reagents and Apparatus provided in the lab for Part II: • Mult-e-cell • 1.0 M ZnSO 4 • 500mL beaker • 0.1 M KNO 3 (Use to make dilutions) • strips of Cu and Zn metal • various concentrations of Cu(NO 3 ) 2 • strips of filter paper • red and black wires to connect to the interface 1. Using a serological pipet and the 1.0 M Cu 2+ solution provided make 10 mL of each of the following Cu 2+ solutions: a 0.10 M, 0.010 M, 0.0010 M and 0.00010 M Cu 2+ solution. 0.1 M KNO 3 is used to dilute the Cu 2+ solution . In your experimental design form (EDG) explain how you would make these solutions and using the solution you made, explain how you would make the next dilution. 2. Set up five Zn/Cu galvanic cells. The Zn 2+ solution concentration is 1.0 M for all five cells, whereas [Cu 2+ ] will range from 0.00010 to 1.0 M. The Zn/Zn 2+ half-cell should be placed in the center well while the five Cu/Cu 2+ half-cells should be assembled in the radial wells, which are number 1 to 5. Before setting up the half cells, rinse each well with the appropriate solution. Use a Pasteur pipet to remove the rinsing solution. Repeat the rinse to increase the accuracy of your measurement. Note that this set-up differs from that of Part I as only one salt bridge is needed to construct each cell. The filter paper must be pre-soaked in 5 mL of KNO 3 in a small beaker and then inserted with one end in the center cup and the other end into one of the outer cups. Set up just one salt bridge at a time to prevent the simultaneous exchange of ions amongst all five of the cells. Clip a straight piece of Zn metal to the black lead and use it as a probe; likewise for a piece of Cu metal and the red lead. 3. In the laboratory, you would measure and record in the second data table of your provided PLA form, the voltage (E cell ) of the most dilute galvanic cell first and work your way to the most concentrated cell so as to minimize the contaminating effect of any residual Cu(NO 3 ) 2 solution that might adhere to the copper electrode. The concentrations of the samples and the corresponding voltages will be sent to columns A and B, respectively, of the Micro Lab spreadsheet. 4. Use your graph paper to produce a plot of E cell versus log[Cu 2+ ] for cells #1-5. (Graph paper can be found at the back of the lab manual.) Draw the best fit line through the points. This graph is a “calibration graph” and serves the same function as the one used in the in-lab activity (electrochemistry activity). 5. Rinse out your used Mult-e-cell and set up a final galvanic cell using the Zn/Zn 2+ half-cell and a half-cell formed between metallic copper and an unknown copper solution . Measure and record E cell . 6. Use the graph to determine the concentration of the unknown copper solution. Show the calculations on your Report Sheet. 7. Using equation (8), an E ° cell value of 1.10 V, and the appropriate [Zn 2+ ] and [Cu 2+ ] values, calculate E cell for cells #1 – 5 , assuming a reaction temperature of 25 ° C. Show a sample calculation and present the five calculated E cell values in the same Data Table which you use to present the measured E cell values.

P OST L ABORATORY A SSESSMENT Complete the provided PLA form on canvas. Include the graph of E cell vs log[Cu 2+ ]. See the detailed Marking Scheme for more information. CHECKLIST Make sure to do the following before you come to the lab to do the experiment:  Read the experiment.  Complete the Experimental Design Form and Pre-lab quiz.  Complete the PLA assessment form and PLA quiz listed on your schedule for this experiment. LAB WEEK REQUIREMENTS DURING THE LAB WEEK …... 1. Attend the online zoom session, discuss pre-lab/ EDF content with instructors, complete H5P activity and participate in the presentation. 2. Complete your Post Lab assessment and upload it on canvas. Some follow up questions will be posted in the PLA quiz as well.    

Experiment 7 Marking Scheme Pre-lab quiz: • Complete the quiz on canvas – you have ONE attempt and this quiz has a time limit as well. Experimental Design Form: • Answer the questions in the canvas EDF form and submit on canvas. You only have one attempt but there is no time limit for this submission. On-Line Introductory Material: • Read the Techniques module “Mult-e-Cell Tutorial” at the Chem 154 Online Laboratory Canvas website. This is part of the EDF and pre-lab work. Laboratory Report using the provided Report Sheet Post Lab Assessment: • Answer the questions listed in the data form. • Complete the calculations for your group (A, B or C). • Fill out the Electrochemical Series Table with reduction potentials and half-cell reaction equations. Graph: • Title and label the graph. Mark where the unknown metal solution should appear on your graph. Include a picture of this graph or the file in your PLA submission. Calculations: • Show the sample calculations requested on the PLA form. Discussion: • Answer the questions listed on the PLA form.

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EXPERIMENT 7B The concept map below summarizes the electrochemistry concepts discussed in this experiment. can be divided into two of are tabulated relatively to a describes on is a difference in can produce from a requires to force a can be a can be an produces Electrochemical Cell Half Cells Reduction Potentials Standard Hydrogen half cell Cell Potential Nernst Equation effects of changes of concentrations Galvanic Cell Useful work spontaneous reaction Electrolytic cell input of work non-spontaneous chemical reaction Can be used to generate Electricity