Lab 8: Electrochemical Cells Introduction: Redox Background: Chemical reactions involving the transfer of electrons from one reactant to another are called oxidation-reduction reactions or redox reactions

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Lab 8: Electrochemical Cells Introduction: Redox Background: Chemical reactions involving the transfer of electrons from one reactant to another are called oxidation-reduction reactions or redox reactions. In a redox reaction, two half-reactions occur; one reactant gives up electrons (undergoes oxidation) and another reactant gains electrons (undergoes reduction). A piece of zinc going into a solution as zinc ions, with each Zn atom giving up 2 electrons, is an example of an oxidation half-reaction. Zn(s) → Zn2+(aq) + 2e- (1) The oxidation number of Zn(s) is 0 and the oxidation number of the Zn2+ is +2. Therefore, in this half-reaction, the oxidation number increases, which is another way of defining an oxidation. In contrast, the reverse reaction, in which Zn2+ ions gain 2 electrons to become Zn atoms, is an example of reduction. Zn2+(aq) + 2e-→Zn(s) (2) In a reduction there is a decrease (or reduction) in oxidation number. Chemical equation representing half-reactions must be both mass and charge balanced. In the half-reactions above, there is one zinc on both sides of the equation. The charge is balanced because the 2+ charge on the zinc ion is balanced by two electrons, 2e-, giving zero net charge on both sides. Another example of reduction is the formation of solid copper from copper ions in solution. Cu2+(aq) + 2e-→ Cu(s) (3) In this half-reaction the oxidation number of the aqueous copper is +2, which decreases to 0 for the solid copper, and again charge and mass are balanced. However, no half-reaction can occur by itself. A redox reaction results when an oxidation and a reduction half-reaction are combined to complete a transfer of electrons as in the following example: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) (4) The two electrons are not shown because they are neither reactants nor products but have simply been transferred from one species to another (from Zn to Cu2+ in this case).In this redox reaction, the Zn(s) is referred to as the reducing agent because it causes the Cu2+ to be reduced to Cu. The Cu2+ is called the oxidizing agent because it causes the Zn(s) to be oxidized to Zn2+. Any half-reaction can be expressed as a reduction as illustrated in the case where equation (1) can be reversed to equation (2). A measure of the tendency for a reduction to occur is its reduction potential, E, measured in units of volts. At standard conditions, 25 ?C and concentrations of 1.0 M for the aqueous ions, the measured voltage of the reduction half-reaction is defined as the standard reduction potential, E?. Standard reduction potentials used in this lab can be found in the simulation, by clicking “Standard Potentials” on the upper right of the screen. For the reduction half-reactions in equations (2) and (3), the standard reduction potentials are –0.76 V for zinc and +0.34 V for copper. The more positive (or less negative) the reduction potential, the greater is the tendency for the reduction to occur. So Cu2+ has a greater tendency to be reduced than Zn2+.Furthermore, Zn has a greater tendency to be oxidized than Cu. The values of E?for the oxidation half-reactions are opposite in sign to the reduction potentials: +0.76 V for Zn and –0.34 V for Cu. Galvanic Cells A galvanic cell or voltaic cell is a device in which a redox reaction, such as the one in equation (4), spontaneously occurs and produces an electric current. In order for the transfer of electrons in a redox reaction to produce an electric current and be useful, the electrons are made to pass through an external electrically conducting wire instead of being directly transferred between the oxidizing and reducing agents. The design of a galvanic cell (shown in Figure 1 for the equation (4) reaction) allows this to occur. In a galvanic cell, two solutions, one containing the ions of the oxidation half-reaction and the other containing the ions of the reduction half-reaction, are placed in separated compartments called half-cells. For each half-cell, the metal, which is called an electrode, is placed in the solution and connected to an external wire. The electrode at which oxidation occurs is called the anode [Zn in equation (4)] and the electrode at which reduction occurs is called the cathode [Cu in equation (4)]. The two half-cells are connected by a salt-bridge that allows a “current” of ions from one half-cell to the other to complete the circuit of electron current in the external wires. When the two electrodes are connected to an electric load (such as a light bulb or voltmeter) the circuit is completed, the oxidation-reduction reaction occurs, and electrons move from the anode (−) to the cathode (+), producing an electric current. Figure 1.Galvanic cell (or battery) based on the redox reaction in equation (4). The cell potential, Ecell, which is a measure of the voltage that the battery can provide, is calculated from the half-cell reduction potentials: Ecell = Ecathode - Eanode (5) At standard conditions, indicated by the superscript o, the standard cell potential, E?cell, is based upon the standard reduction potentials, as shown in equation (5). E?cell = E?cathode–E?anode (5’) Based on the values for the standard reduction potentials for the two half-cells in equation (4) [–0.76 V for zinc anode and +0.34 V for copper cathode], the standard cell potential, E?cell, for the galvanic cell in Figure 1 would be: E?cell = +0.34 V – (–0.76 V) = +1.10 V The positive voltage for Eocell indicates that at standard conditions the reaction is spontaneous. Recall that ?Go = − nFEocell, so that a positive Eocell results in a negative ?Go. Thus, the redox reaction in equation (4) would produce an electric current when set up as a galvanic cell. Finally, when a galvanic cell does not use standard concentrations, Nernst’s equation needs to be used to determine the non-standard cell potential (Ecell): ???????????????????? = ???????????????????? ???? − 0.0257???? ???? ???????????? (6) Here, recall that Q is the reaction quotient (for which you need a balanced equation to write the equilibrium expression and use it to calculate Q with the nonstandard concentrations), and n are the number of transferred electrons (which are the number of electrons that cancel when you add the half reactions to get the balanced equation). Pre-lab assignment(s) 1. For the zinc/copper galvanic cell covered on the introduction: a) Write the spontaneous chemical reaction occurring in this cell. b) Which electrode should be gaining mass? Which should be losing? c) Which electrode is acting as the anode? and write the oxidation half-reaction that occurs. d) Which electrode is acting as the cathode? and write the reduction half-reaction that occurs. 2. Using the Standard Reduction Potential Table, calculate the theoretical EMF for each of the following cells. Pb0,Pb2+(aq)/Cu2+(aq),Cu0 Zn0,Zn2+(aq)/Cu2+(aq),Cu0 Pb0,Pb2+(aq)/Zn2+(aq),Zn0 Procedures The simulation that we will use is found here: Reduction Potentials (The simulation has “levels” Each level will pose a task or a problem that is described in red at the bottom of the screen. We will do levels 0 to 4. There are two more levels that we WILL NOT do for this lab) Part 1 (Level 0): 1. You will construct two galvanic cells combining the zinc half-reaction with two different metal half-reactions (Cu and Pb). You will measure the cell potentials, Eocell, with a multi-meter using 1.0 M solutions for both half-cells, so Q = 1 and ln Q = 0 for the reaction. Thus, as is evident from Nernst’s equation (6), the measured cell potential will be the same as Eocell You will then use your measured Eocell values, the known zinc standard reduction potential, Eo = –0.76 V, and equation (5’) to calculate the Eo values for the two different half-reactions and compare these to the values in Table those on the simulation’s Standard Reduction Potentials Table. 2. Enter the simulation and choose the level as zero. The level is chosen by a dropdown menu at the lower right of the screen. 3. Construct a galvanic cell by adding solutions of 1.0 M Cu(NO3)2 as the aqueous Cu2+ and 1.0 M Zn(NO3)2 as the aqueous Zn2+ solution, and a copper metal electrode in the Cu2+ solution and zinc metal electrode in the Zn2+ solution. Click “Measure Cell Voltage”. Record the POSITIVE cell potential (Take a screen shot of this to include in your report). If the cell voltage is negative, rebuild the cell by switching the left and right half-cells. After you get a positive voltage, which electrode is on the left side, the cathode to the anode? Calculate the half-cell potential of copper and compare to the known value. 4. Repeat steps above to build a galvanic cell of Zn|1.0 M Zn(NO3)2||1.0 M Pb(NO3)2|Pb. Calculate the half-cell potential of lead and compare to the known value (Take a screen shot of your cell measuring a positive voltage). Part 2 (Level 1): 1. Choose Level 1. A cell is presented, and the student is asked to calculate the cell's standard potential (Eo). Record the cell you are presented with and calculate the standard cell potential. To answer this question, you need to look at the half-cell potentials that this simulation uses by clicking at the top left where it states, “Standard Potentials”. Then, you can use equation (5) to calculate E?cell. Take a screen shot here of you getting the right answer to include in your report. Part 3 (Level 2): 1. Choose Level 2. There is a fictitious metal/salt pair, Whodatium (Wd)/Whodatium (II) Nitrate. The user is asked to determine the standard reduction potential (Eo) for this pair. To do this you need to pair the fictitious electrode with a known electrode, and then use equation (5’) to get the fictious electrode standard potential. Do this and check your answer. Take a screen shot here of you getting the answer correct to include in your report. Part 4 (Level 3): 1. Choose Level 3. The student is asked to design a cell which will produce a specified standard potential (E?cell) that will be shown at the bottom of the screen. Keep in mind that the anode needs to be on the left and that you need to use the simulation’s Standard Reduction Potential Table to find the pair of electrodes that will give you the desired E?cell. Take a screen shot here of you getting the answer correct to include in your report. Part 5 (Level 4) 1. Choose Level 4. A cell is presented (with concentrations), and the student is asked to calculate the expected voltage of the cell (Ecell). To solve this problem, you need to use Nernst’s equation (6). Take a screen shot here of you getting the answer correct to include in your report. Lab Report Structure (HAS TO BE TYPED!!!- HANDWRITTEN REPORTS WILL GET A ZERO) ___/10 Pre-lab questions/assignment 1. For the zinc/copper galvanic cell covered on the introduction: a) Write the spontaneous chemical reaction occurring in this cell. b) Which electrode should be gaining mass? Which should be losing? c) Which electrode is acting as the anode and write the oxidation half-reaction that occurs? d) Which electrode is acting as the cathode and write the reduction half-reaction that occurs? 2. Using the Standard Reduction Potential Table, calculate the theoretical EMF for each of the following cells. Pb0,Pb2+(aq)/Cu2+(aq),Cu0 Zn0,Zn2+(aq)/Cu2+(aq),Cu0 Pb0,Pb2+(aq)/Zn2+(aq),Zn0 ___/10 Procedure: Give a brief summary of the procedure that you followed to do each of the five parts of the lab. You can give some general instructions for the use of the simulation and then provide specifics for each level (This should be two to three paragraphs) ___/10 Notes/Observations: Full credit will be given for notes and observations that describe everything you did and observed. You need to have notes for each of the levels that you did. ___/30 Calculations: For the parts that require calculations, you need to show all your calculations here (Typed calculations are preferred. However, you can take a picture of your LEGIBLE calculations and include them on your report) ____/40 Screenshots of results(s) Include the required screenshots that are mentioned in the procedures.