AP Chemistry Laboratory #21: Voltaic Cells. Lab day: Monday, April 21, 2014 Lab due: Wednesday, April 23, 2014

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1 AP Chemistry Laboratory #21: Voltaic Cells Lab day: Monday, April 21, 2014 Lab due: Wednesday, April 23, 2014 Goal (list in your lab book): The goal of this lab is to determine what factors affect the potential in a voltaic cell. Background/Introduction (DON T list in your lab book): Electrical potential is the basis for how all batteries release useful energy. Batteries have a separation of charge in the form of positive and negative electrodes or terminals, and the movement of electrons from one electrode to another can be made to pass through a system that will produce a desired effect such as lighting up a light bulb. The tendency of electrons to pass through this system is referred to as electrical potential and is measured in units called volts. Hence there is a greater tendency for the movement of electrons in a 12 volt battery than in a 6 volt battery. However although potential (or voltage) will measure the tendency for electron movement, voltage has nothing to do with the quantity of electrons that move. This quantity of electrons that move through a system is referred to as current and is measured in amperes which are in units of coulombs of charge per second. So one must consider both electrical potential and current when determining the amount of energy delivered by a system. The tendency for different metals to become oxidized and in turn form metal ions in solution to become reduced is the basis for electrochemical cells. There are two types of electrochemical cells: voltaic cells, which act as batteries and spontaneously release energy, and electrolytic cells, which are nonspontaneous and require outside energy to function. In this investigation, we will be examining galvanic (voltaic) cells only. (Image from Measuring Cell Potentials, ChemFax, Flinn Scientific Inc., Publication 11096, 2014.) In this lab, students will be constructing a series of voltaic cells in microscale. As electrons travel from one half-cell to the other, there would be a tendency for one half-cell to become positive (the one where electrons are leaving) and for one half-cell to become negative (the one where electrons are entering). This situation is prevented by the addition of a salt bridge which contains a salt solution such as potassium nitrate neither of whose ions will react with ions in solution or the electrodes. For every electron that leaves the oxidation half-cell,

2 a negative ion from the salt bridge will take its place, keeping that half-cell electrically neutral. For every electron that enters the reduction half-cell, a positive ion from the salt bridge will take its place, again, keeping that half-cell electrically neutral. The nature of the salt bridge can determine the quantity of current electrons can only leave/enter half-cells as fast as ions from the salt bridge enter half-cells. But the focus of this study is on voltage not current so a crude salt bridge can be constructed. To this end and for this study, a salt bridge can be made by soaking a strip of filter paper in a 1.0 M potassium nitrate solution for a few minutes. This strip can then be folded such that it is in contact with each half-cell solution. Standard reduction potentials commonly found in tables are standardized to a hydrogen gas electrode, and the reduction potential for 1 M hydrogen ions to hydrogen gas at 1 atm is 0.00V. In creating a student-prepared reduction potential table as part of this lab, a hydrogen gas electrode is usually not convenient or available. Hence, some convenient metal in this case, zinc can be used as a reference, keeping in mind that all potentials measured with this reference will theoretically differ from standard reduction potential values found in published tables by the value of the standard reduction potential for zinc. Of course oxidation potentials can be calculated from reversing a reduction half reaction and changing the sign for the potential. Research questions (answer in complete sentences in your lab book please don t write or number the questions): 1) What is a galvanic (or voltaic) cell? 2) What is oxidation? Where does oxidation occur in a galvanic cell? 3) What is reduction? Where does reduction occur in a galvanic cell? The following data were measured using a nickel electrode as the reference standard: Cu 2+ (aq) + 2 e - > Cu (s) V Ni 2+ (aq) + 2 e - > Ni (s) V Fe 2+ (aq) + 2 e - > Fe (s) V Al 3+ (aq) + 3 e - > Al (s) V 4) Which ion is the most easily reduced? 5) Which metal is the most easily oxidized? 6) The aluminum and copper electrodes were connected as part of a voltaic cell to form a battery. (a) Which was the anode? How do you know? (b) Which was the cathode? How do you know? (c) What was the battery s voltage? Show work. (d) Write a balanced net ionic equation for the reaction that took place. (Hint: Requires writing and balancing the half-reactions and then adding them together!.) 7) Most standard reduction potential tables use a hydrogen electrode as a standard/reference. Why won t we? What will we use as our reference instead? 8) Show work to confirm the dilutions of the 1.0 M copper (II) nitrate solution to form M copper (II) nitrate. (Hint: Use the dilution formula!) 9) How do you predict the voltage will change when the concentration of copper

3 (II) nitrate solution is changed in step 8 of the procedure? prediction. Justify your Materials (don t list in your lab book): 10 ml 1.0 M copper (II) nitrate solution 10 ml 1.0 M iron (III) nitrate solution 10 ml 1.0 M lead (II) nitration solution 10 ml 1.0 M magnesium nitrate solution 10 ml 1.0 M silver nitrate solution 10 ml 1.0 M zinc nitrate solution 10 ml 1.0 M potassium nitrate solution 2 pieces copper foil 1 iron nail 1 piece lead foil 1 piece magnesium ribbon 1 piece silver foil 1 piece zinc strip filter paper 9 10 ml beakers sandpaper or steel wool 1 voltmeter wires and alligator clips 7 droppers 3 toothpicks Hazards (list in your lab book): (include the safety contract and the hazards of all of the solutions and lead) Procedure (don t list in your lab book): 1) Physically and chemically clean the beakers 2) Label 6 beakers with the names (or formulas) for each of the metal nitrate solutions (except potassium nitrate) 3) Obtain the solutions in the 10 ml beakers. (Don t obtain potassium nitrate) 4) Prepare a test cell to measure the voltage of the copper and zinc half-cells. a) Polish small strips of zinc and copper metal with sandpaper or steel wool (NOT DIRECTLY ON THE TABLE) b) Place each metal in its appropriate beaker of solution. c) Obtain a small strip of filter paper that has been soaked in potassium nitrate solution. d) Drape the soaked filter paper between the copper (II) nitrate and zinc nitrate beakers so that one end dips in each solution. (This acts as the salt bridge.) NOTE: Use a fresh salt bridge for each new galvanic cell! e) Use a voltmeter to measure the potential difference between the two half-cells. Connect the negative terminal of the voltmeter to the zinc electrode. Use the most sensitive scale practical. NOTE: When the voltmeter reads a positive charge, the electrode connected to the positive terminal is the cathode, and the electrode connected to the negative terminal is the anode. f) Record the two electrodes used and the voltage of the cell. 5) Measure the voltages of the other cells where one of the half-cells is zinc. a) Polish the metals with the sandpaper or steel wool. b) Use a fresh salt bridge for each cell. c) Record the two electrodes used and the voltage of the cell. 6) Measure the voltage (potential difference) between at least six other combinations of the various electrodes (but not using zinc this time). 7) Prepare a M copper (II) nitrate solution

4 a) Count 18 drops of distilled water into the 7th 10 ml beaker b) Add 2 drops of 1.0 M copper (II) nitrate solution c) Stir the mixture well with a toothpick. [This mixture is now 0.10 M.] d) Count 18 drops of distilled water into the 8 th 10 ml beaker e) Add 2 drops of the 0.10 M copper (II) nitrate solution just prepared. f) Stir the mixture well with a toothpick. [This mixture is now M.] g) Count 18 drops of distilled water into the 9 th 10 ml beaker h) Add 2 drops of the M copper (II) nitrate solution just prepared. i) Stir the mixture well with a toothpick. [This mixture is now M.] 8) Measure the voltage of a copper-zinc cell made with the M copper (II) nitrate and 1.0 M zinc nitrate. 9) Clean up! Post-Lab Questions: 1) Write the reduction equations for each metal ion, arranging the equations in decreasing order of measured potential in the table below. Include zinc in the table, using 0.00 V as the 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. A sample table is shown below: Reduction equation Electrode potential using zinc as standard, Zn Accepted electrode potential using hydrogen as standard, Zn - 2) How do the measured potentials vary from the accepted potentials? 3) Use the electrode potentials from Data Table 1 to predict the voltages of the six half-cell combinations selected and recorded in Data Table 2. How did the predicted and measured voltages compare? 4) How did the dilution of the copper (II) nitrate solution affect the voltage? Does the change agree with Le Châtelier s Principle? 5) What does a negative value for a standard potential indicate? Lab handout based on the experiments Electrochemical Cells in Laboratory Experiments for Advanced Placement Chemistry (Second Edition) by S.A. Vonderbrink (Flinn Scientific, 2006) and Electrochemical Cells Guided Inquiry by William B Bond (June 2013)

5 Suggested Data Tables (to be copied/taped into your lab book or design your own!): Data Table 1: Voltage of Each Half-Cell versus Zinc Voltage (V) Name of Anode Name of Cathode Zn vs Cu Zn vs Ag Zn vs Fe Zn vs Mg Zn vs Pb Half-Cell #1 Identity Half-Cell #2 Identity Data Table 2: Predicted and Measured Cell Potentials Measured Potential Balanced Equation for the Cell Reaction Predicted Potential from Experimental Data (using Data Table 1)