EXPERIMENT 16 Electrochemical Cells: A Discovery Exercise 1. Introduction. Discussion

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1 EXPERIMENT 16 Electrochemical Cells: A Discovery Exercise 1 Introduction This lab is designed for you to discover the properties of electrochemical cells. It requires little previous knowledge of electrochemical cells your preparation should concentrate on how you will carry out the experiment, not on theoretical aspects. An electrochemical cell is based on an oxidation-reduction (redox) reaction and consists of two half-cells: an anode half-cell and a cathode half-cell. Oxidation occurs at the anode; reduction occurs at the cathode. An electrochemical cell can produce an electric current, which is driven by an electrical potential difference between the two half-cells. In this experiment you will use a meter to measure and compare the electrical potential differences of several electrochemical cells, some of which will have different concentrations of metal ions. Discussion Electrochemical Half-cell The half-cells constructed in this experiment consist of a piece of metal in contact with an aqueous solution containing ions of the same metal. Standard- State Conditions for such a half-cell require that the metal-ion concentration be 1.0 M. Thus, a piece of pure zinc in a 1.0-M solution of Zn 2+ ions constitutes a standard zinc zinc-ion half-cell; a piece of copper in a 1.0-M solution of Cu 2+ ions constitutes a standard copper copper-ion half-cell, and so on. The piece of metal in a half-cell is referred to as an electrode. An electrode is something that can conduct electrons into or out of the half-cell. Electrochemical Cell If two half-cells are connected by placing a wire between the pieces of metal and by adding a salt bridge between the two solutions, a direct electric current can flow through the circuit. The electric current is generated because metal atoms in the more reducing metal convert to ions and leave one electrode to enter the solution and ions of the less reducing metal accept electrons and plate out on the other electrode. The electrons left behind when positive ions are formed at one electrode pass through the external circuit and into the other electrode. There the electrons combine with ions from the solution to form metal atoms. By measuring the direction of current flow, and the voltage generated in the cell, you can determine which is the more reducing metal (stronger reducing agent), and by how much. 1 This experiment was designed and written by Joe March and revised by Gordon Bain. Further revision by Chad C. Wilkinson and John W. Moore. Adapted by Julie C. Schlenker for use at Harvard. Copyright 2011 by the Department of Chemistry, University of Wisconsin. Experiment 16 1

2 Salt Bridge In order for current to flow, there must be a complete electric circuit. The wire is part of the circuit and the salt bridge completes the circuit. In this experiment, the salt bridge is a porous cylinder soaked with 1.0 M aqueous potassium nitrate. Remember that solutions of salts, such as potassium nitrate, are electrolytes they conduct electrical current by movement of positive and negative ions in the solution. Thus the porous cylinder provides a path for conduction of electricity, just as the wire does, completing the electrical circuit. Because diffusion of the solutions through the porous cylinder is slow, there will be no mixing of the solution of one half-cell with the solution of another on the time scale of the experiment. Thus the half-cells are connected electrically, but not chemically, by the salt bridge. Without a salt bridge a cell will not produce an electric current and you will not be able to measure the electrical potential difference between the two electrodes. Anode and Cathode The half-cell in which oxidation occurs is called the anode. This is the half-cell in which metal atoms lose electrons (are oxidized) to form positively charged ions (which go into solution). The electrons flow into the external circuit from the anode. The half-cell in which reduction occurs is called the cathode. This is the half-cell in which metal ions from the solution gain electrons (are reduced) and plate out onto the electrode as uncharged atoms. The electrons flow out of the external circuit into the cathode. To easily remember what happens at the anode and cathode, note that oxidation and anode both begin with vowels; reduction and cathode both begin with consonants. Measuring Electrical Potential Difference with a Meter You will use a multimeter to measure the electrical potential difference between half-cells. This electrical potential difference is called the cell potential, E cell. The multimeter has two leads (wires connected to it), that you will connect to the two pieces of metal in each pair of half-cells. The reading on the meter is the electrical potential difference between the electrically positive lead (red wire) and the electrically negative lead (black wire). E cell(measured) = E positive lead E negative lead If the meter reading is positive, this means that the positive lead is connected to a metal that has a more positive potential than the metal connected to the negative lead. If the meter reading is negative, this means that the positive lead is connected to a metal that has a more negative potential than the metal connected to the negative lead. Experiment 16 2

3 Because oxidation is defined as loss of electrons (increase in oxidation number), the half-cell where oxidation is taking place generates electrons and causes the piece of metal to become more negative. Thus, when the leads are connected so that the meter reading is positive, the anode (where oxidation is occurring) is the electrode connected to the negative lead of the meter and the cathode is the electrode connected to the positive lead. The meter reads the difference in potential between the positive lead and the negative lead, so, when the meter reading is positive: E cell = E cathode E anode Experimental Procedure Safety in the Laboratory Safety glasses or safety goggles must be worn at all times while in the laboratory. Nitrile gloves must be worn at all times while performing this experiment or handling chemicals. Long sleeves must be worn as students will be working with 1 M silver nitrate that will stain your skin. Do not lean on the lab surface. The aqueous metal salt solutions may irritate your skin. In case of contact with the skin, wash the affected area with water for 15 minutes. Make sure you wash your hands before leaving the laboratory. Waste Disposal and Cleanup All solutions should be poured into the waste collection bucket provided. Metal wire or foil must be washed thoroughly with distilled water, dried, and then returned to the center bench. Do not put any solid metal in the trash! Before You Leave the Lab Have your TF check your lab bench for cleanup. Submit your data and lab report to your TF. Wash your hands before leaving the lab. Equipment and Reagents The following equipment and reagents will be available at the center bench. Multi-EChem module multimeter with leads alligator clips copper strip, Cu (s) zinc strip, Zn (s) nickel, Ni (s) iron nail, Fe (s) silver strip, Ag (s) sandpaper aqueous copper (II) nitrate, 1M Cu(NO 3 ) 2 (aq) aqueous zinc nitrate, 1 M Zn(NO 3 ) 2 (aq) aqueous nickel (II) nitrate, 1 M Ni(NO 3 ) 2 (aq) aqueous iron (II) chloride, 1 M FeCl 2 (aq) aqueous silver nitrate, 1M AgNO 3 (aq) aqueous potassium nitrate, 1.0 M KNO 3 (aq) 1 ml volumetric pipet 10 ml volumetric flask Experiment 16 3

4 Part A: Metal/Metal Ion Cells at 1 M Set Up Each Half-Cell A half-cell consists of a piece of metal partially immersed in a solution of a salt of the same metal. 1. Using a bottle top dispenser, gently dispense 2 ml of solution into wells 1-5 of the Multi-EChem module, according to the table below. (DO NOT dispense the solution too quickly or it will splash into an adjoining well or into your face!) Well # 1 M solution 1 Copper (II) nitrate 2 Zinc nitrate 3 Nickel (II) nitrate 4 Iron (II) chloride 5 Silver nitrate 2. Gently dispense 5 ml of KNO 3 (aq) into the center well of the Multi-EChem Module. 3. Obtain one strip of each type of metal (the electrode) from the center bench. If necessary, use sandpaper to clean the surface of the metal and remove any oxide coating. 4. Place a metal electrode in each corresponding solution (e.g., Zn in the Zn(NO 3 ) 2 (aq) solution, etc.). 5. All half-cells can be prepared in the Multi-EChem Module at the same time. Measure the Electrical Potential Difference between Each Pair of Half-Cells Use the multimeter to measure an electrical potential difference between two half-cells. 1. Connect the (-) lead (black) to the COM port of the multimeter. Connect the (+) lead (red) to the VΩmA port. Turn the dial to 2 V (this means the meter will read a maximum of 2 V). 2. Connect the (+) lead (red) to the metal electrode in one half-cell. Connect the ( ) lead (black) to the metal electrode in a different half-cell. Observe the first steady potential from the meter. (The potential will change slowly, so record the first reading that seems like it is not noise from making the connection.) 3. Record your data in Table 1 of the Data and Analysis portion of the Lab Report. 4. Observe and record the potential difference for all of the possible combinations of half-cells available with the set-up you prepared. (You should have 20 different cell potentials.) Notes: Check the appearance of the metal strip before each measurement. If the strip appears dull you may need to sand it again before taking a measurement in order to have a pure metal surface that will yield good results. Make certain that the alligator clips are firmly attached to the metal strips so that there is a good electrical connection. Make certain there is no mixing of solutions in different cavities of the Multi-EChem Module. This will change concentrations and cause inaccurate results. Make certain there is enough solution in the well to reach the salt bridge. Experiment 16 4

5 Part B: Establishing a Relationship between Cell Potential and Concentration 1. Use a plastic pipet to remove all but the silver-ion solution and the salt-bridge solution out of the Multi-EChem Module. Rinse each empty well by adding distilled water with a squirt bottle and removing it out with the plastic pipet. Do at least three rinses of each well. Put the used chemicals in the Used Chemicals beaker at your bench. 2. Use two 10 ml volumetric flasks and a 1 ml pipet, along with the 1 M copper-ion solution from the center bench, to prepare a solution that is 0.1 M in copper (II) ions. 3. Place the 1 M copper-ion solution and the 0.1 M copper-ion solution into two separate empty wells in the EChem module. Be sure to record which well contains each solution. 4. Repeat your dilution procedure to produce three more successive dilutions and obtain copper-ion solutions that are 0.01 M, M, and M concentrations. As you complete each dilution, add the solution to a different empty well in the EChem module, again recording which well contains each solution. Pour any excess chemicals into the Used Chemicals beaker at your bench. 5. Your Echem module should now contain five copper-ion solutions (1 M, 0.1 M, 0.01 M, M, and M), along with the silver nitrate solution and the salt bridge solution. Measure the potential difference between the silver half-cell and a half-cell made from each of the different copper-ion solutions. Connect the cell in such a way that you get a positive voltage reading. Record the measured potentials in the appropriate data table in your Lab Report. 6. Use a plastic pipet to suck the silver ion solution out of the Multi-Echem Module and rinse out the well three times. Refill the well with a zinc half-cell. (Alternatively, if you still have a clean empty well remaining in your Echem module, you could simply add the zinc solution to that empty well.) Repeat your measurements by connecting the zinc half-cell to each of the different copper solutions. Be sure to connect your cells so you obtain a positive voltage. Record the voltages you obtain in the data table. Waste Disposal and Clean-up Thoroughly rinse all metal strips with distilled water over the Used Chemicals beaker. Dry the metal strips with a paper towel and return them to the appropriate location at the center bench. Thoroughly rinse your EChem module with distilled water, making sure that all rinse water goes into the Used Chemicals beaker. Place your clean EChem module back at the center bench. Rinse out the 10 ml flasks and the pipet with distilled water and put them back at the center bench. Pour all aqueous waste and the contents of your Used Chemicals beaker into the proper waste container in the hood. Put the multimeters and alligator clips back at the center bench. Wipe down your lab bench with a paper towel when you are done. Before You Leave the Lab Have your TF check your lab bench for cleanup. Submit your data and lab report to your TF. Wash your hands before leaving the lab. Experiment 16 5

6 Experiment 16: Data Collection, Lab Report and Prelab Before You Come to Lab: Read the entire lab report, including the previous introduction and discussion, and the entire procedure. Complete the Prelab, which is the last page of the lab report, and turn the prelab into your TF as you enter the lab. Safety in the Laboratory Safety glasses or safety goggles must be worn at all times while in the laboratory. Nitrile gloves must be worn at all times while performing this experiment or handling chemicals. Long sleeves must be worn as students will be working with 1 M silver nitrate that will stain your skin. Do not lean on the lab surface. The aqueous metal salt solutions may irritate your skin. In case of contact with the skin, wash the affected area with water for 15 minutes. Make sure you wash your hands before leaving the laboratory. Waste Disposal and Cleanup All solutions and aqueous waste should be disposed of in the waste bucket in the back hood. Metal wire or foil must be washed thoroughly with distilled water, dried, and then returned to the center bench. Do not put any solid metal in the trash!! The EChem module must be rinsed with distilled water and returned to the center bench. Wipe down your lab bench with a paper towel when you are done. Before You Leave the Lab Have your TF check your lab bench for cleanup. Submit your data and lab report to your TF. This page and all subsequent pages must be stapled and turned in. Wash your hands before leaving the lab. Grading: Prelab: Lab Report: Safety: Cleanup: /10 /20 /3 /2 Total: /35 Experiment 16 6

7 Data and Analysis Part A: Metal/Metal Ion Cells at 1 M 1. Fill in the table of cell potentials that you measured. Include the sign of the meter reading. 1 point Electrochemical cell Negative lead (black, COM port) potential (V) Cu Cu 2+ Zn Zn 2+ Ni Ni 2+ Fe Fe 2+ Ag Ag + Positive lead (red, VΩmA port) Cu Cu Zn Zn Ni Ni Fe Fe Ag Ag Metals are often listed as an electromotive (or electrochemical) series. Such a series places metals that are most easily oxidized at the top. For example, sodium metal is very easily oxidized to sodium ions so easily that it reacts with water to produce sodium ions so sodium is near the top of the electrochemical series. Consider the metals you worked with in this experiment. When a metal is oxidized, ions of another metal are reduced at the same time. Since oxidation occurs at the anode, think about your results for observed cell potentials, and list the metals in order of their ability to be oxidized. Use your data from the table above to place each metal in the correct position in the electromotive series below. Put the most easily oxidized metal at the left of the list and the least easily oxidized metal at the right end. Electromotive Series: 1 point Zn > Fe > Ni > Cu > Ag Part B: Establishing a Relationship between Cell Potential and Concentration 2. Fill in the table below with the results of your experiments with different concentrations of aqueous copper(ii) ions. 2 points Conc. of Cu 2+ (aq) (M) Log Conc of Cu 2+ (aq) (M) E cell (vs Ag Ag + half-cell) (V) E cell (vs Zn Zn 2+ half-cell) 1 M M M M M (V) Experiment 16 7

8 3. On the axes below, plot your data from the previous table( i.e. plot E cell vs. the log of the copper-ion concentration). You should plot the data collected using the Ag Ag + half-cell and the data collected using the Zn Zn 2+ half-cell on the same axes. Using a calculator or software, fit a line to each data set and determine its slope, y-intercept, and R 2 value. 4 points For the Ag Ag + graph: Slope = y-int. = R 2 = For the Zn Zn 2+ graph: Slope = y-int. = R 2 = Experiment 16 8

9 Lab Report 1. Discuss the graphs you obtained for cell potential as a function of the log of the copper (II)-ion concentrations in Part B. How are the graphs of data from the silver half-cell and from the zinc half-cell similar, and how are they different? 2 points They are similar in that the magnitude of the slopes are the roughly the same and respectively because both concentrations are 1M. The differences lie in that Ag has a negative slope and Zn has a positive slope. That means Cu has a different role in chemical reaction for each case. (oxidation and reduction) 2. Oxidation occurs at the anode ( ) and reduction occurs at the cathode (+). Using this information and your data from Part A of the lab, write a complete balanced equation for the reaction that occurs when the half cells are Cu Cu 2+ and Zn Zn 2+, and then write a complete balanced equation for the reaction that occurs when the half-cells are Cu Cu 2+ and Ag Ag +. Explain how you knew which substance was being oxidized and reduced in each case. 2 points Ag/Ag + : (Net reaction) Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2Ag(s) Copper is being oxidized from a formal charge of 0 to 2+ so it loses electrons Silver is being reduced from Ag+ to Ag so it has gained an electron Zn/Zn 2+ : (Net reaction) Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) Copper is being reduced from a 2+ charge to a 0 charge thus it is gaining two electrons. Zinc is being oxidized from Zn charge 0 to a 2+ charge thus it is losing electrons. 3. Using your answers from problem 2, write out the Nernst equation for the reaction when the halfcells are Cu Cu 2+ and Zn Zn 2+, and write out the Nernst equation for the reaction when the halfcells are Cu Cu 2+ and Ag Ag +. Using these equations, explain the reason for the differences in the graphs that you noted in problem 1. 2 points E= (E 0 Ag E 0 Cu) log [Cu 2+ ] 2 [Ag + ] 2 E= (E 0 Cu E 0 Zn) log [Zn 2+ ] 2 [Cu 2+ ] Experiment 16 9

10 4. In the reaction that occurs when the half-cells are Cu Cu 2+ and Zn Zn 2+, is Cu 2+ a reactant or a product? Without referencing the Nernst equation, can you use concepts of equilibrium and LeChatelier s principle to explain the relationship between Cu 2+ concentration and voltage? Note that the greater (more positive) the value of E cell, the greater the tendency of the reaction to run in the forward direction and produce products. 2 points Cu 2+ (aq) is a reactant Observation : [Cu 2+ ] E cell product As [Cu 2+ ] increases, the reaction goes forward. That means the Cu 2+ is a reactant. 5. Describe three possible sources of error in your measurements of cell potentials and how the potentials would be affected by the error. 2 points - The metal strip could be dull and doesn t have a pure metal surface. - The alligator clips could not be firmly attached to the metal strips so that there could be a bad electrical connection. - Mixing of solutions in different cavities of the Multi-EChem Module. This will change concentrations and cause inaccurate results. - If there is not enough solution in the well, the solution could not reach the salt bridge. 6. Predict the effect of replacing the 1 M copper (II) nitrate solution used in this experiment with a 1 M copper (II) chloride solution. Explain the basis for your prediction. 1 point Because both solution has Cu/Cu2+ reaction, there is no difference in the net redox reaction. 7. Suppose that before you began your experiments the supply of potassium nitrate for the salt bridge solution had run out. The only substitute you could find was zinc nitrate solution. Will using the zinc nitrate solution in the salt bridge compartment affect the cell potential you measure? Why or why not 1 point The high reduction potential of Zn 2+ may lead to Zn 2+ reducing, rather than the actual metal cations in the cathode. Moreover, for the Zn 2+ half cell, the potential difference would be changed because [ Zn 2+ ] is changed Experiment 16 10

11 Prelab To be completed and handed in as you enter the lab. 1. Suppose you connect the positive lead to a Zn Zn 2+ half-cell and the negative lead to an Mg Mg 2+ half-cell and obtain a reading of V from the multimeter. Draw a diagram of this electrochemical cell. Label the metals and corresponding solutions and indicate the direction of electron flow. Which metal is the anode and which metal is the cathode? Explain how you know. 7 points Experiment 16

12 2. You will have two 10 ml volumetric flasks and a 1 ml volumetric pipet to use for part B. Explain how you will make solutions with copper-ion concentrations ranging from 0.1 M to M, starting with 1 M copper (II) nitrate. 3 points Experiment 16