EXPERIMENT 16 Electrochemical Cells: A Discovery Exercise1 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 Zn2+ ions constitutes a standard zinc|zinc-ion half-cell; a piece of copper in a 1.0-M solution of Cu2+ 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 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, Ecell. 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). Ecell(measured) = Epositive lead – Enegative 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 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: Ecell = Ecathode – Eanode 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 • aqueous copper (II) nitrate, 1M Cu(NO3)2 (aq) • multimeter with leads • aqueous zinc nitrate, 1 M Zn(NO3)2 (aq) • alligator clips • aqueous nickel (II) nitrate, 1 M Ni(NO3)2 (aq) • copper strip, Cu (s) • aqueous iron (II) chloride, 1 M FeCl2 (aq) • zinc strip, Zn (s) • aqueous silver nitrate, 1M AgNO3 (aq) • nickel, Ni (s) • aqueous potassium nitrate, 1.0 M KNO3 (aq) • iron nail, Fe (s) • 1 mL volumetric pipet • silver strip, Ag (s) • 10 mL volumetric flask • sandpaper Experiment 16 3 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 KNO3(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(NO3)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.