Experiment 10. Electrochemistry: Making a Battery. Lab Report

Please follow the instructions and watch the video that I give you (lab 10 manual and post-lab file), the link for videos are on the file, and complete the Quiz and Post Lab in the file Making Battery Post Lab 2020. If you have any questions, please contact me immediately.

Experiment 10. Electrochemistry: Making a Battery
Terms:
• Alkaline battery
• Anode
• Battery
• Cathode
• Connected in series
• Electric current
• Electrochemistry
• Electrode
• Electroplate
• Half-cell
• Half reaction
• Net cell potential (E0net)
• Oxidation
• Redox reaction
• Reduction
• Salt bridge
• Standard Reduction & Oxidation
§ Potentials (E0red & E0oxid)
• Voltage/Voltmeter
Introduction to Batteries
An electric current is a flow of electrons through a conductor such like copper wire. Onecommon source of electric current which we use every day to run our cell phones, laptops,flashlights, remote controls, game controllers, etc. is the battery. All batteries convert chemicalenergy into electrical energy (an electric current).
Redox Reactions Provide the Chemical Energy for Batteries
The flow of electrons generated by a battery results from a spontaneous redox reaction.Redox reactions involve two reactions which occur together: the oxidation of one reactantwhich loses electrons and the simultaneous reduction of another reactant which gainselectrons. Typically, in a battery, the oxidation and reduction reactions are separated into twoseparate cells - called half-cells. Each half-cell has an electrode - a general term that refers tothe site of the redox reaction where the electrons are released or absorbed. Electrodes areusually a metal submerged in a liquid or it can also be a solid which contains ions able toconduct an electric current.
In batteries, oxidation occurs at the anode when atoms in the metal electrode loseelectrons and become aqueous ions and go into solution (Table 1). Consider the oxidationexample in Table 1 where zinc atoms (Zn0) in the anode dissolve into solution as they areoxidized (lose electrons) and become zinc ions (Zn2+). At the same time, copper ions (Cu2+) fromthe solution are reduced to copper atoms (Cu0) and deposited on the cathode. In this example,the zinc anode dissolves due to oxidation and the cathode increases in mass as layers of copperare added to it. 74Electrode Reaction Electrons Flow ExampleAnode Oxidation AWAY FROM Zn0(s) à Zn2+ (aq) + 2 eCathodeReduction TO Cu2+(aq) + 2e- à Cu0(s)Table 1. An example of a spontaneous redox reaction which produces a flow of electrons. Zincatoms at the anode lose electrons and become zinc ions in solution. These lost electronsflow to the cathode where copper ions in the solution gain electrons and plate out on thecathode as copper atoms.Can any redox reaction be used to generate an electric current? The quick answer is“No”. Only spontaneous redox reactions can be used to make a battery and a flow of electrons.See Table 1 in the Appendix to this lab for a list of the three most common types of batteriesand the redox reactions which power them. The common AA batteries provide ~1.5 volts.Rechargeable batteries use redox reactions which can be run in reverse to replenish the originalreactants lost during the spontaneous redox reactions.How to Put the Flow of Electrons to WorkIf the electrons flow directly from the oxidized ions to the reduced atoms, no usefulelectricity is produced. If you separate the oxidation and reduction into two differentcompartments however, the flow of electrons can be forced to flow through an external circuitand run other devices. These separate compartments are called half-cells and the reaction ineach half-cell is called a half-reaction because it makes up one half of the overall redox reactionin the battery. The solutions in the two cells are connected by a salt bridge soaked with positiveand negative ions not involved in the redox reaction. The movement of the ions from the saltbridge into the half-cells prevents any build-up of charge in the cells. A simple instrument whichmeasures the force of the flowing electrons between the half-cells is a voltmeter, or multimeter(Figure 1).
Figure 1. The basic components used to measurethe voltage generated by a battery. Two electrodesare immersed in solutions and connected by a saltbridge. The K+ and NO3- ions in the salt bridge donot react with any components of the battery, butflow between the two half-cells and keep themelectrically neutral. Not all batteries have a saltbridge. The electrons must flow in the externalcircuit through the voltmeter.http://www.gashalot.com/chem/dbhs.wvusd.k12.ca.us/AP-Chem/AP-Free-Resp98-quest8.GIF75Electroplating Converts Electrical Energy into Chemical EnergyJust as a battery can convert chemical energy into a flow of electrons or electricalenergy, the reverse can occur also. Electroplating uses an electrical current to drive a redoxreaction which is not spontaneous. Often, electroplating produces a thin coating over thecathode’s surface which you will see in a lab demonstration today. For example, steel by itselfcorrodes but if you put a thin coating of zinc on the steel through electroplating and corrosionis largely prevented. The steel is put in a bath of zinc salts (Zn+2) and acts as the cathode. Theanode can be any of a variety of metals. Turn on the power source and electrons flow from thepower source to the steel cathode and reduce the zinc ions (Zn+2) in the solution to zinc atoms(Zn0) which are deposited on the steel cathode. The higher the voltage and the longer you wait,the thicker the deposit of zinc will be on the steel. In electroplating, a power source providesthe flow of electrons to drive a redox reaction which doesn’t occur spontaneously.ExperimentPart 1. Are all redox reactions spontaneous?1. Retrieve a copper (Cu) and an iron (Fe) electrode. If they appear dull, polish them withsandpaper or steel wool. This removes any substances on the surface which might interferewith the redox reaction.2. Place a small amount of copper (II) sulfate (CuSO4) in a 50 mL beaker. Rest the ironelectrode in the beaker. Do you see any signs of a redox reaction? Is there any evidencethat the copper ions (Cu2+) in solution are being reduced (accepting electrons to becomeCu0) and being plated out onto the iron electrode?• If the copper ions are being reduced to copper atoms, this reduction can be written asthe half-reaction: Cu2+(aq) + 2 e- à Cu0(s) If you see evidence of reduction, write theappropriate half-reaction in Table 1 under ‘Reduction’.• If the copper ions are being reduced and gaining electrons, then the electrons must becoming from somewhere. That somewhere can only be the oxidation of the iron atoms(Fe0) to iron ions (Fe2+). Write the half-reaction for the oxidation of the iron atoms inTable 1 on the Data Sheet. Look up the reduction and oxidation potentials in theAppendix and calculate the net cell potential for this redox reaction.• Notice that if you add the two half-reactions together, the same number of electrons islost in the oxidation as is gained in the reduction.3. If the iron electrode in copper sulfate produces a spontaneous reaction, would you expectthe copper electrode in an iron (II) sulfate solution to produce a spontaneous reaction? Tryit and see. Write the appropriate half-reactions in Table 1 and calculate the standard netcell potential. 76Part 2. How to Get Spontaneous Redox Reactions to Do Work (and more work)The First BatteryThe first battery to produce a steady flow of electrons is credited to Count Volta in 1800.He called his creation an ‘artificial electric organ’ since his goal was to reproduce the electricshock given by a particular fish. Volta used silver and zinc plates as the electrodes and NaCl forthe salt bridge. Daniell later improved on Volta’s creation by replacing the silver electrode withcopper and using a saturated copper sulfate solution in place of the salt solution. Try recreatingit to see how many volts it can produce.1. Retrieve a multimeter. Zero it by turning the dial to the lowest voltage range of ‘DCV’ (directcurrent – volts). Hold the two metal leads together and adjust the dial to ‘0’, if needed.You’ll be reading the volts on the dial which corresponds to your voltage range setting. Seeit?2. Get 2 copper and 2 zinc electrodes, and lightly sand them until they’re bright and shiny.3. Let’s see what happens with a dry piece of filter paper and no salt bridge. Lay a copperelectrode on the bench and place a piece of dry filter paper on top of it. Now place the zincelectrode on top. Don’t let the two electrodes touch each other! Touch one lead of thevoltmeter to each of the electrodes in your battery. Record the voltage on the Data Sheet.4. Now let’s create a single-celled battery with a salt bridge. Place 2-3 drops of a 0.1 M coppersulfate (CuSO4) solution in the middle of the filter paper to make a salt bridge between thetwo metal electrodes. Make sure that there’s good contact between the copper sulfatesolution and the copper electrode. Don’t let the two electrodes directly touch each other oryou’ll short out your battery, meaning zero voltage.5. Touch one lead of the voltmeter to each of the electrodes in your battery. Record thevoltage for this single-celled battery on the Data Sheet. If you don’t see a voltage, switchthe leads. In this battery, the zinc atoms in the electrode are being oxidized and the copperions in solution are being reduced. Write the half-reactions for your battery in Table 2.5. Would the voltage be higher if you added another cell on top of the first? Try it! Put asecond copper electrode on top of the zinc electrode, followed by a copper sulfate-soakedfilter paper and a second zinc electrode. Again, make sure the electrodes don’t directlytouch each other. Measure and record the voltage. Cells combined one after another aresaid to be connected in series. Record the voltage under the two-cell battery on the DataSheet. How did it compare to the voltage from the single-celled battery?6. Disassemble your battery and look at the zinc electrodes. Does the surface have somedarker deposits on it? These are copper atoms building up on the zinc electrode. Overtime, they will interfere with the oxidation of the zinc atoms in the electrode. 77Part 3. Understanding Standard Reduction and Oxidation Potentials for Half-ReactionsBy testing different combinations of metals and aqueous solutions of ions, chemistshave been able to rank metals in order of their tendency to be reduced (accept electrons). Thisis called the standard reduction potential (E0red) and is measured in volts under a standard setof conditions - 25oC, 1 atm pressure and a 0.1 M ion concentration. All metals are compared tothe reduction of hydrogen ions (H+) to hydrogen gas (H20) which is arbitrarily set at 0.00 volts.Look at the Appendix to this lab for some standard reduction and oxidation potentials and findthe E0red and E0oxid for zinc. Notice how they are identical in voltage but opposite in sign? TheE0oxid for any atom is identical in magnitude but opposite in sign of the E0red. Let’s see how touse these standard potentials to predict which redox reactions will be spontaneous.1. Using the data in the Appendix, go back to Table 1 on the Data Sheet and fill in the E0redand E0oxid for your first redox reaction with the iron electrode and copper sulfatesolution. Add these two potentials together to get the standard net potential (E0net) forthis cell. A positive E0net indicates a spontaneous redox reaction and will generate avoltage. Does your data agree with what the E0net tells you about the spontaneity of thereaction and the magnitude of the voltage?2. Now use the data in the Appendix to write the standard potentials for the secondreaction in Table 1. This reaction was NOT a spontaneous redox reaction. Notice howthe sum of the E0red for iron ions and E0oxid for the copper atoms is negative (-0.78 volts).When the standard net potential is negative, the redox reaction is NOT spontaneous andcan’t be used to make a battery.3. Now determine the standard net cell potential for the redox reaction in Table 2.• How does this value compare to the voltage you measured for your single cell?• How does this value compare to the voltage you measured with two cells connectedin series? Is it roughly double? Connecting several spontaneous redox reaction halfcells in series can be used to boost the voltage generated by a battery.Part 4. Creating a Battery of Your Own DesignNow that you know a little something about how batteries work, it’s time to be creative.Here’s the premise for the rest of today’s lab. You and your lab partner have crash landed inthe middle of nowhere. The batteries in your cell phone are dead but you have the following 78items with which you hope to create a battery strong enough to power at least one phone callfor help. Design and create your own battery from the materials available to try and power upyour cell phone. We don’t expect you to know exactly what will work ahead of time, but youmay have some ideas. Just try all sorts of different combinations and test them to see whatworks.Here’s a list of some materials available for creating a battery:o Lemonso Electrodes: aluminum, zinc, copper, nickel and irono Pennieso Galvanized nails (covered by zinc oxide)o Voltmeter with alligator clips• Be sure to record what you build with a picture or written description and measure itsvoltage (or lack of voltage). Have your TA verify the voltage and initial it on your DataSheet.• The goal is to make at least one battery that generates a detectable voltage.• In addition, there are some extra points to be had for the group making the battery thatgenerates the highest voltage.Appendix to Electrochemistry LabLead-Acid Battery Half Reactions Rechargeable?Cathode Lead dioxide (Pb4+) Reduction: Pb4+ + 2 e- à Pb2+Anode Lead (Pb0) Oxidation: Pb0 à Pb2+ + 2e- YESElectrodes in: Sulfuric acid (H2SO4)Eocell 2.041 VAlkaline BatteryCathode: MnO2 and graphite Reduction: Mn2+ + 2e- à Mn1+Anode: Zinc and KOH Oxidation: Zn0 àZn2+ + 2 e- NOElectrodes in:Eocell 1.55 VNickel-Cadmium BatteryCathode Nickel oxyhydroxide (Ni4+) Reduction: Ni4+ + 2e- à Ni2+Anode Cadmium (Cd0) Oxidation: Cd0 à Cd2+ + 2e- YESElectrodes in: Potassium hydroxideEocell 1.30 VTable 1. Three common types of batteries and the spontaneous redox reactions that occur inthem. Alkaline batteries may be disposed of in the everyday trash but lead and nickelcadmiumbatteries should be taken to a hazardous waste disposal site. To read more aboutthe hazards of battery disposal read: http://www.ehso.com/ehshome/batteries.php. 79Selected Standard Reduction Potentials for Half-Reactions*Reduction Half-Reaction E0red (volts)Zn2+(aq) + 2e- à Zn0(s) -0.76Fe2+(aq) + 2e- à Fe0(s) -0.44Ni2+(aq) + 2e- à Ni0(s) -0.262H+(aq) + 2e- à H2(g) 0.00Cu2+(aq) + 2e- à Cu0(s) +0.34Table 2. Standard reduction potentials in order of their tendency to be reduced (acceptelectrons) under standard conditions at 25oC, 1 atm of pressure and 0.1 M ion concentration.Selected Standard Oxidation Potentials for Half-Reactions**Oxidation Half-Reaction E0oxid (volts)Zn0(s) à Zn2+(aq) + 2e- +0.76Fe0(s) à Fe2+(aq) + 2e- +0.44Ni0(s) à Ni2+(aq) + 2e- +0.26H2(g) à 2H+(aq) + 2e- +0.00Cu0(s) à Cu2+(aq) +2e- -0.34Table 3. Standard oxidation potentials in order of their tendency to be oxidized (lose electrons)under standard conditions at 25oC, 1 atm of pressure and 0.1 M ion concentration.