INTRODUCTION
Oxidation-reduction reactions may be thought of as a competition for electrons. We describe this competition in different ways at different levels of molecular structure. For isolated atoms, the ionization energy and electron affinity are measures of the tendency of atoms to lose or gain electrons. Atoms that form bonds have an electronegativity that measures the attraction of an atom to the electrons it competes for in its bonds. When we have a reaction, the tendency of chemical species to transfer electrons in oxidation-reduction reactions is measured by the electrochemical potential, E. Thus, when two chemical species compete for electrons, it is the difference between their electrochemical potentials that determines the driving force for electrons to migrate from one to the other.
If the two species are physically touching each other, you will often see one appear to dissolve and the other appear to precipitate – this is not usually a particularly useful reaction. If the species are physically separated so that the electron transfer must occur through an external connection, the reaction driving force will be transferred into a voltage or electromotive force (emf) for charge to flow in the circuit which can be easily measured (and utilized!). When the emf is created by a spontaneous chemical reaction, we call the system a galvanic (or voltaic) cell, and label the potential Ecell. Since it measures the driving force created by the tendency for reaction to occur, the value of Ecell must be related to the Gibbs free energy change, ΔG.
We can derive the relationship between the two functions by considering the units of measurement and the sign convention associated with each. ΔG measures reaction driving force as energy change per molar extent of reaction. Reading the symbol “[=]” as “has units of” we have [ ] rxn J G mol D = where ΔG is defined from the second law such that the sign must be negative for a spontaneous change. The unit for potential is the volt which gives the energy per unit of charge tending to cause the charge to flow. E [ ] J V C = = By historical convention, the voltage is positive when it causes a spontaneous flow of charge. The unit of charge is the coulomb (C) which is a derived S.I. unit equal to the amount of electric charge that would move past a point in a circuit in one second when the rate of current is one ampere (amp). The ampere, in turn, is defined as the rate of current flow such that if that rate were present in two parallel, one-meter lengths of wire separated by one meter, there would be a one Newton magnetic force between the wires. The number of coulombs equal to the charge of one mole of electrons is known as Faraday’s constant, F, which has a value of 96485 Coulombs per mole of electrons.
Now, the equation for relating E to ΔG can be created by starting with E in Joules per coulomb, converting coulombs to moles of electrons, multiplying by the number of moles of electrons, n, transferred in a reaction per molar extent and changing the sign. x x x( ) ( )x( )x( )x( ) rxn J C mol e G sign change E F n C mol e mol − − D = = − Or more conventionally: D = − G nFE Using electrochemical systems gives us two powerful tools for studying redox reactions. Measuring the voltage in the external circuit gives us the driving force or degree of spontaneity of the reaction. Faraday’s constant allows us to measure the extent of the reaction by measuring the current in the circuit and the reaction time. A galvanic cell, when physically constructed, uses a voltmeter which reads the flow of electrons. Voltmeters are set up to expect electrons to flow in a certain direction – if electrons are flowing in the opposite direction, it assigns the voltage as negative. The voltmeter expects electrons to flow from the black colored wire (usually referred to as a “lead”) through the voltmeter to the red wire.
When you are using one in real life, this helps you identify which side of your cell is the anode and which is the cathode, since you can’t physically see the electrons and ions flowing. Remember, the anode is where oxidation occurs, and the cathode is where reduction occurs. For the second part of this experiment, after measuring the potential of two cells containing a common half-cell, we will test whether the cell potentials are additive. That is, we will sum two reactions and their cell potentials then measure the actual potential for the net reaction and compare results. If the test is successful, it may serve as a justification for the use of a relative standard for the assigning of half-reaction potentials. In the experiment, the Cu/Cu2+ half-cell acts as the standard for comparison of the other two half-cells. In practice, the standard hydrogen electrode (SHE) usually serves as the reference half-cell and is assigned a relative potential of zero volts.
If you look up any standard reduction potentials, they will always be referenced to the SHE PROCEDURE
Part A.
Activity Series
This part of the experiment is provided by chemcollective at: http://chemcollective.org/vlab/106 Your first task is to arrange the metals copper, magnesium, zinc, silver and lead in order from strongest to weakest reducing agent. You will do this by performing a series of reactions between each of the metals and the cations of the other metals. You will want to create a table similar to the one given to the right (only likely larger) to track when a reaction occurs or doesn’t occur. You will perform these reactions by adding solid metal granules to solutions of metal nitrates. Do this by removing one metal solid container and the four metal nitrate solutions needed to test. Add 0.01 mol of the metal (indicate the mass of solid added in the column headers of your table – you don’t need to show this calculation, just provide the resulting mass) to the solution. If you follow the component amounts in the information tab, you should be able to tell whether or not a reaction occurred based on whether the solute metal ion changes – indicate this result in your table by writing “reaction” or “no reaction” (don’t count it as a reaction unless it goes to (or nearly to) completion).
Part B.
Galvanic Cells
Your second task it to provide numerical data to back up your data from part using the virtual lab setup found at: https://pages.uoregon.edu/tgreenbo/voltaicCellEMF.html This set up only has silver, zinc, and copper available, so you will only be finding information for those three half reactions. As you set up your cells, you will need to ensure that you set up spontaneous reactions – if you set up the reverse reaction you will need to reconstruct your cell to get a positive voltage. First, construct a standard cell with silver and copper half cells. Remember, a standard half-cell has a metal electrode in a solution of 1M ion. Record the voltage and which lead is on which electrode. You can think through which half-cell should be the anode and which should be the cathode based on your results from Part A if you don’t want to risk having to remake your cell. Next, construct a standard cell with copper and zinc half cells. Record the voltage and which lead is on which electrode. Last, construct a standard cell with silver and zinc half cells. You should be able to determine which lead goes on which half-cell from the prior two cells you built. Record the voltage and which lead is on which electrode.
Part C.
Nonstandard Galvanic Cells
This part uses the same simulator as the last part. Here you will make three cells out of nonstandard half cells, record their spontaneous voltages and note the anode and cathode. The cells you need to construct are: Zn/0.1M Zn2+ with Cu/0.1M Cu2+ Zn/0.1M Zn2+ with Ag/0.1M Ag+ Zn/2M Zn2+ with Cu/0.001M Cu2+ Part D. Concentration Cells This part uses the same simulator as the last part. Here you will make two cells. For the first, construct a standard copper half-cell at each electrode and measure the voltage. For the second, anode, you will use a copper electrode and 0.001 M copper nitrate solution.
For the cathode, you will use a copper electrode and 2 M copper nitrate solution. Record this voltage, but flip the sign – for some reason this simulator has the sign wrong for this cell (you can tell because the electrons are flowing the way they have in every other experiment, but the sign is negative instead of positive).
COMPLETING THE REPORT DATA
You should have 4 tables in your data section. Part A. Activity Series You should have the table you constructed during your experiment here. Part B. Galvanic Cells Organize your data from part B into a table like the one given below, where each line is represents a galvanic cell. This table should have 3 rows of data. Red Lead Black Lead Eo cell(Volts) Part C. Nonstandard Galvanic Cells Organize your data from part C into a table like the one from part B, but be sure to specify the concentration of each solution you used in the table. Part D. Concentration Cells Organize your data from part D into a table like the one from part C.
CALCULATIONS/ANALYSIS
There will be no summary of results for this lab. Instead, include the tables and questions given below in this order. Be sure to number the questions as they are numbered here. Part A. Activity Series Create a new table similar in structure to your data table for Part A of the lab. Include in your table which species was reduced and which was oxidized.
1. Underneath the table, give the ranking of the metals from strongest reducing agent to weakest reducing agent based on your observations. Explain your rankings. Part B. Galvanic Cells Create a table like the following example and fill in the cell-line notation description of the cell (remember, oxidation is always first, then reduction), the balanced redox reaction occurring, and the potential of each. Cell-Line Notation Reaction Potential
2. Which lead is connected to the half-cell undergoing oxidation? Which lead is connected to the half-cell undergoing reduction?
3. Do these reactions follow the same trend of reducing power you found in part A?
4. Do the potentials of these three reactions support that potential is a state function? Why or why not? (Hint: This is like a Hess’s law problem – can you add together 2 of the reactions to get the third reaction? Do the potentials add in the same way?)
5. If the reduction potential of the zinc half-reaction is set to 0V, then what is the reduction potential of the copper half-reaction? Of the silver half-reaction?
6. If the reduction potential of the copper half-reaction is set to 0V, then what is the reduction potential of the zinc half-reaction? Of the silver half-reaction?
7. If the reduction potential of the silver half-reaction is set to 2V, then what is the reduction potential of the copper half-reaction? Of the zinc-half reaction?
Part C.
Nonstandard Galvanic Cells
Create a table like the one in part B and fill in the cell-line notation for the cell (including concentrations since they are nonstandard!), the balanced redox reaction occurring, and the potential of each.
8. Is the potential of the zinc and copper cell with 0.1M concentrations the same as or different from the standard potential of the cell? Explain why this is the case.
9. Is the potential of the zinc and silver cell with 0.1M concentrations the same as or different from the standard potential of the cell? Explain why this is the case.
10. Show how to calculate the potential of the 2M Zn 0.001M Cu cell using the Nernst equation. DON’T FORGET TO USE EQUATION EDITOR. Nonformatted answers will not be graded. Part D. Concentration Cells Create a table like the ones in parts B and C and fill in the cell-line notation for the cell (including concentrations for the nonstandard cell), the balanced redox reaction occurring (include subscripts of “an” for the anode components and “cat” for the cathode components), and the potential of both cells.
11. Explain why the potential of the standard copper/copper cell makes sense.
12. Explain why the potential of the nonstandard copper/copper cell makes sense. Which concentration half-cell is the anode? Which concentration half-cell is the cathode?
13. Calculate the potential of the nonstandard copper/copper cell using the Nernst equation. DON’T FORGET TO USE EQUATION EDITOR. Non-formatted answers will not be graded.