Name (Attach data copies) Section A Describe your observations of any reactions between zinc and cupric sulfate or between copper and zinc sulfate. Write half-reactions describing any chemical changes you observed. Label the half reactions as oxidation or reduction. Indicate which species is the oxidizing agent and which is the reducing agent. Do you think this reaction has a large or small equilibrium constant? Describe the surface of the copper metal in contact with the zinc solution. Does the reaction of copper metal with zinc ion to produce copper (II) ion and zinc have a large or small equilibrium constant? Explain. Provide oxidation numbers to complete the following table. copper(ii) Species Zn ion Zinc ion Sulfate ion Oxidation # S= O= 1
Section B Position of red (+) Zn Position of black (- ) Zn Voltage reading (including sign) Calculate the standard oxidation potential (E o ox)for the oxidation of zinc. Calculate the equilibrium constant for the spontaneous redox reaction between cupric ion and zinc. Section C Position of red (+) Ag Position of black (-) Voltage (V) (including sign) E o oxid (V) E o red (V) 2
Arrange your calculated standard reduction potentials and corresponding half-reactions in order, starting with the most negative at the top of the table and ending with the most positive at the bottom. Include the standard reduction potential provided for. E (V) Half Reaction Compare the standard reduction potentials for and Ag provided in your lab manual to the ones you experimentally determined. Suggest relatively common materials that could be used to make a battery that will produce an approximate voltage of 1 V. Why are there difficulties in making the voltage measurements with magnesium and aluminum metals? Is the multimeter an effective tool in determining standard reduction potentials? Refer to your specific results as support your answer. 3
Section D How did you clean your nickel? What did you observe when the reaction ran for three seconds? What did you observe when dimethylglyoxime was added to the paper? Interpret this result. Which metal is the cathode and which is the anode in this experiment? Write the two half reactions and the chemical equation. Do you think that this electrolysis could be carried out with a standard 1.5 V battery? Whose picture is on the nickel? Section E Part 1. Attach graph of solution voltage vs. number of drops of Ce 4+. Perform a geometric construct on the graph to determine the number of drops to reach the equivalence point. Now utilize the number of drops from the graph and calculate the molarity of the Fe 2+ solution you titrated. Describe the endpoint of the titration (# drops added and color observed). Show calculations to determine the molarity of the Fe 2+ solution you titrated. Write the Nernst equation for the nonstandard conditions in your titration (without numbers). 4
Part II. Attach graph of solution voltage vs. number of drops of Cr 2 O 7 2-. Perform a geometric construct on the graph to determine the number of drops to reach the equivalence point. Now utilize the information from the graph and calculate the molarity of the Fe 2+ solution you titrated. Describe the endpoint of the titration (# drops added and color observed). Show calculations to determine the molarity of the Fe 2+ solution you titrated. Write the Nernst equation for the nonstandard conditions in your titration (without numbers). Section F What did you observe when the leads were attached to the battery and the Pb metal? Theoretically, what would happen if the two lead strips touched when the circuit is complete? What type of cell is a lead-acid battery, voltaic or electrolytic? Explain your answer. Write the half cell reactions for all redox reactions occurring in your lead-acid battery. 5