A. Using the reduction half-reactions in the table below, write a balanced net-ionic equation representing the overall reaction that takes place as the cell operates. B. Determine the value of E° for the cell. C. Determine the value of ΔG° for the cell reaction.

2. Using the available materials, design a spontaneous electrochemical cell at standard conditions with initial E° of +0.90V. a. Label the anode and cathode metals b. Label the anode and cathode solutions c. Draw arrows through the wire to indicate the direction of electron flow d. Indicate which ion from the salt bridge will migrate into each chamber e. Write the balanced half-reactions for each electrode and the overall balanced equation f. Circle whether each electrode will increase or decrease in mass g. Indicate the value of n, the number of electrons/mole rxn, in the balanced equation h. Calculate AG° electron flow KNO anode metal cathode metal anode solution _cathode solution ion migration anode half reaction cathode half reaction electrode mass: increases or decreases electrode mass increases or decreases overall balanced equation: n=

2. For the following concentration cell, use the Nernst equation to determine the expected voltage. Also indicate the oxidation reaction, the reduction reaction, the cell reaction, and the overall expected EMF of the cell using the tables from powerpoint slides 11 and 12 and the Nernst equation. Once you have determined the cell reaction, write the conventional cell notation. Zn and Zn+2 (0.25M) mixed with Cu and Cu+2 (0.35M)

3. Calculate the free energy, AG°, for each electrochemical cell given the following information and determine if the reaction is spontaneous or nonspontaneous. T= 298 K a. E cell =+ 2.10 V, n= 2 b. Каg— 2.2 х 10-3 c. E cell =- 0.677 V, n = 3 =