1. 18.5-18.6 Cell Potential, Free Energy, and the Equilibrium Constant – Cell Potential and Concentration

4. Let’s Try a Practice Problem!

6. Let’s Try Another! In the textbook, there was an example problem which resulted in the fact that the following redox reaction was non-spontaneous: I 2 (s) + 2Br - (aq)  2I - (aq) + Br 2 (l) Based on conceptual reasoning, which of the following best explains why I 2 does not oxidize Br - ? (a)Br is more electronegative than I; therefore, we do not expect Br - to give up electrons to I 2. (b) I is more electronegative than Br; therefore, we do not expect I 2 to give up electrons to Br -. (c)Br - is in solution and I 2 is a solid. Solids to not gain electrons from substances in solution. (a)Br is more electronegative than I. If the two atoms were in competition for the electron, the electron would go to the more electronegative atom (Br). Therefore, I 2 does not spontaneously gain electrons from Br -.

8. Let’s Try a Practice Problem!

10. Let’s Try Another!!!

11. Cell Potential and Concentration So far, we have calculated standard cell potentials. This is when the concentration of the reactants and products are 1 M. Now we must look at conditions that aren’t standard. One way to determine if the cell potential of a system at non-standard conditions will be greater or less than that at standard conditions would be to look use Le Chatelier’s Principle. For example, lets look at the following conditions: Under standard conditions the redox reaction between copper(II) and zinc yields a cell voltage of 1.10 V. Do you think the voltage would be greater or less than 1.10 V if the following changes occur? Zn(s) + Cu 2+ (aq, 2 M)  Zn 2+ (aq, 0.010 M) + Cu(s) E cell = ? The voltage would be greater than 1.10 V because according to Le Chatelier’s Principle, since the reactants are in higher concentration than they are at standard conditions, and products are in lower concentration than they were at standard conditions, the reaction has an even stronger tendency to proceed in the forward direction.

13. From the College Board!

15. Let’s Try a Practice Problem!

16. Concentration Cells As mentioned before, the Nernst equation can be used to calculate the cell potential under non-standard conditions…but since you are not going to be asked to use this equation, know this: A concentration cell is made from two half-cells with the same metal/ion combination, but with different ion concentrations. Electrons flow from the half-cell with lower ion concentration to the cell with the higher ion concentration in order to achieve equilibrium. This occurs because the electrons reduce the metals ions to their solid form, and increase the amount of ions in the dilute cell.

17. 18.5-18.6 pg. 906 #’s 66a., 68a. 72, & 80 Read 18.8 pgs. 890-897