1. Chapter 14 Electrochemistry

3. Basic Concepts Chemical Reaction that involves the transfer of electrons. A Redox reaction. Loss of electrons – oxidation Gain of electrons – reduction Oxidizing agent. A species that takes electrons. Reducing agent. A species that gives electrons.

4. Basics Na(s) + H + -> Na + + H 2 (g) Sodium is a reducing agent Hydrogen ion is the oxidizing agent.

5. Basics We are donating and gaining electrons. If we could use these electrons perhaps we could do some useful work. If we can make the electron travel in an electrical circuit then the amount of current can be measured. Current is related to reaction rate or amount of reaction Potential is related to free energy change of the reaction.

6. Electron Charge q used to denote. Unit is Coulombs (C) Charge on a single electron is 1.602x10 -19 C which will allow us to determine the charge on a mole of electrons. 1.602x10 -19 C * 6.022x10 23 mol -1 ) = 96490 C mol -1 This is called the Faraday Constant q = nF n is the number of moles

7. Current Charge flowing through a circuit One ampere, the charge of one coulomb per second flowing past a given point.

8. Electrodes The interface between a solution and an electrical circuit. Can be actively involved or just serve as a source or sink for electons.

9. Electrical Potential Work required when moving and electric charge from one point to another. Electrical potential (E) is measured in Volts (V). Work is a measure of energy, measured in joules (J). Work = E * q Joules volts coulombs

10. Free Energy Maximum amount of work that can be done on the surroundings is equal to the Gibbs free energy change. then  G = -work = -Eq Or  G = -nFE

11. Ohm’s Law Current is proportional to the potential and and inversely proportional to the resistance. I = E/R

12. Electric Circuit

13. Power Work done per unit time. Unit is the J/s which is know as the watt (W). P = work/sec = Eq/sec = E(q/sec) = EI P = EI = I 2 R = E 2 /R

14. Galvanic Cells Spontaneous chemical reaction used to generate electricity. An example might be

16. Voltmeter A device to measure electrical potential. When electrons tend to flow into the negative terminal then a positive voltage is measured. In this cell 2 AgCl (s) + 2 e - = 2 Ag + 2 Cl - (aq) Red Cd (s) + = Cd 2+ + 2 e - Oxidation Cd (s) + 2 AgCl (s) = Cd 2+ + 2 Cl - Net For this reaction we have a  G of -150 kJ/mole per mole of Cd oxidized.

17. Potential of this System  G = -150 kJ/mole then we have E = -  G/nF = -150 x 10 3 J / (2 mol)(9.649x10 4 C/mol) E = + 0.777 J/C = +0.777 V

18. Cathode/anode Cathode electrode where reduction occurs Anode electrode where oxidation occur Put both terms in alphabetical order to remember

19. Salt Bridge Any bridge in upstate New York in the winter. Used to isolate the half cells so the work can be forced out into an external circuit. The following cell has a problem.

21. What is it? The silver ions in solution can go directly to the cadmium electrode surface and be reduced there. We need to put in a barrier to rapid ionic transfer.

23. What about this cell

24. Isn’t this cute Chemistry paper dolls?

25. Line Notation - Instead of Having to Draw the Cells | phase boundary || salt bridge For First Cell Cd(s) | CdCl 2 (aq) | AgCl(s) | Ag(s) For Second Cell Cd(s) | Cd(NO 3 ) 2 (aq) || AgNO 3 (aq) | Ag(s)

26. A Word of Connectors (Two common in USA)

27. Standard Potential E o The energy to a half cell at standard conditions (1 M and 25 C) Let us look at the reduction of silver ion. Ag + + e - = Ag(s) We will compare this to a fixed reference. That is the SHE or NHE Standard or Normal Hydrogen Electrode. H + (aq, A=1) + e - = ½ H 2 (g, A = 1)

29. SHE - All other redox couples are compared to this half cell. It is assigned a value of 0.000 V In our cell the left side electrode (Pt) is attached to the negative terminal. (Reference) Value of E are collected into Tables (Appendix H)

31. Nernst Equation For the half reaction aA + ne - = bB E o = is the standard Potential R = gas constant (8.314472 (V*C)/(k*mol) T = Temp (K) N = # of electrons in the half reaction F = Faraday A = Activity

32. We will often lump the constants and assume 25 C Nernst equation (25 C and converting to log 10

33. Complete Reaction E = E + - E - for full cell Steps Write both half cells as reductions, make electrons equal Half cell connected to positive terminal is E + Other half cell is E - Net voltage is from the above equation Balance equation (reversing the left half reaction and adding to other half cell) E > 0 spontaneous as written E < 0 spontaneous in reverse

34. E o and K

35. Cells as Chemical Probes Equilibria between the half cells Equilibria within each half cell

36. A Probe Cell

37. Probe Cell Right side: We have our K sp equilibrium The electrochemical reaction under this is AgCl(s) + e - = Ag(s) + Cl - (aq, 0.10 M) E o = 0.222 v Left side: We have our K a for the weak acid. The electrochemical reaction 2 H + (aq) + e - = H 2 (g, 1.00 bar) E = 0.00, but H + is not fixed at 1 M so E varies with H +

38. E o’ Formal Potential Since so many redox couples exist in the body and many have H + we modify the potential that we use to pH 7. (A little more reasonable than 1 M acid.

40. Homework 14- 4 13, 14, 15 and 27