Basics of Electrochemistry: Redox Reactions, Current, and Potential

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+ + H2(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 * 6.022x1023 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 DG = -work = -Eq • Or DG = -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 = I2R = E2/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) + = Cd2+ + 2 e- Oxidation • Cd (s) + 2 AgCl (s) = Cd2+ + 2 Cl- Net • For this reaction we have a DG of -150 kJ/mole per mole of Cd oxidized.

17. Potential of this System • DG = -150 kJ/mole then we have • E = - DG/nF = -150 x 103 J / (2 mol)(9.649x104 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) | CdCl2(aq) | AgCl(s) | Ag(s) • For Second Cell • Cd(s) | Cd(NO3)2(aq) || AgNO3(aq) | Ag(s)

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

27. Standard Potential EoThe 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 H2 (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 Eo = 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 log10

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. Eo 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 Ksp equilibrium • The electrochemical reaction under this is • AgCl(s) + e- = Ag(s) + Cl- (aq, 0.10 M) • Eo = 0.222 v • Left side: • We have our Ka for the weak acid. • The electrochemical reaction • 2 H+(aq) + e- = H2 (g, 1.00 bar) • E = 0.00, but H+ is not fixed at 1 M so E varies with H+

38. Eo’ • 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