1. 1 Electrochemistry

2. 2 Oxidation-Reduction Rxns Oxidation-reduction rxns, called redox rxns, are electron-transfer rxns. So the oxidation states of 1 or more substances changes. Common redox rxns: iron rusting zinc reacting with acid combustion rxns batteries and electrolysis rxns photosynthesis

3. 3 Oxidation States Remember how to determine oxidation numbers for atoms?

4. 4 Oxidation States What are the oxidation states of all the elements in the following rxn? Zn(s) + HCl(aq) → ZnCl 2 (aq) + H 2 (g) What elements are changing oxidation states in the above? If an element loses electrons (oxidation state increases), then it is being oxidized and is the reductant or reducing agent. What element is being oxidized in the above? If an element gains electrons (oxidation state decreases), then it is being reduced and is the oxidant or the oxidizing agent. What element is being reduced in the above?

5. 5 Galvanic or Voltaic Cells As electrons are being transferred, we can use redox rxns to perform electrical work. If we have a spontaneous redox rxn, we can generate an electrical current. To do this, we have to separate the two half-rxns instead of having them take place in the same beaker.

6. 6 The two cells are connected two ways: An electrode in each cell are connected together with a conducting wire. An electrode is simply a piece of metal (or graphite), which may or may not be a reactant or product. A salt bridge allows ions to flow through from one cell to the other, so charge balance is maintained (overall neutral slns).

7. 7 Galvanic Cell Connections

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9. 9 Galvanic Cell

10. 10 Galvanic Cell

11. 11 Shorthand Notation for Galvanic Cells We saw the galvanic cell for the following rxn (Zn anode electrode, Pt cathode electrode): Zn(s) + H + (aq) --> Zn 2+ (aq) + H 2 (g) There is a shorthand notation for this galvanic cell:

12. 12 Cell Potentials What drives the rxn in a galvanic cell? Or what forces the electrons to move from the anode to the cathode? The driving force is an electrical potential called the electromotive force or emf. The emf for a cell is also simply called the cell potential, E, E cell, or the cell voltage.

13. 13 Cell Potentials The E cell potential is read with a voltmeter which is connected to the cell circuit. –The - terminal on the voltmeter must be connected to the cell anode, while the + terminal on the voltmeter must be connected to the cell cathode. –In the lab, this is how the direction of a spontaneous cell rxn is determined: when the voltmeter gives a + reading, the connections are correct, and the anode and cathode are identified. If the connections are incorrect, a 0 or negative voltage reading is obtained (as rxn occurs in reverse).

14. 14 Units for Electrochemistry The unit for the cell voltage is volts, V Also: 1V = 1J/C (C is coulombs, electric charge) 1C = 1amps or 1C = 1As (amp or A is current in amperes) 1watt = 1W = 1J/s (watt is the power) Since the voltage is related to energy as well as electric charge, there MUST be a relationship between a cell potential and energy!

15. 15 Δ G and E There is a mathematical relationship between the 2:  Δ G = -nFE where F = 96,500 C/mol e - and n = # mol e - transferred in cell Notice that if the cell rxn is spontaneous, the E is + but the Δ G is – as we would expect!

16. 16 Δ G and E If the cell is being run under standard state conditions (1 atm for gases, 1M for solutions, and at a specified temperature, usually given as 25°C), then: Δ G° = -nFE° Determine Δ G° for the following cell rxn if E° = 0.92V

17. 17 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red The E° for a cell rxn depends on both the E° for the anode half cell rxn and the E° for the cathode half cell rxn. If we know the E° for both half cells, we can calculate the overall E° for the cell:

18. 18 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red But we can’t measure an E° for a half cell as there’s no rxn unless 2 half cells are coupled! So we have a reference half-cell rxn which has been assigned an E° half-cell potential of 0V. This is the Standard Hydrogen Electrode or SHE. The SHE half-cell may act as the anode or cathode depending on what the other half-cell is.

19. 19 SHE Half-Cell SHE Half-Cell The rxn is and half-cell notations are:

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21. 21 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red So half-cell standard potentials are measured relative to the SHE half-cell. There are Tables of standard potentials for half-cells, readily available (p 830 and Appendix ii). Looking at these Tables you should notice that all of the E° values are for the reduction half- cell rxn, E° red

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23. 23 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red Notice also that some of the values are -, while others are +. This tells you the relative strengths of oxidizing and reducing agents. The higher the E° red value, the stronger the oxidizing agent in the half-cell.

24. 24 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red For example, which is the stronger ox. agent, Fe 3+ or Fe 2+ ?

25. 25 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red The E° red values also gives you the tendency of 2 coupled half-cell rxns to be spontaneous. They also tell you which half-cell will be the anode (oxidation) and which will be the cathode (reduction). The half-cell rxn with the higher E° red value (more + value) WILL be the cathode. So of course, the more negative E° red value will be the anode.

26. 26 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red Let’s show this using Eq 1) and the fact that E° red = - E° ox as they are just the reverse of each other. Rearrange Eq 1) to reflect that Table values are always given as E° red values.

27. 27 Standard Reduction Potentials for Half Cells, E° red Standard Reduction Potentials for Half Cells, E° red We will use Eq 2) to find the overall E° cell

28. 28 Before you start problems, there is a comment to make about E° E° is an INTENSIVE properties, so that means that E° DOES NOT depend on the amount of a substance. This means that E° is the same whether you have 1 mol or 100 mol! What does change is the TIME that the rxn lasts; the more moles, the LONGER the rxn lasts; but the standard voltage is constant. E° cell

29. 29 E° cell Although this seems weird, look at the units of a V: J/C So a V is a ratio of 2 extensive properties; but both change by the SAME amount as the mol changes. So there is NO net change for E° What does this mean for you? You ignore stoichiometry when calculating an E° cell from 2 half-cell E° red values.

30. 30 Non Standard State Conditions What if you aren’t at standard state? You already learned the relationship between Δ G and Δ G°: Δ G = Δ G° + RTlnQ But Δ G = -nFE and Δ G° = -nFE° so: -nFE = -nFE° + RTlnQ Dividing both sides by -nF: E = E° - (RT/nF)lnQ This is the Nernst Equation.

31. 31 Nernst Equation The Nernst Equation is also written in terms of log: And if the temperature is 25°C, this simplifies to: Obviously, you can only use the above at 25°C, otherwise use the complete Nernst Eq.

32. 32 Nernst Equation How often are we at standard state? And if we start at standard state, do we stay at standard state? What changes in the Nernst equation as a reaction progresses? So how does E respond to this change? And when the rxn is complete (or at equilibrium), what will E be?

33. 33 The Relationship between E° and K At equilibrium, E = 0 and Q = K. So the Nernst Equation becomes:

34. 34 The Relationship between E° and K Or at 25°C: Solving for K gives: So, as you would expect, if E° > 0, then K > 1.