1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Electrochemistry The study of the interchange of chemical and electrical energy.

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Review of Terms oxidation-reduction (redox) reaction: involves a transfer of electrons from the reducing agent to the oxidizing agent. oxidation: loss of electrons reduction: gain of electrons

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Half-Reactions The overall reaction is split into two half-reactions, one involving oxidation and one reduction. 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Oxidation: 5Fe 2+  5Fe 3+ + 5e 

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Galvanic Cell A device in which chemical energy is changed to electrical energy.

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Anode and Cathode OXIDATION occurs at the ANODE. REDUCTION occurs at the CATHODE.

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Cell Potential Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Batteries A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Electrolysis...forcing a current through a cell to produce a chemical change for which the cell potential is negative.

9. 8H + + MnO 4- + 5Fe 2+ → 5Fe 3+ + Mn 2+ + 4H 2 O If we place MnO 4- and Fe 2+ in the same container, the electrons are transferred directly when the reactants collide. No useful work is obtained from the chemical energy involved which is instead, released as heat. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 We can harness this energy if we separate the oxidizing agent from the reducing agent, thus requiring the electrons to transfer through a wire.

10. Salt bridge: it’s job is to balance the charge using an electrolyte. It connects the two compartments, ions flow from it, AND it keeps each cell neutral Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Porous Disk or cup: it allows both cells to remain neutral by allowing ions to flow.

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 emf and Work

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Figure 17.6 A Zn/Cu Galvanic Cell

13. ORIGIN OF STANDARD REDUCTION POTENTIALS Each half-reaction has a cell potential Each potential is measured against a standard, which is the standard hydrogen electrode SHE [consists of a piece of inert platinum that is bathed by hydrogen gas at 1 atm]. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13

14. standard conditions 1 atm for gases, 1.0M for solutions and 25°C for all (298 K) That means Ecell, Emf, or εcell become Ecell o, Emf o, or εcell o when measurements are taken at standard conditions. You’ll soon learn how these change when the conditions are non-standard! Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14

15. Interpreting a Table of Standard Electrode Potentials Elements that have the most positive reduction potentials are easily reduced (in general, non- metals) Elements that have the least positive reduction potentials are easily oxidized (in general, metals) The reduction potential table can also be used as an activity series. Metals having less positive reduction potentials are more active and will replace metals with more positive potentials. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16

17. Calculating Standard Cell Potential Symbolized by E°cell 1. Choose the appropriate half-reactions from a table of standard reduction potentials. Decide which element is oxidized or reduced using the table of reduction potentials. THE Metal with the MORE POSITIVE REDUCTION POTENITAL gets to be REDUCED. 2.Write the equation for the above half reaction first, along with it’s potential. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17

18. Calculating Standard Cell Potential Symbolized by E°cell 3.Write the equation for the other half reaction as an oxidation and write it’s oxidation potential as E o oxidation = -E o [this is now E°oxidation] 4.Balance the two half reactions **do not multiply the potentials by the numbers used to balance the electron transfer** Why not? A volt is equivalent to a J/coulomb which is a ratio. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18

19. Calculating Standard Cell Potential Symbolized by E°cell 5.Add the two half reactions and the voltages together. E o will be positive for the overall cell reaction, this indicates that the forward reaction is spontaneous. 6.E°cell = E°oxidation + E°reduction ° means standard conditions: 1atm, 1M, 25°C Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19

20. Egg-sample: copper-zinc cell Cu 2+ │Cu = 0.337 V and Zn 2+ │Zn = -0.763 V Cu 2+ + 2e - → Cu 0.337 V Zn → Zn 2+ + 2e - 0.763 V Electron transfer is balanced! Cu 2+ + Zn → Zn 2+ + Cu E o = 1.100 V Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20

21. Exercise 1 a. Consider a galvanic cell based on the reaction Al 3+ (aq) + Mg(s) → Al(s) + Mg 2+ (aq) Give the balanced cell reaction and calculate E° for the cell. A: 0.71 V Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Corrosion Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize. These are often called noble metals.

23. Origin of standard reduction potentials Each half reaction has a cell potential. Each potential is measured against a standard, which is the standard hydrogen electrode (SHE cell) The SHE consists of a piece of inert platinum that is bathed by hydrogen gas at one atm. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Free Energy and Cell Potential  G  =  nFE  n = number of moles of electrons F = Faraday = 96,485 coulombs per mole of electrons

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Concentration Cell...a cell in which both compartments have the same components but at different concentrations.

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 The Nernst Equation We can calculate the potential of a cell in which some or all of the components are not in their standard states.

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Calculation of Equilibrium Constants for Redox Reactions At equilibrium, E cell = 0 and Q = K.

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Fuel Cells...galvanic cells for which the reactants are continuously supplied. 2H 2 (g) + O 2 (g)  2H 2 O(l) anode: 2H 2 + 4OH   4H 2 O + 4e  cathode: 4e  + O 2 + 2H 2 O  4OH 

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Stoichiometry of Electrolysis 4 How much chemical change occurs with the flow of a given current for a specified time? current and time  quantity of charge  moles of electrons  moles of analyte  grams of analyte

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Standard Reduction Potentials The E  values corresponding to reduction half-reactions with all solutes at 1M and all gases at 1 atm. Cu 2+ + 2e   Cu E  = 0.34 V vs. SHE SO 4 2  + 4H + + 2e   H 2 SO 3 + H 2 O E  = 0.20 V vs. SHE