ELECTROCHEMISTRY

1. ELECTROCHEMISTRY The study of the interchange of chemical and electrical energy http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCell20.html

2. OIL RIG oxidation is loss, reduction is gain (of electrons) Oxidation--the loss of electrons; increase in charge Reduction-- the gain of electrons; reduction of charge

3. Oxidizing agent -the species that is reduced and thus CAUSES oxidation Reducing agent- the species that is oxidized and thus CAUSESreduction

4. Galvanic (voltaic cells) spontaneous chemical reactions (battery) Electrolytic cells non-spontaneous and require external e-source (DC power source) Two types of processes

5. Parts of the voltaic cell Anode- AnOx the electrode where oxidation occurs After a period of time, the anode may appear to become smaller as it falls into solution. Cathode- RedCat the electrode where reduction occurs After a period of time it may appear larger, due to ions from solution plating onto it. Ca+hode is always positive in galvanic cell.

6. Inert Electrodes • used when a gas is involved OR ion to ion involved such as: • Fe3+ being reduced to Fe2+ rather than Fe0 • made of Pt or graphite

7. Salt Bridge A device used to maintain electrical neutrality (balance the charge) in a galvanic cell. This may be filled with agar which contains a neutral salt or it may be replaced with a porous cup.Salt is often KNO3 because it will not form a precipitate.

8. Electron Flow always from anode to cathode (through the wire) FATCAT From the Anode To the CAThode

9. Standard cell notation (line notation) “Ion sandwich” in alphabetical order Anode metal | anode ion || cathode ion | Cathode metal Zn | Zn2+ (1.0M) || Cu2+ (1.0M) | Cu

10. If we place MnO4- and Fe2+ 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!

11. We can harness this energy if we separate the oxidizing agent from the reducing agent, thus requiring the e- transfer to occur through a wire! We can harness the energy that way to run a motor, light a bulb, etc.

12. Sustained electron flow cannot occur in this picture. Why not? As soon as electrons flow, a separation of charge occurs which stops the flow of electrons.

13. Salt Bridge It’s job is to balance the charge using an electrolyte [usually in a U-shaped tube filled with agar that has the salt dissolved into it before it gels].

14. Porous Disk or Cup … also allows both cells to remain neutral by allowing ions to flow.

15. Cell Potential Ecell, Emf, or Ecell — a measure of the electromotive force or the “pull” of the electrons as they travel from the anode to the cathode

16. Standard Reduction Potentials Each half-reaction has a cell potential. Each potential is measured against a standard which is the standard hydrogen electrode [consists of a piece of inert platinum that is bathed by hydrogen gas at 1 atm].

17. Standard Conditions 1 atm for gases 1.0M for solutions 25C for all (298 K) Ecell, Emf, or Ecell become Ecello , Emfo , or Ecello when measurements are taken at standard conditions.

18. Notice that 1.0 M solutions of each salt are used…Notice an overall voltage of 1.10 V for the process…

19. Reading the reduction potential chart 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).

20. The table can also be used to tell the strength of various oxidizing and reducing agents. It 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.

21. The MORE POSITIVE reduction potential gets to indeed be reduced IF you are trying to set up a cell that can act as a battery.

22. Write both equations AS IS from the chart with their voltages. Reverse the equation that will be oxidized and change the sign of the voltage [this is now Eoxidation]. Balance the two half reactions. **do not multiply voltage values**

23. Add the two half reactions and the voltages together. Ecell = Eoxidation + Ereduction ° means standard conditions: 1atm, 1M, 25C

24. Salt Bridge Bridge between cells whose purpose is to provide ions to balance the charge. Usually made of a salt filled agar (KNO3) or a porous cup.

25. ANIONS from the salt move to the anode while CATIONS from the salt move to the cathode!

26. Dependence of Cell Potential on Concentration Voltaic cells at NONstandard conditions -- LeChatlier’s principle can be applied. An increase in the concentration of a reactant will favor the forward reaction and the cell potential will increase. The converse is also true!

27. When cell is not at standard conditions, use Nernst Equation E = Eo – RT ln Q nF R = Gas constant 8.315 J/K mol F = Faraday constant Q = reaction quotient [productscoefficient]/[reactantscoefficient] E = Energy produced by reaction T = Temperature in Kelvins n = # of electrons exchanged in BALANCED redox equation

28. Rearranged, another useful form NERNST EQUATION: E = E° - 0.0592 log Q@ 25°C(298K) n

29. As E declines with reactants converting to products, E eventually reaches zero. Zero potential means reaction is at equilibrium [dead battery]. Also, Q = K AND G = 0 as well.

30. Notice the difference in the concentrations pictured at the left.

31. Because the right compartment contains 1.0 M Ag+ and the left compartment contains 0.10 M Ag+, there will be a driving force to transfer electrons from left to right.

32. Silver will be deposited on the right electrode, thus lowering the concentration of Ag+ in the right compartment. In the left compartment the silver electrode dissolves [producing Ag+ ions] to raise the concentration of Ag+ in solution.

33. Concentration cells e- e- Ag Ag • Electrons will move from 0.1 M Ag+ to 1 M Ag+ • Pull Ag+ from 1M side—plating metal • Taking e- from Ag electrode will produce more Ag+ on 0.1 M side 0.1 M Ag+ 1.0 M Ag+

34. Free Energy G = -nFE G = Gibb’s Free Energy [Reaction is spontaneous if ΔGis negative] n = number of mole of electrons F = 96,500 coulomb/mol e- E = cell potential 1 volt = Joule/coulomb

35. Summary of Gibb’s Free Energy and Cells • -Eo implies NONspontaneous • +Eo implies spontaneous (would be a good battery!) • E = 0, equilibrium reached (dead battery) • the larger the voltage, the more spontaneous the reaction • G will be negative in spontaneous reactions • K>1 are favored

36. Two important equations G = - nFE [“minus nunfe”] G = - RTlnK [“ratlink”] G = Gibbs free energy n = number of moles of electrons. F = Faraday constant 9.6485309 x 104 J/V (1 mol of electrons carries 96,500C ) E = cell potential R = 8.31 J/molK T = Kelvin temperature K = equilibrium constant [products]coeff/[reactants]coeff

37. Favored conditions Ecell > 0 G < 0 K>1

38. ELECTROLYSIS AND ELECTROLYTIC CELLS[NON spontaneous cells]Purpose: To produce purified forms of elements

39. Electrolytic cells Nonspontaneous Usually one cell Two of the same electrodes (graphite is a good choice) Must have a power source Use energy to create a chemical change

40. Voltaic cells are spontaneous Separated into two cells Is a battery Anode is negative Cathode is positive AnOx Red Cat Fat Cat Electrolytic cells are forced to work by using a power source Single container NEEDS a battery Anode is positive Cathode is negative AnOx Red Cat Fat Cat EPA (electrolytic positive anode)

41. If there is no water present and you have a pure molten ionic compound, then… The cation will be reduced (gain electrons/go down in charge). The anion will be oxidized (lose electrons/go up in charge).

42. Diagram an electrolytic cell of a molten solution of KI. Be sure to include anode, cathode, and overall cell reactions. Clearly state any observations and label electron flow.

43. To predict products in aqueous solutions: • No IA or IIA metal will ever be oxidized • No polyatomic ion will ever oxidize Instead– use water Memorize these equations: Oxidation 2H2O  O2 + 4H+ + 4e- Reduction 2H2O + 2e-  H2 + 2OH-

44. If water is present and you have an aqueous solution of the ionic compound, then… You’ll need to figure out if the ions are reacting, or the water is reacting. You can always look at a reduction potential table to figure it out.

45. Draw a diagram of an electrolytic cell containing an aqueous solution of KI. Be sure to label anode, cathode, electron flow and show all half-reactions and the overall reaction. Clearly state any observations at each electrode.

46. Faraday’s Law • The amount of substance being oxidized or reduced at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the cell. • Use dimensional analysis Conversions to memorize: • 1 volt = 1 Joule/Coulomb • 1 Amp = 1 Coulomb/second (current) • Faraday = 96,500 Coulombs/mole of e- • mole of e- come from balanced redox equation

47. A steady current of 1.00 ampere is passed through an electrolytic cell containing a 1 molar solution of AgNO3 and having a silver anode and a platinum cathode until 1.54 grams of silver is deposited. • How long does the current flow to obtain this deposit? • What weight of chromium would be deposited in a second cell containing 1-molar chromium III nitrate and having a chromium anode and a platinum cathode by the same current in the same time as was used in the silver cell? • If both electrodes were platinum in this second cell, what volume of O2 gas measured at STP would be released at the anode while the chromium is being deposited at the cathode? The current and the time are the same as in (a).