1. Chapter 21

2. the study of the production of electricity during chemical rxns and the changes produced by electrical current. Electrochemical reactions are oxidation-reduction reactions. Oilrig: oxidation loss of electrons, reduction gaining of electrons 1. Oxidation = loss of electrons a. the substance oxidized is the reducing agent 2. Reduction = gain of electrons a. the substance reduced is the oxidizing agent

3.  Metals conduct electric currents well.  metallic conduction  Positively charged ions, cations, move toward the negative electrode.  Negatively charged ions, anions, move toward the positive electrode. 3

4. Electrolytic cells - nonspontaneous chemical reactions Voltaic or galvanic cells - spontaneous chemical reactions The two parts of the reaction are physically separated. –oxidation occurs at one cell –reduction occurs in the other cell

5. Cathode - electrode at which reduction occurs (red cat) Anode - electrode at which oxidation occurs (an ox) Inert electrodes do not react with the liquids or products of the electrochemical reaction. Graphite and Platinum are common inert electrodes.

6. Use electrical energy to force nonspontaneous (non thermodynamically favored) chemical reactions to occur. Process called electrolysis. Used in: –plating of jewelry and auto parts –electrolysis of chemical compounds Electrolytic cells consist of a: –container for reaction mixture –electrodes immersed in the reaction mixture –source of direct current

7. In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode). In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

8. The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell. During electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent. Corresponds to the passage of one mole of electrons through the electrolytic cell.

9. Amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.  1 faraday of electricity = 6.022x10 23 e -  1 faraday = 6.022x10 23 e - = 96487 coulombs  1 eq. of oxidizing agent= gain of 6.022x10 23 e -  1 eq. of reducing agent = loss of 6.022x10 23 e -

10. Ex. 1) Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes. (1 ampere = 1 coulomb per second)

11. 11  Ex. 2) Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in Ex. 1.

12.  Electrochemical cells in which a spontaneous chemical reaction produces electrical energy. In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).  Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference Examples: Car & flashlight batteries

13. 13  Half-cell contains the oxidized and reduced forms of an element (or other chemical species) in contact with each other.  Simple cells consist of:  two pieces of metal immersed in solutions of their ions  wire to connect the two half-cells  salt bridge to ▪ complete circuit ▪ maintain neutrality ▪ prevent solution mixing

14. 14  Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Zn strip immersed in 1.0 M zinc (II) sulfate wire and a salt bridge to complete circuit  Initial voltage is 1.10 volts

15. 15

16. 16  In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

17. Short hand notation for voltaic cells Zn-Cu cell example

18. 18  Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Ag strip immersed in 1.0 M silver (I) nitrate wire and a salt bridge to complete circuit  Initial voltage is 0.46 volts

19. 19

20. 20  Compare the Zn-Cu cell to the Cu-Ag cell Cu electrode is cathode in Zn-Cu cell Cu electrode is anode in Cu-Ag cell  Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

21. 21  Demonstrates that Cu 2+ is a stronger oxidizing agent than Zn 2+ Cu 2+ oxidizes metallic Zn to Zn 2+  Ag + is is a stronger oxidizing agent than Cu 2+ Ag + oxidizes metallic Cu to Cu 2+  Arrange these species in order of increasing strengths

22. 22  Establish an arbitrary standard to measure potentials of a variety of electrodes  Standard Hydrogen Electrode (SHE)  assigned an arbitrary voltage of 0.000000… V

23. 23  Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials.  Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials.  Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.

24. 24  Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.  For example, the half-reaction for the standard potassium electrode is:  The large negative value tells us that this reaction will occur only under extreme conditions.

25. 25  Compare the potassium half-reaction to fluorine’s half-reaction:  The large positive value denotes that this reaction occurs readily as written.  Positive E 0 values tell us that the reaction tends to occur to the right  larger the value, greater tendency to occur to the right  Opposite for negative values

26. 1. Choose the appropriate half-reactions from a table of standard reduction potentials. 2. Write the equation for the half-reaction with the more positive E 0 value first, along with its E 0 value. 3. Write the eqn for the other half-reaction as an oxidation with its oxidation potential, reverse the tabulated reduction half-reaction and change the sign of the tabulated E 0. 4. Balance the electron transfer. 5. Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E 0 cell is positive, which indicates that the forward reaction is spontaneous.

27. 27  Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously.  Ex. 3) Will silver ions, Ag +, oxidize metallic zinc to Zn 2+ ions, or will Zn 2+ ions oxidize metallic Ag to Ag + ions? What is the overall value for E o ?  E 0 values are not multiplied by any stoichiometric relationships in this procedure.

28. 28  Ex.4) Will tin(IV) ions oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize tin(II) ions to tin(IV) ions in acidic solution? What is the overall value for rxn?

29. 29  Standard electrode potentials are determined at thermodynamic standard conditions. 1 M solutions 1 atm of pressure for gases liquids and solids in their standard states temperature of 25 0 C  Potentials change if conditions are nonstandard.  Nernst equation describes the electrode potentials at nonstandard conditions.

30. 30

31. 31  Substitution of the values of the constants into the Nernst equation at 25 0 C gives:

32.  used to calculate electrode potentials and cell potentials for concentrations other than standard –state values. (pg 877)  Ex. 5) Calculate the cell potential for the following electrochemical cell. If the [Sn 2+ ] = 4.5 x 10 -1 M and [Ag + ] = 0.110 M  Sn (s) + 2Ag +  Sn 2+ + 2Ag (s)

33. 33  The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes.  Ex. 6) Calculate the initial potential of a cell that consists of an Fe 3+ /Fe 2+ electrode in which [Fe 3+ ]=1.0 x 10 -2 M and [Fe 2+ ]=0.1 M connected to a Sn 4+ /Sn 2+ electrode in which [Sn 4+ ]=1.0 M and [Sn 2+ ]=0.10 M. A wire and salt bridge complete the circuit.

34. 34  From previous chapters we know the relationship of  G 0 and K for a reaction.

35. 35  The relationship between  G 0 and E 0 cell is also a simple one.

36. 36  You can combine these two relationships into a single relationship to relate E 0 cell to K.

37. 37  Ex. 7) Calculate the standard Gibbs free energy change,  G 0, at 25 0 C for the following reaction.

38. 38  Calculate E 0 cell using the appropriate half-reactions.

39. 39  Now that we know E 0 cell, we can calculate  G 0.  The negative value tells us that the reaction is spontaneous as written.

40. 40  Ex. 8) Calculate the thermodynamic equilibrium constant for the reaction in Ex. 7 at 25 0 C.

41. 41  Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO 2, O 2, and H 2 O.

42.  The oxidation of most metals by oxygen is spontaneous. Many metals develop a thin coating of metal oxide on the outside that prevents further oxidation  The presence of a salt accelerates the corrosion process by increasing the ease with which electrons are conducted from anodic to cathodic regions

43. 43  Some examples of corrosion protection. 1 Plate a metal with a thin layer of a less active (less easily oxidized) metal. 2. Connect the metal to a sacrificial anode, a piece of a more active metal.

44. 44 3 Allow a protective film to form naturally.

45. 45 4 Galvanizing, coating steel with zinc, a more active metal. 5. aint or coat with a polymeric material such as plastic or ceramic.

46. 46  As a voltaic cell discharges, its chemicals are consumed.  Once chemicals are consumed, further chemical action is impossible.  Electrodes and electrolytes cannot be regenerated by reversing current flow through cell.

47. 47  One example is flashlight, radio, etc. batteries.  Container is made of zinc  acts as an electrode  Graphite rod is in center of cell  acts as the other electrode  Space between electrodes is filled with a mixture of:  ammonium chloride, NH 4 Cl  manganese (IV) oxide, MnO 2  zinc chloride, ZnCl 2  porous inactive solid

48. 48  As current is produced, Zn dissolves and goes into solution as Zn 2+ ions.  Zn electrode is negative (anode).

49. 49  Secondary cells are reversible, rechargeable.  Electrodes can be regenerated  One example is the lead storage or car battery.

50. 50  Electrodes are two sets of lead alloy grids (plates).  Holes in one grid are filled with lead (IV) oxide, PbO 2.  Other holes are filled with spongy lead.  Electrolyte is dilute sulfuric acid.  When battery is discharging, spongy lead is oxidized to lead ions and the plate becomes negatively charged.

51. 51  Cell reaction for a discharging lead storage battery is

52. 52  This battery can be recharged by passing electricity through the cell in the opposite direction.  Voltaic cell is converted to galvanic cell.  At lead electrode, lead ions are reduced to lead atoms.

53. 53  Rechargeable cell used in calculators, cameras, watches, etc.  Discharge half-reactions are:

54. 54  Fuel cells are batteries that must have their reactants continuously supplied in the presence of appropriate catalysts.  Hydrogen-oxygen fuel cell is used in the space shuttle.  Hydrogen is oxidized at anode  Oxygen is reduced at cathode

55. 55  The reaction combines hydrogen and oxygen to form water.  Drinking water supply for astronauts.  Very efficient energy conversion 60-70%

56. 56 1. What are the explosive chemicals in the fuel cell that were aboard Apollo 13? Which tank exploded? 2. Some of the deadliest snakes in the world, for example the cobra, have venoms that are neurotoxins. Neurotoxins have an electrochemical basis. How do neurotoxins disrupt normal chemistry and eventually kill their prey?