1. Electrochemistry Chapter 17

2. Electrochemistry The study of the interchange of chemical and electrical energy.

3. What are some examples of Redox Reactions? What are some examples of redox reactions? forest fireforest fire rusting steelrusting steel combustion in auto enginecombustion in auto engine metabolism of food in the bodymetabolism of food in the body What everday uses depend on redox reactions? starting a carstarting a car calculatorcalculator digital watchdigital watch Cell phoneCell phone IiPodIiPod

4. 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

5. Redox Oxidation is the loss of electrons--oxidation number becomes more positive. Reduction is the gain of electrons--oxidation number becomes more negative.

6. OIL RIG Oxidation Is Loss. Reduction Is Gain.

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8. Redox metalsnonmetals oxidationreduction reducing agentsoxidizing agents metal ionsnonmetal ions reductionoxidation oxidizing agentsreducing agents

9. Redox Reactions Loss and gain of electrons must be simultaneous. Loss and gain of electrons must be equal. Why must the loss and gain of electrons be equal? Law of Conservation of Matter

10. Redox Reactions Redox reactions are reactions in which electrons are transferred. Decomposition and synthesis reactions may be redox. Single replacement reactions are always redox. Double replacement reactions are never redox. Combustion reactions are always redox.

11. Identifying Oxidation & Reduction in a Reaction Identify the element which is oxidized and the one which is reduced. 1. 2Mg (s) + O 2(g) ---> 2MgO (s) 2. 2Al (s) + 3I 2(s) ---> 2AlI 3(s) 3. 2Cu (s) + O 2(g) ---> 2CuO (s) 4. 2Cs (s) + F 2(g) ---> 2CsF (s)

12. 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 

13. Redox Oxidizing agent is the electron acceptor-- usually a nonmetal. Reducing agent is the electron donor--usually a metal. CH 4(g) + 2O 2(g) ----> CO 2(g) + 2HOH (g) Carbon is oxidized. Oxygen is reduced. CH 4 is the reducing agent. O 2 is the oxidizing agent. -4 +1 0 +4 -2 +1-2+1

14. Redox Reactions Identify the substance oxidized and the substance reduced as well as the oxidizing and reducing agents. PbO (s) + CO (g) ---> Pb (s) + CO 2(g) oxidizedreduced oxidizing agent reducing agent +2 -2 +2 -2 0 +4 -2 Carbon Lead PbO CO

15. Redox Reactions Identify the substance oxidized and the substance reduced as well as the oxidizing and reducing agents. 2PbS (s) + 3O 2(g) ---> 2PbO (s) + 2SO 2(g) oxidizedreduced oxidizing agent reducing agent +2 -2 0 +2 -2 +4 -2 sulfur oxygen O 2 PbS

16. Galvanic Cell A device in which chemical energy is changed to electrical energy.

17. Galvanic Cell A device in which chemical energy is changed to electrical energy. The basic parts are: anodecathode electrochemical solution porous disk or salt bridge

18. Anode and Cathode OXIDATION occurs at the ANODE. REDUCTION occurs at the CATHODE. AN OX RED CAT AN OX RED CAT http://www.mhhe.com/physsci/chemistry/essentialchemist ry/flash/galvan5.swf

19. A galvanic cell (Daniell Cell) involving Zn and Cu electrodes. This cell was the energy source for telegraphy during the War Between the States.

20. Cu-Zn (Daniell Cell) on the microscopic level.

21. Zinc electrode compared to a Standard Hydrogen Electrode (SHE). The Zn has a potential of 0.76 V.

22. Cell Potential Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

23. 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 E  = 0.20 V vs. SHE SHE = Standard Hydrogen Electrode

25. Cell Potential Calculations To Calculate cell potential using Standard Reduction Potentials: 1. One reaction and its cell potential’s sign must be reversed--it must be chosen such that the overall cell potential is positive. 2. The half-reactions must often be multiplied by an integer to balance electrons--this is 2. The half-reactions must often be multiplied by an integer to balance electrons--this is not done for the cell potentials.

26. Cell Potential Calculations Continued Fe 3+ (aq) + Cu (s) ----> Cu 2+ (aq) + Fe 2+ (aq) Fe 3+ (aq) + e - ----> Fe 2+ (aq) E o = 0.77 V Cu 2+ (aq) + 2 e - ----> Cu (s) E o = 0.34 V Reaction # 2 must be reversed.

27. Cell Potential Calculations Continued 2 (Fe 3+ (aq) + e - ----> Fe 2+ (aq) ) E o = 0.77 V Cu (s) ----> Cu 2+ (aq) + 2 e - E o = - 0.34 V 2Fe 3+ (aq) + Cu (s) ----> Cu 2+ (aq) + 2Fe 2+ (aq) E o = 0.43 V

28. Rb p404 2Ag (s) + Cu 2+ (aq) ----> Cu (s ) + 2Ag + (aq)

29. Line Notation solid  Aqueous  Aqueous  solid solid  Aqueous  Aqueous  solid Anode on the left  Cathode on the right Single line different phases. Double line porous disk or salt bridge. If all the substances on one side are aqueous, a platinum electrode is indicated. For the last reaction Cu(s)  Cu +2 (aq)  Fe +2 (aq),Fe +3 (aq)  Pt(s)

30. Free Energy and Cell Potential  G  =  nFE  n = number of moles of electrons Charge is measured in coulombs. F = Faraday = 96,485 coulombs per mole of electrons E= emf = potential (V) = work (J) / Charge(C)

31. Fe 3+ (aq) + Cu (s) ----> Cu 2+ (aq) + Fe 2+ (aq)

32. Calculate delta G for the following reaction Al +3 (aq)  Al(s) Mn +2 (aq)  Mn(s)

33. Common dry cell and its components.

34. Mercury battery used in calculators.

35. Concentration Cell...a cell in which both compartments have the same components but at different concentrations.

37. 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.

38. Calculation of Equilibrium Constants for Redox Reactions At equilibrium, E cell = 0 and Q = K.

39. Calculate Ecell for a galvanic cell based on the following Cd 2+ + 2e - ----> Cd Pb +2 + 2e - ----> Pb Where [Cd 2+ ] =.010 M Where [Cd 2+ ] =.010 M Where [Pb 2+ ] =.100 M Where [Pb 2+ ] =.100 M

42. Electrochemical Electrolytic Spontaneous Nonspontaneous Energy released Energy absorbed Cu 2+ (aq) + Mg (s) Cu (s) + Mg 2+ (aq) Electrochemical cell -- chemical energy to electrical energy. Electrolytic cell -- electrical energy to chemical energy.

43. Electrolysis...forcing a current through a cell to produce a chemical change for which the cell potential is negative.

44. Stoichiometry of Electrolysis -How much chemical change occurs with the flow of a given current for a specified time? 1 amp = 1 C/s

45. Electrolytic Calculations How many grams of copper can be plated out when a current of 10.0 amps is passed through a Cu 2+ solution for 30.0 minutes? (63.5 g/1 mol) (1 mol Cu/2 mole e - ) (1 mol e - /96,485 C)(10.0 C/s) (60 s/1 min) (30.0 min) = 5.94 g Cu

46. How many grams of nickel can be plated out when a current of 3.0 amps is passed through a Ni 2+ solution for 16.2 minutes?

47. Electrolytic Calculations How long must a current of 5.00 A be applied to a solution of Ag 1+ to produce 10.5 g of silver? (10.5 g Ag)(1 mol/107.86 g)(1 mol e - /1 mol Ag) (96,485 C/1 mole e - )(1 s/5.00 C)(1 min/60.0s) = 31.3 min = 31.3 min

48. Electrolysis of Water The electrolysis of water is: DC DC 2HOH (l) -----> 2H 2(g) + O 2(g)

49. Corrosion Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize. These are often called noble metals. About 1/5 of all iron and steel produced each year is used to replace rusted metal.

50. Self-protecting Metals Some metals such as aluminum, copper, and silver form a protective coating that keeps them from corroding further. The protective coating for iron and steel flakes away opening new layers of metal to corrosion.

51. Prevention of Corrosion Coating--painting or applying oil to keep out oxygen and moisture. Galvanizing--dipping a metal in a more active metal -- galvanized steel bucket. Alloying -- mixing metals with iron to prevent corrosion -- stainless steel. Cathodic protection -- attaching a more active metal. Serves as sacrificial metal--used to protect ships, gas lines, and gas tanks.

52. Corrosion of Iron Iron is oxidized at the anodic reaction and oxygen is reduced at the cathodic reaction. Dissolved ions are necessary to transfer electrons between the anodic and cathodic areas.

53. Schematic of Hall-Heroult process. Molten sinks and is tapped off at the bottom. Bauxite-cryolite mixture floats on top and is electrolyzed by the carbon electrodes.

54. Figure 17.11: Charles Martin Hall (1863- 1914)

55. Why did Napoleon III of France served his most honored guests with aluminum silverware?