1. Oxidation Loss of electrons

2. Reduction Gain of electrons

3. Rules for Assigning Oxidation Numbers 1.Free elements = 0 2.Simple ions = charge 3.F always -1 4.O nearly always -2, except when bonded to F, or in a peroxide 5.H nearly always +1, except when bonded to a metal. 6.Sum of the oxidation #’s in a neutral cmpd = 0 7.Sum of the oxidation #’s = charge for a polyatomic ion 8.For a covalent cmpd, the more electronegative element is assigned the negative oxidation # and vice versa

4. Redox Reactions Involve the transfer of electrons

5. Oxidation & Reduction Occur simultaneously. # of electrons lost = # of electrons gained.

6. LEO goes GER LOSS of ELECTRONS = OXIDATION. GAIN OF ELECTRONS = REDUCTION

7. Redox Reactions Single Replacement Synthesis Decomposition

8. Recall Formats Single Replacement: element + compound  new element + new compound Synthesis: 1 product Decomposition: 1 reactant

9. Identifying Redox Reactions Have to assign oxidation numbers to everything in the equation. The ones that change are redox.

10. Half-Reaction Shows either the oxidation or reduction reaction, including the electrons gained or lost.

11. Half-Reactions must obey 1)Conservation of matter 2)Conservation of charge

12. What does conservation of charge mean? Total charge on LHS of equation = Total charge on RHS of equation

13. Oxidation Half-Reation Electron term is on the product side.

14. H 2  2H + + 2e - Oxidation Half-Reaction

15. Reduction Half-Reation Electron term is on the reactant side.

16. Zn 2+ + 2e -  Zn Reduction Half-Reaction

17. Which half-reaction shows conservation of mass & conservation of charge? 1)S  S 2- + 2e - 2)Cl 2  Cl - + e - 3)Mn 7+ + 3e -  Mn 4+ 4)Ca 2+  Ca + 2e -

18. Pesky diatomics H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

19. What’s the problem with the Pesky Diatomics? H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 You have to keep the subscript in the half-reaction!

20. H 2  2H + + 2e - O 2 + 4e -  2O 2- H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 You have to keep the subscript in the half-reaction!

21. Which half-reaction is correct for reduction? a) Fe 2+  Fe + 2e - b) Fe + 2e -  Fe 2+ c) Fe  Fe 2+ + 2e - d) Fe 2+ + 2e -  Fe

22. Which half-reaction is correct for oxidation? a) Ca 2+ + 2e -  Ca b) Ca 2+  Ca + 2e - c) Ca + 2e -  Ca 2+ d) Ca  Ca 2+ + 2e -

23. Which half-reaction is correct for the reduction of O 2 ? a) O 2 + 2e -  O 2- b) O 2  2O 2- + 2e - c) O 2 + 4e -  2O 2- d) O 2  2O 2- + 4e -

24. Oxidizing Agent Helps something else get oxidized by itself being reduced.

25. Reducing Agent Helps something else get reduced by itself being oxidized.

26. 4 3 2 1 0 -2 -3 -4 1) You dig down with an oil rig 2) And if you’re lucky you strike oil & it shoots up

27. Steps in Balancing Redox Eqs. 1. Assign all oxidation #’s. 2. Use oil rig to figure out what’s oxidized & what’s reduced. 3. Write the half-reactions. 4. Add half-reactions, multiplying to adjust electrons if necessary. 5. Transfer coefficients & balance remaining elements.

28. How do you predict if a given redox reaction will occur? Use Table J! If the stand-alone element is above the similar element in Table J, the reaction will occur.

29. Li + AlCl 3  ? Compare Li with Al – both are metals. Li > Al so reaction occurs.

30. I 2 + NaCl  ? Compare I 2 with Cl – both are nonmetals. I 2 < Cl 2 so reaction DOES NOT occur.

31. 2 kinds of cells in electrochemistry?  Galvanic or Voltaic (NYS–electrochemical) Spontaneous rxn  Electrical energy  Electrolytic Electrical energy  Nonspontaneous rxn

32. Galvanic/Voltaic/Electrochemic al (NYS) Cell Uses a spontaneous reaction to produce a flow of electrons (electricity). Exothermic.

33. Galvanic Cell Redox reaction is arranged so the electrons are forced to flow through a wire. When the electrons travel through a wire, we can make them do work, like light a bulb or ring a buzzer. So the oxidation & reduction reactions have to be separated physically.

34. Galvanic/Voltaic/Electrochemical Use a spontaneous single replacement redox reaction to produce a flow of electrons. Electrons flow from oxidized substance to reduced substance.

35. Parts of a Galvanic Cell - 2 half-cells, each with a container, an aqueous solution, & an electrode connected by a - Wire and a - Salt Bridge

36. Electrode Surface at which oxidation or reduction half-reaction occurs.

37. Anode/Cathode in Galvanic Cell The anode is the metal that’s higher in table J. It’s more easily oxidized.

38. Direction of electron flow (wire) Anode to Cathode.

39. Direction of electron flow (wire) Galvanic Cell – remember opposites attract.

40. Direction of positive ion flow (salt bridge) Anode to Cathode.

41. Anode Electrode where oxidation occurs

42. Cathode Electrode where reduction occurs

43. Memory Aid AN Ox ate a RED CAT Works for ALL cells

44. Salt Bridge Allows for migration of ions between half-cells. Necessary to maintain electrical neutrality. Reaction will not proceed without salt bridge.

45. Negative Electrode / Galvanic Cell Electrode where electrons originate. (higher in table J)

46. Positive Electrode / Galvanic Cell Electrode that attracts electrons. (lower in table J)

47. Electron flow  Al = anode Pb = cathode wire Salt bridge Al +3 & NO 3 -1 Pb +2 & NO 3 -1 Positive ion flow  -  Draw galvanic cell with Al & Pb

48. Write the half-reactions for the previous cell Oxidation: Al  Al 3+ + 3e - Reduction: Pb 2+ + 2e -  Pb Metal electrode – Loses mass Aluminum ions in solution – concentration  Lead ions in solution – Concentration  Metal electrode – gains mass

49. Overall Rxn 2(Al  Al +3 + 3e - ) 3(Pb +2 + 2e -  Pb) _____________________________ 2Al + 3Pb +2  2Al +3 + 3Pb

50. What does [Zn 2+ ] mean? Concentration of Zn 2+ ions. Generally [ ] means concentration of whatever is inside [ ].

51. Uses of Galvanic Cells Galvanic Cells provide a voltage. They are a type of battery.

52. Uses of ELECTROLYTIC CELLS 1. Plate metals on other metals 2. Prepare column 1 and column 17 elements from compounds 3. Recharge batteries

53. Electrolytic Cell Uses a flow of electrons (electricity) to force a nonspontaneous reaction to occur. Endothermic.

54. How do you identify an electrolytic cell in a picture or diagram? 1. Its got a power supply – a battery! 2. You don’t have two separate containers.

55. What are the 2 kinds of electrolytic cells you are responsible for? 1) Fused salt cell 2) Plating Cell

56. How do you label the anode & cathode in an electrolytic cell? An Ox ate a Red Cat. Anode is still oxidation. Cathode is still reduction

57. How do you label the positive & negative electrodes in an electrolytic cell? Use the battery. The electrode attached to the positive pole of the battery is positive & vice versa.

58. What’s the memory trick for remembering the polarity of electrolytic cells? A POX on electrolytic cells! Anode Positive Oxidation

59. What is electroplating? Putting a very thin layer of one metal on top of another metal!

60. Battery +- 4) Anode = Oxidation Cu  Cu +2 + 2e - 5) Cathode = Object to be Plated = Reduction Cu +2 + 2e -  Cu Notice: Net reaction is just moving Cu around. Cu +2 and SO 4 -2 + - 1) See battery so it’s electrolytic! 2) Trace the + & - signs back to electrodes. 3) Element to be plated. Gains mass Loses mass

61. What is a fused salt cell? Electrolytic cell used to prepare group 1 & group 17 elements from their compounds (salts).

62. What does fused mean? MOLTEN

63. Fused Salt Cell: Opposites Attract! +- Molten NaCl Na + Cl - Anode = Oxidation. 2Cl -  Cl 2 + 2e - Cathode = Reduction. Na + + e -  Na Anode Cathode