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Oxidation

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Oxidation

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  1. Oxidation Loss of electrons

  2. Reduction Gain of electrons

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

  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 Conservation of matter 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. H2 2H+ + 2e- Oxidation Half-Reaction

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

  16. Zn2+ + 2e- Zn Reduction Half-Reaction

  17. Which half-reaction shows conservation of mass & conservation of charge? S  S2- + 2e- Cl2  Cl- + e- Mn7+ + 3e-  Mn4+ Ca2+  Ca + 2e-

  18. Pesky diatomics H2, N2, O2, F2, Cl2, Br2, I2

  19. What’s the problem with the Pesky Diatomics? H2, N2, O2, F2, Cl2, Br2, I2 You have to keep the subscript in the half-reaction!

  20. H2 2H+ + 2e-O2 + 4e-  2O2- H2, N2, O2, F2, Cl2, Br2, I2 You have to keep the subscript in the half-reaction!

  21. Fe2+ Fe + 2e- • Fe + 2e-  Fe2+ • Fe  Fe2+ + 2e- • Fe2+ + 2e-  Fe Which half-reaction is correct for reduction?

  22. Ca2+ + 2e- Ca • Ca2+  Ca + 2e- • Ca + 2e-  Ca2+ • Ca  Ca2+ + 2e- Which half-reaction is correct for oxidation?

  23. O2 + 2e- O2- • O2  2O2- + 2e- • O2 + 4e-  2O2- • O2  2O2- + 4e- Which half-reaction is correct for the reduction of O2?

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

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

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

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

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

  29. Compare Li with Al – both are metals. Li > Al so reaction occurs. Li + AlCl3 ?

  30. Compare I2 with Cl – both are nonmetals. I2 < Cl2 so reaction DOES NOT occur. I2 + NaCl  ?

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

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

  33. 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. Galvanic Cell

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

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

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

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

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

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

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

  41. Electrode where oxidation occurs Anode

  42. Electrode where reduction occurs Cathode

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

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

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

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

  47. Electron flow  wire Positive ion flow  Pb = cathode Al = anode Salt bridge -  Pb+2 & NO3-1 Al+3 & NO3-1 Draw galvanic cell with Al & Pb

  48. Oxidation: Al  Al3+ + 3e- Reduction: Pb2+ + 2e-  Pb Aluminum ions in solution – concentration  Metal electrode – Loses mass Write the half-reactions for the previous cell 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. Concentration of Zn2+ ions. Generally [ ] means concentration of whatever is inside [ ]. What does [Zn2+] mean?