2. Oxidation and Reduction Reactions and Electrochemistry Oxidation and Reduction Reactions and Electrochemistry “The Ubiquitous Electron”

3. l Redox and Iron in your Body Redox and Iron in your Body

5. Types of Reactions 1. Ions or molecules react w/ no apparent change in electronic structure (ex. Double displacement) 2. Ions or atoms undergo changes of electronic structure, the way e - transfer or the way atoms share e - changes.

6. Oxidation- Reduction Reaction Definition: l Chemical change that occurs when electrons are transferred between reactants l All oxidation reactions are accompanied by reduction reactions l Important: in the corrosion of metals, sources of energy, life processes

7. Oxidation l Part of the redox rxn in which electrons are removed or apparently removed from an atom (loss of electrons  atom gets more positively charged)

8. l Movie Movie

9. Reduction l Part of the redox rxn in which electrons are added or apparently added to an atom (gain of electrons  atoms get more negatively charged)

10. l Movie Movie

11. OIL RIG l Oxidation Is Losing l Reduction Is Gaining

12. LEO the lion goes GER l Loss of Electrons in Oxidation l Gain of Electrons in Reduction

13. Ionization or Solvation = the process of surrounding solute particles with solvent particles to form a solution l Video Video l “like dissolves like”

14. Net Ionic Equations l When reactions take place in water chemists write the equation in ionic form (particles ionize – break into their ions in water) l Chemists only write down the ions that take part in the reaction l Spectator ions- ions that aren’t involved in the reaction (chemists don’t write these) l Makes rxn easier to balance

15. Cu + NO 3 -1  Cu +2 + NO

16. l Show chemistry connections video: 7:36 minutes into video, found in redox folderredox folder

17. Rules for Assigning Oxidation Numbers: l Use oxidation numbers (charges on atoms) to determine which atom underwent reduction and which atom underwent oxidation

18. Rules: 1. The oxidation number for any free element is 0 (zero). Also any diatomic molecule is 0 (zero) H 2, O 2, I 2, Cl 2, F 2, N 2, Br 2 Fe = 0 charge O 2 = 0 charge

19. 2. The oxidation number of any monoatomic ion is equal to the charge written on the ion. Na +1 = +1 Cl -1 = -1

20. 3. Oxidation number of hydrogen in most of its compounds is +1 (except for LiH then H is –1) +1 Ex. HCl

21. 4. Oxidation # of oxygen in most of its compounds is –2.(except peroxides= -1) -2 Ex. H 2 O Ex. H 2 O 2

22. 5. Sum of the oxidation numbers of all of the atoms must equal the apparent charge of that particle.

24. Ex. H 2 SO 4 -zero charge +1 ? -2 H 2 SO 4 +2 +6-8=0 S= +6

25. Ex. NO 3 –1 ? + -2(3) = -1 +5 + (-6) = -1 N= +5

26. 6. Group 1  +1 Group 2  +2 Aluminum & Boron  +3 Group 17  -1

27. Ex. KMnO 4 K= Mn = O = +1 +7 -2

28. l Page 174 #67, 69

29. l Identifying redox, chemistry connections 11:29 minutes in

30. Identifying Redox Reactions l First, figure out the oxidation numbers of all elements in the reaction l If oxidation number changes as you move from reactants to products it is REDOX.

31. This is REDOX, Mg- loss e- (oxidation), H –gained e-(reduction) This is NOT REDOX

32. l P 618 in modern chem- #2, 15

33. Oxidizing & Reducing Agents l Think of these agents as “causers” of redox rxns l Look at reactants l Some substances are better oxidizing or reducing agents

36. l Reducing Agents: substance that donates the electron (contains the atoms that are oxidized- or loss the e-) Causes the reduction to occur l Oxidizing Agent: substance that gains the e- (contains the atoms that are reduced or gains e-) Causes oxidation to occur

38. Ex. 4Al + 3O 2  2Al 2 O 3 0 0 +3 -2 Al- lost e-, oxidized -reducing agent O- gained e-, reduced -O 2 is the oxidizing agent

39. Balancing Redox Reactions --Half Reaction Method l Half Reaction: equation that shows just the oxidation or reduction part of the rxn. l In balancing we balance each of the half rxns first, then add them together & reduce

40. Steps: 1. Place oxidation #’s on everything after it is in the net ionic form. 2. ID the oxidation ½ rxn and the reduction ½ rxn 3. Write out the ½ rxns. 4. Balance the atoms by placing coefficients in front of the atoms  except for H and O Ex. Cl 2  Cl -1 become Cl 2  2Cl -

41. 5. Place the # of electrons lost on the product side of oxidation ½ rxn, place # of electrons gained on reactant side of reduction ½ rxn 6. To balance hydrogens and oxygens: Acidic soln: add H + & H 2 O Basic soln: add OH - & H 2 O

42. 7. Balance the charges (# e- lost must equal # e- gained) by using a least common multiple ( multiply the whole ½ rxn) 8. Add two ½ rxns together and reduce if necessary.

43. l Chemistry connections- balancing with blood alcohol tests (21:00-26:00)balancing with blood alcohol tests

44. Electrochemistry Movie

45. l Because redox reactions involve electron transfer, the release or absorption of energy can occur in the form of electrical energy rather than heat l Electrochemistry is the branch of chemistry that deals w/ electricity related applications of redox reactions

46. Electrochemical Process l Conversion of chemical to electrical energy l Ex. Flashlight batteries, biological systems, electroplating l If the substance that is oxidized is separated from the substance that is reduced you get an energy transfer of electrical energy instead of heat

49. l Electrons can be transferred from one side to the other through a connecting wire l Electric current moves in a circuit (while the electrons are being balanced by the movement of ions in solution)

50. Part of a Cell: l Electrodes: – Conductor in a circuit that carries electrons from one substance to another – Anodes: electrode where oxidation occurs, anions (-) are attracted to this when they are oxidized by losing electrons (the positive electrode) – Cathode: electrode where reduction occurs, cations (+) are attracted to this when they are reduced by gaining electrons (negative electrode)

51. l Salt Bridge: – Porous partition that separates the 2 half reactions – Contains a conducting solution that allows the passage of ions from one compartment to the other w/ out mixing the solutions in the half reactions

52. l Half Cell: – Part of the voltaic cell in which either oxidation or reduction occurs – The two half cells together make a complete electrochemical cell

53. l Ex. Oxidation half cell – Zn  Zn +2 + 2 e- (zinc rod in zinc sulfate) l Reduction half cell – Cu +2 + 2e-  Cu (copper rod in copper sulfate)

54. l Complete Cell Notation Anode electrode |anode solution || cathode solution |cathode electrode (the double line || represents the salt bridge) Ex. Zn (s) | Zn +2 (aq) || Cu +2 (aq) | Cu (s)

55. ZnSO 4 Zn rod Cu rod CuSO 4 Anode-positive electrode, oxid. occurs Cathode- neg. electrode, red. occurs Salt bridge e- Zn(s) |ZnSO 4( aq)||CuSO 4 (aq) | Cu (s)

57. l Fuel Cell

58. Type of Cells l Dry Cell: voltaic cell in which the electrolyte (conducting solution) is a paste – Generates direct current by converting chemical to electrical energy by a spontaneous redox reaction – Also called galvanic cells or voltaic cells – Ex. Batteries (zinc-carbon, alkaline, mercury) – Ex. Flashlight battery (zinc-carbon) Zinc container (anode) filled w/ a moist paste (salt paste) made of MnO 2, ZnCl 2, NH 4 Cl and water w/ a graphite rod (cathode) embedded into it

60. – Alkaline batteries (do not have a carbon rod cathode which allows them to be smaller- uses a graphite/ MnO 2 mix) – Mercury (cathode is HgO/carbon mix)

62. – Lead storage batteries Group of cells that are connected together Can be recharged (use in a car) Ex. 12 V battery- 6 voltaic cells connected together –Each cell contains 2 lead electrodes or grids –Anode- grid packed w/ spongy lead –Cathode – grid packed w/ PbO 2 –Immersed in 5M H 2 SO 4 Recharging occurs whenever the car is running Doesn’t last forever- byproduct PbSO 4 falls from electrodes and collects on bottom (loses too much lead)

63. – Fuel Cells A voltaic cell in which the reactants are being continuously supplied and the product are being continuously removed A fuel substance undergoes oxidation, from which electrical energy is obtained continuously No recharging, no pollution Ex. H-O cell: submarines, military vehicles, Apollo

65. Electrical Potential l In a voltaic cell, the oxidizing agent at the cathode pulls the electrons through the wire away from the reducing agent at the anode l The “pull” on the electrons is called the electric potential l Electrical potential is measured in volts (V)

66. Electrode potential: l The potential difference measure across the complete voltaic cell is easily measured l It equals the sum of the electrode potentials for each of the two half-reactions l The individual electrode potential for a half- reaction cannot be measured directly, but it can be measured by connecting to a standard half-cell as a reference (we use a Hydrogen electrode that is in a 1.0M acidic solution at 1 atm and 25 C)

67. Standard Reduction Potentials (p. 796-book) l Electrode potentials are always written as reductions l The more negative the voltage  oxidation (stronger reducing agent) l The more positive the voltage  reduction (stronger oxidizing agent)

69. Standard Cell Potential (E° cell) l Use this formula: E°cell = E°reduction - E°oxidation or E°cell = E°cathode - E°anode l A spontaneous reaction will have positive value for E° cell

70. Zn (s) | Zn +2 (aq) || Cu +2 (aq) | Cu (s) l Oxidation: Zn +2 + 2 e-  Zn –E° Zn +2 = -.76 V l Reduction: Cu +2 + 2e-  Cu –E° Cu+2 =.34V l E°cell = E°reduction - E°oxidation =.34V - (-.76V) =1.10V Cu +2 + Zn  Cu + Zn +2

71. Zn (s) | Zn +2 (aq) || Fe +2 (aq) | Fe (s) (anode)(cathode) l Oxidation: Zn +2 + 2 e-  Zn –E° Zn +2 = -.76 V l Reduction: Fe +2 + 2e-  Fe –E° Fe+2 = -.44V l E°cell = E°reduction - E°oxidation =-.44V - (-.76V) =.32V Fe +2 + Zn  Fe + Zn +2

72. Practice l Mn| Mn +2 || Br 2 | Br - l H 2 C 2 O 4 | CO 2 || MnO 4 -1 | Mn+2 l Ni | Ni +2 || Hg 2 +2 | Hg l Cu | Cu +2 || Ag +1 | Ag l Pb| Pb +2 || Cl 2 | Cl -

73. l Mn| Mn +2 || Br 2 | Br - – Ecell= 1.07-(-1.18)= 2.25 V – Br 2 +Mn  Mn +2 + 2Br - l H 2 C 2 O 4 | CO 2 || MnO 4 -1 | Mn +2 – E cell= 1.51- (-.49) = 2.00 V – 2 MnO 4 - +6 H + + 5 H 2 C 2 O 4  2Mn +2 + 8H 2 O + 10 CO 2 l Ni | Ni +2 || Hg 2 +2 | Hg – 1.04 V – Ni + Hg 2 +2  Ni +2 + 2Hg l Cu | Cu +2 || Ag +1 | Ag –.46 V – Cu + 2 Ag+  Cu +2 + 2 Ag l Pb| Pb +2 || Cl 2 | Cl - – ` 1.49V – Pb +Cl 2  Pb +2 + 2Cl -

74. l Video Video l How its made nails How its made nails l Corrision Pics Corrision Pics

75. Redox/ Electrochemistry Quest (anode song) (anode song) (anode song) Redox – oxidation #’s – ID if redox or not – Oxidizing or reducing agent (strengths) – Balancing- set up ½ rxns – Balancing oxygens/hydrogens Acids (add H+ and H 2 0) Bases (add OH- and H 2 0)

76. l Electrochem – What is an electrochemical cell – Example – Parts: anode, cathode, salt bridge, what each part does) – Standard cell potential (getting the voltage and write equation)