1. Unit G Oxidation-Reduction

2. Oxidation and Reduction (Redox)
• Redox reactions involve the transfer of electrons from one atom to another.
Example: Magnesium metal in hydrochloric acid:
Mg HCI → MgCl H2
Oxidation: the process of losing electrons.
Mg → Mg2+ + 2e­
Reduction: the process of gaining electrons.
2H++ 2e- →H2
In this case, the Mg was oxidized, while the H+ was reduced!

3. Definitions
Oxidized substance: lost electrons.
Reduced substance: gained electrons.
Oxidizing agent:
Responsible for the oxidation (cause it).
Gains electrons in the reaction (ie.it gets reduced).
Reducing agent:
Responsible for the reduction .
Loses electrons in the reaction (ie.it gets oxidized).

4. LEO the lion goes GER Loss of Electrons Oxidation
Gain of Electrons Reduction

5. Examples of Redox Reactions
Example Half Reactions
oxidation: Zn→Zn e-
Zn(s) + 2 HCI(aq) → reduction: 2 H+ + 2e- → H2
ZnCl2(aq)+ H2(g) overall: Zn +H+ →Zn2+ +H2

6. Examples of Redox Reactions
Example Half Reactions
oxidation: Cu→Cu+ + e-
AgNO3 + Cu → reduction: Ag+ + e- → Ag
CuNO3+ Ag overall: Cu +Ag+ →Cu+ +Ag

7. Examples of Redox Reactions
Example Half Reactions
oxidation: Mg→Mg e-
2Mg+ O2 → 2MgO reduction: O2+ 4e- → 2O2-
overall: 2Mg+ O2 → 2MgO

8. Oxidation Number
The real or apparent charge on an atom or ion when all bonds are assumed to be ionic.

9. The RULES to determine oxidation number:
1. All atoms in a n1olecule of an element have a ZERO value. e.g. O2,H2,Zn, S8,P4,AI, Cu etc.
2. Ions have an oxidation number equal to their charge. e.g. Zn2+ = +2; CI- = -1
3. In compounds, Hydrogen is +1; except in
hydrides of metals (NaH; ZnH2) where it is -1.
4. In compounds, Oxygen is -2; except in peroxides (H2O2) where it is -1.

10. The RULES cont' d:
5. Alkali metals (Group I) are +1.
6. Alkaline earths (Group 2) are +2.
7. In a compound, the sum of all oxidation numbers is 0 (ZERO).
8. In a polyatomic ion, the sum of all oxidation numbers is the charge on the ion.
e.g. NH4+ = +1

11. Examples: Find the oxidation number of the underlined element:
a.MnO2
b.KMnO4
c.KClO3
d.K2Cr2O7
e.ClO3-
f.MnO42-
g.Fe3O4
h.NH3

12. Important!
• if the oxidation number increases during a reaction, the reaction is an oxidation.
• if the oxidation number decreases, the reaction is an reduction.
• if the oxidation number doesn 't change, it is not
a redox reaction.

13. Example (not in notes)
• Which element is oxidized in this reaction?
Fe + H2SO4 → FeSO4 + H2

14. Assignment
Hebden #1-5

15. half- Reactions and Reduction Potential
• all redox reactions can be broken into two half-reactions(or half-cells).
Example 1: Al foil in Cu2+ solution. A reaction occurs.
half reaction for oxidation :_____________________
half reaction for reduction:______________________
Reducing agent:________; oxidizing agent:_________

16. Example 2: Ag in Cu2+ solution.
oxidation half reaction: ______________________
reduction half reaction: _______________________

17. •We can determine which reaction would go without having to do the experiment by referring to the table "Standard Reduction Potentials of Half-Cells".
Any oxidizing agent on the LEFT
will react with any reducing agent on the
RIGHT
that is LOWER in the table.
YES, a reaction will occur

18. Examples:
a. can Cu2+ oxidize Ag?
b. Will F2 oxidize Au?
c. Will I2 oxidize Cu?
d. Will Cu2+ oxidize Mg2+?
e. will Ag+ oxidize Cu?
f. will Cu2+ oxidize Ag+?
g. will Cl- oxidize Ag?
h. give an example of an oxidizing agent that can oxidize Ni but not ?

19. Assignment
Hebden #7-12 (abcd for all)

20. Balancing Redox Equations
A. By Inspection
Just balance each half-reaction, then add and cancel as needed.
Example 1: Cu + Ag+ → Cu Ag
Overall balanced equation : ______________________

21. Balancing Redox Equations
A. By Inspection
Example 2: Al + Sn4+ → Al Sn
Overall balanced equation : ______________________

22. Assignment
Hebden #19 (abcd), 23abc

23. Balancing Redox Equations
B. The Half-Cell Method (in acidic solution)
Example 1: MnO4- + C2O42-→Mn2+ + CO2
Major OH-
Step 1: Divide into 2 half reactions and balance the major atoms.
MnO4- ⇌ Mn2+
C2O42⇌CO2
Step 2: Balance Oxygen atoms by adding H2O.
C2O42⇌2CO2

24. Balancing Redox Equations
B. The Half-Cell Method (in acidic solution)
Example 1: MnO4- + C2O42-→Mn2+ + CO2
Major OH-
Step 3: balance hydrogen atoms by adding H+.
MnO4- ⇌ Mn2+ + 4H2O
C2O42⇌ 2CO2
Step 4: add electrons to balance charge.
8H++ MnO4- ⇌ Mn2+ + 4H2O
C2O42⇌2CO2

25. Balancing Redox Equations
B. The Half-Cell Method (in acidic solution)
Example 1: MnO4- + C2O42-→Mn2+ + CO2
Major OH-
Step 5: Use multipliers to equalize e- in oxidation and reduction half-reaction.
5e-+ 8H++ MnO4- ⇌ Mn2+ + 4H2O
C2O42⇌2CO2 +2e-

26. Balancing Redox Equations
B. The Half-Cell Method (in acidic solution)
Example 1: MnO4- + C2O42-→Mn2+ + CO2
Major OH-
overall reaction:
10e-+ 16H++ 2MnO4- ⇌ 2Mn2+ + 8H2O
5C2O42⇌10CO2 +10e-
Overall ___________________________________

27. Use major hydroxide to help you remember the steps.
Major OH-

28. Balancing Redox Equations
B. The Half-Cell Method (in acidic solution)
Example 2: HNO2 → HNO3 + NO
Major OH-

29. Assignment
Hebden # 24ace

30. Balancing Redox Equations
B. The Half-Cell Method (in basic solution)
Example 3: MnO4- + N2H4→MnO2 + N2
Major OH-
Step 1: Divide into 2 half-reactions and balance the major atoms.
N2H4 ⇌ N2
MnO4- ⇌ MnO2

31. Balancing Redox Equations
B. The Half-Cell Method (in basic solution)
Example 3: MnO4- + N2H4→MnO2 + N2
Major OH-
Step 2: Now balance oxygen, hydrogen, and charge (in that order!)
N2H4 ⇌ N2
MnO4- ⇌ MnO2

32. Balancing Redox Equations
B. The Half-Cell Method (in basic solution)
Example 3: MnO4- + N2H4→MnO2 + N2
Major OH-
Step 3: multiply each half-reaction so that
electrons lost=electrons gained
N2H4 ⇌ N2 +4H++ 4e-
3e- + 4H+ + MnO4- ⇌ MnO2 + 2H2O

33. Balancing Redox Equations
B. The Half-Cell Method (in basic solution)
Example 3: MnO4- + N2H4→MnO2 + N2
Major OH-
Step 4: Add the two half-reactions.
3N2H4 ⇌ 3N2 +12H++ 12e-
12e- + 16H+ + 4MnO4- ⇌ 4MnO2 + 8H2O
3N2H4 +4H MnO4- ⇌ 3N2+4MnO2 + 8H2O

34. Balancing Redox Equations
B. The Half-Cell Method (in basic solution)
Example 3: MnO4- + N2H4→MnO2 + N2
Major OH-
Step 5: Convert to basic conditions by adding sufficient
OH- ions to EACH side to neutralize any H+ ions.
3N2H4 +4H MnO4- ⇌ 3N2+4MnO2 + 8H2O
Overall reaction:_________________________________

35. Assignment
Hebden # 24fgh

36. Redox Titrations
we have seen titrations twice before:
in the Ksp unit (eg. titrating until all the AgCl has ppt.out)
In the acid/base unit( eg. Titrating an acid with a base until an indicator turns colour)
Redox titrations are based on the ability of some elements, like manganese, to change colour when their oxidation number changes. Mn7+ is purple, Mn2+ is colourless, Mn4+ is brown, while Mn6+ is green!

37. Redox Titration Calculations
Example 1: FeSO4 in acid solution is titrated with 27.4 mL M KMnO4. The iron is oxidized to Fe3+, while the MnO4- is reduced to Mn2+ . If25.00 mL ofFeSO4 was used, calculate the original [FeSO4].
•Balance equation: Fe2+ + MnO4- → Fe3+ + Mn2+
Overall:__________________________________
• Solve:

38. Redox Titration Calculations
Example 2: The Sn II in SnSO4 is oxidized to Sn IV in an acid solution by titration with 8.08 mL of M KMnO4 solution. Originally, 25.0 mL of SnSO4 was used. Calculate [SnSO4].
•Balance equation: Sn2+ + MnO4- → Sn4+ + Mn2+
Overall: _________________________________
• Solve:

39. Assignment
Hebden # 26-29
Unit G TEST

40. Unit H
Applications of Redox Reactions

41. Electrochemical Cells
Definition Time!
•Electrochemical Cell:
•Two reaction half-cells, physically separated.
•Electron transfer takes place through a wire between the half cells.
•Spontaneous (occurs without assistance)
•Does useful work.

42. • Electrodes: surfaces on which the oxidation or reduction reactions occur.
• Anode: Electrode where oxidation takes place.
• Cathode: Electrode where reduction takes place.
• Anion: negative ion. Moves to anode.
• Cation: positive ion. Moves to cathode.
• Salt Bridge: electrolyte solution in a tube connecting the two half cells. Ions (not electrons!) move through it to keep solutions neutral. Electrons don' t swim!

43. • Voltage (or Potential Difference): Work done by the cell.
• Voltmeter: Measures voltage difference between half cells.
• Current: Rate of flow of electric charges (amperage)
• Ammeter: Measures current in amps.

44. Label this Diagram

45. Assignment
Draw an electrochemical cell with an iron electrode in Fe(NO3)2 and Mg electrode in Mg(NO3)2; with a BaCl2 salt bridge. Label anode, cathode, motion of ions, and motion of electrons.
Write equations for the half-reactions.

46. Standard Half-Cell Potential (Eo)
• Eo is a measure of the tendency of electrons to flow.
• is evaluated at standard conditions: 25 °C, 1.0 M solutions, kPa pressure.
• the base value against which all other half­ reactions are measured is arbitrarily chosen to be the reduction of hydrogen:
2 H++ 2e- ⇌H2 Eo = 0.00 v

47. Example 1
• Compare the reduction of Cu2+ to H+ to Zn2+:
Cu2+ + 2e- ⇌ Cu Eo = v
2 H++ 2e-⇌H2 Eo = 0.00 V
Zn2+ + 2e- ⇌ Zn Eo = V
• Cu2+ has the highest Eo value of these three, therefore it has the greatest tendency to be reduced.

48. Example 2
• Compare the Cu/Cu2+ half cell with the Zn/Zn2+ half cell: which of these will be oxidized and which will be reduced?
Cu2+ + 2e- ⇌ Cu Eo = v
Zn2+ + 2e- ⇌ Zn Eo = V
the ___ will be reduced, because it has a higher Eo Value.
the ____will be therefore oxidized.
• conclusion: the guy with the lowest Eo value loses (electrons). In other words, he gets oxidized.

49. Assignment
Hebden # 34

50. Predicting Spontaneity
If the E total is POSITIVE, the reaction is SPONTANEOUS

51. Calculate the standard cell potentials for the following reactions:
1. Cu2+ + Na →Cu + Na+
Cu2+ + 2e- ⇌ Cu Eo = v
Na ⇌Na++ e Eo = v
Cu Na→ Cu + 2 Na+ Eo total =________
• Reaction is spontaneous!

52. Calculate the standard cell potentials for the following reactions:
2. Br2+ 2F-→
Br2 + 2e- ⇌ 2Br Eo = v
2F- ⇌ F2+2e Eo = v
Br2 +2F- → 2Br-+ F2 Eo total =________
• Reaction is NOT spontaneous!

53. Find the spontaneous reaction and cell voltage when Al/Al(NO3)3 and Cu/Cu(NO3)2 half-cell are connected
3. Cu2+ + Na →Cu + Na+
Cu2+ + 2e- ⇌ Cu Eo = v
Al⇌Al e Eo = v
3Cu2+ + 2Al→ 3Cu + 2 Al3+ Eo total =________
• Reaction is spontaneous!

54. Effect Of Concentration On Eo
• Consider the Nickel/Cadmium battery:
Ni2+ + Cd →Ni + Cd2+
Ni2+ + 2e- ⇌ Ni Eo = v
Cd⇌Cd e Eo = v
Ni2+ + Cd→ Ni + Cd 2+ Eo total =________
• Reaction is spontaneous!

55. •the [Cd2+] goes____ , and the [Ni2+] goes____.
•As the reaction Ni2+ + Cd→ Ni + Cd 2+
proceeds:
•the [Cd2+] goes____ , and the [Ni2+] goes____.
•the forward reaction slows.
•Eo total drops.
•When equilibrium is reached, the forward and reverse reactions are occuring at the same rate.
Vforward=_____V and Vreverse =______V, so Eototal=_____
• To increase voltage, either increase [Reactant] or
decrease [Products].

56. Effect Of Electrode surface area On Eo
• _______, because no concentrations are changed,
And therefore no shift in equilibrium occurs.

57. Assignment
Hebden # 36 abcd,37,40,43

58. Applications of Electrochemical cells

59. 1. Dry Cell (Zinc-Carbon battery)
• Anode:
• Cathode:
• There is no one simple overall reaction. Many reactions occur simultaneously.

61. 2. Alkaline Cell

62. 2. Alkaline Cell
Use KOH as electrolyte (hence the name alkaline)
Gives more current than standard dry cell.
Resistance does not increase with use.
More expensive
Possible reaction

63. 3 . Lead-acid storage battery (car battery)

64. 3 . Lead-acid storage battery (car battery)
Anode
Cathode
Overall
As battery is discharged,[H2SO4] ____, and so density also_______.
Hydrometer measures charge left by measuring density.

65. 4 . Fuel Cell

66. 4 . Fuel Cell
Anode
Cathode
Overall
Great benefit: produces ____ as a waste product.
Vey efficient
Produces electricity on space shuttle.
Powers hydrogen fuel cell buses and cars (still in experimental stage).

68. 5 . Corrosion of metals
Anode
Cathode
Overall
the number of water molecules associated with the iron oxide varies, and accounts for the different colours of rust ( red, brown,yellow,black).

70. 6 . Cathodic Protection

71. To prevent corrosion of valuable steel structures ( ship hulls, pipelines, cables, etc.) another metal (often Mg) is attached to the structure.
The Mg oxidizes more easily than Fe, making the iron structure the cathode,so it can’t deteriorate into ions.
Anode
Cathode
Result: the Mg corrodes (but can easily be replaced) but the structure itself remains intact.

72. Electrolytic cells
Electrolytic cells are endothermic redox reactions that must absorb electrical energy to proceed.
Compare this to electrochemical cells which are spontaneous, exothermic redox reactions that produce electrical energy.
Definition of electrolysis: when an external source of energy cause a non-spontaneous redox reaction to occur.

73. type 1 Electrolytic cells
contains the fewest chemical species.
Electrodes are inert (carbon or platinum)
Electrolyte is a molten salt.

74. Example: Electrolysis of molten ZnCl2
• Species available:
Anode:
Cathode:
Overall:

76. type 2 Electrolytic cells
Unreactive Electrodes
Electrolyte is an aqueous solution (so water is present)
Find and record the half-reactions for water
Oxidation:________________________
Reduction:_____________________

77. Example 1: predict what happens in the electrolysis of NaI (aq).
Species available:
Possible anode reactions:
Possible cathode reactions:
The ACTUAL reactions are the easiest ones:
reduction: take the HIGHEST in table
oxidation: take Lowest in table
Overall reaction:

78. •Possible Anode reactions:
Example 2: Write the total cell reaction for the electrolysis of CuBr2(aq) .
•Species available:
•Possible Anode reactions:
•Possible Cathode reactions:
•Star the ACTUAL reactions.
•Overall Reaction:

80. •Possible Anode reactions: •Possible Cathode reactions:
Example 3: Predict the reactions involved in the electrolysis of water.
•Species available:
•Possible Anode reactions:
•Possible Cathode reactions:
•Why is the reduction of SO42-
not considered?
•Overall Reaction:

82. Assignment
Hebden # 65abcd,66,67,68,70

83. type 3 Electrolytic cell (Electroplating)
Electrodes will react
Electrolyte is an aqueous solution
The cathode will be plated (covered in a thin coating)

84. •Possible Anode reactions:
Example: The electroplating of copper on to an iron pot from a solution of copper sulphate.
•Species available:
•Possible Anode reactions:

85. • Possible Cathode reactions:
Example: The electroplating of copper on to an iron pot from a solution of copper sulphate.
• Possible Cathode reactions:
Star the actual reactions
• Overall Reaction:

86. Assignment
Hebden # 73-76

87. J. Applications of Electrolytic Cells
Overpotential revisited
In actual practice, electrolysis requires a higher potential(more voltage) than is calculated theoretically.
This is due to the nature of the reactants, resistance to
the current in the wires, and the law of entropy.
For the production of H2 and O2, the overpotential is large
This means that Br- and Cl- will oxidize to Br2 and Cl2 more easily than water will oxidize to O2.
It also means that Zn2+ will reduce to more easily than water will reduce to H2.
These features of the overpotenitial have significant application in the commercial production of many chemicals.

88. •Synthetic rubber is used as a catalyst.
Commercial Applications
1. Production of Chlorine
Uses of Chlorine: plastics, solvents, purify water, bleach, insecticides
a) Down' s Cell (Type 1) • Anode
•Cathode:
•Overall:
•Synthetic rubber is used as a catalyst.

90. b) Electrolysis of Brine (a concentrated solution of NaCI).
1. Production of Chlorine (cont' d)
b) Electrolysis of Brine
(a concentrated solution of NaCI).
• Species available:
• Possible Anode reactions:
• Possible Cathode reactions:
• Overall Reaction:

92. 2.Electrorefining
Pure metal is plated on a cathode from an impure anode.
Lead is refined this way in Trail, BC for use in batteries, paint, gasoline, and crystal.
Anode
Cathode
At the anode, many impurites are left behind. why?
At the cathode, impurities like Zn2+,Mn2+ don’t plate,why?

93. 3.Electrowinning
Pure metal is plated on a cathode from an electrolyte containing the metal ion.
If the plated metal is zinc, this process is called galvnizing.
Anode:
Cathode:

94. Aluminum Refining
Uses: good conductor, one third the density of steel yet strong, corrosion resistant due to oxide coating.
•Aluminum can't be purified by electrorefining or electrowinning. Why?
•Best alternative: Electrolysis of molten Al2O3, but this needs temperature of 2000°C, so the refining of Aluminum involves a process that lowers the melting point.
1. Heat the bauxite ore, Al2O3 • H2O, to drive off the water.
2.Dissolve Al2O3 in cryolite, Na3AIF6, which acts as a catalyst and lowers the m.p. to 1000°C.
3.Electrolyze the molten Al2O3

95. Aluminum Refining cont' d
Anode:
Cathode:
Overall:

99. Breathalyzer Test
Ethanol (in alcoholic beverages )can be oxidized to acetic acid by dichromate in acidic solution.
The breathalyzer measures the degree to which the orange colour moved
[ethanol ] in blood is proportional to [ethanol ] in exhaled breath.