1. Redox Reactions and Electrochemistry
Oxidation Number
Oxidizing and Reducing Reagents
Galavanic or Voltaic Cells
Anode/Cathode/Salt Bridge
Cell Notations
Determining Cell Potential/Cell Voltage/Electromotive force (emf)
Relating Cell Potential to K and DG0
Effect of Concentration on Cell Potential
Corrosion
Batteries
Fuel Cells
Electrolytic Cells
Calculating amounts of substances reduced or oxidized

2. REDOX REACTION

3. DEFINITIONS LEO says GER Oxidation: Gain of oxygen
Oxidation: Loss of electrons.
Reduction: Gain of electrons.
LEO says GER
Oxidation: Gain of oxygen
Reduction: Lost of oxygen
Oxidation: Increasing of oxidation number
Reduction: Reducing of oxidation number

4. CO  CO2 (CO oxidized) why?
CH3COOH  CH3CHO (CH3COOH reduced) why?
H2SO3  H2SO4 ??
HNO3  HNO2 ??
Ca2+  Ca (Ca2+ reduced) why?
Na  Na+ (Na oxidized) why?
Fe3+  Fe2+
Mn2+  MnO4– (Mn2+ oxidized) why?
Cl2  2 Cl– (Cl2 reduced) why?
H2O2  H2O ??
NaH  H2 ??

5. Lose Electrons = Oxidation
LEO says GER :
Lose Electrons = Oxidation
Sodium is oxidized
Na0  Na+1 + 1e-
Gain Electrons = Reduction
Chlorine is reduced
Cl0 + 1e–  Cl–

6. Rules for Assignment of Oxidation Number (ON)
1) The ON of all pure elements is zero. 2) The ON of H is +1, except in hydrides, where it is -1. 3) The ON of O is -2, except in peroxides, where it is -1. 4) The algebraic sum of ON must equal zero for a neutral molecule or the charge on an ion.

7. Variable Oxidation Number of Elements
Sulfur: SO42-(+6), SO32-(+4), S(0), FeS2(-1), H2S(-2)
Carbon: CO2(+4), C(0), CH4(-4)
Nitrogen: NO3-(+5), NO2-(+3), NO(+2), N2O(+1), N2(0), NH3(-3)
Iron: Fe2O3(+3), FeO(+2), Fe(0)
Manganese: MnO4-(+7), MnO2(+4), Mn2O3(+3), MnO(+2), Mn(0)
Copper: CuO(+2), Cu2O(+1), Cu(0)
Tin: SnO2(+4), Sn2+(+2), Sn(0)
Uranium: UO22+(+6), UO2(+4), U(0)
Arsenic: H3AsO40(+5), H3AsO30(+3), As(0), AsH3(-1)
Chromium: CrO42-(+6), Cr2O3(+3), Cr(0)
Gold: AuCl4-(+3), Au(CN)2-(+1), Au(0)

8. BALANCING OVERALL REDOX REACTIONS
Example balance the redox reaction below: Fe + Cl2  Fe3+ + Cl- Step 1: Assign oxidation number, Fe0 + Cl20  Fe3+ + Cl- Step 2: Determine number of electrons lost or gained by reactants.   3e- 2e- Step 3: Cross multiply. 2Fe + 3Cl20  2Fe3+ + 6Cl-

9. may be written as the sum of two half-cell reactions: 2Fe  2Fe3+ + 6e- (oxidation) 3Cl20 + 6e-  6Cl- (reduction) All overall redox reactions can be expressed as the sum of two half-cell reactions, one a reduction and one an oxidation. The overall reaction: 2Fe + 3Cl20  2Fe3+ + 6Cl-

10. Another example - balance the redox reaction: FeS2 + O2  Fe(OH)3 + SO42- Fe+2S20 + O20  Fe+3(OH)3 + S+6O42-   15e- 4e- 4FeS2 + 15O2  Fe(OH)3 + SO42- 4FeS2 + 15O2  4Fe(OH)3 + 8SO42- 4FeS2 + 15O2 +14H2O  4Fe(OH)3 + 8SO H+ This reaction is the main cause of acid generation in drainage from sulfide ore deposits. Note that we get 4 moles of H+ for every mole of pyrite oxidized!

11. Electrochemistry: Interconversion of electrical and chemical energy using redox reactions
Oxidation Half-Reaction: Oxidation Involves Loss
of electrons
Reduction Half-Reaction: Reduction Involves Gain of electrons
2Mg Mg2+ + 4e-
O2 + 4e O2-
Net Redox Reaction: 2Mg + O2  2 Mg O-2

12. Oxidation number The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred
to the more electronegative atom.
Oxidation number equals ionic charge formonoatomic ions in ionic compound
CaBr2; Ca = +2, Br = -1
2. Metal ions in Family A have one, positive oxidation number; Group IA metals are +1, IIA metals are +2
Li+, Li = +1; Mg+2, Mg = +2

13. The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred
to the more electronegative atom.
3. The oxidation number of a transition metal ion is positive, but can vary in magnitude.
Nonmetals can have a variety of oxidation numbers,both positive and negative numbers which can vary in magnitude.
5. Free elements (uncombined state) have an oxidation number of zero. Each atom in O2, F2, H2, Cl2, K, Be has the same oxidation number; zero.

14. 6. The oxidation number of fluorine is always –1
6. The oxidation number of fluorine is always –1. (unless fluorine is in elemental form, F2)
7. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
IF; F= -1; I = +1
8. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 or when it’s in elemental form (H2; oxidation # =0).
HF; F= -1, H= +1
NaH; Na= +1, H = -1

15. 9. The oxidation number of oxygen is usually –2
9. The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1, in elemental form (O2 or O3) it is 0.
H2O ; H=+1, O= -2
SO3; O = -2; S = +6
HCO3-
O = -2
H = +1
3x(-2) C = -1
C = +4
IF7
F = -1
7x(-1) + ? = 0
I = +7
NaIO3
Na = +1
3x(-2) ? = 0
I = +5

16. Determination of Oxidizing and Reducing Agents
Determine oxidation number for all atoms in both the reactants and products.
Look at same atom in reactants and products and see if oxidation number increased or decreased.
If oxidation number decreased:substance reduced  Oxidizing Agent
If oxidation number increased; substance oxidized  Reducing Agent

17. Zn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq)
Spontaneous Redox Reaction
Zn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq) Zn Cu
Cu+2
time
Zn+2

19. Gets Larger
Gets Smaller

20. Types of Electrochemical Cells
Voltaic/Galvanic Cell: Energy released from spontaneous redox reaction can be transformed into electrical energy.
Electrolytic Cell: Electrical energy is used to drive a nonspontaneous redox reaction.

21. Voltaic Cell AnOx or both vowels Anode: Site of Oxidation
Red Cat or both consonants
Anode: Site of Oxidation
Cathode: Site of Reduction
Direction of electron flow: anode to cathode (alphabetical)
Salt Bridge: Maintains electrical neutrality
+ ion migrates to cathode
- ion migrates to anode

22. Cell Notation Anode Salt Bridge Cathode Anode | Salt Bridge | Cathode
| : symbol is used whenever there is a different phase

23. Cell Notation
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)
anode
cathode
Salt bridge
More detail..

24. Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)
K(NO3)
Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

25. Electrochemical Cells
The difference in electrical potential between the anode and cathode is called:
cell voltage
electromotive force (emf)
cell potential
E Units: Volts
Volt (V) = Joule (J)
Coulomb ( C )

30. Standard hydrogen electrode (SHE)
Standard Electrode Potentials
Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction:
2e- + 2H+ (1 M) H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)

32. Determining if Redox Reaction is Spontaneous
+ E°CELL spontaneous reaction
0 E°CELL equilibrium
- E°CELL nonspontaneous reaction
More positive E°CELL stronger oxidizing agent or more likely to be reduced

33. E0 is for the reaction as written
The half-cell reactions are reversible
The sign of E0 changes when the reaction is reversed
Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0
The more positive E0 the greater the tendency for the substance to be reduced

34. Relating E0Cell to G0 Charge = nF Work = (charge)Ecell = -nFEcell
Units
Work: Joule
Charge (Q): Coulomb
Ecell : Volts
Charge = nF
Faraday (F): charge on 1 mole e-
F = C/mole
Work = (charge)Ecell = -nFEcell
G = work (maximum)
G = -nFEcell

35. Relating EoCELL to the Equilibrium Constant, K
G0 = -RT ln K
-RT ln K = -nFE0cell
G0 = -nFE0cell

38. Effect of Concentration on Cell Potential
G = G0 + RTlnQ
-nFEcell = -nFE0cell + RTln Q
G0 = -nFE0cell
Ecell = E0cell - RTln Q nF
Ecell= E0cell ln Q n
Ecell= E0cell – log Q n

39. Corrosion – Deterioration of Metals by Electrochemical Process

43. Cathodic Protection

46. Batteries Dry cell A: Zn (s) Zn2+ (aq) + 2e-
C: 2NH4+ (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

47. Batteries Mercury Battery Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

48. Batteries Lead storage battery Anode:
Pb (s) + SO42- (aq) PbSO4 (s) + 2e-
Cathode:
PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e PbSO4 (s) + 2H2O (l)
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) PbSO4 (s) + 2H2O (l)

49. Fuel Cell vs. Battery Battery: Energy storage device
Reactant chemicals already in device
Once Chemicals used up; discard (unless rechargeable)
Fuel Cell: Energy conversion device
Won’t work unless reactants supplied
Reactants continuously supplied; products continuously removed

50. Fuel Cell
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
2H2 (g) + O2 (g) H2O (l)

52. Faraday’s Constant
Redox Eqn
Molar Mass
Charge =(Current)(Time)