This document discusses redox reactions, electrochemistry, and electrochemical cells. It begins by defining key concepts like oxidation, reduction, oxidizing agents, and reducing agents. It then provides examples of redox reactions and discusses how electrochemical cells work. The rest of the document covers topics like cell notation, standard electrode potentials, how to determine if a redox reaction is spontaneous, the relationship between cell potential and Gibbs free energy, the effect of concentration on cell potential, corrosion, batteries, and different types of electrochemical cells like voltaic cells, electrolytic cells, and fuel cells.

1. Redox Reactions and
Electrochemistry
By Aktr
book 2

4. Loss of e
Gain of e
Reducing Agent
Oxidising agent

5. Oxidation

6. Reduction

7. Electrochemical cell occur both

11. SnCl2(aq) + 2FeCl3(aq) → SnCl4(aq) + 2FeCl2(aq)
CuO(s) + H2(g) → Cu(s) + H2O(l)
Example
+2 +3 +4 +2
+1
0
0
+2

12. Conductance in electrolytic
solution

27. Also act as a conductor

32. Conductance in electrolytic
solutions

36. Metallic conductor depend upon
• Nature and structure of metal.
• No. of valence electron per atom.
• Temperature of the sample.

40. conductivity aqueous solution
depend upon
• Nature of electrolyte.
• Size of ion.
• Solvation of ion.
• Concentration of electrolytic.
• Temperature.

48. One cell const. and resistivity
Known then we can find value
Easily.

49. Molar conductivity

63. Electrochemical cell

65. Chemical to electrical Electrical to chemical

66. 2. Voltaic or Galvanic cells-
An electrochemical cell in which a spontaneous
reaction produces electricity.
Eg. Dry cell, lead storage cell etc.

67. 1. Electrolytic cell-
An electrochemical cell in which a non spontaneous
reaction is forced to occur by passing a direct
current from an external source into the solution.
Eg. Refining metal(purify), electroplating &
production of many chemical substance.

68. Gets Smaller -> <- Gets Larger

69. Cell Notation
1. Anode
2. Salt Bridge
3. Cathode
Anode | Salt Bridge | Cathode
| : symbol is used whenever there is a different phase

70. 19.2
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..

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

72. Electrochemical Cells
19.2
The difference in electrical
potential between the anode and
cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
0
0
0
reduction
oxidation
Cell E
E
E +
=
UNITS: Volts Volt (V) = Joule (J)
Coulomb, C

73. Standard Electrode Potentials
19.3
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.
Ε0
= 0 V
Standard hydrogen electrode (SHE)
2e−
+ 2Η+
(1 Μ) Η2 (1 atm)
Reduction Reaction

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

76. Relating E0
Cell to ∆G0
e
ch
work
ECell
arg
=
Units
work, Joule
charge, Coulomb
Ecell; Volts
charge = nF
Faraday, F; charge on 1 mole e-
F = 96485 C/mole
work = (charge)Ecell = -nFEcell
∆G = work (maximum)
∆G = -nFEcell

77. Relating Εο
CELL to the
Equilibrium Constant, K
∆G0
= -RT ln K
∆G0
= -nFE0
cell
-RT ln K = -nFE0
cell
K
nF
RT
ECell ln
0
=
( )
0257
.
0
96485
298
31
.
8
=






=
mole
C
K
molK
J
F
RT
K
n
K
n
ECell log
0592
.
0
ln
0257
.
0
0
=
=

80. Effect of Concentration on Cell Potential
∆G =∆G0
+ RTlnQ
∆G0
= -nFE0
cell
-nFEcell= -nFE0
cell +RTln Q
Ecell= E0
cell - RTln Q
nF
Ecell= E0
cell - 0.0257ln Q
n
Ecell= E0
cell – 0.0592log Q
n

81. Corrosion – Deterioration of Metals
by Electrochemical Process

82. Corrosion – Deterioration of Metals
by Electrochemical Process

83. Corrosion – Deterioration of Metals
by Electrochemical Process

85. Cathodic Protection

86. Abbreviated Standard Reduction
Potential Table

88. Batteries
19.6
Leclanché cell
Dry cell
Zn (s) Zn2+
(aq) + 2e-
Anode:
Cathode: 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)

89. Batteries
Zn(Hg) + 2OH-
(aq) ZnO (s) + H2O (l) + 2e-
Anode:
Cathode: HgO (s) + H2O (l) + 2e-
Hg (l) + 2OH-
(aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
Mercury Battery
19.6

90. Batteries
19.6
Anode:
Cathode:
Lead storage
battery
PbO2 (s) + 4H+
(aq) + SO4
2-
(aq) + 2e-
PbSO4 (s) + 2H2O (l)
Pb (s) + SO4
2-
(aq) PbSO4 (s) + 2e-
Pb (s) + PbO2 (s) + 4H+
(aq) + 2SO4
2-
(aq) 2PbSO4 (s) + 2H2O (l)

91. 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

92. 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-
4OH-
(aq)
2H2 (g) + 4OH-
(aq) 4H2O (l) + 4e-
2H2 (g) + O2 (g) 2H2O (l)

93. 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.

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