Electrochemistry studies chemical reactions at the interface between an electrode and an electrolyte. Oxidation occurs when an element loses electrons and reduction occurs when an element gains electrons. Galvanic cells produce electrical energy from spontaneous redox reactions. The Nernst equation relates cell potential to concentration. Faraday's laws state that the amount of reaction is proportional to charge and equivalent weights determine amounts deposited. Electrolysis is used industrially to refine and deposit metals.

2. Electrochemistry is a branch
of chemistry that studies chemical
reactions which take place in
a solution at the interface of an
electron conductor (the electrode)
and an ionic conductor
(the electrolyte)

3. Electrochemistry
Electrochemical
Cell
Electrolytic
Cell

4. Oxidation – A process in which an
element attains a more positive
oxidation state
Na(s)  Na+ + e-
Reduction – A process in which an
element attains a more negative
oxidation state
Cl2 + 2e-  2Cl-

5. Gain Electrons = Reduction
An old memory device to remember
oxidation and reduction like this…
LEO says GER
Lose Electrons = Oxidation

6.  Oxidizing agent
The substance that is reduced
is the oxidizing agent
 Reducing agent
The substance that is oxidized
is the reducing agent

7. Galvanic (Electrochemical) Cells
It is an
electrochemical
cell that derives
electrical energy
from spontaneous
redox reaction
taking place with
in a cell.
e-
e- e- e-

8. Zn - Cu
Galvanic
Cell
Zn  Zn2+ + 2e- E0= -0.76V
Cu2+ + 2e-  Cu E0= +0.34V
From the table
of reduction
potentials:

9. Cu2+ + 2e-  Cu E = +0.34V
The less
positive, or
more negative
reduction
potential
becomes the
oxidation…
Zn  Zn2+ + 2e- E = -0.76V
Zn + Cu2+  Zn2+ + Cu E0 = + 1.10 V

10. Line
Notation
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
An abbreviated
representation of
an electrochemical
cell
Anode
solution
Anode
material
Cathode
solution
Cathode
material
| |||

11. Calculating G0 for a Cell
0
(2 )(96485 )(1.10 )
coulombs Joules
G mol e
mol e Coulomb


 
G0 = -nFE0
n = moles of electrons in balanced redox equation
F = Faraday constant = 96,485 coulombs/mol e-
Zn + Cu2+  Zn2+ + Cu E0 = + 1.10 V
0
212267 212G Joules kJ    

12. Nernst Equation
Standard potentials assume a concentration of
1 M. The Nernst equation allows us to calculate
potential when the two cells are not of 1.0 M.
At 25 C (298 K) the Nernst Equation is
simplified this way:
E=E0-RT ln(oxidation)
nF (reduction)
E=E0-0.0591 log(oxidation)
F (reduction)

13. Equilibrium Constants and Cell Potential
At equilibrium, forward and reverse
reactions occur at equal rates, therefore:
1. The battery is “dead”
2. The cell potential,E0 is zero volts
0 0.0591
0 log( )volts E K
n
 
Modifying the Nernst Equation (at 25 C):

14. Zn + Cu2+  Zn2+ + Cu E0 = + 1.10 V
Calculating an Equilibrium Constant from
a Cell Potential
0.0591
0 1.10 log( )
2
volts K 
(1.10)(2)
log( )
0.0591
K
37.2 log( )K
37.2 37
10 1.58 10K x 

15. Conductance and Conductivity
NaCl

16. Conductance and conductivity
• Conductance of a solution is the inverse of its
resistance. Resistance is measured in ohms, W.
Conductance is expressed as the reciprocal of
ohms and is called the Siemens, S (1S = 1W-1)
• Resistance  L
A
where l is the distance between the electrodes
and A is the area of the electrodes
• Resistance =  L where  is resistivity
A
L
A

17. Conductance and conductivity
• Conductivity, k, is the inverse of resistivity so,
• k = 1 = L = c’ = S cm-1
 RA R
• Conductivity is dependent on the number of
ions, so we introduce molar conductivity
m = k ,where C is the molar concentration
C

18. Concentration dependence of
molar conductivities
• For strong electrolytes
molar conductivity
decreases slightly with
increasing concentration
For weak electrolytes, molar
conductivity is small at
moderate concentrations but
high at very low
concentrations. This is
because at low
concentrations the
proportion of dissociated
molecules is high.
m
0

19. • When the ions are far apart their
interactions can be ignored
• Limiting molar conductivity is due to
migration of cations (+ve ions) in one
direction and anions (-ve ions) in the other
under very dilute conditions
• m
0 = ++
0 + --
0
• +
0 and -
0 are respectively the ionic
conductivities of the cations and anions
Kohlrausch’s law gives the empirical relationship between
molar conductivity and concentration for strong electrolytes
Kohlrausch’s law

20. Electrolytic
Processes
A negative
cell potential, (-E0)
A positive free
energy change, (+G)
Electrolytic processes
are NOT spontaneous.
They have:
An electrolytic
process is the use
of electrolysis
industrially to refine
metals or compounds at a
high purity and low cost.

21. Electrolysis of
Water
2 22 4 4H O O H e 
  
2 24 4 2 4H O e H OH 
  
In acidic solution
Anode reaction:
Cathode reaction:
-1.23 V
-0.83 V
-2.06 V2 2 22 2H O H O 

22. Electroplating of
Silver
Anode reaction:
Ag  Ag+ + e-
1. Solution of the plating metal
3. Cathode with the object to be plated
2. Anode made of the plating metal
4. Source of current
Cathode reaction:
Ag+ + e-  Ag
Electroplating requirements:

23. Faraday's laws of electrolysis are
quantitative relationships based on the
electrochemical researches published
by Michael Faraday in 1834.

24. Faraday’s 1st Law Of Electrolysis
The amount of chemical reaction which
occurs at any electrode during
electrolysis by a current is proportional
to the quantity of electricity passed
through the electrolyte.
Amount of
electricity in
coulombs required
to reduce or
oxidized 1 mole of
substance
= n x F x 1mole

25. Faraday’s 2nd Law Of Electrolysis
The amounts of different substances
librated by the same quantity of
electricity passing through the electrolytic
solution are proportional to their chemical
equivalent weight .
Amount deposit (gm) = m x Q
F x nm = atomic mass of substance
Q = Total electric charge passed through substance
F = Faraday’s Constant
n = electrons transferred per ion

26. (1) Determination of equivalent masses of
elements
(2) Electron metallurgy
(3) Manufacture of non-metals
(4) Electro-refining of metals