2. What is
electrochemistry?
 Electrochemistry is the study of
chemical reactions which take place at
the interface of an electrode usually a
solid, metal or semiconductor and an
ionic conductor , the electrolyte.
 Electrochemistry deals with the
interaction between electrical energy
and chemical change.

3. History of electrochemistry
 English chemist john Daniel and physicist
Michael faraday both credited as founders
of electrochemistry today.
 The first germen physicist Otto von
Guericke created the electric
generater,which produced static electricity
by applying friction in the machine.
 The English scientist William Gilbert spent
17 years experimenting with magnetism
and to a lesser extent electricity.
Michael
faraday
john
Daniel

4.  The french chemist charles francois de cisternry du fay
had discovered two types of static electricity.
 William Nicholson and Johann Wilhelm Ritter
succeeded in decomposing water into hydrogen and
oxygen by electrolysis.
 Ritter discovered the process of electroplating.
 William Hyde Wollaston made improvements to the
galvanic cells.
 Orsted’s discovery of the magnetic effect of electrical
currents and further work on electromagnetism to
others.

5.  Michael Faraday's experiments led him to state his two
laws of electrochemistry and john Daniel invented
primary cells.
 Paul Heroult and Charles M.Hall developed an
efficient method to obtain aluminum using
electrolysis of molten alumina.
 Nernst developed the theory of the electromotive force
and his equation known as Nernst equation, which
related the voltages of a cell to its properties.
 Quantum electrochemistry was developed by Revaz
dogonadeze and his pupils.

6. Oxidation-Reduction
 The term redox stands for reduction-oxidation
 It refers to electrochemical processes involving
electron transfer to or from a molecule or iron
changing its states.
 The atom or molecule which loses electrons is known
as the reducing agent.
 The substance which accepts the electrons is called the
oxidizing agent.

7. Balancing redox reactions
 Acidic medium
 Example of manganese reacts with sodium bismuthate
 Unbalanced reaction:
Mn2+
(aq) + NaBiO3(s) → Bi3+
(aq) + MnO4
–
(aq)
 Oxidation:
4 H2O(l) + Mn2+
(aq) → MnO4
–
(aq) + 8 H+
(aq) + 5 e–
 Reduction:
2 e– + 6 H+
(aq) + BiO3
–
(s) → Bi3+
(aq) + 3 H2O(l)
8 H2O(l) + 2 Mn2+
(aq) → 2 MnO4
–
(aq) + 16 H+
(aq) + 10 e–
10 e– + 30 H+
(aq) + 5 BiO3
–
(s) → 5 Bi3+
(aq) + 15 H2O(l)
 Reaction balanced:
14 H+
(aq) + 2 Mn2+
(aq) + 5 NaBiO3(s) → 7 H2O(l) + 2 MnO4
–
(aq) + 5 Bi3+
(aq) +
5 Na+

8.  Basic medium
 Example of reaction between potassium permanganate and
sodium sulfite.
 Unbalanced reaction:
KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH
 Reduction:
3 e– + 2 H2O + MnO4
– → MnO2 + 4 OH–
 Oxidation:
2 OH– + SO3
2– → SO4
2– + H2O + 2 e–
 6 e– + 4 H2O + 2 MnO4
– → 2 MnO2 + 8 OH–
 6 OH– + 3 SO3
2– → 3 SO4
2– + 3 H2O + 6e–
 Equation balanced:
2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH

9.  Neutral medium
 Method to complete combustion of propane.
 Unbalanced reaction:
C3H8 + O2 → CO2 + H2O
 Reduction:
4 H+ + O2 + 4 e– → 2 H2O
 Oxidation:
6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+
 20 H+ + 5 O2 + 20 e– → 10 H2O
 6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+
 Equation balanced:
 C3H8 + 5 O2 → 3 CO2 + 4 H2O

10. Standard electrode potential
To allow prediction of the cell potential,
tabulations of standard electrode potential are available.
Tabulations are referenced to the standard hydrogen
electrode.
The standard hydrogen electrode undergoes the reaction

2 H+
(aq) + 2 e– → H2

11. Standard electrode potentials are usually tabulated
as reduction potentials.
The reactions are reversible and the role of particular
electrode in a cell depends on the relative oxi./red.
Potential of both electrodes.
The cell potential is then calculated as the sum of
reduction potential for cathode and the oxidation
potential for anode.
For example, the standard electrode potential for a
copper electrode is:
Cell diagram
Pt(s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) |
Cu(s)
E°cell = E°red (cathode) – E°red (anode)

12. Gibbs free energy and cell
potential
Though cell potential Cell and get electricity n faraday
in the cell:
= -nFEcell
For standard cell, this equation can we written
0= -RTlnK=-nFE0
G
cell
Though produce of electric energy converted into
electric work,
Wmax= Welectrical= -nFEcell

13. Nernst equation
n+|M)=E0
(M
E(M
n+|M)- ln
But solid M concentrate constant
n+|M)=E0
(M
E(M
n+|M)- ln
Example of Daniel cell
2+|Cu)=E0
(Cu
For cathode : E(Cu
2+|Cu)- ln
2+|Zn)=E0
(Zn
For anode : E(Zn
2+|Zn)- ln
2+|Cu) - E(Zn
Cell Potential : Ecell= : E(Cu
2+|Zn)
= E0
2+|Cu)- ln - E0
(Zn
(Cu
2+|Zn)- ln
= Ecell=E0
cell- ln

14.  Electrical resistivity
It is an intrinsic property that quantities how strongly a
given material opposes the flow of electrical current.
Many resistors and conductors have a uniform cross
section with a uniform flow of electric current and made
of one material
The electrical resistivity defined

15.  Electrical conductivity
The reciprocal of electrical resistivity, and measures a
material’s ability to conduct an electric current.
It is commonly represented by σ
Conductivity is defined as
Conductivity SI units of Siemens per meter.

16. Molar conductivity
Molar conductivity is defined as the conductivity of an
electrolyte solution divided by the molar
concentration of the electrolyte, and so measures the
efficiency with which a given electrolyte conducts
electricity in solution.
From definition, the molar conductivity

17. • Two cases should be distinguished:
Strong eletrolyte and weak electrolyte
 For strong electrolyte
Salts, strong acids and strong bases, the molar
conductivity depends only weakly on concentration.

18.  For weak electrolyte
The molar conductivity strongly depends on
concentration.
The more dilute a solution, the greater its molar
conductivity, due to increased ionic dissociation.
For weak electrolyte obeys Oswald's dilulation law.

19. Kohlrausch’s law of independent
migration of ions
High accuracy in dilute solutions, molar conductivity
is composed of individual contributions of ions.
Limiting conductivity of anions and cations are
additive, the conductivity of a solution of a salt is equal
to the sum of conductivity contributions from the
cation and anion
Λ0
m=v+ Λ0
+ + v- Λ0
-

20. Battery
Many types of battery have been commercialized and
represent an important practical application of
electrochemistry.
Early wet cells powered the first telegraph and
telephone systems, and were the source of current for
electroplating.
The zinc-manganese dioxide dry cell was the first
portable, non-spill able battery type that made
flashlights and other portable devices practical.

21. The mercury battery using zinc and mercuric oxude
provided higher levels of power and capacity than the
original dry cell for early electronic devices.
Lead-acid battery was secondary battery.
The electrochemical reaction that produced current
was reversible, allowing electrical energy and chemical
energy to be interchanged as needed.
Lead-acid cells continue to be widely used in
automobiles.

22. The lithium battery, which does not use water in the
electrolyte, provides improved performance over other
types.
Rechargeable lithium ion battery is an essential part of
many mobile devices.

23. Corrosion
Corrosion is the term applied to steel rust caused by an
electrochemical process.
Corrosion of iron in the form of reddish rust, black
tarnish on silver, red or green may be appear on copper
and its alloys, such as brass.

24. Prevention of corrosion
 Coating
Metals can be coated with paint or other less
conductive metals.
This prevents the metal surface from being exposed to
electrolytes.
Scratches exposing the metal substrate will result in
corrosion.

25. • Sacrificial anodes
The method commonly used to protect a structural
metal is to attach a metal which is more anodic than
the metal to be protected.
This forces the structural metal to be catholic thus
spared corrosion. it is called sacrificial.
Zinc bars are attached to various locations on steel
ship hulls to render the ship hull catholic.
Other metal used magnesium.

26. Electrolysis
The spontaneous redox
reactions of a conventional
battery produce electricity
through the different chemical
potentials of the cathode and
anode in the electrolyte.
Electrolysis requires an
external source of electrical
energy to include a chemical
reaction , and this process
takes place in a compartment
called an electrolytic cell.

27. Electrolysis of molten sodium
chlorine When molten, the salt sodium chloride can be
electrolyzed to yield metallic sodium and gaseous
chlorine.
This process takes place in a special cell named
Down’s cell.
Reactions that take place at Down's cell are the following
Anode (oxidation): 2 Cl– → Cl2(g) + 2 e–
Cathode (reduction): 2 Na+
(l) + 2 e– → 2 Na(l)
Overall reaction: 2 Na+ + 2 Cl–
(l) → 2 Na(l) + Cl2(g)
This process can yield large amounts of metallic
sodium and gaseous chlorine, and widely used on
mineral dressing and metallurgy industries.

28. Quantitative electrolysis and
Faraday’s law
Quantitative aspects of electrolysis were originally
developed by Michel faraday .
Faraday is also credited to have coined the terms
electrolyte.
Electrolysis among many others while studying
analysis of electrochemical reactions.
Faraday advocate of the law of conservation of energy.

29. First law
 The mass of products yielded on the electrodes was
proportional to the the value of current supplied to the cell,
the length of time the current existed, and the molar mass
of the substance analyzed.
 The amount of substance deposited on each electrode of an
electrolytic cell is directly proportional to the quantity of
electricity passed through the cell.
m=

30. Second law
The amounts of bodies which are equivalent to each
other in the ordinary chemical action have equal
quantities of of electricity naturally associated with
them.
The quantities of different elements deposited by a
given amount of electricity are in the ratio of the
chemical equivalent weights

31. Applied aspects of
electrochemistry
Industrial electrolytic processes
Electrochemical Reactors
Batteries
Fuel cells
Some Electrochemical Devices
Electrochemical Methods of Analysis

32. Branch of electrochemistry
 Photo electrochemistry
It is subfield of study within physical chemistry.
The interest in this domain is high in the context of
development of renewable energy conversion and
storage technology.
The effects of luminous radiation on the properties of
electrodes and on electrochemical reactions are the
subject of photo electrochemistry

33.  Semiconductor’s electrochemistry
Semiconductor material has a band gap and generates a
pair of electron and hole per absorbed photon if the
energy of the photon is higher than the band gap of the
semiconductor.
This property of semiconductor materials has been
successfully used to converted solar energy into electrical
energy by photovoltaic devices.
 Semiconductor-electrolyte interface
When a semiconductor comes into contact with a liquid,
to maintain electrostatic equillibrium
There will be a charge transfer between the
semiconductor and liquid phase,if formal redox potential
of redox species lies inside semiconductor band gap.

34. At thermodynamic eqilibrium, the fermi level of
semiconductor and the formal redox potential of redox
species and between interface semiconductor.
This introduce n-type semiconductor and p-type
semiconductor.
This semiconductor used as photovoltaic device similar to
solid state p-n junction devices.
Both n and p type semiconductor can used as photovoltaic
devices to convert solar energy into electrical energy and
are called photoelectrical cells

35.  Boielectrochemistry
It is branch of electrochemistry and biophysical
chemistry concerned with topics like cell electron-proton
transport, cell membrane potentials and
electrode reactions of redo enzymes.
Bioelectrochemistry is a science at the many junctions
of sciences.

36. Nanoelectrochemistry
Nanoelectrochemistry is a branch of electrochemistry
that investigates the electrical and electrochemical
properties of materials at the nanometer size regime.
Nanoelectrochemistry plays significant role in the
fabrication of various sensors, and devices for detecting
molecules at very law concentrations.

37. The term electrochemical nanostructuring can be used
to mean different things.
This term is employed to refer to generation at will of
nanostructure on electrode surface, involving a given
positioning with a certain precision
The term nanostructure is used to describe the
generation of nanometric patterns with move or less
narrow size distribution and a periodic or random
ordering on the surface.
But without control on the spatial location of the
nanostructure.

38. Application of electrochemistry
There are various extremely important electrochemical
processes in both nature and industry.
The coating of objects with metals or metal oxides
through electro deposition and the detection of alcohol in
drunken drivers through the redox reaction of ethanol.
Diabetes blood sugar meters measure the amount of
glucose in the blood through its redox potential.

39. The generation of chemical energy through
photosynthesis in inherently an electrochemical process.
Production of metals like aluminium and titanium from
their ores.
 For Photo electrochemistry
Artificial photosynthesis
Regenerative cell or Dye-sensitized cell
Photo electrochemical splitting of water

40.  For Boielectrochemistry
Some of different experimental techniques that can be
used to study bioelectrochemical problems.
Ampermetic of biosensors
Biofuel cells
Bioelectrosynthesis