This document discusses various electroanalytical techniques. It begins by introducing electroanalytical techniques and their objectives. It then describes different types of techniques including coulometry, amperometry, voltammetry, polarography, potentiometry, conductometry, and others. It also discusses electrochemical cells, potentials in electroanalytical cells using the Nernst equation, and mass transfer processes like migration, convection, and diffusion.

1. Electro-Analytical Techniques
Sushil D.Patil
Bhujbal Knowledge City
MET Institute of Pharmacy,Nashik

2. • understand the basic principles of analytical
techniques
• understand the basic concepts involved in
electro-analytical techniques and its types.
Syllabus:
Introduction & types of Electro-Analytical
Techniques, Electrochemical cell, potentials in
electro-analytical cell& its measurement (Nernst
Equation),current-potential relationships, mass
transfer by migration, convection and diffusion.
Objectives

3. Types of Electro-Analytical Techniques
Introduction
Coulometry-Faraday’s laws of electrolysis
Amperometry
Voltammetry
Polarography-
Potentiometry
Conductometry
Oscillometry-
Condance,dielectric constant by
use of High frequency titration
Chronopotentiometry

4. • Bulk methods- Conductometry
• Interfacial methods-a.Static methods- Potentiometry
b.Dynamic methods -
1.Controlled Potential-
i.Constant Potential coulometry
ii.Votametry
iii.Amperometric Titration
iv.Electrogravimetry.
2.Constant Current- i. Coulometric Titration
ii.Electrogravimetry
Types of Electro-Analytical Techniques

5. Electrochemical Cells
The energy required to effect the
reaction is drawn from the external
power source.
Cu +2 + 2e → Cu
H2O →½ O2 + 2H+ + 2e
Oxidation reaction occur at the anode,
Reduction reaction occur at the cathode,
a. Reduction; Cu +2 + 2e → Cu
b. Oxidation; H2O →½ O2 + 2H+ + 2e

6. Electrochemical
cell
Galvanic cell

7. Potentials in an Electrochemical cell
• Ecell = Oxidation potential + Reduction potential
• Ecell= Ecathode - Eanode + Ejunction
• In order to make it easier to describe a given
electrochemical cell, a special symbolic
notation has been adopted. In this notation
the cell we described above would be
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

8. Mass transfer step:
1. Migration
• The process by which ions move under the
influence of an electric field is called migration
• The rate at which ions migrate to or away from
an electrode surface generally increases as the
electrode potential increases.
• This charge movement constitutes a current,
which also increases with potential.

9. Migration is the movement of ions through a
solution as a result of electrostatic attraction
between the ions and the electrodes.

10. 2. Convection
• Reactants can also be transferred to or from an
electrode by mechanical means.
• Forced convection, such as stirring or agitation,
will tend to decrease the thickness of the
diffusion layer at the surface of an electrode and
thus decrease concentration polarization.
• Natural convection resulting from temperature
or density differences also contributes to the
transport of molecules to and from an electrode.

11. 3. Diffusion
• When there is a concentration difference
between two regions of a solution, ions or
molecules move from the more concentrated
region to the more dilute. This process, called
diffusion, ultimately leads to a disappearance of
the concentration difference.
• The rate of diffusion is directly proportional to
the concentration difference.

13. Nernst Equation
• Nernst Equation is used to calculate the voltage
of an electrochemical cell. Nernst Equation which
relates the cell potential to the standard
potential& to the activities of the electro active
species.
E = E0- RT/nf In Q
where Ecell = Cell Pot.Non std Condition
E0cell = Cell Pot. Std. C.,R= gas con.,T= tem.,n=
number of moles of electrons exchenged in the
electrochemical reaction,F= Faraday’s
constant,Q=reaction quotient
E = E0- 0.0591/n log Q

14. • we reduce the concentration of Zn2+ in the Zn/Cu cell from its
standard effective value of 1M to an to a much smaller value:
• Zn(s) | Zn2+(aq, .001M) || Cu2+(aq) | Cu(s)
• This will reduce the value of Q for the cell reaction
• Zn(s) + Cu2+ → Zn2+ + Cu(s)
• thus making it more spontaneous, or "driving it to the right" as
the Le Châtelier principle would predict, and making its free
energy change ΔG more negative than ΔG°, so that E would be
more positive than E°.
• The relation between the actual cell potential E and the
standard potential E° is developed in the following way.which
relates the standard free energy change (for the complete
conversion of products into reactants) to the standard
potential
• ΔG° = –nFE°
• By analogy we can write the more general equation
• ΔG = –nFE

15. • which expresses the change in free energy for any
extent of reaction— that is, for any value of the
reaction quotient Q. We now substitute these into the
expression that relates ΔG and ΔG° which as chemical
equilibrium:
• ΔG = ΔG° + RT ln Q
• which gives
• –nFE = –nFE° + RT ln Q
• which can be rearranged to
(1)
The Nernst equation is more commonly written in base-
10 log form and for 25°C: