Electroanalytical methods provide several advantages for quantitative analytical chemistry. They involve measuring the electrical properties of analyte solutions in electrochemical cells. Some key points: - Electroanalytical methods allow easy automation through electrical signal measurements. They can also determine low analyte concentrations without difficulty. - Electrochemical processes involve the transfer of electrons between substances during redox reactions. This occurs at the interface between electrodes and solutions in electrochemical cells. - Advantages include low cost compared to spectroscopy and the ability to easily automate measurements and detect low analyte concentrations through electrical signals.

1. Chapter-6
Electroanalytical methods
1

2. 2
Electrochemistry is a branch of chemistry that studies the
relations between chemical reactions and electricity, the
interconversion of chemical energy and electrical energy;
and study of redox reactions
•Electrochemical processes involve the transfer of electrons
from one substance to another
•Electroanalytical chemistry encompasses a group of
quantitative analytical methods that are based upon the
electrical properties of a solution of the analyte when it is
made part of an electrochemical cell.
•Advantages of electroanalytical methods:
• Measurements are easy to automate as they are electrical
signals
• Low concentrations of analytes are determined without
difficulty
• Far less expensive equipment than spectroscopy instruments

3. Redox Reaction
• This is a type of reaction in which electrons are transferred from one
substance to another.
• Oxidation: Loss of electrons or increase in the
oxidation number
Fe 2+ + e-  Fe3+
• Reduction: Gain of electrons or decreases in the
oxidation state
Cu2+ + 2 e-  Cu
• Redox reaction
Zn + Cu2+  Zn2+ + Cu
• Oxidizing agent(oxidant): Species that is being
reduced and causes an oxidation
• Reducing agent (reductant): Species that is being
oxidized and cause a reduction
3

4. Example:
Ce4+ + Fe2+ Ce3+ + Fe3+
Cerium Ce4+: an oxidizing agent/oxidant, electron acceptor.
Iron Fe2+ : an reducing agent/reductant, electron donor.
• Redox equations can be split into two half reactions:
Ce+4 + e-  Ce+3 (reduction reaction)
Fe+2  Fe+3 + e- (oxidation reaction)
----------------------------------
Ce+4 + Fe+2  Ce+3 + Fe+3 (over all reactions
Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq)
- Oxidizing agent
- Reduced species
- Electron gain
- Reducing agent
- Oxidized species
- Electron loss 4

5. Electrochemical Cell
• Oxidation-reduction reaction (redox reaction) can occur in
solution and in the electrochemical cell.
• Ordinary redox reaction in solution:
2Fe3+ + Sn2+  2Fe2+ + Sn4+
5

6. To harvest useful energy, the oxidizing and reducing agent has
to be separated physically in two different compartments so
as to make the electron passing through an external circuit
Reaction takes place at
electrode/solution interface
half-reactions:
oxidation / anode reaction:
Sn2+ - 2e-  Sn4+
reduction / cathode reaction:
2Fe3+ + 2e-  2Fe2+
6

7. Electrochemical Cell
• There are two types of electrochemical cells;
1. primary cell (Galvanic cell)
• It changes chemical energy
into electrical energy
• The reaction is spontaneous
2. Electrolytic cell
• It changes electrical
energy into chemical
energy
• The reaction is
nonspontaneous
7

8. 8

9. Galvanic cell
9

10. Important terms
• Charge (Q ): Results from imbalance between electrons
and protons in a metal, or between anions and cations in
a solution
Charge (q) of an electron = - 1.602 x 10-19 C
Charge (q) of a proton = + 1.602 x 10-19 C
Where ,C = coulombs
• Charge of one mole of electrons = (1.602 x 10-19 C)(6.022 x
1023/mol) = 96,485 C/mol = Faraday constant (F)
• The charge (q) transferred in a redox reaction is given by
q = n x F
• Current (I): The quantity of charge flowing past a point in an
electric circuit per second
I= q/time
Units: Ampere (A) = coulomb per second (C/s); 1A = 1C/s
10

11. • Potential: The potential at a point in space is the work done in
moving a unit charge to that point from infinity.
• Units of volts, V (=J/C); E = W/Q
• Potential Difference (or Voltage): The potential difference or
voltage is the difference between the potentials at two points, and
hence the work done in moving a unit charge from one point to
the other. Its unit is in Volts.
• The amount of energy required to move charged electrons
between two points
• Work done by or on electrons when they move from one point to
another . w = E x Q or E = W/Q
• Units: volts (V or J/C); 1V = 1J/C
• Resistance(R) ; R= E/I ;
• Units: Ω (ohm) or V/A
11

12. Electrode :- it is an electric conductor which conducts electrons into
or out of a redox reaction system. The electrode surface serves as a
junction between an ionic conductor(solutions) and an electronic
conductor(metal wires) in an electrochemical cells. There are two
types. Cathode and anode electrodes where reduction and oxidation
processes takes place respectively.
Salt bridge:- Connects the two half-cells (anode and cathode)
- Filled with gel containing saturated aqueous salt solution such as
KCl
- Ions migrate through to maintain electroneutrality (charge balance)
- Prevents charge buildup that may cease the reaction process
Cell notation: It is a short form of writing that represents a
electrochemical cell
12

13. Phase boundary: represented by one vertical line
Salt bridge: represented by two vertical lines
Zn(s) ZnSO4(aq) CuSO4(aq) Cu(s)
13

14. Electrode potentials
• It is the driving force for either reduction or oxidation half
reaction, when by convention, they are both written as
reductions.
Cu2+ + 2e- ↔ Cu
Ag + + e- ↔ Ag
• We cannot determine absolute electrode potentials but we can
determine relative electrode potentials (cannot just measure half a
cell)
• Therefore, potential of a cell could be calculated first using a
standard reference electrode for one of the half cell.
Potential of cell = Ecathode - Eanode
• Standard electrode potential is the potential of
electrode at standard conditions ( i.e. at 1 bar, 1
M and 25oc)
14

15.  There are different standard reference electrodes.
1. Standard Hydrogen Reference Electrode (SHE)
 This is the standard reference half-cell to measure all other half-
reactions against.
 SHE is a Gas electrode, made up of:
• Metal piece (Pt) coated with platinum black (large surface area). Pt
is in aqueous acid solution (HCl = 1M). Solution is saturated with
H2 (bubble) ;P=1atm. Metal is site of e- transfer only.
 Half reaction for SHE is : 2H+(aq) +2e-  H2(g)
 Shorthand: Pt, H2(p=1.00atm) | ([H+] = 1.00M) || (25C)
 can be the anode or cathode.
 This half-reaction is assigned 0.00V.
 Half-wave potential are always written as reduction reactions.
15

16. 16

17. 2. Standard Calomel Reference electrode
• Saturated Calomel Electrode (SCE)
- Composed of metallic mercury in contact with
saturated solution of mercurous chloride (calomel, Hg2Cl2)
- Pt wire is in contact with the metallic mercury
- Calomel is in contact with saturated KCl solution
E = +0.244 V at 25 oC
3. Silver/Silver Chloride Reference Electrode (Ag/AgCl)
- Consists of silver metal coated with silver chloride paste
- Immersed in saturated KCl and AgCl solution
E = +0.199 V at 25 oC
 etc.
17

18. Standard Reduction Potentials
Reduction Half-Reaction E(V)
F2(g) + 2e-  2F-(aq) 2.87
Au3+(aq) + 3e-  Au(s) 1.50
Cl2(g) + 2 e-  2Cl-(aq) 1.36
Cr2O7
2-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O 1.33
O2(g) + 4H+ + 4e-  2H2O(l) 1.23
Ag+(aq) + e-  Ag(s) 0.80
Fe3+(aq) + e-  Fe2+(aq) 0.77
Cu2+(aq) + 2e-  Cu(s) 0.34
Sn4+(aq) + 2e-  Sn2+(aq) 0.15
2H+(aq) + 2e-  H2(g) 0.00
Sn2+(aq) + 2e-  Sn(s) -0.14
Ni2+(aq) + 2e-  Ni(s) -0.23
Fe2+(aq) + 2e-  Fe(s) -0.44
Zn2+(aq) + 2e-  Zn(s) -0.76
Al3+(aq) + 3e-  Al(s) -1.66
Mg2+(aq) + 2e-  Mg(s) -2.37
Li+(aq) + e-  Li(s) -3.04
Ox.
agent
strength
increases
Red.
agent
strength
increases
18

19. Sign Convention for Electrode Potentials (IUPAC)
Sign of the electrode potential, E0 ,
– is positive when the half-cell behaves spontaneously
as the cathode.
– is negative when the half-cell behaves as an anode.
– is a measure of driving force for the half-reaction.
Positive sign - Cathodic (red) reaction is spontaneous.
19

20. Cell potentials
Cell potential or Cell voltage:
 It is the driving force (or chemical pressure) that pushes
electrons through the external circuit of an electrochemical
cell.
 It is also called electromotive force of the cell.
Potential
Cell
constant
Faraday
trans.
electrons
of
number
Energy
Free
where








cell
cell
E
F
n
G
nFE
G
eq
cell K
RT
nFE
G ln




 

2Ag(s)
Cu
2Ag
Cu(s) 2



 

Ecell is also related to the free energy of the reaction
20

21. – Cell potential is an
electrical potential
difference between the
two electrodes or half-
cells
• Depends on specific
half-reactions,
concentrations, and
temperature
• Under standard state
conditions ([solutes] = 1
M, Psolutes = 1 atm), emf
= standard cell potential,
Ecell
• 1 V = 1 J/C
21

22. Sign Convention
Ag
M)
(0.0200
Ag
M)
(0.0200
Cu
Cu
2Ag(s)
Cu
2Ag
Cu(s)
2
2







Volt Meter
0.412 V
+
-
Cu
M)
(0.0200
Cu
M)
(0.0200
Ag
Ag
Cu(s)
2Ag
Cu
2Ag(s)
2
2







Volt Meter
+
-
Oxid
Anode
Red
cathode
Oxid
Anode
Red
cathode
- 0.412 V
22

23. Half-Cells= Half-cell reaction
cathode
E
anode
E
Ag
M)
(0.0200
Ag
M)
(0.0200
Cu
Cu
2Ag(s)
Cu
2Ag
Cu(s)
2
2







anode
cathode
cell
left
right
cell
E
E
E
E
E
E




Ag
M)
(0.0200
Ag
Ag(s)
Ag




 e
Cu
M)
(0.0200
Cu
Cu(s)
2e
Cu
2
2





By convention, all Half-Cell reactions are
written as reduction half-reactions
23

24. Current Flows and Concentrations Change
24

25. Initial Potential
25

26. Concentrations Change Until Equilibrium is Obtained
26

27.  
Ag
1.00)
(
Ag
SHE
Ag
1.00)
(
Ag
1.00)
(
H
atm
1.0
H
Pt,
2Ag(s)
H
2
2Ag
(g)
H
2
2











a
a
a
V
E
E
E anode
cathode
cell 799
.
0




27

28. Concentration and Ecell
• With the Nernst Eq., we can determine the effect
of concentration on cell potentials.
Ecell = E°cell - (0.0591/n)log(Q)
• Example. Calculate the cell potential for the
following:
Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)
Where [Cu2+] = 0.3 M and [Fe2+] = 0.1 M
28

29. Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)
• First, need to identify the 1/2 cells
Cu2+(aq) + 2e- Cu(s) E°1/2 = 0.34 V
Fe2+(aq) + 2e- Fe(s) E°1/2 = -0.44 V
Fe(s) Fe 2+(aq) + 2e- E°1/2 = +0.44 V
Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s) E°cell = +0.78 V
29

30. • Now, calculate Ecell
Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s) E°cell = +0.78 V
Ecell = E°cell - (0.0591/n)log(Q)

Q 
Fe2
 
Cu2
 

(0.1)
(0.3)
 0.33
Ecell = 0.78 V - (0.0591 /2)log(0.33)
Ecell = 0.78 V - (-0.014 V) = 0.794 V
30

31. • If [Cu2+] = 0.3 M, what [Fe2+] is needed so that Ecell
= 0.76 V?
Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s) E°cell = +0.78 V
Ecell = E°cell - (0.0591/n)log(Q)
0.76 V = 0.78 V - (0.0591/2)log(Q)
0.02 V = (0.0591/2)log(Q)
0.676 = log(Q)
4.7 = Q
31

32. Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)
4.7 = Q

Q 
Fe2
 
Cu2
 
 4.7

Q 
Fe2
 
0.3
 
 4.7
[Fe2+] = 1.4 M
32

33. Current in electrochemical cells
Electroanalytical methods involve electrical currents
and current measurements. We need to consider the
behavior of cells when significant currents are
present.
Electricity is carried within a cell by the movement of
ions. With small currents, Ohm’s law is usually
obeyed, and we may write E = IR where E is the
potential difference in volts responsible for
movement of the ions, I is the current in amperes, and
R is the resistance in ohms of the electrolyte to the
current.
33

34. Types of Electroanalytical Methods
34

35. Potentiometry
• An electroanalytical technique works based on the
measurement of the electromotive force of an
electrochemical cell comprised of a measuring and a
reference electrode to determine the concentration of
analytes. It is without drawing appreciable current.
35

36. A reference electrode is an electrode that has the half-cell
potential known, constant, and completely insensitive to
the composition of the solution under study. In conjunction
with this reference is the indicator or working electrode,
whose response depends upon the analyte concentration.
• Potentiometry is used to;
– locate end points in titrations.
– Determine ion concentrations with ion-selective
membrane electrodes
– Measure the pH
– determine thermodynamic equilibrium constants such as
Ka, Kb,and Ksp.
36

37. In addition to reference and indicator electrodes
potentiometry includes;
• Salt bridge which is used to:
– Preventing components of the analyte solution from
mixing with those of the solution where the reference
electrode is found
– A potential develops across the liquid junctions at
each end of the salt bridge.
– Potassium chloride is a nearly ideal electrolyte for the
salt bridge because the mobility of the K+
ion and the Cl－
ion are nearly equal
37

38. Cu0 / Cu 2 +(0.1M) // Ag+ (0.2M) / Ag 0
Anode Cathode
Cu0 ↔ Cu 2 + + 2e-
E left = E0
Cu2+/Cu0+ 0.0591 / n log [Cu+2] / [Cu0]
= 0.337 + 0.0591 / 2 log (0.1 / 1)
= 0.307 volt.
E0
Cu2+/Cu0 = 0.337volt
38

39. Cu0 / Cu 2 +(0.1M) // Ag+ (0.2M) / Ag 0
Anode Cathode
Ag+ +e- ↔ Ag 0
E right = E0
Ag+/Ag0 + (0.0591 / n) log [Ag+] / [Ag0]
= 0.799 + 0.0591 / 1 log (0.2 / 1)
= 0.757 volt.
E0
Ag+ / Ag0 = 0.799 volt
Cu0 / Cu 2 + // Ag+ (0.2M) / Ag 0
Anode Cathode
E cell = 0.757 – 0.307 = + 0.45 volts
The reaction proceeds in the written direction. 39

40. Ag 0 / Ag+ (0.2M) // Cu 2 +(0.1M) / Cu0
Anode Cathode
E cell = 0.307 –0.757 = - 0.45 volts
The reaction proceeds in the opposite direction.
Cu 0 + 2 Ag + ↔ Cu2+ + 2 Ag0
40

41. • Example: Determination of Ag+
Pt0 / Fe2+(0.05M),Fe3+(0.25) // Ag+ (xM) / Ag0
Ecell = -0.106 volt
E0 Fe3+,Fe2+ = 0.771 volt
E0 Ag+ / Ag0 = 0.799 volt
Fe3+ +e- ↔ Fe2+
E left = E0
Fe3+,Fe2+ + 0.0591 / n log [Fe3+] / [Fe2+]
= 0.771 + 0.0591 / 1 log [0.025/ 0.05]
= 0.8123 volt.
41

42. Ag+ +e- ↔ Ag 0
E right = E0
Ag+/Ag0 + (0.0591 / n) log [Ag+] / [Ag0]
= 0.799 + 0.0591 / 1 log ( x / 1)
Ecell = E right – E left
-0.106 = {0.799 + 0.0591 log x} – 0.8123
Log Ag+ = - 1.56
[Ag+] = 0.027 M
E0
Ag+ / Ag0 = 0.799 volt
42

43. Polarography
Voltammetry is one of the electroanalytical methods which works
based on measurement of current as a function of the potential applied
to a small electrode. Unlike potentiometry measurements, which
employ only two electrodes, voltammetric measurements utilize a
three electrode electrochemical cell. The use of the three electrodes
(working, auxiliary, and reference) along with the potentiostat
instrument allow accurate application of potential functions and the
measurement of the resultant current.
1) working electrode; (2) auxiliary electrode; (3) reference electrode
43

44. • Voltammetry experiments investigate the half cell reactivity of an
analyte. Voltammetry is the study of current as a function of applied
potential. These curves I = f(E) are called voltammograms. The potential
is varied arbitrarily either step by step or continuously, and the actual
current value is measured as the dependent variable. The shape of the
curves depends on the speed of potential variation (nature of driving
force) and on whether the solution is stirred or quiescent (mass
transfer). Most experiments control the potential (volts) of an electrode
in contact with the analyte while measuring the resulting current
(amperes).
• To conduct such an experiment requires at least two electrodes. The
working electrode, which makes contact with the analyte, must apply
the desired potential in a controlled way and facilitate the transfer of
charge to and from the analyte. A second electrode acts as the other
half of the cell. This second electrode must have a known potential with
which to gauge the potential of the working electrode, furthermore it
must balance the charge added or removed by the working electrode.
While this is a viable setup, it has a number of shortcomings. Most
significantly, it is extremely difficult for an electrode to maintain a
constant potential while passing current to counter redox events at the
working electrode
44

45. • To solve this problem, the roles of supplying electrons and
providing a reference potential are divided between two
separate electrodes. The reference electrode is a half cell
with a known reduction potential. Its only role is to act as
reference in measuring and controlling the working
electrodes potential and at no point does it pass any
current. The auxiliary electrode passes all the current
needed to balance the current observed at the working
electrode. To achieve this current, the auxiliary will often
swing to extreme potentials at the edges of the solvent
window, where it oxidizes or reduces the solvent or
supporting electrolyte. These electrodes, the working,
reference, and auxiliary make up the modern three
electrode system.
• Working electrodes used: Hg, Pt, Au, Ag, C or others
• Reference electrode: SCE or Ag/ AgCl;
• Auxiliary electrode: Pt wire
45

46. Polarography
• The difference between polarography and other voltammetry ;In
polarography the working electrode is a dropping mercury
46

47.  Polarography is an voltammetric measurement whose response is
determined by combined diffusion/convection mass transport.
Polarography is a specific type of measurement that falls into the
general category of linear-sweep voltammetry where the electrode
potential is altered in a linear fashion from the initial potential to
the final potential. As a linear sweep method controlled by
convection / diffusion mass transport, the current vs. potential
response of a polarographic experiment has the typical sigmoidal
shape.
 A supporting electrolyte is a salt added in excess to the analyte
solution. Most commonly, it is an alkali metal salt that does not
react at the working electrode at the potentials being used. The
salt reduces the effects of migration and lowers the resistance of
the solution
 Dissolved oxygen is usually removed by bubbling nitrogen
through the solution
47

48. • There are three modes of mass transport to and from the
electrode surface: diffusion, migration, and convection.
• Diffusion from a region of high concentration to a region of low
concentration occurs whenever the concentration of an ion or
molecule at the surface of the electrode is different from that in
bulk solution.
• Convection occurs when a mechanical means is used to carry
reactants toward the electrode and to remove products from the
electrode.
• The most common means of convection is to stir the solution
using a stir bar. Other methods include rotating the electrode and
incorporating the electrode into a flow cell.
• Migration occurs when charged particles in solution are attracted
or repelled from an electrode that has a positive or negative
surface charge.
• Unlike diffusion and convection, migration only affects the mass
transport of charged particles
48

49. • The flux of material to and from the electrode surface is
a complex function of all three modes of mass transport.
• In the limit in which diffusion is the only significant
means for the mass transport of the reactants and
products, the current in a Voltammetric cell is given by
where n is the number of electrons transferred in the redox
reaction, F is Faraday's constant, A is the area of the
electrode, D is the diffusion coefficient for the reactant or
product, CbuIk and Cx=o are the concentration of the analyte in
bulk solution and at the electrode surface, and  is the
thickness of the diffusion layer.
49

50. • For the above equation to be valid, migration and
convection must not interfere with formation of diffusion
layer around the electrode surface.
• Migration is eliminated by adding a high concentration of
an inert supporting electrolyte to the analytical solution.
• The large excess of inert ions, ensures that few reactant
and product ions will move as a result of migration.
• Although convection may be easily eliminated by not
physically agitating the solution, in some situations it is
desirable either to stir the solution or to push the
solution through an electrochemical flow cell.
Fortunately, the dynamics of a fluid moving past an
electrode results in a small diffusion layer, typically of
0.001 - 0.01-cm thickness, in which the rate of mass
transport by convection drops to zero.
50

51. Coulometry
Coulometry is the general name for methods that measure
the amount of electricity required to react exactly with
an analyte. Measure the quantity of electrical charge
(electrons) required to convert a sample of an analyte
quantitatively to a different oxidation state. Coulometry
may be done at either constant potential or constant current
.
If the current is constant, the number of coulombs (Q) is equal
to the product of the current (i) in Amperes and time (sec),
that is
Q (columbs) = i (amps) x t (sec)
If the current varies during the electrolysis,
Q =  idt
51

52. Instead of weighing the substance plated on the electrode, coulometry
is based on measuring the number of electrons that participate in
a chemical reaction.
Coulometry is more versatile than electrodeposition, because they
include both electrochemical reactions in which a gas is formed and
those in which both the reactant and the product are soluble species.
Coulometry is based on Faraday’s law, which states that one faraday
of electricity will react with one equivalent weight of a reactant and
will yield one equivalent weight of a product.
The charge on an electron is defined as 1.6022 × 10–19 coulombs.
Total charge, q, in coulombs, passed during an electrolysis is related to
the amount of analyte by Faraday’s law
q = n · F
where F = 96,485.3415 C/mol,  n = q / F = I · t / F 52

53. Faraday’s law:
Total charge, Q, in coulombs passed during
electrolysis is related to the absolute amount of
analyte:
Q = nFN
n = #moles of electrons transferred per mole of
analyte
F = Faradays constant = 96487 C mol-1
N = number of moles of analyte
53
Example: A 0.3619-g sample of tetrachloropicolinic acid,
C6HNO2CI4, is dissolved in distilled water, transferred to a 1000-ml,
volumetric flask, and diluted to volume. An exhaustive
controlled-potential electrolysis of a 10.00-mL portion of this
solution at a spongy silver cathode requires 5.374 C of charge. What
is the value of n for this reduction reaction?

54. 54

55. Conductometry
 Conductometry means measuring the conductivity of
ionic solutions caused by mobility of ions towards
respective electrodes in presence of an electric field.
 Conductivity is measured by using conductometer.
Units of conductivity is mhos(Ω-1).
Those ionic compounds that are soluble in water and
conduct electric current in aqueous solution are called
electrolytes. The dissolution process consists of complete
dissociation of ionic compounds into mobile cations and
anions. There are many compounds, which though soluble
in water, do not exhibit any conductivity.
55

56. • These are termed nonelectrolytes. There is still another group of
compounds that exhibit conductance in solutions only when that
solution is quite dilute. Such compounds are known as weak
electrolytes. Solutions that contain large numbers of mobile ions
(cations and anions from the soluble ionic compounds) conduct
current well, and solutions that contain only a few ions (acetic
acid) or relatively immobile ions show poor conductivity.
• The conductivity of a solution varies with the number, size, and
charge of the ions constituting the solution. The viscosity of a
solution also affects the conductivity, by affecting the mobility of
the ions. Ions of different species in solution will therefore show
different conductivities. If, by means of a chemical reaction, we
replace one ionic species by another having a different size and/or
charge, we would observe a corresponding change in conductivity
of the resulting solution.
56

57. In general the conductance of the solution depends on:
1. Temperature:
It is increased by increase of temperature.
2. Nature of ions
size, molecular weight, number of charges the ion carries and other factors
3. The concentration of ions:
As the number of ions increases the conductance of the solution increases.
4. The size of the electrodes
L
A
K
G 
A
L
G
K 
L/A is cell constant
K is the specific conductance or conductivity
ohm-1cm-1 or seimen/cm.
57

58. 58

59. • Specific conductivity:-It is conductivity offered by a substance
of 1cm length and 1sq.cm surface area. units are mhos/cm.
• A very useful quantity is the equivalent conductivity. It is
defined as the value of the specific conductivity, k, contributed
by one equivalent of ions of either charge. More specifically, it
is defined as the conductance of a solution containing one
gram-equivalent of an electrolyte placed between electrodes
separated by a distance of 1 cm. If C is the concentration of the
solution in gram-equivalents per liter, the volume of the
solution in cubic centimeters per equivalent (cm3/equiv) is
equal to 1000/C. The equivalent conductance.  , is then given
by;
 = 1000K
C
59

60. • Another frequently used quantity in conductance measurements is
the molar conductance, defined as the conductance of a one cubic
centimeters volume of solution that contains one mole (or formula
weight) of the electrolyte. If M is the concentration of the solution
in moles per liter, then the volume in cubic centimeters per mole is
1000/M. The molar conductance is then given by ;
m = 1000K
M
• NB: Specific conductivity is the reciprocal of specific
resistivity (ρ)
60

61. Molar conductance of various ions at infinite dilution at 25℃
ions molar conductance
K+ 73.52
Na+ 50.11
Li+ 38.69
H+ 349.82
Ag+ 61.92
Cl- 76.34
Br- 78.4
OH- 198
• Total conductance of the solution is directly proportional to the sum of
the n individual ion contributions .
G = Ʃ ciλm 61

62. 62

63. APPLICATIONS OF CONDUCTOMETRY
It can be used for the determination of:-
 Solubility of sparingly soluble salts
 Ionic product of water
 Basicity of organic acids
 Salinity of sea water (oceanographic work)
 Chemical equilibrium in ionic reactions
 Conductometric titration
63

64. CONDUCTOMETRIC TITRATIONS:
 The determination of end point of a titration by means of
conductivity measurements are known as conductometric
titrations.
64

65. END OF THE COURSE !!!
65