Potentiometry:
Is a method of analysis in which we determine
the concentration of an ion or a substance by dipping a
suitable sensor in its solution (indicator electrode).
 The potential of the indicator electrode is measured relative to
reference electrode possessing constant potential .
 The concentration of the ion is determined from Nernest
equation.E25°C= Eo - 0.059/n log 1/ [Mn+]
 The two electrodes form the two half cells of electrochemical
cell , the electromotive force (e.m.f.) of which is the
algebraic difference of the two electrode potential (cathode &
anode ).
 Since the reference electrode potential is constant the change
in e.m.f. of the cell is due to the change of the indicator
electrode potential which in turn is due to the change of the
concentration of the ion to be determined.
1-Galvanic Cell (voltaic cell):
Where the chemical energy is converted to electrical energy which can
be supplied to an external circuit i.e. it produces electrical energy.
e.g. Daniel cell:
Is composed of:
1- A copper electrode dipped in its solution.
2- A zinc electrode dipped in its solution.
3- Both solutions are connected by salt bridge.
4- The electrodes are connected by external circuit.
Zn metal dissolves in its solution and transfers to Zn ion while copper
ions are deposited on the electrode
so a potential difference is produced
leading to the production of current.
Classification of the electrodes
They are classified according to
Type
Function
Reference
First type
Second type
Third type
Redox type
1-Normal
hydrogen
2-Ag/AgCl
3-Calomel
Indicator
21- Redox
Complexometric
reaction
reaction
34- Acid base
Preceptimetric
reaction
reaction
Mode of action
Electron transfer
Ion- Selective
Electrodes used in potentiometric measurements:
Classification of electrodes:
A) According to the type:
1) First type (kind or order):
In this type electrode potential depends on equilibrium between
element and its ions (i.e. two phase electrode).
(Mo / Mn+)
Mo ↔ Mn+ + ne
E25oC = Eo - 0.059/n log 1 / [Mn+]
From Nernest equation E α [Mn+]
e.g. Titration of Cu2+ with EDTA
E
At the begining Cu2+ ions conc is very high so E
is high. as E α [Cu2+], then E decreases with the
E.p.
consumption of copper with EDTA
mls of titrant
It is used for any metal except Fe, Cr, Co, Ni WHY?
Because: 1- They are liable to oxidation by air so they give
irregular potential.
2- They suffer from crystal deformation and prescence
of oxide coating.
2) Electrode of Second type:
 In this type equilibrium is established between three phases.
 The electrode is a metal electrode which is covered with a layer of its
sparingly soluble salt immersed in a solution of electrolyte.
 This electrode is used for determination of anions (non metal ions).
 The electrode potential is indirectly affected by the analyte.
e.g. Ag0 electrode: (Ag0 / AgCl / Cl-)
EAg0 α 1/ [Cl-]
i.e. Ag0 is coated with a layer of silver chloride and dipped in chloride
solution.
Calomel electrode: (Hg0 / Hg2Cl2 / Cl-)
EHg0 α 1/ [Cl-]
Mercury electrode: (Hg0 / Hg2y2- / y4-)
EHg0 α 1/ [Y4-]
 Similar electrodes for bromide, iodide………can be used.
 In Ag electrode; the potential depends on the chloride concentration
in the solution and this is used for its determination. The electrode
potential is:
E25 = E0 - 0.059 log [Cl-]
i.e Potential is inversely proportional to non metal ion concentration.
e.g. titration of Cl- with Ag+
Before titration: conc of Cl- is high
So E is low
During titration:conc of Cl- decreases
So E increases
E
At e.p. conc of Cl is zero
So sharp increase in E
mls of titrant
3) Electrode of Third type:
 In this type equilibrium is established between four phases.
 E.g. We want to determine Ca
2+
ions in CaCl2 but we have Pb rod
 E.g. Pbo / PbC2O4 / CaC2O4 / CaCl2
Represented by the following equations:
PbC2O4
Pb2+ + C2O42Lead acetae and calcium acetate
CaC2O4
Ca2+ + C2O42supresses the ionization of each
CaCl2
Ca2+ + 2Clother by common ion
 Increase in [Ca2+] will supress dissociation of CaC2O4, leading to
lower availability of C2O42- .
 In this case stress of C2O42- on PbC2O4 decreased leading to more
dissociation of PbC2O4 producing more Pb2+ and consequently
potential developed on Pbo is increased.
 EPbo α [Ca2+]
4) Redox electrode:
 In this type equilibrium is established between oxidant and
reductant in the same solution in presence of inert electrode e.g
pt wire (i.e it is one phase system electrode).
 e.g. Fe3+ / Fe2+ in presence of pt
 E25oC = Eo - 0.059/n log Fe2+ /Fe3+
Ce4+/Ce3+
 E25oC α 1/ [red] / [oxid]
 E.g. Determination of Fe2+ with Ce4+
E
Fe3+/Fe2+
E.p.
mls of titrant
B) According to the Function:
Electrodes can be classified according to the function into:
1) reference electrodes and 2) indicator electrodes
Reference electrodes
 In this type electrode must have known and constant potential,
therefore used as a reference to measure the potential of
indicator electrode through galvanic cell.
 e.g a)Normal hydrogen electrode (N.H.E.),
b) calomel electrode,
c) silver, silver Chloride electrode.
A) Normal hydrogen electrode (N.H.E.)
1-Half cell presentation: Pt, H2 / H+ (1.0 N) //
2-Type :1st type
(gas its conc is
3-Electrode reaction: 2H+ + 2e ↔ H2
considered 1)
4- Nernest equation:
E25oC = Eo - 0.059/n log H2 / [H+]2
E25oC = Eo + 0.059 log [H+]
platinum
E α [H+] in N.H.E [H+] = 1 , log 1 is zero
coated with
So E25oC = Eo = zero
Pt black
Salt
bridge
[H+] = 1 M
5-Advantage: It is a primary reference electrode as its potential is
zero.
6-Disadvantages: 1- It is difficult to keep H2 gas at one atmosphere
during all determinations.
2- it needs periodical re-plating of pt sheet with
pt black.
The Secondary Reference Electrodes:
1- The Calomel Electrode:
1-Half cell presentation: Hg0/Hg2Cl2 / KCl (saturated, 1 N or 0.1N) //
2- Type :2nd type
3- Electrode reaction: Hg22+ + 2e ↔ 2Hg0
4-Nernst Equation
:E25oC = Eo - 0.059/2 log Hg0 / [Hg22+]2
From the Nernst equation it is clear that the electrode potential depends
on the [Hg22+]; which is produced from the dissociation of the very small
amount of the sparingly soluble Hg2Cl2.
Hg2Cl2 ↔ 2Hg22+ + 2 Cl
The mercurous ion concentration is controlled
by the solubility product of the salt KspHg2Cl2
KspHg2Cl2 = [Hg22+]2 [Cl-]2
[Hg22+]2 = KspHg2Cl2 / [Cl-]2
E25oC = Eo - 0.059/2 log [Cl-]2 / KspHg2Cl2
E α 1/ [Cl-] electrode potential depends on the [Cl-] **
Although maximum potential is obtained on
using 0.1N KCl; however, we usually use
Sat. KCl WHY?
2- Silver Electrode:
1-Half cell presentation: Ag0/AgCl /,KCl (saturated, 1 N or 0.1N) //
2- Type :2nd type
3- Electrode reaction: Ag+ + e ↔ Ag0
4-Nernst Equation : E25oC = Eo - 0.059/1 log Ag0 / [Ag+]
As explained under calomel electrode, the Ag+ ion concentration is
controlled the solubility product of the salt KspAgCl
Ksp AgCl = [Ag+] [Cl-]
AgCl
[Ag+] = Ksp AgCl / [Cl-]
Ag
E25oC = Eo - 0.059/1 log [Cl-] / Ksp AgCl
E α 1/ [Cl-]
electrode potential depends on the [Cl-] **
Sat KCl solu
Indicator electrodes
An indicator electrode is that which its potential is sensitive to
the concentration of one of the participant or products of reaction.
 Its potential changes rapidly with the change of the particular
ion.
It must:
1- Give rapid response.
2- Its response is reproducible.
Now when it is connected with a reference electrode, the
change in e.m.f. of the formed cell will be due to the change of
indicator electrode potential, which in terns reflects the change of
the ion under investigation.
Indicator electrodes are classified into two classes:
1- Electrodes where electron transfer takes place at the
electrode surface.
i.e. redox reaction occurs (metallic electrodes).
2- Electrodes where (charge) ion exchange takes place at
a specific membrane surface.
i.e. ion selective electrodes or specific ion electrodes.
N.B:
1- We must notice that electrodes potential are related to
activity.
2-The most important application of this method is the
determination of pH.
3- The observed potential of the cell employed during the
measurement is:
Eobs = Eref + Ej – Eind
where Eref , Ej and Eind are the potentials of reference
electrode, liquid junction and indicator electrode.
Liquid junction potential
 The disadvantage of the cell with liquid junction is there is a
potential associated with the liquid junction, called the liquid
junction potential.
 A potential is developed at both boundaries of the junction
 This liquid junction potential arises from difference in rate of
migration of anion and cations of the bridge salt and the
electrolytes in the electrodes solutions. This difference results in
unequal charge distribution at the boundaries, thus a potential
is developed.
 It depends upon:
1- Charge mobilities
2- concentration of the ionic species
3- Nature of the solvent on each side of the boundary
4- pH dependent
 The presence of liquid junction potential limit
the accuracy of potentiometric measurements.
 How to minimize liq.junct. Pot.?
By using a high concentration of a salt, its
ions has nearly equal mobility e.g. KCl or
KNO3 (K+ = 74, Cl- = 73, NO3- = 76)
Use high concentration of salt for preparation of the
bridge to reduce the effect of difference in migration
rates of other ions in the electrode solution.
Ecell = Eind - Eref + Ej
Indicator Electrodes Used For Different
Reactions:
1- Electrodes for Redox Reactions:
These are inert electrodes which acquire the potential of
the system in which they are dipped.
They are made of platinum or gold, in the form of foils, coils
or plates WHY?
to expose a large surface area to the solution
thus attain the equilibrium potential rapidly.
It is one phase system electrode and is connected to
Ag/AgCl or calomel electrode to measure the e.m.f.
produced.
2- Electrodes for complexometric Reactions:
We may use:
A- Electrodes of the First type:
For the determination of a cation (metal ion); where the best electrode
used for a cation is its elemental form.
e.g. in determination of Cu2+ with EDTA , a copper rod is the indicator
electrode, its potential is directly affected by the analyte concentration
as indicated from:
Nernest eq: E25oC = Eo - 0.059/n log 1 / [Cu2+] so E α [Cu2+]
B- Electrodes of the Second type:
For the determination of a mixture of metals
(Cu2+,Cd2+, Bi3+ by EDTA)
We can not dip three different metal rods,
So we use an electrode that is affected by
the titrant instead of the sample.
We use: (Hgo / Hg2y2- / y4-)
EHgo α 1/ [Y4-]
E
E.p.
mls of titrant
3- Electrodes for precipitemetric Reactions:
We may use:
A- Electrodes of the First type:
For the determination of a cation (metal ion); where the best electrode
used for a cation is its elemental form.
e.g.1- in determination of Ag+ with chloride , a silver rod
is the indicator electrode, its potential is directly
E
affected by the analyte concentration as indicated
from:
Nernest eq: E25oC = Eo - 0.059/n log 1/ [Ag+] so
E α [Ag+]
mls of titrant
e.g.2- in determination of mix of Cl- & Br- & I- with AgNO3
So better to choose an indicator electrode which
responds to the titrant i.e. silver rod.
Before titration: there is no Ag+ ions so E is low
At e.p. :conc of Ag+ ions increases
So E increases sharply
E
mls of titrant
3- Electrodes for precipitemetric Reactions:
We may use:
B- Electrodes of the Second type:
For the determination of Cl- or Br- or I- with AgNO3
Det. of anion so use
e.g. Ag0 electrode: (Ag0 / AgCl / Cl-)
EAg0 α 1/ [Cl-]
E
mls of titrant