Potentiometric titration
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Transcript of Potentiometric titration
POTENTIOMETRICTITRATION
Siham AbdounMsc., PhD
1. Introduction:
Potentiometric method include two type of
measurement these are ;
1. Direct measurement of an electrode potential
from which concentration of an active ion may
be found
2. Change of E.M.F. of an electrode cell brought
about by the addition of titrant
Both methods are based on quantitative
measurement of E.M.F. of cell is given by
E cell = E reference + E indicator +E junction
As the reference electrode potential is
independent of solution and junction potential is
constant so the cell potential is measure of
indicator electrode potential and can give
information on the nature and concentration of
substance under test.
In the potentiometric titration the titration
reaction is followed by measurement of
concentration of one or more species
potentiometrically. The beaker or flask
becomes one of the half cells and the
reference electrode is the other half cell.
There is a different between titration reaction
and cell reaction; in titration reaction the
reactant and products be in the same half cell
and the titration reaction always at
equilibrium while in cell reaction is not at
equilibrium.
In potentiometric titration the change in
electrode potential upon the addition of titrant
are noted by the volume of titrant added. At the
end point the rate of change of potential is
maximum
The potentiometric end point has been applied
to all types of chemical reaction. It can be used
with colored or opaque solution
2. Instrumentations:
There are three types of instrumentations
which are used for measurement of
potentials these are:
1. Non – electronic instruments
2. Electronic instruments
3. Automatic instruments
1. Non – electronic instruments
A potentiometer for titrations can be made from
simple instrument of battery for current supply,
two dry cell, a resistance and voltmeter; the
instrument can be operated by dipping the
electrodes in the sample solution and record the
voltmeter reading.
2. Electronic instruments
These instruments have many advantages over
non- electronic instruments
3. Automatic instruments
The use of manual instrument to locate the
end point and to draw a titration curve is
time consuming and boring job; so
automatic instrument for recording and
performing titration curve provides a logical
solution
3. Types of Potentiometric titration:
Potentiometric titrations may be applied to
different type of reactions of these are;
acid – base, oxidation reduction,
precipitation and complexmetirc
1. Acid – Base titrations:
The neutralization of acid or base is always
accompanied by the changes the
concentration of H+ and OH- ions.
In these reactions hydrogen electrode is used
as indicator electrode and N- calomel
electrode as a reference electrode.
A known volume of the acid titrant is kept in a
beaker with continuous stir; the hydrogen and
N- calomel electrode are connected by the salt
bridges and connected to a potentiometer which
record the EMF of the solution, into the beaker
the after the addition of base from burette the
values of EMF are plotted against volume of
titrant added and a curve are obtained.
The potential of a hydrogen electrode is given by :
E+ E⁰-0.0591 log a H+
Where E⁰ is standard electrode potential and pH is - logaH+
E+ E⁰+ 0.0591 pH
As the standard electrode potential is constant so the
cell potential or EMF is proportional to the change of
pH during the reaction.
The point where E.M.F increased rapidly is the end
point.
Amore sensitive and precise method for measure
end point is to plot the slope of curve against
volume as the slope is maximum at equivalence
point, the maximum value give the end point
2. Complexmetric titration
A metal electrode is used whose ions are
involved in the complex formation, example
silver electrode is used to measure cyanide ion
with standard solution of silver
Ag++ 2CN- (Ag (CN)2)-
K= (Ag+)(CN-)2
(Ag(CN)-2)
In this case, solid silver cyanide begins to get
precipitated soon after the equivalence point.
The further addition of silver neither changes
the concentration of the complex nor changes
the silver ion to any extent, so that the curve
has an almost horizontal portion shortly after
the equivalence point
In many Complexmetric reactions the situation
cannot be handled so easily because more than
one complex is formed. Thus, the reactions in
the case of mercuric ion with cyanide are:
Hg2++3CN- Hg (CN)3-
Hg2++4CN- Hg (CN)42-
However, these situations have became widely
used because of the discovery of the metal
chelating agents such as EDTA
3. Oxidation-reduction titrations
Redox reactions can be followed by an inert
indicator electrode. The electrode assumes a
potential proportional to the logarithm of the
concentration ratio of two oxidation states of
the reactant or the titrant whichever is
capable of properly poising the substance
being oxidized to substance being reduced.
For example ,
Ca4+ + Fe2+ Ce3+ + Fe3+
It is generally considered that such a
reaction consists essentially of two half
reactions whose standard potentials may be
used to calculate the standard potential of
the reactions.
Fe+ Fe3+ + e, Eo = -0.67 V ..……. (i)
Ce4+ e Ce3+ , Eo = +1.61 V …….. (ii)
Ce4+ + Fe2+ Fe3++ Ce3+, Eo = +0.85 V… (iii)
The equilibrium constant , K, of any reaction
may be calculated from the following formula :
Eo = log10 K ……. (iv)
Where Eo is the number of equivalent of
electricity associated with one molar unit of
reaction .
If an acidic ferrous solution is titrated with a
standard ceric solution at 25o C, the potential of a
platinum electrode in contact with the solution will
be given by either of the following equations:
E=Eo ce4+/ ce3+ - 0.0591 log10 [ce3+] …. (v)
1 [ce4+]
E=Eo Fe3+/ Fe2+ - 0.0591 log10 [Fe3+] … (vi)
1 [Fe2+]
It would be more convenient to use the
equation (vi) before the equivalence point, as
the right hand term of this equation could be
easily found from the known extent of the
titration . if equation (v) is used , then the
[Ce3+] / Ce4 ) ratio has to be calculated by
means of the equilibrium constant, After the
equivalence point, calculations are done by
means of equation (v).
From equation (vi) it is evident that the
potential that the potential at the start of the
titration should be—co, because Fe3+ ions are
the only ions present and there are no Fe3+
ions.
At the mid-point of the titrations, where
[Fe2+]=[Fe3+], equation (vi) because :
E=Eo Fe3+/ Fe2+ ……(vii)
At the equivalence point, the concentration of
unchanged ferrous ions will be equal the
concentration of the unchanged ceric ions.
Similarly , the concentration of cerous ions will
be equal to the concentration of ferric ions.
Thus, it can be concluded that;
[Fe2+] =[Ce4+] ……(viii)
[Fe3+] [Ce3+]
Now, k = [Fe3+][Ce3+] ……(ix)
[Fe2+][Ce4+]
At equivalence point,
[Fe3+][Ce3+]= k……(x)
[Fe2+][Ce4+]
On combining equations (v) and (vi) with
equation (x) , we get
Eep = EoCe4+ / Ce3+ - log10 K …..(xi)
Eep = EoFe
3+ / Fe2+ + log10 K ……(xii)
On adding, equations (xii), we get
Eep = EoFe
3+/ce3++EoFe
3+/Fe2+
2
where Eep= End point potential .
Oxidation-reduction titration may be used in
procedures such as monitoring of cyanide
wasters from metal plating industries or
chlorine compounds in bleach compounds
manufacturing, and the used of these bleach
compounds in paper man fracturing. They are
also extensively used in water pollution,
sewage treatment, agricultural and biochemical
studies.
4. Precipitation titrations. Any precipitation
titration that involves insoluble salts of metals
such as mercury, silver, lead and copper may be
followed potentiometrically . The indicator
electrode may be made of the metal involved in
the reaction or may be an electrode whose
potential is governed by the concentration of the
anion being precipitated
The magnitude of the potential change at the
end point depends on the solubility of the
substance being precipitated as well as the
concentration involved. The titration of
chloride ions with a standard solution of silver
nitrate using a silver metal indicator electrode
is an example of a precipitation titration.
The other electrode to complete the cell is
unimportant, provided that it is a true reference
electrode, i.e., it maintains a constant potential.
In the above case it will be assumed that the
normal hydrogen electrode (N H E.) is used and
this assumption is convenient because standard
potential may be used directly. The potential of
silver electrode will be governed by the
appropriate Nernst equation :
EAg+/ Ag = Eo
Ag+/ Ag + log10 [Ag+] …..(i)
As soon as enough silver nitrate to precipitate Cl
as AgCl has been added, the following equilibrium
is established,
AgCl Ag+ + Cl- ………….(ii)
The equilibrium constant for the above reaction is
KAgCl = [Ag+][Cl-]=10-10 …………..(iii)
If 0.1 N sodium chloride is titrated against 0.1 N
silver nitrate, the silver ion concentration may
be considered to be 10-9 N as soon as few drops
of silver nitrate have been added. Equation (i)
can be used to calculate the indicator electrode
potential.
EAg+/Ag=0.08 V + 0.0551 log10 10-9 = 0.2681
…….. (iv)
Similarly, half wave through the titration will be
when the chloride ions concentration has been
reduced to 0.033 N.
EAg+/Ag=0.08 V+0.0591 log 10 (3 10-9) = 0.30 V
At the equivalence point, [Ag+]=[Cl-] = 10-5N
EAg+/Ag=0.08 V+0.0591 log 10 10-5 = 0.50 V
4. Non – aqueous titration
The potentiometric method has been found to
be useful for carrying out titrations in non-
aqueous solvents. The ordinary glass- calomel
electrode system can be used
Generally the millivolt scale of the
potentiometer rather than the pH scale should
be employed because the potential in non-
aqueous titration may exceed the pH scale.
Advantages of Potentiometric titrations
over 'classical' visual indicator methods are:
1. Can be used for coloured, turbid or
fluorescent analyte solution.
2. Can be used if there is no suitable indicator
or the colour change is difficult to ascertain.
3. Can be used in the titration of polyprotic
acids, mixtures of acids, mixtures of bases or
mixtures of halides. loured, turbid or
fluorescent analyte solution
4. The apparatus required is inexpensive, reliable
and readily available.
5. It is easy to interpret the titration curve.