Analytical class, potetiometry and conductomtry study

1. ACSS-501
Class - 15
Potentiometric titrations
And
Conductometric titrations
Pabitra Kumar Mani;
Assoc. Prof.;
BCKV;

2. Potentiometry
An electroanalytical technique
based on the measurement of
the electromotive force of an
electrochemical cell comprised
of a measuring and a reference
electrode.
Indicator
electrode
The simplest example of a
measuring electrode is a metal
electrode whose potential
depends on the concentration
of the cation of the electrode
metal.
Electrochemical measuring system.
When a metal M is immersed in a solution containing its own ions Mn+, then
an electrode potential is established, the value of which is given by the
Nernst equation:

3. General Principles
Reference electrode | salt bridge | analyte solution | indicator electrode
Eref
Ej
Eind
Ecell = Eind – Eref + Ej
Reference cell :
a half cell having a known electrode
potential
Indicator electrode:
has a potential that varies in a
known way with variations in
the concentration of an analyte
A cell for potentiometric determinations.

4. Reference electrode
: maintains a fixed potential :
a half cell having a known electrode potential
1) Saturated calomel electrode (S.C.E.)
Hg(l) | Hg2Cl2 (sat’d), KCl (sat’d) | |
electrode reaction in calomel hal-cell
Hg2Cl2 (s) + 2e = 2Hg(l) + 2Cl–
Eo = + 0.268V
E = Eo – (0.05916/2) log[Cl–]2 = 0.244 V
Temperature dependent

5. Fig. 21-2. Diagram of a typical
commercial saturated calomel electrode.
Fig. 21-3. A saturated calomel
electrode made from materials readily
available in any laboratory.

6. 2) Silver-silver chloride electrode
Ag(s) | AgCl (sat’d), KCl (xM) | |
AgCl(s) + e = Ag(s) + Cl–
Eo = +0.244V
E = Eo – (0.05916/1) log [Cl–]
E (saturated KCl) = + 0.199V (25oC)

7. Reference electrodes
Silver-Silver Chloride
Reference Electrode
Calomel electrode

8. 3)
standard hydrogen electrode (SHE)
The most fundamental reference electrode in electrochemistry. "By
definition" its equilibrium potential is considered zero at any temperature,
because this electrode was chosen as an arbitrary zero point for
electrode potentials. A zero point is needed since the potential of a single
electrode cannot be measured, only the difference of two electrode
potentials is measurable.
All electrode potentials are expressed on this "hydrogen scale." It is
a hydrogen electrode with an electrolyte containing unit concentration
of H ions and saturated with H2 gas at unit atmosphere pressure.
This electrode can be somewhat inconvenient to use because of the need
to supply hydrogen gas. Therefore, other reference electrodes (e.g.,
calomel or silver/silver chloride) are often used instead, but the measured
electrode potentials can be converted to the "hydrogen scale." Also called
"normal hydrogen electrode."
Strictly speaking, one must use unit activity rather than concentration of
hydrogen ions and unit fugacity rather than unit pressure of hydrogen gas.
Pt | H2(g, 1.0 atm)|H+(aq, A= 1.0M)
½ H2(g, 1.0 atm) = H+(aq, A= 1.0M) + e
Eo = 0.000 V

9. Liquid-junction potential
A potential difference between two solns of different compositions separated by
a membrane type separator. The simplest example is the case of two solns
containing the same salt in different concns. The salt will diffuse from the
higher concn side to the lower concn side. However, the diffusion rate of
the cation and the anion of the salt will very seldom be exactly the same
(mobility).
Let us assume for this example that the cations move faster;
consequently, an excess positive charge will accumulate on the low
concentration side, while an excess negative charge will accumulate on the
high concn. side of the junction due to the slow moving anions. This sets up
a potential difference that will start an electromigration of the ions that will
increase the net flux of the anions and decrease the net flux of the cations.
In steady-sate conditions, the two ions will move at the same speed
and a potential difference will be created between the two solns. This "steadysate" potential difference seems constant, but this is misleading because it
slowly changes as the concn between the two soln. equalize. The diffusion
process will "eventually" result in equal concn of the salt in the two soln.
separated by the membrane, and the liquid-junction potential will vanish. For a
simple case, the value of the liquid junction potential can be calculated by
the so called "Henderson" equation.

10. Junction potential :
a small potential that exists at the interface between two electrolyte solutions
that differ in composition.
Development of the junction potential caused by unequal mobilities of ions.
Mobilties of ions in water at 25oC:
Na+ : 5.19 × 10 –8 m2/sV
K+ : 7.62 × 10 –8
Cl– : 7.91× 10 –8

11. Fig 21-4 Diagram of a silver/silver chloride electrode
showing the parts of the electrode that produce the
reference electrode potential Eref and the juction
potential Ej
Fig. 21-5. Schematic
representation of a liquid
junction showing the
source of the junction
potential, Ej. The length of
the arrows corresponds to
the relative mobilities of the
ions.

13. Liquid junction potential
Cells without liquid junction
Pt/H2(g), HCl/AgCl/Ag
Rare to have this type of cell
Cells with liquid junction
Glass frit
Salt bridge
Develop a potential by differential migration rates of the cation and anion.
Junction potential
HCl(0.1)/HCl(0.01)
Ej = 40 mV (H+ faster than Cl– )
KCl(0.1)/KCl(0.01)
Ej = –1.0 mV (K+ slower than Cl– )
Usually experimentally determine instrument response

14. Indicator electrodes
Metallic indicator electrode responds to analyte
activity.
Electrode of the first type
Direct equilibrium with analyte
Ag for Ag+, Au for Au3+, etc
Potential described by Nernst equation.
As [M] , E
Xn+(aq) + ne = X(s)
Eind = Eo – (0.5916/n) log (1/
[Xn+])
Note potential linearly related to log of the
concentration !
Remember - indicator BY DEFINITION cathode
measurement theoretically under zero-current (steady
A plot of Equation 21-3 for
an electrode of the first
kind.
state)
Electrode of the second type
Indirect equilibrium with analyte
M/MX/X–
AgCl(s) + e = Ag(s) + Cl–
(aq)
Eind = Eo – 0.5916 log [Cl–]
Silver/Silver chloride for chloride
also Nernstian response
as [X–] , E
Inert Metallic electrode for Redox systems
A plot of Equation 21-4
for an electrode of the
second kind for Cl–.

15. Indicator electrodes
Indicator electrodes for potentiometric measurements are of two basic
types, namely, metallic and membrane.
1) Metallic indicator electrodes :
develop a potential that is determined by the equilibrium position of redox
half-reaction at the electrode surface.
First-order electrodes for cations :
A first order electrode is comprised of a metal immersed in a solution
of its ions, such as silver wire dipping into a silver nitrate solution.
Only a few metals such as silver, copper, mercury, lead , zinc,
bismuth, cadmium and tin exhibit reversible half- reactions with their
ions and are suitable for use as first order electrodes.
Other metals, including iron, nickel, cobalt, tungsten, and chromium,
develop nonreproducible potentials that are influenced by impurities
and crystal irregularities in the solid and by oxide coatings on their
surfaces. This nonreproducible behavior makes them unsatisfactory as
first-order electrodes.

16. Example of First-order electrode
Ag+ + e = Ag(s)
Eo = + 0.800V
E = 0.800 – (0.05916/1) log {1/[Ag+]}

17. Second-order electrodes for anions
A metal electrode can sometimes be indirectly responsive to
the concn of an anion that forms a precipitate or complex ion
with cations of the metal.
Ex. 1.
Silver electrode
The potential of a silver electrode will accurately reflect the
concentration of iodide ion in a solution that is saturated with silver iodide.
AgI(s) + e = Ag(s) + I–
Eo = – 0.151V
E = – 0.151 – (0.05916/1) log [I–]
= – 0.151 + (0.05916/1)pI
2.Mercury electrode :
for measuring the concentration of the EDTA anion Y4–. Mercury
electrode responds in the presence of a small concn of the stable EDTA
complex of mercury(II).
HgY2– + 2e = Hg(l) + Y4–
Eo = 0.21V
E = 0.21 – (0.05916/2) log ([Y4–] /[HgY2–])
K = 0.21 – (0.05916/2) log (1 /[HgY2–])
E = K – (0.05916/2) log [Y4–] = K +(0.05916 / 2) pY

18. Inert electrodes
Chemically inert conductors such as gold, platinum, or carbon that
do not participate, directly, in the redox process are called inert
electrodes.
The potential developed at an inert electrode depends on the
nature and concentration of the various redox reagents in the solution.
Ag(s) | AgCl[sat’d], KCl[xM] | | Fe2+,Fe3+) | Pt
Fe3++e = Fe2+ Eo = +0.770V
Ecell = Eindicator – Ereference
= {0.770 – (0.05916/1) log [Fe2+]/[Fe3+]} – {0.222 – (0.05916/1) log [Cl–]}

19. 2) Membrane indicator electrodes
The potential developed at this type of electrode results from an unequal
charge buildup at opposing surface of a special membrane. The charge
at each surface is governed by the position of an equilibrium involving
analyte ions, which, in turn, depends on the concentration of those ions in
the solution.
The electrodes are categorized according to the type of membrane they
employ :
glass,
polymer,
crystalline,
gas sensor.
The first practical
glass electrode.
(Haber and
Klemensiewcz, Z.
Phys. Chem, 1909,
65, 385.

20. Membrane indicator electrodes
• Glass membrane pH electrodes
The internal element consists of silver-silver chloride electrode immersed in a
pH 7 buffer saturated with silver chloride. The thin, ion-selective glass
membrane is fused to the bottom of a sturdy, nonresponsive glass tube so
that the entire membrane can be submerged during measurements. When
placed in a solution containing hydrogen ions, this electrode can be
represented by the half-cell :
Ag(s) | AgCl[sat’d], Cl–(inside), H+(inside) | glass membrane | H+(outside)
E = Eo – (0.05916/1) log [Cl–] + (0.05916/1) log ([H+(outside)]/[H+(inside)])
E = Q + (0.05916/1) log [H+(outside)]

21. How a pH meter works
When one metal is brought in contact with another, a voltage difference
occurs due to their differences in electron mobility.
When a metal is brought in contact with a solution of salts or acids, a
similar electric potential is caused, which has led to the invention of
batteries.
Similarly, an electric potential develops when one liquid is brought
in contact with another one, but a membrane is needed to keep such
liquids apart.
A pH meter measures essentially the electro-chemical potential
between a known liquid inside the glass electrode (membrane) and an
unknown liquid outside. Because the thin glass bulb allows mainly the agile
and small hydrogen ions to interact with the glass, the glass electrode
measures the electro-chemical potential of hydrogen ions or the
potential of hydrogen.
To complete the electrical circuit, also a reference electrode is
needed. Note that the instrument does not measure a current but only an
electrical voltage, yet a small leakage of ions from the reference electrode is
needed, forming a conducting bridge to the glass electrode.

22. Most often used pH electrodes are glass
electrodes. Typical model is made of glass
tube ended with small glass bubble. Inside of
the electrode is usually filled with buffered
solution of chlorides in which silver wire
covered with silver chloride is immersed. pH
of internal solution varies - for example it can be
1.0 (0.1M HCl) or 7.0 (different buffers used by
different producers).
Active part of the electrode is the glass
bubble. While tube has strong and thick walls,
bubble is made to be as thin as possible.
Surface of the glass is protonated by both
internal and external solution till equilibrium
is achieved. Both sides of the glass are
charged by the adsorbed protons, this charge is
responsible for potential difference. This
potential in turn is described by the
Nernst equation and is directly proportional to
the pH difference between solutions on both
sides of the glass.

23. The majority of pH
electrodes available
commercially are
combination electrodes
that have both glass H+
ion sensitive electrode
and additional reference
electrode conveniently
placed in one housing.

24. In principle it should be possible to determine the H + ion activity or concn. of
a soln by measuring the potential of a Hydrogen electrode inserted in the
given soln. The EMF of a cell, free from liquid junction potential, consisting of
a Hydrogen electrode and a reference electrode, should be given by,
E = E ref – RT/F ln aH+
E = E ref + 2.303 RT/F pH
R= 8.314 J/mol/°K
∴pH = ( E- Eref )F/2.303 RT F= faraday constant ,96,485
T= Kelvin scale
So, by measuring the EMF of the Cell E obtained by combining the H
electrode with a reference electrode of known potential, Eref , the pH of the
soln. may be evaluated.
The electric potential at any point is defined as the work done in
bringing a unit charge from infinity to the particular point
Reduced state ⇋ Oxidised state + n Electron
M
= Mn+ + nE
E(+) = E0 – (RT/F) ln aMn+
Nernst Equation

25. Typical electrode system for
measuring pH. (a) Glass
electrode (indicator) and
saturated calomel
electrode (reference)
immersed in a solution of
unknown pH.
(b) Combination probe
consisting of both an
indicator glass electrode
and a silver/silver chloride
reference. A second
silver/silver chloride
electrode serves as the
internal reference for the
glass electrode.
The two electrodes are
arranged concentrically with the internal reference in the center and the
external reference outside. The reference makes contact with the analyte
solution through the glass frit or other suitable porous medium.
Combination probes are the most common configuration of glass electrode
and reference for measuring pH.

26. Glass electrode: E = E
0
G
+ 0.0591pH
The value of E0 G depends upon the composition of glass membrane and it
includes asymetry potential which is residual emf when identical solns and
electrodes are placed inside and outside the glass membrane.
It varies from day to day on the exposure of the surface to dryness or putting
in very strong acid or alkali.
The glass electrode usually operates in the pH range 1-10 without much
deviation in pH due to asymetry potential

27. To measure the hydrogen ion concentration of a solution the glass electrode must be
combined with a reference electrode, for which purpose the saturated calomel
electrode is most commonly used, thus giving the cell:
Owing to the high resistance of the glass membrane, a simple potentiometer
cannot be employed for measuring the cell e.m.f. and specialised instrumentation
must be used. The e.m.f. of the cell may be expressed by the equation:
or at a temperature of 250C by the expression:
In these equations K is a constant partly dependent upon the nature of the glass used
in the construction of the membrane, and partly upon the individual character of each
electrode; its value may Vary slightly with time. This variation of K with time is related
to the existence of an asymmetry potential in a glass electrode which is
determined by the differing responses of the inner and outer surfaces of the glass bulb
to changes in hydrogen ion activity; this may originate as a result of differing
conditions of strain in the two glass surfaces. Owing to the asymmetry potential, if a
glass electrode is inserted into a test solution which is in fact identical with the interna1
hydrochloric acid solution, then the electrode has a small potential which is found to
Vary with time. On account of the existence of this asymmetry potential of timedependent magnitude, a constant value cannot be assigned to K, and every glass
electrode must be standardised frequently by placing in a solution of known
hydrogen ion activity (a buffer soln)

28. the operation of a glass electrode is related to the situations existing at the
inner and outer surfaces of the glass membrane. Glass electrodes require
soaking in water for some hours before use and it is concluded that a
hydrated layer is formed on the glass surface, where an ion exchange
process can take place. If the glass contains sodium, the exchange process
can be represented by the equilibrium
The concn of the soln within the glass bulb is fixed, and hence on the inner
side of the bulb an equilibrium condition leading to a constant potential
is established. On the outside of the bulb, the potential developed will be
dependent upon the hydrogen ion concentration of the soln in which the
bulb is immersed.
Within the layer of 'dry' glass which exists between the inner and outer
hydrated layers, the conductivity is due to the interstitial migration of
sodium ions within the silicate lattice.

29. Composition of glass membranes
70% SiO2
30% CaO, BaO, Li2O, Na2O,
and/or Al2O3
Ion exchange process at glass
membrane-solution interface:
Gl– + H+ = H+Gl–
(a) Cross-sectional view of a silicate glass struture. In addition to the
three Si│O bonds shown, each silicon is bonded to an additional
oxygen atom, either above or below the plane of the paper. (b) Model
showing three-dimensional structure of amorphous silica with Na + ion
(large dark blue) and several H+ ions small dark blue incorporated.

30. Potentiometric titration:
Ag+/Ag
electrode
 NH 4 NO 3 + Agar  Calomel
Ag+NO3-/Ag ,
Half cell
Indicator Electrode
E Ag
RT
= E Ag −
ln[ Ag + ]
nF
0
E cell = EAg –E SCE =
E
0
Ag
Salt Bridge
to remove ljp
Satd KCl soln, Hg2Cl2,Hg
Reference Electrode
E calomel = -0.246 V
(SCE)
− ESCE
RT
−
ln[ Ag + ]
nF
E cell - E o Ag + E SCE
∴ log [ Ag + ] =
0.591
[Ag+]= 2.51 x 10-3
E cell = -0.4 V
E0 Ag = -0.8 V
ESCE = -0.246 V

31. To determine the Fe+2 /Fe+3 in soln
Fe+2 /Fe+3 , Pt  KCl +Agar 
Half cell
Indicator Electrode
Pt
Salt Bridge
to remove ljp
Calomel electrode
Satd KCl soln, Hg2Cl2,Hg
Reference Electrode
inert electrode, Pt has polarisation effect
E calomel = -0.246 V
E cell = EFe+3/Fe+2 –E SCE
= E0
(SCE)
[ Fe ]
− ESCE − 0.0591log
Fe + 3 / Fe + 2
[ Fe ]
+3
+2
[ Fe
∴ log
[ Fe
[ Fe
[ Fe
- E cell + E o
]=
]
] = 6.3 x10
]
+3
Fe
+2
+3
+2
+ 3
/Fe
0.591
−3
+ 2
+E
SCE
= 42.2
E cell
= - 0.4 V
E0 (Fe+3/Fe+2) = 0.77V
ESCE
= -0.246 V

32. Basics
It is therefore usually considered preferable to employ analytical (or
derivative) methods of locating the end point, these consist in plotting
the first derivative curve (ΔE/ΔV against V), or the second derivative
curve (Δ2E/ΔV2 against V). The first derivative curve gives a
maximum at the point of inflexion of the titration curve, i.e. at the
end point, whilst the second derivative curve (Δ2E/ΔV2) is zero at the
point where the slope of the ΔE/ΔV curve is a maximum.

33. CONDUCTOMETRY

34. Ohm's Law States that the current I (amperes) flowing in a
conductor is directly proportional to the applied electromotive force
E (volts) and inversely proportional to the resistance R (ohms) of
the conductor
The reciprocal of the resistance is termed the conductance (G): this
is measured in reciprocal ohms (or Ω –), for which the name Siemens
(S) is used. The resistance of a sample of homogeneous material,
length l, and cross-section area a, is given by:
where ρ is a characteristic property of the material termed the
resistivity (formerly called specific resistance). In SI units, l and a will be
measured respectively in metres and square metres, so that ρ refers to a
metre cube of the material, and
The reciprocal of resistivity is the conductivity, κ (formerly specific
conductance), which in SI units is the conductance of a one metre cube
of substance and has the units Ω – m - , but if ρ is measured in Ωcm, then
κ will be measured in Ω - cm - '.

35. The conductance of an electrolytic solution at any temperature
depends only on the ions present, and their concentration.
When a solution of an electrolyte is diluted, the conductance will
decrease, since fewer ions are present per millilitre of solution to
carry the current. If all the solution be placed between two
electrodes 1 cm apart and large enough to contain the whole of
the solution, the conductance will increase as the solution is
diluted. This is due largely to a decrease in inter-ionic effects
for strong electrolytes and to an increase in the degree of
dissociation for weak electrolytes.
The molar conductivity (Λ) of an electrolyte is defined as the
conductivity due to one mole and is given by:

36. For strong electrolytes the molar conductivity
increases as the dilution is increased, but it appears to
approach a limiting value known as the molar conductivity at
infinite dilution(Λ∞). The quantity Λ∞ can be determined by
graphical extrapolation for dilute solutions of strong
electrolytes.
For weak electrolytes the extrapolation method cannot
be used for the determination of Λ∞ but it may be calculated
from the molar conductivities at infinite dilution of the respective
ions, use being made of the 'Law of Independent Migration of
Ions‘.
At infinite dilution the ions are independent of each other,
and each contributes its part of the total conductivity, thus:

37. THE BASlS OF CONDUCTIMETRIC TlTRATlONS
principle underlying conductimetric titrations, i.e. The
substitution of ions of one conductivity by ions of another
conductivity.

38. Conductimetric Titrations
The principle of conductimetric titration is based on the fact that during
the titration, one of the ions is replaced by the other and invariably
these two ions differ in the ionic conductivity with the result that
conductivity of the solution varies during the course of titration. The
equivalence point may be located graphically by plotting the change in
conductance as a function of the volume of titrant added.
In order to reduce the influence of errors in the conductometric titration to a
minimum, the angle between the two branches of the titration curve
should be as small as possible . If the angle is very obtuse, a small error in
the conductance data can cause a large deviation. The following
approximate rules will be found useful.
The smaller the conductivity of the ion which replaces the reacting ion,
the more accurate will be the result. Thus it is preferable to titrate a silver
salt with lithium chloride rather than with HCl. Generally, cations should
be titrated with lithium salts and anions with acetates as these ions have
low conductivity.
The larger the conductivity of the anion of the reagent which reacts with
the cation to be determined, or vice versa, the more acute is the angle of
titration curve.

39. The titration of a slightly ionized salt does not give good results, since the
conductivity increases continuously from the commencement. Hence,
the salt present in the cell should be virtually completely dissociated; for a
similar reason; the added reagent should also be as strong electrolyte.
The main advantages to the conductimetric titration are its
applicability to very dilute, and coloured solutions and to system that
involve relative incomplete reactions.
For example, which neither a potentiometric, nor indicator method
can be used for the neutralization titration of phenol (Ka = 10–10) a
conductimetric endpoint can be successfully applied.
The electrical conductance of a solution is a measure of its currents
carrying capacity and therefore determined by the total ionic strength. It is
a non-specific property and for this reason direct conductance
measurement are of little use unless the solution contains only the
electrolyte to be determined or the concentrations of other ionic species
in the solution are known.
 Conductimetric titrations, in which the species in the solution are
converted to non-ionic form by neutralization, precipitation, etc. are of
more value.

40. Consider how the conductance of a solution of a strong
electrolyte A+ B- will change upon the addition of a reagent C+D-,
assuming that the cation A+ (which is the ion to be determined)
reacts with the ion D- of the reagent.
If the product of the reaction AD is relatively insoluble or only
slightly ionised,the reaction may be written:
Thus in the reaction between A+ ions and D- ions, the A+ ions are
replaced by C+ ions during the titration. As the titration proceeds the
conductance increases or decreases, depending upon whether the
conductivity of the C + ions is greater or less than that of the A+ ions.
During the progress of neutralisation, precipitation, etc., changes in
conductance may, in general, be expected, and these may therefore be
employed in determining the end points as well as the progress of the
reactions. The conductance is measured after each addition of a small
volume of the reagent, and the points thus obtained are plotted to
give a graph which ideally consists of two straight lines intersecting
at the equivalence point.

41. Ostwald derived a relationship between the molar
conductivity and limiting molar conductivity.
The molar conductivity of weak electrolyte
can be expressed as the product of degree of
dissociation of the electrolyte and its
limiting molar conductivity.
This relationship is known as Ostwald relation. Substituting this in Eq. 6.3 gives.
Rearrange Eq. (6.10) gives
Thus, we can use this method for the determination Ka. of weak acids and bases. To
further understand this, let us consider the dissociation of nitric acid in methanol over a
wide range of concentration (see Table 6.1).
In methanol nitric acid acts as a weak electrolyte and therefore, we can use Eq. (6.12) to
determine the dissociation constant Ka

42. Strong Acid with a Strong Base, e.g. HCl with NaOH:
 Before NaOH is added, the conductance is high due to the presence of
highly mobile hydrogen ions. When the base is added, the conductance
falls due to the replacement of hydrogen ions by the added cation as H + ions
react with OH− ions to form undissociated water.
This decrease in the conductance continues till the equivalence point.
At the equivalence point, the solution contains only NaCl.
After the equivalence point, the conductance increases due to the large
conductivity of OH- ions.
The conductance first
falls, due to the replacement
of the H+ (Λ∞ 350, Table 13.1)
by the added cation (Λ∞ 40-80)
and then, after the
equivalence point has been
reached, rapidly rises with
further additions of strong
alkali due to the large Λ∞ value
of the hydroxyl ion (198).
Conductimetric titration of a strong acid (HCl) vs. a strong base (NaOH)

43. Strong Acid with a Weak Base, e.g. H2SO4 with dilute NH3 (Kb=10-5 )
Initially the conductance is high and then it decreases due to the
replacement of H+. But after the endpoint has been reached
the graph becomes almost horizontal, since the excess aqueous
ammonia is not appreciably ionised in the presence of
ammonium sulphate.
The first branch of the
graph reflects the
disappearance of the
H ions during the
neutralisation,
Conductimetric titration of a strong acid (H2SO4) vs. a weak base (NH4OH)

44. Weak Acid with a Strong Base, e.g. acetic acid (Ka= 1.8 x10-5 )with NaOH:
Initially the conductance is low due to the feeble ionization
of acetic acid. On the addition of base, there is decrease in conductance not
only due to the replacement of H+ by Na+ but also suppresses the
dissociation of acetic acid due to common ion acetate.
But very soon, the conductance increases on adding NaOH as NaOH
neutralizes the un-dissociated CH3COOH to CH3COONa which is the strong
electrolyte. This increase in conductance continues raise up to the
equivalence point. The graph near the equivalence point is curved due the
hydrolysis of salt CH3COONa. Beyond the equivalence point, conductance
increases more rapidly with the addition of NaOH due to the highly conducting
OH− ions .
As the titration proceeds, a
somewhat indefinite break will
occur at the end point, and the
graph will become linear after all the
acid has been neutralised. Some
curves for acetic acid-sodium
hydroxide titrations are shown in
Fig.clearly it is not possible to fix
an accurate end point

45. Weak acids with weak bases.
The titration of a weak acid and a weak base can be readily carried
out, and frequently it is preferable to employ this procedure
rather than use a strong base.
Curve (c) is the titration curve of 0.003 M acetic acid with
0.0973 M aq. NH3 solution. The neutralisation curve up to the
equivalence point is similar to that obtained with NaOH solution, since
both Na and ammonium acetates are strong electrolytes;
after the equivalence point an
excess of aq. NH3 solution has
little effect upon the conductance,
as its dissociation is depressed
by the ammonium Salt present
in the solution.
The advantages over the use of
strong alkali are that the end
point is easier to detect, and in
dilute solution the influence of
CO2 may be neglected.

46. Mixture of a strong acid and a weak acid with a strong base
Upon adding a strong base to a mixture of a strong acid and a weak acid (e.g. HCl
and acetic acids), the conductance falls until the strong acid is neutralised, then
rises as the weak acid is converted into its salt, and finally rises more steeply as
excess alkali is introduced. Such a titration curve is shown as S in Fig.(d).
The three branches of the curve will be straight lines except in so far as:
(a) increasing dissociation of the weak acid results in a rounding-off at
the first end-point, and
(b) hydrolysis of the Salt of the weak acid causes a rounding-off at the
second end point.
Usually, extrapolation of the straight portions of the three branches leads
to definite location of the end points. Here also titration with a weak base,
such as aq NH3 soln, is frequently preferable to strong alkali for reasons
already mentioned in discussing weak acids:
curve W in Fig. (d) is obtained by substituting
aqueous ammonia solution for the strong alkali.
The procedure may be applied to the
determination of mineral acid in vinegar or
other weak organic acids (K < 4 and can
be used to analyse 'aspirin' tablets.

47. Displacement (or Replacement) Titrations:
When a salt of a weak acid is titrated with a strong acid, the anion of
the weak acid is replaced by that of the strong acid and weak acid
itself is liberated in the undissociated form. Similarly, in the addition
of a strong base to the salt of a weak base, the cation of the weak
base is replaced by that of the stronger one and the weak base itself is
generated in the undissociated form. If for example, 1M-HCl is added
to 0.1 M soln of sodium acetate, the curve shown in Fig. is obtained,
the acetate ion is replaced by the chloride ion after the endpoint.
The initial increase in conductivity is due to the fact that the
conductivity of the Cl- is slightly greater than that of acetate ion. Until
the replacement is nearly complete, the solution contains enough
sodium acetate to suppress the ionization of the liberated acetic
acid, so resulting a negligible increase in the conductivity of the
solution.
However, near the equivalent point, the acetic acid is
sufficiently ionized to affect the conductivity and a rounded portion of
the curve is obtained. Beyond the equivalence point, when excess of
HCl is present (ionization of acetic acid is very much suppressed)
therefore, the conductivity arises rapidly.

48. Care must be taken that to titrate
a 0.1 M-salt of a weak acid, the
dissociation constant should not
be more than 5×10–4, for a 0.01 M
-salt solution, Ka < 5 ×10–5 and for
a 0.001 M-salt solution, Ka < 5
×10–6, i.e., the ionization constant
of the displace acid or base
divided by the original
concentration of the salt must not
exceed above 5 ×10-3.
Fig. 6.6. Also includes the titration
of 0.01M- ammonium chloride
solution versus 0.1 M - sodium
hydroxide solution. The decrease
in conductivity during the
displacement is caused by the
displacement of NH4 ion of grater
conductivity by sodium ion of
smaller conductivity

49. Precipitation Titration and Complex Formation Titration:
A reaction may be made the basis of a conductimetric precipitation titration
provided the reaction product is sparingly soluble or is a stable complex . The
solubility of the precipitate (or the dissociation of the complex) should be less
than 5%. The addition of ethanol is sometimes recommended to reduce the
solubility in the precipitations. An experimental curve is given in Fig. 6.8
(ammonium sulphate in aqueous-ethanol solution with barium acetate).
If the solubility of the precipitate
were negligibly small, the
Precipitation titration. Conductimetric
titration of (NH4)2 SO4 vs. barium
conductance at the e.p should
acetate
be given by AB and not the
observed AC.
The addition of excess of the
reagent depresses the solubility
of the precipitate and, if the
solubility is not too large,
the position of the point B can
be determined by continuing the
Straight portion of the two arms
of the curve until they intersect.