1. The document discusses acid-base theories including Arrhenius, Brønsted-Lowry, and Lewis theories. It also defines terms related to acid-base titrations such as titration, standard solution, titrant, equivalence point, and indicator. 2. Key concepts around acid-base equilibria are explained, including the autoionization of water and the dissociation constants for weak acids and bases (Ka and Kb). The Henderson-Hasselbalch equation relating pH and the concentrations of conjugate acid-base pairs in a buffer solution is also derived. 3. Methods for determining the endpoint in an acid-base titration are described, including the use of acid-

1. 1
Mrs. Prajakta B. Kothawade
Assistant Professor,
PES Modern College of Pharmacy, for ladies, Moshi, Pune

2. Acid Base Titration
2
ACID BASE THEORIES
• Arrhenius Theory
 Acid is defined as a substance that when dissolved in water
gives rise to hydrogen ions.
 Base is defined as the substance that when dissolved in water
ionises to give hydroxyl ions.

3. Acid Base Titration
3
ACID BASE THEORIES
• Lowry Bronsted Theory
 Acid is defined as a species that when can donate protons.
 Base is defined as a species that accept a proton.

4. Acid Base Titration
4
ACID BASE THEORIES
• Lewis Theory
 Acid is defined as a species that accept an electron pair.
 Base is defined as a species that donate an pair of elctron.

5. Definitions
5
 Titration
It consists of determination of volume of solution of accurately
known concentration required to react completely with the
solution of substance to be determined.
 Standard solution
The solution of accurately known concentration
 Titrate
The substance being titrated

6. 6
 Titrant
The solution of known concentration
 Equivalance point
The point at which the reaction between titrant and titrate is
complete.
 End point
The point at which the indicator shows the colour change.
Definitions

7. 7
•Indicator
•It is an auxillary substance which shows clear visual change
after the reaction between titrate and titrant is practically
complete.
•Solution concentration
•Normality
•Molarity
•Molality
•Formal
•Percent
Definitions

8. 8
"The rate of a chemical reactian'is proportional to the active
Masses of the reacting substances.
" In dilute solutions where conditions approach the ideal state,
'active mass' may be represented by the concentration of the
reacting substances i. e. gm-molecules or gm-ions' per liter. The
constant of proportionality is known as 'velocity constant'.
LAW OF MASS ACTION

9. Acid Base Titration
9
Now, let us consider a homogeneous, reversible reaction.
A+B~C+D
According to law of mass action,
Vf = Kl [A] [B]
and Vb = K2[C] [D]
where,
Vf = Velocity of the forward reaction
Vb = Velocity of the backward reaction
[A], [B], [C], [D] = Molar concentration of A, B,C and
D respectively, &.K1and K2are constants.

10. Acid Base Titration
10
At equilibrium
Vf =Vb
i. e. Kl [A] [B] =K2[C] [D]
K1 = [C] [D]
K2 [A] [B]
Since K1 and K2 are constant ratio K1/K2 must also be a
constant

11. Acid Base Titration
11
In extension, the equilibrium constant for .the general
reversible reaction
aA + bB + cC ----------- pP + qQ + rR............
K = [P]p [Q]q [R]r
[A]a [B. ]b [C]c
where, a, b, c , and p, q, r are number of molecules of
reacting species

12. Acid Base Titration
12
ACID BASE EQUILlBRIUM
Consider a reversible chemical reaction
aA + bB.------- cC + dD
From the law of mass action
K =[C]C [D]d
[A]a [B]b
Now, let us consider the dissociation of weak acid in
aqueous
solution.
HA------ H+ +A
where,
H+ represents the hydrated proton.

13. Acid Base Titration
13
The dissociation constant for weak acid. can be denoted
as Ka
and its value can be represented as
Ka = [H+] [A-]
[HA]

14. Acid Base Titration
14
In this system water acts as a second conjugate acid-base
pair
Now consider the dissociation of weak base
B + H2O ------- BH+ + OH
and
the value of dissociation constant (Kb)is
Kb = [BH+] [OH-]
[B]
In this equation the [H2O] is taken to be constant and it is
absorbed in value of Kb.

15. Acid Base Titration
15
Water itself is capable of dissociation
H2O~H+ + OHThe
equilibrium constant
K = [H+] [OH-]
[H2O]
The [H2O] is usually considered to be constant.
Kw = [H+] [OH-]
where, Kw = Ion product of water

16. Acid Base Titration
16
Kw is important in order to derive the relationship between
Ka and Kb.
Ka. = [H+] [A-]
[HA]
Kb= [HA] [OH]
[A-]
Multiplying these two equations
Ka . Kb = [H+] [OH-]
i. e. Kw = Ka.Kb

17. Acid Base Titration
17
pH is the negative logarithm of hydrogen ion
concentration.
PH = - log [H+] = log 1/ [H+]
The ion product of water is
Kw =[H+] [OH-]
i. e. pKw = pH + pOH
Thus, if we know either [H+] or [OH-]; we can calculate the
unknown concentration.

18. Acid Base Titration
STRONG ACID OR BASE
A strong acid is completely dissociated into its component
ions in dilute aqueous solution. In such cases, therefore,
the concept of dissociation constant is not applicable.
e. g. Strong acids: HCl, HClO4, HN03
Strong bases: NaOH, KOH

19. Acid Base Titration
Let, 'c' be the total analytical concentration of a strong acid
[HX], where 'c' is expressed in terms of molarity [M] or
normality [N).
The solution as a whole is electrically neutral
i. e. , [H+] = [OH-] + [X-].
In this equation, the source of OH- is dissociation of water. If
total acid concentration is greater than about 10-6 M, [OH-]
will be very less than [X-] and hence
[H+] = [X-] = C
i. e. The hydrogen ion concentration is equal to the total
strong acid concentration.

20. Acid Base Titration
Similarly for a strong base MOH, the electroneutrality
equation is
[H+] + [M+] =. [OH-]
which can be written in simplified form
C = [OH-]
where, C = analytical concentration of the strong base

21. Acid Base Titration
Weak Acid
A weak acid is one which is incompletely dissociated. The
extent to which it dissociates is characterised by Ka.
Consider a weak acid HA in solution at total solute
concentration C.
The reaction is HA ------------ [H+] +A
The dissociation constant
Ka =[H+] [A-] (1)
[HA]
By electroneutrality rule [H+] = [OH-] + [A-].
Material balance C = [HA] + [A-]

22. Acid Base Titration
We can neglect the dissociation of water giving
[H+] = [A-]
Combining this with above 2 equations
[HA] = C - [H+]
Substituting this in equation (1)
Ka = [H+]2
C- [H+]
Ka = [H+]2
C
... Since [H+] is negligible as compared to C
:. [H+] = Ka.C

23. Acid Base Titration
Weak Base
The reaction may be written as
B + H2O---------- BH+ + OH
And the dissociation constant is
Kb= [BH+] [OH-]
[B]
By using reasoning used in previous case we can write
Kb = [OH-]2
C -[ OH-]
Since [OH-] is less as compared to C
Kb = [OH-]2
C
[OH-] = √Kb. C,

24. Acid Base Titration
24
Buffer Solutions
Buffer solution is a solution of substance or a mixture of
substances which helps in maintaining and establishing
specific pH.
Consider an aqueous solution which is meant to contain 'a'
moles/litre of a weak acid HA and 'b' moles/litre of its
conjugate base A-. But in practice the actual
concentrations of these species will be slightly differing,
because of dissociation equilibrium.

25. Acid Base Titration
25
HA = H+ + A
The concentrations can be considered as
[HA] = a - [H+]
[A-] = b + [H+]
The dissociation constant is
Ka= [H+] [A-]
[HA]
pKa = pH - log [A]
[HA]

26. Acid Base Titration
26
Substituting above values in this equation
pKa = pH - log a - [H+]
b + [H+]
Practically we can consider that [H+] is much smaller as
compared to either 'a' or 'b'.
pKa = pH -log b
a
This is Henderson-Hasselbalch equation which relates pH of
solution containing comparable and appreciable
concentration of a conjugate acid base pair to ratio of their
concentrations.

27. Acid Base Titration
27
Substituting above values in this equation
pKa = pH - log a - [H+]
b + [H+]
Practically we can consider that [H+] is much smaller as
compared to either 'a' or 'b'.
pKa = pH -log b
a
This is Henderson-Hasselbalch equation which relates pH of
solution containing comparable and appreciable
concentration of a conjugate acid base pair to ratio of their
concentrations.
The equation can also be written as
pH = pKa + log [conjugate base]
[conjugate acid]

28. Acid Base Titration
28
Buffer Solutions
Buffer solution is a solution of substance or a mixture of
substances which helps in maintaining and establishing
specific pH.
Consider an aqueous solution which is meant to contain 'a'
moles/litre of a weak acid HA and 'b' moles/litre of its
conjugate base A-. But in practice the actual
concentrations of these species will be slightly differing,
because of dissociation equilibrium.

29. Acid Base Titration
29
HA = H+ + A
The concentrations can be considered as
[HA] = a - [H+]
[A-] = b + [H+]
The dissociation constant is
Ka= [H+] [A-]
[HA]
pKa = pH - log [A]
[HA]

30. Acid Base Titration
30
Substituting above values in this equation
pKa = pH - log a - [H+]
b + [H+]
Practically we can consider that [H+] is much smaller as
compared to either 'a' or 'b'.
pKa = pH -log b
a
This is Henderson-Hasselbalch equation which relates pH of
solution containing comparable and appreciable
concentration of a conjugate acid base pair to ratio of their
concentrations.

31. Acid Base Titration
31
Substituting above values in this equation
pKa = pH - log a - [H+]
b + [H+]
Practically we can consider that [H+] is much smaller as
compared to either 'a' or 'b'.
pKa = pH -log b
a
This is Henderson-Hasselbalch equation which relates pH of
solution containing comparable and appreciable
concentration of a conjugate acid base pair to ratio of their
concentrations.
The equation can also be written as
pH = pKa + log [conjugate base]
[conjugate acid]

32. Acid Base Titration
32
END POINT DETECTION
The object of carrying out an acid-base titration is to determine
the equivalent quantity of the other substance required for
neutralization.
The point at which complete neutralization is achieved is called
as the 'end-point or the 'equivalence point'.
If both the acid and the base are strong electrolytes the
resulting solution will be neutral having pH =7. But if either of
the two is a weak electrolyte the resulting salt formed will
hydrolyse to some extent and solution will possess some
acidic or alkaline properties.
Indicator: It is a substance which exhibits colour change at a
particular stage of a chemical reaction.

33. Acid Base Titration
33
Neutralisation indicator
These are the substances which exhibit different colours at
various values of pH. .Indicators are weak acids and weak
bases which have different colours in their conjugate base
and acid forms.
Most indicators are used in dilute solution form.
For an acid-base titration, we select an indicator which will
show a distinct colour change at pH close to the equivalence
point.
Two mechanisms
1) Ostwald Theory
2) Resonance Theory

34. Acid Base Titration
34
Resonance Theory
All the acid-base indicators which are commonly used are
organic compounds. The difference in colour of same
compound in acid
and base medium is apparently due to difference in structure of
two
forms. Colour shown by the compound is associated with the
capability of the compound to absorb visible light and this
capability is related to the electronic structure. Change in the
electronic features will result in absorption of different colour
components of light with a resultant colour change.

35. Acid Base Titration
Ostwald Theory
The first theory to explain the behavior of indicators was put
forth by W. Ostwald. According to his theory the
undissociated indicator acid [HIn] or a base [lnOH] has a
colour different than its ion.
For an acid indicator equivalence can be written as :
HIn------ [H+] + In-
In an acid solution, there is depression of ionization of indicator
due to common ion effect. Hence initially concentration of
HIn is greater than that of In- and the colour exhibited will be
that of the unionised form. As the titration proceeds the alkali
medium will promote removal of H+ and there is gradual
increase in concentration of ionised form In - and the
solution acquires colour of ionised form.

36. Acid Base Titration
By applying law of mass action.
KIna = [H+] [In-]
[HIn]
log [H+] = log KIna + log [HIn-]
[In-]
Taking negative log
- log [H+] = - log KIna - log [HIn]
[In-]
= - log KIna + log [ln-]
[HIn]

37. Acid Base Titration
pH = pKlna + log [ln-]
[HIn]
KIna is dissociation constant of indicator. The colour of
indicator depends upon the ratio of concentration of ionised
and unionised form and hence directly proportional to pH.
For base indicator
pH = pKw - pKIb - log [lnOH]
[In+]

38. Acid Base Titration
MIXED INDICATORS
In some cases the pH range is very narrow and the colour
change over this range must be very sharp. This is not easily
possible with ordinary acid-base indicators. The result may be
achieved by the
use of the suitable mixture of indicators. These are generally
selected
so that their pKIn values are close together and overlapping
colours

39. Acid Base Titration
MIXED INDICATORS
Example,
A mixture of equal parts of neutral red (0.1 % solution in
alcohol) and
methylene blue (0.1 %solution in alcohol) gives a sharp colour
change form violet-blue to green in passing from acid to
alkaline solution at pH 7.

40. Acid Base Titration
UNIVERSAL OR MULTIPLE RANGE NDICATORS
By suitably mixing certain indicators the colour change may be
made to extend ove.r a considerable portion of the pH range.
Such mixtures are usually called "Universal indicators". They
are not suitable for quantitative titrations but may be employed
for the determination of the approximate pH of a solution by
colorimetric method.

41. Acid Base Titration
UNIVERSAL OR MULTIPLE RANGE NDICATORS
.
Example,
Dissolve 0.1 g of phenolphthalein, 0.2 g of,methyl red, 0.3 g
methyl yellow, 0.4 g of bromothymol blue and 0.5 g thymol blue
in 500 ml of absolute alcohol and add sufficient sodium
hydroxide solution until the colour is yellow. The colour
changes are: pH.2 --red, pH.4 --orange, pH.6 - yellow, pH.8--
green, pH.10--Blue.
[Sarabhai. Universal indicator pH paper]

42. Acid Base Titration
Lecture 9
42
NEUTRALISATION CURVE
If we study. the changes in the hydrogen ion concentration
during the course of a titration we get aclear idea about the
mechanism of neutralisation process. The pH value of greatest
importance is the one near the equivalence point as it gives us
help in , selecting an indicator which will give the smallest
titration error.
The curve obtained by plotting pH as ordinate against the
percentage of acid neutralised [or the number of ml of alkali
(titrant) added] as abscissa during titration, is known as
neutralisation or more generally titration curve.

43. Acid Base Titration
43
STRONG ACID-STRONG BASE TITRATION:

44. Acid Base Titration
44
WEAK ACID-WEAK BASE TITRATION:

45. Acid Base Titration
45
WEAK ACID-WEAK BASE TITRATION:

46. Acid Base Titration
46
WEAK ACID-WEAK BASE TITRATION:

47. Acid Base Titration
Lecture 10
47
STRONG ACID-WEAK BASE TITRATION:

48. Acid Base Titration
48
STRONG ACID-WEAK BASE TITRATION:

49. Acid Base Titration
49
STRONG ACID-WEAK BASE TITRATION:

50. Acid Base Titration
50
WEAK ACID-STRONG BASE TITRATION:

51. Acid Base Titration
51
WEAK ACID-STRONG BASE TITRATION:

52. Acid Base Titration
52
WEAK ACID-STRONG BASE TITRATION:
:
Before titration begins:
At this point sample solution contains only weak acid and the
pH is calculated
[H+] = √Ka.C
During titration
After some strong base has been added, its reaction with
weak acid will producean equivalent amount of conjugate
weak base. The solution now is a mixture of weak acid and its
conjugate base, it is a buffer solution and therefore apply
handerson- hasselbalch equation
pH = pKa + log [conjugate base]
[conjugate acid]

53. Acid Base Titration
53
WEAK ACID-STRONG BASE TITRATION:
At the end point:
Now an amount of strong base has been added that is exactly
equivalent to amount of weak acid initially. present. The solution
now contains only the conjugate base of weak acid and its pH can be
calculated.
Kb=[OH-]2
C - [OH-]
[OH-] = √KbC

54. Acid Base Titration
54
WEAK ACID-STRONG BASE TITRATION:
After the end point:
An excess of titrant has now been added, and pH is determined
essentially by this excess, the appropriate dilution within the solution
being taken into account. It is true that the conjugate weak base
produced also contributes to the solution pOH, but its contribution is
small relative to effect of excess titrant. After equivalence point, the
solution contains excess of OH which will represent the hydrolysis
of salt.

55. Acid Base Titration
POLY ACID TITRATION:
A polybasic acid may be considered as a mixture of acids i. e. it
furnishes more than one proton on dissociation.. Carbonic acid
gives two protons, phosphoric acid gives three protons and hence
are called as polybasic or polyprotic acids. Each stage of
dissociation gives a separate monobasic acid.
H3A [A-] + [H+] Dissoci. const. Kl
H2A  [HA-] + [H+] Dissoci. const. K2

56. Acid Base Titration
POLY ACID TITRATION:
A polybasic acid may be considered as a mixture of acids i. e. it
furnishes more than one proton on dissociation.. Carbonic acid
gives two protons, phosphoric acid gives three protons and hence
are called as polybasic or polyprotic acids. Each stage of
dissociation gives a separate monobasic acid.
H3A [A-] + [H+] Dissoci. const. Kl
H2A  [HA-] + [H+] Dissoci. const. K2

57. Standardization of 0.1 M Sodium
hydroxide
57
Standardization of NaOH,
• (A) Preparation of sodium hydroxide (1 M): Dissolve 42 g of sodium hydroxide in
sufficient carbon dioxide-free water to produce 1000 ml.
• (B) Standardisation: Weigh accurately about 5 g of potassium hydrogen phthalate
previously powdered and dried at 120°C for 2 hours, and dissolved in 75 ml carbon
dioxide free water. Add 0.1 ml phenolphthalein solution and titrate with the sodium
hydroxide solution until a permanent pink colour is produced.
• Factor: 1 ml of 1 M sodium hydroxide is equivalent to 0.2042 g of potassium
hydrogen phthalate.

58. Assay of Ammonium chloride
58
Assay of Ammonium chloride,
•(A)Weigh accurately about 0.1 gm of ammonium chloride. Dissolve in mixture of 20
ml water and 5 ml formaldehyde solution.
•(B) Add 0.1 ml phenolphthalein solution and titrate with the sodium hydroxide
solution until a permanent pink colour is produced.
Factor: 1 ml of 1 M sodium hydroxide is equivalent to 0.005349 g of ammonium
chloride
.

59. 59

60. References
 Vogel’s Text Book of Quantitative
Chemical Analysis, 6/Ed., Pearson
Education, page no:41-50 and 363-383.
 Practical Pharmaceutical Chemistry Part-I
by Beckett A H & Stanlake J B, 4/Ed., CBS
Publisher & Distributors, page no:137-157
and 165.
 Pharmaceutical Analysis Vol. I & K. R.
Mahadik, S.G. Wadodkar, H. N, I. More,
Nirali Prakashan
page no: 52-84.
60