This document provides an overview of acid-base titration and volumetric analysis. It defines key terms like titration, indicator, equivalence point, and standardization. It describes different types of titrations including direct, indirect, and back titration. Acid-base concepts are explained based on Arrhenius, Bronsted-Lowry, and Lewis theories. The document also discusses the ionic product of water, common ion effect, classification of indicators, and theories of indicators including Ostwald and chromophore theories.

1. PHARMACEUTICAL ANALYSIS -I
BY
DIPANKAR NATH.
Assistant Professor.
Department of Pharmaceutical Analysis.
HPI.
ACID BASE TITRATION

2. VOLUMETRIC ANALYSIS:
Volumetric analysis is a mode of quantitative analysis which is based on the determination of
the volume solution of known concentration(Std.)is required to react quantitatively with a
substance to be analyzed(Analyte/Titrate).
Conditions For volumetric analysis:
 Reactions must be simple and should take place quantitatively and there should not form any
side reaction.
 Reaction should be relatively fast.
 An indicator Should be available .
Important terms:
Titration: A process, Operation or method used to determine the concentration of a substance
in a solution by adding a standard reagent of known concentration until the reaction is
complete. Usually the completion of the reaction is observed by colour change(Indicator
method) or by electrical measurement(Instrumental method) and then calculate the unknown
concentration.

3. Indicator: An auxiliary chemical compound that changes color and structure when exposed
to certain conditions and therefore uses to detect the end point of the titration. It can be
used to monitor acidity, alkalinity or the progress of the reaction.
Acidimetry: Volumetric analysis using standard solutions of acid to measure the amount of
alkali present.
Alkalimetry: Volumetric analysis using standard solutions of alkali to measure the amount of
acid present.
Equivalence Point: It is the point or concentration where the addition of titrant is
stoichiometrically equal to the amount of moles of analyte in the given sample. In acid base
titration equivalents refers to the number of moles of H+ and OH-.It is also called as
Theoretical end point or Stoichiometric point.
Standardization: Standardization is the process of determining the exact concentration
(molarity/normality) of a solution. It is done by titration of a solution of unknown strength
against solution of either primary /secondary standard.E.g: A solution of HCl may be
standardized by a standardized solution of NaOH or by Na2CO3.

4. Types of Titration:
(A)Direct Titration: In this method the solution of the substance to be determined
quantitatively is directly titrated with a suitable titrant by using an appropriate indicator or
suitable instrument to locate equivalence point.
E.g: Titration of Strong acids such as HNO3,H2SO4,HCl with strong alkali such as NaOH,KOH
or EDTA.
(B)Indirect Titration: When substance is not directly titrable by employing a chemical
reaction, it can be precipitated or removed from the reaction and the reaction product thus
formed has to be titrated.
(C)Back Titration: This titration is followed when direct titration is not possible. This method
is used for those substances which are not water soluble, or which are weak acids or bases
by nature.e.g: ZnO, alkaloids etc.
In this method the substance to be estimated quantitavely is dissolved in a known sufficient
volume of standard solution of acid or alkali. The excess of acid or alkali remaining in the
solution is back titrated by using a suitable indicator. A blank determination is done after
that.

5. Classification of Titrimetric methods:
Depending upon the nature of the chemical reaction, they are classified as follows:
1.Acid-base titrations: When neutralization reaction is involved then the titration is termed as
acid base or neutralization titration.
NaOH+HCl→NaCl+H2O
2.Complexometric titration: This depend upon the combination of ions other than H+/OH- to
form a soluble slightly dissociated ion or compound.
2CN-+Ag+↔ [Ag(CN)2]-
EDTA largely as disodium salt of EDTA is very important reagent for complexometric titration.
3.Precipitaion titration: These depend upon the combination of ions to form a simple
precipitate.
NaCl +AgNO3→AgCl↓ +NaNO3
4.Nonaqueous titration: When the reaction occurs in the Nonaqueous solvent i.e. Organic
solvent .It is generally used for the analysis of very weak acids or bases. Commonly used
solvents such as Glacial acetic acid, Acetonitrile

6. 5.Redox titration(Oxidation reduction ):This involves the change in oxidation number or
transfer of electrons among the reacting substances.e.g KMnO4 ,K2Cr2O7 ,Iodine ,KI (Oxidizing
agent) Na2s2O3,AS2O3,Hg2(NO3)2(Reducing agent).
6.Diazotisation titration: These titrations involve the diazotization reactions. Generally
primary amines undergoes these type of titrations.
ACID – BASE CONCEPTS
ARRHENIUS CONCEPT: (1887)
According to this Concept, Acids are substances that dissociates in aqueous medium to
give H+ ions and bases are substances that produces OH- ions.
Ionisation of an acid : HCl ----------> H+ + Cl-
Ionisation of a base: NaOH --------> Na+ + OH-
Drawbacks of Arrhenius concept:
 Applicable only to aqueous solutions
 It fails to explain to explain the acidic nature of the substance like CO2,SO2 Which does not
contain H+and for bases which does not contain OH- ion like Na2CO3,NH3
H2O
H2O

7. BRONSTED-LOWRY CONCEPT: (1923)
According to this theory, acid is a substance that is capable of donating a Proton and
bases are substances capable of accepting a Proton. Acids are proton donors and bases
are proton acceptors.
NH3 + H2O -----------> NH4
+ + OH-
Base Acid Conjugate acid Conjugate base
CH3COOH + H2O ---------------> CH3COO- + H3O+
acid base conjugate base conjugate acid
 A conjugate base is an ion or molecule left after an acid donates a proton.
 A conjugate acid is an ion formed after accepting a proton.
 The pairs of Substance which are formed from one another by the loss or gain of a proton
are called as Conjugate acid and base pair.
Limitations:
Substances like BF3, AlCl3 etc. do not have any hydrogen and hence cannot give a
proton but are known to behave as acids.

8. LEWIS CONCEPT:(1938)
According to this theory, An acid is a substance which accepts a pair of an electron and
Base is a substance which donates a pair of electron.
H+ + :NH3 -------------> NH4
+
Limitations:
Relative strength of acids and bases can not be explained by this concept.
Relative Strength of Acids and Bases:
 The strength of an acid depends upon its tendency to provide proton or more precisely
hydroxonium ion in solution.
 Strong acid refers to that the equilibrium lies very much to the right hand side and the
backward reaction takes place almost to a negligible extent.
HCl+H2O→H3O + Cl-
 Weak acid refers to that the equilibrium lies very much to the left hand side and
dissociation of acid to give H+/H3O+ is very small.
CH3COOH+H2O↔H3O+CH3COO-

9.  Strength of base will depend upon the availability of OH- ion.In case of strong base
equilibrium lies very much to the right hand side and the backward reaction takes place
almost to a negligible extent.
NaOH+H2O→ Na++ OH-
 In case of weak base equilibrium mainly lies to the left hand side.
NH3+H2O↔NH4+OH-
 Strength of acids and bases will depend on the degree of dissociation, higher the degree of
dissociation of acids or bases stronger stronger will be the acid or bases.
LAW OF MASS ACTION:
Guldberg and Waage in 1867 stated the law of mass action. According to the law “The
velocity or rate of a chemical reaction is proportional to the product of active masses of the
reacting substances”.

10. Let us consider a simple reversible reaction at constant temperature
A + B ⇌ C + D
Rate of conversion of A and B is proportional to their concentration
r1 ∝ [A] [B]
r1 = K1 [A] [B]
Similarly,
Rate of conversion of C and D is proportional to their concentration
r2 ∝ [C] [D]
r2 = K2 [C] [D]
where, r1 = rate of reaction in forward direction
r2 = rate of reaction in backward direction
K1 and K2 = rate constant or rate coefficient
[ ] = concentration (mol/l)

11. Since the reaction is in equilibrium, the rate of forward reaction becomes equal to the rate
of backward reaction.
i.e. r1 = r2
K1[A][B] = K2[C][D]
K1 = [C][D]
K2 [A][B]
K= [C][D]
[A][B]
Where, K=equilibrium constant at a given temperature
The equilibrium constant K varies with temperature and pressure. Hence, equilibrium is
dynamic in nature.
IONIC PRODUCT OF WATER:
Kohlrausch and Heydweiller found that the most highly purified water posses a small but
definite conductivity. Water must therefore be slightly ionized in accordance with the
equation:
H2O ⇌ H+ + OH-

12. Applying law of mass action,
Ka = [H+] [OH-]
[H2O]
Since, water is slightly ionised in dilute aqueous solution, the conc. of the undissociated
water may be considered as a constant K1
Ka = [H+] [OH-]
[K1]
Ka x K1 = [H+][OH-]
Kw = [H+][OH-]
Kw = 1 x 10-14 at 25 oC, Kw=Ionic Product of water.
COMMON ION EFFECT:
It is the degree of dissociation of an electrolyte suppressed by the addition of another
electrolyte containing a common ion or dissociation of comparatively weak electrolyte
suppressed by a strong electrolyte.

13. E.g.: Dissociation of NH4OH is diminished by the addition of a salt NH4Cl which furnishes the
common ion NH4
+
NH4OH ⇌ NH4
+ + OH-
According to the law of Mass Action,
K = [NH4
+][OH-] Here, K = Equilibrium constant
[NH4OH]
Now, a strong electrolyte NH4Cl dissociates as
NH4Cl -----------> NH4
+ + Cl-
Thus the value of conc. of [NH4
+]will increase and therefore conc. of OH- ion will decrease to
maintain the equilibrium. Finally ionisation of NH4OH is suppressed and hence release of OH-
ion decreases.
The common ion effect provides a valuable tool controlling the conc. of an ion furnished by
weak electrolyte.

14. INDICATORS:
These are the substances which changes colour at the end point and thereby indicate the
completion of the reaction.
These are complex organic compounds, may be acids or bases capable of existing in two
forms of different colour that are mutually convertible one into other at a given hydrogen ion
concentration. They are generally used for:
i) Determination of the end point of a neutralisation process
ii) Determination of the hydrogen ion concentration or pH
Classification of indicators:
i) Natural colouring matter:
These indicators do not give definite end point.
Litmus is a colouring matter obtained from various species of Rocella azolitmin. The
chief principle present is azolitmin. Litmus is used to know acidic or basic pH.
ii) Internal Indicator:
It changes colour at the end point when added to reacting medium.
They are further classified as-

15. (a) Acid – base indicator: These are indicators generally used in acid-base titration.
e.g. Phenolphthalein (acidic indicator),Methyl orange (basic indicator)
(b)Precipitation Indicator: These are the indicators generally used in precipitation titration.
e.g.: K2CrO4(Mohr’s method), FeNH4(SO4)2 (Volhard’s method), Adsorption indicator
(c)Redox Indicator: These are the indicator which shows one colour in the oxidized state
and another colour in the reduced state.
e.g.: O-phenanthrolene ferrous ion (Ferroin)
(d)Non-aqueous Indicator: These are generally used in Non-aqueous titration and prepared
in Non-aqueous solvent. E.g.: Crystal violet in GAA, Oracet blue in GAA etc.
(e)Metallochrome indicator: They gives one colour in presence of Metal and another color in
absence of metal ion(free),e.g.: Mordant black-II(Solochrome blackT/Erichrome blackT),
Murexide (Ammonium purpurate)etc.
iii)External Indicator:
Some indicators cannot be added to the reacting media due to
(a) If the indicator is dark in colour then sudden change in colour cannot be observed
clearly

16. (b) If it forms an insoluble precipitate with the solution to which it has been added
e.g.: Potassium ferricyanide is used as an external indicator in the titration of potassium
dichromate and ferrous sulphate in acidic medium
FeSO4 + 2K3[Fe(CN)6](Potassium ferricyanide) ---------------> Fe3[Fe(CN)6]2(Ferro-ferric cyanide)
+3K2SO4
Thus when a drop of indicator remains unchanged in colour, it indicates that there is no
FeSO4 left unreacted and all of it has reacted with K2Cr2O7, the end point.
iv)Self indicator:
In the titration of oxalic acid or FeSO4 with KMnO4 solution, there is no need to add any
external indicator. Here KMnO4 acts as a self indicator which changes its colour after the
completion of the reaction.
v)Mixed Indicator:
These indicators are generally used to get sharp end point. Here generally two indicators are
mixed together.
e.g.: In the titration of CH3COOH and NH3, sharp control of pH is necessary. A mixture of
equal parts of neutral red(0.1% solution in alcohol) and methylene blue(0.1% solution in

17. alcohol) gives sharp colour change from violet blue to green in passing from acidic to
alkaline solution at pH 7.
vi)Universal Indicator:
By mixing suitable indicators together changes in colour may be obtained over a
considerable portion of pH range, such mixtures are usually called universal indicators.
They are not suitable for quantitative titrations, but may be employed for determination of
the approximate pH of a solution by the colorimetric method.
e.g. One such indicator is prepared by dissolving 0.1g of phenolphthalein, 0.2g of methyl red,
0.3g of methyl yellow, 0.4g of bromothymol blue and 0.5g of thymol blue in 500ml of
absolute ethanol, and adding NaOH solution until the colour is yellow.
Colour changes are:
pH Colour
2 red
4 orange
6 yellow
8 green
10 blue

18. THEORY OF INDICATORS:
1. OSTWALD THEORY
2. QUINONOID/RESONANCE/CHROMOPHORE THEORY
3. PHYSICO-CHEMICAL THEORY
1)Ostwald’s Theory:
According to Ostwald’s theory, the colour change of any indicator is due to its ionisation.
The unionised form of indicator has different colour than its ionised form.
 An indicator is either a weak acid or base, so its ionisation is highly affected in acids and
bases. If an indicator is a weak acid, its ionisation would be very much low in acids due
to common H+ ions while it is fairly ionised in alkalis.
 In the same way, if the indicator is a weak base, its ionisation is large in acids and low in
alkalis due to common OH- ions.
Let’s take examples of two important indicators phenolphthalein which is a weak acid
and methyl orange which is a weak base

23. 2.Chromophore Theory/ Resonance Theory:
 Indicators are the mixtures of two tautomers in equilibrium. One tautomeric form
possesses the Quinonoid structure (highly coloured) whereas other possess simply
benzene ring or Benzenoid form (light coloured).
 Indicators are generally weak acids or bases and they show a change in colour from a
predominantly acidic medium to a predominantly alkaline medium.
Benzenoid form Quinonoid form
 Let us consider the case of phenolphthalein and Methyl orange.
 Phenolpthalein is a most popular diprotic acid indiactor.
 It has benzenoid form in acidic medium and thus, it is colourless while it has quinonoid
form in alkaline medium which has pink colour

24.  In the presence of alkali, the lactone ring of structure (I) opens up to give Triphenyl
carbinol structure (II). This undergoes loss of water to produce resonating structure (III)
which is pink in colour. If phenolphthalein is treated with excess of alcoholic alkali, the
pink colour vanishes due to the formation of structure (Iv).

25. Methyl Orange is a weak base and is yellow in colour in molecular form. Addition of proton
gives a cation which is pink/Red in colour.
3.Physico-chemical Theory:
 Colour changes occur due to the increase in concentration of certain ions. Increase in the
concentration causes the appearance of colour whereas decrease in the conc. of some
ions causes disappearance of colour or appearance of new colour

26. CHOICE OF INDICATOR:
Indicator pH range Colour range
Bromothymol blue 1.2 – 2.8 Red – Yellow
Thymol blue 6.0 – 7.6 Yellow – Blue
Congo red 3.1 – 4.4 Blue – Red
Methyl orange 2.8 – 4.4 Red – orange
Phenolphthalein 8.3 – 10.0 Colourless – pink
pH SCALE:

27. 0 - 2 Strongly acidic
2-4 Moderately acidic
4-7 Weakly acidic
7-10 Weakly basic
10-12 Moderately basic
12-14 Strongly basic
In 1909, Sorensen introduced a term for expressing the concentration of hydrogen
ion which gives an idea about the acid and basic character of aqueous solutions. The
term is called as pH which means power of hydrogen ions. The pH is defined as negative
logarithm of H+ ion concentration in moles per litre.
TITRATION CURVES:
A titration curve is a plot showing the change in pH of the solution in the conical flask
as the reagent is added from the burette.

28. Titration of a strong acid with a strong base: (HCl vs NaOH)
 Let’s assume our strong acid is HCl and
strong Base is NaOH.
HCl + NaOH ---------------> NaCl + H2O
 Point a: No NaOH added yet, so the pH of the
analyte(in conical flask) is low (it
predominantly contains H3O+ from
dissociation of HCl).
 Now we start adding base from the burette
 Point b: This is the pH recorded at a time
point just before complete neutralization
takes place.
 Point c: This is the equivalence point
(halfway up the steep curve). At this point,
moles of NaOH added = moles of HCl in the
analyte. At this point, H3O+ ions are
completely neutralized by OH- ions and
therefore the pH is neutral i.e. pH = 7
Point d: Addition of NaOH continues, pH starts
becoming basic because HCl has been completely
neutralized and now excess of OH- ions are present
in the solution

29. Titration of a weak acid with a strong base(CH3COOH Vs NaOH)
Let’s assume our weak acid is CH3COOH and
strong base is NaOH
CH3COOH + NaOH ---------------> CH3COONa + H2O
Point a: No NaOH added yet, so the pH of the
analyte is low. But acetic acid is a weak acid,
so the starting pH is higher than what we
noticed in case 1 where we had a strong acid
(HCl).
Now we start adding the base dropwise
Point b: This is the pH recorded at a time point
just before complete neutralization takes
place.
Point c: This is the equivalence point (halfway
up the steep curve). At this point, moles of
NaOH added = moles of CH3COOH. The H3O+
ions are completely neutralized by the OH-
ions.
CH3COOH + NaOH ---------------> CH3COONa + H2O
The salt CH3COONa so formed dissociates into
CH3COO- and Na+ ions. This CH3COO- reacts
with H2O to produce OH- ions thus increasing
the pH to ~ 9 at the equivalence point.
CH3COONa + H2O --------------------> CH3COO- + Na+
CH3COO- + H2O ------------------> CH3COOH + OH-
Point d: Beyond the equivalence point (when sodium
hydroxide is in excess) the curve is identical to HCl-
NaOH titration curve

30. Titration of a strong acid with a weak base(HCl Vs NH3)
Suppose our strong acid is HCl and weak
base NH3
NH3 + HCl --------------> NH4Cl
Point a: No NH3 added yet, so the pH of
the analyte is low(in conical flask) is low
(it predominantly contains H3O+ from
dissociation of HCl).
Point b: This is the pH recorded at a time
point just before complete neutralization
takes place.
Point c: This is the equivalence point
(halfway up the steep curve). At this
point, moles of NH3 added = moles of HCl
in the analyte. The H3O+ ions are
completely neutralized by NH3.
NH3 + HCl --------------> NH4Cl
The salt NH4Cl so formed dissociates into
NH4
+ + Cl-. This NH4
+ ions reacts with H2O
to produce H3O+ ions making the solution
acidic decreasing the pH to ~ 5 at
equivalence point.
NH4Cl + H2O ------------------> NH4
+ + Cl-
NH4
+ + H2O -------------------> NH3 + H3O+
Point d: After the equivalence point, NH3 addition
continues and is in excess, so the pH increases.
NH3 is a weak base so the pH is above 7, but is
lower than what we saw with a strong base
NaOH (case 1).

31. Titration of a weak base with a weak acid: (CH3COOH Vs NH3)
Volume of Acid added
Suppose our weak acid if CH3COOH
and weak base is NH3.
There is not any steep in this plot.
Lack of any steep change in pH
throughout the titration renders
titration of a weak base versus a
weak acid difficult, and not much
information can be extracted from
such a curve.