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AN INSTRUCTION MANUAL- STEP BY STEP GUIDE FOR CONDUCTING

TITRIMETRIC ANALYSIS (AS PER PCI SYLLABUS OF B. PHARMACY-I SEM)

Method · October 2021

DOI: 10.13140/RG.2.2.30916.55683

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 AN INSTRUCTION MANUAL

A STEP BY STEP GUIDE FOR

CONDUCTING TITRIMETRIC
ANALYSIS
[STRICTLY AS PER B. PHARMACY-I SEMESTER
PRACTICAL SYLLABUS ]

Dr. Rajesh K. Singh

MPharm, PhD, FIC
Assistant Professor (Stage-III)
Department of Pharmaceutical Chemistry
Shivalik College of Pharmacy, Nangal
(Managed by Local Govt. Department, Punjab)
District: Ropar, 140126, Punjab
Affiliated to IK Gujaral Punjab Technical University, Jalandhar
Mobile: 9417513730, Landline 01887-221276, Fax 01887-221276
Web: www.shivalikpharmacycollege.in

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SYNOPSIS

The materials of this handbook is mainly concerned with various practical involving titration related to
pharmaceutical analysis subject mentioned in the diploma and undergraduate syllabus of various
Universities. However, at appropriate places some information has been provided which may be useful to
post-graduate and students working in industry (Quality control department) as well.

Apart from discussing the theoretical aspects of each practical, a number of experiments have been
described in detail (with calculation and reason behind series of steps) along with solved numerical for
each practical which will be helpful to students in carrying out various determinations. The purpose of
giving experimental details is to acquaint the students with basic concepts behind analytical techniques.
Once these basic concepts are properly understood, student can successfully perform many more
determination.

The major objective of writing this book is to present the information in a lucid, condensed and cohesive
form, to cater specifically the needs of diploma, undergraduate and graduate students of pharmacy and
chemistry. The teachers can choose any practical depending on their syllabi and chemicals available.

The idea of writing this book was conceived when my own students found great difficulty in getting a
pharmaceutical analysis practical book suited to their intellectual needs. It has been noticed that while
carrying out titrimetric analysis, a majority of students follow instruction mechanically without
understanding the theoretical consideration. Most of the students face problem in numerical and
calculating amount of sample required to make desired strength of standard solution. Moreover I felt that
majority of students find this practical dull and boring to perform as they have great difficulty in getting
concepts behind each step involved in titration method.

Although there were several related practical book, they had very high technical standard, complex non-
descriptive calculations and profound language. In most of the book the calculation as well as procedure
is not fully explained and understood. The material in most of them is presented in a diffused form or is
highly specialized and perceptible to those competent in the field. Therefore, to satisfy the needs of
students, I am going to propose a new practical book with optimum mix of theoretical concepts and
descriptive calculations so that students could easily understand and memorize.
Three key features of the book:

2
1. SIMPLICITY AND EYE-CATCHING FLOWCHARTS: This book is not meant for high level
advanced titration method, but well suited for diploma and undergraduate students for best understanding
and for their practical examination point of view. In writing this book much care has been taken to use
simple language to explain complex procedure. As an experiment and to make this book more
understanding, I have tried to draw a flowchart of procedure at the end of every of practical so that
students could have quickly go through the practical. These flowcharts would be easy to understand and
memorize.

2. REASON BEHIND SERIES OF STEPS: This is the most unique part of this book. To clear the
concept, in addition to brief theory, the reason behind each step in procedure have been provided in
tabulated form so that students could easily understand the importance of every steps of titration
method.

3. CALCULATION & NUMERICALS: This is the most important part of this book. Descriptive
and short-cut methods are given to calculate an amount of sample required to make desired strength i.e.
normality or molarity. Descriptive and short-cut methods are also provided to calculate the %age purity of
various drugs. At the end of each practical, related solved numerical have been provided to acquaint the
students with the concept of that practical. The mathematical calculation has been kept simple in order not
to overburden the students with complexities of equation and formulae.

3
LIST OF CONTENTS

1. Acid-base titration

1.1. To prepare and standardize 0.1 N HCl using sodium carbonate as primary standard.
1.2. To prepare and standardize 0.1 N H 2SO4 using sodium carbonate as primary
standard.
1.3. To prepare and standardize 0.1 N NaOH using succinic acid as primary standard.
1.4. To determine the percentage purity (assay) of given sample of ammonium chloride
using standard 0.1 N NaOH.
1.5. To determine the percentage purity (assay) of acetic acid in a given sample of
vinegar standard 0.1 N NaOH.
1.6. To determine the amount of carbonate and hydroxide in a given sample using
standard 0.1 N HCl.
1.7. To determine the percentage purity (assay) of boric acid in a given sample using
standard 0.1 N NaOH.

2. Oxidation-reduction titrations

2.1. To prepare and standardize 0.1 N KMnO 4 solution using oxalic acid as primary
standard.
2.2. To perform the assay of hydrogen peroxides using standard 0.1 N KMnO4 solution.
2.3. To prepare and standardize 0.1 N sodium thiosulphate solution using potassium
iodate as primary standard.
2.4. To perform the assay of copper sulphate using standard 0.1 N sodium thiosulphate
(Iodometry).
2.5. To prepare and standardize 0.1 N ceric ammonium sulphate using arsenic trioxide as
primary standard.
2.6. To perform the assay of ferrous sulphate using standard 0.1 N ceric ammonium
sulpahte.

3. Argentometric titrations

3.1. To prepare and standardize 0.1 N AgNO3 using sodium chloride as primary standard
(Mohr’s method).
3.2. To determine the percentage purity of given sample of sodium chloride injection
using standard 0.1 N AgNO3 (Mohr’s method).
3.3. To prepare and standardize 0.1 N ammonium thiocyanate solution using standard
AgNO3 as secondary standard.
3.4. To determine the percentage purity of given sample of sodium chloride using
standard 0.1 N AgNO3 (Volhard’s method).

4. Complexometric titration

4.1. To prepare and standardize 0.1 N EDTA using granulated zinc as primary standard.

4
 4.2. To determine the percentage purity of given sample of magnesium sulphate using
standard 0.1 N EDTA (Direct titration).
4.3. To perform the assay of calcium gluconate using standard 0.1 N EDTA
(Replacement titration).

5. Non-aqueous titration

5.1. To prepare and standardize 0.1 N perchloric (HClO4) acid using potassium hydrogen
phthalate as primary standard.
5.2. To determine the percentage purity (assay) of given sample of ephedrine using
standard 0.1 N HClO4.
5.3. To determine the percentage purity (assay) of ephedrine hydrochloride using
standard 0.1 N HClO4.
5.4. To determine the percentage purity of sodium benzoate using standard 0.1 N HClO4.
5.5. To prepare and standardize 0.1 N sodium methoxide solution using benzoic acid as
primary standard.
5.6. To determine the percentage purity of given sample of benzoic acid by standard 0.1
N sodium methoxide.

6. Diazotization titration

6.1. To prepare and standardize 0.1 N sodium nitrite solution using sulfanilic acid as
primary standard.
6.2. To find out the percentage purity of sulfanilic acid by diazotization titration.

7. Conductometric and potentiometric titration

7.1. To determine the strength of strong acid using strong base by conductometry.
7.2. To deterimine the strength of strong acid using strong base by potentiometry.

5
1. Acid-base titration

1.1. To prepare and standardize 0.1 N HCl using sodium carbonate as primary standard.
1.2. To prepare and standardize 0.1 N NaOH using succinic acid as primary standard.
1.3. To prepare and standardize 0.1 N H2SO4 using standard sodium hydroxide as
secondary standard.
1.4. To determine the percentage purity (assay) of given sample of ammonium chloride
using standard 0.1 N NaOH.
1.5. To determine the amount of carbonate and hydroxide in a given sample using
standard 0.1 N HCl.
1.6. To determine the percentage purity (assay) of acetic acid in a given sample of
vinegar standard 0.1 N NaOH.
1.7. To determine the percentage purity (assay) of boric acid in a given sample using
standard 0.1 N NaOH.

EXPERIMENT-1.1

6
 TO PREPARE AND STANDARDIZE 0.1 N HCl USING SODIUM CARBONATE
AS PRIMARY STANDARD
THEORY
Laboratory grade hydrochloric acid cannot be used as primary standard as it is not sufficiently
pure because of its gaseous form at room temperature. So the solution of hydrochloric acid needs
to be standardized before any analytical applications. In order to standardize it, one must have an
especially pure reagent which can be accurately weighed out on analytical balance. This pure
reagent is called primary standards which are extremely pure, stable, has no water of hydration,
and has a high molecular weight.

Various primary standards can be used for standardization of hydrochloric acid like anhydrous
sodium carbonate (eq. wt.= 53), potassium bicarbonate (eq. wt.= 100), thallous carbonate (eq.
wt.= 234.40), borax (eq. wt.= 190.70) etc.

In this experiment, a standard solution of sodium carbonate will be used as the primary standard
to determine the exact concentration of a hydrochloric acid solution. The neutralization reactions
that take place in two steps are as follows:

The titration of sodium carbonate involves two end-points. The first corresponding to the
conversion of carbonate to bicarbonate at a pH 8.3. The second corresponding to the conversion
of bicarbonate to carbonic acid at pH 3.93, which show the completion of the reaction. For this
pH, methyl orange indicator solution can be used. At the end-point – when neutralisation just
occurs – the indicator changes colour from yellow to orange.
Note: a) All burette readings must be read and recorded to 2 decimal places. For example, if the volume of the
solution in the burette used was 10.5 ml, record down as 10.50 ml and not 10.5. The zero is significant and must be
included.
b) Round off your answers to 3 significant figures. For example, if your result is 0.12 N, record down as 0.12 N and
not 0.1 N.

7
PROCEDURE

Calculation:
Calculation for making 0.1 N HCl acid: To prepare 500 ml of a 0.1 N HCl, 4.20 ml of conc HCl
will be diluted to 500.0 ml of distilled water. Before making any calculations, use only one
system and one unit of measurement. DO NOT mix measurement systems and units.

Descriptive method Shortcut method 1

1000 ml of 1 N HCl =36.50 g eq. of HCl M= No. of moles or Mass (g) x 1000
500 ml of 0.1 N HCl= 36.5 x 500 x 0.1 Volume (L) Equivalent mass x volume (ml)
1000
= 1.82 g of HCl 0.1 = Mass (g) x 1000 = Mass = 1.82 g of HCl
Density= Mass 36.50 x 500
Volume
Volume=Mass 1.82 g = 1.51 ml of HCl Density= Mass
Density 1.20 g/ml Volume
Volume=Mass 1.82 g = 1.51 ml of HCl
Since the purity of conc. HCl is 36.50 % Density 1.20 g/ml
Therefore, 36.50 % = 1.51 ml
100 % = 1.51 x 100 = 4.20 ml Since the purity of conc. HCl is 36.50 %
36.5 Therefore, 36.50 % = 1.51 ml
100 % = 1.51 x 100 = 4.20 ml
36.5
Shortcut method 2
N1V1 = N2V2
(Desired strength) (Conc. HCl)
V2= N1V1 = 0.1 x 500 = 4.3 ml, N2 =Normality of
N2 11.50 conc. HCl acid
Hence, about 4.2 ml of HCl will be measured and dissolve in 200 ml of distilled water and
volume will be made up to 500 ml to make 0.1 N HCl solution.

Preparation of 0.1 N standard sodium carbonate solution

Descriptive method Shortcut method
1000 ml of 1.0 N Na2CO3 = 53.0 g eq. of M= No. of moles or Mass (g) x 1000
Na2CO3 Volume (L) Equivalent mass x volume (ml)
(Since two mol of HCl will react with one mol
of Na2CO3 )
100 ml of 0.1 N HCl= 53.0 x 100 x 0.1 0.1 = Mass (g) x 1000 = Mass= 0.53 g of HCl
1000 53.0 x 100
= 0.53 g of Na2CO3

Hence, Weigh accurately nearly 0.53 g of sodium carbonate and dissolve in 50 ml of distilled
water and make up the volume to 100 ml.

8
Standardization of HCl solution using sodium carbonate as primary standard.

Preparation of Hydrochloric acid solution:

S. No Steps Reason
1. Measure about 4.20 ml of HCl by measuring cylinder As per above calculation.
and dissolved in 200 ml of distilled water and make
up the volume to 500 ml to make 0.1 N HCl. Never
add water to concentrated acid.
2. Rinse the burette with distilled water and fill the Burette is rinsed to wash out
burette with above prepared hydrochloric acid any impurities left while doing
solution to the zero mark using funnel previous titration.
3. Remove the funnel from the burette and note the So that drops of solution from
reading in note book. Eye must be horizontal to the funnel will not fall into the
meniscus. burette. Eye must be
horizontal to meniscus to
remove parallax erro.

Preparation of standard 0.1 N sodium carbonate solution (standard solution):

S. No Steps Reason
1. Weigh nearly about 0.53 g of sodium carbonate as Sodium carbonate is very pure
primary standard in watch glass and dissolve in and stable and therefore use as
volumetric flask (100 ml) where 50 ml distilled water primary standard.
is already filled.
2. Stopper the flask and mix your solution thoroughly by To ensure proper mixing
inverting the flask. without producing bubbles.
3. Make up the volume to the mark using a dropping To prevent ‘overshooting’ the
pipette to add the last few milliliters of distilled water.
mark resulting in change in
strength. If this occurs, the
experiment will have to be
started again.
4. Transfer the prepared solution to a clean, dry storage As volumetric flask is
bottle and label it if want to store overnight. NEVER transparent, this may degrade
store solutions in a volumetric flask. light sensitive solution so store
in amber coloured bottle.
5. Never hold large volumetric flask by the neck To prevent accidental
alone. Provide support at the bottom. breakage of apparatus.

6. Using the clean pipette (rinsed with distilled water To ensure that all the sodium
and sodium carbonate solution), transfer 10.0 ml of carbonate solution is
sodium carbonate solution into the clean conical flask transferred to the volumetric
(rinsed with distilled water only). flask.
7. When the whole solution has been drained into To allow the last of the liquid

9
 conical flask, touch the tip of the pipette to the side of to drain out. A small amount
the flask. DO NOT blow out the pipette. of solution will always remain
at the pipette’s tip.

8. Add 2-3 drops of methyl orange indicator. Note the Methyl orange is best
colour of the solution. indicator at this pH.

9. Place the conical flask on a white tile or sheet of white To assist in seeing the change
paper. in colour at the end-point.

Titration of hydrochloric acid solution and sodium carbonate solution (standard):

S. No Steps Reason
1. Carry out a rough titration by adding hydrochloric Constant swirling will ensure
acid solution from the burette with constantly complete mixing of reactants.
swirling the flask, until the colour of the solution in
the conical flask changes from yellow to reddish-
yellow.
2. Note the burette reading and calculate how much acid This information enables the
was used. This is only a trial titration and gives the subsequent titrations to be
approximate value of the end point. carried out more quickly.

3. Again, pipette 10.0 ml of sodium carbonate solution Methyl orange is best

into the flask and add 2-3 drops of methyl orange indicator at this acidic pH.
indicator.

4. Add hydrochloric acid solution from the burette To get end point quickly and
rapidly at the start of the titration and then slowly near more accurately.
the rough end point until the colour of the solution
changes from yellow to orange.

5. Repeat the titration until you get the concordant This will minimize error by
results (two readings agree within 0.1 ml) getting accurate readings
within 0.1 ml of each other.
6. Donot try to match the colour of the repeat On long standing, the colour
titration with the previous titration. Just watch for may fade or disaapear.
the change in colour.
7. Take the average of the concordant results To minimize the error.
excluding the rough titration reading and calculate
the concentration of hydrochloric acid.

Observation Table (Specimen reading)

S. Volume of Na2CO3 solution Burette reading Volume of HCl
No (ml) Initial Final used (ml)

10
1. 10.0 (Rough titration) 0.0 9.6 10.10
2. 10.0 10.0 19.90 9.50
3. 10.0 20.0 29.80 9.60

Average of 2 and 3 reading= 9.50 + 9.60 = 9.55

2

Nacid x Vacid = N(sod. carb.) x V(sod. carb.)

Nacid = N(sod. carb.) x V(sod. carb = 0.1 x 10 Volume of Na2CO3 solution in each titration = 10 ml
Vacid 9.55 Concentration of Na2CO3 = 0.1 M

N1 = 0.105

RESULT
The normality of hydrochloric acid solution is 0.105 N.

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. Calculate the approximate weight of sodium carbonate required so that about 15.0 mL of 0.20
N NaOH will be consumed in a titration. (E.W. sodium carbonate = 53 g/equiv.)

11
 N= No. of moles or Mass (g) x 1000
Volume (L) Eq. weight x volume (ml)

Mass= Normality x Volume (ml) x Eq. weight

1000

Mass= 0.20 x 53.0 x 15.0 = 0.159 g of sodium carbonate

1000

Hence, 0.159 g of sodium carbonate is required to titrate with 20 ml of 0.20 N

NaOH

2. If 20 ml of the 0.1 N sodium carbonate solution uses 10 ml of the

hydrochloric acid solution for complete neutralization. Calculate the
normality of the acid. (E.W. sodium carbonate = 53 g/equiv.)

N1 x V1 = N2 x V2 Nacid = 20.0 x 0.1 = 0.2 N

(Acid) (Na2CO3) 10.0

Hence, normality of the hydrochloric acid is 0.20 N

3. If 20 ml of the HCl solution reacts with 25.0 ml of 0.15 N NaOH. Calculate the strength of the
HCl solution in grams of HCl per litre. (E.W. of HCl = 36.5 g/equiv.)

Strength in grams per litre = Normality x Equivalent weight of the solute.

N1 x V1 = N2 x V2 Nacid = 25.0 x 0.50 = 0.625 N

(HCl) (NaOH) 20.0

Strength of HCl solution = 0.625 x 36.5 = 22.8 grams of HCl/litre

4. Titration reveals that 15.50 ml of 0.15 M HCl are required to neutralize 25.0 ml of
NaOH solution. What is the normality and strength of the NaOH solution?

12
Flowchart illustrating the preparation and standardization of 0.1 N HCl.

13
 EXPERIMENT-1.2

TO PREPARE AND STANDARDIZE 0.1 M SULPHURIC ACID USING SODIUM

CARBONATE AS PRIMARY STANDARD
THEORY
Same as previous experiment no 1.
It involves the following reaction:

The end point can be detected by using methyl orange as an indicator.

PROCEDURE
A. Preparation of reagents and solutions:
1. 0.1 M sulphuric acid

Calculation for making 500 ml of 0.1 N sulphuric acid solution:

Descriptive method
1000 ml of 1 M H2SO4 =98 g of H2SO4
500 ml of 0.1 N H2SO4= 98 x 500 x 0.1
1000
= 4.9 g of H2SO4
Density= Mass
Volume
Volume=Mass 4.9 g = 2.66 ml of H2SO4
Density 1.84 g/ml

Shortcut method 2
N1V1 = N2V2
(Desired strength) (Conc. HCl)
V2= N1V1 = 0.1 x 500 = 2.69 ml, N2 =Normality of
N2 18.6 conc. H2SO4 acid

To prepare 500 ml of a 0.1 N H 2SO4, 2.7 ml of conc H2SO4 will be added to 100.0 ml of distilled
water and volume will be made up to 500 ml. Before making any calculations, use only one
system and one unit of measurement. DO NOT mix measurement systems and units.

Preparation of 0.1 N standard sodium carbonate solution

Descriptive method Shortcut method
1000 ml of 1.0 N Na2CO3 = 53.0 g eq. of M= No. of moles or Mass (g) x 1000
Na2CO3 Volume (L) Equivalent mass x volume (ml)
(Since two mol of HCl will react with one mol
of Na2CO3 )

14
100 ml of 0.1 N HCl= 53.0 x 100 x 0.1 0.1 = Mass (g) x 1000 = Mass= 0.53 g of HCl
1000 53.0 x 100
= 0.53 g of Na2CO3

Hence, Weigh accurately nearly 0.53 g of sodium carbonate and dissolve in 50 ml of distilled
water and make up the volume to 100 ml.

Standardization of H2SO4 solution using sodium carbonate as primary standard.

Preparation of Hydrochloric acid solution:

S. No Steps Reason
1. Measure about 2.7 ml of H2SO4 by measuring cylinder As per above calculation.
and dissolved in 100 ml of distilled water and make
up the volume to 500 ml. Never add water to
concentrated acid.
2. Rinse the burette with distilled water and fill the Burette is rinsed to wash out
burette with above prepared sulphuric acid acid any impurities left while doing
solution to the zero mark using funnel previous titration.
3. Remove the funnel from the burette and note the So that drops of solution from
reading in note book. Eye must be horizontal to the funnel will not fall into the
meniscus. burette. Eye must be
horizontal to meniscus to
remove parallax erro.

Preparation of standard 0.1 N sodium carbonate solution (standard solution):

S. No Steps Reason
1. Weigh nearly about 0.53 g of sodium carbonate as Sodium carbonate is very pure
primary standard in watch glass and dissolve in and stable and therefore use as
volumetric flask (100 ml) where 50 ml distilled water primary standard.
is already filled.
2. Stopper the flask and mix your solution thoroughly by To ensure proper mixing
inverting the flask. without producing bubbles.
3. Make up the volume to the mark using a dropping To prevent ‘overshooting’ the
pipette to add the last few milliliters of distilled water.
mark resulting in change in
strength. If this occurs, the
experiment will have to be
started again.
4. Transfer the prepared solution to a clean, dry storage As volumetric flask is
bottle and label it if want to store overnight. NEVER transparent, this may degrade
store solutions in a volumetric flask. light sensitive solution so store

15
 in amber coloured bottle.
5. Never hold large volumetric flask by the neck To prevent accidental
alone. Provide support at the bottom. breakage of apparatus.

6. Using the clean pipette (rinsed with distilled water To ensure that all the sodium
and sodium carbonate solution), transfer 10.0 ml of carbonate solution is
sodium carbonate solution into the clean conical flask transferred to the volumetric
(rinsed with distilled water only). flask.
7. When the whole solution has been drained into To allow the last of the liquid
conical flask, touch the tip of the pipette to the side of to drain out. A small amount
the flask. DO NOT blow out the pipette. of solution will always remain
at the pipette’s tip.

8. Add 2-3 drops of methyl orange indicator. Note the Methyl orange is best
colour of the solution. indicator at this pH.

9. Place the conical flask on a white tile or sheet of white To assist in seeing the change
paper. in colour at the end-point.

Titration of sulphuric acid solution (in burette) and standard sodium carbonate solution (in
beaker):

S. No Steps Reason
1. Carry out a rough titration by adding sulphuric acid Constant swirling will ensure
solution from the burette with constantly swirling the complete mixing of reactants.
flask, until the colour of the solution in the conical
flask changes from yellow to reddish-yellow.
2. Note the burette reading and calculate how much acid This information enables the
was used. This is only a trial titration and gives the subsequent titrations to be
approximate value of the end point. carried out more quickly.

3. Again, pipette 10.0 ml of sodium carbonate solution Methyl orange is best

into the flask and add 2-3 drops of methyl orange indicator at this acidic pH.
indicator.

4. Add hydrochloric acid solution from the burette To get end point quickly and
rapidly at the start of the titration and then slowly near more accurately.
the rough end point until the colour of the solution
changes from yellow to orange.

5. Repeat the titration until you get the concordant This will minimize error by
results (two readings agree within 0.1 ml) getting accurate readings
within 0.1 ml of each other.
6. Donot try to match the colour of the repeat On long standing, the colour
titration with the previous titration. Just watch for may fade or disaapear.

16
 the change in colour.
7. Take the average of the concordant results To minimize the error.
excluding the rough titration reading and calculate
the concentration of sulphuric acid.

Observation Table (Specimen reading)

S. Volume of Na2CO3 solution Burette reading Volume of
No (ml) Initial Final H2SO4 used (ml)
1. 10.0 (Rough titration) 0.0 9.6 10.10
2. 10.0 10.0 19.90 9.50
3. 10.0 20.0 29.80 9.60

Average of 2 and 3 reading= 9.50 + 9.60 = 9.55

2

Nacid x Vacid = N(sod. carb.) x V(sod. carb.)

Nacid = N(sod. carb.) x V(sod. carb = 0.1 x 10 Volume of Na2CO3 solution in each titration = 10 ml
Vacid 9.55 Concentration of Na2CO3 = 0.1 M

N1 = 0.105

RESULT

The normality of sulphuric acid solution is 0.105 N.

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. Calculate the approximate weight of sodium carbonate required so that about 15.0 mL of 0.20
N NaOH will be consumed in a titration. (E.W. sodium carbonate = 53 g/equiv.)

N= No. of moles or Mass (g) x 1000

Volume (L) Eq. weight x volume (ml)

Mass= Normality x Volume (ml) x Eq. weight

1000

Mass= 0.20 x 53.0 x 15.0 = 0.159 g of sodium carbonate

17
 1000

Hence, 0.159 g of sodium carbonate is required to titrate with 20 ml of 0.20 N

NaOH

2. If 20 ml of the 0.1 N sodium carbonate solution uses 10 ml of the

hydrochloric acid solution for complete neutralization. Calculate the
normality of the acid. (E.W. sodium carbonate = 53 g/equiv.)

N1 x V1 = N2 x V2 Nacid = 20.0 x 0.1 = 0.2 N

(Acid) (Na2CO3) 10.0

Hence, normality of the hydrochloric acid is 0.20 N

3. If 20 ml of the HCl solution reacts with 25.0 ml of 0.15 N NaOH. Calculate the strength of the
HCl solution in grams of HCl per litre. (E.W. of HCl = 36.5 g/equiv.)

Strength in grams per litre = Normality x Equivalent weight of the solute.

N1 x V1 = N2 x V2
(HCl) (NaOH)

Nacid = 25.0 x 0.50 = 0.625 N

20.0

Strength of HCl solution = 0.625 x 36.5 = 22.8 grams of HCl/litre

4. Titration reveals that 15.50 ml of 0.15 M HCl are required to neutralize 25.0 ml of
NaOH solution. What is the normality and strength of the NaOH solution?

18
Flowchart illustrating the preparation and standardization of 0.1 N HCl

19
 EXPERIMENT-1.3

TO PREPARE AND STANDARDIZE 0.1 N NaOH USING SUCCINIC ACID AS

PRIMARY STANDARD

THEORY
Sodium hydroxide cannot be used as a primary standard because of its hygroscopic properties as
a solid. Because it is so prone to absorbing water, it is impossible to accurately measure the mass
of a solid sample, so instead it must be put into solution and titrated with a known acidic
solution.
Solid sodium hydroxide cannot be used as a primary standard as it is not available in pure form
because it absorbs atmospheric moisture and carbon dioxide during storage and also during a
weighing operation.

Since hydroxide ion is consumed by this reaction, the concentration of a standard sodium
hydroxide solution will be changed.
Another reason for not using sodium hydroxide solutions as primary standard is that it often
contains significant amounts of impurities as carbonates and bicarbonates, and is highly
hygroscopic. Therefore the solution of sodium hydroxide needs to be standardized before any
analytical applications.
Various primary standards can be used for standardization of sodium hydroxide like succinic
acid (eq. wt. = 59), potassium hydrogen phthalate (eq. wt. = 204.0), benzoic acid (eq. wt.=
122.0), sulfamic acid (eq. wt.= 97.0) etc.

In the present experiment, a standard solution of succinic acid is used as the primary standard to
determine the exact concentration of a sodium hydroxide solution. The succinic acid is stable,
pure, well characterized material which dissolves in water to produce H+ and succinate ions.
Succinic acid contains two titrable acidic hydrogens, which react according to the equations as
follows:

20
The pH at the completion of this reaction is around 8.0 (basic). For this pH, phenolphthalein
indicator solution can be used. At the end-point – when neutralisation just occurs – the indicator
changes colour from colourless to faint pink.

PROCEDURE

A. Preparation of reagents and solutions:

1. 0.1 M sodium hydroxide:

Calculation for making 100 ml of 0.1 N NaOH solution:

Descriptive method Shortcut method

1000 ml of 1 N NaOH = 40 g of NaOH N= No. of moles or Mass (g) x 1000
100 ml of 0.1 N NaOH= 40 x 100 x 0.1 Volume (L) Molar mass x volume (ml)
1000
= 0.4 g of NaOH 0.1 = Mass (g) x 1000 = Mass = 0.4 g of NaOH
40 x 100

Hence, 0.4 g of NaOH will be weighed and dissolve in 50 ml of distilled water and volume will
be made up to 100 ml to make 0.1 M NaOH solution.

2. 100 ml of 0.1 M standard succinic acid solution

Descriptive method Shortcut method

1000 ml of 1 N succinic acid =59.0 g eq. of N= No. of moles or Mass (g) x 1000
succinic acid Volume (L) Equivalent mass x volume (ml)
(Since two mol of NaOH will react with one
mol of succinic acid )
100 ml of 0.1 N HCl= 59.0 x 100 x 0.1 0.1 = Mass (g) x 1000 = Mass= 0.59 g of HCl
1000 59.0 x 100
= 0.59 g of Na2CO3

Hence, weigh nearly about 0.59 g of succinic acid and dissolved in 50 ml of distilled water and
make up the volume to 100 ml.

3. Methyl orange indicator

B. Titration steps
Standardization of NaOH solution using succinic acid as primary standard

Preparation of 0.1 N NaOH solution:

.

21
S. No Steps Reason
1. Weigh about 2.0 g of NaOH and dissolve in 200 ml of Tap and even deionised water
freshly prepared distilled water and make up the may contain dissolved carbon
volume to 500 ml to make 0.1 N NaOH. dioxide.

2. Fill the burette with as much NaOH solution as To prevent exposing NaOH to
needed for titration. PREPARE THE BURETTE the atmospheric moisture and
WHEN IT NEEDED. CO2.
3. Remove the funnel from the burette and note the So that drops of solution from
reading in note book. the funnel will not fall into the
burette

Preparation of standard solution of 0.1 N succinic acid:

S. No Steps Reason
1. 0.53 g of succinic acid as primary standard will be Succinic acid is very pure and
weighed in watch glass and dissolve in 100 stable and therefore use as
volumetric flask where 50 ml distilled water is already primary standard.
filled.
2. Stopper the flask and mix your solution thoroughly by To ensure proper mixing
inverting the flask. without producing bubbles.
3. Make up the volume to the mark using a dropping To prevent ‘overshooting’ the
pipette to add the last few milliliters of distilled water.
mark resulting in change in
strength. If this occurs, the
experiment will have to be
started again.
4. Transfer the prepared solution to a clean, dry storage As volumetric flask is
bottle and label it if want to store overnight. NEVER transparent, this may degrade
store solutions in a volumetric flask. light sensitive solution so store
in amber coloured bottle.
5. Never hold large volumetric flask by the neck To prevent accidental
alone. Provide support at the bottom. breakage of apparatus.

6. Using the clean pipette (rinsed with distilled water To ensure that all of the
and succinic acid solution), transfer 10 ml of succinic succinic acid solution is
into the clean conical flask (rinsed with distilled water transferred to the volumetric
only). flask.
7. When the whole solution has been drained, touch the To allow the last of the liquid
tip of the pipette to the side of the flask. to drain out. DO NOT blow
out the pipette.

8. Add 2-3 drops of phenolphthalein indicator. Note the Phenolphthalein (transition pH

colour of the solution. range 8.0-9.6) is the
appropriate indicator at this
pH.

22
Titration of sodium hydroxide solution and succinic acid solution (standard)

S. No Steps Reason
1. Carry out a rough titration by adding NaOH solution Constant swirling will ensure
from the burette with constantly swirling the flask, complete mixing of reactants.
until the colour of the solution in the conical flask
changes from colorless to faint pink.
2. Note the burette reading and calculate how much base This information enables the
was used. This is only a trial titration and gives the subsequent titrations to be
approximate value of the end point. carried out more quickly.

3. Again, pipette 10 ml of succinic acid solution into the Phenolphthalein is best

flask and add 2-3 drops of phenolphthalein indicator. indicator at this pH.
Fill the burette with fresh sodium hydroxide solution.

4. Add sodium hydroxide solution from the burette To get end point quickly and
rapidly at the start of the titration and then slowly near more accurately.
the rough end point until the colour of the solution
changes from colourless to faint pink.
5. Repeat the titration until you get the concordant This will minimize error by
results (two readings agree within 0.1 ml) getting accurate readings
within 0.1 ml of each other.
6. Donot try to match the colour of the repeat On long standing, the colour
titration with the previous titration. Just watch for may fade or disappear.
the change in colour.
7. Take the average of the concordant results To minimize the error.
excluding the rough titration reading and calculate
the normality of sodium hydroxide.

Observation Table (Specimen reading)

S. Volume of succinic acid Burette reading Volume of NaOH
No solution (ml) Initial Final used (ml)
1. 10.0 (Rough titration) 0.0 9.6 10.1
2. 10.0 10.0 19.9 9.50
3. 10.0 20 29.8 9.60

Average of 2 and 3 cocordant readings= 9.50 + 9.60 = 9.55

2
NNaOH x VNaOH = N(succ. acid.) x V(succ.acid.)

NNaOH = N(succ. acid) x V (succ. acid) = 0.1 x 10 Volume of succinic acid solution in each titration = 10 ml
VNaOH 9.55 Normality of succinic acid = 0.1 M

NNaOH = 0.105 N
RESULT
Concentration of sodium hydroxide solution is 0.105 M.

23
Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. Calculate the approximate weight of succinic required so that about 20.0 ml of 0.1 N sodium
hydroxide will be used in a titration. (E.W. succinic acid= 59.0 g/equiv.)

N= No. of moles or Mass (g) x 1000

Volume (L) Eq. weight x volume (ml)

Mass= Normality x Volume (ml) x Eq. weight

1000

Mass= 0.1 x 59.0 x 20.0 = 0.118 g of succinic acid

1000

Hence, 0.118 g of succinic acid is required to titrate with 20 ml of 0.1 N NaOH

2. Calculate the normality of a solution of potassium hydrogen phthalate (KHP) prepared by

mixing a 0.50 g in 50.0 ml of water. (E.W. KHP = 204.0 g/equiv.)

N= No. of moles or Mass (g) x 1000

Volume (L) Eq. weight x volume (ml)

N = 0.50 x 1000 = 0.049 N

204 x 50

Hence, the normality of 0.049 N will be prepared by mixing 0.5 g of KHP in

50.0 ml of water.

3. Calculate the approximate weight of potassium hydrogen phthalate (KHP) require to prepare
100 ml of 0.1 N KHP solutions for standardization of sodium hydroxide solution. (E.W. KHP =
204 g/equiv.)

24
Flowchart illustrating the preparation and standardization of 0.1 N NaOH

25
 EXPERIMENT-1.4

TO DETERMINE THE PERCENTAGE PURITY OF GIVEN SAMPLE OF

AMMONIUM CHLORIDE USING STANDARD 0.1 N NaOH

THEORY

% purity is the percentage of the material which is the actually desired chemical in a sample of it.
In pharmaceutical industry, it would not be acceptable to manufacture a drug with impurities in it
that may be harmful to health. However in any chemical process it is almost impossible to get
100.00 % purity and so sample should be analyzed in industry for % purity to monitor the quality
of the product.
Ammonium chloride is also known as the salt of ammonia. It is represented by a chemical
formula NH4Cl. Ammonium chloride when dissolve in water form acidic solution. Reaction
between ammonium chloride and sodium hydroxide produces some new compounds like
ammonia, water and sodium chloride. Ammonia gas liberated may combine with hydrochloric
acid to form ammonium chloride and hence direct titration of ammonium chloride with sodium
hydroxide produce erroneous results.

So for the titration of ammonia chloride with base, the addition of formaldehyde would improve
the titration. The ammonium chloride reacts with formaldehyde to form hexamethylene
tetramine. Because the weak acid ammonium (pKa 9.3) is converted to the stronger
hexamethylene tetramine ion (pKa 4.9). This improves the end point.

26
PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 M sodium hydroxide: See experiment 1.3
B. Titration Procedure:

S. No Steps Reason
1. 0.1 N NaOH will be prepared and standardized using As per the procedure
succinic acid. explained previously in exp.
no.
2. Weigh accurately about 0.1 g of ammonium This helps to make sure that
chloride in triplicate in analytical balance using results are consistence with
weighing by difference method directly into a conical each other. Scientific evidence
flask and add 20 ml of water. must be reproducible.

3. Add a mixture of 5.0 ml of formaldehyde solution Formaldehyde improves the

(previously neutralize to dilute phenolphthalein) and titration by forming
20 ml of water. hexamethylenetetramine.

4. After two minutes, add 2-3 drops of phenolphthalein To get end point quickly and
indicator and titrate solution slowly against 0.1 N more accurately.
NaOH till faint pink colour appear.

5. Note down the burette reading. Repeat the titration Averaging of the results will
with other two flasks and take the average readings of minimize the random error.
sample taken and volume of NaOH used.

Note: Neutralization of formaldehyde: Formaldehyde is neutralized by taking in a conical flask

20 ml formaldehyde solution and two drops of phenolphthalein then titrate against 0.1 N NaOH
till faint pink colour appear.

Observation Table (Specimen reading)

S. Amount of sample taken (g) Burette reading Volume of NaOH

No Initial Final used (ml)
1. 0.095 0.0 17.0 17.0
2. 0.11 0.0 18.2 18.2
3. 0.10 0.0 17.8 17.8

Total 0.305 Total 53.0

Mean of sample taken 0.305/3=0.101 g Mean of volume used 53.0/3= 17.6 ml

Calculation for percentage purity:

Descriptive method Short cut method

27
1000 ml of 1 N NaOH=53.5 g eq. of NH4Cl Percentage purity= V x N x 53.5 x 100
17.6 ml of 0.1 N NaOH= 53.5 x 17.6 x 0.1 W (gm) x 1000
1000
= 0.0942 g V=Volume of NaOH (ml)
0.101 g of sample contains= 0.0942 g of N= Normality of NaOH
NH4Cl W= Weight of sample (gm)
100 gm of sample contain = 0.0942 x 100 17.6 x 0.1 x 53.50 x 100
0.101 0.1 x 1000
= 93.28 %
= 93.28 %

RESULT
The given sample of ammonium chloride contain 93.28 % of pure ammonium chloride.

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. A 0.5 gm sample of ammonium chloride required 12 ml of 0.1 N for titration. Calculate the
percentage purity (assay) of the ammonium chloride.

Percentage purity= V x N x 53.5 x 100

W (gm) x 1000

V=Volume of NaOH (ml)

N= Normality of NaOH
W= Weight of sample (gm)
12 x 0.1 x 53.5 x 100
0.5 x 1000
= 12.84 %

2. Suppose 30 ml of 0.2 M NaOH is required to titrate 1.5 g of a sample containing ammonium

chloride. What percentage ammonium chloride is present in the sample?

21.4 %

28
Flowchart illustrating the assay (% purity) of ammonium chloride
EXPERIMENT-1.5

29
 TO DETERMINE THE PERCENTAGE PURITY OF ACETIC ACID IN A GIVEN
SAMPLE OF VINEGAR USING 0.1 N NaOH

THEORY

Acetic acid (vinegar) is the main component of Vinegar. It is a carbon based compound with a
single ionizable proton with a –COOH functional moiety

According to the reaction equation

PROCEDURE
A. Preparation of reagents and solutions:
2. Preparation and standardization of 0.1 M sodium hydroxide: See experiment no.
3
B. Titration Procedure:
C.
PROCEDURE
S. No Steps Reason
1. 0.1 N NaOH will be prepared and standardized using As per the procedure
succinic acid. explained previously.
2. Weigh accurately about 2 g or 2 ml (density of acetic This helps to make sure that
acid is ~ 1.0, so you can take 2 ml) in triplicate of results are consistence with
vinegar in volumetric flask and dilute with 50 ml of each other. Scientific evidence
water. must be reproducible.

3. Add 2-3 drops of phenolphthalein indicator and titrate Phenolphthalein is the

solution slowly with 0.1 N NaOH. appropriate indicator at this
pH.
4. Note down the burette reading. Repeat the titration Averaging of the results will

30
 with other two flasks and take the average readings of minimize the random error.
sample taken and volume of NaOH used.

Observation Table (Specimen reading)

S. Amount of sample taken (g or Burette reading Volume of NaOH
No ml) Initial Final used (ml)
1. 2.0 0.0 16.40 16.40
2. 2.5 0.0 19.70 19.70
3. 1.5 0.0 11.60 11.60
Total 6.0 Total 47.70
Mean of sample taken 6.0/3= 2.0 g Mean of volume used 47.70/3= 15.90 ml
Calculation

Descriptive method Short cut method

1000 ml of 1 N NaOH=60 g eq. of CH3COOH Percentage purity= V x N x 60 x 100
15.9 ml of 0.1 N NaOH= 60 x 15.90 x 0.1 W (gm) x 1000
1000
= 0.0954 g of CH3COOH V=Volume of NaOH (ml)
2.0 g of sample contains = 0.096 N= Normality of NaOH used
g of CH3COOH W= Weight of sample (gm)
100 gm of sample contain = 0.0954 x 100 15.90 x 0.1 x 60 x 100
2 2 x 1000
= 4.77 %
= 4.77 %

RESULT
Percentage purity of given sample of acetic acid contain 4.77 % of pure acetic acid.

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. A vinegar sample is titrated with 0.54 M NaOH. In one trial, 2.5 g vinegar requires 11.0 ml of
NaOH to reach a phenolphthalein endpoint. Find the % purity of acetic acid in this vinegar.

31
Descriptive method Short cut method
1000 ml of 1 N NaOH=60 g eq. of CH3COOH Percentage purity= V x N x 60 x 100
7.0 ml of 0.54 N NaOH= 60 x 11.0 x 0.54 W (gm) x 1000
1000
= 0.3564 g V=Volume of NaOH (ml)
2.5 g of sample contains=0.3564 g of CH3COOH N= Normality of NaOH
100 gm of sample contain = 0.3564 x 100 W= Weight of sample (gm)
2.5 11 x 0.54 x 60 x 100
2.5 x 1000
= 14.2 % = 14.2 %

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. 1.5 g acid sample contaminated with some impurities was taken in 100 ml volumetric
flask and volume was made up to the mark. To titrate 10 ml of this acid sample need 20
ml of 0.1N NaOH solution. What is the percentage purity of the sample?

1000 ml of 1 N NaOH=60 g eq. of CH3COOH ≡ 60 g of CH3COOH acid

20.0 ml of 0.1 N NaOH= 60 x 20.0 x 0.1
1000
= 0.12 g of CH3COOH acid
i.e. 10 ml of acid sample contain 0.12 g of CH3COOH
100 ml of acid sample will have 1.2 g of CH3COOH acid
That mean, 1.5 g of sample contain 1.2 g of CH3COOH acid
100 gm of sample contain = 1.2 x 100
1.5
= 80.0 %
2. 3.0 g sample of unknown acid was dissolved in distilled water and diluted to 50 ml. What was
the percentage of acetic acid in the sample, if titration of 10.0 ml aliquot of the diluted solution
requires on average 25.0 ml of 0.12 N solution of NaOH.
30.0 %

32
Flowchart illustrating the assay (% purity) of acetic acid in a given sample of vinegar

EXPERIMENT 1.6

33
 TO DETERMINE THE AMOUNT OF CARBONATE AND HYDROXIDE IN A GIVEN
SAMPLE USING STANDARD 0.1 N HCl

THEORY

Sodium hydroxide or caustic soda - both solid and dissolved – is the most typical of the strong
alkalis. It is highly hygroscopic and easily reacts with atmospheric carbon dioxide. It is also
highly deliquescent and absorbs moisture to form an aqueous solution. That means it is usually
contaminated with sodium carbonate Na2CO3. It is not a problem to determine sum of hydroxide
and carbonates concentration by titration with a strong acid.

Solution of caustic soda contains three bases - OH-, CO32- and HCO3-. The stronger the base, the
easier it react with acid. Of the three bases present, NaOH is the strongest, so it will be
neutralized first corresponding to reaction A. The change is pH near the equivalence point is 4.5
to 9.5 and phenolphthalein can be used as an indicator. Next one is CO32- . Sodium carbonate
reacts with HCl in two stages by reaction B and C.

There are three reactions taking place during titration:

When all CO32- is converted to HCO3- (corresponding to reaction B), pH of the solution is 8.35
and and completion of the reaction B (HCO3-), the pH of the solution is 3.90. The first end point
can, therefore be detected by means of phenolphthalein and the second end point by using methyl
orange as indicator.

PROCEDURE

34
S. No Steps Reason
1. 0.1 N HCl will be prepared and standardized using As per the procedure
sodium carbonate. explained previously.
2. Weight 0.1 g of sample of caustic soda in triplicate This helps to make sure that
in analytical balance using weighing by difference results are consistence with
method directly into conical flask and add 10 ml of each other. Scientific evidence
distilled water. must be reproducible.

3. Add 2 to 3 drops of phenolphthalein indicator and The change is pH near the

titrate against standard 0.1 N HCl till the colour equivalence point is 4.5 to 9.5
changes from faint pink to colourless. and phenolphthalein is
appropriate indicator.
4. Note the burette reading which give the volume of Since Na2CO3 has been
HCl used for the complete neutralization of NaOH converted to bicarbonates.
and half neutralization of Na2CO3.
5. After that add 2-3 drops of methyl orange indicator Methyl orange is the best
(colour changes to orange) in the same conical flask indicator at this pH range.
and again continue the titration with standard 0.1 N
HCl till the colour changes from red to yellow.
6. Note down the burette reading. This give the volume of HCl
used for the complete
neutralization of Na2CO3.
7. Repeat the titration with other two flasks and take the Averaging the results will
average readings of sample taken and volume of minimize the random error.
NaOH used.

Calculation (Specimen reading)

Volume of 0.1 N HCl used for the neutralization of ½ Na2CO3 and NaOH (X) = 18.0 ml
Volume of 0.1 N HCl used for the neutralization of complete NaOH + ½Na 2CO3 + ½Na2CO3 (Y)
ml= 20 ml
Volume of 0.1 N HCl used for the neutralization of ½Na2CO3 = (Y-X) = 20 - 18 ml
= 2.0 ml
Volume of 0.1 N HCl used for the neutralization of complete Na2CO3 = 2(Y-X) = 4.0 ml
Volume of 0.1 N HCl used for the neutralization of NaOH = X - 2(Y-X) = 18 – 4 = 16.0 ml

To calculate the % purity of Na2CO3

Descriptive method Short cut method
1000 ml of 1 N HCl=53.0 g eq. of Na2CO3 Percentage purity= V x N x 53 x 100
4.0 ml of 0.1 N HCl = 53 x 4.0 x 0.1 W (gm) x 1000
1000
= 0.0212 g V=Volume of HCl (ml)
0.1 g of sample contains= 0.0212 g of N= Normality of HCl
Na2CO3 W= Weight of sample (gm)

35
100 gm of sample contain = 0.0212 x 100
0.1 Percentage purity= 4.0 x 0.1 x 53.5 x 100
0.1 x 1000
= 2.12 % = 2.12 %

To calculate the % purity of NaOH

Descriptive method Short cut method
1000 ml of 1 N HCl=40 g eq. of NaOH Percentage purity= V x N x 40 x 100
16 ml of 0.1 N HCl = 40 x 16 x 0.1 W (gm) x 1000
1000
= 0.064 g V=Volume of HCl (ml)
0.1 g of sample contains= 0.064 g of Na2CO3 N= Normality of HCl used
100 gm of sample contain = 0.064 x 100 W= Weight of sample (gm)
0.1
Percentage purity= 16 x 0.1 x 40 x 100
= 64.0 % 0.1 x 1000
= 64.0 %

RESULT
The given sample of caustic soda contains 2.12 % of sodium carbonate and 64.0 % of sodium
hydroxide.

Using the equations attached to the experiment and all of your knowledge about reactions
and statistics answer the following questions.

1. A water sample contains OH- and CO32- ions and HCO3- are absent. On titrating 50 ml of the
sample with 0.1 N HCl, the X and Y values were found to be 12 and 15 respectively. Calculate
the amounts of NaOH and Na2CO32- present per litre of the sample.

36
X ≡ NaOH + ½ Na2CO3 = 12.0 ml of 0.1 N HCl
Y ≡ NaOH + ½Na2CO3 + ½Na2CO3 = 15.0 ml of 0.1 N HCl

Na2CO3 ≡ 2 (Y-X) = 2 x (15-12) = 6.0 ml.

NaOH ≡ Y-2(Y-X) = 9.0 ml

Calculation of the amount of NaOH

1000 ml of 1 N HCl ≡ 40 g of NaOH

9.0 ml of 0.1 N HCl = 40 x 9.0 x 0.1
1000
= 0.034 g or 34.0 mg NaOH
Now, 100 ml of the sample contains 34.0 mg of NaOH
1000 ml of the sample contains = 34.0 x 1000 = 340.0 mg of NaOH/litre
100

Calculation of the amount of Na2CO3

1000 ml of 1 N HCl=53 g eq. of Na2CO3 ≡53.0 g of Na2CO3

6.0 ml of 0.1 N HCl = 53 x 6.0 x 0.1
1000
= 0.0318 g or 31.8 mg Na2CO3
Now, 100 ml of the sample contains 31.8 mg Na2CO3
1000 ml of sample contain = 31.8 x 1000
100
= 318 mg of Na2CO3 /litre

37
Flowchart illustrating the determination of NaOH and Na2CO3 in a given sample.

38
 EXPERIMENT 1.7

TO DETERMINE THE PERCENTAGE PURITY OF GIVEN SAMPLE OF BORIC

ACID USING STANDARD 0.1 N NaOH

THEORY

% purity is the percentage of the material which is the actually desired chemical in a sample of it.
In pharmaceutical industry, it would not be acceptable to manufacture a drug with impurities in it
that may be harmful to health. However in any chemical process it is almost impossible to get
100.00 % purity and so sample should be analyzed in industry for % purity to monitor the quality
of the product.
Boric acid is a very weak acid and cannot be titrated accurately with alkali. However, when
dissolved in a solution of glycerin or other organic polyhydroxy compounds such as mannitol),
it acts as a storng acid and can be titrated with standard alkali using phenolphthalein as indicator.
End point is indicated by appearance of pink colour. Boric acid reacts with glycerin to form
glyceroboric acid, a strong monobasic acid. As glycerin is slightly acidic in nature, its solution in
water is first neutralized with dilute alkali to phenolphthalein.

PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 M sodium hydroxide: See experiment 1.3

B. Titration Procedure:

39
S. No Steps Reason
1. 0.1 N NaOH will be prepared and standardized using As per the procedure explained
succinic acid. previously in exp. no.
2. Weigh accurately about 0.2 g of boric acid in triplicate This helps to make sure that
in analytical balance using weighing by difference method results are consistence with each
directly into a conical flask and add 10 ml of water. other. Scientific evidence must
be reproducible.

3. Add 10 ml of glycerine, previously neutralized to Formaldehyde improves the

phenolphthalein/ titration by forming
hexamethylenetetramine.
4. After two minutes, add 2-3 drops of phenolphthalein To get end point quickly and
indicator and titrate solution slowly against 0.1 N NaOH more accurately.
till faint pink colour appears.

5. Note down the burette reading. Repeat the titration with Averaging of the results will
other two flasks and take the average readings of sample minimize the random error.
taken and volume of NaOH used.
Note: Neutralization of glycerine: Glycerine is neutralized by taking in a conical flask 30 ml
formaldehyde solution and two drops of phenolphthalein then titrate against 0.1 N NaOH till
faint pink colour appear.

Observation Table (Specimen reading)

S. Amount of sample taken (g) Burette reading Volume of NaOH
No Initial Final used (ml)
1. 0.20 0.0 25 25.0
2. 0.21 0.0 27 27.0
3. 0.19 0.0 23 23.0
Total 0.305 Total 75.0
Mean of sample taken 0.6/3=0.2 g Mean of volume used 75/3= 25 ml

Calculation for percentage purity:

Descriptive method Short cut method

1000 ml of 1 N NaOH=61.83 g eq. of H3BO3 Percentage purity= V x N x 53.5 x 100
17.6 ml of 0.1 N NaOH= 61.83 x 25 x 0.1 W (gm) x 1000
1000
= 0.154 g V=Volume of NaOH (ml)
0.2 g of sample contains= 0.154 g of H3BO3 N= Normality of NaOH
100 gm of sample contain = 0.154 x 100 W= Weight of sample (gm)
0.2 25 x 0.1 x 61.83 x 100
0.2 x 1000
= 77.3 % = 93.28 %

RESULT
The given sample of ammonium chloride contain 93.28 % of pure ammonium chloride.

40
2. Oxidation-reduction titrations

2.1. To prepare and standardize 0.1 N KMnO 4 solution using oxalic acid as primary
standard.
2.2. To perform the assay of hydrogen peroxides using standard 0.1 N KMnO4 solution.
2.3. To prepare and standardize 0.1 N sodium thiosulphate solution using potassium
iodate as primary standard.
2.4. To perform the assay of copper sulphate using standard 0.1 N sodium thiosulphate
(Iodometry).
2.5. To prepare and standardize 0.1 N ceric ammonium sulphate using arsenic trioxide as
primary standard.
2.6. To perform the assay of ferrous sulphate using standard 0.1 N ceric ammonium
sulphate.

41
 EXPERIMENT NO 2.1

TO PREPARE AND STANDARDIZE 0.1 N POTASSIUM PERMANGANATE USING

OXALIC ACID AS PRIMARY STANDARD

THEORY

It is based on the redox titrations in which strength of oxidizing agent (KMnO 4) is estimated by
titrating with a reducing agent (oxalic acid) and vice-versa. Potassium permanganate acts as an
strong oxidizing agent in acidic medium that oxidizes oxalic acid into manganese sulphate and
carbon-dioxide.
Potassium hydroxide is available in high purity but it cannot be used as a primary standard
because it slowly decomposes into manganese dioxide (colourless) in neutral solution. Moreover,
freshly prepared boiled and cooled water is used because the reducing impurities present in
ordinary water react with KMnO4 to produce MnO2 by reaction…1.
Therefore, standard oxalic acid is used as primary standard to standardize potassium
permanganate in which potassium permanganate act as self indicator. Reaction involved in this
titration is….2.
Sulphuric acid is most suitable acid because it has no action upon permanganate ion. In the
presence of hydrochloric acid, some permanganate ions are consumed in production of chlorine
by the following reaction…3:
Electronically, permanganate ion is converted to manganous ion in the redox process by the gain
of 5 electrons, its equivalent weight is one-fifth of its moleculal weight, 158/5=31.6 ….4.
Oxalic acid which is used as primary standard is oxidized to carbon dioxide by the acidified
KMnO4 by the loss of two electrons. Its equivalent weight is half of its molecular weight,
126/2=63. …..5

In addition to oxalic acid, various other primary standards can be used for standardization of
KMnO4 like arsenic trioxide (As2O3), sodium oxalate etc.

PROCEDURE

42
 A. Preparation of reagents and solutions:

1. Preparation of 0.1 N potassium permanganate:

Calculation for making 100 ml of 0.1 N potassium permanganate solution:

1000 ml of 1 N KMnO4=158/5=31.6 g/eq of KMnO4
100 ml of 0.1 N NaOH= 31.6 x 100 x 0.1
1000
= 0.316 g of KMnO4

Hence, weigh nearly about 0.32 g of KMnO4, dissolved in 100 ml of distilled water in beaker.
Boil the solution for 15 to 30 minutes. Cool, and filter through funnel containing a plug of glass
wool or asbestos. Collect the filtrate and store in dark-brown glass bottle.

2. Preparation of 0.1 M standard oxalic acid solution (Primary standard)

1000 ml of 1 N oxalic acid dihydrate=126/2=63 g eq. of

oxalic acid dihydrate

100 ml of 0.1 N oxalic acid dihydrate = 63 x 100 x 0.1

1000
= 0.63 g of oxalic acid dehydrate

Weigh nearly about 0.63 g of oxalic acid dihydrate and dissolved in 50 ml of distilled water in
100 ml volumetric flask and make up the volume to 100 ml.

B. Standardization of KMnO4 solution using oxalic acid as primary standard

1. The burette is filled with 0.1M potassium permanganate.

2. Pippete out 20 ml of prepared 0.1 N standard oxalic acid solution in an conical flask. Add
5 ml of concentrated sulphuric acid and warm the solution to 50-60 0C.

3. Start titration with the potassium permanganate with constant shaking. The temperature
should be between 50-60 0C throughout the course of titration.

4. End point is reached when a faint pink colour persists for 15 seconds.

5. Repeat the titration at least three times to get concordant reading. Calculate the normality.

Note: Formation of brown colour during the titration is caused by insufficient acid, by
using too high temperature or by the use of dirty flask.

Observation Table (Specimen reading)

S. Volume of succinic acid Burette reading Volume of NaOH
No solution (ml) Initial Final used (ml)
1. 10.0 (Rough titration) 0.0 9.6 10.1

43
2. 10.0 10.0 19.9 9.5
3. 10.0 20.0 29.8 9.6

Average of 2 and 3 cocordant readings= 9.50 + 9.60 = 9.55

2
NNaOH x VNaOH = N(succ. acid.) x V(succ.acid.)

NNaOH = N(succ. acid) x V (succ. acid) = 0.1 x 10 Volume of succinic acid solution in each titration = 10 ml
VNaOH 9.55 Normality of succinic acid = 0.1 M

NNaOH = 0.105 N

RESULT

Concentration of potassium permanganate solution is 0.105 M.

44
 EXPERIMENT NO 2.2

TO PEFORM THE ASSAY OF HYDROGEN PEROXIDE USING STANDARD 0.1 N

POTASSIUM PERMANGANATE

THEORY

Hydrogen peroxide solution (I.P.) is an aqueous solution of hydrogen peroxide containing not
less than 5.0 % w/v and not more than 7.0 % w/v of H 2O2. It is colourless and odourless liquid.
Hydrogen peroxide can act as an oxidizing as well as a reducing agent. If it is titrated with a
stronger oxidizing agent like potassium permanganate in acidic solution, it acts as reducing
agents and oxidized to water and oxygen. This forms the basis of the assay of hydrogen peroxide
solution. A high concentration of sulphuric acid (5N) is used to prevent the conversion of
permanganate ion to manganese dioxide which is an active catalyst for the decomposition of
hydrogen peroxide.

Hence, H2O2 can accept two electrons, forming neutral water and an oxide. So equivalent weight
of H2O2 is M.wt/2 or 68/2=17
So,
1000 ml 1 N of potassium permanganate= 17 g/eq. of H2O2

PROCEDURE

A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 potassium permanganate solution: See
experiment no. 2.1
2. Sulphuric acid solution (5 N): Refer experiment no. 1.2

B. Titration Procedure:
1. The burette is filled with 0.1M potassium permanganate solution.

45
 2. Take 10 ml of H2O2 and dilute to 250 ml with water. Take 25 ml of dilution in
a conical flask.

3. Add 5 ml of 5N sulphuric acid.

4. Start titration with the potassium permanganate until appearance of pink

colour that should persist for 30 seconds signals the actual end point.

5. Repeat the titration three times to get precise readings. Take mean and
calculate the percentage purity of hydrogen peroxide.

Note: The solution may lose its colour after 30 seconds or more, but it is not
considered.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of

No Initial Final KMnO4
used (ml)
1. 25 ml of H2O2 + 5 ml sulphuric 0.0 42 42
acid (5N)
2. As above 0.0 40 40
3. As above 0.0 39 39

Average of 2 and 3 concordant readings= 40 + 39 = 39.5

2

Calculation for percentage purity:

Descriptive method Short cut method

1000 ml 1 N KMnO4= 17 g/eq. of H2O2 %age purity (g/100 ml)= V x N x E.wt x 100 x10 (D.F.)
17.5 ml of 0.1 N KMnO4= 17 x 39.5 x 0.1 x 10 (D.F.) V1 (ml) x 1000
1000
= 0.67 g E.wt= Equivalent weight of H2O2
10 ml of sample contains = 0.67 g of H2O2 V=Volume of potassium permanganate used
100 gm of sample contain = 0.67 x 100 N= Normality of potassium permanganate
1.0 V1= Volume of sample (ml) taken
D.F.= Dilution factor
= 6.7 % 39.5 x 0.1 x 17 x 100x 10
10.0 x 1000
= 6.7 %

RESULT
The percentage purity of given sample of hydrogen peroxide is found to be 6.7 %.

46
 EXPERIMENT NO 2.3

TO PREPARE AND STANDARDIZE 0.1 N SODIUM THIOSULPHATE USING

POTASSIUM IODATE AS PRIMARY STANDARD

THEORY

It is based on the principal of Iodometry in which iodine (oxiding agent) is liberated which can
be titrated with sodium thiosulphate (reducing agent).
Sodium thiosulphate is efflorescent in nature because it comes in pentahydrate form having five
molecules of water of crystallization whose composition changes with temperature and humidity,
i.e. in high humidity, crystals absorb water and at low humidity and high temperature there is
loss of water of crystallization. Therefore, exact composition to get an exact molecular weight
(essential condition for primary standard) for thiosulphate is very difficult to obtain. So it is
standardized against standard potassium iodate in presence of sulphuric acid and excess of
potassium iodide. Potassium iodate oxidizes the iodide ion in acidic medium with liberation of
equivalent amount of iodine (Iodometry). The liberated iodine is titrated directly with sodium
thiosulphate giving sodium tetrathionate ion and sodium iodide. The end point is detected by
starch solution which shows disappearance of permanent blue colour due to the conversion of
iodine into sodium iodide.
Potassium iodate is an oxidizing agent and is used as a source of iodine which liberated in the
presence of KI and sulphuric acid. One molecule of potassium iodate is liberating 3 molecules of
iodine and one molecule of iodine is reacting with 2 molecules of sodium thiosulphate i.e.
equivalent weight of potassium iodate is one-sixth of the molecular weight. Two thiosulphate
(S2O3--) ion in the redox reaction by the loss of two electrons, hence the equivalent weight of
sodium thiosulphate (used as pentahydrate salt, Na2S2O3.5H2O), is equal to its molecular weight,
248.17.

47
The freshly boiled and cooled water should be used for the preparation of solution of solution of
sodium thiosulphate to expel carbon dioxide. The latter is present as carbonic acid (CO 2 + H2O
= H2CO3) will cause a slow decomposition of the thiosulphate leading to the formation of
sulphure. The decomposition of the salt may also be caused by bacterial action which can be
prevented by adding 3 drops of chloroform or 10 mg of mercuric iodide per liter of the solution
and by keeping the pH of solution between 9 and 10.

PROCEDURE

A. Preparation of reagents and solutions:

1. 0.1 N sodium thiosulphate:

Calculation for making 100 ml of 0.1 N sodium thiosulphate solution:

Descriptive method Shortcut method

1000 ml of 1 N Na2S2O3.5H2O =248 g of Na2S2O3.5H2O N= No. of moles or Mass (g) x 1000
100 ml of 0.1 N Na2S2O3.5H2O = 248 x 100 x 0.1 Volume (L) Molar mass x volume (ml)
1000
= 2.48 g of Na2S2O3.5H2O 0.1 = Mass (g) x1000 =2.48 g of Na2S2O3.5H2O
248 x 100

Hence, 2.48 g of Na 2S2O3.5H2O will be weighed and dissolve in 50 ml of CO 2 free distilled water and
volume will be made up to 100 ml to make 0.1 M sodium thiosulpahte solution.

2. Preparation of 0.1 M standard potassium iodate solution

Descriptive method Shortcut method

1000 ml of 1 N KIO3 =214/6=35.6 g eq. of N= No. of moles or Mass (g) x 1000
KIO3 Volume (L) Equivalent mass x volume (ml)

48
100 ml of 0.1 N KIO3= 35.6 x 100 x 0.1
1000 0.1 = Mass (g) x 1000 = Mass=0.356 g of KIO3
= 0.356 g of KIO3 35.6 x 100

Hence, weigh nearly about 0.356 g of KIO3and dissolved in 50 ml of distilled water and make up the
volume to 100 ml.

3. Dilute sulphuric acid: Dilute 57 ml of sulphuric acid to 1000 ml with distilled water.
4. Starch indicator solution: Take 1 g of soluble starch and triturate it with 5 ml of water and
add it to 100 ml of boiling water containing 10 mg of mercuric iodide with continuous
stirring.

B. Standardization of 0.1 M Sodium thiosulphate

1. The burette is filled with 0.1M Na2S2O3 (sodium thiosulphate).

2. Take 10 ml of prepared 0.1 N standard potassium iodate solution in an iodine flask. Add
2 g of potassium iodide followed by 5 ml of dilute sulphuric acid. Stopper the flask and
keep the flask in dark for 5-10 minutes.

3. Start titration with the sodium thiosulphate until colour changed to yellowish-green. Add
3-5 drops of starch indicator solution. The solution becomes blue in colour.

4. Continue the titration until the blue colour disappears and a green colour appears. This is
the end pint.

5. Repeat the titration at least three times to get concordant reading. Calculate the normality.

Observation Table (Specimen reading)

S. Volume of std. potassium Burette reading Volume of Na2S2O3
No iodate (0.1 N) (ml) Initial Final (ml)
1. 10.0 (Rough titration) 0.0 9.6 10.1
2. 10.0 10.0 19.9 9.50
3. 10.0 20 29.8 9.60

Average of 2 and 3 cocordant readings= 9.50 + 9.60 = 9.55

2
NNaOH x VNaOH = N(succ. acid.) x V(succ.acid.)

NNaOH = N(succ. acid) x V (succ. acid) = 0.1 x 10 Volume of succinic acid solution in each titration = 10 ml
VNaOH 9.55 Normality of succinic acid = 0.1 M

NNaOH = 0.105 N

RESULT

Concentration of sodium hydroxide solution is 0.105 M.

49
50
 Flowchart illustrating the preparation and standardization of 0.1 N Sodium thiosulphate
EXPERIMENT NO 2.4

TO PEFORM THE ASSAY OF COPPER SULPHATE USING STANDARD 0.1 N

SODIUM THIOSULPHATE (IODOMETRY)

THEORY

Copper sulphate (CuSO4.5H2O) occurs as blue triclinic prisms or a blue crystalline powder.
Assay of copper sulphate is based on the iodometric titration in which liberated iodine is titrated
with sodium thiosulphate. In the presence of acetic acid, copper sulphate reacts with potassium
iodide and gives cupric iodide (1).
Cupric iodide due to its instability decomposes into cuprous idide and liberate equivalent amount
of free iodine (2).
The liberated free iodine is titrated with known strength of sodium thiosulphate using starch as
indicator (3).
Decomposition of cupric iodide to cuprous iodide is reversible. There is chance that near the end
point when all the free iodine is reacted, cuprous iodide undergoes backward reaction to cupric
iodide giving erroneous results. Therefore, potassium thiocyanate is added towards the end of
reaction to convert cuprous iodide into cuprous thiocyanate which is more sparingly soluble than
cuprous iodide and no back reaction occur (4).

1000 ml 1 N of sodium thiosulphate= 249.7 g/eq. of CuSO4.5H2O

51
 (Equivalent weight of CuSO4.5H2O is same as molecular weight )
PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 sodium thiosulphate solution: See
experiment no. 2.3

B. Titration Procedure:
1. The burette is filled with 0.1M sodium thiosulphate solution.

2. Weigh accurately about 0.1 g of copper sulphate (FeSO4.5H2O) sample and

transfer it to conical flask.

3. Add 50 ml of water, 3 g of potassium iodide and 5 ml of acetic acid. Start

titration with the ceric ammonium sulphate until red colour disappear
suggesting the end point.

4. Add 2 to 3 drops of starch solution as indicator. Blue colour appears.

5. Start titration with the sodium thiosulphate till blue colour changes to
light/faint blue colour.

6. Add 2 g of potassium thiocyanate, stir well, and continue the titration until the
blue colour disappears.

7. Repeat the titration three times to get precise readings. Take mean and
calculate the percentage purity of copper sulphate.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of ceric

No Initial Final sulphate used (ml)
1. 1 g CuSO4.5H2O + 3 g KI + 5 ml acetic 0.0 40 40
acid + 2 g KSCN + 50 ml H 2O + 2 to 3
drops starch indicator
2. As above 0.0 38 38
3. As above 0.0 37 37

Average of 2 and 3 concordant readings= 38 + 37 = 37.5

2
Calculation for percentage purity:
Descriptive method Short cut method

52
1000 ml 1 N ceric sulphate= 249.7 g/eq. of CuSO4.5H2O Percentage purity= V x N x 53.5 x 100
17.5 ml of 0.1 N ceric sulphate =249.7 x 37.5 x 0.1 W (gm) x 1000
1000
= 0.936 g V=Volume of sodium thiosulphate used
1.0 g of sample contains= 0.936 g of CuSO4.5H2O N= Normality of sodium thiosulphate
100 gm of sample contain = 0.936 x 100 W= Weight of sample (gm) taken
1.0 37.5 x 0.1 x 249.7 x 100
1.0 x 1000
= 93.6 % = 93.6 %

RESULT. The percentage purity of given sample of copper sulphate is found to be 93.6 %.
EXPERIMENT NO 2.5

TO PREPARE AND STANDARDIZE 0.1 N CERIC AMMONIUM SULPHATE USING

ARSENIC TRIOXIDE AS PRIMARY STANDARD

THEORY

The experiment is based upon the redox titration. Cerium sulphate or ceric ammonium sulphate
is a powerful oxidizing agent in acid solution. It is bright yellow color and corresponding cerium
salt formed by reduction is colourless. Strong solutions are self indicating. However since dilute
solutions are used hence indicators are used for observation of end point. Ferroin is a suitable
indicator. Arsenic trioxide is used as primary standard in the presence of sulphuric acid and
osmic acid using ferroin sulphate as an indicator for standardization of solution.
Reaction involved in this titration is as follow.

As the reaction of ceric sulphate with arsenic trioxide is very slow at room temperature therefore
osmic acid is used as catalyst. Sodium hydroxide is added to dissolve arsenic trioxide.

PROCEDURE
A. Preparation of reagents and solutions:
1. 0.1 N ceric ammonium sulphate

53
Calculation for making 100 ml of 0.1 N ceric ammonium sulphate
1000 ml of 1 N (NH4)Ce(SO4).2H2O=632.5 g/eq of KMnO4
100 ml of 0.1 N NaOH= 632.5 x 100 x 0.1
1000
= 6.325 g of KMnO4

Hence, weigh accurately about 6.5 g of ceric ammonium sulphate, dissolve it in a mixture of 3 ml
sulphuric acid and 50 ml water with gently heat. Cool and filter the solution, then make up the
volume upto 100 ml with water.

2. Preparation of sodium hydroxide (8.0 % w/v): Weight accurately about 8 g of

sodium hydroxide and transfer it in a 100 ml volumetric flask and make up the
volume to the mark.
3. Dilute sulphuric acid: Dilute 57 ml of sulphuric acid to 1000 ml with distilled
water.
4. Osmic acid solution: Weigh accurately about 1 g of osmic acid and dissolve it in
100 ml distilled water.
5. Ferroin sulphate indicator: Weight accurately about 0.7 g of ferrous sulphate and
1.5 g 1,10-phenanthroline hydrochloride in 70 ml of distilled water and make up the
volume to 100 ml.

B. Standardization of ceric sulphate solution using arsenic trioxide as primary

standard
1. The burette is filled with 0.1M ceric ammonium sulphate.

2. Weigh accurately about 0.2 g of arsenic trioxide (previously dried at 105 oC for 1
hour) and transfer it to 500 ml conical flask.

3. Add 25 ml of sodium hydroxide solution (8.0 % w/v) through inner wall of the flask
and swirl to dissolve. Add 100 ml of distilled water and mix properly.

4. Add 30 ml dilute sulphuric acid, 0.15 ml osmic acid solution and 0.1 ml ferroin
sulphate solution (as indicator) and mix properly.

5. Start titration with the ceric ammonium sulphate until pink colour changes to very
pale blue colour suggesting the end point.

6. Repeat the titration three times to get precise readings. Calculate the normality for
each reading and take the mean of normality.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of ceric
No Initial Final sulphate used (ml)
1. 0.2 g As2O3 + 25 ml NaOH + 30 ml 0.0 42 42

54
 dil. H2SO4 + 0.15 ml osmic acid + 0.1
ml ferroin sulphate
2. 0.190 g As2O3 + 25 ml NaOH + 30 ml 0.0 40 40
dil. H2SO4 + 0.15 ml osmic acid + 0.1
ml ferroin sulphate
3. 0.180 g As2O3 + 25 ml NaOH + 30 ml 0.0 39 39
dil. H2SO4 + 0.15 ml osmic acid + 0.1
ml ferroin sulphate

Normality (N)= No. of moles or Weight (g) of primary standard x 1000

Volume (L) Eq. weight of primary standard x volume (ml) of titrant used

As we know,
1000 ml 1 N of ceric ammonium sulphate= 198/4 g/eq. of arsenic trioxide.
1000 ml 1 N of ceric ammonium sulphate= 49 g/eq. of arsenic trioxide.

N1 = 0.2x1000
49 x 42
N2 = 0.19x1000
49 x 40
N3 = 0.18x1000
49 x 39

N=N1 + N2 + N3
3
= 0.097 + 0.097 + 0.094 = 0.096
3

RESULT
The strength of the prepared ceric ammonium sulphate solution was found out to be 0.096 N

55
 EXPERIMENT NO 2.6

TO PEFORM THE ASSAY OF FERROUS SULPHATE USING STANDARD 0.1 N

CERIC AMMONIUM SULPHATE

THEORY

Ferrous sulphate is mainly used as hematinic and is a reducing agent. It effloresces in dry air. On
exposure to moist air, the crystals rapidly oxidize and become coated with brownish-yellow
ferric sulphate. Assay of ferrous sulphate is based on redox titration (cerimetry). Known strength
(standard) ceric ammonium sulphate is titrated with ferrous sulphate using ferroin solution as an
indicator. Ferrous sulphate is oxidized quantitatively to ferric sulphate when titration is carried
out with ceric ammonium sulphate in acidic media.

In this redox titration, Fe2+ is oxidized to Fe3+ and Ce4+ get reduced to Ce3+
1000 ml 1 N of ceric ammonium sulphate= 278 g/eq. of FeSO4.7H2O
(Equivalent weight of FeSO4.7H2O is same as molecular weight as only one electron is involved
in converting Fe2+ ot Fe3+).

PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 ceric ammonium sulphate: See
experiment no. 2.5

56
 2. Ferroin sulphate indicator: Weight accurately about 0.7 g of ferrous sulphate and
1.5 g 1,10-phenanthroline hydrochloride in 70 ml of distilled water and make up the
volume to 100 ml.
3. Sulphuric acid (1M): See experiment no 1.2

B. Titration Procedure:
1. The burette is filled with 0.1M ceric ammonium sulphate.

2. Weigh accurately about 0.5 g of ferrous sulphate (FeSO4.7H2O) sample and transfer
it to conical flask.

3. Add 30 ml of water and 20 ml of sulphuric acid (1M). Add 0.1 ml ferroin sulphate
solution and mix properly. Red colour appear.

4. Start titration with the ceric ammonium sulphate until red colour disappear
suggesting the end point.

5. Repeat the titration three times to get precise readings. Take mean and calculate the
normality of ceric ammonium sulphate.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of

No Initial Final ceric sulphate
used (ml)
1. 0.5 g FeSO4.7H2O + 20 ml dil. H2SO4 + 30 0.0 20 20
ml H2O + 0.1 ml ferroin sulphate indicator
2. As above 20 38 18
3. As above 20 37 17

Average of 2 and 3 concordant readings= 18 + 17 = 17.5

2

Calculation for percentage purity:

Descriptive method Short cut method

1000 ml 1 N ceric sulphate= 278 g/eq. of FeSO4.7H2O Percentage purity= V x N x 53.5 x 100
17.5 ml of 0.1 N ceric sulphate =278 x 17.5 x 0.1 W (gm) x 1000
1000
= 0.486 g V=Volume of ceric sulphate used
0.5 g of sample contains= 0.486 g of FeSO4.7H2O N= Normality of ceric sulphate used
100 gm of sample contain = 0.486 x 100 W= Weight of sample (gm) taken
0.5 17.5 x 0.1 x 278 x 100
0.5 x 1000
= 97.3 % = 97.3 %

57
RESULT
The percentage purity of given sample of ferrous sulphate is found to be 97.3 %.

3. Argentometric titrations

3.1. To prepare and standardize 0.1 N AgNO3 using sodium chloride as primary standard
(Mohr’s method).
3.2. To determine the percentage purity of given sample of sodium chloride injection
using standard 0.1 N AgNO3 (Mohr’s method).
3.3. To prepare and standardize 0.1 N ammonium thiocyanate solution using standard
AgNO3 as secondary standard.
3.4. To determine the percentage purity of given sample of sodium chloride using
standard 0.1 N AgNO3 (Volhard’s method).

58
 EXPERIMENT NO 3.1

TO PREPARE AND STANDARDIZE 0.1 N SILVER NITRATE USING SODIUM

CHLORIDE AS PRIMARY STANDARD (MOHR’S METHOD)

THEORY

This titration is based on the argentometric titration (type of titration which involves silver ions).
Typically it is most often used for determination of chloride present in a sample but they can be
used also for other halides (bromide, iodide) and some pseudohalides (thiocyanate).
Silver nitrate is titrated directly against sodium chloride and form white precipitate with silver
nitrate.
Silver nitrate reacts with chloride ions in sodium chloride to form a white precipitate of silver
chloride. Once all the chloride ions are consumed the next drop of silver ion reacts with the
indicator potassium chromate to form a brick red colour precipitate at the end point.

AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)

2AgNO3 (aq) + K2CrO4 (aq)  2KNO3 (aq) + Ag2CrO4 (s)
(Potassium chromate ) (Silver chromate)
.PROCEDURE
A. Preparation of reagents and solutions:
1. 0.1 N silver nitrate:

59
Calculation for making 100 ml of 0.1 N silver nitrate solution:
1000 ml of 1 N AgNO3=169.9 g/eq of AgNO3
100 ml of 0.1 N AgNO3= 169.9 x 100 x 0.1
1000
= 1.69 g of AgNO3
Hence, weigh accurately about 1.7 g of AgNO 3, dissolve in 70 ml of distilled water in beaker and
make up the volume to 100 ml.

2. Preparation of 0.1 M standard sodium chloride solution

1000 ml of 1 N sodium chloride =58.4 g eq. of sodium chloride

100 ml of 0.1 N sodium chloride = 58.4 x 100 x 0.1
1000
= 0.584 g of sodium chloride

Weigh accurately about 0.584 g of sodium chloride (dried at 300 0C for 2h and cool in a
desiccator) and dissolved in 50 ml of distilled water in 100 ml volumetric flask and make up the
volume to 100 ml.
3. Preparation of 5 % w/v potassium chromate solution (As indicator): Weigh
accurately about 5 g of potassium chromate and dissolve in 100 ml of distilled water.

B. Standardization of AgNO3 solution using sodium chloride as primary standard.

1. The burette is filled with 0.1M silver nitrate solution.

2. Pippete out 10 ml of prepared 0.1 N standard sodium chloride solution in an conical flask.

3. Add 1 ml of potassium chromate solution as an indicator.

4. Start titration with the silver nitrate solution with constant shaking. So that, coagulation of
precipitated silver chloride occurs just before the end-point. Easy to detect colour change.

5. End point is suggested by appearance of permanent brick red colour.

6. Repeat the titration at least three times to get concordant reading. Calculate the normality.

Sr. Steps Reason

No.
1. The burette is filled with 0.1M silver nitrate Silver nitrate is acting as titrant.
solution.
2. Pippete out 10 ml of prepared 0.1 N standard Sodium chloride acts as a primary standard.
sodium chloride solution in an conical flask.
3. Add 1 ml of potassium chromate solution as It is the best indicator for titration by Mohr’s
an indicator. method.
4. Start titration with the silver nitrate solution So that, coagulation of precipitated silver
with constant shaking. chloride occurs just before the end-point. Easy
to detect colour change.
5. End point is suggested by appearance of Due to the formation of silver chromate.
permanent brick red colour.

60
 6. Repeat the titration three times to get precise Averaging the results will minimize the
readings. Take concordant reading or mean random error.
and calculate the normality of AgNO 3.

Observation Table (Specimen reading)

S. Volume of silver chloride Burette reading Volume of AgNO3
No (ml) Initial Final used (ml)
1. 10.0 (Rough titration) 0.0 9.6 9.6
2. 10.0 10.0 19.5 9.50
3. 10.0 20 29.6 9.60

Average of 2 and 3 cocordant readings= 9.50 + 9.60 = 9.55

2
NAgNO3 x VAgNO3 = N(NaCl.) x V(NaCl)

NAgNO = N(NaCl) x V (NaCl) = 0.1 x 10 Volume of sodium chloride soln. in each titration = 10 ml
VAgNO3 9.55 Normality of sodium chloride = 0.1 M

NAgNO3 = 0.105 N

RESULT
Concentration of silver nitrate solution is 0.105 M.
EXPERIMENT NO 3.2

TO PERFORM THE ASSAY OF SODIUM CHLORIDE USING STANDARD SILVER

NITRATE SOLUTION (MOHR’S METHOD)

THEORY

Sodium chloride is an electrolyte replenisher. The principle involved in the assay of sodium
chloride is Argentimetric titration by Mohr’s method named after Karl Friedrich Mohr. Silver
nitrate reacts with chloride ions in sodium chloride to form a white precipitate of silver chloride.
Once all the chloride ions are consumed the next drop of silver ion reacts with the indicator
potassium chromate to form a brick red colour precipitate of silver chromate at the end point.
2Ag+ (aq) + CrO2−4 (aq) → Ag2CrO4 (s) (Ksp = 1.1 × 10−12)
The solution needs to be near neutral, because silver chromate forms at high pH, while the
chromate forms H2CrO4 at low pH, reducing the concentration of chromate ions, and delaying
the formation of the precipitate. Carbonates and phosphates precipitate with silver, and need to
be absent to prevent inaccurate results.
AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)

61
 2AgNO3 (aq) + K2CrO4 (aq)  2KNO3 (aq) + Ag2CrO4 (s)
(Potassium chromate ) (Silver chromate)

PROCEDURE

A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 N silver nitrate: Described in
experiment no. 3.1
2. Preparation of 5 % w/v potassium dichromate solution (as indicator): Weigh
accurately about 5 g of potassium chromate and dissolve in 100 ml of distilled water.

B. Titration steps
1. The burette is filled with 0.1M silver nitrate solution.

2. Weigh accurately about 0.1 g of sodium chloride sample and transfer it to conical flask.

3. Add 3 to 4 drops of potassium chromate solution as an indicator.

4. Start titration with the silver nitrate solution with constant shaking.

5. End point is suggested by appearance of brick red colour.

6. Repeat the titration at least three times to get concordant reading. Calculate the
percentage purity.

Sr. No. Steps Reason

1. The burette is filled with 0.1M silver Silver nitrate acts as titrant
nitrate solution.
2. Weigh accurately about 0.1 g of sodium It acts as primary standard.
chloride sample and transfer it to conical
flask.
3. Add 3 to 4 drops of potassium chromate It is the best indicator for titration by
solution as an indicator. Mohr’s method.
4. Start titration with the silver nitrate So that, coagulation of precipitated silver
solution with constant shaking. chloride occurs just before the end-point.
Easy to detect colour change.
5. End point is suggested by appearance of Due to the formation of silver chromate.
permanent brick red colour.
6. Repeat the titration three times to get Averaging the results will minimize the
precise readings. Take the mean and random error.
calculate the percentage purity

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of ceric

62
 No Initial Final sulphate used (ml)
1. 0.2 g sodium chloride + 10 ml water 0.0 36 36
+ 2-3 drops of potassium chromate
2. As above 0.0 34 34
3. As above 0.0 33 33

Average of 2 and 3 concordant readings= 34 + 33 = 33.5

2

Calculation for percentage purity:

Descriptive method Short cut method
1000 ml 1 N silver nitrate= 58.4 g/eq. of sodium chloride Percentage purity= V x N x 58.4 x 100
17.5 ml of 0.1 N silver nitrate = 58.4 x 33.5 x 0.1 W (gm) x 1000
1000
= 0.195 g of sodium chloride V=Volume of silver nitrate used
0.2 g of sample contains= 0.195 g of sodium chloride N= Normality of silver nitrate
100 gm of sample contain = 0.195 x 100 W= Weight of sample (gm) taken
10.2 33.5 x 0.1 x 58.4 x 100
0.2 x 1000
= 97.8 % of sodium chloride = 97.8 % of NaCl

RESULT
The percentage purity of given sample of sodium chloride is found to be 97.8 %.

EXPERIMENT NO 3.3

TO PREPARE AND STANDARDIZE 100 ML OF 0.1 N AMMONIUM THIOCYANATE

USING STANDARD 0.1 N SILVER NITRATE AS SECONDARY STANDARD

THEORY

Ammonium thiocyanate reacts with silver nitrate in nitric acid solution. The thiocyanate solution
is always used in the burette and is run into the silver nitrate which has been acidified with nitric
acid (nitrous acid-free, because nitrous acid gives a red colour with thiocyanate ion). The end
point is detected by use of ferric ammonium sulphate as indicator since ferric ions give a deep
red colour (ferric thiocyanate) with a trace of thiocyanate ion. The temperature of the solution
must be kept below 25 °C because at higher temperature the colour of the ferric thiocyanate
complex fades.
AgNO3 + NH4SCN AgSCN + NH4NO3
NH4SCN + Fe(NH4)(SO4)2 Fe(SCN)3 + 2(NH4)2SO4
Ammonium thiocyanate is a deliquescent substance and a slight amount in excess of the
theoretical quantity is required for preparation of its solution.

PROCEDURE

63
 A. Preparation of reagents and solutions:
1. 100 ml of standardized 0.1 N silver nitrate: As described in experiment no. 3.1

2. Preparation of 0.1 N ammonium thiocyanate solution

1000 ml of 1 N NH4SCN =76.12 g eq. of NH4SCN

100 ml of 0.1 N NH4SCN = 76.12 x 100 x 0.1

1000
= 0.76 g of NH4SCN

Weigh accurately about 0.8 g of NH4SCN (dried at 300 0C for 2h and cool in a desiccator) and
dissolved in 50 ml of distilled water in 100 ml volumetric flask and make up the volume to 100
ml. Ammonium thiocyanate is a deliquescent substance and a slight amount in excess of the
theoretical quantity is required for preparation of its solution.

3. Preparation of 10 % w/v ferric ammonium sulphate or ammonium iron (II)

sulphate: Dissolve 10 g of ferric ammonium sulpahte in a 50 ml of water and make
up the volume to 100 ml.

B. Standardization of NH4SCN using standard silver nitrate as secondary standard

1. The burette is filled with 0.1M ammonium thiocyanate solution.

2. Pipette out 10 ml of prepared 0.1 N standard silver nitrate solution into a conical flask.

3. Add 3 ml of concentrated nitric acid and 2 ml of 10 % w/v ferric ammonium sulphate as

an indicator. Dilute with about 25 ml of water. The titration must be done in acidic
medium to prevent the precipitate of Iron(III) as hydrated oxide (iron hydroxide). Iron
(III) is an indicator (this sharpen the end point).

4. Start titration with ammonium thiocyanate solution with constant shaking. A white
precipitate of silver thiocyanate is formed which disappears upon shaking.

5. End point is suggested by appearance of permanent red brown colour which doesnot
disappear upon shaking..

6. Repeat the titration at least three times to get concordant reading. Calculate the normality.

Sr. No. Steps Reason

1. The burette is filled with 0.1M ammonium
thiocyanate solution.
2. Pippete out 10 ml of prepared 0.1 N standard It acts as a secondary standard because its
silver nitrate solution into a conical flask. strength has already been determined in
previous experiment using primary
standard (Sodium chloride).
3. Add 3 ml of concentrated nitric acid and 2 ml of Nitric acid prevents the precipitation of
10 % w/v ferric (II) ammonium sulphate as an Iron (III) as hydrated oxide (iron

64
 indicator. Dilute with about 25 ml of water. hydroxide). Iron (II) is an indicator.
Hence, this overall sharpens the end point).
4. Start titration with ammonium thiocyanate A white precipitate is due to formation of
solution with constant shaking. A white silver thiocyanate.
precipitate is formed which disappears upon
shaking.
5. End point is suggested by appearance of Red brown colour precipitate is of ferric
permanent red brown colour which doesnot thiocyanate.
disappear upon shaking.
6 Repeat the titration three times to get precise Averaging the results will minimize the
readings. Take concordant reading or mean and random error.
calculate the normality of NH4SCN.

OBSERVATION TABLE (SPECIMEN READING)

S. Volume of NAgNO3 (ml) Burette reading Volume of AgNO3
No Initial Final used (ml)
1. 10.0 (Rough titration) 0.0 9.6 10.1
2. 10.0 10.0 19.9 9.50
3. 10.0 20 29.8 9.60

Average of 2 and 3 cocordant readings= 9.50 + 9.60 = 9.55

2
NNH4SCN x VNH4SCN = NAgNO3 x VAgNO3

NNH4SCN = N(AgNO3) x V (AgNO3) = 0.1 x 10 Volume of sodium chloride soln. in each titration = 10 ml
VNH4SCN 9.55 Normality of sodium chloride = 0.1 M

NNH4SCN = 0.105 N
RESULT. Concentration of ammonium thiocyanate solution is 0.105 M.
EXPERIMENT NO 3.4

TO PERFORM THE ASSAY OF SODIUM CHLORIDE USING STANDARD

AMMONIUM THIOCYANATE BY VOLHARD METHOD

THEORY

It is not always possible to use Mohr method to determine concentration of chlorides. For
example, Mohr method requires neutral solution, but in many cases solution has to be acidic, to
prevent precipitation of metal hydroxides (like in the presence of Fe 3+). In such cases, we can use
Volhard method, which is not sensitive to low pH.
In the Volhard method chlorides are first precipitated with excess silver nitrate, then excess
silver is back titrated with ammonium thiocyanate. To detect end point, we use Fe 3+ cations,
which easily react with the thiocyanate, creating distinct wine red complex.

65
There is a problem though. Silver thiocyanate solutility is slightly lower than solubility of silver
chloride, and during titration thiocyanate can replace chlorides in the existing precipitate:
AgCl(s) + SCN- → AgSCN(s) + Cl-
To avoid problems, we can filtrate precipitated AgCl before titration. However, there exists
much simpler and easier procedure that gives the same result. Before titration, we add some
small volume of a heavy organic liquid that is not miscible with water (like nitrobenzene,
chloroform or carbon tetrachloride). These liquids are better at wetting precipitate than water.
Once the precipitate is covered with non polar liquid, it is separated from the water and unable to
dissolve.
Precipitate solubility is not a problem during determination of I- and Br-, as both AgBr and AgI
have much lower solubilities than AgSCN.

PROCEDURE

A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 N silver nitrate: Described in
experiment no. 3.1

66
 2. Preparation and standardization of 0.1 N ammonium thiocyanate: Described in
experiment no. 3.3
3. Preparation of 10 % w/v ferric ammonium sulphate or ammonium iron (II)
sulphate: Dissolve 10 g of ferric ammonium sulpahte in a 50 ml of water and make
up the volume to 100 ml.

B. Steps
1. The burette is filled with 0.1M ammonium thiocyanate solution.

2. Weigh accurately equivalent to about 0.25 g of sodium chloride sample and transfer it to
conical flask. Add 50 ml of water.

3. Add 50 ml of 0.1 N AgNO3, 3 ml of nitric acid and 5 ml of nitrobenzene.

4. Add 2 ml of ferric ammonium sulphate solution as indicator and shake well.

5. Start titration with the 0.1 N NH4SCN solution with continuous shaking until the colour
becomes reddish yellow which should persist for 5 minutes.

6. Repeat the titration at least three times to get concordant reading. Calculate the
percentage purity.

Sr. No. Steps Reason

1. The burette is filled with 0.1M Ammonium thiocyanate acts as titrant.
ammonium thiocyanate solution.
2. Weigh accurately equivalent to about ---
0.25 g of sodium chloride sample and
transfer it to conical flask. Add 50 ml of
water.
3. Add 50 ml of 0.1 N AgNO3, 3 ml of Nitric acid prevents the precipitation of Fe
nitric acid and 5 ml of nitrobenzene. (III) as hydrated oxide.
Nitrobenzene being immiscible, mask Ag+
of AgCl (form film), otherwise, Ag+ will
react with thiocyanate to give AgSCN and
released Cl- which result in
overconsumption of thiocyanate.
4. Add 2 ml of ferric ammonium sulphate It is the best indicator for Volhard method
solution as indicator and shake well. as it is stable in acidic solution.
5. Start back titration with the 0.1 N ----
NH4SCN solution with continuous
shaking until the colour becomes
reddish yellow which should persist for
5 minutes.
6. Repeat the titration three times to get Averaging the results will minimize the
precise readings. Take the mean and random error.
calculate the percentage purity

67
 OBSERVATION TABLE (SPECIMEN READING)

S. Content in flask Burette reading Volume of

No Initial Final NH4SCN (ml)
1. 0.0 11 11
0.25 g sodium chloride sample +
50 ml of 0.1 N AgNO3 + 3 ml of
nitric acid +5 ml of nitrobenzene
+ 2 ml of ferric ammonium
sulphate solution
2. As above 0.0 9 9
3. As above 0.0 9 9

Volume of NH4SCN used= 9.0 ml (Concordant reading)

Va= V of AgNO3 added = 50

Vb= V of NH4SCN used= 9 ml
V=Volume of actual titrant reacted with sodium chloride (Va-Vb) = 50-9= 41 ml

CALCULATION FOR PERCENTAGE PURITY:

Descriptive method Short cut method

1000 ml 1 N AgNO3= 58.4 g/eq. of sodium chloride Percentage purity= V (Va-Vb) x N x Eq. wt. x 100
W (gm) x 1000
41 ml of 0.1 N AgNO3=58.4 x 41 x 0.1
1000 V=Volume of AgNO3 used
= 0.24 g of sodium chloride Va= V of AgNO3 added = 50
0.25 g of sample contains= 0.24 g of sodium chloride Vb= V of NH4SCN used= 9 ml
100 gm of sample contain = 0.24 x 100 V=Volume of actual titrant reacted with sodium
0.25 chloride (Va-Vb) = 50-9= 41 ml
Eq. wt.= Equivalent weight of sodium chloride
= 95.8 % N= Normality of AgNO3
W= Weight of sample (gm) taken
41 x 0.1 x 58.4 x 100
0.25 x 1000
= 95.8 %

RESULT
The percentage purity of given sample of sodium chloride is found to be 95.8 %.

68
4. Complexometric titration

4.1. To prepare and standardize 0.1 N EDTA using granulated zinc as primary standard.
4.2. To determine the percentage purity of given sample of magnesium sulphate using
standard 0.1 N EDTA (Direct titration).
4.3. To perform the assay of calcium gluconate using standard 0.1 N EDTA
(Replacement titration).

69
 EXPERIMENT NO 4.1

TO PREPARE 100 ML OF 0.05 N EDTA SOLUTION AND STANDARDIZE IT USING

GRANULATED ZINC AS PRIMARY STANDARD

THEORY

EDTA is ethylene diamine tetraacetic acid is commonly used complexing agent. It is not very
water soluble so disodium edetate salt is used. A complexing agent is an electron-donating ion
(Lewis base) which is usually called a ligand. Metal ion is electron deficient species (Lewis
acid).
EDTA has the ability to “wrap” itself around positive metal ions in water solution. This process
is called chelation or complex formation. The chelation reaction between EDTA and many metal
ions has a very large equilibrium constant. The reaction is always 1 mol of EDTA to 1 mol of
metal ion.
Since EDTA is an acid substance with four weak acid dissociations, the reactions with metal ions
are pH dependent. The metal ions that react most strongly with EDTA can be titrated in acidic
solution. Zinc is an example of a metal ion that is titrated in acidic solution. The metals that react
more weakly with EDTA must be titrated in alkaline solution.
For best results it is good to standardize EDTA solution against the same cation and using the
same method as will be later used during sample analysis. Note, that EDTA solution can be
prepared without a need for standardization, as EDTA itself can be obtained in form pure
enough. Titration is done in pH 10 solution.
The standardization of EDTA is based on the titration of the disodium edetate solution with a
standard zinc chloride solution prepared from a known weight of granulated zinc. Since EDTA is
insoluble in water, the disodium salt of EDTA is used for this experiment.

i.e. 1000 ml of 1N EDTA= Eq./wt of Zn

1000 ml of 1N EDTA= 65.38 g of Zn

70
PROCEDURE
A. Preparation of reagents and solutions:
1. 0.05 N EDTA

Calculation for making 100 ml of 0.05 N EDTA solution:

1000 ml of 1 N Disodium edentate dehydrate = 372 g/eq of Disodium edetate
100 ml of 0.05 N Disodium edentate dihydrate = 378 x 100 x 0.05
1000
= 1.89 g of Disodium edetate
Hence, weigh accurately about 1.9 g of disodium edentate. Dissolve in 50 ml
water then make up the volume upto 100 ml with water.

2. 0.1 N standard granulated zinc solution

Calculation for making 250 ml of 0.05 N granulated zinc solution
1000 ml of 1 N zinc =65.38 g of granulated zinc
100 ml of 0.05 N zinc = 65.38 x 250 x 0.05
1000
= 0.81 g of zinc

3. Dilute HCl (10 % w/v of HCl): 12 ml prepare by diluting 10 ml conc. HCl to 50 ml

of water and volume will be made up to 100 ml.
4. Bromine water: 3 ml bromine in 100 ml water.
5. Sodium hydroxide (2N): 8 g NaOH in 100 ml water
6. Preparation of ammonia (5 M): Dissolve 38.5 g of strong ammonia solution to a
volumetric flask and make upto 1000 ml.
7. Preparation of ammonia buffer (pH 10 M): Dissolve 5.4 g of ammonium chloride
in 70 ml of 5 N ammonia and dilute with water to 100 ml.

B. Standardization of EDTA using zinc metal as primary standard:

S. Steps Reason
No
1. 0.05 N EDTA will be prepared and fill in the As per the calculation explained previously.
burette.
2. Weigh accurately about 0.8 g of granulated 0.8 g obtained as per the calculation explain
zinc. above.

3. Dissolve by gentle heating in the minimum

volume (12 ml) of hydrochloride acid (10% As granulated zinc is not soluble in simple
w/v). water. It is converted into zinc chloride by
reacting with HCl which is easily soluble.
4. Add 0.2 ml or 5 drops of bromine water
To ensure oxidation of trace iron impurity to
iron (III) which forms a much less stable
edetate complex than iron (II).
5.
Gently boil to remove excess bromine. Cool, THE NORMALITY OF THIS ZINC

71
 add DW to produce 250 ml. SOLUTION BECAME 0.05 N.
6.
Pipette 20 ml of the resulting solution into a Acidic pH decreases the stability of the
flask and neutralize with sodium hydroxide complex.
(2N) until almost neutral (small amount of
white precipitate is formed).
7.
Dilute to about 150 ml with DW and add Alkaline pH condition make stable complex,
ammonia buffer pH 10 until the precipitate is prevent backward reaction and hence sharp
just dissolved and then 5 ml in excess. end-point.
8. Mordant black II is the best indicator at this pH
Add a mixture of mordant black II and range.
sodium chloride (1:99; 50 mg) as indicator. This give the volume of HCl used for the
complete neutralization of Na2CO3.
9. This give the volume of EDTA used for the
Titrate with 0.05 N EDTA solution until the complete neutralization of Zn.
solution become green.
10. Averaging the results will minimize the
Repeat the titration three times to get precise random error.
readings. Take mean and calculate the
normality of EDTA.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of EDTA
No Initial Final (ml)
1. 0.8 g of granulated zinc + HCl 0.0 25.5 25.5
acid+ Bromine water+ ammonia
buffer + distilled water
2. As above 26 50 24
3. As above 0.0 23 23

Average of 2 and 3 concordant readings= 24 + 23 = 23.5

2

NEDTA x VEDTA = NZinc x VZinc

NEDTA = N(zinc) x V ( zinc)= 0.05 x 25 Volume of zinc solution in each titration = 25 ml

VEDTA 23.5 Normality of zinc acid = 0.05 N

NEDTA = 0.051 N

Normality (N)= No. of moles or Weight (g) of primary standard x 1000

Volume (L) Eq. weight of primary standard x volume (ml)

Normality of zinc N = 0.8x1000 = 0.05 N

72
 65.38 x 250

RESULT

The strength of the prepared EDTA solution was found out to be 0.103 N

EXPERIMENT NO 4.2

TO DETERMINE THE PERCENTAGE PURITY OF GIVEN SAMPLE OF

MAGNESIUM SULPHATE USING STANDARD 0.1 N EDTA (DIRECT
TITRATION)

THEORY

Magnesium sulphate is assayed by titration with 0.05 N disodium edetate using mordant black II
mixture as indicator. The pH of the solution is kept 10 by the addition of ammonia buffer
solution. The colour change at the end-point is from pink to blue. The reaction can be
represented by

PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.05 disodium edetate solution: See
experiment no. 4.1
2. Ammonia buffer pH 10: See experiment no. 4.1

B. Titration Procedure:
1. The burette is filled with 0.05M standard disodium edetate solution.

73
 2. Weigh accurately about 0.3 g of magnesium sulphate sample and transfer it to
conical flask.

3. Dissolve in 50 ml of water. Add 10 ml ammonia buffer pH 10.

4. Add 0.5 g of mordant black II mixture as indicator. Pink colour appears.

5. Start titration with the 0.05 N disodium edetate solution until the solution
becomes full blue.

6. Repeat the titration three times to get precise readings. Take mean and
calculate the percentage purity of magnesium sulphate.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of ceric

No Initial Final sulphate used (ml)
1. 0.3 g MgSO4 sample + 10 ml 0.0 40 40
water+ 10 ml ammonia buffer pH
10 + 0.5 g mordant black II
indicator
2. As above 0.0 38 38
3. As above 0.0 37 37

Average of 2 and 3 concordant readings= 38 + 37 = 37.5

2

Calculation for percentage purity:

Descriptive method Short cut method

1000 ml 1 N disodium edetate= 120.4 g/eq. of MgSO4 Percentage purity= V x N x 120.4 x 100
37.5 ml of 0.05 N Dis. Edetate =120.4 x 37.5 x 0.05 g of MgSO4 W (gm) x 1000
1000
= 0.225 g V=Volume of disodium edetate used
0.3 g of sample contains= 0.225 g of MgSO4 N= Normality of disodium edetate prepared
100 gm of sample contain = 0.225 x 100 W= Weight of sample (gm) taken
0.3 37.5x0.05 x 120.4 x 100
= 75.25 % 0.3 x 1000
= 75.25 %

RESULT

74
The percentage purity of given sample of magnesium sulphate is found to be 75.25 %.

EXPERIMENT NO 4.3

TO PEFORM THE ASSAY OF CALCIUM GLUCONATE USING STANDARD 0.1 N

EDTA (REPLACEMENT TITRATION)

EDTA is a stronger complexing agent than the indicator, and displaces the indicator from the
metal ion allowing the indicator to return (through shades of violet) to a pure blue color,
indicating the end of the reaction.
Calcium ion (Ca+2) does not form a stable red complex with the EBT indicator; therefore the
direct titration of Ca+2 by EDTA may not cause a sharp color change of EBT indicator at the end
point. The magnesium complex with EBT is stable and thus, a displacement titration of Ca +2 by
the mixture of Mg+2 and EDTA will help to determine the end point with the following
mechanism.
To accomplish this displacement titration, a small amount of Mg+2 will be mixed with the EDTA
solution.
This is based on the replacement complexometric titration in which Mg +2 ions are replaced from
its [Mg-Ind] complex. Mordant black II indiator donot give distinct colour change at the end
point with calcium gluconate. Therefore a known volume of 0.05 N MgSO4 is added to calcium
gluconate. Magnesium forms complex with mordant black II indicator which shows first colour
of pink (1).

75
Magnesium-indicator complex is much more stable than calcium-indicator complex therefore
calcium has not any effect on magnesium-indicator complex (1). On titration against disodium
edetate complex of Ca-EDTA is formed (2).
When calcium is totally consumed, next drop of disodium edetate breaks the [Mg-Ind] complex
and make complex with magnesium by liberating free indicator. End point is detected by
observing second colour of blue at that time (3).

76
PROCEDURE
This experiment is divided into two parts:
1. Preparation and standardization of EDTA solution (0.05 N): As discussed in Exp No.
4.1
2. Performing assay of calcium gluconate

A. Preparation of reagents and solutions:

1. Preparation of 0.05 N EDTA solution:
2. Preparation of 0.05 MgSO4:
Calculation for making 100 ml of 0.05 N MgSO4 solution.
1000 ml of 1 N Magnesium sulphate=120 g eq. of magnesium sulphate
100 ml of 0.05 N Magnesium sulphate = 120 x 100 x 0.05
1000
= 0.6 g of Magnesium sulphate

Hence, weigh accurately about 0.6 g of magnesium sulphate. Dissolve in 50 ml

water then make up the volume upto 100 ml with water.

3. Ammonia buffer pH 10: Dissolve 5.4 g of ammonium chloride in 70 ml of 5N

ammonia and dilute with water to 100 ml.

B. Titration Procedure:

Sr. Steps Reason

No.
1. The burette is filled with 0.05 N std. Prepared as directed.
EDTA solution
2. Weigh accurately about 0.5 g of calcium Calcium gluconate is completely solubilize
gluconate sample in 50 ml warm water in warm water.
and allow to cool.
3. Add 5 ml of 0.05 N magnesium sulphate. As Ca+2 doesnot give sharp end point with
mordant black. Mg-In complex will form and
give pink colour.
4. Add 10 ml of ammonia buffer pH 10. Basic pH is required to enhance the stability
of complex by preventing the backward
reaction.
5. Add small amount of mordant black II Pink colour is form by Mg-In complex.
mixture as an indicator. Pink colour
appears.
6. Start titration with the std. EDTA until When all Ca+2 consumes, the extra EDTA
blue colour appear suggesting the end drop will make complex with Mg+2 to form
point. Mg-EDTA and free the indicator which gives
blue colour.
7. Repeat the titration three times to get Averaging the results will minimize the

77
 precise readings. Take mean and random error.
calculate the percentage purity of
calcium gluconate.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of ceric

No Initial Final sulphate used (ml)
1. 0.5 g calcium gluconate + 50 ml 0.0 22 20
water + 5 ml MgSO4 + 10 ml
ammonia buffer pH 10
2. As above 22 42 21
3. As above 0.0 19 20

Average of 2 and 3 concordant readings= 20 + 21 + 20 = 20

3
Volume of EDTA used= 20 ml
Va= V of EDTA used with Ca gluconate = 20 ml
Vb= V of EDTA used without Ca gluconate i.e volume of MgSO4 used.= 5.0 ml
V=Volume of EDTA used (Va-Vb) = 20-5= 15 ml

Calculation for percentage purity:

Descriptive method Short cut method

1000 ml 1 N EDTA= 448 g/eq. of Cal. Glu. Monohyd Percentage purity= V (VA-VB) x N x Eq. wt. x 100
(C12H22CaO14· H2O) W (gm) x 1000

15.0 ml of 0.05 N EDTA =448 x 15 x 0.05 g of CG. Monohyd V=Volume of EDTA used
1000 Va= V of EDTA used with Ca gluconate
= 0.336 g Vb= V of EDTA used without Ca gluconate i.e
0.5 g of sample contains= 0.336 g of Cal. Glu. Monohyd. volume of MgSO4 used.
100 gm of sample contain = 0.336 x 100 Eq. wt.= Equiv weight of Cal Gluconate
0.5 N= Normality of EDTA
W= Weight of sample (gm) taken
= 67.2 % 15 x 0.05 x 448 x 100
0.5 x 1000
= 67.2 %

RESULT
The percentage purity of given sample of calcium gluconate is found to be 67.2 %.

78
5. Non-aqueous titration

5.1. To prepare and standardize 0.1 N perchloric (HClO4) acid using potassium hydrogen
phthalate as primary standard.
5.2. To determine the percentage purity (assay) of ephedrine using standard 0.1 N HClO4.
5.3. To determine the percentage purity (assay) of ephedrine hydrochloride using
standard 0.1 N HClO4.
5.4. To determine the percentage purity of sodium benzoate using standard 0.1 N HClO4.
5.5. To prepare and standardize 0.1 N sodium methoxide solution using benzoic acid as
primary standard.
5.6. To perform the assay of allopurinol using standard 0.1 N sodium methoxide.

79
 EXPERIMENT NO 5.1

TO PREPARE AND STANDARDIZE 0.1 N PERCHLORIC ACID USING POTASSIUM

HYDROGEN PHTHALATE AS PRIMARY STANDARD

THEORY

This titration is based on the non-aqueous method of titration in which no water is involved.
Perchloric acid is usually available as a 70 to 72% mixture with water (sp. gr. 1.6). It usually
undergoes a spontaneous explosive decomposition and, therefore, it is available always in the
form of a solution. Perchloric acid is not only a powerful oxidising agent but also a strong acid.
Hence, it must be handled very carefully.
In usual practice, potassium hydrogen phthalate (or potassium biphthalate, KHC 8H4O4) is
employed as a standardizing agent for acetous perchloric acid. KHP is available in puriest form
99.9%. It is stable in normal storage conditions and therefore used as primary standard for
perchloric acid. Actually KHP nature is acidic (pH:4.5). So we use KHP for the standardization
of NaOH but in HClO4, we standardize in presence of glacial acetic acid. Acetic acid & KHP
combined to form strong conjugate base (CH 3COO-K+). That strong conjugate base reacts with
strong conjugate acid i.e. onium ion formed by HClO 4 in glacial acetic acid during
standardization of perchloric acid. End point is detected by crystal violet solution as an indicator.
The reaction may be expressed as follows:

80
.

PROCEDURE
A. Preparation of reagents and solutions:
1. 0.1 N Perchloric acid:

Calculation for making 100 ml of 0.1 N perchloric acid solution:

1000 ml of 1 N HClO4=100 g/eq of HClO4
100 ml of 0.1 N HClO4= 100 x 100 x 0.1
1000
= 1.0 g of HClO4
Density= Mass Since the purity of HClO4is 70 %
Volume Therefore, 70 % = 0.59 ml
Volume=Mass 1.0 g = 0.59 ml of HClO4 100 % = 0.59 x 100 = 0.84 ml
Density 1.7 g/ml 70

Hence, measure accurately about 0.84 ml of HClO4, dissolved in 70 ml of glacial acetic acid. Add
3 ml of acetic anhydride, cool and make up the volume to 100 ml with glacial acetic acid.

Note: a. The acetic anhydride reacts with the water (approx. 30%) in perchloric acid and
some traces in glacial acetic acid thereby making the resulting mixture practically
anhydrous. Thus, we have :

b. Conversion of acetic anhydride to acetic acid requires 40-45 minutes for its
completion. It being an exothermic reaction, the solution must be allowed to
cool to room temperature before adding glacial acetic acid to volume.

81
 c. Perchloric acid is not only a powerful oxidising agent but also a strong acid.
Hence, it must be handled very carefully.

2. Preparation of crytal violet indicator solution: Dissolve 500 mg of crystal violet

in 100 ml of anhydrous glacial acetic acid.

B. Standardization of HClO4 solution using KHP as primary standard

1. The burette is filled with 0.1M perchloric acid solution.

2. Weigh accurately about 0.35 g of potassium hydrogen phthalate powder in a 100 ml

conical flask.

3. Add 25 ml of anhydrous glacial acetic acid and 2-3 drops of crystal violet indicator.

4. Start titration with the perchloric acid solution with constant shaking.

5. End point is suggested by the change of violet colour to emerald green

6. Repeat the titration three times to get precise readings.

7. Take one reading of blank determination i.e. do titration according to above procedure
without potassium hydrogen phthalate.

8. Calculate the normality for each reading (after subtracting blank reading) and take the
mean of normality.

S. Steps Reason
No.
1. The burette is filled with 0.1M 0.1 N is prepared as mentioned above.
perchloric acid solution.
2. KHP is available in puriest form 99.9%. It is
Weigh accurately about 0.35 g of stable in normal storage conditions and therefore
potassium hydrogen phthalate (KHP) used as primary standard for perhloric acid.
in a 100 ml conical flask.
3.
Add 25 ml of anhydrous glacial acetic Glacial acetic acid increases the basicity of KHP.
acid and 2-3 drops of crystal violet
indicator.
4.
Start titration with the perchloric acid
solution with constant shaking.
5.
End point is suggested by the change
of violet colour to emerald green.
6.
Repeat the titration three times to get -----

82
 precise readings.
7.
Take one reading of blank
determination i.e. do titration
according to above procedure without
potassium hydrogen phthalate.
8. Averaging the results will minimize the random
Calculate the normality for each error.
reading (after subtracting blank
reading) and take the mean of
normality.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.35 g KHP + 50 ml glacial 0.0 18.8 18.8
acetic acid + 2-3 drops of crystal
violet indicator
2. 0.36 g KHP + 50 ml glacial 20 40 20.0
acetic acid + 2-3 drops of crystal
violet indicator
3. 0.33 g KHP + 50 ml glacial 0.0 16.9 16.9
acetic acid + 2-3 drops of crystal
violet indicator

Blank titration:

Content in flask Burette reading Volume of

S. Initial Final perchloric acid
No (ml)
1. 50 ml glacial acetic acid + 2-3 0.0 0.4 0.4
drops of crystal violet indicator

Normality (N)= No. of moles or Weight (g) of primary standard x 1000

Volume (L) Eq. weight of primary standard x volume (ml) (reading-blank)

As we know,
1000 ml 1 N of perchloric acid= 204 g/eq of potassium hydrogen phthalate

83
 N1 = 0.35x1000
204 x 18.4
N2 = 0.36x1000
204 x 19.6
N3 = 0.33x1000
204 x 16.5

N=N1 + N2 + N3
3
= 0.093 + 0.090 + 0.098 = 0.094
3
RESULT
Normality of HClO4 solution is found to be 0.094

EXPERIMENT NO 5.2

TO PERFORM THE ASSAY OF EPHEDRINE USING STANDARD 0.1 N

PERCHLORIC ACID

THEORY

Ephedrine is very weakly basic drug. So it cannot be directly titrated with perchloric acid. So we
use non-aqueous method of titration for the estimation of ephedrine in sample. Ephedrine
basicity is increased by adding glacial acetic acid to form strong conjugate base.

1000 ml 1 N perchloric acid= 165 g/eq. of ephedrine

84
PROCEDURE

A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 N perchloric acid: Described in
experiment no. 5.1
2. Preparation of crytal violet indicator solution: Dissolve 500 mg of crystal violet
in 100 ml of anhydrous glacial acetic acid.

B. Assay of ephedrine using 0.1 N perchloric acid

1. The burette is filled with 0.1M perchloric acid solution.

2. Weigh accurately about 0.4 g of ephedrine sample in a 100 ml conical flask.

3. Add 25 ml of anhydrous glacial acetic acid and 2-3 drops of crystal violet indicator.

4. Start titration with the standard perchloric acid solution with constant shaking.

5. End point is suggested by the change of violet colour to emerald green

6. Repeat the titration three times to get precise readings. Take one reading of blank
determination i.e. do titration according to above procedure without potassium
hydrogen phthalate.

7. Calculate the percentage purity of each reading (after subtracting blank reading) and
take the mean of percentage purity.

85
 Observation Table (Specimen reading)
S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.4 g Ephedrine + 25 ml glacial 0.0 25.0 25.0
acetic acid + 2-3 drops of crystal
violet indicator
2.
3.

Blank titration:
Content in flask Burette reading Volume of
S. Initial Final perchloric acid
No (ml)
1. 25 ml glacial acetic acid + 2-3 0.0 0.2 0.2
drops of crystal violet indicator
Volume of perchloric acid used 25.0 - 0.2 = 24.8 ml

Calculation
Descriptive method Short cut method
1000 ml 1 N perchloric acid= 165 g/eq. of ephedrine Percentage purity= V x N x 165 x 100
17.5 ml of 0.094 N perchloric acid = 165 x 24.8 x 0.094 W (gm) x 1000
1000
= 0.384 g of ephedrine V=Volume of perchloric acid used
0.4 g of sample contains= 0.384 g of ephedrine N= Normality of silver nitrate
100 gm of sample contain = 0.384 x 100 W= Weight of sample (gm) taken
0.4 24.8 x 0.094 x 165 x 100
0.4 x 1000
= 96.2 % of ephedrine = 96.2 % of ephedrine

RESULT
The percentage purity of given sample of ephedrine is found to be 96.2 %.

EXPERIMENT NO 5.3

TO PERFORM THE ASSAY OF EPHEDRINE HYDROCHLORIDE USING

STANDARD 0.1 N PERCHLORIC ACID

THEORY

86
Ephedrine hydrochloride is very weakly basic drug. The halide ions, namely chloride, bromide
and iodide are very weakly basic in nature so much so that they cannot react quantitatively with
acetous perchloric acid. In order to overcome this problem, mercuric acetate is usually added (it
remain undissociated in acetic acid solution) to a halide salt thereby causing the replacement of
halide ion by an equivalent amount of acetate ion, which serve as a strong base in acetic acid as
shown below:

Mercuric acetate is added to prevent the interference of the hydrochloric acid displaced through
the formation of the relatively un-ionized HgCl2 , thereby making a predominant shift in the
equilibrium so that the titrimetric reaction is quantitative.
Blank titration is usually carried out to account for the possible reaction of atmospheric moisture
with the titrant perchloric acid and also to check the titrant being employed to bring about the
blue-green end point.

PROCEDURE

87
 A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 N perchloric acid: Described in
experiment no. 5.1
2. Preparation of crytal violet indicator solution: Dissolve 500 mg of crystal violet
in 100 ml of anhydrous glacial acetic acid.

B. Assay of ephedrine hydrochloride using 0.1 N perchloric acid

1. The burette is filled with 0.1M perchloric acid solution.

2. Weigh accurately about 0.4 g of ephedrine hydrochloride sample in a 100 ml

conical flask.

3. Add 50 ml of anhydrous glacial acetic acid and 10 ml of mercuric acetate solution.

4. Add 2-3 drops of crystal violet indicator.

5. Start titration with the standard perchloric acid solution with constant shaking.

6. End point is suggested by the change of violet colour to emerald green

7. Repeat the titration three times to get precise readings. Take one reading of blank
determination i.e. do titration according to above procedure without potassium
hydrogen phthalate.

8. Calculate the percentage purity of each reading (after subtracting blank reading) and
take the mean of percentage purity.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.4 g ephedrine hydrochloride + 0.0 19.0 19.0
50 ml glacial acetic acid + 10 ml
of mercuric acetate + 2-3 drops
of crystal violet indicator
2.
3.

Blank titration:
Content in flask Burette reading Volume of
S. Initial Final perchloric acid
No (ml)

88
 1. 50 ml glacial acetic acid + 10 ml 0.0 0.2 0.2
of mercuric acetate + 2-3 drops
of crystal violet indicator
Volume of perchloric acid used 19.0 - 0.2 = 18.8 ml
Calculation

Descriptive method Short cut method

1000 ml 1 N perchloric acid= 201.7 g/eq. of ephedrine HCl Percentage purity= V x N x 165 x 100
17.5 ml of 0.094 N perchloric acid = 201.7 x 18.8 x 0.094 W (gm) x 1000
1000
= 0.356 g of ephedrine HCl V=Volume of perchloric acid used
0.4 g of sample contains= 0.356 g of ephedrine HCl N= Normality of perchloric acid
100 gm of sample contain = 0.356 x 100 W= Weight of sample (gm) taken
0.4 18.8 x 0.094 x 201 x 100
0.4 x 1000
= 89.1 % of ephedrine HCl = 89.1 % of ephedrine HCl

RESULT
The percentage purity of given sample of ephedrine hydrochloride is found to be 89.1 %.

EXPERIMENT NO 5.4

TO PERFORM THE ASSAY OF SODIUM BENZOATE USING STANDARD 0.1 N

PERCHLORIC ACID

89
THEORY

Sodium benzoate is a white crystalline or granular powder which is mainly used as a

preservative. Sodium benzoate is a weak base and is dissolved in glacial acetic acid. During the
titration with strong acid, acetic acid behaves like a base and accurate end point is determined.
End point is detected by using 1-naphtholbenzein solution as an indicator.

PROCEDURE
A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 N perchloric acid: Described in

experiment no. 5.1
2. Preparation of 1-naphtholbenzein indicator: Dissolve 0.2 g of 1-naphtholbenzein
in 100 ml of anhydrous glacial acetic acid.

B. Assay of sodium benzoate

1. The burette is filled with 0.1M perchloric acid solution.

2. Weigh accurately about 0.25 g of sodium benzoate sample in a 100 ml conical flask.

3. Add 20 ml of anhydrous glacial acetic acid. Warm to 50 °C if needed.

4. Add 2-3 drops of 1-naphtholbenzeine indicator solution.

90
 5. Start titration with the standard perchloric acid solution with constant shaking.

6. End point is suggested by the change of violet colour to emerald green

7. Repeat the titration three times to get precise readings. Take one reading of blank
determination i.e. do titration according to above procedure without potassium
hydrogen phthalate.

8. Calculate the percentage purity of each reading (after subtracting blank reading) and
take the mean of percentage purity.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.25 g sodium benzoate + 20 ml 0.0 15.0 15.0
glacial acetic acid + 2-3 drops of
1-naphtholbenzein indicator
2.
3.

Blank titration:
Content in flask Burette reading Volume of
S. Initial Final perchloric acid
No (ml)
1. 20 ml glacial acetic acid + 2-3 0.0 0.2 0.2
drops of 1-naphtholbenzein
indicator
Volume of perchloric acid used 15.0 - 0.2 = 14.8 ml

Calculation
Descriptive method Short cut method
1000 ml 1 N perchloric acid= 144 g/eq. of sodium benzoate Percentage purity= V x N x 165 x 100
17.5 ml of 0.094 N perchloric acid = 144 x 14.8 x 0.094 W (gm) x 1000
1000
= 0.384 g of sodium benzoate V=Volume of perchloric acid used
0.25 g of sample contains= 0.384 g of sodium benzoate N= Normality of perchloric acid
100 gm of sample contain = 0.384 x 100 W= Weight of sample (gm) taken
0.25 14.8 x 0.094 x 144 x 100
0.25 x 1000
= 80.1 % of sodium benzoate =80.1% of sodium benzoate

RESULT
The percentage purity of given sample of ephedrine is found to be 80.1 %.

91
 EXPERIMENT NO 5.5

TO PREPARE AND STANDARDIZE 0.1 N SODIUM METHOXIDE USING BENZOIC

ACID AS PRIMARY STANDARD

THEORY

.PROCEDURE
A. Preparation of reagents and solutions:
1. 0.1 N sodium methoxide:

Calculation for making 100 ml of 0.1 N sodium methoxide solution:

1000 ml of 1 N sodium methoxide =23 g/eq of sodium metal
100 ml of 0.1 N sodium methoxide = 23 x 100 x 0.1
1000
= 0.23 g of sodium metal
Hence, weight accurately about 0.25 g of sodium metal and add in small proportions to 15 ml of
anhydrous methanol in beaker. When all metal dissolve, make up the volume to 100 ml with
toluene.
2. Preparation of thymolphthalein indicator solution:

B. Standardization of sodium methoxide solution using KHP as primary standard

1. The burette is filled with 0.1M sodium methoxide solution.

2. Weigh accurately about 0.4 g of benzoic acid in a 100 ml conical flask.

3. Add 80 ml of dimethyl formamide.

92
 4. Add 2-3 drops of thymolphthalein indicator.

5. Start titration with the sodium methoxide solution with constant shaking.

6. End point is suggested by the change of violet colour to emerald green.

7. Repeat the titration three times to get precise readings.

8. Take one reading of blank determination i.e. do titration according to above procedure
without potassium hydrogen phthalate.

9. Calculate the normality for each reading (after subtracting blank reading) and take the
mean of normality.

Observation Table (Specimen reading)

S. No Content in flask Burette reading Volume of
Initial Final perchloric acid (ml)
1. 0.4 g benzoic acid + 80 ml DMF + 0.0 31.8 31.8
2-3 drops of thymolphthalein
indicator
2. 0.38 g benzoic acid + 80 ml DMF + 0.0 30.6 30.6
2-3 drops of thymolphthalein
indicator
3. 0.39 g benzoic acid + 80 ml DMF + 0.0 31.2 31.2
2-3 drops of thymolphthalein
indicator

Blank titration
Content in flask Burette reading Volume of
S. Initial Final perchloric acid
No (ml)
1. 80 ml DMF + 2-3 drops of 0.0 0.6 0.6
thymolphthalein indicator

Normality (N)= No. of moles or Weight (g) of primary standard x 1000

Volume (L) Eq. weight of primary standard x volume (ml) (reading-blank)

As we know,
1000 ml 1 N of sodium methoxide= 122 g/eq of benzoic acid

N1 = 0.4x1000
122 x 31.2
N2 = 0.38x1000
122 x 30.0
N3 = 0.39x1000
122 x 29.6

93
 N=N1 + N2 + N3
3
= 0.105 + 0.10 + 0.108 = 0.104
3
RESULT
Normality of sodium methoxide solution is 0.104
EXPERIMENT NO 5.6

TO PERFORM THE ASSAY OF ALLOPURINOL USING STANDARD 0.1 N

SODIUM METHOXIDE SOLUTION

THEORY

Allopurinol is very weakly acidic drug which cannot be titrated directly with sodium methoxide.

PROCEDURE
A. Preparation of reagents and solutions:
1. Preparation and standardization of 0.1 N sodium methoxide: Described in
experiment no…
2. Preparation of methyl red indicator solution: Dissolve 200 mg of methyl red in
100 ml of dioxan as solvent to form 0.2 % w/v solution.

B. Assay of Allopurinol

1. The burette is filled with 0.1M standardized sodium methoxide solution.

94
 2. Weigh accurately about 0.2 g of allopurinol sample in a 100 ml conical flask.

3. Add 50 ml of dimethyl formamide as solvent and dissolve with gentle heating

4. Add 2-3 drops of methyl red indicator.

5. Start titration with the standard perchloric acid solution with constant shaking.

6. End point is suggested by the change of yellow to red.

7. Repeat the titration three times to get precise readings. Take one reading of blank
determination i.e. do titration according to above procedure without allopurinol sample

8. Calculate the percentage purity of each reading (after subtracting blank reading) and take
the mean of percentage purity.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.2 g allopurinol + 50 ml DMF+ 0.0 14.0 14.0
2-3 drops of methyl red
2.
3.

Blank titration:
Content in flask Burette reading Volume of
S. Initial Final perchloric acid
No (ml)
1. 50 ml DMF+ 2-3 drops of 0.0 0.2 0.2
methyl red
Volume of perchloric acid used 14.0 - 0.2 = 13.8 ml
Calculation
Descriptive method Short cut method
1000 ml 1 N sodium methoxide= 136 g/eq. of allopurinol Percentage purity= V x N x 136 x 100
10.5 ml of 0.104 N sodium methoxide = 136 x 13.8 x 0.104 W (gm) x 1000
1000
= 0.195 g of allopurinol V=Volume of sodium methoxide used
0.2 g of sample contains= 0.195 g of allopurinol N= Normality of sodium methoxide
100 gm of sample contain = 0.195 x 100 W= Weight of sample (gm) taken
0.2 13.8 x 0.104 x 136 x 100
97.6 g of allopurinol 0.4 x 1000
i.e. 97.6 % allopurinol =97.6 % allopurinol

95
RESULT
The percentage purity of given sample of allopurinol is found to be 97.6 %.

6. Diazotization titration

6.1. To prepare and standardize 0.1 N sodium nitrite using sulfanilic acid as primary
standard.
6.2.To determine the percentage purity (assay) of sulfanilic acid using standard by
diazotization titration.

96
 EXPERIMENT NO 6.1

TO PREPARE AND STANDARDIZE 0.1 N SODIUM NITRITE SOLUTION USING

SULPHANILIC ACID AS PRIMARY STANDARD

THEORY

The diazotization titration is carried out for the estimation of drug containing primary aromatic
group. Several drugs containing either primary aromatic group or they can be converted to have
such group by simple reaction for instance by hydrolysis, reduction etc can easily be estimated
by this method.
The principle involved in this method is that the primary aromatic amine present in the sample
reacts with the sodium nitrite in the presence of acid such as hydrochloric acid to obtain a
diazonium salt at 0–5 °C.
In titration process, when sodium nitrite is added from burette to conical flask containing
sulphanilic acid in acidic solution, sodium nitrite react with hydrochloric acid forming nitrous
acid (HONO). The nitrous acid react with the primary aromatic group (sulphanilic acid)
quantitatively resulting into the formation of an unstable nitrite that decompose ultimately with
the formation of diazonium salt. After the end point (when sulphanilic acid consumes), the
excess nitrous acid leading to the release of iodine (react with HI) that react with external
indicator (starch iodide paper) to form blue colour.
The diazonium salt produced is not stable and if the reaction is not maintained between 0-5 °C, it
shall undergo decomposition thereby forming phenol products which may react further with
nitrous acid leading to wrong results.
At the equivalence point a slight excess of nitrous acid

97
 0.1 N NaNO2

NaNO2 + HCl

0-5 0C
HO3S NH2 HONO Ar-N NCl + NaCl + H2O
Sulphanilic acid

Starch + KI Starch-iodide paper

Starch-iodide paper as
KI + HCl KCl + HI external indicator

2HI + 2HONO I2 + 2NO + 2H2O Colourless to blue colour

This iodine react with starch of

starch-iodide paper to give
blue colour. 0.3 g sulphanilic acid + 50 ml 2M water+
3 g of KBr

Requirements

1. The temperature should be maintained at 0-5 C as diazonium salts are not stable at
elevated temperatures. They are readily decomposed at high temperature.
2. KBr is added to catalyse the reaction and ultimately speed up the rate of reaction between
sodium nitrite and analyte drugs because some of the drugs are slow diazotizable like
sulphanilic acid and anthranilic acid and some are fast diazotiazable like aniline
aminophenol and toluidine.
KBr increases the electrophilicity of NO+. This is due to higher electronegativity of Br-,
as it pulls the electron cloud toward itself.

PROCEDURE

A. Preparation of reagents and solutions:

1. 0.1 N sodium nitrite:

Calculation for making 100 ml of 0.1 N sodium nitrite solution:

1000 ml of 1 N NaNO2=69g/eq of NaNO2
100 ml of 0.1 N NaNO2 = 69 x 100 x 0.1
1000
= 0.69 g of NaNO2
Hence, weight accurately about 0.7 g of NaNO2, dissolved in 70 ml of distilled water and
make up the volume to 100 ml with glacial acetic acid.

98
 Preparation of starch-iodide indicator solution: Add 1 gram of starch (either corn or
potato) and 0.25 g of potassium iodide (KI) into 10 mL of distilled water, shake well, and
pour into 100 mL of distilled water. Stir thoroughly and boil for a 1 minute. Leave to cool
down. If the precipitate forms, decant the supernatant and use as the indicator solution. A
fresh solution should be used always.

2. Preparation of 2M HCl: See Exp No 1.1

B. Standardization of NaNO2 solution using sulphanilic acid as primary standard.

1. The burette is filled with 0.1M sodium nitrite solution.

2. Weigh accurately about 0.3 g of sulphanilic acid in a 100 ml conical flask.
3. Add 50 ml of 2 M hydrochloric acid solution
4. Add 3 g of potassium bromide.
5. Keep the conical flask in bowl containing crushed ice to maintain temperature (0-5 oC)
6. Start titration with the sodium nitrite solution with constant shaking.
7. After addition of titrant, until the tip of the glass rod dipped into the conical flask
immediately produces a distinct blue ring on being touched to starch-iodide solution or
paper.
8. Repeat the titration three times to get precise readings.
9. Calculate the normality for each reading and take the mean of normality.

S. Steps Reason
No.
1. The burette is filled with 0.1M Sodium nitrite solution acts as titrant.
sodium nitrite solution.
2. Weigh accurately about 0.3 g of 0.3 g weighed as per the pharmacopoeia.
sulphanilic acid in a 100 ml
conical flask.
3. Add 50 ml of 2 M
hydrochloric acid solution. HCl acid helps in forming nitrous acid
(HONO) by reacting with sodium nitrite.
4. Add 3 g of potassium bromide. KBr increases the electrophilicity of
NO+. This is due to higher
electronegativity of Br-, as it pulls the
electron cloud toward itself.
5. Keeps the conical flask in bowl The temperature should be maintained at
containing crushed ice to maintain 0-5 C as diazonium salts are not stable at
o
temperature (0-5 C). elevated temperatures. They are readily
decomposed at high temperature.

6. Start titration with the sodium

nitrite solution as titrant with -----
constant shaking.

99
 7. After addition of titrant, until the
tip of the glass rod dipped into the After the end point (when sulphanilic acid
conical flask immediately consumes), the excess nitrous acid leading to
produces a distinct blue ring on the release of iodine (react with HI) that react
being touched to starch-iodide with external indicator (starch iodide paper)
solution or paper. to form blue colour.
8. Repeat the titration three times to Averaging the results will minimize the
get precise readings. Take mean random error.
and calculate the normality of
sodium nitrite.

Observation Table (Specimen reading)

S. Content in flask Burette reading Volume of
No Initial Final perchloric acid
(ml)
1. 0.3 g sulphanilic acid + 50 ml 0.0 16.0 16.0
2M HCl + 3 g of potassium
bromide
2. 0.31 g sulphanilic acid + 50 ml 16.0 33 17.0
2M HCl + 3 g of potassium
bromide
3. 0.29 g sulphanilic acid + 50 ml 30.0 45.0 15.0
2M HCl + 3 g of potassium
bromide

Normality (N)= No. of moles or Weight (g) of primary standard x 1000

Volume (L) Eq. weight of primary standard x volume (ml)

As we know,
1000 ml 1 N of sodium nitrite= 173 g/eq of sulphanilic acid

N1 = 0.30x1000
173 x 16.0
N2 = 0.31x1000
173 x 17.0
N3 = 0.29x1000
173 x 15.0

N=N1 + N2 + N3
3
= 0.108 + 0.105 + 0.11 = 0.107
3

RESULT

100
Normality of sodium nitrite solution is 0.107.

EXPERIMENT NO 6.2

TO FIND OUT THE PERCENTAGE PURITY OF SULFANILIC ACID BY

DIAZOTIZATION TITRATION

THEORY

The principle involved in this method is that the primary aromatic amine present in the sample
reacts with the sodium nitrite in the presence of acid such as hydrochloric acid to obtain a
diazonium salt at 0–5 °C.
In titration process, when sodium nitrite is added from burette to conical flask containing
sulphanilic acid in acidic solution, sodium nitrite react with hydrochloric acid forming nitrous
acid (HONO). The nitrous acid react with the primary aromatic group (sulphanilic acid)
quantitatively resulting into the formation of an unstable nitrite that decompose ultimately with
the formation of diazonium salt. After the end point (when sulphanilic acid consumes), the
excess nitrous acid leading to the release of iodine (react with HI) that react with external
indicator (starch iodide paper) to form blue colour.
The diazonium salt produced is not stable and if the reaction is not maintained between 0-5 °C, it
shall undergo decomposition thereby forming phenol products which may react further with
nitrous acid leading to wrong results.
At the equivalence point a slight excess of nitrous acid

101
 0.1 N NaNO2

NaNO2 + HCl

0-5 0C
HO3S NH2 HONO Ar-N NCl + NaCl + H2O
Sulphanilic acid

Starch + KI Starch-iodide paper

Starch-iodide paper as
KI + HCl KCl + HI external indicator

2HI + 2HONO I2 + 2NO + 2H2O Colourless to blue colour

This iodine react with starch of

starch-iodide paper to give
blue colour. 0.3 g sulphanilic acid + 50 ml 2M water+
3 g of KBr

REQUIREMENTS

1. The temperature should be maintained at 0-5 C as diazonium salts are not stable at
elevated temperatures. They are readily decomposed at high temperature.
2. KBr is added to catalyse the reaction and ultimately speed up the rate of reaction between
sodium nitrite and analyte drugs because some of the drugs are slow diazotizable like
sulphanilic acid and anthranilic acid and some are fast diazotiazable like aniline
aminophenol and toluidine.
KBr increases the electrophilicity of NO+. This is due to higher electronegativity of Br-,
as it pulls the electron cloud toward itself.

PROCEDURE

A. Preparation of reagents and solutions:

1. Preparation and standardization of 0.1 N sodium nitrite solution: Described in
experiment no. 6.1
2. Preparation of starch-iodide solution: Add 1 gram of starch (either corn or potato)
and 0.25 g of potassium iodide (KI) into 10 mL of distilled water, shake well, and
pour into 100 mL of distilled water. Stir thoroughly and boil for a 1 minute. Leave to
cool down. If the precipitate forms, decant the supernatant and use as the indicator
solution. A fresh solution should be used always.

B. TITRATION STEPS

1. The burette is filled with 0.1M sodium nitrite solution.

102
2. Weigh accurately about 0.3 g of sample in a 100 ml conical flask.
3. Add 50 ml of 2 M hydrochloric acid solution.
4. Add 3 g of potassium bromide.
5. Keep the conical flask in bowl containing crushed ice to maintain temperature (0-5 oC).
6. Start titration with the sodium nitrite solution with constant shaking.
7. After addition of titrant, until the tip of the glass rod dipped into the conical flask
immediately produces a distinct blue ring on being touched to starch-iodide solution or
paper.
8. Repeat the titration three times to get precise readings.
9. Calculate the normality for each reading and take the mean of normality.

S. Steps Reason
No.
1. The burette is filled with 0.1M Sodium nitrite solution acts as titrant.
sodium nitrite solution.

2. Weigh accurately about 0.3 g of sample 0.3 g weighed as per the pharmacopoeia.
in a 100 ml conical flask.

3. Add 50 ml of 2 M hydrochloric
acid solution. HCl acid helps in forming nitrous acid
(HONO) by reacting with sodium nitrite.
4. Add 3 g of potassium bromide. KBr increases the electrophilicity
of NO+. This is due to higher
electronegativity of Br-, as it
pulls the electron cloud toward
itself.
5. Keeps the conical flask in bowl The temperature should be
containing crushed ice to maintain maintained at 0-5 C as diazonium
temperature (0-5 oC). salts are not stable at elevated
temperatures. They are readily
decomposed at high temperature.
6. Start titration with the sodium nitrite
solution as titrant with constant -----
shaking.
7. After addition of titrant, until the tip of
the glass rod dipped into the conical After the end point (when sulphanilic
flask immediately produces a distinct acid consumes), the excess nitrous acid
blue ring on being touched to starch- leading to the release of iodine (react
iodide solution or paper. with HI) that react with external
indicator (starch iodide paper) to form
blue colour.
8. Repeat the titration three times to get Averaging the results will minimize the
precise readings. random error.
9. Calculate the percentage purity of the
given sample of sulphanilic acid.

103
 OBSERVATION TABLE (SPECIMEN READING)

S. Content in flask Burette reading Volume of

No Initial Final perchloric acid
(ml)
1. 0.3 g sample + 50 ml 2M HCl + 0.0 11.0 11.0
3 g of potassium bromide

CALCULATION

Descriptive method Short cut method

1000 ml 1 N sodium nitrite = 173 g/eq. of sulphanilic acid Percentage purity= V x N x 173 x 100
17.5 ml of 0.107 N sodium nitrite = 173 x 11.0 x 0.107 W (gm) x 1000
1000
= 0.20 g of sulphanilic acid V=Volume of perchloric acid used
0.3 g of sample contains= 0.20 g of sulphanilic acid N= Normality of silver nitrate
100 gm of sample contain = 0.20 x 100 W= Weight of sample (gm) taken
0.3 11.0 x 0.107 x 173 x 100
0.3 x 1000
= 67.8 % of sulphanilic acid = 67.8 % of sulphanilic acid

RESULT. The percentage purity of given sample of sulphanilic acid is found to be 67.8 %.

7. Conductometric and potentiometric titration

7.1. To determine the strength of strong acid using strong base by conductometry.
7.2. To deterimine the strength of strong acid using strong base by potentiometry.

104
 EXPERIMENT NO 7.1

TO DETERMINE THE STRENGTH OF STRONG ACID USING STRONG BASE BY

CONDUCTOMETRY
PRINCIPLE
Hydrochloric acid is a strong acid and sodium hydroxide is a strong base. The strength of
hydrochloric acid can be determined by titrating it directly with sodium hydroxide. When
electrode is immersed in a beaker containing HCl, it conductivity is high which is called initial
conductivity. This is because strong acid completely dissociates into H + of HCl and produce
water which results in decrease of conductivity on every addition. At the end-point all the H + of
HCl reacts with OH- of NaOH resulting water. After that point, further addition of NaOH
increase conductance due to OH- which results in ‘V’ shaped graph.

105
PROCEDURE
A. Preparation of reagents and solution
1. Preparation and standardization of hydrochloric acid solution (1M): See experiment
no. 1.1
2. Preparation and standardization of sodium hydroxide (1M): See experiment no 1.3
B. Titration
1. Clean and dry all glassware as per standard laboratory procedure.
2. Rinse the burette with distilled water. Then fill it with standard NaOH solution.
3. Switch on the conductometer and allow it to stabilize for 20 to 30 minutes.
4. Connect calomel electrode to positive and glass electrode to negative terminal and
rinse both the electrodes with distilled water.
5. Calibrate the electrodes by using standard buffer solutions of pH 4.0 to 7.0
6. Add 50 ml of HCl acid in a beaker and immersed the platinum electrode in it.
7. Note down the initial conductance.
8. Add 1 ml NaOH at a time from burette and note down the readings.
9. Continue until the end point is observed (at the end point there is sharp increase in
conductivity) and take few more readings even after the end point.
10. Plot the graph between conductivity (mho) Vs volume of NaOH

OBSERVATION TABLE

106
 Sr. No Volume of NaOH Conductivity
(ml) (mho)
1
2
3
4
5
6
7
So on..

CALCULATIONS

M1V1 = M2V2
M2 = M1V1/V2
Where,
M2 = molarity of HCl
V1 = volume of NaOH (ml)
M1 = molarity of NaOH
V2 = volume of HCl (ml)

RESULTS: The strength of hydrochloric acid was found to be -----M

EXPERIMENT NO 7.2

TO DETERMINE THE STRENGTH OF STRONG ACID USING STRONG BASE BY

POTENTIOMETRY

PRINCIPLE

Potentiometry is an electrochemical method. Sulphuric acid is a strong acid and sodium

hydroxide is an strong base. The strength of sulphuric acid can be determined by titrating it
directly with sodium hydroxide. End point can be detected by sharp change in potential across
the electrodes and determined by plotting the graph between volume of base Vs potential
(E.M.F.).

107
PROCEDURE

A. Preparation of reagents and solution

1. Preparation and standardization of sulphuric acid solution (1M): See experiment no.
1.2
2. Preparation and standardization of sodium hydroxide (1M): See experiment no 1.3
B. Titration
1. Clean and dry all glassware as per standard laboratory procedure.
2. Rinse the burette with distilled water. Then fill it with standard NaOH solution.
3. Switch on the potentiometer and set the temperature knob at room temperature.
4. Connect calomel electrode to positive and glass electrode to negative terminal and
rinse both the electrodes with distilled water.
5. Calibrate the electrodes by using standard buffer solutions of pH 4.0 to 7.0
6. After calibration, rinse the electrodes with distilled water.
7. Add 50 ml of sulphuric acid in a beaker placed on a stirrer and immerse both the
electrodes in it.
8. Note down the reading, i.e. e.m.f. in millivolts (mV).
9. Switch on the stirrer and add 1 ml NaOH at a time from burette and note down the
readings in mV.
10. Continue until the end point is observed (at the end point there is sharp increase in
emf) and take few more readings even after the end point.
11. Plot the graph between:
a. Emf Vs volume of NaOH added (Normal curve)
b. ∆E/∆V Vs volume of NaOH added (1st derivative curve)
c. ∆2E/∆V2 Vs volume of NaOH added (2nd derivative curve).

108
 OBSERVATION TABLE

Sr. No Volume of NaOH EMF (mV) ∆E/∆V = E2-E1/V2-V1 ∆2E/∆V2

(ml)
1
2
3
4
5
6
7
So on..

CALCULATIONS

M1V1 = M2V2
M2 = M1V1/V2
Where,
M2 = molarity of H2SO4 (ml)
V1 = volume of NaOH (ml)
M1 = molarity of NaOH
V2 = volume of H2SO4 (ml)

RESULTS
Normal, 1st and 2nd derivative curves are plotted for sulphuric acid and its strength was found to
be __ M.

109

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