Category Archives: Pharmaceutical Analysis – I (Practical)

PRACTICAL-ASSAY OF AMMONIUM CHLORIDE BY ACID BASE TITRATION

Object – To perform the assay of Ammonium chloride by acid base titration

Reference – 1. Siddiqui A.A. & Ali M., Practical Pharmaceutical Chemistry, CBS publishers and distributors, New Delhi, I edition, 1997, 22-23

  1. Rao G.D., Practical pharmaceutical analysis, Birla publications Pvt. Ltd., Delhi, II edition, 2007-08, 33

Glasswares & apparatus required – digital balance, pipette, burette, burette stand, conical flask

Chemicals required – ammonium chloride, distill water, formaldehyde solution, phenolphthalein indicator, sodium hydroxide solution

Theory – Acid base titration or neutralization titration involves neutralization of acid with base or base with acid and the end point is determined by means of indicator.

Acidimetry titration – it is a direct or residual volumetric analysis of a base with a standard acid. Direct titration is performed by introducing a standard acid solution gradually from a burette into a solution of base being assayed till the end point is obtained e.g. assay of sodium bicarbonate. Residual titration is used when the rate of reaction between a basic compound with an acid is slow. In this, the solution of the base is treated with an excess of accurately measured standard acid and excess acid is subsequently titrated with standard base e.g. assay of zinc oxide.

Alkalimetry titration – it is an estimation of acid/acidic drugs by titration with standard alkali. It also includes direct titration and residual/back titration methods as in acidimetry. Assay of ammonium chloride comes under direct titration method of Alkalimetry titration.

Ammonium chloride is an example of diuretic and systemic acidifier. The assay of ammonium chloride is based on the principle of formal titration. The formal titration is done in the presence of formaldehyde. When ammonium chloride is treated with formaldehyde (HCHO), HCl is liberated. The liberated hydrochloric acid is btitrated with a standard solution of sodium hydroxide using phenolphthalein indicator.

Procedure

  1. 1 gm of ammonium chloride was accurately weighed and dissolved in 20 ml water.
  2. 5 ml formaldehyde solution was added.
  3. After two minutes, titration was started slowly with 0.1 N Sodium hydroxide (given below) solution by using phenolphthalein indicator till pink color was obtained.
Preparation of 100 ml 0.1 N NaOH
Normality Volume Amount Dissolved in
1 N 1000 ml 40 gm 1000 ml
0.5 N 1000 ml 20 gm 1000 ml
0.1 N 1000 ml 4 gm 1000 ml
0.1 N 500 ml 2 gm 500 ml
0.1 N 100 ml 0.4 gm = 400 mg 100 ml

 

The reaction involved is-

HCl + NaOH           →      NaCl + H2O

Equivalent Factor – 1 ml of 0.1 N NaOH is equivalent to 0.005349 g of NH4Cl

Observation table –

S.NO. VOLUME OF AMMONIUM CHLORIDE SOLUTION IN ML INITIAL BURETTE READING FINAL BURETTE READING VOLUMJE OF 0.1 N SODIUM HYDROXIDE CONSUMED IN ML
1 25 50 33 (suppose) 17 (suppose)

 

 

Calculation –

Percentage purity of ammonium chloride =

Volume of 0.1 N sodium hydroxide x equivalent factor x 100 x normality of sodium hydroxide (actual)

Weight of ammonium chloride in grams x normality of sodium hydroxide (expected)

17 x 0.005349 x 100 x 0.1

=   ———————————-

  • x 0.1

= 0.05349 x 17 / 0.01 = 0.909 / 0.01 = 90.9%

Result – Assay of Ammonium chloride by acid base titration was carried out and the percentage purity was found 90.9%.

PRACTICAL-TITRATION OF STRONG ACID (0.1 N HCl) WITH STRONG BASE (0.1 N NaOH) BY POTENTIOMETRY

Object – Determination of Normality by electrochemical methods [Potentiometric titration of strong acid (0.1 N HCl) against strong base (0.1 N NaOH)]

Reference – Rao G.D., Practical pharmaceutical analysis, Birla publications Pvt. Ltd., Delhi, II edition, 2007-08, 136

Materials required-

Chemicals – 0.1 N HCl, 0.1 N NaOH, Distill water

Glassware & Apparatus/Instumentation – Potentiometer or PH meter, Magnetic stirrer, burette, beaker, pipette.

Theory – Potentiometric determination of the end point depends on the fact that the potential cross the two electrodes (reference and indicator) immersed in the solution changes sharply at the equivalence or end point. This change is similar to the color change by an indicator in usual method. But the potentiometric method is more accurate. These ti0trations are useful when no suitable color indicators are available. Equivalence point can be accurately found out after plotting normal plot (i.e. volume of Titrant Vs potential).

Procedure –

  1. In a 250 ml beaker, 10 ml of 0.1 N HCl was taken and then 100 ml of water was added to it, so that the electrodes can dip properly.
  2. The beaker was now kept on magnetic stirrer.
  3. Potential was noted down without adding alkali.
  4. Now potential was noted down with adding known volume of 0.1 N NaOH Solution.
  5. Now from the obtained data, a graph was plotted between volume of titrant and potential.
  6. From the graph end point was calculated.

Observation –

S.NO. VOLUME OF TITRANT ADDED POTENTIAL
1 1 ML
2 1 ML
3 1 ML
4 1 ML
5 1 ML
6 1 ML
7 1 ML
8 1 ML
9 1 ML
10 1 ML

 

Result – The end point in the titration of strong acid (0.1 N HCl) with strong base (0.1 N NaOH) was found ……………………….

PRACTICAL-PREPARATION & STADARDIZATION OF 0.1 N SULPHURIC ACID

Object – Preparation & standardization of Sulphuric acid (0.1 N)

Reference – 1. Kasture A.V. etal, Practical Pharmaceutical Chemistry-I, Nirali Prakashan, Pune, XI edition, 2005, 47

  1. Rao G.D., Practical pharmaceutical analysis, Birla publishers Pvt. Ltd., Delhi, II edition, 2007-08, 25

Glasswares & apparatus required – Volumetric flask, burette, burette stand, pipette, conical flask, digital balance, heating mantle.

Chemicals required – standard 0.1 N sodium carbonate solution (prepared by dissolving 530 mg of sodium carbonate in 100 ml of distill water), conc. Sulphuric acid, methyl orange solution

Theory – Sulphuric acid is a diprotic acid and 1 N solution contain 98.08/2 = 49.04 g H2SO4. Taking into consideration specific gravity (1.83) of sulphuric acid about 49.0 ml of conc. Sulphuric acid is required to prepare 1000 ml solution.

It is an example of alkalimetry. When a strong acid is titrated with a strong base, the salt produced in the reaction is not hydrolysed and therefore the ph of the resultant solution at the end point is exactly 7.0. sulphuric acid is a strong acid, is standardized by titrating with a strong base i.e. sodium carbonate (primary standard). The following reaction takes place when sodium carbonate is titrated with sulphuric acid. In this titration, end point detection is carried out by using methyl orange indicator.

Procedure –

In a volumetric flask, 4.9 ml of conc. Sulphuric acid was taken and 900 ml water was slowly added, cooled and then the volume was made upto 1000 ml with water.

Standardization –

  1. 10 ml of 0.1 N Sodium carbonate solution was pipette out into a clean and dried conical flask.
  2. 2 drops of methyl orange indicator was added to it.
  3. The contents of the flask was now titrated with sulphuric acid until red color was obtained.
  4. Burette reading was taken.

Observation table –  

S.NO. VOLUME OF 0.1 N SOD. CARBONATE SOLUTION TAKEN INITIAL BURETTE READING FINAL BURETTE READING VOLUME OF 0.1 N SULPHURIC ACID CONSUMED
1 10 ML 50 38 12 (SUPPOSE)

 

 

Calculation – Normality of H2SO4  is calculated by-

N1V1 = N2V2

N1 = 0.1 N = Normality of Na2CO3 Solution, N2 = ? = Normality of H2SO4

V1 = 10 ml = Volume of Na2CO3 Solution, V2 = 30 ml = Volume of H2SO4

0.1 X 10 = N2  X 12

Or N2 = 1 / 12 = 0.08 N

Result – Sulphuric acid (0.1 N) was prepared and standardized. The exact normality was found to be 0.08 N

PRACTICAL PHARMACEUTICAL ANALYSIS – PREPARATION AND STANDARDIZATION OF 1 N SODIUM HYDROXIDE

Object – Preparation & standardization of Sodium hydroxide (1 N solution)

Reference – 1. Siddiqui A.A. & Ali M., Practical Pharmaceutical Chemistry, CBS publishers and distributors, New Delhi, I edition, 1997, 6-7

  1. Rao G.D., Practical pharmaceutical analysis, Birla publishers Pvt. Ltd., Delhi, II edition, 2007-08, 25-26

Glasswares & apparatus required – digital balance, volumetric flask, pipette, conical flask, burette, burette stand

Chemicals required – sodium hydroxide pellets, distill water, phenolphthalein solution, hydrochloric acid, methyl orange solution.

Theory – A neutralization reaction involves the titration of the free bases with a standard acid (acidimetry), and the titration of free acid with a base (alkalimetry). These reactions include the combination of hydrogen and hydroxide ion to form water. The point at which this is reached is the equivalence point or theoretical end point. If both the acid and base are strong electrolytes, the resultant solution will be neutral and have a pH of 7. If either the acid or the base is a weak electrolyte, the salt will be hydrolyzed to some extent and the solution at the equivalence point will be either slightly alkaline or slightly acid.

The acid base indicators possess different colors according to the hydrogen ion concentration of the solution and the position of the color change interval in the pH scale varies widely with different indicators.

In aqueous solution an acid is dissociated into hydrogen ion and anion-

HCl → H+ + Cl

The hydrogen ion combines with water to give hydronium ion.

An alkali on dissociation in water gives hydroxyl ions,

NaOH ↔ Na+ + OH

The neutralization reaction is the reaction between an acid and a base to give a salt.

HCl + NaOH → NaCl + H2O

Procedure –

Preparation of 1 N Sodium Hydroxide-

Accurately weighed 40 gm sodium hydroxide pellets was taken in a 1000 ml volumetric flask and then 50 ml water was added, shaked and then the volume was made upto 1000 ml with water.

Standardization –

  1. 20 ml of 1 N Sodium hydroxide was pipette and taken in a conical flask.
  2. Then 2 drops of phenolphthalein was added to it.
  3. Now it was titrated with 1 N Hydrochloric acid until the color became colorless.

Observation – Suppose Burette reading = 16 ml

Calculation –                                  N1V1 = N2V2

Where N1 = Normality of unknown, V1 = Volume of unknown = 20 ml, N2 = Normality of known (1 N HCl) = 1, V2 = Volume of known (1 N HCl)= 16 ml

                                           N1 = N2V2 / V1 = 1 X 16 / 20 = 0.8 N

Result – Sodium hydroxide (1 N solution) was prepared and standardized. The normality was found to be 0.8 N.