Physical Pharmacy-I Lab, Manik

Course Code: PHRM 1206
Name: :
Full ID :
Section : 2A 2B

Physical Pharmacy-I Lab
Laboratory Notebook
Department of Pharmacy; Northern University Bangladesh.
ACETATE BUFFER SOLUTIONS PH 3.6 - 5.6 ------------------------------------------------------------------------------------------------------1
------------------------------------------------------------------------------------------------------------------------------------------- 2
INDEX
EXPERIMENT NO. 01----------------------------------------------------------------------------------------------------------------------------- 3
NAME OF THE EXPERIMENT: PREPARATION OF SOLUTIONS OF DIFFERENT CONCENTRATIONS. ------------------------------------------------- 3
1.1. PRINCIPLE: -------------------------------------------------------------------------------------------------------------------------------- 3
EXPERIMENT NO. 02 ---------------------------------------------------------------------------------------------------------------------------- 8
NAME OF THE EXPERIMENT: PREPARATION OF SOLUTIONS OF DIFFERENT CONCENTRATIONS BY DILUTION.------------------------------------ 8
2.1. PRINCIPLE:-------------------------------------------------------------------------------------------------------------------------------- 8
EXPERIMENT NO. 03 ----------------------------------------------------------------------------------------------------------------------------12
NAME OF THE EXPERIMENT: STANDARDIZATION OF NAOH SOLUTION BY POTASSIUM HYDROGEN PHTHALATE (KHP). 12
INTRODUCTION ---------------------------------------------------------------------------------------------------------------------------------12
TERMINOLOGIES USED IN TITRIMETRIC ANALYSIS ---------------------------------------------------------------------------------------------------12
PRINCIPLE-------------------------------------------------------------------------------------------------------------------------------------12
EXPERIMENT NO. 04 ----------------------------------------------------------------------------------------------------------------------------17
NAME OF THE EXPERIMENT: STANDARDIZATION OF NAOH SOLUTION BY OXALIC ACID SOLUTION. ----------------------------------------------17
PRINCIPLE-------------------------------------------------------------------------------------------------------------------------------------17
EXPERIMENT NO. 05 --------------------------------------------------------------------------------------------------------------------------- 20
NAME OF THE EXPERIMENT: STANDARDIZATION OF HCL SOLUTION BY NAOH SOLUTION. ------------------------------------------------------ 20
PRINCIPLE------------------------------------------------------------------------------------------------------------------------------------ 20
EXPERIMENT NO. 06 --------------------------------------------------------------------------------------------------------------------------- 25
NAME OF THE EXPERIMENT: PREPARATION OF CONSTANT PH BUFFER AND ASSESSMENT OF ITS ACTIVITY ----------------------------------- 25
PRINCIPLE:----------------------------------------------------------------------------------------------------------------------------------- 25
EXPERIMENT NO. 07---------------------------------------------------------------------------------------------------------------------------- 30
NAME OF THE EXPERIMENT: DETERMINATION OF THE ACID DISSOCIATION CONSTANT (KA) OF A WEAK ACID (ACETIC ACID).------------------ 30
PRINCIPLE------------------------------------------------------------------------------------------------------------------------------------ 30
BIBLIOGRAPHY--------------------------------------------------------------------------------------------------------------------------------- 34

Laboratory Notebook Page 1
Acetate buffer solutions pH 3.6 - 5.6
The stock solutions are as followings:
(1) 0.1M acetic acid
(2) 0.1M sodium acetate (tri-hydrate) (13.6g /L)
Mix in the following proportions to get the required pH [Source: Pearse et al 1980]
pH Vol. of 0.1M acetic acid Vol. of 0.1M sodium acetate
3.6 185 mL 15 mL
3.8 176 mL 24 mL
4.0 164 mL 36 mL
4.2 147 mL 53 mL
4.4 126 mL 74 mL
4.6 102 mL 98 mL
4.8 80 mL 120 mL
5.0 59 mL 141 mL
5.2 42 mL 158 mL
5.4 29 mL 171 mL
5.6 19 mL 181 mL
Md.
Imran
Nur
Manik
Scan QR code

Sl. Date Name of the experiment Remarks
01.
Preparation of Solutions of Different
Concentrations.
02.
Preparation of Solutions of Different
Concentrations by Dilution Method.
03.
Standardization of NaOH solution by
Potassium Hydrogen Phthalate (KHP).
04.
Standardization of NaOH solution by
Oxallic Acid Solution.
05.
Standardization of HCl Solution by
NaOH Solution.
06.
Preparation of constant pH buffer and
assessment of it’s activity.
07.
Determination of the acid dissociation
constant (Ka) of a weak acid (acetic
acid).
Edition: November, 2021

Experiment No. 01 Date:
Name of the experiment: Preparation of solutions of different concentrations.
1.1. Principle:
A solution may be defined as a homogeneous mixture of two or more components that form a single phase.
The component that determines the phase of the solution is termed the solvent and usually constitutes the
largest proportion of the system. The (other) components dispersed as molecules or ions throughout the solvent
are termed solutes.
The concentration of a solution is defined as the amount of solute present in a specific (fixed) amount of solvent
or solution.
The solution that contains a precisely known concentration of an element or a substance is known as a
standard solution. For example, a solution of HCl with a concentration of 0.1025 M is a standard solution of
HCl.
In a standard solution, the amount of solute and the volume of solvent are accurately known.
Standard solutions are used to determine the concentrations of other substances, such as solutions in titration.
The concentrations of standard solutions are expressed in various units such as molarity, molality, normality,
mole fraction, etc.
Molarity is the number of moles of solute dissolved per liter of solution. It can be mathematically expressed as
solution
of
Litre
solute
of
Moles
)
Molarity(M 
Normality is the number of gram equivalent weights of solute per liter of solution.
solution
of
Litre
solute
of
weight
Equivalent
N)
Normality( 
Relationship between Normality and Molality
The relationship between molarity and normality for the same solute in the same solution is expressed by the
following equation:
N=nM
Where, N=Normality M=Molarity and n=Equivalence Number
Equivalent mass (Eq. M.) / Gram equivalent Wight:
The equivalent mass of a substance can be calculated by dividing its molar mass by the number of active units in
one molecule of this substance thus:
solution
of
molecule
one
in
unit
active
of
No
Mass
Molecular
Mass
Euivalent 
The active unit in the acid-base reactions is the number of hydrogen ions liberated by a single molecule of acid
or reacted with a single molecule of a base.
For example, 1 mole of NaOH (m.wt. = 40) can combine with 1 mole of hydrogen ion, therefore the equivalent
weight of NaOH is 40÷1=40 gm. Thus the concentration of 1N, NaOH is the same as 1M, NaOH.
On the other hand, H2SO4 has two ionizable hydrogen atoms; its equivalent mass will be, g
49
2
98
Eq.M. 

Thus, the concentration of 1M, H2SO4 is the same as 2N, H2SO4.

1.2. Preparation of the standard solution
The standard solution can be prepared in various ways.
Method-1
1. At first, the amount of solute required to prepare the standard solution of the desired concentration is
calculated accurately.
2. Later the calculated amount of solute is measured with a high degree of precision.
Note: If the solute is a pure solid material capable of being weighed accurately, then an analytical
balance must be used to weigh it.
3. Then the measured amount of solute is transferred into a volumetric flask (which is an accurate,
high-precision, very common type of flask) of appropriate size. Often a funnel is used as an aid (If
needed).
4. Afterward, a small quantity of solvent is added to the volumetric flask and the contents are swirled
gently until the substance is completely dissolved.
5. More solvent is added until the meniscus of the liquid reaches the calibration mark on the neck of the
volumetric flask (a process called “diluting to volume”).
6. Finally, the volumetric flask is capped and inverted several times until the contents are mixed and
completely dissolved.
Disadvantages
The disadvantages of this method are as follows:
1. Some of the weighed solid may adhere to the original container, weighing paper, or funnel.
2. Solid may be spilled when it is transferred into the slim neck of the volumetric flask.
Method-2
1. At first, the amount of solute required to prepare the standard solution of the desired concentration is
calculated accurately.
2. Later the solute is measured with a high degree of precision.
Note: If the solute is a pure solid material capable of being weighed accurately, then an analytical
balance must be used to weigh it.
3. The weighed solute is first taken in a small beaker.
4. A small amount of solvent is added to the beaker and the solution is stirred until the solid is dissolved.
5. Then the solution is transferred into the volumetric flask of the intended volume. Again, a funnel may
be needed to insert into the slim neck of the volumetric flask.
Note: Before adding additional solvent to the flask, the beaker, stirring rod, and funnel must be rinsed
carefully to ensure that, all the remaining traces of the solution have been transferred.
6. Afterward, the volumetric flask is diluted to the intended volume by adding additional solvent. (Solvent
is added to the flask until the liquid level reaches the calibration mark).
7. Finally, the flask is capped and inverted until the contents are thoroughly mixed.
Disadvantages
The disadvantages of this method are as follows:
1. Some of the solutions may adhere to the beaker, stirring rod, or funnel if not washed thoroughly.
2. There is a possibility of contamination exists from the beaker, rod, or funnel if they have not been
washed carefully.

Problem: How would you prepare mL, M/N
solution using method -1 and Method-2?
Problem: How would you prepare mL, M/N
solution using method -1 and Method-2?
Chemicals/Reagents:
1. Pure solvent
2. Pure solvent
3. Distilled water
Apparatus:
1. Electrical balance
2. Pipette and pipette filler
3. Funnel
4. Volumetric flask
5. Beaker
6. Measuring cylinder
7. Spatula
8. Stirrer
9. Glue /Permanent marker for labeling
1.3.1. Calculation: Solid reagents
A. Preparation of N/M mL solution:
The molecular weight of = gram/mol.
And equivalent weight = gram/mol.
Hence, we can say that,
gm
1000
mL
gm
mL
1000






Since the purity of is = %
 
gm
100
in
present
is
gm
gm
100
in
present
is
gm
Thus,





1.3.2. Calculation: Liquid Reagents
A. Preparation of N/M mL solution:
gm
1000
mL
gm
mL
1000






Since the density of the supplied solution is kg/L = gm/mL

 
mL
100
1
of
gm
100
mL
1
is
of
gm
Hence,





On the other hand the purity of is = %
 
mL
in
present
is
gm
mL
100gm
in
present
is
gm
Thus,






1.4. Experimental procedures:
Preparation of standard solution: According to the Method-01
1. gm/mL was weigh out accurately.
2. Then it was transferred into a mL volumetric flask.
3. A small quantity of solvent was added to the volumetric flask and the content (solute) was swirled gently
until the substance was completely dissolved.
4. More solvent was added so as to reach the bottom of the solution meniscus to the calibration mark of
the volumetric flask.
5. Finally, the volumetric flask was capped and inverted several times whilst the content (solute) get mixed
and completely dissolved.
6. After complete dissolution, the final solution was labeled appropriately.
Preparation of standard solution: According to the Method-02
1. gm/mL was weigh out accurately.
2. Then it was transferred into a mL beaker.
3. A small quantity of solvent was added to the beaker and the solution was stirred until the solid was
dissolved.
4. Then the solution was transferred to a mL volumetric flask using a funnel.
5. Afterward, the beaker, stirring rod, and funnel were rinsed carefully and the washings were added to
the volumetric flask. Ensuring that all the remaining traces of the solution have been transferred.
6. Later supplementary solvent was added to the volumetric flask until the liquid level reaches the
calibration mark.
7. Finally, the flask was capped and inverted multiple times so that the contents became thoroughly
mixed.
8. The final solution was labeled appropriately.
Precautions:
1. The solute should be pure.
2. All the apparatus must be thoroughly cleaned and dry.
3. The weighed amount of solute must be completely dissolved.
4. The bottom of the meniscus of the final solution should reach the calibration mark of the volumetric
flask.

Fig-1.1: Making the solution by Method-01 [Using Methanol as solvent]
Fig-1.2: Reading the calibration mark Fig-1.3: Making the solution by Method-02
I verify that the student performed
his/her lab work following rules and
regulations.
The student performed the experiment,
collected data & completed the
Laboratory Notebook:
Lab Demonstrator
Department of Pharmacy
Northern University Bangladesh
Course instructor
Date:__________________ Date:___________________

Name of the experiment: Preparation of solutions of different concentrations by dilution.
2.1. Principle:
A solution is a homogeneous mixture created by dissolving one or more solutes in an appropriate solvent.
The chemical present in a smaller amount, known as the solute, is soluble in the solvent (the chemical present
in a larger amount). Solutions with accurately known concentrations can be referred to as standard (aka stock)
solutions. The stock solution is obtained by dissolving the calculated amount of solute in a volumetric flask of a
specific volume.
Stock solutions are frequently diluted to solutions of lesser concentration for experimental use in the laboratory.
Dilution is the addition of more solvents to produce a solution of reduced concentration. Most often a diluted
solution is created from a small volume of a more concentrated stock solution.
To make dilute a solution, an exactly calculated amount of the stock solution is taken into a clean volumetric
flask, which is then diluted to the intended volume. The amount of stock solution required for a particular
dilution is calculated from the relationship between strength and total quantity.
Relationship between Strength and Total Quantity
If the amount of solute remains the same and the volume of solution (solvent) gets larger, then the
concentration of the solution gets smaller and the solution becomes a diluted solution. Likewise, if the amount
of solute remains the same and the volume gets smaller, the concentration of the solution increases, and the
solution becomes a concentrated solution. Thus there is an inverse proportion between the strength and total
quantity (volume) of the solution.
)
1
...(
..........
..........
solution
of
Volume
1
Strength 
The strength of the solution can be expressed in several ways. The concentrations or strength of standard
solutions are normally expressed in units of moles per liter (mol/L, often abbreviated to M for molarity). Since
molarity changes in accordance with the increment or decrement of the solution volume, for a particular
amount of solute at a definite temperature, thus the equation (1) can be written as
constant)
ality
Proportion
K
Where,
(
V
1
K.
M
Or,
V
1
M
Or,
solution
of
Volume
1
Molarity




Or, MV=K
If M1, V1 are the initial molarity and volume of a given stock solution and M2, V2 the changed molarity and
volume for the intended dilute solution, we can write;
M1V1 = k = M2V2
Or, M1V1 = M2V2------------------------------------------------------(2)
By applying equation (2), a dilute solution can be prepared from a concentrated solution.
In doing so, the targeted concentration and volume are assumed to be M2 & V2 and the concentration and
volume of the stock solution be M1 & V1 respectively. Later on, based on the intended dilute concentration,
M2 and volume, V2 the required volume of the stock solution are calculated. The calculated volume of the stock
is taken into a clean volumetric flask, which is then diluted to the intended volume.

Problem-01: How would you prepare mL, N/M solution
from N/M solution using dilution method?
Problem-02: How would you prepare mL, N/M solution
from N/M solution using dilution method?
Chemicals/Reagents:
1. Pure Solute (A.) and (B.) .
2. Stock Solution (A.) N/M mL solution.
and (B.) N/M mL solution.
3. Distilled Water
Apparatus:
3. Funnel
4. Volumetric flask
5. Beaker
7. Spatula
8. Stirrer
9. Glue /Permanent marker for labeling
2.2. Preparation of Stock solution
2.2.1. A: Solid reagents
Preparation of N/M mL solution:
gm
1000
mL
gm
mL
1000






Since the purity of is = %
 
gm
100
in
present
is
gm
gm
100
in
present
is
gm
Thus,





2.2.2. B: Liquid Reagents
Preparation of N/M mL solution:
gm
1000
mL
gm
mL
1000







 
mL
100
1
of
gm
100
mL
1
is
of
gm
Hence,





On the other hand the purity of is = %
 
mL
in
present
is
gm
mL
100gm
in
present
is
gm
Thus,






The stock solution was prepared with the calculated amounts of solute using the method(s) as per
Experiment No.-01
2.3. Preparation of N/M mL solution from the stock solution.
Preparation of N/M mL solution from the stock solution.
2.3.1. Calculation: Solid Reagents
We know that,
M1V1 = M2V2
M
V
M
V
Or,
1
2
2
1



Or, V1= mL
Here,
The volume of diluted Solution (Final volume), V2 = mL
Strength of diluted solution (Final Strength), M2 = N/M
Volume of Stock Solution (Volume to be taken), V1 = ? mL
Strength of Stock solution, M1 = N/M
Result: The required volume of Stock Solution = mL.
2.3.2. Calculation: Liquid Reagents
We know that,
M1V1 = M2V2
M
V
M
V
Or,
1
2
2
1



Or, V1= mL
Here,
The volume of diluted Solution (Final volume), V2 = mL
Strength of diluted solution (Final Strength), M2 = N/M
Volume of Stock Solution (Volume to be taken), V1 = ? mL
Strength of Stock solution, M1 = N/M
Result: The required volume of Stock Solution = mL.
2.4. Experimental procedures:
Dilution procedures-01:
1. At first mL clean and dry volumetric flask(s) were taken.
2. Later mL stock solutions were pipette out with a high degree of precision.
3. Then the measured volume was transferred into the mL volumetric flasks
using a funnel.
4. Afterward, a small quantity of solvent (DW) was added to the volumetric flask and swirled gently.
5. More solvent was added until the meniscus of the liquid reaches the calibration mark of the volumetric
flask.
6. Finally, the volumetric flask was capped and inverted several times to get a completely mixed
homogenous solution.

Dilution procedures-02:
1. At first mL clean and dry volumetric flask(s) were taken.
2. Later mL stock solutions were pipette out with a high degree of precision.
3. Then the measured volume was transferred into the mL volumetric flasks
using a funnel.
4. Afterward, a small quantity of solvent (DW) was added to the volumetric flask and swirled gently.
5. More solvent was added until the meniscus of the liquid reaches the calibration mark of the volumetric
flask.
6. Finally, the volumetric flask was capped and inverted several times to get a completely mixed
homogenous solution.
Precautions:
1. All solutes should be pure.
2. All the apparatus must be thoroughly cleaned and dry.
3. The weighed amount of Sodium carbonate must be completely dissolved.
4. The bottom of the meniscus for the final solution should reach the calibration mark of the
volumetric flask.
5. DW should be added slowly as the meniscus reaches the calibration mark in the volumetric flask.
6. The final solution must be shaken well to get a uniform solution
Fig-2.1.: Preparation of stock solution Fig-2.2.: Dilution of a solution
regulations.
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________

Name of the experiment: Standardization of NaOH solution by Potassium Hydrogen
Phthalate (KHP).
Introduction
Standardization refers to an experiment in which the concentration of a solution is determined to the
desired accuracy. Titration is one type of analytical procedure often used in standardization. In titration, an
exact volume of one substance is reacted with a known amount of another substance.
The main difference between standardization and titration is that standardization processes essentially
use primary standard solutions whereas titrations do not essentially use primary standard solutions.
Terminologies Used in Titrimetric Analysis
The term 'titrimetric analysis' refers to quantitative chemical analysis. It is carried out by determining
the volume of a solution of accurately known concentration, required to react quantitatively with a measured
volume of a solution of the substance to be determined. The solution of accurately known strength is called the
standard solution.
In titrimetric analysis, the reagent of known concentration is called the titrant and the substance being
titrated is termed the titrand or analyte.
The process of adding the standard solution until the reaction is just complete is termed a titration, and
the substance to be determined is titrated. The point at which this occurs is called the equivalence point or the
theoretical (or stoichiometric) end point.
The completion of the titration is detected more usually, by the addition of an auxiliary reagent, known
as an indicator. After the reaction between the substance and the standard solution is practically complete, the
indicator should give a clear visual change (either a color change or the formation of turbidity) in the liquid
being titrated.
The point at which this occurs is called the end point of the titration.
Principle
In titrimetry, chemicals used to make reference solutions are two types: namely primary standards and
secondary standards.
Primary Standards
A primary standard is a compound of sufficient purity from which a standard solution can be prepared by
direct weighing of a quantity of it, followed by dilution to give a defined volume of solution.
The solution produced is called a primary standard solution.
A primary standard should satisfy the following requirements.
1. It must be easy to obtain, purify, dry (preferably at 110°C-120°C), and preserve in a pure.
(This requirement is not usually met by hydrated substances, since it is difficult to remove surface moisture completely without effecting partial
decomposition.) (The total amount of impurities should not, in general, exceed 0.01-0.02%)
2. It should be unaltered in the air during weighing; that means, it should not be hygroscopic, oxidized by air, or
affected by carbon dioxide.
3. It should be capable of being tested for impurities by qualitative and other tests of known sensitivity.
4. It should be readily soluble under the conditions in which it is employed.
5. The reaction with the standard solution should be stoichiometric and practically instantaneous.
6. The titration error should be negligible, or easy to determine accurately by experiment.
In practice, an ideal primary standard is difficult to obtain, and a compromise between the above ideal
requirements is usually necessary.

The substances commonly employed as primary standards are indicated below:
(a) Acid-base reactions: Sodium carbonate (Na2CO3); sodium tetraborate (Na2B4O7),
potassium hydrogen phthalate (C8H5O4K), potassium hydrogen iodate KH(IO3)2, benzoic acid (C6H5COOH)
etc.
(b) Complex formation reactions: Silver, Silver nitrate (AgNO3), Sodium chloride (NaCl), various metals
(e.g. Spectroscopically pure Zinc, Mg, Copper, and Manganese) and salts, depending upon the reaction used.
(c) Precipitation reactions: Silver, Silver nitrate (AgNO3), Sodium chloride(NaCl), Potassium chloride (KCl), and
Potassium bromide(KBr) (prepared from potassium bromate).
(d) Oxidation-reduction reactions: Potassium dichromate (K2Cr2O7), Potassium bromate (KBrO3),
Potassium iodate (KIO3), Potassium hydrogen iodate KH(IO3)2, Sodium oxalate (Na2C2O4),
Arsenic(III) oxide (As2O3), and pure Iron.
Secondary Standards
Solutions that are prepared by standardization against a primary standard are referred
to as secondary standards.
It follows that a secondary standard solution is a solution in which the concentration of
dissolved solute has not been determined from the weight of the compound dissolved, but by the reaction
(titration) of a volume of the solution, against a measured volume of a primary standard solution.
Examples: Sodium tetraborate Na2B4O7.10H2O , Copper sulfate Cu2SO4.5H2O etc.
In the present experiment, NaOH solution is to be used secondary standard chemical and needs to be
standardized by the primary standard chemicals. Here the primary standard chemical is KHP.
This is a 1:1 titration therefore, one mole of the base, will titrate one mole of acid.
The end point of the solution would be determined by an indicator, phenolphthalein.
Reaction:
NaOH(aq) + KHC8H4O4(aq) = KNaC8H4O4(aq) + H2O(l)
Indicator
Phenolphthalein: Phenolphthalein exists in two tautomeric forms. Phenolphthalein has benziod form in an
acidic medium and thus, it is colorless while it has quinonoid form in an alkaline medium which has a pink
color.
Fig 3.1. Two tautomer’s of Phenolphthalein.

3.1. Chemicals / Reagents:
1. 0.1 N Sodium hydroxide (NaOH) solution
2. 0.1 N Potassium hydrogen phthalate (C8H5O4K)
3. 0.5% Phenolphthalein indicator
4. Distilled Water
3.2. Apparatus:
2. Spatula
3. Volumetric flask
5. Burette
7. Conical flask
8. Beaker
9. Funnel
3.3. Reagents and their preparations:
A. N/M mL Potassium Hydrogen Phthalate solution:
The molecular weight of Potassium hydrogen phthalate (KHP) is gram/mol.
The equivalent weight of KHP = gram/mol.
Hence, we can say that:
gm
1000
(KHP)
phthalate
hydrogen
Potassium
N/M
mL
KHP
gm
(KHP)
phthalate
hydrogen
Potassium
N/M
mL
1000






Since Potassium hydrogen phthalate is %pure
 
gm
100
in
present
is
phthalate
hydrogen
Potassium
gm
KHP
gm
100
in
present
is
phthalate
hydrogen
Potassium
gm
Thus,





Preparation
1. gm KHP was weighed out in a mL volumetric flask;
2. Then it was completely dissolved with a small portion of DW.
3. Finally the volume was adjusted to mL with DW Q.S.
B. N/M mL NaOH solution:
The molecular weight of Sodium Hydroxide (NaOH) is gram/mol.
The equivalent weight of NaOH = gram/mol.
NaOH
gm
1000
Hydroxide
Sodium
N/M
mL
NaOH
m
g
Hydroxide
Sodium
N/M
mL
1000






Since Sodium Hydroxide is %pure
 
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
Thus,





Preparation
1. gm NaOH was weighed out in a mL volumetric flask.

C. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 gm Phenolphthalein+ 50 mL DW.
Standardization of NaOH solution by Potassium Hydrogen Phthalate (KHP) solution
3.4. Experimental Procedures:
1. A clean and dry burette was taken.
2. It was filled with the prepared N/M Solution.
3. The initial volume of solution in the burette (IBR) was recorded.
4. Then 10 mL N/M solution was pipetted out into a conical flask with great
accuracy.
5. 01 drop of phenolphthalein (HPh) indicator was added to it.
6. Afterward, the burette's stopcock was opened and the standard solution was allowed to run down into the
conical (Erlenmeyer) flask.
7. The conical flask was swirled while adding solution from the burette.
8. The stopcock was closed at the very close to the end point, and then the stopcock was only opened
enough to allow one drop at a time to fall from the burette.
9. The stopcock was closed as soon as the reaction reached its equivalence point when the indicator
permanently changed its color.
10. The final volume of alkali used to neutralize the acid (FBR) was recorded (as shown on the burette).
11. The titration was repeated two times and the titer was calculated.
Table-3.1: Data for the standardization of NaOH solution:
No. of
observations
Volume of
.
(V1 mL)
Volume of solution (mL) Difference
(FBR-IBR)
(mL)
Mean
volume
(V2 mL)
IBR FBR
1 10
2 10
3 10
3.5. Calculation of the strength of NaOH solution:
We know that,
V[ ]S[ ]= V[ ]S[ ]
 
   
 
 
  N/M
S
Or,
N/M
S
Or,
V
S
V
S
Or,




Here,
Volume of , V1 = 10 mL
Strength of , S1 =
Volume of , V2 = mL (Mean)
Strength of , S2=
3.6. Result:
The standardized concentration of the prepared NaOH solution was N/M.

Fig-3.1.: Acid-base Titration
Precautions:
1. All weights must be taken carefully. Particularly, the weight of KHP must be precise to
ensure precise concentration.
2. Only 1 drop of the indicator solution should be used.
3. All volumetric readings (e.g. of the volumetric flask and the burette) must be taken
considering the meniscus.
4. Colour change must be monitored very carefully.
regulations.
Laboratory notebook:
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________

Name of the experiment: Standardization of NaOH solution by Oxalic acid Solution.
Principle
An acid-base titration involves the addition of a titrant solution to an analyte solution. In titrimetry,
chemicals used as reference solutions are known as primary standards or secondary standards.
A primary standard is a reagent that is extremely pure, stable, has no water of hydration, and has a
high molecular weight. A secondary standard solution refers to a solution that has its concentration measured
by titration with a primary standard solution.
The secondary standard chemicals are not absolutely pure and may contain some impurities that come
during their synthesis. Therefore the secondary standard chemical should standardize by using the primary pure
chemicals.
standardized by the primary standard chemicals. Here the primary standard chemical is Oxalic acid. The end
point of the solution will be determined by an indicator, phenolphthalein.
Reaction:
1. H2C2O4 (aq) + 2NaOH (aq) ⇌ Na2C2O4 (aq) + 2H2O (l)
4.1. Chemicals / Reagents:
2. 0.1 N Oxalic acid (dihydrate) [H2C2O4.2H2O]
4. Distilled Water
4.2. Apparatus:
2. Spatula
3. Volumetric flask
5. Burette
7. Conical flask
8. Beaker
9. Funnel
4.3. Reagents and their preparations:
A. N/M mL Oxalic acid solution:
The molecular weight of Oxalic acid (H2C2O4.2H2O) is gram/mol.
The equivalent weight of Oxalic acid is (126.08 ÷ 2) gm = gram/mol.
O
.2H
O
C
H
gm
1000
acid
Oxalic
N/M
mL
O
.2H
O
C
H
gm
acid
Oxalic
N/M
mL
1000
2
4
2
2
2
4
2
2






Since Oxalic acid is %pure
 
O
.2H
O
C
H
gm
100
in
present
is
acid
Oxalic
gm
O
.2H
O
C
H
gm
100
in
present
is
acid
Oxalic
gm
Thus,
2
4
2
2
2
4
2
2






Preparation
1. gm Oxalic acid was weighed out in a mL volumetric flask;
The equivalent weight of Sodium Hydroxide (NaOH) = 40 gram/mol.
NaOH
gm
1000
40
Hydroxide
Sodium
N/M
mL
NaOH
gm
40
Hydroxide
Sodium
1N
mL
1000






 
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
Thus,





Preparation
2. Then it was completely dissolved with small portion of DW.
B. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 gm Phenolphthalein+ 50 mL DW.
Standardization of NaOH solution by Oxalic acid solution
4.4. Experimental Procedures:
2. It was filled with the prepared N/M solution.
3. The initial volume of solution in the burette (IBR) was recorded. (The Initial Burette Reading).
4. Then 10 mL N/M solution was pipetted out into a conical flask with great
accuracy.
5. 01 drop of phenolphthalein (HPh) indicator was added to it.
9. The stopcock was closed as soon as the reaction reached its equivalence point when the indicator
permanently changed its color.
10. The final volume of alkali used to neutralize the acid (FBR) was recorded (The final burette reading).
11. The titration was repeated two times and the titer was calculated.

Table-4.1.: Data for the standardization of NaOH solution:
No. of
observations
Volume of
.
(V1 mL)
Volume of solution (mL) Difference
(FBR-IBR)
(mL)
Mean
volume
(V2 mL)
IBR FBR
1 10
2 10
3 10
4.5. Calculation of the strength of NaOH solution:
We know that,
V[ ]S[ ]= V[ ]S[ ]
 
   
 
 
  N/M
S
Or,
N/M
S
Or,
V
S
V
S
Or,




Here,
Strength of , S1 =
Strength of , S2=
4.6. Result:
The standardized concentration of the prepared NaOH solution was N/M.
Precautions:
1. Run a little amount of solution from the burette through the tap to make sure that no air is trapped.
2. Ensure that the bottom of the meniscus is level with the zero line at the top of the burette.
3. Always read the volume at eye level and from the bottom of the meniscus.
4. Don’t overshoot the end point, so don't try to get too close on this first careful titration. Partially open
the stopcock so that solution falls drop by drop into the conical flask. Continue swirling the flask.
5. When you are very close to the end point, a drop of solution from the burette will change the colour
(pink) of the solution, but when you swirl the flask, this pink colour will disappear.
6. The end point of the titration has been reached when 1 drop changes the colour of the solution in the
conical flask, and, swirling the flask does not cause this colour to disappear. (Referred to as a "permanent
colour change")
regulations.
Laboratory notebook:
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________

Name of the experiment: Standardization of HCl Solution by NaOH Solution.
Principle
Titration (aka titrimetry) is a common laboratory technique used for quantitative chemical analysis.
Titration is a process of neutralization whereby a titrant (a solution of known concentration) is added into an
analyte (a solution of unknown concentration) until the unknown solution is completely neutralized.
Because volume measurement plays a key role in titration, it is also known as volumetric analysis.
An indicator can signal the completion of the reaction (by colour change) and the concentration of the unknown
solution can be determined.
Any chemicals that react in a solution can be titrated with each other. Since acids and bases are usually found
in solution, they are commonly involved in titrations. Titrations of a strong acid or a strong base involve the
neutralization reaction between hydrogen ions and hydroxyl ions. These ions combine to form the neutral
water molecule: H2O⇌ H+ + OH–
The point when the titration is theoretically stopped, where the number of moles of titrant is equal to the
number of moles of analyte, or some multiple thereof (as in di- or tri- protic acids).
An acid-base indicator is an organic dye that signals the end-point by a visual change in colour. It changes
colour over a certain pH range and indicates the end of the titration.
standardized by the primary standard chemicals. Here the primary standard chemical is Oxalic acid. After being
standardized by the Oxalic acid, NaOH is to be used for the standardization of HCl.
The end point of the solution will be determined by an indicator, phenolphthalein.
Reaction:
1. H2C2O4 (aq) + 2NaOH (aq) ⇌ Na2C2O4 (aq) + 2H2O (l)
2. HCl (aq) + NaOH (aq) ⇌H2O (l) + NaCl (aq)
5.1.Apparatus:
2. Spatula
3. Volumetric flask
5. Burette
7. Conical flask
8. Beaker
9. Funnel
5.2.Chemicals / Reagents:
2. 0.1 N Oxalic acid (dihydrate) [H2C2O4.2H2O]
3. 0.1 N Hydrochloric acid (HCl) solution
5. Distilled Water

5.3.Reagents and their preparations:
A. N/M mL Oxalic acid solution:
The molecular weight of Oxalic acid (H2C2O4.2H2O) is 126.08 gram/mol.
The equivalent weight of Oxalic acid is (126.08 ÷ 2) gm = 63.040 gram/mol.
O
.2H
O
C
H
gm
1000
acid
Oxalic
N/M
mL
O
.2H
O
C
H
gm
acid
Oxalic
1N/M
mL
1000
2
4
2
2
2
4
2
2






Since Oxalic acid is %pure
 
O
.2H
O
C
H
gm
100
in
present
is
acid
Oxalic
gm
O
.2H
O
C
H
gm
100
in
present
is
acid
Oxalic
gm
Thus,
2
4
2
2
2
4
2
2





Preparation
1. gm Oxalic acid was weighed out in a mL volumetric flask;
The molecular weight of Sodium Hydroxide (NaOH) = gram/mol.
The equivalent weight of Sodium Hydroxide = gram
NaOH
gm
1000
Hydroxide
Sodium
N/M
mL
NaOH
gm
Hydroxide
Sodium
1N/M
mL
1000






 
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
Thus,





Preparation
C. 0.1 N mL HCl solution:
The molecular weight of HCl = 36.46 gram/mol.
And equivalent weight =(36.46÷1)=36.46 gram/mol.
HCl
gm
1000
36.46
0.1
Acid
ic
Hydrochlor
0.1N
mL
HCl
gm
46
.
36
Acid
ic
Hydrochlor
1N
mL
1000






Since the density of the supplied HCl= kg/L = gm/mL

 
mL
100
1
acid
ic
Hydrochlor
of
gm
100
mL
1
is
acid
ic
Hydrochlor
of
gm
Hence,





On the other hand the purity of HCl is = %
 
HCl
mL
100
in
present
is
acid
ic
Hydrochlor
gm
HCl
mL
100gm
in
present
is
HCl
gm
Thus,






D. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 gm Phenolphthalein+ 50 mL DW.
5.4.1. Standardization of NaOH solution by Oxalic acid solution
5.4.1. 1. The main points are outlined below
a) The burette was filled with the prepared 0.1 N NaOH solution.
b) 10 mL standard 0.1N Oxalic acid solution was taken in a conical flask.
c) The initial volume of solution in the burette (IBR) was recorded. (The Initial Burette Reading).
d) 1 drop Phenolphthalein indicator was added to the acid solution and titrated with the NaOH solution,
until the colour of the acid solution changed from colourless to faint pink.
e) The stopcock was closed as soon as the reaction reached its equivalence point, when the indicator
permanently changed its colour.
f) The final volume of alkali used to neutralize the acid (FBR) was recorded (The final burette reading).
g) Another two titrations were performed and the result was calculated.
Table-5.1.: Data for the standardization of NaOH solution:
No. of
observations
The volume of
Oxalic acid
(V1 mL)
Volume of NaOH solution (mL) Difference
(FBR-IBR)
(mL)
Mean
volume
(V2 mL)
IBR FBR
1 10
2 10
3 10
5.4.1.2.Calculation of the strength of NaOH solution:
We know that,
V[ ]S[ ]= V[ ]S[ ]
 
   
 
 
  N/M
S
Or,
N/M
S
Or,
V
S
V
S
Or,




Here,
Strength of , S1 =
Strength of , S2=
5.4.1.3. Result:
The standardized concentration of the prepared sodium hydroxide solution was N.

5.4.2. Standardization of HCl solution by NaOH solution
5.4.2.1. Experimental Procedures:
2. It was filled with the prepared 0.1 N NaOH solution.
3. The initial volume of solution in the burette (IBR) was recorded. (The Initial Burette Reading).
4. Then 10 mL 0.1N HCl acid solution was pipetted out into a conical flask with great accuracy.
5. 01 drop phenolphthalein (HPh) indicator was added to it.
9. The stopcock was closed as soon as the reaction reached its equivalence point, when the indicator
permanently changed its colour.
10. The final volume of alkali used to neutralize the acid (FBR) was recorded (The final burette reading).
11. The titration was repeated two times and the titre was calculated.
Table-5.2.: Data for standardization of HCl solution:
No. of
observations
Volume of HCl
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(FBR-IBR)
(mL)
Mean volume
(V2 mL)
IBR FBR
1 10
2 10
3 10
5.4.2.2. Calculation of the strength of HCl solution:
We know that,
V[ ]S[ ]= V[ ]S[ ]
 
   
 
 
  N/M
S
Or,
N/M
S
Or,
V
S
V
S
Or,




Here,
Strength of , S1 =
Strength of , S2=
5.4.2.3. Result:
The standardized concentration of the prepared hydrochloric acid solution was N.

Precautions:
1. Rapid addition of NaOH from the burette is strictly prohibited.
4. HCl is corrosive. Be careful while dealing with acids.
5. NaOH is corrosive. Be careful while dealing with strong bases.
6. The difference of volume of NaOH between two successive titrations should not be more than
0.1 mL.
7. The conical Flask should be continuously swirled while adding the NaOH from the burette.
8. Place a piece of white paper below the conical flask during titration to aid the determination of
the colour change at the end point.
regulations.
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________
Fig. 5.1. Titration process.

Name of the experiment: Preparation of constant pH buffer and assessment of its activity.
Principle:
A solution that resists the changes in pH, when small quantities of an acid or an alkali are added to it, is called a
buffer solution. Most of the buffer solutions usually consist of a mixture of a weak acid and one of its salts or a
weak base and one of its salts. The resistance to a change in pH is known as buffer action.
A buffer solution contains large and equal concentrations of an acid and its conjugate base.
The pH of this solution is approximately equal to the pKa of the acid. The addition of a small amount of acid or
base results in the mopping up or the release of protons by the conjugate base or the acid is necessary, which
keeps the solution pH constant.
A solution of acetic acid and sodium acetate (CH3COOH + CH3COONa) (pH range is 3.7-5.6) is an example
of a buffer that consists of a weak acid and its salt.
How does buffer work
A mixture of acetic acid and sodium acetate
is acidic because the Ka of acetic acid is
greater than the Kb of its conjugate base
acetate. It is a buffer because it contains
both the weak acid and its salt. Hence, it
acts to keep the hydronium ion
concentration (and the pH) almost constant
by the addition of either a small amount of a
strong acid or a strong base. Figure 6.1. Mechanism of Action of Acetate Buffer.
If we add a base such as sodium hydroxide, the hydroxide ions react with the few hydronium ions present.
Then more of the acetic acid reacts with water, restoring the hydronium ion concentration almost to its original
value:
The pH changes are very little. If we add an acid such as hydrochloric acid, most of the hydronium ions from
the hydrochloric acid combine with acetate ions, forming acetic acid molecules: Thus, there is very little
increase in the concentration of the hydronium ion, and the pH remains practically unchanged.
TYPES OF BUFFERS
Buffers are classified into the following types
1. Simple Buffers:
Simple buffers are categorized into three different ways.
A) Salts of a weak acid and a weak base: CH3COONH4, NH4CN.
B) Proteins and amino acids.
C) A mixture of acid salt and a normal salt formed from polybasic acid: Na2HPO4 and Na3PO4
2. Mixed Buffers
Acidic Buffer: A solution of a weak acid and its salt with a strong base.
Example: CH3COOH +CH3COONa, H2CO3 +Na2CO3
Basic Buffer: A solution of a weak base and its salt with strong acid.
Example: NH4OH+NH4Cl, NH4OH+NH4NO3
3. Natural Buffers
A solution is said to be naturally buffered if it contains buffering compounds as it exists in nature.
Blood is an example of a naturally buffered solution.

Apparatus:
2. Spatula
3. Volumetric flask
4. Conical flask
5. Beaker
7. pH meter
8. Stirrer
Chemicals/Reagents:
1. 0.1 M Na-Acetate (CH3COONa) solution
2. 0.1M Acetic Acid solution
3. 0.1 M HCl Solution
4. 0.1 M NH4OH Solution
5. Distilled Water
Reagents and their preparations:
A. 0.1 M mL Acetic acid solution:
The molecular weight of Acetic acid (CH₃COOH) is= gram/mol.
COOH
CH
gm
1000
0.1
acid
Acetic
0.1M
mL
COOH
CH
gm
acid
Acetic
1M
mL
1000
3
3






As the density of the supplied solution is kg/L = gm/mL
 
COOH
CH
mL
100
1
acid
Acetic
of
gm
100
COOH
CH
mL
1
acid
Acetic
of
gm
Hence,
3
3





Since Acetic acid is %pure
 
COOH
CH
mL
in
present
is
acid
Acetic
gm
COOH
CH
mL
100gm
in
present
is
acid
Acetic
gm
Thus,
3
3






Preparation
1. mL Acetic acid was weighed out in a mL volumetric flask.
B. 0.1 M mL Na-Acetate (CH3COONa) solution:
The molecular weight of Sodium Acetate (CH3COONa) is = gram/mol.
gm
1000
0.1
Acetate
Sodium
0.1M
mL
COONa
CH
gm
Acetate
Sodium
1M
mL
1000 3






On the other hand, Sodium Acetate is= %pure

 
COONa
CH
gm
100
in
present
is
COONa
CH
gm
COONa
CH
gm
100
in
present
is
COONa
CH
gm
Thus,
3
3
3
3





Preparation
1. gm Sodium Acetate was weighed out in a mL volumetric flask;
C. 0.1 M mL HCl solution:
HCl
gm
1000
0.1
Acid
ic
Hydrochlor
0.1M
mL
HCl
gm
Acid
ic
Hydrochlor
1M
mL
1000






Since the density of the supplied HCl= kg/L = gm/mL
 
mL
100
1
acid
ic
Hydrochlor
of
gm
100
mL
1
is
acid
ic
Hydrochlor
of
gm
Hence,





On the other hand, the purity of HCl is = %
 
HCl
mL
in
present
is
HCl
gm
HCl
mL
100gm
in
present
is
HCl
gm
Thus,






Preparation
1. mL HCl was weighed out in a mL volumetric flask.
D. 0.1 M mL NH4OH solution: (From NH3 solution)
The molecular weight NH3 solution= gram/mol.
gm
1000
0.1
solution
Ammonia
0.1M
mL
NH
gm
solution
Ammonia
1M
mL
1000 3






Since the density of the supplied NH3= kg/L = gm/mL
 
mL
100
1
solution
Ammonia
of
gm
100
mL
1
is
solution
Ammonia
of
gm
Hence,





On the other hand, the purity of NH3 is = %

  3
3
3
3
mLNH
in
present
is
NH
gm
NH
mL
100gm
in
present
is
NH
gm
Thus,






Preparation
1. mL Ammonia was weighed out in a mL volumetric flask;
3. Finally the volume was adjusted to mL with DW Q.S
Calibration of pH meter:
1. Rinse the pH meter electrode with distilled water. Rinse and pat dry with a lint-free tissue.
2. Place the electrode in the appropriate standard buffer.
3. Take pH meter reading.
4. Accurate reading of the standard buffer indicates the pH meter is OK.
Experimental Procedures:
A. Preparation of 200 mL buffer solution with pH 4.6
1. 102 mL 0.1 M CH3COOH solution was taken in a 250 mL beaker.
2. Then 98 mL 0.1M sodium acetate solution was added to it slowly with continuous stirring.
3. The detector of the pH meter was submerged carefully into the solution.
4. Simultaneously the pH reading was taken while adding CH3COONa.
5. Afterward, the pH of the final solution was adjusted to 4.6 using HCl and/or NaOH with
continuous stirring by the stirrer.
B. Assessment of buffer action of the prepared buffer solution:
Addition of small amount of weak acid:
1. 50 mL of the prepared buffer solution was taken in a 100 mL beaker.
2. Then 2 mL 0.1M HCl solution was added to it slowly with continuous stirring.
3. The detector of the pH meter was submerged carefully into the solution.
4. Afterward, the pH reading was taken simultaneously while adding HCl Solution.
Addition of small amount of weak base:
1. 50 mL of the prepared buffer solution was taken in a 100 mL beaker.
2. Then 2 mL 0.1M NH4OH solution was added to it slowly with continuous stirring
3. The detector of the pH meter was submerged carefully into the solution
4. Afterward, the pH reading was taken simultaneously while adding NH4OH Solution.

Table-6.1: Data for assessment of buffer capacity of the solution:
Volume of Buffer
solution (mL)
Volume of acid and base Solution added
(mL)
pH of final solution
(After addition of the acid/Base)
50
0.1M HCl
0.1 M NH4OH
DW
Result:
The pH of the acetate buffer :
The pH of the acetate buffer after adding acid :
The pH of the acetate buffer after the addition of base :
Inference: The pH of the buffer solution did not change after the addition of a small amount of acid or base.
Precautions:
1. The weighing should be done with high accuracy.
2. The calibration mark in the volumetric flask should not exceed while preparing the solution.
3. Pipette should be used to get accurate volume.
4. The pH meter must be calibrated each time before use.
Figure 6.2. (a) The buffered solution on the left and the un-buffered solution on the right have the same pH (pH 8);
they are basic, showing the yellow colour of the indicator methyl orange at this pH.
(b) After the addition of 1 mL of a 0.01-M HCl solution, the buffered solution has not detectably changed its pH but the
un-buffered solution has become acidic, as indicated by the change in colour of the methyl orange, which turns red
at a pH of about 4. (credit: modification of work by Mark Ott)
regulations.
Laboratory Notebook.
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________

Name of the experiment: Determination of the acid dissociation constant (Ka) of a weak acid
(acetic acid).
Principle
Based on the degree of dissociation and release of protons into solution, acids are divided into two
types – strong acids and weak acids. Strong acids completely dissociate in solution and release protons, but
weak acids only partially dissociate. For weak acids, the dissociation of acid to release proton reaches
equilibrium as shown below:
HA⇌H++A—
The acid dissociation constant (Ka) can be defined as: ..(i)
..........
[HA]
]
][A
[H
Ka
_


Henderson-Hasselbach equation describes the relationship between the pH of a solution made of a weak acid
and its acid dissociation constant as : ..(ii)
..........
[HA]
]
[A
log
pKa
P
_
H


Here, [HA] refers to the concentration of the weak acid at equilibrium, and [A ] refers to the concentration of
the conjugate base at equilibrium, pH must also be measured at the equilibrium. Now, when[ ] [ ],
equation (2) becomes:
pH=pKa
Now, if a titration curve is prepared by plotting pH along the y-axis and volume of titrant added along the x-
axis, the half-equivalence point is the point where one-half of the end-point volume has been added. At the
half-equivalence point, [ ] [ ], so the pH at the half-equivalence point is the pKa of the weak acid. From
pKa, Ka can be calculated as follows:
Ka=10–pKa
Apparatus:
2. Spatula
3. Volumetric flask
5. Burette
7. Conical flask
8. Beaker
9. Funnel
10.pH meter
11.Dropper
Chemicals / Reagents:
2. 0.1N Acetic acid (CH3COOH) solution
4. Distilled Water

Reagents and their preparations:
B. 0.1 N mL Acetic acid solution:
The molecular weight of Acetic acid (CH₃COOH) is= gram/mol.
The equivalent weight of Acetic acid is = gram/mol.
COOH
CH
gm
1000
0.1
acid
Acetic
0.1N
mL
COOH
CH
gm
acid
Acetic
1N
mL
1000
3
3






 
COOH
CH
mL
100
1
acid
Acetic
of
gm
100
COOH
CH
mL
1
acid
Acetic
of
gm
Hence,
3
3





Since Acetic acid is %pure
 
COOH
CH
mL
in
present
is
acid
Acetic
gm
COOH
CH
mL
100gm
in
present
is
acid
Acetic
gm
Thus,
3
3






Preparation
1. mL Acetic acid was weighed out in a mL volumetric flask;
A. 0.1 N mL NaOH solution:
The equivalent weight of Sodium Hydroxide (NaOH) = 40 gram/mol.
NaOH
gm
1000
40
0.1
Hydroxide
Sodium
0.1N
mL
NaOH
gm
40
Hydroxide
Sodium
1N
mL
1000






 
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
NaOH
gm
100
in
present
is
Hydroxide
Sodium
gm
Thus,





Preparation
1. gm Sodium Hydroxide (NaOH) was weighed out in a mL volumetric flask.
D. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 gm Phenolphthalein+ 50 mL DW.

Experimental Procedures:
1. A clean and dry burette was taken and filled with the prepared 0.1 N NaOH solution.
2. Afterwards 25 mL of the acetic acid solution was taken into a beaker.
3. Latter, 2 drops of phenolphthalein (PhI) indicator was added into it.
4. A pH meter was placed into the beaker and the pH of the solution was noted down.
5. Then the burette's stopcock was opened and the solution was allowed to run down into the beaker.
6. After the addition of each 1 mL of the titrant, the pH was measured and noted down.
8. The process was continued until the end point was achieved i.e. colour changed from colourless to pink.
9. 04 more readings were taken after achieving the end-point.
Table-7.1.: Titration curve data:
SL. Cumulative volume of titrant added (mL) pH
01
02
03
04
05
06
07
08
09
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35

Preparation of the titration curve
Plot the pH values along the y-axis and the cumulative volume values along the x-axis in MS Excel and connect
the points by smooth curve (as opposed to a straight line).
Fig.7.1. Titration curve
Calculation of Ka
Find the end-point volume. The end-point volume i.e. titration volume is _____________ mL.
So, the half-equivalence point is ___________ mL.
The pH that corresponds to the volume of the half-equivalence point was found out.
The pH of the half-equivalence point = __ ______
So, PKa of Acetic acid= pH of the half-equivalence point= ___ _____
Now, Ka=10–pKa =10– =
Result
The acid dissociation constant (Ka) of acetic acid was found to be ___________ __________.
Precautions:
1. All weights must be taken carefully.
2. pH must be carefully monitored
3. Please handle the pH meter with utmost care.

regulations.
Lab Demonstrator
Course instructor
Date:__________________ Date:___________________
Bibliography
1. Jeffery, G. H., Bassett, J., Mendham, J., & Denney, R. C. (1989). Vogel's Textbook of
Quantitative Chemical Analysis, 5th edn. Longman Scientific and Technical.
2. Kenkel, J. (1992). Analytical Chemistry refresher manual. CRC Press.
3. Kenkel, J. (2002). Analytical chemistry for technicians. CRC Press.
4. Arun Bahl, Bahl B.S, and Tuli G.D., (2011). Essentials of Physical Chemistry,
Multi colour edn. New Delhi, S.Chand and Company Limited.
5. Cairns, D. (Ed.). (2008). Essentials of pharmaceutical chemistry. 3rd edn.
Pharmaceutical Press, UK.
Md.
Imran
Nur
Manik

Physical Pharmacy-I Lab, Manik

Physical Pharmacy-I Lab, Manik

More Related Content

What's hot

Similar to Physical Pharmacy-I Lab, Manik

More from Imran Nur Manik

Recently uploaded

Physical Pharmacy-I Lab, Manik