Physical Pharmacy-I lab + Data Sheet

Physical Pharmacy-I lab + Data Sheet
1. Standardization of Acids and bases.
2. Determination of pKa and pKb values
3. Preparation of solutions of different pH & buffer capacities.
4. Determination of phase diagram of binary systems.
5. Determination of distribution coefficients.
6. Determination of molecular weight by Victor Meyer’s Method.
7. Determination of heats of solutions by measuring solubility as a function of
temperature (Van’t Hoff equation.)

1. Laboratory Manual
Course Code: PHRM 1206



Prepared By
Md. Imran Nur Manik
Lecturer
Department of Pharmacy
Northern University Bangladesh

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Sl. No. Date Name of the experiment Page No.
01.
Preparation of Solutions of Different Concentrations.
01 – 03
02.
Preparation of Solutions of Different Concentrations by
Dilution Method.
04 – 05
03.
Standardization of NaOH solution by Potassium Hydrogen
Phthalate (KHP).
06 – 09
04. Standardization of NaOH solution by Oxallic Acid Solution.
10 – 12
05.

Standardization of HCl Solution by NaOH Solution.
13 – 16
06.

Determination of pKa value of weak acid.
17 – 20
07.

Preparation of constant pH buffer and assessment of it’s
activity.
21 – 23
Bibliography and Some commonly used Lab Equipment.
24



Edition: April 2019



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Experiment No. 01 Date:
Name of the experiment: Preparation of solutions of different concentrations.

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
solutionofLitre
soluteofMoles
)MMolarity(
Normality is the number of gram equivalent weights of solute per liter of solution.
solutionofLitre
soluteof weightEquivalent
)NNormality(


Relationship between Normality and Molality
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=Valence

Equivalent mass (Eq. M.) / Gram equivalent Wight:
Equivalent mass of a substance can be calculated by dividing its molar mass per the number of active units in one
molecule of this substance thus: solution of molecule onein unit active of No
MassMolecular
MassEuivalent 

The active unit in the acid-base reactions is the number of hydrogen ions liberated by a single molecule of an 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 same as 1M NaOH.

On the other hand H2SO4 has two ionisable hydrogen atoms; its equivalent mass will be,g49
2
98
Eq.M. 
Thus, the concentration of 1M H2SO4 is same as 2N H2SO4.



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Preparation of a standard solution
Standard solution can be prepared in various ways.
Method-1
1. At first the amount of solute required to prepare the standard solution of 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. Then calculated amount 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. Afterwards 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 disadvantage of this method is that some of the weighed solid may adhere to the original container, weighing
paper, or funnel. Also, 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 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 then transferred to the volumetric flask (of intended volume). Again, a funnel may
need to be inserted 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 and the washings added to the volumetric flask making sure all remaining traces of the solution
have been transferred.
6. Finally, the volumetric flask is diluted to volume (additional solvent is added to the flask until the liquid
level reaches the calibration mark).
7. The flask is capped and inverted as before until the contents are thoroughly mixed.

Disadvantages
The disadvantage to this method is that some of the solution may adhere to the beaker, stirring rod, or
funnel if not washed thoroughly. Also, a possibility of contamination exists from the beaker, rod, or
funnel if they have not been washed carefully.

 : How would you prepare 250mL, 1M Na 2CO3 solution using method -1 and Problem
Method-2?

Chemicals/Reagents:
1. Pure Na2CO3 2. Distilled water Md. Imran Nur Manik

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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

Calculation:
We know that, 1000 mL 1M Na
2CO
3 solution =106 gm Na
2CO
3
250 mL 1M Na
2CO
3 solution ={(106×250)÷ 1000}gm Na
2CO
3 = 26.5 gm Na
2CO
3

Result: The required amount of Na2CO3 = 26.50 gm

Experimental procedures: (Specimen according to the Method-01)

Preparation of 1M 250 mL Na2CO3 solution:
1. Weigh out accurately 26.5 gm Na2CO3.
2. Transfer it to a 250 mL volumetric flask.
3. Add a small quantity of solvent to the volumetric flask and swirl the content (Na2CO3) gently until the
substance is completely dissolved.
4. Add more solvent until the meniscus of the liquid reaches the calibration mark on the neck of the
volumetric flask.
5. Finally cap the volumetric flask and invert several times until the content (Na2CO3) is mixed and
completely dissolved.
6. After complete dissolution label the final solution appropriately.

Preparation of 1M 250 mL Na2CO3 solution: (Specimen according to the Method-02)
1. Weigh out accurately 26.5 gm Na2CO3.
2. Transfer it to a 500 mL beaker.
3. Add a small amount of solvent is to the beaker and the solution is stirred until the solid is dissolved.
4. Then transfer the solution to a 250 mL volumetric flask using a funnel.
Note: Before adding additional solvent to the flask, rinse the beaker, stirring rod, and funnel carefully and
add the washings to the volumetric flask. Ensure that all remaining traces of the solution have been
transferred.
5. Dilute the volumetric flask to the volume by adding supplementary solvent until the liquid level reaches
the calibration mark.
6. Finally the cap the flask and invert until the contents are thoroughly mixed. Label the solution
appropriately.


Precautions:
1. Sodium Carbonate 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 meniscus of the final solution should reach the calibration mark of the volumetric flask.
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Experiment No. 02 Date:
Name of the experiment: Preparation of solutions of different concentrations by dilution.

Principle:
A solution is a homogeneous mixture created by dissolving one or more solutes in a 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 (stock) solutions.
Stock solution is obtained by dissolving the desired 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 solvent 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 exact calculated amount of the stock solution is taken into a clean volumetric flask,
which is then diluted to intended volume. The amount of stock solution required for 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 (diluted solution). Likewise, if the amount of solute remains the same and the volume
gets smaller, the concentration increases. Thus there is an inverse proportion between the strength and total
quantity (volume) of solution. )1...(....................
solution of Volume
1
Strength 

The strength of solution can be expressed in several ways. The concentrations (strength) of standard solutions are
normally expressed in units of moles per litre (mol/L, often abbreviated to M for molarity). Since molarity
changes in accordance with the increment or decreasement of the solution volume, for a particular amount of
solute at a definite temperature, thus the equation (1) can be written as constant)ality ProportionK Where,(
V
1
K. M Or,
V
1
M Or,
solution of Volume
1
Molarity




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 the equation (2) the, a dilute solution can be prepared from a concentrated solution.
In doing so, the targeted concentration and volume is 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 and volume
(M2 & V2) the required volume of the stock solution is calculated. The calculated volume of stock is taken into a
clean volumetric flask, which is then diluted to intended volume.


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 : How would you prepare 10 0mL, 0.2M Na2CO3 solution from 1M Na2CO3 Problem
solution using dilution method?

Chemicals/Reagents:
1. Pure Na2CO3
2. 1M 250 Na2CO3 solution (Stock Solution)
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

Preparation of Stock solution (1M 250 mL Na2CO3 solution)

Calculation:
We know that, 1000 mL 1M Na
2CO
3 solution =106 gm Na
2CO
3
250 mL 1M Na
2CO
3 solution ={(106×250)÷ 1000}gm Na
2CO
3
=26.5 gm Na
2CO
3

Experimental procedures:
Prepare the solution using either the method-1 or method-2 as per Experiment No.-01
NOTE: Detailed procedure for the preparation of reagents must be written in the laboratory note book.

Preparation of 0.2M 100 mL Na2CO3 solution from 1M Na2CO3 solution

Calculation:
We know that,
M1V1 = M2V2 1
1002.0
M
VM
V Or,
1
22
1



= 20 mL
Here, Volume of diluted Na2CO3 Solution (Final volume), V2 = 100 mL
Strength of diluted solution (Final Strength), M2 = 0.2 M
Volume of Stock Solution (Volume to be taken), V1 = ? mL
Strength of Stock solution, M1 = 1M


Result: The required volume of Stock Na2CO3 Solution = 20 mL.

Experimental procedures:

1. At first take a 100 mL clean and dry volumetric flask.
2. Later pipette out the calculated volume (20mL) of stock solution with a high degree of precision.
3. Then transfer the measured volume into the 100mL volumetric flask. Use a funnel if needed.
4. Afterwards add a small quantity of solvent (DW) is added to the volumetric flask and swirl the contents
gently.
5. Add more solvent until the meniscus of the liquid reaches the calibration mark on the neck of the
volumetric flask.
6. Finally cap the volumetric flask and invert several times until the contents are completely mixed and a
homogenous solution is obtained.
Precautions:
1. Sodium Carbonate 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 meniscus of the final solution should reach the calibration mark of the volumetric flask.
5. Add DW slowly as reaches the calibration mark in the volumetric flask. 6. The final solution must be shaken well to get a uniform solution.
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Experiment No. 03 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 a 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
uses 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 colour 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 solution 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 110-120°C), and to 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 air during weighing; that means, it should not be hygroscopic, oxidised 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.
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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 hydrogeniodate 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 hydrogeniodate KH(IO3)2, Sodium oxalate (Na2C2O4),
Arsenic(III) oxide (As2O3), and pure Iron.

Secondary Standards
Solutions that are prepared by standardisation 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 sulphate 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 base, will titrate one mole of acid.

The end point of the solution would be determined by an indicator, phenolphthalein.

Reaction:

+Na
+
, K
+




Phenolphthalein: Phenolphthalein exists in two tautomeric forms. Phenolphthalein has benziod form in
acidic medium and thus, it is colourless while it has quinonoid form in alkaline medium which has pink colour.

Fig. Two forms of Phenolphthalein.

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 Md. Imran Nur Manik

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Apparatus:
1. Electrical balance
2. Spatula
3. Volumetric flask
4. Measuring cylinder
5. Burette
6. Pipette and pipette filler
7. Conical flask
8. Beaker
9. Funnel

Calculation (Sample): Full calculations required to make the reagents, for the experiment must be written down
in the laboratory note book.

Gram Equivalent weight of NaOH = (40÷ 1)= 40 gm
250 mL 0.1N NaOH = {(40250 0.1) ÷1000} gm = 1gm (for 100% pure NaOH)

Gram Equivalent weight of KC8H5O4 = ( 204.22÷ 1) gm = 204.22 gm
50 mL 0.1N KHC8H4O4 = {(204.22500.1) ÷1000} gm = 1.021 gm (for 100% pure KHC8H4O4)

Reagents and their preparations:
Prepare the solution using either the method-1 or method-2 as per Experiment No.-01

NOTE: Detailed procedure for the preparation of reagents must be written in the laboratory note book.

1. 0.1 N 50 mL Potassium Hydrogen Phthalate solution: Weigh out 1.021gm Potassium hydrogen
phthalate (KHP) in a 50 mL volumetric flask; then completely dissolve it with small portion of DW. Finally add
distilled water Q.S. to make 50 mL solution.

2. 0.1 N 250 mL NaOH solution: Weigh out 1 gm NaOH in a 250 mL volumetric flask. Completely
dissolve it with small portion of DW (e.g.125 mL). Finally adjusted the volume up to the 250 mL mark by Q.S. of
distilled water.
3. 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
Experimental Procedures:
1. Take a clean and dry burette.
2. Fill a burette with the prepared 0.1 N NaOH (secondary standard substance, exact concentration of which to
be determined)
NOTE: Run a little through the tap to make sure that no air is trapped. Ensure that the bottom of the meniscus is level with the zero
line at the top of the burette.
3. Record the initial volume of solution in the burette (IBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
4. Pipette out 10 mL acid solution (0.1N KHP solution; solution of standard substance)
into a conical flask with great accuracy.
5. Add 2-3 drops of phenolphthalein (HPh) indicator. (The indicator will show its acidic
colour. (HPh is colourless in acid solution)
6. Open the burette's stopcock and allow standard solution to run down into the conical
Erlenmeyer) flask.
NOTE: Be careful while adding the standard solution so that it does not get adhered to the side wall of
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7. Swirl the conical flask while adding solution from the burette.
8. When you are close to (say within 80% of) the expected end point based on the rough titration, then close
the stopcock.
Caution: 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.
Note: 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.
9. At the very close to the end point this point, close the stopcock, and then only open the stopcock enough to
allow one drop at a time to fall from the burette.
Note: 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")
10. Close the stopcock as soon as the reaction reaches its equivalence point. When the indicator changes colour
permanently, the end point has been reached.
11. Record the final volume (of alkali used to neutralise the acid) as shown on the burette (FBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
Attention: Calculate the titre using data table. {(The volume of solution added to the conical flask from the
burette: Titre = final burette reading (FBR) - initial burette reading (IBR)}
12. Clean and rinse your conical (Erlenmeyer) flask, add another aliquot of solution and indicator, refill the
burette (burette) if required, then repeat this titration. (At least two times) to ensure that the results are
reliable.

Table-1: Data for the standardization of NaOH solution:
No. of
observations
Volume of KHP
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(FBR-IBR)
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

2. Calculation of strength of NaOH solution:

We know that, V1S1 = V2S2 Here, Volume of KHP, V1 = 10 mL
S2 =2
11
V
SV Strength of KHP,S1 = 0.1 N
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2= ?

= N

Result: The strength of NaOH= N.

Precautions: Write down Precautions as depicted in the Experiment No.01 & Experiment No. 02

Comments:
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Experiment No. 04 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 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 absolute pure and may contain some impurities that come during their
synthesis. Therefore the secondary standard chemical should standardize by using the primary pure chemicals.

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 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)
Chemicals / Reagents:
1. 0.1 N Sodium hydroxide (NaOH) solution
2. 0.1 N Oxalic acid (dihydrate) [H2C2O4.2H2O]
3. 0.5% Phenolphthalein indicator
4. Distilled Water

Apparatus:
1. Electrical balance
2. Spatula
3. Volumetric flask
4. Measuring cylinder
5. Burette
6. Pipette and pipette filler
7. Conical flask
8. Beaker
9. Funnel

Calculation (Sample): Full calculations required to make the reagents, for the experiment must be given in the
laboratory note book.

Gram equivalent weight of H2C2O4 = (126.08 ÷ 2) gm = 63.040 gm
100 mL 0.1N H2C2O4 = {(63.0401000.1)÷ ÷1000 }gm = Y gm (for 100% pure H2C2O4)

Reagents and their preparations:
Prepare the solution using either the method-1 or method-2 as per Experiment No.-01

NOTE: Detailed procedure for the preparation of reagents must be written in the laboratory note book.

1. 0.1 N 100 mL Oxalic acid solution: Weigh out 0.636 gm of Oxalic acid dehydrate (99%) in a 100 mL
volumetric flask and completely dissolve it with a small volume of DW. Finally adjust the volume up to
100 mL calibration mark by adding distilled water Q.S.

2. 0.1 N 1000 mL NaOH solution: Weigh out 4.12 gm of NaOH (97%) in a1000 mL volumetric flask with
a small volume of DW. Finally adjust the volume up to 1000 mL mark by adding distilled water Q.S.

3. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 g Phenolphthalein+ 50 mL DW.
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Experimental Procedures:
1. Take a clean and dry burette.
2. Fill a burette with the prepared 0.1 N NaOH (secondary standard substance, exact concentration of which to
be determined)
NOTE: Run a little through the tap to make sure that no air is trapped. Ensure that the bottom of the meniscus is level with the zero
line at the top of the burette.
3. Record the initial volume of solution in the burette (IBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
4. Pipette out 10 mL acid solution (0.1N C2H2O4 solution; solution of standard substance) into a conical flask
with great accuracy.
5. Add 2-3 drops of phenolphthalein (PhI) indicator. (The indicator will show its acidic colour. PhI is olourless in
acid solution)
6. Open the burette's stopcock and allow standard solution to run down into the conical Erlenmeyer) flask.
NOTE: Be careful while adding the standard solution so that it does not get adhered or added by the side wall of the conical flask.
7. Swirl the conical flask while adding solution from the burette.
8. When you are close to (say within 80% of) the expected end point based on the rough titration, then close
the stopcock.
Caution: 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.
Note: 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.
9. At the very close to the end point this point, close the stopcock, and then only open the stopcock enough to
allow one drop at a time to fall from the burette.
Note: 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")
10. Close the stopcock as soon as the reaction reaches its equivalence point. When the indicator changes colour
permanently, the end point has been reached.
11. Record the final volume (of alkali used to neutralise the acid) as shown on the burette (FBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
Attention: Calculate the titre using data table. {(The volume of solution added to the conical flask from the
burette: Titre = final burette reading (FBR) - initial burette reading (IBR)}
12. Clean and rinse your conical (Erlenmeyer) flask, add another aliquot of solution and indicator, refill the burette
(burette) if required, then repeat this titration. (At least two times) to ensure that the results are reliable.

Table-1: Data for standardization of NaOH solution:

No. of
observations
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
Md. Imran Nur Manik

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2. Calculation of strength of NaOH solution:
We know that, V1S1 = V2S2 Here, Volume of Oxalic acid, V1 = 10 mL
S2 =2
11
V
SV Strength of Oxalic Acid, S1= 0.1 N
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2 = ?

= N
Result: The strength of NaOH= N.

Precautions: Write down Precautions as depicted in the Experiment No.01 & Experiment No. 02

Comments:



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Experiment No. 05 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 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.

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 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)


Apparatus:
1. Electrical balance
2. Spatula
3. Volumetric flask
4. Measuring cylinder
5. Burette
6. Pipette and pipette filler
7. Conical flask
8. Beaker
9. Funnel
Chemicals / Reagents:
5. 0.1 N Sodium hydroxide (NaOH) solution
6. 0.1 N Oxalic acid (dihydrate) [H2C2O4.2H2O]
7. 0.1 N Hydrochloric acid (HCl) solution
8. 0.5% Phenolphthalein indicator
9. Distilled Water



Calculation (Sample): Full calculations required to make the reagents, for the experiment must be given in the
laboratory note book.

Gram equivalent weight of H2C2O4 = (126.08 ÷ 2) gm = 63.040 gm
100 mL 0.1N H2C2O4 = {(63.0401000.1)÷ ÷1000 }gm = Y gm (for 100% pure H2C2O4)

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Reagents and their preparations:
Prepare the solution using either the method-1 or method-2 as per Experiment No.-01

NOTE: Detailed procedure for the preparation of reagents must be written in the laboratory note book.

1. 0.1 N 100 mL Oxalic acid solution: Completely dissolve 0.636 gm of Oxalic acid dihydrate (99%) in a
100 mL volumetric flask with a small volume of DW. Finally adjust the volume up to 100 mL mark by
adding distilled water Q.S.
2. 0.1 N 1000 mL NaOH solution: Completely dissolve 4.12 gm of NaOH (97%) in a1000 mL volumetric
flask with a small volume of DW. Finally adjust the volume up to 1000 mL mark by adding distilled water
Q.S.
3. 0.1N 100 mL HCl solution: Take 1.14 mL of HCl (32%) and completely dissolve it in 100 mL
volumetric flask in 100 mL DW.
4. 0.5% 100 mL Phenolphthalein solution: 50 mL Ethanol+ 0.5 g Phenolphthalein+ 50 mL DW.

Experimental Procedures:

A. Standardization of NaOH solution by Oxalic acid Solution:
The standardization of the NaOH solution by oxalic acid solution should be done as per Experiment No.-04

NOTE: Detailed procedure for the titration of NaOH by C2H2O4.2H2O must be written in the laboratory note book.

(The main points are outlined below)
a) Fill the burette with the prepared 0.1 N NaOH solution.
b) Take 10 mL standard 0.1N Oxalic acid solution in a conical flask.
c) Add 1-2 drops of Phenolphthalein indicator in the acid solution and titrate it with the NaOH solution,
until the colour of acid solution changes from colourless to faint pink.
d) Perform another two titrations and calculate the result.

Table-1: Data for standardization of NaOH solution:
No. of
observations
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

B. Calculation of strength of NaOH solution:
We know that, V1S1 = V2S2 Here, Volume of Oxalic acid, V1 = 10 mL
S2 =2
11
V
SV Strength of Oxalic Acid, S1= 0.1 N
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2 = ?

= N
The strength of NaOH= N.

C. Standardization of HCl solution by NaOH solution:

1. Take a clean and dry burette.
2. Fill a burette with the standardized 0.1 N NaOH (exact concentration of which has been determined in A) Md. Imran Nur Manik

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NOTE: Run a little through the tap to make sure that no air is trapped. Ensure that the bottom of the meniscus is level with the zero
line at the top of the burette.
3. Record the initial volume of solution in the burette (IBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
4. Pipette out 10 mL 0.1N HCl acid solution (secondary standard substance, exact concentration of which to be
determined) into a conical flask with great accuracy.
5. Add 2-3 drops of phenolphthalein (PhI) indicator. (The indicator will show its acidic colour. PhI is olourless in
acid solution)
6. Open the burette's stopcock and allow standard solution to run down into the conical
Erlenmeyer) flask.
NOTE: Be careful while adding the standard solution so that it does not get adhered or added by the
side wall of the conical flask.
7. Swirl the conical flask while adding solution from the burette.
8. When you are close to (say within 80% of) the expected end point based on the rough
titration, then close the stopcock.
Caution: 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.
Note: 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.
9. At the very close to the end point this point, close the stopcock, and then only open the stopcock enough to
allow one drop at a time to fall from the burette.
Note: 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")
10. Close the stopcock as soon as the
reaction reaches its equivalence point. When
the indicator changes colour permanently,
the end point has been reached.
11. Record the final volume (of alkali used
to neutralise the acid) as shown on the
burette (FBR).
Note: Always read the volume at eye level and from
the bottom of the meniscus.
Attention: Calculate the titre using data
table. {(The volume of solution added to the
conical flask from the burette: Titre = final
burette reading (FBR) - initial burette reading
(IBR)}
12. Clean and rinse your conical
(Erlenmeyer) flask, add another aliquot of
solution and indicator, refill the burette
(burette) if required, then repeat this
titration. (At least two times) to ensure that
the results are reliable.
Table-2: Data for standardization of HCl solution: Md. Imran Nur Manik

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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

D. Calculation of strength of HCl solution:
We know that, V1S1 = V2S2 Here, Volume of HCl, V1 = 10 mL
S1 =1
22
V
SV Strength of HCl, S1 = ?
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2 = N (Known Earlier)

= N

Result: The strength of HCl= N.

Precautions:
1. Rapid addition of NaOH from the burette is strictly prohibited.
2. HCl is corrosive. Be careful while dealing with acids.
3. NaOH is corrosive. Be careful while dealing with strong bases.
4. The difference of volume of NaOH between two successive titrations should not be more than
0.1 mL.
5. The conical Flask should be continuously swirled while adding the NaOH from the burette.
6. Place of a piece of white paper below the conical flask during titration to aid the determination
of the colour change at the end point.

Comments:


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Experiment No. 06 Date:
Name of the experiment: Determination of pKa value of weak acid.

Principle
A buffer solution is one which resists (or buffers) a change in its pH. If acid is added then, within reason, the pH
does not fall; if base is added, the pH does not rise. Buffers are usually composed of a mixture of weak acids or
weak bases and their salts and function best at a pH equal to the pKa of the acid or base involved in the buffer.
The equation that predicts the behaviour of buffers is known as the Henderson–Hasselbalch equation
(named after chemists Lawrence Joseph Henderson and Karl Albert Hasselbalch). It is derived as follows, by
considering a weak acid that ionises in solution. [A cid]
[Salt]
logpHpK a or,
[A cid]
[Salt]
logpK apH



Where, pKa= –log Ka; and Ka=Decomposition constant of acid.
An example of a buffer is a mixture of acetic acid and sodium acetate, which ionises and acts as follows:

CH3COOH⇌CH3COO
–
+H
+
CH3COO
–
Na
+
⇌CH3COO
–
+Na
+
Apparatus:
1. Electrical balance
2. Spatula
3. Volumetric flask
4. Measuring cylinder
5. Burette
6. Pipette and pipette filler Fig. Mechanism of Buffer action of an acid buffer.
7. Conical flask
8. Beaker
9. Funnel

Chemicals / Reagents:
1. 0.1 N Sodium hydroxide (NaOH) solution
2. 1% Acetic acid (CH3COOH) solution
3. 0.1 N Oxalic acid (dihydrate)[ H2C2O4.2H2O]
4. 0.5% Phenolphthalein indicator
5. Distilled Water

Calculation: Full calculations required to make the reagents, for the experiment must be written down.

Preparation of Reagents:
A. 1% 100 mL Acetic Acid Solution: Take 1 mL glacial acetic acid (100%) in a 100 mL volumetric flask, and
then adjust the volume up to 100 mL mark by adding distilled water.

B.0.1 N 1000 mL NaOH Solution: Completely dissolve 4.12 gm NaOH (97%) pellets in a 1000 mL volumetric
flask at first with a small portion of DW, and finally adjust the volume up to the mark by adding DW Q.S.

C.0.1 N 100 mL Oxalic Acid Solution: Completely dissolve 0.636 gm oxalic acid (99%) in a 100 mL volumetric
flask at first with a small portion of DW, and finally make the volume up to mark by adding distilled water Q.S.

D.0.5% Phenolphthalein: 50 mL Ethanol + 0.5 gm phenolphthalein + 50 mL distilled water. Md. Imran Nur Manik

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Experimental Procedures:
A. Standardization of NaOH solution by Oxalic acid Solution:

The standardization of the NaOH solution by oxalic acid solution should be done as per Experiment No.-04

NOTE: Detailed procedure must be written in the laboratory note book.

(The main points are outlined below)

a) Fill the burette with the prepared 0.1 N NaOH solution.
b) Take 10 mL standard 0.1N Oxalic acid solution in a conical flask.
c) Add 1-2 drops of Phenolphthalein indicator in the acid solution and titrate it with the NaOH solution,
until the colour of the acid solution changes from colourless to faint pink.
d) Perform another two titrations and calculate the result.

Table-1: Data for standardization of NaOH solution:

No. of
observations
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

B. Calculation of strength of NaOH solution:
We know that, V1S1 = V2S2 Here, Volume of Oxalic acid,V1 = 10 mL
S2 =2
11
V
SV Strength of Oxalic acid, S1 = 0.1 N
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2= ?

= N
The strength of NaOH = N.
C. Determination of the strength of CH3COOH solution by NaOH solution:

1. Take a clean and dry burette.
2. Fill a burette with the standardized 0.1 N NaOH (exact concentration of which has been determined in A)
NOTE: Run a little through the tap to make sure that no air is trapped. Ensure that the bottom of the meniscus is level with the zero
line at the top of the burette.
3. Record the initial volume of solution in the burette (IBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
4. Pipette out 10 mL 1% CH3COOH acid solution (secondary standard substance,
exact concentration of which to be determined) into a conical flask with great
accuracy.
5. Add 2-3 drops of phenolphthalein (PhI) indicator. (The indicator will show its acidic
colour. PhI is olourless in acid solution)
6. Open the burette's stopcock and allow standard solution to run down into the
conical Erlenmeyer) flask.
NOTE: Be careful while adding the standard solution so that it does not get adhered or added by the side wall of the conical flask. Md. Imran Nur Manik

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7. Swirl the conical flask while adding solution from the burette.
8. When you are close to (say within 80% of) the expected end point based on the rough titration, then close
the stopcock.
Caution: 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.
Note: 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.
9. At the very close to the end point this point, close the stopcock, and then only open the stopcock enough to
allow one drop at a time to fall from the burette.
Note: 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")
10. Close the stopcock as soon as the reaction reaches its equivalence point. When the indicator changes colour
permanently, the end point has been reached.
11. Record the final volume (of alkali used to neutralise the acid) as shown on the burette (FBR).
Note: Always read the volume at eye level and from the bottom of the meniscus.
Attention: Calculate the titre using data table. {(The volume of solution added to the conical flask from the
burette: Titre = final burette reading (FBR) - initial burette reading (IBR)}
12. Clean and rinse your conical (Erlenmeyer) flask, add another aliquot of solution and indicator, refill the burette
(burette) if required, then repeat this titration, (at least two times) to ensure that the results are reliable.

Table-2: Data for standardization of CH3COOH solution:

No. of
observations
Volume of CH3COOH
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(FBR-IBR)
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

D. Calculation of strength of Acetic acid solution:
We know that, V1S1 = V2S2 Here, Volume of CH3COOH, V1 = 10 mL
S1 =1
22
V
SV Strength of CH3COOH, S1 = ?
Volume of NaOH, V2 = mL (Mean)
= Strength of NaOH, S2 = N (Known Earlier)

= N
The strength of CH3COOH = N.
E. Determination of pH of Buffer solution:
1) Take 40 mL 1% Acetic acid solution and 20 mL standardized NaOH solution in a 100 mL beaker and
mix them very well.

2) Then measure the pH of the buffer solution using a pH meter.

Result: The pH of the solution=
Md. Imran Nur Manik

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F. Determination of pKa of the solution:
Now calculate pKa value is by the following equation, volumeTotal
COOHCH ofStrength COOHCH Remaining of Volume
Volume Total
NaOH ofStrength NaOH of Volume
logpHpKa
getw)equation(1thefromThus
VolumeTotal
COOHCHofStrengthCOOHCH Remaining of Volume
[Acid]Acid,ofionconcentratexperimentpresenttheIn
VolumeTotal
COOHCHofStrengthCOOHCHofVolume
[Acid]Acid,ofionconcentrattheAnd
VolumeTotal
NaOHofStrengthNaOHofVolume
[Salt]Salt,ofionconcentrattheHere,
)1..(..............................................................................................................
[Acid]
[Salt]
logpHpKa
33
33
33










e


Result: The pKa value of the solution =
Precautions:

Comments:

Md. Imran Nur Manik

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Experiment No. 07 Date:
Name of the experiment: Preparation of constant pH buffer and assessment of its activity.

Principle:
A solution which resists changes in pH when small quantities of an acid or an alkali are added to it, is called
buffer solution. Most of the buffer solutions usually consist of a mixture of 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. Addition of small amounts of acid or base results in the mopping up
or the release of protons by the conjugate base or the acid as 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.

Figure 1. (a) The buffered solution on the left and the unbuffered solution on the right have the same pH (pH 8); they are basic, showing the yellow
color 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 unbuffered solution has become acidic, as indicated by the change in color of the methyl orange, which turns red at a pH of
about 4. (credit: modification of work by Mark Ott)

How Buffers 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 2. 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 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 (Figure 2).
TYPES OF BUFFERS
Buffers are classified into following types

1. Simple Buffers:
Simple buffers are categorized into three different ways.
A) Salts of weak acid and a weak base: CH3COONH4, NH4CN.
B) Proteins and amino acids.
C) A mixture of an acid salt and a normal salt formed from polybasic acid: Na2HPO4 and Na3PO4
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2. Mixed Buffers
Acidic Buffer: A solution of weak acid and its salt with strong base.
Example: CH3COOH +CH3COONa, H2CO3 +Na2CO3
Basic Buffer: A solution of 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:
1. Electrical balance
2. Spatula
3. Volumetric flask
4. Conical flask
5. Beaker
6. Measuring cylinder
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

Calculation: Full calculations required to make the reagents, for the experiment must be written down.
Preparation of Reagents:
Prepare the solution using either the method-1 or method-2 as per Experiment No.-01

NOTE: Detailed procedure for the preparation of reagents must be written in the laboratory note book.
1. 0.1 M 100 mL CH3COONa Solution: Completely dissolve 1.374 gm CH3COONa (99%) in a 100 mL
volumetric flask with a small amount of DW. Later on adjust the volume up to the mark by adding distilled water
Q.S.

2. 0.1 M 100 mL CH3COOH Solution: Completely dissolve 0.605 gm CH3COOH (99%) in a 100 mL
volumetric flask with a small amount of DW. Later on adjust the volume up to the mark by adding distilled water
Q.S.

3. 0.1 M 50 mL HCl Solution: Completely dissolve 0.42 mL HCl (Purity 37% & density 1.19 g/mL) in a 50 mL
volumetric flask with a small amount of DW. Later on adjust the volume up to the mark by adding distilled water
Q.S.

4. 0.1 M 50 mL NH4OH Solution: Completely dissolve 0.3786 mL NH3 (Purity 25% & density 0.9 g/mL) in a
50 mL volumetric flask with a small amount of DW. Later on adjust the volume up to the mark by adding distilled
water 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. Md. Imran Nur Manik

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Experimental Procedures:
A. Preparation 0f 200 mL buffer solution with pH 4.6
1. Take 102 mL 0.1 M CH3COOH solution in a 250 mL beaker.
2. Then slowly add 98 mL 0.1M sodium acetate solution with continuous stirring.
3. Carefully submerge the detector of pH meter into the solution.
4. Take the reading of pH meter simultaneously while adding CH3COONa.
5. The pH of the solution should be 4.6.
6. Now adjust solution to final desired pH using HCl or NaOH (if needed) with continuous stirring by
the stirrer.
B. Assessment of buffer action of the preapared buffer solution:
Addition of small amount of acid:
1. Take 50 mL prepared buffer solution in a 100 mL beaker.
2. Then slowly add 2-3mL 0.1M HCl solution with continuous stirring.
3. Carefully submerge the detector of pH meter into the solution.
4. Take the reading of pH meter simultaneously while adding HCl Solution.
Addition of small amount of base:
1. Take 50 mL prepared buffer solution in a 100 mL beaker.
2. Then slowly add 2-3mL 0.1M NH4OH solution with continuous stirring.
3. Carefully submerge the detector of pH meter into the solution.
4. Take the reading of pH meter simultaneously while adding NH4OH Solution.

Table-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 addition of base:

Inference: The pH of the buffer solution did not chanced after the addition of 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.

Comments: Md. Imran Nur Manik

26. Manual: Physical Pharmacy-I Lab


Prepared By: Md. Imran Nur Manik Page 24
Lecturer; Department of Pharmacy; Northern University Bangladesh.

TOP

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. 3
rd edn. Pharmaceutical Press, UK.
Md. Imran Nur Manik

27. Manual: Physical Pharmacy-I Lab


Prepared By: Md. Imran Nur Manik Page 25
Lecturer; Department of Pharmacy; Northern University Bangladesh.

TOP

Appendix
Acetate buffer solutions pH 3.6 - 5.6
Make up the following solutions:
(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 [Pearse 1980]
ph Vol. of 0.1M acetic acid Vol. of 0.1M sodium acetate
3.6 185 mls 15 mls
3.8 176 mls 24 mls
4.0 164 mls 36 mls
4.2 147 mls 53 mls
4.4 126 mls 74 mls
4.6 102 mls 98 mls
4.8 80 mls 120 mls
5.0 59 mls 141 mls
5.2 42 mls 158 mls
5.4 29 mls 171 mls
5.6 19 mls 181 mls

What is the molarity of the following?

(a) 0.694 mol dissolved in 3.55 L of solution
(b) 2.19 mol of NaCl dissolved in 700.0 mL of solution
(c) 0.3882 g of KCl dissolved in 0.5000 L of solution
(d) 1.003 g of CuSO4·5 H2O dissolved in 250.0 mL of solution
(e) 30.00 mL of 6.0 M NaOH diluted to 100.0 mL of solution
(f) 0.100 L of 12.0 M HCl diluted to 500.0 mL of solution

Tell how you would prepare each of the following:

(a) 500.0 mL of a 0.10 M solution of KOH from pure solid KOH
(b) 250.0 mL of a 0.15 M solution of NaCl from pure solid NaCl
(c) 100.0 mL of a 2.0 M solution of glucose from pure solid glucose (C6H12O6)
(d) 500.0 mL of a 0.10 M solution of HCl from concentrated HCl, which is 12.0 M
(e) 100.0 mL of a 0.25 M solution of NaOH from a solution of NaOH that is 2.0 M
(f) 2.0 L of a 0.50 M solution of sulfuric acid from concentrated sulfuric acid, which is 18.0 M

1. How many milliliters of a KOH solution, prepared by dissolving 60.0 g of KOH in 100.0 mL of solution, are needed to
prepare 450.0 mL of a 0.70 M solution?
2. How many milliliters of a KCl solution, prepared by dissolving 45 g of KCl in 500.0 mL of solution, are needed to
prepare 750.0 mL of a 0.15 M solution?
3. A solution of KNO3, 500.0 mL with 0.35 M concentration, is needed. Tell how you would prepare this solution
(a) From pure, solid KNO3.
(b) From a solution of KNO3 that is 4.5 M.
Md. Imran Nur Manik

28. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name
Standardization of NaOH solution by
Potassium Hydrogen Phthalate (KHP) solution.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3


Procedures:
Standardization of NaOH solution by Potassium Hydrogen Phthalate Solution:
A. Fill the burette with the prepared NaOH solution.
B. Take 10mL standard 0.1N KHP in a conical flask.
C. Add 1-2 drops of Phenolphthalein indicator in the acid solution.
D. Now titrate it with the NaOH solution, until the colour of the acid solution changes from colourless to faint pink.
E. Perform another two titrations and calculate the result.

Table: Data for standardization of NaOH solution:
No. of
observations
Volume of KHP
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

Abbreviation:
IBR= Initial Burette Reading (Before starting the titration)
FBR= Final Burette Reading (After completion of one successful titration)
Difference= (IBR-FBR) Or (FBR-IBR). Whatever applicable.

Calculation of strength of NaOH solution:

We know that, V1S1 = V2S2 Here, Volume of KHP, V 1 = 10 mL

Or, S2 = 2
11
V
SV Strength of KHP, S1 = 0.1 N
Volume of NaOH, V2 = mL

= Strength of NaOH, S2 = ?

= N



Result: The strength of NaOH= N.









Checked By
Remarks

29. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name
Standardization of NaOH solution by
Oxalic acid Solution.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3


Procedures:
Standardization of NaOH solution by Oxalic acid Solution:
A. Fill the burette with the prepared NaOH solution.
B. Take 10mL standard 0.1N Oxalic acid Solution in a conical flask.
C. Add 1-2 drops of Phenolphthalein indicator in the acid solution.
D. Now titrate it with the NaOH solution, until the colour of the acid solution changes from colourless to faint pink.
E. Perform another two titrations and calculate the result.

Table: Data for standardization of NaOH solution:
No. of
observations
Volume of Oxalic acid
Solution
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

Abbreviation:
IBR= Initial Burette Reading (Before starting the titration)
FBR= Final Burette Reading (After completion of one successful titration)
Difference= (IBR-FBR) Or (FBR-IBR). Whatever applicable.

Calculation of strength of NaOH solution:

We know that, V1S1 = V2S2 Here, Volume of C2H2O4, V1 = 10 mL

Or, S2 = 2
11
V
SV Strength of C2H2O4, S1 = 0.1 N
Volume of NaOH, V2 = mL

= Strength of NaOH, S2 = ?

= N



Result: The strength of NaOH= N.






Remarks

Checked By

30. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name Standardization of HCl solution by NaOH solution.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3


Procedures:
Standardization of NaOH solution by Oxalic acid Solution:
a) Fill the burette with the prepared NaOH solution.
b) Take 10mL of the standard 0.1N Oxalic acid in a conical flask.
c) Add 1-2 drops of Phenolphthalein indicator in the acid solution and titrate it with the NaOH solution, until the colour
of the acid solution changes from colourless to faint pink.
d) Perform another two titrations and calculate the result.

Table-1: Data for standardization of NaOH solution:

No. of
observations
Volume of Oxalic acid
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

Calculation of strength of NaOH solution:
We know that, V
1S
1 = V
2S
2 Here, Volume of Oxalic acid ,V
1 = 10 mL
S
2 = 2
11
V
SV Strength of Oxalic acid S
1 = 0.1 N
Volume of NaOH, V
2 = mL
= Strength of NaOH, S
2 = ?

= N

Standardization of HCl solution by NaOH solution:
a) Fill the burette with the standardized NaOH solution.
b) Take 10mL of the prepared HCl solution in a conical flask.
c) Add 1-2 drops of Phenolphthalein indicator in the acid solution and titrate it with the NaOH solution, until the
colour of the acid solution changes from colourless to faint pink.
d) Perform another two titrations and calculate the result.
Table-2: Data for standardization of HCl solution:
No. of
observations
Volume of HCl
(V1 mL)
Volume of NaOH solution
(mL)
Difference
(mL)
Mean
volume
(V2 mL) IBR FBR
1 10
2 10
3 10

Calculation of strength of HCl solution:
We know that, V
1S
1 = V
2S
2 Here, Volume of HCl, V
1 = 10 mL
S
1 = 1
22
V
SV Strength of HCL,S
1 = ?
Volume of NaOH, V
2 = mL
= Strength of NaOH, S
2 = N (Known Earlier)

= N

Result: The strength of HCl = N.
Remarks

Checked By

31. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name Determination of pKa value of a weak acid
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3


Procedures:
1. Standardization of NaOH solution by Oxalic acid Solution:
Table-1: Data for standardization of NaOH solution:
No. of
observations
Volume of Oxalic acid
(V1 mL)
Volume of NaOH solution (mL)
Difference
(mL)
Mean volume
(V2 mL)
IBR FBR
1 10
2 10
3 10

2. Calculation of strength of NaOH solution:
We know that, V
1S
1 = V
2S
2 Here, Volume of Oxalic acid,V
1 = 10 mL
S
2 = 2
11
V
SV Strength of Oxalic acid, S
1 = 0.1 N
Volume of NaOH, V
2 = mL
= Strength of NaOH, S
2 = ?

= N

3. Determination of the strength of CH3COOH solution by NaOH solution:
Table-2: Data for standardization of CH3COOH solution:
No. of
observations
Volume of CH3COOH
(V1 mL)
Volume of NaOH solution (mL)
Difference
(mL)
Mean volume
(V2 mL)
IBR FBR
1 10
2 10
3 10

4. Calculation of strength of Acetic acid solution:
We know that, V
1S
1 = V
2S
2 Here, V
1 = 10 mL
S
1 = 1
22
V
SV S
1 = ?
V
2 = mL
= S
2 = N (NaOH Known Earlier)

= N
The strength of CH
3COOH = N.
5. Determination of pH of Buffer solution:
The pH of the solution=
6. Determination of pKa of the solution:
Now calculate pKa value is by the following equation, 


 p Kaor
v o lu meTo tal
CO O HCH ofStren g th CO O HCH Remain in g of Vo lu me
Vo lu me To tal
N aO H ofStren g th N aO H of Vo lu me
lo gpHp K a,
33

Result: The pKa value of the solution=



Remarks

Checked By

32. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name
Preparation of constant pH buffer and assessment of
its activity.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3


Assessment of buffer action of the prepared buffer solution:
Addition of small amount of acid:
1. Take 50 mL prepared buffer solution in a 100 mL beaker.
2. Then slowly add 2-3mL 0.1M HCl solution with continuous stirring.
3. Carefully submerge the detector of pH meter into the solution.
4. Take the reading of pH meter simultaneously while adding HCl Solution.
Addition of small amount of base:
1. Take 50 mL prepared buffer solution in a 100 mL beaker.
2. Then slowly add 2-3mL 0.1M NH4OH solution with continuous stirring.
3. Carefully submerge the detector of pH meter into the solution.
4. Take the reading of pH meter simultaneously while adding NH4OH Solution.

Table-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 addition of base :


Inference:













Remarks

Checked By

33. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name Preparation of Solutions of Different Concentrations.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3









Remarks

Checked By

34. Excellent Very Good Good Nice Fine Not Well
Need Improvement Bad Worst Need More Attention
Exp. No. Exp. Name
Preparation of Solutions of Different Concentrations by
Dilution Method.
Batch: Date: D D M M Y Y Y Y

ID: 0 5 1 6 0 1 0 1 2 5 3



Preparation of M mL Solution From M Solution

Calculation
We know that, M1V1=M2V2 )...(..........................................................................................
M
VM
V Or,
1
22
1 i


Here,
The volume of diluted Solution (Final volume), V2= mL
The Strength of diluted Solution (Final strength), M2= M
The volume of Stock Solution (Volume to be taken), V1= ? mL
The Strength of Stock Solution M1= M

Now,
Putting these values in the equation (i) we get
OR mL

M
VM
V
1
22
1 



Result:
The volume of the stock solution to be taken= mL








Remarks
Checked By