chemistry practical

LABORATORY MANUAL

CHE102

CHEMISTRY LAB

LMCHE 102: CHEMISTRY LAB

Some rules to be followed in the Lab:

General Rules:

1. Entry without lab coat in chemistry lab is strictly prohibited.

2. Mobile phones should be switched off and kept in the bag during lab hours.

3. No group discussions are allowed in the lab.

4. Clean the apparatus as well slab after your experiment is finished.

5. Do not do any indiscipline activity in lab as you are under strict cc-TV surveillance.

6. Always bring the lab manuals with you.

7. Do not use laptop while performing the experiments.

8. Switch off electrical apparatus after their use.

9. Do not throw filter papers in sink, dispose all solid waste in dustbin.

10. Liquid waste must be deposited in the waste containers designated for Chlorinated

organic solvents, Organic solvents and aqueous waste

Precautionary Rules:

1. Never pipette out strong acids and bases with your mouth, it can be dangerous,

therefore use measuring cylinders for such chemicals.

2. Never try to smell the chemicals as it can be dangerous for you.

3. Cap the bottles after taking chemical as uncovered bottles can be a source of harmful

fumes.

4. In case of any accidental spill over of any chemical on you, report your teacher or lab

technician immediately.

5. Report your lab technician if any breakage of glass apparatus takes place.


Table of Contents

S.No Title of Experiment Page No.

1 With the help of complexometeric titration how you will determine

the hardness of given hard water sample by using EDTA. Provided

standard hard water.(1ml of S.H.W.=1mg of CaCO3)

4-6

2 Identification of elements present in given compound. 7-9

3 Qualitatively analysis of the given organic compound for Carboxylic

acids

10-12

4 Determination of the dissociation constant of acetic acid using pH-

meter.

13-16

5 Determination of Strength of hydrochloric acid

solution(approximately N/10) by titrating it against sodium hydroxide

solution conductometrically

17-20

6 To test the validity of Beer-Lambert‟s law using colorimeter and to

determine unknown concentration of solution

21-24

7 Estimation of nickel in the given sample using dimethyl glyoxime. 25-26

8 Determination of the rate constant of hydrolysis in case of ethyl

acetate using an alkali.

27-30

9 Determine the strength of given solution of ferrous ammonium

sulphate by titrating against potassium dichromate solution

31-32

10 Separation of a mixture of organic compounds by thin layer

chromatography.

33-35

11 Some important instruments in chemistry lab 36-38


EXPERIMENT NO. 1

EXPERIMENT: To determine the hardness of the given hard water sample by EDTA

method. Provided Standard Hard Water (1 ml of SHW =1 mg of CaCO3).

EQUIPMENTS REQUIRED: Burette, Burette stand, Titration flask, Pipette, Beakers,

funnel etc.

MATERIALS REQUIRED: Standard Hard water, EDTA solution, Eriochrome black-T

(EBT), Buffer Solution, Sample hard water.

LEARNING OBJECTIVES:

i. The purpose of this experiment is to determine the hardness of water by measuring the

concentrations of calcium and magnesium in water samples by titration.

ii. To know about use of buffer solution: The buffer being used has composition NH4Cl

and NH4OH. Its pH is the order of 10.5.

iii. How the indicator works: When indicator is added to hard water it combines with free

metal ions present in water.

HIn-2

+ M+2

→ MIn- + H

+ {M = Mg or Ca}

(Wine red)

When EDTA solution is added to the titration flask it combines with the free metal ions

giving metal EDTA complex, which is stable and colorless.

H2Y2-

+ M+2

→ MY-2

+ 2H+

When all the free metal ions are exhausted, next drop of EDTA removes the metal ion

engaged with indicator and the original blue color is restored.

H2Y2-

+ Min- → MY

-2 + HIn

-2 + H

+

OUTLINE OF PROCEDURE:

Standardisation of EDTA solution:

(a) Pipette out 10ml of standard hard water in the titration flask. Add to it 2-3ml of buffer

solution and two drops of Eriochrome Black-T indicator. A wine red color appears.

(b) Titrate this solution against EDTA solution taken in a burette till wine red colour

changes to blue color.

(c) This is the end point. Recovered the volume of EDTA consumed as A ml. Repeated

the procedure to get at least three concordant readings.


Table1: Standardization of EDTA

S. No. Burette readings Volume of EDTA

Consumed (R2 - R1) mL Initial (R1) Final (R2)

1.

2.

3.

4.

5.

Determination of Total Hardness:-

(a) Pipette out 10ml of sample hard water in the titration flask. Add to it 2-3ml of buffer

solution and two drops of Eriochrome Black-T indicator.

(b) A wine red colour appears. Titrate this solution against EDTA solution taken in a

burette till wine red colour changes to blue color.

(c) This is the end point. Recovered the volume of EDTA consumed as B ml. Repeated

the procedure to get at least three concordant readings

Table2: Determination of Total Hardness

S. No. Burette readings Volume of EDTA

Consumed (R2 - R1) mL Initial (R1) Final (R2)

1.

2.

3.

4.

5.

Calculations:

(a).Standardisation of EDTA solution:-

1 ml of standard hard water = 1 mg of CaCO3

10 ml of S.H.W. = 10 mg of CaCO3 = A ml of EDTA

A ml of EDTA = 10 mg of CaCO3

1 ml of EDTA = 10/A mg of CaCO3


(b).Calculation of total hardness:

10 ml of hard water sample = B ml of EDTA

Now 1 ml of EDTA = 10/A mg of CaCO3

10 ml of hard water sample = B x 10/A mg of CaCO3.

1 ml of hard water sample = B x 10/A x 1/10 mg of CaCO3.

1000 ml of hard water sample = B x 10/A x 1/10 x 1000 mg of CaCO3.

Hence total hardness = 1000 x B/A ppm

REQUIRED RESULT

PARAMETERS: The hardness of water is...........ppm

RELATIONSHIPS TO BE DETERMINED: How much EDTA is consumed for 1 mL of

SHW, which is then, can be used for the calculation of unknown hardness.

GRAPHS AND PLOTS: NA

ERROR ANALYSIS: NA

CAUTIONS:

1. The burette, pipette and conical flask should be washed and then rinsed with distilled

water.

2. Redistilled water should be employed for preparing the EDTA solution.

3. The colour change near the end point is very slow and thus should be observed very

carefully.


EXPERIMENT-2

EXPERIMENT: Identification of elements present in given compound.

EQUIPMENTS REQUIRED: Test tube, China dish, fusion tubes, glass rod, Bunsen burner,

tripod stand, funnel, beaker, filter papers, beakers etc.

MATERIAL REQUIRED: Sodium metal, Ferrous sulphate (freshly prepared), sodium

hydroxide, dilute HCl/H2SO4, Tollen‟s reagent, lead acetate solution, Sodium nitroprusside

solution, Conc. HNO3, NH4OH


(i) To know about presence of various elements present in the given compound

(ii) Presence of nitrogen and sulphur in the given compound can be detected.

(iii) Presence of halogens present in the compound can be detected.

THEORY:

Chemical equations:

FeSO4 + 2NaOH Fe(OH)2 + Na2SO4

Fe(OH)2 + 2NaCN Fe(CN)2 +2NaOH

Fe(CN)2 + 4NaCN Na4(FeCN)6

( sodium ferrocyanide)

3Na4(FeCN)6 + 4Fe3+

Fe4[Fe(CN)6]3 + 12Na+

( ferric ferrocyanide)

(prussian blue)

Na2S + Pb(CH3COO)2 PbS (ppt) + 2 CH3COONa

NaCN + HNO3 NaNO3 + HCN (g)

Na2S + HNO3 2NaNO3 + H2S (g)

NaX + AgNO3 AgX + NaNO3 (X = Cl, Br, I)


Preparation of Lassaigne’s extract:

1. In a dry fusion tube take a small piece of sodium metal, and heat the fusion tube in the

flame so that the sodium metal is melted completely.


2. To this heated tube add carefully a small amount of given compound, and heat the

fusion tube again in the flame till it gets red hot.

3. Pour the red hot fusion tube immediately into a china dish containing 20 mL of distilled

water (care should be taken that china dish should not contain any impurity or chemical

which could interfear in the detection)

4. Repeat the procedure for at least seven times with fusion tube (add the fusion tube to

the same china dish having distilled water)

5. Crush all the fusion tube in china dish with the help of a clean glass road and boil the

solution in china dish to evaporate so that the volume is reduced nearly half.

6. Filter the content through filer paper and collect the solution obtained in a clean beaker,

the solution obtained will be Lassaigne‟s extract.

Test for nitrogen.

1. In a clean test tube take about 2 mL of the Lassaigne‟s extract and add a freshly

prepared solution of ferrous sulphate, the dirty green precipitate of Fe(OH)2 will appear

in the solution.

2. To this solution add a small amount of dilute sodium hydroxide solution and heat it

gently in the flame.

3. Dirty green precipitate may disappear to this solution add a small amount of dilute HCl,

if the solution turns out be prussian blue to prussian green, nitrogen is present.

Test for Sulphur.

1. In a clean test tube take about 2 mL of the Lassaigne‟s extract and add one mL of acetic

acid and then add few drops of lead acetate, if black coloured precipitates appear in the

solution, sulphur is present in the compound.

2. In a clean test tube take about 2 mL of the Lassaigne‟s extract and add few drops of

Sodium nitroprusside, if the purple or violet colour appears, it shows the presence of

sulphur in the compound.

Test for Halogens.

1. In a clean test tube take about 2 mL of the Lassaigne‟s extract and add one mL Conc.

HNO3 to this solution add 1 mL of tollen‟s reagent the formation of precipitate

indicates the presence of halogens in the given compound

Identification of halogen

1. If the precipitate formed are white in colour which are readily soluble in NH4OH, these

indicates the presence of chloride in the given compound


2. If the precipitate formed are pale yellow in colour which are partially soluble in

NH4OH, these indicates the presence of bromine in the given compound

3. If the precipitate formed are yellow in colour which are insoluble in NH4OH, these

indicates the presence of iodine in the given compound

RESULTS REQUIRED:

PARAMETERS: Analysis of various elements present in the given compound

RELATIONSHIPS TO BE DETERMINED:

Which test is applicable to confirm the presence of an element?

GRAPHS/PLOTS: NA

ERROR ANALYSIS: NA

CAUTIONS:

1. Sodium metal should be handled very carefully

2. Sodium metal should never be allowed to come in contact with water.

3. On addition of Fe(OH)2 a dirty green precipitate comes out which are not true indicator

for the presence of nitrogen in the compound, these appear due to the formation of

ferric hydroxide which is soluble in dilute sodium hydroxide.


EXPERIMENT-3

EXPERIMENT: Qualitatively analysis of the given organic compound for Carboxylic acids.

EQUIPMENTS REQUIRED: Test tube, Test tube holder, Bunsen burner, spatula, dropper,

Blue litmus solution, water bath

CHEMICALS REQUIRED: Sodium hydroxide, Blue Litmus solution, Ethanol,

Phenolphthalein, sodium bicarbonate, H2SO4, conc. HCl.


1. To understand the solubility behaviour of the aromatic acids

2. To understand the differentiation between aromatic and aliphatic compounds

3. To confirm about the presence of carboxylic acids in the given organic compound.


Physical Examination

Colour of the given organic compound: White/Cream/Brown/Pale Yellow.

Physical state of the compound: Solid/Powder/Amorphous/Liquid.

Solubility Tests

Take a pinch of the given organic compound in the test tube and add about one mL of water

to it shake it and check for the solubility of the compound, if compound is insoluble in cold

water warm the water and check the solubility, if compound is still insoluble dissolve the

compound in 5% NaOH solution, if the compound dissolves it confirms the presence of

carboxylic acid in the given organic compound.

Test for Aliphatic/Aromatic nature

Take a pinch of the compound on the clean nickel spatula and heat it in the Bunsen burner

flame, if the compound burns with black sooty flame it is an aromatic compound/ if it buns

with non luminous flame its aliphatic in nature.

Litmus Paper Test

In a clean test tube add a small portion of compound; now add 2-3 mL of distilled water and

add to it 2-3 drops of blue litmus solution, if the blue litmus turns to be red, then an acid

group is present.


Sodium Hydroxide-phenolphthalein Test

To one mL of dilute sodium hydroxide, add 2-3 drops of phenolphthalein so that the solution

turns out to be pink in colour. In another test tube dissolve a pinch of given organic

compound in water or alcohol, to this test tube add sodium hydroxide-phenolphthalein

reagent and shake well. Disappearance of pink colour in the solution indicates the presence of

carboxylic acid.

Sodium Bicarbonate Test

To 3 mL of cold saturated sodium bicarbonate solution (made by dissolving excess of sodium

bicarbonate in water) add a pinch of given organic compound (if compound is not soluble in

water than compound must be dissolved in alcohol before adding to the sodium bicarbonate

solution), vigorous evolution of carbon dioxide indicates the presence of carboxylic acid

group.

Confirmatory Test for Carboxylic Acids.

1. Ester Test.

In a clean test tube take about 2 mL of ethyl alcohol and add 2-3 drops of

concentrated H2SO4 and then add a pinch of the given compound in the test tube. Heat

the test tube over the water bath gently and blow the air from test tube towards your

nose, a pleasant fruity smell due to formation of ester of carboxylic acid is confirmed.

2. Sodium Hydroxide Test.

Take the given organic compound in a clean test tube and add sodium hydroxide

solution in the test tube and shake it well to dissolve the compound, if the compound

dissolves it conforms the presence of aromatic acid which has been dissolved due to

the formation of sodium salt of the given aromatic acid. Now add a few drops of conc.

HCl to the test tube immediately white colours precipitates are formed indicating the

reformation of aromatic acid and sodium chloride.

RESULTS REQUIRED:

PARAMETERS: Analysis of carboxylic acid group present in the given compound.

RELATIONSHIPS TO BE DETERMINED: Presence of acid functional group in given

compound

GRAPHS/PLOTS: NA

ERROR ANALYSIS:


CAUTIONS:

While burning the compound in the flame spatula needs to be very clean.

Heating of the alcohols should always be done over the water bath, as heating on the

naked flame it catch fire

Concentrated acids must be added very slowly to the aqueous solution drop wise.


EXPERIMENT NO. 4

EXPERIMENT: Determination of the dissociation constant of acetic acid using pH-meter.

EQUIPMENTS REQUIRED: pH meter, 100 mL beaker, pH electrode, pipette, burette,

funnel etc.

CHEMICAL REQUIRED: Acetic acid, sodium hydroxide, buffer solution of pH 4 and pH

7.


i. Students will learn the basics of pH meter and how to use pH meter.

ii. To monitor the total pH of a solution and to determine equivalence point of titrations

that involves ions.

iii. Students will learn how to use a calibration curve to determine the unknown strength

of a solution.

iv. To calculate the dissociation constant of weak acid.

Glass electrode

To set temperature To set knob at pH

pH reading Buffer Solution of pH 4 & 7

Fig: pH meter

THEORY: The strength of an acid is experimentally measured by determining its

equilibrium constant or dissociation constant (K). Since strong acids are strong electrolytes,

they are ionized almost completely in aqueous solutions. It is not meaningful to study the

ionic equilibrium of strong acids and calculate their equilibrium constants as the unionized


form is present to such a small extent. Hence, the study of ionic equilibrium and calculation

of K is applicable only to weak acids.

e.g. Acetic acid ionizes feebly as,

CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO

- (aq)

K = [H3O+] [CH3COO

-]/[ CH3COOH]

pKa is a modern method of expressing acid strengths. pKa is determined by measuring the

changes in pH of acid solution at different amounts of the base added.

During the titration of an acid with a base, the pH of the solution rises gradually at first, then

more rapidly and until at the equivalence point, there is a very sharp increase in pH for a very

small quantity of added base. Once past the equivalence point, the pH increases only slightly

on addition of excess base. The titration curve is obtained by plotting changes in pH at

different amounts of the base added and the equivalence point is determined.


1. Switch on the pH meter after connecting the pH electrode to it .

2. With the help of temperature knob shown in the Fig set the temperature to the room

temperature.

3. Make sure that pH knob is pointing towards pH as shown in Fig.

4. Take buffer solutions of pH 4 and pH 7 which will be provided by the lab technician.

5. Put the pH electrode in pH 4 solution and set the pH on screen to 4 with calibration

knob present on the pH meter.

6. Then wash the electrode and put it in pH 7 solution and set 7 on the screen.

7. Your pH meter is now ready to to take readings of unknown solutions. Don‟t touch

any of the button now onwards till the end of the experiment.

8. Pipette out 50 mL of the given weak acid into a 100 cm3 beaker. Immerse electrode

assembly into the acid. Measure the pH of the acid.

9. Fill a burette with the base (0.1 N sodium hydroxide).


10. Add 0.5 mL of the base from burette to the acid. Stir the solution thoroughly and

measure the pH after addition.

11. Continue adding 0.5 mL of base and noting down pH at each successive addition.

12. When the pH begins to show a tendency to increase rapidly (e.g. from 6-9), add only

small increments (say 0.1 mL) of the base and measure the pH after each addition.

Continue till there is only a slight increase in pH on the addition of the base.

13. Again add 1mL of base for 3-4 time more and note down the change in pH.

14. Plot a graph of pH (ordinate) against the volume of sodium hydroxide added

(abscissa). Determine the equivalence point and hence the pH at half equivalence

point. This gives the pKa value of the acid.

OBSERVATIONS AND CALCULATIONS:

S. No Volume of NaOH added

(ml.)

Observed pH

1.

2.

3.

4.

-

0

0.5

1.0

1.5

-

1) Equivalence point =........................(on x-axis i.e. vol. of NaOH in mL)

2) Half equivalence point =............... (take half of the vol. of NaOH at Eq. Pt. in mL)

3) pH at half equivalence point =………………(on y-axis)

pKa of the given weak acid = pH at half equivalence point =....................

pKa = - log Ka

Ka = Antilog10 (-pH at half equivalence point)

Ask for the original value from your instructor for error calculations.

REQUIRED RESULT:

PARAMETER USED: pKa = - log Ka, hence the value of dissociation constant (Ka) can be

calculated


RELATIONSHIPS TO BE DETERMINED: Effect of addition of a base to an acid.

Graph:

Equivalence point

pH

Volume of alkali added(ml)

With the help of plot, Equivalence point is calculated.

ERROR ANALYSIS: To obtain the error bars.

CAUTION:

1. Handle the glass electrode very carefully.

2. Switch on the pH meter at least 10 minutes before the start of the measurements.

3. Stir the solution thoroughly before taking the reading.

4. Let the reading stabilize for tome time (15 seconds) before taking the reading.

Half equivalence point


EXPERIMENT NO. 5

EXPERIMENT: To find the strength of hydrochloric acid solution (approximately N/10) by

titrating it against sodium hydroxide solution conductometrically.

EQUIPMENTS REQUIRED: Conductivity Bridge, conductivity cell, beaker, funnel,

burette and pipette.

MATERIALS REQUIRED: 00.1 N NaOH solution and approximately HCl solution.


1. Students will learn the basics of Conductometer and how to use Conductometer.

2. To monitor the total conductance of a solution and to determine the end points of

titrations that involve ions.

3. Students will learn how to use a calibration curve to determine the unknown strength of a

solution.

4. Students will get knowledge about conductometric titrations.

To set Cond. or Cell constant

Fig.: Conductometer

THEORY: Conductometry can be used to detect the equivalence point (end point) of a

titration. This method is based upon the measurement of conductance during the course of

titration. The conductance varies differently before and after the equivalence point. This is

due to the reason that electrical conductance of a solution depends upon the number of icons

Conductivity Electrode

To set at the value of cell

constant

Range

To set temperature

Conductivity Electrode

Value of Cell Constant

http://www.britannica.com/EBchecked/topic/186728/end-point


present and their ionic mobilities i.e. speeds. When conductance values are plotted against

volume of titrant added, two straight lines are obtained; the point of intersection of the lines

gives the end point. For studying HCl vs NaOH titration, a know volume of HCl is taken in a

beaker and NaOH solution in the burette. The conductance of acid solution is noted initially

as well as after successive additions of small amounts of NaOH solution.Conductance of acid

solution in the beginning is very high due to presence of highly mobile H+

ions. On adding

NaOH solution, the H+

ions are replaced by slow moving Na ions, decreasing the

conductance of solution.

[H+

+ Cl -] + [Na

+ + OH

-] Na

++ Cl

- + H2O

When neutralization is complete, further addition of NaOH will cause the

conductance to increase due to excess of highly mobile OH- ions. The conductance will thus

be minimum at the equivalence point. Thus if conductance values are plotted against the

volume of NaOH added, a curve of the type xyz is obtained.

The point of intersection (i.e. point Y) corresponds to the end point.


Determine the cell constant of the given conductivity cell which is written on the neck

of the cell.

Connect the conductivity cell to the conductometer.

Set the function switch to check position. Display must read 1.000. If it does not, set it

with CAL control at the back panel.

Put the „Function Switch‟ to „Cell Constant‟ and set the value of the cell constant

determined in step-1 with the help of cell constant Knob shown in the Fig.

Set the temperature control to the actual temperature of the solution under test.

Rinse the conductivity cell with the solution whose conductivity is to be measured.

Take 50 ml of the given HCl in a 100 ml beaker.

Wash the conductivity cell with distilled water and then rinse it with it with the given

HCl solution. Dip the cell in the solution taken in the beaker.

Set the range with the help of “range” knob shown in the Fig to 200.

Set the „Function Switch‟ to „Conductivity‟ and read the display. This will be the exact

conductivity note it down.


Take alkali (NaOH) in the burette and add 0.2 mL of it into the beaker containing HCl.

Stir and determine the conductivity.

Repeat the procedure of addition of 0.2 mL of NaOH and noting down the

conductivity in the observation table.

Take 25-30 readings in the ways. After each addition, stir the solution gently.

Plot a graph between observed conductivity value along Y-axis against the volume of

alkali added along x-axis. The point of intersection gives the amount of alkali required

for neutralization of acid.

OBSERVATIONS AND CALCULATIONS:

Volume of HCl taken = 50 ml

Normality of NaOH solution = 0.1 N

OBSERVATION TABLE

S. No Volume of NaOH added

(ml.)

Observed conductivity

(mmoh/cm)

1.

2.

3.

4.

5

-

22.

23.

24.

0.2

0.4

0.6

0.8

1.0

-

-

-

-

From graph the volume of NaOH used is (calculated by drawing perpendicular on X-axis

from the point of intersection) = A ml (also called as equivalence point).

Applying normality equation


N1V1 = N2V2

(HCl) (NaOH)

N1 x 50 = 0.1 x A

N1 = a N

We know strength in grams per litre = Normality x Eq. Wt.

Therefore, strength of acid = a x 36.5 g/litre = Y g/litre.

REQUIRED RESULT: Strength of given HCl solution = Y g/litre.

PARAMETERS: Conductometric titrations are used to calculate conductance and strength

of solutions. Conductivity meters are used in conjunction with water purification systems,

such as stills or deionizers, to indicate the presence or absence of ion-free water.Normality

equation

RELATIONSHIP TO BE DETERMINED: Effect of adding a base to an acid on

conductivity and strength of an acid

PLOT: When a graph is potted between volume of the alkali added and conductance then a V

– shaped graph is obtained. The point of intersection will give the end point.

X z

Y end point

Volume of alkali added (ml)

ERROR ANALYSIS: To obtain the error bars.

% error: Ask your instructor for the actual value and the calculate error in your result.

CAUTIONS:

1. The solution taken in the burette should be about ten times stronger than that taken in

the beaker so that the volume change of latter solution is negligible on the addition of

the former solution.

2. After every addition of NaOH solution, the solution must be stirred thoroughly.

Co

nd

uct

ance

Co

nd

uct

ance

Co

nd

uct

ance

(oh

m)

http://www.britannica.com/EBchecked/topic/637217/water-purification


EXPERIMENT NO. 6

EXPERIMENT: To test the validity of Beer-Lambert‟s law using colorimeter and to

determine unknown concentration of solution.

EQUIPMENTS REQUIRED: Colorimeter, test tubes, burette, 50 cc measuring flask,

MATERIALS REQUIRED: Distilled Water, 0.01 M Potassium permangnate.


1. Students will learn the basics of colorimetry and how to use colorimeters

2. Students will gain practice in preparing solutions through dilution and in calculating

solution concentrations

3. Students will use algebraic representations to describe data

4. Students will learn how to use a calibration curve to determine the unknown

concentration of a solution

Fig: PhotoColorimeter

THEORY: When a monochromatic light of intensity I is incident on a transparent medium, a

part of it is absorbed, a part, I, is reflected and the remaining part, I, is transmitted.

In case of aqueous solutions, is negligible as compared to and .

According to Beers Lambert‟s law the decrease in the intensity of the incident light is

proportional to the thickness of the absorbing medium and the concentration of the solution.

(1)

Where C is concentration of solute expressed in mole/litre, l is the length of the cell and is a

constant characteristic of the solute called molar extinction coefficient or molar absorptivity.

To set O.D at zero

Filter (to set at λ)


Further also called as optical density (OD) or absorbance (A). Since absorbance A of the

medium is given by

(2)

From equation (1) and (2)

A = Ɛ Cl

Transmittance, T of a solution is the ratio of i.e., the fraction of incident light transmitted

by the solution.

A plot between absorbance and concentration is expected to the linear. Such a straight line

plot, passing through the origin, shows that Beer- Lambert‟s law is obeyed. This plot, known

as calibration curve can also be also employed in finding the concentration of a given solution.

Outline of procedure:

Connect the instrument to the mains and put on the power switch.

Adjust the wavelength knob to the 40 wavelength region on scale (approximately).

Open the lid on the cell compartment and insert a cuvette containing the distilled water.

Close the lid.

Adjust the reading on the digital screen to zero optical density with the knob shown in

the Fig.

Remove the cuvette and close the lid tightly again. Empty the cuvette and rinse it with

standard solution of KMnO4 (0.001 IM) [which will be provided to you by lab

technician]. Fill the cuvette this solution and note the optical density.

Change the wave length to the next high value using set wavelength knob every time

and note down corresponding optical density. Make table with wavelength on LHS

and OD at RHS. Table 1:

S.No. Wavelength (λ) in nm O.D.

1.

2.

3.


Plot a graph between wavelength on the x-axis and O.D. on the y-axis. Find the value of

λmax ( O.D is maximum)

O.D

λmax

Wavelength (nm)

Now set the λmax value in the colorimeter with the help of the knob used for setting

wavelength and place the cuvette containing Distilled water in the cell compartment. Set

O.D to zero again.

Prepare KMnO4 solution in water with composition 10%, 20%, 30%, 40%,------------

100%. 10% composition means 10ml of KMnO4 and 90ml of water or 1ml of KMnO4

and 9ml of water. Make table with Concentration on LHS and OD on RHS.

Table 2:

S.No. Composition (%) O.D.

1. 10

2. 20

3. 30

4. 40

5. 50

6. 60

7. 70

8. 80

9. 90

10. Note down the absorbance (OD) of series of solution of KMnO4 prepared above (from

10% to 100%) by the method described above. Do not change wavelength now.

11. Plot a graph between O.D against composition. (If a straight line is obtained Lambert -

Beer‟s a law is verified)

12. Now take a solution of a unknown concentration and note down optical density. Find

out the concentration of the unknown solution from graph.


Table 3:

Composition (%) O.D

Unknown

Required Result: Report the results of unknown solution in gram/litre.

Parameters: λmax, absorbance and concentration.

Relationships to be determined: Between concentration and absorbance.

Graph: A plot between absorbance and concentration is expected to the linear. Such a

straight line plot, passing through the origin, shows that Beer- Lambert‟s law is obeyed. This

plot, known as calibration curve can also be also employed in finding the concentration of a

given solution.

O.D

Composition of KMnO4 solution(%)

Scope of result: This experiment is used to study the absorbance power of different solutions

and also to find the unknown concentration of solution

Error Analysis: To obtain the error bars.

% error: Ask your instructor for the actual value and the calculate error in your result.

Cautions::

1. Handle the glass cuvettes very carefully.

2. Switch on the colorimeter at least 10 minutes before the start of the measurements.

3. There should be no air drop outside the cuvette.

4. Let the reading stabilize for tome time (15 seconds) before taking the reading.


EXPERIMENT NO. 7

EXPERIMENT: Estimation of nickel in the given sample using DMG.

EQUIPMENTS REQUIRED: Beakers, Suction pump, sintered glass crucible, oven, glass

rod etc.

CHEMICAL REQUIRED: Nickel ammonium sulphate, Dimethyl glyoxime, concentrated

ammonia, Ethanol, HCl etc.


To know about gravimetric analysis of ions

The students will learn how the transition complex formation is helpful to determine

amount of particular ion in given salt.

How to use sintered glass crucible and suction pump.

CHEMICAL REACTION:

Red ppts

N

N

H3C

H3C

OH

OH

NiSO4 2NH4OH Ni

N

NNC

O

O

H3C

H3CC

CCH3

CH3

O

O

H

H

NC

(NH4)2SO4 2H2O2

THEORY:

Nickel dimethyl glyoxime is prepared by the action of alcoholic solution of dimethyl

glyoxime on soluble nickel salts such as Nickel chloride or Nickel ammonium

sulphate in presence of NH4OH solution or alkaline medium. Dimethyl glyoxime is a

chelating agent. It forms a coordination complex with Ni2+

ions. The coordination number of

the central metal atom is 4. The oxidation number of Ni is +2. The complex has square planar

geometry.

Procedure: Note: Apparatus should be cleaned properly prior to use, glass beakers shall be

first washed with some organic solvent like acetone/methanol and subsequently rinsed with

conc. HCl in order to remove traces of ammonia solution



1. Dissolve 1.0 g of nickel ammonium sulphate or Nickel chloride in distilled water in a

beaker and dilute it with 20 ml of distilled water. Add 1.0 ml of concentrated HCl.

2. Dissolve 0.6 g of dimethyl glyoxime in 15 ml of ethyl alcohol in a seaparate conical

flask.

3. Add dimethyl glyoxime solution to nickel ammonium sulphate solution along with

stirring.

4. Heat the mixture solution to 60-70oC on water bath.

5. Add 6N NH4OH solution (1:1 NH3) or Ammonia solution slowly with constant

stirring till precipitation starts. Add excess of 6N NH4OH solution (means a few drops

more even after precipitation)

6. Allow the reaction mixture to stand for about 20 minutes so that mixture to settle

down.

7. Separate the precipitate by filtration through a sintered glass crucible (ask for this

from your lab technician) under suction and wash with cold water.

8. Remove the brilliant red precipitate formed and dry in air oven.

9. Note the colour and weight of the product formed.

Calculation:

288.69 58.69

Hence, weight of Nickel = 0.2033 X weight of the precipitate

RESULTS REQUIRED: Color of the compound……….

Weigh of the ppt. …………….

PARAMETERS: color and weight

RELATIONSHIP TO BE DETERMINED: weight of the samples with the molecular

weight of known molecules.

PLOTS/GRAPHS: NA

ERROR ANALYSIS:

% error: Calculate the actual amount of Ni in starting Nickel chloride taken for the

preparation and calculate the % error.


EXPERIMENT NO. 8

EXPERIMENT: Determination of the rate constant of hydrolysis in case of ethyl acetate

catalyzed by HCl.

EQUIPMENTS REQUIREMENTS: Six conical flasks, burette, pipette,

MATERIALS REQUIRED: 0.1 NaOH, methyl acetate, 0.5 N HCl, Stop watch, Water bath,

Phenolphthalein.

LEARNING OBJECTIVE:

(1) To gain knowledge about chemical kinetics of psuedounimolecular reactions.

(2) To determine the rate constant of a reaction

(3) Student will learn how to find the order of reaction with the help of rate constant at

different time intervals.

(4) To prove that order of reaction is experimental concept.

THEORY: The reaction is catalysed by H+ ions of an acid (HCl). This reaction is an example

of psuedounimolecular reactions. Since water is present in large excess, its concentration is

practically constant throughout the reaction. The concentration of HCl (catalyst) also remains

constant. Therefore, the rate of reaction depends upon only on the concentration of ester.

Rate = -dx/dt = k [CH3COOC2H5]. Hence reaction is of first order.

During the hydrolysis of ester, acetic acid is produced. Therefore, the progress of reaction is

followed by determining the amount of acetic acid formed at different time intervals.

CH3COOC2H5 H2OH

CH3COOH C2H5OH+ +

A definite quantity of the reaction mixture is withdrawn after different time intervals and is

titrated against a standard solution of alkali. The amount of alkali used is equivalent to the

total amount of HCl present initially and the amount of acetic acid formed. The volume of

alkali used at the start of reaction is equivalent to amount of HCl alone. Hence, the amount of

acetic acid formed (x) after different intervals of time can be calculated. The amount of acetic

acid formed at the end of reaction is equivalent of initial concentration of ester (a). Suppose

the volumes of alkali used required for the reaction, at the start, after time t and the end of

reaction are V0,Vt,Vα respectively, then initial concentration of ester (a) is proportional to V∞-

V0. The concentration of ester after time t is V∞ - Vt


K = 2.303 log V∞ - V0

t V∞- Vt

If the value of K comes out constant during different intervals of time, then order of reaction

will first order.


Take 50ml of 0.5 N HCl in a clean dry 250mL conical flask and about 10mL of pure

ethyl acetate in a test tube, cork both of them and place them in a thermostat or water

bath at 45-50 C. (Take water in a big beaker and put conical flask and test tube in

it).

Keep the 0.5N HCl and ethyl acetate in the thermostat or water bath for about 10 min to

allow them to acquire the temperature of the bath (40-45 oC).

In the mean time, fit the burette properly and fill it with 0.1N NaOH solution.

At about 9 minutes add 25mL of ice cold water in a separate conical flask.

After 10 minutes pipette out 5mL of the ethyl acetate from the test tube and add it to the

flask containing 50 mL of 0.5 N HCl. Start stop watch at this moment from zero again.

Shake the contents for 2-3 seconds and immediately pipette out 5 mL of reaction

mixture and transfer it at once to first conical flask containing ice cold water.

Keep the flask containing reaction mixture in the water bath again, it should be noted

that reaction mixture will be kept in the water bath at constant temperate throughout the

experiment.

Titrate the solution in conical flask containing 25 mL ice cold water and 5 mL of

reaction mixture against 0.1 N NaOH taken in burette by using phenolphthalein as

indicator. Appearance of pink colour is end point. The volume of 0.1 N NaOH used

against the withdrawn sample of the ester and dil. HCl mixture is taken as V0.

After about 9 minutes add 25mL of ice cold water in a separate conical flask.

Pipette out 5 mL of mixture and add it to the conical flask containing ice cold water

after 10 min. Titrate it against 0.1 N NaOH . This gives Vt after 10 min.

Repeat the above procedure after every 10 min for taking readings upto 60 minutes.

Place the remaining reaction mixture in the separate water bath at 60-70oC for about half

to one hour time. Pipette out 5 mL of mixture and titrate it against alkali solution. This

gives V∞.


Observations and Calculations:

S. No Time

(min)

Volume of

NaOH (ml)

V∞ - V0 V∞ - Vt log (V∞ - Vt) log(V∞ - V0)

1. 0 V0

2. 10 V10 V∞ - V10

3. 20 V20 V∞ - V20

4. 30 V30 V∞ - V30

5. 40 V40 V∞ - V40

6. 50 V50 V∞ - V50

7. 60 V60 V∞ - V60

8. ∞ V∞

Calculations: Calculation of K at time 10, 20, 30 and 40 min will be done by the following

formula:

K = 2.303 log V∞- V0 min -1

t V∞- Vt

log(V∞- Vt ) = - Kt/2.303 +log (V∞- V0) is equation of straight line.

REQUIRED RESULT: Rate constant of ethyl acetate at given temperature is ..........

PARAMETER USED: Rate constant (K): The value of K comes out constant during

different intervals of time, then order of reaction will first order.

RELATIONSHIP TO BE DETERMINED: between K, t and volumes of NaOH

used.

GRAPH

log(V∞ - Vt)

t (min)

slope = - K/2.303 With the help of slope, the value of rate constant can

be calculated.


ERROR ANALYSIS:NA

Scope of result: We study the kinetics of hydrolysis of esters; hydrolysis of ethyl acetate has

a very rapid rate that could be carried in short time.

CAUTIONS:

Use the ice cold water only.

Perform the titrations properly.

Always take alkali in burette.


EXPERIMENT NO. 9

EXPERIMENT: Determine the strength of given solution of ferrous ammonium sulphate

(Mohr‟s salt) by titrating against potassium dichromate solution.

EQUIPMENT REQUIRED: Volumetric flask, 50 mL Burette, pipette, funnel etc.

MATERIAL REQUIRED:

Potassium dichromate K2Cr2O7 reagent grade, Ferrous ammonium sulphate, Sulphuric Acid

(concentrated).


Student will learn how to calculate the exact normality of Ferrous ammonium sulphate

(Mohr‟s salt) by titrating with potassium dichromate solution.

To gain knowledge about redox titration.

Theory: This experiment is an example of redox titration. The loss of electrons is oxidation;

the gain of electrons is reduction. Reduction/oxidation (redox) processes occur when

electrons are transferred from a donor species (the reducing agent 2FeSO4(NH4)2SO4) to

another acceptor species (the oxidizing agent K2Cr2O7).

Reactions: The reaction between Potassium dichromate and Mohr‟s salt can be represented

as:

Molecular equations:

K2Cr2O7 + 4H2SO4 → K2SO4 + Cr2(SO4)3 + 3[O]

2FeSO4(NH4)2SO4 +H2SO4 +O → Fe2(SO4)3 +2(NH4)2SO4 + H2O

Ionic Equations:

2Cr2O72–

+ 14H+ + 6e

– → 2Cr

3+ + 7H2O

Fe2+

→ Fe3+

+ e-


Titration of K2Cr2O7 soln. with Mohr’s salt solution

Pipette out 10 mL of Mohr‟s salt solution in the conical flask.

Add approximately 4 mL of Conc. H2SO4 to the same flask.

Add 2-4 drops of the Ferroin indicator.

Titrate the reaction mixture with potassium dichromate solution taken in burette till a

colour change from Wine red - green is obtained.

Repeat the titration for three concordant readings.

Observations and Calculations

Volume of Mohr‟s salt solution used in each titration = 10 mL

Indicator used = Ferroin indicator


Color change at end point = Wine red - green

Equivalent weight of Mohr‟s salt = 392

Table-1

S. No. Burette readings Volume of K2Cr2O7

Consumed (R2 - R1) mL Initial Final

1.

2.

3.

4.

5.

6.

Thus applying the normality relation NMohr VMohr = NdichrVdichr

Thus NMohr = NdichrVdichr/VMohr

Thus strength (g/L) of Mohr‟s salt solution = Normality x Eq. Wt

REQUIRED RESULT: Strength of given Ferrous ammonium sulphate solution= … g/litre.

PARAMETERS USED: Volumetric parameters, Normality equation

RELATIONSHIP TO BE DETERMINED: Redox titration are used to calculate strength

of solution, based on an oxidation-reduction reaction between analyte and titrant. Many

common analytes in chemistry, biology, environmental and materials science can be

measured by redox titrations.

PLOT/GRAPH: NA

ERROR ANALYSIS:

CAUTIONS:

Always take Potassium dichromate solution in burette.

Potassium dichromate acts as oxidizing agent in acidic medium. Therefore always

add dil. H2SO4 in the reducing agent.

Read the upper meniscus while taking burette readings because K2Cr2O7 is coloured.


EXPERIMENT NO. 10

EXPERIMENT: Separation of amino acids by thin layer chromatography.

EQUIPMENTS REQUIRED: Glass plates, beaker, Glass rod,

MATERIAL REQUIRED: Silica gel (for TLC), Glycine, Leucine, acetic acid, n-butanol,

water, alcoholic solution of ninhydrin


Students will learn the basis of TLC and how to prepare the TLC plates.

To calculate retention factor (Rf).

Students will learn how to separate the mixture of organic compounds on the basis of

Rf value

Theory: Thin Layer Chromatography (TLC) is used extensively for qualitative analysis

(tentative identification of mixture of two or more organic compounds). In this technique a

small amount of the material (to be separated), dissolved in an appropriate solvent, is spotted

near one edge of the plate covered with thin layer of adsorbent.

The adsorbent is usually a thin layer of alumina or silica gel. After the sample has

been deposited on the adsorbent, the coated plate is placed in a beaker containing small

amount of solvent in such a manner that the lower end of plate dips 1-2 cm below the surface

of solvent. The solvent rises through the adsorbent by capillary action and the various

components of the mixture ascends at different rates, depending on their different affinities to

the adsorbent.

This results in the separation from one another. When the solvent front has almost

reached the top of the adsorbent layer or three-fourth of it, the plate is removed from the

beaker, dried and examined.

TLC involves the following steps:

(a) Preparation of a thin layer plate

(b) Application of the materials to be separated on the plate

(c) Development of the chromatogram plate in a solvent

(d) Visualization or Location of components

(e) Calculation of Rf values.

Retention Factor: The movement of any substance relative to the solvent front in a given

chromatographic system is constant and characteristic of a substance.


Rf value =

= b/a


Take small amount of silica gel in a beaker and dissolve it in distilled water with

constant stirring by a glass rod. Continue to stir until a uniform paste free from air

bubbles is formed. Add some more water to obtain slurry of suitable consistency. (OR

this slurry may be provided to you by the lab technician).

Mark the base line on the glass plate about 1 cm from the bottom edge of the glass

plate. (Mark only at edges with pencil). This line is just to take an idea that where 1

cm distance lies from the bottom where sample mixture is to be applied.

Pour the slurry on to the clean and dry plate and prepare a uniform thin layer by glass

rod.

Allow the layer to dry for 5-10 minutes and then heat the plates in an electric oven at

100-120oC for about 20 min.

Prepare solution by mixing two or more different organic compounds. (Given to you

by lab technician).


At base line and apply small sample of the mixture with the help of thin capillary

tube in the centre. Take care the spots must be as small as possible.

Allow the spot to dry. Place the glass plate in a beaker containing solvent to a depth of

about 1 cm or less than that and allow the solvent to flow up until it nearly reaches the

top of the plate. (take due care that spot must not touch the solvent)

Remove the plate from the beaker, mark the position of the solvent front (on edge of

the plate with pencil) and allow the solvent to evaporate.

Spray with alcoholic ninhydrin and dry TLC plate in the oven for 2 min.

Spots will develop on the plate. Take measurements of the distance moved by

solvents and each component from the 1 cm mark.

Calculate the Rf values of the components in the mixture.

REQUIRED RESULT:

Rf value of Glycine= …..

Rf value of Leucine= …

PARAMETER USED: Retention Factor (Rf): The movement of any substance.

RELATIONSHIP: determination of relative Rf value of mixture of compounds

PLOT/GRAPH: Draw slide with observed lengths of solvent and two components

ERROR ANALYSIS: NA

Scope of Result: Thin Layer Chromatography (TLC) is used extensively for qualitative

analysis and tentative identification of mixture of two or more organic compounds.TLC is a

useful screening technique in clinical chemistry; for example, it can be used to detect the

presence of drugs in urine.

Cautions:

Make the slurry very carefully; it should not be very thick or very thin.

Always prepare fresh silica slurry.

Make sure that sample dot is always outside the solvent layer.

Spots of mixture must be as small as possible.

Dry TLC plates carefully.


Some Common Apparatus used in Chemistry lab:

Apparatus for TLC TLC plate in solvent in development phase

TLC plate in solvent in development phase Ninhydrin Spray apparatus


Burette with stand How to read the burette

Graduated Pipette

Pipette with filler


Conical Flasks

Measuring Cylinders

Droppers

Beakers

Spatula

Water bath

chemistry practical

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