LMCHE102

LMCHE 102: CHEMISTRY LAB

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LABORATORY MANUAL

CHE102

CHEMISTRY LAB


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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)

3-4

2 Preparation of sodium trioxalatoferrate(III) Na2[Fe(C2O4)3] .9H2O

compound. 5-6

3 To find the distribution coefficient of benzoic acid between benzene and

water. 7-9

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

10-12

5 Determination of Strength of hydrochloric acid solution(approximately N/10)

by titrating it against sodium hydroxide solution conductometrically 13-16

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

determine unknown concentration of solution 17-20

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

21-22

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

using an alkali. 23-26

9 Determine the strength of given solution of ferrous ammonium sulphate by

titrating against potassium dichromate solution 27-28

10 Separation of a mixture of organic compounds by thin layer

chromatography. 29-31

11 To determine the surface area of given solid by adsorption of acetic acid from its aqueous solution.

32-36

Textbook: 1. J. Mendham, R. C. Denney, J. D. Barnes, and R. C. Denney, Vogel’s

Quantitative Chemical Analysis, 6th Edition, Prentice Hall, 2000. (Physical

Chemistry)

Other Reading:

2. J. B. Yadav, Advance Practical chemistry, Krishna Publications, Merrut 12th

Revised Edition 3. Basset J Denny. R. C., Jeffery C. H. and Mendham J., Vogel’s Textbook of

Quantitative Inorganic analysis, ELBS, 1978. (Inorganic) 4. N. K. Vishnoi Advanced Practical Organic Chemistry, Vikas Publications, New

Delhi, 2nd

edition 5. M.S. Saini, Senior practical chemistry-Vol-III, Modern Publishers-2006


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Experiment No. 1

Title: TO DETERMINE THE HARDNESS OF THE GIVEN HARD WATER SAMPLE

BY EDTA METHOD. PROVIDED STANDARD HARD WATER. (1 mL of S.H.W.=1

mg of CaCO3).

Equipments to be used: Burette, Titration flask, Pipette, Beakers, funnel

Chemicals Used: Standard Hard water, EDTA solution, EBT, Buffer Solution

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 colour is restored.

H2Y2- + MIn+ → MY-2 + HIn+2 + H+

Procedure:

(a).Standardisation of EDTA solution:

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 colour appears.

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

changes to blue colour. This is the end point. Recovered the volume of EDTA

consumed as A ml. Repeated the procedure to get at least three concordant

readings.

(b). Determination of Total Hardness:-


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Pipette out 50ml of tap water in the titration flask. Add to it 5-6 ml of buffer solution

and four drops of Eriochrome Black-T indicator. A wine red colour appears. Titrate

this solution against EDTA solution taken in a burette till wine red colour changes to

blue colour. This is the end point. Recovered the volume of EDTA consumed as B

ml. Repeated the procedure to get at least three concordant readings

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

I ml of EDTA = 10/A mg of CaCO3

(b).Calculation of total hardness:

50 ml of hard water sample = B ml of EDTA

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

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

1 ml of hard water sample = B x 10/A x 1/50 mg of CaCO3.

1000 ml of hard water sample = B x 10/A x 1/50 x 1000 mg of CaCO3.

Hence total hardness = 200 x B/A ppm

Result: The hardness of water is...........

Scope of result: The determination of water hardness is a useful test that provides a measure of quality of water for households and industrial uses. Originally, water hardness was defined as the measure of the capacity of the water to precipitate soap.


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Experiment No. 2

Title: Preparation of Sodium trioxalato ferrate(III)

EQUIPMENTS TO BE USED: China dish, water bath, glass rod, Buchner funnel, glass funnel, weighing machine. CHEMICAL REQUIRED: Mohr’s Salt ((NH4)2SO4•FeSO4•6H2O), 3M H2SO4, oxalic acid, water, ethanol, 30% H2O2, acetone. LEARNING OBJECTIVES: (i) To know about gravimetric analysis of ions

(ii) The students will learn how the transition complex formation is helpful to

determine amount of particular ion in given salt.

(iii) How to use sintered glass crucible and suction pump.

THEORY: Sodium trioxalato ferrate (III) is a coordination compound; with a coordination number of six and oxidation state of +3. The shape of the complex is octahedral and it is sp3d2 hybridized. Na3[Fe(C2O4)3] •3H2O is highly photosensitive. The bright green crystals on exposure soon become covered with a yellow powder of ferrous oxalate. 2Na3[Fe(C2O4)3]•3H2O →2FeC2O4•2H2O + 2CO2 + 3Na2C2O4 + 2H2O

PROCEDURE:

Preparation of Ferrous Oxalate:

Weigh ferrous ammonium sulphate (Mohr’s Salt) (NH4)2SO4•FeSO4•6H2O.Transfer

the ferrous ammonium sulphate to a beaker containing 25 mL of warm water

acidified with 1 mL dilute 3M H2SO4. Dry the vial and weigh the empty vial. Record

the mass of ferrous ammonium sulphate used. To this solution add 25 mL of oxalic

acid solution (5 g/50mL). Cautiously heat the solution to boiling on a hot plate,

continuously stirring the solution to avoid bumping. Remove and then allow the

precipitate to settle. Decant and discard the clear supernatant liquid (avoid loss of

precipitate). Stir the remaining precipitate with 25 mL of hot water, the solubility of

FeC2O4•2H2O is .022 g/100 mL in cold water and .026 g/100 mL in hot water. Decant

and discard the wash liquid. Add 25 mL of hot water to the precipitate and stir, then

filter by suction using a small Buchner funnel and wash the precipitate thoroughly

with small portions (10-20 mL) of hot water. Finally rinse with 20 mL of acetone.

Weigh the final product.


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Preparation of Sodium trioxalato ferrate(III): Suspend the washed and dried ferrous oxalate in 30 mL of warm solution containing

3.5 g of Sodium oxalate (Na2C2O4•H2O). Place this solution in an ice bath and add

slowly with continuous stirring 7-10 mL of 30% H2O2 in very small portions, then heat

the mixture to boiling and dissolve the precipitate by adding in one portion 7 mL of a

solution of oxalic acid (containing 2.5 g/25 mL). A further 3 mL is then added drop by

drop using an eye dropper (an excess of oxalic acid is to be avoided). The liquid

should be near the boiling point while these additions are being made. Filter through

a Buchner funnel. Transfer the hot solution to a clean beaker and while still hot, add

to the filtrate 15 mL of 95% ethanol and redissolve any precipitated crystals by gently

heating. Then store the solution in a dark cupboard for crystallization. After

crystallization has occurred, filter with suction and wash the product on the filter

paper with 20 mL of an equi-volume mixture of ethanol and water and finally with 20

mL of acetone. Draw air through the precipitate for several minutes. Weigh the dry

crystals.

RESULT: Colour of the compound……….

Weight of the ppt. …………….

SCOPE OF THE RESULT: Gravimetric analysis is helpful in estimation of amount of Fe in given complex.


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Experiment-3

Title: To study the distribution of benzoic acid between benzene and water. EQUIPMENTS REQUIRED: Pipette, Burette, titration flask, eight washed empty bottles, stoppers, separating funnel. CHEMICAL REQUIRED: Benzene, Benzoic acid, water, NaOH, HCl, Phenolphthalein. LEARNING OBJECTIVE: In this experiment we will use the volumetric technique of acid-base titration to determine the distribution of benzoic acid in the water-benzene system. From these distribution measurements we will be able to examine the two equilibria: (1) dimerization of benzoic acid in benzene and, (2) distribution of benzoic acid monomers between water and benzene. INTRODUCTION: Distribution coefficient (D) is the ratio of concentrations of a compound in the two phases of a mixture of two immiscible solvents at equilibrium. Hence these coefficients are a measure of differential solubility of the compound between these two solvents. A general treatment by Moelwyn-Hughes1 allows examination of the above two equilibria from distribution measurements. The benzoic acid water-benzene system has been studied by Huq and Lodhi. This treatment has been extended to other acids (e.g. acetic and propionic acids) distributed between water and other organic solvents. In the benzoic acid(HBz )-water-benzene system, there are three equilibria which must be considered. In the aqueous layer, benzoic acid, being a weak acid will dissociate: HBz ↔ H+(aq) + Bz-(aq) (1) and we can write the acid dissociation constant, Ka for the equilibrium as Ka = [H+] [Bz-] / [HBz] (2) In the benzene layer, benzoic acid exists both as a monomer and dimer (i.e. two benzoic acid molecules bonded together by hydrogen bonding to form a new species). This dimerization is also an equilibrium: 2 HBz↔ (HBz)2 KD = [(HBz)2]/[(HBz)2]b (3) where [HBz]b, is the equilibrium molar concentration of the monomer in benzene and [(HBz)2]b is the equilibrium molar concentration of the dimer in benzene. The third equilibrium is the equilibrium constant for the distribution of benzoic acid monomers between water and benzene: (HBz)W ↔ (HBz)b KM = [HBz]b/[HBz]W (4) The total concentrations of benzoic acid in water and benzene, CW and Cb respectively, is obtained by titration of the appropriate phase with sodium hydroxide. CW = [HBz]W + [Bz-]W = CW(1 - α) + α CW (5) where α is the degree of dissociation of benzoic acid; and

http://en.wikipedia.org/wiki/Concentration

http://en.wikipedia.org/wiki/Chemical_compound

http://en.wikipedia.org/wiki/Immiscible

http://en.wikipedia.org/wiki/Solvent

http://en.wikipedia.org/wiki/Partition_equilibrium

http://en.wikipedia.org/wiki/Solubility


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Cb = [HBz]b + 2 [(HBz)2]b (6) Using Eqs.(3) and (4) as well as noting that [HBz]W = CW(1 - α) allows Eq.(6) to be rearranged to,(7) Cb/CW(1 - α)= KM + 2KM

2KDCW(1 - α) A plot of Cb/CW(1 - α) vs. CW(1 - α) will be linear with a slope of 2KM 2KD and an intercept of KM. The value of α used in Eq.(7) is obtained from the known value of Ka. Ka = α2CW/1 - α= 6.3 x 10-5 for benzoic acid at 20oC pKa= 4.1 (benzoic acid) Cb and Cw are calculated from the titration of the respective layers. PROCEDURE: A. Preparation of the benzene/benzoic acid/water mixtures: Eight mixtures of benzene/benzoic acid/water are to be studied so that the distribution can be observed over a wide range of concentration. These mixtures are to be made according to Table I, using the burettes for benzene and the benzoic acid solution in benzene, the water should be added using the 50.00 mL pipette. Make up the mixtures in the stoppered bottles and allow the equilibrium distribution to be attained over a period of about 30 minutes, shaking frequently. During the equilibration period, practice the two titrations involved on samples of benzoic acid in benzene and in water as supplied in the laboratory. Separate the organic and the aqueous layer using separating funnel. TABLE 1: Mixture Benzoic acid Benzene Water in benzene

Mixture Benzoic acid in Benzene

Benzene Water

1 50 +0 +50

2 45 +5 +50

3 40 +10 +50

4 35 +15 +50

5 30 +20 +50

6 25 +25 +50

7 20 +30 +50

8 15 +35 +50

At the end of the equilibration period, allow the mixture to separate into two layers. Decant the upper benzene layer into a large, clean, dry test tube, and cork it at once. (This separation need not be perfect). B. Titration of the benzene layer Pipette 15.00 mL of the benzene solution from the upper layer in the test tube into a clean (not necessarily dry) Erlenmeyer flask. WIPE THE OUTSIDE OF THE PIPETTE BEFORE ADJUSTING THE TRANSFER VOLUME TO AVOID


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CARRYING OVER ANY EXTRA, UNMEASURED SOLUTION. Pipette 25.00 mL of the 0.030 M NaOH and titrate with 0.0200 M HCl until the pink color of the phenolphthalein indicator disappears (this is the end point for the titration. Since neutralization of the benzoic acid in the benzene layer with the NaOH in aqueous solution is slow because of the immiscibility of the two layers, the two layers of solution must be stoppered and well shaken before the back titration with the 0.020 M HCl. Repeat for all the eight mixtures. Cb is obtained. C. Titration of the aqueous layer Titrate 20.00 mL of the lower, aqueous layer with the 0.006 M NaOH. The lower, aqueous layer must be reached through the upper, toluene layer. This is best done by blowing gently through the pipette as it is inserted to ensure that none of the toluene layer enters. Transfer the 20.00 mL to a clean Erlenmeyer flask; wipe the pipette before adjusting the transfer volume as before. Since the concentration of the benzoic acid in the water layer will be low, expect small amounts of titrant to be used in this titration. Repeat the titration. Follow the same procedure in duplicate for the other eight solutions and record your titration values. Cw is obtained. CALCULATION:

Bottle No. Cw Cb α= (Ka/Cw)1/2 Cb/CW(1 - α) CW(1 - α)

1.

2.

A graph between Cb/CW(1 - α) Vs CW(1 - α) is plotted. The graph gives a straight line with a slope of 2KM

RESULT: The distribution coefficient of benzoic acid in benzene and water is............. PRECAUTIONS:

1. Titration should be carefully performed. 2. The bottles should be shaked properly. 3. Separating funnel should be handled by care.


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Experiment No. 4

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

Equipments to be used: pH meter, 100 cm3 beaker, glass electrode, calomel

electrode,

Chemicals to be used: Acetic acid, sodium hydroxide.

Learning objectives: (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 involve ions.

(ii) Students will learn how to use a calibration curve to determine the unknown strength of a solution.

(iii) To calculate the dissociation constant of weak acid.

Glass electrode

To set temperature To set knob at pH

pH reading

pH meter


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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.

PROCEDURE: Pipette out 50 cm3 of the given weak acid into a 100 cm3 beaker.

Immerse a glass electrode calomel electrode assembly into the acid and connect the

cell to a pH meter. Measure the pH of the acid. Fill a burette with the base (sodium

hydroxide). In the beginning, add large increments of (say 1cm3) of the base to the

acid. Stir the solution thoroughly and measure the pH after each addition. When the

pH begins to show a tendency to increase rapidly, add only small increments (say

0.1 cm3) 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.

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:


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1) Equivalence point ........................

2) Half equivalence point ........................

3) pH at half equivalence point ………………

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

pKa of acetic acid at 25oC= 4.26

Scope of the result:

Parameter used: pKa = - log Ka, hence the value of dissociation constant(Ka) can

be calculated

Plot:

Equivalence point

pH

Volume of alkali added(ml)

With the help of plot, Equivalence point is calculated.

CAUTIONS:

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


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Experiment No. 5

Title: 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, burette and pipette

Chemicals required: 0.01 N KCl solution, 0.1 N NaOH solution and approximately

0.01 N HCl.

Learning objectives: (i) Students will learn the basics of Conductometer and how to use Conductometer

(ii) To monitor the total conductance of a solution and to determine the end points of

titrations that involve ions.

(ii) Students will learn how to use a calibration curve to determine the unknown strength of a solution.

(iii) Students will get knowledge about conductometric titrations.

To set the knob at the

conductance or Cell constant

To set at the value of cell constant Range To set temperature

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

Glass Electrode

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


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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 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.

PROCEDURE:

1. Determine the cell constant of the given conductivity cell. 2. Rinse the conductivity cell with the solution whose conductivity is to be

measured. 3. Take 10 ml of the given HCl in a 100 ml beaker. 4. 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. 5. Connect the conductivity cell to the conductometer. 6. 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. 7. 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. 8. Set the temperature control to the actual temperature of the solution under test. 9. Set the ‘Function Switch’ to ‘Conductivity’ and read the display. This will be the exact conductivity note it down. 10. Take alkali (NaOH) in the burette and 1 ml of it into the beaker containing HCl. Stir and determine the conductivity. 11. Set the temperature control to the actual temperature of the solution under test.

12. Repeat the procedure in addition of 1 ml of NaOH and noting down the

conductivity in the observation table. Take 12-15 readings in the ways. After each

addition, stir the solution gently

13. Plot a graph between observed conductance 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.


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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 conductance

(ohm-1)

1.

2.

3.

4.

-

-

22.

23.

24.

0.5

1.0

1.5

2.0

-

-

11.0

11.5

12.0

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 = 1 x A

= X N

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

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

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


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Scope of result: 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.

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)

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

C

on

du

ctan

ce (

oh

m)

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




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Experiment No. 6

AIM: To test the validity of Beer-Lambert’s law using colorimeter and to determine

unknown concentration of solution.

Equipments to be used: Colorimeter, test tubes, burette, 50 cc measuring flask,

Chemical required: Distilled Water, 0.1 M Potassium dichromate.

Learning Objectives:

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

Filter (to set at λmax)

Colorimeter

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.

To set O.D at zero

Give the valve of O.D


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

According to Beers Lambert’s law the intensity of the incident light is proportional to

the length of 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. 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 = ЄCl

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.

PROCEDURE:

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

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

Open the lid on the cell compartment and insert a cuvette containing the blank solvent (distilled water). Close the lid.

Adjust the needle to 100% transmittance or zero optical density.

Remove the cuvette and close the lid tightly again. Empty the cuvette and rinse it with standard solution of KMnO4 (0.01 IM). Fill it with standard solution. Change the wave length by about 20 nm every time and note down corresponding optical density. Repeat step 4 to 8 for each


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wavelength measurements. 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 place the cuvette containing the standard solution in the cell compartment. Note the O.D and transmittance.

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.

Note down the absorbance (OD) of series of solution of KMnO4 prepared above by the method described above.

Plot a graph between O.D against composition. (If a straight line is obtained Lambert - Beer’s a law is verified)

10. Now take a solution of a unknown concentration and note down optical

density. Find out the concentration of the unknown solution from graph.

Report the results of unknown solution in gram/litre.

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

solutions and also to find the unknown concentration of solution

Plot: 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.


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O.D

Composition of KMnO4 solution(%)

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.


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Experiment No. 7

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

Equipments to be used: Beakers, Suction pump, sintered glass crucible, oven.

Chemical required: Nickel ammonium sulphate, Dimethyl glyoxime, concentrated

ammonia

Learning objectives:

(i) To know about gravimetric analysis of ions

(ii) The students will learn how the transition complex formation is helpful to

determine amount of particular ion in given salt.

(iii) 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 sulphate in

presence of NH4OH solution or alkaline medium. Soluble Nickel salt is dissolved in

water and the solution acidified with dil HCl Solution is heated to 60-70oC and then

treated with alcoholic solution of dimethyl glyoxime. Precipitate of nickel dimethyl

glyoxime is carried out by adding 6N NH4OH solution (1:1) till it gives smell of

ammonia. A brilliant red precipitate of nickel dimethyl glyoxime separates out. It is

allowed to stand for 20 minutes so that precipitates settle down.

The precipitate is separated by filtration through Buchner funnel and washed. It is

dried in air oven at 110-120oC. Its color and weigh is noted.

Procedure:


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1. Dissolve 4.0 g of nickel ammonium sulphate in distilled water in a 100 ml

beaker and dilute to 75 ml. Add 1.0 ml of concentrated HCl.

2. Dissolve 2.4 g of dimethyl glyoxime in 50 ml of ethyl alcohol in a 250 ml

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) slowly with constant stirring till precipitation

starts.

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 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

Result: Color of the compound……….

Weigh of the ppt. …………….

Scope of result: Gravimetric analysis is helpful in estimation of amount of Ni in

given salt.


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Experiment No. 8

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

catalyzed by HCl at room temperature

Equipments requirements: Six conical flasks, burette, pipette,

Chemical required: 0.1 NaOH, methyl acetate, 0.5 N HCl, Stop watch, Water bath,

Phenolphthalein.

Learning Objective: (i) 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


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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.

Procedure : 1. 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 or near room temperature. 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.

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

add 25ml of ice cold water in six conical flasks.

3. Pipette out 5ml of the ester from the conical flask and add it to the flask containing

50 ml of 0.5 N HCl.

4. Shake the contents, pipette out 10ml of reaction mixture and transfer it at once to

first conical flask containing ice cold water. Titrate it 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

5. Pipette out 10 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.

6. Repeat the above procedure after every 10 min.

7. Place the remaining reaction mixture it in the separate water bath at 60-70oC for

about one hour. Pipette out 10ml of mixture and titrate it against alkali solution. This

gives V∞.


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Observations;

S. No Time (min) Volume of NaOH (ml)

1.

2.

3.

4.

5.

0

10

20

30

40

V0

V10

V20

V30

V40

Calculations: 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.

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.

Parameter used: Rate constant (K): The value of K comes out constant during

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


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Plot:

log(V∞ - Vt ) slope = - K/2.303. With the help of slope,

the value of rate constant can

be calculated.

T (min)

Cautions: (i) Use the ice cold water only.

(ii) Perform the titrations properly.

(iii)Always take alkali in burette.


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Experiment No. 9 Title: Determine the strength of given solution of ferrous ammonium sulphate(Mohr’s

salt) by titrating against potassium dichromate solution.

Equipment Required: Volumetric flask, 250 mL Beakers, 50 mL Burette, pipette

Material Required:

Potassium dichromate K2Cr2O7 reagent grade, Ferrous ammonium sulphate,

Sulphuric Acid (concentrated).

Learning Objectives: (i) Student will learn how to calculate the exact normality of Ferrous ammonium sulphate (Mohr’s salt) by titrating with potassium dichromate solution. (ii) 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: 2Cr2O7

2– + 14H+ + 6e– → 2Cr3+ + 7H2O

Fe2+ → Fe3+ + e-

Procedure: Titration of K2Cr2O7 soln. with Mohr’s salt solution (a) Pipette out 10 mL of Mohr’s salt solution in the conical flask. (b) Add approximately 10 mL of dil. H2SO4 to the same flask. (c) Add 2-4 drops of the indicator N-phenyl anthranilic acid (d) Titrate the reaction mixture with potassium dichromate solution taken in burette

till a colour change from light green to blue- violet is obtained. (e) Repeat the titration for three concordant readings. Observations and Calculations

Titration Observation:

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

Indicator used = N-phenyl anthranilic acid


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Colour change at end point = green to violet

Equivalent weight of Mohr’s salt = Mol. Wt./ no. of equivalents

= 392/ 1 = 392

S. No. Burette readings Volume of K2Cr2O7

Initial R1 Final R2 Consumed (R2 - R1)

mL

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

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

Scope of the results expected: 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.

Cautions:

(i) Always take Potassium dichromate solution in burette.

(ii) Potassium dichromate acts as oxidizing agent in acidic medium. Therefore

always add dil. H2SO4 in the reducing agent.

(iii) Read the upper meniscus while taking burette readings because K2Cr2O7 is

coloured.


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Experiment No. 10

Title: Separation of amino acids by thin layer chromatography.

Equipments Required: Glass plates, beaker, Glass rod,

Chemicals Required: Silica gel (for TLC), Glycine, Leucine, acetic acid, n-butanol,

water, alcoholic solution of ninhydrin

Learning Objectives: (1) Students will learn the basis of TLC and how to prepare

the TLC plates..

(2)To calculate retention factor (Rf).

(3) 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.


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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

Procedure: 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. Pour the slurry on to the clean and dry plates 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 half an hour. Prepare solution by


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mixing two or more different organic compounds. Mark the base line on the glass

plate about 2 cm from the bottom edge of the glass plate. At about 1 cm from the

end of the plate make 3-4 equally spaced marks on the base line and apply small

samples of the mixture with the help of capillary tube. 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 and allow the solvent to flow up until it nearly reaches the top of

the plate. Remove the plate from the beaker, mark the position of the solvent front

and allow the solvent to evaporate.

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

Calculate the Rf values of the components in the mixture,

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.

Parameter used: Retention Factor (Rf): The movement of any substance

(2) Spots of mixture must be as small as possible.

(3) Dry TLC plates carefully.


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Experiment No. 11

Title: Determine the adsorption of acetic acid from aqueous solution by active charcoal at room temperature. Equipment Required: Conical flasks with corks, beakers, pipette, funnel, filter paper, glazed paper. Material Required: Activated charcoal, 0.5 M acetic acid, 0.1 M NaOH Learning Objectives:

1. Students will learn about the adsorption phenomenon.

2. Students will gain practice of preparing volumetric solutions.

3. Students will learn how to verify Freundlich isotherm.

4. Students will get knowledge about volumetric titrations.

Theory: Adsorption is phenomenon of higher concentration in the surface as compared to bulk. The substance which adsorbs is called the adsorbent while the one which gets adsorbed is called adsorbate.

A graph between the amount of substance adsorbed (adsorbate) per gram of the adsorbent versus equilibrium concentration at the particular temperature is known as adsorption isotherm. According to Freundlich isotherm:

x = k C1/n where x = weight of the adsorbed substance

m m = weight of the adsorbent

C = equilibrium concentration of the solute

The value of k and n depends upon the nature of solute, solvent and adsorbent. Taking the log of the above equation:

log x/m = log k+ 1/n log C. It is an equation of the straight line between logx/m vs. logC with slope 1/n and intercept log k.


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Procedure:

Take six clean and dry stoppered conical flasks and label them 1 to 6.

Fill the flasks as given below:

Stopper the flaks and shake the solutions thoroughly.

Weigh 2g of activated charcoal each on six different glazed papers. Transfer

the contents of one paper each, to each of the conical flasks without much

loss of time.

Cork the flasks, shake them well and keep them at room temperature for 30-

40 minutes. As far possible, each flask should be shaken for equal time.

Slope = 1/n

Intercept= log k

Flask No. Volume of water (ml) Volume of acetic acid

(ml)

1

2

3

4

5

6

---

10

20

30

40

50

50

40

30

20

10

---


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Filter the solutions in flasks separately but simultaneously. Reject the first 4-5

ml of the filtrate in each case and receive the rest in six other clean and

numbered flasks (or beakers).

Now titrate 10ml of each filtered solution against 0.1 M NaOH using

phenolphthalein as indicator. Repeat the experiment to get a set of

concordant readings.

Observations:

Room temperature = -----oC

Mass of activated charcoal in each flask, m = 2g

Volume of acetic acid solution in each case = 10ml

Bottle No.

Initial reading

Final reading

Volume of 0.1 M NaOH used (ml)

Mean Value

(ml)

1 (i)

(ii)

V1

2 (i)

(ii)

V2

3 (i)

(ii)

V3

4 (i)

(ii)

V4

5 (i)

(ii)

V5

6 (i)

(ii)

V6

Calculation of equilibrium concentration, C:

Volume of 0.1N NaOH used for 10 ml of filtrate taken from bottle 1= V1 ml

Applying normality equation,

N1 V1 = N2 V2


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(oxalic acid) (0.1N NaOH)

N1 X 10 = 0.1 V2

N1 = V2/100

i.e. equilibrium concentration in bottle 1 = V2/100 g/litre

Equilibrium concentration (C) in bottles 2-6 can be calculated similarly.

Various observations and calculations may be tabulated as follows:

Flask

No.

Original conc.

Of oxalic acid

(Co)

Equilibrium

concentration

(C)

Amount of acetic

acid adsorbed in

moles/litre

x=(Co-C)x50/1000

x/m

logc

logx/m

1 0.5 N 0.01V1

2 0.5x40/50

=0.4N

0.02 V2

3 0.5x30/50

=0.3N

0.03 V3

4 0.5x20/50

=0.2N

0.04 V4

5 0.5x10/50

=0.1N

0.05 V5

6 0.5x0/50

=0

0.06 V6

Plot of graph: To verify Freundlich isotherm, plot logx/m vs log C.


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A straight line is obtained which proves the validity of Freundlich isotherm.

Cautions:

1. Handle the apparatus carefully.

2. Alkali should be taken in the burette.

3. Cork the flask properly while shaking solutions.

LMCHE102

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