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LABORATORY MANUAL

CHE102

ENGINEERING CHEMISTRY LABORATORY

LMCHE 102: ENGINEERING CHEMISTRY LABORATORY

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 because 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 To determine the hardness of the given hard water sample by EDTA

method. You are provided with Standard Hard Water (1 ml of SHW

=1 mg of CaCO3).

4-7

2 To determine the rate constant of hydrolysis of ethyl acetate

catalyzed by HCl.

8-11

3 To determine the strength of given solution of ferrous ammonium

sulphate by titrating it against potassium dichromate solution.

12-14

4 To estimate the nickel content in the given sample using dimethyl

glyoxime.

15-17

5 To Separate the mixture of amino acids by thin layer

chromatography.

18-22

6 To Identify the elements present in given organic compound.

23-26

7 To determine the dissociation constant of acetic acid using pH-meter.

27-31

8 To determine the Strength of hydrochloric acid solution by titrating it

against sodium hydroxide solution conductometrically

32-36

9 To test the validity of Beer-Lambert’s law and also determine

unknown concentration of solution using colorimeter

37-42

Appendix

Common Apparatus/Instruments used in Chemistry lab 43-45


Experiment No. 1

1. Experiment: To determine the hardness of the given hard water sample by EDTA

method. You are provided with Standard Hard Water (1 ml of SHW =1 mg of CaCO3).

Equipment required: Burette, Burette stand, Titration flask, Pipette, Beakers, Funnel

etc.

Materials required: Standard Hard water, Sample hard water, EDTA solution,

Eriochrome black-T (EBT), Buffer Solution.

2. Learning Objectives:

I) The purpose of this experiment is to determine the hardness of water by measuring the

concentrations of calcium in water samples by titration.

II) To gain knowledge about complexometric titration.

III) To know the purpose of EDTA used

IV) To know about use of buffer solution: The buffer being used has composition

NH4Cl and NH4OH. Its pH is the order of 10.5.

V) To know the purpose of indicator

THEORY

Complexometric titrations are mainly used to determine the concentration of divalent

cations such as calcium, magnesium, zinc, copper, lead etc. Hard water is the water that has

high concentration of calcium and magnesium ions. The Ethylenediaminetetraacetic acid

(EDTA) is the most commonly used complexant. EDTA is a chelating agent. It is able to form

coordination complex with metal ions present in hard water. Although these reactions are easy

to perform, it is necessary to maintain well defined pH.

How does the indicator work?

The indicator used here is organic molecule capable of forming a colored complex

with the metal cations to be determined. The indicator-metal ion complex would be less

stable than the complex formed by the cation with the corresponding titrant (EDTA). The

indicator used is Eriochrome black T.

(i) When indicator is added to hard water it combines with free metal ions present in water

giving metal indicator complex which is wine red in colour.

HIn2- + M2+ → MIn- + H+ (M=Ca2+, Mg2+)

(Wine red)


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

giving metal EDTA complex, which is stable and colourless.

H2Y2-

+ M2+

→ MY2-

+ 2H+

(Colourless)

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

engaged with indicator and the original blue colour of indicator is restored.

H2Y2- + MIn- → MY2- + HIn2- + H+

(Blue)

3. Outline of Procedure:

A. Standardization of EDTA solution:

a) Fill the burette with EDTA solution.

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

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

appears.

c) Titrate this solution against EDTA solution taken in a burette till wine red color

changes to blue.

d) This is the end point. Recovered the volume of EDTA consumed as A ml. Repeat 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.

B. Determination of Total Hardness:

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

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


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

changes to blue.

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

procedure to get at least three concordant readings.

Table 2: 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 sample hard water = B ml of EDTA

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

B ml of EDTA = B × 10/A mg of CaCO3

10 ml of sample hard water = 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


4. Required Results:

Parameters: Volumetric analysis.

Relationships: determine how much EDTA is consumed for 1 mL of Standard Hard

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

Graphs: NA

Error Analysis: Ask for the actual hardness and calculate the % error.

% ��

��

Results: The hardness of given sample hard water is =….......... ppm.

% error =…………%.

Scope of result

The determination of water hardness is a useful test to measure the quality of water for

households and industrial uses. Hard water can cause serious problems in industrial setting,

where hard water is monitored to avoid costly breakdown in boilers, cooling towers and

other equipments that handles water.

5. Cautions:

i. The burette, pipette and conical flask should be washed properly and then rinsed

with distilled water.

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

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

carefully.

6. Learning outcomes: to be written by the students in 50-70 words.


Experiment No. 2

1. Experiment: To determine the rate constant of hydrolysis of ethyl acetate catalyzed by

HCl.

Equipments requirements: Six conical flasks, burette, pipette, Stop watch, Water bath

etc.

Materials required: 0.1N NaOH, ethyl acetate, 0.5 N HCl, Phenolphthalein.


I) To make the concepts of chemical kinetics understandable.

II) To know how to determine the rate constant of a reaction.

III) To gain knowledge about chemical kinetics of pseudounimolecular reactions.

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

different time intervals.

V) To prove that order of reaction is an experimental concept.

VI) To know the effect of temperature on rate of reaction.

VII) To gain knowledge about acid base titration.

THEORY

This reaction is an example of psuedounimolecular reactions. Since water is present in large

excess, its concentration is practically constant throughout the reaction. The reaction is

catalysed by H+ ions of an acid (HCl). 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. Suppose the

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

reaction are V0,Vt,V∞ respectively, then initial concentration of ester is proportional to V∞-V0.

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

� � 2.303

��

!" � !#

!" � !$

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

will first order.

3. Outline of procedure:

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

acetate in a test tube, cork both of them. Keep both the solutions separately under

ambient condition.

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

c) Take 25 mL of ice cold water in a separate conical flask.

d) Prepare a reaction mixture, by adding 10 ml of ethyl acetate from the test tube to the

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

e) 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.

f) Add 2-3 drops of phenolphthalein indicator into a conical flask (containing 25 ml ice

cold water and 5ml reaction mixture).

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

reaction mixture and indicator) against 0.1 N NaOH taken in burette. 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.

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


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

after 10 min. Add indicator and titrate it against 0.1 N NaOH . This gives Vt after 10

min.

j) Repeat the above procedure after every 10 mins for taking readings upto 60 minutes.

k) Place the remaining reaction mixture on a water bath at 60-70oC for about half to one

hour. Pipette out 5 ml of mixture. Add it to the conical flask containing 25 ml ice cold

water. Add indicator and titrate it against 0.1 N NaOH. 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:

� � 2.303

��

!" � !#

!" � !$

�� %!" � !$ & � �'$

(.)#)* �� %!" � !#& is equation of straight line.

4. Required result:


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

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

Relationship to be determined: Between rate constant, time and volumes of NaOH

used.

Graph

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

the value of rate constant can

be calculated.

t (min)

Error Analysis:

Result: Rate constant of ethyl acetate at given temperature is.............

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.

5. Cautions:

i. Use the ice cold water only.

ii. Shake the reaction mixture properly at regular intervals.

iii. Perform the titrations properly.

iv. Always take alkali in burette.

v. Read lower meniscus while taking reading of the burette.



Experiment No. 3

1. Experiment: To determine the strength of given solution of ferrous ammonium sulphate

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

Equipment Required: Titration flask, 50 mL Burette, pipette, funnel etc.

Material Required: Potassium dichromate, Mohr’s salt, conc. Sulphuric Acid, Ferroin

indicator.


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).

During titration of Fe2+ ions against K2Cr2O7, ferrous ions are oxidized to ferric ions by

potassium dichromate in acidic medium. The completion of oxidation reaction is marked by

the appearance of green colour by oxidized ferroin indicator.

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

as:

Chemical Reactions

K2Cr2O7 + 4H2SO4 → K2SO4 + Cr2(SO4)3 + 4H2O + 3[O]

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

K2Cr2O7 + 6FeSO4 + 7H2SO4 → K2SO4 + Cr2(SO4)3 + 3Fe2(SO4)3 + 7H2O

Ionic Equations

Cr2O72–

+ 14H+ + 6e

– → 2Cr

3+ + 7H2O

Fe2+ → Fe3+ + e- × 6

Cr2O72– + 6Fe2+ + 14H+ → Fe3+ + 2Cr3+ + 7H2O



a) Fill the burette with potassium dichromate solution.

b) Pipette out 10 mL of Mohr’s salt solution in the conical flask.

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

d) Add 2-4 drops of the Ferroin indicator.

e) Titrate this reaction mixture against potassium dichromate solution till the colour

change from wine red to green.

f) 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 to green

Equivalent weight of Mohr’s salt = 392

Observation Table:

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 × Eq. Wt


4. Required Result:

Parameters used: Volumetric parameters, Normality equation

Relationship to be determined: determine the amount of oxidizing agent required for

Fe2+ ions solution that can be used to calculate the concentration of the latter.

Plot/Graph: NA

Error Analysis: Ask for the actual normality of Fe2+ ions solution and calculate the %

error.

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

Scope of result: Redox titration is 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.

5. 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.



Experiment No. 4

1. Experiment: To estimate the nickel content in the given sample using dimethyl

glyoxime.

Equipments required: Beakers, Suction pump, sintered glass crucible, oven, glass rod

etc.

Chemical required: Nickel ammonium sulphate, Dimethyl glyoxime, conc. ammonia,

Ethanol, conc. HCl etc.

2. Learning objectives:

I) To know about gravimetric analysis technique.

II) The students will learn how the transition metal complex formation is helpful to

determine amount of particular metal ions in given salt.

III) How to use sintered glass crucible and suction pump.

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.

Chemical Reaction:

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


Ni(DMG)2 complex (red ppt.)


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

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

b. Dissolve 0.6 g of dimethyl glyoxime in 15 ml of ethyl alcohol in a separate conical flask.

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

stirring.

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

e. 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)

f. Allow the precipitates to settle down for about 20 minutes.

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

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

h. Remove the brilliant red precipitate formed and dry in oven.

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

Observation

Colour of the compound = …………………

Weight of the ppt = ……………….

Calculations

(C4H8O2N2)2Ni ≡ Ni

288.69 58.69

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


Parameters: Colour and weight

Relationship: weight of the samples with the molecular weight of known molecules.

Graphs: NA

Error Analysis:

% error: Calculate the actual amount of Ni present in starting nickel ammonium


sulphate salt taken for the preparation of complex and calculate the % error.

Result: The Nickel content present in Ni(DMG)2 complex = ……………….

% error…………

Scope of result

The characteristic colour of the complex is used for qualitative analysis of nickel. Gravimetric

analysis is helpful in estimation of amount of Ni in given salt.

5. Cautions:

i. Apparatus should be cleaned properly. Rinse the apparatus with Conc. HCl then wash

with water properly.

ii. The acid should be handled with care and dropper or a measuring cylinder should be

used for addition of it.

iii. Ammonia solution should be added after heating the reaction mixture.



Experiment No. 5

1. Experiment: To Separate the mixture of amino acids by thin layer chromatography.

Equipments required: Glass plates, beaker, Glass rod, watch glass.

Material required: Silica gel (for TLC), Glycine, Leucine, acetic acid, n-butanol, water,

alcoholic solution of ninhydrin.


I) Students will learn the basis of chromatography technique and preparation of TLC

plates.

II) To calculate retention factor (Rf).

III) Students will learn how to identify the number of components present in a 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). Chromatography is

based on the general principle of distributing the components of a mixture of organic

compounds between two phases - a stationary phase and a mobile phase. In thin layer

chromatography the stationary phase is a solid supported on a glass plate, while the mobile

phase is a liquid (solvent). As the stationary phase is a solid, the basis of separation of

different components is adsorption.

Hence, the chromatography can be defined as the technique for the separation of a mixture of

compounds where the separation is brought about by the differential movement of the

individual compounds through a porous medium under the influence of a moving solvent.

Thin layer chromatography (TLC) is type of adsorption chromatography, which involves

separation of the components of a mixture over a thin layer of an adsorbent. A thin layer

(about 0.2 mm thick) of an adsorbent (Silica gel or alumina) is spread over a glass plate of

suitable size. The plate is known as thin layer chromatography (TLC) plate. In this technique

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

applied as a small spot near one edge (about one cm above) of the TLC plate covered with

thin layer of adsorbent. 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 level of solvent must be lower than the marked spot. As the solvent in a beaker moves up

through the adsorbent by capillary action, the various components of mixture move up along

with the solvent to different distances depending on this degree of adsorption. This results in

the separation of various components of mixture. The relative adsorption of each component

of the mixture is expressed in terms of its retention factor ie., Rf Value.

+, � Distance moved by the substance from base line @bA

Distance moved by the solvent from base line @aA

RETENTION FACTOR: The movement of any substance relative to the solvent front is

called retention factor. It is constant for a given chromatographic system and

characteristic of a substance.

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.


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).

b) 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.


c) Pour the slurry on to the clean and dry plate and prepare a uniform thin layer of silica

gel by glass rod.

d) 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.

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

by lab technician).

f) 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.

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

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

3/4 of the plate. (take due care that spot must not touch the solvent)

h) 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.

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

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

solvents and each component from the 1 cm mark.

k) Calculate the Rf values of the components in the mixture.

OBSERVATION:

Distance moved by first component Glycine (b1) = ……… cm

Distance moved by second component Leucine (b2) = ……… cm

Distance moved by solvent (a) = ……… cm

CALCULATIONS:

+, � Distance moved by the substance from base line @bA

Distance moved by the solvent from base line @aA


+,B�C BDCE� F�GH�IJI� � KL

M

+,B�C EJF�IN F�GH�IJI� � K(

M


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

a

b1

b2

Error Analysis: NA

Result:

Rf value for first component (Glycine) = …..

Rf value for second component (Leucine) = …

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.

5. Cautions:

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

ii. Always prepare fresh silica slurry.


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

iv. Spots of mixture must be as small as possible.

v. Dry TLC plates carefully.



Experiment No. 6

1. Experiment: To identify the elements (Nitrogen and Sulphur) present in given organic

compound.

Equipments required: Test tube, China dish, fusion tubes, glass rod, Bunsen Burner,

tripod stand, funnel, beaker, filter papers etc.

Material Required: Sodium metal, Ferrous sulphate (freshly prepared), sodium

hydroxide, dilute HCl/H2SO4, acetic acid, lead acetate solution, Sodium nitroprausside

solution.

2. Learning objective:

I) How to determine the presence of various elements (nitrogen and sulphur) in any organic

compound?

THEORY:

The main constituents of organic compounds are carbon and hydrogen. Besides this Nitrogen,

Sulphur and halogen may also present in organic compounds.

Nitrogen and sulphur in any organic compound are detected by 'Lassaigne's test'. The organic

compounds are covalent compounds. The fusion of organic compound with sodium metal

convert covalently bonded nitrogen and sulphur into ionic compounds (NaCN and Na2S). The

ionic compounds formed during the fusion are extracted in aqueous solution, and can be

detected by simple chemical tests. The aqueous solution obtained by extracting the fused

mass in water is called sodium extract or Lassaigne's extract.

Chemical equations:

Detection of Nitrogen

6NaCN + FeSO4 → Na4[Fe(CN)6] + Na2SO4

(Sodium ferrocyanide)

3Na4[Fe(CN)6] + 4Fe3+

→ Fe4[Fe(CN)6] 3 + 12Na+

(Ferric ferrocyanide)

(Prussian blue/green)


Detection of Sulphur

Na2S + Pb(CH3COO)2 → PbS↓ + 2CH3COONa

(Black ppt)

Na2S + Na2[Fe(CN)5NO] → Na4[Fe(CN)5NOS]

(Sodium nitroprausside) (Sodium thionitroprausside)

(Purple)

3. Outline of Procedure:

(a) Preparation of Lassaigne’s extract:

• 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.

• 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.

• 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 interfere in the detection)

• Repeat the procedure for at least three times with fusion tube (add the fusion tube to the

same china dish having distilled water)

• 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.

• Filter the content through filter paper and collect the solution obtained in a clean beaker,

the solution obtained will be Lassaigne’s extract (LE).

(b) Test for nitrogen:

• 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.

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

gently in the flame.

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

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


(c) Test for Sulphur:

• 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.

• 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.

Observation Table:

Sr. No. Test Observation Inference

1 Detection of Nitrogen: 2 ml LE +

FeSO4 + NaOH + Dil HCl

2 Detection of Sulphur: 2ml LE +

Acetic acid + Sodium acetate

3 Detection of Sulphur: 2 ml LE +

sodium nitroprausside

4. Results Required:

Parameters: Qualitative analysis of various elements present in the given organic

compound

Relationships To Be Determined: Which test is applicable to confirm the presence

of an element?

Graphs/Plots: NA

Error Analysis: NA

Result: _____________ is present in given organic compound.

Scope of Result: Experiment is helpful for the detection of nitrogen and sulphur in any

organic compound.


CAUTIONS:

i. Sodium metal should be handled very carefully.

ii. China dish and test tubes should be properly cleaned.

iii. Take distilled water in china dish.

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

v. On addition of FeSO4 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 No. 7

1. Experiment: To determine the dissociation constant of acetic acid using pH-meter.

Equipments required: pH meter, 100 mL beaker, pH electrode, 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) 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.


a) Switch on the pH meter after connecting the pH electrode to it.

b) With the help of temperature knob shown in the Fig set the temperature to the room

temperature.

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

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

e) Put the pH electrode in pH 7 solution and set the pH on screen to 7 with calibration

knob present on the pH meter.

f) Then wash the electrode and put it in pH 4 solution and set 4 on the screen.

g) Your pH meter is now ready to take readings of unknown solutions. Don’t touch any

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


h) 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.

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

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

measure the pH after addition.

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

l) 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.

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

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.

-

-

39.

40.

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)

4. 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: Ask for the original value from your instructor for % error calculations.

Result: The dissociation constant of acetic acid is …………..

5. Caution:

1. Handle the glass electrode very carefully.

Half equivalence point


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.



Experiment No. 8

1. Experiment: To find the strength of hydrochloric acid solution by titrating it against

sodium hydroxide solution conductometrically.

Equipments required: Conductivity Bridge, conductivity cell, beaker, funnel,

burette.

Materials required: 00.1 N NaOH solution and approximately HCl solution.


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.

III) 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

Conductivity Electrode

To set at the value of cell

constant

Range

To set temperature

Conductivity Electrode

Value of Cell Constant


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

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 known 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.


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

of the cell.

b) Connect the conductivity cell to the conductometer.

c) 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.

d) 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.

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

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

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

h) 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.

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

j) Set the ‘Function Switch’ to ‘Conductivity’ and read the display. This will be the exact

conductivity note it down.


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

Stir and determine the conductivity.

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

conductivity in the observation table.

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

n) 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.

4. Required result:

Parameters: Normality equation

Relationship to be determined: Effect of adding a base to an acid on conductivity

and strength of an acid

Graph: 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.

Equivalence point

A ml

Volume of NaOH added

Conductance

X

Y

Z


Error Analysis:

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

result.

Result: Strength of given HCl solution = ________ g/litre.

% error = __________%

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.

5. Cautions:

a) 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.

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



Experiment No. 9

1. Experiment: To test the validity of Beer-Lambert’s law and also determine unknown

concentration of solution using colorimeter.

Equipments required: Colorimeter, test tubes, burette, pipette, beakers etc

Materials required: Distilled Water, 0.01 M Potassium permangnate.


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

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

solution concentrations

III. Students will use algebraic representations to describe data

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

concentration of a solution

Fig: Photograph of 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.

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

To set O.D at zero

Filter (to set at λ)


Beers Lambert’s Law: It states that the decrease in intensity of incident light with thickness

of absorbing medium is directly proportional to the intensity of incident light and

concentration of absorbing medium.

(1)

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

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 = εCl

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.


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

b. Adjust the wavelength knob to the 40 wavelength region on scale

(approximately).

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

distilled water. Close the lid.


d. Adjust the reading on the digital screen to zero optical density with the knob

shown in the Fig.

e. 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.

f. 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.

g. 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)


h. 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.

i. 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.

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

j. Make table with Concentration on LHS and OD on RHS.

k. 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.

l. Plot a graph between O.D against composition. (If a straight line is obtained

Lambert - Beer’s a law is verified)

m. 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

4. 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.

Result: The concentration of unknown solution = …………..

% error = …………..


5. 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.



Appendix

Common Apparatus/Instruments 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