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Humanities & Sciences Engineering Chemistry Lab Narsimha Reddy Engineering College Page -1 Vision: To emerge as a destination for higher education by transforming learners into achievers by creating, encouraging and thus building a supportive academic environment. Mission: To impart Quality Technical Education and to undertake Research and Development with a focus on application and innovation which offers an appropriate solution to the emerging societal needs by making the students globally competitive, morally valuable and socially responsible citizens.

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Page 1: Humanities & Sciences Engineering Chemistry Lab Master.pdf · Humanities & Sciences Engineering Chemistry Lab Narsimha Reddy Engineering College

Humanities & Sciences Engineering Chemistry Lab

Narsimha Reddy Engineering College Page -1

Vision:

To emerge as a destination for higher education by transforming learners into

achievers by creating, encouraging and thus building a supportive academic

environment.

Mission:

To impart Quality Technical Education and to undertake Research and

Development with a focus on application and innovation which offers an

appropriate solution to the emerging societal needs by making the students

globally competitive, morally valuable and socially responsible citizens.

Page 2: Humanities & Sciences Engineering Chemistry Lab Master.pdf · Humanities & Sciences Engineering Chemistry Lab Narsimha Reddy Engineering College

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GENERAL PRECAUTIONS TO BE TAKEN IN THE LABORATORY

1. Never work in the laboratory unless a demonstrator or teaching assistant is present.

2. Do not throw waste such as match stems filter papers etc. into the sink.

They must be thrown into the waste jars.

3. Keep the water and gas taps closed expect when these utilities are needed.

4. Never taste any chemical unless instructed to do so and don’t allow chemicals to come in

contact with your skin.

5. While working with gases, conduct the experiment in a fume hood.

6. Keep all the doors and windows open while working in the laboratory.

7. You should know about the hazards and properties of every chemical which you are going

to use for the experiment. Many chemicals encountered in analysis are poisonous and must be

carefully handle

8. Sulphuric acid must be diluted only when it is cold .This should be done by adding it

slowly to cold water with stirring ,and not vice versa.

9. Reagent bottles must never be allowed to accumulate on the work bench. They should be

placed back in the shelves as and when use.

10. Containers in which reaction to be performed a little later should be labeled. Working

space should be cleaned immediately.

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INSTRUCTIONS FOR RECORD WRITING

1. Write on the right hand page the following order:

a) Serial number and date of performance (in the margin)

b) Name and number of the experiment as given in the list.

c) Aim of the experiment.

d) Description of the apparatus.

e) Procedure including sources of error and precautions taken to eliminate or to minimize

them

f) Inference or Result.

g) Explanation, if necessary of any divergence in the expected result.

2. Left hand page should contain the following in their proper places.

a) Neat diagram of the main apparatus.

b) Observation in tabular form.

c) Calculation in tabular form.

d) Graph sheets and other papers to be attached.

3. Students should submit a record of the previous experiments when they come for practical

work.

4. An experiment is deemed to be complete when it is satisfactorily performed and recorded.

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Objectives:

1. To analyze water for industrial purpose.

2. To characterize lubricating oils.

3. To know the synthetic methods of polymers and other compounds.

4. To analyze the materials and estimate various metals.

5. To use various instruments in analytical methods

Outcomes:

a)Graduate will able to apply chemical principles in Science& Technology.

b) Graduates will demonstrate the ability to design and conduct experiments,

interpret and analyze data, and report results.

c) Graduates will demonstrate the ability to design engineering & Technology

systems that meet desired specifications and requirements.

d) Graduates will demonstrate the ability to function in engineering and

Science laboratory teams, as well as on multidisciplinary design teams.

e) Graduate will be able to demonstrate instrumental, wet chemical methods and

analytical methods.

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LIST OF EXPERIMENTS

I. Volumetric Analysis

1. Estimation of ferrous ion by Dichrometry

2. Estimation of hardness of water by complexometric method using EDTA

3. Estimation of ferrous and ferric ion in a given mixture by Dichrometry

4. Estimation of ferrous ion by Permangnometry

5. Estimation of copper by Iodometry

6. Estimation of percentage of MnO2 in pyrolusite

7. Determination of percentage of available chlorine in bleaching powder

8. Determination of salt concentration by ion-exchange resin

II. Instrumental Methods of Analysis 9. Estimation of HCl by Conductometry

10. Estimation of ferrous ion by Potentiometry

11. Determination of ferrous ion in cement by Colorometric method

12. Determination of viscosity of an oil by Ostwald’s viscometer

13. Estimation of manganese in KMnO4 by Colorometric method

14. Estimation of HCl and acetic acid in a given mixture by Conductometry

15. Estimation of HCl by Potentiometry

III. Preparation of Polymers

16. Preparation of Bakelite

17. Preparation of urea-formaldehyde

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EXPERIMENT-1

ESTIMATION OF FERROUS ION BY DICHROMETRY

AIM: To estimate the amount of ferrous ion (Fe

+2) present in given solution by using

standard dichromate.

APPARATUS: Beakers, conical flask, burette, pipette, measuring jar.

CHEMICALS REQUIRED: K2Cr2O7, Mohr’s salt, Conc.H2SO4, conc.H3PO4

INDICATOR: Diphenyl amine

END POINT: Colorless to blue-violet

PRINCIPLE: Potassium dichromate, K2Cr2O7, is an inorganic chemical reagent, most

commonly used as an oxidizing agent. In acidic medium K2Cr2O7 oxidizes the ferrous ion

(Fe+2

) to ferric ion (Fe+3

). The completion of the oxidation reaction is marked by appearance

of blue-violet color.

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

PROCEDURE:

Part-1: Standardization of Standard Potassium Dichromate solution:

1. Pipette out 10mL of standard Mohr’s salt solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4, 2mL of Conc. H3PO4 and 2drops of Diphenyl amine

indicator. Shake the contents well.

3. Fill the burette with the given K2Cr2O7 solution.

4. Titrate the contents of conical flask with the K2Cr2O7 solution till the color of the solution

turns blue-violet.

5. Repeat the titration to get concurrent readings.

Part-2: Estimation of Ferrous Ion (Fe+2

):

1. Pipette out 10mL the given unknown ferrous solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4, 2mL of Conc. H3PO4 and 2drops of Diphenyl amine

indicator. Shake the contents well.

3. Fill the burette with the standard K2Cr2O7 solution.

4. Titrate the contents of conical flask with the K2Cr2O7 solution till the color of the solution

turns blue-violet.

5. Repeat the titration to get concurrent readings.

OBSERVATIONS AND CALCULATIONS:

Part-1: Standardization of K2Cr2O7:

S.NO

Volume of Mohr’s salt sol.

(in mL)

Burette reading Volume of

K2Cr2O7 run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 10.6 10.6

2. 10 0 10.6 10.6

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= Normality of standard Mohr’s salt solution = 0.012N

= Volume of standard Mohr’s salt solution = 10mL

= Normality of K2Cr2O7 solution =?

= Volume of K2Cr2O7 solution = 10.6mL

Part-2: Estimation of Ferrous Ion (Fe+2

):

S.NO

Volume of unknown Fe+2

sol.

(in mL)

Burette reading Volume of

K2Cr2O7 run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 21.2 26 4.8

2. 10 26 30.8 4.8

= Normality of standard K2Cr2O7 solution = 0.0113N

= Volume of standard K2Cr2O7 solution = 4.8mL

= Normality of unknown Fe+2

solution =?

= Volume of unknown Fe+2

solution = 10mL

Amount of Fe+2

present in the given solution is

RESULT: The amount of ferrous ion present in the given solution is 0.303744g/L.

VIVA QUESTIONS:

1. Define Morality. Give its formula.

2. Name the indicator used in this experiment?

3. What type of titration is involved in this experiment?

4. What are the molecular and equivalent weights of K2Cr2O7?

5. Write the chemical reactions involved in this experiment?

6. What is the equivalent weight of iron in this experiment?

7. What is the equivalent ratio of K2Cr2O7 and FeSO4?

8. What is Oxidation?

9. What is Reduction?

10. Why do you observe a green colour before the end point?

11. How do you detect the end point?

12. Define an oxidizing and a reducing agent.

13. What are oxidizing & reducing agents in this experiment?

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EXPERIMENT-2

ESTIMATION OF HARDNESS OF WATER BY EDTA METHOD

(COMPLEXOMETRIC TITRATION)

AIM: To estimate the amount of total, permanent and temporary hardness present in the

given water sample.

APPARATUS: Conical flask, burette, funnel, pipette and measuring cylinder.

CHEMICALS REQUIRED: EDTA, Magnesium sulphate (MgSO4).

INDICATOR: Eriochrome Black-T

BUFFER: Ammonium chloride-ammonium hydroxide (NH4Cl - NH4OH)

END-POINT: Wine-red to Blue color

PRINCIPLE:

Calcium (Ca+2

) or Magnesium ions (Mg+2

) form a wine red color unstable complex with EBT

indicator.

Ca+2

/ Mg+2

+ EBT Ca-EBT / Mg-EBT (wine red color)

(Metal-EBT complex)

On titrating the above complex against EDTA, a stable metal EDTA complex is formed by

releasing the free indicator, due to the liberation of indicator, blue color will appear at end

point.

EDTA forms stable complex with calcium and magnesium ions present in hard water at pH =

10.

EDTA + M-EBT ----------> M-EDTA + EBT (Blue color)

(Metal-

EDTA complex)

PROCEDURE:

Part-A: Standardization of EDTA solution:

1. Pipette out 20mL of standard magnesium sulphate into a clean conical flask.

2. Add 2mL of buffer solution and 2drops of EBT indicator.

3. Fill the burette with EDTA solution and titrate against the standard magnesium sulphate

solution till color changes from wine-red to blue.

4. Repeat the titration to get concurrent readings.

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Part-B: Estimation of Total Hardness in given Water Sample:

1. Pipette out 20 ml of water sample into a clean conical flask. Add 3 ml of buffer 3drops of

EBT indicator.

2. Fill the burette with EDTA solution.

3. Now titrate the water sample till the color of the solution turns wine-red to blue color.

4. Repeat the titration to get concurrent readings.

Part-C: Estimation of Permanent Hardness in given Water Sample:

1. Measure 100mL of sample water into a beaker.

2. Boil this water until the volume is reduced to 50mL.

3. Cool and filter the water into a 100mL standard flask and make up the solution using

distilled water.

4. Pipette out 20mL of this water into a conical flask. Add 2mL of buffer and 2drops of

indicator.

5. Titarate this solution against EDTA taken in burette until the color changes from wine-red

to blue.

OBSERVATIONS AND CALCULATIONS:

Part-A: Standardization of EDTA solution:

S.NO

Volume of MgSO4 solution

(in mL)

Burette reading Volume of

EDTA run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 5.5 5.5

2. 10 5.5 11 5.5

M1= Molarity of MgSO4 solution = 0.01M

V1= Volume of MgSO4 solution = 10mL

M2= Molarity of EDTA solution =?

V2= Volume of EDTA solution used =5.5mL

Part-B: Estimation of Total Hardness of Given Water Sample:

S.NO

Volume of Water sample

(in mL)

Burette reading Volume of

EDTA run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 7 7

2. 10 0 14 7

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M2= Molarity of EDTA solution = 0.018M

V2= Volume of EDTA solution used = 7mL

M3= Molarity of given Water sample =?

V3= Volume of given Water sample = 10mL

The amount of Total hardness present in the given water sample =

M3 X 100 X 1000 = 0.0126*100*1000=1260ppm

Part-C: Estimation of permanent hardness of given water sample:

S.NO

Volume of Boiled Water

sample

(in mL)

Burette reading Volume of

EDTA run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 4 4

2. 10 4 8 4

M2= Molarity of EDTA solution = 0.018M

V2= Volume of EDTA solution used = 4mL

M4= Molarity of Boiled water sample = ?

V4= Volume of Boiled Water sample = 10mL

The amount of Permanent hardness present in the given water sample =

M4 X 100 X 1000 = 0.0072*100*1000=720ppm

The amount of temporary hardness water present in given water sample =

Total hardness – Permanent Hardness = 1260-720=540ppm

RESULT:

1. The amount of total hardness present in the given water sample is 1260 ppm.

2. The amount of permanent hardness present in the given water sample is 720ppm.

3. The amount of temporary hardness present in the given water sample is 540ppm.

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VIVA QUESTIONS:

1. Mention the salts causing temporary and permanent hardness.

2. How is hardness of water tested?

3. How is hardness expressed? What are its units?

4. What is Normality? How can we calculate the weight of solute using

normality of the solution?

5. What is the indicator used in this experiment?

6. What is a standard solution?

7. What is a buffer solution and which buffer is used in this titration?

8. What is the colour of the solution at the end point, in this titration?

9. What is the colour of metal indicator complex & M-EDTA complex?

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EXPERIMENT-3

ESTIMATION OF FERROUS AND FERRIC IONS IN A GIVEN MIXTURE BY

DICHROMETRY

AIM: To estimate the amount of ferrous (Fe+2

) and ferric (Fe+3

) ions present in given mixture

using standard dichromate.

APPARATUS: Beakers, conical flask, burette, pipette, measuring jar.

CHEMICALS: K2Cr2O7, FeSO4, FeCl3, stannous chloride (SnCl2), mercuric

chloride(HgCl2), Conc.H2SO4, Conc.H3PO4, HCl

INDICATOR: Diphenyl amine

END POINT: blue-violet

PRINCIPLE:

The Fe+3

ions present in the sample are estimated by reducing them to Fe+2

ions using SnCl2.

Then the sample is estimated for the amount of Fe+2

present using K2Cr2O7. The difference of

the two values i.e., before and after reduction gives the amount of ferric ions.

2FeCl3 + SnCl2 2FeCl2 + SnCl4

SnCl2 + 2HgCl2 Hg2Cl2 + SnCl4

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

PROCEDURE:

Part-1: Standardization of Standard Potassium Dichromate solution:

1. Pipette out 10mL of standard Mohr’s salt solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4, 2mL of Conc. H3PO4 and 2drops of Diphenyl amine

indicator. Shake the contents well.

3. Fill the burette with the given K2Cr2O7 solution.

4. Titrate the contents of conical flask with the K2Cr2O7 solution till the color of the solution

turns blue-violet.

5. Repeat the titration to get concurrent readings.

Part-2: Estimation of Total Ions (Fe+3

and Fe+2

):

1. Pipette out 10mL of the given solution-A into a clean conical flask.

2. To it add 20mL of conc. HCl solution and boil till the color changes to clear yellow.

3. To the hot solution add stannous chloride drop by drop till the yellow color disappears.

4. Cool the solution and add 10ml of HgCl2. A silky white precipitate is obtained.

5. Then to it add 3mL of conc. H2SO4, 3mL of Conc. H3PO4 and 2drops of diphenyl amine

indicator. Shake the contents well.

6. Now, titrate this solution against standard K2Cr2O7 till the blue color changes to blue

violet.

7. Repeat the titration to get concurrent readings.

Part-3: Estimation of Ferrous Ion (Fe+2

):

1. Pipette out 10mL the given unknown ferrous solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4, 2mL of Conc. H3PO4 and 2drops of diphenyl amine

indicator. Shake the contents well.

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3. Fill the burette with the standard K2Cr2O7 solution.

4. Titrate the contents of conical flask with the K2Cr2O7 solution till the color of the solution

turns blue-violet.

5. Repeat the titration to get concurrent readings.

OBSERVATIONS AND CALCULATION:

Part-1: Standardization of K2Cr2O7 :

S.NO

Volume of Mohr’s salt sol.

(in mL)

Burette reading Volume of

K2Cr2O7 run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 10.2 10.2

2. 10 10.2 20.4 10.2

= Normality of standard Mohr’s salt solution = 0.012N

= Volume of standard Mohr’s salt solution = 10mL

= Normality of K2Cr2O7 solution =?

= Volume of K2Cr2O7 solution = 10.2mL

Part-2: Estimation of Total Ions (Fe+3

and Fe+2

):

S.NO

Volume of total Fe+2

and Fe+3

ions solution

(in mL)

Burette reading Volume of

K2Cr2O7 run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 6.4 6.4

2. 10 6.4 12.8 6.4

= Normality of standard K2Cr2O7 solution = 0.0117N

= Volume of standard K2Cr2O7 solution = 6.4mL

= Normality of total Fe+2

and Fe+3

ions solution =?

= Volume of total Fe+2

and Fe+3

ions solution = 10mL

Amount of total Fe+2

and Fe+3

ions present in the given solution is

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Part-3: Estimation of Ferrous Ion (Fe+2

):

S.NO

Volume unknown Fe+2

solution (in mL)

Burette reading Volume of

K2Cr2O7 run

down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 6.4 6.4

2. 10 6.4 12.8 6.4

= Normality of standard K2Cr2O7 solution = 0.0113N

= Volume of standard K2Cr2O7 solution =5.2mL

= Normality of unknown Fe+2

solution =?

= Volume of unknown Fe+2

solution = 10mL

Amount of Fe+2

present in the given solution is

Amount of Fe+3

present in the given solution is

= amount of total Fe+2

and Fe+3

ions - amount of Fe+2

=0.420141-0.3278=0.092341g/L

RESULT:

1. The amount of ferric ion present in the given solution is 0.3278g/L.

2. The amount of ferrous ion present in the given solution is 0.092341g/L.

VIVA QUESTIONS:

1. Define Morality. Give its formula.

2. Name the indicator used in this experiment?

3. What type of titration is involved in this experiment?

4. What are the molecular and equivalent weights of K2Cr2O7?

5. Write the chemical reactions involved in this experiment?

6. What is the equivalent weight of iron in this experiment?

7. What is the equivalent ratio of K2Cr2O7 and FeSO4?

8. What is Oxidation?

9. What is Reduction?

10. Why do you observe a green colour before the end point?

11. How do you detect the end point?

12. Define an oxidizing and a reducing agent.

13. What are oxidizing & reducing agents in this experiment?

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EXPERIMENT-4

ESTIMATION OF FERROUS ION BY PERMANGANOMETRY

AIM: To estimate the amount of ferrous ion (Fe+2

) present in given solution by using

standard permanganate solution.

APPARATUS: Beakers, conical flask, burette, pipette, measuring jar.

CHEMICALS: Potassium permanganate (KMnO4), Mohr’s salt solution, Conc.H2SO4,

END POINT: Colorless to pale pink

PRINCIPLE:

Potassium permanganate (KMnO4) is strong oxidizing agent. In acidic medium KMnO4

oxidizes ferrous sulphate present in Mohr’s salt to ferric sulphate. The completion of the

oxidation reaction is marked by appearance of pale pink color.

2 KMnO4+ 10FeSO4 + 8H2SO4 K2SO4 + 2MnSO4 + 5Fe2(SO4)3 + 8H2O

PROCEDURE:

Part-1: Standardization of Standard KMnO4 Solution:

1. Pipette out 10mL of standard Mohr’s salt solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4 and shake well.

3. Fill the burette with the given KMnO4 solution.

4. Titrate the contents of conical flask with the KMnO4 solution till the color of the solution

turns pale pink.

5. Repeat the titration to get concurrent readings.

Part-2: Estimation of Ferrous Ion (Fe+2

):

1. Pipette out 10mL the given unknown ferrous solution into a clean conical flask.

2. To it add 2mL of conc. H2SO4 and shake the contents well.

3. Fill the burette with the standard KMnO4 solution.

4. Titrate the contents of conical flask with the KMnO4 solution till the color of the solution

turns pale pink.

5. Repeat the titration to get concurrent readings.

OBSERVATIONS AND CALCULATIONS:

Part-1: Standardization of KMnO4:

S.NO

Volume of Mohr’s salt sol.

(in mL)

Burette reading Volume of KMnO4

run down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 12.5 12.5

2. 10 12.5 25 12.5

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= Normality of standard Mohr’s salt solution = 0.012N

= Volume of standard Mohr’s salt solution = 10mL

= Normality of KMnO4 solution =?

= Volume of KMnO4 solution = 12.5mL

Part-2: Estimation of Ferrous Ion (Fe+2

):

S.NO

Volume unknown Fe+2

solution

(in mL)

Burette reading Volume of KMnO4

run down

(in mL)

Initial

(in mL)

Final

(in mL)

1. 10 0 6 6

2. 10 6 12 6

= Normality of standard KMnO4 solution = 0.0096N

= Volume of standard KMnO4 solution = 6mL

= Normality of unknown Fe+2

solution =?

= Volume of unknown Fe+2

solution = 10mL

Amount of Fe+2

present in the given solution is

RESULT: The amount of ferrous ion present in the given solution is 0.03271g/L

VIVA QUESTIONS

1. Define Molarity. Give its formula.

2. Name the indicator used in this experiment?

3. What type of titration is involved in this experiment?

4. What are the molecular and equivalent weights of KMnO4?

5. Write the chemical reactions involved in this experiment?

6. What is the equivalent weight of iron in this experiment?

7. What is the equivalent ratio of KMnO47 and FeSO4?

8. What is Oxidation?

9. What is Reduction?

10. Why do you observe a green colour before the end point?

11. How do you detect the end point?

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EXPERIMENT 5

ESTIMATION OF COPPER BY IODOMETRY

AIM: To estimate the amount of Copper present in the given solution using a standard

solution of Potassium Dichromate and Hypo.

APPARATUS: Funnel, burette, Iodometric flask, conical flask, pipette

CHEMICALS REQUIRED: K2Cr2O7, Hypo (Na2S2O3), KI, Acetic acid, Na2CO3,

ammonium thiocyanate, CuSO4, conc. ammonia, HCl, Starch & distilled water.

PRINCIPLE:

Iodometric titration of copper is based on the oxidation of iodides to iodine by Cu (II) which

gets reduced to Cu (I).

Any cupric salt in neutral medium when treated with Potassium Iodide forms a white

precipitate of cuprous iodide and iodine is set free quantitatively. The liberated Iodine is

treated against a reducing agent Hypo using starch as the indicator. Following reaction take

place:

2CuSO4. 5H2O + 2KI 2CuI2 + 2K2 SO4 + 5H2O

2CuI2 Cu2I2 + I2

I2 + 2Na2S2O3 Na2S4O6 + 2NaI

The titration fails in presence of any acid so, before commencing the titration the acid present

in copper solution is neutralized with ammonia followed by acetic acid.

PROCEDURE: Step –1: Standardization of Hypo:

1. Rinse and fill the burette with hypo solution without any air bubbles.

2. Pipette out 10 ml of 10%KI solution in a clean conical flask and add 0.1g of Na2CO3

followed by 3ml of concentrate HCl gently rotate the flask for mixing the liquids.

3. Pipette out 10mL K2Cr2O7 to the above solution, shake it well, stopper it, and keep it in

dark place for 5 minutes.

4. Titrate the liberate iodine by running down hypo from the burette with constant stirring.

5. When the solution attains a pale yellow color add 2 ml of freshly prepared starch solution.

The color changes to blue.

6. Continue the titration drop-wise till the color changes from blue to colorless indicating the

end point

7. Repeat the titration for concurrent values.

Step-2: Estimation of copper:

1. Pipette out 10mL of the given copper solution into a clean conical flask

2. Add few drops of conc. ammonia till bluish white precipitate is obtained.

3. Dissolve the precipitate in acetic acid. Now add 10 ml of 10% KI, when iodine is liberated

giving a brown color.

4. Titrate this solution against standard hypo solution till light yellow color is obtained.

5. Now add 2 ml of starch solution and 1spatula of ammonium thiocynate continue the

titration till blue color changes to creamy white, which is the end point.

6. Repeat the titration for concurrent values and calculate the amount of copper.

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OBSERVATIONS AND CALCULATION:

Step –1: Standardization of Hypo:

S.No

Volume of the standard K2Cr2O7

solution

(V1 in mL )

Burette Reading

(in mL )

Volume of hypo

rundown(V2 in

mL)

Initial Final

1 10 0 3.4 3.4

2 10 3.4 6.8 3.4

N1 V1= N2 V2

N1 = Normality of K2Cr2O7 solution= 0.01N

V1 = Volume of K2Cr2O7 solution = 10ml

N2 = Normality of Hypo =?

V2 = Volume of Hypo = 3.4

N2 = N1 V1 / V2 = 0.01*10/3.4=0.0294N

Step-2: Estimation of copper:

S.No Volume of the Copper solution (in mL)

Burette Reading

(in mL)

Volume of hypo

Run down (in mL)

Initial Final

1 10 0 3.4 3.4

2 10 3.4 6.8 3.4

N3 V3= N2 V2

N3 = Normality of the Copper solution =?

V3 = Volume of the Copper solution = 10ml

N2 = Normality of Hypo =0.0294N

V2 = Volume of Hypo =3.4 ml

N3 = N2 V2 / V3 =0.294*3.4/10=0.00996 N

Amount of Copper present in the whole of the given solution

RESULT: Amount of Copper present in the whole of the given solution = g/L

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VIVA QUESTIONS:

1. Why KI is added to K2Cr2O7 or Cu2+ solution in the experiment

2. Why blue colour appears when starch solution is added to the contents of conical

flask during the experiment?

3. Why is conical flask covered with watch glass? Why the flask is kept in the dark

4. Place?

5. What is GEW & GMW of K2Cr2O7?

6. What is GEW of Cu2+?

7. What is the Indicator used in the experiment?

8. What is the role of starch in the experiment?

9. Why Na2CO3 & HCl are added to K2Cr2O7 solution?

10. Why NH4CNS is added to Cu2+

?

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EXPERIMENT-6

ESTIMATION OF PERCENTAGE PURITY OF MnO2 IN PYROLUSITE

AIM: To estimate the amount of MnO2 present in the given sample of pyrolusite

APPARATUS: pipette, burette, conical flask, beaker, funnel

CHEMICALS REQUIRED: Pyrolusite, oxalic acid, dilute H2SO4, KMnO4

Principle: MnO2 occurs in the nature in the form of pyrolusite. Pyrolusite is not only used as a

source of mangane4se but also as an oxidizing agent in many industrial processes. For such

purposes, ore is graded on the basis of its available oxygen content rather than on its

percentage of manganese.

The percentage of MnO2 is usually determined by treatment with an excess of acidified

solution of a reducing agent, such as sodium oxalate. The excess of reducing agent is

determined by titration with standard KMnO4 where KMnO4 acts as a self indicator with pale

pink color as the end point.

MnO2 + H2C2O4 + H2SO4 MnSO4 + 2CO2 + 2H2O

PROCEDURE:

1. Weigh out accurately about 0.1g of finely powdered dry pyrolusite into a conical flask

and add 25mL of standard N/10 oxalic acid.

2. Add 25mL of 10% H2SO4 and place a short funnel over the conical flask.

3. Boil the contents of the flask gently until no black particles are visible in the flask.

4. Allow it to cool about 60-70C and titrate the excess oxalate with standard N/10 KMnO4.

5. Repeat the above process with 2 other samples of similar weight.

6. Calculate the amount of oxalic acid consumed in the reaction and calculate the percentage

of MnO2 in the pyrolusite ore.

OBSERVATION AND CALCULATIONS:

S.No

Volume of the oxalic acid

( in mL )

Burette Reading

(in mL )

Volume of

KMnO4

consumed (in

mL) Initial Final

1 25 0 13 13

2 25 13 26 13

Weight of ore taken = 0.1gm

Volume of N/10 Oxalic acid used =25mL

Normality of KMnO4 = 0.1N

Volume of KMnO4 consumed = 13mL

1mL of KMnO4 = 0.004346g of MnO2

13mL of KMnO4= 13*0.004346=0.056498g of MnO2

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Percentage of MnO2 in 100g of ore

RESULT: Percentage purity of the given Pyrolusite sample is 56.49%

VIVA QUESTIONS:

1. What is the difference between mineral and an ore?

2. What is an alloy? Give examples.

3. What is Pyrolusite?

4. What happens when MnO2 is heated with excess of Na2C2O4?

5. Write all the chemical reactions involved in the experiment.

6. What is the titrant used for the estimation of unreacted (remaining) sodium oxalate

present in the solution?

7. How is the end point in this titration detected?

8. If the initial burette reading of the titration is 16ml, find out the volume of KMnO4

for used up Na2C2O4

9. What is the formula used for calculating the purity of the pyrolusite sample?

10. To what extent do we heat the solution of Na2C2O4 & MnO2?

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EXPERIMENT-7

DETERMINATION OF PERCENTAGE OF AVAILABLE CHLORINE IN

BLEACHING POWDER

AIM: To estimate the amount of available chlorine in bleaching powder by iIodometry.

APPARATUS: Funnel, burette, Iodometric flask, conical flask, pipette

CHEMICALS REQUIRED: K2Cr2O7, Hypo (Na2S2O3), KI, glacial acetic acid, Na2CO3,

HCl, Starch & distilled water.

Principle:

Bleaching powder [CaOCl2] is used as a bleaching agent and also as a disinfectant. Bleaching

powder supplies chlorine with dilute acids.

CaOCl2+ 4HCl CaCl2 + 2Cl2+ 2H2O

The available chlorine is defined as the percentage of chlorine made available by bleaching

powder when treated with dilute acids. The available chlorine present in the bleaching sample

is determined iodometrically by treating with excess of KI in acidic medium.

OCl-

+ 2H+ + 2I

- I2 + Cl

- + 5H2O

The liberated iodine is treated with hypo using freshly prepared starch solution as indicator.

I2 + 2Na2S2O3 Na2S4O6 + 2NaI

PROCEDURE: Step –I: Standardization of Hypo:

1. Rinse and fill the burette with hypo solution without any air bubbles.

2. Pipette out 10 ml of 10%KI solution in a clean conical flask and add 0.1g of Na2CO3

followed by 3ml of concentrate HCl gently rotate the flask for mixing the liquids.

3. Pipette out 10mL K2Cr2O7 to the above solution, shake it well, stopper it, and keep it in

dark place for 5 minutes.

4. Titrate the liberate iodine by running down hypo from the burette with constant stirring.

5. When the solution attains a pale yellow color add 2 ml of freshly prepared starch solution.

The color changes to blue.

6. Continue the titration drop-wise till the color changes from blue to colorless indicating

the end point

7. Repeat the titration for concurrent values.

Step-II Estimation of available chlorine in bleaching powder:

1. Fill the burette with standard hypo solution.

2. Pipette out 10mL of 10%KI solution in a clean conical flask

3. Add 10mL of the given bleaching powder solution and 5mL glacial acetic acid to the

conical flask.

4. Cover the conical flask and keep it in dark for 5minutes, a dark brown colored solution is

obtained.

5. Now titrate this solution against hypo taken in the burette till a pale yellow color solution

is obtained.

6. Now add 2ml starch and titrate against hypo till the solution turns colorless indicating the

end point.

7. Repeat the titrations for concurrent values.

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OBSERVATIONS AND CALCULATIONS:

Step –I: Standardization of Hypo:

S.No

Volume of the standard K2Cr2O7

solution

(in mL )

Burette Reading

(in mL )

Volume of hypo

rundown(in mL)

Initial Final

1 10 0 3.4 3.4

2 10 3.4 6.8 3.4

N1 V1= N2 V2

N1 = Normality of K2Cr2O7 solution= 0.01N

V1= Volume of K2Cr2O7 solution = 10ml

N2 = Normality of Hypo =?

V2= Volume of Hypo = 3.4ml

N2 = N1 V1/ V2= 0.01*10/3.4=0.0294N

Step-II Estimation of available chlorine:

S.No

Volume of the bleaching powder solution

(in mL)

Burette Reading

(in mL)

Volume of hypo

Run down (in mL)

Initial Final

1 10 0 4.8 4.8

2 10 4.8 9.6 4.8

N3 V3= N2 V2

N3= Normality of the bleaching powder =?

V3= Volume of the bleaching powder solution = 10ml

N2 = Normality of Hypo =0.0294N

V2= Volume of Hypo = 4.8ml

N3 = N2 V2/ V3= 0.0294*4.8/10=0.0141N

Amount of available chlorine present in bleaching powder is

RESULT:

Amount of available chlorine present in the given bleaching powder solution is g/L

VIVA QUESTIONS:

1. What is the formula of bleaching powder?

2. What is the need to addition of bleaching powder to water?

3. What is disinfection?

4. What are another chemicals added to water to kill the pathogenic bacteria.

5. What is reaction take place by the addition of bleaching powder to water?

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EXPERIMENT-8

DETERMINATION OF SALT CONCENTRATION BY ION-EXCHANGE RESIN

AIM: To determine the concentration of Na2SO4 by ion-exchange resin.

APPARATUS: Funnel, conical flask, chromatography column

CHEMICALS REQUIRED: resin, HNO3, sodium nitrate, NaOH, burette

PRINCIPLE:

Ionic Exchange is the process of exchange of mobile ions into other ions bearing the same

sign, which occur on special resins. Cations are exchanged in cation resins, anions in anion

resins. The combination of these resins, cation and anion, allows to demineralize water; this

is possible if all cations are exchanged by H+

ions, while all anions by the OH-. The result is

chemically pure water.

Ionic exchangers, known also as ions exchanging resins, are water insoluble and are able to

exchange positively or negatively charged ions in solution. They are divided into:

1. Cation Resins (exchange cations, their character is similar to salts are acidic). They can be

neutral or slightly acidic, when containing weakly dissociated functional groups (-OH, -

COOH, -SH, -CH2SH). Or are more acidic, able to exchange all the cations, with strongly

dissociated groups (-SO3H, -CH2SO3H).

2. Anion Resins (exchange anions, their character is similar to salts or is alkaline)

The process of ionic exchange can be schematically presented as follows:

RH+ + C

+ RC + H

+ [C

+ is cation]

ROH- + A

- RA + OH

- [A

- is anion]

PROCEDURE: 1. Weigh and transfer 1g of dried cation exchange resin into a column, through a funnel.

2. Add sufficient distilled water to cover the resin. Tap the column gently to remove any air

bubbles.

3. Run down 100ml of given NaNO3 solution into the column carefully at 2ml per minute.

4. Collect the effluent into a clean and dry conical flask.

5. When all the solution has passed through the column, titrate the Na+

effluent against NaOH

solution using phenolphthalein as indicator. Pink color is the end point.

6. Repeat the same process using anion exchange resin.

OBSERVATIONS AND CALCULATIONS:

1. Titration –I: Estimation of cation present in the resin:

S.No

Volume of Na+

effluent

( in mL )

Burette Reading

(in mL )

Volume of standard

NaOH rundown

(in mL)

Initial Final

1

2

3

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N1 V1= N2 V2

N1 = Normality of Na+

effluent =?

V1 = Volume of Na+

effluent =______mL

N2 = Normality of NaOH solution= 0.1N

V2 = Volume of NaOH solution = ______mL

N2 = N1 V1 / V2 = __________N

Amount of Na+

effluent present in given solution

Step-II Estimation of copper:

S.No

Volume of NO3-

effluent

( in mL )

Burette

Reading

(in mL )

Volume of standard NaOH

rundown

(in mL)

Initial Final

1

2

3

N3 V3= N2 V2

N3 = Normality of NO3-

effluent =?

V3 = Volume of NO3-

effluent = _____ml

N2 = Normality of NaOH solution= 0.1N

V2 = Volume of NaOH solution =

N3 = N2V2 / V3 = __________N

Amount of NO3-

effluent present in given solution

RESULT:

1. Amount of Na+

effluent present in the given solution = ______ g/L

2. Amount of NO3-

effluent present in the given solution = ______ g/L

VIVA QUESTIONS:

1. What is ion exchange resin?

2. What is cation and anion resin?

3. What are the examples for cation exchange resins?

4. What are the examples for anion exchange resins?

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EXPERIMENT -9

CONDUCTOMETRIC TITRATION OF STRONG ACID VS STRONG BASE

AIM: To determine the neutralization point of the titration of HCl against NaOH

Conductometrically.

APPARATUS: Conductometer, conductivity cell, beaker, pipette, burette

CHEMICALS REQUIRED: HCl, 1N NaOH.

PRINCIPLE:

Conductometric titration is the volumetric analysis based upon the measurement of the

conductance during the course of titration. If we consider the titration of strong acid like HCl

VS strong base like NaOH.

NaOH + HCl NaCl + H2O

The initial conductivity of HCl (strong acid) is high because it completely dissociates in to H+

ions. When NaOH is added as titrant, the OH- ions and H

+ ions react to produce water and the

number of H+

ions decreases thereby conductivity gradually decreases after every addition.

After the neutralization point, when all the H+ has reacted, the addition of NaOH causes

increases in the number of OH- ions and hence the conductivity starts to increase.

GRAPH: A plot of conductivity VS volume of NaOH added will consists of two straight line

branches intersecting at the neutralization point like V shape.

MODEL GRAPH:

PROCEDURE:

1. Take a clean beaker and add 40mL of 0.1 M HCl.

2. Fill the burette with 1M NaOH.

3. Take a clean conductivity cell washed with distilled water and place the cell in the beaker.

4. Now add 1mL of NaOH from burette drop wise, stir the solution gently and note down the

change in conductance.

5. Repeat the same procedure and note down the measured conductance.

6. Plot the graph between conductivity against volume of base added, the intersection of two

straight lines gives the end point.

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OBSERVATIONS:

S.No.

Volume of NaOH

(in mL)

Conductance

(in milliSiemens)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

24

21

19

18

15

13

11

9

7

7

9

11

12

13

14

15

16

18

19

20

RESULT:

The volume of 1M of NaOH required to neutralize 0.1M of HCl is 4mL

VIVA QUESTIONS

1. What is conductance? What are its units?

2. Define specific conductance and equivalent conductance? What are its units?

3. How do specific and equivalent conductances vary with dilution?

4. What is cell constant and what are its units?

5. Define Kholraush Law?

6. Is Conductivity cell an electrolytic or an electrochemical cell?

7. How does the conductance vary in this titration ?

8. How do you detect the neutralization point in Conductometric titration?

9. What are the advantages of Conductometric titration over volumetric titrations?

10. Draw a rough graph for this experiment?

11. What is the effect of temperature on the conductance of the electrolyte?

12. Write the relationship between conductance and specific conductance of an

electrolyte?

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EXPERIMENT -10

ESTIMATION OF FERROUS ION BY POTENTIOMETRY

AIM: To estimate the amount of ferrous present in the given sample by Potentiometry

APPARATUS: Potentiometer, platinum electrode, calomel electrode beaker, burette, pipette,

glass rod

CHEMICALS REQUIRED: K2Cr2O7, FeSO4, Conc.H2SO4

PRINCIPLE:

The titration of ferrous ions with chromium (VI) is an example of oxidation-reduction

titration.

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

The ferrous ion concentration in the given solution can be obtained by titration with standard

K2Cr2O7 by Potentiometry. The combination of the saturated calomel electrode as the

reference electrode and the platinum electrode as the indicator electrode is used. Using

potential of the indicator electrode obtained concentration of the solution can be calculated.

The reference electrode has constant potential.

PROCEDURE:

1. Calibrate the potentiometer before starting the experiment.

2. Take 25mL of FeSO4 in a beaker, to it add 1mL of conc.H2SO4 and immerse the platinum

electrode and calomel electrode into the solution.

3. Fill the burette with 0.1N K2Cr2O7 solution.

4. Carry out the titrations by adding 2mL of K2Cr2O7 and measure the EMF at every stage.

5. Plot a graph by taking volume of K2Cr2O7 on x-axis and EMF on y-axis. The end point of

the reaction is halfway between the jump in EMF.

MODEL GRAPH:

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OBSERVATION AND CALCULATIONS:

S. No. Vol. of K2Cr2O7 added

(V in mL)

EMF

(E in mV)

1

2

3

4

5

6

7

8

9

10

11

12

13

0

2

4

6

8

10

12

14

16

18

20

22

24

320

395

407

510

531

602

789

810

815

824

830

830

833

N1 = Normality of FeSO4 =?

V1 = Volume of FeSO4 taken = 25mL

N2 = Normality of K2Cr2O7 = 0.1N

V2 = Volume of K2Cr2O7 from the graph = 10mL

Amount of Fe+2

present in the given solution is

/L

RESULT:

The amount of ferrous present in the given sample is g/L.

VIVA QUESTIONS:

1. Define electrochemical cell?

2. What is SCE? Give its EMF value?

3. What is the role of quinhydrone?

4. Write the cell notation of the cell used in the experiment.

5. How do you get EMF of a cell consisting of two electrodes?

6. Give the relation between PH and EMF of quinhydrone electrode?

7. What are the advantages of Potentiometric titrations over Conductometric titrations?

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EXPERIMENT -11

ESTIMATION OF FERROUS ION IN CEMENT BY COLORIMETRIC METHOD

AIM: To estimate the amount of ferrous present in the given sample of cement by

colorimetery.

APPARATUS: Colorimeter, cuvettes, standard flask and graduated pipette.

CHEMICALS REQUIRED: cement sample, Mohr’s salt, ammonium thiocyanate, conc.

H2SO4, HCl, KMnO4, HNO3

PRINCIPLE:

The analysis which involves the measurement of absorption of light radiation in

visible region is known as colorimetry. Knowing the intensity of light absorbed and by

making suitable calibration, it is possible to find out concentration of any given solute.

Colorimeter works on the Beer-Lamberts law.

Beer-Lambert’s Law: According to Beer-lamberts law the absorbance of light by a colored

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

A= εCl

Where, A = Absorbance (or) Optical Density

C = Concentration of the solution (in M)

l = thickness of the medium

ε = Molar absorption coefficient

Molar absorption coefficient (ε) is the absorbance of a solution having unit concentration

placed in a cell of unit thickness.

Ammonium thiocyanate yields a blood red color with ferric ion and the color produced is

stable in nitric acid medium. Its optical density is measured in a photo colorimeter and the

concentration of ferric iron is found from the standard calibration curve which is obtained by

plotting a graph between optical density and concentration. A straight line curve is obtained

passing through the origin.

PROCEDURE:

Part-1: To Obtain Standard calibration curve:

1. Measure accurately 0.722g of Mohr’s salt and transfers it to a 100ml standard flask.

2. To it add 5mL conc. H2SO4 of and make up to mark using distilled water.

3. Titrate this solution against standard KMnO4 taken in a burette until a pale pink colour is

obtained.

4. Now transfer 1mL to 8mL of this pale pink solution into 100mL standard flask and label

them as A,B,C,D,E,F,G,H.

5. To each of the standard flask add 1ml of HNO3 and 5mL of 40% ammonium thiocyanate

solution, a blood red color is obtained. Make up the solution to mark using distilled water.

6. Calibrate the colorimeter using blank and measure the optical densities of all the above

solutions.

7. Plot a graph by taking amount of ferrous on x-axis and optical density on y-axis.

8. The curve obtained is called standard calibration curve.

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Part-2: Estimation of Ferrous in Cement:

1. Measure accurately 0.1g of cement sample into a beaker and add 5ml of distilled water to

moisten the cement.

2. Add conc. HCl drop wise along the walls of the beaker and close it immediately with a

watch glass for some time.

3. Remove the watch glass and heat it gently. Then add 20mL of distilled water for complete

dissolution.

4. Transfer the entire solution into a 100mL standard flask. Wash the beaker with distilled

water and transfer it to standard flask and make upto the mark.

5. Pipette out 10ml of the made up solution into a 100ml standard flask and add 1ml of HNO3

and 5mL of 40% ammonium thiocyanate solution to obtain a blood red color solution.

6. Make up the solution to the mark and measure its optical density using colorimeter.

OBSERVATION AND CALCULATIONS:

Part-1: To Obtain Standard calibration curve:

0.722g of Mohr’s salt dissolved in 100ml of distilled water = 0.0184N

Amount of ferrous present in given sample= 0.0184 X 55.8 = 0.1mg/ml

Thus, 1mL of solution (pale pink) contains 0.1mg of ferrous ions

Similarly, 2ml of same solution contains 0.2mg of ferrous ions and so on.

S.No. Concentration of Fe+2

(in mg/mL)

Absorbance/O.D.

(in nm)

1 0.10 0.30

2 0.20 0.35

3 0.30 0.90

4 0.40 0.95

5 0.50 1.0

6 0.60 1.05

7 0.70 1.10

8 0.80 1.15

9 unknown 0.90

Part-2: Determination of Fe+2

in cement:

Percentage of Fe+2

in cement

RESULT: The percentage of ferrous present in the given sample of cement is 3%

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MODEL GRAPH:

O.D.

(in nm)

Concentration (in M)

VIVA QUESTIONS:

1. Why H2SO4 is not used in this titration?

2. Why is the red colour is obtained by adding KSCN reagent to ferric ion solution?

3. Why dil.HCl or dil.HNO3 should be present?

4. Why excess of thiocyanate should be used in colorimetric determination?

5. What should be done, if some interfering substances are present?

6. What is the use of standard calibration curve in this experiment?

7. Define Lambert-Beer’s Law?

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EXPERIMENT -12

DETERMINATION OF VISCOSITY OF OIL BY OSTWALD’S VISCOMETER

AIM: To determine the viscosity of a liquid by using Ostwald`s viscometer.

APPARATUS REQUIRED: Ostwald`s viscometer, stop watch, beakers

CHEMICALS REQUIRED: Distilled water and oil sample.

PRINCIPLE:

When a liquid flows, it has an internal resistance to flow, an internal friction, which is

called its viscosity. Viscosity is a characteristic property of all liquids. Example- Honey flows

less readily than water because honey has a higher viscosity than water (honey has more

viscosity than water).

The Ostwald’s Viscometer method relates the flow of a liquid through a capillary tube

with the coefficient of viscosity.

PROCEDURE:

1. Take a clean and dry Ostwald`s viscometer and introduce a measured volume (10mL)

of distilled water into the bulb B (on left hand side).

2. Then suck the liquid into the bulb A (on right hand side) till the level of the liquid

rises above the mark X.

3. At that marked X point, ensure that there is no air bubble along the bulb as well as

bulb should be full of liquid.

4. Allow the liquid to flow through the capillary tube from X to Y and note the time of

flow`t1 (by using stop watch) in seconds.

5. epeat the process at least 3 times to get an average value of t1 .

6. Now dry the viscometer and repeat the total process for the given unknown sample

and values are tabulated to get an average value of .

DIGRAM-

OSTWALD’S

VISCOMETER

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OBSERVATIONS AND CALCULATIONS:

S.No Liquid Time flow in seconds

Average Trail-1 Trail-2 Trail-3

1.

Distilled Water

32 32 32

(

32

2. Oil Sample 2040 2040 2040

2040

= viscosity of the oil sample =? centipoise

= viscosity of distilled water = 0.815 centipoise

= density of the oil sample = 0.919g/cm3

= density of distilled water = 0.99594 g/cm3

= time flow of the oil sample = 2040s

= time flow of distilled water = 32s

RESULT:

Viscosity of the given unknown liquid centipoise.

VIVA QUESTIONS:

1. What is viscosity? Discuss its significance for a lubricant.

2. Discuss the effect of temperature and pressure on the viscosity of lubricating oil?

3. What are the differences between Redwood viscometer No 1 & No 2?

4. Define (i) Absolute viscosity & (ii) viscosity index?

5. How viscosity index is determined.

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EXPERIMENT -13

ESTIMATION OF MANGANESE IN KMnO4 BY COLORIMETRIC METHOD

AIM: To estimate the amount of manganese present in the given KMnO4 sample by

Colorimetry.

APPARATUS REQUIRED: Colorimeter, cuvettes, standard flask and graduated pipette.

CHEMICALS REQUIRED: KMnO4

PRINCIPLE:

The analysis which involves the measurement of absorption of light radiation in visible

region is known as colorimetry. Knowing the intensity of light absorbed and by making

suitable calibration, it is possible to find out concentration of any given solute. Colorimeter

works on the Beer-Lamberts law.

BEER-LAMBERT΄S LAW:

According to Beer-lamberts law the absorbance of light by a colored solution directly

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

A= εCl

Where, A = Absorbance (or) Optical Density

C = Concentration of the solution (in M)

l = thickness of the medium

ε = Molar absorption coefficient

A set of standard solutions of KMnO4 of known concentration are prepared and intensity of

color is measured using colorimeter. The concentration of test sample is calculated from

calibration curve, with optical density on y-axis and volume of KMnO4 on x-axis

PROCEDURE:

Part-1: Preparation of Various KMnO4 Solutions:

1. Pipette out 1mL to 8ml of standard 0.1 N KMnO4 solutions into 100ml standard flask and

label them as A,B,C,D,E,F,G,H.

2. Dilute the solutions and make up to the mark with distilled water and shake well.

3. Measure the absorbance of each solution using colorimeter.

Part-2: Estimation of Mn in Standard Solution:

1. Calibrate the colorimeter by placing a cuvette containing distilled water in and selecting

the appropriate filter.

2. Take the prepared KMnO4 solutions in to cuvettes one by one and place them in to the

colorimeter and note the reading in terms of optical density.

OBSERVATION AND CALCULATIONS:

Part-1: Preparation of Various KMnO4 Solutions:

N2 = Normality of standard KMnO4 =0.1N

V2= Volume of standard KMnO4 = ---mL

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N1 = Normality of prepared KMnO4 samples = 1N

V1 = Volume of prepared KMnO4 samples = 100ml

Part-2: Estimation of Mn in Standard Solution:

S.No. Flask

Label

Volume of

KMnO4

(in mL)

Concentration of

KMnO4

(in N)

Absorbance/O.D.

(in nm)

1 A 1 0.001 0.40

2 B 2 0.002 0.46

3 C 3 0.003 0.50

4 D 4 0.004 0.56

5 E 5 0.005 0.60

6 F 6 0.006 0.74

7 G 7 0.007 0.67

8 H 8 0.008 0.70

9 Unknown 0.011 0.63

Weight of Mn in KMnO4:

According to formul;a weight,

158g of KMnO4 contains 55g of Mn

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RESULT: The amount of Mn present in the given KMnO4 sample is 3.85g/100mL

MODEL GRAPH:

O.D.

(in nm)

Concentration (in M)

VIVA QUESTIONS:

1. Write Lambert’s Law?

2. What is Beer’s Law?

3. What is Optical density?

4. What is lmax?

5. What is the use of calibration curve?

6. What is the wave length range of U.V. and visible light?

7. What is monochromator?

8. What is blank solvent?

9. Why some compounds exhibit colors?

10. In this experiment, a graph is plotted by taking which parameters?

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EXPERIMENT -14

CONDUCTOMETRIC TITRATION OF MIXTURE OF ACIDS VS STRONG BASE

AIM: To determine the Composition of mixture of acids (acetic acid +HCl) by titrating

against NaOH Conductometrically.

APPARATUS REQUIRED: Conductometer, conductivity cell, beaker, pipette, and burette

CHEMICALS REQUIRED: HCl Solution, acetic acid, 1N NaOH.

PRINCIPLE:

When a mixture containing acetic acid and HCl is titrated against NaOH, strong acid (HCl) is

neutralized first. The neutralization of weak acid (CH3COOH) starts only after complete

neutralization of strong acid (HCl). Thus the conductance titration curve will be marked by

two breaks. The first one corresponds to the neutralization point of strong acid (HCl) and

second to that of weak acid (CH3COOH).

NaOH + HCl NaCl + H2O

NaOH + CH3COOH CH3COONa + H2O

PROCEDURE:

1. Take a clean beaker and add 20mL of HCl and 20mL of CH3COOH.

2. Fill the burette with 1N NaOH.

3. Take a clean conductivity cell washed with distilled water and place the cell in the beaker.

4. Now add 1mL of NaOH from burette drop wise, stir the solution gently and note down the

change in conductance.

5. Repeat the same procedure and note down the measured conductance.

6. Plot the graph between conductance and volume of NaOH.

MODEL GRAPH:

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OBSERVATION AND CALCULATIONS:

S.No.

Volume of NaOH

(in mL)

Conductance

(in milliSiemens)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

12

10

8

6

4

4

4

5

5

7

8

10

11

12

13

14

15

16

18

18

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From the graph, V1=2.0ml,

V2-V1=3.0-2.0=1.0ml

Normality of strong acid=V1 X conc.of NaoH/vol.of Hcl=2X1/20=0.1N

Normality of weak acid= V2-V1 X conc.of NaoH/vol.of CH3COOH=1X1/20=0.05N

RESULT:

1. The normality of HCl is 0.1N.

2. The normality of CH3COOH is 0.05N

VIVA QUESTIONS:

1. Define an electrolyte.

2. Define strong and weak electrolyte.

3. How does conductance vary in this experiment?

4. What are strong and weak electrolytes in this experiment?

5. Write the reaction between Oxalic acid & NaOH solution

6. Write the chemical reaction involved in the experiment

7. What is pH? Give the pH range of phenolphthalein indicator

8. Define basicity. What is the basicity of CH3COOH?

9. Draw the rough graph for this experiment?

10. In the given acid mixture, which one initially reacted with NaOH solution?

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EXPERIMENT -15

ESTIMATION OF HCl BY POTENTIOMETRY

AIM: To determine the normality of HCl by titrating it against NaOH using potentiometer.

APPARATUS: Potentiometer, glass electrode, beaker, burette, pipette, glass rod

CHEMICALS REQUIRED: NaOH, HCl

PRINCIPLE:

When HCl is titrated against NaOH, the change in pH is reflected as change in EMF (E). The

change in pH/EMF depends upon the no. of H+

ions removed during the course of titration as

the equivalence point reaches the no. of H+ ions decreases thereby causing a rapid change in

EMF. After the equivalence point there is a small change in EMF due to addition of excess of

NaOH.

HCl + NaOH H2O + NaCl

Thus, if the EMF of the cell is plotted against the volume of NaOH a curve is obtained. The

point of intersection in curve gives the equivalence point. The change in EMF is much more

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rapid near the equivalence point. This principle applies to all the Potentiometric titrations, the

titration curve being the same for all. When the titration curve does not show a sharp

intersection point, a precise method where, values is plotted against the volume of

NaOH.

PROCEDURE:

1. Calibrate the potentiometer before starting the experiment.

2. Take 40mL of HCl in a beaker and immerse the glass electrode and the reference electrode

into the solution.

3. Fill the burette with 1N NaOH solution.

4. Carry out the titrations by adding 0.5mL of NaOH and measure the EMF at every stage.

5. Calculate the values.

6. Plot against the volume of NaOH. The maximum of the curve represents the

equivalence point.

OBSERVATION AND CALCULATIONS:

S. No.

Vol. of NaOH

added

(V in mL)

EMF

(E in mV)

(in mV)

(in mL)

(mV/mL)

MODEL GRAPH:

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1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

21

0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

347

336

333

328

325

322

318

317

302

278

-45

-97

-116

-135

-145

-145

-177

-202

-231

-267

-298

-

11

3

5

3

3

4

6

10

24

323

52

19

15

14

15

15

17

35

29

36

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

1

11

3

5

3

3

4

6

10

24

323

52

19

15

14

15

15

17

35

29

36

N1 = Normality of HCl =?

V1 = Volume of HCl taken in beaker = 40mL

N2 = Normality of NaOH = 1N

V2 = Volume of NaOH from graph = 8mL

RESULT:

The normality of HCl by titrating with NaOH using potentiometer is 0.2N.

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VIVA QUESTIONS:

1. Define electrochemical cell?

2. What is SCE? Give its EMF value?

3. What is the role of quinhydrone?

4. What is cell notation? Write the cell notation of the cell used in the experiment.

5. Write Nernst’s equation.

6. How do you get EMF of a cell consisting of two electrodes?

7. Give the relation between PH and EMF of quinhydrone electrode?

8. What are the advantages of Potentiometric titrations over Conductometric titrations?

9. Define single electrode potential?

10. What is the role of salt bridge?

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EXPERIMENT-16

PREPARATION OF BAKELITE

AIM: To Prepare Bakelite (phenol formaldehyde resin) in the laboratory.

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APPARATUS: Beaker, Conical flask, glass rod, measuring cylinder

CHEMICALS REQUIRED: Glacial acetic acid, 40 % formaldehyde solution, phenol,

conc. HCl and distilled water

PRINCIPLE: Phenol resins are condensation polymerization product of phenolic derivative

with aldehyde (like formaldehyde). It is prepared by condensing phenol with formaldehyde in

presence of acid or alkaline catalyst.

Step I:- Formation of Novolac:

Step-2: Formation of Bakelite from Novolac:

Novolac + Hexamethylene Tetraamine Bakelite

PROCEDURE:

1. Place 5ml of glacial acetic acid and 2.5ml of formaldehyde solution in 500 ml beaker.

2. Add 2g of phenol and 1ml of conc. HCl solution in it.

3. Heat the solution slowly with constant stirring for 5minutes.

4. A large mass of pink Plastic is formed.

5. The residue obtained is washed several times with distilled water.

6. Dry the product and calculate the yield accurately.

RESULT: The weight of obtained Bakelite is 4.3g.

VIVA QUESTIONS:

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1. What are the monomers of Bakelite?

2. What are properties of Bakelite?

3. What type of polymerization involved preparing Bakelite?

4. What are the applications of Bakelite?

EXPERIMENT-17

PREPARATION OF UREA FORMALDEHYDE

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AIM: To prepare urea formaldehyde resin in the laboratory.

APPARATUS: Beaker, Conical flask, glass rod, measuring cylinder

CHEMICALS REQUIRED: 40%formaldehyde solution, urea, conc. H2SO4

PRINCIPLE: Urea formaldehyde is also known as urea methanol so named for its common

synthesis pathway. The plastic is made from urea and formaldehyde heated in presence of a

base.

PROCEDURE:

1. Place 5 ml of 40% formaldehyde solution in 100ml beaker.

2. Add 2 .5g of urea with constant stirring till saturated solution is obtained.

3. Now add few drops of conc. H2SO4 with constant stirring.

4. A large mass of white solid formed.

5. Wash with distilled water and dry in folds of filter paper.

6. Calculate the yield accurately.

RESULT: The weight of obtained urea formaldehyde resin is 5.3g.

VIVA QUESTIONS:

1. What are the monomers of urea-formaldehyde resin?

2. What are properties of urea-formaldehyde resin?

3. What type of polymerization involved preparing urea-formaldehyde resin?

4. What are the applications of urea-formaldehyde resin?