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Engineering Chemistry Lab Manual
2019
ACADEMIC CALENDAR
EVEN SEMESTER, 2019
Note:
1. *Additional lab will start from the next day of last semester examination
Department of Chemistry, SMIT
Page 1
Engineering Chemistry Lab Manual
2019
EVALUATION SCHEME
THEORY PAPER EVALUATION SCHEME
LABORATORY EVALUATION SCHEME
TOTAL MARKS (100) = INTERNAL (60 MARKS) + EXTERNAL (40 MARKS)
DAILY LAB EVALUATION = 10 marks
Marks
Experiment
set up/
Preperation
Execution of
Experiment
Data
Calculation
And Result
Knowledge of
Student
Lab File
2
2
2
2
2
10 marks
Department of Chemistry, SMIT
Page 2
Engineering Chemistry Lab Manual
2019
TIME TABLE
Engineering Chemistry Lab (CH 1163)
Day
Department of Chemistry, SMIT
Time(FN)
Time(AN)
Page 3
Engineering Chemistry Lab Manual
2019
LABORATORY MANUAL OF
I Year
ENGINEERING CHEMISTRY
LABORATORY
(CH1163)
Department of Chemistry, SMIT
Page 4
Engineering Chemistry Lab Manual
2019
SUBJECT CODE: CH 1163
SUBJECT TITLE: ENGINEERING LABORATORY
No of Credit: 1.5
Contact Hours / Week: 3hrs
OBJECTIVES

The main aim of the course is to teach basic computer programming concepts and apply them
to computer based problem solving methods.

To teach the student problem solving using C.

To introduce the students to the field of programming using C language.

To introduce the student to data structures such as arrays, lists, stacks etc.
PREREQUISITES
There are no specific prerequisites for this lab.
LEARNING OUTCOMES
Upon completion of the course, students will be able to:

Solve moderately difficult problem using C language.

The students will be able to enhance their analyzing and problem solving skills and use the
same for writing programs in C.

Write error free code in C

Debug syntax errors prompted by the C compiler.
Lab Incharge
Even Semester (Jan-June 2019)
Name & Signature 1:
Name & Signature 2:
Department of Chemistry, SMIT
Page 5
Engineering Chemistry Lab Manual
2019
STANDARDIZATION OF KMnO4 SOLUTION
Experiment No: 1
Date: - ……………
AIM: -To standardize the supplied potassium permanganate solution with the help of prepared
standard solution of oxalic acid.
APPARATUS REQUIRED: Burette, Pipette, Standard flask, conical flask, Funnel, Test tube.
REAGENT REQUIRED: Oxalic acid, Dil.H2SO4, Potassium permanganate.
PRINCIPLE:
Acidified KMnO4 oxidizes oxalic acid as follows:
2KMnO4 + 3H2SO4
3H2O + 5(O)
5H2C2O4 + 5(O)
K2SO4 + 2MnSO4 +
10CO2 + 5H2O
5 molecules of oxalic acid = 5 atoms of (O)
∴ Equivalent weight of H2C2O4.2H2O = Mol. Wt / 2 = 126/2 = 63
.
A standard solution of oxalic acid is prepared by accurately weighing the crystals. A definite volume
of this solution is acidified and titrated against KMnO4 for the standardization.
PROCEDURE:
1. Preparation of Standard Solution:About 1.6gms of oxalic acid crystals are accurately weighed into a 250ml standard flask. It is first
dissolved in a little distilled water and then made up to the mark with distilled water and shaken well
for uniform concentration.
2. Standardization of KMnO4 Solution:Fill the burette with the supplied KMnO4 solution ensuring that there is no air gap in the burette. 25ml
of the prepared oxalic acid solution is pipette out into a conical flask and 1 test tube of dilute H2SO4 is
added. The product is heated to 70o to 80o C and the hot solution is titrated against KMnO4 solution in
the burette until a permanent pale pink colour is obtained.
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
2019
OBSERVATION AND CALCULATIONS:
Weight of oxalic acid taken
W=__________gms.
Strength of the solution prepared =
(N1) Oxalic acid
𝑊𝑡 𝑥 1000
𝐸𝑞.𝑤𝑡 𝑥 𝑉𝑜𝑙𝑢𝑚𝑒
𝑊(___________) 𝑥 1000
=
63 𝑥 𝑉(___________ )
= _______________ N.
1. Standardization of KMnO4 solution:Solution taken in the burette
= given supplied KMnO4 solution
.
Solution taken in the conical flask = 25ml of oxalic acid soln + 1t.t of dil.
H2SO4. (Heat upto 70o to 80o C)
Indicator used …………………… = KMnO4 ( self indicator).
End point …………………………..= from colourless to pale pink.
Oxalic acid X KMnO4
Trial No.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
1.
2.
25
25.1
3.
25.2
Agreeing value = V1of KMnO4 = ___________ml.
∴ Strength of the given KMnO4 solution (N2) = 25 X N1of oxalic acid = 25 X(
V1 of KMnO4
(
)
)
N2 of KMnO4 = ___________________ N.
RESULT: Strength of KMNO4 solution =____
N
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/ pipette
while taking readings.
6. Ensure that there are no air bubbles in the burette.
7. For each titration, use same number of drops of indicator.
8. Read lower meniscus in case of colorless solution.
Date:
Department of Chemistry, SMIT
Signature of teacher/technician:
Page 7
Engineering Chemistry Lab Manual
2019
Viva:
1.
What is the principle of volumetric analysis ?
2.
What is titration ?
3.
What is indicator ?
4.
What is end point ?
5.
Why a titration flask should not be rinsed ?
6.
What are primary and secondary standard substances ?
7.
Burette and pipette must be rinsed with the solution with which they are filled, why ?
8.
It is customary to read lower meniscus in case of colourless and transparent solutions and
upper meniscus in case of highly coloured solutions, why ?
9.
What is a normal solution ?
10. Why the last drop of solution must not be blown out of a pipette ?
11. Pipette should never be held from its bulb, why ?
12. What is acidimetry and alkalimetry ?
13. What do you mean by 1.0 M solution ?
14. What is meant by the term ‘concordant readings’ ?.
15. Can one take oxalic acid solution in the burette and sodium hydroxide solution in the
titration flask ?
16. What is the difference between an end point and an equivalence point ?
17. What is basicity of an acid ?
18. What is the relation between equivalent mass of acid and its molecular mass ?
Department of Chemistry, SMIT
Page 8
Engineering Chemistry Lab Manual
2019
ESTIMATION OF MOHR’S SALT
Experiment No: 2
Date: - ……………
AIM: - To estimate the weight of Mohr’s salt (Ferrous ammonium sulphate) crystal dissolved in 250ml
using approximately decinormal potassium permanganate solution and pure oxalic acid crystals.
APPARATUS REQUIRED: Burette, Pipette, Standard flask, conical flask, Funnel, Test tube.
REAGENT REQUIRED: Mohr salt, Oxalic acid, Dil.H2SO4, Potassium permanganate.
PRINCIPLE:
Acidified KMnO4 oxidizes oxalic acid and ferrous ammonium sulphate as follows:
2KMnO4 + 3H2SO4
5H2C2O4 + 5(O)
10FeSO4 + 5H2SO4 +5(O)
K2SO4 + 2MnSO4 + 3H2O + 5(O)
10CO2 + 5H2O
5Fe2(SO4)3 + 5H2O.
5 molecules of oxalic acid = 10 molecules of ferrous sulphate.
= 5 atoms of (O) = 10 equivalents.
∴ Equivalent weight of H2C2O4.2H2O = Mol. Wt / 2 = 126/2 = 63
∴ Equivalent weight of FeSO4 (NH4)2SO4.6H2O = Its Mol. Wt = 392.
A standard solution of oxalic acid is prepared by accurately weighing the crystals. A definite volume
of this solution is acidified and titrated against KMnO4 for the standardization.
A definite volume of Mohr’s salt solution is acidified and titrated against KMnO4 solution. Thus
Mohr’s salt solution is standardized. For both titrations pink coloured KMnO4 acts as self indicator.
PROCEDURE:
1. Preparation of Standard Solution :About 1.6gms of oxalic acid crystals are accurately weighed into a 250ml standard flask. It is first
dissolved in a little distilled water and then made upto the mark with distilled water and shaken well
for uniform concentration.
2. Standardization of KMnO4 solution :Fill the burette with the supplied KMnO4 solution ensuring that there is no air gap in the burette.
25ml of the prepared oxalic acid solution is pipette out into a conical flask and 1 test tube of dilute
H2SO4 is added. The product is heated to 70o to 80o C and the hot solution is titrated against KMnO4
solution in the burette until a permanent pale pink colour is obtained.
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
2019
3. Estimation of Mohr’s salt :25ml of the supplied Mohr’s salt solution is pipette out into a conical flask and 1 test tube of dil. H2SO4
is added and titrated against KMnO4 solution in the burette. Appearance of permanent pale pink colour
marks the end point. Titration is repeated to get agreeing values.
NOTE: The above experiment may also be given as: Estimate the weight of oxalic acid crystals present
in the whole of the solution given in 250ml standard flask. You are provided with approximately 0.1N
KMnO4 solution and pure crystals of Mohr’s salt.
OBSERVATION AND CALCULATIONS:
Weight of oxalic acid taken
Strength of the solution prepared =
W=__1.6________gms.
𝑊𝑡 𝑥 1000
𝐸𝑞.𝑤𝑡 𝑥 𝑉𝑜𝑙𝑢𝑚𝑒
(N1) Oxalic acid
=
𝑊(___________) 𝑥 1000
63 𝑥 𝑉(___________ )
= _____0.1__________ N.
1. Standardization of KMnO4 solution:Solution taken in the burette
= given supplied KMnO4 solution.
Solution taken in the conical flask = 25ml of oxalic acid soln + 1t.t of dil.
H2SO4. (Heat upto 70o to 80o C)
Indicator used …………………… = KMnO4 (self indicator).
End point …………………………= from colourless to pale pink.
Oxalic acid X KMnO4
Trial No.
1.
Final Burette Reading (ml)
26.2
26.2
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
Agreeing value = V1of KMnO4 = ___________ml.
2.
3.
26.4
∴ Strength of the given KMnO4 solution (N2) = 25 X N1of oxalic acid = 25 X N1(
V1 of KMnO4
V1 (
)
)
N2 of KMnO4 = _______0..095____________ N
Department of Chemistry, SMIT
Page 10
Engineering Chemistry Lab Manual
2019
2. Estimation of Mohr’s salt:Solution taken in the burette
= given supplied KMnO4 solution.
Solution taken in the conical flask = 25ml of Mohr’s salt sol + 1t.t of dil. H2SO4
Indicator used …………………… = KMnO4 (self-indicator).
End point …………………………. = from colorless to pale pink.
Mohr’s salt X KMnO4
Trial No.
1.
2.
Final Burette Reading (ml)
23.9
23.8
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
3.
23.7
Agreeing value = V2 of KMnO4 = ___________ml.
∴ Strength of Mohr’s salt solution (N3) = V2 (
) of KMnO4 X N2(
25
N3 of Mohr’s salt = ________0.09___________ N.
) of KMnO4
Weight of Mohr’s salt present in the whole of
the solution given in the 250ml standard flask = N3 (
)of Mohr’s salt x 392 x Volume (250)
1000
= _______________gms.
RESULT: Weight of Mohr’s salt dissolved in whole solution = _8.82___
gms.
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/ pipette while taking
readings.
6. Ensure that there are no air bubbles in the burette.
7. For each titration, use same number of drops of indicator.
8. Read lower meniscus in case of colourless solution.
Date:
Department of Chemistry, SMIT
Signature of teacher/technician:
Page 11
Engineering Chemistry Lab Manual
2019
VIVA:
1.
What kind of titration is, oxalic acid/ mohr salt with KMnO4?
2.
Why is it called a redox titration?
3.
What is the oxidizing and reducing agent in the titration? Name the substance oxidised and
reduced.
4.
What is the formula of mohr salt?
5.
What is the molecular mass of mohr salt?
6.
What is the change in oxidation state of Mn in KMnO4 during the titration?
7.
What is the oxidation state of Mn in KMnO 4 ?
8.
W h a t i s t h e u s e o f H 2 S O 4 in the titration?
9.
Why reaction mixture is heated in oxalic acid titrations.?
10.
What is the indicator used in KMnO4 titrations?
11.
Why is KMnO4 is called a self indicator?
12.
What is a double salt?
13.
What is the difference between double salt and complex salt?
14.
What kind of salt is mohr salt?
15.
What is water of crystallisation?
Department of Chemistry, SMIT
Page 12
Engineering Chemistry Lab Manual
2019
ESTIMATION OF IODINE
Experiment No: - 3
Date:-……………
AIM: - To estimate the weight of iodine per liter of the given solution, using decinormal Na2S2O3
solution and also the pure crystals of K2Cr2O7.
APPARATUS REQUIRED: Burette, Pipette, Standard flask, conical flask, Funnel, Test tube,
measuring cylinder.
REAGENT REQUIRED: Sodium thiosulphate, Potassium dichromate, Iodine, 10% Potassium iodide
solution, Conc. HCl, Starch.
PRINCIPLE :
Acidified K2Cr2O7 oxidizes KI to iodine according to the following equations:
K2Cr2O7+ 8HCl
4H2O + 3(O)
6KI + 3H2O + 3(O)
2KCl + 2CrCI3 +
6KOH + 3I2
Sodium thiosulphate reacts with iodine forming tetarthionate and sodium iodide according to the
following equation:
2 Na2S2O3 + I2
Na2S4O6 + 2NaI.
From the above equation,
2 molecules of Na2S2O3
∴ 1 molecule of Na2S2O3
∴ Equivalent weight of Iodine
∴ Equivalent weight of K2Cr2O7
= 1 molecule of Iodine.
= 2 atoms of Iodine.
= 1 atoms of Iodine.
= It’s atomic Wt = 127
= Mol.Wt /6 = 294.2/6 = 49.03
A standard solution of K2Cr2O7 is prepared in a 250ml standard flask by weighing 1.225 gms of the
salt. A known volume of this solution is titrated with KI in presence of HCl. Equivalent quantity of
iodine set free is titrated against given Na2S2O3 solution. Starch as indicator is to be added just prior to
get the end point. Thus Na2S2O3 solution is standardized.
A known volume of iodine solution is then titrated against the standardized sodium thiosulphate using
the same indicator. By this titration normality of iodine solution and hence grams per liter of iodine
are calculated.
Department of Chemistry, SMIT
Page 13
Engineering Chemistry Lab Manual
2019
PROCEDURE:
1.
Preparation of Standard K2Cr2O7 Solution :K2Cr2O7 crystals given in the weighing bottle is weighed out into a 250ml standard flask. It is dissolved
in a little distilled water and then made upto the mark with distilled water and shaken well for uniform
concentration.
2. Standardization of Na2S2O3 solution:Fill the burette with the supplied Na2S2O3 solution ensuring that there is no air gap in the burette. 25ml
of the prepared K2Cr2O7 solution is pipette out into a clean conical flask. 1/3 t.t of conc.HCl and 10ml
of 10% KI solution are added. The solution becomes reddish brown due to the liberation of Iodine. It
is titrated against the sodium thiosulphate solution taken in the burette, till the reddish brown colour
turns to yellowish green. Now 1ml of freshly prepared starch solution is added to the solution. It turns
deep blue. The addition of sodium thiosulphate solution is continued drop by drop, until the blue colour
of the solution is completely discharged to get transparent green. The volume of thiosulphate solution
added is noted. The experiment is repeated to get agreeing values.
3. Estimation:25ml of the iodine solution is pipette out into a clean conical flask. It is diluted with two test tube of
distilled water. The sodium thiosulphate solution is added from the burette till the colour of the solution
in the conical flask changes from reddish brown to pale yellow. Now 1ml of freshly prepared starch
solution is added. The solution becomes deep blue. Titration is continued until the blue colour just
disappears to colourless. Titration is repeated to get concordant result.
OBSERVATION AND CALCULATIONS:
Weight of K2Cr2O7 taken
W = _______gms.
∴ Strength of the solution prepared (N1) K2Cr2O7 =
𝑊(___________) 𝑥 1000
49.03 𝑥 𝑉(___________ )
= ______________ N.
1. Standardisation of Na2S2O3 solution:Solution taken in the burette
= given Na2S2O3 solution.
Solution taken in the conical flask = 25ml of K2Cr2O7 + 1/3 t.t of conc.HCl
+ 10ml of 10% KI solution.
Indicator used …………………… = 1ml starch solution.
End point …………………………= disappearance of dark blue colour
(Transparent Green Remains).
K2Cr2O7X Na2S2O3
Trial No.
1.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of Na2S2O3 (ml)
Agreeing value = V1of Na2S2O3 = ___________ml.
Department of Chemistry, SMIT
2.
3.
Page 14
Engineering Chemistry Lab Manual
2019
∴Strength of the given Na2S2O3 solution (N2) = 25 X N1(
) of K2Cr2O7
V1 (
)f Na2S2O3
N2 of Na2S2O3 = ___________________ N.
2. Estimation of Iodine:Solution taken in the burette
= given Na2S2O3 solution.
Soln taken in the conical flask
= 25ml of Iodine soln + 2 t.t of distilled water.
Indicator used ……………………..= 1ml starch solution.
End point ………………………….= from dark blue to colourless.
Iodine X Na2S2O3
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of Na2S2O3 (ml)
Agreeing value = V2of Na2S2O3 = ___________ml.
∴ Strength of Iodine solution
N3 of Iodine
= V2 (
) of Na2S2O3 X N2 (
25
= ___________________ N.
∴ Weight of Iodine per liter of the given solution = N3 (
) of Na2S2O3
) of Iodine X 127
= ____________gms.
RESULT: - Weight of Iodine in a liter of the solution = ____________gms.
Precaution:
1.
2.
3.
4.
5.
6.
7.
8.
9.
Clamp burette vertically.
Before use, rinse the burette with given prepared solution.
Do not hold the pipette from the bulb.
Do not blow the last drop from the pipette.
Always keep your eye at the same level as the level of label of liquid in burette/pipette while taking
readings.
Ensure that there are no air bubbles in the burette.
For each titration, use same number of drops of indicator.
Read lower meniscus in case of colourless solution.
Do not rinse the conical flask.
Date:
Department of Chemistry, SMIT
Signature of teacher/technician:
Page 15
Engineering Chemistry Lab Manual
2019
Viva :
1. Why iodine value is important?
2. What does iodine value indicate?
3. What does a high iodine number mean?
4. Why is starch added towards the end point?
5. Why starch is used as indicator in iodine value analysis?
6. Why we use freshly prepared starch solution?
7. What is difference between iodometric and Iodimetric titration?
8. How does iodine work as an indicator?
9. Why is KI used in iodometric titration?
10. Why does iodine turn blue black in the presence of starch?
11. What is the principle of iodometric titration?
12. What will react with iodine?
13. Which indicator is used in iodometric titration?
14. Why sodium thiosulphate is used in titration?
15. Why is starch not added at the beginning of titration?
Department of Chemistry, SMIT
Page 16
Engineering Chemistry Lab Manual
2019
ESTIMATION OF TOTAL HARDNESS OF WATER
Experiment No: - 4
Date:-……………
AIM:- To estimate the total hardness of the given sample of water using 0.02M E.D.T.A.
solution and solid calcium carbonate, measuring cylinder.
APPARATUS REQUIRED:
Test tube, measuring cylinder .
Burette,
Pipette,
Standard
flask,
Conical
flask,
Funnel,
REAGENTS REQUIRED: EDTA, Calcium Carbonate, Dil HCl, NaOH solution, Buffer solution,
Erichrome Black-T,
PRINCIPLE:
Ethylene diamine tetra acetic acid (E.D.T.A) or its sodium salt (Na2H2Y) forms soluble complex with
hardness causing Mg+2 and Ca+2 ions. (Y4-  Ethylene diamine tetra acetate ion)
Na2H2Y + Ca+2
Na2H2Y + Mg+2
Na2CaY + 2H+
Na2MgY + 2H+
If a small amount of a dye Eriochrome Black-T is added to a solution containing Mg+2 and Ca+2 ions
at a pH of 10, the solution attains wine red colour. If E.D.T.A is added as titrant, it forms a complex
with dissolved Mg+2 and Ca+2 ion. After the sufficient addition of E.D.T.A to complex Mg+2 and Ca+2
ions the solution turns to blue from wine red. It is the end point. Hardness is expressed in parts by
weight CaCO3 per million parts by weight of water (ppm CaCO3 or its equivalent).
A standard solution of CaCO3 is prepared by weighing 0.5gms of the substance. A known volume of
this solution is titrated against given E.D.T.A solution using Eriochrome Black-T indicator. A buffer
solution of NH4Cl – NH4OH is added to maintain the pH 10. Thus E.D.T.A is standardized. Experiment
is repeated with water sample. From the titrate value of molarity of water sample and hence hardness
of water is calculate,
(Molecular Weight of CaCO3 = 100).
PROCEDURE:
1. Preparation of Standard CaCO3 Solution:CaCO3 solid given in weighing bottle is accurately weighed into 250ml standard flask. Dilute HCl is
added slowly till the solid is completely dissolved. The solution is just neutralized with NaOH solution
till a white ppt. is formed. Then the ppt. is dissolved in a minimum amount of dil. HCl. The solution is
made up to the mark with distilled water and shaken well for uniform concentration.
2. Standardisation of E.D.T.A solution:Department of Chemistry, SMIT
Page 17
Engineering Chemistry Lab Manual
2019
Fill the burette with the supplied EDTA solution ensuring that there is no air gap in the burette.
25ml of the prepared CaCO3 solution is pipette into a clean conical flask. 2ml of (NH4Cl-NH4OH)
buffer solution is added followed by three drops of E.B.T indicator. The solution turns to violet in
colour. This is titrated against E.D.T.A solution taken in a burette. The end point is indicated when the
solution turns to blue colour without reddish tinge. Titration is repeated to get concordant values.
3. Estimation:25ml of the given hard water is pipette out into a clean conical flask. 2ml of (NH4Cl-NH4OH) buffer
solution is added followed by three drops of E.B.T indicator. The solution turns to violet in colour.
This is titrated against E.D.T.A solution taken in a burette. The end point is indicated when the solution
turns to blue colour without reddish tinge. Titration is repeated to get agreeing values.
OBSERVATION AND CALCULATIONS:
Weight of CaCO3
W = _______gms.
1. Standardisation of EDTA Solution:Solution taken in the burette = given E.D.T.A solution.
Soln taken in the conical flask = 25ml of CaCO3 + 2ml of (NH4Cl-NH4OH)
Buffer soln
Indicator used ………………...= 3drops of EBT.
End point ……………………..= change of colour from violet to blue.
CaCO3 dissolved solutionVs E.D.T.A solution
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of E.D.T.A (ml)
Agreeing value = V1of E.D.T.A = ___________ml.
2. Estimation of Hardness of water:Solution taken in the burette
= given E.D.T.A solution.
n
Sol taken in the conical flask
= 25ml of hard water + 2ml of buffer soln
Indicator used ……………………...= 3drops of EBT.
End point …………………………..= change of colour from violet to blue.
Hard Water vs E.D.T.A solution
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of E.D.T.A (ml)
Agreeing value = V2of E.D.T.A = ___________ml.
Department of Chemistry, SMIT
Page 18
Engineering Chemistry Lab Manual
2019
CALCULATIONS:
V1 ml of E.D.T.A solution reacts completely with 25ml of Ca+2 solutions.
∴V1 ml of E.D.T.A solution = 25ml of Ca+2 solution =
𝑊
10
gms of CaCO3.
(note: W=weight of calcium carbonate dissolved in 250 ml of water)
∴1 ml of E.D.T.A solution =
𝑊
10 𝑥 𝑉1
gms of CaCO3
V2 ml of E.D.T.A solution reacts completely with Ca+2 and Mg+2 ions present in 25ml of water sample.
∴25ml of water sample = V2 ml of E.D.T.A solution
=
𝑊(
)𝑥 𝑉2 (
10 𝑥 𝑉1 (
)
gms of CaCO3
)
= ____________ gms of CaCO3
25ml of water sample contains (x) = ________________gms of CaCO3
∴ 1 million i.e. 106 gms of water contains =
) 𝐗 𝟏𝟎𝟔
x(
25
ppm CaCO3
= _________ppm CaCO3
∴ Hardness of water = ________________________ ppm CaCO3
RESULT:
Total hardness of the given sample of water = _____
PRECAUTION:
1.
2.
3.
4.
5.
6.
7.
8.
9.
ppm. CaCO3 or its equivalent.
Clamp burette vertically.
Before use, rinse the burette with given prepared solution.
Do not hold the pipette from the bulb.
Do not blow the last drop from the pipette.
Always keep your eye at the same level as the level of label of liquid in burette/pipette while
taking readings.
Ensure that there are no air bubbles in the burette.
For each titration, use same number of drops of indicator.
Read lower meniscus in case of colourless solution.
Do not rinse the conical flask.
Date:
Department of Chemistry, SMIT
Signature of teacher/technician:
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Engineering Chemistry Lab Manual
2019
Viva:
1. How do you measure water hardness?
2. What is hardness water?
3. What are the different types of hardness water?
4. What are the causes of hardness of water?
5. How can we remove the temporary hardness of water?
6. Which indicator is used in estimation of hardness of water?
7. What do you mean by hardness of water?
8. How will you distinguish between permanent and temporary hardness of water?
9. What is complexometric titration?
10. Why di-sodium salt of EDTA is chosen for determination or hardness?
11. What is the indicator used in EDTA titration?
12. Why is NH3-NH4Cl buffer solution added during EDTA titration?
13. Why does the wine red colour obtained when EBT is added to buffered hard water?
14. Why does the colour of the solution change from wine red to blue at the end point?
15. Why does phenolphthalein and methyl orange not used in EDTA titration?
16. What type of indicator is EBT ?
17. What will be the colour of the solution at the end point if buffer solution is not added?
Department of Chemistry, SMIT
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2019
ESTIMATION OF FERRIC CHLORIDE
Experiment No: - 5
Date:-…………….
AIM: - To estimate the weight of Ferric chloride in a given sample using a decinormal
K2Cr2O7 and pure crystals of Mohr’s salt.
APPARATUS REQUIRED:
Test tube, Measuring cylinder.
Burette,
Pipette,
Standard
flask,
conical
flask,
Funnel,
REAGENTS REQUIRED: Mohrs salt, Potassium dichromate, Diphenyl amine, mixture of Sulphuric
acid and phosphoric acid, Ferric chloride, Conc. HCl, Mercuric chloride stannous chloride.
PRINCIPLE:
Acidified K2Cr2O7 oxidizes Mohr’s salt as follows:K2Cr2O7+ 4H2SO4
6FeSO4 + 3H2SO4 + 3(O)
K2SO4 + Cr2 (SO4)3 + 4H2O + 3(O)
3Fe2 (SO4)3 + 3H2O
6 molecules of Mohr’s salt = 3 atoms of (O) = 6 equivalents.
∴ Equivalent weight of Mohr’s salt = it’s molecular Wt = 392
Ferric chloride is not a reducing agent. It does not react with K2Cr2O7 solution. Ferric chloride can be
reduced to ferrous chloride with stannous chloride in presence of HCl.
2FeCl3 + SnCl2
2FeCl2 + SnCl4
FeCl2 is a reducing agent and it reacts with K2Cr2O7 in presence of dil. HCl acid as
K2Cr2O7 + 8HCl
[2FeCl2 + 2HCl + (O)
On adding,
K2Cr2O7+ 14HCl + 6FeCl2
From the equation,
2 molecules of FeCl3
2KCl + 2CrCl3 + 4H2O + 3(O)
2FeCl3 + H2O] X 3
2KCl + 2CrCl3 + 6FeCl3 + 7H2O
= 2 molecules of FeCl2
= 1 atom of oxygen
= 2 equivalents of oxygen
∴ Equivalent weight of FeCl3 = molecular weight = 162.2
∴ Equivalent weight of FeCl2 = molecular weight = 126.8
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(Note: equivalent weight is calculated with respect of change of Oxidation Number)
A slight SnCl2 should be added to ensure the complete reduction FeCl3 to FeCl2. Unreacted SnCl2 also
reacts with K2Cr2O7. Unreacted SnCl2 must be destroyed by the addition of saturated mercuric chloride
solution.
SnCl2 + 2HgCl2
Hg2Cl2 (Silky white ppt) + SnCl4
At this stage silky white precipitate should be obtained. A black precipitate of finely divided mercury
may be produced if too much of SnCl2 is present. So, if black ppt. is obtained, the reduction must be
repeated. If there is no ppt. also the reduction must be repeated. The ferrous salt in the solution is then
estimated by titration against the given K2Cr2O7 solution using diphenylamine as internal indicator.
Given K2Cr2O7 solution is standardized by titration against a standard Mohr’s salt solution prepared.
PROCEDURE:
1. Preparation of Mohr’s salt Solution:About 10gms of Mohr’s salt crystals kept in a weighing bottle are transferred into a 250ml standard
flask after weighing accurately. The salt is dissolved in a little (1t.t) dilute H2SO4. The solution is made
upto the mark with distilled water and shaken well.
2. Standardization of K2Cr2O7:Take the supplied K2Cr2O7 solution and fill the burette ensuring that there is no air gap in the burette.
25ml of Mohr’s salt solution is pipette into a conical flask. About 8 drops of 1% solution of
diphenylamine is added as the indicator. Then about 10ml of Sulphuric acid-Phosphoric acid mixture
is added. The solution is titrated slowly with constant stirring against K2Cr2O7 solution in the burette.
Near the end point the solution becomes bluish violet and remains permanent. This marks the end
point. The titration is repeated to get agreeing values.
3. Estimation:25ml of the FeCl3 solution is pipette out into a conical flask. 1/3 test tube of conc. HCl is added and
heated just to boiling. To the hot solution SnCl2 solution is dropped from a dropping bottle until the
yellow colour of the solution just discharges. One or two drops are added in excess. The hot solution
is rapidly cooled under tap water to room temperature with protection from air. 5ml of saturated
mercuric chloride solution is rapidly added in one portion and with through mixing. A light silky white
precipitate should be obtained. If a black ppt. obtained, it should be rejected and the process must be
repeated with another 25ml of FeCl3 solution. About 8 drops of 1% solution of diphenylamine is added
as indicator. Then about 10ml Sulphuric acid- Phosphoric acid mixture is added and titrated against
K2Cr2O7 solution of the burette with constant stirring till the solution turns to intense purple or bluish
violet colour. This marks the end point. Titration is repeated to get agreeing values.
OBSERVATION AND CALCULATIONS:
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∴ Weight of Mohr’s salt taken
∴ Strength of the solution prepared (N1)
W = _________gms.
= W(
)X4
392
= ________________ N.
(N1) of Mohr’s salt
1. Standardization of K2Cr2O7 solution:Solution taken in the burette
= given K2Cr2O7 solution.
Soln taken in the conical flask
= 25ml of Mohr’s salt + 10 ml (H2SO4 +H3PO4)
Indicator used ……………………. = 8 drops of 1% diphenylamine.
End point …………………………..= Appearance of bluish violet colour.
Mohr’s Salt vs K2Cr2O7
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of K2Cr2O7 (ml)
Agreeing value = V1of K2Cr2O7 = ___________ml.
Strength of the given K2Cr2O7 solution = 25 X N1 (
V1 (
) of Mohr’s salt
) of K2Cr2O7
N2 of K2Cr2O7 = ___________________ N.
2. Estimation of FeCl3:Solution taken in the burette
= given K2Cr2O7 solution.
Solution taken in the conical flask = 25ml of FeCl3 soln + 1/3 conc. HCl (heated to
boiling) + SnCl2 (cooled) + 5ml of saturated
HgCl2 +10 ml (H2SO4 +H3PO4) mixture.
Indicator used ……………………. = 8 drops of 1% diphenylamine.
End point …………………………..= Appearance of bluish violet colour.
FeCl3 vs K2Cr2O7
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of K2Cr2O7 (ml)
Agreeing value = V2of K2Cr2O7 = ___________ml.
∴ Strength of FeCl3 solution = V2 (
) of K2Cr2O7 X N2 (
25
) of K2Cr2O7
N3 of FeCl3 = ___________________ N.
∴ Weight of FeCl3 present in the given 250ml solution = N3 (
Department of Chemistry, SMIT
) of (FeCl3) X 162.2 x 250
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1000
= _____________ gms.
RESULT:
Weight of FeCl3 dissolved in the whole of the solution kept in 250ml standard
Flask = __________________ gms.
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/pipette
while taking readings.
6. Ensure that there are no air bubbles in the burette.
7. For each titration, use same number of drops of indicator.
8. Read lower meniscus in case of colourless solution.
9. Do not rinse the conical flask.
Date:
Viva:
1.
2.
3.
4.
Signature of teacher/technician:
Why ferric chloride is converted into ferrous chloride?
How can you reduce ferric chloride?
Why excess SnCl2 is added during the reaction?
Hoe can you destroy unreacted SnCl2 ?
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
2019
ESTIMATION OF MnO2 IN PYROLUSITE
Experiment No: - 6
Date:- …………….
AIM:-To estimate the percentage of MnO2 in the given sample of Pyrolusite using approximately 1N
solution of oxalic acid and a standard solution of Potassium Permanganate Solution.
APPARATUS REQUIRED:
Test tube, Beaker.
Burette,
Pipette,
Standard
flask,
conical
flask,
Funnel,
REAGENTS REQUIRED: Potassium permanganate, Oxalic acid, Dil.H2SO4, Magnesium oxide.
PRINCIPLE:
A known weight of finely powdered sample is heated with known excess of Oxalic acid solution in
presence of dilute Sulphuric acid. Oxalic acid reduces MnO2 of the sample to manganese salt.
MnO2 + H2C2O4 + H2SO4
MnSO4 + 2CO2 + 2H2O
Acidified KMnO4 oxidizes excess oxalic acid as follows:
2KMnO4 + 3H2SO4
K2SO4 + 2MnSO4 + 3H2O + 5(O)
5H2C2O4 + 5(O)
10CO2 + 5H2O
Excess oxalic acid left over is determined by back titration against standard KMnO4 solution. From
this amount of Oxalic acid consumed by the sample is determined and hence the percentage of MnO2
in the sample is calculated.
Equivalent weight of MnO2 = Mol.Wt
2
= 86.94 = 43.47
2
PROCEDURE :1. Standardization of KMnO4 solution:Fill the burette with supplied KMnO4 solution ensuring that there is no air gap in the burette.
25ml of given 1N oxalic acid is pipette out in to 250ml standard flask, made up to the mark with
distilled water and shaken well. 25ml of this solution is pipette out into a conical flask. About 1 test
tube of dil. H2SO4 is added and heated gently nearby 70o to 80oC. Hot solution is titrated against the
standard solution of KMnO4 until the permanent pink colure is obtained. Titration is repeated to get
agreeing values.
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2. Estimation:About 0.6 gms of powdered sample is weighed accurately and transferred into a conical flask. Exactly
25ml of given 1N oxalic acid solution is pipette and 50ml of (4N) dil. sulfuric acid by measuring
cylinder are added to the conical flask. Mouth of the conical flask is covered with a glass funnel. It is
gently boiled until no black particles remains and becomes clear brown solution is obtained. Cool and
the cleared brown solution is transferred quantitatively into a 250ml standard flask. It is made up to
the mark with distilled water and shaken well. 25ml of the solution is pipette out into a conical flask.
1 test tube of dil. H2SO4 is added, heated gently nearby 70o to 80oC and titrated against standard KMnO4
sol. till a pale pink colour is obtained. Titration is repeated to get agreeing values. From this value
excess of oxalic acid left over is calculated.
OBSERVATION AND CALCULATIONS
1. Weight of pyrolusite transfer
W = ____________
_gms.
2. Standardization of KMnO4 Solution:Solution taken in the burette = 0.1N KMnO4 solution.
Solution taken in the flask = 25ml of Oxalic acid solution + 1t.t H2SO4
(heated up to 70o to 80oC )
Indicator used
= KMnO4 (self indicator.)
End point
= Change of colour from colorless to pale pink.
Oxalic Acid vs KMnO4
Trial No.
1.
2.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
3.
Agreeing value = V1of KMnO4 = ___________ml.
3. Estimation of MnO2:- Titration of ore solution :( Excess of Oxalic acid)
Solution taken in the burette = 0.1N KMnO4.
Solution in the flask
= 25 ml of experimental solution + 1 t.t of dilute
H2SO4. (heated up to 70o to 80oC).
Indicator used……………… = KMnO4 self indicator.
End point ………………
= change of colour from colorless to pale pink.
Ore solution vs KMnO4
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
Agreeing value = V2of 0.1 N KMnO4 = ___________ml.
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CALCULATION:
25ml of dilute oxalic acid = V1 ml of 0.1N KMnO4 = ________________ml.
Excess of oxalic acid
[25ml of ore solution]
= V2 ml of 0.1N KMnO2 = ________________ml.
∴ Volume of 0.1N KMnO4 equivalent to oxalic acid reacted = V1 – V2 ml.
Volume of 0.1N KMnO4 equivalents to oxalic acid reacted by the entire ore = 10 (V1 – V2) ml.
(Note: Ore solution after heating diluted to 10 time.)
1000ml of 1N KMnO4
= 43.47 gm of MnO2 (Eq. Wt)
∴ 10 (V1 – V2) ml of 0.1N KMnO4 (x) = 10 X [V1(
) – V2(
)] X 43.47 X 0.1
1000
(x) = _____________ gms of MnO2.
∴ Percentage of MnO2 in the given sample of Pyrolusite ore = x(
W(
) X 100
)
= ___________%.
RESULT: - Percentage of MnO2 in given sample (Pyrolusite) = ____________%.
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/ pipette while taking
readings.
6. Ensure that there are no air bubbles in the burette.
7. For each titration, use same number of drops of indicator.
8. Read lower meniscus in case of colourless solution.
9. Do not rinse the conical flask.
Date:
Signature of teacher/technician:
Viva:
1. Why do we use back titration?
2. 1 ml of 1N KMnO4 is equal to how much amount of MnO2?
3. What is the molecular weight of MnO2?
4. In this experiment do-not required any indicator, why?
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
2019
TITRATION CURVE THROUGH pH METER
Experiment No: - 7
Date:- …………….
AIM:-Determination of pK values of phosphoric acid by titration curve for phosphoric acid and
sodium hydroxide using pH Meter.
APPARATUS REQUIRED: pH meter, Burette, Pipette, Beaker, Measuring cylinder.
REAGENTS REQUIRED: H3PO4, NaOH , Buffer Solution (pH- 4, 7, 10).
PRINCIPLE:A pH meter is an instrument which is used for the direct measurement of pH of unknown solution. It
is a solid state device employing a high resistance field effect transistor or an operational amplifier.
The titration of a triprotic acid with a strong base will have three equivalence points one for each of
the three neutralization points:
In practice, HPO42- is too weak an acid for feasible titration in aqueous solution because its dissociation
constant is so small and close to ionic product of water, that the pH change at the equivalence point
can scarcely be distinguished from the titration of pure water. In general for the titration curve of a
polyprotonic acid to show separate well defined equivalence point breaks, the individual Ka values all
must be larger than 10-9 and differ from one another by at least a factor of 10-3 .
Department of Chemistry, SMIT
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In the given fig. the rate of pH change is smallest where the Vol. of NaOH added equals a/2 and a +
[(b-a)/2]. These points are halfway to the 1st and 2nd equivalence point respectively where the buffer
capacity of the titrand solution is greatest. The rate of change of pH is greatest at the equivalence points
a and b, where the buffer capacity is small.
Thus the slope of the pH vs. added base curve is smallest at the titration points halfway
to equivalence, where the buffer capacity is at maximum and largest at the equivalence points where
the buffer capacity is minimum.
The point where [H2PO4-]/[H3PO4] = 1 occurs halfway to the first equivalence point, when
half of the NaOH moles needed to complete the neutralization of the first acidic proton have been
added into the solution.
At this point, we have for Ka1
Thus, a measurement of pH at ½ eq. point from graph gives Ka1 directly.
Similarly, a measurement of pH half-way between the first and second equivalence point
gives Ka2 directly.
PROCEDURE:
Stage 1: Calibration of pH instrument
1. Connect pH electrode to the input B.N.C. Socket on the Rear Panel.
2. Before operating the instrument be sure that the function switch should always be kept at
stand by position.
3. Switch on the instrument and allow at least 15 minutes for warm-up.
4. Clean the electrode with distilled water and soak/dry it using tissue paper.
5. Dip the electrode in standard Buffer solution of 7.00 pH value.
6. Measure the temperature of solution and place the temperature knob accordingly.
Department of Chemistry, SMIT
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7. Bring the function switch to pH mode.
8. Adjust the “Calibrate” control so that display reads 7.00.
9. Now again turn the function switch into standby mode.
10. Remove the electrode from 7 pH buffer solution and wash it with distilled water.
11. Put the electrode in 4 pH buffer solution.
12. Bring the function switch in pH mode and adjust the “slope Z” control (Right side of the
instrument) so that the display reads 4.00.
13. Remove the electrode from 4 pH buffer solution and wash it with distilled water.
14. Always keep the function switch at stand by mode after measuring the pH value.
15. Repeat steps 5 to 13.
Stage 2: pH Measurements
1.
2.
3.
4.
5.
Calibrate the instrument as describe above.
Set the temperature control to the temperature of the test solution.
Dip the electrode in the solution under test.
Allow the reading to stabilize.
The display shows the pH value of the solution.
Stage 3: Preparation of buffer solution
(I) Take the buffer capsule of different pH values (like pH 4, 7 etc.) and mix with distilled
Water.
(II) Determine the pH values of prepared buffer solutions with the help of 1st
and 2nd stage instruction.
Stage 4: Phosphoric acid-sodium hydroxide titration
1. Prepare 0.1 M phosphoric acid and 0.1M NaOH solution by dissolving 9.8 ml of phosphoric
acid and in 1000 ml of distilled water and 4gm of NaOH in 1000ml of distilled water.
2. Take phosphoric acid solution in beaker and after dipping in pH electrode, measure pH and
Potential (in mV).
3. From the burette keep on adding 5ml of 0.1N NaOH solution in the beaker. At each
successive addition, measure pH and Potential.
4. From these observations, draw a graph between pH or Potential Vs. volume of NaOH.
5. From the graph, find out the Ka1 and Ka2 values and record in the result.
OBSERVATION:
Department of Chemistry, SMIT
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Sl. No.
Volume of NaOH added (ml.)
2019
pH
GRAPH:
Draw a graph between pH and Volume of NaOH added as shown in principle.
RESULT:
The pKa1 and pKa2 of phosphoric acid were experimentally found to be _______.
Date:
Signature of teacher/technician:
Viva:
1. Define pH.
2. What does pH of a solution signify?
3. What is pH of a solution if it is acidic?
4. What is the effect of dilution on pH of (i) an acidic solution (ii) a basic solution?
5. Give an example of each of tribasic acid and dibasic acid.
Department of Chemistry, SMIT
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CONDUCTOMETRIC TITRATION
Experiment No: - 8
Date:- …………….
AIM:- Determination of strength and weight of HCl in given HCl solution by titrating against N/10
standard Sodium Hydroxide solution, Conductometrically.
APPARATUS REQUIRED: Conductometer, Beaker, Burette, Measuring Cylinder, Pipette.
REAGENT REQUIRED:
HCl solution, N/10 NaOH solution.
PRINCIPLE:
Electrolyte conductivity is a measure of the ability of a solution to carry electric current solutions
of electrolytes conduct an electric current by the migration of ions under the influence of an electric field.
They obey Ohm’s Law. According to this, current I flowing through a conductor is directly proportional
to the resistance R of the conductor ,i.e.,
𝐸
𝐼=𝑅
Where, I= Current, E= Potential difference, R= Resistance.
Resistance of the conductor depends upon :
𝑅 ∝ 𝑙 … … … … . (𝑖)
1
𝑅 ∝ 𝐴 … … … … (𝑖𝑖)
R ∝ l/A
( is length)
( A is area of cross section)
𝑹=𝝆
𝒍
𝑨
Where, ρ is the specific resistance depends upon nature of the conductor.
If,
𝑙 = 1, 𝐴 = 1
𝝆=𝑹
Inverse of resistivity is known as conductivity or specific conductance.
1
𝑙
1
k=𝜌=𝐴×𝑅
𝑙
k = 𝐴 × conductance
k = Cell constant × C
If,
𝑙 = 1 cm, A = 1 cm2 then k = C
∴ Specific conductance is the conductance of one cm cube of the material.
Department of Chemistry, SMIT
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It means we measure the conductivity of the solution (due to moving ions) by conductometer.
Thus, the titration of strong acid like HCl with strong base NaOH, HCl is taken in conductivity
vessel and NaOH is drop wise added from the burette into the solution.
The conductance of HCl is due to the conductance of H+ ions and Cl- ions. As NaOH is gradually
added, H+ ions by combining with OH- will form unionized water. Whereas, slow moving Na+ ion
concentration will increase.
Therefore, on adding more of NaOH, the conductance will go on decreasing until whole of acid
has been neutralized by the base. Further, addition of NaOH will increase the conductance value.
Therefore, OH- ions are as fast moving ions. Minimum value of the conductance in the graph will
correspond to the point or equivalence point. a graph of the following type will be obtained on plotting
conductance versus vol. of NaOH added. Point of intersection will give value required for
neutralization.
PROCEDURE :
1. Take 20ml of the HCl solution in a 100ml beaker.
2. Calibrate the instrument after switch on the instrument with the help of
𝑁
calibrating solution or by preparing 10 KCl solution (Standard).
3. Dip the cell into the beaker containing HCl solution.
4. Note the conductance of this solution.
𝑁
5. Fill the burette with 10 NaOH and add drop wise into the beaker with equal
distance (interval).
6. Note the conductance for each interval and plot a graph b/w conductance v/s vol.
of NaOH added.
OBSERVATION TABLE :
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Sl.No.
Vol. of NaOH added
(ml)
1.
0
2.
2
3.
4
4.
6
5.
8
6.
10
7.
12
8.
14
9.
16
10.
18
11.
20
12.
22
13.
24
14.
26
15.
28
16.
30
17.
32
18.
34
19.
36
20.
38
Department of Chemistry, SMIT
Conductometer Reading
(S/cm)
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CALCULATION :
According to normality equation,
N1V1 = N2V2
(HCl)
(NaOH)
𝑁1 𝑉2 =
1
10
1
N x 𝑉2 (V2 obtained from graph and 𝑁2 = 10 N )
1
Strenght of HCl (𝑁1 ) = (10 x
𝑉2
) N or
𝑉2 (
20
)
200
N
= _______ N
∴ Weight of HCl per liter = Normality x Eqt. Wt.
=
N1(_______) X 36.5 = ______ g / Lit.
RESULT : The Weight of given HCl acid is __ __ __ g / Lit and strength of HCl is _______N.
Precaution:
1. NaOH is strong base and is very corrosive to your eyes.
2. Handle all reagents very carefully and avoid contact with skin.
3. Handle the electrode very carefully and avoid bumping them against the beaker or stirrer.
Date:
Signature of teacher/technician:
Viva:
1.
2.
3.
4.
5.
What is normality?
How can prepare 0.1 N solution of NaOH in 250ml.
What is conductivity?
What is the unit of conductance?
Conductivity will increase or decrease on dilution?
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
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DETERMINATION OF KMnO4 CONCENTRATION BY UVVISIBLE SPECTROPHOTOMETER
Experiment No: 9
Date: - …………….
AIM:- Determination of concentration of KMnO4 solution by Spectrophotometer.
APPARATUS REQUIRED:
U.V- visible Spectrophotometer, Beaker (100ml), distilled water.
REAGENT REQUIRED:
Solution of different concentration of KMnO4 (0.2%, 0.5%, 1.0%, 2.0%, 2.5% and 3.0% etc.).
PRINCIPLE:
The study of variation in intensity of a given coloured solution with the change in
concentration of the given coloured component is termed as calorimetric analysis.
A spectrophotometer is a device which detects the percentage transmittance of light radiation,
when light of certain intensity and frequency range is passed through the sample. Thus, the instrument
compares the intensity of the transmitted light with that of incident light.
Spectroscopy is the branch of science which deals with the interaction of matter with
electromagnetic radiations. Electromagnetic radiations consist of waves of energy. It covers a wide
range of wavelength or energies and visible light is part of electromagnetic radiation. When
monochromatic light falls upon a homogeneous medium, a portion of the incident light is reflected, a
portion is absorbed with in the medium and the remainder is transmitted.
Io = Ia + It + Ir
Io = Intensity of incident light
Ia = Intensity of absorbed light
It = Intensity of transmitted light
Ir = Intensity of reflected light
Where,
In case of aqueous solution, Ir is negligible as compared to Io and It
∴
Io = Ia + It
According to Beer-Lambert’s law the intensity of the incident light is proportional to the length
of thickness of the absorbing medium and the concentration of the solution.
It = Io 10-εcl
Department of Chemistry, SMIT
………………………. (1)
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Where, c is concentration of the 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,
A = log Io/ It
From eqn. (1)
A = εcl (Absorbance)
…………………….. (2)
Transmittance T of a solution is the ratio of It/ Io .
∴ T = It / Io
=≫ - log T = log It / Io = A
=≫
A= -log T = εcl
Thus, if a graph is plotted b/w A and C, we get a straight line for solution obeying BeerLambert’s Law. This is known as Calibration Curve.
This calibration curve is then used for measuring the concentration of unknown solution.
PROCEDURE:
Initial setting of Spectrophotometer:
1. Switch on the equipment and ensure that light glows on the screen.
2. Adjust the wavelength knob to the required wavelength region on scale.
3. Adjust the “set zero” knob so that meter reads zero on T scale and 100 on O.D.Scale.
Final setting of Instrument:
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4. Open the lid of the sample housing and insert a cuvette containing blank solution (distilled
water). Close the lid so that it fits properly.
5. Adjust the control knob (set 100) in appropriate direction to bring 100% transmittance or zero
optical density.
6. Open the lid and remove the cuvette. Close the lid tightly again.
7. Check zero on the meter. Adjust zero if disturbed.
8. Repeat 3 and 7 till zero and 100% transmittance.
9. Make the standard solution of KMnO4 with 0.5%, 1.0%, 1.5%, 2.0%, 2.5% and 3.0%
concentration respectively (20ml each).
10. Then measure the absorbance values of all the prepared solutions.
11. Now take the solution of unknown conc. of KMnO4 and find out the optical density. Find out
the concentration of the unknown solution from the graph corresponding to the optical
density of the solution.
OVSERVATION :
Concentration of KMnO4 Soln
Absorbance
0.5
1.0
1.5
2.0
2.5
3.0
Unknown
RESULT :
The concentration of KMnO4 solution = ………………….. g/lit.
Precaution:
1. Cuvette is fragile and it should be used with great care.
2. Cuvette should be properly cleaned before and after use and should be wiped with tissue
paper gently.
Date:
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CONCENTRATION DETERMINATION BY FLAME PHOTOMETER
Experiment No: - 10
Date:- …………….
AIM:- To determine the amount of Sodium and Potassium in a given sample by flame photometer.
APPARATUS REQUIRED:
Flame Photometer, Electric balance, beakers, measuring cylinders, funnel, measuring flask, etc.
REAGENT REQUIRED:
NaCl and KCl
PRINCIPLE:
Flame Photometry is based on the measurement of the light emitted when a metal is introduced into
a flame. It is also known as flame emission spectroscopy because flame is used to provide the energy
of excitation to atoms introduced into the flame. It is simple, rapid and reliable method for the routine
analysis of the elements like (Li, Na, K, Ca, Mg, etc.) which have an easily excited flame spectrum of
sufficient intensity for detection by photocell. The measurement of these elements is very useful in
medicine, agriculture and plant science.
In flame photometry, the following sequence of events takes place:
(i)
Aspiration of liquid sample (containing element) into a flame,
(ii)
Formation and evaporation of liquid droplets resulting in the formation of residue,
(iii)
Decomposition of residue into neutral atoms,
(iv)
Excitation of atoms and emissions of radiation from excited atoms,
(v)
Measurement of wavelength and intensity of emitted radiation by flame photometer.
Instrumentation of flame photometer
A block diagram of a flame photometer is given below:
Fuel
Oxidan
t
Flame
Excitation
Unit
____
Sampl
e
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Spectral
Insulator
(Optical
Filter)
___
Light
Sensitive
Detector
(Photocell)
___
___
Amplifier
Digital
Output
Block Diagram of flame photometer
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PROCEDURE :
1. Prepare stock solution of Na and K as follows:
(a) A stock solution of sodium (1000 ppm) is made by dissolving 2.542 gm pure and dry NaCl in
one litre of distilled water.
(b) A stock solution of potassium (1000ppm) is made by dissolving 1.909 gm of pure and dry
KCl in one litre of distilled water.
2. With the help of stock solution, prepared in step (1) make four standard solutions of 10, 5, 2.5
and 1ppm of sodium and potassium respectively.
3. Pass air into atomizer, the suction produced draws solution of the sample into the atomizer and
mixes with air steams. Supply fuel gas under pressure. Burn the mixture in the burner. The
radiations resulting from the flame pass through a spectral isolator which permits only the
radiation characteristics of the element under investigation to pass through the photocell.
4. Measure the output from the digital scale.
5. Repeat the same process (step 3 and 4) with all the standard solutions and unknown solutions
of sodium (at 589 nm) and potassium (at 766 nm).
6. Plot a calibration curve (one each for Na and K) between the optical readings (along y-axis)
and concentration in ppm (along x-axis).
7. With the help of calibration curve, find the concentration of unknown samples of Na and K.
OBSERVATIONS:
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(i)
2019
Determination of Sodium
(Absorption intensity measurements are carried out at 589 nm)
(ii)
Sl.NO
Concentration (ppm)
1.
2.
3.
4.
5.
10.0
5.0
2.5
1.0
Unknown
Absorption intensity of
sodium
A
B
C
D
E
Determination of Potassium
(Absorption intensity measurements are carried out at 766 nm)
Sl.NO
Concentration (ppm)
1.
2.
3.
4.
5.
10.0
5.0
2.5
1.0
Unknown
Absorption intensity of
sodium
A1
B1
C1
D1
E1
RESULT:
1. The concentration of unknown solution of sodium was found to be __________ppm.
2. The concentration of unknown solution of potassium was found to be __________ppm.
Precaution:
1. Handle the reagents carefully and avoid contact with skin.
Date:
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Signature of teacher/technician:
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POTENTIOMETRIC TITRATION
Experiment No. 11
Date:………………
AIM:To determine the amount of Iron in the given solution by Potentiometric titration.
APPARATUS REQUIRED: Pipette, Beaker, and Potentiometer.
CHEMICALS REQUIRED: Ferrous ammonium sulphate, Dil.H2SO4, K2Cr2 07 (0.1N)
PRINCIPAL :
For any redox reaction
a A + b B → c C + d D,
The potential is given by Nernst equation
Ecell = Eocell + 2.303 RT log (C)c (D )d
nF
(A )a (B)b
Where Eocell is the standard potential of the cell. The potential of the system is thus controlled by the
ratio of the concentration of the oxidized to the reduced species present. As the reaction proceed the
ratio and the potential changes more rapidly in the vicinity of the end point of titration.
This is thus followed potentiometrically.
In the determination of Fe2+ by potentiometric titration, the reaction that takes place is
Fe2+ → Fe3+ eCr2 072-. + 14H ++ 6 e -→2 Cr 3+ + 7H2O
Before the equivalence point the ratio Fe3+ / Fe2 +determines the potential
Ecell = EFe2+ = E 0Fe2+ 0.0591
1
3+
log {Fe }
2+
{Fe }
=0.75 V+0.0591 log { Fe3+}
{ Fe2+}
The potential of the solution will be around 0.75 V because the contribution to the potential by the
second term is negligible.
At the equivalence point, the potential is determined by both E 0Fe2 + and E 0Cr2O72- and
is given by
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SYNTHESIS OF IODOFORM
Experiment No: - 12
Date: - …………….
AIM: - Synthesis of Iodoform.
APPARATUS REQUIRED: Beaker, Measuring cylinder, Water bath, Oven, Petri dish.
REAGENT REQUIRED: Acetone, Methanol, 10%NaOH, Iodine solution.
Laboratory Preparation:
It is prepared in the laboratory by the action of iodine on ethyl alcohol or acetone,in the presence of
alkali.This is called haloform reaction.
Preparation of Iodoform from acetone:
The reaction taking place are similar to the above reaction, finally giving rise to iodoform
CH3COCH3 + 3I2 + 4NaoH
CHI3 + 3NaI + CH3COONa + 3H20.
Propanone
Procedure:
A)Synthesis:
To a soluiton of acetone (3ml),water(30ml),10% sodium hydroxide(15ml) is added.Then the supplied
iodine solution is added dropwise with constant shaking till the colour of the iodine persists.The
mixture is heated in water bath at 60⁰C.More iodine is added if the colour disappears while heating.The
mixture is heated till the precipitate settles down.The yellow precipiate of iodoform is filtered and
crystallized from methanol.The yield is calculated.
From acetone
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B)Melting point:
After calculating the yield the small amount of prepared iodoform is put in a melting point capillary
nad with the help of melting point apparatus the melting point is calculated.
Properties of Iodoform:
Physical properties:
It is yellow colored crystalline solid with melting point 392 K.
It has characteristic unpleasant odor.
It is insoluble in water but readily dissolves in ethyl alcohol and ether.
Due to the liberation of free iodine it has an antiseptic action.
Chemical properties:
Stability
On heating iodoform decomposes to give iodine vapour.This reaction is accelerated by moisture air
or light.
Reduction
Iodoform can be reduced with P and HI to give methylene iodide.
CHI3 + 2H
CH2I2
Methylene Iodide
+ HI
Hydrolysis
On boiling with aqueous or alcholic KOH, iodoform gives potassium formate.
CHI3 + 3KOH
HCOOH + 3KI + H2O
HEAT
HCOOH + KOH
HCOOK
Potassium Formate
+ H2O
Carbylamine reaction
When iodoform is warmed with primary amine and alcholic KOH, it forms isocyanide or
carbylamine, which has very unpleaseant smell.
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With Silver powder
Iodofom, when heated with silver powder gives acetylene.
Uses:
As an antiseptic and this nature is due to iodine that it liberates.However, because of its unpleasant
smell, it has now been replaced by better antiseptics.
In the manufacture of pharmaceuticals.
Yield:
________ gm.
Precaution :
1. NaOH is strong base and is very corrosive to your eyes.
2. Handle all reagents very carefully and avoid contact with skin.
Date:
Viva:
1.
2.
3.
4.
5.
Signature of teacher/technician:
What is an important medical use for Iodofom?
In which alcohol Iodofom will not dissolve?
Tell two chemical properties of iodoform.
What will be the product if Iodofom is heated with silver powder?
What are starting chemicalsrequired to prepare Iodoform.
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DETERMINATION OF PERCENTAGE PURITY
OF AMMONIUM SULPHATE
Experiment No:-13
Date :-……….
AIM:-To determine the percentage purity of the given sample of ammonium sulphate using
approximately 1N NaOH solution and a standard 0.1 N H2SO4 solution.
APPARATUS REQUIRED: Burette,
Test tube, Beaker, Measuring cylinder.
Pipette,
Standard
flask,
conical
flask,
Funnel,
REAGENTS REQUIRED: NaOH(1N), Ammonium Sulphate, Methyl red, 0.1(N) H2SO4, Litmus
paper.
PRINCIPLE:
A known weight of given sample of ammonium sulphate is boiled with known excess of NaOH
solution when the following reaction takes place.
(NH4)2SO4 + 2NaOH
2NH3 + 2H2O + Na2SO4
Ammonia is expelled and no more ammonia escape with steam; the excess of NaOH left over is back
titrated with standard H2SO4 solution using methyl red indicator. From the above equation:2NaOH (excess) + H2SO4
Na2SO4+2 H2O
From the above equation:
Equivalent weight of (NH4) 2 SO4= Mol. Wt = 132 =66
2
2
The given NaOH solution is standardized with standard H2SO4 solution using methyl red indicator.
PROCEDURE: 1.Standardization of NaOH solution: Take the supplied H2SO4 solution and fill the burette ensuring that there is no air gap in the burette.
25ml of given 1N NaOH solution is pipette out into a 250ml standard flask and made up to the mark
with distilled water and shaken well for uniform concentration. 25ml of this solution is pipette out into
a conical flask and three drops of methyl red indicator is added. It is titrated against the std. 0.1N H2SO4
solution in the burette to red end point. Titration is repeated to get agreeing values.
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2. Estimation: About 0.8gms of the given sample of (NH4)2SO4 is weighed accurately into a conical flask. 25ml of
1N NaOH solution is pipette out into this and a glass funnel is placed in the neck of the flask and boiled
until all ammonia is expelled. (Moist red litmus paper held at the escaping steam does not turn blue or
moist mercurous nitrate paper does not turn black). Volume of the solution is maintained during boiling
by adding distilled water. The solution is cooled and transferred quantitatively into a 250ml standard
flask and made up to the mark with distilled water and shaken well for uniform concentration. 25ml of
this solution is pipette out into a conical flask .3 drops of methyl red indicator is added and titrated
against standard 0.1N H2SO4 solution in the burette to red end point. Titration is repeated to get
agreeing values.
OBSERVATION AND CALCULATION
Weight of (NH4)2SO4 W = ______________ gms.. . .
1. Standardization of NaOH solution: Solution taken in the burette = standard dil. H2SO4 solution.
Solution taken in the flask
= 25ml of NaOH solution.
Indicator used ……………..= 3 drops of methyl red.
End point …………………. = Change of colour from yellow to red.
NaOH X H2SO4
Trail No.
1
2
3
Final Burette Reading
Initial burette reading
Vol. of H2SO4(ml)
Agreeing values (V1) of H2SO4 = ____________ ml.
3. Estimation: (Back titration of excess of NaOH solution)
Solution in the burette = standard H2SO4 solution.
Solution in the flask = 25ml of NaoH solution.
Indicator used …….= 3 drops of methyl red.
End point …………= change of colour from yellow to red.
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(NH4)2 SO4 X H2SO4
Trail No.
Final Burette Reading
Initial burette reading
Vol. of H2SO4(ml)
2019
1
2
3
Agreeing values V2 of H2SO4=__________ml
Calculation:
Volume of 0.1N H2SO4 equivalent to NaOH reacted = V1 – V2 ml. = _______ ml
.
Strength of (NH4)2SO4 solution N = (V1-V2) X 0.1 = (
) X 0.1
25
25
=___________N
Weight of (NH4)2SO4 / 250 ml of the solution (x)
= N (NH4)2SO4 X 66 = (
) X 66
4
4
x = _____________gms.
.
. .
Percentage purity of (NH4) 2 SO4 sample
= x X100 = (
) X100
W
(
)
=______________ %
RESULT: Percentage purity of (NH4) 2 SO4 in the sample= ________%
Precaution:
1.
2.
3.
4.
5.
Clamp burette vertically.
Before use, rinse the burette with given prepared solution.
Do not hold the pipette from the bulb.
Do not blow the last drop from the pipette.
Always keep your eye at the same level as the level of label of liquid in burette/pipette
while taking readings.
Date:
Viva:
1.
2.
3.
4.
5.
Signature of teacher/technician:
Which indicator is used in this experiment.
How can you confirm that (NH4)2SO4 is consumed in NaOH solution.
other than litmus which chemical can be used to test ammonia?
Give name of one external indicator.
Give name of one internal indicator.
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DETERMINATION OF PERCENTAGE PURITY OF ZINC
Experiment No: - 14
Date: ………
AIM :-To determine the percentage purity of the given sample of Zinc using (0.02M) EDTA solution.
APPARATUS REQUIRED: Burette, Pipette,
Test tube, Beaker, Glass rod, Measuring cylinder.
Standard
flask,
conical
flask,
Funnel,
REAGENTS REQUIRED: EDTA, Zinc pieces, ZnSO4, Buffer solution, EBT, Conc.HCl, NH4OH
solution.
PRINCIPLE : The disodium salt of Ethylene – diamine tetra – acetic acid (EDTA – Na2H2Y) reacts with the bivalent
Zn++ ions as follows:
Zn+++ Na2H2Y
Na2Zn Y+ 2H+
When all the Zn++ ions react, Eriochrome black- T indicator added at a pH of 10, change the colour
from wine red to blue.
A known weight of the given sample of Zinc is dissolved in dilute hydrochloric acid. The solution is
transferred quantitatively into a standard flask after neutralization. A known volume of this solution is
titrated against EDTA solution using Eriochrome black-T indicator. A buffer solution of (NH4Cl –
NH4OH) is added to maintain pH value at 10. From the titer value molarity of Zn solution and hence
the percentage purity of Zn is calculated. The given standard solution of ZnSO4 is used for the
standardization of EDTA solution.
PROCEDURE:1. Standardization of EDTA solution:Take the supplied EDTA solution and fill the burette ensuring that there is no air gap in the burette.
25ml of ZnSO4 solution is pipetted into a conical flask. 2ml of NH4OH buffer solution and 3-4 drops
of Eriochrome Black –T indicator are added. The solution turns wine red in colour. This is titrated
against EDTA solution taken in the burette till the solution turns perfect blue colour. This is the end
point. Titration is repeated to get agreeing values.
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2. Estimation:
About 0.3 gm of the given sample of Zinc is accurately weighed and transferred to a 250ml beaker.
Conc. Hydrochloric acid is added drop wise till all the Zinc pieces are completely dissolved. Dilute it
with 4t.t. of water. The solution is just neutralized with NH4OHsolution till white precipitate is
obtained. Then the precipitate is dissolved in minimum amount of dil HCl. The solution is then
transferred to a 250ml standard flask quantitatively using a funnel & a glass rod. The beaker and glass
rod are washed several times with distilled water and washing are also transferred to the same and
shaken well for uniform concentration.
25ml of this solution is pipetted put into a conical flask. 2ml of NH4Cl-NH4OH buffer solution and 34 drops of Eriochrome Black- T indicator are also added. The solution turns wine red colour. This
titration against EDTA solution turns prefect blue colour. This marks end point. Titration is repeated
to get agreeing values.
OBSERVATION AND CALCULATIONS
I.
Standardisation of EDTA solution:
Solution in the burette
= EDTA solution.
Solution taken in the flask
= 25ml of made up ZnSO4 solution +2ml of NH4ClNH4OH buffer solution.
Indicator used ………………..
=3-4 drops of (Eriochrome Black-T).
End point…………………….. =Change of colour from wine red to perfect blue.
ZnSO4 X EDTA
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of EDTA(ml)
Agreeing value= V1 of EDTA=__________ml.
.
. .
Strength of EDTA solution(N1) = 25
V1 (
N1 (EDTA) solution
X
0.02
) of EDTA
=_____________N.
2. Estimation of Zinc:
Weight of Zinc taken W =__________________gm.
Solution in the burette
= EDTA solution.
Solution taken in the flask
= 25ml of made up Zn solution +2ml of NH4ClNH4OH buffer solution.
Indicator…………………..
= 3-4 drops of Eriochrome Black-T
End point………………… .
= Change of the colour from wine red to Purple blue.
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Zn solution X E.D.T.A
Trial No.
1
2
3
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of EDTA (ml)
Agreeing value= V2 of EDTA =________ml.
Strength of Zinc solution (N2) = V2 (
) of EDTA X N1(
25
N2=_______M.
) of EDTA
) of Zn solun. X 65.37
4
x =____________ gms.
Weight of Zinc/ 250ml of the Zinc solution x = N2 (
.
. .
Percentage purity of Zinc
= x X 100 = (
) X 100
W
(
)
=________________%
RESULT: Percentage purity of the given sample of the Zinc=________%
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/pipette
while taking readings.
Date:
Signature of teacher/technician:
Viva:
1. Why is EBT indicator used for EDTA complexometric titration in determining water hardness?
2. Why do we add buffer solution in EDTA titration?
3. Which buffer is used in EDTA titration?
4. Why are metal EDTA complexes highly stable?
5. Why ammonia buffer is used in EDTA titration?
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DETERMINATION OF AMOUNT OF SODIUM CARBONATE AND
SODIUM HYDROXIDE IN A MIXTURE
EXPERIMENT NO: 15
DATE: ………………
AIM: To Determine the amount of Na2CO3 and NaOH presence in a given test solution by indicator
method.
APPARATUS REQUIRED: Burette, Pipette, Standard flask, conical flask, Funnel, Test tube.
REAGENTS REQUIRED: Sodium carbonate, Methyl orange, Phenolphthalein, 1.0(N)HCl Solution.
PRINCIPLE:
The amount of NaOH and Na2CO3 can be estimated quantitatively from a mixture by using a strong
acid like HCl. As it is a strong acid Vs strong base titration as well as strong acid Vs weak base titration
process, so we get two end point value separately, using appropriate indicator.
If the mixture containing phenolphthalein as a indicator is titrated against 0.1 (N) HCl, the pink colour
of the indicator gets discharged when whole NaOH and half of Na2CO3 is neutralised.
NaOH + HCl
NaCl + H20
Na2CO3 + HCl
NaCl + NaHCO3
Na2CO3
CO32- + H2O
HCO3- + H20
2Na+ + CO32HCO3- + OHH2CO3 + OH-
So, in the above neutralization solution 2-3 drops of methyl orange is added and further HCl is added
quantitatively till the red colour is obtained. Thus, quantity of acid added will correspond to the
quantity required to neutralize the remaining part or half of the carbonate.
HCl + NaHCO3
NaCl + H2O + CO2
The amount of acid required for equation must be same because one molecule of NaHCO3 has been
produced from one molecule of Na2CO3.
PROCEDURE:
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Preparation of Standard solution:
About 1.325gm of given Sodium carbonate is accurately weighed and transferred to a 250 ml
standard flask. It is first dissolved in a little distilled water and then made up to the mark with
distilled water.
1. Standardization of Sodium bi carbonate Solution:
Fill the burette with supplied HCl solution ensuring that there is no air gap in the burette.
25 ml of prepared Sodium carbonate is pipette out in a conical flask and to it 2-3 drops of methyl red
indicator is added and titrate it till the colour changes from yellow to red. Titration is repeated to get
the agreeing values.
2. Estimation:
25 ml of given mixture is taken in a 250 ml standard flask and to it 2-3 drops of phenolphthalein
indicator is added. It is titrated against HCl till the colour changes from pink to colourless. When the
solution becomes colourless burette reading is noted and to it 1-2 drops of methyl orange is added
and continue the titration until the solution turns red. Titration is repeated to get the agreeing value.
OBSERVATION AND CALCULATION:
Weight of Sodium carbonate (W) =_________gms
Strength of the solution prepared (N1)
= (W) X 4
53
(N1) of Sodium carbonate =_________N
Standardization of HCl solution
Solution taken in the burette
Solution taken in a conical flask
Indicator used
End point
= given HCl solution
= 25 ml of Sodium carbonate (V1) + Methyl Orange
= Methyl Orange
= from yellow to red.
Sodium carbonate vs HCl
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Trial No.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of HCl(V2) (ml)
1.
2.
3.
Agreeing value (V2) = __________ml
(N2) = 2N1V1 = 2(
V2
N2 of HCl solution = ________N
)x (
(
)__
)
Estimation:
Solution taken in the burette
= given HCl solution
Solution taken in a conical flask = 25 ml of given mixture (V4) + Phenolphthalein
Indicator used
= (Phenolphthalein)
st
1 end point
= from pink to colourless.
nd
2 end point
= from colourless to red after adding methyl orange
HCl solution vs mixture of sodium carbonate and sodium hydroxide:
Sl.
no
Burette reading
With
phenolphthalein
With
methyl
orange
Volume of HCl
used in titration
of NaOH +half
of Na2CO3 (Vc )
Volume of
HCl used in
titration of
NaOH +
= Va
Na2CO3 (Vd)
= Vb
Volume of
HCl used in
titration of
HCO3(Ve) = Vd Vc
Volume of
HCl used in
titration of
Na2CO3
(V3) = 2Ve
Volume of
HCl used in
titration of
NaOH (Vf) =
Vd - 2 V e
1
2
3
Agreeing value with phenolphthalein (Va) =
Agreeing value with methyl orange
(Vb) =
CALCULATION
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Normality of sodium carbonate in the solution (N3) = N2 (
) x V3(
2 x V4(
)
)
(N3) = -------N
Strength of sodium carbonate present in the solution = N3 (
) x 53
=_______g/lit
Normality of sodium hydroxide in the solution (N4) = N2 (
) x Vf(
V4
)
=________N
Strength of sodium hydroxide present in the solution= N4 X 40
=____________g/lit
RESULT:
In the given test solution Normality of Na2CO3 is.______N and NaOH is______N.
Strength of Na2CO3 in the given test solution______g/lit and Strength of NaOH is ______g/lit.
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/pipette
while taking readings.
6. Ensure that there are no air bubbles in the burette.
Date:
Viva:
1.
2.
3.
4.
Signature of teacher/technician:
What are the commercial names of NaHCO3 and Na2CO3?
Which indicator was used in the experiment?
What are the equivalent weights of NaOH, NaHCO3, Na2CO3?
Draw the structure of phenolpthalein.
ESTIMATION OF OXALIC ACID SALT
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Experiment No: 16
Date: - ……………
AIM: - To estimate the weight of oxalic acid crystals dissolved in 250ml using approximately
decinormal potassium permanganate solution and pure Mohr’s salt (Ferrous ammonium sulphate)
crystal.
APPARATUS REQUIRED: Burette, Pipette, Standard flask, conical flask, Funnel, Test tube.
REAGENT REQUIRED: Mohr salt, Oxalic acid, Dil.H2SO4, Potassium permanganate.
PRINCIPLE:
Acidified KMnO4 oxidizes oxalic acid and ferrous ammonium sulphate as follows:
2KMnO4 + 3H2SO4
10FeSO4 + 5H2SO4 +5(O)
5H2C2O4 + 5(O)
K2SO4 + 2MnSO4 + 3H2O + 5(O)
5Fe2(SO4)3 + 5H2O.
10CO2 + 5H2O
10 molecules of ferrous sulphate = 5 molecules of oxalic acid .
= 5 atoms of (O) = 5 equivalents.
∴ Equivalent weight of FeSO4 (NH4)2SO4.6H2O = Its Mol. Wt = 392.
∴ Equivalent weight of H2C2O4.2H2O = Mol. Wt / 2 = 126/2 = 63
A standard solution of Mohr’s salt is prepared by accurately weighing the crystals. A definite volume
of this solution is acidified and titrated against KMnO4 is standardized.
A definite volume of oxalic acid solution is acidified and titrated against KMnO4 solution. Thus oxalic
acid solution is standardized. For both titrations pink coloured KMnO4 acts as self indicator.
PROCEDURE:
4. Preparation of Standard Solution :About 9.8 gms of Mohr’s salt crystals are accurately weighed into a 250ml standard flask. . The salt is
dissolved in a little (1t.t) dilute H2SO4. The solution is made upto the mark with distilled water and
shaken well.
5. Standardization of KMnO4 solution :Fill the burette with the supplied KMnO4 solution ensuring that there is no air gap in the burette. 25ml
of the prepared Mohr’s salt solution is pipette out into a conical flask and 1 test tube of dil. H2SO4 is
added and titrated against KMnO4 solution in the burette. Appearance of permanent pale pink colour
marks the end point. Titration is repeated to get agreeing values.
6. Estimation of Oxalic acid :-
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Engineering Chemistry Lab Manual
2019
25ml of the supplied oxalic acid solution is pipette out into a conical flask and 1 test tube of dilute
H2SO4 is added. The product is heated to 70o to 80o C and the hot solution is titrated against KMnO4
solution in the burette until a permanent pale pink colour is obtained.
OBSERVATION AND CALCULATIONS:
Weight Mohr’s salt taken (W)
∴ Strength of the solution prepared
(N1) of Mohr’s salt solution
= __________gms.
= [(W) X 4] / 392
= _______________ N.
2. Standardisation of KMnO4 solution:Solution taken in the burette
= given supplied KMnO4 solution.
Solution taken in the conical flask = 25ml of Mohr’s salt soln + 1t.t of dil. H2SO4
Indicator used …………………… = KMnO4 – self indicator.
End point …………………………..= from colourless to pale pink.
Mohr’s salt X KMnO4
Trial No.
1.
2.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
Agreeing value = V1of KMnO4 = ___________ml.
∴ Strength of the given KMnO4 solution (N2) = 25 X N1(
V1 (
3.
) of Mohr’s salt
) of KMnO4
N2 of KMnO4 = ___________________ N.
3. Estimation of Oxalic acid:Solution taken in the burette
= given supplied KMnO4 solution.
Solution taken in the conical flask = 25ml of Oxalic acid sol + 1t.t of dil. H2SO4.
(Heat upto 700 to 800 C)
Indicator used …………………… = KMnO4 (self-indicator).
End point …………………………= from colorless to pale pink.
Oxalic acid X KMnO4
Trial No.
1.
2.
3.
Final Burette Reading (ml)
Initial Burette Reading (ml)
Volume of KMnO4 (ml)
Agreeing value = V2of KMnO4 = ___________ml.
Department of Chemistry, SMIT
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Engineering Chemistry Lab Manual
2019
∴ Strength of Oxalic acid solution (N3) = V2(
)of KMnO4 x N2(
25
) of KMnO4
N3 of Oxalic acid = ___________________ N.
Weight of Oxalic acid present in the whole of
the solution given in the 250ml standard flask
= N3 (
) of Oxalic acid X 63
4
= _______________gms.
RESULT: Weight of Oxalic acid dissolved in whole solution = ____gms.
Precaution:
1. Clamp burette vertically.
2. Before use, rinse the burette with given prepared solution.
3. Do not hold the pipette from the bulb.
4. Do not blow the last drop from the pipette.
5. Always keep your eye at the same level as the level of label of liquid in burette/ pipette while taking
readings.
6. Ensure that there are no air bubbles in the burette.
7. For each titration, use same number of drops of indicator.
8. Read lower meniscus in case of colourless solution.
Date:
Signature of teacher/technician:
Viva:
1. What is acidimetry and alkalimetry ?
2. What indicator is used in the titration of oxalic acid with sodium hydroxide ? Which solution
is taken in the burette and what is the end point ?
3. What is basicity of an acid ?
4. Which is an oxidising agent and a reducing agent in the reaction between KMnO4 and FeSO4?
5. What is the indicator used in KMnO4 titration ?
6. Why are a few drops of dilute sulphuric acid added while preparing a standard solution
of Mohr’s salt ?
Department of Chemistry, SMIT
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