1
Engineering chemistry
Laboratory Manual
Gandhi Institute of Technology
GITAM UNIVERSITY
Visakhapatnam – 530 045
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INDEX
1. Calibration of Volumetric Apparatus.
2. Determination of Sodium Carbonate in soda ash.
3. Estimation of Iron as Ferrous Iron in an Ore Sample.
4. Estimation of Calcium in Portland Cement.
5. Estimation of Volume Strength of Hydrogen Peroxide.
6 a) Estimation of Active Chlorine Content in Bleaching Powder.
b) Determination of Hardness of a Ground Water Sample.
7. Determination of Chromium (VI) in Potassium Dichromate.
8. Determination of Copper in a Copper Ore.
9. a) Determination of Viscosity of a Liquid
b) Determination of Surface Tension of a Liquid.
10. a) Determination of Mohr’s Salt by Potentiometric Method.
b) Determination of Strength of an Acid by pH metric Method
***
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INTRODUCTION
Analytical chemistry plays an important role in recognizing different
substances and determining their constituents. The chemical analysis is mainly
divided into
(1) Qualitative analysis and (2) Quantitative analysis.
The aim of qualitative analysis is the detection and identification of the constituents of
a compound or a mixture of compounds or elements, where as the aim of quantitative
analysis is the determination of the percentage or molecular composition of a sample
(amount of constituents). Quantitative analysis is divided into (1) Gravimetric
analysis and (2) Volumetric or Titrametric analysis. Volumetric analysis has greater
advantages over the gravimetric analysis. The solution of accurately known strength is
called standard solution.
A standard solution is one which is prepared by dissolving an exactly weighed
(accurate to 0.1 mg.) quantity of a primary standard substance. (such as K2Cr2O7,
Na2CO3, Na2C2O4 etc.,) in a known volume of distilled water.
A primary standard substance should confirm to the following requirements.
i) It should be stable, non-hygroscopic and must be of fixed composition.
ii) It should not gain or lose weight.
iii) It should be non corrosive.
Solutions can also be prepared by dissolving an approximate weight of a secondary
standard substances (like mineral acid dilute solutions, hypo, sodium hydroxide,
Mohr’s salt, oxalic acid etc.,) in a known volume of distilled water. However,
strength of these solutions should be ascertained by titrating against a standard
solution referred above, before using.
The concept of equivalent weight changes according to the type reaction involved in
the volumetric titration. In an acid base titration, the equivalent weight is calculated
by dividing the molelcular weight by the number of replaceable H+ ions or OH ions .In a redox reaction, the equivalent weight is calculated by dividing the molecular
weight by the total no. of electrons gained or lost.For example in the estimation of
ferrous iron with standard dichromate solution,
Cr2O72- + 14H+ + 6e-  2Cr3+ + 7H2O, 6Fe2+  6Fe3+ + 6eThe equivalent weight of the oxidising agent (potassium dichromate) is 1/6th of its
molecular weight since the number of electrons gained by Cr(VI) are six.In the
titration of oxalic acid with potassium permanganate,
2KMnO4 + 3 H2SO4  K2SO4 + 2MnSO4 + 3H2O + 5 (O)
5H2C2O4 + 5(O)  10CO2 + 5H2O
the equivalent weight of permanganate is 158/5 = 31.6. (since no. of electrons
gained by Mn(VII) are five). The equivalent weight of oxalic acid is 126.07 / 2 =
63.035. (Since total no. of electrons lost by oxalic acid are two).A molar solution is
one which contains, one mole(Gram molecular weight) of the substance dissolved in
one liter of the solution.
In a redox reaction, oxidation is the process in which electrons are lost (or) the
process in which oxidation number increases by loosing electrons. The substance
which oxidises the other substance is known as oxidising agent or oxidant. An oxidant
always oxdises the reductant but itself undergoes reduction. Reduction is the process
in which oxidation number decreases by gain of electrons. The substance which
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reduces the other substances is called a reducing agent (or) reductant. A reductant
always reduces the oxidant and itself undergoes oxidation. The process in which both
oxidation and reduction occurs simultaneously is known as redox process or redox
reaction . In redox reaction electrons are transferred from reductant to oxidant.
A standard solution is the solution whose strength is exactly known.The
reagent of known concentration is called the titrant and the substance being titrated is
termed as the titrand / analyte . The process of adding the standard solution to the
unknown solution until the reaction is just completed is known as titration, and the
substance to be determined is titrated. The point at which this occurs is called the
equivalence point or stoichiometric end-point. Equivalence point/End–point is a stage
at which the amount of reagent added is exactly and stoichiometrically equivalent to
the amount of the reacting substance in the titrated solution. The end - point and
equivalence point may not be identical. End – point is usually detected only after
adding slight excess of the titrant.
The completion of reaction is usually judged by adding auxilary reagents which give a
colour change at the end-point of the titration. These reagents are known as
indicators. We can see the exact colour change at the end – point with the help of
white glazed tile.
There are four types of reactions in titrimetric analysis.
1) Neutralization reactions or acidimetry and alkalimetry : These reactions involve
the combination of H+ and OH- ions to form water. That means simple acid - base
reactions will come under this class.
HCl + NaOH  NaCl + H2O; H2C2O4 + 2NaOH  (COONa)2 + 2H2O
2) Complexometric reactions : These depend upon the combination of ions other than
H+ and OH- ions, to form a soluble complex.
2CN- + Ag+  [Ag(CN)2]- ; EDTA + [Ca / Mg ]+2  [Ca EDTA] or [Mg EDTA]
3) Precipitation reactions: These depend upon the combination of ions to form a
simple precipitation.In this no change in oxidation states among reacting substances
Eg. Ag NO3 + NaCl  AgCl  + NaNO3 .
4) Oxidation – Redox reactions : These reactions involve the change in oxidation
number or transfer of electrons among the reacting substances.
2KMnO4 + 3 H2SO4 + 5H2C2O4  K2SO4 + 2MnSO4 + 8H2O + 10CO2
A reaction involved in volumetric analysis must fulfil the following requirements.
i) The two solutions should react completely in stoichiometric proportions.
ii) The reaction should be simple and take place instantaneously.
iii) There must be a definite change in physical or chemical property of a solution.
iv) An indicator should be available for the reaction.
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1. Calibration of volumetric apparatus :
Apparatus :
In titrimetric analysis, while using the following apparatus, we should take some
precautions to minimize the practical errors.
1) Burette : Generally the burette has the capacity of holding 50 ml and graduated
1/10 of a milliliter and it will be graduated from the top to bottom and at the lower
end it is provided with stop cock to regulate the flow of solution.
a) The burette should always be kept in non – greasy condition. For this
condition, the burette is cleaned with chromic acid and then thoroughly washed
with tap water and then distilled water.Then the burette should be thoroughly
rinsed with little of solution which is to be filled in it and reject the rinsed
solution.
b) The burette is filled with the standard solution. The burette readings should be
noted with out parallax error and then it should be vertically clamped in a
burette stand.
c) While taking the readings it is better to hold a white paper behind the burette
readings and eye must be placed in the same line of the meniscus. Generally
the reading tangential to lower meniscus in the case of colourless or light
coloured solutions and upper meniscus incase of concentrated or deeply
coloured solutions, is taken as the burette reading.
2) Volumetric pipette : The pipette is a long narrow tube having cylindrical bulb in
the middle and tapping to a fine nozzle at its lower end and carrying a mark round
the glass tube above the bulb. The volume delivered between the upper mark and
lower tip is the volume being titrated. (Caution ! Don’t blow the last drop in the
nozzle of the pipette). With this we can transfer a definite volume of solution
from one flask to another. The pipette should be washed repeatedly with tap water
and then distilled water and then rinsed with a little of solution which is to be
transferred.
3) Graduated pipette : Another type of pipette used in the laboratory is the graduated
pipettes for transferring approximate volumes of reagents (such as acids, alkalis,
indicators etc.,) . These are graduated in 1/10 of a milliliter, for every milliliter.
These are available to measure 1.0 ml , 2.0 ml, 5.0 ml , 10.0 ml and 20.0 ml. An
advantage of these graduated pipettes is fraction of a ml can be measured and
transferred to the conical flask.
4) Measuring jar : This is cylindrical jar and provided with a round base at its one
end. It will be graduated from the bottom to top and has the capacity of holding
50 ml or 100 ml. and graduated in 1.0 ml units . This is only used for measuring
and transferring the approximate volumes of solutions (generally reaction medium
and other reagents), where accuracy is not important. This should be thoroughly
washed with chromic acid followed by tap water and distilled water. They are
also available in the following capacities 10 ml , 20 ml, 25 ml, 250 ml, 500ml and
1000 ml.
5) Conical flask or Erlenmayer flask : It is cone shaped flask and is used in titrations
as reaction flask. It is available in the volume range from 10 ml , 25 ml, 50 ml,
100 ml, 250 ml, 500 ml and 1000 ml. Generally 250 ml conical flasks are
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recommended for class work. While titration is on, the neck of the conical flask
should hold using the right hand. It helps to make swirling motion of the solution
convenient by rotating the flask.
6) Distilled water bottle / Wash bottle : It is readily available in the range 500 ml,
750 and 1000 ml. it is made of pyrex glass or corning glass or polyethene. The
importance of using wash bottles is to get fine streams of water to wash down the
adhering droplets of the titrant in to the bulk of the titrand in the conical flask.
7) White glazed tile : It is white ceramic tile and helps to see the exact colour change
at the end point of the reaction by providing a glossy white back ground.
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GENERAL INSTRUCTIONS
 Before coming to the laboratory, understand the theory behind the reaction you are
going to carry out.
 Keep the work - bench and sink neat and clean. Don’t allow filter paper, broken
bits of glass, sticks of matches etc., to lie on the table or in sink. Put these things
in the dustbin placed at the worktable.
 Apparatus should be in non-greasy condition.
 Keep the apparatus clean and properly arranged on the work - bench.
 If any piece of apparatus is broken, report at once to the staff members / Lab
Assistant.
 Once you have transferred a reagent from a reagent bottle, never pour it back even
if there is some excess.
 Arrange reagent bottles in their proper places after use and see that they are
properly stoppered.
 Handle reagent bottles / chemicals carefully.
 Use only minimum possible quantity of chemicals / Reagents for any reaction.
 Close the water tap immediately after use; do not waste water.
 Precautions should be taken to avoid fire accidents.
 After the class, before you leave the laboratory, wash the apparatus clean, wipe the
table and keep the apparatus in proper place.
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2. Estimation of Sodium carbonate in Soda Ash.
Aim : Determination of the amount of sodium carbonate present in commercial soda
ash by using Standard Hydrochloric acid.
Theory : Sodium carbonate is a weak base has the formula Na2 CO3 and Hydrochloric
acid is a strong acid with the formula HCl. Sodium carbonate can be estimated
with a standard hydrochloric acid solution using acid-base indicator. These
two reacts as follows:
Na2CO3 + 2HCl  2NaCl + H2O + CO2 
In this reaction salt and water are formed, which is an example of
neutralization reaction. All the reactants and products are colourless, so Methyl
Orange indicator is used to locate the end point of the reaction. The colour change of
methyl–orange is from yellow in alkaline medium to orange–red in acidic medium
(pH range 3.3 to 4.3). Sodium carbonate is supposed to decompose in aqueous
solution according to the following equation.
Na2CO3 + 2H2O  2NaOH + H2CO3
Methyl orange is not affected by the very weak carbonic acid formed in
solution.According to the above reaction equivalent weight of sodium carbonate is
equal to half of its molecular weight.
Procedure :
The given soda ash solution is made upto the mark of volumetric flask with
distilled water carefully. The flask is stoppered and shaken thoroughly about 3 to 5
minutes for complete homogenization.10.0 ml of the above solution is pipetted out
into a clean conical flask carefully and 50.0 ml of distilled water is added with
measuring jar. Two drops of methyl orange indicator is added directly to the of
contents of conical flask. The conical flask contents are titrated with standard
hydrochloric acid solution after noting the initial reading. The titration is continued
until the colour changes from yellow to pale red. The final reading of burette is noted.
A number of titrations are repeated for getting concurrent results. The results are
tabulated in Table No. I.
Table – I
Titration of Soda ash solution with standard hydrochloric acid solution.
Normality of standard hydrochloric acid solution ______________ N
Indicator : Methyl orange
Colour change at the end point : Yellow to pale red.
Volume of Soda ash
S.No.
solution taken in ml. (V1)
Burette readings
Initial
Final
Volume of hydrochloric
acid solution consumed in
(V2) ml.
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Calculations :
By the law of equivalence
V1N1 = V2N2
N 1 = Normality of Soda ash solution
V 1 = Volume of Soda ash solution
N 2 = Normality of Hydrochloric acid solution
V2 =Volume of Hydrochloric acid solution
 N1= V2N2 / V1
=
=
=
=
?
ml.
N
ml.
 The Normality of Soda ash solution _____________ N.
Equivalent weight of sodium carbonate =Molecular weight / 2=106/2=53
Amount of Sodium carbonate present in 1000 ml of the Soda ash solution
=
Normality of Soda carbonate ash solution X equivalent weight of Sodium carbonate =
gm.
Amount of Sodium carbonate present in the given 100 ml of solution =
Amount present in 1000 ml of solution / 10 =
gm
Table– II
Percentage error table
Roll No. /
Regd. No.
Flask
No.
Amount of sodium carbonate present in
100 ml of the soda ash solution. Grams.
Reported
Given
Percentag
e error
Report:
Amount of sodium carbonate present in the given 100ml. of solution ----- gm.
***
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3. Estimation of Ferrous Iron (Fe+2) in an ore sample
Aim
: Estimation of ferrous iron in a given 100 ml of ore solution by titration
against a standard solution of potassium permanganate.
Theory
:
The Potassium permanganate is a powerful oxidizing agent has the formula
KMnO4 in which Manganese is in +7 oxidation state and Mohr’s salt with the
formula FeSO4. (NH4)2 SO4. 6H2O in which Iron (II) is acts as a reducing agent.
These two reacts in presence of sulphuric acid medium as follows. In this reaction
Mn+7 is reduced to Mn+2 and Fe+2 is oxidized to Fe+3.
2KMnO4 + 3 H2SO4  K2SO4 + 2MnSO4 + 3H2O + 5 (O)
10 FeSO4 + 5H2 SO4 + 5 (O)  5 Fe2 (SO4)3 + 5H2O
-------------------------------------------------------------------------------------2KMnO4 + 10 FeSO4 + 8 H2SO4  K2SO4 + 2MnSO4 + 5 Fe2 (SO4)3 + 8H2O
-------------------------------------------------------------------------------------According to this one mole of potassium permanganate is reacting with five
moles of Iron (II). In this all the reactants and products are colourless except
potassium permanganate. So the colour of the permanganate is used to locate the end
point. When once all the Fe(II) ions are completely reacted the excess trace amount
of potassium permanganate can give sufficient colour to the solution i.e. colour
change at the end point is colourless to pale pink. Potassium permanganate acts as a
self indicator, so it is called self indictor reaction.
Part – I :
Standardisation of potassium permanganate solution by using standard
solution of Mohr’s Salt.
Procedure : 10.0 ml of standard Mohr’s salt solution is pipetted out into a clean
conical flask carefully. 40.0 ml of distilled water and 5.0 ml of 1:1 dilute sulphuric
acid are added to the contents of the conical flask with a measuring jar. Then the
burette is filled with potassium permanganate solution after rinsing with the same
solution. Now the colourless conical flask contents are titrated with permanganate
solution after noting the initial reading of the burette, which coincides with the upper
meniscus. The titration is continued drop wise with constant shaking until a drop of
permanganate solution gives a colour change from colourless to pale pink, which is
the end point of the titration. The final reading of burette is noted with out parallax
error and results are tabulated in Table – I. A number of titrations are carried out until
2 or 3 concurrent readings are obtained.
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Table – I
Titration of potassium permanganate solution with a standard
Mohr’s salt solution.
Normality of standard solution of Mohr’s salt ____________ N
Indicator : Self Indicator (Colour of the Potassium permanganate)
Colour change at the end point : Colourless to pale pink.
S.No.
Volume of Mohr’s salt
solution taken in ml.
(V1)
Calculations :
By the law of equivalence
Burette readings
Initial
Final
Volume of potassium
permanganate solution
consumed in ml.( V2)
V1 N1 = V2 N2
N1 = Normality of Mohr’s salt solution
=
V1 = Volume of Mohr’s salt solution
=
N2 = Normality of potassium permanganate solution =
V2 = Volume of potassium permanganate solution =
 N2 = V1 N1 / V2
N
ml.
?
ml.
The Normality of potassium permanganate solution is_____________ N.
Part – II :
Estimation of Ferrous iron present in a given 100 ml of ore solution by
titrating against a standard solution of potassium permanganate.
Procedure :
The given iron ore solution is diluted up to the mark of volumetric flask with
distilled water carefully and the flask is stoppered tightly. Then the solution is shaken
throughly about 3 to 5 minutes for complete homogenisation. 10.0 ml of above Iron
ore solution is pipetted out into a clean conical flask carefully. 40.0 ml of distilled
water and 5.0 ml of 1:1 dilute sulphuric acid are added to conical flask contents with a
measuring jar. Then the burette is filled with standard potassium permanganate
solution. Now the conical flask contents are titrated with permanganate solution after
noting the initial reading of the burette, which coincides with the upper meniscus.
The titration is continued by drop wise with constant shaking until a drop of
permanganate solution gives a clear cut colour change from colourless to pale pink,
which is the end point of the titration. The final reading of burette is noted with out
parallax error and the results are tabulated in Table – II. A number of titrations are
carried out until 2 or 3 concurrent readings are obtained.
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Table – II
Titration of standard potassium permanganate solution with a unknown Iron
ore solution.
Normality of standard Potassium permanganate solution = ____________ N
Indicator : Self Indicator (Colour of the Potassium permanganate)
Colour change at the end point : Colourless to pale pink.
Volume of Iron ore solution
S.No.
taken in ml.( V3)
Calculations :
By the law of equivalence
N3
V3
N4
V4
Burette readings
Initial
Final
Volume of potassium
permanganate solution
consumed in ml.( V4)
V3 N 3 =V4 N4
= Normality of an Iron ore solution
= Volume of an Iron ore solution
= Normality of potassium permanganate solution
= Volume of potassium permanganate solution
=
=
=
=
?
ml.
N
ml.
 N3 = V 4 N 4 / V 3
The Normality of the given Iron ore solution = _____________ N.
Equivalent weight of ferrous Iron = atomic weight / 1 = 55.85 / 1 = 55.85
Amount of Ferrous Iron present in 1 liter of solution = Normality of ferrous iron ore
solution X equivalent weight of ferrous iron =
g
Amount of ferrous iron present in the given 100 ml of solution = amount present in 1
liter of solution / 10 =
g
Table – III
Percentage Error Table
Roll No. /
Regd. No.
Flask
No.
Amount of Ferrous Iron present in
100 ml of Iron ore solution. Grams.
Reported
Given
Percentage
error
Report: The amount of ferrous iron present in the given 100ml. unknown
solution is
g
***
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4. Estimation of Calcium in port land cement
Aim
: Estimation of calcium (as oxide) in the given sample of portland
cement by using potassium permanganate solution.
Theory
:
Metals like as calcium, copper, lead and zinc which give sparingly soluble
oxalates may be determined by dissolving the washed precipitate in dilute sulphuric
acid and titrating with standard potassium permanganate solution. This method is
widely used for the determination of calcium.
Calcium is precipitated as oxalate by the addition of ammonium oxalate
solution to a dilute Hydrochloric acid solution of the cement followed by the
neutralisation of the acid with dilute ammonia solution. The washed precipitate is
dissolved in dilute sulphuric acid and the liberated oxalic acid titrated with the
standard potassium permanganate solution.
Cement + HCl  CaCl2+ Excess HCl
Excess HCl + NH4OH  NH4Cl + H2O (Neutralization)
CaCl2 + (NH4)2 C2O4  CaC2O4 + NH4Cl
CaC2O4 + H2SO4  CaSO4 + H2C2O4
2KMnO4 + 3 H2SO4 + 5H2C2O4  K2SO4 + 2MnSO4 + 8H2O + 10CO2
In this reactions all the reactants and products are colourless except potassium
permanganate. The colour of potassium permanganate is used to locate the end point
and one of the products i.e., Mn+2 ions act as catalyst to catalyze the preceding
reaction. So this is an example of self indicator and auto - catalytic reaction.
Preparation of Cement Solution :
About 5.0 grms. of portland cement are accurately weighted into a 600 ml
beaker, covered with a watch glass, and dissolved in 1:4 hydrochloric acid. The
contents of the beaker are heated to boiling. The solution is now filtered and the
residue on the filter paper is washed a number of times with a little acid. The filtrate
and washings are boiled and made ammonical with 1:4 ammonia. The precipitate
formed is allowed to settle and the solution is filtered. Now the precipitate is washed
with hot water and the solution is diluted to 1000 ml. Then the solution is boiled and
from this hot solution calcium is precipitated as the oxalate by the addition of strong
ammonium oxalate solution. Boiling is continued for 10 minutes and the precipitate
is allowed to settle. A drop of ammonium oxalate is added to the supernatant solution
to ensure that precipitation is completed. Now the solution is filtered through
whatmann filter paper followed by washing with cold water until it is free from
oxalate and chloride. Filtrate is tested for the presence of calcium with ammonium
oxalate solution. A hole is pierced in the filter paper with a pointed glass rod, and the
bulk of precipitate is washed with little amount of hot water into the volumetric flask.
The filtrate is treated with dilute sulphuric acid (1:8) and finally filter paper is washed
with hot water. If necessary some more dilute sulphuric acid is added to dissolve
completely.
Theory : Potassium permanganate is an oxidising agent with the formula KMnO4
and oxalic acid is a reducing agent with the formula H2C2O4.2H2O . Solutions of
oxalic acid can be estimated by titrating against a standard solution of potassium
permanganate in acidic medium (with dilute sulphuric acid) and heating to about 70-
14
0
80 C and titrating while hot. These two react in sulphuric acid medium as per the
following reaction.
2KMnO4 + 3 H2SO4  K2SO4 + 2MnSO4 + 3H2O + 5 (O)
5H2C2O4 + 5(O)  10CO2 + 5H2O
-------------------------------------------------------------------------------------2KMnO4 + 3 H2SO4 + 5H2C2O4  K2SO4 + 2MnSO4 + 8H2O + 10CO2
-------------------------------------------------------------------------------------In the beginning stages the reaction is very slow, even though it is heated to
0
70 C but it is catalyzed by Mn+2 ions which are formed in the same reaction. Here all
the reactants and products are colourless except potassium permanganate. When all
the oxalic acid has completely reacted, the excess trace amount of permanganate will
give the pink colour to the solution, which is the end point of titration. In other words
in this reaction no indicator is necessary, because potassium permanganate itself acts
as an indicator. Also one of the products that is Mn+2 ions are acting as a catalyst,
which catalyzes the preceding reaction. So it is an example of self indicator coupled
with auto catalytic type of reaction.
Part – I :
Standardisation of potassium permanganate solution by titrating against a
standard oxalic acid solution
Procedure :
10.0 ml of standard oxalic acid is pipetted out into a clean conical flask
carefully and to this 45.0 ml of distilled water followed by 5.0 ml of 1:1 dilute
sulphuric acid are added with a measuring jar. The burette is filled with potassium
permanganate solution after rinsing with the same. Now the conical flask contents are
heated to just boiling which will be indicated by commencement of bubbles from the
bottom of the flask. After noting the initial reading of the burette a small portion ( ~
0.5 ml ) of potassium permanganate is added to the hot solution of conical flask and
the flask is kept in undisturbed position until the permanganate solution is
decolourised to produce the sufficient amount of Mn+2 ions, which catalyse the
preceding reaction. Now the titration is continued while hot with permanganate
solution until the pale pink colour is obtained. The final reading of the burette is
noted which coincides with upper meniscus. A number of titrations are carried out
until 2 or 3 concurrent readings are obtained. The results are tabulated in table No - I.
15
Table – I
Titration of potassium permanganate solution with standard
Oxalic acid solution.
Normality of standard solution of Oxalic acid ____________ N
Indicator : Self Indicator (Colour of the Potassium permanganate)
Colour change at the end point : Colourless to pale pink.
S.No.
Volume of Oxalic acid
solution taken in ml.
( V1 )
Burette readings
Initial
Final
Volume of potassium
permanganate solution
consumed in ml. ( V2)
Calculations :
By the law of equivalence,V1 N 1 = V2 N 2
N1 = Normality of Oxalic acid solution
=
V1 = Volume of Oxalic acid solution
=
N2 = Normality of potassium permanganate solution =
V2 = Volume of potassium permanganate solution =
N
ml.
?
ml
 N2 = V1 N 1 / V2
The Normality of potassium permanganate solution = ___________N.
Procedure :
The Potassium permanganate solution is standardised using standard oxalic
acid solution (refer Part-I of experiment No. 4). The given cement solution is diluted
upto the mark of volumetric flask with distilled water carefully. The flask is
stoppered tightly and is shaken thoroughly about 3 to 5 minutes for complete
homogenisation. 10.0 ml of cement solution is pipetted out into a clean conical flask
carefully and to this 45.0 ml of distilled water followed by 5.0 ml of 1:1 dilute
sulphuric acid are added. The burette is filled with standard potassium permanganate
solution. Now the conical flask contents are heated to just boiling. After noting the
initial reading of the burette a small portion( ~ 0.5 ml) of potassium permanganate is
added to the hot solution and kept it in undisturbed position until the decolourisation
of the permanganate solution. Now the titration is continued while hot with
permanganate solution until the colour changes from colourless to pale pink colour.
The final reading of the burette is noted. A number of titrations are carried out until 2
or 3 concurrent readings are obtained. The results are tabulated in Table – I.
16
Table – I
Titration of standard potassium permanganate solution with a unknown
cement solution.
Normality of standard solution of potassium permanganate _________ N
Indicator : Self Indicator (Colour of the Potassium permanganate)
Colour change at the end point : Colourless to pale pink.
S.No.
Volume of unknown
Cement solution taken in
ml. ( V1)
Burette readings
Initial
Final
Volume of potassium
permanganate solution
consumed in ml. ( V2)
Calculations :
By the law of equivalence, V1N 1 = V2 N 2
N1 = Normality of unknown Cement solution
=
V1 = Volume of unknown Cement solution
=
N2 = Normality of potassium permanganate solution =
V2 = Volume of potassium permanganate solution =
?
ml.
N
ml.
 N 1 = V 2 N 2 / V1
The normality unknown cement solution = ________________ N
Equivalent weight of calcium in cement = Molecular weight / 2= 40 / 2=20
Amount of calcium present in 1 liter of solution = Normality of unknown calcium in
cement solution X equivalent weight of calcium.
g
Amount of calcium present in a given 100 ml of cement solution = Amount present in
1 liter solution / 10
g
Table – II
Percentage Error Table
Roll No. /
Regd. No.
Flask
No.
Amount of calcium present in 100 ml of
cement solution. Grams.
Reported
Given
Percentage
error
Report: The amount of calcium present in the 100 ml. of the given unknown cement
solution is
g
***
17
5. Estimation of Hydrogen peroxide
Aim
: Estimation of Hydrogen peroxide by using potassium
permanganate solution.
Theory
:
Hydrogen peroxide is usually available in the form of an aqueous solution
containing about 3% , 6%, 12% and 30%. Hydrogen peroxide. It is frequently
marketed in four strengths, 10 Vol.,20 vol.,40 vol., and 100 volume concentrations.
In sulphuric acid medium hydrogen peroxide is converted into oxygen and
water when treated with potassium permanganate. The oxidation of hydrogen
peroxide with permanganate in acidic solution proceeds slowly at first, but it is
catalyzed by the manganese (II) ions formed in the reaction (like oxalate and
permanganate).
2KMnO4 + 3 H2SO4  K2SO4 + 2MnSO4 + 3H2O + 5 (O)
5H2O2 + 5(O)  5H2O + 5O2
------------------------------------------------------------------------------------------------2KMnO4 + 3 H2SO4 + 5H2O2  K2SO4 + 2MnSO4 + 8H2O + 5O2
-------------------------------------------------------------------------------------------------
In this reaction manganese in permanganate is reduced from Mn+7 to Mn+2 and
hydrogen peroxide is oxidised to oxygen. Potassium permanganate is the only
coloured solution, so the colour of permanganate solution will give pink colour to the
solution when peroxide is completely oxidised, which shows end of the reaction.
Procedure :
The given hydrogen peroxide solution is diluted up to the mark of volumetric
flask with distilled water carefully. The flask is stoppered tightly and shaken
thoroughly about 3 to 5 minutes for complete homogenization. 10.0 ml of the above
solution is pipetted out into a 250 ml of conical flask carefully. 40.0 ml of distilled
water and 5.0 ml of dilute sulphuric acid are added to the conical flask contents with
a measuring jar . The burette is filled with standard potassium permanganate solution
after rinsing with the same. Now conical flask contents are titrated with potassium
permanganate solution, after noting the initial reading of the burette with constant
swirling until the colour changes from colourless to pale pink. The final reading of the
burette is noted . A number of titrations are carried out until 2 or 3 concurrent
readings are obtained. The results are tabulated in Table No I.
18
Table – I
Titration of an unknown hydrogen peroxide solution
with standard potassium permanganate solution
Normality of standard solution of potassium permanganate = 0.1 N
Indicator : Self Indicator (Colour of the Potassium permanganate)
Colour change at the end point : Colourless to pale pink.
S.No.
Volume of unknown
hydrogen peroxide solution
taken in ml.
Burette readings
Initial
Final
Volume of potassium
permanganate solution
consumed in ml.
Calculations : ( By weight )
Equivalent weight of hydrogen peroxide = Mol.wt./ 2= 34.02 / 2= 17.01
1 gr. Eq.weight of potassium permanganate= 1 gr.eq.weight of hydrogen peroxide. 1
gm. Eq.wt. of potassium permanganate dissolved in 1 lit.solution gives 1 normal
solution
1000 ml of 1 N solution of potassium permanganate = 17.01gm.of hydrogen peroxide
1ml of 0.1N solution of potassium permanganate = 0.001701 gm of hydrogen peroxide
10ml of hydrogen peroxide consumes‘X’ml of 0.1N potassium permanganate solution
Amount of hydrogen peroxide present in 1 lit. of solution = 100 x X x 0.001701
…………..gm.
Amount of hydrogen peroxide present in a given 100 ml of solution
=10x X x 0.001701
gm.
Table – II
Percentage Error Table
Roll No. /
Regd. No.
Flask
No.
Amount of Hydrogen peroxide present
in 100 ml of solution, Grams
Reported
Given
Percentage
error
Report: The amount of hydrogen peroxide present in the 100 ml. of the given
solution is
g
19
6 (a) . Estimation of Active Chlorine Content in a Disinfectant Bleaching powder
Aim
:Determination of the amount of available chlorine in the given sample of
Bleaching powder by using standard Hypo solution through Iodometric
method.
Theory
:
Bleaching powder consists essentially of a mixture of Ca(OCl) Cl and basic
chloride [CaCl2, Ca(OH)2. H2O]. The active constituent of the bleaching powder is
hypo chlorite, which is responsible for the bleaching action of the material.This is
very good disinfectant in water, which is due to the formation of nascent oxygen.
CaOCl2 + H2O  Ca(OH)2 + Cl2
Cl2 + H2O  HOCl + HCl.
(HOCl is germicide)
HOCl  HCl + [O] (nascent oxygen is responsible for bleaching action)
In the evaluation of bleaching powder the result is usually expressed in terms of
available chlorine, which is the chlorine liberated in acidic the solution. Available
chlorine can be determined iodometrically by adding potassium iodide and acetic acid
to the suspension of the material.
OCl- + Cl- + 2H+  H2O +Cl2
Cl2 + 2KI  2KCl + I2
In acetic acid medium the reaction proceeds slowly. The liberated chlorine in
presence of acid liberates iodine quantitatively from the solution of potassium iodide.
This liberated iodine can be titrated with the standard solution of sodium thiosulphate.
I2 + Na2S2O3  2NaI + Na2S4O6
In iodometric titrations starch is used as an indicator to locate the end point.Starch
forms soluble blue coloured complex with iodine even at low concentrations.
Preparation of Bleaching powder solution : The given sample of bleaching powder is
weighed out accurately into clean glass mortar. A little water is added and the
mixture is ground in to a smooth paste with a pistle. Then little more water is added to
this paste and mixed thoroughly, allowed to settle, and the milky liquid is poured into
a 100 ml volumetric flask. Then the residue is grounded well with a little more water
and the operations are repeated until the whole of the sample has been transferred into
the flask either in solution or in a state of very fine suspension. The flask is then
filled up to the mark with distilled water and is shaken thoroughly.
Procedure :
The given bleaching powder solution is diluted upto the mark with distilled
water carefully and stoppered well. The solution is shaken thoroughly for 3 to 5
minutes for complete homogenisation. 10.0 ml of well shaken bleaching powder
solution or suspension is pipetted out into a clean conical flask and to this 50.0 ml of
distilled water, 10.0 ml of 10% potassium Iodide solution are added using a
measuring Jar. Then the above solution is acidified by adding 10.0 ml of glacial
acetic acid and the flask is covered with watch glass for about 5 minutes for complete
liberation of iodine. The burette is filled with standard hypo solution and the initial
reading is noted. The brown coloured solution (liberated iodine) is then titrated with
standard hypo solution drop wise with constant shaking until the brown colour
becomes pale yellow. At this stage 1.0 ml of 1% freshly prepared starch solution is
20
added to this solution and the colour changes to blue. The titration is continued until
an excess drop of hypo solution changes the colour of the solution from blue to
colourless (disappearance of blue colour), which is the end point of the reaction. The
final reading of the burette is noted. A number of titrations are repeated until 2 or 3
concurrent readings are obtained. The results are tabulated in Table – I.
Table – I
Titration of bleaching powder suspension with a standard hypo solution.
Normality of standard solution of Hypo = ____________ N
Indicator : 1 % freshly prepared starch solution
Colour change at the end point : Blue to Colourless
(disappearance of blue)
Volume of bleaching
S.No.
powder solution taken in ml.
Burette readings
Initial
Final
Volume of Hypo solution
consumed in ml.
Calculations :
According to the law of equivalence
1000 ml of 1 N Hypo solution
= 35.46 gr. of chlorine
1.0 ml of 1 N Hypo solution
= 0.03546 gr. of chlorine
1.0 ml of 0.1 N Hypo solution
= 0.003546 gr. of chlorine
10.0 ml of bleaching powder suspension of consumed ‘X’
ml of 0.1 N hypo
solution.
Amount of chlorine present in the given 100 ml of solution = 10 x X x 0.003546 gr.
% of available chlorine =
10 x X x 0.003546
weight of bleaching powder taken
x 100
Report: The percentage of available chlorine in given bleaching powder sample is
21
6 (b). Determination of Hardness of an Under ground water sample
Aim
:Determination of the total hardness of the given sample of water by
titrating against a standard EDTA solution using Eriochrome Black- T as indicator.
Theory
:Hardness present in a given sample of water can be determined by
using the complexometric method, in which the disodium salt of EDTA is employed
(soluble in water) and it can be represented as follows.
(EDTA – Ethylene Diamine Tetra Acetic acid)
HOOCH2C
CH2COOH
N-CH2 - CH2 -N
NaOOCH2C
CH2COONa
EDTA forms complexes with calcium and magnesium when the pH is in the range of
around 9.5 to 10.5 and to maintain the pH, a basic buffer solution is used (NH 4OH +
NH4Cl buffer serves pH 9.5 to 10.5) . The complexes of calcium and magnesium
with EDTA are colourless, therefore it is necessary to use indicator to locate the end
point. In this titration Eriochrome black – T is used as indicator, which forms an
unstable wine red coloured complex with calcium and magnesium. Calcium ions
complexed first with EDTA, but the colour change does not occur until all the
magnesium has also completely reacted. It is thus possible to determine the total
amount of these metals in the solution and the total hardness can be calculated.
Calcium itself does not give a satisfactory end point with Eriochrome black – T
indicator unless the solution also contains magnesium. When once all the calcium
and magnesium ions are completely removed by EDTA,free indicator is left in the
solution which imparts blue colour to the solution. So the colour change at the end
point is wine red to blue.
Ca 2
pH 10
+ 2In-2 
 In2
2
Mg
Ca
Ca
pH 10
+ EDTA 
+ 2In-2
 EDTA
Mg
Mg
unstable complex
Stablecomplex
Free indicator
Wine red in colour
Colourless
Blue in colour
The di-sodium salt of EDTA solution can be standardized by using standard zinc
sulphate solution in presence of ammonia – ammonium chloride buffer (pH 10) using
Eriochrome Black – T as indicator.
Zn+2 + H2 Y2-  ZnY2- + 2H+
Procedure :
50.0 ml of sample of hard water is pipetted out into a clean conical flask. To this
2 or 3 ml of ammonia - ammonium chloride buffer solution (pH 9.5 – 10.5) and 2
or 3 drops of Eriochrome Black – T indicator are added. (if necessary 2 drops of
0.1M – Mg – EDTA complex is added in order to get the clear end point). The
burette is filled with 0.01 M EDTA solution, after rinsing with same and the
initial reading is noted. Now the contents are titrated with EDTA solution until
the colour changes from wine red to blue which is the end point of the reaction.
The final reading of the burette is noted. A number of titration are carried out
until 3 or 4 concurrent readings are obtained. The results are tabulated in Table– I.
22
Table – I
Titration of standard EDTA solution with unknown water sample
Morality of standard solution of EDTA ____________ M
Indicator : Eriochrome Black – T.
Colour change at the end point : Wine red to blue
S.No.
Volume of Water sample
taken in ml.
Burette readings
Initial
Final
Volume of EDTA
solution consumed in ml.
Calculations :
1 gram mol.weight of EDTA = 1 gram mol.weight of calcium carbonate
1 gram mol.weight of EDTA dissolved in 1000 ml of solution gives one
molar solution
1000ml of 1 M EDTA
=
100 gr. of calcium carbonate
1.0 ml of 1 M EDTA
=
0.1 gr of calcium carbonate
1.0 ml of 0.01 M EDTA
=
0.001 gr of calcium carbonate
1.0 ml of 0.01 M of EDTA equivalent to 1.0 mg of calcium carbonate
equivalent hardness
50.0 ml of hard water samples consumes ‘x’ ml of 0.01 M EDTA solution
 Total hardness of water sample per liter = X x 1000 / 50 = X x 20 mg/litre
or ppm.
Report: The total hardness of the given water sample is
***
mg/l or ppm.
23
7. Estimation of Chromium (VI) in potassium dichromate
Aim
: Estimation of Chromium (VI) in potassium dichromate by titration
against standard solution of Mohr’s salt.
Theory
:
Potassium dichromate (K2Cr2O7) acts as an oxidising agent in the presence of
sulphuric acid or hydrochloric acid oxidising ferrous iron to ferric iron, getting it self
reduced to a green chromic (Cr + 3) salt. and Mohr’s salt is a reducing agent with the
formula (NH4)2 SO4. FeSO4. 6H2O . In acid solution, the reduction of potassium
dichromate may be represented as :
K2Cr2O7 + 4H2SO4  K2SO4 + Cr2 (SO4)3 + 4H2O + 3 (O)
K2Cr2O7 + 8HCl  2KCl + 2CrCl3 + 4H2O + 3 (O)
From either of these equations it follows that the equivalent weight of potassium
dichromate is one sixth of the molecular weight. The reaction between potassium
dichromate and Ferrous iron is as follows .
K2Cr2O7 + 4H2SO4  K2SO4 + Cr2 (SO4)3 + 4H2O + 3 (O)
6 FeSO4 + 3H2 SO4 + 3 (O)  3 Fe2 (SO4)3 + 3H2O
-----------------------------------------------------------------------------------------K2Cr2O7 + 7H2SO4 + 6 FeSO4  K2SO4 + Cr2 (SO4)3 + 3 Fe2 (SO4)3 + 7H2O
-------------------------------------------------------------------------------------------As potassium dichromate solution is not suitable to locate the end point,
diphenyl amine is used as an external indicator to locate the end point of the reaction.
This reaction imparts green colour to the iron (II) solution due to the formation of
chromic sulphate, which deepens to blue green colour shortly before the end point of
the titration and at the end point intense purple or bluish violet colour is obtained.
Due to weak oxidising nature of dichromate when compared to permanganate,
the reaction becomes slow at the end point. So syrrupy phosphoric acid should be
added in order to increase the rate of reaction at the end pint, by removing the Ferric
ions as Feric phosphate from the sphere of the reaction. According to Lechatlier’s
principle the equilibrium shifts to forward direction. More over also the phosphoric
acid reduces the reduction potential of Fe(II)–Fe(III) system by 0.15 – 0.3 volts,
there by reducing capacity is increased, when compared to diphenylamine indicator.
When once all the ferrous iron ions completely oxidized with dichromate an excess
drop of dichromate will oxidise the indicator from bluish green to bluish violet.
Part – I :
Standardisation of Mohr’s salt solution by titration against a standard solution
of potassium dichromate.
Procedure :
10.0 ml of Mohr’s salt solution is pipetted out into a clean conical flask
carefully. 40.0 ml of distilled water, 5.0 ml of 1:1 dilute sulphuric acid and 3.0 ml of
syrupy phosphoric acid are added with measuring cylinder. One or Two drops
diphenyl amine indicator is added directly to the conical flask contents. The burette
is rinsed and then filled with standard potassium dichromate solution, the initial
reading is noted, which coincides with the lower meniscus, after removing the air gap
from the nozzle. Now the contents are titrated with potassium dichromate solution
with constant thorough shaking until the colour changes from colourless - pale green,
24
- dark green – bluish green to bluish violet, which is the end point of the reaction.
The final reading of the burette is noted without parallax error. A number of titrations
are carried out until 2 or 3 concurrent results are obtained. The results are tabulated in
Table No. 1.
Table – I
Titration of standard potassium dichromate solution with Mohr’s salt
solution.
Normality of standard solution of potassium dichromate ____________ N
Indicator : 1 % Diphenyl amine solution
Colour change at the end point : Bluish green to Bluish violet
S.No.
Volume of Mohr’s salt
solution taken in ml.
(V1)
Burette readings
Initial
Final
Volume of potassium
dichromate solution
consumed in ml. (V2)
Calculations :
By the law of equivalence, V1 N 1= V2 N 2
N1 = Normality of Mohr’s salt solution
V1 = Volume of Mohr’s salt solution
N2 = Normality of potassium dichromate solution
V2 = Volume of potassium dichromate solution
=
=
=
=
?
ml.
N
ml.
 N1 = V2 N 2 / V1
The Normality of Mohr’s Salt solution = _____________ N.
Part – II :
Estimation of chromium (VI) in the given 100 ml of potassium dichromate
solution by titrating against a standard solution of Mohr’s salt.
Procedure :
The given potassium dichromate solution is diluted upto the mark of
volumetric flask with distilled water carefully. The flask is stoppered tightly and is
shaken thoroughly about 3 to 5 minutes for complete homogenization. 10.0 ml of
standard Mohr’s salt solution is pipetted out into a clean conical flask carefully. 40.0
ml of distilled water, 5.0 ml of 1:1 dilute sulphuric acid and 3.0 ml of syrupy
phosphoric acid are added with a measuring cylinder. One or Two drops diphenyl
amine indicator is added directly to the conical flask contents. The burette is filled
with the above potassium dichromate solution after rinsing with the same solution and
then initial reading is noted. Now the conical flask contents are titrated with
potassium dichromate solution with constant thorough shaking until the colour
changes from colourless - pale green, - dark green – bluish green to bluish violet,
which is the end point of the reaction. The final reading of the burette is noted
without parallax error. A number of titrations are carried out until 2 or 3 concurrent
results are obtained. The results are tabulated in Table No. II.
25
Table – II
Titration of potassium dichromate solution with a standard solution of Mohr’s
salt.
Normality of standard solution of Mohr’s salt ____________ N
Indicator : 1 or 2 drops of Diphenyl amine solution
Colour change at the end point : Bluish green to Bluish violet
Volume of Mohr’s salt
Burette readings
Volume of potassium
S.No.
solution taken in ml.
dichromate solution
Initial
Final
(V3)
consumed in ml. (V4)
Calculations :
By the law of equivalence
V3 N3 =V4 N4
N3 = Normality of Mohr’s salt solution
V3 = Volume of Mohr’s salt solution
N4 = Normality of potassium dichromate solution
V4 = Volume of potassium dichromate solution
=
=
=
=
N
ml
?
ml
 N4 = V 3 N3 / V 4
Normality of unknown potassium dichromate solution =___________ N
Eq.wt.of chromium (VI) in potassium dichromate = At. wt./6 =104/6 =17.33
Amount of chromium (VI) present in 1 lit. of solution = Normality of an unknown
dichromate solution X equivalent weight of chromium (VI) =
g
Amount of chromium (VI) present in the given 100 ml solution = Amount present in 1
lit. of solution / 10 =
g
Table – III
Percentage error table
Roll No. /
Regd. No.
Flask
No.
Amount of chromium (VI) present in the
given 100 ml of solution, Grams.
Reported
Given
Percentage
error
Report: The amount of Chromium(VI) present in a given 100ml of an unknown
solution is
g.
26
Eq.wt.of potassium dichromate = Mol.wt./6 =
294.18
= 49.03
6
Amount of potassium dichromate present in 1 lit. of solution =
Normality of an un known potassium dichromate solution X equivalent weight of potassium
dichromate =
g
Amount of potassium dichromate present in the given 100 ml solution = Amount
present in 1 lit. of solution / 10 =
g
Table – III
Percentage error table
Roll No. /
Regd. No.
Flask
No.
Amount of potassium dichromate
present in the given 100 ml of solution,
Grams.
Reported
Given
Percentage
error
Report: The amount of potassium dichromate present in a given 100ml unknown
solution is
g.
27
8. Estimation of Copper in Copper ore
Aim
: Determination of Copper in an ore by titration against a standard
solution of hypo through Iodometric method.
Theory
:Copper sulphate has the formula CuSO4.5H2O and hypo with the
formula Na2S2O3.5H2O . Copper sulphate cannot be titrated directly with hypo,
(because there is no direct reaction between them) it can be estimated indirectly via
the liberation of iodine from potassium iodide. Copper sulphate liberates Iodine
quantitatively from the solution of potassium Iodide when freed from mineral acids.
The liberated Iodine reacts with hypo. This method is known as indirect Iodometric
titration method. (Iodometry). In Iodometric titrations starch is used as an indicator
which gives blue coloured complex even at very low concentrations of Iodine. The
two redox reactions involved are given below.
CuSO4 + 4KI  2 CuI + 2K2SO4 + I2
1st redox reaction
I2 + 2Na2S2O3  2NaI + Na2S4O6
2nd redox reaction
2 CuSO4  I2  2Na2 S2 O3
The Copper sulphate solution may contain mineral acid impurities and these
acids interfere with the above reaction making Iodides oxidised by using atmospheric
Oxygen. Also it may some times interfere with starch indicator making it hydrolysed.
4I- + O2 + 4H+  2I2 + 2H2O .
So the above acids can be neutralized by adding sodium carbonate. Sodium carbonate
also precipitates some of copper sulphate as copper carbonate. It can be re-dissolved
with dilute CH3 COOH (1:1) (non-mineral acid).
Part – I
:
Standardisation of Hypo solution by titrating against a standard solution of
copper sulphate.
Procedure :
10.0 ml of standard copper sulphate solution is pipetted out into a clean conical
flask carefully. To this sodium carbonate solution is added drop wise until the solution
becomes turbid and then dilute acetic acid (1:1)is added by drop wise to re-dissolve
the precipitate obtained and the solution becomes clear . Then 10.0 ml of 10%
potassium iodide solution is added to the conical flask contents with gentle swirling
and it is covered with watch glass immediately. After 1 or 2 minutes,the solution is
diluted with 30.0 ml of distilled water and the resultant solution is dark brown in
colour. The burette is filled with the given hypo solution after rinsing with the same
solution and initial reading of the burette is noted. Now the dark brown conical flask
contents are titrated with hypo solution drop wise with constant shaking until the
colour of solution changes from dark brown – dark yellow – yellow – pale yellow to
wheatish yellow / straw yellow. At this stage 1.0 ml of freshly prepared starch
solution is added. Then the solution becomes blue in colour. The titration is
continued drop wise until the blue colour just disappears due to cuprous iodide
formation (flesh white). Which is the end point of the titration. The final reading of
the burette is noted. A number of titrations are carried out until 2 or 3 concurrent
readings are obtained. The results are tabulated in Table No-I.
28
Table – I
Standardization of Hypo solution with standard copper sulphate solution.
Normality of standard copper sulphate solution ____________ N
Indicator : 1 % freshly prepared starch solution
Colour change at the end point : Blue to flesh white
S.No.
Volume of copper sulphate
solution taken in ml. (V2)
Burette readings
Initial
Final
Calculations :
By the law of equivalence
V1 N1 = V 2 N2
N1 = Normality of Hypo solution
=
V1 = Volume of Hypo solution
=
N2 = Normality of copper sulphate solution =
V2 = Volume of Copper sulphate solution
=
Volume of Hypo solution
consumed in ml. (V1)
?
ml
N
ml
 N1 = V 2 N2 / V 1
The Normality of Hypo solution = _____________ N.
Part II :
Estimation of copper in a copper ore by titrating against a standard
solution of Hypo via the liberation of Iodine through potassium Iodide.
Procedure :
The given copper ore solution (as sulphate) is diluted upto the mark of
volumetric flask with distilled water carefully. The flask is stoppered tightly and is
shaken thoroughly about 3-5 minutes for complete homogenization. 10.0 ml of the
copper ore solution is pipetted out into a clean conical flask carefully. To this sodium
carbonate solution is added by drop wise until the solution becomes turbid. Then
dilute acetic acid (1:1) is added by drop wise until the solution becomes clear. Then
10.0 ml of 10% potassium iodide solution is added to the conical flask contents with
gentle swirling and then flask is covered with watch glass immediately for one or two
minutes for complete liberation of iodine. The solution is then diluted with 30.0 ml of
distilled water and the resultant solution is dark brown in colour. The burette is filled
with standard Hypo solution and initial reading of the burette is noted. Now the dark
brown conical flask contents are titrated with Hypo solution by drop wise with
constant shaking until the colour of solution changes from dark brown – dark yellow –
yellow – pale yellow to wheatish yellow / straw yellow. At this stage 1.0 ml of
freshly prepared starch solution is added to wheatish yellow coloured conical flask
contents, the solution becomes blue in colour. Now the titration is continued drop
wise with thorough shaking until the blue colour just disappears (flesh white) which is
the end point of the reaction. The final reading of the burette is noted. A number of
titrations are carried out until 2 or 3 concurrent readings are obtained. The results are
tabulated in Table No-II.
29
Precautions :
1. After the addition of potassium iodide the flask is covered with watch glass
for about one / two minutes for complete liberation of Iodine.
2. Starch indicator should be added before the end point.
3.The minimum quantities of Na2 CO3 solution and CH3 COOH are to be added.
Table – II
Titration of standard Hypo solution with copper ore solution.
Normality of standard solution of Hypo ____________ N
Indicator : 1 %freshly prepared starch solution
Colour change at the end point : Blue to flesh white
S.No.
Volume of copper ore
solution
taken in ml. (V4)
Burette readings
Initial
Final
Volume of Hypo solution
consumed in ml. (V3)
Calculations :
By the law of equivalence , V3 N3 = V4 N4
N3 = Normality Hypo solution
V3 = Volume of Hypo solution
N4 = Normality of copper ore solution
V4 = Volume of Copper ore solution
=
=
=
=
N
ml
?
ml
 N4 = V 3 N3 / V 4 .
The Normality of copper ore solution _____________ N.
Equivalent weight of copper = atomic weight / 1 = 63.6 / 1 = 63.6
Amount of copper present in 1 liter of solution = normality of copper ore solution x
equivalent weight of copper =
g
Amount of copper present in the given 100 ml of copper ore solution = amount
present in 1 liter of the solution / 10 =
g
Table – III
Percentage Error Table
Roll No. /
Regd. No.
Flask
No.
Amount of Copper present in the
given100 ml of ore solution, Grams.
Reported
Given
Percentage
of error
Report: The amount of copper present in the given 100 ml of an unknown solution
is
g.
30
Equivalent weight of copper sulphate = molecular weight / 1 =249.7 / 1 =249.7
Amount of copper sulphate present in 1 liter of solution = normality of copper ore
solution x equivalent weight of copper sulphate =
g
Amount of copper sulphate present in the given 100 ml of copper ore solution =
amount present in 1 liter of the solution / 10 =
g
Table – III
Percentage Error Table
Roll No. /
Regd. No.
Flask
No.
Amount of Copper present in the
given100 ml of ore solution, Grams.
Reported
Given
Percentage
of error
Report: The amount of copper sulphate present in the given 100 ml of an unknown
solution is
g.
31
9 (a). DETERMINATION OF VISCOSITY OF THE GIVEN LIQUID
Aim
To determine the viscosity of the given liquid using Ostwald’s viscometer.
Apparatus and Chemicals
Ostwald’s viscometer, stop watch, specific gravity bottles, pipette, rubber
tubing, water, organic liquid.
Theory
The Ostwald’s viscometer method is based on Poseuille’s equation. This
relates the rate of flow of a liquid through a capillary tube with the coefficient of
viscosity expressed by the equation
 =
 r4 t
8vl

Where r = Volume of the liquid of viscosity  flowing in time ‘t’ through a
capillary tube of radius ‘r’ and length ‘l’. p- hydrostatic pressure of the
liquid.
The determination of absolute viscosity by means of poseuille’s expression
which involves the determination of v, r, t, l and p. But in practice this method is
tedious. Hence a simpler method of comparing the viscosities of two liquids can be
followed.
If t1 and t2 are the flow times required to flow for equal volumes of two liquids
through same length of capillary tube, then
1
Pt
 1 1
2
P2 t 2

P = hdg
here h & g are same for the two liquids.
1
dt
 11
2
d2t 2
1 
d1 t1
d2 t2
x 2
Procedure
The viscometer is cleaned first with chromic acid, water and then with distilled
water. It is finally washed with alcohol and ether and then dried. A piece of clean
rubber tube is attached to the end ‘C’ of viscometer and is clamped vertically in air.
32
A sufficient volume of distilled water is introduced in one of the bulbs (B) so that the
bend portion of tube and half or a little more than half bulb (B) are filled up. With
the help of rubber tube attached to the upper arm of bulb (A) water is sucked until it
raises above the upper mark ‘C’ and is allowed to flow under its own weight. The
time of flow of water from ‘C’ to ‘D’ is counted by starting stopwatch as the meniscus
just passes lower mark ‘D’. The same procedure is repeated three times and the mean
values is determined. The viscometer is cleaned and dried. The same procedure is
repeated with the same volume of given liquid and the time of flow of liquid is
measured and the values are recorded in Table-1. The relative density of a given
liquid is measured using pycnometer.
Observations
S.No.
Water
Time of flow
Given liquid
Time of flow Mean (t1)
Mean (t2)
1
2
3
4
1
dt
 11
2
d2t 2
Where
1 =
2 =
d1 =
d2 =
Viscosity of the given liquid = ?
Viscosity of water=
poise
density of the given liquid = gm/cm3
density of water = gm/cm3
t1 = time of flow of given liquid = sec
t2 = time of flow of water
= sec
(Absolute viscocity)
1 
d1 t1
d2 t2
x 2
Precautions
1
The viscometer should be thoroughly cleaned.
2
3
Viscometer must be strictly kept in vertical position
Same volumes of liquid and water are to be taken while performing the
experiment..
Report
The relative viscosity of the given liquid with respect to water at room
temperature is ________poise.
33
9 (b). DETERMINATION OF SURFACE TENSION OF A GIVEN LIQUID
Aim :
To determine the surface tension of the given liquid at room temperature by
stalagmometer.
Apparatus and chemicals :
Stalagmometer, beaker, rubber tubing, pinch cork, relative density bottle,
thermometer, water, given organic liquid.
Theory :
When a liquid is allowed to flow through a capillary tube, a drop is formed at
its lower end. It increases to a certain size and falls off. The size of the drop depends
on the radius of the capillary and the surface tension. The surface tension acting
along the circumference of the capillary tube supports the drop in the upward
direction.
The measurement of surface tension of a liquid is based on the fact that the
drop of the liquid at the lower end of capillary falls down when the weight of the drop
becomes equal to the surface tension. The surface tension of the given liquid is
determined relative to water at room temperature by using stalagmometer. The
number of drops for the same volume of water and the given liquid are counted and
let there be n1 and n2 respectively. Now if the densities of water and given liquid at
room temperature as determined separately using specific gravity bottle, then the
surface tension  2 of the given liquid can be calculated using the relationship.
 1 n 2 d1
dynes/cm
 x
 2 n1 d 2
Procedure
The stalagmometer is cleaned thoroughly first with chromic acid solution and
finally with distilled water and then dried. The lower end of stalagmometer is
immersed in a beaker containing distilled water. The water is sucked until the level
rises above the mark ‘C’ and the screw is tightened. The liquid is allowed carefully
so that the liquid drops start falling at an interval of about 2-3 sec. in successive drops.
Counting of the drops is started when the meniscus just reaches the upper mark ‘C’
and stopped when the meniscus just passes the lower meniscus ‘D’. The same
procedure is repeated thrice and the mean value is considered. The values obtained
are noted in Table 1. The stalagmometer is cleaned and dried. It is then filled until the
level rises above the upper mark ’C’ and the number of drops are counted as
described earlier. A specific gravity bottle is cleaned and dried. The density of the
given liquid is measured using the specific gravity bottle.
34
Table – I
Determination of Surface Tension of the given liquid
S.No.
Water (n1)
No. of drops
Mean
Given liquid (n2)
No. of drops
Mean
1
2
3
 1 n 2 d1
 x
 2 n1 d 2
1
2
n1
n2
d1
d2
= Surface tension of water = …….. dynes/cm
= Surface tension of the given liquid = ?
= No. of drops of water =
= No. of drops of the given liquid =
= density of water = grams / cm3
= density of the given liquid = grams / cm3
 2 =
n1 d 2
 x xd
x
x 1  1 1 2
n 2 d1
n2 x d2
Precautions
1. The stalagmometer and the specific gravity bottle should be cleaned properly
and dried before use.
2. The stalagmometer should be fixed vertically.
3. The No. of drops must be between 15-20 per 3 minutes.
Report
The relative surface tension of the given liquid with respect to water at room
temperature is ___
35
10 (a) . ESTIMASTION OF MOHR’S SALT BY POTENTIOMETRIC METHOD
Aim
To estimate the amount of given Mohr’s salt by titrating against standard
potassium dichromate solution potentiometrically.
Apparatus and solutions required
Potentiometer, calomel electrode, platinum electrode, beaker, salt bridge, N/20
dichromate solution, unknown Mohr’s salt solution.
Theory
It is very well known that the electrode potential of the electrode depends upon
the concentration of its ions in the solution so the potential of an indicator electrode
goes on changing with respect to a standard (reference) calomel electrode by the
change of concentration of ions during the titration. The equivalence point is
indicated by fairly a large change in electrode potential value. This can be found by
plotting a graph between the emf of the cell on Y axis and the volume of titrant and
added on X axis.
0.591
[oxidant]
E = Eo +
log
n
[reductant]
Procedure
10.0 ml of unknown Mohr’s salt solution is pipetted out in to a 100 ml beaker.
5.0 ml of 1:1 sulphuric acid and 35 ml of distilled water are taken in a measuring
cylinder and are transferred in to the beaker. The potentiometer is connected with an
indicator electrode (platinum) and a reference electrode (calomal) in proper direction.
The contents of the beaker are then titrated against standard dichromate solution. A
pilot titration is carried out by adding 1 ml portion of dichromate solution each time.
The solution is thoroughly mixed and the corresponding emf values are noted in table
No. 1. At the equivalence point a large change in potential is noticed. Similarly an
accurate titration is carried out by adding 0.1 ml portions of dichromate solution at the
vicinity of the equivalence point. The results are presented in table No. II. A graph is
drawn between volume of dichromate solution on X-axis and potentials on Y-axis for
the accurate titration.
Table I
Pilot titration of Mohr’s salt solution with standard solution of potassium
dichromate
10.0 ml Mohr’s Salt Solution + 5 ml of 1:1 H2SO4 + 35 ml water
S.No.
Volume of dichromate
solution (ml)
Potentials (mV)
36
Table II
Accurate titration of Mohr’s salt solution with standard solution of potassium
dichromate
10.0 ml Mohr’s Salt Solution + 5 ml of 1:1 H2SO4 + 35 ml water
S.No.
Volume of dichromate
solution (ml)
Calculations
According. to law of equivalence, V1N1
V1= Volume of Mohr’s salt solution
N1 = Normality of Mohr’s salt solution
V2 = Volume of dichromate solution
N2 =Normality of dichromate solution
Potentials (mV)
= V2N2
= 10.0 ml
=?
=
ml
=
0.05N
 N1 
V2 N 2
V1
Normality of Mohr’s salt solution =
N
Amount of Mohr’s salt in 1lit = Normality x Eq. wt.
= N1 x 392.2
N1 x 392.2
Amount of Moh r’s salt present in 100 ml =
=
10
gm.
Table– III
Percentage error table
Roll No. /
Regd. No.
Flask
No.
Amount of Mohr’s salt present in 100
ml of the solution. Grams.
Reported
Given
Report
Amount of Mohr’s salt present in 100 ml of solution
=
Percentage
error
g.
Precautions
1. The platinum electrode should be activated using spirit lamp before starting the
titration.
2. The solution should be thoroughly shaken with a stirrer each time.
37
10 (b). DETERMINATION OF STRENGTH OF HYDROCHLORIC ACID BY PH
METRIC METHOD
Aim
To determine the strength of hydrochloric acid with standard sodium hydroxide
by pH metric method.
Apparatus &Chemicals required:
pHmeter, Combined glass electrode, 250 ml beaker, semimicro burette, distilled
water bottle,. 0.1N NaOH, unknown HCl solution, Buffer solution of pH 4 and 9.2. .
Theory
Hydrochloric acid can be estimated using sodium hydroxide (following pH
metric method) as per the neutralisation reaction. The pH meter will give good result
between the pH ranges 2 and 10.
The pH meter consists of combined glass electrode and indicator electrode
which responds to H+ conc. and a calomel electrode (reference electrode). The
electrochemical cell is
H2(Pt) | HCl | Salt bridge | (N) KCl, Hg2Cl2(S) | Hg
1 atm
(Sat.KCl)
The emf of the cell is given by
0.0591
E = Eo +
log [H+] at 27oC
n
Procedure
The unknown solution of hydrochloric acid is made up to the mark with
distilled water. 10.0 ml of the solution is pipetted out in to a clean beaker. 40 ml of
distilled water is added to this solution. Meanwhile the pH meter is switched on and
is allowed to stabilize for about 10 to 15 min. The instrument is calibrated with
standard buffers of pH 4 and 9.2. A standard solution of sodium hydroxide is taken in
a micro burette. A pilot titration is carried out by titrating the contents of the beaker
with 1 ml portions of sodium hydroxide each time. The solution is thoroughly mixed
and the corresponding pH values are noted in table I. Accurate titration is carried out
similarly with 0.1 ml portions of sodium hydroxide near the neutralization point. The
corresponding pH values are recorded in Table – II
Table I
Pilot titration of unknown hydrochloric acid solution with standard sodium
hydroxide solution
10 ml of hydrochloric acid + 40 ml of distilled water
S.No.
Volume of NaOH Solution
1
2
3
0.5
1.0
1.5
pH
38
Table II
Accurate titration of unknown hydrochloric acid solution with standard sodium
hydroxide solution
10 ml of hydrochloric acid + 40 ml of distilled water
S.No.
Volume of NaOH Solution
1
2
3
0.5
1.0
1.5
pH
The corresponding pH values are recorded in Table II. 10 ml Hydrochloric acid + 40
ml distilled water.
Calculations
According to law of Equivalence, V1N1 = V2N2
V1 - Volume of sodium hydroxide =
ml
N1 - Normality of sodium hydroxide = N
V2 - Volume of hydrochloric acid =
ml
N2 - Normality of hydrochloric acid = ?
Amount of hydrochloric acid present in 100 ml of the given solution
Normalityof HCl x Eq. Wt.of HCl (36.5)

10
=
gm.
Table– III
Percentage error table
Roll No. /
Regd. No.
Flask
No.
Amount of Hydrochloric acid present in
100 ml of the solution. Grams.
Reported
Given
Percentage
error
Report : Amount of Hydrochloric acid present in 100 ml of the solution -- Grams.
***