Analytical_Chemistry_Lab_manual_Bsc_Chemistry.pdf

This is manual which consist of analytical methods including acid-base titrations, redox titrations, complexometric titrations.

1. LAB MANUAL
ANALYTICAL CHEMISTRY

FOR

B. Sc. CHEMISTRY

2. Index
1. PREPARATION AND STANDARDIZATION OF SODIUM HYDROXIDE
SOLUTION
2. ESTIMATION OF FERROUS ION BY DICHROMATE METHOD
3. ESTIMATION OF HARDNESS OF WATER BY EDTA
4. DETERMINATION OF CHLORIDE ION IN GIVEN WATER SAMPLE
5. ESTIMATION OF OXALIC ACID USING KMnO4 SOLUTION
6. ESTIMATION OF COPPER USING THIOSULPHATE (HYPO) SOLUTION
7. ESTIMATION OF ALKALI CONTENT - ANTACID TABLET USING HCl
8. ESTIMATION OF CARBONATE AND BICARBONATE PRESENT TOGETHER
IN A MIXTURE

3. EXPERIMENT- I
PREPARATION AND STANDARDIZATION OF SODIUM HYDROXIDE
SOLUTION
AIM
To prepare and standardize the solution of sodium hydroxide (N/10) against standard solution
of oxalic acid (N/5) by acid-base titration
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula
Chemicals: NaOH, oxalic acid, phenolphthalein
THEORY
Standard solution is one in which exact amount of a substance is present in a definite volume
of the solution, or a solution whose concentration (strength) is known to us. Volumetric
solutions are classified as: Primary Standard Solution and Secondary Standard solution.
Primary Standard Solution: The substance whose standard solution is prepared by
dissolving directly its known amount in a definite volume of solvent or solution is known as a
primary standard substances & the solution is called as primary standard solution. Commonly
used primary standard substances are anhydrous Na2CO3, Oxalic acid etc.
Secondary Standard Solution: The substance whose solution cannot be prepared directly by
weighing its definite amount and then dissolving in definite volume of solvent is called
secondary standard substance & the solution is called as secondary standard solution. The
solution of this type of substance firstly prepared is of approximate strength which is then
standardized with a standard solution of a primary standard substance. The common
secondary standard substances are alkali hydroxides, inorganic acids and KMnO4 etc.
Standardization of sodium hydroxide is done by acid base titration. Oxalic acid is
allowed to react with sodium hydroxide in the stoichiometric ratio 1:2 to produce sodium
oxalate and water as products. In this titration, phenolphthalein indicator is used to locate the
end point. Appearance of the pale pink colour which persists for 30 sec is the endpoint.

4. (COOH)2.2H2O + 2NaOH → Na2C2O4 + 4H2O
Normality (N): The normality of a solution is the number of gram-equivalents of the solute
per litre of the solution.
N =



PROCEDURE:
Preparation: 0.4 gm of NaOH is accurately weighed and transferred in 100 mL volumetric
flask. The volume is made up to mark with distilled water to get 0.1 N solution of NaOH.
Standardization: Weigh accurately 1.26 g of oxalic acid into a 100 mL volumetric flask and
make up to the mark with distilled water. Pipette out 10 mL of prepared NaOH solution in the
conical flask. Add 2-3 drops of phenolphthalein as indicator. The colour of the solution
becomes pink. Titrate against 0.5 N oxalic acid solution taken in the burette till the pink
colour of the solution just disappears. Repeat the titration to get the concordant value.
OBSERVATION AND CALCULATIONS
Weight of oxalic acid taken, x = gm
Normality of oxalic acid, NOx =






=
Standardization of NaOH solution
NaOH vs Oxalic acid solution (phenolphthalein indicator)
Sl
No:
Vol of NaOH solution
taken (VNaOH mL)
Burette Reading (mL) Vol of oxalic acid
used (VOx mL) Initial Final
1
2
3


NNaOH VNaOH = NOx VOx
Normality of NaOH (NNaOH) =

5. RESULTS
1. Standard NaOH solution is prepared.
2. Strength of prepared standard NaOH solution =
Precautions
1. Solutions should be made up to the mark after dissolving the solute completely in
small amount of water.
2. During titration the solution should be stirred thoroughly
3. Do not take mean of the burette readings
Viva-Voce Questions
1. What is a standard solution?
2. Define normality
3. Define molarity
4. Difference between primary and secondary standard solutions
5. What is molality?
6. What is the colour change of phenolphthalein indicator in acidic and basic medium?
7. Write down the structures of phenolphthalein in acidic and basic medium
8. What is the endpoint?
9. Why should the last drop of the solution not be blown out of a pipette?
10. What type of reaction is an acid-alkali reaction?

6. Experiment 2
ESTIMATION OF FERROUS ION BY DICHROMATE METHOD
AIM
To determine the ferrous content in the supplied sample of iron ore by titrimetric analysis
against standard potassium dichromate (N/20) solution using potassium ferricyanide
K3[Fe(CN)6] as an external indicator
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula, glass rod, white glazed
tile
Chemicals: standard K2Cr2O7 solution, Potassium Ferricyanide indicator, iron sample
THEORY
Acidic K2Cr2O7 is a strong oxidizing agent. When it is added to ferrous ammonium sulphate
((NH₄)₂Fe(SO₄)₂·6H₂O) solution containing dilute H2SO4, only FeSO4 is oxidized and (NH4)2SO4 remains
unchanged hence not shown in the equation. The reactions taking place are as follows:

The complete reaction is

Ionically,

The end point of potassium dichromate titration is determined by using external indicator. Here the
indicator cannot be added to the solution to be titrated, but is placed in the depressions of a white glazed groove
tile. About 0.1% solution of potassium ferricyanide acts as external indicator. Freshly prepared pot.assium
Ferricyanide indicator is used because the old solution of it might get contaminated with potassium ferrocyanide.
The use of potassium ferricyanide as an external indicator is the best-known example in the titration of iron with
dichromate. As long as ferrous ion is present in the titration mixture, the indicator will turn blue on the transfer
of solution. The appearance of blue colour is due to the formation of ferroferricyanide (Turnbull’s blue)
according to flowing reaction.

Usually at the end point (all Fe
2+
ions are oxidized to Fe
3+
) the colour of the indicators drop becomes
light brownish yellow due to reaction of indicator with Fe
3+
ions to produce brown coloured ferric ferricyanide
complex.

Thus, at the end point, the indicator fails to produce blue color when treated witha drop of titration
mixture.

7. PROCEDURE
1. Rinse and fill the burette with N/20 K2Cr2O7.
2. Place small drops of potassium ferricyanide K3[Fe(CN)6] solution indicator on a dry white
tile with the help of glass rod.
3. Pipette out 10 ml of iron ore solution in a conical flask and add K2Cr2O7 solution from the
burette. After adding one ml solution of K2Cr2O7, withdraw a drop of this solution from the
conical flask with the help of glass rod and mix it with one drop of potassium ferricyanide
indicator on the white tile. Appearance of greenish-blue color indicates that the end point has
not yet reached.
4. Continue adding K2Cr2O7 and repeat the process of withdrawing a drop of solution and
mixing it with one drop of potassium ferricyanide indicator on the white tile. When the
indicator drop does not give blue color (or remains yellow), the end point is reached. Note this
reading.
5. Repeat the above titration by reducing the volume interval of K2Cr2O7 to get the exact end
point.
6. Repeat the process until two concordant readings are obtained.
OBSERVATION AND CALCULATIONS
Sl.
No
Volume of sample solution
taken (mL)
Burette Reading Volume of
K2Cr2O7 used
(mL)
Initial Final
1
2
3


Nsample Vsample = Ndichromate Vdichromate
Normality of sample, Nsample =
??????
�??????�ℎ���??????�� ??????
�??????�ℎ���??????��
??????
�??????����

=
Concentration of ferrous content = Nsample x Eq. wt of Fe2+
= Nsample x 55.84 g/L
=
RESULTS
The ferrous ion content in the given sample = ………………. g/L

Precautions
1. Do not take mean of the burette readings
2. Burette and pipette should be rinsed properly before starting the titration

8. 3. The volume of solution taken should be maintained properly.

Viva Questions
1. What is another name of ferrous ammonium sulphate?
2. What is external indicator?
3. What will happen if indicator is added into titration flask?
4. What is the formula of ferrous ammonium sulphate?
5. Why dil H2SO4 is added during titration?
6. Why K2Cr2O7 is used as an oxidizing agent only in presence of an acid and not in
presence of a base?
7. What are the main ores of Iron?
8. What internal indicators can be used in this titration?
9. Which indicator is used in this titration?
List of indicators
1. N-Phenylanthranilic acid Green to violet-red
2.Diphenylamine in Conc. H₂SO4 Colorless to violet
3.Diphenyl benzidine in Conc. H2SO 4 Colorless to violet
4.Ferroin in Conc. H2SO4 Red to faint blue

9. EXPERIMENT 3
ESTIMATION OF HARDNESS OF WATER BY EDTA
AIM
To estimate the temporary, permanent and total hardness of water present in the given water
sample by titrating it against standard EDTA solution (N/20) using Eriochrome black T as an
indicator
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, spatula, beaker, measuring cylinder, tripod stand,
wire gauze, filter paper, dropper, measuring flask
Chemicals: Ammonium buffer solution, inhibitor, Eriochrome black T indicator, standard
EDTA solution, hard water (given water sample)
THEORY
Hardness of water mainly occurs due to the presence of carbonates (temporary) and non-
carbonates (permanent). Carbonate hardness is due to the amount of carbonates and
bicarbonates present in the solution which can be removed or precipitated by boiling.

Non carbonate hardness is due to the presence of sulphate, chloride or nitrate salts and cannot
be removed by boiling.
Total hardness is defined as the sum of the calcium and magnesium concentrations, both
expressed as calcium carbonate in mg/L. Hardness is most commonly measured by titration
with an EDTA solution. At pH 10, Eriochrome black T dye gives an unstable wine red
colored complex with Ca
2+
and Mg
2+
present in the hard water. Only at this pH such a
complexation is possible. The titration against EDTA solution of known strength changes the
color of complex from wine red to original blue color indicating the end point. When this
solution is titrated against EDTA, it replaces the indicator from the weak complex of metal-
EBT to form stable EDTA complex. When all the hardness causing ions are complexed by
EDTA,the indicator is set free. The color of the free indicator is steel blue. Thus the end point
is the change of color from wine red to steel blue.

10. Since EDTA is insoluble in water, the disodium salt of EDTA is taken for this
experiment. In aqueous solution, EDTA ionizes to give 2 Na
+
ions and acts as strong
chelating agent. EDTA forms colorless stable complexes with Ca
2+
and Mg
2+
ions present in
water at pH = 9-10.

Eriochrome Black T (EBT): a complex organic compound (sodium-1-(1-hydroxy 2-
naphthylazo)-6-nitro-2-naphthol-4-sulphate) act as metal ion indicator. The optimum pH for

11. the determination of hardness of water is 10.0 + 0.1 and is adjusted by NH4OH – NH4Cl
buffer.

Structure of EBT
PROCEDURE
1. Clean the burette thoroughly and rinse with N/20 EDTA.
2. Pipette out 10 mL (V1) of given hard water sample into a conical flask.
3. Add 2-3 mL of basic buffer solution.
4. Add two drops of EBT indicator solution. The solution turns wine red in color.
5. Add the standard EDTA titrant drop by drop with continuous stirring until the last
reddish tinge disappears from the solution. The color of the solution at the end
point is blue under normal conditions.
6. Note down the volume of EDTA added (V2).
7. Repeat for concordant values.
8. Take 50 mL of hard water sample in a beaker and boil it until the volume reduces
to nearly half.
9. Filter the solution and make it to 50 mL by adding distilled water.
10. Pipette out 10 mL of this solution and titrate with EDTA as mentioned above.
11. Repeat for concordant values.
OSERVATIONS AND CALCULATIONS
Water Sample vs EDTA (EBT indicator)
Sl No
Volume of water
sample (Vnormal mL)
Burette Reading (mL) Volume of EDTA
added (VEDTA-N mL) Initial Final
1
2
3

12. Boiled Water Sample vs EDTA (EBT indicator)
Sl No
Volume of water
sample (Vboiled mL)
Burette Reading (mL) Volume of EDTA
added (VEDTA-B mL) Initial Final
1
2
3

1 mL 1 M EDTA ≡ 100 mg of CaCO3
1 mL 1 N EDTA ≡ 50 mg of CaCO3
1) Total hardness as CaCO3
NEDTA x VEDTA-N = Nnormal x Vnormal
Nnormal = NEDTA x VEDTA-N / Vnormal
=
Strength = Nnormal x Eq. wt of CaCO3 x 1000 ppm
= …….ppm
2) Permanent hardness as CaCO3
NEDTA x VEDTA = Nboiled x Vboiled
Nbolied = NEDTA x VEDTA / Vboiled
=
Strength = Nboiled x Eq. wt of CaCO3 x 1000 ppm
= …….ppm
3) Temporary hardness = Total hardness as CaCO3 - Permanent hardness as CaCO3
= …………. mg/L

RESULT
1. Total hardness present in given water sample = ………….
2. Permanent hardness present in given water sample = …………
3. Temporary hardness present in given water sample = …………
Questions:
1 How can you differentiate between hard water and soft water?
2 How the hardness of water is expressed?

13. 3 What does EDTA stands for and its chemical structure
4 Why disodium salt of EDTA is used instead of EDTA in determining hardness of
water?
5 Why buffer solution is used in the experiment?
6 A common example each for acidic and basic buffer

14. EXPERIMENT 4
DETERMINATION OF CHLORIDE ION IN GIVEN WATER SAMPLE
AIM
To the chloride ion content in given water sample by Argentometric method (Mohr’s method)
using potassium chromate as an indicator
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, beaker, dropper, measuring flask
Chemicals: Standard silver nitrate solution (N/50), Potassium chromate, given water sample,
distilled water
THEORY
Chlorine in the form of chloride (Cl
-
) ion is one of the major inorganic anions in water and
wastewater. The chloride concentration is higher in wastewater than in raw water because
sodium chloride is a common article of diet and passes unchanged through the digestive
system. Along the sea coast chloride may be present in high concentration because of leakage
of salt water into the sewage system. It also may be increased by industrial process. In
potable water, the salty taste produced by chloride concentration is variable and depends on
the chemical composition of water. Some waters containing 250 ppm Cl
-
ions may have a
detectable salty taste if sodium cation is present. On the other hand, the typical salty taste
may be absent in waters containing Cl
-
ions when the predominant cations are calcium and
magnesium.
Argentometric titration, also known as silver nitrate titration, is a type of precipitation
titration that involves the use of AgNO3 as the titrant. It is often used for the determination of
halide ions, such as chloride, bromide, and iodide. The Mohr Method uses silver nitrate for
titration, if the pH of solution is neutral or slightly basic. Volhard’s method is used when the
pH of the solution, after the sample has been prepared, is acidic. The Mohr titration should be
carried out under conditions of pH 6.5 – 9. At higher pH silver ions may be removed by
precipitation with hydroxide ions, and at low pH chromate ions may be removed by an acid-
base reaction to form hydrogen chromate ions or dichromate ions, affecting the accuracy of
the end point. During the titration, chloride ion is precipitated as white silver chloride.
Ag
+
+ Cl
-
⇌ AgCl (Solubility product constant, Ksp = 3×10
-10
)
(White precipitate)

15. The indicator (potassium chromate) is added to visualize the endpoint, demonstrating
presence of excess silver ions. Cl
–
ions are more reactive than CrO4
2-
so the Ag
+
ions
preferentially react with Cl
–
ions first to form a white precipitate before reacting with CrO4
2–
.
In the presence of excess silver ions, solubility product of silver chromate exceeded and it
forms a reddish-brown precipitate. This stage is taken as evidence that all chloride ions have
been consumed and only excess silver ions have reacted with chromate ions:
2Ag
+
+ CrO4
2-
⇌ Ag2CrO4 (Ksp = 5×10-
12
)
(Brick red precipitate)
PROCEDURE
1. Blank Titration
 Pipette out 10 mL distilled water in a conical flask and add 2-3 drops of K2CrO4
indicator
 Titrate the solution against N/50 AgNO3 solution taken in the burette till brick red
colour appears in the solution
 Note down the reading
 Repeat the experiment until two concordant readings are obtained
2. Sample Titration
 Pipette out 10 mL sample water in a conical flask and add 2-3 drops of K2CrO4
indicator
 Titrate the solution against N/50 AgNO3 solution taken in the burette till brick red
colour appears in the solution
 Note down the reading
 Repeat the experiment until two concordant readings are obtained
OSERVATIONS AND CALCULATIONS
Distilled Water vs AgNO3 solution (K2CrO4 indicator)
Sl No
Volume of distilled
water (Vdistilled mL)
Burette Reading (mL) Volume of EDTA
added (V1 mL) Initial Final
1
2
3

16. Water Sample vs AgNO3 solution (K2CrO4 indicator)
Sl No
Volume of water
sample (Vsample mL)
Burette Reading (mL) Volume of EDTA
added (V2 mL) Initial Final
1
2
3

Volume of AgNO3 used = (V2 - V1)
NAgNO3 x VAgNO3 = Nsample x Vsample
Nsample = NAgNO3 x VAgNO3 / Vsample
=
Strength = Nsample x Eq. wt of Cl
-
x 1000 ppm
= …….ppm
RESULT
Chloride content present in the given water sample = ………….
Precautions:
1. The whole apparatus should be washed with distilled water before the start of the
experiment
2. The reaction mixture should be briskly shaken during the titration
3. The endpoint of the reaction should be carefully observed
4. The volume of the indicator should be same all the times
5. The pH of the sample solution should be adjusted to 7-8 ranges by adding acidic/basic
solution
Questions:
1 Why the method is called Argentometric method?
2 What kind of titration method is used in estimation of Cl
-
ions?
3 Which indicator is used in titration?
4 What are the two precipitates formed in the titration?
5 How will you determine the end point?
6 What is solubility product?
7 Why Ag2CrO4 is precipitated after AgCl in this titration?
8 Name other indicators used in precipitation titration

17. 9 Name other method used used in estimation of Cl
-
ions
10 Why Mohr’s method is carried out in neutral or slightly basic medium?
11 Name the sources of Cl
-
in water

List of indicators (For your reference)

18. EXPERIMENT 5
ESTIMATION OF OXALIC ACID USING KMnO4 SOLUTION
AIM
To estimate the amount of oxalic acid in given sample using KMnO4 solution
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, beaker, dropper, measuring flask, hot plate
Chemicals: Standard oxalic acid solution (N/20), KMnO4 solution, dil. H2SO4
THEORY
The titration of potassium permanganate (KMnO4) against oxalic acid (C2H2O4) is an
example of redox titration. Potassium permanganate is a strong oxidising agent and in the
presence of sulfuric acid it acts as a powerful oxidising agent.
In dilute H2SO4 medium MnO4
-
quantitatively oxidizes C2O4
2-
to CO2 and itself is reduced to
Mn
2+
:

It is an example of autocatalytic reaction in which Mn
2+
, a product of the reaction, acts as the
catalyst.
KMnO4 solution must be standardized against standard oxalic acid solution in 4 N
H2SO4 medium at 70-80˚C. Purple coloured KMnO4 acts as a self-indicator. Solution
containing MnO4
–
ions are purple in colour and the solution containing Mn
2+
ions are
colourless and hence permanganate solution is decolourised when added to a solution of a
reducing agent. The moment there is an excess of potassium permanganate, the solution
becomes purple. Thus, KMnO4 serves as self indicator in acidic solution.
Reduction Half reaction:- 2KMnO4 + 3H2SO4 → K2SO4 + 2MnSO4 + 3H2O + 5[O]
Oxidation Half reaction:- 5(COOH)2 + 5[O] → 5H2O + 10CO2↑
The overall reaction takes place in the process:
2KMnO4 + 3H2SO4 + 5(COOH)2 → K2SO4 + 2MnSO4 + 8H2O + 10CO2↑

19. This titration cannot be carried out in the presence of acids like nitric acid or
hydrochloric acid because they act as an oxidising agent. So hydrochloric acid chemically
reacts with KMnO4 solution forming chlorine which is also an oxidising agent.
PROCEDURE
Standardization of KMnO4
1. Pipette out an aliquot of 10 mL of N/20 standard oxalic acid in a conical flask.
2. Add 10 mL of 4 N H2SO4 and heat to about 70-80˚C.
3. Titrate the hot solution with supplied KMnO4 solution until the solution turns light
pink colour that is stable for ~30 seconds.
4. Repeat the titration to have a concordant reading.
5. Calculate the strength of KMnO4 solution.
Estimation of Oxalic Acid
1. Pipette out 10 mL given unknown oxalic acid solution in a conical flask.
2. Add 10 mL of dil. H2SO4 and heat to about 70-80˚C.
3. Titrate the hot solution with standardized KMnO4 solution until the solution turns
light pink colour that is stable for ~30 seconds.
4. Repeat the titration to have a concordant reading.
5. Calculate the amount of the supplied oxalic acid in gram per litre.
OSERVATIONS AND CALCULATIONS
Standardization of KMnO4
N/20 Oxalic acid vs KMnO4 solution (Self indicator)
Sl No
Volume of oxalic
acid (VOA mL)
Burette Reading (mL) Volume of KMnO4
added (VKMnO4 mL) Initial Final
1
2
3
NOA x VOA = NKMnO4 x VKMnO4
NKMnO4 = NOA x VOA / VKMnO4
=

20. Estimation of Oxalic Acid
Unknown Oxalic acid vs KMnO4 solution (Self indicator)
Sl
No
Volume of unknown
oxalic acid (VUOA mL)
Burette Reading (mL) Volume of KMnO4
added (VKMnO4 mL) Initial Final
1
2
3
NUOA x VUOA = NKMnO4 x VKMnO4
NUOA = NKMnO4 x VKMnO4/ VUOA
=
Strength = NUOA x Eq. wt of oxalic acid
= ……. gm/L
RESULT
Amount of oxalic acid present in the given sample = ………….
Precautions:
1. Always read the upper meniscus for KMnO4 solution
2. Clean all the apparatus with distilled water before starting the experiment and then rinse
with the solution to be taken in them.
Questions:
1. What kind of titration method is used in estimation of oxalic acid vs. KMnO4?
2. Which indicator is used in titration?
3. What is self indicator?
4. How will you determine the end point?
5. Why is sulfuric acid used here and not hydrochloric acid or nitric acid?
6. Equivalent weight of oxalic acid
7. What will happen if titration is performed without heating?
8. What is oxidising agent and reducing agent?
9. Name the oxidising agent and reducing agent in this titration

21. EXPERIMENT 6
ESTIMATION OF COPPER USING THIOSULPHATE (HYPO) SOLUTION
AIM
Determine the strength per litre of the given solution of CuSO4.5H2O using N/20 potassium
dichromate solution.
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, beaker, dropper, measuring flask
Chemicals: Potassium dichromate solution (N/20), Hypo solution (N/20), dil. H2SO4, 10%
KI solution, Sodium bicarbonate, Copper sulphate solution, Starch solution.
THEORY
Iodometric titrations are those in which some oxidising agent liberates iodine from an iodide
and the liberated iodine is titrated against standard solution of reducing agent added from a
burette. In such titrations a neutral or acidic solution of oxidising agent is used. The amount
of iodine liberated from the iodide is proportional to the quantity of the oxidising agent
present. Evolution of iodine by interaction between an iodide and the oxidising agent is slow.
So, 1-2 minutes should be allowed for the completion of the reaction.
Cu (II) is Iodometrically estimated by treating with an excess of KI solution and
titrating the liberated iodine with a sodium thiosulphate (Na2S2O3.5H2O) solution which is
standardized against a standard K2Cr2O7 solution.
Cu
2+
+ 2I
-
→ CuI2
2CuI2 → Cu2I2 + I2

Sodium thiosulphate solution is standardized against a standard K2Cr2O7 solution:

HI (excess from KI and acid) is readily oxidised by air especially in the presence of
Cr
3+
salts. To overcome this difficulty, about 1 g of solid NaHCO3 should be added to the
solution. CO2 evolved keeps the atmospheric oxygen off.
2KI + H2SO4 → K2SO4 + 2HI
2HI + (O) from air → H2O + I2↑
On adding freshly prepared starch solution, a blue colour is obtained. The indicator should be
added near the end point otherwise a stable starch iodide complex is formed and blue colour
does not disappear on adding more of hypo solution.

22. PROCEDURE
Standardization of Hypo solution
1. Pipette out 10 mL of N/20 standard potassium dichromate solution in to a conical
flask. Add 10 mL of dil. H2SO4 and 5 mL 10% KI solution and 1-2 g solid sodium
bicarbonate.
2. Cover the flask with a watch glass for two minutes. Dilute with 50 mL of distilled
water and titrate with hypo solution from the burette with constant shaking.
3. Add 1-2 mL starch solution when the colour of solution becomes pale yellow.
4. Titrate with hypo solution drop wise till the blue colour changes into light green
colour (due to the presence of Cr
3+
ions).
5. Repeat the titration to have a concordant reading.
Estimation of Copper in given sample
1. Pipette out 10 mL given sample in to a conical flask. Add 5 mL 10% KI solution.
2. Add 10 mL of dil. H2SO4 and heat to about 70-80˚C.
3. Cover the flask with a watch glass for two minutes and dilute with 50 mL of distilled
water.
4. Titrate with hypo solution from the burette with constant shaking. Add 1-2 mL starch
solution when the colour of solution becomes pale yellow. Blue colour solution is
obtained.
5. Titrate with hypo solution drop wise till the blue colour disappears and a white
precipitate is obtained due to the formation of cuprous iodide.
6. Repeat the titration to have a concordant reading.
7. Calculate the amount of copper in given sample in gram per litre.
OSERVATIONS AND CALCULATIONS
Standardization of Hypo solution
N/20 potassium dichromate vs Hypo solution (Starch)
Sl No
Volume of K2Cr2O7
solution (VPD mL)
Burette Reading (mL) Volume of Hypo
solution added
(VHypo mL)
Initial Final
1
2
3

23. NPD x VPD = NHypo x VHypo
NHypo = NPD x VPD / VHypo
=


Estimation of Copper
Sample vs Hypo solution (Starch)
Sl
No
Volume of sample
(Vsample mL)
Burette Reading (mL) Volume of Hypo
solution added
(VHypo mL)
Initial Final
1
2
3
Nsample x Vsample = NHypo x VHypo
Nsample = NHypo x VHypo / Vsample
=
Strength = Nsample x Eq. wt of CuSO4.5H2O (249.5 g)
= ……. gm/L
RESULT
Strength of CuSO4.5H2O in the given sample = ………….
Precautions:
1. Always take hypo solution in the burette
2. Add starch solution towards the end point of the titration
3. The solution should be diluted before titration to check the volatility of iodine duringthe
titration.
4. The indicator should be freshly prepared since on keeping, it is spoiled on account of
bacterial attack.
Questions:
1. What is the difference between iodimetry and iodometry?
2. Why is recommended to add NaHCO3 in iodometric titration?
3. Which indicator is used in titration?
4. Why we need to use freshly prepared starch solution?
5. How will you determine the end point?

24. 6. Why starch is added towards the end point?
7. Equivalent weight of hydrated copper sulphate
8. What is oxidising agent and reducing agent?
9. Name the oxidising agent and reducing agent in this titration

25. EXPERIMENT 7
ESTIMATION OF ALKALI CONTENT - ANTACID TABLET USING HC l
AIM
Determine the alkali content present in the given antacid tablet solution using HCl.
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, beaker, dropper, measuring flask
Chemicals: Standard (0.05 N) Na2CO3 solution, HCl solution, Antacid, Methyl Orange
THEORY
Antacids are a combination of various compounds with various salts of calcium, magnesium,
and aluminum as active ingredients. The antacids act by neutralizing the acid in the stomach
and by inhibiting pepsin, which is a proteolytic enzyme. Commonly present substances in
antacid are aluminium hydroxide Al(OH)3, magnesium hydroxide Mg(OH)2 and sodium
hydrogen carbonate NaHCO3. Commonly found antacid tablets are digene, gelucil etc.
The neutralization reaction occurs during the titration are:
Al(OH)3 + 3HCl → AlCl3 + 3H2O
Mg(OH)2+ 2HCl → MgCl2 + 2H2O
NaHCO3 + HCl → NaCl + CO2 + H2O
PROCEDURE
Standardization of HCl: Pipette out 10 mL of N/10 Na2CO3 solution in 100 mL conical
flask and add 2-3 drops of methyl orange indicator. Titrate it against the HCl solution taken
in the burette, until the solution becomes orange or faint pink. This is the end point. Repeat
thrice for concordant values.
Estimation of alkali content in antacid tablet
Pipette out 10 mL given sample in to a conical flask. Add 1-2 drops of phenolphthalein
indicator to the conical flask. Titrate it against standardized HCl solution taken in the burette,
until the solution becomes colourless. Repeat the titration to have a concordant reading.
OSERVATIONS AND CALCULATIONS
Standardization of HCl solution
Weight of sodium carbonate taken =

26. N/10 Na2CO3 solution vs HCl solution (Methyl Orange)

Estimation of alkali content in antacid tablet
Antacid solution vs HCl solution (Phenolphthalein)
Sl No
Volume of antacid
solution (VA mL)
Burette Reading (mL) Volume of HCl
solution added
(VHCl mL)
Initial Final
1
2
3

NA x VA = NHCl x VHCl
NA = NHCl x VHCl / VA
=
RESULT
Normality of OH
-
ions (alkali content) in the given sample = ………….
Precautions:
1. Solutions should be made up to the mark after dissolving the solute completely in small
amount of water.
2. Stomach antacid tablets used in this experiment have been stored under laboratory
conditions and may be contaminated. Do not taste any materials used in this experiment.
3. When diluting the hydrochloric acid, remember to add the concentrated acid to water to
avoid splattering. Take care in handling the container as the dilution will generate heat.
Questions:

1. Which is appropriate in the dilution of acid and Why? sWater to acid or acid to water

27. 3. Write down the reactions involved.
4. Difference between normality and molarity
5. What is the colour change shown by phenolphthalein and methyl orange in acid-base media
6. Why a standard solution of HCl cannot be prepared directly?

28. EXPERIMENT 8
AIM
To determine the strength and percentage composition of sodium carbonate and sodium
bicarbonate in a given mixture
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula
Chemicals: NaHCO3, Na2CO3, HCl solution, phenolphthalein, methyl orange
THEORY
Neutralisation of Na2CO3 solution by strong acid (HCl) occurs in two steps:
Na2CO3 + HCl = NaHCO3 + NaCl (pH = 8.3 at the equivalence point)
NaHCO3 + HCl = NaCl + CO2 + H2O (pH = ~ 4 at the equivalence point)
The acid base indicator is to be so selected that its pH range for the colour change coincides
with the sudden sharp change of pH at the equivalence point. So at the first neutralisation
point (pH = 8.3), phenolphthalein shows its colour change from pink to colourless. At this
stage Na2CO3 consumes only half the amount of HCl required for complete neutralisation. If
methyl orange is added to this titrated solution and the titration with HCl is continued up to
second equivalence point, then this titre value corresponds to the amount of HCl required to
convert NaHCO3 to NaCl (i.e. NaHCO3 derived from Na2CO3 plus the amount of NaHCO3
present in the original mixture.
PROCEDURE
Estimation of Na2CO3 and NaHCO3 present: Weigh accurately about 1 gm of given
Na2CO3 and NaHCO3 mixture in 100 mL measuring flask and prepare a homogeneous
solution by dissolving in 100 mL distilled water. Pipette out 10 mL of alkali mixture solution
in 100 mL conical flask and add 2-3 drops of phenolphthalein indicator. Titrate it against
standardized N/10 HCl solution taken in the burette, till the pink colour of the solution
disappears. This gives the first end point of the titration. Now add 2-3 drops of methyl orange
indicator to the same titrating mixture, the solution turns yellow. Titrate again with
standardized HCl till the solution becomes pink or red colour. This gives the second end
point. Repeat thrice for concordant values.
OBSERVATION AND CALCULATIONS
Normality of HCl (NHCl) = 0.05 N
Estimation of Na2CO3 and NaHCO3 present in given mixture
Weight of sample taken = 1 gm

29. First: Alkali mixture vs HCl solution (phenolphthalein indicator)

Second: Same alkali mixture vs HCl solution (methyl orange indicator)

Volume of N/10 HCl used during first titration, x =
Volume of N/10 HCl used during second titration, y =
(Give proper justification and equations wherever necessary)
Na2CO3 ≡ 2x mL
NaHCO3 ≡ y – x mL
Na2CO3 in alkali mixture
NNa2CO3 × VNa2CO3 = NHCl × 2x
NNa2CO3 =
Amount of Na2CO3 in 100 mL =
NaHCO3 in alkali mixture
NNaHCO3 × VNaHCO3 = NHCl × (y – x)
NNaHCO3 =
Amount of NaHCO3 in 100 mL =
Percentage of Na2CO3 in given mixture =
Percentage of NaHCO3 in given mixture =
RESULTS
1 Strength of Na2CO3 and NaHCO3 in given mixture =
2 Percentage composition, Na2CO3 and NaHCO3 =

30. Precautions
1. Rinse the burette and pipette with the solutions to be used in them, to avoid the dilution
with distilled water
2. Never suck a strong acid or an alkali with the pipette, use a pipette bulb.
3. While transferring the solution to the flask, do not blow out the last drop of the solution
from the jet of the pipette.
Viva Questions
1. Reactions corresponds to the neutralization
2. How to distinguish the end points
3. Chemical structure of indicator and its changes upon the colour change
4. How to prepare 500 mL 0.2 N H2SO4 from concentrated acid (12 N).
5. Note down the equivalent mass of H2SO4 and Ca(OH)2.