Physical_Chemistry_Lab_manual_for_BSc_Chemistry.pdf

Physical Chemistry Lab manual for BSc Chemistry students. This includes acid-base titrations, redox titrations, Argentometric titrations. A brief description of theory and calculation are provided.

1. LAB MANUAL
PHYSICAL CHEMISTRY

FOR

B. Sc. CHEMISTRY

2. Index
1. Estimation of carbonate and hydroxide present together in a mixture
2. Estimation of carbonate and bicarbonate present together in a mixture
3. Estimation of free alkali present in different soaps/detergents
4. Estimation of oxalic acid using standardized KMnO4 solution
5. Estimation of Cl
-
by Mohr’s method
6. Estimation of Cu (II) using sodium thiosulphate solution (iodimetrically)
7. Estimation of Fe (II) with K2Cr2O7 using diphenylamine as internal indicator

3. EXPERIMENT 1
AIM
To determine the strength and percentage composition of sodium carbonate and sodium
hydroxide in a given mixture
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula
Chemicals: NaOH, Na2CO3, HCl solution, phenolphthalein, methyl orange
THEORY
The carbonate and hydroxide ions react with hydrogen ions as follows:
NaOH + HCl → NaCl + H2O
Na2CO3 + HCl → NaHCO3 + NaCl ; NaHCO3 + HCl → NaCl + H2O + CO2
The pKa1 and pKa2 values of H2CO3 are quite distinct and there occur a sharp pH
change at these two distinct regions. The first corresponding to the formation of bicarbonate
(pH 8 to 10) and the second due to complete neutralization at pH 4-6. The first is roughly in
the pH range in which colour of phenolphthalein changes from red to colourless and the
second is that at which methyl orange changes from yellow to orange red.
When both sodium carbonate and sodium hydroxide present together in a solution, a
titration using phenolphthalein gives the volume at the equivalence point corresponding to
sodium hydroxide plus half the carbonate and the volume obtained with methyl orange
corresponds to the total alkali. The individual sodium carbonate and hydroxide concentrations
can be calculated from the data.
PROCEDURE
1. Standardization of HCl: Pipette out 25 mL of N/10 Na2CO3 solution in 250 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.
2. Estimation of Na2CO3 and NaOH present: Weigh accurately about 1 gm of given
NaOH and Na2CO3 mixture in 100 mL measuring flask and prepare a homogeneous

4. solution by dissolving in 100 mL distilled water. Pipette out 25 mL of alkali mixture
solution in 250 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
1. Standardization of HCl
Weight of empty weighing bottle, m =
Weight of weighing bottle + sodium carbonate, n =
Weight of sodium carbonate, m-n =
Normality of sodium carbonate, N1 =






=

Sl
No:
Vol of Na2CO3 solution
taken (V mL)
Burette Reading (mL) Vol of HCl used
(V2 mL) Initial Final
1
2
3

V2 × NHCl = VN1 (Na2CO3)
Normality of HCl (NHCl) =






2. Estimation of Na2CO3 and NaOH present in given mixture
Weight of sample taken = 1 gm
First: Alkali mixture vs HCl solution (phenolphthalein indicator)
Sl
No:
Vol of alkali mixture
solution taken (V3 mL)
Burette Reading (mL) Vol of standardized
HCl used (V4 mL) Initial Final
1
2
3

5. Second: Same alkali mixture vs HCl solution (methyl orange indicator)
Sl
No:
Vol of alkali mixture
solution taken (V3 mL)
Burette Reading (mL) Vol of standardized
HCl used (V5 mL) Initial Final
1
2
3


Volume of N/10 HCl used during first titration, V4 =
Volume of N/10 HCl used during second titration, V5 =
(Give proper justification and equations wherever necessary)
V4 mL ≡ NaOH +


Na2CO3
V5 mL ≡


Na2CO3
Na2CO3 ≡ 2V5 mL
NaOH ≡ V4 - V5 mL
NaOH in alkali mixture
NNaOH × V3 = NHCl × (V4 - V5)
NNaOH =
MNaOH =
Amount of NaOH in 100 mL =
Na2CO3 in alkali mixture
NNa2CO3 × V3 = NHCl × 2V5
NNa2CO3 =
MNa2CO3 =
Amount of Na2CO3 in 100 mL =
Percentage of NaOH in given mixture =
Percentage of Na2CO3 in given mixture =
RESULTS
1 Strength of NaOH and Na2CO3 in given mixture =
2 Percentage composition, NaOH =
Na2CO3 =
Notes:
1. Preparation of standard N/10 sodium carbonate solution: Weigh accurately 1.4 gm of
AR grade sodium carbonate and transfer it into 250 mL measuring flask. Dissolve it
completely and make up to the volume using distilled water

6. 2. Preparation of standard N/10 HCl solution: To a 1000 mL measuring flask of 100 mL
distilled water, add 8.4 mL conc HCl drop wise. Shake the flask thoroughly and then
make up to the volume by adding distilled water.
Questions:
1. Which is appropriate in the dilution of acid and Why? Water to acid or acid to water
2. What is the purpose of adding two different indicators?
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?

7. EXPERIMENT 2
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
1. Standardization of HCl (Refer to Experiment 1)
2. 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 25 mL of alkali mixture
solution in 250 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

8. 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
1. Standardization of HCl
Weight of empty weighing bottle, m =
Weight of weighing bottle + sodium carbonate, n =
Weight of sodium carbonate, m-n =
Normality of sodium carbonate, N1 =






=

Sl
No:
Vol of Na2CO3 solution
taken (V mL)
Burette Reading (mL) Vol of HCl used
(V2 mL) Initial Final
1
2
3

V2 × NHCl = VN1 (Na2CO3)
Normality of HCl (NHCl) =





2. Estimation of Na2CO3 and NaHCO3 present in given mixture
Weight of sample taken = 1 gm
First: Alkali mixture vs HCl solution (phenolphthalein indicator)
Sl
No:
Vol of alkali mixture
solution taken (V3 mL)
Burette Reading (mL) Vol of standardized
HCl used (V4 mL) Initial Final
1
2
3



Second: Same alkali mixture vs HCl solution (methyl orange indicator)
Sl
No:
Vol of alkali mixture
solution taken (V3 mL)
Burette Reading (mL) Vol of standardized
HCl used (V5 mL) Initial Final
1
2
3

9. Volume of N/10 HCl used during first titration, V4 =
Volume of N/10 HCl used during second titration, V5 =
(Give proper justification and equations wherever necessary)
Na2CO3 ≡ 2V4 mL
NaHCO3 ≡ V5 – V4 mL
Na2CO3 in alkali mixture
NNa2CO3 × V3 = NHCl × 2V4
NNa2CO3 =
MNa2CO3 =
Amount of Na2CO3 in 100 mL =
NaHCO3 in alkali mixture
NNaHCO3 × V3 = NHCl × (V5 – V4)
NNaHCO3 =
MNaHCO3 =
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 =
NaHCO3 =
Notes:
Preparation of standard N/10 sodium carbonate solution and Preparation of standard N/10
HCl solution (Refer Experiment 1).
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.

10. EXPERIMENT 3
AIM
To determine the percentage of sodium carbonate present in washing soda
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula
Chemicals: Washing soda, Na2CO3, HCl solution, phenolphthalein, methyl orange
THEORY
Washing soda is chemically the hydrated sodium carbonate (Na2CO3.10H2O). The sodium
carbonate content can be determined by titrating against standard HCl solution using methyl
orange as indicator.
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
The completion of the reaction is indicated by the generation of red colour due to the
presence of indicator.
PROCEDURE
1. Standardization of HCl (Refer to Experiment 1)
2. Estimation of Na2CO3 in washing soda: Weigh accurately 2 gm of washing soda
and transfer it to 100 mL measuring flask with the aid of distilled water. Stopper the
flask and shake thoroughly to dissolve the sample completely. Make up to the
volume. Pipette out 10 mL of the solution to conical flask and add 2-3 drops of
methyl orange indicator. Titrate against standard N/10 HCl solution taken in the
burette till the yellow colour of the solution changes to pink or red colour. This is the
endpoint. Note down the volume of HCl used and repeat thrice for concordant values.
OBSERVATION AND CALCULATIONS
1. Standardization of HCl
Weight of sodium carbonate, x =
Normality of sodium carbonate, N1 =






=

11. Sl
No:
Vol of Na2CO3 solution
taken (V mL)
Burette Reading (mL) Vol of HCl used
(V2 mL) Initial Final
1
2
3

V2 × NHCl = VN1 (Na2CO3)
Normality of HCl (NHCl) =





2. Estimation of Na2CO3 present in washing soda
Weight of sample taken = 2 gm
alkali mixture vs HCl solution (methyl orange indicator)
Sl
No:
Vol of washing soda
solution taken (V1 mL)
Burette Reading (mL) Vol of standardized
HCl used (V2 mL) Initial Final
1
2
3


Nwashing soda × V1 = NHCl × V2
Nwashing soda =





Strength of Na2CO3 =
Amount of Na2CO3 present in 100 mL washing soda solution =

Therefore, % of Na2CO3 in washing soda =

RESULT
The percentage of Na2CO3 in given washing soda sample =

12. Notes:
Preparation of standard N/10 sodium carbonate solution and Preparation of standard N/10
HCl solution (Refer Experiment 1).
Questions:
1. Define percentage composition of ‘a’ in a mixture of ‘a’ and ‘b’
2. How to determine the strength of a solution, if N is provided
3. Why some solutions have to be standardized before using it for titration?
4. Write the chemical formula for washing soda
5. Is NaOH a primary standard? Justify your answer

13. EXPERIMENT 4
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↑

14. 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
=

15. 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

16. EXPERIMENT 5
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)

17. 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

18. 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

19. 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)

20. 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.

21. 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

22. 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?

23. 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

24. EXPERIMENT 7
AIM
To estimate the amount of ferrous ions present in 100 ml of the given solution by using
approximately 0.05 N potassium dichromate solution.
APPARATUS AND CHEMICALS
Apparatus: Conical flasks, burette, pipette, measuring flask, spatula
Chemicals: Potassium dichromate, Sulphuric acid, Syrupy phosphoric acid, Diphenylamine
indicator, Ferrous ion sample
THEORY
The estimation is based on redox titration. Ferrous iron is estimated by dichrometry in
presence of H3PO4 using diphenylamine as internal indicator. Potassium dichromate oxidises
ferrous iron to ferric iron in an acidic medium.

Complete ionic equation:

Under this condition K2Cr2O7 quantitatively oxidizes Fe
2+
to Fe
3+
:
 1/6 mol Cr2O7
2
 1 mol Fe
2+
= 1 equivalent
Diphenylamine is the internal indicator as it is added to the reaction mixture. The end
point is marked with an intense blue violet colouration. The reduction potential value of this
system is E°red = +0.76 V which is very near to ferrous-ferric system (E°red = + 0.77 V). Thus
phosphoric acid is required in this titration as it reacts with the product Fe
3+
ion to form the
complex ion [Fe(HPO4)]
+
, thus lowering the formal potential of the Fe(III)/Fe(II) system
thereby increasing its reducing power. Thus the end point is sharper.
The indicator action of diphenylamine can be understood as

25. PROCEDURE
1. Preparation of standard potassium dichromate: Weigh out approximately 0.245 g
of potassium dichromate and transfer into 100 ml standard (volumetric) flask.
Dissolve the dichromate in a small quantity of distilled water, and make up to the
mark. The contents in the flask are shaken well for uniform concentration. Calculate
the normality of potassium dichromate.
2. Estimation of Ferrous ion: Rinse the pipette with given ferrous solution and pipette
out 10 ml into a clean conical flask add 10ml of the acid mixture (sulphuric acid and
phosphoric acid), and 3-4 drops of diphenylamine indicator. Fill the burette with the
prepared potassium dichromate solution after rinsing it, with the same. Titrate the
solution in the conical flask against the standard potassium dichromate from the
burette till the colour changes to blue violet. Repeat the titrations for concordant
values.
OBSERVATION AND CALCULATIONS
1. Preparation of Standard K2Cr2O7 solution
Weight of K2Cr2O7, x =
Normality of K2Cr2O7, N1 =






=
2. Estimation of Ferrous ion
Ferrous sample vs K2Cr2O7 solution (diphenylamine indicator)
Sl
No:
Vol of sample solution
taken (V1 mL)
Burette Reading (mL) Vol of K2Cr2O7
used (V2 mL) Initial Final
1
2
3


N1 × V1 (Sample) = N2 × V2 (K2Cr2O7 solution)
Nsample =



Amount of Ferrous iron present in the whole of the given solution = N2 x 55.85 =
RESULT
Amount of Ferrous iron present in the whole of the given solution =

26. Precautions
1. Prepare indicator solution in concentrated sulphuric acid
2. Freshly prepare standard dichromate solution should be used
3. Dichromate is toxic to the environment. After the experiment, excess dichromate and
the analyte are to be discarded in the designated waste bottles located on the reagent
bench.
Questions:
1. Name other internal indicators used
2. Example of external indicator
3. What is the colour change at the endpoint of titration?
4. What are the causes of series of colour change from initial to end of the titration?
5. Why Syrupy phosphoric acid is added to the reaction media?
6. Generally, diphenylamine is not a suitable indictor for the determination of ferrous
ion using dichrometry. Why?
7. How to prepare diphenylamine indicator solution?
8. How is it possible to make diphenylamine as a suitable indicator for the determination
of ferrous ion using dichrometry?
9. Equivalent weight of Fe.