Download CHEMISTRY LABORATORY-II

The Great Chemist
ALFRED NOBEL
CHEMISTRY
LABORATORY-II
-1-
WORK SHEET
Titration 1 : Standardization of EDTA
Burette solution : EDTA solution
Pipette solution : Standard hard water
Addition solution : 5 ml of ammonia buffer solution
Indicator
: Eriochrome Black -T
End point
: Change of colour from wine red to steel blue
Standard hard water Vs EDTA solution
Sl.
No
Volume of
Std. hard
water (mL)
Burette readings (mL )
Initial
Final
1 mL of standard hard water
Volume of
EDTA
(V1 mL )
Concordant
Value
= 1 mg of CaCO3
20 mL of standard hard water contains 20 mg of CaCO3
20 mL of standard hard water consumes --------------- (V1) mL of EDTA
Hence,
------------- (V1) ml of EDTA solution
1 mL of EDTA Solution
= 20 mg of CaCO3
= (20/V1) mg of CaCO3
= 20 / ……..( V1)
= …………. mg of CaCO3
-2-
Ex.No.
Date:
ESTIMATION OF HARDNESS OF WATER SAMPLE
BY EDTA METHOD
Aim
To Estimate the amount of total, permanent and temporary hardness in
the given sample of water provided that Standard hard water and a link
solution of Ethylene Diamine Tetra Acetic acid (EDTA).
Principle
The total hardness of water is estimated by titrating against EDTA using
Eriochrome Black –T indicator.
Structure of disodium salt of EDTA
NaOOCH2C
CH2COOH
N-CH2-CH2-N
HOOCH2C
CH2COONa
Calcium (Ca2+) and Magnesium (Mg2+) ions present in the hard water
form a stable complex ions with EDTA. The complexation of the reaction is
indicated by Eriochrome Black-T indicator. EBT forms an unstable wine-red
colour complex with Ca and Mg ions at pH 9-10.
Ca
2+
/ Mg
2+
NH4Cl -NH4OH buffer
[Ca2+ / Mg2+ +EBT]
less stable
(Wine-red colour complex)
+ EBT
pH 9-10
-3-
WORK SHEET
Titration 2 : Estimation of Total hardness
Burette solution : EDTA solution
Pipette solution : Sample hard water
Addition solution : 5 ml of ammonia buffer solution
Indicator
: Eriochrome Black -T
End point
: Change of colour from wine red to steel blue
Sample Water Vs Standard EDTA solution
Volume of
Volume of
Burette readings (mL )
Sl.
EDTA
Water
No
Initial
Final
(V2 mL )
Sample(mL)
Volume of EDTA consumed in titration 2
Concordant
Value
= ---------- (V2) mL
Total hardness of sample
water in 1 litre
= ( 20 / V1) xV2 x 1000 mg of CaCO3
20
= [20 / …….(V1 ) x …….( V2 ) x 1000 ]
20
= ----------- mg of CaCO3
Total hardness
= ---------- mg / L (or) ppm
-4-
When EDTA is added into the hard water, the metal ions form a stable
metal complex with EDTA by leaving the indicator. When all the metal ions
are taken by EDTA from the indicator metal ion complex, the wine red
colour changes into steel blue, which indicates the end point
[Ca2+ / Mg 2+ + EBT]+ EDTA
[Ca2+ / Mg 2+ +EDTA] +EBT
More stable
steel blue
(colourless complex)
colour
Procedure
Titration 1 Standardization of EDTA
A 50 mL burette is washed with distilled water and then rinsed with the
EDTA solution. It is then filled with the same EDTA solution up to zero
level without air bubbles. Initial reading of the burette is noted. 20 mL of
standard hard water solution is pipetted out into clean conical flask.5 mL of
ammonia buffer solution and two drops of EBT indicator is added. The
solution turns into wine- red colour. The solution is titrated against EDTA
solution taken in the burette. The end point is change of colour from wine
red to steel blue. The final burette reading is noted. The titration is repeated
to get concordant values. Let the volume of EDTA be V1 mL
Titration 2 Estimation of Total Hardness
The burette is filled with the same EDTA solution. 20 mL of the given
water sample is pipetted out into a clean conical flask. 5 mL of ammonia
buffer and two drops of EBT indicator is added. The solution is titrated
against standard EDTA solution taken in the burette. The end point is the
change of colour from wine red to steel blue. The final burette reading is
noted. The titration is repeated to get concordant values. Let the volume of
EDTA be V2 mL
-5-
WORK SHEET
Titration 3 : Estimation of Permanent hardness
Burette solution
Pipette solution
Addition solution
Indicator
End point
: EDTA solution
: Boiled Sample hard water
: 2 ml of ammonia buffer solution
: Eriochrome Black -T
: Change of colour from wine red to steel blue
Boiled Sample water Vs Standard EDTA solution
Sl.
Volume of
No
Boiled Sample
water (mL)
Burette readings (mL )
Initial
Final
Volume of EDTA consumed in titration 3
Volume of
Concordant
EDTA
Value
(V3 mL )
= ---------- (V3) mL
Permanent hardness of
Sample water in 1 litre = [(20 / V1) x V3 x 1000]
20
= [(20 / .……(V1) x …….(V3 )x 1000]
20
= ---------- mg of CaCO3
So,
Permanent hardness
= ---------- mg / L (or) ppm
Temporary hardness
= Total hardness – Permanent hardness
= -------------(ppm) – -------------(ppm)
= ------------- (ppm)
-6-
Titration 3 Estimation of Permanent Hardness
The burette is filled with same EDTA solution. 20 mL of the given
sample water taken in a 100 mL beaker and boiled until the volume is
reduced to 10 mL. Then it is cooled and filtered using filter paper. To the
filtrate, 2 mL of ammonia buffer and two drops of EBT indicator are added
and titrated against EDTA solution taken in the burette. The end point is the
change of colour from wine red to steel blue. The final Burette reading is
noted. Let the volume of EDTA be V3 mL
Result
(1) Total hardness of the given sample water
= ---------------- ppm
(2) Permanent hardness of the given sample water = ---------------- ppm
(3) Temporary hardness of the given sample water = ---------------- ppm
Viva voce questions
1. What is potable water?
2. Why disodium salt of EDTA is chosen for determination of hardness?
3. Why does the colour of the solution change from wine red to steel
blue?
4. Why does hard water not lather with soap?
5. Which is the best method of hardness determination and why?
6. What is buffer solution? How and why the pH of the solution is
maintained at 10?
7. Which salt produces temporary and permanent hardness?
8. Draw the structure of EDTA and EDTA – metal ion complex.
9. What are units of hardness?
10. Draw the structure of EBT
11. What is the relationship between mg/l and ppm?
-7-
WORK SHEET
Standard H2SO4 Vs Water Sample :
Burette readings (mL)
Vol.of
Final
Sample
Sl.
Water
No
(mL)
Initial
At Phenol- At Methyl
phthalein
orange
end point
end point
(P)
(M)
-8-
Volume of
Concordant
H2SO4
value
(P)
(M)
(P)
(M)
mL
mL
mL
mL
Ex.No.
Date:
DETERMINATION OF ALKALINITY
OF A WATER SAMPLE
Aim
To determine the different types of alkalinity of a given sample of
water.
Principle
Alkalinity in water is due to the presence of OH-, CO32-, HCO3- ions.
The alkalinity of a given sample of water can be obtained by neutralizing the
above mentioned ions with standard H2SO4. Titrating given sample of water
at a pH of 8.3 or till the decolourization of phenolphthalein indicator will
indicate complete neutralization of OH- ions and half of CO32- ions. Titrating
the same sample of water at pH of 4.4 or till a sharp color change from
yellow to pink on methyl orange indicator, indicates the total alkalinity i.e.
the amount of
OH-, CO32- and HCO3- present in the given sample.
OH- + H+
H2O
CO32- + H+
HCO3- + H+
HCO3–
(H2CO3)
H2O + CO2
OH- and HCO3- ions cannot exist in water together because they combine
instantaneously to form CO32-ions.
OH- + HCO3-
CO32- + H2O
NaOH + NaHCO3
Na2CO3 +H2O
It is for this same reason, the three ions OH-, CO32- and HCO3- cannot exist
together.
-9-
WORK SHEET
Calculation:
Phenolphthalein alkalinity (mg/ L) as CaCO3
P ml of H2SO4 X Normality of H2SO4 X 50 X 1000
= ----------------------------------------------------------------Volume of water sample taken
…… X …… X 50 X 1000
= -----------------------------------……..
= --------------------- ppm
Calculation:
Methyl orange alkalinity (mg/ L) as CaCO3 (Total alkalinity)
M ml of H2SO4 X Normality of H2SO4 X 50 X 1000
= ------------------------------------------------------------------Volume of water sample taken
…… X …… X 50 X 1000
= -----------------------------------……..
= --------------------- ppm
-10-
Procedure
The burette is washed with the distilled water and then rinsed with
H2SO4. It is then filled with H2SO4 upto the zero mark. 20 mL of the given
water sample is pipetted out into a clean conical flask. To this, 2 drops of
phenolphthalein indicator is added and the solution is titrated against H2SO4
taken in the burette. The end point is just disappearance of pink colour. The
corresponding burette reading is noted and it is denoted as phenolphthalein
end point (P).
To the same solution, two drops of methyl orange indicator is added
and the titration is continued. The end point is colour change from pale
yellow to pale pink. This end point indicates complete neutralization of
alkalinities present in water sample and is known as methyl orange end point
(M)
Result
a) Amount of Phenolphthalein alkalinity = ----------- ppm
b) Amount of Methyl orange alkalinity
-11-
= ----------- ppm
WORK SHEET
Observation
Volume of
K2Cr2O7
solution
added (mL)
Potential of
the test cell
Ecell (mV)
∆E
∆V
∆E/∆V
Average
volume of
K2Cr2O7 (mL)
Where, ∆V = difference in two consecutive volumes of titrant added in mL
∆E = difference in two consecutive reading in millivolts
-12-
Ex.No.
Date:
POTENTIOMETRIC TITRATION
(ESTIMATION OF FERROUS ION)
Aim
To estimate the amount of ferrous ions present in the whole of the given
solution potentiometrically by titrating it against standard potassium
dichromate solution (0.05N).
Principle
When ferrous ion is titrated against potassium dichromate in acidic
medium, it gets oxidized into ferric ion.
6FeSO4 +K2Cr2O7+ 7H2SO4
K2SO4 + Cr2 (SO4)3+ 3 Fe2 (SO4)3 + 7 H2O
The ferrous- ferric system establishes equilibrium as
Fe2+
Fe3+ + e-
and the electrode potential is given by the Nernst equation
E= EO – (0.0591/1) log (Fe2+/ Fe3+)
The reduction potential value depends on the concentration of ferrous
and ferric ions present. While adding dichromate, ferrous ions are converted
into ferric ions and hence the potential increase gradually and then steeply at
the end point.
A plot of potential Vs. Volume of potassium dichromate is used to
assess the end point. More accurate titre value can be derived from
derivative graph.
-13-
WORK SHEET
Calculation:
Volume of K2Cr2O7
V1
Strength of K2Cr2O7
N1 = ------------- (N)
Volume of Fe2+ taken V2
= ------------- (mL) from graph
= 20 mL
Strength of Fe2+ taken
N2 = ------?------ (N)
Strength of Fe2+
N2 = [-------- (V1) x -------- (N1)] /--------- (V2)
Strength of Fe2+
= ----------------- N
Amount of ferrous ion present in the whole of the given solution
= [Normality x Eq.weight of Fe2+ x 100] / 1000
= [----------- (N2) x ----------- x 100] / 1000
= ------------------- g
(Equivalent Weight of Fe2+ = 55.84)
-14-
Procedure
The given ferrous ion solution is made up to 100mL in a standard
flask using distilled water. 20 mL of the made up solution is pipetted out into
a clean beaker and a test tube full of dilute sulphuric acid is added. Platinum
(inert indicator electrode) and calomel (secondary reference electrode)
electrodes are immersed in solution and connected to the potentiometer.
Potassium dichromate solution is added from the burette in 1 mL portions.
The solution is stirred well and after each addition, the emf value is noted.
Near the end point, emf changes rapidly. Therefore, add 0.2 mL increments
near the end point and note the sudden increase in emf. Continue the titration
with 1 mL addition and take 3 to 4 readings.
The potential (E) is plotted against the volume of potassium dichromate
added. The end point of the titration is obtained from the midpoint of
inflection. The accurate end point can be obtained from the peak point by
plotting ∆E/∆V against the average volume of K2Cr2O7 added. The volume
of K2Cr2O7 corresponding to the peak of the curve is the end point of the
titration.
-15-
WORK SHEET
-16-
For better accuracy, the end point obtained from the differential plot is
taken for calculation. From the titre value, the normality of ferrous ion can
be calculated. From this value, the amount of ferrous ion present in the
whole of the given solution is calculated.
Experimental Setup
Note
At the start of the experiment (Before adding K2Cr2O7), theoretically
there should not be any Fe3+ions in the solution. So, no potential is
developed for the Fe2+/ Fe3+equilibrium. However the meter reads some
sensible potential even before the addition of K2Cr2O7 because of the
presence of Fe3+ions as impurities in the ferrous sample. So, the zero should
not be recorded in a potentiometric titration.
Result
The amount of the ferrous ion present in the
whole of the given solution is ------------ g
-17-
WORK SHEET
Observation
BaCl2 Vs Na2SO4
Sl.
No
Volume of Na2SO4 added
(mL)
Observed Conductance (ohm-1)
-18-
Ex.NO.
Date:
CONDUCTOMETRIC TITRATION USING BaCl2 Vs Na2SO4
Aim
To determine the amount of BaCl2 present in one litre of the given
solution by conductometric titration using standard Na2SO4 of -----------N
Principle
Solution of electrolytes conducts electricity due to the presence of ions.
Since specific conductance of a solution is proportional to the concentration
of ions in it, conductance of the solution is measured during titration.
In the precipitation titration, the ions are converted to insoluble
precipitate, which will not contribute in the conductance.
When Na2SO4 is added slowly from the burette to the solution of BaCl2,
BaSO4 gets precipitated while the chloride ions remain unchanged.
[Ba2+ + 2Cl] + [2Na+ SO42-]
BaSO4
+ 2Na+ + 2Cl
Unionized
The Ba2+ ions in the solution are replaced by free Na+ ions. Since the
mobility of Na+ ions is less than that of Ba2+ ions the conductance of the
solution decrease.
After the end point, when all the Ba2+ ions are replaced, further addition
of Na2SO4 increases the conductance. This is due to the increase of Na+ and
SO42- ions in the solution.
-19-
WORK SHEET
Calculation
V1N1 = V2N2
Volume of Na2SO4
V1 = ---------- mL (from graph)
Strength of Na2SO4
N1 = ---------- N
Volume of BaCl2
V2 = --------- mL
Strength of BaCl2
N2 = -----?----- N
N2
= (V1 X N1) / V2
= [……..( V1) X ……..( N1)] / …….( V2)
Strength of unknown BaCl2 (N2) = --------- N
Total amount of BaCl2 present in
the given solution = [N2 x Equivalent of BaCl2 x 50]/1000
= [………..X ……… X 50]/1000
= ………..g
(Equivalent Weight of BaCl2 = 122.14)
-20-
Procedure
The burette is filled with Standard N/10 Na2SO4 solution upto the zero
level. 50mL of the given BaCl2 solution is pipetted out into a clean 100mL
beaker. The conductivity cell is placed in it and then diluted to 50mL by
adding conductivity water. The two terminals of the cell are connected with
a conductivity bridge.
Now 1ml of Standard N/10 Na2SO4 from the burette is added to the
solution, taken in the beaker, stirred, and then conductivity is measured. This
is continued upto the end point. (The conductivity is going on decreasing
upto to the end point). After the end point, again Standard N/10 Na2SO4 is
gradually added and few more reading are noted.
Thus the conductivity is continuously measured for each addition of
Standard N/10 Na2SO4 and are tabulated. Now the graph is plotted between
the volume of Standard N/10 Na2SO4 and conductivity. From the graph, end
point is noted and hence amount of BaCl2 present in 1 litre is calculated.
Result
The amount of BaCl2 present in 1litre of the given solution ---------g
-21-
WORK SHEET
Titration I : Standardization of EDTA
Burette solution : EDTA
Pipette solution : Standard ZnSO4
Addition solution : 5 ml of ammonia buffer solution
Indicator
: Eriochrome Black -T
End point
: Change of colour from wine red to steel blue
Standard ZnSO4 Vs EDTA solution
Sl. Volume of
No ZnSO4 (mL)
Burette readings (mL )
Initial
Final
Volume of
EDTA
(mL )
Concordant
Value (mL)
Calculations :
Volume of ZnSO4
V1 = 20 ml
Strength of ZnSO4
N1 = 0.01 N
Volume of EDTA
V2 = ……
Strength of EDTA
N2 = ……?
= (V1 x N1) / V2
= (…… x …….) / …..
= ……..N
-22-
Ex.No.
Date:
ESTIMATION OF COPPER IN BRASS BY EDTA METHOD
Aim
To estimate the amount of copper present in the given solution being
supplied with a standard zinc sulphate and EDTA solution
Principle
Brass is an alloy of Cu and Zn. It also contains small amounts of lead,
tin or aluminium. In this method, brass is dissolved in concentrated nitric
acid so as to bring copper into cupric ions. Fast sulphon black- F is added to
this solution to impart a purple color by the formation of a weak complex.
Cu2+ + FSB-F [ Cu2+ - FSB –F]
Purple complex
When EDTA is adder to this, FSB- F in the complex is replaced by EDTA as
per the following reaction
[ Cu2+ - FSB –F] + EDTA [ Cu2+ - EDTA] + FSB –F
Green Color
Procedure
Titration I :
Standardisation of EDTA solution
The given EDTA solution was taken in the burette. 20 mL of the
standard zinc sulphate solution was pipetted out into a clean and dry conical
flask. To this added a 5 ml of ammonia buffer and 2 drops of EBT indicator.
The solution was then titrated against the EDTA solution taken in the burette
till the color changes from wine red to steel blue. Titrations were repeated
till a concordant value was obtained.
-23-
WORK SHEET
Titration II : Estimation of Copper
Burette solution : EDTA
Pipette solution : Copper Sulphate
Addition solution : 5 ml of ammonia buffer solution
Indicator
: FSB -F
End point
: Change of colour from Purple to Green
Standard EDTA Vs CuSO4 Solution
Burette readings (mL )
Sl. Volume of
No CuSO4 (mL)
Initial
Final
Volume of
EDTA
(mL )
Concordant
Value (mL)
Calculations :
Volume of EDTA
V1
= ……..
Strength of EDTA
N1
= ……..
Volume of CuSO4
V2
= 20 mL
Strength of CuSO4
N2
= ……? (V1 x N1) / V2
= (…… x …….) / …..
= ……..N
Hence, Weight of copper present in the
whole of the given solution
= 63.54x (----------N2) = -----------g
10
The percentage of copper in brass
= [Wt.of copper / Wt.of brass] x 100
= [………… / 1] X 100
= -------------%
(Equivalent weight of copper = 63.54)
-24-
Titration II :
Estimation of copper in brass
20 mL of the given copper sulphate solution was pipetted out into a
clean and dry conical flask. To this added a 5 mL of ammonia buffer and
two drops of Fast sulphon black - F indicator. The solution was then titrated
against the EDTA solution taken in the burette till the color changes from
purple to green. Titrations were repeated till a concordant value was
obtained.
Result
Amount of copper present in the whole of the given solution = ……
Percentage of copper in the given brass sample
-25-
= ……%