1- Neutralization Reactions
Experimental (1)
Preparation of approximately 0.1 N of HCl and standardization by 0.1 N Na2CO3.
A) Pre-lab assignments:
Answer the following questions:
1- What the difference between primary and secondary standard substances?
2- Calculate the volume of conc. HCl required for preparing 250 ml 0.1 N.
3- Calculate the weight of Na2CO3 required for preparing 100 ml 0.1 N.
(10 minutes)
B) Theory:
Sodium carbonate can be estimated by displacement titration with a mineral acid like HCl:
Na2CO3 + 2HCl
→
2NaCl + CO2 + H2O
This reaction proceeds into two stages:
Na2CO3 + HCl
→
NaHCO3 + NaCl
NaHCO3 + HCl
→
H2O + CO2 + NaCl
pH (11.5 – 8.31) (ph.ph)
pH (8.31 – 3.8)
(M.O)
It thus follows that when phenolphthalein is used in the neutralization of HCl with sodium carbonate, the
volume of acid used will be equivalent to half of the carbonate. When methyl orange is used in this titration,
the volume of acid used will be equivalent of all the carbonate. Methyl orange is generally used in this case
as phenolphthalein is sensitive to carbon dioxide.
C) Chemicals required:
1- Na2CO3 solid.
2- Conc. HCl.
3- Methyl Orange indicator.
D) Procedure:
1- Transfer known volume of the sodium carbonate solution, with a pipette, to a conical flask then add
one or two drops of methyl orange to this solution.
2- Add the acid from the burette gradually with continuous swirling of the solution in the conical flask,
and near the end point, the acid is added drop by drop. Continue the addition of the acid until the
color of the solution passes from yellow to faint red.
3- Repeat the experiment three times and tabulate your results then take the mean of the three readings
E) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
1
F) Calculations:
2
Experimental (2)
Determination of normality and strength of sodium hydroxide solution by a standard
solution of hydrochloric acid.
A) Theory:
Hydrochloric acid reacts with sodium hydroxide according to the equation:
HCl + NaOH → NaCl + H2O
The equivalent weight of both HC1 and NaOH is equal to their molecular weights and as both the acid
and alkali are strong. any indicator may be used.
C) Chemicals required:
1- Hydrochloric acid solution (standard).
2- Sodium hydroxide solution of unknown normality.
D) Procedure:
1- Transfer known volume of the sodium hydroxide solution, with a pipette, to a conical flask then add
one or two drops of phenolphthaline. The solution has the pink color.
2- Add the acid from the burette gradually with continuous swirling of the solution in the conical flask,
and near the end point, the acid is added drop by drop. Continue the addition of the acid until the
color of the solution discharged.
3- Repeat the experiment three times and tabulate your results then take the mean of the three readings.
E) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
F) Calculations:
3
Experimental (3)
Determination of the normality and strength of sodium carbonate and sodium
bicarbonate mixture by standard HCl.
A) Theory:
-
Total carbonate and bicarbonate is determined by titration with standard acid using M.O. as
indicator.
V1(acid) ≡ all CO3-- + all HCO3-
-
(M.O.)
The carbonate is determined by similar titration but using Ph.Ph. as indicator
V2(acid) ≡ half CO3--
volume of acid ≡ CO3-- = 2V2
(Ph.Ph.)
-
volume of acid ≡ HCO3 = V1 - 2V2
B) Chemicals required:
1- Hydrochloric acid solution (standard).
2- A mixture of carbonate and bicarbonate.
C) Procedure:
1- Dilute 10 ml of the mixture with distilled water and titrated with 0.1 N HCl using methyl orange as
indicator. This is a measure of Na2CO3 and NaHCO3.
2- Repeat the experiment three times and tabulate your results then take the mean of the three readings.
3- Dilute another 10 ml of the mixture with distilled water and titrated with 0.1 N HCl using Ph.Ph. as
indicator. Repeat the experiment three times and tabulate your results then take the mean of the three
readings. Multiply the volume of standard acid by two, and this would be a measure of the carbonate.
5- Subtract volume of standard acid equivalent to carbonate from the used in step 1, and this would be a
measure of the bicarbonate content.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
4
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
E) Calculations:
5
Experimental (4)
Determination of the normality and strength of sodium carbonate and sodium hydroxid
mixture by standard HCl.
A) Theory:
- Total carbonate and hydroxide is determined by titration with standard acid using M.O. as indicator.
V1(acid) ≡ all CO3-- + all OH-
(M.O.)
- The carbonate is determined by similar titration but using Ph.Ph. as indicator
V2(acid) ≡ half CO3-- + all OH-
(Ph.Ph.)
--
- Difference between the two reading ≡ half CO3 = V1 – V2
volume of acid ≡ all CO3-- = 2(V1 - V2)
volume of acid ≡ all OH- = V1 - 2(V1 - V2)
B) Chemicals required:
1- Hydrochloric acid solution (standard).
2- A mixture of carbonate and hydroxide.
C) Procedure:
1- Dilute 10 ml of the mixture with distilled water and titrated with 0.1 N HCl using methyl orange as
indicator. This is a measure of Na2CO3 and NaOH.
2- Repeat the experiment three times and tabulate your results then take the mean of the three readings.
3- Dilute another 10 ml of the mixture with distilled water and titrated with 0.1 N HCl using Ph.Ph. as
indicator. Repeat the experiment three times and tabulate your results then take the mean of the three
readings.
4- Multiply the difference between the two reading by two, and this would be a measure of the carbonate.
5- Subtract volume of standard acid equivalent to carbonate from the used in step 1, and this would be a
measure of the hydroxide content.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
6
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
E) Calculations:
7
Experimental (5)
Determination of ammonia in ammonium salt.
A) Theory:
When an ammonium salt is heated with sodium hydroxide, ammonia gas evolves according to the
following equation: NH4Cl + NaOH → NaCl + H2O + NH3
Ammonia may be directly estimated by passing the gas in a standard solution of sulphuric acid then
titrating the excess acid with standard sodium hydroxide solution. This gives us the amount of sulphuric acid
which reacts with ammonia, or in other words, the amount of ammonia that evolved in the reaction.
Ammonia may be estimated by an indirect method by adding an excess of a measured volume of sodium
hydroxide solution to a known volume of the ammonium salt solution then boiling the mixture until the
evolution of ammonia gas ceases, The residual sodium hydroxide is then titrated with a standard acid.
B) Chemicals required:
1- Ammonium chloride solution
2- Sodium hydroxide solution 0.1 N
3- Hydrochloric acid solution 0.1 N
C) Procedure:
1- Transfer with a pipette 10 ml of the ammonium chloride solution to the conical flask.
2- Add 20 ml of sodium hydroxide solution then cover the neck of the flask with a small funnel to avoid
splashing of the solution during boiling.
3- Boil the mixture until the evolution of ammonia gas ceases; this is indicated by exposure of a filter
paper wetted with mercurous nitrate solution to the evolved vapors. If the paper does not turn to black,
it indicates that no ammonia vapors are evolved.
4- Cool the conical flask and add to its contents one or two drops of methyl orange then titrate with the
standard HCl.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
8
E) Calculations:
9
Experimental (6)
Analysis of a commercial sample of phosphoric acid
A) Theory:
- Phosphoric acid is a tribasic acid which dissociates in solution in three stages:
H3PO4
H+ + H2PO4-
K1 = 7.5 x10-3
H2PO4-
H+ + HPO4--
K2 = 6.2 x 10 -8
HPO4--
H+ + PO4---
K3 = 5 x 10 -13
- It is obvious that the difference between the dissociation constants K1, K2 and K3 is too 1arge.
Accordingly, the acid may behave as a mixture of three monobasic acids with different
dissociation constants as cited above, and the process of neutralization may take place in three stages,
where each has a separate end point. Every such end point is characterized by a certain pH value.
The first end point pH = 4.6
(methyl orange range).
The second end point pH = 9.7
(phenolphthalein range).
The third end point pH = 12.6
(no indicator could be used) .
B) Chemicals required:
1- A sample of commercial phosphoric acid.
2- Sodium hydroxide solution 0.1 N
C) Procedure:
1- Weight accurately about 2 gm, of commercial phosphoric acid in a small weighing bottle.
2- Transfer the weighed sample quantitatively to a 250 ml. measuring flask with distilled water and
complete to the mark then shake the solution carefully to get thorough mixing.
3- Transfer 25 ml with the aid of a pipette, to a conical flask then titrate with 0.1 N sodium hydroxide
solution in presence of methyl orange.
4- Repeat the experiment several times and tabulate your results. The amount of alkali used in this case
(V1) equals 1/3 the acid and the amount of alkali required to neutralize all the acid = 3 V1 ml.
5- In another experiment, titrate 25 ml. of phosphoric acid solution using phenolphthalein. The amount of
alkali used in this case (V2) equals 2/3 the acid and the amount required to neutralize all the acid
= 3/2 V2 ml.
10
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
E) Calculations:
11
2- Precipitation Reactions
Experimental (1)
Determination of chloride
1- By Mohrʼs method:
A) Theory:
This method is used for the determination of halides in neutral medium using potassium chromate
as indicator. Owing to the solubility product of Ag2CrO4 being more than that of silver halides. when
a neutral halide solution is titrated with standard silver nitrate solution, all the halide is first
precipitated as its silver salt then any excess of silver nitrate will produce a brick red precipitate of
silver chromate.
B) Chemicals required:
1- Sodium chloride solution of unknown normality.
2- Silver nitrate solution 0.05 N
3- potassium chromate indicator.
C) Procedure:
1- Prepare 100 ml of a standard silver nitrate solution 0.05 N.
2- Pipette 10 ml of sodium chloride solution into a 250 ml conical flask.
3- Add 1 ml of a 2% neutral K2CrO4 solution and titrate with AgNO3 solution
4- Continue the titration, dropwise, until a faint, but distinct, brick red color is formed and does not
disappear on vigorous shaking.
5- Repeat the experiment several times and tabulate your results.
6- A blank experiment must be done, as an excess silver nitrate must be added to produce detectable
quantity of precipitate to human eye.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
12
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
E) Calculations:
13
Experimental (2)
Analysis of a mixture of sodium hydroxide and sodium chloride.
A) Theory:
Sodium hydroxide is estimated in the mixture by direct titration against standard hydrochloric acid.
Sodium
chloride
may
be
estimated
by
Volhard's
method
by
addition
of
nitric
acid to the mixture then adding a known excess of standard silver nitrate to precipitate silver chloride
followed
by
titration
of
the
excess
of
silver
nitrate
against
standard
thiocyanate
solution in the presence of ferric alum indicator.
B) Chemicals required:
1- A mixture of sodium hydroxide and sodium chloride.
2- Standard solution of hydrochloric acid.
3- Standard solution of silver nitrate.
4- Standard solution of potassium thiocyanate.
5- Dilute nitric acid
C) Procedure:
1 - Transfer 10 ml. of the mixture to a conical flask and add, one drop of methyl orange then titrate against
hydrochloric acid till the end point.
2- Transfer another, 10 ml of the mixture to a conical flask and add one drop of phenolphthalein then add
dilute nitric acid till the solution becomes acidic (discharge of red color). Add 30 ml of silver nitrate
solution (excess).
3- Add 1 ml. of ferric alum indicator potassium thiocyanate solution till brownish-red color appears
4- Repeat the experiment several times and tabulate your results.
6- A blank experiment must be done, as an excess silver nitrate must be added to produce detectable
quantity of precipitate to human eye.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
14
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
E) Calculations:
15
3- Complex Formation Titration
Experimental (1)
Determination of Zinc
A) Theory:
B) Chemicals required:
1- EDTA solution 0.01 M
2- Solution of zinc sulphate.
3- Ammonia buffer solution.
4- EBT indicator
C) Procedure:
1- prepare 100 ml of a standard EDTA solution 0.01M.
2- Pipette 10 ml of Zn+2 into a 250 ml conical flask, add 2 ml of ammonia buffer solution.
3- Add few specks of EBT indicator and titrate with EDTA solution.
4- Continue the titration, dropwise, until color change from wine red to blue.
5- Repeat the experiment several times and tabulate your results.
D) Results:
Exp.
Start
End
Volume
No.
Point
point
consumed
1
2
3
Average =
E) Calculations:
1 ml of 0.01 M EDTA =
mol.wt.of ZnSO4.7H2O
= 0.002874 g
1000 × 100
Conc. of ZnSO 4 .7H 2 O =
V × f × F × 1000
= gl −1
volume of sample
16
Experimental (2)
Determination of copper
A) Theory:
Cu-ind. complex (yellow)
Cu2+ + ind.
CuY-- + ind. (purple) + 2H+
Cu-ind. complex (yellow) + H2Y-B) Chemicals required:
1- Standard EDTA solution.
2- Copper solution.
3- Dilute ammonium hydroxide solution.
4- Murexide indicator.
C) Procedure:
1- Pipette 10 ml of the sample into a 250 ml conical flask, add dilute ammonium hydroxide solution until
The precipitate first formed is just redissolved.
2- Add few specks of murexide indicator and titrate with standard EDTA solution.
3- Continue the titration, dropwise, until color change from yellow to purple.
4- Repeat the experiment several times and tabulate your results.
D) Results:
Exp.
Start
End
Volume
No.
point
point
consumed
1
2
3
Average =
17
E) Calculations:
Na2CuY + H2 SO4
Na2 H2Y.2H2O + CuSO4 .5H2O
+ 7H2O
Na2H2Y.2H2 O = CuSO4 .5H2 O
1 ml of 0.01 M EDTA =
mol.wt.of CuSO 4 .5H 2 O
= 0.00297 g
1000 × 100
Conc. of CuSO 4 .5H 2 O =
V × f × F × 1000
= gl −1
volume of sample
18
Experimental (3)
Determination of Nickel.
A) Theory:
Nickel can be determined by direct titration as in copper determination.
B) Chemicals required:
1- Standard EDTA solution.
2- Nickel solution.
3- Ammonia buffer solution.
4- Murexide indicator.
C) Procedure:
1- Pipette 10 ml of the sample into a 250 ml conical flask; add 2 ml ammonia buffer solution.
2- Add few specks of murexide indicator and titrate with standard EDTA solution.
3- Continue the titration, dropwise, until color change from yellow to violet.
4- Repeat the experiment several times and tabulate your results.
D) Results:
Exp.
Start
End
Volume
No.
point
point
Consumed
1
2
3
Average =
E) Calculations:
1 m l of 0.01 M EDTA =
Conc.of NiSO 4 =
mol.wt.of NiSO 4
1000 × 100
V × f × F × 1000
= gl −1
volume of sample
19
Experimental (4)
Determination of Aluminum.
A) Theory:
Aluminum can be determined by indirect or back titration.
B) Chemicals required:
1- Standard EDTA solution.
2- Aluminum solution.
3- Ammonia buffer solution.
4- EBT indicator.
C) Procedure:
1- Pipette 10 ml of the sample into a 250 ml conical flask; add 25 ml of 0.01 M EDTA and add 2 ml
ammonia buffer solution.
2- Boil for 10 minute, cool to room temperature, add 2 ml ammonia buffer and few specks of EBT
indicator.
3- Titrate excess EDTA with 0.01 M ZnSO4 until color change from blue to wine red.
4- Repeat the experiment several times and tabulate your results.
D) Results:
Exp.
Start
End
Volume
No.
point
point
Consumed
1
2
3
Average =
E) Calculations:
1 m l of 0.01 M EDTA =
Conc.of Al3 + =
26.98 Al 3+
= 0.0002698 Al 3+ g
1000 × 100
(25 − V) × f × F × 1000
= gl−1
volume of sample
20
Experimental (5)
Determination of Ca2+ and Mg2+ mixture.
A) Theory:
EDTA form colorless complex compound with Ca2+ and Mg2+ those of Ca2+ are more stable than those of
Mg2+. E.B.T form complexes with Ca2+ and Mg2+ but those of Mg2+ are more stable than those of Ca2+ and
have wine red color. So, when E.B.T is added to solution containing Ca2+ and Mg2+ and titrated with EDTA,
Ca2+ will be extracted first from its dye complex, then Mg2+, (color change from wine red to blue). This is
measure of total Ca2+/Mg2+. But at pH = 12 in presence of murexide, Ca2+ only form pink color complex with
murexide. So, when the solution containing both calcium and magnesium is treated with murexide and then
titrated with EDTA, calcium form colorless complex and the color change into purple (the color of the free
dye). This will be measure of calcium only.
B) Chemicals required:
1- Standard EDTA solution.
2- Calcium and magnesium mixture.
3- Ammonia buffer solution.
4- E.B.T indicator.
5- Murexide indicator.
C) Procedure:
* For total calcium and magnesium.
1- Pipette 10 ml of the sample into a 250 ml conical flask; dilute to 50 ml with distilled water and add
2 ml ammonia buffer solution.
2- Add few specks of E.B.T indicator and titrate with standard EDTA solution.
3- Continue the titration, dropwise, until color change from wine red to blue.
4- Repeat the experiment several times and tabulate your results.
* For calcium only
1- Pipette 10 ml of the sample into a 250 ml conical flask; dilute to 50 ml with distilled water.
2- Add 1 ml of 8% NaOH, while mixing, add few specks of murexide indicator and titrate with standard
EDTA solution.
3- Continue the titration, dropwise, until color change from pink to purple.
4- Repeat the experiment several times and tabulate your results.
21
D) Results:
Exp.
Start
End
Volume
No.
point
point
Consumed
1
2
3
Average =
Exp.
Start
End
Volume
No.
point
point
Consumed
1
2
3
Average =
E) Calculations:
1 m l of 0.01 M EDTA =
1 m l of 0.01 M EDTA =
mol.wt.of CaCl 2
= 0.001 g CaCl2
1000 × 100
mol.wt.of MgSO 4 .7H 2 O
= 0.002463 g MgSO4.7H2O
1000 × 100
22
4- Oxidation-Reduction Titration
Experimental (1)
Determination of oxalic acid.
A) Theory:
Oxalic acid is oxidized by potassium permanganate, in acid solution, to carbon dioxide and water.
2MnSO4 +
2KMnO4 + 3H2 O + 5H2 C2O4
K2SO4 + 10CO2
+
8H2O
The reaction is complete at a temperature of about 60- 90°C
B) Chemicals required:
1- Oxalic acid solution of unknown normality .
2- Dilute sulphuric acid.
3- Potassium permanganate 0.1N.
C) Procedure:
1- Transfer 10 ml. of oxalic acid solution to the conical flask and add an equivalent amount of dilute
sulphuric acid.
2- Warm the solution gently until the temperature of the solution reaches 60 - 80°C, then add the
permanganate solution slowly from the burette till the solution acquires a light rose color. Keep the
solution hot during the titration.
If a brown precipitate is formed during titration, this may be due to one of the following reasons:
a) The temperature of the solution may be below 60°C.
b) The addition of permanganate solution was carried out rapidly.
c) The amount of sulphuric acid is insufficient,
3- Repeat the experiment three times and take the mean value of your readings.
D) Results:
Exp.
Start
End
Volume
No.
Point
Point
consumed
1
2
3
Average =
23
E) Calculations:
24
Iodimetry and Iodometry
The oxidation reduction reactions which involve the use of iodine may be divided into two classes as
follows:
a- Iodimetry (direct method): is the use of I2 as standard oxidizing agent for direct determination of reducing
agent (e.g. S 2O32-, Sn2+, ……etc).
b- Iodometry (indirect method): is the use of I- for indirect determination of oxidizing agent (e.g. IO3-,
Cr2O72-, MnO4-, H2O2 …….etc). By addition of known excess of I- where oxidizing agent oxidizes iodide and
liberating equivalent amount of iodine which is titrated with standard solution of thiosulphate.
Experimental (2)
Standardization of approximately 0.1 N sodium thioslphate with potassium iodate.
A) Theory:
This experimental is an example of iodometry where potassium iodate reacts in acid medium
with potassium iodide with the liberation of equivalent amount of iodine which is titrated with
thiosulphate.
3K2SO4 + 3I2 + 3H2O
KIO3 + 3H2SO4 + 5KI
2Na2S2O3
Na2S4O6 + 2NaI
+ I2
B) Chemicals required:
1- Potassium iodate solution 0.1 N.
2- Potassium iodide
3- Dilute sulphuric acid.
4- Thiosulphate solution of unknown normality.
C) Procedure:
1- Transfer 10 ml of potassium iodate solution, to a conical flask.
2- Add 1 gm. of potassium iodide then add 2 ml of dilute sulphuric acid.
3- Dilute the solution with disti1led water then titrate the liberated iodine against sodium thiosulphate
solution till the color becomes pale yellow.
4- Add 1 ml of starch solution and continue titration with thiosti1phate solution till the color changes from
blue to colorless.
25
D) Results:
Exp.
Start
End
Volume
No.
Point
Point
consumed
1
2
3
Average =
E) Calculations:
26
Experimental (3)
Standardization of approximately 0.1 N iodine with sodium thiosulphate.
A) Theory:
This experimental is an example of iodimetry where iodine oxidizes thioslphate and reduced to
iodide.
2Na2S2O3
Na2S4O6 + 2NaI
+ I2
B) Chemicals required:
1- Sodium thiosulphate solution 0.1 N.
2- Iodine solution of unknown normality.
C) Procedure:
1- Fill the burette with sodium thiosulphate solution..
2- Transfer, with the aid of a pipette, 10 ml. of the iodine solution to a conical flask then add about 10 ml.
of distilled water.
3- Add the thiosulphate solution till the color of the solution is pale yellow, and then add 1 ml of starch
solution.
4- Complete titration with thiosulphate till the blue color is discharged.
5- Repeat the experiment three times.
D) Results:
Exp.
Start
End
Volume
No.
Point
Point
consumed
1
2
3
Average =
E) Calculations:
27