QUANTITATIVE ANALYSIS
The term quantitative analysis refers to the analysis carried out by reacting a
standard solution (a solution of accurately known concentration) with a solution
whose concentration is to be determined. It’s usually called volumetric analysis.
Standard solution.
This is a solution whose concentration is accurately known. The concentration can
be in moles per litre (Molarity) or grams in a given volume like 1 litre, 500cm3 etc
Primary standard.
This is a substance which is analytically pure and whose composition is accurately
known and is used to prepare a standard solution. E.g Anhydrous sodium carbonate,
borax, potassium iodate, oxalic acid e.t.c
Properties of a good primary standard.
 It must be readily available in a state of high purity.
 It should be readily soluble in water.
 It should have a high molecular mass to reduce errors during weighing.
 It must be stable in air (have a constant composition) i.e should not be
hygroscopic, deliquescent or efflorescent.
The following substances cannot be used as primary standards;
(a) Sodium hydroxide.
This is because;
 It is deliquescent (absorbs water from the atmosphere and dissolves in it)
 It absorbs carbon dioxide from air forming a solution of sodium carbonate
solution.
2NaOH(aq) + CO2(g)
Na2CO3(aq) + H2O(l)
This reduces the concentration of sodium hydroxide solution.
When excess carbon dioxide comes into contact with sodium hydroxide
solution, a white precipitate is formed.
Na2CO3(aq) + H2O(l) + CO2(g)
2NaHCO3(s)
(b) Concentrated Sulphuric acid.
It is hygroscopic hence does not have a constant composition.
(c) Ammonia solution.
It is highly volatile.
(d) Concentrated nitric acid.
It is a strong oxidizing agent.
(e) Sodium thiosulphate
It is efflorescent (loses its water of crystallization to the atmosphere)
When left exposed in air, it absorbs carbon dioxide from air forming carbonic
acid. The carbonic acid decomposes the thiosulphate ions to form sulphur
dioxide, sulphur and water. This reduces the concentration of sodium
thiosulphate.
H2O(l) + CO2(g)
H2CO3(s)
2+
S2O3 (aq) + 2H (aq)
S(s) + SO2(g) + H2O(l)
 Distinguish between S2O32- and SO32Reagent: Dilute Sulphuric acid or dilute hydrochloric acid.
S2O32- - A yellow precipitate and bubbles of a colourless gas.
SO32- - Bubbles of a colourless gas.
1
(f) Potassium mangante(VII)
It is impure, contaminated with manganese(IV) oxide, thus its solution
should be filtered first before use.
In the neutral or alkaline medium, it is reduced to a dark brown solid of
manganese(IV) oxide.
MnO-4(aq) + 2H2O(l) + 3e
MnO2(s) + 4OH-(aq)
ACID BASE TITRATIONS.
EXPERIMENT 1.
You are provided with the following;
FA1, which is a solution of hydrochloric acid
Solid A, which is anhydrous sodium carbonate.
You are required to standardize FA1 using sodium carbonate.
Procedure.
Weigh accurately about 2.7g of A and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA2.
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of methyl
orange indicator and titrate with FA1 from the burette until the endpoint. Repeat
the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + A=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of A alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
Calculate the;
(i) Number of moles of sodium carbonate in FA2 that reacted with hydrochloric acid
in FA1.
2
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(ii) Molarity of hydrochloric acid in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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Experiment 2
You are provided with the following;
Solid B, which is borax.
FA1, which is a solution of hydrochloric acid
You are required to standardize of FA1 using borax.
Theory
Borax(Na2B4O7.10H2O) is a salt of a strong base and a weak acid.
Borax undergoes partial hydrolysis to form boric acid and sodium hydroxide.
Na2B4O7(aq) + 7H2O(l)
2NaOH(aq) + 4H3BO3(aq)
The sodium hydroxide imparts on the alkaline reaction to the solution.
Hence borax reacts with hydrochloric acid according to the equation below
Na2B4O7.10 H2O(aq) + 2HCl(aq)
2NaCl(aq) + 4H3BO3(aq) + 5H2O(l)
Procedure.
Weigh accurately about 4.8g of B and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA2.
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of methyl
orange indicator and titrate with FA1 from the burette until the endpoint. Repeat
the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + B=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of B alone=……………………………………………………………g
3
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………………….. cm3
Questions;
Calculate the concentration of;
(i) Borax in FA2 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
(ii) Hydrochloric acid in FA1 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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Experiment 3
You are provided with the following;
FA1, which is approximately 0.1M solution of sodium hydroxide.
FA2, which is 0.1M hydrochloric acid
Solid C, which is an organic acid of formula, (COOH)2. xH2O
4
You are required to standardize FA1 and use it to determine the number of moles of
water of crystallization, x of an organic acid, (COOH)2. xH2O
Procedure I
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA2 from the burette until you
reach the end point. Repeat the titration 2-3 times until you obtain consistent
results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA2 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3
Average volume of FA2 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of FA1.
……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 1.6g of C and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask. Add 2-3 drops of
5
phenolphthalein indicator and titrate with FA3 from the burette until the endpoint.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + C=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of C alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA3 used (cm3)
Titre values used to calculate average volume of FA3.
………………………………………………………………………………………………………cm3
Average volume of FA3 used …………………………………………………………..cm3
Questions;
(b) Calculate the;
(i) Number of moles of sodium hydroxide in FA1 that reacted with FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Molecular mass of the organic acid (COOH)2.xH2O.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(c) Determine the value of x in the organic acid, (COOH)2.xH2O
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
6
Experiment 4
You are provided with the following;
FA1, which is a solution of sodium hydroxide.
FA2, which is 0.1M hydrochloric acid
Solid D, which is an organic acid of formula, (CH2)n(COOH)2.
You are required to standardize FA1 and use it to determine the value, n in the
organic acid (CH2)n(COOH)2.
Procedure I
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA2 from the burette until you
reach the end point. Repeat the titration 2-3 times until you obtain consistent
results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA2 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3.
Average volume of FA2 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of sodium hydroxide solution in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………
7
Procedure II
Weigh accurately about 1.6g of D and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and titrate with FA3 from the burette until the endpoint.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + D=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of D alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA3 used (cm3)
Titre values used to calculate average volume of FA3.
………………………………………………………………………………………………………cm3
Average volume of FA3 used ………………………………………………… cm3
Questions;
(b) Calculate the;
(i) Number of moles of sodium hydroxide in FA1 that reacted with FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of FA3 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(iii) Determine the value of n in the organic acid, (CH2)n(COOH)2
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
8
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TRIAL QUESTIONS
1. 20cm3 of 0.15M sodium hydroxide solution required 20cm3 of a solution of an
acid H2X for complete reaction which was made by dissolving 2.45g of the acid
in 250cm3. Calculate the relative molecular mass of the acid.
2. 25cm3 of 0.04M sodium carbonate solution was neutralized by 6.67cm3 of
0.15M solution of an acid HX. Determine the mole ratio for the reaction that
took place.
3. 20cm3 of 0.2M potassium hydroxide solution required 25cm3 of hydrochloric
acid for complete reaction. Calculate the concentration of hydrochloric acid in
grams per litre.
4. 2.0g of an impure acid H2X was dissolved in water and the solution made to
250cm3 with distilled water. 25cm3 of this solution required 25cm3 of 0.1M
sodium hydroxide solution for complete reaction. Calculate the percentage
purity of the acid. [X=124]
5. 4.0g of hydrated sodium carbonate, Na2CO3.XH2O was dissolved in water and
the solution made to 250cm3 in a volumetric flask. 25cm3 of this solution
required 21cm3 of 0.2M hydrochloric acid for complete reaction. Determine the
value of X.
6. 5.0g of an acid YCO2H was dissolved in water and the solution made to 250cm3
with distilled water in a volumetric flask. 25cm3 of this solution required
20.5cm3 of 0.4M sodium hydroxide solution for complete reaction. Determine
the value of Y in the acid.
7. 20cm3 of an acid HOOC(CH2)nCOOH made by dissolving 10g in a litre required
33.9cm3 of 0.1M sodium hydroxide solution for complete neutralization.
Determine the value of n in the acid.
9
BACK TITRATION.
Back titration is a technique of analyzing the composition of a substance by reacting
it with an excess reagent of known concentration and then titrating the unreacted
reagent with another standard solution.
Theory
An excess standard solution X is reacted with a substance Y whose concentration is
to be determined.
Y+X
Product + Unreacted X
The unreacted X is diluted to a given volume i.e 250cm3. A given volume of the
diluted solution is then titrated against another standard solution Z.
Unreacted X + Z
Product
By back calculation, the amount of Y that reacted with X can be determined.
Procedure
 Calculate the number of moles of Z that reacted with unreacted X.
 Calculate the number of unreacted X using the mole ratios.
 Calculate the number of moles of unreacted X in the diluted volume i.e
250cm3
 Calculate the original moles of X that was added.
 By subtraction, calculate the number of moles X that reacted with Y.
 Calculate the number of moles of Y that reacted with X using moles ratios.
Worked examples
4.0g of impure calcium carbonate was added to 150cm3 of 0.48M hydrochloric acid
and the mixture shaken until effervescence stopped. The resultant solution was
made to a 250cm3 mark in a volumetric flask using distilled water. 25cm3 of this
solution required 22.5cm3 of 0.2M sodium hydroxide solution for complete
neutralization.
Calculate the;
(a) Number of moles hydrochloric acid that reacted with sodium hydroxide.
(b) Number of moles of hydrochloric acid that reacted with calcium carbonate.
(c) Number of moles of calcium carbonate that reacted with hydrochloric acid.
(d) Percentage purity of calcium carbonate in the sample.
Solution
3
(a)
1000cm of solution contains 0.2 moles of NaOH
22.5cm3 of solution contains
0.2×22.5
1000
moles of NaOH
=0.0045 moles
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
Mole ratio NaOH: HCl = 1:1
Moles of HCl that reacted = 0.0045 moles
(b)
25cm3 of solution contains 0.0045 moles of HCl
250cm3 of solution contains
0.0045×250
25
=0.045 moles
Original moles of HCl
1000cm3 of solution contains 0.48 moles of HCl
10
150cm3 of solution contains
0.48×150
1000
moles of HCl
=0.072 moles
Moles of HCl that reacted with CaCO3 = 0.072- 0.045 = 0.027 moles
(c)
CaCO3(s) + 2HCl(aq)
CaCl2(aq) + CO2(g) + H2O(l)
Mole ratio CaCO3: HCl = 1:2
1
Moles of CaCO3 that reacted = 2 × 0.027 = 0.0135 moles
(d)
Rmm of CaCO3 = 40 + 12 + (3×16) = 100
1 mole of CaCO3 weighs 100g
0.0135 moles of CaCO3 weighs
Percentage purity of CaCO3 =
0.0135×100
1
1.35×100
4.0
= 1.35g
= 33.75%
Experiment 5
You are provided with the following;
FA1, which is 1M sodium hydroxide solution
FA2, which is 1M hydrochloric acid
Solid E, which is a metal carbonate, XCO3.
You are required to determine the relative atomic mass of X in XCO3.
Procedure
Weigh accurately about 1.0g of E and place it in a beaker. Add 100cm3 of FA2 and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and titrate with FA1 from the burette until the endpoint.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + E=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of E alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………cm3
11
Questions;
(a) Calculate the number of moles of;
(i) Excess hydrochloric acid that reacted with sodium hydroxide.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Hydrochloric acid that reacted with E.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(iii) E that reacted with hydrochloric acid.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(b) Determine the relative atomic mass of X in XCO3.
……………………………………………………………………………………………………………………………
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12
Experiment 6
You are provided with the following;
FA1, which is 0.1M sodium hydroxide solution
FA2, which is 2M hydrochloric acid
Solid F, which is impure calcium carbonate.
You are required to determine the percentage impurity of solid F.
Procedure
Weigh accurately about 1.5g of F and place it in a beaker. Add 20cm3 of FA2 and
carefully stir to dissolve. Add 80cm3 of distilled water and label the solution FA3.
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and titrate with FA3 from the burette until the endpoint.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + F=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of F alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA3 used (cm3)
Titre values used to calculate average volume of FA3.
………………………………………………………………………………………………………cm3
Average volume of FA3 used …………………………………………………………..cm3
Questions;
(a) Calculate the number of moles of;
(i) Hydrochloric acid in 100cm3 of FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………...
(ii) Hydrochloric acid that reacted with F.
13
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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(b) Determine the percentage impurity of solid F. (Ca=40, C=12, 0=16)
……………………………………………………………………………………………………………………………
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Experiment 7
You are provided with the following;
FA1, which is 0.1M sodium hydroxide solution
FA2, which is 0.05M sulphuric acid
Solid G, which is an organic acid, (COOH)2.nH2O .
You are required to determine the number of moles of water of crystallization, n in
the organic acid.
Procedure
Weigh accurately about 1.0g of G and place it in a beaker. Add 200cm3 of FA1 and
carefully and stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and titrate with FA2 from the burette until the endpoint.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + G=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of G alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA2 used (cm3)
14
Titre values used to calculate average volume of FA2.
………………………………………………………………………………………………………cm3
Average volume of FA2 used …………………………………………………………..cm3
Questions;
(a) Calculate the number of moles of;
(i) Excess sodium hydroxide that reacted with sulphuric acid.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium hydroxide that reacted with G.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………...
(iii) G that reacted with sodium hydroxide.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Determine the value of n in (COOH)2. nH2O.
……………………………………………………………………………………………………………………………
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……………………………………………………………………………………………………………………………
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15
Experiment 8
You are provided with the following;
FA1, which is 1M sodium hydroxide solution
FA2, which is 0.2M hydrochloric acid
Solid H, which is impure ammonium chloride.
You are required to determine the percentage purity of solid H.
Procedure
Weigh accurately about 4.0g of H and transfer it into a conical flask and add 100cm3
of FA1 and carefully boil the mixture for about 10 minutes. Cool and transfer the
contents into a 250cm3 volumetric flask and add water to make up to the mark.
Label the solution FA3. Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask.
Add 2-3 drops of phenolphthalein indicator and titrate with FA2 from the burette
until the endpoint. Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + H=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of H alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA2 used (cm3)
Titre values used to calculate average volume of FA2.
………………………………………………………………………………………………………cm3
Average volume of FA2 used …………………………………………………………..cm3
Questions;
(a) Calculate the number of moles of;
(i) Excess sodium hydroxide that reacted with FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium hydroxide that reacted with H.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
16
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iii) H that reacted with FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Determine the percentage purity in H.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 9
You are provided with the following;
FA1, which is a 1M hydrochloric acid
FA2, which is sodium hydroxide solution
FA3, which is 0.1M Sulphuric acid
Solid I, which is an acid H3X
You are required to standardize sodium hydroxide solution in FA2 and use to
determine the value of X in H3X.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA1 from the burette until you
reach the end point. Repeat the titration 2-3 times until you obtain consistent
results.
Enter your results in the table below.
Results
17
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of sodium hydroxide solution in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 3.4g of I and place it in a beaker. Add 100cm3 of FA2 and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA4.
Pipette 25cm3 of FA4 into a clean conical flask. Add 2-3 drops of phenolphthalein
indicator and titrate with FA3 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + I=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of I alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
18
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA3 used (cm3)
Titre values used to calculate average volume of FA3.
………………………………………………………………………………………………………cm3
Average volume of FA3 used …………………………………………………………..cm3
Questions;
(b) Calculate the number of moles of;
(i) Sulphuric acid in FA3 that reacted with sodium hydroxide in FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Excess sodium hydroxide in 250cm3 of FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
(iii) Sodium hydroxide that reacted with I.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iv) I that reacted with sodium hydroxide.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
19
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(c) Determine the value of X in H3X.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
TRIAL QUESTIONS
1. To 2.0g of impure calcium carbonate in a beaker was added 50cm3 of 1M
hydrochloric acid. The resultant solution was made to 250cm3 with distilled
water. 25cm3 of this solution required 20cm3 of 0.1M hydrochloric acid for
complete reaction. Determine the percentage of pure calcium carbonate in the
sample.
2. To 1.6g of ammonium sulphate was added 50cm3 of 1M sodium hydroxide
solution and the mixture boiled until all the ammonia was expelled. The
resultant solution was made to 250cm3 with distilled water. 25cm3 of this
solution required 20cm3 of 0.1M hydrochloric acid for complete reaction.
Determine the percentage of ammonia expelled.
3. 4.0g of ammonium chloride was added to 200cm3 of 0.17M sodium hydroxide
solution. The mixture was boiled until all the ammonia was expelled and the
resultant solution made to 250cm3 with distilled water. 25cm3 of this solution
required 20cm3 0f 0.05M Sulphuric acid for complete neutralization using
phenolphthalein indicator. Calculate the percentage of ammonia expelled.
4. 10g of ammonium chloride was added to 200cm3 of 1M sodium hydroxide
solution. The mixture was boiled until ammonia ceased to be evolved. The
resultant solution was titrated with 1M hydrochloric acid and 11.8cm3 of
hydrochloric acid was required for complete reaction. Calculate the percentage
of ammonia in the salt.
20
DOUBLE INDICATOR TITRATIONS.
INTRODUCTION
A double indicator titration is a technique used for analyzing the composition of a
mixture of two different bases that react with an acid in the same reaction using two
different indicators.
In one type of analysis, the indicators are added separately i.e each indicator is used
in separate titrations. [Separate titrations]
In another type of titration, the indicators are added simultaneously i.e one
indicator (usually phenolphthalein) is added to reach the first end point and the
second indicator (usually methyl orange) is added to complete the reaction and
reach the second end point. (Continuous titrations)
The mixtures considered include;
(a) Sodium carbonate and sodium hydrogencarbonate.
(b) Sodium carbonate and sodium hydroxide
(c) Sodium hydrogencarbonate and sodium hydroxide
Sodium carbonate Vs phenolphthalein indicator
Phenolphthalein indicator changes colour to pink under alkaline conditions i.e pH
range of 8.2-10.0
Sodium carbonate solution is alkaline (pH=9.0) is able to turn the indicator pink.
During the titration with an acid, the carbonate is converted to a hydrogencarbonate
which is an acid salt (pH =3.7), the indicator turns colourless to mark the end point
for half neutralization.
If the volume of the acid needed is V1, therefore for complete reaction of the
carbonate, the volume of the acid needed is twice V1 i.e 2V1
Sodium hydrogencarbonate Vs phenolphthalein indicator
A solution of sodium hydrogencarbonate is acidic (pH =3.7). A drop of
phenolphthalein indicator to this solution remains colourless and therefore cannot
be used to mark the end point for the titration with an acid.
Sodium carbonate Vs methyl orange indicator
Methyl orange indicator turns yellow in a solution of sodium carbonate. During the
titration with an acid, the carbonate reacts completely to form a salt, water and
carbon dioxide. Sodium hydrogencarbonate also reacts completely.
21
ANALYSIS OF A MIXTURE OF SODIUM CARBONATE AND SODIUM
HYDROGENCARBONATE.
Method I (continuous titration)
(a) The mixture is pipetted, phenolphthalein indicator added and titrated against a
standard solution of hydrochloric acid until the end point. The volume of the acid, V1
used is recorded.
Sodium carbonate undergoes half neutralization to form sodium hydrogencarbonate
while the original sodium hydrogencarbonate does not react.
Na2CO3(aq) + HCl(aq)
V1
NaHCO3(aq) + NaCl(aq)
Therefore for the complete neutralization of sodium carbonate alone, the volume of
the acid would be 2V1.
(b) To the resultant solution in (a) is added a few drops of methyl orange indicator
and the titration continued until the end point. The volume of the acid, V2, used is
recorded.
Both the formed sodium hydrogencarbonatein (a) and the original sodium
hydrogencarbonate react with the acid.
NaHCO3(aq) + HCl(aq) V1
NaCl(aq) + CO2(g) + H2O(l)
(Formed)
V2
NaHCO3(aq) + HCl(aq)
NaCl(aq) + CO2(g) + H2O(l)
(Original)
The volume of the acid that reacted with the NaHCO3 formed=V1
The volume of the acid that reacted with the original NaHCO3 =V2-V1
Method II (separate titration)
(a) The mixture is pipetted, phenolphthalein indicator added and titrated against a
standard solution of hydrochloric acid until the end point. The volume of the acid, V1
used is recorded.
Sodium carbonate undergoes half neutralization to form sodium hydrogencarbonate
while the original sodium hydrogencarbonate does not react.
Na2CO3(aq) + HCl(aq) V1
NaHCO3(aq) + NaCl(aq)
Therefore for the complete neutralization of Na2CO3 alone, the volume of the acid
would be 2V1.
(b) A fresh sample of the mixture is pipetted, a few drops of methyl orange indicator
is added and then titrated with the acid until the end point. The volume of the acid,
V2, used is recorded.
Both the sodium carbonate and the sodium hydrogencarbonate react with the acid
completely.
Na2CO3(aq) + 2HCl(aq) 2V1
2NaCl(aq) + CO2(g) + H2O(l)
V2
NaHCO3(aq) + HCl(aq)
NaCl(aq) + CO2(g) + H2O(l)
(original)
22
The volume of the acid that reacted completely with Na2CO3 alone = 2V1
The volume of the acid that reacted with the original NaHCO3 =V2-2V1
Worked example.
25cm3 of a mixture containing sodium carbonate and sodium hydrogencarbonate
required 12cm3 of 0.1M hydrochloric acid using phenolphthalein indicator. When
three drops of methyl orange indicator are added to the resultant mixture and the
titration continued, 21.3cm3 of the acid was used. Determine the percentage
composition of the mixture.
Solution
Sodium carbonate
The volume of the acid that reacted with Na2CO3 alone = 2V1=2 × 12 = 24 cm3
1000cm3 of solution contains 0.1 moles of HCl
24cm3 of solution contains
0.1×24
1000
moles of HCl
=0.0024 moles
Na2CO3(aq) + 2HCl(aq)
2NaCl(aq) + CO2(g) + H2O(l)
Mole ratio Na2CO3: HCl = 1:2
1
Moles of Na2CO3 that reacted = 2 × 0.0024 = 0.0012 moles
25cm3 of solution contains 0.0012 moles of Na2CO3
1000cm3 of solution contains
0.0012×1000
25
moles of Na2CO3
=0.048 moles
Molarity of Na2CO3 is 0.048M
RFM of Na2CO3= (2 × 23) + 12 + (3 × 16) = 106
1 mole of Na2CO3 weighs 106g
0.048 moles of Na2CO3 weighs (106 × 0.048) = 5.088𝑔
Concentration of Na2CO3 is 5.088gl-1
Sodium hydrogencarbonate
The volume of the acid that reacted with NaHCO3 = V2-V1 =21.3 − 12 = 9.3 cm3
1000cm3 of solution contains 0.1 moles of HCl
9.3cm3 of solution contains
0.1×9.3
1000
moles of HCl
=0.00093 moles
NaHCO3(aq) + HCl(aq)
NaCl(aq) + CO2(g) + H2O(l)
Mole ratio NaHCO3: HCl = 1:1
Moles of NaHCO3 that reacted = = 0.00093 moles
25cm3 of solution contains 0.00093 moles of NaHCO3
1000cm3 of solution contains
0.00093×1000
25
moles of NaHCO3
=0.0372 moles
Molarity of NaHCO3 is 0.0372M
RFM of NaHCO3= 23 + 1 + 12 + (3 × 16) = 84
1 mole of NaHCO3 weighs 84g
0.0372 moles of NaHCO3 weighs (84 × 0.0372) = 3.1248𝑔
23
Concentration of NaHCO3 is 3.1248gl-1
Total concentration = 5.088+3.1248= 8.2128gl-1
Hence, percentage of Na2CO3 =
Percentage of NaHCO3 =
5.088×100
8.2128
3.1248×100
8.2128
= 61.95%
= 38.05%
ANALYSIS OF A MIXTURE OF SODIUM CARBONATE AND SODIUM HYDROXIDE.
Method I (continuous titration)
(a) With phenolphthalein indicator, Na2CO3 undergoes half neutralization to form
NaHCO3. The sodium hydroxide also reacts. The volume of the acid, V1 used is
recorded.
Na2CO3(aq) + HCl(aq)
NaHCO3(aq) + NaCl(aq)
V1
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
(b) To the resultant solution in (a) is added a few drops of methyl orange indicator
and the titration continued until the end point. The volume of the acid, V2, used is
recorded.
The formed NaHCO3 in (a) reacts with the acid.
NaHCO3(aq) + HCl(aq) V2
NaCl(aq) + CO2(g) + H2O(l)
(formed)
The volume of the acid needed for the complete neutralization of Na2CO3 =2V2
The volume of the acid needed to react with NaOH alone =V1-V2
Method II (separate titration)
(a) With phenolphthalein indicator, Na2CO3 undergoes half neutralization to form
NaHCO3. The sodium hydroxide also reacts. The volume of the acid, V1 used is
recorded.
Na2CO3(aq) + HCl(aq)
NaHCO3(aq) + NaCl(aq)
V1
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
(b) A fresh sample of the mixture is pipetted, a few drops of methyl orange indicator
is added and then titrated with the acid until the end point. The volume of the acid,
V2, used is recorded.
Both the sodium carbonate and the sodium hydroxide react with the acid
completely.
Na2CO3(aq) + 2HCl(aq)
2NaCl(aq) + CO2(g) + H2O(l)
V2
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
The volume of the acid for half neutralization of Na2CO3 = V2 - V1
24
The volume of the acid for complete neutralization of Na2CO3 =2(V2 - V1)
The volume of the acid that reacted with NaOH alone = V2 - 2(V2 - V1)
= 2V1-V2
ANALYSIS OF A MIXTURE OF SODIUM HYDROGENCARBONATE AND SODIUM
HYDROXIDE.
Method I (continuous titration)
(a) With phenolphthalein indicator, only NaOH reacts with the acid to reach the first
end point. The volume of the acid, V1 used is recorded.
NaOH(aq) + HCl(aq) V1
NaCl(aq) + H2O(l)
(b) To the resultant solution in (a) is added a few drops of methyl orange indicator
and the titration continued until the end point. The volume of the acid, V2, used is
recorded.
The NaHCO3 reacts with the acid.
NaHCO3(aq) + HCl(aq) V2
NaCl(aq) + CO2(g) + H2O(l)
(original)
The volume of the acid needed for the complete neutralization of NaHCO3 =V2
The volume of the acid needed to react with NaOH alone =V1
Method II (separate titration)
(a) With phenolphthalein indicator, only NaOH reacts with the acid to reach the first
end point. The volume of the acid, V1 used is recorded.
NaOH(aq) + HCl(aq) V1
NaCl(aq) + H2O(l)
(b) A fresh sample of the mixture is pipetted, a few drops of methyl orange indicator
is added and then titrated with the acid until the end point. The volume of the acid,
V2, used is recorded.
Both the sodium hydrogencarbonate and the sodium hydroxide react with the acid
completely.
NaHCO3(aq) + HCl(aq)
NaCl(aq) + CO2(g) + H2O(l)
V2
NaOH(aq) + HCl(aq) V1
NaCl(aq) + H2O(l)
The volume of the acid needed to react with NaOH alone =V1
The volume of the acid for complete neutralization of NaHCO3 =V2 - V1
25
Experiment 10
You are provided with the following;
FA1, which is 0.1M hydrochloric acid
FA2, which is a solution containing sodium carbonate and sodium hydroxide.
You are required to determine the molar concentration of sodium carbonate and
sodium hydroxide.
Procedure
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA1 from the burette until you
reach the end point. Record your results in table I.
Without pouring away the mixture, add 2-3 drops of methyl orange indicator and
continue to titrate with FA1 from the burette until you reach the second end point.
Record your results in table II.
Repeat the titration 2-3 times until you obtain consistent results
Note:
The final reading with phenolphthalein indicator is the initial reading with methyl
orange indicator.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used……………………………………………………………cm3
Table II
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
26
Average volume of FA1 used……………………………………………………………cm3
Questions;
Calculate the molar concentration of;
(a) Sodium carbonate solution in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Sodium hydroxide solution in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Experiment 11
You are provided with the following;
FA1, which is a solution of hydrochloric acid.
FA2, which is a solution containing sodium carbonate and sodium
hydrogencarbonate.
FA3, which is a solution containing 19.10g of borax, Na2B4O7.10H2O in 1dm3 of
solution.
You are required to standardize FA1 and use it to determine the composition of the
mixture in FA2
Borax reacts with hydrochloric acid according to the equation:
Na2B4O7.10 H2O(aq) + 2HCl(aq)
2NaCl(aq) + 4H3BO3(aq) + 5H2O(l)
27
Procedure I
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 2-3 drops of methyl
orange indicator and then titrate with FA1 from the burette until you reach the end
point. Repeat the titration 2-3 times until you obtain consistent results.
Record your results in the table I.
Results
Volume of pipette used=…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3.
Average volume of FA2 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of hydrochloric acid in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA1 from the burette until you
reach the end point. Record your results in the appropriate section of table II (using
phenolphthalein indicator).
Without pouring away the mixture, add 2-3 drops of methyl orange indicator and
titrate with FA1 from the burette until you reach the second end point. Record your
results in the appropriate section table II (using methyl orange indicator).
28
Repeat the titration 2-3 times until you obtain consistent results
Note:
The final reading with phenolphthalein indicator is the initial reading with methyl
orange indicator.
Phenolphthalein indicator Methyl orange indicator
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1 using phenolphthalein
indicator.
…………………………………………………………………………………………………………………….. cm3
Average volume of FA1 used with phenolphthalein indicator ………………………… cm3
Titre values used to calculate average volume of FA1 using methyl orange indicator
……………………………………………………………………………………………………………………… cm3
Average volume of FA1 used with methyl orange indicator……………………………… cm3
Questions;
(b) Calculate the concentration of;
(i) Sodium carbonate solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium hydrogencarbonate solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
29
Experiment 12
You are provided with the following;
FA1, which is 0.1M hydrochloric acid
FA2, which is a solution containing sodium hydroxide and sodium
hydrogencarbonate.
You are required to determine the percentage composition of sodium hydroxide and
sodium hydrogencarbonate.
Procedure
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA1 from the burette until you
reach the end point. Record your results in table I.
Pipette another 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
methyl orange indicator and then titrate with FA1 from the burette until you reach
the end point. Record your results in table II.
Repeat the titration 2-3 times until you obtain consistent results
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Table II
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3
Average volume of FA1 used……………………………………………………………cm3
Questions;
(a)Calculate the concentration of;
30
(i) Sodium hydroxide solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium hydrogencarbonate solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Calculate the percentage composition of sodium hydroxide and sodium
hydrogencarbonate in the mixture.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
31
Experiment 13
You are provided with the following;
FA1, which is 0.1M sulphuric acid
FA2, which is a solution containing sodium carbonate and sodium
hydrogencarbonate.
You are required to determine the composition of sodium carbonate and sodium
hydrogencarbonate.
Procedure
Using a measuring cylinder, measure 100cm3 of FA2 into a 250cm3 volumetric flask
and make it to the mark with more distilled water. Label the solution FA3
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and then titrate with FA1 from the burette until you
reach the end point. Record your results in the appropriate section of table I (using
phenolphthalein indicator).
Pipette another 20cm3 or 25cm3 of FA3 into a clean beaker. Add 2-3 drops of methyl
orange indicator and titrate with FA1 from the burette until you reach the second
end point. Record your results in the appropriate section table I (using methyl
orange indicator).
Repeat the titration 2-3 times until you obtain consistent results
Results
Volume of pipette used…………………………………..cm3
Phenolphthalein indicator Methyl orange indicator
(cm3)
Final burette reading
Initial burette reading (cm3)
Volume of FA1 used (cm3)
(i) Titre values used to calculate average volume of FA1 using phenolphthalein
indicator.
………………………………………………………………………………………………………cm3
Average volume of FA1 used with phenolphthalein indicator ………………………… cm3
(ii) Titre values used to calculate average volume of FA1 using methyl orange
indicator
………………………………………………………………………………………………………cm3
Average volume of FA1 used with methyl orange indicator……………………………… cm3
Questions;
Calculate the concentration of;
(i) Sodium carbonate solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
32
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Sodium hydrogencarbonate solution in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
TRIAL QUESTIONS
1. 25cm3 of a solution containing a mixture of sodium carbonate and sodium
hydrgencarbonate required 7.1cm3 of 0.1M hydrochloric acid using
phenolphthalein indicator. Another 25cm3 of the solution needed 21.6cm3 of
0.1M hydrochloric acid using methyl orange indicator. Determine the percentage
of each component in the mixture.
2. 25cm3 of a mixture containing sodium carbonate and sodium hydroxide
required 27.1cm3 of 0.1M hydrochloric acid using phenolphthalein indicator.
Without pouring away the solution, a further 8.3cm3 of 0.1M hydrochloric acid
33
was used using methyl orange indicator. Determine the concentration of each
component in the mixture.
3. 25cm3 of a solution containing a mixture of sodium hydroxide and sodium
hydrgencarbonate was diluted to 250cm3. 25cm3 of the diluted solution required
12.5cm3 of 0.1M hydrochloric acid using phenolphthalein indicator. Another
25cm3 of the diluted solution needed 32.4cm3 of 0.1M hydrochloric acid using
methyl orange indicator. Determine the concentration of each component in the
original undiluted solution.
REDOX TITRATIONS.
PERMANGANATE TITRATIONS.
Experiment 15
You are provided with the following;
FA1, which is a solution of potassium manganate(VII).
1M sulphuric acid
Solid J, which is ammonium iron(II) sulphate-6-water, (NH4)2SO4FeSO4.6H2O
You are required to determine the molar concentration of potassium
managanate(VII) in FA1.
Procedure.
Weigh accurately about 4.9g of J and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add an equal volume of 1M
sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + J=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of J alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
34
Questions;
Calculate the molar concentration of FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Experiment 16
You are provided with the following;
FA1, which is a solution of potassium manganate(VII).
FA2, which is a solution containing 5.6g of sodium ethanedioate, Na2C2O4 per litre.
2M sulphuric acid
Solid K, which is ammonium iron(II) sulphate-6-water, (NH4)2SO4FeSO4.nH2O
You are required to standardize FA1 and use it to determine the value, n in
(NH4)2SO4FeSO4.nH2O
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and heat the mixture to about 700C. Titrate the hot mixture with FA1
from the burette until you reach the end point. Repeat the titration 2-3 times until
you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
35
Titre values used for calculating average volume of FA1 used
……………………………………………………………………………………………………………...cm3
Average volume of FA1 used…………………………………………………………………...cm3
Questions;
(a) Calculate the molar concentration of FA1. [Na=23, C=12, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure I
Weigh accurately about 4.4g of K and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + K=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of K alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
36
Questions;
(b) Calculate the;
(i) Molar concentration of iron(II) ions in FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Molar mass of K and hence the value of n. [N=14, H=1, S=32, O=16, Fe=56]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 17
You are provided with the following;
FA1, which is a solution of potassium manganate(VII).
FA2, which is a 0.05M solution ammonium iron(II) sulphate-6water,(NH4)2SO4FeSO4.nH2O
2M sulphuric acid
Solid L, which is an ethanedioate (oxalate) with the formula, CO2Y
CO2Y
You are required to standardize FA1 and use it to determine the relative atomic
mass of Y.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid. Titrate the hot mixture with FA1 from the burette until you reach the
end point. Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
37
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3
Average volume of FA1 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 1.0g of L and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of
FA3 and heat the mixture to about 600C. Titrate the hot mixture with FA1 from the
burette until the endpoint. Repeat the titration 2-3 times until you obtain consistent
results.
Enter your results in the table below.
Results;
Mass of empty bottle + L=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of L alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
38
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the;
(i) Number of moles oxalate ions in 250cm3 of FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Molar mass of L.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Relative atomic mass of Y. [O=16, C=12]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
39
Experiment 18
You are provided with the following;
FA1, which is a 0.02M solution of potassium manganate(VII).
FA2, which is a solution containing a mixture of sodium ethanedioate and
ethanedioic acid.
FA3, which is a 0.1M solution of sodium hydroxide.
2M sulphuric acid
You are required to determine the percentage composition of sodium ethanedioate
in FA2.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add 2-3 drops of
phenolphthalein indicator and titrate with FA3 from the burette until you reach the
end point. Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA3 used(cm3)
Titre values used for calculating average volume of FA3 used
………………………………………………………………………………………………………cm3.
Average volume of FA3 used …………………………………………………………………...cm3
Procedure II
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and heat the mixture to about 600C. Titrate the hot mixture with FA1
from the burette until the endpoint. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
40
Questions;
(a) Calculate the concentration of;
(i) Ethanedioic acid in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium ethanedioate in FA2 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Calculate the percentage composition of sodium ethanedioate in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
41
Experiment 19
You are provided with the following;
FA1, which is a solution of potassium manganate(VII).
FA2, which is a solution containing 2.8g of sodium oxalate, Na2C2O4 in 500cm3.
FA3, which is a solution of ammonium iron(III) sulphate-24-water,
(NH4)2SO4Fe2(SO4)3.24H2O
2M sulphuric acid.
Zinc powder.
You are required to standardize FA1 and use it to determine the percentage of iron
in FA3.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and heat the mixture to about 700C. Titrate the hot mixture with FA1
from the burette until you reach the end point. Repeat the titration 2-3 times until
you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the molar concentration of FA1. [Na=23, C=12, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
42
Procedure II
Measure 100cm3 of FA3 using a measuring cylinder into a conical flask and add 3g
of zinc powder followed by 5 cm3 of concentrated Sulphuric acid. Heat the mixture
until a pale green solution is obtained and allow the mixture to cool. Label the
solution FA5.
Pipette 20cm3 or 25cm3 of FA5 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
Questions;
(b) Calculate the;
(i) Concentration of iron(III) ions in FA3 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of iron(III) ions in FA3 in grams per litre.
.……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
43
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iii) The percentage of iron in FA3. [N=14, H=1, S=32, O=16, Fe=56]
.……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
Experiment 20
You are provided with the following;
FA1, which is a solution 2.41g of manganate(VII) ions per litre.
FA2, which is a solution containing a mixture of ammonium iron(III) sulphate-24water, (NH4)2SO4Fe2(SO4)3.24H2O and ammonium iron(II) sulphate-6-water,
(NH4)2SO4FeSO4.6H2O
2M sulphuric acid
Zinc powder.
You are required to determine the percentage of iron(III) in FA2.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and titrate with FA1 from the burette until you reach the end point.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used…………………………………………………………...cm3
44
Procedure II
Measure 100cm3 of FA2 using a measuring cylinder into a conical flask and add 3g
of zinc powder followed by 5 cm3 of concentrated Sulphuric acid. Heat the mixture
until a pale green solution is obtained and allow the mixture to cool. Label the
solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
Questions;
(b)Calculate the;
(i) Molar concentration of FA1. [Mn=55, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of iron(II) ions in FA2 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
45
(iii) Concentration of iron(III) ions in FA2 in moles per litre.
.……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iv) The percentage of iron(III) ions in FA2. [N=14, H=1, S=32, O=16, Fe=56]
.……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
Experiment 21
You are provided with the following;
FA1, which is 0.02M solution of potassium manganate(VII).
FA2, which is a 0.08M solution of sodium hydroxide.
2M sulphuric acid
Solid M, which is an oxalate with the formula, HX(C2O4)Y.nH2O
You are required to determine the value of X, Y and n in compound M.
Procedure I
Weigh accurately about 1.6g of M and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask, add 2-3 drops of
phenolphthalein indicator and titrate the hot mixture with FA2 from the burette
until you reach the end point. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
46
Results;
Mass of empty bottle + M=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of M alone=……………………………………………………………g
Volume of pipette used=…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA2 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3
Average volume of FA2 used…………………………………………………………...cm3
Procedure I
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and heat the mixture to about 600C. Titrate the hot mixture with FA1
from the burette until the endpoint. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(a) Calculate the number of moles of;
(i) Hydrogen ions per litre in FA4 that reacted with hydroxide ions in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
47
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Oxalate ions per litre in FA4 that reacted with manganate(VII) ions in FA1
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Determine the mole ratio of hydrogen ions to oxalate ions and hence the value of
X and Y.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(c) Determine the molar mass of M and hence the value of n. [H=1, C=12, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
48
Experiment 22
You are provided with the following;
FA1, which is a solution of potassium manganate(VII) of unkown concentration.
FA2, which is a solution made by dissolving 2.016g of anhydrous sodium sulphite in
0.5dm3 of solution.
FA3, which is a 0.3M ethanedioic acid solution.
2M sulphuric acid
Solid N, which is a sample of pyrolusite.
You are required to standardize FA1 and use it to determine the percentage of
manganese(IV) oxide in pyrolusite.
THEORY
Pyrolusite is impure form of manganese(IV) oxide. When heated with excess
ethanedioic acid, it reacts according to the following equation.
MnO2(s) + H2C2O4(aq) + 2H+(aq)
Mn2+(aq) + 2CO2(g) + 2H2O(l)
Sulphite ions and oxalate ions react with manganate(VII) ions according to the
equations;
2MnO4- (aq) + 5C2O42-(aq) + 16H+(aq)
2Mn2+(aq) + 10CO2(g) + 8H2O(l)
2MnO4- (aq) + 5SO32-(aq) + 6H+(aq)
2Mn2+(aq) + 5SO42- (aq) + 3H2O(l)
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
sulphuric acid and titrate with FA1 from the burette until you reach the end point.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3
Average volume of FA1 used……………………………………………………………cm3
Questions;
(a) Calculate the molar concentration of potassium manganate(VII) in FA1.
(Na=23, S=32, O=16)
49
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 0.2g of N and place it in a beaker. Add 100cm3 of FA3 and
heat the mixture until it just dissolves. Cool the mixture and transfer the contents
into a 250cm3 volumetric flask and add distilled water to make up to the mark. Label
the solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid and heat the mixture to about 700C. Titrate the hot mixture with FA1
from the burette until the endpoint. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + N=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of N alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the Number of moles of;
(i) Excess ethanedioic acid in FA4 that reacted with FA1.
50
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Ethanedioic acid that reacted with the manganese(IV) oxide in the pyrolusite.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Percentage of manganese(IV) oxide in pyrolusite. [Mn=55, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Experiment 23
You are provided with the following;
FA1, which is a solution of potassium manganate(VII).
FA2, which is a solution of hydrogen peroxide
Solid P, which is sodium ethanedioate.
2M Sulphuric acid
You are required to standardize FA1 and use it to determine the volume strength of
hydrogen peroxide in FA2.
51
Procedure I
Weigh accurately about 1.3g of P. Add 100cm3 of water and carefully stir to dissolve.
Transfer the contents into a 250cm3 volumetric flask and add water to make up to
the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid and heat the mixture to about 600C. Titrate the hot mixture with FA1
from the burette until the endpoint. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + P=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of P alone=……………………………………………………………g
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(a) Calculate the molar concentration of FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
52
Procedure II
Using a measuring cylinder, measure 10cm3 of FA2 into a 250cm3 volumetric flask
and add distilled water to make up to the mark. Label the solution FA4.
Pipette 10cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the;
(i) Number of moles hydrogen peroxide in 250cm3 of FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of hydrogen peroxide in moles per litre in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
53
(iii) Volume strength of hydrogen peroxide in FA2. [1 mole of a gas occupies
22.4dm3 at s.t.p]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 24.
You are provided with the following;
FA1, which is a solution containing 3.16g of potassium manganate(VII) per litre of
solution.
FA2, which is a solution of diammonium iron(II) sulphate-6-water,
(NH4)2SO4FeSO4.6H2O.
Solid X, which is impure potassium chlorate(V), KClO3
2M Sulphuric acid.
A thermometer.
You are required to standardize FA2 and use it to determine the percentage purity
of potassium chlorate(V) in solid X.
THEORY.
In acidic medium, chlorate(V) ions reacts with iron(II) ions according to the
equation.
ClO3-(aq) + 6H+(aq) + 6Fe2+(aq)
6Fe3+(aq) + Cl-(aq) + 3H2O(l)
Procedure I
Pipette 10cm3 of FA2 into a clean conical flask, add an equal volume of 2M Sulphuric
acid and titrate the mixture with FA1 from the burette until the end point.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
54
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the molar concentration of diammonium iron(II) sulphate-6-water in
FA2. (K=39, Mn=55, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 0.5g of X and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and make up the solution to the mark using distilled water. Label
the solution FA3.
Pipette 10cm3 of FA3 into a clean conical flask. Add 25cm3 of FA2 using a measuring
cylinder followed by an equal volume of 2M Sulphuric acid. Heat the mixture to
about 850C and then cool for about 3 minutes. Titrate the cold mixture with FA1
from the burette until the end point.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + X=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of X alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
55
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the number of moles of;
(i) Excess iron(II) ions in FA1 that reacted with manganate(VII) ions.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Iron(II) ions that reacted with 10cm3 of chlorate(V) ions in FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Determine the percentage purity of potassium chlorate(V) in solid X.
(K=39, O=16, Cl=35.5)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
56
Experiment 25
You are provided with the following;
FA1, which is approximately 0.02M potassium manganate(VII) solution.
FA2, which is a solution of containing 5.2gdm-3 of a metal persuphate, molecular
mass=270
Solid P, which is diammonium iron(II) sulphate-6-water, (NH4)2SO4FeSO4.6H2O
2M Sulphuric acid
You are required to standardize FA1 and use it to determine the mole ratio for the
reaction between aqueous iron(II) ions and the persulphate ions.
Procedure I
Weigh accurately about 6.0g of P. Dissolve it 150cm3 of 2M Sulphuric acid and
transfer the contents into a 250cm3 volumetric flask and add water to make up to
the mark. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask and titrate with FA1 from
the burette until the endpoint. Repeat the titration 2-3 times until you obtain
consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + P=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of P alone=……………………………………………………………g
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
……………………………………………………………………………………………………………….cm3
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the molar concentration of iron(II) ions in FA4.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
57
(b) Determine the concentration of FA1 in moldm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask followed by 10cm3 of 2M
Sulphuric acid and titrate with FA1 from the burette until the endpoint. Repeat the
titration 2-3 times until you obtain consistent results.
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
Questions;
(b) Calculate the number of moles of;
(i) iron(II) ions that reacted with manganate(VII) ions in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) iron(II) ions that reacted with persulphate ions.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
58
(c) Determine the;
(i) Number of moles of persulphate ions that reacted.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Reaction mole ratio between persulphate ions and iron(II) ions.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 26.
You are provided with the following;
DA1 which Is 19.6g ammonium ferrous sulphate, (NH4)2SO4.FeSO4.6H2O in 500cm3
of solution.
DA2 which is a solution of potassium manganate (VII) of unknown concentration
Solid T which is an impure metal persulphate, YS2O8 (Y=24).
1M sulphuric acid solution
You are required to determine;
(i) The molar concentration of potassium manganate(VII) in DA2
(ii) The percentage of purity of YS2O8 in T.
Procedure I
Pipette 25cm3 (or 20cm3) of DA1 into a clean conical flask, then add 20cm3 of 1M
sulphuric acid and titrate with solution DA2 from the burette until the end point.
Repeat the titration 2-3 times until you obtain consistent results. Enter your results
in the table below.
Volume of pipette used = ……………………………………………………………….. cm3
Titration number
1
2
3
Final burette reading(cm3)
Initial burette reading(cm3)
Volume of DA2 used(cm3)
59
Titre values for calculating average volume of DA2.
……………………………………………………………………………………………………………..
Average volume of DA2 …………………………………………………………………. Cm3
Questions
(b) Calculate the molar concentration of manganate (VII) ions in DA2
(H = 1, N = 14, O = 16, S = 32, Fe = 56)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………..
……………………………………………………………………………………………..………………………………
……………………………………………………………………………..………………………………………………
……………………………………………………………..………………………………………………………………
……………………………………………..………………………………………………………………………………
……………………………..………………………………………………………………………………………………
……………..……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..………………
Procedure II
Weigh accurately about 0.5g of solid T and place it in a beaker. Add to it about
50cm3 of water and stir to dissolve. Transfer the contents of a beaker into a 250cm3
volumetric flask. Add exactly 150cm3 of DA1 to the solution in a volumetric flask
and top up with distilled water to the mark. Shake and allow to stand for 4-5
minutes. Label the resultant solution DA3.
Pipette 25cm3 (or 20cm3) of DA3 into a clean conical flask followed by 10cm3 of 1M
sulphuric acid and then titrate with DA2 from the burette until the end point. Repeat
the titration 2-3 times to obtain consistent results. Enter your results in the table
below.
Mass of T + empty bottle
= ……………………………………………. g
Mass of empty bottle
= ……………………………………………. g
Mass of T alone
= ……………………………………………. g
Volume of pipette used
= ………………………………………….. cm3
Titration number
1
2
3
Final burette reading(cm3)
Initial burette reading(cm3)
Volume of DA2 used(cm3)
60
Titre values for calculating average volume of DA2
………………………………………………………………………………………………………….. cm3
Average volume of DA2 used ……………………………………………………………… cm3
Questions
(b) Calculate the moles of;
(i) Excess iron(II) ion that reacted with the MnO4- in DA2
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
(ii) Excess iron(II) ions contained in 250cm3 of DA3
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
(iii) Iron (II) ions that reacted with persulphate ions in T.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
61
(c) Determine the percentage purity YS2O8 in T.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
THIOSULPHATE TITRATIONS.
EXPERIMENT 27.
You are provided with the following;
FA1, which is a solution of iodine.
Solid Q, which is sodium thiosulphate-5-water, Na2SO3.5H2O.
You are required to standardize FA1 using sodium thiosulphate-5-water.
Procedure
Weigh accurately about 6.2g of Q and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA2.
Pipette 20cm3 or 25cm3 of FA1 into a clean conical flask and titrate with FA2 from
the burette until the solution just turns pale yellow. Add 2cm3 of starch indicator
and continue titrating until the blue-black solution turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + Q=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of Q alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA2 used (cm3)
62
Titre values used to calculate average volume of FA2.
………………………………………………………………………………………………………cm3
Average volume of FA2 used ………………………………………………… cm3
Questions;
Calculate the number of moles of;
(i) Sodium thiosulphate in FA2 that reacted with FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of iodine in FA1 in grams per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
EXPERIMENT 28.
You are provided with the following;
FA1, which is a solution of sodium thiosulphate.
Solid T, which is potassium iodate, KIO3
2M Sulphuric acid.
10% potassium iodide.
You are required to standardize FA1 using potassium iodate.
Procedure.
Weigh accurately about 0.9g of T and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and add water to make up to the mark. Label the solution FA2.
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution and titrate the
liberated iodine with FA1 from the burette until the solution just turns pale yellow.
Add 2cm3 of starch indicator and continue titrating until the blue-black solution
turns colourless.
63
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + T=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of T alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
Calculate the concentration of;
(i) Sodium thiosulphate in FA1 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium thiosulphate in FA1 in grams per litre.
.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
64
EXPERIMENT 29.
You are provided with the following;
FA1, which is a solution of sodium thiosulphate.
FA2, which is a 0.02M solution of potassium manganate(VII).
Solid U, which is copper(II) sulphate-5-water.
2M Sulphuric acid.
10% potassium iodide.
You are required to standardize FA1 and use it to determine the percentage of
copper in the sample U.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette until the solution just turns pale yellow.
Add 2cm3 of starch indicator and continue titrating until the blue-black solution
turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the concentration of FA1 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
65
Procedure II
Weigh accurately about 6.0g of U and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask. Add sodium carbonate solution drop-wise until a slight blue
precipitate has formed and make up the solution to the mark using distilled water.
Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 10cm3 of 10%
potassium iodide solution and titrate the liberated iodine with FA1 from the burette
until the solution just turns pale yellow. Add 2cm3 of starch indicator and continue
titrating until the blue-black solution turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + U=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of U alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the number of moles of;
(i) Iodine liberated by FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Copper(II) ions in FA3 in 250cm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
66
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(b) Determine the percentage of copper in U.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 30.
You are provided with the following;
FA1, which is approximately a 0.1M solution of sodium thiosulphate.
FA2, which is a solution containing 2.4gdm3 of potassium iodate.
Solid V, which is a salt containing dichromate ions.
2M Sulphuric acid.
5% potassium iodide.
Starch solution.
You are required to standardize FA1 and use it to determine the percentage by mass
of chromium in V.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 5% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
67
Questions;
(a) Calculate the number of moles of iodine liberated by FA2. (K=39, I=127, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(b) Determine the concentration of FA1 in moles per dm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 1.2g of V and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and make up the solution to the mark using distilled water. Label
the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 5% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + V=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of V alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
68
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(c) Calculate the number of moles of iodine liberated by FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(d) Determine the;
(i) Concentration of FA3 in moles per dm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Mass of chromium in V and hence its percentage. (Cr = 52)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
69
EXPERIMENT 31.
You are provided with the following;
FA1, which is a solution of sodium thiosulphate.
FA2, which is a solution containing 2.0g of potassium dichromate(V) in 500cm3 of
solution.
Solid W, which is potassium halate(V), KYO3.
2M Sulphuric acid.
10% potassium iodide.
Starch solution.
You are required to standardize FA1 and use it to determine the relative atomic
mass of element Y in W.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the concentration of FA1 in moles per litre. (K=39, Cr=52, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
70
Procedure II
Weigh accurately about 1.0g of W and place it in a beaker. Add 100cm3 of water and
carefully stir to dissolve. Transfer the contents of the beaker into a 250cm3
volumetric flask and make up the solution to the mark using distilled water. Label
the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + W=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of W alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
Questions;
(b) Calculate the number of moles of iodine liberated by FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Determine the;
(i) Concentration of FA3 in moles per dm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
71
(ii) Relative molecular mass of V and hence the relative atomic mass of Y in KYO3.
(K=39, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 32.
You are provided with the following;
FA1, which is approximately a 0.1M solution of sodium thiosulphate.
FA2, which is a solution containing 2.41g of manganate(VII) ions per litre.
FA3, which is a solution of hydrogen peroxide.
2M Sulphuric acid.
10% potassium iodide.
You are required to standardize FA1 and use it to determine the concentration of
hydrogen peroxide in FA3.
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette until the solution just turns pale yellow.
Add 2cm3 of starch indicator and continue titrating until the blue-black solution
turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
72
Questions;
(a) Calculate the concentration of FA1 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Using a measuring cylinder, measure 10cm3 of FA3 into a 250cm3 volumetric flask
and add distilled water to make up to the mark. Label the solution FA4.
Pipette 20cm3 or 25cm3 of FA4 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution and titrate the
liberated iodine with FA1 from the burette until the solution just turns pale yellow.
Add 2cm3 of starch indicator and continue titrating until the blue-black solution
turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the;
(i) Number of moles hydrogen peroxide in 250cm3 of FA4.
73
3
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Concentration of hydrogen peroxide in moles per litre in FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Volume strength of hydrogen peroxide in FA3. [1 mole of a gas occupies
22.4dm3 at s.t.p]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 33.
You are provided with the following;
FA1, which is approximately a 0.1M solution of sodium thiosulphate.
FA2, which is a solution prepared by dissolving 1.0g of potassium iodate in 250cm3
of distilled water.
Liquid X, which is a commercial bleaching agent containing sodium hypochlorite.
FA3; which is 0.5M potassium iodide.
2M Sulphuric acid.
You are required to standardize FA1 and use it to determine the percentage of
chlorine in the commercial bleaching agent.
74
THEORY
In acidic medium, iodate ions react with iodide ions according to the equation;
IO3-(aq) + 6H+(aq) + 5I-(aq)
3I2(aq) + 5H2O(l)
In acidic medium, hypochlorite ions produce chlorine according to the equation;
Cl-(aq) + OCl-(aq) + 2H+(aq)
H2O(l) + Cl2(g)
The chlorine produced displaces iodine from potassium iodide according to the
equation;
Cl2(g) + 2I-(aq)
I2(aq) + 2Cl-(aq)
The iodine liberated in both cases reacts with thiosulphate ions according to the
equation;
I2(g) + 2S2O32-(aq)
2I-(aq) + S4O62-(aq)
Procedure I
Pipette 10cm3 of FA2 into a clean conical flask, add an equal volume of 2M Sulphuric
acid followed by 10cm3 of FA3. Titrate the liberated iodine with FA1 from the
burette until the solution just turns pale yellow. Add 2cm3 of starch indicator and
continue titrating until the blue-black solution turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the concentration of FA1 in moles per litre.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
75
Procedure II
Using a measuring cylinder, measure 30cm3 of liquid X into a 250cm3 volumetric
flask and add distilled water to make up to the mark. Label the solution FA5.
Pipette 10cm3 of FA5 into a clean conical flask. Add an equal volume of 2M
Sulphuric acid followed by 10cm3 of FA3 and titrate the liberated iodine with FA1
from the burette until the solution just turns pale yellow. Add 2cm3 of starch
indicator and continue titrating until the blue-black solution turns colourless.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used ………………………………………………… cm3
Questions;
(b) Calculate the;
(i) Number of moles sodium thiosulphate in FA1 that reacted.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Number of moles of iodine liberated by 10cm3 of FA5.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Number of moles of chlorine in 30cm3 of X.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
76
(iv)The mass of chlorine in 30cm3 of X. (Cl=35.5)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(V) The percentage by mass of chlorine in X.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 34.
You are provided with the following;
FA1, which is a solution of sodium thiosulphate.
FA2, which is a solution containing 2.94g of dichromate(V) ions per litre.
Solid Y, which is calcium hypochlorite (bleaching powder).
2M Sulphuric acid.
2M Ethanoic acid.
10% potassium iodide.
Starch indicator.
You are required to standardize FA1 and use it to determine the relative atomic
mass of element Y in W.
THEORY
In acidic medium, dichromate(VI) ions react with iodide ions according to the
equation;
Cr2O72-(aq) + 14H+(aq) + 6I-(aq)
2Cr3+(aq) + 3I2(aq) + 7H2O(l)
In acidic medium, hypochlorite ions produce chlorine according to the equation;
Cl-(aq) + OCl-(aq) + 2H+(aq)
H2O(l) + Cl2(g)
The chlorine produced displaces iodine from potassium iodide according to the
equation;
Cl2(g) + 2I-(aq)
I2(aq) + 2Cl-(aq)
The iodine liberated in both cases reacts with thiosulphate ions according to the
equation;
I2(aq) + 2S2O32-(aq)
2I-(aq) + S4O62-(aq)
Procedure I
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
77
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3.
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the concentration of FA1 in moles per litre. (Cr=52, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 1.3g of Y and place it into a clean beaker, add little water
and grind to form a thin paste. Add more little water and carefully pour the milky
liquid into a 250cm3 volumetric flask and make up the solution to the mark using
distilled water. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask, add an equal volume of 2M
ethanoic acid followed by 10cm3 of 10% potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + Y=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of Y alone=……………………………………………………………g
78
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the
(i) Number of moles of iodine liberated.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
(ii) Concentration of chlorine in FA3 in moles per dm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iii) Mass of chlorine and hence the percentage of chlorine in Y. (Cl=35.5)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
79
EXPERIMENT 35.
You are provided with the following;
FA1, which is a solution of sodium thiosulphate.
FA2, which is a solution containing 10.7g of a mixture of potassium iodate and
potassium iodide per litre.
Solid Y, which is potassium manganate(VII) crystals.
2M Sulphuric acid.
O.5M potassium iodide solution.
Starch indicator.
You are required to standardize FA1 and use it to determine the percentage of
potassium iodate in FA2.
Procedure I
Weigh accurately about 0.8g of Y and place it into a clean beaker, add little water to
dissolve and transfer the contents into a 250cm3 volumetric flask and make up the
solution to the mark using distilled water. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask, add an equal volume of 2M
Sulphuric acid followed by 10cm3 of 0.5M potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results
Mass of empty bottle + Y=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of Y alone=……………………………………………………………g
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
…………………………………………………………………………………………………………………..cm3
Average volume of FA1 used …………………………………………………………………...cm3
Questions;
(a) Calculate the concentration of FA1 in moles per litre. (K=39, Mn=55, O=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
80
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Pipette 20cm3 or 25cm3 of FA2 into a clean conical flask, add an equal volume of 2M
ethanoic acid followed by 10cm3 of 0.5M potassium iodide solution. Titrate the
liberated iodine with FA1 from the burette using starch indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA1 used (cm3)
Titre values used to calculate average volume of FA1.
………………………………………………………………………………………………………cm3
Average volume of FA1 used …………………………………………………………..cm3
Questions;
(b) Calculate the
(i) Number of moles of iodine liberated.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
(ii) Concentration of potassium iodate in FA2 in moles per dm3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
81
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(iii) Percentage of potassium iodate in FA2. (K=39, I=127, 0=16)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
EXPERIMENT 36.
You are provided with the following;
FA1, which is a solution of iodine.
FA2, which is a 0.1M solution of sodium thiosulphate.
Solid R, which is a reducing agent of molecular mass 126.
Solid S *NaHCO3
You are required to standardize FA1 and use it to determine the mole ratio for the
reaction between iodine and R.
Procedure I
Pipette 10cm3 of FA1 into a clean conical flask and titrate with FA2 from the burette
using starch indicator. Repeat the titration 2-3 times until you obtain consistent
results.
Enter your results in the table below.
Results
Volume of pipette used…………………………………..cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA2 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………………….cm3
Average volume of FA2 used …………………………………………………………………...cm3
82
Questions;
(a) Calculate the concentration of iodine in moles per litre of FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Procedure II
Weigh accurately about 1.6g of R and place it in a beaker. Add 100cm3 of distilled
water and carefully stir to dissolve. Transfer the contents of the beaker into a
250cm3 volumetric flask and add water to make up to the mark. Label the solution
FA3.
Pipette 10cm3 of FA3 into a clean conical flask. Add 20cm3 of FA1 using a measuring
cylinder followed by two spatula end-fuls of solid S and shake the mixture well.
Titrate the excess iodine in the mixture with FA2 from the burette using starch
indicator.
Repeat the titration 2-3 times until you obtain consistent results.
Enter your results in the table below.
Results;
Mass of empty bottle + R=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of R alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading (cm3)
Volume of FA2 used (cm3)
Titre values used to calculate average volume of FA2.
………………………………………………………………………………………………………cm3
Average volume of FA2 used …………………………………………………………..cm3
Questions;
(b) Calculate the number of moles of;
83
(i) Excess iodine in FA1 that reacted with the thiosulphate ions in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Iodine that reacted with X.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(c) Determine the;
(i) The number of moles of X that reacted.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
(ii) Reaction mole ratio of X and iodine.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
84
PHYSICAL CHEMISTRY PRACTICALS
COLLIGATIVE PROPERTIES
Experiment 37
You are provided with the following;
Naphthalene (C10H8)
Camphor (C10H16O)
A thermometer.
You are required to determine the freezing point constant, Kf per kg of naphthalene.
Procedure.
(a) Pour 200cm3 of water into a 250cm3 beaker.
Heat the water to boiling point on a tripod stand.
Weigh separately and accurately about 5.0g of naphthalene and 1.0g of camphor.
Record the results in the spaces provided below.
Results.
Mass of container + Naphthalene = ……………………………………………………..g
Mass of empty container
= ……………………………………………………..g
Mass of naphthalene
= ……………………………………………………..g
Mass of container + Camphor
= ……………………………………………………..g
Mass of empty container
= ……………………………………………………..g
Mass of Camphor
= ……………………………………………………..g
(b) Transfer the whole of naphthalene into a clean dry boiling tube. Immerse the
boiling tube containing naphthalene into the beaker of hot water and continue
heating the water until the whole naphthalene melts.
Insert a thermometer in the liquid formed and heat to about 890C. Remove the
boiling tube from the hot water and start a stop clock when its temperature drops to
850C. Allow the liquid to cool while stirring with a thermometer and record its
temperature after every half a minute for three minutes. Enter your results in the
table below.
(c) Transfer the whole of camphor into a boiling tube containing naphthalene.
Immerse the boiling tube into the beaker of hot water and continue heating the
water until the mixture melts. Continue heating until the temperature of the molten
mixture is about 890C.
Remove the boiling tube from the hot water and start a stop clock when its
temperature of the mixture drops to 850C. Allow the mixture to cool while stirring
with a thermometer and record its temperature after every half a minute for three
minutes. Enter your results in the table below.
Time (minutes)
0.0 0.5 1.0 1.5 2.0 2.5 3.0
Temperature of naphthalene(0C)
Temperature of mixture of
naphthalene and camphor (0C)
Questions:
(a) Plot on the same axes, a graph of temperature;
85
(i) Pure naphthalene against time.
(ii) Mixture of naphthalene and camphor against time.
(b) From the graphs, read off the temperature after 2.5 minutes of;
(i) Pure naphthalene.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Mixture of naphthalene and camphor.
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………......
(c) Use the temperatures you have obtained in (b) above to determine the
depression in freezing point of naphthalene.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(d) Calculate the freezing point depression constant, Kf per kg of naphthalene.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
86
87
Experiment 38.
You are provided with the following;
Naphthalene
Camphor
A thermometer.
You are required to determine the relative molecular mass of camphor using
freezing point depression method.
Procedure.
Weigh accurately about 5.0g of naphthalene into a dry boiling tube, then insert a
thermometer into the tube.
Put the boiling tube and its contents in a beaker of boiling water. Stir continuously
with a thermometer until the whole solid melts.
Remove the beaker of boiling water from the heat source while keeping the boiling
tube in the water and allow the melted solid to cool while stirring gently with a
thermometer.
Record the steady temperature, T10C when the crystals start to appear.
Weigh accurately about 1.0g of camphor, carefully add it to naphthalene in the
boiling tube.
Put back the boiling tube into a beaker of boiling water. Stir continuously with a
thermometer until the whole mixture melts.
Remove the beaker of boiling water from the heat source while keeping the boiling
tube in the water and allow the melted mixture to cool while stirring gently with a
thermometer.
Record the steady temperature, T20C when the crystals start to appear.
Record your results in the spaces below.
Results.
Mass of naphthalene used = ………………………………………..g
Mass of camphor used
= ………………………………………..g
Freezing point of naphthalene = ………………………………….g
Freezing point of the mixture = …………………………………..g
Questions:
Calculate the;
(a) Freezing point depression, T30C
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(b) Relative molecular mass of camphor given that the freezing point constant of
naphthalene is 6.90C/mol/kg
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
88
89
THERMOCHEMISTRY
Experiment 39
You are provided with the following:
FA1, which is a solution of hydrochloric acid.
Metal X
Substance Y, which is an oxide of X, with a formula XO
You are required to determine the enthalpy change for the reduction of the oxide of
X and comment on your answer.
PART I
Procedure:
Weigh accurately about 1.2g of X.
Using a measuring cylinder, transfer 100cm3 of FA1 into a plastic beaker. Read and
record its initial temperature, in the table I.
Add at once the 1.2g of X into FA1 into the plastic beaker and at the same time start
the stop clock or watch.
Stir gently with the thermometer and record the temperature of the solution after
every half-minute in table I, up to the four minute.
Results
Mass of the weighing container + X = …………………………… g
Mass of the weighing container alone = …………………………… g
Mass of X used
= ……………………........... g
Table I
Time (minutes)
0
½
1
1½ 2
2½ 3
3½ 4
Temperature of
solution (0C)
PART II
Procedure
Weigh accurately about 2.0g of Y.
Using a measuring cylinder, transfer 100cm3 of FA1 into a plastic beaker. Read and
record its initial temperature, in the table II.
Add at once the 2.0g of Y into FA1 into the plastic beaker and at the same time start
the stop clock or watch.
Stir gently with the thermometer and record the temperature of the solution after
every half-minute in table II, up to the four minute.
Results
Mass of the weighing container + Y = …………………………… g
Mass of the weighing container alone = …………………………… g
Mass of Y used
= ……………………........... g
Table II
Time (minutes)
0
½
1
1½ 2
2½ 3
3½ 4
Temperature of
solution (0C)
90
(a) (i) Plot on the same axes, graphs of temperature against time for results
obtained in both Part I and Part II.
(ii) From your graphs, determine the maximum temperature change for the
reaction.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Calculate the heat of reaction in Part I and Part II. [Specific heat capacity of the
solution=4.2 Jg-1K-1 and its density=1gcm-3 in each case, equations of the reactions of
X and Y are as follows]
X(s) + 2HCl(aq)
XCl2(aq) + H2(g)
XO(s)+ 2HCl(aq)
XCl2(aq) + H2O(l)
Part I
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Part II
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iv) Calculate the molar heat of reactions in Part I and Part II. [X=24, O=16]
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
91
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(b) Determine the heat energy change for the reaction.
XO(s) + H2 (g)
X(s) + H2O(l)
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Comment on your answer in (b).
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
92
93
Experiment 40
You are provided with the following;
FA1; which is 2M sodium hydroxide solution
FA2; which is 2M solution of a strong acid
You are required to determine the basicity and heat of neutralization of the acid in
FA2
Procedure.
Pour 50cm3 of FA2 into a plastic beaker using a measuring cylinder. Using the same
measuring cylinder, add 10cm3 of water to the plastic beaker. Stir and record the
temperature of the mixture, T10C in the table below.
Using a clean measuring cylinder, measure 60cm3 of FA1 into another plastic beaker
and record its temperature, T20C in the table below.
Pour the 60cm3 of FA1 into the plastic beaker containing FA2 and stir gently with a
thermometer. Record the maximum temperature of the mixture, T30C.
Repeat the procedures above with different volumes of water, FA1 and FA2 given in
the table below. Make sure the plastic beakers and the thermometer used between
each experiment is rinsed.
Record your results in the table below.
Experiment number
1
2
3
4
5
6
7
8
Volume of FA2 (cm3)
50
40
35
30
25
20
15
10
Volume of water (cm3)
10
20
25
30
35
40
45
50
Volume of FA1 (cm3)
60
60
60
60
60
60
60
60
Temperature, T1 (0C)
Temperature, T2 (0C)
Average temperature,
𝑇1 + 𝑇2 0
( C)
2
Maximum temperature, T3 (0C)
Temperature rise, T3 -
𝑇1 + 𝑇2
2
(0C)
Questions
94
(a) Plot a graph of temperature rise against volume of FA2 used.
(b) From the graph, determine the;
(i) Maximum temperature rise
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Volume of FA2 required to produce the maximum temperature rise.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Calculate the number of moles of the;
(i) Acid in FA2 that completely reacted with sodium hydroxide in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Sodium hydroxide in FA1 that completely reacted with the acid in FA2.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(d) Calculate the basicity of the acid.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(e) Calculate the heat of neutralization of the acid. (Specific heat capacity of solution
is 4.2j/g/0C and the density of solution is 1gcm-3)
95
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
96
97
Experiment 41
You are provided with the following;
FA1; which is 0.25M copper(II) sulphate solution.
Solid X, which is zinc powder.
You are required to determine the heat of displacement of copper(II) sulphate
solution using zinc.
Procedure
Weigh roughly 1.3g of X.
Place 50cm3 of FA1 into a plastic beaker and record the temperature of the solution
every minute for four minutes.
On exactly the fifth minute, add solid X to the solution and stir continuously with the
thermometer. Record the temperature after every minute until 11 minutes.
Record your results in the table below.
Results
Mass of the weighing container + X = …………………………… g
Mass of the weighing container alone = …………………………… g
Mass of X used
= ……………………........... g
Table
Time (min)
0
1
2
3
4
5
6
7
8
9
10
11
Temperature
(0C)
(a) Plot a graph of temperature against time.
(b) From the graph, determine the;
(i) Experimental (observed) temperature rise.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Corrected (theoretical) temperature rise.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Hence calculate the;
(i) Experimental enthalpy of displacement of copper(II) sulphate solution.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
98
(ii) Heat of displacement of copper(II) sulphate solution using the corrected
temperature rise.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………...........................................
99
100
PARTITION COEFFICIENT
Experiment 42
You are provided with the following;
FA1, which is 1.5M aqueous ammonia.
FA2, which is trichloromethane
FA3, which is 0.5M hydrochloric acid.
FA4, which is 0.05M hydrochloric acid.
You are required to determine the partition coefficient of ammonia between water
and trichloromethane.
Procedure.
Add 50cm3 of FA1 to 50cm3 of FA2 in a conical flask. Shake vigorously for about 5
minutes and allow to stand for about 10 minutes.
Carefully decant the upper (aqueous) layer into a beaker. Transfer the lower
(organic) layer into another beaker and cover.
Read and record the room temperature.
Pipette 10cm3 of the aqueous layer into a conical flask, add 2-3 drops of methyl
orange indicator and titrate with FA3 from the burette. Repeat the titration 2-3
times to obtain consistent results. Record your results in table I below.
Results
Room temperature = …………………………………………………………………………..0C
Volume of pipette used=……………………………………………………………………... cm3
Table I
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA3 used(cm3)
Titre values used for calculating average volume of FA3 used
………………………………………………………………………………………………………cm3
Average volume of FA3 used …………………………………………………………………...cm3
Pipette 10cm3 of the organic layer into a conical flask, add 2-3 drops of methyl
orange indicator and titrate with FA4 from the burette. Repeat the titration 2-3
times to obtain consistent results. Record your results in table II below.
Results
Volume of pipette used=…………………………………………….cm3
Table II
101
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA4 used(cm3)
Titre values used for calculating average volume of FA4 used
………………………………………………………………………………………………………cm3
Average volume of FA4 used …………………………………………………………………...cm3
Questions
(a) Calculate the;
(i) Molar concentration of ammonia in the aqueous layer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………..…………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………..
(ii) Molar concentration of ammonia in the organic layer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
102
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………..…………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………….
(b) Determine the value of the distribution coefficient of ammonia between water
and trichloromethane at the temperature you have recorded.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………….......................................................................................
Experiment 43
You are provided with the following;
FA1, which is 0.5M sodium thiosulphate solution.
FA2, which is 0.005M sodium thiosulphate solution.
FA3, which is 0.1M potassium iodide solution.
Solid D, which is iodine crystals.
You are required to determine the partition coefficient of iodine between
tretrachloromethane and water.
Procedure
Weigh accurately 1.0g of D and dissolve it in tetrachloromethane to make 50cm3 of
solution.
Transfer the solution into a separating funnel and add about 200cm3 of distilled
water. Cork the funnel and shake the mixture vigorously for every half a minute for
about 10-15 minutes and allow to stand for about 10 minutes.
Carefully tap off the lower (organic) layer into a beaker.
Pipette 10cm3 of the organic layer into a conical flask, add 20cm3 of FA3 and titrate
with FA1 from the burette using starch indicator. Repeat the titration 2-3 times to
obtain consistent results. Record your results in table I below.
103
Results
Mass of empty bottle + D=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of D alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………cm3
Average volume of FA1 used…………………………………………………………………...cm3
Pipette 25cm3 of the aqueous layer into a conical flask and titrate with FA2 from the
burette using starch indicator. Repeat the titration 2-3 times to obtain consistent
results. Record your results in table II below.
Results
Volume of pipette used=…………………………………………….cm3
Table II
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA2 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3
Average volume of FA2 used …………………………………………………………………...cm3
Questions
(a) Calculate the;
(i) Molar concentration of ammonia in the organic layer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
104
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………..…………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Molar concentration of ammonia in the aqueous layer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
…………………………………………………………………..…………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(b) Determine the value of the distribution coefficient of iodine between
tetrachloromethane and water at room temperature.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
105
SOLUBILITY
Experiment 44
You are provided with the following;
FA1, which is approximately a 0.1M solution of hydrochloric acid.
FA2, which is a solution containing 19.10g of borax, Na2B4O7.10H2O in 1dm3 of
solution.
Solid W, which is calcium hydroxide
You are required to standardize FA1 and use it to determine the solubility product
of calcium hydroxide.
Borax reacts with hydrochloric acid according to the equation:
Na2B4O7.10 H2O(aq) + 2HCl(aq)
2NaCl(aq) + 4H3BO3(aq) + 5H2O(l)
Procedure I
Pipette 20cm3 or 25cm3 of FA3 into a clean conical flask. Add 2-3 drops of methyl
orange indicator and then titrate with FA1 from the burette until you reach the end
point. Repeat the titration 2-3 times until you obtain consistent results.
Record your results in the table I.
Results
Volume of pipette used=…………………………………..cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA2 used
………………………………………………………………………………………………………cm3
Average volume of FA2 used …………………………………………………………………...cm3
Questions;
(a) Calculate the molar concentration of hydrochloric acid in FA1.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
106
Procedure
Weigh accurately 2.0g of W and transfer the whole mass into a conical flask.
Add 100cm3 of distilled water, cork the flask and shake the mixture vigorously for
15 minutes and filter. Label the solution FA3.
Pipette 20cm3 or 25cm3 of FA3 into a conical flask, add 2-3 drops of phenolphthalein
indicator and titrate with FA1 from the burette. Repeat the titration 2-3 times to
obtain consistent results. Record your results in table I below.
Results
Mass of empty bottle + W=……………………………………………..g
Mass of empty bottle alone=………………………………………….g
Mass of W alone=……………………………………………………………g
Volume of pipette used=…………………………………………….cm3
Table 1
Titration number
1
2
3
Final burette reading (cm3)
Initial burette reading(cm3)
Volume of FA1 used(cm3)
Titre values used for calculating average volume of FA1 used
………………………………………………………………………………………………………………….cm3
Average volume of FA1 used …………………………………………………………………...cm3
Questions
(b) Calculate the;
(i) Number of moles of hydrochloric acid in FA1 that reacted with FA3.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
(ii) Molar concentration of hydroxide ions.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
107
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Solubility product, Ksp of calcium hydroxide.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
CHEMICAL KINETICS
Experiment 45
You are provided with the following;
FA1, which is a solution of iodine.
FA2, which is a solution of propanone.
FA3, which is a solution of sodium thiosulphate.
1M Sulphuric acid.
Sodium hydrogencarbonate solution.
You are required to determine the order of reaction with respect to iodine and the
rate constant. (Assume that the acid and propanone are in large excess)
Propanone reacts with iodine in the presence of an acid according to the following
equation.
CH3COCH3(aq) + I2(aq)
H+(aq)
CH3COCH2I(aq) + HI(aq)
Procedure
(a) Using a measuring cylinder, measure 25cm3 of FA2 into a beaker followed by
25cm3 of 1M Sulphuric acid.
(b) Using a measuring cylinder, 50cm3 of FA1 is added into the beaker and at the
same time the stop clock is started. The mixture is shaken well for about 1 minute.
Label the resultant mixture FA4.
108
(c) After exactly five minutes, pipette 10cm3 of FA4 into a conical flask followed by
an equal volume of sodium hydrogencarbonate solution and shake well until
evolution of a gas stops. Titrate the mixture with FA3 from the burette using starch
indicator.
(d) Repeat procedure (c) at intervals of five minutes and enter your results in the
table below.
Table
Time (minutes)
5
10
15
20
25
Final burette reading(cm3)
Initial burette reading (cm3)
Volume of FA3 used (cm3)
Questions
(a) Plot a graph of volume of FA3 against time.
(b) From the graph,
(i) Deduce the order of reaction with respect to iodine and give a reason for answer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Determine the rate constant for the reaction.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(c) State the role of;
(i) Sodium hydrogencarbonate solution.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Sulphuric acid.
……………………………………………………………………………………………………………………………
109
Experiment 46
You are provided with the following;
FA1, which is a 0.02M solution of hydrogen peroxide.
FA2, which is prepared by dissolving 3.2g of potassium manganate(VII) in water to
make 1 litre of solution.
2M sodium hydroxide solution.
1M Sulphuric acid.
1M iron(III) chloride solution.
The decomposition of hydrogen peroxide catalyzed by iron(III) ions under alkaline
conditions proceeds according to the following equation.
2H2O2(aq)
Fe3+(aq)
2H2O(l) + O2(g)
The undecomposed hydrogen peroxide can then be titrated with potassium
manganate(VII) solution.
2MnO4-(aq) + 6H+(aq) + 5H2O2(aq)
2Mn2+(aq) + 8H2O(l) + 5O2(g)
The volume of the manganate(VII) used is directly proportional to the amount of
hydrogen peroxide undecomposed.
You are required to determine the order of reaction with respect to hydrogen
peroxide and the rate constant for the reaction.
Procedure.
(a) Pipette 10cm3 of FA1 into a conical flask. Add 1cm3 of sodium hydroxide solution
followed by 1cm3 of iron(III) chloride solution and the stop clock started at the same
time. Shake the flask well and leave it to stand for 2 minutes.
(b) After 2 minutes, add 20cm3 of 1M Sulphuric acid and titrate the mixture with
FA2 from the burette until the end point.
(c) Repeat procedure (a) and (b) but allow the hydrogen peroxide to decompose for
4, 6, 8 and 10 minutes before adding 1M Sulphuric acid.
(d) Record the results in the table below.
Time (minutes)
2
Volume of FA2 used, Vt (cm3)
logVt
110
4
6
8
10
Questions
(a) Plot a graph of logVt against time.
(b) Use the graph to;
(i) Deduce the order of reaction with respect to hydrogen peroxide and give a
reason for your answer.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………...
(ii) Determine the rate constant and state its units. Hence determine the half-life for
the reaction.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
………………………………………………………………………………………………..……………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) State the role of Sulphuric acid.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
Experiment 47
You are provided with the following;
FA1, which is 0.12M sodium thiosulphate solution.
FA2, which is 1.0M hydrochloric acid
Distilled water.
You are required to determine the rate constant and hence the half-life of the
reaction.
Theory
Sodium thiosulphate reacts with hydrochloric acid according to the equation below;
111
Na2S2O3(aq) + 2HCl(aq)
S(s) + 2NaCl(aq) + H2O(l) + SO2(g)
Procedure
(a) Make a small cross on a white paper using a pen.
(b) Using a measuring cylinder, measure 50cm3 of FA1 into a conical flask.
(c) Place the conical flask on the cross.
(d) Using a clean measuring cylinder, add 10cm3 of FA2 to the flask and
simultaneously start the stop clock. Shake the mixture to mix the solutions
thoroughly. Look at the cross through the solution from above and stop the clock
when the cross just becomes invisible.
(e) Record the time taken in the table below.
(f) Repeat the procedure (b) to (e) by measuring 40, 30, 20 and 10cm3 of FA1 into
the conical flask. In each case make up the total volume of solution to 50cm3 by
adding distilled water before adding 10cm3 of FA2.
(g) Record your results in the table below.
Volume of FA1, V
Volume of distilled
(cm3)
water added (cm3)
50
0
40
10
30
20
20
30
10
40
Time (s)
Log10V
Questions
(a) Plot a graph of Log10V against time.
(b) Deduce the order of reaction with respect to sodium thiosulphate.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(c) Determine the;
(i) Slope of the graph.
112
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(ii) Rate constant, K for the reaction.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
(iii) Half-life for the reaction.
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
……………………………………………………………………………………………………………………………
113
114
QUALITATIVE INORGANIC ANALYSIS
This involves identifying cation(s) and anion(s) in the given substance or mixture
of substances. The cations normally tested can be grouped according to their
colour in solution as follows:
𝐍𝐇𝟒+ (aq)
NON – COLOURED CATIONS IN SOLUTION
1.
It forms no precipitate with sodium hydroxide solution.
NB:
 Ammonium ions are confirmed by adding sodium hydroxide solution and
warming the mixture.
Observation
A colourless gas with a chocking smell that turns damp red litmus paper blue and
forms dense white fumes with concentrated hydrochloric acid.
−
NH4+ (aq) + OH(aq)
NH3(g) + H2O(l)
Note;
When the ammonium salt is heated with sodium hydroxide , ammonia gas is given
off . This is the confirmatory test for ammonium cation.
2. Ba 2+(aq), Ca2+(aq) and Mg2+(aq)
(a) Addition of sodium hydroxide solution.
These form a white precipitate insoluble in excess.
Ba2+(aq) + 2OH-(aq)
Ba(OH)2(s)
Ca2+(aq) + 2OH-(aq)
Ca(OH)2 (s)
Mg2+(aq) + 2OH − (aq)
Mg(OH)2 (s)
(b) Addition of ammonia solution.
(i) Ba 2+(aq) and Mg2+(aq)
They form a white precipitate insoluble in excess.
Ba2+(aq) + 2OH-(aq)
Ba(OH)2(s)
Mg2+(aq) + 2OH − (aq)
Mg(OH)2 (s)
(ii) Ca2+(aq)
It gives no observable change.
Note.
 Magnesium and barium ions can be distinguished from calcium ions by using
ammonia solution.
 Magnesium ions are confirmed as follows.
Test.
Add solid ammonium chloride followed by disodium hydrogenphosphate
solution followed by excess ammonia solution.
Observation.
A white precipitate insoluble in ammonia solution.
Mg2+(aq) + Na2HPO4(aq) + NH4+(aq)
MgNH4PO4(s) + 2Na+(aq) + H+(aq)
 Barium ions are confirmed by;
115
 Adding potassium chromate(VI) solution followed by excess sodium
hydroxide solution..
Observation.
A yellow precipitate insoluble in excess sodium hydroxide solution.
Ba2+(aq) + CrO42-(aq)
BaCrO4(s)
 Adding ammonium oxalate solution followed by excess ethanoic acid.
Observation
A white precipitate soluble in ethanoic acid.
Ba2+(aq) + C2O42-(aq)
BaC2O4(s)
BaC2O4(s) + 2CH3COOH(aq)
(CH3COO)2Ba(aq) + H2C2O4(aq)
 Calcium ions are confirmed by;
Test
Adding ammonium oxalate solution followed by excess ethanoic acid.
Observation
A white precipitate insoluble in excess ethanoic acid.
Ca2+(aq) + C2O42-(aq)
CaC2O4(s)
3. Zn2+(aq) , Al3+(aq), Pb 2+(aq) and Sn2+(aq)
(a) Addition of sodium hydroxide solution
These form a white precipitate soluble in excess to form a colourless solution.
i.e They form amphoteric hydroxides.
Zn2+(aq) + 2OH-(aq)
Zn(OH)2(s)
Zn(OH)2(s)+ 2OH (aq)
Zn(OH)42-(aq)
Al3+(aq) + 3OH-(aq)
Al(OH)3(aq) + OH-(aq)
Al (OH)3(s)
Al(OH)4-(aq)
Pb2+ + 2OH-(aq)
Pb(OH)2(s) + 2OH-(aq)
Pb(OH)2(s)
Pb(OH)42-(aq)
Sn2+(aq) + 2OH-(aq)
Sn(OH)2(s)
Sn(OH)2(s) + 2OH-(aq)
Sn(OH)42-(aq)
(b) Addition of ammonia solution
(i) Zn2+(aq)
It forms a white precipitate soluble in excess to form a colourless solution.
Zn2+(aq) + 2OH-(aq)
Zn(OH)2(s)
Zn(OH)2(s)+ 4NH3(aq)
Zn(NH3)42+(aq) + 2OH-(aq)
(ii) Al3+(aq), Pb 2+(aq) and Sn2+(aq)
They form a white precipitate insoluble in excess.
Al3+(aq) + 3OH-(aq)
Al(OH)3(s)
Pb2+ + 2OH-(aq)
Pb(OH)2(s)
2+
Sn (aq) + 2OH (aq)
Sn(OH)2(s)
Note.
 Zn2+(aq) are distinguished from Al3+(aq) Pb2+(aq) and Sn2+(aq) by using ammonia
solution.
116
 Zn2+(aq) are confirmed by adding solid ammonium chloride followed by disodium
hydrogenphosphate solution and excess ammonia solution.
Observation.
A white precipitate soluble in excess ammonia solution is observed.
 Pb2+ (aq) ions can be confirmed by;
 Adding potassium iodide solution.
Observation
A yellow precipitate.
Pb2+(aq) + 2I-(aq)
PbI2 (s)
 Adding potassium chromate(VI) solution followed by excess sodium
hydroxide solution.
Observation
A yellow precipitate soluble in excess sodium hydroxide solution.
Pb2+(aq) + CrO42-(aq)
PbCrO4(s)
PbCrO4(s) + 4OH (aq)
Pb(OH)42-(aq) + CrO42-(aq)
 Adding dilute hydrochloric acid
Observation
A white precipitate
Pb2+(aq) + 2Cl- (aq)
PbCl2(S)
Lead(II) chloride is soluble on heating and recrystallizes on cooling.
 Adding dilute Sulphuric acid/ sodium sulphate solution.
Observation
A white precipitate
Pb2+(aq) + SO42-(aq)
PbSO4(s)
 With Al3+(aq) and Sn2+(aq) there is no observable change.
 Aluminum ions can be confirmed by;
 Adding litmus solution followed by excess ammonia solution.
Observation
A blue “lake” is formed.
 Adding a few drops of Alizarin solution followed by ammonia solution.
Observation
A pink precipitate in a colourless solution.
 Sn2+(aq) ions can be distinguished from Al3+ (aq) ions by using an oxidizing agent
such as;
 Acidified potassium manganate(VII) solution which will oxidize it to
tin(IV) ions and itself reduced to manganese(II) ions
Observation.
The purple solution turns colourless.
Sn2+(aq) + 2MnO-4(aq) + 16H+(aq)
5Sn4+(aq) + 2Mn2+(aq) + 8H2O(aq)
 Mercury(II) chloride solution.
Observation
A white precipitate is formed and it slowly turns grey.
Sn2+(aq) + 2HgCl2 (aq)
Sn4+(aq) + Hg2Cl2 (s) + 2Cl- (aq)
117
COLOURED CATIONS IN SOLUTION.
GREEN IONS IN SOLUTION.
These include; Fe2+(aq), Cr 3+(aq) , Ni2+(aq) and Cu2+(aq)
1. Cu2+(aq)
(a) With sodium hydroxide solution.
Observation
It forms a blue precipitate insoluble in excess.
Cu2+(aq) + 2OH-(aq)
Cu(OH)2(S)
(b) With ammonia solution.
Observation
A blue precipitate soluble in excess to form a deep blue solution.
Cu2+(aq) + 2OH-(aq)
Cu(OH)2(S)
Cu(OH)2(s) + 4NH3(aq)
Cu(NH3) 42+(aq) + 2OH-(aq)
NB:
 Copper(II) ions are confirmed by;
 Adding potassium iodide solution.
Observation
A white precipitate in a brown solution.
2Cu2+(aq) + 4 I-(aq)
Cu2I2(s) + I2(aq)
 Adding potassium hexacyanoferrate(II) solution.
Observation
A brown precipitate.
2Cu2+(aq) + [Fe(CN)6]4-(aq)
Cu2Fe(CN)6(S)
2+
2. Ni (aq)
(a) With sodium hydroxide solution
Observation
A green precipitate insoluble in excess.
Equation
Ni2+(aq) + 2OH-(aq)
Ni (OH)2(s)
(b) With ammonia solution
Observation
A green precipitate soluble in excess to form a blue solution.
Ni2+(aq) + 2OH-(aq)
Ni (OH)2(s)
Ni (OH)2 (s) + 6NH3(aq)
Ni(NH3)62+(aq) + 2OH-(aq)
NB:
 Ni2+(aq) ions are confirmed by
 Adding ammonia solution and then 2-3 drops of dimethylglyoxime.
Observation.
A red precipitate.
2+
3. Fe (aq)
(a) With sodium hydroxide solution
Observation
A green precipitate insoluble in excess that turns brown on standing.
Fe2+(aq) + 2OH-(aq)
Fe (OH)2 (S)
4Fe(OH)2(s) + O2(g)
2Fe2O3 . 2H2O(s)
or 4Fe(OH)2(s) + O2(g) + 2H2O(l)
4 Fe(OH)3(s)
118
(b) With ammonia solution
Observation
A green precipitate insoluble in excess that turns brown on standing.
Fe2+(aq) + 2OH-(aq)
Fe(OH)2 (S)
4Fe(OH)2(s) + O2(g)
2Fe2O3 . 2H2O(s)
or 4Fe(OH)2(s) + O2(g) + 2H2O(l)
4 Fe(OH)3(s)
NB:
 Fe2+ ions are confirmed by;
 Adding potassium hexacyanoferrate(III) solution
Observation
A dark- blue precipitate.
Fe2+(aq) + K+(aq) + Fe(CN)63- (aq)
KFe[Fe(CN)6](s)
3+
4. Cr (aq)
(a) With sodium hydroxide solution
Observation
A green precipitate soluble in excess to form a green solution.
Equations
Cr3+(aq) + 3OH-(aq)
Cr(OH)3 (s)
Cr(OH)3(s) + OH-(aq)
Cr(OH)4-(aq)
Tetrahydroxochromate(III) ions
(b) With dilute ammonia solution
Observation
A green precipitate insoluble in excess.
Cr3+(aq) + 3OH-(aq)
Cr(OH)3 (s)
NB:
 If concentrated ammonia solution is used, then a green precipitate soluble in
excess to form a pink solution is observed.
Cr3+(aq) + 3OH-(aq)
Cr(OH)3 (s)
Cr(OH)3(s) + 6NH3(aq)
Cr(NH3)63+(aq) + 3OH-(aq)
Hexaamminechromium(III) ions
 Chromium(III) ions can be confirmed by;
 Adding excess sodium hydroxide solution followed by hydrogen
peroxide solution and boiling the resultant mixture for a few
minutes.
Observation
A green precipitate soluble in excess to form a green solution which
turns yellow on addition of hydrogen peroxide solution.
̅ H(aq)
Cr3+(aq) + 3O
Cr(OH) 3 (s)
̅
Cr(OH)3 (s)+ OH(aq)
Cr(OH) (aq)
̅ H(aq)+3H202(aq)
2Cr(OH) (aq)+2O
2CrO (aq)+8H20(l)
 When lead(II) acetate solution is added to the resultant solution
(yellow solution), a yellow precipitate is formed.
Pb2+(aq) + 2Cr0 (aq)
PbCr04 (s)
119
PINK IONS IN SOLUTION
1.
Mn2+(aq)
These include Mn2+(aq) and Co2+(aq)
(a) With sodium hydroxide solution
Observation
A white precipitate insoluble in excess that turns brown on standing.
Mn2+(aq) + 2OH-aq)
Mn(OH)2(S)
Or 4Mn(OH)2(s) + O2(g)
2Mn2O3.2H2O(s)
(b) With ammonia solution
Observation
A white precipitate insoluble in excess that turns brown on standing.
Mn2+(aq) + 2OH-aq)
Mn(OH)2(S)
Or 4Mn(OH)2(s) + O2(g)
2Mn2O3.2H2O(s)
NB:
 Mn2+ ions can be confirmed by;
 Adding concentrated nitric acid followed by either solid lead(IV) oxide or
sodium bismuthate solution and the mixture boiled.
Observation
A purple solution is formed.
2Mn2+(aq) + 5BiO3- (aq) + 14H+(aq)
2MnO-4(aq) + 5Bi3+(aq) + 7H2O(l)
2Mn2+(aq) + 5 PbO2(s) + 4 H+(aq)
2MnO-4(aq) + 5 Pb2+(aq) + 2H2O(l)
2+
2. Co (aq)
(a) With sodium hydroxide solution
Observation
A blue precipitate insoluble in excess that turns brown on standing.
Co2+(aq) + 2OH-aq)
Co(OH)2(S)
4Co(OH)2(s) + O2(g)
2Co2O3 .2H2O(s)
(b) With ammonia solution
Observation
A blue precipitate soluble in excess to form a brown solution that turns reddish
brown on standing.
Co2+(aq) + 2OH-(aq)
Co(OH)2(S)
Co(OH)2(s) + 6NH3(aq)
Co(NH3)62+(aq) + 2OH-(aq)
4Co(NH3)62+(aq) + O2(g) + 2H2O(l)
4Co(NH3)63+(aq) + 4OH-(aq)
NB:
 Co2+ ions can be confirmed by;
 Adding potassium thiocyanate solution.
Observation
A blue solution is formed.
Co2+(aq) + 2SCN-(aq)
Co(SCN)2(aq)
 Adding concentrated hydrochloric acid
Observation
The pink solution turns blue.
120
CO(H2O)62+(aq) + 4 Cl-(aq)
CoCl42-(aq) + 6 H2O (l)
 However, when excess water is added to the resultant solution,
the blue solution turns pink
CoCl42-(aq) + 6 H2O (l)
CO(H2O)62+(aq) + 4 Cl-(aq)
YELLOW IONS IN SOLUTION.
This only includes Fe3+(aq)
(a) With sodium hydroxide solution
Observation
A brown precipitate insoluble in excess.
Fe3+(aq) + 3OH-(aq)
Fe(OH)3 (S)
(b) With ammonia solution
Observation
A brown precipitate insoluble in excess.
Fe3+(aq) + 3OH-(aq)
Fe(OH)3 (S)
NB:
 Fe3+ ions are confirmed by;
 Adding potassium hexacyanoferrate(II) solution
Observation
A dark- blue precipitate.
Fe3+(aq) + K+(aq) + Fe(CN)64- (aq)
KFe[Fe(CN)6](s)
 Adding potassium thiocyanate solution(KSCN),
Observation
A blood red colouration.
Fe3+(aq) + SCN-(aq)
[FeSCN]2+(aq)
ANIONS
The anions include; CO32-(aq), HCO3-(aq), C2O42-(aq), CH3 COO-, SO42-, SO32-, Cl-(aq), Br-(aq),
I-(aq), NO3-(aq), NO2-(aq) and PO43- ions
1. C2O42-(aq)
(a) Effect of heat on the solid substance.
All oxalates decompose to form oxides, carbon dioxide and carbon monoxide except
those of group(I) which form carbonates and carbon monoxide.
XC2O4(s)
XO(s) + CO2(g) + CO(g)
Observation.
A colourless gas that turns damp blue litmus paper red and lime water milky. (The
gas evolved is carbon dioxide hence C2O42-(aq) should be suspected.)
(b) Addition of lead(II) nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
Pb2+(aq) + C2O42-(aq)
PbC2O4(s)
121
(c) Addition silver nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid and
ammonia solution)
2Ag+(aq) + C2O42-(aq)
Ag2C2O4(s)
(d) Addition of barium nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
Ba2+(aq) + C2O42-(aq)
BaC2O4(s)
NB:
 C2O42-(aq) are confirmed by;
 Adding acidified potassium manganate(VII) solution and warm.
Observation.
The purple solution turns colourless.
2MnO4-(aq) + 16H+(aq) + 5C2O42-(aq)
2Mn2+(aq) + 8H2O(l) + 10CO2(g)
2. CO32-(aq) and HCO3-(aq)
(a) Effect of heat on the solid substance.
Observation.
A colourless gas that turns damp blue litmus paper red and lime water milky. (The
gas evolved is carbon dioxide hence CO32-(aq) and HCO3-(aq) should be suspected.)
XCO3(s)
XO(s) + CO2(g)
X(HCO3)2(s)
XCO3(s) + CO2(g) + H2O(l)
(b) Addition of dilute acid on the solid substance.
Observation.
Bubbles of a colourless gas that turns damp blue litmus paper red and lime water
milky. (The gas evolved is carbon dioxide hence CO32-(aq) and HCO3-(aq) should be
suspected.)
XCO3(s) + 2H+(aq)
X2+(aq) + CO2(g) + H2O(l)
X(HCO3)2(s) + 2H+(aq)
X2+(aq) + CO2(g) + 2H2O(l)
NB:
 CO32-(aq) and HCO3-(aq) can be distinguished from C2O42- by using a dilute acid,
both carbonates and hydrogen carbonates liberate carbon dioxide gas on
addition of dilute acid. Oxalates do not liberate carbon dioxide gas.
 C2O42-(aq) can be distinguished from CO32-(aq) using silver nitrate solution or
lead(II) nitrate solution or barium nitrate solution followed by dilute nitric acid.
 C2O42-(aq), a white precipitate soluble in the acid is observed.
 CO32-(aq), a white precipitate soluble in the acid with bubbles of a
colourless gas.
2 A CO3 (aq) is distinguished from HCO-3(aq) using magnesium sulphate
solution
 With a CO32-(aq) - A white precipitate is observed.
Mg2+(aq) + CO32-(aq)
MgCO3(s)
 With a HCO3 (aq) - No observable change.
But a white precipitate is observed on heating.
Mg(HCO3)2(aq)
MgCO3(s) + CO2(g) + H2O(l)
122
 When a dilute acid is added to a residue obtained after filtering and there is
bubbles of a colourless gas that turns damp blue litmus paper red and lime
water milky, a CO32-(aq) should be confirmed since all HCO3-(aq) are soluble.
3. SO32(a) Addition of lead(II) nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
Pb2+(aq) + SO32-(aq)
PbSO3(s)
(b) Addition of barium nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
Ba2+(aq) + SO32-(aq)
BaSO3(s)
NB:
 SO32-(aq) are confirmed by;
 Adding acidified potassium manganate(VII) solution.
Observation.
The purple solution turns colourless.
2MnO4-(aq) + 6H+(aq) + 5SO32-(aq)
2Mn2+(aq) + 3H2O(l) + 5SO42-(aq)
 Adding iodine solution or bromine solution.
Observation.
The brown solution turns colourless.
I2(aq) + H2O(l) + SO32-(aq)
2I-(aq) + 2H+(aq) + SO42-(aq)
2Br2(aq) + H2O(l) + SO3 (aq)
2Br-(aq) + 2H+(aq) + SO42-(aq)
4. SO42-(aq)
(a) Addition of lead(II) nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is insoluble in dilute nitric acid and
on heating)
Pb2+(aq) + SO42-(aq)
PbSO4(s)
NB:
 SO42-(aq) are confirmed by;
 Adding barium nitrate solution followed by dilute nitric acid.
Observation.
A white precipitate is formed insoluble in the acid.
Ba2+(aq) + SO42-(aq)
BaSO4(s)
35. PO4
(a) Addition of lead(II) nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
3Pb2+(aq) + 2PO43-(aq)
Pb3(PO4)2( (s)
(b) Addition of barium nitrate solution.
Observation.
A white precipitate is formed. (The precipitate is soluble in dilute nitric acid)
3Ba2+(aq) + 2PO43-(aq)
Ba3(PO4)2( (s)
123
(c) Addition silver nitrate solution.
Observation.
A yellow precipitate is formed. (The precipitate is soluble in dilute nitric acid and
ammonia solution)
3Ag+(aq) + PO43-(aq)
Ag3PO4( (s)
(d) Addition iron(II) chloride solution.
Observation.
A brown precipitate is formed.
3Fe2+(aq) + 2PO43-(aq)
Fe3(PO4)2(s)
NB:
 PO43-(aq) are confirmed by;
 Adding dilute nitric acid followed by ammonium molybdate solution.
Observation.
A yellow precipitate is formed. (The precipitate is soluble in ammonia
solution)
6. Cl-(aq), Br-(aq), I-(aq)
(a) Effect of heat on the solid sample.
(i) Cl-(aq)
Observation.
White fumes which turn damp blue litmus red and form dense white fumes
with concentrated ammonia solution or form a white precipitate with silver
nitrate solution. (HCl gas is evolved hence a Cl- should be highly suspected).
(ii) Br-(aq)
Observation.
Brown fumes. (Br2 gas evolved, hence Br- should be highly suspected)
(iii) I-(aq)
Observation.
Purple fumes. (I2 gas evolved, hence I- should be highly suspected)
(b) Addition of concentrated Sulphuric acid to a solid sample and heat.
(i) Cl-(aq)
Observation.
White fumes which form dense white fumes with concentrated ammonia
solution. (HCl gas is evolved hence a Cl- present).
(ii) Br-(aq)
Observation.
Brown fumes. (Br- oxidized to Br2)
(iii) I-(aq)
Observation.
Purple fumes which form a black sublimate. (I- oxidized to I2)
(c) Addition of lead(II) nitrate solution.
(i) Cl-(aq)
Observation.
A white precipitate. (The precipitate is insoluble in dilute nitric acid)
Pb2+(aq) + 2Cl-(aq)
PbCl2(s)
 The white precipitate is soluble on heating and recrystallizes on cooling.
124
(ii) Br-(aq)
Observation.
A white precipitate.
Pb2+(aq) + 2Br-(aq)
PbBr2(s)
(iii) I-(aq)
Observation.
A yellow precipitate.
Pb2+(aq) + 2I-(aq)
PbI2(s)
(d) Addition of silver nitrate solution [Confirmatory test]
(i) Cl-(aq)
Observation.
A white precipitate. (The precipitate is insoluble in dilute nitric acid)
Ag+(aq) + Cl-(aq)
AgCl(s)
 The white precipitate is soluble in ammonia solution.
AgCl(s) + 2NH3(aq)
Ag(NH3)2+(aq) + Cl-(aq)
(ii) Br (aq)
Observation.
A pale yellow precipitate. (The precipitate is insoluble in dilute nitric acid)
Ag+(aq) + Br-(aq)
AgBr(s)
(iii) I-(aq)
Observation.
A yellow precipitate. (The precipitate is insoluble in dilute nitric acid)
Ag+(aq) + I-(aq)
AgI(s)
 The yellow precipitate is insoluble in ammonia solution.
NB:
 Br-(aq) and I-(aq) can also be confirmed by adding bleaching powder/jik
followed by concentrated nitric acid followed by carbon tetrachloride.
Observation
 With Br-(aq), a reddish-brown colouration in the organic layer.
 With I-(aq), a purple colouration in the organic layer.
7. CH3COO-(aq)
(a) Effect of heat on the solid.
Metal oxalates decompose on heating to form acetone (propanone) which burns on
application of heat, carbon dioxide and a metal oxide.
O
(CH3COO)2X(s)
CH3CCH3(g) + CO2(g) + XO(s)
Observation
White fumes with a sweet odour that form a yellow precipitate with Brady’s
reagent.
NB:
 Carbon dioxide is also produced and should therefore be tested for. [ A
colourless gas that turns damp blue litmus paper red and lime water
milky]
125
(b) Addition of concentrated Sulphuric acid to a solid sample and heat.
Observation
White fumes with a vinegar smell (smell of acetic acid) which turns damp
blue litmus paper red.
CH3COO-(aq) + H+(aq)
CH3COOH(aq)
NB: Ethanoate ions are confirmed by;
 Adding an alcohol (ethanol) followed by 2-3 drops of concentrated
Sulphuric acid and the mixture warmed.
Observation
A sweet fruity smell. [An ester formed]
CH3COOH(aq) + CH3CH3OH(l)
CH3COOCH2CH3 (aq) + H2O(l)
 Adding iron(III) chloride solution and the mixture warmed.
Observation
A reddish-brown precipitate
3CH3COO-(aq) + Fe3+(aq)
(CH3COO)3Fe(s)
8. NO3 (aq)
(a) Effect of heat on the solid.
Most nitrates when heated produce nitrogen dioxide except those of sodium
and potassium.
Observation
Brown fumes. [Nitrogen dioxide evolved]
(b) Addition of concentrated Sulphuric acid to a solid sample and heat.
Observation
Brown fumes. [Nitrogen dioxide evolved]
NB: Nitrate ions can be confirmed by;
 Adding Devarda’s alloy (A mixture of Al, Cu and Zn) or Zinc or Aluminium
powder followed by sodium hydroxide solution and warm. [Ammonia
gas is evolved]
Observation
A colourless gas with a chocking smell that turns damp red litmus
paper blue and forms dense white fumes with concentrated
hydrochloric acid.
 Adding copper turnings followed by 2-3 drops of concentrated Sulphuric
acid and heat.
Observation
Brown fumes
 Adding freshly prepared iron(II) sulphate solution followed by
concentrated Sulphuric acid down the sides of the test tube carefully.
Observation
A brown ring is formed at the junction of the two layers.
9. NO2 (aq)
(a) Addition of concentrated/dilute Sulphuric acid to a solid sample.
Observation
Brown fumes. [Nitrogen dioxide evolved]
 Nitrate ions do not react with dilute Sulphuric acid.
126
NB: Nitrite ions can be confirmed by;
 Adding Devarda’s alloy (A mixture of Al, Cu and Zn) or Zinc or Aluminium
powder followed by sodium hydroxide solution and warm. [Ammonia
gas is evolved]
Observation
A colourless gas with a chocking smell that turns damp red litmus
paper blue and forms dense white fumes with concentrated
hydrochloric acid.
 Adding freshly prepared iron(II) sulphate solution followed by
concentrated Sulphuric acid down the sides of the test tube carefully.
Observation
A brown ring is formed at the junction of the two layers.
 Adding acidified potassium manganate(VII) solution.
Observation.
The purple solution turns colourless.
2MnO4-(aq) + 6H+(aq) + 5NO2-(aq)
2Mn2+(aq) + 3H2O(l) + 5NO3-(aq)
Experiment 48
You are provided with substance A, which contains two cations and two anions. You
are required to carry out the following tests on A and identify the cations and anions
in A. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of A strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of water,
add two spatula end-fuls of
A and shake well. Filter the
mixture and keep both the
filtrate and residue.
Divide the filtrate into six
portions.
127
Deductions
(i) To the first portion of
the filtrate, add dilute
sodium hydroxide dropwise until in excess.
(ii) To the second portion
of the filtrate, add dilute
ammonia solution dropwise until in excess.
(iii) To the third portion of
filtrate, add 3-4 drops of
dilute Sulphuric acid.
(iv) To the fourth portion
of the filtrate, add 3-4
drops of potassium
chromate(VI) solution
followed by excess sodium
hydroxide solution.
(v) To the fifth portion of
filtrate, add lead (II) nitrate
solution and heat.
(v) Use the sixth portion of
the filtrate to carry out a
test of your own to confirm
one of the first anion in A.
(c) Wash the residue and
then add dilute nitric acid
until no further change.
Divide the resultant
solution into four portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
128
(iii) To the third portion,
add 3-4 drops of dilute
Sulphuric acid.
(iv) Use the fourth portion
to carry out a test of your
own to confirm the second
cation in A.
(d) Identify the;
(i) Cations in A …………………………………… and ………………………………………………………….
(ii) Anions in A …………………………………… and ………………………………………………………….
Experiment 49
You are provided with substance B, which contains two cations and two anions. You
are required to carry out the following tests on B and identify the cations and anions
in B. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of B strongly in a dry
test tube until there is no
further change.
(b) To a spatula end-full of
B, add 3-4 drops of
concentrated Sulphuric
acid and heat.
(c) To about 6cm3 of water,
add two spatula end-fuls of
B and shake well. Filter the
129
Deductions
mixture and keep both the
filtrate and residue.
Divide the filtrate into five
portions.
(i) To the first portion of
the filtrate, add dilute
sodium hydroxide dropwise until in excess.
(ii) To the second portion
of the filtrate, add dilute
ammonia solution dropwise until in excess.
(iii) Use the third portion of
filtrate to carry out a test of
your own choice to confirm
the first cation in B.
(iv) To the fourth portion
of filtrate, add 3 drops of
neutral iron(III) chloride
solution and heat.
(v) To the fifth portion of
the filtrate, add 1cm3 of
ethanol followed 3 drops of
concentrated Sulphuric
acid and heat. Pour the hot
mixture in beaker of cold
water.
(d) Wash the residue and
then add dilute nitric acid
until no further change.
Divide the resultant
solution into four portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
130
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) To the third portion,
add 3-4 drops of dilute
Sulphuric acid.
(iv) Use the fourth portion
to carry out a test of your
own to confirm the second
cation in B.
(e) Identify the;
(i) Cations in B …………………………………… and ……………………………………….
(ii) Anions in B …………………………………… and ……………………………………….
Experiment 50
You are provided with substance C, which contains two cations and two anions. You
are required to carry out the following tests on C and identify the cations and anions
in C. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of C strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of water,
add two spatula end-fuls of
C and shake well. Filter the
131
Deductions
mixture and keep both the
filtrate and residue.
Divide the filtrate into six
portions.
(i) To the first portion of
the filtrate, add dilute
sodium hydroxide dropwise until in excess.
(ii) To the second portion
of the filtrate, add dilute
ammonia solution dropwise until in excess.
(iii) To the third portion of
filtrate, add 3-4 drops of
dilute Sulphuric acid.
(iv) To the fourth portion
of the filtrate, add litmus
solution followed by dilute
ammonia solution.
(v) To the fifth portion of
filtrate, add lead (II) nitrate
solution.
(v) Use the sixth portion of
the filtrate to carry out a
test of your own to confirm
one of the first anion in C.
(c) Wash the residue and
then add dilute nitric acid
until there is no further
change. Divide the
resultant solution into four
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
132
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own to confirm the second
cation in C.
(d) Identify the;
(i) Cations in C …………………………………… and ……………………………………….
(ii) Anions in C …………………………………… and ……………………………………….
Experiment 51
You are provided with substance D, which contains two cations and two anions. You
are required to carry out the following tests on D and identify the cations and anions
in D. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of D strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of water,
add two spatula end-fuls of
D and shake well. Filter the
mixture and keep both the
filtrate and residue.
Divide the filtrate into four
portions.
133
Deductions
(i) To the first portion of
the filtrate, add dilute
sodium hydroxide solution
and warm.
(ii) To the second portion
of the filtrate, add 3-4
drops of lead(II) nitrate
solution followed by dilute
nitric acid.
(iii) To the third portion of
filtrate, add 3-4 drops of
silver nitrate solution
followed by dilute nitric
acid.
(iv) To the fourth portion
of the filtrate, add acidified
potassium manganate(VII)
solution and heat.
(c) Wash the residue and
then add dilute
hydrochloric acid until
there is no further change.
Divide the resultant
solution into three
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own to confirm the second
cation in D.
134
(d) Identify the;
(i) Cations in D …………………………………… and ……………………………………….
(ii) Anions in D …………………………………… and ……………………………………….
Experiment 52
You are provided with substance E, which contains two cations and two anions. You
are required to carry out the following tests on E and identify the cations and anions
in E. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of E strongly in a dry
test tube until there is no
further change.
(b) To a spatula end-full of
E, add 3-4 drops of
concentrated Sulphuric
acid and heat.
(c) To about 6cm3 of water,
add two spatula end-fuls of
E and shake well. Filter the
mixture and keep both the
filtrate and residue.
Divide the filtrate into five
portions.
(i) To the first portion of
the filtrate, add dilute
sodium hydroxide dropwise until in excess.
(ii) To the second portion
of the filtrate, add dilute
135
Deductions
ammonia solution dropwise until in excess.
(iii) To the third portion of
the filtrate, add 3-4 drops
of sodium sulphate
solution.
(iv) Use the fourth portion
of filtrate to carry out a test
of your own choice to
confirm the first cation in
E.
(v) To the fifth portion of
filtrate, add a few copper
turnings followed by 3
drops of concentrated
Sulphuric acid and heat.
(d) Wash the residue and
then add dilute nitric acid
until no further change.
Divide the resultant
solution into four portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own to confirm the second
cation in E.
136
(e) Identify the;
(i) Cations in E …………………………………… and ……………………………………….
(ii) Anions in E …………………………………… and ……………………………………….
Experiment 53
You are provided with substance F, which contains two cations and two anions. You
are required to carry out the following tests on F and identify the cations and anions
in F. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of F strongly in a dry
test tube until there is no
further change.
(b) To a spatula end-full of
F, add 3-4 drops of
concentrated Sulphuric
acid and warm.
(b) To two spatula end-fuls
of F, add dilute nitric acid
drop-wise until there is no
further change.
(c) To the resultant
solution in (b), add dilute
sodium hydroxide solution
drop-wise until in excess
and then filter. Keep both
the filtrate and residue.
137
Deductions
(d) To the filtrate from (c),
add dilute nitric acid dropwise until the solution is
just acidic. Divide the acidic
solution into five portions.
(i) To the first portion of
the acidic solution, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic solution, add
dilute ammonia solution
drop-wise until in excess.
(iii) Use the third portion of
acidic solution to carry out
a test of your own to
confirm one of the cations
in F.
(iv) To the fourth portion
of acidic solution, add lead
(II) nitrate solution and
heat.
(v) Use the fifth portion of
the acidic solution to carry
out a test of your own
choice to confirm one of the
anions in F.
138
(e) Wash the residue and
then add dilute nitric acid
solution until there is no
further change. Divide the
resultant solution into four
parts.
(i) To the first part, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second part, add
aqueous ammonia dropwise until in excess.
(iii) To the third part, add
3-4 drops of dilute
Sulphuric acid.
(iv) Use the fourth part to
carry out a test of your own
to confirm the second
cation in F.
(f) Identify the;
(i) Cations in F …………………………………… and ……………………………………….
(ii) Anions in F …………………………………… and ……………………………………….
139
Experiment 54
You are provided with substance G, which contains three cations and one anion. You
are required to carry out the following tests on G and identify the cations and anion
in G. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of G strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of water,
add two spatula end-fuls of
G and shake well.
To the resultant solution,
add dilute sodium
hydroxide solution dropwise until there is no
further change, warm and
filter. Keep both the filtrate
and residue.
(c) To the filtrate in (b)
above, add dilute nitric acid
drop-wise until the
solution is just acidic.
Divide the acidic solution
into five portions.
(i) To the first portion of
the acidic solution, add
dilute sodium hydroxide
140
Deductions
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic solution, add
dilute ammonia solution
drop-wise until in excess.
(iii) Use the third portion of
acidic solution to carry out
a test of your own choice to
confirm the one of the
cations in G.
(iv) To the fourth portion,
add lead(II) nitrate
solution.
(v) Use the fifth portion to
carry out a test of your own
choice to confirm the anion
in G.
(d) Wash the residue and
then add dilute
hydrochloric acid until
there is no further change.
Divide the resultant
solution into three
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
141
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) To the third portion,
add little solid ammonium
chloride followed by
disodium
hydrogenphosphate
solution followed by excess
ammonia solution.
(d) Identify the;
(i) Cations in G ………………………, ……………………… and ……………………………………..
(ii) Anion in G ……………………………………
Experiment 55
You are provided with substance H, which contains two cations and two anions. You
are required to carry out the following tests on H and identify the cations and anions
in H. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of H strongly in a dry
test tube until there is no
further change.
142
Deductions
(b) To a spatula end-full of
H, add 3-4 drops of
concentrated Sulphuric
acid and heat.
(c) To about 6cm3 of
water, add two spatula
end-fuls of H and shake
vigorously.
(d) To the resultant
solution in (b) above, add
dilute sodium hydroxide
solution drop-wise until in
excess and filter. Keep
both the filtrate and
residue.
(e) To the filtrate, add
dilute nitric acid dropwise until the solution is
just acidic. Divide the
acidic filtrate into eight
portions.
(i) To the first portion of
the acidic filtrate, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic filtrate, add
dilute ammonia solution
drop-wise until in excess.
(iii) Use the third portion
of acidic filtrate to carry
out a test of your own
choice to confirm one of
the cations in H.
143
(iv) To the fourth portion
of the acidic filtrate, add 45 drops of silver nitrate
solution.
(v) To the fifth portion of
the acidic filtrate, add 3-4
drops of lead(II) nitrate
solution.
(vi) Use the sixth portion
of the acidic filtrate to
carry out a test of your
own choice to confirm one
of the anions in H.
(vii) To the seventh
portion of the acidic
filtrate, add 3 drops of
neutral iron(III) chloride
solution and heat.
(viii) To the eighth portion
of the acidic filtrate, add
1cm3 of ethanol followed 3
drops of concentrated
Sulphuric acid and heat.
Pour the hot mixture in
beaker of cold water.
(f) Wash the residue and
then add dilute
hydrochloric acid until no
further change. Divide the
resultant solution into
three portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
144
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own to confirm the second
cation in H.
(e) Identify the;
(i) Cations in H …………………………………… and ……………………………………….
(ii) Anions in H …………………………………… and ……………………………………….
Experiment 56
You are provided with substance I, which contains two cations and two anions. You
are required to carry out the following tests on I and identify the cations and anions
in I. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of I strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of
water, add two spatula
end-fuls of I, shake
vigorously and filter. Keep
both the filtrate and the
residue.
Divide the filtrate into
three portions.
145
Deductions
(i) To the first portion of
the filtrate, add lead(II)
nitrate solution.
(ii) To the second portion
of the filtrate, add barium
nitrate solution followed
by dilute nitric acid.
(iii) Use the third portion
of filtrate to carry out a
test of your own choice to
confirm one of the anions
in I.
(c) Dissolve the residue in
(b) in minimum dilute
nitric acid.
To the resultant solution,
add dilute sodium
hydroxide solution dropwise until in excess and
filter. Keep both the
filtrate and residue.
(d) To the filtrate in (c)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the acidic filtrate
into five portions
(i) To the first portion of
the acidic filtrate, add
dilute sodium hydroxide
146
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic filtrate, add 34 drops of dilute Sulphuric
acid.
(iii) To the third portion of
the acidic filtrate, add
dilute ammonia solution
drop-wise until in excess.
(iv) Use the fourth portion
of the acidic filtrate to
carry out a test of your
own choice to confirm one
of the cations in I.
(e) Wash the residue with
dilute sodium hydroxide
solution and dissolve it in
dilute nitric acid.
Divide the resultant
solution into three
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
147
(iii) To the third portion,
add 3-4 drops of dilute
Sulphuric acid.
(iv) Use the fourth portion
to carry out a test of your
own to confirm the second
cation in I.
(f) Identify the;
(i) Cations in I …………………………………… and ……………………………………….
(ii) Anions in I …………………………………… and ……………………………………….
Experiment 57
You are provided with substance J, which contains two cations and two anions. You
are required to carry out the following tests on J and identify the cations and anions
in J. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of J strongly in a dry
test tube until there is no
further change.
148
Deductions
(b) To a spatula end-full of
J, add half a spatula endfull of manganese(IV)
oxide followed by 4-5
drops of concentrated
Sulphuric acid and heat
gently.
(c) To three spatula endfuls of J, add dilute nitric
acid drop-wise until there
is no further change.
Warm gently.
(d) To the resultant
solution in (c) above, add
dilute ammonia solution
drop-wise until in excess
and filter. Keep both the
filtrate and the residue.
(e) To the filtrate in (d)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the acidic filtrate
into five portions.
(i) To the first portion of
the acidic filtrate, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic filtrate, add
dilute ammonia solution
drop-wise until in excess.
149
(iii) Use the third portion
of acidic filtrate to carry
out a test of your own
choice to confirm one of
the cations in J.
(f) Wash the residue and
dissolve it in dilute
Sulphuric acid. Divide the
resultant solution into
three portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own choice to confirm the
second cation in J.
(g) To two spatula endfuls of J, add 4cm3 of
water, shake vigorously
and filter. Divide the
filtrate into three portions.
(i) To the first portion, add
3-4 drops of silver nitrate
solution followed by
ammonia solution dropwise until in excess.
150
(ii) To the second portion,
add 1-2 drops of
concentrated nitric acid
followed by sodium
thiosulphate solution
drop-wise until in excess.
(iii) To the third portion,
add 3-4 drops of lead(II)
nitrate solution.
(f) Identify the;
(i) Cations in J …………………………………… and ……………………………………….
(ii) Anions in J …………………………………… and ……………………………………….
Experiment 58
You are provided with substance K, which contains three cations and one anion. You
are required to carry out the following tests on K and identify the cations and anion
in K. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of K strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of
water, add two spatula
end-fuls of K and shake
thoroughly.
To the resultant solution,
add dilute sodium
hydroxide solution dropwise until there is no
further change, warm and
filter. Keep both the
filtrate and residue.
151
Deductions
(c) To the filtrate in (b)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the acidic solution
into six portions.
(i) To the first portion of
the acidic solution, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion
of the acidic solution, add
dilute ammonia solution
drop-wise until in excess.
(iii) To the third portion of
the acidic solution, add 2-3
drops of potassium iodide
solution.
(iv) Use the fourth portion
of acidic solution, add 2-3
drops of litmus solution
followed by dilute
ammonia solution dropwise until in excess.
(v) To the fifth portion of
the acidic solution, add 2-3
drops of lead(II) nitrate
solution and heat.
(vi) Use the sixth portion
of the acidic solution to
carry out a test of your
own choice to confirm the
anion in K.
(d) Wash the residue and
dissolve it in dilute
hydrochloric acid. Divide
the resultant solution into
three portions.
152
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) To the third portion,
add 3-4 drops of
potassium thiocyanate
solution.
(e) Identify the;
(i) Cations in K ………………………, ……………………… and ……………………………………..
(ii) Anion in K ……………………………………
Experiment 59
You are provided with substance L, which contains two cations and two anions. You
are required to carry out the following tests on L and identify the cations and anions
in L. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
Tests
Observations
(a) Heat two spatula endfuls of L strongly in a dry
test tube until there is no
further change.
(b) To about 6cm3 of
water, add two spatula
end-fuls of L and shake
vigorously. Filter and keep
both the filtrate and
residue.
Divide the filtrate into five
portions.
153
Deductions
(i) To the first portion of
the filtrate, add 2-3 drops
of lead(II) nitrate solution.
(ii) To the second portion
of the filtrate, add silver
nitrate solution followed
by dilute nitric acid.
(iii) To the third portion of
filtrate, add 2-3 drops of
iron(II) chloride solution.
(iv) To the fourth portion
of the filtrate, add
concentrated nitric acid
followed by ammonium
molybdate solution
followed by excess dilute
ammonia solution.
(c) To the residue from (b)
above, add dilute nitric
acid until there is no
further change.
To the resultant solution,
add dilute sodium
hydroxide solution dropwise until in excess and
filter. Keep both the
filtrate and residue.
(d) To the filtrate in (c)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the acidic filtrate
into three portions.
(i) To the first portion of
the acidic filtrate, add
dilute sodium hydroxide
solution drop-wise until in
excess.
154
(ii) To the second portion
of the acidic filtrate, add
dilute ammonia solution
drop-wise until in excess.
(iii) Use the third portion
of the acidic filtrate to
carry out a test of your
own choice to confirm one
of the cations in L.
(e) Wash the residue from
(c) above and dissolve it in
minimum amount of dilute
nitric acid.
Divide the resultant
solution into three
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own to confirm the second
cation in L.
(f) Identify the;
(i) Cations in L …………………………………… and ……………………………………….
(ii) Anions in L …………………………………… and ……………………………………….
155
Experiment 60
You are provided with substance M, which contains two cations and two anions. You
are required to carry out the following tests on M and identify the cations and
anions in M. Identify any gas(es) evolved. Record your observations and deductions
in the table below.
Tests
Observations
(a) Heat two spatula endfuls of M strongly in a dry
test tube until there is no
further change.
(b) To a spatula end-full of
M, add 4-5 drops of
concentrated Sulphuric
acid and heat.
(c) To two spatula end-fuls
of M in a test tube, add
5cm3 of water, shake well
and filter.
Keep both the filtrate and
residue.
Divide the filtrate into
three portions.
(i) To the first portion, add
3-4 drops lead(II) nitrate
solution.
(ii) To the second portion,
add silver nitrate solution
followed by excess
ammonia solution.
156
Deductions
(iii) To the third portion,
add bleaching powder
followed by concentrated
nitric acid followed by
3cm3 of carbon
tetrachloride.
(d) Dissolve the residue in
(c) above in minimum
amount of dilute nitric
acid.
To the resultant solution,
add dilute sodium
hydroxide solution dropwise until in excess and
filter. Keep both the
filtrate and the residue.
(e) To the filtrate in (d)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the acidic filtrate
into five portions.
(i) To the first portion of
the acidic filtrate, add
dilute ammonia solution
drop-wise until in excess.
(ii) To the second portion
of the acidic filtrate, add 34 drops of dilute
hydrochloric acid.
(iii) To the third portion of
the acidic filtrate, add
dilute sodium hydroxide
solution drop-wise until in
excess.
157
(iv) Use the fourth portion
of acidic filtrate to carry
out a test of your own
choice to confirm one of
the cations in M.
(f) Wash the residue and
dissolve it in dilute nitric
acid. Divide the resultant
solution into three
portions.
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own choice to confirm the
second cation in M.
(g) Identify the;
(i) Cations in M …………………………………… and ……………………………………….
(ii) Anions in M …………………………………… and ……………………………………….
Experiment 61
You are provided with substance N, which contains two cations and two anions. You
are required to carry out the following tests on N and identify the cations and anions
in N. Identify any gas(es) evolved. Record your observations and deductions in the
table below.
158
Tests
Observations
(a) Heat two spatula endfuls of N strongly in a dry
test tube until there is no
further change.
(b) To a spatula end-full of
N, add 4-5 drops of
concentrated Sulphuric
acid and warm.
(c) To two spatula end-fuls
of N in a test tube, add
5cm3 of water, shake well
and filter.
Keep both the filtrate and
residue.
Divide the filtrate into four
portions.
(i) To the first portion, add
3-4 drops concentrated
nitric acid followed by
sodium thiosulphate
solution.
(ii) To the second portion,
add silver nitrate solution
followed by excess
ammonia solution.
(iii) To the third portion,
add copper(II) sulphate
solution followed by 4
drops of starch indicator.
159
Deductions
(iv) To the third portion,
add 2cm3 of jik followed
by concentrated nitric acid
followed by 3cm3 of
carbon tetrachloride.
(d) Wash the residue in (c)
and dissolve it in
minimum amount of dilute
nitric acid.
(Do not divide the
resultant solution)
(i) To 1cm3 of the
resultant solution, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To 1cm3 of the
resultant solution, add 3-4
drops of potassium iodide
solution.
(iii) To 1cm3 of the
resultant solution, add
dilute ammonia solution
drop-wise until in excess.
(e) To the remaining
solution in (d) above, add
dilute ammonia solution
drop-wise until in excess.
Filter and keep the filtrate.
(f) To the filtrate in (e)
above, add dilute nitric
acid drop-wise until the
solution is just acidic.
Divide the resultant
solution into three
portions.
160
(i) To the first portion, add
dilute sodium hydroxide
solution drop-wise until in
excess.
(ii) To the second portion,
add dilute ammonia
solution drop-wise until in
excess.
(iii) Use the third portion
to carry out a test of your
own choice to confirm the
second cation in N.
(g) Identify the;
(i) Cations in N …………………………………… and ……………………………………….
(ii) Anions in N …………………………………… and ……………………………………….
161
ORGANIC QUALITATIVE ANALYSIS
Organic qualitative analysis involves identifying the nature of the functional
group of a given organic compound.
The nature of the organic compound can be obtained when the organic
compound is burnt i.e when an organic compound burns with a non – sooty
flame, the organic compound is aliphatic and when it burns with a sooty flame,
the organic compound is aromatic.
Functional groups are identified using different reagents after dissolving the
organic compound in water and testing the resultant solution with a litmus paper.
 When the resultant solution is acidic (turns blue litmus paper red), then a
carboxylic acid, an acid chloride and acid anhydride, should be suspected.
 When the resultant solution is neutral (has no effect on litmus paper), then an
alcohol, carbonyl compound and ester should be suspected.
 When the resultant solution is alkaline (turns red litmus paper blue), then an
amine should be suspected.
The following are tests for some common functional groups.
1. Test for carboxyl group (-COOH)
To the suspected organic compound,
 Add sodium carbonate solution or sodium hydrogencarbonate solution
or magnesium metal.
When effervescence of a colourless gas occurs, then a carboxylic acid is
confirmed.
2CH3COOH(l) + Na2CO3(aq)
2CH3COONa(aq) + CO2(g) + H2O(l)
CH3COOH(l) + NaHCO3(aq)
CH3COONa(aq) + CO2(g) + H2O(l)
2CH3COOH(l) + Mg(s)
(CH3COO)2Mg(aq) + H2(g)
 Add an alcohol (ethanol) followed by a few drops of concentrated
Sulphuric acid and heat the mixture.
When a sweet fruity smell is observed (an ester) then a carboxylic acid is
confirmed.
CH3COOH(aq) + CH3CH3OH(l)
CH3COOCH2CH3 (aq) + H2O(l)
 Add phosphorous(V) chloride .
When white fumes of the Hydrogen chloride gas is produced. Then a
carboxylic acid is highly suspected.
CH3COOH(l) + PCl5(g)
CH3 COCl(l) + POCl3(s) + HCl(g)
 Using sulphur dichloride oxide (thionyl chloride)
When white fumes of the Hydrogen chloride gas is produced. Then a
carboxylic acid is highly suspected.
CH3COOH + SOCl2 (s)
CH3COCl + SO2 + HCl
162
2. Test for carbonyl group.
To the suspected organic compound, add Brady’s reagent (2,4dinitrophenylhydrazine solution)
When a yellow precipitate is formed, then an aldehyde or ketone should be
suspected.
Differentiating between aldehydes and ketones
Aldehydes are easily oxidized to corresponding carboxylic acids while ketones are
resistant to oxidation. (cannot be oxidized)
Therefore aldehydes are differentiated from ketones by using strong and mild
oxidizing agents such as
(a) Acidified potassium dichromate(VI) solution and the mixture warmed.
With an aldehyde, the orange solution turns green.
With a ketone, there is no observable change.
(b) Acidified potassium manganate(VII) solution and the mixture warmed.
With an aldehyde, the purple solution turns colourless.
With a ketone, there is no observable change.
(c) Tollen’s reagent (Ammoniacal silver nitrate solution)
With aldehydes, a silver mirror is observed on heating.
With ketones, there is no observable change.
CH3CHO(l) Ag(NH3)2+(aq)
Ag(s) + CH3COO-(aq)
(d) Fehling’s solution (Ammoniacal copper(II) sulphate solution)
With aldehydes, a reddish-brown precipitate is observed on heating.
With ketones, there is no observable change.
CH3CHO(l) Fehling’s solution Cu2O(s) + CH3COO-(aq)
Note
Carbonyl compounds with a methyl group attached to carbonyl carbon form a
yellow precipitate with Iodine solution in presence of sodium hydroxide
solution.
CH3CHO(l) I2(aq)/NaOH(aq) CHI3(s) + HCOONa(aq)
3. Test for hydroxyl (-OH) group.
The following reagents can be used to confirm whether the organic compound
is an alcohol.
(a) Using sodium metal
Effervescence of a colourless gas is observed.
CH3CH2OH(l) + Na(s)
CH3CH2ONa(aq) + ½ H2 (g)
(b) Using phosphorous(V) chloride .
White fumes of the hydrogen chloride gas is produced.
CH3CH2OH(l) + PCl5
CH 3CH2Cl(l) + POCl3(s) + HCl(g)
(c) Using sulphur dichloride oxide (thionyl chloride)
White fumes of the hydrogen chloride gas is produced.
CH3CH2OH + SOCl2
CH3CH2Cl + SO2 + HCl
(d) Using a carboxylic acid (ethanoic acid) followed by a few drops of concentrated
Sulphuric acid and the mixture heated.
A sweet fruity smell is observed (an ester is formed).
CH3COOH(l) + CH3CH3OH(l)
CH3COOCH2CH3 (l) + H2O(l)
163
Distinguishing between classes of alcohols.
Primary, secondary and tertiary alcohols can be distinguished using;
1. Anhydrous zinc chloride and concentrated hydrochloric acid. (Lucas reagent)
With a primary alcohol, there is no observable change at room temperature.
With a secondary alcohol, a cloudy solution forms between 5-10 minutes.
With a tertiary alcohol, a cloudy solution forms immediately.
2. Using an oxidizing agent such as acidified potassium dichromate(VI)
solution.
 With tertiary alcohols there is no observable change. [They are
resistant to oxidation]
 With primary and secondary alcohols, the orange solution turns
green.
However, on addition of Brady’s reagent to the resultant solution, a
yellow precipitate is formed for the case of secondary alcohols
and no observable change for primary alcohols.
Note:
Secondary and primary alcohols with a methyl group attached to the
carbon atom carrying the hydroxyl group form a yellow precipitate with
iodine solution in sodium hydroxide solution.
CH3CH2OH(l) I2(aq)/NaOH(aq) CHI3(s) + HCOONa(aq)
4. Test for an amine group.
Amines are identified using a mixture of sodium nitrite and concentrated
hydrochloric acid (Nitrous acid) below 50C.
 With a primary aromatic amine, there is no observable change, but
effervescence of a colourless gas occurs on warming.
 With a primary aliphatic amine, effervescence of a colourless gas occurs
below 5°C.
 With a secondary amine, a yellow oily liquid is observed.
 With a tertiary amine, there is no observable change.
SUMMARY
TESTS ON ORGANIC SAMPLES.
1. Burning /combustion.
Observation
Deduction.
Burns with yellow non-sooty
Saturated aliphatic compound of low C:H ratio
flame
Burns with yellow sooty flame Aromatic compound
2. Solubility in water
Observation
Miscible/soluble in water.
Deduction
Polar compound of low molecular mass.
164
Sparingly miscible/soluble in
water.
Polar compound of high molecular mass.
3. Solubility in NaOH solution. (If compound was immiscible with water)
Observation
Deduction
Soluble
Acidic compound, phenol or carboxylic acid present.
Insoluble
Basic compound, amine present.
4. Testing the solution with litmus
Observation
Solution has no effect on litmus
paper.
Solution turns blue litmus paper
red.
Solution turns red litmus paper to
blue.
Deduction
Neutral compound, alcohol or carbonyl
compound present
Acidic compound, carboxylic acid or phenol
present
Basic compound, amine or salt of
carboxylic acid.
5. Action of sodium hydrogencarbonate or sodium carbonate (solid or solution) or
magnesium ribbon.
Observation
Deduction
Bubbles of a colourless gas.
Carboxylic acid present
No observable change.
Carboxylic acid absent
6. Alcohol with conc. H2SO4 (esterification)
Observation
Deduction
Sweet fruity smell
Ester formed. Therefore, carboxylic acid present.
No observable change
Carboxylic acid absent.
7. Action of acidified potassium dichromate(VI) solution
Observation
Deduction
Orange solution turns
Reducing agent present, primary alcohol, secondary
green.
alcohol, aldehyde or methanoic acid present.
No observable change
Reducing agent absent, primary alcohol, secondary
alcohol, aldehyde or methanoic acid absent.
8. Action of sodium metal.
Observation
Deduction
Bubbles of a colourless
Alcohol present.
gas.
9. Brady’s reagent. (2,4–dinitrophenylhydrazine solution)
Observation
Deduction
Yellow precipitate
Carbonyl compound present
No observable change
Carbonyl compound absent
165
10. Sodium hydrogensulphite solution.
Observation
White precipitate
No observable change
Deduction
Carbonyl compound present
Carbonyl compound absent
11. Lucas reagent (Anhydrous zinc chloride and concentrated hydrochloric acid)
Observation
Deduction
Cloudy solution formed immediately.
Tertiary alcohol
Cloudy solution formed between 5-10
Secondary alcohol
minutes
No observable change
Primary alcohol
12. Tollen’s Reagent (Ammoniacal silver nitrate solution)
Observation
Deduction
Silver mirror
Aldehyde present (if carboxylic acid, then methanoic acid)
No observable
Aldehyde absent.
change
13. Fehling’s solution
Observation
Deduction
A reddish-brown precipitate. Aldehyde present. (If carboxylic acid, then methanoic
acid)
No observable change.
Aldehyde absent.
14. Iodine solution and sodium hydroxide solution. (Iodoform)
Observation
Deduction
Yellow
CH3CH2OH, CH3CH(OH)R, CH3CHO or CH3COR
precipitate
Note: Ethanol is the only primary alcohol and ethanal is the only aldehyde that gives
positive iodoform test.
15. Neutral iron(III) chloride solution.
Observation
Violet/purple colouration.
No observable change.
Brown precipitate.
16. Bromine water
Observation
White precipitate.
Deduction
Phenol present
Phenol absent
Benzoic acid present
Deduction
Phenol (also aromatic amine, in case solution was alkaline)
166
Commenting
 Aliphatic (or aromatic)
 Low C:H ratio ie saturated (or high C:H ratio i.e unsaturated)
 Functional group alcohol, aldehyde, ketone, carboxylic acid or phenol
 Primary or secondary or tertiary (where applicable)
 CH3CH2OH, CH3CH(OH)R, CH3CHO or CH3COR (where applicable)
Experiment 62
You are provided with an organic compound A.
You are required to carry out tests below on A and describe the nature of A.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of A on a spatula end.
(b) To about 0.5cm3 of A,
add 1 cm3 water, shake
and test the mixture with
litmus paper.
(c) To about 1cm3 of A,
add 2-3 drops of neutral
iron(III) chloride
solution.
(d) To about 1cm3 of A,
add 3- 4 drops of Brady’s
reagent.
(e) To about 1cm3 of A,
add 4-5 drops of acidified
potassium
dichromate(VI) solution
and heat.
(f) To about 1cm3 of A,
add 4 cm3 of iodine
solution followed by
sodium hydroxide
solution drop-wise until
the solution is paleyellow. Warm and allow
to cool.
(g) Describe the nature of A.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
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Experiment 63
You are provided with an organic compound B.
You are required to carry out tests below on B and describe the nature of B.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
B on a spatula end
(b) To about 0.5cm3 of B,
add 1 cm3 water, shake and
test the mixture with
litmus paper.
(c) To about 1cm3 of B, add
2-3 drops of neutral
iron(III) chloride solution.
(d) To about 1cm3 of B, add
3- 4 drops of Brady’s
reagent.
(e) To about 1cm3 of B, add
4-5 drops of acidified
potassium dichromate(VI)
solution and heat. Allow to
cool and use it in part (f).
(f) To the mixture from (e),
add 3-4 drops of Brady’s
reagent.
(g) To about 1cm3 of B, add
2cm3 of Lucas reagent.
(h) Describe the nature of B.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
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Experiment 64
You are provided with an organic compound C.
You are required to carry out tests below on C and describe the nature of C.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
C on a spatula end
(b) To about one spatula
end-full of C, add 1cm3
water, shake and test the
mixture with litmus paper.
Divide the resultant
solution into five portions.
(i) To the first portion, add
2-3 drops of neutral
iron(III) chloride solution.
(ii) To the second portion,
add half a spatula end-full
of sodium carbonate.
(iii) To the third portion,
add 3- 4 drops of 2,4dinitrophenylhydrazine
solution.
(iv) To the fourth portion,
add 4-5 drops of acidified
potassium dichromate(VI)
solution and heat.
(v) To the fifth portion, add
an equal volume of ethanol
followed by 2-3 drops of
concentrated Sulphuric
acid and heat the mixture.
(c) Describe the nature of C.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
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Experiment 65
You are provided with an organic compound D.
You are required to carry out tests below on D and describe the nature of D.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
D on a spatula end.
(b) To about 1cm3 of D, add
2 cm3 water, shake and test
the mixture with litmus
paper.
Divide the resultant
solution into four portions.
(i) To first portion, add 3- 4
drops of 2,4dinitrophenylhydrazine
solution.
(ii) To the second portion,
add half a spatula end-full
of sodium
hydrogencarbonate.
(iii) To the third portion,
add 2-3 drops of neutral
iron(III) chloride solution.
(iv) To the fourth portion,
add 4-5 drops of acidified
potassium dichromate(VI)
solution and heat.
(c) To 1cm3 of D, add an
equal volume of ethanoic
acid followed by 2-3 drops
of concentrated Sulphuric
acid and heat the mixture.
Pour it in a beaker of cold
water and allow to stand.
(d) To 1cm3 of D, add 2-3
drops of concentrated
Sulphuric acid and heat.
Pass the vapour evolving
into a test tube containing
acidified potassium
manganate(VII) solution.
170
(e) To 1cm3 of D, add 2cm3
of Luca’s reagent.
(f) Describe the nature of D.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
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Experiment 66
You are provided with an organic compound E.
You are required to carry out tests below on E and describe the nature of E.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
E on a spatula end
(b) To half a spatula endfull of E in a test tube, add
sodium hydroxide solution
and shake vigorously.
(c) To about one spatula
end-full of E in a boiling
tube, add about 5cm3 of
water, shake. Heat and test
the resultant solution with
litmus paper.
Divide the resultant
solution into three
portions.
(i) To the first portion of
the hot solution, add 2-3
drops of neutral iron(III)
chloride solution.
(ii) To the second portion
of the hot solution, add half
a spatula end-full of
sodium carbonate.
(iii) To the third portion of
the hot solution, add 3- 4
drops of Brady’s reagent.
171
(d) To half a spatula endfull of E in a test tube, add
about 2cm3 of methanol.
Shake to dissolve and then
add 2-3 drops of
concentrated Sulphuric
acid and heat the mixture.
Pour the contents into a
beaker of cold water.
(e) Describe the nature of E.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 67
You are provided with an organic compound F.
You are required to carry out tests below on F and describe the nature of F.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of F on a spatula end
(b) To about 1cm3 of F,
add 1 cm3 water, shake
and test the mixture with
litmus.
(c) To about 1 cm3 of F,
add 2-3 drops of neutral
iron(III) chloride
solution.
(d) To about 1 cm3 of F,
add 3- 4 drops of Brady’s
reagent.
(e) To about 1cm3 of F,
add 2-3 drops of sodium
carbonate solution.
172
(f) To about 1 cm3 of F,
add 4-5 drops of acidified
potassium
dichromate(VI) solution
and heat. Then add 1cm3
of ethanol followed by 23 drops of concentrated
Sulphuric acid and heat.
Pour the mixture into a
small beaker of cold
water.
(g) To about 1cm3 of F,
add 4 cm3 of iodine
solution followed by
sodium hydroxide
solution drop-wise until
the solution is paleyellow. Warm and allow
to cool.
(h) To about 1cm3 of F,
add about 1cm3 of
Tollen’s reagent and heat
gently.
(i) Describe the nature of F.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 68
You are provided with an organic compound G.
You are required to carry out tests below on G and describe the nature of G.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of G on a spatula end
(b) To about 1cm3 of G,
add 1 cm3 water, shake
and test the mixture with
litmus paper.
173
(c) To about 1 cm3 of G,
add 2-3 drops of neutral
iron(III) chloride
solution.
(d) To about 1 cm3 of G,
add 3- 4 drops of Brady’s
reagent.
(e) To about 1cm3 of G,
add half a spatula endfull of magnesium
powder.
(f) To about 2cm3 of G,
add 2cm3 of sodium
hydroxide solution,
shake well and boil the
mixture. Cool for 1
minute and add 2cm3 of
dilute nitric acid followed
by 3 drops of silver
nitrate solution and filter.
(g) To the filtrate in (f),
add an equal volume of
ethanol followed by 2-3
drops of concentrated
Sulphuric acid and heat.
Pour the mixture into a
small beaker of cold
water.
(h) Describe the nature of G.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 69
You are provided with an organic compound H.
You are required to carry out tests below on H and describe the nature of H.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
H on a spatula end
174
(b) To about 2cm3 of H, add
4 cm3 water, shake and test
the resultant solution with
litmus paper.
Divide the resultant
solution into five portions.
(i) The first portion, add 34 drops of sodium
carbonate solution.
(ii) To the second portion,
add 3-4 drops of acidified
potassium dichromate(VI)
solution and heat.
(iii) To the third portion,
add 2-3 drops of 2,4dinitrophenylhydrazine
solution.
(iv) To fourth portion, add
ammoniacal silver nitrate
solution and heat.
(v) To the fifth portion, add
an equal volume of ethanol
followed by 2-3 drops of
concentrated Sulphuric
acid and heat the mixture.
Pour it in a beaker of cold
water.
(f) Describe the nature of H.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
175
Experiment 70
You are provided with an organic compound I.
You are required to carry out tests below on I and describe the nature of I.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of I on a spatula end
(b) To about 1cm3 of I,
add 1 cm3 water, shake
and test the mixture with
litmus paper.
Divide the mixture into
two portions.
(i) To the first portion,
add 2-3 drops of neutral
iron(III) chloride
solution.
(ii) To second portion,
add 3- 4 drops of Brady’s
reagent.
(c) To about 1cm3 of I,
add 2-3 drops of
concentrated Sulphuric
acid. Heat the mixture
and pass the gas through
acidified potassium
manganate(VII) solution.
(d) To about 1 cm3 of I,
add 4-5 drops of acidified
potassium
dichromate(VI) solution
and heat. Divide the
resultant solution into
two portions.
(i) To the first portion,
add 3- 4 drops of Brady’s
reagent.
(ii) To the second
portion, add about 1cm3
of Fehling’s solution and
heat.
176
(e) To about 1cm3 of I,
add 4 cm3 of iodine
solution followed by
sodium hydroxide
solution drop-wise until
the solution is paleyellow. Warm and allow
to cool.
(f) To about 1cm3 of I,
add 1cm3 of Lucas
reagent.
(i) Describe the nature of I.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 71
You are provided with an organic compound J.
You are required to carry out tests below on J and describe the nature of J.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of J on a spatula end
(b) To about 2cm3 of J,
add 2 cm3 water, shake
and allow to stand. Test
the mixture with litmus
paper.
Divide the mixture into
two portions.
(i) To the first portion,
add 2-3 drops of neutral
iron(III) chloride
solution.
(ii) To second portion,
add a spatula end-full of
sodium carbonate.
177
(c) To about 1 cm3 of J,
add 4-5 drops of acidified
potassium
dichromate(VI) solution
and heat
(d) To about 1 cm3 of J,
add 3- 4 drops of Brady’s
reagent.
(e) To about 1cm3 of J,
add Tollen’s reagent and
heat.
(f) To about 1cm3 of J,
add 4 cm3 of iodine
solution followed by
sodium hydroxide
solution drop-wise until
the solution is paleyellow. Warm and allow
to cool.
(g) Describe the nature of J.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 72
You are provided with an organic compound K.
You are required to carry out tests below on K and describe the nature of K.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of K on a spatula end
(b) To about 2cm3 of K,
add 2 cm3 water, shake
and test the mixture with
litmus paper.
Divide the mixture into
three portions.
(i) To the first portion,
add 2-3 drops of neutral
iron(III) chloride
solution.
178
(ii) To second portion,
add 2-3 drops of Brady’s
reagent.
(c) To the third portion of
the mixture from (b), add
4-5 drops of acidified
potassium
manganate(VII) solution
and heat gently. Allow to
cool and divide the
resultant solution into
three portions.
(i) To the first portion,
add 3- 4 drops of Brady’s
reagent.
(ii) To the second
portion, add 2-3 drops of
acidified potassium
dichromate(VI) solution
and heat.
(iii) To the third portion,
add 4-5 drops of Fehling’s
solution and heat.
(d) To about 1cm3 of K,
add 1cm3 of Lucas
reagent.
(e) To about 1cm3 of K,
add 4 cm3 of iodine
solution followed by
sodium hydroxide
solution drop-wise until
the solution is paleyellow. Warm and allow
to cool.
(f) Describe the nature of K.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
179
Experiment 73
You are provided with an organic compound L.
You are required to carry out tests below on L and describe the nature of L.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of L on a spatula end
(b) To one spatula endfull of Z in a test tube, add
about 2cm3 of sodium
hydroxide solution and
shake.
(c) Shake one spatula
end-full of L with about
4cm3 of water and test
the mixture with litmus
paper.
Divide the mixture into
five portions.
(i) To the first portion,
add 2-3 drops of neutral
iron(III) chloride
solution.
(ii) To second portion,
add 2-3 drops of Brady’s
reagent.
(iii) To the second
portion, add 1cm3 of
sodium carbonate
solution.
(iv) To the fourth portion,
add soda lime and heat.
(d) To the fifth portion of
the solution from (c), add
4-5 drops of acidified
potassium
manganate(VII) solution
and heat. Divide the
resultant solution into
two portions.
180
(i) To the first portion,
1cm3 of Brady’s reagent.
Shake and allow to stand
for about 2 minutes.
(ii) To the second
portion, add Tollen’s
reagent and heat.
(e) Dissolve a spatula
end-full of L in water and
add Lucas reagent.
(f) Describe the nature of L.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 74
You are provided with an organic compound M.
You are required to carry out tests below on M and describe the nature of M.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount of
M on a spatula end
(b) To about 0.5 cm3 of M,
add 1 cm3 water, shake and
test the mixture with
litmus paper.
(c) To about 1 cm3 of M,
add 2-3 drops of neutral
iron(III) chloride solution.
(d) To about 1cm3 of M,
add sodium carbonate
solution.
(e) To about 1 cm3 of M,
add 4-5 drops of acidified
potassium dichromate(VI)
solution and heat.
181
(f) To about 1 cm3 of M,
add ammoniacal silver
nitrate solution and heat.
(h) Describe the nature of M.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
Experiment 75
You are provided with an organic compound N.
You are required to carry out tests below on N and describe the nature of N.
Record your observations and deductions in the table below.
TESTS
OBSERVATIONS
DEDUCTIONS
(a) Burn a small amount
of N on a spatula end.
(b) Shake one spatula
end-full of N with about
6cm3 of water and test
the mixture with litmus
paper.
Divide the mixture into
five portions.
(i) To the first portion,
add sodium hydroxide
solution.
(ii) To the second
portion, add a spatula
end-full of sodium
carbonate.
(iii) To third portion, add
4-5 drops of acidified
potassium
manganate(VII) solution
and heat.
182
(iv) To the fourth portion,
add 2-3 drops of Brady’s
reagent.
(v) To the fifth portion,
add 2-3 drops of neutral
iron(III) chloride
solution..
(f) Describe the nature of N.
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
…………………………………………………………………………………………………………………………
183