NAME: WAREN OTIENO 11588

ADM: ____________CLASS: 3 WEST

__________DATE: 11/04/2025
____________

FORM 3 FORM 3
Kenya Certificate of Secondary Education (K.C.SE)

CHEMISTRY
FORM 3 HOLIDAY ASSIGNMENT
THE MOLE

Revision Questions
1. When 94.5g of hydrated barium hydroxide, Ba(OH)2.nH2O were heated to constant mass, 51.3g of anhydrous
barium hydroxide were obtained. Determine the empirical formula of the hydrated barium hydroxide. (3 mks)

2. An alkanol has the following composition by mass: hydrogen 13.5%, oxygen 21.6% and carbon 64.9%.
Determine the empirical formula of the alcohol (C=12.0; H=1.0) =16.0). (2 marks)

3. 6.84g of aluminium sulphate were dissolve in 150cm3 of water. Calculate the molar concentration of the
sulphate ions in the solution.
(Relative formula mass of aluminium sulphate is 342) (3 marks)

4. When a hydrated sample of calcium sulphate CaSO 4.xH2O was heated until all the water was lost, the following
data recorded;
Mass of crucible = 30.296 g
Mass of crucible +hydrated salt = 33.111 g
Mass of crucible + anhydrous salt = 32.781 g
Determine the empirical formula of the hydrated salt.
(Relative formula mass of CaSO4 =136, H2O =18). (3 marks)

5. Phosphoric acid is manufactured from calcium phosphate according to the following equation.
Ca3(PO4)2(s) + 3H2SO4(l) → 2H3PO4(aq) + 3CaSO4(s)
Calculate the mass (in kg) of phosphoric acid that would be obtained if 155 kg of calcium phosphate reacted
completely with the acid
(Ca=40, P=31, S=32, O=16, H=1) (2 marks)

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 233 Chemistry Form 2
6. In an experiment to determine the percentage of magnesium hydroxide in an anti-acid, a solution containing
0.50 g of the anti-acid was neutralized by 23.0 cm3 of 0.010M hydrochloric acid. (Relative formula mass of
magnesium hydroxide =58)
Determine the:
(a) Mass of magnesium hydroxide in the anti-acid; (2 marks)

(b) Percentage of magnesium hydroxide in the anti-acid. (1 mark)

7. When 8.53g of sodium nitrate were heated in an open tube, the mass of oxygen gas produced was 0. 83g.Given
the equation of the reaction is:
ℎ𝑒𝑎𝑡
2NaNO3(s) → 2NaNO2(s) + O2(g)
Calculate the percentage of sodium nitrate that was converted to sodium nitrite. (Na= 23.0, N
= 14.0, O = 16.0) (3 marks)

8. Aluminium oxide reacts with both acids and bases.

(a) Write an equation for the reaction between aluminium oxide and hydrochloric acid.
(1 mark)

(b) Using the equation in (a) above, calculate the number of moles of hydrochloric acid that would react
completely with 153.0g of aluminium oxide. (Al = 27.0, O= 16.0) (2 marks)

9. The pressure of nitrogen gas contained in a 1 dm3 cylinder at -196 °C was 107 Pascals.
Calculate the:
(a) Volume of the gas at 25 °C and 10 5 Pascals. (1½ marks)

(b) Mass of nitrogen gas. (Molar volume of gas is 24 dm3, N = 14.0)

(1 ½ marks)

10. Analysis of a compound showed that it had the following composition:

69.42% carbon, 4.13% hydrogen and the rest oxygen.
(a) Determine the empirical formula of the compound. (C = 12.0, H = 1.0, O = 16.0) (2 marks)

(b) If the mass of one mole of the compound is 242, determine its molecular formula. (1 mark)

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 233 Chemistry Form 2
11. Use the information in the table below to answer the questions that follow. The letters do not represent the
actual symbols of the elements.
Element Atomic number Melting point (°C)
R 11 97.8
S 12 650.0
T 15 44.0
U 17 -102
V 18 -189
W 19 64.0

(a) Give the reasons why the melting point of:

(i) S is higher than that of R; (1 mark)

(ii) V is lower than that of U. (2 marks)

(b) How does the reactivity of W with chlorine compare with that of R with chlorine? Explain. (2 marks)

(c) Write an equation for the reaction between T and excess oxygen. (1 mark)

(d) When 1.15g of R were reacted with water, 600cm3 of gas was produced. Determine the relative atomic mass of R.
(Molar gas volume = 24000 cm3) (3 marks)

(e) Give one use of element V. (1 mark)

12. When lead (II) nitrate is heated, one of the products is a brown gas.
(a) Write the equation of the reaction that occurs. (1 mark)

(b) If 0.290 dm3 of the brow n gas was produced, calculate the mass of the lead (II) nitrate that was
heated. (R.F.M of lead (II) nitrate = 331; Molar gas volume = 24 dm 3). (2 marks)

13.
(a) State the Gay Lussac's Law. (1 mark)

(b) 10 cm3 of a gaseous hydrocarbon, C2HX required 30cm3 of oxygen for complete combustion. If steam and
20 cm3 of carbon (IV) oxide were produced, what is the value of x? (2 marks)

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 233 Chemistry Form 2
14. The data given below was recorded when metal M was completely burnt in air. M is not the actual symbol of
the metal. (R.A.M; M=56, O=16)
Mass of empty crucible and lid =10.240g
Mass of crucible, lid and metal M =10.352g
Mass of crucible, lid and metal oxide = 10.400g
(a) Determine the mass of:
(i) Metal M (½ mark)

(ii) Oxygen (½ mark)

(b) Determine the empirical formula of the metal oxide. (2 marks)

15. 10cm3 of concentrated sulphuric (VI) acid was diluted to 100 cm 3. 10 cm3 of the resulting solution was
neutralised by 36 cm3 of 0.1M sodium hydroxide solution. Determine the mass of sulphuric (VI) acid that was
in the concentrated acid. (S = 32.0; H= 1.0; O = 16.0). (3 marks)

16. The empirical formula of A is CH2Br. Given that 0.470g of A occupies a volume of 56 cm3 at 546K and 1
atmosphere pressure, determine its molecular formula.
(H = 1.0, C= 12.0, Br = 80.0, molar gas volume at STP = 22.4 dm3) (3 marks)

17. Describe how the percentage by mass of copper in copper carbonate can be determined. (3 marks)

18. When 15 cm3 of a gaseous hydrocarbon, P, was burnt in 100cm 3 of oxygen, the resulting gaseous mixture
occupied 70 cm3 at room temperature and pressure.
When the gaseous mixture was passed through potassium hydroxide solution, its volume decreased to 25 cm 3.

✓ What volume of oxygen was used during the reaction? (1 mark)

✓ Determine the molecular formula of the hydrocarbon (2 marks)

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 233 Chemistry Form 2
19. A solution was made by dissolving 8.2g of calcium nitrate to give 2 litres of solution. (Ca= 40.0; N=14.0; O=
16.0)
Determine the concentration of nitrate ions in moles per litre. (3 marks)

20. The grid given below represents part of the periodic table. Study it and answer the questions that follow. The
letters do not represent the actual symbol of the element.

M N P T

✓ Select a letter which represents an element that losses electrons most readily. Give a reason for your
answer. (2 marks)

✓ Explain why the atomic radius of P is found to be smaller than that of N (2 marks)

✓ Element M reacts with water at room temperature to produce 0.2 dm 3 of gas. Determine the mass of M
which was reacted with water.
(Molar gas volume at room temperature is 24 dm 3; Relative atomic mass of M=7)
(3 marks)

21. 100cm3 of 0.05 M sulphuric (VI) acid were placed in a flask and a small quantity of anhydrous sodium
carbonate added. The mixture was boiled to expel all the carbon (IV) oxide. 25cm 3 of the resulting solution
required 18 cm3 of 0.1 M sodium hydroxide solution to neutralize it. Calculate the mass of sodium carbonate
added.
(Na = 23.0; O=16.0; C=12.0) (3 marks)

22. An organic compound had the following composition 37.21% carbon,

7.75% hydrogen and the rest chlorine. Determine the molecular formula of the compound, given that the
molecular mass of the compound is 65.
(C=12.0; H=1.0; Cl=35.5) (3 marks)

23. Calculate the mass of Zinc oxide that will just neutralize dilute nitric (V) acid containing 12.6 g of nitric (V)
acid in water. (Zn = 65.0; O =16.0, H = 1.0, N = 14.0).
(3 marks)

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 233 Chemistry Form 2
24. A hydrocarbon contains 14.5% of hydrogen. If the molar mass of the hydrocarbon is 56, determine the
molecular formula of the hydrocarbon.
(C = 12.0; H = 1.0) (3 marks)

25. Zinc oxide can be obtained by heating zinc nitrate. A student heated 5.76 g of zinc nitrate.
(a) Write an equation for the reaction that occurred. (1 mark)

(b) Calculate the total volume of gases produced. (Molar gas volume is 24 dm3; Zn = 65.4; O = 16.0; N =
14.0). (4 marks)

(c) Identify the element that is reduced when zinc nitrate is heated. Give a reason.
(2 marks)

26. The empirical formula of lead(II) oxide was determined by passing excess dry hydrogen gas over 6.69g of
heated lead(II) oxide.
(a) What was the purpose of using excess dry hydrogen gas? (2 marks)

(b) The mass of lead was found to be 6.21g. Determine the empirical formula of the oxide. (Pb = 207.0 0 =
16.0) (2 marks)

27. 20 cm3 of ethanoic acid was diluted to 400 cm3 of solution. Calculate the concentration of the solution in moles
per litre. (C = 12.0; H = 1.0; 0 =16.0) (Density of ethanoic acid = 1.05 g/cm 3) (3 marks)

28. W is a colourless aqueous solution with the following properties:

- It turns blue litmus paper red.
- On addition of cleaned magnesium ribbon, it gives off a gas that burns with a pop sound.
- On addition of powered sodium carbonate, it gives off a gas which forms a precipitate with calcium
hydroxide solution.
- When warmed with copper(II) oxide powder, a blue solution is obtained but no gas is given off.
- On addition of aqueous barium chloride, a white precipitate is obtained.

(a) (i) State what properties (I) and (III) indicate about the nature of W. (1 mark)

(ii) Give the identity of W. (1 mark)

(iii) Name the colourless solution formed in (II) and (III). (2 marks)
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 233 Chemistry Form 2

(iv) Write an ionic equation for the reaction indicated in (V). (1 mark)

(b) Element V conducts electricity and melts at 933 K. When chlorine gas is passed over heated V, it forms a
vapour that solidifies on cooling. The solid chloride dissolves in water to form an acidic solution. The
chloride vapour has a relative molecular mass of 267 and contains 19.75% of V. At a higher temperature, it
dissociates to a compound of relative molecular mass 133.5. When aqueous sodium hydroxide is added to
the aqueous solution of the chloride, a white precipitate is formed which dissolves in excess alkali. (V =
27.0; CI = 35.5)
- Determine the:
I. empirical formula; (2 marks)

II. molecular formula. (2 marks)

- Draw the structure of the chloride vapour and label the bonds. (1 mark)

- Write an equation for the reaction that form a white precipitate with sodium hydroxide.
(1 mark)

29.
(a) When 0.048 g of magnesium was reacted with excess dilute hydrochloric acid at room temperature and
pressure, 50 cm3 of hydrogen gas was collected. (Mg = 24.0; Molar gas volume = 24.0 dm3)
(i) Draw a diagram of the apparatus used to carry out the experiment described above. (3 marks)

(ii) Write the equation for the reaction. (1 mark)

(iii) Calculate the volume of hydrogen gas produced. (2 marks)

(iv) Calculate the volume of 0.1 M hydrochloric acid required to react with 0.048 g of magnesium.
(3 marks)

30. 30.0 cm3 of aqueous sodium hydroxide containing 8.0 g per litre of sodium hydroxide were completely
neutralised by 0.294 g of a dibasic acid. Determine the relative formula mass of the dibasic acid. (Na = 23.0; O
= 16.0; H 1.0) (3 marks)

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 233 Chemistry Form 2
31. Distinguish between empirical and molecular formula of a compound. (1 mark)

32. A solution contains 40.3g of substance XOH per litre. 250.0 cm3 of this solution required 30.0cm3 of 0.3M
sulphuric(VI)acid for complete neutralisation.
(a) Calculate the number of moles of XOH that reacted. (½ mark)

(b) Determine the relative atomic mass of X. (1½ mark)

33. 5 g of calcium carbonate was strongly heated to a constant mass.

Calculate the mass of the solid residue formed. (Ca = 40.0; C = 12.0; 0 = 16.0). (2 marks)

34. Name an appropriate apparatus that is used to prepare standard solutions in the laboratory. (1 mark)

35. X g anhydrous sodium carbonate reacted completely with 45cm 3 hydrochloric acid giving 1.08 dm3 carbon
(IV) oxide at r.t.p. (molar gas volume at r.t.p = 24dm 3)
a) Give a balanced equation for the reaction. (1mk)

b) Calculate the value of X. (2mks)

c) Find out the molarity of the HCl solution. (2mks)

36. 20cm3 of a gaseous hydrocarbon, C2HX Required 70cm3 oxygen for complete combustion, forming 40cm3
carbon (IV) oxide. Assuming that all the volumes were measured at the same temperature and pressure
determine:

a) The value of X. (2mks)

b) The volume of steam formed. (2mks)

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 233 Chemistry Form 2
37. When 10g of ZnCO3 was heated, 6.48g of residue was left.
a) What observations are made during and after the reaction. (1mk)

b) Give a balanced equation for the reaction. (1mk)

c) Calculate the moles of CO2 formed in the reaction .(2mks)

d) If the residue was reacted with excess hydrochloric acid, calculate the mass of zinc chloride formed. Zn =
65, C = 12, O = 16, Cl = 35.5) (2mks)

38. 25cm3 of a 1M sulphuric acid solution required 2.8g of the base XOH(s) for complete neutralization.
a) Give a balanced equation for the reaction. (1mk)

b) Calculate the relative atomic mass of X. (2mks)

39. 25cm3 of 2M hydrochloric acid was reacted with 1.3g zinc metal.

a) Give a balanced equation for the reaction (1mk)

b) Using an appropriate explanation, state which of the reactants is in excess? (2mks)

c) Calculate the volume of hydrogen gas formed in the reaction at room temperature and pressure. (2mks)
(Zn = 65, molar gas volume at r.t.p = 24dm 3)

40. The diagram below shows an experiment in which carbon (II) oxide was reacted completely with a heated
oxide of iron. Study the diagram and the data shown below it and answer the questions that follow:

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 233 Chemistry Form 2
Mass of porcelain boat = 13.23g

Mass of porcelain boat + oxide of iron = 16.71g

Mass of porcelain boat + residue = 15.75g

a) Determine the empirical formula of the oxide. (2mks)

b) If the molecular mass of the oxide is 232 determine its molecular formula. (2mks)
( Fe = 56, O = 16)

41. Given:

- Solid A (Xg of KHCO3)

- 2.0M HCl solution B
- 0.1M NaOH(aq) ;solution D
20cm3 of solution B was measured into a 100ml beaker and all of solid A added to it. The resulting mixture was then
transferred into a 250ml volumetric flask and diluted up to the mark. The resulting solution was labelled solution C.

25cm3 of the solution C was titrated with Sodium Hydroxide solution and the results recorded as shown in table 1
below:

Table 1:

Titration Number I II III

Final Burette reading(cm3) 28.4 28.3

Initial Burette reading (cm3) 0.0 0.0

Volume of Sodium Hydroxide used 28.4 28.3

i) Complete table 1 (3mks)

ii) Work out the average titre (1mk)

iii) Calculate the number of moles of:

I) Sodium Hydroxide in the volume of solution D used (1mk)

II) Hydrochloric acid in 25cm3 of solution C (2mks)

III) Hydrochloric acid in 250cm3 of solution C (1mk)

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 233 Chemistry Form 2

IV) Hydrochloric acid in 20cm3 of solution B (1mk)

V) Hydrochloric acid that reacted with Solid A (2mks)

VI Solid A and its mass X. (2mks)

42. The formula of a hydrated salt of manganese is MnSO₄ • XH₂O. Given that the salt contains 24.7% manganese,
determine the value of X. (Mn = 55.0; S = 32.0; O = 16.0; H = 1.0) (3 marks)

43. Copper can be obtained from copper (II) oxide using carbon (II) oxide or coke.
(a) Name another reagent that can be used to obtain copper from copper (II) oxide. (1 mark)

(b) The equation for the reaction with carbon (II) oxide is:
CuO (s) + CO (g) → Cu (s) + CO₂ (g)
Calculate the maximum mass of copper that would be obtained using 200 dm³ of carbon (II) oxide (Cu = 63.5; Molar
volume of gas = 24.0 dm³). (2 marks)

44. An experiment was carried out to prepare crystals of magnesium sulphate. Excess magnesium powder was
added to 100 cm³ of dilute sulphuric (VI) acid in a beaker and warmed until no further reaction took place. The
mixture was filtered and the filtrate evaporated to saturation, then left to cool for crystals to form.
(a) (i) Write an equation for the reaction. (1 mark)

(ii) Explain why excess magnesium powder was used. (1 mark)

(iii) State how completion of the reaction was determined. (1 mark)

(iv) What is meant by a saturated solution? (1 mark)

(v) Explain why the filtrate was not evaporated to dryness. (2 marks)

(b) When bleaching powder, CaOCl₂, is treated with dilute nitric(V) acid, chlorine gas is released. This reaction can be
used to determine the chlorine content of various samples of bleaching powders and liquids.
(i) Write an equation for the reaction of nitric(V) acid with bleaching powder. (1 mark)
(ii) Calculate the volume of chlorine produced when 10 g of CaOCl₂ is treated with excess nitric(V) acid. (Ca = 40.0;
O = 16.0; Cl = 35.5; 1 mole of gas occupies 22.4 dm³ at s.t.p) (3 marks)

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 233 Chemistry Form 2
(c) Apart from use of chlorine gas in bleaches and water treatment, state two other uses of chlorine gas. (1 mark)

45. A student titrates 3 samples of 0.800 mol / dm³ aqueous sodium hydroxide, NaOH(aq), with aqueous
ethanedioic acid. In titration 1 the student:
• rinses and fills a burette with aqueous ethanedioic acid
• uses a volumetric pipette to add 25.0 cm³ of NaOH(aq) to a conical flask
• adds thymolphthalein indicator to the conical flask
• places the conical flask on a white tile
• adds aqueous ethanedioic acid from the burette while swirling the flask, adding drop by drop

Near the end-point, until the solution just changes colour. The student repeats the titration 2 more times.

(a) Fig. 2.1 shows the burette readings for two of the titrations.

a) Complete the titration table below using the above results (3 marks)
Titration 1 2 3
Final burette reading ( cm³) 20.1
Initial burette reading (cm³) 0.0
Volume of ethanedioic acid used (cm³) 20.1

(b) Calculate the average volume of aqueous ethanedioic acid needed to neutralise 25.0 cm³ of the aqueous sodium
hydroxide. Volume ............................................... cm³ - (1 mark)
(c) Calculate the number of moles of NaOH in 25.0 cm³ of 0.800 mol / dm³ NaOH(aq). Number of moles
............................................................................................................................................. (1 mark)

d) One mole of ethanedioic acid is neutralised by two moles of sodium hydroxide. Use your answers to (c) and (d) to
calculate the concentration, in mol / dm³, of ethanedioic acid. Give your answer to three significant figures.
Concentration ........................ mol / dm³ (2 marks)

(e) The formula of ethanedioic acid is C₂H₂O₄.nH₂O.

(i) 100 cm³ of the aqueous ethanedioic acid contains 6.3 g of C₂H₂O₄.nH₂O. Use your answer from (e) to calculate the
relative formula mass of C₂H₂O₄.nH₂O. (2 marks)

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 233 Chemistry Form 2
(ii) Use your answer from (e)(i) to deduce the value of n in C₂H₂O₄.nH₂O. Give your answer to the nearest whole
number. [H, 1; C, 12; O, 16] n ............................................................................................................................. (1 mark)

46. The following experiment was used to determine the percentage of copper (II) oxide in a mixture C, containing
copper and copper (II) oxide.
A sample of C was added to a previously weighed beaker, which was then reweighed.
• Mass of beaker + C = 32.65 g
• Mass of beaker = 27.80 g
(a) Calculate the mass of C used in the experiment. (1 mark)

(b) 50.0 cm³ of 1.00 mol/dm³ sulphuric (VI) acid (excess) was transferred to a beaker containing the sample of C. The
mixture was warmed gently while being stirred for a few minutes. The unreacted copper settled at the bottom of the
beaker leaving a colourless solution.
(i) What was the purpose of adding sulphuric (VI) acid? (1 mark)

(ii) What colour was the solution? (1 mark)

(iii) What volume of the sulphuric (VI) acid did copper not react with? (1 mark)

(iv) Write an equation for the reaction between copper (II) oxide and sulphuric (VI) acid. (1 mark)

The solution which remained after the copper was removed, was transferred to a volumetric flask and made up to 250
cm³ with distilled water. This was solution D.
c) 25.0 cm³ of D was transferred into a conical flask and A few drops of indicator was added. A burette was filled
with solution of 0.100 mol/dm³ sodium hydroxide. This was run into the conical flask containg D and the indicator
until the end point was reached. In this experiment the indicator is blue in acid and green in alkali
(i) State the indicator used in this experiment. (1 mark)

(ii) What was the colour change at the end point? (1 mark)

(iii) Why was it necessary to add a few drops of indicator? (1 mark)

The solution was run into the conical flask containing D and the titration was done.
(v) At what point are the titrations complete? (1 mark)

The diagrams below show parts of the burette with the liquid levels at the beginning and end of each titration.

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 233 Chemistry Form 2

(d) Use the diagrams to complete the following results table. (3 marks)

Titration 1 2 3
Final reading / cm³
Initial reading / cm³
Volume of sodium hydroxide used / cm³
Best titration results (√)

Summary
Tick (√) the best titration results. Using these results, the average volume of 0.100 mol / dm³ sodium hydroxide was
_____ __ cm³.
(e) Calculate the number of moles of sodium hydroxide in the average volume of 0.100 mol / dm³ sodium hydroxide
in (d). (2 marks)

Sodium hydroxide reacts with sulphuric acid according to the following equation.
2NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2H₂O(l)
(f) Calculate the number of moles of sulphuric (VI) acid which reacted with the sodium hydroxide in (e).(2 marks)

(g) Using your answer in (f), calculate the number of moles of sulphuric (VI) acid in 250 cm³ of solution D. (2 mks)

(h) Calculate the number of moles of sulphuric (VI) acid in 50.0 cm³ of 1.00 mol / dm³ sulphuric (VI) acid. (2 marks)

(i) By subtracting your answer in (g) from your answer in (b), calculate the number of moles of sulphuric (VI) acid
which reacted with the copper (II) oxide in C. (2 marks)

(j) Using your equation in (b)(iii), deduce the number of moles of copper (II) oxide in the sample of C. (2 marks)

(k) Using your answers in (a) and (j) calculate:

(i) the mass of copper (II) oxide in the sample of C. [Cu = 63.5, O = 16.0] (1 mark)

(ii) the percentage of copper (II) oxide in the sample of C. (1 mark)

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THE END