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Calculations – Past papers questions T.S.ELHAGE www.chem-exptc.com Page 1 of 42 TAREK S. ELHAGE 2010 IGCSE - CHEMISTRY Calculations Explanation & past papers questions www.chem-exptc.com UAE ABU DHABI www.chem-exptc.com

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Page 1: Calculations Questionss

Calculations – Past papers questions

T.S.ELHAGE www.chem-exptc.com Page 1 of 42

TAREK S. ELHAGE

2010

IGCSE - CHEMISTRY

Calculations Explanation & past papers questions

www.chem-exptc.com

U A E – A B U D H A B I www.chem-exptc.com

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Terms Atom, element, ion, molecule, compound, empirical and molecular formulae

°°°° Atom: the smallest particle of an element that retains all properties of that element

°°°° Subatomic particles: An atom is made up of protons, neutrons and electrons

⋅ Protons: positively charged particles present in the nucleus

⋅ Neutrons: neutral particles present in the nucleus

⋅ Electrons: negatively charged particle found around the nucleus

°°°° Nucleons: protons and neutrons are called nucleons

“one proton or one neutron is a nucleon”

°°°° Element: A Substance that cannot be split up (broken) chemically into simpler substances

“An element is made up of only one type of atoms”

°°°° Compound: A Substance that can be broken down chemically into simpler substances

“A compound is made up of different types of atoms that are chemically combined”

Chemical means are 1heat,

2light and

3electrolysis

°°°° Molecule: A Substance made up of two or more non-metal elements covalently bonded

°°°° Atomic number (Z): The number of protons in the nucleus of an atom of an element

°°°° Mass number (A): The sum of protons and neutrons in the nucleus of an atom

°°°° Isotopes: Atoms of the same element with the same number of protons but different

number of neutrons.

“Atoms with the same atomic number (Z) but different mass numbers”

"Isotopes have the same chemical properties BECAUSE they have the same number of

protons and electrons ( it is the electrons that determine the chemical properties)"

"Isotopes have different physical properties such as density, melting and boiling

points, BECAUSE they have a different masses"

°°°° Relative atomic mass ( )rA : The weighted average mass of all isotopes of an element

compared to 121 of the mass of one atom of carbon-12 isotope

°°°° Relative formula mass: The sum of relative atomic masses of all atoms in the chemical formula

°°°° Relative molecular mass ( )rM : the sum of relative atomic masses of all atoms in one molecule

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°°°° Amount of substance (mole): The amount of substance that contains exactly as many

particles ( atoms, ions or molecules) as there are atoms in 12 grams of carbon-12.

“12 grams of carbon-12 contain exactly ×02.6 1023 atoms”

°°°° Molar mass ( M ): The mass of mole of a substance

Equals to the relative atomic or molecular / formula mass in grams

°°°° Molar volume ( mV ): The volume occupied by 1 mole of any gas under the same conditions of

temperature and pressure

°°°° At standard temperature and pressure, 1 atm pressure and 298 K (25o

C ), the volume of 1 mole

of any gas is 3dm 24

°°°° Avogadro's constant (((( ))))AL or N : Number of particles in exactly 12 g carbon-12

AN = ×02.6 1023 particles

One mole of any substance contains exactly ×02.6 1023 particles

°°°° Avogadro's law: “equal volumes of all gases contain equal number of molecules (moles)

under the same temperature and pressure”.

"One mole of any gas occupy the same volume under the same conditions of

temperature and pressure"

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Formulas

1. Empirical formula (simplest formula): The formula which shows the smallest whole-number

ratio of the different atoms present

2. Molecular formula: The formula which shows the actual number of atoms of every element in

one molecule

Molecular Empirical Molecular Empirical

42HC 2CH 84HC 2CH

62HC 3CH 422 OCH 2HCO

6126 OHC OCH2 22ONa NaO

62ClAl 3AlCl OH 2 OH 2

3. Structural formula: the formula which shows the number of atoms in one molecule and the way

they are arranged relative to each other. It may show some of the bonds

2323 )(CHCHCHCH or 3

3

23|

CH

CH

HCCHCH −−

323 )( CHOHCHCHCH or 323|

CH

OH

HCCHCH −−

COOHCHCH 23 or OH

O||C2CH3CH −−

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Calculations

���� Calculating the amount (mole) of substance from a given mass and vice versa

massmolar

massGiven moles ofnumber = , applicable for substances in all states ( s ,l & g)

or rA

mn = for atoms /

rMm

n = for molecules

���� Calculating the amount (mole) of substance in a solution of known concentration in

( )3. −dmmol and vice versa

)dm(in volumemolarity moles ofnumber 3×= or )dm(in volume

mole)matter(in ofamount molarity

3=

Unit of molarity is 3mol.dm−

���� Calculating the mass of a substance in a given volume of a solution of known concentration in ( )3. −dmmol

massmolar )dm(in volumemolarity mass 3 ××=

or Mv c m ××=

���� Calculating the amount (mole) of substance of a gas and vice versa

24

)(dm gas of volumemoles ofnumber

3

= , applicable to gases only

Sample question:

Calculate the amount of substance (in mol) of NaCl in 250cm3 of 0.2 mol.dm-3 NaCl solution.

Answer:

⋅ 3dm 0.2501000250

volume ==

⋅ )volume(dm)ion(mol.dmconcentratn 33 ×=

∴ mol 0.050.250.2NaC of moles =×=l

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Writing the balanced equation

To balance an equation ensure

1st: Mass balance (by putting the appropriate coefficient in front of symbol /formula)

2nd: Charge balance (by adding electrons to either of the two sides, in redox reactions)

- Write the formula/ symbol of every reactant and product including the physical state of each. then

- Never change the subscript of any formula

Calculations using chemical equations 1. Change the given quantities to mole, when moles are not given

Determine the limiting reactant (only when the amounts of more than one reactant are given)

2. Write the balanced chemical equation

3. Write the mole stoichiometric ratio,

4. Put the given and required moles under the mole ratio and do the calculation

Calculating the reacted and the produced masses during a chemical reaction

Sample question:

In a reaction 5.4 grams of Aluminium reacted with enough amount of hydrochloric acid according to the following equation of reaction

)(3H (aq)2AlCl (aq) 6HC 2Al(s) 23 gl +→+

(a) Calculate the minimum mass of HCl needed to react with all the mass of Al

(b) Calculate the maximum mass of AlCl3 that could be obtained in this reaction.

(c) Calculate the maximum volume of 2H gas could be collected in this reaction.

Given molar masses in 1. −molg : [Al=27, H=1, Cl=35.5], molar volume of gases is 24 3dm

subscriptO2H 2←coefficent

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Answer:

(a) ).(

)()(

1−=molgM

gmAlmoles mol 0.2

27

5.4 ==

)(3H (aq)2AlCl (aq) 6HC 2Al(s) 23 gl +→+

⋅ Stoichiometric: mol2 mol6 mol2 mol3 mole ratio

⋅ From the balanced equation: mol2 Al : mol6 HCl Given / required mol2.0 : mol?

⇒ mol 0.6n :HCl of moles minimum 260.2 == ×

-136.5g.mol35.51 :)M(HC =+l

∴ g 25.1836.50.6m :HCl of mass minimum =×=×= Mn

(b) From the balanced equation: mol2 Al : mol2 3AlCl

Given / required mol2.0 : mol?

⇒ 0.2moln :CA of moles maximum 3 =ll

-1g.mol5.33135.5)3(27 :)CM(A =×+ll

∴ g 7.265.1330.2m :CofA mass maximum 3 =×=×= Mnll

(c) From the balanced equation: mol2 Al : mol3 2H

mol2.0 : mol?

⇒ 0.3mol2

30.2n :H of moles maximum 2 =×=

∴ 332 2.7240.324V :H of volumemaximum dmdmn =×=×=

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Calculating the reacted and the produced volumes of gases during a chemical reaction

Sample question:

In a reaction 1.25 dm3 of 83HC gas burnt with enough amount of oxygen according to the following equation of reaction

O(g)4H (g)3CO (g)5O (g)HC 22283 +→+

(a) Calculate the minimum volume of 2O gas used in this reaction.

(b) Calculate the maximum volume of 2CO gas produced in this reaction.

Given molar volume of gases is 243dm

Answer:

(a) O(g)4H (g)3CO (g)5O (g)HC 22283 +→+

⋅ Stoichiometric: mol1 mol5 mol3 mol4 mole ratio

⋅ Stoichiometric: 31dm 35dm 33dm 34dm for gases only volume ratio

⋅ From the balanced equation: 31dm 83HC : 35dm 2O

1.25 : 3?dm

∴minimum 32 25.6525.1V :O of volume dm=×=

(b) From the balanced equation: 31dm 83HC : 33dm 2CO

1.25 : 3?dm

∴maximum 32 75.3325.1V :CO of volume dm=×=

Percentage yield

100yield tal)(experimen Actual

)yield d(calculate ltheoreticayield % ×=

Where

⋅ Theoretical yield: amount of the product calculated using the given amount of the limiting reactant and the mole stoichiometric ratio of the balanced equation. Sometimes called the maximum obtained amount

⋅ Actual yield: amount of the product obtained from lab experiment / industrial process

⋅ Always actual yield < theoretical yield

Therefore, % yield < 100

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Sample question:

In a reaction 5.4 grams of Aluminium reacted with enough amount of hydrochloric acid according to the following equation of reaction

)(3H (aq)2AlCl (aq) 6HC 2Al(s) 23 gl +→+

(a) Calculate the maximum mass of AlCl3 that could be obtained in this reaction.

(b) What is the percentage yield of 3AlCl if the obtained amount of 3AlCl was 24g

(c) Suggest possible reasons for loss in the yield.

Given molar masses in 1. −molg : [Al=27, H=1, Cl=35.5]

Answer:

(d) From the balanced equation: mol2 Al : mol2 3AlCl

mol2.0 : mol?

⇒Theoretical (maximum) 0.2moln :CA of moles 3 =ll

-1g.mol5.33135.5)3(27 :)CM(A =×+ll

⇒ g 7.265.1330.2m :CofA mass (maximum) lTheoretica 3 =×=×= Mnll

100yield ltheoretica

yield actual%yield ×=

%89.8910026.7

24yield % =×=∴

���� Perform calculation using Avogadro’s constant

AN

molecules ofnumber molecules of moles ofnumber =

Example:

molecules moles ofnumber Xatoms of moles ofnumber x=

Where X: number of atoms of an element in one molecule/formula

molecules moles ofnumber Xions of moles ofnumber x=

Where X: number of ions in one formula of an ionic compound

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1. The diagram shows a model of a molecule of an organic acid.

What is the relative molecular mass of this acid?

A. 11 B. 40 C. 58 D. 74

2. The compound ethyl mercaptan, C2H5SH, has a very unpleasant smell.

What is its relative molecular mass?

A. 34 B. 50 C. 61 D. 62

3. Students are asked to state

• the number of atoms in one molecule of ethanoic acid,

• the relative molecular mass, Mr, of this acid. Which line is correct?

Number of atoms Mr

A.

B.

C.

D.

8

8

9

9

32

60

26

46

4. Which compound has the largest relative molecular mass, Mr?

A. CO2 B. SiO2

C. NO2 D. SO2

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5. The relative atomic mass of oxygen is 16 and that of hydrogen is 1.

This means that … (i) … of oxygen has the same mass as … (ii) … of hydrogen.

Which words correctly complete the gaps?

gap (i) gap (ii)

A.

B.

C.

D.

an atom

an atom

a molecule

a molecule

thirty-two molecules

eight molecules

sixteen atoms

eight atoms

6. How many oxygen atoms and double bonds are there in one molecule of ethanoic acid?

number of oxygen atoms

number of double bonds

A.

B.

C.

D.

1

1

2

2

0

1

0

1

7. The diagram shows a model of a molecule containing carbon, hydrogen and oxygen.

How many atoms of each element are in the molecule?

carbon hydrogen oxygen

A.

B.

C.

D.

1

2

2

6

6

5

6

2

2

1

1

1

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8. For complete combustion, one molecule of an organic compound needs 8 molecules of oxygen.

What could the formula of this compound be?

A. C5H11OH

B. C6H9OH

C. C6H11OH

D. C6H12

9. Water is formed when 48 g of oxygen combine with 6 g of hydrogen.

What mass of oxygen combines with 2 g of hydrogen?

A. 12g B. 16g C. 96g D. 4g

10. A substance X is heated in an evaporating basin until there is no further change.

Mass of basin and contents

Before heating 25.52 g

after heating 26.63 g

What could X be?

A. copper

B. copper(II) carbonate

C. copper(II) oxide

D. hydrated copper(II) sulphate

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11. A sample of clean, dry air is passed over hot copper until all the oxygen in the air reacts with the copper.

The volume of air decreases by 30 cm3.

What was the starting volume of the sample of air?

A. 60 cm3 B. 100 cm3 C. 150 cm3 D. 300 cm3

12. Two gases react as shown.

X2+ Y2 → 2XY reactants product

When measured at the same temperature and pressure, what is the value of

?reactants of volume

products of volume

A. 2

1

B. 1

C. 2

D. 4

heat

copperairdry clean

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13. (a) Benzene contains 92.3% of carbon and its relative molecular mass is 78. (i) What is the percentage of hydrogen in benzene?

……………………………………………………………………………………..……….

(ii) Calculate the ratio of moles of C atoms: moles of H atoms in benzene.

……………………………………………………………………………………..……..….

……………………………………………………………………………………..……..….

(iii) Calculate its empirical formula and then its molecular formula.

• The empirical formula of benzene is ………………………………………………….…

• The molecular formula of benzene is ……………….………………………………..…..

14. The structural formula of Vitamin C is drawn below.

(a) What is its molecular formula?

……………………………………………………………………………………..……….

(b) Name the two functional groups which are circled.

……………………………………………………………………………………..……….

C C

C C

CC

H H H

H

OH

OH

OHHO

OO

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15. The alkanes are generally unreactive. Their reactions include combustion, substitution and racking.

(a) The complete combustion of an alkane gives carbon dioxide and water. (i) 20 cm3 of butane is mixed with 150 cm3 of oxygen, which is an excess. The mixture is

ignited. What is the volume of unreacted oxygen left and what is the volume of carbon dioxide formed?

C4H10(g) + 2

16 O2 (g) → 4CO2(g) + 5H2O (l)

Volume of oxygen left = ……………………………………………………cm3

Volume of carbon dioxide formed = ……………………………………… cm3

16. (a) One piece of marble, calcium carbonate, 0.3 g, was added to 5 cm3 of hydrochloric acid, concentration 1.00 mol/ dm3.

(i) Which reagent is in excess? Give a reason for your choice.

mass of one mole of CaCO3 = 100 g

number of moles of CaCO3 = ………………………………………………..

number of moles of HCl = …………………………………..………………

reagent in excess is …………………………………………..………………

reason …………………………………………………………………………

Use your answer to(i) to calculate the maximum volume of carbon dioxide produced

measured at r.t.p.

………….…………………………..,,………………………..……….…..……..…..

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17. Crystals of sodium sulphate-10-water, Na2SO4.10H2O, are prepared by titration.

(a) 25.0 cm3 of aqueous sodium hydroxide is pipetted into a conical flask. A few drops of an indicator are added. Using a burette, dilute sulphuric acid is slowly added until the indicator just changes colour. The volume of acid needed to neutralize the alkali is noted.

(b) Using 25.0 cm3 of aqueous sodium hydroxide, 2.24 mol / dm3, 3.86 g of crystals were obtained. Calculate the percentage yield.

2NaOH + H2SO4 → Na2SO4 + 2H2O

Na2SO4 + 10H2O → Na2SO4.10H2O

Number of moles of NaOH used = ………………………………………………

Maximum number of moles of Na2SO4.10H2O that could be formed =…………

Mass of one mole of Na2SO4.10H2O = 322 g

Maximum yield of sodium sulphate-10-water = …….…………………………. g

Percentage yield = …….…………………………………………………………%

burette filled with sulphuric acid

25.0 cm3 of sodium hydroxide(aq) concentration 2.24 mol/dm3

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18. A better way of measuring the degree of unsaturation is to find the iodine number of the unsaturated compound. This is the mass of iodine that reacts with all the double bonds in 100 g of the fat.

Use the following information to calculate the number of double bonds in one molecule of the fat.

Mass of one mole of the fat is 884 g.

One mole of I2 reacts with one mole

The iodine number of the fat is 86.2 g.

Complete the following calculation.

100 g of fat reacts with 86.2 g of iodine.

884 g of fat reacts with ……………………………………………… g of iodine.

One mole of fat reacts with…………………………………… moles of iodine molecules.

Number of double bonds in one molecule of fat is …………………………………

End of question

19. Propene reacts with hydrogen iodide to form 2 - iodopropane.

CH3–CH=CH2 + HI → CH3–CHI–CH3

1.4 g of propene produced 4.0 g of 2 - iodopropane.

Calculate the percentage yield.

moles of CH3–CH=CH2 reacted = …………………….…………..………

maximum moles of CH3–CHI–CH3 that could be formed = ………...……

mass of one mole of CH3–CHI–CH3 = 170 g

maximum mass of 2 - iodopropane that could be formed = ………...……

percentage yield = ………………………………………………………… %

C C

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20. When calcium carbonate is heated strongly, it decomposes.

CaCO3 → CaO + CO2

(i) Calculate the relative formula mass of:

CaCO3 …………………………………………………………………………

CaO …………..………………………………………………………………

(ii) 7.00 kg of calcium oxide was formed. What mass of calcium carbonate was heated?

……………………………………………………………………………………

…………………………………………………………………………………………

21. An ore of copper is the mineral, chalcopyrite. This is a mixed sulphide of iron and copper.

(a) Analysis of a sample of this ore shows that 13.80 g of the ore contained 4.80 g of copper, 4.20 g of iron and the rest sulphur.

Complete the table and calculate the empirical formula of chalcopyrite.

copper iron sulphur

composition by mass / g

4.80 4.20

number of moles of atoms

simplest mole ratio of atoms

The empirical formula is …….………………………………….………………

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22. (a)The following method is used to make crystals of hydrated nickel sulphate.

An excess of nickel carbonate, 12.0 g, was added to 40 cm3 of sulphuric acid, 2.0

mol/dm3. The unreacted nickel carbonate was filtered off and the filtrate

evaporated to obtain the crystals.

NiCO3 + H2SO4 → NiSO4 + CO2 + H2O

NiSO4 + 7H2O → NiSO4.7H2O

Mass of one mole of NiSO4.7H2O = 281 g

Mass of one mole of NiCO3 = 119 g

(i) Calculate the mass of unreacted nickel carbonate.

Number of moles of H2SO4 in 40 cm3 of 2.0 mol/dm3 acid = ………………….

Number of moles of NiCO3 reacted = ……………….

Mass of nickel carbonate reacted = ………………… g

Mass of unreacted nickel carbonate =………………. G

(ii)The experiment produced 10.4 g of hydrated nickel sulphate.

Calculate the percentage yield.

The maximum number of moles of NiSO4 .7H2O that could be formed =

…………..………………………………………………………………….………

The maximum mass of NiSO4.7H2O that could be formed = ……………………. g

The percentage yield = …………………………………………………………… %

(b) *Gypsum is hydrated calcium sulphate, CaSO4.xH2O. It contains 20.9% water by mass. Calculate x.

Mr: CaSO4, 136; H2O, 18.

79.1 g of CaSO4 = ……………………………………….……… moles

20.9 g of H2O = ……………………………………………..…… moles

x = ……………….

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23. Chemists use the concept of the mole to calculate the amounts of chemicals involved in a reaction

(a) 3.0 g of magnesium was added to 12.0 g of ethanoic acid.

Mg + 2CH3COOH → (CH3COO)2Mg + H2

The mass of one mole of Mg is 24 g.

The mass of one mole of CH3COOH is 60 g.

(i) Which one, magnesium or ethanoic acid, is in excess? You must show your reasoning

.........................................................................................................................................

.........................................................................................................................................

.........................................................................................................................................

(ii) How many moles of hydrogen were formed?

..........................................................................................................................................

(iii) Calculate the volume of hydrogen formed, measured at r.t.p.

...........................................................................................................................................

(b) In an experiment, 25.0 cm3 of aqueous sodium hydroxide, 0.4 mol / dm3, was neutralized by 20.0 cm3 of aqueous oxalic acid, H2C2O4.

2NaOH + H2C2O4 → Na2C2O4 +2H2O

Calculate the concentration of the oxalic acid in mol / dm3.

(i) Calculate the number of moles of NaOH in 25.0 cm3 of 0.4 mol / dm3 solution.

..................................................................................................................................................

(ii) Use your answer to (i) and the mole ratio in the equation to find out the number of moles of H2C2O4 in 20 cm3 of solution.

....................................................................................................................................................

(iii) Calculate the concentration, mol / dm3, of the aqueous oxalic acid.

.....................................................................................................................................................

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24. (a)Silicon reacts with chlorine to produce silicon tetrachloride, as shown in the

following equation Si(s) + 2Cl2(g) → SiCl4(l)

(i) Calculate the mass of silicon tetrachloride obtained from 10.0 g of silicon.

(ii) Calculate the minimum volume of chlorine that would be required to react completely

with 10.0 g of silicon.

[1 mol of gas occupies 24.0 dm3 under the conditions of the experiment]

3. (a) Sodium chloride, NaCl, can be made by the reaction of sodium with chlorine.

2Na(s) + Cl2(g) → 2NaCl(s)

(i) Calculate the maximum mass of sodium chloride which could be obtained

from 92 g of sodium

*(ii) Calculate the concentration of the solution obtained when this mass of

sodium chloride is dissolved in water and made up to a volume of 10 dm3 with

distilled water

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(iii) Calculate the volume of chlorine gas required to react with 92 g of sodium.

[1 mol of gas occupies 24 dm3 under the conditions of the experiment]

25. In a second experiment, the student reacted 1.20 g of magnesium with 2.00 mol dm-3

hydrochloric acid. [Ar: Mg, 24.0; Cl, 35.5].

Mg(s) + 2HCl (aq) → MgCl2 (aq) + H2(g)

(i) How many moles of Mg were used in the experiment?

(ii) Calculate the minimum volume of 2.00 mol dm-3 hydrochloric acid needed to react

completely with this amount of magnesium.

(iii) Calculate the volume of H2 gas that would be produced at room temperature and

pressure (r.t.p.). [1 mole of gas molecules occupies 24.0 dm3 at r.t.p.]

26. The sulphur dioxide content of a wine can be found by titration . An analyst found that the sulphur dioxide in 50.0 cm

3 of white wine reacted with exactly 16.4 cm

3 of 0.01 mol. dm-3

aqueous iodine.

(a) How many moles of iodine, I2, did the analyst use in the titration?

(b) How many moles of sulphur dioxide were in the 50.0 cm3 of wine?

(c) What was the concentration of sulphur dioxide in the wine

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(i) in mol dm-3;

(ii) in g dm-3?

27. Balance the equation for the oxidation of PbO.

PbO (s) + O2(g) → Pb3O4 (s)

(a) What is the molar mass of Pb3O4? [Ar: O, 16.0; Pb, 207.0]

(b) Calculate the mass of Pb3O4 that could be formed from 0.300 mol of PbO.

28. A compound A is formed when chlorine is bubbled through hot concentrated potassium

hydroxide solution. (a) Ananlysis of A shows that it contains 31.84 % potassium, 28.98% chlorine and the

remainder is oxygen.

How that the empirical formula of A is 3KClO

(b) On being heated strongly solid A decomposes completely to give oxygen gas and solid potassium chloride

3KClO → KCl2 + 23O

If 1.00 g of solid A decomposed completely in this way, calculate the volume of

oxygen gas produced at room temperature and pressure.

[1 mole of gas at room temperature and pressure occupies 24.0 dm3 ]

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29. Potassium forms a superoxide, KO2. This reacts with carbon dioxide according to the equation:

4KO2(s) + 2CO2(g) → 2K2CO3(s) + 3O2(g)

(a) Carbon dioxide gas was reacted with 4.56 g of potassium superoxide.

(i) Calculate the amount, in moles, of KO2 in 4.56 g of potassium superoxide.

(ii) Calculate the amount, in moles, of carbon dioxide that would react with 4.56 g of

potassium superoxide.

(iii)Calculate the volume of carbon dioxide, in dm3, that would react with 4.56 g of

potassium superoxide. Assume that 1.00 mol of a gas occupies 24 dm3 under the conditions of the experiment.

(iv) What volume of oxygen gas, in dm3, measured under the same conditions of pressure

and temperature would be released?

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Percentage composition of a compound by mass, atoms (moles)

The percentage composition of a mixture( or a compound) by mass: ( )mm% = 100×

samplewhole

part

m

m

For Compounds

Generally, the percentage composition of a compound by mass: ( )mm% = 100×

compound

element

m

m

When only the chemical formula is given then: ( )mm% = 100×

compound

element

M

xM,

where x represents the number of atoms of each element in one molecule (formula)

For compounds only, the percentage composition by atoms and the percentage composition by moles are equal.

%by atoms 100×=Atomicity

elementeachofatomsofnumber=%by mole

Atomicity: the sum of all atoms in one molecule (formula) of a compound

Example: the atomicity of 43POH is 3+1+4=8, Hence, %(H) by atoms= %5.371008

3 =×

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Mass Percentage

A. From the chemical formula (molecular and structural)

1. Calculate the percentage composition of each of the following by mass and by moles (atoms) a) 3FeCl b) 43PONa c) 4KHSO d) 434)( PONH

e) 66HC f) OHHC 52 g) 722 OCrK h) 3CaCO

2. The molecular formula of Vitamin A is OHC 3020 . Determine its percentage composition by mass

3. Ethylene glycol, used as permanent antifreeze, has the structural formula

(a) Write its molecular formula, ………………………………………

(b) calculate its percentage composition by mass.

% O

% C

% H

HO

H

H

C

H

H

COH −−−−−|

||

|

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4. Glycine (glycerol) is a tri-alcohol of the structural formula

(a) Write its molecular formula, ………………………………………

(b) calculate its percentage composition by mass.

_ % O

_ % C

_ % H

(i) Determine mass percentage composition of the following compound of formula i. Phosphoric acid, 43POH

_ % O

_ % P

_ % H

5. Nitroglycerin: violent explosive

_ % O

_ % N

_ % H

OHCH

OHHC

OHHC

2

2|

2|

22|||

HCHCHC −−

2O

O|N 2O

O|N 2O

O|N

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6.Aspirin (analgesic and anti pyretic medicinal drug)

_ % O

_ % C

_ % H

B. From masses of the constituents elements

(i) 12.5 grams sample of a compound, containing only phosphorous and sulfur, analyzed and found to contain 7.04 grams of phosphorous. What is the percentage composition of this compound by mass?

(ii) A sample of an air pollutant, composed of sulfur and oxygen, found to contain 1.40 grams of sulfur and 2.10 grams of oxygen. Determine the percentage composition of the compound by mass.

−O

O||C 3CH−

OH−

O||C

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Empirical and Molecular formulae

1. This question is about some of the chemicals used in car engines and their reactions. (a) A Compound X, shown below, is one component of petrol.

(i) Give the Molecular formula of X.........................................................

(ii) Give the empirical formula of X .........................................................

(b) When potassium is burnt in excess oxygen, a compound is produced that contains 54.9

% potassium Calculate the percentage of oxygen present and hence calculate the empirical formula of this compound.

2. Potassium superoxide contains 54.9 % potassium by mass. Show that the empirical formula of this compound is KO2.

3. Show that the following data are consistent with the empirical formula CaN2O6.

Symbol of element % by mass

Ca 24.4

N 17.1

O 58.5

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4. What is the empirical formula of a compound that contains 30.4% nitrogen and 69.6%

oxygen by mass?

5. (a)Lead compounds are extensively used to provide the colour in paints and pigments. ‘White lead’, used for over 2000 years as a white pigment, is based upon lead carbonate. Analysis shows that lead carbonate has the following percentage composition by mass: Pb, 77.5%; C, 4.5%; O, 18.0%.

Calculate the empirical formula of lead carbonate. [Ar: C, 12.0; O, 16.0; Pb, 207.0]

(b) Phosgene is a compound of chlorine, carbon and oxygen, used to make polyurethanes and dyes. Phosgene has the percentage composition by mass: Cl, 71.7%; C, 12.1%; O, 16.2%. Show that

the empirical formula of phosgene is COCl2 .

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(c) Copper occurs naturally as the mineral malachite. The composition, by mass, of malachite is as follows:

[Cu = 57.5% C = 5.4% O = 36.2% H = 0.9%] Calculate its empirical formula.

(d) A compound F has the composition Cu 49.4%, S 12.5%, O 37.4%, H 0.78% by mass.

Calculate its empirical formula.

(e) A hydrocarbon has the empirical formula CH2 and a relative molecular mass

of 70. Write the molecular formula of the hydrocarbon.

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1. student added 30 cm3of 1.5 mol/dm3aqueous silver nitrate to a beaker containing 50 cm3 of 1.0 mol/dm3aqueous sodium bromide. A precipitate of silver bromide was produced.

(a) (i) What colour was the precipitate?

...............................................................................................................................................................................

(ii) Name the method by which this precipitate was separated from the mixture.

................................................................................................................................................................................ [2]

(b) (i)Calculate the number of moles of silver nitrate contained in 30 cm3of 1.5 mol/dm

3 aqueous

silver nitrate.

................................................................................................................................................................... moles

(ii) Calculate the number of moles of sodium bromide contained in 50 cm3of 1.0 mol/dm3aqueous

sodium bromide.

................................................................................................................................................................. moles [2]

Sodium bromide reacts with silver nitrate according to the equation below.

AgNO3+ NaBr →AgBr + NaNO3

(c) Using this equation and your answers to (b), calculate the mass of silver bromide produced in this experiment.

[Ar: Ag, 108; Br, 80]

................................................................................................................................................................. g [2]

(d) The student repeated the experiment using 40 cm3of 1.5 mol/dm3aqueous silver nitrate with 50 cm3of 1.0

mol/dm3sodium bromide. Calculate the mass of silver bromide produced in this experiment. ................................................................................................................................................................ g [2]

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2. A student was given some hydrated sodium carbonate crystals, Na2CO3.xH2O, where x is a whole number. They were placed in a previously weighed container, which was reweighed.

mass of container + sodium carbonate crystals = 9.87 g mass of container = 5.83 g

(a) Calculate the mass of sodium carbonate crystals used in the experiment.

.......................................g [1]

The container and crystals were heated to remove the water of crystallisation and then reweighed. This process was repeated until there was no further change in mass.

(b) Describe the appearance of the sodium carbonate crystals after heating.

mass of container + sodium carbonate after heating = 7.35 g

(c) (i)Calculate the mass of sodium carbonate which remained after heating.

.......................................g [1]

(ii)Calculate the mass of water which was lost from the crystals.

.......................................g [1] (d) (i) Calculate the relative formula mass of sodium carbonate, Na2CO3, and the relative

formula mass of water. [Ar: Na, 23; C, 12; O, 16; H, 1]

relative formula mass of sodium carbonate .......................................... [1]

relative formula mass of water ............................................................. [1]

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(e) Using your answers to (c) and (d), calculate

(i)the number of moles of sodium carbonate which remained after heating,

....................................... [1]

(ii)the number of moles of water which were lost on heating.

....................................... [1]

(f) Using your answers to (e) calculate the value of x in the formula Na2CO3.xH2O.

x =....................................... [2]

[Total: 10]

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3. A student found the composition of air using the apparatus shown below.

Syringe A contained 80 cm3 of air. The air was forced over heated copper into syringe B. The air was then forced back into syringe A.

The process was repeated several times until the volume of gas forced back into syringe A was constant.

The diagram below shows the volume of gas in syringe A after the experiment was finished.

(a) (i)Name the major component of the gas remaining in syringe A.

.............................................................[1]

(ii)What is the volume of gas remaining in syringe A? ....................................................................[1]

(iii) Calculate the percentage of oxygen in the original sample of air.

.......................................[1]

(b) The copper reacted with oxygen in the air to produce copper(II) oxide.

(i) Write the equation for this reaction.

............................................................................................................................................................... [1]

(ii) What colour is copper(II) oxide? ..........................................................................................[1]

copper

A B heat

100 80 60 40 20 20 40 60 80 100

100 80 60 40 20

A

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(c) In another experiment 0.16 g of copper was placed in the tube.

(i) Calculate the number of moles of copper in the tube. [Ar: Cu, 64]

.......................................[1]

(ii) Using your equation in (b)(i) deduce the number of moles of oxygen required to react with 0.16 g of copper.

............................................[1]

(iii) Using your answer to (c)(ii) calculate the volume of oxygen required to react with 0.16 g of

copper. [1 mol of a gas measured at 25 °C occupies a volume of 24 dm3.]

............................................cm3 [1]

(iv) Using your answers to (a)(iii) and (c)(iii) calculate the volume of air required to react with 0.16 g of

copper.

............................................cm3 [1]

[Total: 9]

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(d) The nitrates of metallic elements also decompose when heated. Calcium nitrate decomposes to form calcium oxide, nitrogen dioxide and oxygen.

2Ca(NO3)2(s) → 2CaO(s) + 4NO2(g) + O2(g) A 0.010 mol sample of calcium nitrate is heated. Calculate the number of moles of gas produced when this sample is completely decomposed.

4. Verdigris has the formula [Cu(CH3CO2)2]2.Cu(OH)2.xH2O. It has a relative formula mass of 552. Calculate the value of x in the formula.

x is …………………. [2]

[Total: 5]

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5. Tartaric acid can be extracted from grape juice. The structure of tartaric acid is shown below.

(a) Deduce the empirical formula of tartaric acid. ......................................................................................................................................[1]

(b) A solution of tartaric acid was titrated with 0.100 mol/ dm3 potassium hydroxide.

C2H2(OH)2(CO2H)2 + 2KOH → C2H2(OH)2(CO2K)2 + 2H2O tartaric acid

It required 6.00 cm3 of the potassium hydroxide solution to neutralise 20.0 cm

3 of

tartaric acid. Calculate the concentration, in mol / dm3, of the tartaric acid solution.

....................................mol/dm3 [3]

(c) Tartaric acid is purified by recrystallisation. On analysis, 8.00 g of impure tartaric acid was found to contain 7.40 g of pure

tartaric acid. Calculate the percentage purity of the impure tartaric acid.

..............................................% [1]

OHOH

HCO22HCO

HH

CC

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6. Iron(II) sulphate, FeSO4, is easily oxidised to iron(III) sulphate.

(a) Calculate the percentage by mass of iron in iron(II) sulphate.

........................................ % [2]

(b) A sample of iron(II) sulphate is dissolved in water. Describe a test to show the presence of sulphate ions in this solution.

reagents ...........................................................................................................................

observation ............................................................................................................................. [2]

(c) In the presence of aqueous hydrogen ions and dissolved oxygen, aqueous iron(II) ions are oxidised to form iron(III) ions and water. Write an ionic equation for this reaction.

(d) Aqueous iron(II) ions can also be oxidised by reaction with acidified potassium dichromate(VI),

K2Cr2O7. At the same time aqueous dichromate(VI) ions are reduced.

(i) Describe the colour change of the chromium-containing species during the reaction.

........................................................................................................................................... [1]

(ii) Describe the colour change of the iron-containing species during the reaction.

............................................................................................................................................ [1]

(e) An impure sample of iron(II) sulphate was analysed by titration.

The sample was dissolved in 25.0 cm3 of dilute sulphuric acid and then titrated against 0.0400 mol/dm3 potassium dichromate(VI) solution.

19.0 cm3 of potassium dichromate(VI) solution was required to reach the end-point.

(i) Calculate the number of moles of potassium dichromate(VI) used in the titration.

........................................ moles [1]

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(ii) One mole of potassium dichromate(VI) reacts with six moles of iron(II) ions. Calculate the mass, in grams, of iron(II) ions in the sample analysed.

mass of iron(II) ions........................................ g [2]

[Total: 11]

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7. Dilute ethanoic acid and dilute hydrochloric acid both react with magnesium ribbon to form hydrogen.

(a) Give the formula of one ion found in both of these dilute acids. [1]

(b) Magnesium ribbon reacts with hydrochloric acid as shown in the equation.

Mg + 2HCl → MgCl2 + H2

A 0.24 g sample of magnesium ribbon is added to 5.0 cm3 of 2.0 mol/dm3 hydrochloric acid.

(i) Which reactant, magnesium or hydrochloric acid, is in excess? Use calculations to explain your answer. [2]

(iii) Calculate the maximum mass of magnesium chloride that can be formed in this reaction.

[2]

(iv) A 0.24 g sample of magnesium ribbon is added to 5.0 cm3 of 2.0 mol/dm3 ethanoic acid. Explain why this reaction forms the same volume of hydrogen but takes place much more slowly than the reaction of the same mass of magnesium with 5.0 cm3 of 2.0 mol/dm3 hydrochloric acid. [3]

(c) (i) Write an equation for the reaction between dilute ethanoic acid and sodium carbonate. [1] ………………………………………………………………………………………………………….

(ii) What observations would be made during this reaction? [1]

[Total: 10]

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8. Brass is an alloy of zinc and copper. (a) Describe, with the aid of a labelled diagram, the structure of a metal such as copper. [2]

(b) Explain, in terms of their structures, why both zinc and copper are good conductors of electricity. [1]

(c) A 1.2 g sample of powdered brass was analysed by reaction with excess dilute sulphuric acid. The zinc reacts as shown in the equation to form 0.072 dm3 of hydrogen measured at room temperature and pressure.

Zn + 2H+→ Zn2+ + H2

(i) Suggest why brass was used in a powdered rather than lump form. [1]

(ii) Calculate the mass of zinc in the sample of brass. [2]

(iii) Calculate the percentage of zinc in the sample of brass. [1]

(d) Describe how aqueous ammonia can be used to show that only the zinc in the sample reacted with the acid. [3]

[Total: 10]