3 Chemical Formulae and Equations Edited

FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________

CHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS

A RELATIVE ATOMIC MASS AND RELATIVE MOLECULAR MASS

Learning Outcomes You should be able to:

state the meaning of relative atomic mass based on carbon-12 scale state the meaning of relative molecular mass based on carbon-12 scale state why carbon-12 is used as a standard for determining relative atomic mass and relative molecular mass calculate the relative molecular mass of substances

Activity 1

1. The relative atomic mass (Ar) of an element is the ……………………. mass of one atom of the ………………. when compared with 1/12 of the mass of an atom of carbon-12.

2. By comparing relative atomic masses, we can determine the ratio of the actual masses of atoms. Example:

(i) Calculate how many times heavier are 3 calcium atoms compared to 5 carbon atoms. [Relative atomic mass: C, 12; Ca, 40]

(ii) How many magnesium atoms will have the same mass as two silver atoms? [Relative atomic mass: Mg, 24; Ag, 108]

Relative atomic mass of an element= the average mass of one atom of an element 1/12 x the mass of a carbon-12 atom

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________3. The relative molecular mass (Mr) of a substance is the average ……………… of a ……………. of the substance when compared with 1/12 of the mass of one carbon-12 atom.

4. A molecule is made up of a number of ………………. Therefore, the relative molecular mass of a substance is calculated by adding up the ………………………….. of all the atoms present in a molecule of the substance. Example:

(i) Calculate the relative molecular mass of ammonia. [Relative atomic mass: H, 1; N, 14]

(ii) Calculate the relative molecular mass of water. [Relative atomic mass: H, 1; O, 16]

5. The term ‘relative molecular mass’ can only be used for substances that are made up of .................... For ionic compounds, the term ……………………………… (Fr) is used instead.

Relative molecular mass of a substance= the average mass of one molecule of a substance 1/12 x the mass of a carbon-12 atom

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________ 6. The relative formula mass of an ionic compound is calculated by adding up the ……………………… of all atoms in its formula. Example:

(i) Calculate the relative formula mass of sodium chloride. [Relative atomic mass: Na, 23; CI, 35.5]

(ii) Calculate the relative formula mass of potassium oxide. [Relative atomic mass: K, 39; O, 16]

(iii) Calculate the relative formula mass of copper(II) chloride. [Relative atomic mass: Cu, 64; CI, 35.5]

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________B THE MOLE AND THE NUMBER OF PARTICLES


define a mole state the meaning of Avogadro constant relate the number of particles in one mole of a substance with the Avogadro constant solve numerical problems to convert the number of moles to the number of particles of a given substance and vice versa

Activity 2

1. In chemistry, we use the unit …………………… to measure the amount of substance. It has the symbol ………………………………

2. One ……………….. is defined as the amount of substance that contains as many particles as the number of atoms in exactly 12 g of carbon-12. One mole of substance contains ………………………… particles.

3. The Avogadro constant, NA is defined as the number of………………….. in one mole of a substance. The value of the Avogadro constant is ……………………………..

4. 1 mole of atomic substance contains ……………………………. atoms.

5. 1 mole of molecular substance contains ……………………………. molecules.

6. 1 mole of ionic substance contains ……………………………… formula units.

7. Relationship between number of moles and number of particles:

x NA

÷ NA

8. Find the number of particles for the substances given the number of moles.

Number of moles Number of particles

(i) 0.5 mole of carbon

Number of moles Number of particles

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(ii) 0.2 mole of hydrogen gas

(iii) 2 mole of carbon dioxide gas

(iv) 0.3 mole of zinc bromide

(v) 0.5 mole of iodine molecule

(vi) 0.5 mole potassium bromide

9. Find the number of moles for the substances given the number of particles.

Number of particles Number of moles

(i) 3.01 x 1022 of water molecules

(ii) 3.01 X 1024 of hydrogen molecules

(iii) 1.806 x 1023 of oxygen molecules

(iv) 1.505 x 1024 of bromine molecules

(v) 9.03 x 1023 of carbon dioxide molecules

(vi) 1.806 x 1023 formula units of zinc bromide

C THE MOLE AND THE MASS OF SUBSTANCES


state the meaning of molar mass relate molar mass to the Avogadro constant relate molar mass of a substance to its relative atomic mass, relative molecular mass or

relative formula mass

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solve numerical problems to convert the number of moles of a given substance to its mass and vice versa

Activity 3

1. Molar mass is the mass of one ………………………… of a substance. It has unit of …………………….

2. The molar mass of a substance is numerically equal to its ………………………, ………………………… or ………………………………

3. The relationship between the number of moles and the mass of a substance:

x Molar mass

÷ Molar mass

3. Find the mass for the substances given the number of moles.

Number of moles Mass

(i) 0.1 mole of magnesium

(ii) 0.5 mole of copper

(iii) 1.5 moles of carbon dioxide

(iv) 0.01 mole of ammonia gas

(v) 0.3 mole aluminium

(vi) 0.05 mole of sodium hydroxide

4. Find the number of moles of the substances given the mass.

Mass Number of moles

(i) 49.2 g of calcium nitrate

Number of moles Mass (g)

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(ii) 57.5 g of sodium

(iii) 4.04 g of potassium nitrate

(iv) 70 g of carbon monoxide gas

(v) 4 g of hydrogen gas

(vi) 11.2 g of iron

D NUMBER OF MOLES AND VOLUME OF GAS


state the meaning of molar volume of a gas relate molar volume of a gas to the Avogadro constant make generalisation on the molar volume of a gas at a given temperature and pressure calculate the volume of gases at STP or room conditions from the number of moles

and vice versa solve numerical problems involving number of particles, number of moles, mass of substances and volume of gases at STP or room conditions

Activity 4

1. The molar volume of a gas is defined as the ................................ of one ................................... of the gas.

2. One mole of any gas always has the …………........................ under the same temperature and pressure.

3. The molar volume of any gas is ……………………… at STP or ……………………….. at room conditions.

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Number of moles Volume of gas (dm3)


4. STP refers to standard temperature of ………………. and pressure of …………………. Room conditions refer to the temperature of ………………… and pressure of …………………..

5. The relationship between the number of moles and volume of gas:

x Molar volume x 22.4/24 dm3

÷ Molar volume

4. Calculate the volume of gas given the number of moles at STP.

22.4 dm3

at STP

1 mole of hydrogen gas, H2

22.4 dm3

at STP

1 mole of oxygen gas, O2

22.4 dm3

at STP

1 mole of nitrogen gas, N2

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5.

Calculate the volume of gas given the number of moles at room conditions.

6.

The relationships between the number of moles, number of particles, mass and volume of gases:

Activity 5

1. What is the volume of 12.8 g of oxygen gas, O2, in cm3, at STP? [Relative atomic mass: O, 16. Molar volume: 22.4 dm3 mol-1 at STP]

Volume (dm3)

Number of molesMass (g) Number of particles

÷ Molar volume

x Molar volume÷ Molar

mass x NA

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Number of moles Volume of gas

(i) 0.25 mole of oxygen gas

(ii) 2.5 moles of chlorine gas

(iii) 0.4 mole of carbon dioxide gas

(iv) 4 moles of helium gas

Number of moles Volume of gas

(i) 0.3 mole of oxygen gas

(ii) 1.2 moles of ammonia gas

(iii) 1.5 moles of hydrogen gas

(iv) 0.4 mole of nitrogen gas


2. How many molecules of carbon dioxide, CO2 are produced when 120 cm3 of the gas is released during a chemical reaction between an acid and a carbonate at room conditions? [Molar volume: 24 dm3 mol-1 at room conditions. Avogadro constant: 6.02 x 1023 mol-1]

3. What is the mass of 0.6 dm3 of chlorine gas, CI2, at room conditions? [Relative atomic mass: CI, 35.5. Molar volume: 24 dm3 mol-1 at room conditions]

4. A sample of nitrogen gas, N2 has a volume of 1800 cm3 at room conditions. What is the mass of the sample and how many molecules of nitrogen gas, N2 are in it?

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5. 1.12 dm3 of hydrogen gas, H2 and 1.12 dm3 of oxygen gas, O2 are mixed together in a closed container at STP. What is the total number of molecules in the container? What is the total mass of the gases in the container?

E CHEMICAL FORMULAE


state the meaning of chemical formula state the meaning of empirical formula state the meaning of molecular formula determine empirical and molecular formula of substances compare and contrast empirical formula with molecular formula solve numerical problems involving empirical and molecular formulae write ionic formulae of ions construct chemical formulae of ionic compounds state names of chemical compounds using IUPAC nomenclature use symbols and chemical formulae for easy and systematic communication in the field

of chemistry

Activity 6

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________1. A chemical formula is a representation of a chemical substance using letters for ………………………

and subscript numbers to show the ………………….. of each type of atoms that are present in the

substance.

2. The chemical formula

of a compound shows all the ……………………. that are present in the

compound and the ………………………… of atoms of each element.

3. There are 2 types of chemical formulae:

(i) ……………………………………

(ii) …………………………………...

4. Empirical formula of a compound shows the simplest whole number ………………… of atoms of each

………………………… present in the compound.

5. Molecular formula of a compound gives the …………………. number of atoms of each element

present in one molecule of the compound.

6.

Compound Molecular formula Simplest ratio of

atoms of elements

Empirical formula

Water H2O H : O = 2 : 1 H2O

Ethene C2H4 C : H = 1 : 2 CH2

Benzene C6H6 C : H = 1 : 1 CH

Vitamin C C6H8O6 C : H : O = 3 : 4 : 3 C3H4O3

7. The steps in determining the empirical formula of a compound:

(i) ……………………………………………………………………………………..

(ii) ……………………………………………………………………………………..

(iii) …………………………………………………………………………………….

H2The subscript 2 shows that there are ……………….. hydrogen atoms in a molecule of hydrogen gas, H2

The letter ‘H’ shows the symbol of ……………...atom

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8. To determine the molecular formula of a compound, we need to know the ………………………………..

and …………………………. of the compound.

Activity 7

1. 1.08 g of aluminium combines chemically with 0.96 g of oxygen to form an oxide. What is the empirical

formula of the oxide? [Relative atomic mass: O, 16; Al, 27]

Element

Mass of element(g)

Number of moles of

atoms

Ratio of moles

Simplest ratio of

moles

Empirical formula

2. Copper (II) iodide contains 20.13 % copper by mass. Find its empirical formula.

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________ [Relative atomic mass: Cu, 64; I, 127]

Element

Mass of element(g)

Number of moles of

atoms

Ratio of moles

Simplest ratio of

moles

Empirical formula

3. Phosphoric acid has the percentage composition as follows.

What is the empirical formula of the acid?

[Relative atomic mass: H, 1; O, 16; P, 31]

Element

Mass of element(g)

Number of moles of

atoms

Ratio of moles

Simplest ratio of

moles

Empirical formula

H, 3.06 %; P, 31.63 %; O, 65.31 %

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________4. A carbon compound has an empirical formula of CH2 and a relative molecular mass of 70. Find the

molecular formula of the compound.

[Relative atomic mass: H, 1; C, 12]

5. 8.5 g of hydrogen peroxide contains 0.5 g of hydrogen. If the molar mass of hydrogen peroxide is 34

34 g mol-1, find its molecular formula.

[Relative atomic mass: H, 1; O, 16]

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6. Ethanoic acid is an important ingredient of vinegar. The empirical formula of this acid is CH2O. Given

that its molar mass is 60 g mol-1, find its molecular formula.

[Relative atomic mass: H, 1; C, 12; O, 16]

Activity 8

1. Ionic compounds consist of …………………….. and …………………………

2. To construct the chemical formulae of ionic compounds, we need to know the formulae of cations

and anions.

3.

Cation Formula Anion Formula

Sodium ion Fluoride ion

Potassium ion Chloride ion

Silver ion Bromide ion

Hydrogen ion Iodide ion

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Ammonium ion Hydroxide ion

Copper (II) ion Nitrate ion

Calcium ion Ethanoate ion

Magnesium ion Manganate(VII) ion

Aluminium ion Oxide ion

Zinc ion Carbonate ion

Barium ion Sulphate ion

Iron(II) ion Thiosulphate ion

Iron(III) ion Chromate(VI) ion

Lead(II) ion Dichromate(VI) ion

Lead(IV) ion Phosphate ion

Tin(II) ion

Tin(IV) ion

Chromium(III) ion

4. Chemical formula of an ionic compound by exchanging the charges of the ions.

Example:

(i) Construct the chemical formula of iron(II) chloride.

Ion

Charge of ion

Number of ions

Chemical formula

(ii) Construct the chemical formula of aluminium oxide.

Ion

Charge of ion

Number of ions

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Chemical formula

(iii) Construct the chemical formula of zinc sulphate.

Ion

Charge of ion

Number of ions

Simplest ratio of ions

Chemical formula

5. Construct the chemical formula for each of the following ionic compounds:

Ionic compound Chemical compound

(i) Magnesium chloride

(ii) Potassium carbonate

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(iii) Calcium sulphate

(iv) Copper(II) oxide

(v) Silver iodide

Activity 9

1. Chemical compounds are named systematically according to the guidelines given by the International

Union of Pure and Applied Chemistry (IUPAC).

2. For ionic compounds, the name of the cation comes first, followed by the name of the anion.

Cation Anion Name of ionic compound

Sodium ion Chloride ion

Calcium ion Carbonate ion

Barium ion Sulphate ion

3. Certain metals can form more than one type of ……………………… Thus, Roman numerals are 19

FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________ are used in their naming to distinguish the different types of ions. For example, iron can form 2

cations, namely iron(II) ion and iron(III) ion. Thus, the names of the compounds formed by these

ions with chlorine would be ………………………. and ………………………………

4. For simple molecular compounds, the more ………………… element is written last and is added with

an ‘ide’. The name of the ……………….. element is maintained. For example, a molecular compounds

consisting of hydrogen and chlorine is given the name ………………………………….

5. Greek prefixed are used to show the …………………….. of atoms of each element in a compound.

Prefix Meaning Example

Mono- 1 Carbon monoxide

Di- 2 Sulphur dioxide

Tri- 3 Sulphur trioxide

Tetra- 4 Carbon tetrachloride

Penta- 5 Phosphorus pentachloride

F CHEMICAL EQUATIONS

Learning Outcomes

You should be able to:

state the meaning of chemical equation identify the reactants and products of a chemical equation write and balance chemical equations interpret chemical equations quantitatively and qualitatively solve numerical problems using chemical equations identify positive scientific attitudes and values practiced by scientists in doing research justify the need to practise positive scientific attitudes and good values in doing research use chemical equations for easy and systematic communication in the field of chemistry

Activity 10

1. Chemical equation is a precise …………………………. of a chemical reaction.

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________2. The chemical equation can be written in word, but it is more convenient and quicker to use

………………………………

3. The starting substances are called ………………………….. They are shown on the left-hand side of

the equation.

4. The new substances formed are called …………………………….. They are shown on the right-hand

side of the equation.

5. Based in the law of conservation of mass, matter can neither be ……………………. nor

…………………. in a chemical reaction. This means that the number of atoms before and after a

chemical reaction are the …………………… Therefore, a chemical equation must be

…………………….

6. Write a balanced equation for each of the following reactions.

(i) Carbon monoxide gas + oxygen gas carbon dioxide gas ………………………………………………………………………………………………….

(ii) Hydrogen gas + nitrogen gas ammonia gas ………………………………………………………………………………………………….

(iii) Aluminium + iron(III) oxide aluminium oxide + iron …………………………………………………………………………………………………

(iv) Ammonia gas reacts with oxygen gas to yield nitrogen monoxide gas and water. …………………………………………………………………………………………………

(v) Silver nitrate solution is added to calcium chloride solution. Silver chloride precipitate and calcium nitrate solution are produced.

C (s) + O2 (g) CO2 (g) Reactants Product

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FORM 4 CHEMISTRYCHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS__________________________________________________________________________________ …………………………………………………………………………………………………..

(vi) When solid zinc carbonate is heated, it decomposes into zinc oxide powder and carbon dioxide gas. ………………………………………………………………………………………………….

7. Chemical equations give us the following qualitative information: (i) ………………………………………………………............................ (ii) …………………………………………………………………………

8. Quantitatively, the ………………….. in a balanced equation tell us the exact ……………………… of reactants and products in a chemical reaction.

9. ……………………… is the study of quantitative composition of substances involved in the chemical reactions. A balanced equation can be used to calculate ………………………, ………………………, and ………………………….. or ………………………….. of a reactant or product.

Activity 11

1. Copper(II) oxide, CuO reacts with aluminium according to the following equation.

3CuO (s) + 2Al (s) Al2O3 (s) + 3Cu (s)

Calculate the mass of aluminium required to react completely with 12 g of copper(II) oxide, CuO. [Relative atomic mass: O, 16; Al, 27; Cu, 64]

2H2 (g) + O2 (g) 2H2O (l) 2 molecules 1 molecule 2 molecules or or or2 mol 1 mol 2 mol

2Cu(NO3)2 (s) 2CuO (s) + 4NO2 (g) + O2 (g) 2 formula units 2 formula units 4 molecules 1 molecule or or or or 2 mol 2 mol 4 mol 1 mol

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2. A student heats 20 g of calcium carbonate, CaCO3 strongly. It decomposes according to the equation below. ∆ CaCO3 (s) CaO (s) + CO2 (g)

If the carbon dioxide produced is collected at room conditions, what is its volume? [Relative atomic mass: C, 12; O, 16; Ca, 40. Molar volume: 24 dm3 mol-1]

3. Hydrogen peroxide, H2O2 decomposes according to the following equation.

2H2O2 (l) 2H2O (l) + O2 (g)

Calculate the volume of oxygen gas, O2 measured at STP that can be obtained from the decomposition of 34 g of hydrogen peroxide, H2O2. [Relative atomic mass: H, 1; O, 16. Molar volume: 22.4 dm3 mol-1 at STP]

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4. Ethene gas burns in excess oxygen according to the following equation.

C2H4 (g) + 3O2 (g) 2CO2 (g) + 2H2O (l)

Find the volume of carbon dioxide released as STP if 42 g of ethene is burnt completely. [Relative atomic mass: H, 1; C, 12. Molar volume: 22.4 dm3 mol-1 at STP]

5. 16 g of copper(II) oxide, CuO is reacted with excess methane, CH4. Using the equation below, find the mass of copper that is produced. [Relative atomic mass: O,16; Cu, 64]

4CuO (s) + CH4 (g) 4Cu (s) + CO2 (g) + 2H2O (l)

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6. Zn (s) + 2HNO3 (aq) Zn(NO3)2 (aq) + H2 (g) What is the mass of zinc needed to produce 2.4 dm3 of hydrogen gas at room conditions? [Relative atomic mass: Zn, 65. Molar volume: 24 dm3 mol-1 at room conditions]

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3 Chemical Formulae and Equations Edited