1. Chapter 3 : Chemical Formulae and
Equations
3.1 Relative Atomic Mass & Relative Molecular Mass
3.2 Relationship between the Number of Moles and the Number
of Particles
3.3 Relationship between the Number of Moles of a Substance
and its Mass
3.4 Relationship between the Number of Moles of a Gas and its
Volume
3.5 Chemical Formulae
3.6 Chemical Equations
3.7 Scientific Attitudes and Values in Investigating Matter

2. 3.1 Relative Atomic Mass and Relative
Molecular Mass
 Relative
Atomic Mass (Ar) – the number of
times one atom of the element is heavier
than one twelfth of the mass of a carbon-12
atom

3. Example:
Sodium atom , Na is 23 time heavier than onetwelfth of the mass of one carbon-12 atom.
Thus the relative atomic mass of Na is 23.

Mass of one Na atom (23)
1
× mass of one carbon - 12 atom
12
= 23

4.  Relative
Molecular Mass (Mr) – the number of
one molecule of the compound is heavier
than one-twelfth of the mass of a carbon-12
atom

5. Example:
A methane molecule , CH4 is 16 time heavier
than one-twelfth of the mass of one carbon-12
atom. Thus the relative molecular mass of CH4
is 16.

Mass of one CH 4 molecule (16)
= 16
1
× mass of one carbon - 12 atom
12

6. 3.2 Relationship between the Number
of Moles and the Number of Particles
 Mole
– the amount of substance which
contains the same number of particles
(atoms/ions/molecules) as there are in 12
grams of carbon-12.
 The number of atoms in 12grams of carbon12 is 6.02 x 1023 (Avogadro’s Number or
Avogadro’s Constant (NA)

7. 
Example:
- 1 mol of gold contains 6.02 x 1023 of gold
atoms
- 1 mol of magnesium ions contains 6.02 x
1023 Mg2+ ions
- 1 mol of magnesium chloride (MgCl2)
contains 6.02 x 1023 Mg2+ ions and 2 x 6.02 x
1023 Cl- ions.
- 1 mol of carbon dioxide contains 6.02 x 10 23
CO2 molecules
CO2 is a covalent compounds; chemical bond that involves the
sharing of electron pairs between atoms.

8.  Conversion
of the number of moles to the
numbers of particles and vice versa:-
Number of Particles = Number of mole x NA
Number of Moles = Number of particles ÷ NA
 Example:-
Calculate the number of particles in 0.75 mol
of aluminium atoms,Al.
Solution:
0.75 mol x 6.02 x 1023 Al atoms = 4.52 x 1023 Al
atoms.

9. 3.3 Relationship between the Number
of Moles of a Substance and Its Mass
 Molar
mass – mass of a substance that
contains one mole of the substance
 The molar mass of any substance contains
6.02 x 1023 particles
 The mole atom = relative atomic mass of an
atom but expressed in gram.
Eg: 1 mole atom of Al = 27g
 The mole molecule = relative molecular mass
of a compund expressed in gram.
Eg: 1 mole molecule of water ,H2O = 18g

10.  Conversion
of the number of moles of a
substance to its mass and vice versa:Number of mole-atom = mass in gram ÷
relative atomic mass
Number of mole-molecule = mass in
gram ÷ relative molecular mass
Mass in gram = Number of mole x relative
atomic mass or relative molecular mass

11.  Example:
Calculate the number of moles of 23.5g of
copper (II) nitrate , Cu(NO3)2.
[ RAM: Cu = 64, N=14, O=16]
Solution:
1 mol of Cu(NO3)2 = 64 + 2[14+3(16)] g
= 188 g  RMM
23.5 g of Cu(NO3)2 = 23.5 × 1 mol
188
= 0.125 mol

12.  Example:
Determine the mass for 0.08 mol of ascorbic
acid , C6H8O6.
[RAM: C=12,H=1,O=16]
Solution:
1 mol of C6H8O6
= 6(12) + 8(1)+ 6(16)
0.08 mol of C6H8O6
= 176g
= 0.08 x 176g
=14.08g

13. 3.4 Relationship between the Number
of Moles of a Gas and Its Volume
 One
mole of any gas at room temperature
and 1 atm presure occupies a volume of
24dm3 ( 24000 cm3)
 At standard temperature and pressure,s.t.p
( 0oC and 1 atm), one mole of gas occupies a
volume of 22.4dm3 ( 22400 cm3).
 Molar volume – volume occupied by one
mole of any gas.

14.  Conversion
of the number of moles of a
gas to its volume and vice versa:Number of mole of a gas = volume of gas
÷ molar volume
Volume of gas = Number of mole of a gas
x molar volume

15.  Example:
Calculate the number of moles of
4.8dm3 of chlorine gas at room
temperature.
[1 mol of gas occupies a volume of
24dm3 at room temperature]
Solution:
3
4.8dm
Number of moles=
×1 mol
3
24dm
= 0.2 mol

16.  Example:
Calculate the volume of 0.75 mol of
nitrogen gas at s.t.p.
[ 1 mol of gas occupies a volume of
22.4dm3 at s.t.p]
Solution:
Volume of nitrogen gas = 0.75 mol x
22.4dm3
= 16.8dm3

18. 3.5 Chemical Formulae
 Used
to represent a chemical compound
 It shows:- the elements (denoted by symbols)
- the relative numbers (indicated by subscript
after the symbol)
 Example:-
H2 O

19. Chemical formulae of some covalent
compounds
Name of
compound
Chemical
formula
Number of each
element in the
compound
Oxygen
O2
2 oxygen atoms
Water
H2O
2 hydrogen atoms
1 oxygen atom

20. Chemical formulae of some ions
(cations)
Charge
Cation
Symbol
+1
Sodium ion
Na+
+2
Magnesium ion
Mg 2+
+3
Iron(III) ion
Fe3+

21. Chemical formulae of some ions
(anions)
Charge
Anion
Symbol
-1
Fluoride ion
F-
-2
Oxide ion
O2-
-3
Nitride ion
N3-

22.  To
write the chemical formula of an ionic
compounds:- write the formula of the ions involved in
forming the compound
- balance the positive and negative charge
(use subscript)
- Write the chemical formula of the ionic
compound without the charges.

23. Example:-

24. Formulae of some ionic compounds
Cation
Anion
Chemical
formula
Na+
Cl-
NaCl
Ca2+
Cl-
CaCl2
Al3+
N3-
AlN

25.  Empirical
formulae of a compound:- shows the simplest ratio of the atoms of the
elements that combine to form a compound
- steps to determine the empirical formula of a
compound:1) write the mass / percentage of each
element in the compound
2) calculate the number of moles for each
element
3) Divide each number by the smallest
number to obtain simplest ratio
4) Write the empirical formula of the
compound

26. Example:-

27.  Molecular
formulae of the compound:- shows the actual numbers of the atoms of the
elements that combine to form the compound
Compound
Molecular
formula
Simplest
ratio of the
elements
Empirical
formula
Water
H2O
H:O = 2:1
H2O
Ethene
C2H4
C:H=1:2
CH2
Glucose
C6H12O6
C:H:O=1:2:1
CH2O

28. 3.6 Chemical Equations
 Chemical
reaction can be represented by a
chemical equation
 Reactants – chemicals that are reacting.
Written on LHS.
 Products – chemicals formed in the reaction.
Written on the RHS

29.  Writing
a chemical equation:1) write the correct formulae of all reactants on
the LHS of the equation
2) write the correct formulae of all products on
the RHS of the equation
3) the equation is then balanced.
4) make sure the number of atoms before and
after reaction are the same
5) Write the physical state of each reactants
and products

30. Example:-