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REVISION
All matter is composed of tiny invisible particles called
atoms.
The many different kinds of atoms are referred to as
different elements. Different kinds of matter are different
because they contain different combinations of elements.
Physical change involves a change in the state or condition
of matter but not a change in its chemical identity:
Example: Salt, NaCl is chemically unchanged - i.e. it
remains NaCl - when the solid is melted to form a liquid or
if the solid is dissolved in water.
Chemical change, i.e. a chemical reaction involves the
conversion of one type of matter into another:
Example:
Na + Cl  NaCl
The total mass remains constant during a chemical reaction
(Law of Conservation of Mass).
A pure compound always contains constant proportions of
the elements by mass (Law of Constant Composition)
The Structure of Chemical Compounds
Atoms in chemical compounds, i.e. in compound matter,
are not just bundled together like a collection of billiard
balls, but are held together by forces called chemical
bonding.
Two types of chemical compound can be distinguished:
Molecular and Non-molecular. Molecular substances are
defined as substances that contain molecules.
A molecule is a distinct finite number of atoms held
together (bonded) by chemical forces and capable of
behaving as a single unit.
A
molecule
has
distinct
physical
and
chemical
properties that are characteristic of the molecular unit
as a whole and which are different from those of the
constituent atoms. If the molecule is broken down into its
constituent atoms its characteristic properties are lost. The
molecule is therefore the smallest unit of a compound
that retains the chemical properties of the substance as
a whole.
Examples of molecular substances:
H2
Hydrogen molecule
O2
Oxygen molecule
CO2
Carbon dioxide molecule
SO3
Sulphur trioxide molecule
N2O4
Dinitrogen Tetroxide molecule
C76H52O46
Tannic acid molecule
Molecular formula: The composition of a molecule is
given in the molecular formula by the symbol for each of
the constituent elements followed by the number of atoms
of that element involved as a subscript. For instance, the
formula of tannic acid above specifies that the molecule
contains 76 carbon atoms, 52 hydrogen atoms and 46
oxygen atoms.
By convention, in molecules containing carbon and
hydrogen together with other elements, the formula is
written with carbon and hydrogen first followed by the
other elements in alphabetical order. Examples: tannic
acid (see above) or nitrobenzene, C6H5NO2.
There are three kinds of non-molecular substances:
1. Mono-atomic substances: e.g. He, Ne, Kr, Ar and Xe
(known as ‘Noble Gas’ elements).
2. Giant molecules: e.g. silicon carbide i.e. SiC or (SiC)n.
In this compound there are equal numbers of silicon and
carbon atoms. Each silicon atom is surrounded by and
bonded to four carbon atoms, while each carbon is
similarly
attached
to
four
silicon atoms. The size of the
C
particle of solid silicon carbide
Si
Si
C
is the only limitation on the
Si
C
C
number of atoms in the network
– which is so large as, for all
intents and purposes, to be
Si
effectively infinite.
3. Ionic solids: In ionic solids the component atoms (or
group of atoms) bear positive or negative electrical
charges and are called IONS. The ions are held together
by electrostatic attraction between the opposite charges.
Since an ionic compound must be neutral overall there
must be equal numbers of charges on the positive ions
(cations) as there are on the negative ions (anions) and this
determines the formula for that compound.
Example: Sodium chloride (NaCl, 'table salt') is composed
of sodium cations (Na+) and chloride anions (Cl-). For the
charges to balance there must be equal numbers of cations
and anions so that the formula must be NaCl. However
there is no distinct, finite NaCl molecule – in the solid no
single Na+ cation can be considered to be associated with
just one Cl- anion - each cation is surrounded by six
identical anions and each anion by six identical cations.
Atomic Mass: The atoms of the ca. 118 known elements
each have a characteristic mass – the atomic mass (aka
atomic weight). The masses of atoms are conveniently
expressed in atomic mass units (abbreviated ‘amu’). The
1
amu is defined as exactly 12 of the mass of the most
common form of the carbon atom, i.e. 126C.
On this scale hydrogen, H, has a mass of 1 amu and the
heaviest presently known element has a mass of around
290 amu.
Listings of atomic masses are found in the inside cover of
most First Year chemistry textbooks.
The given atomic masses of the naturally-ocurring
elements are actually average figures. This is because
naturally occurring elements generally exist as a mixture of
so-called isotopes.
Isotopes are atoms with identical chemical properties but
different masses which are present in nature in different
relative ratios depending on the element involved.
Example: Magnesium, Mg.
Naturally-occurring magnesium contains three different
isotopes:
Atomic mass
23.98 amu
78.70%
24.98
10.13%
25.98
11.17%
Consider 10000 magnesium atoms:
7870 will have a mass of 23.98 amu = 188722.6 amu
1013 will have a mass of 24.98 amu = 25304.7 amu
1117 will have a mass of 25.98 amu = 29019.6 amu
10000 atoms have a total mass of
243046.9 amu
Therefore the weighted average mass of a single
magnesium atom will be 24.304 amu.
Molecular Mass (aka Molecular Weight): The mass of a
molecule is simply the sum of the atomic masses of the
constituent atoms:
H2:
2 x 1.008 = 2.016 amu
CO2: 12.011 + (2 x 15.99) = 43.99 amu
C6H12O6: (6 x 12.011) + (12 x 1.008) + (6 x 15.999) =
180.00 amu
For non-molecular substances a formula mass (aka
formula weight) is determined from the empirical formula
(i.e. the smallest whole-number ratio of atoms involved):
Silicon carbide, SiC: 28.086 + 12.011 = 40.10 amu
Calcium Chloride, CaCl2: 40.08 + (2 x 35.453) = 110.99
amu.
The Mole Concept:
So far, atomic mass units (amu) have been used for atomic
and molecular masses, i.e. the masses of 1 atom, 1
molecule or 1 formula unit. These mass values, however,
have practical limitations, because chemical reactions are
not normally carried out on individual atoms or molecules.
In practice, chemical reactions are carried out on many
billions of atoms or molecules at a time. For this reason a
more practically useful unit, the mole, abbreviated as mol
and defined as 6.02 x 1023 (known as Avogadro's
Number, symbol N) of atoms, molecules or formula units,
was introduced.
Why did Avogadro choose the number 6.02 x 1023?
Because 6.02 x 1023 is the number of atomic mass units
(amu) in 1g. Therefore the mass in grams of a mole of
atoms, molecules or formula units is numerically the same
as the corresponding atomic, molecular or formula unit
masses in amu.
Example:
Mass of one Sodium (Na) atom = 22.99 atomic mass units
Mass of one mole of Na atoms = 22.99 x 6.02 x 1023 amu
Converting the mass of a mole of Na atoms in amu to mass
in grams using the conversion factor 6.02 x 1023 amu = 1g:
22.99 x 6.02 x 10 23 amu
6.02 x 10 23 amu g -1
= 22.99 grams
Example:
HCl, hydrogen chloride:
Molecular mass = 1.008 + 35.453 = 36.45 amu
Mass of 1 mol of HCl molecules = 36.45 g
When using moles in chemical calculations it is
essential to be clear about the precise chemical species
that you are dealing with.
Example:
Cl, chlorine atom:
Mass of one chlorine atom = 35.45 amu
Mass of one mol of Cl atoms = 35.45 g
Cl2, chlorine molecule:
Mass of one chlorine molecule = 70.9 amu
Mass of one mol of Cl2 molecules = 70.9 g
REVISION
1. One mole of any chemical species (atom, molecule,
formula unit) contains Avogadro's Number, i.e. 6.02 x
1023, of that species.
2. The mass in atomic mass units (amu) of any individual
chemical species (i.e. atom, molecule or formula unit) is
numerically equal to the mass in grams of one mole of
that chemical species (i.e. atom, molecule or formula
unit).
3. In working with moles in chemical calculations it is
essential to specify precisely the chemical species that
you are dealing with.
4. The mass of a mole of atoms in grams is called the gram
atomic weight, the mass of a mole of molecules in
grams is called the gram molecular weight and the
mass of a mole of formula units is called the gram
formula weight. These are commonly - though less
precisely - referred to simply as atomic weight,
molecular weight and formula weight.