- LAWS OF CHEMICAL COMBINATIONS
The
combination of elements to form compounds is governed by the
following five basic laws.
- Law of Conservation of Mass
It
states that matter
can neither be created nor destroyed.
- Law of Definite Proportions
Joseph
Proust stated that a
given
compound always
contains
exactly the same
proportion
of elements by
weight.
It
is sometimes also referred to as Law
of definite composition.
- Law of Multiple Proportions
If
two elements can combine to form more than one compound, the masses
of one element that combine with a fixed mass of the other element,
are in the ratio of small whole numbers.
For
example, hydrogen combines with oxygen to form two compounds, namely,
water and hydrogen peroxide. Here, the masses of oxygen (i.e. 16 g
and 32 g) which combine with a fixed mass of hydrogen (2g) bear a
simple ratio.
- Gay Lussac’s Law of Gaseous Volumes
When gases
combine to form gaseous products their exists a simple ration between
their volumes at the same temperature and pressure.
- Avogadro Law
Equal
volumes of gases at the same temperature and pressure should contain
equal number of molecules.
- DALTON’S ATOMIC THEORY
1.
Matter consists of indivisible atoms.
2. All the
atoms of a given element have identical properties
including
identical mass. Atoms of different elements differ in mass.
3. Compounds
are formed when atoms of different elements combine in a fixed
ratio.
4. Chemical
reactions involve reorganization of atoms. These are neither
created nor destroyed in a chemical reaction.
ATOMIC
AND MOLECULAR MASSES
- Atomic Mass
One
atomic
mass unit is
defined as a mass exactly equal to onetwelfth
the
mass of one carbon - 12 atom.
1
amu = 1.66056×10–24
g
Mass
of an atom of hydrogen = 1.6736×10–24
g
The
mass of hydrogen atom = 1.6736×10–24
g
1.66056×10–24
g
=
1.0080 amu
Today,
‘amu’
has
been replaced by ‘u’
which
is known as unified
mass.
- Molecular Mass
Molecular
mass is the sum of atomic masses of the elements present in a
molecule.
Molecular
mass of methane, (CH4) = (12.011 u) + 4 (1.008 u) = 16.043 u
- MOLE CONCEPT
One
mole is the amount of a substance that contains as many particles or
entities as there are atoms in exactly 12 g (or 0.012 kg) of the 12C
isotope.
1
mol of hydrogen atoms = 6.022×1023
atoms
1
mol of water molecules = 6.022×1023
water molecules
The
mass of one mole of a substance in grams is called its molar mass.
Molar
mass of water = 18.02 g
Molar
mass of sodium chloride = 58.5 g
- PERCENTAGE COMPOSITION
Mass
% of an element = Mass
of that element in the compound
Molar
mass of the compound
- Empirical Formula for Molecular Formula
An
empirical
formula represents
the simplest whole number ratio of various atoms present in a
compound whereas the molecular
formula
shows
the exact number of different types of atoms present in a molecule of
a compound.
- Limiting Reagent
The
reactant which consumes completely in a chemical reaction is called
limiting reagent.
- Reactions in Solutions
The
concentration of a solution or the amount of substance present in its
given volume can be expressed in any of the following ways.
1.
Mass per cent or weight per cent (w/w %)
2.
Mole fraction
3.
Molarity
4.
Molality
1.
Mass per cent
Mass per cent = Mass
of solute
X 100
Mass
of solution
2.
Mole Fraction
It
is the ratio of number of moles of a particular component to the
total number of moles of the
solution.
Mole
fraction of A= No.
of moles of A
No.
of moles of solution
Mole
fraction of B = No.
of moles of B
No. of
moles of solution
3.
Molarity
It
is defined as the number of moles of the solute in 1 litre of the
solution.
Molarity
(M) = No.
of moles of solute
Volume
of solution in litres
4.
Molality
It
is defined as the number of moles of solute present in 1 kg of
solvent.
Molality (m) =
No.
of moles of solute
Mass
of solvent in kg
The
molality
of a solution does not change with temperature since mass remains
unaffected with temperature. Therefore, the molality is preferred
than molarity.
