Periodic table :
periodicity

The Modern Periodic Table
• The periodic table is an arrangement of the
chemical elements, organized on the basis of
their atomic numbers, electron configurations
and recurring chemical properties.
• There are 92 naturally occurring elements.
• The Modern Periodic Table is made up of 18
groups and 7 period.

The modern periodic table

The physical properties of
elements of period 2 and 3

Atomic radius (atomic radii)
• The size of an atom cannot be measured
exactly, However we can measure the size of
atom in terms of its atomic radius.
• The atomic radius is half the distance between
the nuclei of two closest and identical atoms

• We can classify atomic radius into three
a) Covalent radius
b) Metallic radius
c) Van der Waals radius

Covalent radius
• Definition : half the length from one nuclei to
another atom’s nuclei bonded covalently.

Metallic radius
• Definition : half of the distance between the
two adjacent metal atoms in the metallic
lattice.

NOTE
#For metallic elements it may refer to covalent
radius or metallic radius.
>Metallic used for non metal
>Covalent for metals

• For metallic elements which consists of
covalent radius or metallic radius they have
generally smaller atomic radius .
• Covalent bonds and lattice are very strong
bonds that pull the shells closer together,
causing the radius to decrease.

Van der Waals radius
• Definition : half the distance between two
neighbouring atoms which are not chemically
bonded in solid state.
• Appear like touching, less
attractive force.

Factors affecting Atomic radius
• There are factors that affect the atomic radius
(size of atom ) which causes atoms having
different sized across the group and down the
period.
a) Screening effect
b) Nuclear charge

Screening effect
• The decrease in attraction between an
electron and the nucleus of an atom with
more then one electron.
• Its caused by mutual repulsion between
electrons in the inner shell with those at the
outer shell.
• This repulsion (screening ) causes the size of
atom to increase.

+
Mutual repulsion

• The increase number of electrons the higher
the repulsion force.
• The inner electrons shield the outer electrons
from the nucleus pull
• greater the screening effect, easier the
removal of electron.

• However if the inner electronic shells have
electrons , the attraction forces between
nucleus and outermost electron will not be
strong.
• The outermost electron is shield from the
nucleus by the inner electrons.

11+

• Eg : magnesium atom has larger atomic radius
then berylium. Although Mg has a larger
nuclear charge it has more occupied shell then
berylium.
• The more the number of electrons, less the
attraction force , smaller the size
(refer periodic table )
• The lower in the period the more screening
effect the larger the atomic radius.

Nuclear charge
• The nucleus charge is the total charge of all the
protons in the nucleus.
• It has the same value as the atomic number.
• The nuclear charge increases as you go across the
periodic table.
• The nuclear charge pulls the electrons closer to
the nucleus and causes the atomic size to
decrease.
• The stronger the nucleus charge
(atomic number ) the smaller the atomic radius

• Eg : across period 2 carbon has atomic
number of 6 and nitrogen 7. this means
nitrogen has higher nuclear charge hence
smaller atomic radius.
• The effective nuclear charge : the difference
between the screening constant and the
actual nuclear charge.
• The higher the effective nuclear charge the
smaller the atomic radius

Comparing =)
Screening effect
• Down the group the size
increases
• The number of shells with
occupied electrons increase
• The attraction force
decreases
• The size increases
Nuclear charge
• Across the period the size
decreases
• The proton number
increases across the period
• The attraction force
increases
• The size decreases

Atomic radii (radius ) across period 2
and 3
Across the period 2 and 3 (from left to right )
there is an decreases in atomic size.
This is due to the increase in nuclear charge
across these periods
Hence increasing its electrostatic pull between
electrons and nucleus, resulting in decrease in
atomic size.

• The screening effect will remain almost
unchanged as the electrons across the same
period are added to the same quantum shell
which are 2s (period 2 ) and 2p( period 3).
• This will cause the effective nuclear charge to
increase
• Due to the outermost electrons being pulled
closer to the nucleus hence, decreasing its
atomic radius.

Atomic radii down a group.
• Going down a group the atomic radius will
increase.
• This is because the increase in proton number
that results in increase number of shells. This
causes the attraction forces between nucleus
and outermost electrons to decrease.
• Increase in screening effect ( repulsion )
• This causes the atomic size to increase.

Ionic radii
• The radius of a atom’s ion (cation or anion)
• Ionic radii of the cation and anion gives the
distance between the ions in a crystal lattice.
• The higher the nuclear charge, the higher the
forces of attraction and hence larger ionic
radius

Ionic radii across period 2 and 3
(The size of cations and anions decrease)
Cations decrease with increasing proton number
a) The increase in proton number increases the
attraction force between nucleus and electron
hence smaller atomic size to decrease.
b) In period 2,all cations have valence of 2 but as
the proton number increases as for Li+ (3) and
(5) for B3+ so does the strength.
c) The attraction between B3+ is stronger then
Li+

The anion size decreases with increasing proton
number
a) The proton number increasing causes the
attraction between the nucleus and electrons to
increase hence the atomic size decrease.
b) All anions in period 2 has 8 electrons but the
nuclear charge increases across the period
causes the attraction to increase
• Eg : N3+ (7) and F-(9)

The ionic radius down the group
• Going down the group the size of ions
increases
• The screening effect increases with the
addition of extra shell.
• The more the number of shells the less the
attraction force ( increase in repulsion ) hence
the atomic size increases

• Neutral atoms or ions of same number of
electrons are said to be isoelectronic.
• Positive ions (cations) are smaller than its
neutral atom.
• Na+ ion is smaller than Na atom. This is
because Na+ is more stable by donating one
electron hence the attraction force is stronger
• Negative ions (anions) are larger than its
neutral atoms .
• Cl- is larger than Cl atom as Cl- is more stable
by accepting one electron making its
attraction force between nucleus weaker

Melting/boiling point and the enthalpy
of vaporisation.
• Melting point is the temperature when a solid
changes into liquid
• Boiling point is the temperature when the
vapor pressure of the liquid is equal of that
atmospheric pressure.
• Enthalpy of vaporisation is the heat energy
required to covert 1mole of liquid into vapour
at its boiling point.Also it measures the
strength of the intermolecular forces between
particles in its liquid state.

• The melting points depend on two factors:
a) The bonds involved (ionic, covalent , metallic
or Van der Waals )
b) The structure/particle arrangements ( giant
covalent , simple molecular or metallic s)

CHEMICAL BONDING
• For elements with strong bonding like ionic
solids or giant covalent molecules their
melting and boiling points are very high
• Because their molecules are held by strong
attraction forces that makes it harder to break
• More energy is needed to break the bonds
hence making the melting and boiling point
higher

MOLECULAR STRUCTURE
• A covalently bonded molecule between
atoms in giant crystal lattice.
• Metallic bonds in metal lattice, the positively
charges metal ions are attracted to the could
of electron. As the valence electron increase
the attraction force increases
• Simple molecular structure are for non
metallic elements that form simple structure.
•

Based on the graph,
• sodium, magnesium and aluminium are giant
molecules with metallic bonds
• Silicon has giant covalent molecule.
• Phosphorus , sulphur , chlorine and argon are
simple covalent molecule.

• Hence the boiling point increases gradually for Na,
Mg ,Al and Si as they have a strong
covalent/metallic bond (intermolecular forces )that
requires more energy to break the bonds causing a
steep increase in boiling/melting point
• As for the simple molecules P,S,Cl they require less
energy hence the boiling/melting point experience
a decrease.
• The strength of metallic/covalent bond increases
with the increase in valence electron.

• For the enthalpy of vaporisation it is directly
proportional to the melting/boiling.
• As the boiling/melting increases the enthaply
of vaporisation also increases as more heat is
needed to convert 1mole of liquid into vapour.

Electrical Conductivity

• Metals are good conductors of electricity .This
is due to the presence of mobile electrons
• Non-metals do not conduct electricity.
• As the number of delocalised electrons
increase the electric conductivity increases.

• Li and Be are metals that conduct good
electricity in solid and molten state due to
their delocalised electrons that move freely
across the metal.

Electronegativity

• Is the ability of an atom in a covalent bond to
attract shared electrons to itself
• Tendency to attract electrons.
• The greater its electronegativity the greater its
tendency to attract electron.
• Is effected by the
>atomic number(nuclear charge) or
>distance of valence electron from
nucleus(atomic radius)

The electronegativity increases across
the period
• As the atomic radius decreases across the
period its electronegativity increases across
the period
• Its attraction forces between nucleus and
electron will increase hence the tendency to
attract electrons are higher

Electronegativity decreases down the
group.
• As going down the group the nuclear charge
and screening effect increases but, the
screening effect has a bigger increase than the
nuclear charge
• Hence the atomic radius increases
• Resulting in an decrease in effective nuclear
charge causing the attraction between nucleus
and electrons to decrease.
• Causing the tendency to attract electrons to
decrease.