Topic 3: Periodicity
3.2
Physical properties
Periodicity
refers to the pattern of physical and chemical properties that repeat at regular intervals in the Periodic
Table.
Physical Properties
Factors that affect the trends in physical properties – these factors MUST be used to explain the trend

Nuclear charge
the number of protons in the nucleus (atomic number)
Electrostatic attraction between the nucleus and valence electrons increases with nuclear charge

Period
the energy level of the valence electrons
Electrostatic attractions between the nucleus and valence electrons decrease as the period
increases because the valence electrons are further away from the nucleus.

Electron repulsion electrons in the same energy level repel each other
Electrostatic repulsions between valence electrons cause the space between electrons to
increase as the number of valence electrons increases

Intermolecular force the strength of attractions between two molecules or parts of two molecules
These forces affect the physical state, the melting point and the boiling point
First Ionization Energy the energy required to remove ONE electron from an ATOM in its GASEOUS state
X (g) → X+ (g) + e–
Ca(g) →Ca+(g) + e–
not Ca2+
N(g) → N+(g) + e–
not N2 not N3–
Valence electrons that are held more tightly require more energy to remove.
1A – Alkali metals ( Li → K )
7A – Halogens ( F → Br )
Li > Na > K
F > Cl > Br
decreasing electrostatic attraction
increasing energy level
Period 3 (Na → Cl)
Na < Mg < Al < . . . < S < Cl
increasing electrostatic attraction
increasing nuclear charge
Electronegativity
a measure of the relative attraction of one atom for a PAIR of BONDING electrons
Pairs of bonding valence electrons are more strongly attracted when:
 there is more nuclear charge (more protons)
 the valence electrons are closer to the nucleus (lower energy level)
Fluorine has the highest electronegativity value (4.0)
1A – Alkali metals ( Li → K )
7A – Halogens ( F → Br )
Li > Na > K
F > Cl > Br
decreasing electrostatic attraction
increasing energy level
Period 3 (Na → Cl)
Na < Mg < Al < . . . < S < Cl
increasing electrostatic attraction
increasing nuclear charge
Atomic Radius
the distance from the centre of the nucleus to the outermost (valence) electron
half of the distance between two nuclei of bonded atoms
Valence / bonding electrons are:
 further away as the energy level increases.
 closer as nuclear charge increases
1A – Alkali metals ( Li → K )
7A – Halogens ( F → Br )
Li < Na < K
F < Cl < Br
increased distance
increasing energy level
Period 3 (Na → Cl)
Na > Mg > Al > . . . > S > Cl
decreased distance
increasing nuclear charge
Ionic Radius
Cations (positive ions)
lose all valence electrons → outermost electrons of the ion are at a lower energy level.
Ionic radius of a metal cation is much smaller than the atomic radius of the same metal atom.
 same nuclear charge, different energy levels
1A – Alkali metals ( Li → K )
Li+< Na+ < K+
increased distance
increasing energy level
Period 3 (Na+ → Al3+)
Na+ < Mg2+ < Al3+
decreased distance
same number of valence electrons
(complete)
increasing nuclear charge
Anions (negative ions)
gain valence electrons to have complete energy level (like nearest noble gas)
Ionic radius of a non-metal anion is larger than the atomic radius of the same atom.
 same nuclear charge and energy level
 more valence electrons repel more
7A – Halogens ( F → Br )
F < Cl < Br
increased distance
increasing energy level
Period 3 (P3– → Cl-)
P3- > S2– > Cl–
decreased distance
same number of valence electrons
(complete)
increasing nuclear charge
Metal cations (+) are much smaller than non-metal ions (anions) of the same period.
 lower energy level (lost valence electrons)
Period 3 (Na → Cl)
Na+ < Mg2+ < Al3+ <<
P3- > S2– > Cl–
increased distance
decreased energy level of metals