Periodic trends

Atomic Radius
§ The electron cloud doesn’t have a definite edge (no sharp boundaries).
§ Atomic radius (r) is measured as half the distance between neighboring nuclei.

Ø Group trends
The atomic radii increase
1. Although the nuclear charge increases ,but
2. Shielding also increases , as electrons are added to a higher energy level, so
3. Decreases the electrostatic attraction between outer electrons added and nucleus

Ø Period trends
The atomic radii decrease
1. Shielding is constant, as electrons are added the same energy level, but
2. Nuclear charge increases
3. Electrostatic attraction between outer electrons added and nucleus increases

Ionic Radius:
Ø Cations: (positive ions) are smaller than their corresponding atoms
§ Cations form by losing electrons.
§ Cations are smaller than the atom they come from, why? Because the cation has less
electrons, less
repulsion and more electrostatic attraction between outer valence electrons and nucleus
(protons )
§ The effective nuclear charge is greater for cations, and the ionic radius is smaller than
the atomic
radius.
§ This effect increases with the charge magnitude, so larger positive charges lead to
smaller radii.

Ø Anions: (negative ions) are larger than their corresponding atoms

§ Anions form by gaining electrons.
§ Anions are bigger than the atom they come from. Why? Because the anion has more
electrons , more
repulsion and less electrostatic attraction between outer valence electrons and nucleus
(protons )
§ The force of attraction between the nucleus and the additional electrons is smaller,
and the ionic
radius is greater than the atomic radius.
§ The greater the magnitude of negative charge, the larger the ionic radius.

Ø Group trends:
The ionic radius increases down a group
1. Atomic radii increase
2. Nuclear charge increases ,but
3. Shielding also increases , as electrons are added to a higher energy level, so
4. Decreases the electrostatic attraction between outer electrons added and nucleus.

Ø Periodic Trends :
The ionic radius decrease from group 1 to 14 for the positive ions
1. They are isoelectronic (Na+ to Si4+) have same number of electrons (10e-)
2. nuclear charge (protons) increases
3. Electrostatic force of attraction of protons on the valence shell electrons increases

The ionic radius decrease from group 14 to 17 for the negative ions
1. They are isoelectronic (S4- to Cl-) have same number of electrons (18e-)
2. nuclear charge (protons) increases
3. Electrostatic force of attraction of protons on the valence shell electrons increases

Ionization energy

First ionization energy: The energy required to remove one mole of electrons from one
mole of gaseous
atoms.
X (g) → X⁺(g) + e-

Ø Group Trend:
Ionization energy decreases down a group
1. although the nuclear charge increases ,but
2. Atomic radii increase
3. Shielding also increases , as electrons are added to a higher energy level, so
4. Decreases the electrostatic attraction between outer electrons added and nucleus

Ø Period Trend:
Ionization energy increase across a period
1. Atomic radii decrease
2. Nuclear charge increases
3. Shielding is constant, as electrons are added the same energy level, so
4. Electrostatic attraction between outer electrons added and nucleus increases

Electron affinity
Electron affinity: it is the energy changed (energy released) when one mole of electron
is added to one
mole of gaseous atoms.
X (g) + e- → X-(g)

Ø Group Trend:
Electron affinity decreases down a group
1. Although the nuclear charge increases ,but
2. Atomic radii increase
3. Shielding also increases , as electrons are added to a higher energy level, so
4. Decreases the electrostatic attraction between outer electrons added and nucleus

Ø Period Trend:
Electron affinity increases across a period
1. Atomic radii decrease
2. Nuclear charge increases
3. Shielding is constant, as electrons are added the same energy level, so
4. Electrostatic attraction between outer electrons added and nucleus increases

Electronegativity: (ability)

It is the ability of an atom to attract a pair of bonded electrons in a covalent bond

(molecule).
§ An element with high electronegativity has a strong electron pulling power while the
element with
low electronegativity has a weak electron pulling power.

Ø Group Trend
Electronegativity decreases a down group
1. although the nuclear charge increases ,but
2. Atomic radii increase
3. Shielding also increases , as electrons are added to a higher energy level, so
4. Decreases the electrostatic attraction between outer electrons added and nucleus

Ø Period Trend
Electronegativity increases across a period
1. Atomic radii decrease
2. Nuclear charge increases
3. Shielding is constant, as electrons are added the same energy level, so
4. Electrostatic attraction between outer electrons added and nucleus increases

Note:
§ The most electronegative element is fluorine and the least is Francium
§ The noble gases has no electronegativity value because they have completely filled
energy levels and
highly stable they don`t attract electrons or form compounds

Melting points:
It depends on both the type of the bond and the structure.

Ø Group trend
 For Group 1 (alkali metals):
 Melting point decreases down group 1.
 Why?
§ They are metals , the type of bond controlling the melting point is the Metallic bond .
§ The metallic bond strength depends on:
(i) the number of valence electrons.
(ii) Charge of cation
(iii) Size of cation

 For Group 17 (halogens)

 Melting point increases down group 17.
 Why?
§ The halogens are non-metals diatomic molecules, so F2, Cl2, Br2, I2
§ They are non polar covalent
§ higher Mr ,means stronger London`s intermolecular attractions (I.A) between
molecules.
§ Higher m.p.

Ø Period trend
 Melting point generally rises across a period. It reaches a maximum at group 14,
and then falls to
reach a minimum at group 18.
 In period 3 for example, bonding changes from metallic bond (Na, Mg and Al) to
giant covalent
(Si) to simple molecules with weak London`s attraction between them(P4, S8, Cl2) and
single
atoms (Ar).
 Explain why sulfur (S8) has higher melting point than phosphorous(P4).
Sulfur exists as a molecule of 8 atoms while phosphorous exists as a molecule of 4
atoms, they are
nonpolar covalent, larger Mr , stronger London` s force of attraction and higher m.p