's' block: ns¹ or².
'p' block: ns² np^1-6
'd' block: (n - l)d^1-10 ns^{1 or 2}
'I' block: (n - 2}f^1-14 (n- l)d^Oor2 ns²

Lanthanides: [₅₄Xe] 4f^1-14 5d¹ 6s²
Actinides : [₈₆Rn] 5f^0-146d^0-2 7s²

The total number of electrons are less in the cation than the neutral atom while the nuclear charge remains the same. Therefore the effective nuclear charge of the cation is higher than the corresponding neutral atom. Thus the successive ionisation energies, always increase in the following order
IE₁ < IE₂ < IE₃ < .....

It is defined as the relative tendency of an element present in a covalently bonded molecule, to attract the shared pair of electrons towards itself.

(i) The number of electrons increases by the same number as the increase in the atomic number.
(ii) As the number of electrons increases, the electronic structure of the atom changes.
(iii) Electrons in the outermost shell of an atom (valence electrons) determine the chemical properties of the elements.

Newlands law of Octaves:
When elements are arranged in increasing order of their atomic weights every eighth element resembles its properties with the first one just like the eight note of a musical table.

Limitations: Failed for heavier elements beyond calcium.

(i) According to aufbau's principle 5th period has nine orbital (one 5s, five 4d and three 6p) to be filled.
(ii) Nine orbitals can accommodate a maximum of 18 electrons. Hence fifth period of the periodic table should has 18 elements from rubidium (2 = 37) to Xenon (Z = 54).

Z = 118; [₈₆Rn] 5f¹⁴ 6d¹⁰ 7s² 7p⁶
In the periodic table the element with Z = 118 is located in p-block,
Period no = 7 (as n = 7 for valence shell)
Group no. = 18 (group no = 10+ ns electrons + np electrons) (n - outer most shell)

Characteristics of f-block elements:
(i) The elements in which the last electron enters in f-orbital are called f-block elements.
(ii) This block consists of 4f block elements and 5f block elements.
Lanthanoid series (4f) ₅₈Ce^- ₇₁Lu
Actinoid series (5f) ₉₀Th^- ₁₀₃Lr
(iii) Their general electronic configuration is (n - 2) f^1-14 (n - 1)d^0-1 ns².
Lanthanoid - 4f^1-14 5d^0-1 6s²
Actinoid - 5f^1-14 6d^0-1 7s²

(i) Variation of Electronegativity in a period: The electronegativity generally increases across a period from left to right. The atomic radius decreases in a period, as the attraction between the valence electron and the nucleus increases. Hence the tendency to attract shared pair of electrons increases. Therefore, electronegativity also increases in a period.
(ii) Variation of Electronegativity in a group: The electronegativity generally decreases down a group. As we move down a group the atomic radius increases and the nuclear attractive force on the valence electron decreases. Hence, the electronegativity decreases.

The electronic configuration of Sodium (Z = 11) Is²2s²2p⁶3s¹.
Magnesium (Z = 12) 1s²2s²2p⁶3s²
Magnesium atom has a smaller radius and higher nuclear charge than a sodium atom, thus more energy will be required to remove the electron from the same orbital (3s), making the first ionisation energy of magnesium higher than that of sodium.
However, the second ionization enthalpy of sodium is higher than that of magnesium. This is because after losing 1 electron, sodium attains the stable noble gas configuration of neon (1s²2s²2p⁶). On the other hand, magnesium, after losing 1 electron still has one electron in the 3s-orbital(1s²2s²2p⁶3s¹). In order to attain the stable noble gas configuration, Thus, the energy required to remove the second electron in case of sodium is much higher than that required in case of magnesium. Hence, the second ionization enthalpy of sodium is higher than that of magnesium.

On moving diagonally across the periodic table, the second and third period elements show certain similarities. It is quite pronounced in the following pair of elements.

The similarity in properties existing between the diagonally placed elements is called diagonal relationship.