Samacheer Class 11 Chemistry - Periodic Classification of Elements

Two ions having the same number of electrons are called isoelectronic ions.
Example: Na + (1s 2 2s 2 2p 6 ) and F – (1s 2 2s 2 2p 6 ). Both these ions contain eight electrons.

The net nuclear charge experienced by valence electrons in the outermost shell is called the effective nuclear charge.
Z eff = Z – S
Where Z = Atomic number
S = Screening constant calculated by using Slater’s rules.

“Ionisation enthalpy is defined as the energy required to remove the most loosely bound electron from the valence shell of an atom”
The given statement is not correct. Ionization energy is defined as the minimum amount of energy required to remove the most loosely bound electron from the valence shell of the isolated neutral gaseous atom in its ground state.

Magnesium loses electrons successively in the following steps,
Step 1:
Mg (g) + IE 1 → Mg + (g) + 1e, Ionisation energy = I.E 1
Step 2:
Mg + (g) + IE 2 → Mg 2+ (g) + 1e, Ionisation energy = I.E 2
Step 3:
Mg 2+ (g) + IE 3 → Mg 3+ (g) + 1e, Ionisation energy = I.E 3
The total number of electrons is 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 ionization energies, always increase in the following order I.E 1 < I.E 2 < I.E 3. Thus, Step-3 will have the ionization energy.

• The second ionization potential is always higher than the first ionization potential.
• Removal of one electron from the valence orbit of a neutral gaseous atom is easy so first ionization energy is less. But from a uni positive ion, removal of one more electron becomes difficult due to the more forces of attraction between the excess of protons and less number of electrons.
• Due to greater nuclear attraction, second ionization energy is higher than first ionization energy.

Energy of an electron in the ground state of the hydrogen atom is – 2.8 × 10 -8.
The ionization energy of atomic hydrogen is 2.8 × 10 -18 × 6.023 × 10 23 J/mol
= 16.86 × 10 5 J/mol
= 1686 kJ/mol.

* The electronic configuration of an atom affects the value of ionization potential and electron gain enthalpy.
* Half-filled valence shell electronic configuration and completely filled valence shell electronic configuration are more stable than partially filled electronic configuration.
* For e.g. Beryllium (Z = 4) 1s 2 2s 2 (completely filled electronic configuration)
Nitrogen (Z = 7) 1s 2 2s 2 2p x 1 2p y 1 2p z 1 (half-filled electronic configuration) Both beryllium and nitrogen have high ionization energy due to more stable nature.
* In the case of beryllium (1s 2 2s 2 ), nitrogen (1s 2 2s 2 2p 3 ) the addition of extra electrons will disturb their stable electronic configuration and they have almost zero electron affinity.
* Noble gases have stable ns 2 np 6 configuration and the addition of further electrons is unfavorable and they have zero electron affinity.

The fifth period of the periodic table has 18 elements. 5 th period starts from Rb to Xe (18 elements).
5 th period starts with principal quantum number n = 5 and l = 0, 1, 2, 3 and 4.
When n = 5, the number of orbitals = 9.
1 for 5s
5 for 4d
3 for 5p
The total number of orbitals = 9.
Total number of electrons that can be accommodated in 9 orbitals = 9 x 2 = 18.
Hence the number of elements in the 5th period is 18.

The elements ‘a’ [Ne (Z= 10), 1s 2, 2s 2, 2p 6 ] and ‘c’ [Ar (Z = 18), 1s 2, 2s 2, 2p 6, 3s 2, 3p 6 ] have the same valence electronic configuration and hence, belongs to the same group, i.e., group 18 of the periodic table. Similarly, the elements ‘b’ [ Al (Z = 13), 1s 2, 2s 2, 2p 6, 3s 2, 3p 1 ] and ‘d’ [B (Z= 5), 1s 2, 2s 2, 2p 1 ] have the same valence electronic configuration and hence, belongs to the same group, i.e., group 13 of the periodic table.

Halogens are having the general electronic configuration of ns 2, np 5 and readily accept an electron to get the stable noble gas electronic configuration. Therefore, halogens have high electron affinity. Hence, halogens act as oxidizing agents.

• In the 1 st group, lithium forms compounds with more covalent character while the other elements of this group form only ionic compounds.
• In the 2 nd group, beryllium forms compounds with more covalent character while the other elements of this family form only ionic compounds.

Ionic radius is defined as the distance from the centre of the nucleus of the ion upto which it exerts its influence on the electron cloud of the ion. The ionic radius of a uni-univalent crystal can be calculated using Pauling’s method from the interionic distance between the nuclei of the cation and anion. Pauling assumed that ions present in a crystal lattice are perfect spheres, and they are in contact with each other and therefore,
d = r c+ + r A- …………..(1)
where ‘d’ is the distance between the centre of the nucleus of the cation C + and A –. r c+ and r A- are the radius of the cation and anion respectively.
Pauling also assumed that the radius of the ion having noble gas electronic configuration (Na + and Cl – having 1s 2, 2s 2, 2p 6 configuration) is inversely
proportional to the effective nuclear charge felt at the periphery of the ion.
i.e., r c+ ∝ \(\frac{1}{\left(Z_{e f f}\right)^{C+}}\) ………(2)
and r A- ∝ \(\frac{1}{\left(Z_{e f f}\right)^{A-}}\) ………….(3)
where Z eff is the effective nuclear charge.
Z eff = Z – S.
Dividing the equation (2) by (3)
\(\frac{r_{c^{+}}}{r_{A^{-}}}=\frac{\left(Z_{e f f}\right)^{A-}}{\left(Z_{e f f}\right)^{C+}}\)
On solving the equations (1) and (4), the ionic radius of cation and anion are calculated.

(a) The energy required to remove the most loosely held electron from an isolated gaseous atom is called ionization energy.
(b) Variation in a period:
Ionization energy is a periodic property. On moving across a period from left to right, the ionization enthalpy value increases. This is due to the following reasons.
* Increase of nuclear charge in a period
* Decrease of atomic size in a period
Because of these reasons, the valence electrons are held more tightly by the nucleus. Therefore, ionization enthalpy increases.
(c) Variation in a group:
As we move from top to bottom along a group, the ionization enthalpy decreases. This is due to the following reasons.
* A gradual increase in atomic size
* Increase of screening effect on the outermost electrons due to the increase of the number of inner electrons.
Hence, ionization enthalpy is a periodic property.

On moving diagonally across the periodic table, the second and third-period elements show certain similarities. Even though the similarity is not the same as we see in a group, it is quite pronounced in the following pair of elements.
The similarity in properties existing between the diagonally placed elements is called ‘diagonal relationship’.

The 1st ionization enthalpy of magnesium is higher than that of Na due to the higher nuclear charge and slightly smaller atomic radius of Mg than Na. After the loss of the first electron, Na + formed has the electronic configuration of neon (2, 8). The higher stability of the completely filled noble gas configuration leads to a very high second ionization enthalpy for sodium. On the other hand, Mg + formed after losing the first electron still has one more electron in its outermost (3s) orbital. As a result, the second ionization enthalpy of magnesium is much smaller than that of sodium.

d = r K+ + r Cl- = 3.14 Å
\(\frac{r_{K+}}{r_{C l-}}=\frac{\left(Z_{e f f}\right)^{C l-}}{\left(Z_{e f f}\right)^{K+}}\)
(Z eff ) Cl- = Z – S = 17 – 10.9 = 6.1
(Z eff ) K+ = Z – S = 19 – 16.8 = 2.2
\(\frac{r_{K+}}{r_{C l-}}=\frac{6.1}{2.2}\) = 2.77
r K+ = 2.77 r Cl-
2.77 r Cl- + r Cl- = 3.17 Å
3.77 r Cl- = 3.17 Å
r Cl- = 0.83 Å
r K+ = (3.14 – 0.83) Å = 2.31 Å
The ionic radius of the K + ion is 2.31 Å and Cl – ion is 0.83 Å.

Nitrogen with 1s 2, 2s 2, 2p 3 electronic configuration has higher ionization energy than oxygen. Since the half-filled electronic configuration is more stable, it requires higher energy to remove an electron from the 2p orbital of nitrogen. Whereas the removal of one 2p electron from oxygen leads to a stable half-filled configuration. This makes it comparatively easier to remove 2p electron from oxygen.
(ii) First ionisation potential of the C-atom is greater than that of the B atom, whereas the reverse is true is for the second ionisation potential.
The first ionization potential of the C atom and B-atom are as follows:
C(1s 2, 2s 2, 2p 2 ) + IE 1 → C + (1s 2, 2s 2, 2p 1 )
B(1s 2, 2s 2, 2p 1 ) + IE 1 → B + (1s 2, 2s 2 )
The ionization energy usually increases along a period. Hence, the first ionization energy of carbon is greater than that of Boron.
The second ionization potential of the C atom and B atom is as follows:
C + (1s 2, 2s 2, 2p 1 ) + IE 2 → C + (1s 2, 2s 2 )
B + (1s 2, 2s 2 ) + IE 2 → B + (1s 2, 2s 1 )
B + has completely filled 2s orbital which is more stable than the partially filled valence shell electronic configuration of the C + atom. Hence, the second ionization energy of Boron is greater than that of carbon.
(iii) The electron affinity values of Be, Mg, and noble gases are zero, and those of N (0.02 eV) and P (0.80 eV) are very low.
Be, Mg and noble gases have completely filled stable configuration and the addition of further electron is unfavourable and requires energy. The addition of extra electrons will disturb their stable electronic configuration and hence, they have almost zero electron affinity.
Nitrogen and Phosphorus have a half-filled stable configurations and the addition of further electrons is unfavourable and requires energy. The addition of extra electrons will disturb their stable electronic configuration and hence, they have very low zero electron affinity.
(iv) The formation of from F – (g) from F(g) is exothermic while that of O 2- (g) from O(g) is endothermic.
The sizes of oxygen and fluorine atoms are comparatively small and they have high electron density. The extra electron added to fluorine has to accommodate in the 2p orbital which is relatively compact. Hence, the formation of F- from F is exothermic. In the case of oxygen, the formation of O 2- from O is endothermic due to extra stability of the completely filled 2p orbital of O 2- formation.

The repulsive force between inner shell electrons and the valence electrons leads to a decrease in the electrostatic attractive forces acting on the valence electrons by the nucleus. Thus the inner shell electrons act as a shield between the nucleus and the valence electrons. This effect is called the shielding effect (or) screening effect.

Pauling’s scale:
* Electronegativity is the relative tendency of an element present in a covalently bonded molecule to attract the shared pair of electrons towards itself.
* Pauling assigned arbitrary values of electronegativities for hydrogen and fluorine as 2.2 and 4, respectively.
* Based on this the electronegativity values for other elements can be calculated using the following expression.
(X A -X B ) = 0.182 √E AB – (E AA E BB )
Where E AB, E AA, and E BB are the bond dissociation energies of AB, A 2, and B 2 molecules respectively.
X A and X B are electronegativity values of A and B.

Variation of Electronegativity in a period:
The electronegativity generally increases across a period from left to right. As discussed earlier, 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.
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. Noble gases are assigned zero electronegativity. The electronegativity values of the elements of 5-block show the expected decreasing order in a group. Except for 13 th and 14 th groups, all other p-block elements follow the expected decreasing trend in electronegativity.