Classification of Elements and Periodicity in Prop

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The radius of Na⁺ ion is smaller than that of Na is due to the following reasons:-

(i) The effective nuclear charge of Na⁺

(ii) The disappearance of 3s orbital from its outermost shell electronic configuration.

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Metallic character- Tendency to lose electrons.

Non-Metallic character- Tendency to accept electrons.

Across the period the metallic character decreases whereas non-metallic character increases across the period.

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(a) As the effective nuclear charge increases and shielding effect decreases across the periods thus the electronegativity increases on moving from left to right in the periodic table.

(b) As the number of shells increases down the group thus its ionisation enthalpy decreases down the group.

The deviation in the ionization enthalpy of some elements from the general trend can be explained by the points as given below:-

(i) The fully filled and half filled orbital provide extra stability due to the symmetry

(ii) The effective nuclear charge

(iii) The e^- - e^- repulsion which lead to instability

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(i) S

The ionization enthalpy increases across the period and decreases down the group. N possesses a half filled 2p orbital which provides it extra stability due to symmetry.

(ii)P

The non-metallic character increases across the period

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Exothermic reaction- The reaction in which heat is released

For example- CaO + CO₂ → CaCO₃ + Heat

Endothermic reaction- The reaction in which heat is absorbed.

For example- 2NH₃ + heat → N₂ +3H₂

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The electronic configuration of Na is [Ne]3s¹ while that of Mg is [Ne]3s² as Mg possesses a fully filled 3s orbital thus its first ionisation enthalpy is higher than that of Na.

After losing one electron Na obtained the electronic configuration of Ne while Mg acquired the electronic configuration of Na thus the second ionisation energy of Na is higher than that of Mg.

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In the p block, some elements are metallic some elements are non-metallic while some elements are metalloids in nature. The oxides of metals are basic in nature and that of oxides of nonmetals are acidic in nature.

Acidic oxide SO₂ + H₂O →H₂SO₃

Basic oxide  Tl₂O + 2HCl → 2TlCl +H₂O

Amphoteric oxide

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Al₂O₃ +2 NaOH → 2NaAlO₂ + H₂O

Reaction with water

B₃O₂ +3H₂O → 2H₃BO₃

P₄O₁₁  + 6H2O  → 4H₄PO₃

Cl₂ O₇ + H2O → HClO₄

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The anomalous behaviour of the first member of each group of representative elements i.e. of second period can be attributed to their small size, high charge/radius ratio, high electronegativity and absence of vacant d-orbitals to expand their oxidation state.The first member of each group of p-Block elements displays their greater ability to form multiple bond with itself, e.g. C=C, O=O, N=N and to other second periodic elements, e.g. C=O, C=N, N=N.

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The N possesses a half filled p-orbital which provides it extra stability due to symmetry due to which its electron gain enthalpy is positive and its ionization enthalpy is larger than that of O.