Скачать или смотреть Given below are two statements: Statement (I): The metallic radius of Al is less than that of Ga.

Atomic radius, ionic radius, periodic trends, and d-block contraction are the key concepts involved in this question. Understanding the periodic trends in the periodic table is essential to compare the sizes of elements like Aluminium (Al) and Gallium (Ga), both of which belong to group 13. Atomic size generally increases down a group due to the addition of electron shells. However, exceptions exist due to d-block contraction or inert pair effect.

Metallic radius refers to half the distance between the nuclei of two adjacent metal atoms. Aluminium (Al) comes above Gallium (Ga) in the periodic table. Even though Ga is below Al and should ideally have a larger metallic radius, Ga has additional d-electrons, which do not shield the nuclear charge effectively. This leads to a stronger effective nuclear charge experienced by the outer electrons, thereby pulling them closer and reducing the size. This phenomenon is known as d-block contraction, which makes the metallic radius of Ga comparable or sometimes smaller than expected.

Ionic radius is the radius of a charged atom. Al³⁺ and Ga³⁺ are both +3 cations. Due to the higher nuclear charge and smaller number of electrons, Al³⁺ is very small. However, Ga³⁺, having more protons and more inner electrons (including d-electrons), also experiences poor shielding, which enhances the effective nuclear charge and results in a smaller ionic radius. But still, due to its position in the periodic table and d-block contraction, Ga³⁺ ends up having a slightly larger ionic radius than Al³⁺. This trend is crucial in advanced inorganic chemistry, particularly when analyzing trends across transition and post-transition elements.

Question Statement:
Given below are two statements:
Statement (I): The metallic radius of Al is less than that of Ga.
Statement (II): The ionic radius of Al³⁺ is less than that of Ga³⁺.

In the light of the above statements, choose the most appropriate answer from the options given below:
(1) Both Statement I and Statement II are incorrect
(2) Statement I is incorrect but Statement II is correct
(3) Statement I is correct but Statement II is incorrect
(4) Both Statement I and Statement II are correct