MCQ Detailed View

In the periodic table, elements are arranged in groups (columns) that share similar chemical properties. One important property that varies within a group is the atomic radius, which is the distance from the nucleus to the outermost electron.

The atomic radius increases down a group because each successive element has an additional electronic shell. As more electron shells are added, the outer electrons are farther from the nucleus. Although the nuclear charge (number of protons) also increases, the shielding effect of inner electrons reduces the effective nuclear attraction on the outermost electrons, allowing the atom to expand.

Other options are incorrect. There are no “proton shells” or “neutron shells”; the nucleus contains protons and neutrons but does not have shells. The nucleus itself does not directly cause the increase in atomic size. The key factor is the addition of electronic shells, which determines the physical size of atoms in a group.

Understanding the trend in atomic radii is important in inorganic chemistry for predicting chemical reactivity, bond formation, and periodic behavior of elements. For example, larger atomic radii in group 1 metals explain their high reactivity with water, as the outer electron is farther from the nucleus and more easily removed.

The trend of increasing atomic radius down a group is a fundamental concept in atomic structure and periodic properties. It is closely related to ionization energy, electronegativity, and metallic character, which all influence chemical behavior.

Thus, in a group of the periodic table, atomic radii increase due to the successive addition of electronic shells, making this a key principle in understanding element trends and periodicity.