Explanation:

* Atomic radius is the distance from the nucleus to the outermost electron shell of an atom.

* Alkali metals have only one valence electron in their outermost shell. This single electron is loosely held and experiences a weaker effective nuclear charge (the net positive charge experienced by an electron), resulting in a larger atomic radius.

* Noble gases have a full outer shell of electrons, which are tightly held by the nucleus. The strong effective nuclear charge attracts the electrons closer to the nucleus, leading to smaller atomic radii.

Example:

In the third period, the atomic radius of sodium (Na) is larger than that of argon (Ar). This is because sodium has one valence electron, while argon has a full outer shell of electrons.

Summary:

The trend of atomic radii across a period is decreasing from left to right. This is due to the increasing effective nuclear charge, which pulls the electrons closer to the nucleus. Therefore, alkali metals have larger atomic radii than noble gases in the same period.