The reactivity of alkali metals increases as one moves down Group 1 of the Periodic Table.

Alkali metals, which are located in Group 1, include elements such as lithium ($\text{Li}$), sodium ($\text{Na}$), potassium ($\text{K}$), rubidium ($\text{Rb}$), cesium ($\text{Cs}$), and francium ($\text{Fr}$). A notable characteristic of these metals is their high reactivity, which tends to increase as you descend the group. This trend can be attributed to two primary factors: the increasing atomic radius and the decreasing ionization energy of the elements.

The atomic radius refers to the distance from the nucleus of an atom to the outermost shell containing the valence electrons. As we progress down Group 1, the atomic radius grows larger because each successive element has an additional electron shell. Consequently, the outermost electrons are situated further from the nucleus and are held less tightly by the positive charge of the protons.

Ionization energy is defined as the energy required to remove an electron from an atom. As an electron’s distance from the nucleus increases, the energy needed to detach it diminishes. Thus, as the atomic radius enlarges down the group, the ionization energy decreases. This reduction in ionization energy facilitates the loss of the outermost electron, which is essential for the reactivity of these metals.

For instance, lithium ($\text{Li}$), positioned at the top of the group, possesses a smaller atomic radius and a higher ionization energy compared to sodium ($\text{Na}$), which is located directly below it. As a result, sodium is more reactive than lithium. This pattern continues down the group, culminating with francium ($\text{Fr}$), which is the most reactive alkali metal.

In conclusion, the increasing reactivity of alkali metals down Group 1 is primarily due to the growing atomic radius and the declining ionization energy of the elements. These changes result in the outermost electrons being more readily lost, thereby enhancing the overall reactivity of the metals.