How does the reactivity of alkali metals change down the group?

The reactivity of alkali metals increases as you move down the group.

Alkali metals, found in Group 1 of the Periodic Table, are known for their high reactivity. This group includes elements like lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). As you move down the group, the reactivity of these metals increases. This trend is due to the increasing atomic radius and the decreasing ionisation energy of the elements.

The atomic radius of an atom is the distance from the nucleus to the outermost shell where the valence electrons are located. As you move down the Group 1, the atomic radius increases because an extra electron shell is added for each new element. This means that the outermost electrons are further away from the nucleus and are less tightly held by the positive charge of the protons in the nucleus.

Ionisation energy is the energy required to remove an electron from an atom. The further an electron is from the nucleus, the less energy is needed to remove it. Therefore, as the atomic radius increases down the group, the ionisation energy decreases. This makes it easier for the atom to lose its outermost electron and form a positive ion, which is a key aspect of reactivity.

For example, lithium, at the top of the group, has a smaller atomic radius and higher ionisation energy than sodium, which is directly below it. This means that sodium is more reactive than lithium. This trend continues down the group, with francium, at the bottom, being the most reactive of the alkali metals.

In summary, the reactivity of alkali metals increases down the group due to the increasing atomic radius and decreasing ionisation energy of the elements. This results in the outermost electrons being more easily lost, leading to a higher reactivity.

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