This family of elements includes Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr). They are exceptionally reactive metals, so much so that they are never found as uncombined elements in nature but always as compounds. The defining characteristic of all alkali metals is that they each possess precisely one valence electron in their outermost electron shell. Their strong tendency is to readily lose this single valence electron to achieve a stable electron configuration, resembling the noble gases, which have a full outer shell. This ease of electron loss makes them incredibly reactive. As one descends Group 1, from Lithium down to Francium, the reactivity of the alkali metals progressively increases. This trend can be explained by considering the increasing atomic radius. As more electron shells are added with increasing atomic number down the group, the single valence electron becomes progressively further away from the positively charged nucleus. This increased distance, coupled with the shielding effect of the inner electron shells, weakens the electrostatic attraction between the nucleus and the valence electron. Consequently, less energy is required to remove this electron, making the element more eager to react and form a positive ion. This is why potassium (K) reacts far more vigorously with water, often producing a flame, than sodium (Na), even though both reactions are exothermic.

This group comprises Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). In stark contrast to the alkali metals, the halogens are highly reactive non-metals. All halogens share the common feature of having seven valence electrons in their outermost electron shell. Their strong chemical desire is to gain just one additional electron to complete their outer shell and achieve a stable octet (a configuration of eight valence electrons, like the noble gases). This strong electron affinity drives their high reactivity. As one descends Group 17, from Fluorine down to Astatine, the reactivity of the halogens progressively decreases. This trend, too, can be understood through the lens of atomic structure. As the atomic number increases down the group, the atomic radius increases due to the addition of more electron shells. While the nuclear charge increases, the outermost electron shell, where a new electron would be accepted, is further from the nucleus and shielded by a greater number of inner electrons. This diminished effective nuclear charge experienced by an incoming electron means that the attraction for an additional electron decreases. Therefore, fluorine (F), being the smallest and having the strongest attraction for an incoming electron, is the most reactive halogen, readily accepting an electron to form a negative ion.