Alkaline earth metals and alkali metals differ in their electronic configurations, which determine the number of valence electrons they have.
Alkali metals, such as lithium (Li), sodium (Na), and potassium (K), belong to Group 1 of the periodic table. They have a single electron in their outermost shell (valence shell), which makes them highly reactive. These elements tend to lose this single valence electron to achieve a stable electron configuration with a full lower shell. By losing this electron, they attain a noble gas-like configuration and become positively charged ions (cations).
On the other hand, alkaline earth metals, such as beryllium (Be), magnesium (Mg), and calcium (Ca), belong to Group 2 of the periodic table. They have two valence electrons in their outermost shell. These two electrons are more strongly attracted to the nucleus compared to the single valence electron of alkali metals. As a result, alkaline earth metals have higher ionization energies (energy required to remove an electron) than alkali metals. Though alkaline earth metals also tend to lose their valence electrons to achieve a stable configuration, it requires more energy to remove two electrons.
The difference in the number of valence electrons between alkali metals and alkaline earth metals arises due to their position in the periodic table and the organization of electron shells and subshells. The periodic table is structured in a way that elements in the same group tend to have similar chemical properties because they have the same number of valence electrons.