Valence electrons are the outermost electrons of an atom, occupying the highest energy level shell. Their number is a primary factor in an element’s chemical behavior. The periodic table organizes elements into families based on these shared electron characteristics, such as the Alkaline Earth Metals, whose distinct chemical properties stem from their specific outer shell structure.
Defining Group 2: The Alkaline Earth Metals
The Alkaline Earth Metals are the elements found in the second column of the periodic table, officially known as Group 2. This family includes beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These elements share several physical characteristics, including a shiny, silvery-white appearance when freshly cut. They are relatively soft metals with low densities, melting points, and boiling points compared to many other metals. Although they are quite reactive, they are less so than their neighbors in Group 1, the Alkali Metals. The name “alkaline earth” originates from their oxides, which form basic (alkaline) solutions when dissolved in water and were historically termed “earths.”
Electron Configuration and the Factual Answer
Alkaline Earth Metals have a consistent number of electrons in their outer shell: they each possess exactly two valence electrons. This number is directly correlated with their position in Group 2 of the periodic table. The electron arrangement for any element in this group ends with the general configuration \(ns^2\), where ‘n’ represents the highest occupied principal energy level.
This configuration means the two valence electrons completely fill the outermost \(s\)-orbital. For example, magnesium’s configuration is \([Ne]3s^2\), and calcium’s is \([Ar]4s^2\). This \(ns^2\) arrangement is relatively unstable compared to the full shell of a noble gas. The atom achieves a more stable configuration by losing the two outermost electrons.
How Two Valence Electrons Drive Reactivity
The chemical behavior of Alkaline Earth Metals is governed by their tendency to shed those two valence electrons. They exhibit low ionization energy, which is the energy required to remove an electron from an atom. This low energy requirement allows the atom to readily give up both electrons during a chemical reaction.
Losing these two electrons results in the formation of a cation, or positively charged ion, with a characteristic charge of \(+2\) (\(M^{2+}\)). For instance, magnesium becomes \(Mg^{2+}\) and calcium becomes \(Ca^{2+}\). This ability to easily form a stable \(+2\) ion makes them strong reducing agents and explains their high reactivity with other elements. Their compounds are predominantly ionic, as seen in substances like calcium chloride (\(CaCl_2\)).