The chemical elements known as Alkaline Earth Metals (AEMs) occupy the second column, or Group 2, of the modern periodic table. This group includes beryllium, magnesium, calcium, strontium, barium, and radium. These metallic elements share a defining characteristic that dictates their chemical behavior: they possess exactly two valence electrons. This consistent count explains their uniform properties and reactivity across the group.
Understanding Valence Electrons
Valence electrons are the electrons located in the outermost electron shell of an atom. These are the electrons farthest from the nucleus and interact with other atoms during chemical reactions. They determine the atom’s capacity to form chemical bonds and dictate whether an atom will tend to lose, gain, or share electrons to achieve a stable configuration.
The inner electrons, referred to as core electrons, are shielded closer to the nucleus and are not involved in bonding. An atom’s chemical properties are almost entirely dependent on the electrons in that outermost shell. Elements with similar numbers of valence electrons therefore exhibit similar behaviors, which is the foundational principle for the organization of the periodic table.
Why Alkaline Earth Metals Have Two
The reason Alkaline Earth Metals consistently possess two valence electrons is directly linked to their placement in Group 2 of the periodic table. For main group elements, the group number corresponds to the number of electrons in the outermost shell. Thus, the electron configuration for all Group 2 elements ends with two electrons occupying the outermost \(s\)-orbital.
For instance, magnesium has an electron configuration summarized as \([Ne]3s^2\). This notation indicates magnesium has the stable electron arrangement of neon, followed by two valence electrons in the \(3s\) subshell. Similarly, calcium’s configuration is \([Ar]4s^2\), with its two valence electrons residing in the \(4s\) orbital.
This consistent pattern is represented by the general configuration \(ns^2\), where ‘n’ is the principal quantum number. The \(s\)-orbital can hold a maximum of two electrons, meaning the valence shell of Group 2 metals is exactly full. This full \(s\)-orbital contrasts sharply with the Group 1 Alkali Metals, which have only one electron in their \(s\)-orbital.
The Resulting Chemical Identity
The two valence electrons define the chemical identity of the Alkaline Earth Metals, primarily by influencing their path to achieving atomic stability. Atoms tend toward the stability of having a full outer shell, often explained by the octet rule which suggests eight electrons are ideal. Since Group 2 elements only have two valence electrons, it is energetically much easier for them to lose those two electrons than to gain six more.
By losing the two outermost electrons, the metal atom achieves the stable electron configuration of the preceding noble gas. This loss creates a positively charged ion, specifically a cation with a \(+2\) charge, represented as \(M^{2+}\). For example, magnesium metal readily forms the \(Mg^{2+}\) ion.
The drive to shed these two electrons makes Alkaline Earth Metals highly reactive, though they are generally less reactive than the Group 1 metals. Their tendency to form \(+2\) cations means they predominantly participate in ionic bonding, combining with nonmetals to form neutral compounds, such as \(CaCl_2\) or \(MgO\). Calcium’s chemical identity allows it to form strong ionic compounds like calcium phosphate, which provides structural rigidity for bones and teeth in biological systems.