The atomic structure of a chemical element dictates its behavior and how it interacts with the world. Understanding this structure allows scientists to predict an element’s properties and categorize it within the periodic table. Alkaline Earth Metals (AEMs) are a family of elements whose atoms share a defining structural feature that governs their distinct chemical properties. This group forms the second column on the periodic table and exhibits a predictable pattern of organization and reactivity.
Identifying Alkaline Earth Metals
The Alkaline Earth Metals are the six elements located in Group 2 of the periodic table: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). Elements are arranged in groups because their atoms share similar outer electron configurations, which leads to similar chemical characteristics. Moving down this column, the atoms increase in size and mass, but their fundamental nature remains consistent.
These elements are all metals that share physical properties, such as a shiny, silvery-white appearance and a relatively low density. They are also somewhat reactive, meaning they do not exist freely in nature but are found bound in various minerals and ores. Calcium and Magnesium are widely known examples, being significant components in bone structure and lightweight alloys, respectively.
The Defining Atomic Structure
The common thread linking all Alkaline Earth Metals lies in the specific arrangement of electrons in their outermost shell, known as the valence shell. Every atom in this group possesses exactly two valence electrons. These two electrons reside in the outermost \(s\) subshell, represented by the general electron configuration \(ns^2\), where ‘n’ denotes the principal energy level.
The consistency of having two valence electrons is a direct consequence of their Group 2 placement. Although the number of total electron shells increases as you descend the group, the number of electrons in the outermost shell does not change. For example, the valence configuration progresses from \(2s^2\) for Beryllium to \(6s^2\) for Barium.
This specific \(ns^2\) structure means that the outermost \(s\) orbital is completely filled. The two valence electrons are held less tightly than the electrons in the inner, filled shells, making them the primary participants in chemical interaction. The energy needed to remove these two electrons is relatively low, which explains the atoms’ chemical activity.
How the Atom Achieves Stability
The atomic structure of two valence electrons dictates the chemical behavior of Alkaline Earth Metals, as every atom strives to achieve maximum energetic stability. Stability is reached by mimicking the electron structure of the nearest noble gas, which requires a full outermost shell. It is energetically favorable for the AEM atom to lose its two valence electrons rather than attempt to gain six more.
The mechanism for achieving stability involves the atom readily transferring its two outermost electrons during a chemical reaction. Upon losing these two negatively charged particles, the atom forms a positively charged ion, or cation. This cation always carries a charge of \(+2\), a characteristic feature that defines the chemical reactivity of the entire group.
This eagerness to shed electrons makes Alkaline Earth Metals highly reactive, although they are less reactive than the Group 1 Alkali Metals. Their reactivity is demonstrated by their tendency to react with oxygen to form metal oxides and with water to form alkaline solutions. The resulting \(+2\) charged ion, such as \(\text{Ca}^{2+}\) or \(\text{Mg}^{2+}\), is structurally stable and is the form in which these elements are most commonly found in compounds.