Use the electron configurations of the alkaline earth metals to explain why they tend to form \(2+\) ions.

At the very heart of understanding the behavior of alkaline earth metals lies the concept of electron configurations. These are like addresses that tell us where the electrons live within an atom. For alkaline earth metals, the electron configurations end in 'ns^2', indicating that the outer shell has exactly two electrons in the s-subshell. Since electrons are arranged in shells and subshells based on increasing energy levels, the outermost electrons—also known as valence electrons—play a pivotal role in chemical bonding.

These two valence electrons are held relatively loosely compared to those in inner shells, which makes them relatively easy to remove. The ease with which these electrons are removed plays directly into the tendency of these metals to form cations—positively charged ions. Understanding this basic principle is crucial for predicting the reactive nature of these elements.

Diving into the specifics, the s-subshell mentioned earlier is one of several types of subshells within an atom. It has the capacity to hold up to two electrons and is uniquely spherical in shape. The significance of the s-subshell in the context of alkaline earth metals cannot be overstated. It's the outermost subshell for these elements and having it completely filled results in a higher level of stability.

However, while this filled s-subshell denotes stability, it also makes these metals highly reactive. This is because the fully occupied s-subshell can become completely empty by shedding its pair of electrons, leading to what is known as a 'fully-stable' or 'noble gas' electron configuration, which is a highly favored state for an atom.

The concept of ionic charge refers to the electric charge an atom holds when it gains or loses electrons to form an ion. For alkaline earth metals, when they lose their two outer s-subshell electrons, they form 2+ ions. This positive charge is a direct result of the atom having more protons than electrons after the loss of the two valence electrons.

To put it simply, these metals prefer to give up their two loosely-held valence electrons rather than attempt to gain more to fill up the next shell. The resulting ionic charge is integral to the chemical behavior of these metals, dictating how they bond with other elements and the types of compounds they form.

Chemical stability is akin to an atom’s desire to be in a low-energy, settled state—much like a person relaxing after a long day. Alkaline earth metals reach a state of chemical stability when they form 2+ ions, emulating the electron configuration of noble gases which have full outer shells and are known for their low reactivity.

The journey to stability involves the metals losing their two valence electrons, which, while creating a positive ion, also leads to a noble gas configuration that is energetically favorable. In essence, the more stable an atom or ion, the less reactive it is. For this group of metals, achieving the greatest stability means shedding those two outer electrons and taking on a 2+ charge, which is why you’ll find Be2+, Mg2+, and so on in chemical compounds.