Helium contains the same number of electrons in its outer shell as do the alkaline earth metals. Explain why helium is inert whereas the Group \(2 \mathrm{~A}\) metals are not.

When examining the periodic table, helium stands out due to its striking chemical inactivity, often termed 'inertness'. Unlike other elements that readily react and participate in chemical bonding, helium exhibits an aloof character. This behavior is rooted in its unique electron configuration, which is the key to unlocking the secret of its inertness.

The electron configuration for helium is denoted as 1s², indicating that both of its electrons occupy the lowest energy level (the 1s orbital) and that this level is completely filled. This fulfills the 'duet rule', an analogy to the more commonly known 'octet rule', which suggests elements are most stable when they have a full valence shell. For helium, with only one shell and two electrons filling it, this condition is satisfied. In contrast, elements with unfilled outer shells or excess valence electrons tend to be more reactive as they aim to achieve a stable electronic arrangement, either by gaining, losing, or sharing electrons.

Moving on from the tranquility of noble gases, we delve into the energetically ambitious world of the alkaline earth metals. This group of elements is characterized by their propensity to part ways with the two outermost electrons. This behavior can be ascribed to their distinct electron configurations, which house an s² configuration in their outermost shell, similar to helium, but with significant differences that drive their reactivity.

For example, beryllium, with its electron configuration 1s² 2s², appears similar to helium at a glance. However, the presence of available p orbitals in the same energy level as the s subshell leaves room for these metals to aspire towards a more stable, lower-energy state. By shedding their two valence electrons, these metals adopt the electron configuration closer to that of the noble gases — a fully-filled s and p subshells. This electron donation translates into reactivity, positioning these metals starkly apart from the unyielding inertness of helium.

Finally, we approach the noble gases, holding court at the far right of the periodic table. Their 'noble gas configuration' is a coveted electronic state that many elements seek to imitate through chemical reactions. The noble gases, consisting of helium, neon, argon, and their kin, maintain a full complement of electrons in their outermost s and p subshells, expressed as s²p⁶, apart from helium which uniquely has the s² configuration.

The achievement of this configuration is seen as a zenith of electronic stability. Elements such as the alkaline earth metals, with a s² outer shell electron configuration forge ahead to emulate the noble gas status by losing two electrons. In this way, they leave behind a fully-filled shell, reverting to the stable electron configuration of the noble gas that precedes them in the periodic sequence. This drive for stability underpins the reactivity of the alkaline earth metals and the inertness of the noble gases, illustrating the profound influence that electron configuration exerts on chemical behavior.