The tendency of main group elements to achieve a stable electron configuration, typically characterized by having eight valence electrons, drives their chemical reactivity. Main group metals, such as sodium, tend to lose electrons to resemble the noble gas preceding them in the periodic table. For instance, sodium (atomic number 11) aims to lose one electron to match the electron configuration of neon (atomic number 10), resulting in a stable configuration with 10 electrons.
Conversely, main group nonmetals, like sulfur, generally gain electrons to emulate the noble gas that follows them. Sulfur, with an atomic number of 16, seeks to gain two electrons to achieve the electron configuration of argon (atomic number 18). This process leads to fully filled s and p subshells, which enhances stability and reduces chemical reactivity.
For example, consider lithium and fluorine. Lithium has the electron configuration of 1s² 2s¹, with one electron in its outer shell. Fluorine, on the other hand, has the configuration 1s² 2s² 2p⁵, indicating it has seven electrons in its second shell. Lithium will lose its one outer electron, filling its s subshell and achieving a configuration similar to helium. Meanwhile, fluorine gains that electron, becoming negatively charged and filling its p subshell to reach a configuration akin to neon (2p⁶). This electron transfer results in both lithium and fluorine forming ions that mimic the stable configurations of nearby noble gases.
The stability of noble gases, attributed to their filled outer shells, is the ultimate goal for these elements during chemical reactions. By achieving similar electron configurations, main group elements can attain greater stability and lower their reactivity.