Final answer:
Silicon (Si) can achieve stability by forming four covalent bonds, thus completing its valence shell of eight electrons following the octet rule. This typically results in a tetrahedral geometry and sp³ hybridization, allowing silicon to act as a Lewis acid in certain chemical contexts.
Step-by-step explanation:
In order for silicon (Si) to achieve a full outer energy level and become stable, it needs to either gain or share electrons to complete its valence shell. Silicon has the electron configuration 3s²3p², meaning it has four valence electrons in its third shell. Since the third shell can hold up to eight electrons, silicon often forms bonds to achieve a full outer shell of eight electrons, a state that is energetically favorable according to the octet rule.
Silicon typically does this by forming four covalent bonds, sharing each of its four valence electrons with other atoms (such as oxygen or carbon) to fill its outer shell. This formation leads to a common tetrahedral shape seen in many silicon compounds. For instance, in silicon dioxide (SiO₂), each silicon atom is covalently bonded to four oxygen atoms, which fulfills the stability requirement by completing the octet of valence electrons.
Another important concept in silicon chemistry is its hybridization state. In most stable silicon compounds, it is sp³ hybridized, contributing to the tetrahedral geometry. This allows silicon to act as a Lewis acid due to its empty d orbitals, enabling it to accept electron pairs and form additional bonds if necessary.