Question

Which of the following compounds or ions do not have an octet of electrons surrounding the central atom: \(\mathrm{BF}_{4}^{-}, \mathrm{SiF}_{4}, \mathrm{SeF}_{4}, \mathrm{BrF}_{4}^{-}, \mathrm{XeF}_{4} ?\)

Short answer

The compounds that do not follow the octet rule are \\(\mathrm{SeF}_{4}, \mathrm{BrF}_{4}^{-}, \mathrm{XeF}_{4}\\).

Step-by-step solution

Understand the Octet Rule

The octet rule states that atoms are generally most stable when they have eight electrons in their valence shell. We will determine if each compound adheres to this rule by examining the number of valence electrons around the central atom.

Check \\(\mathrm{BF}_{4}^{-}\\)

Boron (B) is the central atom. Normally, boron forms bonds with fewer than 8 valence electrons due to its electron deficiency. In \(\mathrm{BF}_{4}^{-}\)\, the boron shares electrons with four fluorine atoms. Despite this collaboration, boron ends up having more than 8 electrons around it in this case due to the negative charge.

Check \\( \mathrm{SiF}_{4}\\)

Silicon (Si) is the central atom, which has 4 valence electrons, and forms four bonds with fluorine. With each F contributing 1 electron, silicon achieves an octet (8 electrons total). Thus, \(\mathrm{SiF}_{4}\)\ satisfies the octet rule.

Check \\( \mathrm{SeF}_{4}\\)

Selenium (Se) has 6 valence electrons. In \(\mathrm{SeF}_{4}\)\, it forms 4 bonds with fluorine atoms, using 4 electrons and leaving 2 as a lone pair. This results in 10 electrons around Se, which exceeds the octet rule. Therefore, \(\mathrm{SeF}_{4}\)\ does not follow the octet rule.

Check \\( \mathrm{BrF}_{4}^{-}\\)

Bromine (Br) has 7 valence electrons, and due to the negative charge, there are 8 electrons to consider. Br forms 4 bonds with fluorine atoms, using these 4, leaving 4 electrons as 2 lone pairs. The presence of 12 electrons around Br indicates it does not follow the octet rule.

Check \\( \mathrm{XeF}_{4}\\)

Xenon (Xe) has 8 valence electrons. It forms 4 bonds with fluorine atoms, using 4 electrons and leaving 4 as 2 lone pairs. This leads to a total of 12 electrons around Xe, indicating it violates the octet rule.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical reactions. These electrons determine how atoms will interact and bond with each other. In a Lewis structure, these electrons are represented as dots. Understanding the number of valence electrons helps us predict how an atom will form chemical bonds.

For example, boron has 3 valence electrons, silicon has 4, and xenon has 8. Knowing these numbers allows us to visualize how many bonds each atom can form with others. It's essential for predicting the stability of molecules and ions. Therefore, valence electrons serve as the foundation for understanding chemical bonding and can help determine whether a structure satisfies the octet rule.

Chemical Bonding

Chemical bonding occurs when atoms share or transfer valence electrons. This process allows atoms to achieve greater stability, often reaching an electron configuration akin to the noble gases. Several types of bonds exist, with covalent bonds being the most common, formed by sharing electrons.

In the case of \(\mathrm{SiF}_{4}\), silicon forms four covalent bonds with fluorine, achieving a stable octet. However, not all molecules strictly follow the octet rule. Molecules like \(\mathrm{SeF}_{4}\) surpass the octet by accommodating additional electrons through expanded valence shells. This can occur in elements present in or beyond the third period of the periodic table.

Understanding chemical bonding helps clarify why certain elements can form stable or unstable compounds, with the octet rule as a guiding principle for many but not all scenarios.

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule. These structures help visualize the arrangement of atoms and predict how they will bond.

Creating a Lewis structure involves counting the total number of valence electrons, arranging atoms to show bonds (usually with the least electronegative atom in the center), and distributing remaining electrons as lone pairs. For example, in \(\mathrm{XeF}_{4}\), xenon's central position and its additional lone pairs illustrate an expanded octet, which is clearly depicted in its Lewis structure.

Lewis structures are invaluable for understanding molecular geometry, predicting the number of bonds an atom can form, and identifying formal charges. This visual tool simplifies complex molecules, making it easier to analyze their chemical properties and behaviors.