Question

Draw an acceptable Lewis electron dot diagram for these molecules that violate the octet rule. a) \(\mathrm{SF}_{4}\) b) \(\mathrm{BeH}_{2}\)

Short answer

SF4 has an expanded octet, while BeH2 is electron-deficient.

Step-by-step solution

Understanding the Octet Rule Violations

The molecules \(\text{SF}_4\) and \(\text{BeH}_2\) do not obey the octet rule. In \(\text{SF}_4\), sulfur can have more than 8 electrons due to the availability of 3d orbitals. In \(\text{BeH}_2\), beryllium does not have enough electrons to form 4 bonds as it does not fill the p orbitals.

Drawing the Lewis Structure for \\( ext{SF}_4\\)

1. Count the valence electrons: Sulfur has 6 and each fluorine has 7, totaling 34 electrons.2. Place sulfur in the center and form single bonds to each of the four fluorine atoms.3. Assign remaining electrons to fluorine atoms to fulfill their octets.4. Place any remaining electrons as non-bonding pairs on sulfur. \[\begin{array}{c} \begin{array}{cccccc} F & & & & F & \ & & :S: & & & \ F & & & & F & \ \end{array}\end{array}\]

Drawing the Lewis Structure for \\( ext{BeH}_2\\)

1. Count the valence electrons: Beryllium has 2 electrons and each hydrogen has 1, totaling 4 electrons.2. Place beryllium in the center and form single bonds with the two hydrogen atoms.3. Beryllium has only 4 electrons in its valence shell, which is typical for an electron-deficient molecule like \(\text{BeH}_2\).\[\begin{array}{c} H - \, Be \, - H \\end{array}\]

Key concepts

Octet Rule Violation

The octet rule is a simple guideline in chemistry that suggests atoms often form bonds to surround themselves with eight electrons, akin to the noble gases. This rule, while generally useful, has its exceptions. For instance, molecules like \( \mathrm{SF}_4 \) and \( \mathrm{BeH}_2 \) do not follow this rule.

In \( \mathrm{SF}_4 \), sulfur can accommodate more than eight electrons due to the presence of 3d orbitals, making it one of the exceptions. This is known as an expanded octet. On the other hand, \( \mathrm{BeH}_2 \) is electron-deficient. Beryllium has only four electrons in its valence shell, as it forms bonds with two hydrogen atoms, each sharing one electron. Understanding these exceptions is crucial for mastering Lewis structures and predicting the shapes of molecules.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in bond formation. Counting valence electrons is the first step in drawing Lewis structures.

For \( \mathrm{SF}_4 \), sulfur has 6 valence electrons, and each fluorine has 7. Totaling the electrons, we get 34. Placing sulfur as the central atom, we use these electrons to form single bonds with each of the four fluorines, fulfilling their octets first.

• Beryllium in \( \mathrm{BeH}_2 \), however, has only 2 valence electrons.
• Each hydrogen atom has 1 valence electron, making the total 4 electrons available.

This highlights the unique electronic configuration beryllium has in \( \mathrm{BeH}_2 \), which stands as an electron-deficient molecule, leading to its octet rule violation.

Electron Dot Diagrams

Also known as Lewis structures, electron dot diagrams represent atoms and their bonding within a molecule.

Let's look at \( \mathrm{SF}_4 \):

• Sulfur is central, surrounded by four fluorine atoms.
• Single bonds connect sulfur with each fluorine, and lone pairs of electrons are positioned accordingly.

For \( \mathrm{BeH}_2 \):

• Beryllium connects with two hydrogen atoms with only two single bonds total.
• No lone pairs are available on beryllium once bonding is complete.

These diagrams visually indicate how valence electrons are shared, and thus understanding them helps clarify molecular interactions and potential reactive sites.

Molecular Geometry

Molecular geometry describes the shape of a molecule, which arises from the spatial arrangement of its atoms. This aspect is derived from the Lewis structures and the electron repulsions described by the VSEPR theory.

For \( \mathrm{SF}_4 \), the molecular geometry is often described as a seesaw shape. The central sulfur atom is bonded to four fluorines with one lone pair, causing distortion from the ideal tetrahedral geometry.

In the case of \( \mathrm{BeH}_2 \), the molecule is linear. Beryllium forms single bonds with two hydrogens, and there are no lone pairs on the beryllium atom. The minimal repulsion between the bonds leads to this straight-line shape.

These geometries influence the physical and chemical properties of molecules, affecting everything from boiling points to reactivity in complex reactions.