Question

\(\mathrm{SF}_{6}, \mathrm{ClF}_{5},\) and \(\mathrm{XeF}_{4}\) are three compounds whose central atoms do not follow the octet rule. Draw Lewis structures for these compounds.

Short answer

The Lewis structures for the given compounds are as follows: For \(\mathrm{SF}_{6}\): \[ \chemfig{F-[:30]S(-[:90]F)(-[:150]F)(-[:210]F)(-[:270]F)-[:330]F} \] For \(\mathrm{ClF}_{5}\): \[ \chemfig{F-[:18]Cl(-[:90]F)(-[:162]F)(-[:234]F)(-[:306]F)<:18(F)} \] For \(\mathrm{XeF}_{4}\): \[ \chemfig{F-Xe(-[:90]F)(-[:180]F)(<:0|--[:30]F)(-[:270]F)} \]

Step-by-step solution

1. Calculate the total number of valence electrons

First, determine the number of valence electrons for each atom in the compound and add them together to find the total number of valence electrons. The number of valence electrons for an element corresponds to its group number in the periodic table. For \(\mathrm{SF}_{6}\): Sulfur (S): 6 valence electrons (Group 16) Fluorine (F): 7 valence electrons (Group 17) Total = 6 + 6(7) = 48 valence electrons For \(\mathrm{ClF}_{5}\): Chlorine (Cl): 7 valence electrons (Group 17) Fluorine (F): 7 valence electrons (Group 17) Total = 7 + 5(7) = 42 valence electrons For \(\mathrm{XeF}_{4}\): Xenon (Xe): 8 valence electrons (Group 18) Fluorine (F): 7 valence electrons (Group 17) Total = 8 + 4(7) = 36 valence electrons

2. Arrange atoms in a suitable structure

Next, arrange the atoms in the compound. The central atom should be the one that does not follow the octet rule. For each compound, arrange the remaining atoms evenly around the central atom. For \(\mathrm{SF}_{6}\): S is the central atom with six F atoms surrounding it. For \(\mathrm{ClF}_{5}\): Cl is the central atom with five F atoms surrounding it. For \(\mathrm{XeF}_{4}\): Xe is the central atom with four F atoms surrounding it.

3. Connect the atoms using covalent bonds

Connect the atoms using single covalent bonds, which consist of one pair of shared electrons. We will also subtract the electrons involved in these bonds from our total number of valence electrons for each compound. For \(\mathrm{SF}_{6}\): Connect each F atom to the S atom, using 2(6) = 12 electrons, leaving 48 - 12 = 36 electrons. For \(\mathrm{ClF}_{5}\): Connect each F atom to the Cl atom, using 2(5) = 10 electrons, leaving 42 - 10 = 32 electrons. For \(\mathrm{XeF}_{4}\): Connect each F atom to the Xe atom, using 2(4) = 8 electrons, leaving 36 - 8 = 28 electrons.

4. Add lone pairs to complete the Lewis structure

Finally, add lone pairs (pairs of electrons not involved in covalent bonding) to each atom to complete the Lewis structure, using the remaining valence electrons calculated in step 3. For \(\mathrm{SF}_{6}\): Each F atom receives 3 lone pairs (using 6(6) = 36 electrons). Lewis structure for \(\mathrm{SF}_{6}\): \[ \chemfig{F-[:30]S(-[:90]F)(-[:150]F)(-[:210]F)(-[:270]F)-[:330]F} \] For \(\mathrm{ClF}_{5}\): Each F atom receives 3 lone pairs (using 6(5) = 30 electrons). The Cl atom receives 1 lone pair (using 2 electrons). Lewis structure for \(\mathrm{ClF}_{5}\): \[ \chemfig{F-[:18]Cl(-[:90]F)(-[:162]F)(-[:234]F)(-[:306]F)<:18(F)} \] For \(\mathrm{XeF}_{4}\): Each F atom receives 3 lone pairs (using 6(4) = 24 electrons). The Xe atom receives 2 lone pairs (using 4 electrons). Lewis structure for \(\mathrm{XeF}_{4}\): \[ \chemfig{F-Xe(-[:90]F)(-[:180]F)(<:0|--[:30]F)(-[:270]F)} \]

Key concepts

Octet Rule

In chemistry, the octet rule is a guideline that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell. This configuration is particularly stable because it resembles the electron configuration of the noble gases. Most atoms aim to complete their outer shell by sharing, gaining, or losing electrons through chemical bonding.

However, the compounds

• \( \text{SF}_6 \)
• \( \text{ClF}_5 \)
• \( \text{XeF}_4 \)

are exceptions to the octet rule. These molecules have central atoms that hold more than eight electrons in their valence shell. This typically happens because these atoms belong to lower periods of the periodic table, where they have access to d orbitals. This allows them to expand their octet and accommodate more electrons.

When drawing Lewis structures, these exceptions showcase the limitations and flexibility of the octet rule, demonstrating that while the rule is a useful guideline, there are valid exceptions.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how the atom will bond with others. In a Lewis structure, dots representing these electrons are placed around the elemental symbol to show their availability for bonding. The number of valence electrons generally corresponds to an element's group number in the periodic table.

For instance:

• Sulfur (\( \text{S} \)) in group 16 has 6 valence electrons.
• Fluorine (\( \text{F} \)) in group 17 has 7 valence electrons.
• Xenon (\( \text{Xe} \)) in group 18 has 8 valence electrons.

Understanding how to count and utilize valence electrons is essential in predicting how atoms will interact and bond to form molecules. For molecules that don’t follow the octet rule, understanding their valence electrons helps us see why these elements can form stable compounds even when their electron shells are expanded, as seen in \( \text{SF}_6 \), \( \text{ClF}_5 \), and \( \text{XeF}_4 \).

Covalent Bonds

Covalent bonds are formed when atoms share pairs of electrons. This type of bonding typically occurs between non-metal atoms that have similar electronegativities. It allows each atom to attain the electron configuration of a noble gas, achieving a stable electron arrangement.

In the context of the Lewis structures for \( \text{SF}_6 \), \( \text{ClF}_5 \), and \( \text{XeF}_4 \), covalent bonds are represented by lines between the central atom (sulfur, chlorine, xenon) and the surrounding fluorine atoms. Each line is equivalent to a pair of shared electrons, forming a strong linkage between the atoms.

• In \( \text{SF}_6 \), each sulfur-fluorine bond involves the sharing of one pair of electrons.
• For \( \text{ClF}_5 \), the central chlorine atom shares electrons with each fluorine.
• In \( \text{XeF}_4 \), xenon shares electrons with the surrounding fluorines, creating four single covalent bonds.

The strength and directionality of these bonds play a significant role in determining the shape and stability of the molecules.

Lone Pairs

Lone pairs are pairs of valence electrons that are not shared with another atom in a covalent bond. They are also known as non-bonding pairs. These electrons fill space around an atom and take part in determining the geometry and polarity of a molecule.

In Lewis structures, lone pairs are shown as pairs of dots around atoms. They influence the shape of a molecule by exerting repulsive forces on bonding pairs, thus affecting the bond angles.

For example:

• In \( \text{SF}_6 \), fluorine atoms possess three lone pairs each, contributing to the molecule's overall stability and geometry.
• In \( \text{ClF}_5 \), aside from the fluorines having lone pairs, the chlorine atom itself holds an additional lone pair, affecting the molecule's shape.
• In \( \text{XeF}_4 \), the central xenon atom has two lone pairs, resulting in a square planar shape as these pairs impose additional repulsion on the bonding pairs.

Lone pairs are an essential consideration in predicting the shape and reactivity of molecules, often having a profound influence despite not contributing to bonding directly.