Question

Sulfur forms a range of compounds with fluorine. Draw Lewis electron dot structures for \(\mathrm{S}_{2} \mathrm{F}_{2}\) (connectivity is FSSF), \(\mathrm{SF}_{2}, \mathrm{SF}_{4}, \mathrm{SF}_{6},\) and \(\mathrm{S}_{2} \mathrm{F}_{10} .\) What is the oxidation number of sulfur in each of these compounds?

Short answer

In S2F2, sulfur has +1; in SF2, +2; in SF4, +4; in SF6, +6; in S2F10, +5.

Step-by-step solution

Draw Lewis Structure for S2F2

To draw the Lewis structure for \( \mathrm{S}_2\mathrm{F}_2 \), start by arranging the atoms based on the provided connectivity: FSSF. Sulfur typically forms two bonds, and fluorine forms one bond as it needs only one electron to complete its octet. Place each fluorine atom with a single bond to one sulfur. Connect the two sulfur atoms with a single bond as well. Make sure each fluorine atom has 3 lone pairs and each sulfur atom has 2 lone pairs. Check the octet rule for all atoms, adjusting lone pairs if necessary.

Determine Oxidation Number for S2F2

In \( \mathrm{S}_2\mathrm{F}_2 \), fluorine is more electronegative than sulfur. Fluorine will have an oxidation number of -1. With two fluorine atoms, the total negative charge is -2. Therefore, for the compound to be neutral, the two sulfur atoms together must have an oxidation number of +2. Dividing this oxidation number equally between the two sulfurs, each sulfur has an oxidation number of +1.

Draw Lewis Structure for SF2

For \( \mathrm{SF}_2 \), sulfur is the central atom bonded to two fluorine atoms. Sulfur forms a single bond with each fluorine. Fluorine atoms will each have 3 lone pairs. Sulfur will have 2 lone pairs to satisfy the octet rule.

Determine Oxidation Number for SF2

In \( \mathrm{SF}_2 \), each fluorine atom has an oxidation number of -1, accounting for a total of -2 for both. Therefore, sulfur must have an oxidation number of +2 to maintain a neutral overall charge.

Draw Lewis Structure for SF4

To draw \( \mathrm{SF}_4 \), position sulfur at the center, connecting it to four fluorine atoms via single bonds. Each fluorine will have 3 lone pairs. Sulfur will have one lone pair, achieving an expanded octet which is possible for sulfur.

Determine Oxidation Number for SF4

In \( \mathrm{SF}_4 \), fluorine atoms collectively have an oxidation number of -4. Sulfur must have an oxidation number of +4 for the compound's neutrality.

Draw Lewis Structure for SF6

For \( \mathrm{SF}_6 \), sulfur is the central atom and forms six single bonds with fluorine atoms. Each fluorine atom has 3 lone pairs, achieving full octets, while sulfur has no lone pairs but uses its d orbitals to accommodate 12 electrons around it.

Determine Oxidation Number for SF6

In \( \mathrm{SF}_6 \), with six fluorine atoms, their total oxidation number is -6. Sulfur must thus have an oxidation number of +6 to balance.

Draw Lewis Structure for S2F10

To construct \( \mathrm{S}_2\mathrm{F}_{10} \), place two sulfur atoms bonded to each other. Each sulfur connects to five fluorine atoms. Fluorines will have 3 lone pairs each. The sulfur atoms take advantage of their ability to expand the octet.

Determine Oxidation Number for S2F10

In \( \mathrm{S}_2\mathrm{F}_{10} \), ten fluorine atoms contribute an oxidation number of -10. Thus, each sulfur atom needs an oxidation number of +5 to ensure the molecule's neutrality.

Key concepts

Lewis Structures

Lewis structures are diagrams that represent the arrangement of valence electrons in a molecule. These structures help us understand how atoms are bonded and how electrons are distributed within a compound. To draw Lewis structures for sulfur fluoride compounds like \( \mathrm{S}_2\mathrm{F}_2, \mathrm{SF}_2, \mathrm{SF}_4, \mathrm{SF}_6,\ \text{and}\ \mathrm{S}_2\mathrm{F}_{10} \), we follow a specific set of steps.

• Identify the central atom, which is usually the less electronegative element, typically sulfur in these compounds.
• Arrange the other atoms around the central atom based on the given formula or connectivity, such as FSSF for \( \mathrm{S}_2\mathrm{F}_2 \).
• Connect the atoms using single bonds initially, respecting the typical bonding preferences, such as fluorine forming one bond.
• Distribute the remaining electrons as lone pairs, starting with the most electronegative atoms until each has a complete octet, if possible.
• Ensure all atoms, especially sulfur, have their requisite number of electrons, adjusting for any expanded octets possible for sulfur.

The Lewis structure helps visualize the molecular geometry and electron distribution, which is crucial for understanding chemical reactions and bonding.

Oxidation Number

Oxidation numbers are assigned to atoms to help track electron transfer in chemical reactions. They can be considered as the hypothetical charge an atom would have if bonding were completely ionic. Calculating the oxidation number of sulfur in sulfur fluoride compounds involves some basic principles:

• Fluorine, which is highly electronegative, always has an oxidation number of -1.
• Sulfur's oxidation number is determined by balancing the total oxidation numbers to make the compound electrical neutral.
• In compounds like \( \mathrm{S}_2\mathrm{F}_2 \), with each fluorine contributing -1, the sum for fluorine atoms is -2, hence sulfur balances with +2 shared among the two sulfur atoms.
• The sum of oxidation numbers in a neutral compound is always zero, as seen in \( \mathrm{SF}_6 \) where +6 for sulfur balances the total -6 from the six fluorine atoms.

Understanding oxidation numbers is important in especially determining the electron flow in reactions, which is foundational for redox chemistry.

Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in compounds. In sulfur fluoride compounds, covalent bonding is predominant, characterized by the sharing of electrons between atoms.

• In covalent bonds, atoms share electrons to achieve a full outer shell, aligning with the octet rule, which holds for most atoms except for possible exceptions like expanded octets for sulfur.
• Each bond in these compounds generally involves two electrons, with sulfur potentially forming multiple bonds to fluorine given its higher number of valence electrons.
• Sulfur has the unique ability to expand its valence shell because it can utilize d-orbitals, which is why structures like \( \mathrm{SF}_6 \) can exist with six bonds.
• The bonding determines not only structure and stability but also physical and chemical properties of the molecules.

By examining these bonding scenarios, we can predict the behavior and reactivity of these compounds, aiding in applications in chemical synthesis and industrial processes.

Sulfur Oxidation States

The oxidation states of sulfur in its compounds with fluorine demonstrate sulfur's versatility in bonding. Depending on the compound, sulfur can exhibit a range of oxidation states:

• In \( \mathrm{SF}_2 \), sulfur typically shows an oxidation state of +2, balancing out the -2 from the two fluorine atoms.
• Advancing to \( \mathrm{SF}_4 \), the oxidation state of sulfur increases to +4, corresponding to the four bonds with fluorine atoms.
• In \( \mathrm{SF}_6 \), sulfur reaches its maximum common oxidation state of +6, indicative of sulfur’s ability to form bonds with six fluorine atoms.
• In the dimeric structure \( \mathrm{S}_2\mathrm{F}_{10} \), each sulfur contributes to a unified +10 oxidation state when shared, with each sulfur technically holding a +5 state.

These varying states of sulfur reflect its ability to accommodate different numbers of electrons, leading to diverse chemical behaviors and molecule configurations.