Question

The \(\mathrm{SF}_{5}^{-}\) ion is formed when \(\mathrm{SF}_{4}(g)\) reacts with fluoride salts containing large cations, such as \(\mathrm{CsF}(s)\). Draw the Lewis structures for \(\mathrm{SF}_{4}\) and \(\mathrm{SF}_{5}^{-}\), and predict the molecular structure of each.

Short answer

The Lewis structures for SF4 and SF5- are completed with a central sulfur atom connected to 4 and 5 fluorine atoms using single bonds, respectively, plus a lone pair on the sulfur atom in each case. SF4 has a see-saw molecular structure, while SF5- has a square pyramidal molecular structure.

Step-by-step solution

Find the total number of valence electrons for each molecule

We need to find the total number of valence electrons for SF4 and SF5-. We do this by identifying the number of valence electrons each atom brings to the molecule, then add them together. For SF4: - Sulfur has 6 valence electrons. - Each fluorine atom has 7 valence electrons, so we have a total of 4 x 7 = 28 valence electrons from fluorine. Adding the valence electrons together: 6 (S) + 28 (F) = 34 valence electrons For SF5-: - Sulfur has 6 valence electrons. - Each fluorine atom has 7 valence electrons, so we have a total of 5 x 7 = 35 valence electrons from fluorine. - Since it is a negatively charged ion, we add an extra electron to the total count: 35 (F) + 1 = 36 valence electrons. Adding the valence electrons together: 6 (S) + 36 (F) = 42 valence electrons

Draw the Lewis structures for both molecules

To draw the Lewis structures, we start by connecting the central atom (Sulfur in both cases) to the surrounding atoms (Fluorine) by single bonds. Then we fill the remaining electron pairs on the surrounding atoms to satisfy the octet rule. Finally, we can assign any leftover electrons as lone pairs on the central atom. For SF4: 1. Connect Sulfur to each of the 4 Fluorine atoms using single bonds. We have used 4 * 2 = 8 electrons so far. 2. Complete the octets of each Fluorine atom. We have used another 24 electrons here (4 * 6 = 24). 3. Assign the remaining 2 electrons as a lone pair on the central Sulfur atom. The Lewis structure of SF4 is completed. For SF5-: 1. Connect Sulfur to each of the 5 Fluorine atoms using single bonds. We have used 5 * 2 = 10 electrons so far. 2. Complete the octets of each Fluorine atom. We have used another 30 electrons here (5 * 6 = 30). 3. Assign the remaining 2 electrons as a lone pair on the central Sulfur atom. The Lewis structure of SF5- is completed.

Predict the molecular structure of each molecule

To predict the molecular structure, we must identify the arrangement of atoms and electron pairs around the central atom. For this, we can use the VSEPR (Valence Shell Electron Pair Repulsion) theory, which states that electron pairs around the central atom repel each other and try to be as far apart as possible. For SF4: There are 4 bonded pairs and 1 lone pair around the central Sulfur atom. This results in a see-saw molecular structure. For SF5-: There are 5 bonded pairs and 1 lone pair around the central Sulfur atom. This results in a square pyramidal molecular structure. In conclusion, SF4 has a see-saw molecular structure, while SF5- has a square pyramidal molecular structure.

Key concepts

Lewis Structures

Lewis structures are a way to represent molecules by showing the arrangement of valence electrons around the atoms. For the molecule SF ₄, sulfur is the central atom connected to four fluorine atoms. Each single bond between sulfur and fluorine accounts for two electrons. After drawing these bonds, it's important to ensure that the remaining valence electrons satisfy the octet rule for each fluorine while also accounting for any extra electrons, such as lone pairs on sulfur.

In SF ₄, sulfur is surrounded by four bonded pairs and a lone pair, which gives it a total of 34 valence electrons. Similarly, for SF ₅^-, sulfur is at the center of five fluorine atoms with a negative charge adding an extra electron. This results in a total of 42 valence electrons.

By properly arranging the electrons and bonds, Lewis structures help visualize how electrons are shared or distributed, offering a foundational insight into the molecule's electron configuration.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding. Understanding how many valence electrons an atom possesses allows us to predict how it will interact and bond with other atoms.

Each element in the periodic table follows a pattern, with elements in the same group having the same number of valence electrons. For example, sulfur in SF ₄ and SF ₅^- has six valence electrons. Fluorine, on the other hand, has seven valence electrons. Counting the total number of valence electrons in a molecule is crucial to creating its Lewis structure.

In molecular structures like SF ₅^-, the negative charge indicates an additional valence electron, changing the balance of electrons. By correctly counting and distributing these valence electrons, we can predict molecular shapes and understand bonding patterns.

VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) theory is used to predict the three-dimensional shape of molecules. It starts with the idea that electron pairs around a central atom repel each other and will position themselves as far apart as possible to minimize repulsion.

For instance, in SF ₄, there are four bonded pairs and one lone pair around the sulfur atom. These arrange themselves into a see-saw shape to reduce repulsion. In SF ₅^-, five bonded pairs and one lone pair create a square pyramidal shape.

By applying VSEPR theory, we can understand why molecules have certain shapes, impacting their reactivity and properties. This theory is essential in predicting molecular geometry and understanding the spatial arrangement of atoms in a molecule.