Question

The \(\mathrm{SF}_{5}^{-}\) ion is formed when \(\mathrm{SF}_{4}(\mathrm{~g})\) 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 obtained by distributing the 34 and 42 valence electrons, respectively, around the central Sulfur atom. For SF₄, there are 4 bonding pairs and 1 lone pair, giving it a "see-saw" molecular shape based on the trigonal bipyramidal arrangement. For SF₅⁻, there are 5 bonding pairs and 1 lone pair, resulting in a square pyramidal molecular shape based on the octahedral arrangement.

Step-by-step solution

Determine the total number of valence electrons

For SF₄: Sulfur (S) belongs to Group 16, so it has 6 valence electrons. Fluorine (F) belongs to Group 17, so it has 7 valence electrons. Since there are 4 Fluorine atoms, we multiply 7 by 4. Total valence electrons = 6 (from S) + 4 * 7 (from four F atoms) = 6 + 28 = 34 For SF₅⁻: Sulfur (S) has 6 valence electrons. Each of the 5 Fluorine atoms contributes 7 valence electrons. Since it's a negatively charged ion, we add one electron. Total valence electrons = 6 (from S) + 5 * 7 (from five F atoms) + 1 (extra electron due to the negative charge) = 6 + 35 + 1 = 42

Draw the Lewis structures

For SF₄: 1. Place Sulfur as the central atom. 2. Arrange the 4 Fluorine atoms around the Sulfur atom. 3. Distribute the 34 valence electrons as bonding pairs (single bonds) between the central atom and surrounding atoms. 4. Distribute the remaining electrons as lone pairs. For SF₅⁻: 1. Place Sulfur as the central atom. 2. Arrange the 5 Fluorine atoms around the Sulfur atom. 3. Distribute the 42 valence electrons as bonding pairs (single bonds) between the central atom and surrounding atoms. 4. Distribute the remaining electrons as lone pairs. 5. Add brackets and indicate the charge.

Predict the molecular structures

For SF₄: Applying VSEPR theory, we see that SF₄ has 4 bonding pairs and 1 lone pair on the central Sulfur atom. It has a total of 5 electron domains, which means it will have a trigonal bipyramidal arrangement, but due to the lone pair, it adopts a "see-saw" molecular shape. For SF₅⁻: Applying VSEPR theory, we see that SF₅⁻ has 5 bonding pairs and 1 lone pair on the central Sulfur atom. It has a total of 6 electron domains, which means it will have an octahedral arrangement, but due to the lone pair, it adopts a square pyramidal molecular shape.

Key concepts

Valence Electrons

Valence electrons are the electrons present in the outermost shell of an atom. They play a key role in determining how atoms interact and bond with each other. Understanding the number of valence electrons is crucial for drawing Lewis structures, which depict the bonding between atoms in a molecule.

Consider sulfur (\(S\) in \(\mathrm{SF}_{4}\)) and fluoride (\(\mathrm{F}\)) as examples:

• Sulfur is in Group 16 of the periodic table and has 6 valence electrons.
• Fluorine, found in Group 17, has 7 valence electrons.

In the molecule \(\mathrm{SF}_{4}\), the total valence electrons can be calculated by adding the valence electrons from sulfur and four fluorine atoms: 6 from \(\text{S}\) + 28 from \(4 \times 7\text{F}\) = 34 valence electrons.

When adding an extra fluoride ion to form \(\mathrm{SF}_{5}^{-}\), an additional electron is included because of the negative charge, resulting in a total of 42 valence electrons.

VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) theory is an essential concept for predicting the three-dimensional shape of molecules. According to this theory, electron pairs, whether they are in bonds or lone pairs, repel each other and, hence, arrange themselves as far apart as possible around the central atom.

For instance, in \(\mathrm{SF}_{4}\), the central sulfur atom has 4 bonding pairs and 1 lone pair. These 5 electron domains adopt a spatial arrangement that minimizes repulsion, resulting in a trigonal bipyramidal geometry. Due to the presence of the lone pair, however, the molecular shape becomes a "see-saw."

In contrast, \(\mathrm{SF}_{5}^{-}\) with 5 bonding pairs and 1 lone pair on sulfur, has 6 electron domains that exhibit an octahedral arrangement. The lone pair introduces a deviation leading to a square pyramidal shape. VSEPR theory thus helps in visualizing how lone pairs can affect the geometry of the molecule by altering idealized shapes.

Molecular Geometry

Molecular geometry describes the three-dimensional shape of a molecule, which is determined by the arrangement of its atoms in space. The shape of a molecule can influence its chemical behavior and physical properties, such as polarity and reactivity.

For \(\mathrm{SF}_{4}\), despite having a trigonal bipyramidal electronic geometry due to five electron domains, the lone pair present on sulfur shifts the actual molecular shape to a "see-saw" structure. This irregular shape results from the lone pair occupying more space and altering bond angles.

Conversely, \(\mathrm{SF}_{5}^{-}\) exhibits a square pyramidal molecular geometry. Although the basic electronic geometry is octahedral, one face of the octahedron is missing due to the lone pair, leading to the characteristic square pyramidal shape. Understanding these geometrical shifts is vital for predicting how molecules interact with each other.

Chemical Bonding

Chemical bonding is the process where atoms combine to form compounds by sharing or exchanging electrons. In the case of \(\mathrm{SF}_{4}\) and \(\mathrm{SF}_{5}^{-}\), these are largely covalent bonds where electrons are shared between sulfur and fluorine atoms.

The Lewis structures for these molecules are integral for illustrating these shared bonds. For \(\mathrm{SF}_{4}\), sulfur forms four single covalent bonds with each fluorine atom, sharing its valence electrons. \(\mathrm{SF}_{5}^{-}\) involves an additional bond due to the presence of an extra fluorine atom, facilitated by an extra electron because of its negative charge.

These covalent interactions are central to understanding the stability and formation of such molecules. Moreover, the introduction of an extra electron in \(\mathrm{SF}_{5}^{-}\) alters its electron cloud, affecting the geometry and consequently, its chemical properties. By analyzing these bonding interactions, one can deduce the reactivity and potential interactions of a given molecule.