Question

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. \begin{equation}\begin{array}{lllll}{\text { a. } \mathrm{PF}_{5}} & {\text { b. } \mathrm{I}_{3}^{-}} & {\text { c. }} & {\mathrm{SF}_{4}} & {\text { d. GeF }_{4}}\end{array}\end{equation}

Short answer

PF5: Phosphorus with five single bonds to fluorine, expanded octet. I3-: Linear, two terminal I atoms single-bonded to a central I atom, with three lone pairs on each terminal I and two on the central I, overall negative charge. SF4: Sulfur with four single bonds to fluorine, one lone pair giving a seesaw shape, expanded octet. GeF4: Germanium with four single bonds to fluorine.

Step-by-step solution

Write the Lewis Structure for PF5

Phosphorus has five valence electrons and each fluorine atom has seven valence electrons. Phosphorus can expand its octet to accommodate 10 electrons. Place P in the center and distribute the five F atoms around it, each with a single bond to P. Complete the octets of the F atoms. P uses 10 electrons in bonding, displaying an expanded octet.

Write the Lewis Structure for I3-

Each iodine atom has seven valence electrons, and the ion has an extra electron due to its negative charge, totaling 22 electrons to be placed. Two I atoms will be bonded to a central I atom. Place brackets around the structure and add a negative sign to indicate the charge of the ion. The central I will have an expanded octet.

Write the Lewis Structure for SF4

Sulfur has six valence electrons, and each fluorine atom has seven valence electrons, totaling 34 electrons. Place S in the center, with four F atoms surrounding it. Each F will be bonded with a single bond to S. Sulfur will have an expanded octet, using 10 electrons in bonding.

Write the Lewis Structure for GeF4

Germanium has four valence electrons, and each fluorine has seven valence electrons, totaling 32 electrons. Place Ge in the center, with four F atoms around it, each with a single bond to Ge. Complete the octets of the F atoms. Ge uses eight electrons in bonding.

Key concepts

Valence Electrons and Lewis Structures

The key to understanding molecules begins with valence electrons. These are the electrons located in the outermost shell of an atom and are instrumental in bond formation. In Lewis structures, we represent an atom's valence electrons as dots surrounding the element's symbol. For example, in the molecule PF₅, phosphorus (P) is surrounded by five fluorine (F) atoms. Phosphorus has five valence electrons, while each fluorine atom brings seven to the table. The Lewis structure for this molecule shows each fluorine with a single bond to the central phosphorus, resulting in P sharing its valence electrons with the Fs to complete their octets. However, P itself ends up with an expanded octet, housing 10 electrons.

This visualization allows students to see where electrons are placed and how atoms bond in a molecule. The goal is not just to follow the rules, but to understand how electrons are shared or transferred between atoms to achieve stable configurations, typically by completing octets. However, atoms in the third period of the periodic table, such as phosphorus, sulfur, and even iodine in the case of I₃^-, can hold more than eight electrons, leading to structures with expanded octets.

Expanded Octets

While the 'octet rule' is a fundamental concept in chemistry, which posits that atoms tend to form bonds until they have eight electrons surrounding them, there are exceptions. Notably, certain atoms can have expanded octets, meaning they can accommodate more than eight electrons. Elements in or beyond the third period of the periodic table have d-orbitals in addition to s and p-orbitals, which allows them to hold additional electrons.

In the exercise's molecules, both P in PF₅ and S in SF₄ exhibit expanded octets. Phosphorus, a third-period element, can expand its octet to hold 10 electrons, while sulfur can do the same, utilizing its d-orbitals. This concept is crucial when drawing Lewis structures for molecules containing these elements, as recognizing the possibility of expanded octets can help avoid errors in predicting molecular structures and properties.

Molecular Geometry

The shape of a molecule, known as its molecular geometry, is critical to understanding its function and properties. Upon establishing the Lewis structure, molecular geometry can be predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory states that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion.

For instance, PF₅ has a trigonal bipyramidal geometry due to its five pairs of bonding electrons. SF₄, with four bonding pairs and one lone pair of electrons, adopts a see-saw shape. It is essential to recognize that these geometries directly influence a molecule's polarity, reactivity, and interaction with other molecules. Thus, properly deducing the molecular geometry aids in predicting chemical behavior and is a vital skill for students to develop.