Problem 37 - Chemistry: An Atoms-Focused Approach

Question

For many years, it was believed that the noble gases could not form covalently bonded compounds. However, xenon does react with fluorine. One of the products is the pentafluoroxenate anion, \(\mathrm{XeF}_{5}^{-} .\) Draw the Lewis structure of \(\mathrm{XeF}_{5}^{-}\) and predict its geometry.

Step-by-Step Solution

Answer

Answer: The Lewis structure of the pentafluoroxenate anion (XeF5-) has a central Xenon (Xe) atom surrounded by five Fluorine (F) atoms and one lone pair of electrons. The molecular geometry of XeF5- is square pyramidal.

Step 1: Count the valence electrons

We will count the number of valence electrons in the pentafluoroxenate anion by considering the number of valence electrons for each atom involved. Xenon (Xe) has 8 valence electrons (from the 5th period, 8A group) and each fluorine atom (F) has 7 valence electrons (from the 2nd period, 7A group). For the anion charge, we will add one extra electron. Total valence electrons = 8 (Xe) + 5 * 7 (F) + 1 (anion charge) = 8 + 35 + 1 = 44 electrons

Step 2: Draw the Lewis structure

To draw the Lewis structure, we will first position the atoms, showing that xenon is the central atom, surrounded by the five fluorine atoms. Xe--F | F Now, we will distribute the 44 valence electrons by forming the bonds and completing the octets around each fluorine atom. Xe-F || F-|-F | | F-|-F

Step 3: Add extra electrons on the central atom

As xenon (Xe) can hold more than 8 electrons (due to available d-orbitals), let's place the remaining two electrons on Xe. Xe-F || F-|-F . . F-|-F Now we have the Lewis structure of the pentafluoroxenate anion (XeF5-).

Step 4: Predict the molecular geometry

To predict the molecular geometry, we will apply the VSEPR (Valence Shell Electron Pair Repulsion) theory. Here, xenon has 5 bond pairs and 1 lone pair of electrons. According to VSEPR theory, based on the electron pair repulsion, the molecular geometry of XeF5- will be a distorted octahedron or square pyramid. In conclusion, the Lewis structure of the pentafluoroxenate anion (XeF5-) is given by the representation above, and its molecular geometry is square pyramidal.

Key Concepts

Chemical Bonding

In chemistry, chemical bonding refers to the force that holds atoms together in compounds. There are different types of bonds, such as covalent, ionic, and metallic bonds. In the case of noble gas compounds, we often deal with covalent bonds, where atoms share electrons to attain stability. For example, despite the inert nature of noble gases like xenon, they can still form compounds with highly electronegative elements like fluorine.

• Covalent bonds involve sharing of electrons.
• Noble gases usually have a full octet, but can expand their octet to form bonds.
• Xenon can form stable compounds due to its ability to hold more than 8 electrons.

These unique bonding situations, particularly with xenon, lead to interesting chemistry that expands our understanding of the reactive possibilities for traditionally nonreactive elements.

Lewis Structure

A Lewis structure is a simple way to represent the arrangement of valence electrons in a molecule. It shows how atoms are bonded and any lone pairs of electrons. For the pentafluoroxenate anion (\(\mathrm{XeF}_{5}^{-}\)), we start by calculating the total number of valence electrons to be distributed.

• Xenon contributes 8 valence electrons.
• Each of the five fluorine atoms contributes 7 electrons.
• One extra electron is added as indicated by the anion charge.

This leads to 44 electrons. In the structure, xenon is the central atom bonded to five fluorine atoms, with the final two electrons forming a lone pair on xenon. Through Lewis structures, we visualize not only the atom connections but how electrons are paired and distributed.

Molecular Geometry

Understanding the molecular geometry of a compound involves recognizing the three-dimensional arrangement of atoms. For \(\mathrm{XeF}_{5}^{-}\), this involves predicting how its atoms are arranged in space based on its bonding and electronic structure.

• Molecular geometry is essential for understanding how a molecule interacts with its environment.
• In the \(\mathrm{XeF}_{5}^{-}\) anion, the molecular geometry results from five bonded fluorine atoms and one lone pair on xenon.

This results in a square pyramidal shape, where four fluorine atoms form the base of the pyramid, while the fifth fluorine atom occupies the top vertex. This geometry is influenced by the lone pair that slightly distorts the shape from a perfect octahedron to a square pyramid.

VSEPR Theory

The VSEPR (Valence Shell Electron Pair Repulsion) theory is a key concept for predicting the shapes of molecules. It is based on the idea that electron pairs around a central atom will repulse each other and thus arrange themselves as far apart as possible in 3D space.

• VSEPR theory helps determine the molecular geometry.
• Electron pairs include both bonding pairs and lone pairs.

For \(\mathrm{XeF}_{5}^{-}\), the VSEPR model shows that the electron pairs—5 bonding and 1 lone—drive the shape into a square pyramidal geometry. The lone pair exerts a repulsive force differing from bonding pairs, leading to this unique structure. Understanding VSEPR theory is crucial for anyone studying molecules, as it provides a framework for predicting and explaining molecular shapes, influencing both chemical reactivity and properties.