Question

Of the noble gases, only \(\mathrm{Kr}, \mathrm{Xe},\) and \(\mathrm{Rn}\) are known to form a few compounds with \(\mathrm{O}\) and/or \(\mathrm{F}\). Write Lewis structures for these molecules: (a) \(\mathrm{XeF}_{2},\) (b) \(\mathrm{XeF}_{4}\) (c) \(\mathrm{XeF}_{6},\) (d) \(\mathrm{XeOF}_{4},\) (e) \(\mathrm{XeO}_{2} \mathrm{~F}_{2}\). In each case Xe is the central atom.

Short answer

The Lewis structures for each molecule would have Xe as the central atom bonded to F or O atoms by single bonds, with multiple lone pairs of electrons on the Xe atom and all negatively charged lone pairs on the fluorine and oxygen atoms. The exact configuration depends on the number of the particular atoms in each molecule.

Step-by-step solution

Recognizing valence electrons

First step is to count the valence electrons of each atom to be involved in the Lewis structure. Xenon (Xe) belongs to Group 18 and has 8 valance electrons. Fluorine (F) belongs to Group 17, containing 7 valence electrons. Oxygen (O) is in Group 16 and contains 6 valence electrons.

Lewis Structure for XeF2

XeF2 has a total of 22 valence electrons (8 from Xenon and 7 each from two fluorine atoms). Draw a structure where both fluorine atoms form a single bond with Xenon, leaving 3 lone pairs of electrons on the Xenon atom and 6 lone pairs total electronegatively charged on the two fluorine atoms.

Lewis Structure for XeF4

XeF4 has a total of 36 valence electrons (8 from Xenon and 7 from each of the 4 fluorine atoms). Draw a structure where each fluorine atom forms a single bond with Xenon, leaving 2 lone pairs of electrons on the Xenon atom, and 8 total lone pairs on the four fluorine atoms.

Lewis Structure for XeF6

XeF6 has a total of 50 valence electrons (8 from Xenon and 7 from each of the 6 fluorine atoms). Draw a structure where each fluorine atom forms a single bond with Xenon, leaving 1 lone pair of electrons on the Xenon atom and 12 total lone pairs on the six fluorine atoms.

Lewis Structure for XeOF4

XeOF4 has a total of 42 valence electrons (8 from Xenon, 6 from Oxygen, and 7 from each of the 4 fluorine atoms). Draw a structure where each Fluorine atom and the Oxygen atom form single bonds with Xenon, and leaving 1 lone pair of electrons on the Xenon atom and 10 total lone pairs on the four fluorine atoms and 2 lone pairs on Oxygen atom.

Lewis Structure for XeO2F2

XeO2F2 has a total of 34 valence electrons (8 from Xenon, 6 each from the 2 Oxygen atoms, and 7 each from the 2 fluorine atoms). Draw a structure where each Fluorine atom and the Oxygen atom form single bonds with Xenon, and leaving 2 lone pairs of electrons on the Xenon atom, 8 total lone pairs on the two oxygen atoms and 4 total lone pairs on the two fluorine atoms.

Key concepts

Valence Electrons

Understand the basics of chemical bonding; we've got to start with valence electrons. These are the electrons located in the outermost shell of an atom, and they play a pivotal role in how atoms connect with each other. In the case of Xenon (Xe), a noble gas, it has 8 valence electrons—this follows the 'octet rule,' which says that atoms are most stable with eight electrons in their valence shell.

However, Xenon is a bit of a rebel among its noble gas friends as it doesn't mind sharing or borrowing electrons. Fluorine (F), on the other hand, has 7 valence electrons, making it one electron shy of the perfect octet and very eager to form bonds. Similarly, Oxygen (O) has 6 valence electrons and is also looking to complete its octet. Recognizing valence electrons is not just counting; it's about understanding how these electrons shape the molecular architecture through bonding.

Lewis Structure Representation

Visualizing molecules is key, and Lewis structures are the sketches of the atomic world. They're a way to represent molecules showing how valence electrons are arranged among the atoms. Drawing them involves dots around atomic symbols to embody lone pairs, or single, double, or triple lines representing bonds between atoms.

For molecules like Xenon fluoride compounds, the struggle is to accommodate all those electrons while keeping an eye on the noble gas's comforts. This involves meticulous placement of dots and lines to match the total valence electrons with those accounted for in the structure. When learning how to draw these structures, imagine them as puzzles where electrons pair up like dancers, and the atoms are the ballroom.

Xenon Fluoride Compounds

XeF2, XeF4, and XeF6

Imagine the locker room of valence electrons, and Xenon's got 8 in its locker, while that Fluorine is looking for an electron to complete its set. Now, combine them—what do you get? A series of stable xenon fluoride compounds such as XeF2, XeF4, and XeF6. The trick to creating these compounds is in the sharing. The electrons between Xenon and Fluorine atoms share in just the right way to satisfy everyone's octet needs.

The final structure depends on the number of Fluorine atoms tagging along. The distribution of electrons will differ, increasing the lone pairs on Xenon as you move from XeF2 to XeF6. This results in different molecular shapes as well—a straight line for XeF2 and more complex geometries for XeF4 and XeF6.

Chemical Bonding of Noble Gases

Noble gases were once thought to be chemical loners, too aloof to bond with other elements. Xenon turned that idea on its head when it decided to interact with Fluorine and Oxygen, showing that under the right conditions, even the nobility could mingle. This rare bonding behavior stems from the fact that these elements have full valence shells but can be coaxed into bonding through the right energetic encounters—typically high pressures or the presence of a highly electronegative element like Fluorine.

Xenon's willingness to form compounds isn't just a quirk of chemistry; it's a window into the flexibility of chemical rules. It shows that under high enough pressure or with sufficiently persuasive elements, even the most unwilling of atoms can form lasting connections, resulting in a beautiful array of compounds that defy the old rules.