Question

The discovery in the 1960 s that xenon and fluorine react to form a molecular compound was a surprise to most chemists, because it had been thought that noble gases could not form bonds. (a) Why was it thought that noble gases could not form bonds? (b) Draw a Lewis structure of \(\mathrm{XeF}_{4}\) in which \(\mathrm{Xe}\) is the central atom. How many electron clouds are there on the central atom? (c) What type of bonds are the \(\mathrm{Xe}-\mathrm{F}\) bonds? Explain.

Short answer

(a) Noble gases have a stable electronic configuration. (b) XeF4 has 6 electron clouds. (c) Xe-F bonds are polar covalent.

Step-by-step solution

Understanding Noble Gases

Noble gases were believed to be inert because they have complete valence electron shells, giving them a stable electronic configuration. This stability minimizes their tendency to form chemical bonds.

Lewis Structure of XeF4

To draw the Lewis structure of \( \text{XeF}_4 \), start by counting the valence electrons. Xenon (Xe) has 8 valence electrons and each fluorine (F) has 7 valence electrons. For 4 fluorines, the total is \(8 + 4 \times 7 = 36\) valence electrons. Xe is the central atom, with 4 F atoms surrounding it. Each F shares a pair of electrons with Xe, forming 4 single bonds.

Electron Clouds on the Central Atom

The central atom Xe in \( \text{XeF}_4 \) has 6 electron clouds: 4 bonding pairs (from the bonds with fluorine) and 2 lone pairs.

Type of Bonds in XeF4

The \( \text{Xe-F} \) bonds in \( \text{XeF}_4 \) are polar covalent bonds. This is because Xe and F share electrons, but F is more electronegative and attracts the shared electrons more strongly, leading to a dipole.

Key concepts

Noble Gases

Noble gases are a group of elements found in the far right column of the periodic table, including helium, neon, argon, krypton, xenon, and radon. These elements are known for their lack of reactivity under normal conditions.
Unlike other elements, noble gases possess a full set of electrons in their outer shell, typically comprising 8 electrons. This full valence shell creates a very stable electron configuration, making them generally uninterested in forming bonds with other elements.
Because of this stability, noble gases were historically considered "inert" and were thought not to form any chemical compounds. The discovery of noble gas compounds challenged this view, expanding our understanding of chemical bonding and reactivity.

Lewis Structures

Drawing Lewis structures is a crucial way to represent molecules and show how atoms are bonded together and where lone pairs of electrons reside. To create a Lewis structure, you start by counting all the valence electrons in the molecule.
For example, in the molecule XeF\(_4\), xenon (Xe) holds 8 valence electrons, and each of the four fluorine (F) atoms has 7 valence electrons. This results in a total of 36 valence electrons.
Here’s how you can distribute these electrons:

• Xenon serves as the central atom, while the four fluorine atoms surround it.
• Each fluorine atom forms a single bond with xenon, sharing a pair of electrons (a bonding pair).
• After placing the bonding pairs, xenon’s structure still has two pairs of lone electrons, which are not shared with any fluorine atom.

The Lewis structure thus reveals the arrangement of electrons and the type of chemical bonds in a molecule.

Polar Covalent Bonds

Polar covalent bonds occur when two atoms share electrons unequally, leading to partial electric charges on either side of the bond. This typically happens when one atom is more electronegative than the other, meaning it attracts the shared electrons more strongly.
In the case of the compound XeF\(_4\), the bonds between xenon (Xe) and fluorine (F) are polar covalent. Although xenon is a noble gas, it bonds with fluorine because fluorine is highly electronegative and can pull the shared electrons closer to itself more effectively.
This difference in electronegativity between xenon and fluorine causes the electrons in the bond to skew towards fluorine, creating a dipole moment. The result is a molecule in which each Xe-F bond has partial positive and negative charges at the ends, contributing significant characteristics to the chemical properties and behavior of the compound.