The inert gas xenon and the halogen fluorine combine to generate XeF2, a covalent inorganic halide. This is an active solvent that is soluble in a variety of fluorides, including HF and bromine pentafluoride.
Consider the following equation for the synthesis of xenon difluoride:
Xe + F2 ——Heat——> XeF2
XeF2 is a fluorinating and oxidising agent that can be used to oxidise a variety of hydrocarbons, including aromatic and acyclic molecules.
Furthermore, without the usage of external energy, this fluoride molecule can be used to etch silicon to generate silicon tetrafluoride (SiF4).
When it comes to the appearance of XeF2, it is a colourless to white crystalline solid with a density of roughly 4.32 g/cc.
This halide can produce serious side effects such as skin burns and severe eye impairment. Not only that, but if inhaled or ingested, it is lethal.
Two or more atoms react and combine to generate homogeneous and heterogeneous molecules. Molecules are formed through the establishment of certain bonds that hold the atoms together based on their strength.
Chemical bonding is the backbone that defines the internal structure and nature of a molecular complex, as well as the properties it possesses ( both physical and chemical).
Before we go into the chemical bonding of XeF2 in this article, there are a few terms and concepts that you should familiarise yourself with.
Electrons of Valence
When it comes to chemical bonding, the concept of valence electrons is the first thing to grasp. The outermost shell electrons of an atom that govern its valency are known as valence electrons.
When looking at the periodic table, the groups will assist us in determining the valence electron number of a particular atomic element.
Carbon, for example, belongs to group 4 (also known as Group XIV), hence its valence shell has four electrons.
Lewis Structure of XeF2
Lewis Structure, also known as electron dot structure, is a fundamental model of chemical bonding in which the valence electron concept is used to sketch a two-dimensional figure of a given molecule schematically.
The outer shell electrons are represented by dots, and the bond type is represented by lines.
The element xenon is an inert gas. It possesses eight valence electrons as a result. Because fluorine is a halogen in group VII, it has a valency of seven.
The total number of valence electrons is equal to 8 + 7*2 = 22.
Because xenon is less electronegative than fluorine, it will take up the position of a central atom, according to the periodic table.
We’ll find out where the electrons are located to reach the octet configuration now that we have the central atom and the total valence electron number ( i.e. eight electrons outside each atom in its outer shell)
As you can see, we’ve completed the octet fulfilment for the three atoms inside XeF2 and formed a single bond.
However, counting the total valence electrons in the figure reveals that the dots sum up to 20, not 22.
Instead of single bonds, the remaining two valence electrons should be used to build a double bond.
Because fluorine is so electronegative, it rarely forms double bonds. So we calculate the formal charge by keeping the valence electrons near xenon.
We can assign a formal charge to an atom if we suppose that every electron is shared equally among atoms.
This is an important topic in chemical bonding, particularly when determining the Lewis Structure, where we must check the least possible formal charges of each combining atom to provide the appropriate diagrammatic representation.
Charge in writing
The formal charge of Xe is equal to 8 (valence electron number) – 0.5*4 (number of bound electrons) – 6 (number of lone pair electrons) = 0.
As a result, we get the best Lewis Structure because both elements have the lowest possible formal charge values.
Note that xenon possesses more than eight electrons in its valence shell cloud, hence it is an exception to the octet rule.
Geometry of Molecular Structure
We need to study about molecular geometry if we want to learn more about how the inside of a molecule looks.
This is a transition from 2D to 3D structure representation, allowing us to visualise how a molecule stays in bonding nature in reality.
We employ a theory called the Valence Shell Electron Pair Repulsion Model to help us calculate the correct molecular shape, as well as bond lengths and angles (VSEPR).
This hypothesis works by minimising the like charge repulsion that occurs between negative electron clouds surrounding atomic nuclei in order to obtain a near-stable composite.
Because xenon is a noble gas in group 8 (group XVIII), it does not tend to form bonds. It serves as the centre element in XeF2, as well as forming two single bonds with the two fluorine atoms.
As we previously established, molecule geometry can be predicted using VSEPR theory. Take a look at the following diagram:
So, according to their steric number and lone pair number, we may determine the various forms of molecules here.
Lewis Structure reveals that Xenon in XeF2 has three lone pairs.
To get the steric number, multiply the number of bound atoms in central Xe by the lone pair of electrons.
As a result, the Xe steric number in xenon difluoride is 3+2 = 5.
Looking at the illustration, we can see that the shape will be linear, and the bond angle will be 180 degrees.
If we go into further depth, we can claim that such a structure is constructed to reduce electron repulsion. At initially, the structure is linear due to the single bonds.
Each lone pair then follows, pushing the lateral atoms away from Xe to a certain extent. Finally, the final structure is linear due to the combined impact of all three lone pairs.
Atoms have a probability function called Atomic Orbitals (AO) that tells us where the constituent electrons are located. Atomic orbitals are s, p.d, and f.
Do you aware that inside a molecule, AOs do not form bonds?
They combine to generate hybrid orbitals (for example, sp, sp3), which are involved in bond formation in chemistry.
Hybridization is the term for this procedure.
If we want to learn more about xenon difluoride hybridization, we need to look at the electrical configurations.
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 Xe: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6
Alternatively, Xe: [Kr] 4d10 5s2 5p6
1s2 2s2 2p5 F: 1s2 2s2 2p5
Alternatively, F: [He] 2s2 2p5
As a result, we may examine all of the atomic orbitals of the XeF2 molecule’s atoms.
We must now determine which AO coupled with the other to generate hybridised orbitals.
According to Lewis Structure, xenon contains more than eight electrons in its valence shell (unpaired). As a result, the atom becomes excited, and the conformation becomes 5s2 5p5 5d1.
So this is an sp3d hybridization. In XeF2, two hybrid orbitals are employed to produce sigma bonds (single bonds) (F-Xe-F).
Orbital Diagram of a Molecular
We learn about the Molecular Orbital Theory, a quantum mechanics idea, if we delve a little deeper into chemical bonding and hybridization.
This offers us a notion of the MO diagram, which does not include linear bonds. Rather, we work with the spatial and energetic features of electrons that move around due to the impact of nuclei.
Bonding and antibonding orbitals, sigma and pi bonds, as well as HOMO (Highest Occupied MO) and LUMO (Least Occupied MO) are all discussed here (Lowest Unoccupied MO).
The chemical XeF2 is a polyatomic compound. As a result, we have a lot of atomic orbitals to deal with, which leads to a lot of MOs.
We have 10 electrons around Xe in XeF2, which can lead to resonance and delocalized bonding.
Molecular orbital theory, on the other hand, shows that Xenon Difluoride has eight filled MOs:
Each orbital has two bonding and non-bonding orbitals, as well as two bonding and anti-bonding orbitals.
When MOs are mixed, anti-bonding orbitals form, giving us 11 filled molecular orbitals. We also know that XeF2 follows a 3-center-4-electron bond model.
The noble gas xenon difluoride is a hypervalent halogen compound with an octet rule exception and no net dipole moment.
It is necessary to understand chemical bonding in order to learn about the many qualities and characteristics. The Lewis Structure, Molecular Geometry, Hybridization, and Molecular Orbital Theory-based Diagram of XeF2 have all been explored in this article. Now is the time for you to go over this and gain a better understanding of chemistry.
Good luck with your studies!
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