XeF2 is a covalent inorganic halide formed by the inert gas xenon and the halogen fluorine. This is an active solvent and is found to be soluble in different fluorides like HF and bromine pentafluoride.
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If we look at the process of synthesis of xenon difluoride, here’s the equation:
Xe + F2 ——Heat——> XeF2
XeF2 acts as an oxidizing and fluorinating agent and is used to oxidize different hydrocarbons including both aromatic and acyclic compounds.
Not only this, but this fluoride compound can also be used to etch silicon to form silicon tetrafluoride (SiF4) without any external energy application.
If you are thinking about what XeF2 looks like, it appears as a colorless-to-white crystalline solid with a density of around 4.32 g/cc.
This halide can cause some serious hazards like skin burns and major eye damage. Not only this, if inhaled or swallowed, it turns out to be fatal.
When two or more atoms come together they react and combine to form homogeneous and heterogeneous molecules. This formation of molecules happens via the creation of certain bonds which hold the atoms together according to their strength.
This is known as chemical bonding which is the backbone to define the internal structure and nature of a given molecular compound including the properties it exhibits( both physical and chemical).
Before we jump directly into the chemical bonding of XeF2 in this article, we would like you to learn and recapitulate certain important terminologies and concepts.
The very first thing to learn when we are into chemical bonding is the concept of valence electrons. Valence electrons signify the outermost shell electrons of an atom that determine its valency.
If you go through the periodic table, the groups will help us find out the valence electron number of a certain atomic element.
For example, Carbon is in group 4 (also called Group XIV), hence the valence shell consists of four electrons.
XeF2 Lewis Structure
Lewis Structure, also known as electron dot structure, is an essential model of chemical bonding where we use the valence electron concept to schematically sketch a two-dimensional figure of a given molecule.
We use dots to represent outer shell electrons and lines to represent the bond type.
Xenon is an inert gas element. Therefore, it has eight valence electrons. Fluorine is a halogen belonging to group VII, therefore it has a valency of seven.
Total number of valence electrons =8 + 7*2 = 22
If you go by the periodic table, we know that xenon is less electronegative than fluorine, hence it will take up the position of a central atom.
Now that we have the central atom and the total valence electron number, we will find out how the electrons are located to reach the octet configuration( i.e. eight electrons outside each atom in its outer shell)
Now, as we can see we have achieved the octet fulfillment for the three atoms inside XeF2 and single bond formation has been done.
But if you count the total valence electrons in the diagram, we can see the dots add up to 20 and not 22.
Now, the obvious choice is to form a double bond with the remaining two valence electrons instead of single bonds.
Fluorine is so electronegative it doesn’t usually form double bonds. So, we keep the valence electrons around xenon and calculate the formal charge.
If we get to assume that every electron is shared equally among atoms, we can assign a formal charge to an atom.
This is a necessary topic for chemical bonding, especially for Lewis Structure determination where we need to check the least possible formal charges of each combining atom to get the perfect diagrammatic representation.
For Xe, formal charge= 8 (valence electron number) – 0.5*4 (number of bonded electrons) – 6 (no of lone pair electrons) = 0
For each fluorine atom, formal charge= 7 – 0.5*2 – 6 = 0
Therefore, since both the elements are in their least possible formal charge values, we have got our most suitable Lewis Structure.
Note: This is an exception to the octet rule since xenon has more than eight electrons in this compound in its valence shell cloud.
If you want to dive deeper into how the inside of a molecule looks like, we need to learn about molecular geometry.
This is a jump from 2D to 3D structure representation where we can visually picture how a molecule actually remains in bonding nature in reality.
Here, to help us determine the correct molecular shape as well as the bond lengths and angles, we use a theory called Valence Shell Electron Pair Repulsion Model (VSEPR).
This theory deals with the like charge repulsion occurring between negative electron clouds surrounding atomic nuclei and operates to minimize them to get a near stable compound.
Xenon, being a noble gas and belonging to group 8 ( group XVIII), does not tend to form bonds as such. Here, in XeF2, it acts as the central atom and also forms two single bonds with the two fluorine atoms.
Now, we have discussed earlier that molecular geometry can be predicted via VSEPR theory. Let us look at this diagram:
Well, here we can find the different shapes of molecules according to their steric number and lone pair number.We can see via Lewis Structure that Xenon in XeF2 has three lone pairs.
Now for steric number, we need to add the number of bonded atoms to central Xe as well as the lone pair of electrons.
Therefore, Steric number of Xe in xenon difluoride = 3+2 = 5.
If we look at the picture, we can find out that the shape will be linear and the bond angle therefore 180 degrees.
If we discuss in detail, we can say that to minimize electron repulsion, such a structure is formed. The single bonds make the structure linear at first.
This is followed by each lone pair that pushes the lateral atoms away from Xe to a certain degree. Finally, the combined impact of all the three lone pairs makes the resultant structure linear.
Atoms have a probability function called Atomic Orbitals (AO) which gives us an idea about the location of constituent electrons. We have s,p.d,f atomic orbitals.
Do you know that AOs do not form bonds inside a molecule?
They fuse to form what is known as hybrid orbitals ( example sp, sp3) which take part in bond formation in chemistry.
This process is known as hybridization.
If we need to learn about the hybridization of xenon difluoride, we have to look at the respective electronic configurations.
Xe: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6Or, Xe:
F: 1s2 2s2 2p5Or, F:
So, we can look at all the atomic orbitals of the atoms of the XeF2 molecule.
Now, we have to find out which AO combined with the other AO to form hybridized orbitals.
Now, here xenon has more than eight electrons in its valence shell( unpaired) as per Lewis Structure. So, this makes the atom excited and the configuration now has 5s2 5p5 5d1.
So, the hybridization here is sp3d. Two hybrid orbitals are used for sigma bond formation( single bond) in XeF2 (F-Xe-F).
Molecular Orbital Diagram
If we go a little further into chemical bonding and hybridization, we get to know about the Molecular Orbital Theory, a concept of quantum mechanics.
This gives us an idea about the MO diagram where we do not consider linear bonds as such. Instead, we work with the spatial and energetic properties of electrons who happen to move around under nuclei’s influence.
Here, we involve the concepts of bonding and antibonding orbitals, sigma and pi bonds as well as HOMO (Highest Occupied MO) and LUMO (Lowest Unoccupied MO).
XeF2 is a polyatomic molecule. Therefore, we have many atomic orbitals to deal which results in a large number of MOs.
In XeF2 we have ten electrons around Xe which result in possible resonance and delocalized bonding.
However, when we are looking at molecular orbital theory, we can see that Xenon Difluoride has 8 filled MOs:2