Xenon difluoride (XeF2) Lewis dot structure, molecular geometry
or shape, electron geometry, bond angle, formal charge,
hybridization
XeF2 is the chemical formula for xenon difluoride. It appears as a colorless-to-white crystalline solid at r.t.p. This chemical compound is not so commonly discussed yet it holds importance in the chemistry laboratory as an oxidizing and fluorinating agent.
If you were searching about the chemistry of XeF2 and couldn’t find satisfaction anywhere online, then this is the right place for you.
In this article, we have discussed the Lewis structure of XeF2, its molecular geometry or shape, electron geometry, bond angle, formal charges, hybridization, etc., In short, everything you need to know about the xenon difluoride (XeF2) molecule.
The Lewis structure of xenon difluoride (XeF2) consists of a xenon (Xe) atom at the center. It is bonded to 2 atoms of fluorine (F), one on each side of the molecule. There are a total of 5 electron pairs in this Lewis structure. Out of the 5 electron pairs, 2 are bond pairs while there are 3 lone pairs present on the Xe atom.
Drawing the Lewis structure of XeF2 is quite easy if we follow all the simple steps given below one at a time.
Steps for drawing the Lewis dot structure of XeF2
1. Count the total valence electrons in XeF2
The Lewis dot structure of a molecule is referred to as a simplified representation of all the valence electrons present in it. Therefore, the very first step while drawing the Lewis structure of XeF2 is to count the total valence electrons present in the concerned elemental atoms.
The valence electrons present in an atom of an element can be easily determined by identifying the element in the Periodic Table. In XeF2, there are two different elemental atoms i.e., xenon (Xe) and fluorine (F).
If you go through the Periodic Table of elements, you will readily identify that Xenon (Xe) is a noble gas present in Group VIII A. So, Xe has a total of 8 valence electrons. On the other hand, Fluorine (F) is a halogen, located in Group VII A of the Periodic Table thus it has a total of 7 valence electrons.
∴ The XeF2 molecule is made up of 1 Xe atom and 2 atoms of F. Hence, the total valence electrons available for drawing the Lewis structure of XeF2 = 8 + 2(7) = 22 valence electrons.
2. Find the least electronegative atom and place it at the center
The second step is to identify a central atom that is most likely to share its electrons with the atoms in its surroundings. Electronegativity is defined as the ability of an atom to attract a shared pair of electrons towards itself.
Therefore, the least electronegative or most electropositive atom is chosen as the central atom considering its greater tendency to share its electrons with the surrounding atoms.
In the case of XeF2, there are two different atoms involved. Fluorine (F) is the most electronegative element of the Periodic Table while xenon (Xe) is less electronegative.
So Xe is placed at the center of the XeF2 Lewis structure while the F atoms are situated in its surroundings, as the outer atoms. Refer to the figure below.
3. Connect outer atoms with the central atom
In this step, we join the outer atoms present in the Lewis structure of a molecule with the central atom using single straight lines.
As we already identified, Xe is the central atom in XeF2 while F atoms are the outer atoms. So, both the F atoms are joined to the central Xe using straight lines.
Each straight line represents a single covalent bond i.e., a bond pair containing 2 electrons. There are a total of 2 single bonds in the above diagram, so the total valence electrons used so far are 2(2) = 4 valence electrons.
Total valence electrons available – electrons used till step 3 = 22 – 4 = 18 valence electrons.
This means 18 valence electrons are still available which we need to accommodate somewhere in the Lewis structure of XeF2.
4. Complete the octet of outer atoms
Now we need to complete the octet of the atoms surrounding the central Xe atom. The F atoms are the outer atoms surrounding Xe in the XeF2 molecule. Each fluorine (F) atom requires a total of 8 valence electrons to achieve a stable octet electronic configuration.
Both the F atoms are bonded to the Xe atom using single bonds. This means each F atom already contains 2 electrons and it has a deficiency of 6 electrons which are required to complete its octet shell.
Thus, 6 valence electrons are placed as 3 lone pairs around each F atom, as shown in the figure below.
5. Complete the octet of the central atom
Let’s again count the total valence electrons used so far.
Total valence electrons used till step 4 = 2 single bonds + 2 (electrons placed around each F atom, shown as dots) = 2 (2) + 2(6) = 16 valence electrons.
Total valence electrons available – electrons used till step 4 = 22 -16 = 6 valence electrons.
These 6 valence electrons are consequently placed as 3 lone pairs on the central xenon atom. In this way, the Xe atom has a total of 10 valence electrons (2 single bonds + 3 lone pairs) around it.
Xenon (Xe) does not follow the octet rule. It has more than 8 valence electrons in the Lewis structure of XeF2. This is possible because the Xe atom has an expanded octet owing to the availability of d-subshells in its atomic structure.
During chemical bonding, the incoming electrons have access to the 4d subshell which facilitates the accommodation of more than 8 valence electrons.
In the final step, we just need to check the stability of the Lewis structure obtained till step 5. Let’s do that using the formal charge concept.
6. Check the stability of Lewis’s structure with the help of the formal charge concept
The less the formal charge on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charge can be calculated using the formula given below.
Let’s use this formula and the Lewis structure obtained in the previous step to determine whether any formal charges are present on the XeF2 atoms or not.
Zero formal charges present on all the atoms in the XeF2 molecule ensure the stability of its Lewis structure. In conclusion, we have drawn it correctly therefore we are good to proceed to the next section of this article.
What are the electron and molecular geometry of XeF2?
The ideal electron geometry of the xenon difluoride (XeF2) molecule is trigonal bipyramidal. But the lone pairs of electrons situated on the central Xe atom make the molecule adopt a different shape or molecular geometry from its electron geometry i.e., linear.
The lone pairs occupy positions on the molecule farthest away from one another in order to minimize their repulsive effect in the linear shape of XeF2.
Molecular geometry of XeF2
The XeF2 molecule has a symmetrical linear shape and molecular geometry. There are 2 bond pairs around the central Xe atom in the XeF2 molecule. 3 lone pairs are also present on this Xe atom. Therefore, lone pair-lone pair and lone pair-bond pair repulsions exist in the molecule in addition to bond pair-bond pair repulsions.
But as lone pair-lone pair repulsions > lone pair-bond pair repulsions > bond pair-bond pair repulsions thus the F-Xe-F bonded atoms lie on a straight line while the 3 lone pairs are situated above and below the molecule, as far apart from one another as possible.
It is important to note that the molecular geometry or shape of a molecule depends on the different number of lone pairs and bond pairs present on the central atom in a molecule.
Contrarily, the ideal electron geometry of a molecule stays unaffected by any such distinction, it merely depends on the total electron density regions (bond pairs + lone pairs) around the central atom.
Let’s discuss how this concept applies to the XeF2 molecule.
Electron geometry of XeF2
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electronic geometry of a molecule containing 5 regions of electron density around the central atom is trigonal bipyramidal.
As there are 5 electron pairs around the central Xe atom in the XeF2 molecule including the 2 bond pairs and 3 lone pairs thus the electron geometry of XeF2 is trigonal bipyramidal.
The trigonal bipyramidal geometry represents a triangular base and two pyramids, one on top and the other below the triangular base as shown in the figure below.
A more straightforward way of finding the electron and the molecular geometry or shape of a molecule such as XeF2 is to use the AXN method.
AXN is a simple formula to represent the number of atoms bonded to the central atom in a molecule and the lone pairs present on it.
It is used to predict the shape and geometry of a molecule based on the VSEPR concept.
AXN notation for XeF2 molecule
The symbol A represents the central atom present in a molecule. In the XeF2 molecule, xenon (Xe) is present at the center thus A=Xe for XeF2.
X denotes the atoms directly bonded to the central atom. As per the Lewis structure of XeF2, 2 fluorine (F) atoms are bonded to the central Xe atom so X=2 for XeF2.
N stands for the number of lone pairs situated on the central atom. The Lewis structure of XeF2 tells us that N=3 for XeF2.
In short, the AXN generic formula for the XeF2 molecule is AX2N3.
Now, use the VSEPR chart given below to determine the molecular and electron geometry assigned against the AX2N3 generic formula.
The VSEPR chart given above indicates that an AX2N3 formula represents molecules that have a trigonal bipyramidal electron geometry while their molecular geometry or shape is linear, in line with what we noted down for the XeF2 molecule.
Hybridization of XeF2
The central Xe atom is sp3d hybridized in the XeF2 molecule.
The electronic configuration of xenon (Xe) is 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6.
During chemical bonding, one of the three 5p electron pairs gets unpaired. Consequently, one 5p electron gets excited and shifts to the empty 5d atomic orbital. The 5s orbital, three 5p orbitals, and one 5d atomic orbital mix to yield five sp3d hybrid orbitals.
Three of these five sp3d hybrid orbitals contain paired electrons. These electron pairs are situated as 3 lone pairs on the central Xe atom. The remaining two sp3d hybrid orbitals contain one electron each. These sp3d hybrid orbitals overlap with the p orbitals of F on each side of the molecule to form 2 sigma bonds, as shown below.
A shortcut for finding the hybridization present in a molecule is to memorize the table given below. You can use the steric number of the molecule against this table to determine the required hybridization.
There are a total of 5 regions of electron density around the central Xe atom in the XeF2 molecule, so its steric number is 5. Hence, the table shown below tells us that XeF2 has sp3d hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The XeF2 bond angles
All three atoms involved in bond formation in the XeF2 molecule lie on a straight line in a perfectly symmetrical manner. Thus, the F-Xe-F bond angle is 180°.
The sum of all angles on a straight line is equal to 180°.
Both the Xe-F bond lengths are equal in the XeF2 molecule i.e., approximately 198 pm.
A large electronegativity difference of 1.38 units exists between the extremely electronegative fluorine (E.N = 3.98) atom and the less electronegative xenon (E.N = 2.60) atom.
According to Pauling’s electronegativity scale, a polar bond has an electronegativity difference of greater than 0.5 units between its bonded atoms.
Thus, both the Xe-F bonds present in the XeF2 molecule are individually polar. Each Xe-F bond possesses a specific dipole moment value (symbol μ).
However, it is due to the symmetrical linear shape of the molecule that the dipole moments of Xe-F bonds get canceled in opposite directions. In conclusion, the xenon difluoride molecule overall is non-polar with a net μ=0. The electron cloud stays evenly distributed over the molecule.
How many bond pairs and lone pairs are present in the Lewis structure of XeF2?
The Lewis structure of XeF2 consists of a total of 22 valence electrons i.e., 11 electron pairs.
Out of these 11 electron pairs, 2 are bond pairs present between the bonded Xe and F atoms on each side of the Lewis structure.
There are a total of 9 lone pairs. Three lone pairs are present on the central Xe atom while 3 lone pairs are present on each F atom surrounding the central atom.
Is the shape and electron geometry of XeF2 the same?
No, the xenon difluoride (XeF2) molecule has a different shape and electron geometry.
The ideal electronic geometry of XeF2 is trigonal bipyramidal while its shape is linear.
Can the geometry of XeF2 be called distorted?
The molecular geometry or shape of XeF2 is different from its ideal electron geometry therefore we can call it distorted.
The presence of lone pairs on the central Xe atom distorts the geometry of the molecule and makes it occupy a shape different from that which is expected.
What is the bond angle of XeF2?
The XeF2 molecule has an F-Xe-F bond angle of 180°. The three bonded atoms lie on a straight line in a linear arrangement and the sum of all angles on a straight line is always equal to 180°.
How many sigma bonds does xenon form with fluorine atoms in the XeF2 molecule?
There are two Xe-F single bonds, one on each side of the XeF2 molecule. Each of these two single bonds is a sigma bond formed by sp3d-p overlap thus there are 2 sigma bonds in XeF2.
The total valence electrons available for drawing the XeF2 Lewis structure are 22.
The molecular geometry or shape of XeF2 is linear.
The ideal electron geometry of XeF2 is trigonal bipyramidal.
The presence of three lone pairs on the central Xe atom in the XeF2 molecule heavily influences its shape and consequently makes the molecule adopt a different shape from its ideal electronic geometry.
The central Xe atom is sp3d hybridized in the XeF2
The F-Xe-F bond angle is 180° due to the linear arrangement of atoms in the molecule.
Both the Xe-F bond lengths are equal i.e., 198 pm.
Each Xe-F bond is individually polar due to an electronegativity difference > 0.5 units between the bonded atoms. However, it is due to the symmetrical linear arrangement of atoms in the molecule that Xe-F dipole moments get canceled and the molecule overall is non-polar (net μ=0).
Zero formal charges present on both Xe, and F atoms reaffirm the extraordinary stability of the xenon difluoride Lewis structure.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
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