Hybridization, Molecular Geometry, and Polarity of KrF2 Lewis Structure

KrF2, or Krypton difluoride, is one of Krypton’s initial compounds, consisting of Krypton and Fluorine. It’s a colourless, extremely volatile, and thermally unstable solid.

Although it decomposes at normal temperature, it may be kept at -78 degrees Celsius indefinitely. KrF2 forms KrF+ and Kr2F3+ salts when it reacts vigorously with Lewis acids.

KrF2 has a molar mass of 121.795 gmol1.

KrF2 has a density of 3.24 g cm3.

Because of its propensity to oxidise even gold to its +5 oxidation state, KrF2 is largely used as an oxidising and fluorinating agent.

The following methods can be used to make it:

  1. Electrical Discharge – This was Turner and Pimentel’s first method.

Bombardment with Protons No. 2

  1. Using a hotwire
  2. Synthesis of photochemicals

Now that you know what KrF2 is and what its Lewis structure, hybridization, and molecular shape are, it’s time to learn about its Lewis structure, hybridization, and molecular shape.

Lewis Structure of KrF2

We’ll take a look at how the Lewis structure of KrF2 should look before we set down the procedures to create it.

Valence electrons are the electrons in an atom’s shell that are furthest from the nucleus. We’ll draw the Lewis structure using valence electrons as our major guide.

K is surrounded by three lone pairs of electrons in the Lewis structure of KrF2, and it forms solitary bonds with each of the F atoms.

We’ll now look at the stages involved in designing the Lewis structure:-

  1. The total number of valence electrons in the molecule is counted. You can use the periodic table as a reference if you forget the valence electrons of a specific atom.
  2. After that, we look for the molecule’s core atom.
  3. We now begin to arrange these electrons as lone pairs, which represent a chemical link between each atom.
  4. Next, double-check that each atom has completed its octet/duplet. So, bearing the aforementioned point in mind, we arrange the remaining valence electrons.
  5. We can transform lone pairs into double or triple bonds to make the molecule more stable.

While doing so, we should always check each atom’s formal charge to ensure that it is as low as possible.

Each atom in the optimal Lewis structure should have a neutral charge (0). The following is the formula for determining formal charge:-

  1. For the molecule KrF2, count the total number of valence electrons.

Kr is a member of the 8th group. Noble gases in Group 8 have 8 valence electrons because they are highly stable. F belongs to group 7, and there are 72=14 valence electrons in F because it has two atoms.

As a result, the total number of valence electrons in the molecule KrF2 is 8+14= 22.

  1. A centre atom should have the following characteristics:

The valence factor with the highest value.

The place with the most bonding sites.

The core atom in KrF2 is unmistakably Kr.

  1. We now begin to arrange the electrons on each atom as lone pairs in order to build a chemical connection. Only four valence electrons are used up in F because there are only two atoms.
  2. At this point, we begin arranging the remaining valence electrons around each atom to complete its octet. When we’ve completed the preceding process, we’ll find that just 16 valence electrons have been consumed.

As a result, the remaining 6 valence electrons on the core atom Kr will function as lone pairs. This is an example of an octet rule exception.

Kr has more than 8 valence electrons, which is due to the fact that elements with periods less than 3 can have an expanded octet (more than 8 valence electrons), which is an exception to the octet rule. As a result, Kr contains three lone pairs and can hold more than eight valence electrons.

  1. You might be wondering why we didn’t double or triple bond the lone pairs. That is an excellent question!

Remember that in the ideal Lewis structure, each atom should have a charge of 0. When we look at the formal charge of each atom in KrF2, we find that it is zero.

However, when lone pairs are converted to double or triple bonds, the formal charge is not as low as it may be. As a result, the molecule will be extremely unstable.

With each atom having a formal charge of 0, this Lewis structure of KrF2 with three lone pairs provides the best stability.

Hybridization of KrF2

When it comes to understanding the nature of a molecule’s chemical bonds, hybridization is crucial. Hybridization aids in the discovery of a more stable molecule by lowering the molecule’s energy.

KrF2 has a Sp3d hybridization.

A molecule’s hybridization can be discovered using two methods:-

  1. Theoretical approach:

Any molecule’s hybridization can be calculated by multiplying the number of bonded sites by the number of lone pairs of the central atom.

The following factors influence the value of Hybridization (H):

When H=2, the organism is sp hybridised.

If H=3, the hybridization is sp2.

H=4 indicates sp3 hybridization.

H=5 indicates sp3d hybridization.

And H=6 denotes Sp3d2 hybridization.

Kr is the core atom of KrF2, as we well know. It forms one sigma bond with each F atom and is surrounded by three lone pairs. By multiplying the amount of bonds and lone pairs, we get 2+3 = 5, indicating that KrF2 is Sp3d hybridised.

  1. The portion about the formula. Although the theoretical portion is much easier to memorise, the formula can be used to corroborate your solution.

The following is the formula for calculating the Hybridization of any molecule:-


H = Hybridization of the molecule’s core atom

The number of Valence electrons in the core atom is denoted by V.

The number of monovalent atoms bound to the core atom is denoted by the letter M.

C = Charge on a cation or an atom with a higher electropositivity.

A= Charge on an anion or an atom with a higher electropositivity.

We know that Kr is the core atom in KrF2. Thus, V =8. (valence electrons of Kr). M = 2 since F is a monovalent atom and there are two of them.

Both C and A will be 0 because the molecule’s total charge is neutral. As a result of the formula,


KrF2 is Sp3d hybridised, as indicated by H=5.

Using two easy approaches, we were able to determine the Hybridization of KrF2.

Molecular Geometry of KrF2

The molecular shape is used to determine the form of a molecule and its bond angles, as the name implies. The shape of a molecule differs from the geometry of a molecule.

When establishing the geometry of molecules, the electrons are also taken into account.

KrF2 has a trigonal bipyramidal geometry.

The lone pairs on the core atom are taken into consideration by the molecular form. As a result, KrF2 has a linear molecular form. Each atom forms a 180-degree connection with the other.

Any molecule’s molecular shape can be determined using the symbol AXN. The letter A stands for the number of central atoms. The number of atoms bound to the core atom is denoted by the letter X.

Finally, the number of lone pairs or non-bonding electrons in the core atom is denoted by N.

Because Kr is the only central atom in KrF2, A=1.

Because there are two F atoms linked to the centre atom, X=2.

Because there are three lone pairs on the core atom Kr, N will be three. As a result of the preceding formula, the shape of KrF2 is AX2N3.

We can see that KrF2 has a linear shape if we look at it in the VSEPR graphic.

Polarity of KrF2

As previously stated, the core atom of the KrF2 molecule is connected linearly to two F atoms on both sides, giving a symmetrical shape. Furthermore, the electronegativity of both F atoms is the same.

As a result, the charge is pulled equally by both F atoms, resulting in an equal charge distribution. Both dipoles cancel each other out by acting in opposite directions. As a result, the dipole moment is net-zero.

As a result, the KrF2 molecule is non-polar.


The Lewis structure, hybridization, and molecular shape of KrF2 have all been explored in this article. So now you should have a good understanding of the KrF2 molecule’s fundamentals.

Please feel free to contact me if you have any questions about any of the points.

Study hard!

Read more: Bohr Model of Oxygen (Diagram, Steps To Draw)

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.


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