MO Diagram, Molecular Geometry, Hybridization, and SF6 Lewis Structure

SF6, or sulphur hexafluoride, is an inorganic gas that is one of the most stable in chemistry. The density of this gas is higher than that of air. Because the gas is odourless and colourless, it cannot be identified in a generic sense. There is also no discernible flavour of gas.

In nature, SF6 is noncombustible and nonflammable. It might, however, break out of its storage container and rocket into the air under extreme heat and pressure.

SF6 can dissociate further and participate in subsequent processes by reacting with a few molecules.

SF6    +    4H2O     ——–>    H2SO4   +    6HF

Sulfuric acid and hydrogen fluoride are formed when SF6 reacts with water.

SF6    +     8NaOH     ——–>     Na2SO4    +    6NaF   +    4H2O

Here, sodium sulphate, sodium fluoride, and water are formed when SF6 interacts with sodium hydroxide.

Otherwise, SF6 is insoluble in water and can only be dissolved in non-polar organic solvents.

Human Reaction:

When humans inhale this gas, it is transported to the lungs, and any skin or eye contact can result in frostbite.

The different properties of SF6 are listed below.

146.06 g/mol molecular weight

-63.8 deg C is the boiling point.

-83 degrees Fahrenheit is the melting point.

The gas is transported as a liquefied gas, compressed by its own pressure, with no extra handling or storage requirements.

Lewis Structure of SF6

By knowing about the Lewis structure of any molecule, we may delve into the depths of understanding how things work behind the scenes.

We examine the nature of bonds, types of bonds, and how these bonds appear while constructing the Lewis structure of any compound.

To create a Lewis structure, you must understand the octet law, valence electrons, and lone pairing thoroughly.

Do you recall what these terms mean?

If not, let’s get together and modify them.

Valence electrons are the electrons that are present in an atom’s outer shell, to put it simply.

These electrons are organised into four energy shells: s, p, d, and f. Valence electrons play a crucial role in shaping the characteristics of atoms.

The Octet Rule: To become stable, every atom tries to follow the octet rule. In nature, you need 8 electrons in the outermost shell to be stable.

This is what every atom strives towards by forming bonds with its neighbours. As a result, various compounds with a wide range of chemical and physical properties are generated.

Returning to the SF6 Lewis structure, let’s locate the valence electrons before drawing the final Lewis dot structure.

SF6 Valence Electrons

In SF6, there are two types of atoms: sulphur and fluorine.

Sulfur has only one atom, but fluorine has six atoms.

Let’s start with the valence electrons of Sulfur and Fluorine to build the Lewis structure.

Sulfur’s valence electrons are 6, and Fluorine’s valence electrons are 7.

Because there are 6 fluorine atoms in this compound, the total valence electrons are = 7*6 = 42 valence electrons.

As a result, the total valence electrons of SF6 are = 42+6 = 48.

Has the thought occurred to you as to why sulphur has six valence electrons?

If that’s the case, then let us provide you with an answer.

Sulfur has an atomic number of 16. This number represents the total number of electrons in this element. If we use the electronic configuration approach to fill these electrons, we will end up with the following arrangement:

1s2 2s2 2p6 3s2 3p4

In the case of Fluorine, we have 6 electrons in the outermost shell, commonly known as valence electrons.

SF6 Lewis Dot Structure

Sulfur is the centre atom because it is less electronegative than Fluorine. This is due to the fact that Fluorine’s outer shell has 5 electrons, and it only requires one more electron to achieve stability, which is easier to achieve.

Because Fluorine has six atoms, there will be six bonds formed between Sulfur and Fluorine. Sulfur will share six electrons with each fluorine atom, one for each.

Here’s a visual illustration of what I’m talking about.

A single straight line represents a bond creation, while two dots represent lone pairs of electrons.

Now that we know how many bonds are produced, let’s look at how many lone pairs are formed.

When six bonds are established, 12 electrons are consumed, leaving 36 valence electrons available.

Did you notice that Sulfur’s outer shell has more than 8 electrons?

This is because some atoms have the ability to enlarge their valence shells. These kinds of situations arise when the outermost shell can accommodate additional electrons.

The remaining electrons on the Fluorine atom will now be lone pairs of electrons. There will be three lone pairs of electrons on each Fluorine atom.

There will be 18 lone pairs of electrons in all, for a total of 36 electrons.

The octet of all Fluorine atoms is complete, with the exception of Sulfur, whose octet is stretched.

Let’s look at the hybridization and molecular geometry of SF6 now that we’ve covered the Lewis Structure.

Hybridization of the SF6 molecule

The Lewis structure is a model of bond creation. However, if you want to know how a compound looks on a plane, you’ll need to know about its hybridization and molecular geometry.

The filling of the electrons in the various energy orbitals is depicted in this diagram.

Because Sulfur has a total of 16 electrons, the shells are filled in different energy levels depending on their capacity and hierarchy level. As a result, the hybridization of Sulfur in the ground state is 3s2 3p4.

The hybridization of fluorine in the ground state is 2s2 2p5.

The electrons of the Sulfur migrate from energy level p to energy level d when these two atoms link together and the combination enters an excited state.

As a result, the configuration is now sp3d2. As a result, the SF6 hybridization is sp3d2.

This occurs because Sulfur’s valence shells can extend to form bonds in order to achieve stability.

SF6’s Molecular Geometry and MO Diagram

When we look at Sulfur on a planar level, there are Fluorine atoms all around it, giving the compound symmetry.

We can see a 3-D image of how the atoms are placed when we look at the molecular geometry of any compound, which we can’t see when building a Lewis Structure.

We can also determine bond angles using molecular geometry. Because the atoms are equally dispersed around the centre element Sulfur, the bond angle for SF6 is 90 degrees.

The VSEPR theory can be used to determine the form of a chemical. This theory is concerned with electron repulsion and the need for compounds to take on a form in order to achieve stability.

There are six sigma bonds between Sulfur and Fluorine in SF6, as well as three lone pairs on each Fluorine atom. These lone electrons oppose each other and keep the centre atom symmetric.

That is why SF6 has an octahedral form, which can also be referred to as its molecular geometry.

The SF6 MO diagram is shown below.

To learn how to build a MO diagram step by step, watch this video.

SF6’s polarity

When we examine a compound’s structure and properties in depth, the polarity of the compound is a prominent topic of debate.

The charge distribution is uniform across the SF6 molecule, resulting in a dipole moment of 0. The dipole moment is cancelled due to the symmetric configuration.

As a result, SF6 is non-polar in nature.

You should also read the page about the polarity of SF6.


Now that we’ve reached the end of this lesson, let’s review what we’ve learned so far about the SF6 chemical in this article:

SF6 is a colourless, odourless gas that is non-flammable and non-combustible.

Sulfur is the core element, which is connected to six fluorine atoms.

On fluorine, the Lewis dot structure has 6 sigma bonds and rests lone pairs.

The SF6 hybridization is sp3d2.

SF6 is non-polar and has an octahedral molecular shape.

We hope you enjoyed learning about this fascinating molecule.

If you have any questions, please contact us and we will gladly answer them.

Thank you for taking the time to read this.

Read more: Molecular Geometry, Hybridization, and Polarity of PBr5 Lewis Structure

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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