Molecular Form, Lewis Structure, Hybridization, and Polarity of IF5

In the realm of chemistry, IF5 is a crucial molecule. Several examinations have included questions about the molecule. Among the most frequently asked questions is, “How is it formed?” Is this molecule enduring? Exist any applications for this compound? Draw this molecule’s form, etc. Numerous facts make it an exception in numerous instances.

I am here to describe all of these details and to teach you everything there is to know about this molecule, including its structure, hybridization, and form, among other things.

I hope that at the conclusion of this post, you will be able to answer any question about IF5!

IF5 (Iodine pentafluoride) is a colourless liquid, however some impure samples may have a yellow hue. Henri Moissan discovered it in 1891 by burning solid iodine in fluorine gas. The following is the response:

I2 + 5F2 ——> 2IF5

It is an interhalogen molecule utilised in organic synthesis as a fluorination reagent. Iodine pentafluoride (IF5) is widely employed in the synthesis of fluoride-containing alkyl iodides, which serve as intermediates in the production of perfluoro-organic combinations.

These are then used to produce water- and oil-repellent emulsions for textile treatment and fire extinguishing foams.

IF5 possesses a molar mass of 221.89 g/mol and a density of 3.250 g/cm3.

Boiling Point- 97.85 °C

Point de fusion – 9.43 °C

Lewis Structure IF5

Before we begin constructing the Lewis structure for IF5, there are several factors to consider.

A Lewis structure essentially represents the number of an atom’s valence electrons.

Counting the number of columns of a periodic table from left to right, ignoring transition elements, is the simplest approach to determine the number of valence electrons of an atom.

For instance, Carbon (C) has four valence electrons, but Fluorine (F) has seven.

In this situation, the single exception is Helium (He), which possesses two valence electrons. Atoms strive for stability through the octet rule, which states that each atom should be surrounded by eight electrons.

Valence electrons are depicted by dots in lewis diagrams. Therefore, while drawing the Lewis diagram for the Cl atom, seven dots are drawn around it.

Now we will discuss the stages involved in making a Lewis diagram:

The initial step is to count each molecule’s valence electrons.

With regard to IF5,

Iodine contains seven valence electrons. F has seven valence electrons as well. However, since there are 5 F atoms, 75 = 35 valence electrons.

Adding both we get 35+7= 42.

Therefore, IF5 has a total of 42 valence electrons.

Identifying the core atom

Typically, each atom is bonded to a centre atom. Typically, there is only one of it in a molecule. In the case of IF5, the core atom is I (Iodine).

Draw individual bonds to the core atom. A single bond reflects the sharing of two electrons.

  1. Place all remaining valence electrons as lone pairs on the atoms (2 dots).
  2. Transform the lone pairs into double or triple bonds so that each atom has an octet.

Examine each atom’s formal charge as a final resort. It should be as low as feasible and can be computed using the following formula.

In the case of IF5, it is evident that I had not only completed its octet but also surpassed it by two electrons. We shall reach that section shortly.

Now, the remaining valence electrons are arranged around each atom to complete their octet. We then discover that there are still two valence electrons remaining.

We position these as lone pairs on the centre atom I. Consequently, the Lewis structure of IF5 has been completed. Observe that the atom of iodine is surrounded by 12 valence electrons.

This is an exception to the octet rule known as the expanded octet rule, which stipulates that central atoms in the third period or below that are connected to strongly electronegative terminal atoms (F) can have an extended octet with up to 12 valence electrons.

Hence I can store a maximum of 12 valence electrons. Now, if you examine the formal charge of each atom, you will find that it is 0, which is the minimum conceivable value.

Consequently, the Lewis structure of any molecule may be created using these straightforward procedures!

IF5 Hybridization

Sp3d2 is the hybridization of IF5.

A molecule’s hybridization can be understood in two ways:-

The theoretical portion – Iodine’s core atom possesses seven valence electrons. In its ground state, its electrical configuration consists of –

I = 5s2, 5p5, 5d10

Seven valence electrons are evenly divided between the s and p orbitals.

Now in its excited state, because I must establish 5 bonds with F, two valence electrons from the p orbital jump (are promoted) to the d orbital, thereby forming 5 bonds with F, while the electrons remaining in the s orbital act as lone pairs (non-bonding), resulting in the hybridization of all electrons.

Consequently, hybridization is determined by tallying the number of formed bonds and lone pairs.

It is Sp3d2 hybridised due to the lone pair in the s orbital, three sigma bonds in the p orbital, and two sigma bonds in the p orbital.

Following is the formula for determining hybridization:

H= 1/2[V+M-C+A]

H= Hybridization

V equals the number of valence electrons

Charge on a cation or electropositive element.

A equals the charge on an anion or a more electronegative element.

Now, if H equals 2, it demonstrates Sp hybridization.

H= 3, which suggests Sp2 hybridization.

H= 4 will demonstrate Sp3 hybridization

H=5 indicates Sp3d hybridization.

H=6, the molecule will hybridise with Sp3d2.

With regard to IF5,

V = 7 (valence electrons of central atom)

M= 5 (5 monovalent atoms of F)

Since IF5’s total charge is zero, C and A will also be zero.



H=6, which indicates that Sp3d2 hybridised. Consequently, we can easily identify the hybridization of IF5 with these two techniques.

ICl5 and IF5 are extremely similar molecules. Additionally, I have produced an article regarding ICl5. Consult the article on ICl5’s Lewis structure.

IF5 Molecular Geometry

The atomic structure of IF5 is square pyramidal. The produced bond angles are close to 90 degrees.

Using VSEPR theory, the molecular geometry of IF5 may be derived. Below is a chart showing the VSEPR.

A denotes the number of central atoms, X represents the number of atoms bound to the central atom (5), and E represents the non-bonding electrons (lone pair).

Consequently, its molecular form is pyramidal and its electron geometry is octahedral.

When drawing the shape, molecular shape displays only the atoms, whereas electron geometry depicts all electron pairs, making it octahedral.

Due to the presence of a lone pair, IF5 has a slightly bowed square pyramidal shape, as depicted in the image below.

Polarity of

IF5 is a Molecule of Polarity.

Polarity develops when a difference in the electronegativity of two linked atoms causes an electric dipole moment.

This is evident while examining its Molecular Geometry. Due to lone pair and bond pair repulsion, IF5 has a square pyramidal form that is bowed.

Observe that the four bonds of IF5 cancel each other out, leaving only one bond.

Fluorine being more electronegative than I induces a dipole moment, making IF5 a polar molecule.

For additional information, please consult IF5 Polarity.


As stated previously, in order to answer any question pertaining to IF5, it is necessary to understand its Lewis Structure, Hybridisation, Polarity, and Molecular Shape. These themes have all been presented in a straightforward and concise manner.

You should now feel comfortable answering any queries concerning IF5 that may arise!

I hope all your questions have been answered, but if you have any questions on any of the aforementioned principles, feel free to contact me. Happy Studying!

Read more: Polarity, HCN Lewis Structure, Molecular Geometry, Hybridization, and MO Diagram

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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