IF3 Lewis structure

IF₃ (iodine trifluoride) has one iodine atom and three fluorine atoms.

In the IF₃ Lewis structure, there are three single bonds around the iodine atom, with three fluorine atoms attached to it. Each fluorine atom has three lone pairs, and the iodine atom has two lone pairs.

Alternative method: Lewis structure of IF₃

Rough sketch

• First, determine the total number of valence electrons

In the periodic table, both iodine and fluorine lie in group 17.

Hence, both iodine and fluorine have seven valence electrons.

Since IF₃ has one iodine atom and three fluorine atoms, so…

Valence electrons of one iodine atom = 7 × 1 = 7
Valence electrons of three fluorine atoms = 7 × 3 = 21

And the total valence electrons = 7 + 21 = 28

Learn how to find: Fluorine valence electrons

• Second, find the total electron pairs

We have a total of 28 valence electrons. And when we divide this value by two, we get the value of total electron pairs.

Total electron pairs = total valence electrons ÷ 2

So the total electron pairs = 28 ÷ 2 = 14

• Third, determine the central atom

We have to place the least electronegative atom at the center.

Since iodine is less electronegative than fluorine, assume that the central atom is iodine.

Therefore, place iodine in the center and fluorines on either side.

• And finally, draw the rough sketch

Lone pair

Here, we have a total of 14 electron pairs. And three I — F bonds are already marked. So we have to only mark the remaining eleven electron pairs as lone pairs on the sketch.

Also remember that iodine is a period 5 element, so it can keep more than 8 electrons in its last shell. And fluorine is a period 3 element, so it can keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are fluorines.

So for each fluorine, there are three lone pairs, and for iodine, there are two lone pairs.

Mark the lone pairs on the sketch as follows:

Formal charge

Use the following formula to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

For iodine atom, formal charge = 7 – 4 – ½ (6) = 0

For each fluorine atom, formal charge = 7 – 6 – ½ (2) = 0

Here, both iodine and fluorine atoms do not have charges, so no need to mark the charges.

Final structure

The final structure of IF₃ contains a central iodine atom linked to three fluorine atoms through single covalent bonds. In this configuration, the iodine atom utilizes an expanded valence shell to accommodate ten electrons, which include three bonding pairs and two lone pairs. Within this layout, each of the three fluorine atoms satisfies the octet rule by maintaining three lone pairs alongside its single shared bond. This arrangement represents the most stable state for the molecule because it results in a formal charge of zero for every atom involved. Accordingly, this specific electronic distribution serves as the definitive and most accurate Lewis representation of iodine trifluoride.

Next: XeO₄ Lewis structure

External video

• IF3 Lewis Structure: How to Draw the Lewis Structure for IF3 – YouTube • Wayne Breslyn

External links

• https://techiescientist.com/if3-lewis-structure/
• https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/IF3-lewis-structure.html
• https://topblogtenz.com/if3-lewis-structure-molecular-geometry-hybridization-polar-or-nonpolar/
• https://sciedutut.com/if3-lewis-structure/