IF3 Lewis structure

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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.

Steps

Use these steps to correctly draw the IF₃ Lewis structure:

#1 First draw a rough sketch
#2 Mark lone pairs on the atoms
#3 Calculate and mark formal charges on the atoms, if required

Let’s discuss each step in more detail.

#1 First draw a 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

#2 Mark lone pairs on the atoms

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:

#3 Calculate and mark formal charges on the atoms, if required

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.

In the above structure, you can see that the central atom (iodine) forms an octet. And the outside atoms (fluorines) also form an octet. Hence, the octet rule is satisfied.

Therefore, this structure is the stable Lewis structure of IF₃.

Next: XeO₄ Lewis structure

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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/