The **Lewis structure of IF _{3}** contains three single bonds, with iodine in the center, and three fluorines on either side. There are three lone pairs on each fluorine atom, and two lone pairs on the iodine atom.

## Steps

By using the following steps, you can easily draw the Lewis structure of IF_{3}.

#1 Draw skeleton

#2 Show chemical bond

#3 Mark lone pairs

#4 Calculate formal charge and check stability (if octet is already completed on central atom)

Let’s one by one discuss each step in detail.

### #1 Draw skeleton

In this step, first calculate the total number of valence electrons. And then, decide the central atom.

- Let’s calculate the total number of valence electrons

We know that… both iodine and fluorine are the group 17 elements. Hence, both iodine and fluorine have **seven** valence electrons.

Now IF_{3} has one iodine atom and three fluorine atoms.

So the total number of valence electrons = valence electrons of iodine atom + (valence electrons of fluorine atom × 3)

Therefore, the **total number of valence electrons** = 7 + 21 = 28

- Now decide the central atom

The atom with the least electronegative value is placed at the center. By looking at the periodic table, we get the electronegativity values for iodine and fluorine as follows:

Electronegativity value of iodine = 2.66

Electronegativity value of fluorine = 3.98

Obviously, iodine is less electronegative than fluorine. Hence, assume that **iodine is the central atom**.

So now, put iodine in the center and fluorines on either side. And draw the rough skeleton structure for the Lewis structure of IF_{5} something like this:

**Also read:** How to draw Lewis structure of CNO^{–} (5 steps)

### #2 Show chemical bond

Place two electrons between the atoms to show a chemical bond. Since iodine is surrounded by three fluorines, use six electrons to show **three chemical bonds** as follows:

**Also read:** How to draw Lewis structure of N_{2}O_{4} (5 steps)

### #3 Mark lone pairs

As calculated earlier, we have a total of 28 valence electrons. And in the above structure, we have already used six valence electrons. Hence, twenty-two valence electrons are remaining.

Two valence electrons represent one lone pair. So twenty-two valence electrons = **eleven lone pairs**.

Note that iodine is a period 5 element, so it can keep more than 8 electrons in its last shell. And fluorine is a period 2 element, so it can not keep more than 8 electrons in its last shell.

Also, make sure that you start marking these lone pairs on outside atoms first. And then, on the central atom.

The outside atoms are fluorines, so each fluorine will get three lone pairs. And the central atom (iodine) will get two lone pairs.

So the Lewis structure of IF_{3} looks something like this:

In the above structure, you can see that the octet is completed on the central atom (iodine), and also on the outside atoms. Therefore, the octet rule is satisfied.

After completing the octet, one last thing we need to do is, calculate the formal charge and check the stability of the above structure.

**Also read:** How to draw Lewis structure of PF_{6}^{–} (4 steps)

### #4 Calculate formal charge and check stability

The following formula is used to calculate the formal charges on atoms:

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

Collect the data from the above structure and then, write it down below as follows:

- For
**iodine**atom

Valence electrons = 7

Nonbonding electrons = 2

Bonding electrons = 10

Formal charge = 7 – 2 – ½ (10) = 0

- For
**each fluorine**atom

Valence electrons = 7

Nonbonding electrons = 6

Bonding electrons = 2

Formal charge = 7 – 6 – ½ (2) = 0

Mention the formal charges of atoms on the structure. So the Lewis structure of IF_{3} looks something like this:

In the above structure, you can see that the formal charges of both (iodine and fluorine) are zero. Therefore, this is the **stable Lewis structure of IF**** _{3}**.

And each horizontal line drawn in the above structure represents a pair of bonding valence electrons.

## Related

- Lewis structure of CNO
^{–} - Lewis structure of N
_{2}O_{4} - Lewis structure of PF
_{6}^{–} - Lewis structure of XeO
_{4} - Lewis structure of CO
_{3}^{2-}

## External links

- IF3 Lewis Structure in 5 Steps (With Images) – Pediabay
- IF3 Lewis Structure, Hybridization, Molecular Geometry, and Polarity – Techiescientist
- IF3 Lewis structure, molecular geometry, hybridization, polar or nonpolar – Topblogtenz
- How to Draw the Lewis Dot Structure for IF3 – The Geoexchange
- Who do you determine the Lewis Dot structure of IF3? – Quora
- Draw the Lewis structure for IF3 – Homework.Study.com
- How to draw IF3 Lewis Structure? – Science Education and Tutorials
- Draw the Lewis structure of iodine trifluoride (IF3) that has the least amount of formal charge on the atom – Chegg
- Lewis Dot of Iodine Trifluoride IF3 – Kent’s Chemistry
- Is IF3 a polar or a non-polar molecule? – ECHEMI
- IF3 Lewis Structure (Iodine Trifluoride) – Pinterest
- given the lewis structure of if3 below, what are the approximate bond angles in the molecule? – Brainly
- IF3 Lewis Acid/Base – Reddit

Deep

Rootmemory.com was founded by Deep Rana, who is a mechanical engineer by profession and a blogger by passion. He has a good conceptual knowledge on different educational topics and he provides the same on this website. He loves to learn something new everyday and believes that the best utilization of free time is developing a new skill.