IF5 Lewis Structure in 5 Steps (With Images)

So you have seen the above image by now, right?

Let me explain the above image in short.

IF5 lewis structure has an Iodine atom (I) at the center which is surrounded by five Fluorine atoms (F). There are 5 single bonds between the Iodine atom (I) and each Fluorine atom (F). There are 3 lone pairs on all the Fluorine atoms (F) and 1 lone pair on the Iodine atom (I).

If you haven’t understood anything from the above image of IF5 (iodine pentafluoride) lewis structure, then just stick with me and you will get the detailed step by step explanation on drawing a lewis structure of IF5.

So let’s move to the steps of drawing the lewis structure of IF5.

Steps of drawing IF5 lewis structure

Step 1: Find the total valence electrons in IF5 molecule

In order to find the total valence electrons in a IF5 (iodine pentafluoride) molecule, first of all you should know the valence electrons present in iodine atom as well as fluorine atom.
(Valence electrons are the electrons that are present in the outermost orbit of any atom.)

Here, I’ll tell you how you can easily find the valence electrons of iodine as well as fluorine using a periodic table.

Total valence electrons in IF5 molecule

→ Valence electrons given by iodine atom:

Iodine is a group 17 element on the periodic table. ^[1] Hence the valence electrons present in iodine is 7.

You can see the 7 valence electrons present in the iodine atom as shown in the above image.

→ Valence electrons given by fluorine atom:

Fluorine is group 17 element on the periodic table. ^[1] Hence the valence electron present in fluorine is 7.

You can see the 7 valence electrons present in the fluorine atom as shown in the above image.

Hence, 

Total valence electrons in IF5 molecule = valence electrons given by 1 iodine atom + valence electrons given by 5 fluorine atoms = 7 + 7(5) = 42.

Step 2: Select the central atom

For selecting the center atom, you have to remember that the atom which is less electronegative remains at the center.

Now here the given molecule is IF5 (iodine pentafluoride) and it contains iodine atom (I) and fluorine atoms (F).

You can see the electronegativity values of iodine atom (I) and fluorine atom (F) in the above periodic table.

If we compare the electronegativity values of iodine (I) and fluorine (F) then the iodine atom is less electronegative.

So here the iodine atom (I) is the center atom and the fluorine atoms (F) are the outside atoms.

Step 3: Connect each atoms by putting an electron pair between them

Now in the IF5 molecule, you have to put the electron pairs between the iodine atom (I) and fluorine atoms (F).

This indicates that the iodine (I) and fluorine (F) are chemically bonded with each other in a IF5 molecule.

Step 4: Make the outer atoms stable. Place the remaining valence electrons pair on the central atom.

Now in this step, you have to check the stability of the outer atoms.

Here in the sketch of IF5 molecule, you can see that the outer atoms are fluorine atoms.

These outer fluorine atoms are forming an octet and hence they are stable.

Also, in step 1 we have calculated the total number of valence electrons present in the IF5 molecule.

The IF5 molecule has a total 42 valence electrons and out of these, only 40 valence electrons are used in the above sketch.

So the number of electrons which are left = 42 – 40 = 2.

You have to put these 2 electrons on the central iodine atom in the above sketch of IF5 molecule.

Now let’s proceed to the next step.

Step 5: Check the stability of lewis structure

Now you have come to the final step in which you have to check the stability of lewis structure of IF5.

The stability of lewis structure can be checked by using a concept of formal charge.

In short, now you have to find the formal charge on iodine (I) atom as well as fluorine (F) atoms present in the IF5 molecule.

For calculating the formal charge, you have to use the following formula;

Formal charge = Valence electrons – (Bonding electrons)/2 – Nonbonding electrons

You can see the number of bonding electrons and nonbonding electrons for each atom of IF5 molecule in the image given below.

For Iodine (I) atom:
Valence electrons = 7 (because iodine is in group 17)
Bonding electrons = 10
Nonbonding electrons = 2

For Fluorine (F) atom:
Valence electrons = 7 (because fluorine is in group 17)
Bonding electrons = 2
Nonbonding electrons = 6

Formal charge	=	Valence electrons	–	(Bonding electrons)/2	–	Nonbonding electrons
I	=	7	–	10/2	–	2	=	0
F	=	7	–	2/2	–	6	=	0

From the above calculations of formal charge, you can see that the iodine (I) atom as well as fluorine (F) atoms have a “zero” formal charge.

This indicates that the above lewis structure of IF5 is stable and there is no further change in the above structure of IF5.

In the above lewis dot structure of IF5, you can also represent each bonding electron pair (:) as a single bond (|). By doing so, you will get the following lewis structure of IF5.

I hope you have completely understood all the above steps.

For more practice and better understanding, you can try other lewis structures listed below.

Try (or at least See) these lewis structures for better understanding:

SCN- lewis structure	ClF3 lewis structure
Cl2 lewis structure	HF lewis structure
SCl2 lewis structure	PF5 lewis structure