SI4 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.

SI4 lewis structure has a Sulfur atom (S) at the center which is surrounded by four Iodine atoms (I). There are 4 single bonds between the Sulfur atom (S) and each Iodine atom (I). There is 1 lone pair on the Sulfur atom (S) and 3 lone pairs on all the four Iodine atoms (I).

If you haven’t understood anything from the above image of SI4 lewis structure, then just stick with me and you will get the detailed step by step explanation on drawing a lewis structure of SI4.

So let’s move to the steps of drawing the lewis structure of SI4.

Steps of drawing SI4 lewis structure

Step 1: Find the total valence electrons in SI4 molecule

In order to find the total valence electrons in SI4 molecule, first of all you should know the valence electrons present in sulfur atom as well as iodine 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 sulfur as well as iodine using a periodic table.

Total valence electrons in SI4 molecule

→ Valence electrons given by sulfur atom:

Sulfur is a group 16 element on the periodic table. ^[1] Hence the valence electrons present in sulfur is 6.

You can see the 6 valence electrons present in the sulfur atom as shown in the above image.

→ Valence electrons given by iodine atom:

Iodine is a group 17 element on the periodic table. ^[2] 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.

Hence, 

Total valence electrons in SI4 molecule = valence electrons given by 1 sulfur atom + valence electrons given by 4 iodine atoms = 6 + 7(4) = 34.

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 SI4 and it contains sulfur atom (S) and iodine atoms (I).

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

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

So here the sulfur atom (S) is the center atom and the iodine atoms (I) are the outside atoms.

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

Now in the SI4 molecule, you have to put the electron pairs between the sulfur atom (S) and iodine atoms (I).

This indicates that the sulfur (S) and iodine (I) are chemically bonded with each other in a SI4 molecule.

Step 4: Make the outer atoms stable

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

Here in the sketch of SI4 molecule, you can see that the outer atoms are iodine atoms.

These outer iodine 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 SI4 molecule.

The SI4 molecule has a total 34 valence electrons and out of these, only 32 valence electrons are used in the above sketch.

So the number of electrons which are left = 34 – 32 = 2.

You have to put these 2 electrons on the central sulfur atom in the above sketch of SI4 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 SI4.

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 sulfur (S) atom as well as iodine (I) atoms present in the SI4 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 SI4 molecule in the image given below.

For Sulfur (S) atom:
Valence electrons = 6 (because sulfur is in group 16)
Bonding electrons = 8
Nonbonding electrons = 2

For Iodine (I) atom:
Valence electron = 7 (because iodine is in group 17)
Bonding electrons = 2
Nonbonding electrons = 6

Formal charge	=	Valence electrons	–	(Bonding electrons)/2	–	Nonbonding electrons
S	=	6	–	8/2	–	2	=	0
I	=	7	–	2/2	–	6	=	0

From the above calculations of formal charge, you can see that the sulfur (S) atom as well as iodine (I) atom has a “zero” formal charge.

This indicates that the above lewis structure of SI4 is stable and there is no further change in the above structure of SI4.

In the above lewis dot structure of SI4, you can also represent each bonding electron pair (:) as a single bond (|). By doing so, you will get the following lewis structure of SI4.

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:

NSF Lewis Structure	C2H4Br2 Lewis Structure
CO2 lewis structure	SO2 lewis structure
NO2 lewis structure	NH3 lewis structure