SI4 Lewis structure

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SI₄ has one sulfur atom and four iodine atoms.

In SI₄ Lewis structure, there are four single bonds around the sulfur atom, with four iodine atoms attached to it. Each iodine atom has three lone pairs, and the sulfur atom has one lone pair.

Steps

Use these steps to correctly draw the SI₄ 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, sulfur lies in group 16, and iodine lies in group 17.

Hence, sulfur has six valence electrons and iodine has seven valence electrons.

Since SI₄ has one sulfur atom and four iodine atoms, so…

Valence electrons of one sulfur atom = 6 × 1 = 6
Valence electrons of four iodine atoms = 7 × 4 = 28

And the total valence electrons = 6 + 28 = 34

Learn how to find: Sulfur valence electrons

• Second, find the total electron pairs

We have a total of 34 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 = 34 ÷ 2 = 17

• Third, determine the central atom

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

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

Therefore, place sulfur in the center and iodines on either side.

• And finally, draw the rough sketch

#2 Mark lone pairs on the atoms

Here, we have a total of 17 electron pairs. And four S — I bonds are already marked. So we have to only mark the remaining thirteen electron pairs as lone pairs on the sketch.

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

So for each iodine, there are three lone pairs, and for sulfur, there is one lone pair.

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 sulfur atom, formal charge = 6 – 2 – ½ (8) = 0

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

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

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

Therefore, this structure is the stable lewis structure of SI₄.

Next: GaCl₃ Lewis structure

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