SI4 Lewis structure

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.

Alternative method: Lewis structure of SI₄

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

Lone pair

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:

Formal charge

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.

Final structure

The final structure of SI₄ contains a central sulfur atom connected to four iodine atoms through single covalent bonds. Within this layout, the sulfur atom serves as an exception to the octet rule, utilizing an expanded valence shell to accommodate ten electrons, which include four bonding pairs and one lone pair. Each iodine atom fulfills its octet by maintaining three lone pairs of its own alongside the single shared bond. This arrangement is the most stable because it optimizes the formal charge distribution; all atoms in the structure—both the central sulfur and the four iodine atoms—carry a formal charge of zero. Accordingly, this specific electronic distribution serves as the definitive and most accurate Lewis representation of SI₄.

Next: GaCl₃ Lewis structure