ICl2- Lewis structure

ICl₂^– (iodine dichloride) has one iodine atom and two chlorine atoms.

In the ICl₂^– Lewis structure, there are two single bonds around the iodine atom, with two chlorine atoms attached to it, and each atom has three lone pairs.

Also, there is a negative (-1) charge on the iodine atom.

Alternative method: Lewis structure of ICl₂^–

Rough sketch

• First, determine the total number of valence electrons

In the periodic table, both iodine and chlorine lie in group 17.

Hence, both iodine and chlorine have seven valence electrons.

Since ICl₂^– has one iodine atom and two chlorine atoms, so…

Valence electrons of one iodine atom = 7 × 1 = 7
Valence electrons of two chlorine atoms = 7 × 2 = 14

Now the ICl₂^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 7 + 14 + 1 = 22

Learn how to find: Chlorine valence electrons

• Second, find the total electron pairs

We have a total of 22 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 = 22 ÷ 2 = 11

• Third, determine the central atom

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

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

Therefore, place iodine in the center and chlorines on either side.

• And finally, draw the rough sketch

Lone pair

Here, we have a total of 11 electron pairs. And two I — Cl bonds are already marked. So we have to only mark the remaining nine electron pairs as lone pairs on the sketch.

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

So for each atom, there are three lone pairs.

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 iodine atom, formal charge = 7 – 6 – ½ (4) = -1

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

Here, the iodine atom has a charge, so mark it on the sketch as follows:

Final structure

The final structure of ICl₂^– comprises a central iodine atom connected to two chlorine atoms through single covalent bonds. Within this layout, the iodine atom serves as an exception to the octet rule, utilizing an expanded valence shell to accommodate ten electrons through two bonding pairs and three lone pairs. Each chlorine atom satisfies its octet by retaining three lone pairs alongside its single shared bond with the central iodine. This arrangement is the most stable because it minimizes formal charges across the structure, representing the most energetically favorable state for the species. Thus, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the ion.

To properly represent this as a polyatomic ion, the entire Lewis structure is enclosed within square brackets. The overall charge of 1- is then written as a superscript outside the brackets at the top right, indicating that the structure possesses one additional electron beyond the valence count of the neutral atoms.

Next: ICl₃ Lewis structure

External video

• ICl2- Lewis Structure: How to Draw the Lewis Structure for ICl2- – YouTube • Wayne Breslyn

External links

• https://topblogtenz.com/icl2-lewis-dot-structure-molecular-electron-geometry-bond-angle/
• https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/ICl2-lewis-structure.html
• https://lambdageeks.com/icl2-lewis-structure/