Hybridization of I₃⁻

Triiodide in Chemistry usually refers to the Triiodide ion, I₃^- This anion, one of the polyhalogen ions, is composed of 3 iodine atoms and is formed by combining the aqueous solution of iodine and iodide salts. A few salts of the anion have been isolated, including ammonium Triiodide ([NH4]+[I3]− and thallium(I) Triiodide (Tl+[I3]−)). The Triiodide is recognized to be red in solution.

What is the Hybridization of Triiodide Ion?

To determine the Hybridization of the Triiodide ion, we can use the following simple Hybridization formula:

Valence electron + monovalent + (negative charge) – (positive charge)/2 = Number of Hybridization

Iodine atoms have seven valence electrons in their outer shell, as well as two monovalent atoms. Furthermore, when Iodine is combined with the other two Iodine atoms, the centre atom receives a negative charge, which is equal to 1.

If we use the formula to place or substitute the values, we get

7+1+2/2

=10/2

=5

As a result, the hybridisation number is 5. Hybridisation is now classified as sp³d.

Alternatively, knowing the number of valence electrons and lone pairs and computing their sum can be used to calculate the Hybridization of I₃^-. In this scenario, there are three lone pairs, although the number of atoms giving valence electrons is only two. When these values are added together, we get 5, which indicates sp³d hybridisation.

Another way to find the Hybridization of a given molecule is by taking the help of lone pairs and valence electrons. Here, the number of lone pairs in the molecule is 3 and the number of atoms sharing valence electrons is 2. 

\[I^{-}_{3}\] Hybridization 

I₃ (-) Hybridization is a term used to describe the process of combining two or more. In this section, we'll learn how to determine the Hybridization of I₃^-. I₃^- is a linear anion, which is important for students to memorize. The bonding of I₂ with the I ion creates it. Where electrons are generally accommodated in the vacant d orbitals, the I ion will be the donor and the I₂ molecule will be the acceptor. The i₃^- ion is essentially an sp³d hybridized ion.

Henceforth, 3+2=5, which also determines the same sp³d Hybridization.

Structure of Lewis

The electrons of molecules are represented by the Lewis structure. Lone pairs and valence electrons aid in the determination of the molecule's Hybridization and shape. One molecule of Iodine will be in the centre because there are so many of them. In addition, iodine belongs to the periodic table's seventh group and possesses seven valence electrons in its outer orbit.

Lewis structure is nothing but the electron representation of the molecules. There are lone pairs and valence electrons that help in determining the shape and Hybridization of the molecule. As there are Iodine molecules, one molecule of iodine will present in the centre. Iodine also lies in the seventh group of the periodic table and has 7 valence electrons in its outer orbit.

Here, we have 3 molecules of iodine, which along with an extra electron that gives it a negative charge. So, the total valence electrons count is given by 7×3 + 1= 22.

Moreover, in this molecule, there are 22 valence electrons in total. Now, there is an octet rule followed by every atom. According to this rule, every atom should have 8 electrons in their outer orbit. So, if there are 8 electrons in the central atom's outer shell, there exist 2 other atoms required to complete their octet. As all the atoms need 8 electrons in their outer shell for octet completion, 1 central iodine's electron atom will be taken by both neighbouring iodine atoms. It means 8-1-1=6 because both the atoms will take the electrons.

So, the valence electrons on the central atom of iodine will become 6. These 6 electrons will form the electron lone pairs that do not bond. As it forms the pair of electrons now, there will be 3 lone pairs, and 2 bond pairs of the electrons because each Iodine atom has a bond with the central atom sharing 1 electron each. So, in total, there are 3 lone pairs and 2 bond pairs on the central atom.

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\[I^{-}_{3}\] Bond Angles and Molecular Geometry

The Geometry of I₃^- molecules is linear. Because one of the three Iodine atoms has a negative charge, there are three lone pairs of electrons and two bond pairs. It will have a steric number of 5. The three solitary pairs will resist one another and take up equatorial positions. The remaining two iodine atoms are at an Angle of 180 degrees to one another.

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The Molecular Geometry of I₃^- is linear. While there are 3 Iodine atoms, one of the atoms has a negative charge, which further gives 2 bond pairs and 3 lone pairs of electrons. Its steric number will be 5. The 3 lone pairs will repel each other and take up the equatorial positions. The remaining 2 Iodine atoms present at 180° from each other.

Polarity of \[I^{-}_{3}\]Ion

So, there is a tricky part of this ion here. Firstly, we can call the charge present on it as a polyatomic ion instead of a molecule. Ions are the charges that we see on the molecules. As I₃^- has only 1 electron, this Ion has an overall negative charge. Whereas, the molecules have Polarity because it has both the charges which are charged partially positively and negatively. There exists a dipole moment on the molecules based upon the charges separation on the molecule. If the distance between both charges is large, then the dipole moment will also become larger.

However, when we discuss I₃^- ion, it is a negatively charged ion. Even while drawing its Lewis structure, we do not see any dipole moment of the polar bonds in it, because the overall charge itself is negative on the ion. So, it is neither polar nor nonpolar. However, if we have to describe the ion, we can use the phrase "like a polar molecule" because I₃^- is soluble in water.

Important Things to Keep in Mind

• The bonding of I₂ with the I ion produces I₃^-.

• The centre atom develops a negative charge with a value of 1 during the combining of Iodine atoms.

• The donor is an ion, while the acceptor is an I₂ molecule. The vacant d orbitals are mostly occupied by electrons.