ICl5is iodine pentachloride, which has a molecular mass of 304.40 g/mol. It is a useful compound in scientific laboratories to conduct research work. Although, unlike IF5, it is not a very popular compound because of its unstable nature.
This instability occurs mainly due to the larger size of the chlorine atom in comparison to fluorine. This is why ICl5is freshly prepared in labs whenever it is required to be used in a reaction.
ICl5 falls under the category of interhalogen compounds where iodine is larger in size and more electropositive than chlorine, and also acts as the central atom in this compound.
In simple language, an interhalogen compound is a type of compound in which two different halogens react with each other. These are of the type ABx where x can take the values: 1,3,5 & 7. Interhalogens in which x = 3,5 & 7 are known as polyhalides.
Therefore, ICl5is an example of a polyhalide.
To know more about this compound, it is necessary to get acquainted with its bonding characteristics as well as polarity.
For this, it is required to understand Lewis structure, VSEPR theory, and Hybridization concepts.
Lewis Structure of ICl5
The lewis dot structure gives an idea of the bonding in the ions of a compound based on the octet rule, in terms of shared electron pairs, including bonding as well as the non-bonding ones.
The lewis structure of ICl5 can be made by following the following steps:
Step 1: Calculate the total valence electrons in the compound.
For the first step, it is required to know the electronic configuration of iodine and chlorine. The electronic configuration of iodine is [Kr]4d105s25p5.
As we can see, there are seven electrons in the valence shell of iodine.
Secondly, the electronic configuration of chlorine is 1s22s22p63s23p5. Since chlorine and iodine belong to the same group i.e. group 17, there are seven electrons in the valence shell of chlorine as well.
The total no. of valence electrons in the molecule will be equal to the sum of the valence electrons of iodine and chlorine.
It is equal to ((7*1)+ (7*5)) which is equal to a total of 42 valence electrons.
Step 2: Sketching out a rough Lewis diagram
For our own feasibility, a skeletal diagram is drawn by just placing the iodine and chlorine atoms in a way where five chlorine atoms surround the iodine atom.
The following image represents the rough lewis diagram of ICl5.
Step 3: Completing the Lewis structure by fulfilling the octet rule criteria for every atom in the compound.
In the Lewis structure (a), it can be seen that there are 8 valence electrons around each chlorine atom and a total of 12 valence electrons around the iodine atom (elements belonging to period >= 3 can extend their octet to more than 8 electrons).
In the Lewis structure (b), it can be clearly seen that ten out of twelve electrons around iodine are involved in bonding (represented by the single bond) and one electron pair acts as the lone pair.
The lewis structure is complete after the formal charge on the compound gets known.
For this, it is required to know the formula to calculate the formal charge. It can be calculated by the formula:
Formal charge = V – B/2 – N
V = no. of valence electrons
B = no. of bonding electrons
N = no. of non-bonding electrons
The formal charge of a compound is a summation of the formal charge of each bonded atom in the compound. In simple language, it is needed to first calculate the formal charge of each atom individually.
Formal charge of I = 12 – 10/2 – 2 = 5
Formal charge of Cl (a) = 8 – 2/2 – 6 = 1
Similarly formal charge of Cl (b,c,d,e) = 1
Since chlorine is more electronegative than iodine, it tends to attract the shared electron pair towards itself and therefore, induces a negative charge.
Therefore, the formal charge on each chlorine atom will be -1 and not 1.
formal charge on ICl5 = (formal charge of Iodine + formal charges of chlorine atoms) = ( 5 + ( -1 + -1 + -1 + -1 + -1 ))
Hence, the formal charge on ICl5is 0, which means that the compound is neutral.
It is agreeable that a Lewis structure gives a lot of information about a certain compound, although it is difficult to tell a compound’s geometry and hybridization exactly by this concept.
Therefore, it becomes necessary to know the VSEPR theory and hybridization concept.
ICl5 Molecular Geometry
The geometry of a compound can be well explained by the VSEPR Theory which is based on the presumption that a molecule takes a shape in which it is the most stable and has the minimum electron-electron repulsions.
While predicting the shape of a molecule, the least electronegative element is the central atom of the compound.
The theory postulates that the shape of a molecule is not the same as its geometry if the no. of bond pairs electrons is not equal to the steric number which is a certain number assigned to each type of geometry according to the theory and it can be calculated by the following formula:
Steric number (S) = (A + n + |x| – y)/2
A = no. of valence electrons of the central atom
n = no. of monoatomic side atoms
x = negative charge on the molecule
y = positive charge on the molecule
The following table shows the assigned steric number to each geometry.
The steric number of ICl5 according to the VSEPR theory is calculated as follows: A = no. of valence electrons of Iodine, which is = 7
n = no. of chlorine atoms, which is = 5
Since the ICl5 compound is neutral (as discovered earlier during the lewis concept), x or y will be equal to 0 here.
S = (7 + 5) / 2 = 6
Therefore, S = 6 and so, the geometry of ICl5is octahedral.
The no. of lone pairs in a compound can be easily calculated by subtracting the no. of bond pairs from the steric number.
Depending on the no. of bond pairs and lone pairs, compounds with octahedral geometry can take up three types of shapes;
1) When the no. of bond pairs equals 4 and no. of lone pairs = 2 then the shape is square planar.
2) When the no. of bond pair equals 5 and no. of lone pair = 1 then the shape is square pyramidal.
3) When the no. of bond pairs equals the steric number i.e. 6, then the shape is the same as the geometry of the compound, which is octahedral.
Since there is 1 lone pair and 5 bond pairs in ICl5, therefore its shape must be square pyramidal.
Although due to the presence of a lone pair, the compound experiences lone pair – bond pair repulsions which have a high intensity.
In order to minimize these repulsions, the compound tends to stabilize itself by taking a stable form. To do this, it distorts its structure slightly from the square pyramidal shape.
Therefore the actual shape of ICl5is distorted square pyramidal.
One similar molecule is IF5, having the same geometry and structure. You must also check out the lewis structure of IF5.
Hybridization is a concept of the formation of hybrid orbitals from the mixing of pure atomic orbitals identical in energy and shape.
It is very important to note that the no. of hybrid orbitals thus formed must always be equal to no. of atomic orbitals mixed together.
While VSEPR gives a method to predict the geometry of a compound almost correctly, the concept of hybridization helps in understanding the bonding behind these geometrical shapes.
Although the concept of mixing of orbitals remains similar for all types of compounds, their hybridization varies since every compound arranges itself in the most stable shape in order to have minimum electron-electron repulsions.
Now let us take a look at the hybridization in ICl5,
The electronic configuration of I in the ground state is [Kr]4d105s25p5(note that we always consider the valence shells electrons only).
Below is the ground state of the Iodine atom.
To combine with five chlorine atoms, Iodine will require to extend its octet by exciting two of its valence electrons in the 5p orbital to the 5d orbital.
The electronic configuration of iodine in the excited state is as follows:
The six atomic orbitals i.e. one 5s, three 5p, and two 5d orbitals mix with each other to form six sp3d2 hybrid orbitals. Therefore, the hybridization of ICl5is sp3d2.
The electron pair present originally in the 5s orbital and finally in one hybrid orbital acts like the lone pair in the compound.
The five remaining hybrid orbitals combine with the five chlorine atoms by sigma bonding as can be seen in the following image.
However, hybridization can be easily calculated from the steric number (discussed in the VSEPR theory) itself.
The following table shows the hybridizations for respective steric numbers.
Any of the two methods can be used to figure out the hybridization of a compound according to one’s own convenience.
A compound is said to be polar in nature if it has a dipole moment i.e μ ≠ 0.
Two bonded atoms have a dipole moment if there is an electronegativity difference between them. The electronegativity of iodine is 2.66 and that of chlorine is 3.16. An electronegativity difference hence occurs between I and Cl.
It is feasible to calculate the dipole moment of a compound from its geometrical shape, which is square pyramidal in the case of ICl5.
Chlorine is more electronegative than Iodine, indicating the direction of the dipole moment is pointed towards itself.
A dipole moment with the same value in an opposite direction induces a negative sign.
The four chlorine atoms (b.c.d.e) lying in the same square plane in the geometry cancel out each other’s dipole moment as can be seen in the diagram.
In the diagram, the arrows depict the direction of the dipole moment.
One chlorine atom is left (a) which does not lie in the square plane. The bond between this chlorine atom (a) and iodine is responsible for the dipole moment of ICl5. Since in ICl5, μ ≠ 0, therefore it is polar in nature.
To summarise, in ICl5, I and Cl are bonded by five sigma bonds and the molecule has a single lone pair. ICl5is a neutrally charged compound that has an octahedral geometry and a distorted square pyramidal shape.
It has a steric number of 6 and a hybridization of sp3d2. The compound has a non-zero dipole moment and is therefore polar in nature.