Lewis Structure, Molecular Geometry, Hybridization and MO Diagram for CH2Cl2

Methylene chloride (CH2Cl2), also known as chloromethane or CH2Cl2, is a colourless, volatile liquid with a boiling point of 39.6°C. and melting point of -96.7°C. It is also known as methylene chloride. It is a common solvent in chemistry labs around the world.

In spite of its polarity and the presence of two chloro groups, chloroform is not miscible with water. However, chloroform is miscible with a variety of organic solvents, such as carbon tetrachloride and several of the above-mentioned organic solvents.

Chlorine gas is used to treat methane or chloromethane at high temperatures in order to produce CH2Cl2.

Overexposure to CH2Cl2 through inhalation can cause symptoms such as nausea, dizziness, numbness, and weakness. Carbon monoxide is the byproduct of its metabolism, and as a result, it is toxic.

In the chemistry of CH2Cl2

Chemical bonding can be better understood using the Lewis theory of chemical bonding, which is a primitive and extremely limited theory of electronic structure.

Using this theory, you can better understand how atoms in molecules form chemical bonds.

A line is drawn for each pair of electrons that form a link between two atoms. Lewis (dot) structures are the names given to the structures drawn using this approach.

One thing to note is that, as can be seen in Lewis structure, there are various atoms that follow the octet rule, which means that they tend to achieve eight electrons in their valence shell through chemical bonding.

Because hydrogen has only one electron in its K shell, it tends to be a duplet rather than an octet, because it needs only one more electron to reach the maximum capacity of its K shell.

Let’s take a closer look at the Lewis structure of CH2Cl2’s chemical bonds.

Step 1: To determine the distance between an atom and an octet, we will first calculate the amount of valence electrons present in each CH2Cl2 atom (or duplet in the case of hydrogen).

i. Because carbon has an atomic number of six, it has six electrons in its neutral state. The K shell has two electrons, and the L shell has four. Consequently, there are four valence electrons. Carbon needs an additional four electrons to complete the octet.

Hydrogen, like oxygen, has an atomic number of 1. As a result, each H possesses a single electron and needs another to form a duplex.

Chlorine has an atomic number of 17. There are two electrons in the K shell, eight in the L shell, and seven in the M shell. Because there are seven valence electrons in total, Cl just requires one more to complete the octet.

The next step is to determine the core atom to which the rest of the atoms will be bound. The shortest atom in the octet, the centre atom, has the most bonding ability. This requirement is met by carbon in CH2Cl2 (4 electrons short of the octet versus 1 for chlorine).

Step 3: Using carbon as the centre atom, we’ll build the molecule’s backbone. The octet of carbon still lacks four electrons. To help carbon accomplish this feat, two hydrogen and two chlorine atoms can be added.

The circumstance will be a win-win for all five atoms because both hydrogen and chlorine atoms will acquire their respective duplets and octets at the same time.

As depicted in the Lewis structure, carbon will be linked to H, H, Cl, and Cl.

There is no charge on the molecule, hence it is considered neutral. Let’s figure out the formal charges of each atom. The formal charge formula follows.

In chemistry, the formal charge (FC) is calculated as follows:

As for carbon, FC = 0, as for hydrogen, FC = 0, as well as Cl.

Using CH2Cl2 as a template

When two atoms have orbitals that overlap one other, they create a bond because these orbitals share electrons.

Take a peek at the orbitals to see how each CH2Cl2 atom’s electrical arrangement looks in its ground state.

One thousand twenty-two thousand twenty-two twenty-two twenty-two twenty-two

The first letter of the first number of the alphabet

1.S.22.PXYZ3.S.23.PXYZ1.Cl:

The electrical configuration of carbon changes to 1s22s22px12py12pz1 when it is stimulated because one of the 2s electrons has been promoted to the 2p state. This means that the carbon atom has only half of its two-sigma, two-px, two-py and two-pz orbitals.

They generate four identical sp3 orbitals with the same energy when they hybridise together. Each of these hybrid orbitals already has one electron, but it is capable of taking on another one if necessary.

Each atom of H, H, Cl, and Cl contributes a single electron: 1s1 of each H and 3pz1 of each Cl atom. As a result, four carbon sp3 hybrid orbitals create four single bonds (also known as sigma bonds).

The following formula can be used to figure out how much the core atom has hybridised.

A + (VE – V – C)/2 = Hybridization.

where

To determine how many atoms are bonded to the centre atom, use the following formula:

The number of valence electrons on the centre atom is referred to as VE.

The centre atom has a valency of V.

The core atom’s charge is designated as C.

How much is H worth?

Species of crossbreeding

2 sp

3 sp2

4 sp3

5 sp3d

6 sp3d2

sp3 is represented by Hyb = 4, which is equal to a value of A = 4, VE = 4, V = 4, and C = 0.

Molecular Arrangements of CH2Cl2

The hybridization of the centre atom can be used to identify the molecule’s shape. When the central atom has no lone pairs of electrons, the geometry is tetrahedral, as predicted by sp3 hybridization.

Based on stoichiometry, number of bond pairs, and number of lone pairs on the core atom, the VSEPR theory aids in determining a molecule’s geometry.

According to this theory, molecules adopt an atomic structure that minimises the repulsion caused by the electrons in their valence shells on all of their constituent atoms.

A list of molecular stoichiometries is shown below in accordance with VSEPR theory.

In accordance with the data in the preceding table, the crystal structure of CH2Cl2 is tetrahedral, as predicted by the equation for AX4.

CH2Cl2’s tetrahedral structure is not as flawless as that of CH4. This is because the hydrogen atoms in CH4 are all the same, whereas the hydrogen atoms in CH2Cl2 are different.

The molecular form of the latter is slightly asymmetrical as a result. There are some differences between CH2Cl2 and CH4 in terms of bond angles and bond lengths.

Polarity of CH2Cl2

There’s something about CH2Cl2 that makes it polar.

Because of the molecule’s tetrahedral structure and the difference in electronegativity between carbon and chlorine, Polarity is determined by the asymmetric form and electronegativity differences between atoms.

Due to its polar nature, the C-Cl bond results in a non-uniform charge distribution throughout the molecule. You may learn more about the polarity of CH2Cl2 by reading this article.

Orbital Diagram of the Molecule CH2Cl2

Molecular orbital (MO) theory assumes that all atoms in a molecule contribute to the production of molecular orbitals, which are a linear combination of atomic orbitals. This is the premise of MO theory. Molecular orbitals are used to assign electrons in molecules instead of atomic orbitals.

CH2Cl2’s MO diagram is available here. Carbon’s 2s and 2p orbitals combine (to varying degrees) with the 1s and 2pz orbitals of the hydrogen and chlorine atoms.

8 atomic orbitals combine to generate 8 molecular orbitals, as can be seen in the image below. According to the relative energy alignment of atomic orbitals, mixing and the contribution of individual atomic orbitals to a molecule orbital can vary greatly.

The most energetic molecular orbital is the first to be filled with electrons. A bonding orbital is one that has a molecule’s electrons attached to it, while an anti-bonding orbital is one that has their electrons removed.

Non-bonding molecular orbitals are formed when a lone pair of an atomic orbital does not combine with any other orbital.

Schematic of the CH2Cl2 MO reaction

Conclusion

Chemical bonding is used in this article to provide a rudimentary grasp of CH2Cl2’s structure.

If you have any additional queries, please do not hesitate to contact me.

Learning is a joy, so have fun with it!

Read more: Diagram and Instructions for Drawing the Lithium Bohr Model

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.

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