Lewis Structure, Molecular Geometry, Hybridization, Polarity, and the MO Diagram are all examples of CH3Cl Lewis Structure.

Chloromethane, or CH3CL, is a highly reactive and combustible haloalkane chemical. Researchers discovered that the average life of chloromethane in the air is ten months, and that it can readily reach the stratosphere within this timeframe.

Around 2 x 106 tonnes of chloromethane enter the stratosphere each year, accounting for nearly 25% of total chlorine emissions.

Chloromethane is hazardous to the environment because it combines with a variety of natural sinks to reach all ecosystems on land, air, and water. Although chloromethane is a stable molecule, its high concentration and the ever-changing weather pattern have combined to make it a hazardous molecule to deal with.

Chloromethane Lewis Structure (CH3Cl)

The Lewis structure is a diagram that depicts the electrons that make up an atom’s valence shell. The diagram is made to show how different atoms’ valence electrons collaborate in bond creation to build a molecule.

The valence electrons are shown by the ‘dots’ near the symbol of an element in the illustration. You may notice that these valence electrons are always written in pairs on each side of the symbol to indicate whether or not all paired electrons exist.

Furthermore, the Lewis structure influences whether a single, double, or triple bond forms between the contacting atoms.

What is the best way to sketch the Lewis Structure of CH3Cl?

To begin with the Lewis structure of CH3Cl, we must first establish each involved atom’s electrical configuration.

Hydrogen (H) has an atomic number of one, hence its electronic configuration is 1s1. Because the s shell requires two electrons, there is a one-electron vacancy, resulting in a total of one valence electron in one Hydrogen (H) atom.

Carbon (C) has an atomic number of 6, which means its electronic configuration is 1s2 2s2 2p2. This gives each Carbon (C) a total of four valence electrons.

Finally, the atomic number of chlorine (Cl) is 17, and its electronic configuration is 1s2 2s2 2p6 3s2 3p5. So, for Chlorine (Cl), the number of valence electrons in the outermost shell is 7.

Let’s look at the approach for sketching the Lewis structure of chloromethane step by step (CH3Cl).

The first step is to locate the core atom: The centre atom is usually the single atom with the least electronegativity. There are only two single atoms in CH3Cl, C and Cl, with electronegativity values of 2.6 and 3.2, respectively. The Carbon (C) atom will be the core atom because it has a lower value.

Step 2: Determine the total number of valence electrons involved: This will be calculated by adding all of the atoms’ valence electrons. In the case of CH3Cl, there will be a total of 14 valence electrons.

Step 3: Determine the number of additional electrons required to stabilise one CH3Cl molecule: To properly stabilise the structure of a CH3Cl molecule, it requires an additional 8 electrons.

Step 4: Examine a CH3Cl molecule for the number and type of bonds that form: Only single covalent bonds develop between the involved atoms in the instance of Ch3Cl.

Step 5: Assemble the Lewis diagram using the procedures listed above.


CH3Cl Molecular Geometry

The CH3Cl molecule is a Penta atomic molecule with a 109.5° bond angle, giving it a bent shape.

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that chloromethane (CH3Cl) has a tetrahedral structure because the bond angle is 109.5° with Carbon (C), always as the centre atom.

Furthermore, because there is no lone pair in the CH3Cl molecule, each link is of equal angle and present at equal distance from one another, resulting in no structural deformation because all single bonds contribute to equal repulsion.

The CH3Cl molecule is depicted in three dimensions below. Take a peek around!

Is CH3Cl (Chloromethane) Polar or Non-Polar?

Polar molecules have a large difference in electronegativity between the two participating atoms, resulting in charge separation.

As a result, one end retains a positive charge while the other gains a negative charge. Non-polar molecules, on the other hand, are those in which the electronegativity difference between the participating molecules is small, if at all.

Under typical circumstances, this does not result in charge separation, and the molecule remains in a stable form with no intention of bonding with other nearby atoms.

To find the answer to this question, all of the participating atoms’ electronegativity values must be determined and compared. Because Cl (3.2) is more electronegative than C (2.6), the dipole behaviour of Carbon (C) and Chlorine (Cl) atoms shifts from downward to upward.

Furthermore, the electronegativity values of Hydrogen (2.20) and Carbon are so similar that the difference between them is minimal, making the H-C bond non-polar. However, because the C-Cl link is polar, the entire CH3Cl molecule has a net dipole moment, making it polar.

The CH3Cl molecule is polar because of the electronegativity migrating upward between the C-Cl link.

You should read the specific article on the polarity of CH3Cl at least once.

Hybridization of CH3Cl

Hybridization is a concept in which the atomic orbitals of distinct chemical elements mix and intermix to form new hybrid molecules that influence the molecule’s molecular geometry as well as its chemical bonding nature.

The Valence Bond Theory can be used to learn more about how the central atom of a molecule, such as sp, sp2, sp3, and others, undergoes hybridization.

In the CH3Cl molecule, the core element, Carbon, is sp3 hybridised. Because the CH3Cl molecule has four single bonds but no lone pair of electrons, this is the case.

When we talk about finding a molecule’s hybridization, we always start with the shells that are far away from the nucleus. In the case of carbon, these shells are 2s2 and 2p2, with paired electrons filling the 2s shell first, followed by the 2px and 2py shells.

However, this is not possible since the four single bonds must fit into all of the 2s, 2px, 2py, and 2pz shells.

To do so, each electron hybridises and fills all of the 2s, 2px, 2py, and 2pz shells, giving four single bonds the opportunity to fill the area.

The central carbon atom becomes sp3 hybridised, with the three hydrogen atoms becoming s-sp3 hybridised and the solitary chlorine atom becoming sp3-p hybridised as a result.

CH3Cl Molecular Orbital Diagram

The molecular orbital diagram is a picture of chemical bonding within a molecule that follows the rules of molecular orbital theory and the linear combination of atomic orbitals.

The theory behind the molecular orbital diagram is that some atomic orbitals combine to produce new molecular orbitals with the same number of atoms.

The involved electrons shift themselves into other orbitals, resulting in hybridised orbitals.

The basic goal of molecular orbital diagrams is to compare the energy levels of the participating orbitals in order to figure out how a molecule’s bonding occurred.

Only the sigma bond () is present in the case of a single bond, with no pi () bond. The figures above depict how hybridization occurs within a molecule, and they explain why one molecule chose to interact with the other exclusively.

Even when it comes to the energy levels of their respective orbitals, Carbon and Hydrogen atoms with almost comparable electronegativity values tend to behave similarly.

Carbon and chlorine, on the other hand, have a larger disparity in their electronegativity values, resulting in various diameters of lobes and a noticeable difference in their energy levels.


Under normal conditions, chloromethane (CH3Cl) is a stable molecule with stable atoms that do not easily react with other elements. For the CH3Cl molecule, the Lewis structure can be used to predict molecular geometry, hybridization, polarity, and a molecular orbital diagram.

The CH3Cl molecule is one of the easiest Penta-atomic structures to study and learn how to forecast different energy levels of the orbitals using a molecular orbital diagram since it is highly stable with minimal distortion in the structure.

Read more: Molecular Geometry, Hybridization, and Polarity of CH2F2 Lewis Structure

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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