CH2F2 Lewis Structure, Molecular Geometry, Hybridization, and Polarity


CH2F2 or difluoromethane or difluoromethylene is an organic compound of the haloalkane family. Haloalkanes or alkyl halides are organic compounds, which contain at least one halogen atom bonded to the carbon atom.

It is a colorless gas at standard temperature and pressure. It has high thermal stability and slight solubility in water owing to its polar nature. It is utilized in the refrigeration, air conditioner, and fire extinguishers process due to its common use in endothermic processes.

The molar mass of difluoromethane is 52.024 g/mol. Its melting point and boiling points are -136 °C and -52 °C, respectively. The low melting and boiling points of difluoromethane result in frostbite.

CH2F2 contains fluorine, an electronegative element, but still, it does not show hydrogen bonding owing to the absence of an H-F bond.

The key features of chemical bonding in difluoromethane are shown in the following table for quick learning. We will understand every feature in detail step by step.

Compound CH2F2
Molecular Geometry Tetrahedral
Hybridization sp3
Polarity Polar


CH2F2 Lewis Structure

The chemical bond is an attractive force between two atoms and atoms form a chemical bond to achieve stability as most of the elements are unstable in their atomic form.

Kossel and Lewis were the first to explain the formation of chemical bonds in terms of electrons especially valence electrons.

According to Lewis, every atom tends to achieve a stable octet when they form chemical bonds except hydrogen and helium. Only the valence electrons, which are present in the outermost shell of an atom participate in the bond formation.

In CH2F2, carbon, hydrogen, and fluorine are from groups 14, 1, and 17, respectively. Hence, the number of valence electrons in carbon, hydrogen, and fluorine is 4, 1, and 7, respectively.

Group valence of carbon, hydrogen, and fluorine is 4, 1, and 1, respectively. The group valence indicates the number of chemical bonds an atom can form with other atoms.

Hence, carbon can form four chemical bonds whereas hydrogen and fluorine can form only one chemical bond. Therefore, carbon will be the central atom in difluoromethane.

The total number of valence electrons in difluoromethane are 4 + 1(2) + 7(2) = 20 electrons.

In the Lewis structure of the molecule, valence electrons are represented as dots.

The important conditions for the Lewis structures are that every combining atom has to contribute at least one electron for sharing with other atoms and a chemical bond is formed owing to the sharing of electron pairs between atoms.

The combining atom will achieve the nearest noble gas configuration by sharing electrons.

Now, we have to arrange 20 valence electrons in the Lewis structure of CH2F2 with carbon as a central atom.

The valence electron in carbon is four and hence, it will share its two valence electrons with two hydrogen atoms and the remaining two electrons with two fluorine atoms, resulting in the formation of an octet around it.

Therefore, the Lewis structure of CH2F2 can be represented as:

CH2F2 lewis structure

In the Lewis structure of CH2F2, both Hydrogen and Fluorine atoms are sharing only one electron with the carbon atom.

Hence, there is a formation of the single bond between carbon and hydrogen as well as carbon and fluorine.

In brief, a carbon atom, a central atom, will form four single bonds without any lone pair on it.

But we cannot limit ourselves to the only Lewis structure for understanding the chemical reactions of CH2F2, which involves bond breaking and bond formation.

Therefore, there is a need to understand the molecular geometry of CH2F2.


CH2F2 Molecular Geometry

The molecular geometry of CH2F2 can be predicted by valence shell electron pair repulsion (VSEPR), a theory given by Sidgwick and Powell. This theory is based on repulsion between valence electrons of the atom.

These valence shell electrons may be bonded (bond pair) or non-bonded (lone pair). These pairs of electrons occupy the position around the central atom to minimize the repulsion and maximize the stability of the molecule.

For example, if there are two chemical bonds around the central atom, then these bonds will arrange themselves in such a way that the bond angle would be 180 °and hence, linear geometry will form.

In difluoromethane, the carbon atom is a central atom that has four bond pairs of electrons. These four chemical bonds will arrange in tetrahedral geometry with a bond angle of 109.5 ° for minimum repulsion between any two bond pairs.

The shape of difluoromethane can also be predicted by the following table, which is based on VSEPR theory.

General formula Number of bond pairs Molecular shape/geometry
AX 1 Linear
AX2 2 Linear
AX3 3 Trigonal planar
AX4 4 Tetrahedral
AX5 5 Trigonal bipyramidal
AX6 6 Octahedral

The general formula for difluoromethane will be AX4. Hence, difluoromethane will have a tetrahedral geometry. The tetrahedral geometry of difluoromethane leads to a bond angle (F-C-F, H C-H, and F-C-H) of 109.5 °.

CH2F2 bond angle

But, the F-C-F bond angle is slightly less than 109.5 ° and the H-C-H bond angle is slightly greater than 109.5 ° in CH2F2.

Also, the C-H bond length of CH2F2 is smaller than the C-H bond length of methane (CH4). These two anomalies can be explained by the Bent rule of hybridization.

The bent rule of hybridization can be understood once we know the hybridization of the carbon atom in difluoromethane.


CH2F2 Hybridization

The hybridization of the carbon atom in difluoromethane is predicted by the Valence bond theory (VBT). The terms hybridization and hybrid orbitals were given by Pauling.

Hybrid orbitals are formed from the combination of atomic orbitals of similar energy. All hybrid orbitals are of the same shape and energy. The number of hybrid orbitals will be equal to the number of atomic orbitals.

For example, two atomic orbitals combine to form two hybrid orbitals of equivalent energy and shape. These hybrid orbitals participate in bond formation.

The ground state electronic configuration of carbon is [He] 2s22p2.

According to the Lewis structure, carbon forms four single bonds, and hence, we need four unpaired electrons. Now, one of the electrons from the 2s orbital will excite to the 2p orbital of the carbon atom. It leads to the excited-state electronic configuration of carbon as [He] 2s12p3.

One 2s orbital and three 2p orbitals of the carbon atom will combine to form four sp3 hybrid orbitals of the equivalent energy and shape. Every hybrid orbital has 25 % s-character and 75 % p character.

These hybrid orbitals are directed toward the four corners of the tetrahedron and the angle between them is 109.5 °.

CH2F2 hybridization

Two of the sp3 hybrid orbitals of the carbon atom will overlap with the 1s atomic orbital of the hydrogen atom and the remaining two will overlap with the 2p atomic orbital of the fluorine atom.

It results in the formation of four sigma bonds due to the end-to-end overlapping of atomic orbitals. The orbital diagram of difluoromethane can be represented as:


CH2F2 hybridization

According to Bent’s rule, if a central atom is bonded to multiple atoms then it will hybridize in such a way that hybrid orbitals with more s-character will point towards the more electropositive element and hybrid orbitals with more p-character will direct towards the more electronegative element.

In CH2F2, Fluorine is more electronegative than hydrogen. Hence, Fluorine will increase the p character and decrease the s-character in the C-F bond due to which the C-F bond length increases. Similarly, the C-H bond length will decrease due to an increase in s-character in the C-H bond.

Due to a decrease in s-character, the F-C-F bond angle is slightly shorter than 109.5 ° whereas the H-C-H bond angle is greater than 109.5 °.


CH2F2 Polarity

The polarity of the molecule depends upon its shape, net dipole moment, and distribution of charges.

The shape of CH2F2 is tetrahedral, which is a symmetrical shape and hence, symmetric distribution of the atoms around the carbon atom.

The dipole moment depends upon the difference in electronegativity, which should be greater than 0.4 to the chemical bond to be polar. The electronegativity values of carbon, hydrogen, and fluorine atom are 2.55, 2.20, and 3.98, respectively.

This large difference in electronegativity generates a dipole. Fluorine being the electronegative element will pull the electron pair towards it and hence, it will have a partial negative charge.

For detailed information, Go through the article I already wrote on CH2F2 Polarity.

CH2F2 polarity

Similarly, hydrogen will have a partial positive charge due to its electropositive nature.

Therefore, difluoromethane is a polar molecule. It will have dipole-dipole intermolecular forces owing to polarity, which holds the molecules together.

CH2Cl2 is a similar compound having polar nature as CH2F2. For detailed information regarding factors affecting polarity in such a compound, you must refer to the polarity of CH2Cl2.

It results in the gaseous nature of difluoromethane with a boiling point of -51˚C at standard temperature and pressure.



The difluoromethane is an alkyl halide, which exists in gaseous form at standard temperature and pressure.
It has tetrahedral geometry with sp3 hybridization of the carbon atom, which is predicted from valance shell electron pair repulsion theory, and valance bond theory, respectively.

Difluoromethane is a polar molecule and hence, it shows dipole-dipole forces to hold the molecules together.

I hope you have enjoyed the chemical bonding aspects of the difluoromethane molecule.

Happy Learning.

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