Molecular Geometry, Hybridization, and Polarity of CH2F2 Lewis Structure

CH2F2, also known as difluoromethane or difluoromethylene, is a haloalkane chemical molecule. Alkyl halides, also known as haloalkanes, are organic compounds with at least one halogen element linked to the carbon atom.

At ordinary temperature and pressure, it is a colourless gas. Because of its polar nature, it has a high thermal stability and a low water solubility. Because of its widespread application in endothermic processes, it is used in refrigeration, air conditioning, and fire extinguishers.

Difluoromethane has a molar mass of 52.024 g/mol. It has a melting point of -136 °C and a boiling point of -52 °C. Frostbite is caused by difluoromethane’s low melting and boiling temperatures.

Despite the fact that CH2F2 includes fluorine, an electronegative element, it lacks hydrogen bonding due to the absence of an H-F bond.

For easy understanding, the important properties of chemical bonding in difluoromethane are listed in the table below. We’ll go through each feature in depth, one by one.

CompoundCH2F2
Molecular GeometryTetrahedral
Hybridizationsp3
PolarityPolar

CH2F2 Lewis Structure

Because most elements are unstable in their atomic form, atoms create chemical bonds to acquire stability.

The creation of chemical bonds was first explained in terms of electrons, particularly valence electrons, by Kossel and Lewis.

Except for hydrogen and helium, every atom in a chemical bond prefers to form a stable octet, according to Lewis. Only the valence electrons, which are found in an atom’s outermost shell, are involved in bond formation.

Carbon, hydrogen, and fluorine belong to groups 14, 1, and 17, respectively, in CH2F2. As a result, carbon, hydrogen, and fluorine have 4, 1, and 7 valence electrons, respectively.

Carbon, hydrogen, and fluorine have group valences of 4, 1, and 1, respectively. The number of chemical bonds that an atom can create with other atoms is determined by its group valence.

As a result, carbon can make four chemical bonds, while hydrogen and fluorine can only form one. As a result, the centre atom in difluoromethane will be carbon.

In difluoromethane, the total number of valence electrons is 4 + 1(2) + 7(2) = 20.

Valence electrons are shown as dots in the Lewis structure of the molecule.

The Lewis structures require that each combining atom donate at least one electron for sharing with other atoms, and that the sharing of electron pairs between atoms results in the formation of a chemical bond.

By sharing electrons, the combining atom will obtain the closest noble gas structure.

With carbon as the core atom, we must now arrange 20 valence electrons in the Lewis structure of CH2F2.

Because carbon has a valence electron of four, it will share two of its valence electrons with two hydrogen atoms and the remaining two electrons with two fluorine atoms, forming an octet surrounding it.

As a result, the Lewis structure of CH2F2 is as follows:

Both the hydrogen and fluorine atoms share only one electron with the carbon atom in the Lewis structure of CH2F2.

As a result, the single bond between carbon and hydrogen, as well as carbon and fluorine, is formed.

In a nutshell, a carbon atom, which is the core atom, will produce four single bonds with no lone pair.

However, comprehending the chemical processes of CH2F2, which entail bond breaking and bond formation, cannot be limited to the sole Lewis structure.

As a result, it’s critical to comprehend CH2F2’s molecular geometry.

Molecular Geometry of CH2F2

The valence shell electron pair repulsion (VSEPR) theory proposed by Sidgwick and Powell can be used to predict the molecular geometry of CH2F2. This theory is based on the atom’s valence electrons repelling each other.

These electrons in the valence shell might be bonded (bond pair) or non-bonded (lone pair). To minimise repulsion and increase molecule stability, these pairs of electrons occupy the position around the centre atom.

If there are two chemical bonds surrounding the core atom, for example, the bonds will arrange themselves in such a way that the bond angle is 180 degrees, resulting in linear geometry.

The carbon atom in difluoromethane is a core atom with four electron bond pairs. For minimal repulsion between any two bond pairings, these four chemical bonds will form a tetrahedral geometry with a bond angle of 109.5 degrees.

The following table, which is based on VSEPR theory, can also predict the form of difluoromethane.

General formulaNumber of bond pairsMolecular shape/geometry
AX1Linear
AX22Linear
AX33Trigonal planar
AX44Tetrahedral
AX55Trigonal bipyramidal
AX66Octahedral

AX4 will be the generic formula for difluoromethane. As a result, difluoromethane has a tetrahedral shape. The tetrahedral geometry of difluoromethane results in a bond angle of 109.5 degrees (F-C-F, H C-H, and F-C-H).

In CH2F2, however, the F-C-F bond angle is slightly less than 109.5 °, while the H-C-H bond angle is somewhat larger.

Furthermore, the C-H bond length of CH2F2 is shorter than that of methane (CH4). The Bent rule of hybridization can explain these two abnormalities.

Once we understand the hybridization of the carbon atom in difluoromethane, we may understand the bent rule of hybridization.

Hybridization of CH2F2

The Valence bond theory predicts that the carbon atom in difluoromethane will hybridise (VBT). Pauling coined the words hybridization and hybrid orbitals.

Hybrid orbitals are created by combining two atomic orbitals with similar energies. The form and energy of all hybrid orbitals are the same. The total number of hybrid orbitals equals the total number of atomic orbitals.

Two atomic orbitals, for example, can be combined to generate two hybrid orbitals with the same energy and shape. These hybrid orbitals are involved in the creation of bonds.

Carbon’s electrical ground state configuration is [He] 2s22p2.

Carbon creates four single bonds according to the Lewis structure, so we’ll need four unpaired electrons. One of the electrons from the 2s orbital will now excite to the carbon atom’s 2p orbital. It gives rise to [He] 2s12p3 as the excited-state electrical configuration of carbon.

The carbon atom’s one 2s orbital and three 2p orbitals will combine to generate four sp3 hybrid orbitals with the same energy and shape. Every hybrid orbital is made up of 25% s-character and 75% p-character.

The four corners of the tetrahedron are targeted by these hybrid orbitals, and the angle between them is 109.5 degrees.

Two of the carbon atom’s sp3 hybrid orbitals will overlap with the hydrogen atom’s 1s atomic orbital, while the other two will overlap with the fluorine atom’s 2p atomic orbital.

The end-to-end overlapping of atomic orbitals results in the production of four sigma bonds. The orbital diagram for difluoromethane is as follows:

If a central atom is connected to numerous atoms, it will hybridise in such a way that hybrid orbitals with more s-character point towards the more electropositive element while hybrid orbitals with more p-character point towards the more electronegative element, according to Bent’s rule.

Fluorine is more electronegative than hydrogen in CH2F2. Fluorine enhances the p character while decreasing the s character in the C-F bond, resulting in a longer C-F bond. The length of the C-H bond will also shorten when the s-character of the C-H bond increases.

The F-C-F bond angle is somewhat shorter than 109.5 ° due to a decrease in s-character, whereas the H-C-H bond angle is bigger than 109.5 °.

Polarity of CH2F2

The polarity of a molecule is determined by its shape, net dipole moment, and charge distribution.

CH2F2 has a tetrahedral form, which means it has a symmetrical shape and thus a symmetric distribution of atoms around the carbon atom.

The dipole moment is determined by the difference in electronegativity between the chemical bond and the dipole moment, which must be greater than 0.4 for the chemical bond to be polar. Carbon, hydrogen, and fluorine have electronegativity values of 2.55, 2.20, and 3.98, respectively.

A dipole is created by the huge disparity in electronegativity. Because fluorine is an electronegative element, it attracts electron pairs and thus has a partial negative charge.

Check out the post I wrote about CH2F2 Polarity for more details.

Due to its electropositive nature, hydrogen will also have a partial positive charge.

Difluoromethane is thus a polar compound. Due to polarity, which holds the molecules together, it will have dipole-dipole intermolecular forces.

CH2Cl2 is a polar chemical that is related to CH2F2. Refer to the polarity of CH2Cl2 for further information on the factors that determine polarity in such a molecule.

It causes difluoromethane to become gaseous, with a boiling point of -51 C at standard temperature and pressure.

Conclusion

The difluoromethane is an alkyl halide that occurs as a gas at room temperature and pressure.

It has tetrahedral geometry and sp3 carbon atom hybridization, as predicted by valance shell electron pair repulsion theory and valance bond theory, respectively.

As a polar molecule, difluoromethane exhibits dipole-dipole forces to keep the molecules together.

I hope you enjoyed learning about the chemical bonds in the difluoromethane molecule.

Good luck with your studies.

Read more: Is NF3 a polar or a nonpolar protein?

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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