Geometry, Hybridization, and Polarity of N2F2 Lewis Structure

The inorganic compound dinitrogen difluoride, N2F2, has a molecular mass of 66.01 g/mol. At typical room temperature, it is a colourless gas.

N2F2 is regarded as one of the most powerful and stable halides. It is available in both cis and trans isomers.

N2F2 Lewis Structure

Lewis structure is a diagram that depicts the distribution of valence electrons in a molecule for chemical bonding, particularly covalent bonds.

Lewis dot structure or electron dot structure/diagram are other names for it.

The dots represent the distribution and location of electrons in a molecule. Gilbert N. Lewis first introduced it in 1916.

N2F2’s Lewis structure is represented as:

To satisfy the octet rule, we already know that the elements engage in chemical bonding.

The octet of both Nitrogen and Fluorine atoms is completed by the production of a double bond between both Nitrogen atoms and a single bond between each of the nitrogen and fluorine atoms, as shown in the Lewis diagram above.

N2F2 Lewis Structure Diagram

Step 1: To sketch the Lewis structure of N2F2, we must first figure out how many valence electrons there are in the molecule.

Nitrogen is a group 15 element with 5 valence electrons and requires 3 electrons to complete its octet, whereas Fluorine is a group 17 element with 7 electrons in its outermost shell and thus only requires one electron to complete its octet.

Now, for N2F2, count the total number of electrons:

Nitrogen has 5 Valence electrons, hence 2 X 5 = 10 for 2 Nitrogen atoms.

Fluorine has 7 valence electrons, hence 2 X 7 = 14 for two fluorine atoms.

As a result, N2F2 has a total of 24 valence electrons.

Step 2: Now we’ll sketch the structure’s outline, with all of the atoms connected by a single link.

This phase examines how many more electrons are needed to complete the octet of all the atoms in the molecule.

F – N – N – F

Step 3: Because nitrogen is the least electronegative atom in the molecule, it is positioned in the middle to prevent undesired electron pull.

Furthermore, because each bond represents one shared pair of electrons, the given structure clearly shows that Fluorine has received its desired electron and has achieved a stable state by sharing one electron with Nitrogen to complete its octet.

Step 4: Despite this, both Nitrogen atoms are still missing an electron.

As a result, they form a double bond amongst themselves, leaving only one pair of lone electrons.

On the other hand, fluorine has three pairs of lone electrons.

Step 5: Finally, the Lewis structure of N2F2 looks like this after completing the octet of both the Nitrogen and Fluorine atoms:

N2F2 Molecular Geometry

The electrons inside a molecule have a natural inclination to arrange themselves to prevent inter-electronic repulsion, according to the Valence Shell Electron Pair Repulsion (VSEPR) Theory.

This repulsion can occur between lone pairs of electrons and electrons engaged in the formation of covalent bonds; however, because lone pairs are free in space, their repulsion is stronger than that of electrons involved in chemical bond formation.

The gap in electronegativity between the core atom and other atoms also influences the magnitude of repulsion.

The amount of lone pairs in a molecule, as well as the extent of electronic repulsion, affect the shape of the molecule.

The distortion of the bond angle between the central atom and adjacent atoms is also determined by the lone pair of electrons present on the central atom.

We must first identify a core atom in order to comprehend the molecular geometry of N2F2.

All of the other atoms in a molecule are thought to be related to the centre atom, according to VSEPR Theory.

Both Nitrogen atoms are positioned at the core of the molecule, as in N2F2, and any one of them might be chosen as the central atom.

Assuming that one Nitrogen atom is the centre element, it is attached to one Fluorine atom by a single bond and to another Nitrogen and Fluorine atom by a double bond.

As a result of this assumption, the atoms tend to organise themselves as far as possible to avoid inter-electronic repulsion, resulting in a linear geometry (180°).

However, as previously stated, the Nitrogen atom has one lone pair of electrons bonded to it, which drives both groups on the central atom downward.

This causes the molecule to distort, resulting in a bond angle of approximately 118° and the formation of a trigonal planar geometry.

N2F2 hybridization

Hybridization was first proposed by Linus Pauling in 1931.

As the name implies, it is the process of combining two or more atomic orbitals, such as s, p, d, and so on, to generate a new hybrid orbital with similar energies.

The atomic orbitals involved in the production of hybrid orbitals are given names. One s and three p orbitals, for example, were joined in the creation of sp3.

In the instance of N2F2, hybridization can be estimated using either the Highest Occupied Molecular Orbital (HOMO) or the Lowest Unoccupied Molecular Orbital (LUMO) calculations (LUMO). P-orbitals are a part of both HOMO and LUMO for cis-N2F2.

Furthermore, the p-orbitals are perpendicular to the plane, a feature of sp2 hybridization.

According to density functional theory, the bond order for N=N is close to 2.

Trans-N2F2 looks to be sp3 hybridised since its HOMO is somewhat different from cis-N2F2, however sp3 hybridised molecules tend to arrange their lone pair in a tetrahedral configuration.

However, as previously stated, N2F2 has a trigonal planar shape, indicating that cis-N2F2 is sp2 hybridised.

Another way to tell if N2F2 is hybridised is to look at its Lewis structure:

The core atom, Nitrogen, is linked to another Nitrogen atom with a double bond, i.e. 1 and 1 orbital, and to one Fluorine atom with a single bond, i.e. 1 orbital, on one side.

This proposes three hybrid orbitals for a single N-atom, i.e. sp2 hybridization, one for forming N–F bonds, another for forming N=N bonds, and the third for holding the lone pair.

As a result, N2F2 hybridization is sp2.

N2F2 has both a Cis and a Trans form.

The two geometrical forms, Cis and Trans, are symmetrically distinct and have differing boiling points, melting points, and other attributes.

In compared to the trans form, the cis-form is more thermodynamically stable, averaging over 12.5 KJ/mol.

Some of the differences in attributes between these two isomers due to different symmetries are as follows:

S. No.PropertiesCis FormTrans form
 Boiling Point-105.75° C-111.45° C
 Melting Point< -195° C-172° C
 SymmetryC2vC2h
 Bond Angle114°106°

Polarity of N2F2

The existence of two opposed charges, positive and negative, on distinct atoms of the same molecule is referred to as polarity.

The formation of polar bonds is mainly caused by a difference in electronegativity between the joining atoms.

Because polarity is determined by the arrangement of atoms in a molecule, the dipole moments of the two forms of N2F2 vary.

Both Fluorine atoms are positioned on the same side in cis-N2F2, resulting in a dipole moment of 0.12-0.16D, however in trans-N2F2, where the Fluorine atoms are arranged on different sides, the charges cancel out, and the dipole moment of trans-N2F2 becomes 0.

Conclusion

  1. N2F2 has a Lewis structure.
  1. N2F2 has a trigonal planar molecular geometry.
  2. The N2F2 hybridization is sp2.
  3. The cis-N2F2 dipole moment is 0.12-0.16 D, while the trans-N2F2 dipole moment is 0 D.

Read more: Are Gold Bars Flexible? (An Accurate Scientific Justification)

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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