NO2F, an inorganic compound, is named Nitryl Fluoride. It is a colorless gas and acts as a strong oxidizing agent. So, it can be used in the replacement of liquid oxygen, an oxidant in propellants of the rocket. NO2F can easily release its fluoride ion to other species in the reaction.
The boiling point of nitryl fluoride is low i.e., -72 °C, indicating its molecular nature instead of the ionic nature.
The nitryl fluoride is prepared by fluorination of nitrogen dioxide (NO2) by cobalt trifluoride (CoF3). Here, NO2 is oxidized to NO2F, and CoF3 is reduced to CoF2.
NO2 + CoF3 —–> NO2F + CoF2
The molecular weight of nitryl fluoride is 65 g/mol.
Here, we are going to learn about the chemical bonding of Nitryl Fluoride.
NO2F Lewis Structure
We need to depict the Lewis structure of nitryl fluoride to understand the nature of chemical bonding in it.
The Drawing of the Lewis structure involves step by step procedure for the two-dimensional representation of the molecule. The Lewis structure is based on the concept of valence electrons. The valance electrons are the electrons that are present in the outermost shell of the atom and prone to form chemical bonds.
The nitryl fluoride consists of one nitrogen atom, one fluorine, and two oxygen atoms. The first step would be to compute the total number of valence electrons in the nitryl fluoride.
The nitrogen, oxygen, and fluorine atoms belong to group 15, group 16, and group 17 of the modern periodic table, respectively. Therefore, Nitrogen, oxygen, and fluorine atoms have 5, 6, and 7 valence electrons, respectively.
Hence, NO2F molecule consists of 5 + (2*6) + 7 = 24 valence electrons.
The second step consists of finding out the central atom of the molecule. It will help in drawing the Lewis structure. Generally, the least electronegative element acts as a central atom.
As we know that electronegativity increases across the period. So, the increasing order of electronegativity of three elements is N < O < F.
It gives us a clear indication that the nitrogen atom is a central atom and oxygen and fluorine atoms will surround it. The skeletal structure of NO2F would be:
Now, the next step would be to place the valence electrons around the nitrogen atom in the nitryl fluoride.
Here, we need to understand the octet rule. The octet rule tells us that every atom tends to achieve the nearest noble gas configuration i.e., eight electrons in the valence shell of the atom. Therefore, these three elements tend to achieve neon gas configuration on bonding. The 24 valence electrons can be arranged like the following diagram:
Here, oxygen and fluorine atoms have eight electrons in their valance shell but the nitrogen atom has only six electrons. To complete its octet, let us move two valance electrons (lone pair) from the oxygen atom to form a double bond or share four electrons with the nitrogen atom. The oxygen atom will share its electrons instead of the fluorine atom as it is less electronegative than the fluorine atom.
The Lewis dot structure or electron dot structure of the nitryl fluoride would be:
The Lewis structure of NO2F in the bond notation form would be:
There will be two resonance structures of nitryl fluoride as any oxygen atom can share its electron pair. Now, every atom has eight electrons but nitrogen can form only three bonds owing to the absence of d-orbitals. Hence, there is a need to estimate the formal charge on every atom in the molecule.
The formal charge is defined as:
Formal charge = valence electrons − 1/2(bonding electrons) − nonbonding electrons
The formal charge on N = 5 −12(8) − 0 = 1
The formal charge on F = 7 −12(2) − 6 = 0
The formal charge on O, which forms a double bond with N atom = 6 −12(4) − 4 = 0
The formal charge on the remaining O = 6 −12(2) − 6 = −1
Therefore, the +1 formal charge is present on the nitrogen atom and the oxygen atom, which does not form a double bond with the nitrogen atom, has a negative formal charge.
Hence, the most desirable Lewis structure of nitryl fluoride would be:
NO2F Molecular Geometry
For getting further insight into the nature of chemical bonding in nitryl fluoride, we need to understand its molecular geometry, which can be anticipated from Valence Shell Electron Pair Repulsion (VSEPR) theory.
This theory has been established the concept of repulsion of valence shell electrons, as they create a negatively charged atmosphere around the atom. These valence shell electrons can be bonding as well as nonbonding electrons.
According to VSEPR theory,
1. The lone pair – lone pair repulsion is higher than lone pair-bond pair repulsion, which in turn higher than bond pair-bond pair repulsion.
3. The molecule becomes more stable by minimizing the repulsion between valence electrons and hence, increasing the distance between them.
As the nitrogen atom does not have nonbonding electrons, hence, the shape of the nitryl fluoride can be predicted from the following table.
|General formula||Number of bond pairs||Molecular shape/geometry|
Here, the nitrogen atom is a central atom with three bond pairs of electrons, as can be observed from the Lewis structure. Hence, the general formula of nitryl fluoride will be AX3 and it will have a trigonal planar geometry around the nitrogen atom.
The ideal trigonal planar geometry leads to the bond angle of 120 °.
All X atoms are not the same in NO2F and the nitrogen atom forms one double bond with the oxygen atom and two single bonds with another oxygen atom and a fluorine atom. These factors lead to slightly distorted trigonal planar geometry and bond angle are slightly different from 120 °.
The nitryl fluoride has trigonal planar geometry, which can be represented as:
Hybridization of the nitrogen atom, a central atom, in Nitryl Fluoride is determined by Valence bond Theory (VBT).
According to this theory, the atomic orbitals of the central atom cannot overlap directly with other orbitals because they cannot explain the directional property of the structure. Hence, these atomic orbitals combine and form hybrid orbitals of equivalent energy. Then, the covalent bond is formed by overlapping these hybrid orbitals with the atomic orbitals of the surrounding atoms.
The steps for the determination of hybridization of the nitrogen atom in Nitryl Fluoride by VBT are given below.
The nitrogen atom has ground state electronic configuration: [He] 2s22p3
As nitrogen atom has a positive charge, therefore, electronic configuration of nitrogen ion (cation) in its ground state: [He] 2s22p2
The electronic configuration of the nitrogen ion by excitation of 2s electron to 2p orbital i.e., in the excited state: [He] 2s12p3
According to the VSEPR theory as well as the lewis structure, nitryl fluoride constitutes one pi bond three sigma bonds from the nitrogen atom. Hence, the one 2s orbital and two 2p orbitals of the nitrogen atom will fuse, which take place in the following manner:
Hence, three sp2 hybrid orbitals are formed whereas the 2pz orbital remains unchanged and hence, 2pz orbital is used for the formation of a pi bond with the oxygen atom. The sideways overlapping of atomic orbitals results in the formation of a pi bond. The orbital diagram of nitryl fluoride, representing only the sigma bonds, is shown below.
Therefore, the hybridization of the nitrogen atom in Nitryl Fluoride is sp2 hybridization with trigonal planar geometry.
Nitryl fluoride is a polar molecule.
The polarity of any compound depends upon the net dipole moment and distribution of charges around the central atom in that compound.
If the compound has a net dipole moment and asymmetric charge distribution around the central atom, then it is a polar molecule.
From the Pauling Electronegativity chart, the electronegativity values of the nitrogen atom, fluorine atom, and oxygen atom are 3.04, 3.98, and 3.44, respectively. These electronegativity values confirm that electronegativity increases across the period.
Let us calculate the electronegativity difference for the bonds.
For N-O bond, 3.44- 3.04 = 0.40, Polar covalent bond
For N-F bond, 3.98 – 3.04 = 0.94, Polar covalent bond
Hence, the N-F bond is polar and acts as a dipole with a partial positive charge and a partial negative charge on the nitrogen atom and the fluorine atom, respectively.
Likewise, the N=O bond has a dipole moment and it is a polar bond.
However, the N-O bond has a completely negative and positive charge on the oxygen atom and nitrogen atom, respectively.
The electronegativity difference generates the net dipole moment in the molecule and therefore, nitryl fluoride is a polar molecule.
The nitryl fluoride has trigonal planar geometry with asymmetric charge distribution around the nitrogen atom and hence, it is a polar molecule.
The polar nature of nitryl fluoride makes it soluble in polar solvents.
The bonding nature of one of the inorganic compounds, the nitryl fluoride, NO2F, is discussed here.
The nitrogen atom is present at the center of the structure of nitryl fluoride. In the Lewis structure of the nitryl fluoride, there is a positive and a negative formal charge on the nitrogen atom and the oxygen atom, respectively. The nitryl fluoride has trigonal planar geometry with asymmetric charge distribution. Hence, it is a polar molecule. The nitryl fluoride shows sp2 hybridization at the nitrogen atom.
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