Who has not heard about NO2, Nitrogen Dioxide?
It is one of the most common gaseous molecules having a reddish-brown hue. It can be cooled and compressed into a yellowish-brown liquid for tendshipping and transport.
A highly toxic poisonous chemical compound, NO2 is a major air pollutant and belongs to the group of oxides of nitrogen.
This chemical can be used for bleaching and sterilization purposes.
It has applications in tobacco and explosive industries. NO2 acts as an intermediate in HNO3 acid manufacturing as well as an oxidizer for fuels in rockets and space probes.
Below are the laboratory preparation methods of Nitrogen Dioxide.
A Brief Intro
Inside an atom, we have the positively charged nucleus surrounded by electrons in their shells forming a negatively charged cloud.
The electrons in the outermost shell which determine the valency are known as valence electrons.
Lewis Structure is a diagrammatic representation of any given molecule with the help of the constituent atoms and the position and arrangement of electrons to form bonds and lone pairs.
This is a limited theory on chemical bonding nature and electronic structure but provides a simple viewpoint towards the formation of any molecular composition.
Let us talk about drawing the Lewis Structure for Nitrogen Dioxide ( NO2 ).
Lewis Structure of NO2
A molecule of nitrogen dioxide consists of one nitrogen atom and two atoms of oxygen.
Let us look at the periodic table.
Nitrogen belongs to group 15( or group 5) and has an atomic number of 7, therefore has a valency of 5. Oxygen belongs to group 16 ( or group 6) and has an atomic number of 8, therefore a valency of 6.
Total valence electron number in NO2 = 5 + 6*2
Now, to find out which atom will act as a central one, we have to look at Pauling’s Electronegativity chart.
Nitrogen is the least electronegative element between the two, hence N will be the central atom.
This is how we have placed the nitrogen in the center surrounded by the two O atoms.
Now, with the help of dot structures, we will surround the atoms by their valence electrons.
We have followed the octet rule here.
Do you know that the elements present in the main groups of the periodic table have a tendency to form an eight-electron arrangement in their valence shell following noble gas configurations?
This is known as the octet rule and is an important concept to sketch the most probable Lewis Structure of a molecule.
Look at the diagram again. We have already made the two O atoms achieve the Neon configuration.
Now, the Nitrogen atom has only five electrons surrounding it and the total number of valence electrons has been used up already.
For the Nitrogen atom to have a near octet configuration, we are going to shift two electrons from one O atom and form a double bond.
Here we go.
Nitrogen now has 7 electrons and achieved a near octet configuration.
Let us check the formal charge.
Formal Charge of O ( in a single bond with N )
= 6 – 0.5*2 – 6 = 6 – 1 – 6 = -1.
Formal Charge of O ( in a double bond with N )
= 6 – 0.5*4 – 4 = 6 – 2 – 4 = 0.
The formal charge of central N atom
= 5 – 0.5*6 – 1 = 5 – 3 – 1 = 1.
If we now find the summation of the total formal charges we get a net result of 0.
Hence, the molecule is neutral and the elements are present in their least possible formal charge values.
Therefore, the most suitable Lewis Structure of NO2 is:
Let us look at some of the drawbacks of Lewis Structure:
- It can only provide a 2D representation and not a 3D one.
- Lewis Structure can only talk about the arrangement of electrons but not how the electrons are shared.
- It can only identify the type of bond formation but cannot elaborate on how the bond formation takes place.
This is where the role of VSEPR theory comes:
VSEPR stands for Valence Shell Electron Pair Repulsion model.
According to this theory, since electrons carry like charges ( ‘-’ ), they are bound to experience repulsion and this strength of repulsion varies across bonded and lone pair electrons.
So, this repulsion force needs to be minimized to form a stable or balanced polyatomic molecular or ionic structure.
VSEPR thus helps in predicting the 3D molecular geometry or shape of a molecule along with the bond angles.
What is the Molecular Geometry of NO2?
In the VSEPR model, we have the AXnEx notation.
A stands for the central atom, here we have nitrogen ( N ).
X stands for the surrounding atoms, here we have two O atoms.
n stands for the number of atoms around the central element, we have the value of 2.
E stands for non-bonded electrons ( usually a lone pair), x stands for the number.
Here, since we have just one lone electron, Let us take it as 1.
We have an AX2E1 notation.
We can see that NO2 has a bent molecular geometry and the angle is around 120 degrees.
But here we have some exceptions. In NO2, we have 2 Bond Pairs and 1 lone electron.
If we look at the nitrite ion NO2-, we have 2 Bond Pairs and 1 Lone pair of electrons.
There we also have a bent structure but since repulsion strength is of the order of LP-LP > LP-BP > BP-BP, the bond angle order is as below
NO2+ > NO2 > NO2-.
Therefore, the bond angle of NO2 is around 134 degrees. The bond length is around 1.20 Å.
Do you know that an electron can be described with the help of its wave function?
We are aware of the concept of atomic orbitals, the mathematical probability functions indicating the presence of electrons in any regional space.
When wave functions of atomic orbitals combine or form a fusion, it results in hybrid orbitals and the process is known as orbital hybridization.
Hybridization of NO2
The single bond in N-O has one sigma bond and the double bond N=O has one sigma bond and one pi bond. Pi bond doesn’t take part in hybridization.
Now, in NO2 we have 17 valence electrons, therefore this is an odd-electron system. For single electron species, we have the following rule to follow:
If the oxidation state of the central atom is found positive then the electron will participate in the hybridization process but if the oxidation state is negative, it will not participate.
In nitrogen dioxide,
Steric No = 2 sigma + 1 lone electron.
The central N is less electronegative than O. Therefore, the charge on N is ‘+’.
So, the oxidation state is found positive and the lone electron will take part in hybridization.
steric no = 2 + 1 = 3. Therefore, the hybridization of NO2 is sp2.
What is Polarity?
Polarity is an important concept of chemistry. Every molecular composition has a property of polarity which determines whether the said molecule is polar or non-polar. This depends on the nature of chemical bonding which influences the shape and geometry.
A diatomic homogeneous molecule that consists of like atoms always results in zero dipole moment and hence non-polar.
However, inside a heterogenous polyatomic molecule, we have several atomic elements with different varying electronegativity values. This induces partial charges on atoms which if not canceled out results in polar molecules.
Also, the irregularity and asymmetry of molecules with the help of lone pairs and bonds cause polarity due to the uneven distribution of charges.
What makes NO2 a polar molecule?
Nitrogen has an electronegativity value of 3.04 and oxygen has that of 3.44.
This means there is an electronegativity difference between the two atomic elements. Although the difference is quite less, we have an asymmetrical bent molecular structure that induces the net dipole moment and makes NO2 a polar molecule in reality.
For detailed information, you should also once read out an article on the Polarity of NO2.
Molecular Orbital (MO) Diagram
Electrons can have both particle and wave-like nature. Molecular Orbital theory is a concept of quantum mechanics that attempts to explain the chemical bonding inside any molecule.
In this theory, we get to know that valence electrons can be shared amongst all constituent atoms and atomic orbitals from different atoms combine to form molecular orbitals ( MOs ).
Here, we will incorporate certain terminologies like anti-bonding, non-bonding, and bonding orbitals. Also, we have the concept of HOMO ( Highest Occupied Molecular Orbital) and LUMO ( Lowest Unoccupied Molecular Orbital).
MO Diagram for NO2
Let us look at the electronic configuration of both N & O.
N: 1s2 2s2 2p3
O: 1s2 2s2 2p4 ( we have two O atoms in NO2 )
The six electrons present in the 1s orbital do not take part in bonding, therefore will play the role of non-bonding orbitals.
The two electrons of 1s2 in the Nitrogen atom take part in σ2s MO. The oxygen atoms contribute to 2 lone pairs each.
The remaining electrons in p orbitals of N and O form the σ2px, 𝜋2py, 𝜋2pz, and σ*2s.
NO2 is one of the most common heteronuclear diatomic molecules. We, in this article, have discussed in detail the nature of the chemical bonding of the molecule.
We hope you have gone through the steps of forming a perfect Lewis structure, usage of notations of VSEPR theory to calculate the bond angles and predict the molecular geometry, the process of hybridization, polarity, and also the molecular orbital concept of bonding.