Nitrous oxide (N2O) Lewis dot structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polar vs non-polar
N2O is the chemical formula for nitrous oxide or dinitrogen oxide. It is a colorless, sweet-tasting, water-soluble gas. It is famously known as laughing gas. But that is not something to laugh at, because continuous inhalation of N2O induces a euphoric effect and can even impair our decision-making ability.
However, N2O is used as a valuable medicinal agent due to its anesthetic effects.
Let’s discuss through this article, the chemistry behind the N2O molecule including how to draw its Lewis dot structure, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity nature, etc.
|Name of Molecule||Nitrous oxide|
|Molecular geometry of N2O||Linear|
|Electron geometry of N2O||Linear|
|Total Valence electron in N2O||16|
|Overall Formal charge in N2O||0|
How to draw lewis structure of N2O?
The Lewis structure of nitrous oxide (N2O) is made up of 2 nitrogen (N) atoms and one atom of oxygen (O). One of the 2 nitrogen atoms is present at the center of the molecule while the oxygen and the other nitrogen atom occupy terminal positions.
There are a total of 2 electron density regions around the central N atom, defined by an O-N single bond and an N-N triple covalent bond. There is no lone pair on the central N-atom.
If you want to draw the Lewis dot structure of N2O so immediately grab a piece of paper and a pencil and draw it with us using the simple steps given below.
Steps for drawing the Lewis dot structure of N2O
1. Count the total valence electrons in N2O
The Lewis dot structure of a molecule is referred to as a simplified representation of all the valence electrons present in it. Therefore, the very first step while drawing the Lewis structure of N2O is to calculate the total valence electrons present in the concerned elemental atoms.
The valence electrons present in an elemental atom can be determined by identifying its position in the Periodic Table of elements.
For example, nitrogen (N) is present in Group V A of the Periodic Table so it has a total of 5 valence electrons while oxygen (O) is situated in Group VI A so it has a total of 6 valence electrons in each atom.
∴ The N2O molecule consists of 1 O-atom and 2 N-atoms. Therefore, the total valence electrons available for drawing the Lewis dot structure of N2O = 2(5) + 1(6) = 16 valence electrons.
2. Choose the central atom
In the second step of drawing the Lewis structure of a molecule, we need to place the least electronegative atom at the center.
As electronegativity refers to the ability of an elemental atom to attract a shared pair of electrons from a covalent chemical bond therefore the least electronegative atom is the one that is most likely to share its electrons with other atoms.
As a nitrogen atom is less electronegative than an oxygen atom so the nitrogen (N) atom is selected as the central atom in the N2O Lewis structure. Both the N-atoms present in N2O are identical so any one N atom can be placed at the center, while the other N and O atoms are placed around it as outer atoms.
3. Connect outer atoms with the central atom
In the third step, we will join the outer N and O atoms with the central N-atom using single straight lines.
Each straight line represents a single covalent bond i.e., a bond pair containing 2 electrons. There are a total of 2 single bonds in the above diagram which means a total of 2(2) = 4 valence electrons are used till this step, out of the 16 initially available.
- Total valence electrons available – electrons used till step 3 = 16-4 = 12 valence electrons.
- This means we still have 12 valence electrons to be accommodated in the Lewis dot structure of N2O.
4. Complete the octet of outer atoms
As we already identified, one nitrogen and one oxygen atom act as the outer atoms in the N2O Lewis structure, and both N and O atoms need a total of 8 valence electrons to achieve a stable octet electronic configuration. Single bonds with the central N-atom on each side of the molecule show both N and O atoms already have 2 electrons. So, both outer N and O atoms require 6 more valence electrons to complete their octet.
Thus, these 6 electrons are placed as 3 lone pairs around each outer atom in the N2O Lewis structure, as shown in the figure below.
5. Complete the octet of the central atom and make a covalent bond if necessary
- Total valence electrons used till step 4 = 2 single bonds + electrons placed around N-atom + electrons placed around O-atom = 2(2) + 6 + 6 = 16 valence electrons.
- Total valence electrons available – electrons used till step 4 = 16-16= 0 valence electrons.
As all the 16 valence electrons initially available are already used up in drawing the Lewis dot structure of N2O so there is no lone pair on the central N-atom.
But the problem here is that there are only 2 single bonds around the central N-atom which mean there are only 4 valence electrons around it. This denotes that this nitrogen has an incomplete octet and it still needs 4 more electrons to achieve a stable octet configuration.
We can solve this problem by converting the lone pairs of electrons present on an outer atom into covalent bonds between the central N-atom and the concerned outer atom.
Now if we have to choose which of the two outer atoms (N or O) will agree to share its electrons with the central atom to fulfil its deficiency? Then, we can again take into account the electronegativity factor.
Both nitrogen and oxygen are highly electronegative atoms but in comparison, the electronegativity of N (E.N = 3.04) is less than that of O (E.N = 3.44). Hence, nitrogen is more likely to convert its non-bonded electrons into bond pairs in the N2O Lewis structure. So, 2 lone pairs on the outer N-atom are converted into bond pairs between the outer N-atom and the central N atom.
In this way, the central N-atom has a complete octet with a triple bond and a single bond. The octet of the outer N-atom is also complete with 1 triple bond and 1 lone pair while the octet of the outer O-atom is complete with 1 single bond and 3 lone pairs of electrons.
The final step is to check the stability of the Lewis structure obtained till step 5. Let us do that using the formal charge concept.
6. Check the stability of the N2O Lewis structure using the formal charge concept
The less the formal charge on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charge can be calculated using the formula given below.
- Formal charge = [ valence electrons – nonbonding electrons- ½ (bonding electrons)]
Now let us use this formula and the Lewis structure obtained in step 5 to determine the formal charges on nitrous oxide (N2O).
For the central nitrogen atom
- Valence electrons of nitrogen = 5
- Bonding electrons = 1 triple bond + 1 single bond = 3(2) + 2 = 8 electrons
- Non-bonding electrons = no lone pair = 0 electrons
- Formal charge = 5-0-8/2 = 5-0-4 = 5-4 = +1
For the outer nitrogen atom
- Valence electrons of nitrogen = 5
- Bonding electrons = 1 triple bond = 3 (2) = 6 electrons
- Non-bonding electrons = 1 lone pair = 2 electrons
- Formal charge = 5-2-6/2 = 5-2-3 = 5-5 = 0
For oxygen atom
- Valence electrons of oxygen = 6
- Bonding electrons = 1 single bond = 2 electrons
- Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons
- Formal charge = 6-6-2/2 = 6-6-1 = 6-7 = -1
This calculation shows that a zero formal charge is present on the outer nitrogen while the central N atom and the outer O-atom carry a +1 and a -1 formal charge respectively.
So, look at the best and most stable lewis structure of N2O.
An interesting fact is that there are three possible resonance structures of N2O, as shown below. The actual structure of nitrous oxide is a hybrid of these resonance structures. Each resonance structure is a way of representing the Lewis structure of a molecule.
The above resonance structures show that the non-bonded and pi-bonded electrons in N2O are delocalized. These electrons move from one position to another on the molecule followed by a change in the formal charges present on the bonded atoms.
Resonance structure 2 is the best possible Lewis representation of N2O as the more electronegative O-atom bears a -1 charge while the less electronegative N-atom carries a +1 charge in this structure.
However, the positive and negative formal charges cancel out to give an overall zero charge on the molecule which accounts for its stability.
Now, that we have discussed everything about N2O Lewis structure, we are good to proceed to the next section of this article where we will discuss the shape and geometry of N2O.
Also check –
What are the electron and molecular geometry of N2O?
Nitrous oxide (N2O) molecule has an identical electron and molecular geometry or shape i.e., linear. There is no lone pair present on the central N-atom in N2O therefore there is no distortion in the shape and geometry of the molecule.
Molecular geometry of N2O
The nitrous oxide (N2O) molecule has a linear shape and molecular geometry. Both the nitrogen and oxygen atoms bonded to the central N-atom lie on a straight line, in a planar arrangement.
There is no lone pair of electrons on the central nitrogen atom therefore no bond pair-lone pair and lone pair-lone pair electronic repulsions exist in the molecule. N-N and N-O bond pair-bond pair electronic repulsions exist which push the bond pairs as far apart from each other as possible and the O and N atoms occupy the terminals, as shown in the figure below.
Electron geometry of N2O
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule containing a total of 2 electron density regions around the central atom is linear.
In N2O there is 1 single bond and 1 triple bond around the central nitrogen atom which makes a total of 2 electron density regions. Thus, its electron geometry is also linear.
An easy way to find the shape and geometry of the molecule is to use the AXN method.
AXN is a simple formula to represent the number of atoms bonded to the central atom in a molecule and the number of lone pairs present on it.
It is used to predict the geometry or shape of a molecule using the VSEPR concept.
AXN notation for N2O molecule
- A in the AXN formula represents the central atom. In the N2O molecule, a nitrogen (N) atom is present at the center so A = Nitrogen.
- X denotes the atoms bonded to the central atom. In N2O, 1 nitrogen (N) atom and 1 oxygen (O) atom are bonded to the central N so X = 1+1 =2.
- N stands for the lone pairs present on the central atom. As per the Lewis structure of N2O, there is no lone pair on central nitrogen so N=0.
Thus, the AXN generic formula for the N2O molecule is AX2.
Now, you may have a look at the VSEPR chart below.
According to the VSEPR chart above, the ideal electron pair geometry and molecular geometry or shape of a molecule with an AX2 generic formula are linear, as we already noted down for N2O.
Hybridization of N2O
Both the nitrogen atoms in N2O are sp hybridized while the oxygen atom is sp3 hybridized. As the hybridization present in a molecule is determined as per the central atom so the nitrous oxide (N2O) molecule has sp hybridization.
The electronic configuration of a nitrogen (N) atom is 1s2 2s2 2p3.
During chemical bonding, the 2s electrons of nitrogen get unpaired, consequently, one 2s and one 2p orbital mix to yield two equivalent sp hybrid orbitals. Each sp hybrid orbital contains a single electron only and possesses a 50% s-character and a 50% p-character.
One sp hybrid orbital of central N forms the N-N sigma (σ) bond while the other sp hybrid orbital forms the N-O sigma (σ) bond by sp-sp and sp-sp3 overlap with nitrogen and oxygen atoms respectively.
The second sp hybrid orbital of outer nitrogen accommodates the lone pair of electrons. Similarly, the remaining three sp3 hybrid orbitals of oxygen contain the three lone pairs of electrons present on it.
A shortcut to finding the hybridization present in a molecule is by using its steric number against the table given below. The steric number of central N in N2O is 2 so it has sp hybridization.
The N2O bond angle
As all three bonded atoms (O, N, and N) lie on a straight line in the linear N2O molecule therefore they form a mutual bond angle of 180°. The N-N triple covalent bond is stronger and shorter in length i.e., 112.6 pm while the N-O single covalent bond is longer i.e., 118.6 pm.
Also check:- How to determine bond angle?
Is N2O polar or nonpolar?
Nitrous oxide (N2O) is a polar molecule overall. No electronegativity difference exists between the identical nitrogen atoms in N2O while a specific electronegativity difference of 0.4 units exists between the central N (E.N = 3.04) and the outer O-atom (E.N = 3.44). Thus, the N-O bond is polar.
Due to its high electronegativity, the oxygen atom not only attracts the shared electron cloud between the N-O bond but also attracts N-N electrons. The electron cloud stays non-uniformly distributed in the molecule overall thus N2O is polar (net dipole moment, µ = 0.160 D).
Read in detail–
How many total lone pairs are there in the Lewis structure of N2O?
There are a total of 16 valence electrons i.e., 8 electron pairs in the Lewis structure of N2O. Out of the 8 electron pairs, there are 4 bond pairs and 4 lone pairs.
There is 1 lone pair on an outer N-atom and three lone pairs on the oxygen atom which means there is no lone pair on the central N-atom in the N2O Lewis structure, as shown below.
Why does the Lewis dot structure of N2O have Nitrogen (N) as the central atom and not Oxygen (O) as the central atom?
While drawing the Lewis structure of a molecule, usually the less electronegative atom is chosen as the central atom. This is because the less electronegative atom is more likely to share its electrons with other atoms in its surroundings.
Nitrogen (E.N = 3.04) is less electronegative than oxygen so it is placed as the central atom in the N2O Lewis dot structure.
What is the molecular geometry of N2O?
The N2O molecule has a linear shape and molecular geometry which is identical to its electron pair geometry owing to the presence of 2 regions of electron density around the central N-atom.
How is the molecular geometry of three greenhouse gases different i.e., O3, N2O, and H2O?
The ozone (O3) molecule has a trigonal planar electron geometry while its molecular geometry or shape is angular or bent due to a lone pair of electrons present on the central O-atom. [Check in detail about – O3 molecular and electron geometry]
The nitrous oxide (N2O) molecule has a linear electron and molecular geometry or shape. There is no lone pair on the central N-atom in N2O thus no distortion is present in the shape and geometry of the molecule.
The water (H2O) molecule has a tetrahedral electronic geometry while its molecular geometry or shape is also bent or angular. There are 2 lone pairs of electrons on the central O atom in H2O. [Check in detail about – H2O molecular and electron geometry]
The difference in these molecular geometries shows that the lone pairs of electrons present on the central atom strongly influence the shape of a molecule.
Compare the molecular geometry of N2O with NO2?
The nitrous oxide or dinitrogen oxide (N2O) molecule has a linear shape and molecular geometry while the nitrogen dioxide (NO2) molecule has a bent shape while its ideal electron pair geometry is trigonal planar.
There is no lone pair of electrons on the central N atom in N2O while a lone pair of electrons is present on the central N-atom in NO2 which distorts its shape and geometry.
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- The total number of valence electrons available for drawing the nitrous oxide (N2O) Lewis structure is 16.
- N2O molecule has an identical electron and molecular geometry or shape i.e., linear.
- In the Lewis dot structure of N2O, there are a total of 4 bond pairs and no lone pair around the central nitrogen atom.
- The central N atom in N2Ois sp
- The N-N-O bond angle is 180° due to the linear arrangement while the N-N and N-O bond lengths are 112.6 pm and 118.6 pm respectively.
- N2O is a polar molecule (net μ= 0.160 D).
- The overall formal charge on N2 O is zero. +1 formal charge on the central N-atom cancels with -1 formal charge on the outer O-atom which accounts for the exceptional stability of the molecule.
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