CH3NO2 Lewis structure, molecular geometry or shape, electron geometry, bond angle, hybridization, resonance structure, polar or nonpolar
CH3NO2 is the chemical formula for nitromethane, a colorless, highly flammable, pungent-smelling and toxic chemical compound. It is mainly used as a soil fumigant for synthesizing industrial and pharmaceutical antimicrobial products. Unprotected, long-term exposure to CH3NO2 proves harmful to the human nervous system.
In this article, you will learn a lot of interesting facts regards CH3NO2 chemistry, including how to draw its Lewis dot structure, what is its molecular geometry or shape, electron geometry, bond angles, hybridization, formal charges, polarity, etc.
So what are we waiting for? Let’s start reading!
|Name of Molecule||Nitromethane|
|Molecular geometry of CH3NO2||Tetrahedral (w.r.t C-atom), Trigonal planar (w.r.t. N-atom)|
|Electron geometry of CH3NO2||Tetrahedral (w.r.t C-atom), Trigonal planar (w.r.t. N-atom)|
|Hybridization||sp3 (C-atom), sp2 (N-atom)|
∠O=N-O (126.2°), ∠O=N-C (116.9°) , ∠H-C-H (112.6°)
|Total Valence electron in CH3NO2||24|
|Overall Formal charge in CH3NO2||Zero|
How to draw lewis structure of CH3NO2?
The Lewis structure of CH3NO2 consists of a carbon (C) atom at the center. It is bonded to three H-atoms and a nitro (NO2) functional group. In this way, the central C-atom is composed of a total of 4 electron density regions or electron domains.
All 4 electron density regions are constituted of bond pairs; thus, there is no lone pair of electrons on the central C-atom in the CH3NO2 Lewis structure.
Drawing the Lewis dot structure of CH3NO2 is not difficult at all, especially if you follow the simple steps given below.
Steps for drawing the Lewis dot structure of CH3NO2
1. Count the total valence electrons in CH3NO2
The very first step while drawing the Lewis structure of CH3NO2 is to calculate the total valence electrons present in its concerned elemental atoms.
As four different elemental atoms are present in CH3NO2, so you first need to look for the position of these elements in the Periodic Table.
Carbon (C) belongs to Group IV A (or 14), so it has a total of 4 valence electrons. Nitrogen (N) is present in Group V A (or 15), so it has 5 valence electrons, oxygen (O) is situated in group VI A (or 16), so it has 6 valence electrons, while hydrogen (H) lies at the top of the Periodic Table containing a single valence electron only.
- Total number of valence electrons in hydrogen = 1
- Total number of valence electrons in carbon = 4
- Total number of valence electrons in nitrogen = 5
- Total number of valence electrons in oxygen = 6
∴ The CH3NO2 molecule consists of 1 C-atom, 1 N-atom, 2 O-atoms, and 3 H-atoms. Therefore, the total valence electrons available for drawing the Lewis dot structure of CH3NO2 = 1(4) + 1(5) + 2(6) + 3(1) = 24 valence electrons.
2. Choose the central atom
In this second step, usually the least electronegative atom out of all the concerned atoms is chosen as the central atom.
This is because the least electronegative atom is the one that is most likely to share its electrons with the atoms spread around it.
Oxygen (E.N = 3.44) and nitrogen (E.N = 3.04) is more electronegative elements than both carbon and hydrogen. So neither O nor an N-atom can be selected as the central atom in the CH3NO2 Lewis dot structure.
Hydrogen (E.N = 2.20) is less electronegative than carbon (E.N = 2.55). Still, it cannot be chosen as the central atom because a hydrogen (H) atom can accommodate only 2 electrons which denotes it can form a bond with a single adjacent atom only. This denotes that H is always placed as an outer atom in a Lewis structure.
Consequently, the C-atom is placed at the center of the CH3NO2 Lewis structure, while 2 O-atoms, 1 N-atom, and 3 H-atoms, occupy terminal positions, as shown below.
3. Connect outer atoms with the central atom
This step is unique in drawing the Lewis structure of CH3NO2.
An H-atom can only form a single bond with its adjacent atom because it can only accommodate a total of 2 valence electrons.
Therefore, all three H-atoms are directly joined to the central C-atom. The single N-atom is also joined to the central C-atom, usually a single straight line.
However, as the O-atoms are present next to the N-atom, therefore these two O-atoms are joined to the N-atom and not to the C-atom directly.
Nitrogen is comparatively more electronegative than carbon, so it readily forms a bond by accepting electrons from its adjacent O-atom, not allowing the central C-atom a chance to form a bond with the latter.
A single straight line represents a bond pair containing 2 electrons. There are a total of 6 single bonds in the above structure representing 6(2) = 12 valence electrons.
∴ Hence, 12 valence electrons are already consumed out of the 24 initially available for drawing the CH3NO2 Lewis structure.
4. Complete the duplet and/or octet of the outer atoms
As we already identified, the hydrogen, oxygen, and nitrogen atoms are the outer atoms in the Lewis dot structure of CH3NO2.
Each hydrogen (H) atom requires a total of 2 valence electrons in order to achieve a stable duplet electronic configuration.
A C-H bond represents 2 valence electrons around each H-atom. This means all three H-atoms already have a complete duplet in the Lewis structure drawn till yet. Thus, we do not need to make any changes with regard to the hydrogen atoms in this structure.
In contrast to that, an N-atom and an O-atom need a total of 8 valence electrons each to achieve a stable octet electronic configuration.
Each O-atom has only an N-O single bond in the above structure, representing a total of 2 valence electrons only. It is thus deficient in 6 valence electrons in order to gain a full octet. So 6 valence electrons are placed around each O-atom as 3 lone pairs, respectively, as shown in the figure below.
Now before completing the octet of the outer N-atom, let’s see how many valence electrons are consumed out of the 24 initially available.
- Total valence electrons used till step 4 = 6 single bonds + 2 (electrons placed around oxygen atom, shown as dots) = 6(2) +2 (6) = 24 valence electrons.
- Total valence electrons – electrons used till step 4 = 24– 24 = 0 valence electrons.
As per the above calculation, all the valence electrons are already consumed; thus, we cannot place a lone pair on the N-atom to complete its octet.
But don’t worry; we can easily solve this issue by converting a lone pair from anyone O-atom adjacent to the N-atom into a covalent chemical bond.
In this way, the N-atom also has a complete octet with 2 single bonds + 1 double bond, in addition to the complete octet of both O-atoms in the CH3NO2 Lewis dot structure.
5. Complete the octet of the central atom
As seen in the above step, all 24 valence electrons initially available for drawing the CH3NO2 Lewis structure are already consumed; hence there is no lone pair of electrons on the C-atom at the centre.
The central C-atom already has 4 single bonds (3 C-H bonds + 1 C-N bond) around it. 4 single bonds mean 8 valence electrons, i.e., a complete octet of the central C-atom. So we need not make any further changes in the Lewis structure obtained below.
As a final step, we just need to check the stability of this Lewis structure. Let’s do that using the formal charge concept.
6. Check the stability of Lewis’s structure using the formal charge concept
The fewer formal charges present on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charges 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 present on CH3NO2 atoms.
For hydrogen atoms
- Valence electrons of hydrogen = 1
- Bonding electrons = 1 single bond = 2 electrons
- Non-bonding electrons = no lone pairs = 0 electrons
- Formal charge = 1-0-2/2 = 1-0-1 = 1-1 = 0
For carbon atom
- Valence electrons of carbon = 4
- Bonding electrons = 4 single bonds = 4(2) = 8 electrons
- Non-bonding electrons = no lone pairs = 0 electrons
- Formal charge = 4-0-8/2 = 4-0-4 = 4-4 = 0
For nitrogen atom
- Valence electrons of nitrogen = 5
- Bonding electrons = 2 single bonds + 1 double bond = 2(2) + 4 = 8 electrons
- Non-bonding electrons = no lone pair = 0 electrons
- Formal charge = 5-0-8/2 = 5-0-4 = 5-4 = +1
For double-bonded oxygen atom
- Valence electrons of oxygen = 6
- Bonding electrons = 1 double bond = 4 electrons
- Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons
- Formal charge = 6-4-4/2 = 6-4-2= 6-6 = 0
For single-bonded 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
The above calculation shows that zero or no formal charges are present on the double-bonded oxygen atom, the carbon atom, and three hydrogen atoms in the CH3NO2 Lewis structure. However, the nitrogen atom and single-bonded oxygen atom carry +1 and -1 formal charges, respectively.
The +1 charge cancels with -1 to yield an overall zero charge on the nitromethane molecule.
You may also note that unbonded and pi-bonded electrons present in CH3NO2 are delocalized. Anyone O-atom can give its lone pair to form a N=O double bond. This results in the following two resonance forms of CH3NO2.
Each resonance structure is a way of representing the Lewis structure of a molecule. The actual CH3NO2 structure is a hybrid of the two resonance forms.
Now that we have learned everything about CH3NO2 Lewis dot structure, we are good to proceed forward and discuss its electron geometry and molecular geometry or shape.
Also check –
What are the electron and molecular geometry of CH3NO2?
The nitromethane (CH3NO2) molecule can adopt two different shapes and electron geometries, respectively. With respect to the central carbon (C) atom, the electron geometry, and molecular geometry or shape of CH3NO2 are tetrahedral. However, with respect to the nitrogen (N) atom, its electron geometry, and molecular geometry or shape is trigonal planar.
In either case, there is no lone pair of electrons on the C or the N-atom; thus, no distortion is witnessed in the shape and/or geometry of the nitromethane molecule.
Molecular geometry of CH3NO2
The molecular geometry or shape of CH3NO2 w.r.t C-atom is tetrahedral.
To one C-atom at the center, three H-atoms and one NO2 functional group are attached along the four vertices of a regular tetrahedron. There is no lone pair of electrons on the central C-atom, so no lone pair-lone pair or lone pair-bond pair repulsions exist in the molecule.
Only a bond pair-bond pair repulsive effect exists that pushes the bonded atoms maximally apart and results in a tetrahedral shape.
However, if we consider the shape of CH3NO2 w.r.t N-atom, then that is trigonal planar. The nitrogen atom has one C-atom and two O-atoms directly bonded to it and has no lone pair of electrons. The bonded atoms occupy positions along the three vertices of an equilateral triangle, forming a trigonal planar shape, as shown below.
Electron geometry of CH3NO2
Absence of a lone pair of electrons on either the C or the N-atom results in CH3NO2 having an electron pair geometry identical to its molecular geometry or shape w.r.t both atoms, i.e., tetrahedral and trigonal planar, respectively.
A shortcut to finding the electron and the molecular geometry of a molecule is by using 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 shape and geometry of a molecule based on the VSEPR concept.
As we considered the less electronegative C-atom as the undisputed central atom while drawing the CH3NO2 Lewis structure, so here we will explain the AXN formula of CH3NO2 considering C-atom as the central atom only.
AXN notation for CH3NO2 molecule
- A in the AXN formula represents the central atom. In CH3NO2, a carbon (C) atom is present at the center, so A = C.
- X denotes the atoms bonded to the central atom. In CH3NO2, three hydrogens (H) atoms and a nitro (NO2) group are directly bonded to the central C-atom. The NO2 group is considered 1 region of electron density. In short, X = 3+1 = 4 for CH3NO2.
- N stands for the lone pairs present on the central atom. As per the Lewis structure of CH3NO2, there are no lone pairs of electrons on the central carbon; hence N = 0.
As a result, the AXN generic formula for CH3NO2 is AX4.
Now, you may have a quick look at the VSEPR chart below.
The VSEPR chart confirms that the molecular geometry or shape of a molecule with an AX4 generic formula is identical to its ideal electron pair geometry, i.e., tetrahedral, as we already noted down for the nitromethane (CH3NO2) molecule.
Hybridization of CH3NO2
The C-atom is sp3 hybridized, while the N-atom is sp2 hybridized in CH3NO2.
The electronic configuration of carbon is 1s2 2s2 2p2.
During chemical bonding, the 2s atomic orbital of carbon hybridizes with its three 2p atomic orbitals to produce four sp3 hybrid orbitals.
Each sp3 hybrid orbital is equivalent and contains a single electron only. It possesses a 25% s-character and a 75% p-character.
An sp3 hybrid orbital of carbon overlaps with the sp2 hybrid orbital of nitrogen to form the C-N sigma (σ) bond by sp3-sp2 overlap.
The other three sp3 hybrid orbitals overlap with the s-orbitals of adjacent hydrogen atoms to form the C-H sigma bonds by sp3-s overlap.
Contrarily, the remaining two sp2 hybrid orbitals of nitrogen overlap with the respective atomic orbitals of oxygen to form N-O bonds. The unhybridized p orbital of nitrogen forms the N=O pi bond by p-p overlap.
CH3NO2 hybridization can also be determined from its steric number. The steric number of the central C-atom in CH3NO2 is 4, so it has sp3 hybridization.
The CH3NO2 bond angle
It is due to the different types of bonds present in CH3NO2 that more than one different types of bond angle are present in it.
In ascending order, the H-C-H bond angle is 112.6°, and the O=N-C and O=N-O bond angles are 116.9° and 126.2°, respectively.
Similarly, the different bond lengths present in CH3NO2 are 110 pm (C-H), 123 pm (N-O), and 149 pm (C-N).
Also check:- How to find bond angle?
Is CH3NO2 polar or nonpolar?
Pauling’s electronegativity scale states that a covalent chemical bond is polar if an electronegativity difference of 0.5 to 1.6 units exists between the bonded atoms.
A small electronegativity difference of 0.35 units exists between the carbon and hydrogen atoms in each C-H bond.
However, oxygen and nitrogen are highly electronegative elements. So a comparatively higher electronegativity difference of 0.4 units is present between a nitrogen (E.N = 3.04) and an oxygen (E.N = 3.44) atom in each of the N-O and N=O bonds.
Similarly, an electronegativity difference of 0.49 units exists between a carbon (E.N = 2.55) and a nitrogen (E.N = 3.04) atom in the C-N bond. So the N-O, N=O, and C-N bonds are all strongly polar.
Oppositely charged poles develop in the molecule. The dipole moments of individually polar bonds are not canceled equally on each side of the molecular shape.
The electron cloud stays non-uniformly distributed over the molecule. Thus CH3NO2 is overall polar (net µ> 0).
Read in detail–
What is the Lewis structure for CH3NO2?
The single-bonded O-atom carries 3 lone pairs of electrons, while the double-bonded O-atom carries 2 lone pairs of electrons, respectively.
How can CH3NO2 adopt a tetrahedral and a trigonal planar shape at the same time?
The nitromethane (CH3NO2) molecule has a tetrahedral molecular shape w.r.t C-atom. It is bonded to three H-atoms and a nitro (NO2) functional group making a total of 4 electron density regions around it, thus creating a tetrahedral shape.
Contrarily, CH3NO2 has a trigonal planar shape w.r.t N-atom. The nitrogen atom is bonded to one C-atom and two O-atoms, thus forming an equilateral triangle.
Is the molecular geometry or shape of CH3NO2 the same as its ideal electron geometry w.r.t central C-atom?
Yes, the central C-atom in CH3NO2 has a total of 4 electron density regions or electron domains around it.
All four electron domains are constituted of bond pairs, and there is no lone pair of electrons present on the central C-atom.
Thus, no lone pair-lone pair or lone pair-bond pair electronic repulsions exist in the molecule; hence no distortion is witnessed in its shape and/or geometry.
How is the shape of CH3NO2 different from that of CH3NH2?
Both nitromethane (CH3NO2) and methylamine (CH3NH2) have an identical molecular geometry or shape, i.e., tetrahedral w.r.t central C-atom.
However, CH3NO2 has a trigonal planar shape w.r.t N-atom while CH3NH2 has a trigonal pyramidal shape w.r.t N-atom owing to the presence of a lone pair of electrons on the nitrogen atom in this case.
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- The total number of valence electrons available for drawing nitromethane (CH3NO2) Lewis structure is 24.
- The molecular and electron geometry of CH3NO2 w.r.t C-atom is tetrahedral.
- The molecular and electron geometry of CH3NO2 w.r.t N-atom is trigonal planar.
- The carbon atom is sp3 hybridized, while the nitrogen atom is sp2 hybridized in CH3NO2.
- CH3NO2 has two different resonance forms.
- There are multiple bond angles and bond lengths present in the nitromethane molecule.
- CH3NO2 is overall polar due to strongly electronegative nitrogen and oxygen atoms present in the molecule resulting in an unbalanced charge distribution.
- The N-atom carries a +1 formal charge, while the single-bonded O-atom carries a -1 formal charge.
- +1 formal charge cancels with -1, so the overall charge present on CH3NO2 is zero, which accounts for the incredible stability of the nitromethane Lewis structure obtained in this article.
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