Nitrite [NO2]- Lewis dot structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polar or non-polar
Are you familiar with the nitrite [NO2]– ion? Quite similar in chemical formula to the nitrate [NO3]– ion, with just one less oxygen atom. But still quite different, especially in terms of Lewis structure, molecular geometry, shape, etc. The nitrite ion is an intermediate in the ecosystem’s nitrogen cycle. It is used as a reagent in the chemical and pharmaceutical industries.
In this article, we will discuss important concepts, such as how to draw the Lewis dot structure of NO2–, what is its molecular geometry or shape, electron geometry, bond angle, formal charges, hybridization, polarity, etc.
So if you are curious to know all this, then continue reading!
|Name of Molecular ion||Nitrite|
|Molecular geometry of [NO2]–||Bent, angular or V-shaped|
|Electron geometry of [NO2]–||Trigonal planar|
|Total Valence electron in [NO2]–||18|
|Overall Formal charge in [NO2]–||-1|
How to draw lewis structure of NO2-?
The Lewis structure of a NO2– consists of a nitrogen (N) atom and two oxygen (O) atoms. The nitrogen (N) atom is present at the center of the molecular ion, while the two O-atoms occupy terminal positions. Meanwhile, there is a lone pair of electrons on the central N-atom.
In this way, there are a total of 3 electron density regions around the central N-atom in the Lewis structure of [NO2]–, comprised of 2 bond pairs and 1 lone pair.
Drawing NO2– Lewis structure is not that difficult at all. You just need to grab a piece of paper and a pencil and follow the simple instructions given below to learn how to draw the Lewis dot structure of NO2–.
Steps for drawing the Lewis dot structure of [NO2]–
1. Count the total valence electrons in [NO2]–
The very first step while drawing the Lewis structure of [NO2]– is to calculate the total valence electrons present in the concerned elemental atoms.
There are two different elemental atoms present in the nitrate [NO2]– ion, i.e., a nitrogen (N) atom, and an oxygen (O) atom. The nitrogen atom contains a total of 5 valence electrons, while 6 valence electrons are present in each atom of oxygen.
The [NO2]– ion consists of 1 N-atom, 2 O-atoms, and it also carries a negative one (-1) charge, which means 1 extra valence electron. Thus, the valence electrons in the Lewis structure of [NO2]– = 1(5) + 2(6) + 1 = 18 valence electrons.
2. Choose the central atom
Electronegativity is defined as the ability of an atom to attract a shared pair of electrons from a covalent chemical bond. So, the atom which is least electronegative or most electropositive is placed at the center of a Lewis structure. This is because this atom is most likely to share its electrons with the more electronegative atoms surrounding it.
As nitrogen (N) is less electronegative than oxygen (O) so, an N-atom is placed at the center of the [NO2]– Lewis structure while the two O-atoms are spread around it, as shown below.
3. Connect outer atoms with the central atom
At this step of drawing the Lewis structure of a molecule or molecular ion, we need to connect the outer atoms with the central atom using single straight lines.
As the O-atoms are the outer atoms in the Lewis structure of the nitrite [NO2]–, ion, so both the oxygen atoms are joined to the central N-atom using straight lines, as shown below.
Each straight line represents a single covalent bond containing 2 electrons.
Now, if we count the total valence electrons used till this step out of the 18 available initially, there are a total of 2 single bonds in the structure above. Thus, 2(2) = 4 valence electrons are used till step 3.
- Total valence electrons available – electrons used till step 3 = 18-4 = 14 valence electrons.
- This means we still have 14 valence electrons to be accommodated in the Lewis dot structure of [NO2]–.
4. Complete the octet of outer atoms
There are two O-atoms present as outer atoms in the Lewis structure of [NO2]–. Each O-atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
Each N-O bond already represents 2 electrons; therefore, both the O-atoms require 6 more electrons each to complete their octet. Thus, these 6 valence electrons are placed as 3 lone pairs around each O-atom, as shown below.
5. Complete the octet of the central atom
- Total valence electrons used till step 4 = 2 single bonds + 2 (electrons placed around O-atom, shown as dots) = 2(2) +2(6) = 16 valence electrons.
- Total valence electrons available – electrons used till step 4 = 18-16 = 2 valence electrons.
These 2 valence electrons are placed as a lone pair on the central N-atom in NO2– Lewis structure.
In this way, in the above structure, the central N-atom has 2 single bonds +1 lone pair around it. This makes a total of 6 valence electrons, which means it is still short of 2 electrons in order to complete its octet.
So, to solve this issue, a lone pair present on any one of the two outer O-atoms is converted into an additional covalent bond between the central N-atom and the respective O-atom, as shown below.
Finally, the central N-atom has a complete octet with 1 single bond + 1 double bond + 1 lone pair. Also, as you can see, the octet of each outer O-atom is complete with 1 double bond + 2 lone pairs and 1 single bond + 3 lone pairs, respectively.
So, let’s check the stability of this Lewis structure using the formal charge concept.
6. Check the stability of the NO2– Lewis structure using the formal charge concept
The less the formal charge on the atoms of a molecule or molecular ion, 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 on the nitrite [NO2]– ion.
For nitrogen atom
- Valence electrons of nitrogen = 5
- Bonding electrons =1 double bond + 1 single bond = 4 + 2 = 6 electrons
- Non-bonding electrons = 1 lone pair = 2 electrons
- Formal charge = 5-2-6/2 =5-2-3 = 5-5 = 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
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
The above calculation shows that zero formal charges are present on the central N-atom and the double-bonded O-atom while the single-bonded O-atom carries a -1 formal charge which is also the charge present on the nitrite ion overall.
Thus, the NO2– Lewis structure is enclosed in square brackets, and a negative one formal charge is placed at the top right corner, as shown below.
Another important point to remember is that two distinct resonance structures are possible for representing the nitrite (NO2– ) ion. This is because anyone O-atom out of the two available can form a double covalent bond with the central N-atom.
Each resonance structure is a way of representing the Lewis structure of a molecule or molecular ion. Non-bonded electrons keep revolving from one position to another, along with a consequent movement of the double bond and thus, the formal charges on the bonded atoms keep changing.
In accordance with that, the actual NO2– structure is a hybrid of the following resonance structures.
Also check –
What are the electron and molecular geometry of NO2-?
The electron geometry of the [NO2]– ion is trigonal planar. However, it is due to the presence of a lone pair of electrons on the central N-atom that lone pair-bond pair repulsions exist in the molecular ion. Thus NO2– adopts a shape or molecular geometry different from its ideal electron pair geometry, i.e., bent or V-shaped.
Molecular geometry of [NO2]–
The molecular geometry or shape of the nitrite [NO2]– ion is bent, also known as angular or V-shaped. The presence of a lone pair of electrons on the central N-atom leads to lone pair-bond pair repulsions in addition to bond pair-bond pair electronic repulsions. This strong repulsive effect pushes the bonded atoms away from the lone pair at the center.
The N-O and N=O bonds tilt inwards, forming an inverted V (the alphabet) and consequently adopting a bent shape. Refer to the figure drawn below.
Always keep in mind that the molecular geometry or shape of a molecule or molecular ion is strongly influenced by the distinction between lone pairs and bond pairs around the central atom.
However, the ideal electron geometry only depends on the total number of electron density regions or electron domains around the central atom, regardless of the fact whether it’s a bond pair or a lone pair.
Electron geometry of [NO2]–
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule containing 3 regions of electron density around the central atom is trigonal planar.
In NO2–, a double bond around the central N-atom is considered one region of electron density. This N=O double bond, the N-O single bond, and 1 lone pair make a total of 3 electron density regions around the central N-atom; thus, the ideal electron pair geometry of NO2– is trigonal planar.
A simpler way of finding the electron and molecular geometry of NO2– 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 or molecular ion and the number of lone pairs present on it.
It is used to predict the shape and geometry of a molecule or molecular ion based on the VSEPR concept.
AXN notation for [NO2]– molecular ion
- A in the AXN formula represents the central atom. In [NO2]–, a nitrogen (N) atom is present at the center, so A =N for [NO2]–.
- X denotes the atoms bonded to the central atom. In [NO2]–, a total of two oxygen (O) atoms are bonded to the central nitrogen atom, so X = 2.
- N stands for the lone pairs present on the central atom. As per the Lewis structure of [NO2]– there is one lone pair of electrons on the central nitrogen atom, so N = 1.
Hence, the AXN generic formula for the nitrite [NO2]– ion is AX2N1.
Now, you may have a look at the VSEPR chart given below.
The VSEPR chart confirms that the molecular geometry or shape of a molecule or molecular ion with an AX2N1 generic formula is bent or V-shaped while its ideal electron geometry is trigonal planar, as we already noted down for the nitrite [NO2]– ion.
Hybridization of [NO2]–
The central N-atom has sp2 hybridization in NO2–.
The electronic configuration of nitrogen (N) is 1s22s22p3.
During chemical bonding, the 3s atomic orbital of nitrogen mixes with two half-filled 3p orbitals to produce three sp2 hybrid orbitals of equal energy.
Each sp2 hybrid orbital of nitrogen possesses a 33.3% s-character and a 67.7% p-character. One of the three sp2 hybrid orbitals contains paired electrons which are situated as a lone pair on the central N-atom in NO2–.
The other two sp2 hybrid orbitals contain a single electron each which they use for sigma (σ) bond formation on each side of NO2– by sp2-sp2 overlap (in N=O) and sp2-p orbital overlap (in N-O).
The unhybridized p-orbital of nitrogen forms the required pi (π) bond in N=O by overlapping with the p-orbital of oxygen, as shown below.
A shortcut to finding the hybridization present in a molecule or a molecular ion is by using its steric number against the table given below.
The steric number of central N-atom in [NO2]– is 3, so it has sp2 hybridization.
The [NO2]– bond angle
The O=N-O bond angle is experimentally determined to be 134° in the nitrite (NO2–) ion. It is due to the resonance present in NO2– that each N-O bond length is approx. 150 pm, as opposed to the expectation of a shorter N=O double bond and a longer N-O single bond.
Also check:- How to find bond angle?
Is NO2- polar or nonpolar?
A specific electronegativity difference of 0.4 units exists between the bonded nitrogen (E.N = 3.04) and oxygen (E.N =3.44) atoms in both the N-O and N=O bonds.
Thus, both N-O and N=O bonds are polar in the nitrite [NO2]– ion and possess a specific dipole moment value (symbol μ). The central N-atom gains a partial positive (δ+) charge while both the terminal O-atoms contain a partial negative (δ–) charge, respectively.
The bent shape of the nitrite ion further enhances the polarity effect. The dipole moments of the N-O and N=O bonds do not get canceled.
The electron cloud stays non-uniformly distributed. Thus, NO2– is overall polar (net dipole moment, µ > 0).
Read in detail–
What is the Lewis structure of NO2–?
Out of the 6 lone pairs, there are 2 lone pairs on the double-bonded O-atom, 3 lone pairs on the single-bonded O-atom, and 1 lone pair on the central N-atom in NO2–.
What is the electron geometry of NO2– according to the VSEPR theory?
The AXN generic formula for NO2– is AX2N1. Two atoms are bonded to the central N-atom, and it has a lone pair of electrons. Thus, according to the VSEPR theory, NO2– ion has a trigonal planar electron geometry.
What is the molecular geometry or shape of NO2–?
The molecular geometry or shape of NO2– is bent or V-shaped. It is due to a lone pair of electrons on the central N-atom that lone pair-bond pair repulsions exist in NO2– in addition to bond pair-bond pair repulsions.
Thus, NO2– adopts a different shape from its ideal electronic geometry.
How is the electron geometry and shape of NO2– different from that of NO3–?
Both the nitrite (NO2–) and the nitrate (NO3–) ions possess an identical trigonal planar electron geometry, as there are a total of 3 electron domains around the central N-atom in each case.
The molecular shape of NO2– is bent as opposed to that of NO3–, which is identical to its electronic geometry, i.e., trigonal planar.
There is no lone pair on the central N-atom in NO3– while there is one lone pair on the central N-atom in NO2– hence its shape or molecular geometry experience distortion.
How is the shape of NO2– different from that of NO2?
Both the nitrite [NO2]– ion as well as the nitrogen dioxide (NO2) molecule possesses the same shape, i.e., bent. The central N-atom is bonded to two O-atoms, and it has a lone pair of electrons.
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- The total valence electrons available for drawing nitrite [NO2]– ion Lewis structure are 18.
- The molecular geometry or shape of NO2– is bent or V-shaped.
- The ideal electron geometry of NO2– is trigonal planar.
- The central N-atom has sp2 hybridization in NO2–.
- The O=N-O bonded atoms form a mutual bond angle of 134°.
- NO2– is overall polar (net µ > 0).
- A -1 formal charge is present on the single-bonded O-atom in NO2– which accounts for the negative one charge on the monovalent anion overall, i.e., nitrite [NO2]–.
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