Nitrous acid (HNO2) Lewis structure, molecular geometry or shape, resonance structure, bond angle, hybridization, polar or nonpolar
Is nitrous acid (HNO2) the twin sister of the very well-known nitric acid (HNO3)? Well, you could say so, as the oxidation of HNO2 leads to the formation of HNO3. Nitrous acid (HNO2) is a weak, monoprotic acid.
In this article, our focus will be on discussing how to draw the Lewis dot structure of HNO2, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity, etc.
So without further delay, let us start reading to explore all these valuable facts.
The Lewis structure of HNO2 consists of three different elemental atoms. A nitrogen (N) atom is present at the center of the molecule. It is directly bonded to an oxygen (O) atom and a hydroxyl (OH) functional group, one on each side. The central N-atom also consists of a lone pair of electrons.
In this way, there are a total of 3 electron density regions around the central N-atom in the HNO2 Lewis structure.
You can easily draw the Lewis dot structure of HNO2, and to do so, you just need to grab a piece of paper and a pencil and follow the simple steps given below.
In HNO2, there are atoms from three different elements of the Periodic Table. So you need to look for the group numbers of these elements in the Periodic Table.
Nitrogen (N) belongs to Group V A (or 15), so it has a total of 5 valence electrons. Oxygen (O) is present 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.
∴ The HNO2 molecule consists of 1 N-atom, 1 H-atom, and 2 O-atoms. Therefore, the total valence electrons available for drawing the Lewis dot structure of HNO2 = 1(5) + 1 + 2(6) = 18 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 is more electronegative than both nitrogen and hydrogen, so it cannot be chosen as the central atom. Hydrogen is less electronegative than nitrogen, but it can also not be chosen as the central atom because a hydrogen (H) atom can accommodate only 2 electrons; hence 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.
As a result, the nitrogen (N) atom is chosen as the central atom in the Lewis dot structure of HNO2, while one H-atom and two O-atoms are placed around it as outer atoms, as shown in the figure below.
3. Connect outer atoms with the central atom
In this step, all the outer atoms are joined to the central atom using single straight lines. But as we discussed already, hydrogen can form a single bond with one adjacent atom only.
So H is first connected to its adjacent O-atom. Contrarily, both the outer O-atoms are directly joined to the central N-atom using straight lines, as shown below.
Each straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons. There are a total of 3 single bonds in the above diagram.
As 3(2) = 6, that means 6 valence electrons are already consumed out of the 18 initially available. But we still have 18 – 6 = 12 valence electrons to be accommodated in the Lewis dot structure of HNO2.
4. Complete the duplet and/or octet of the outer atoms
As we already identified, the hydrogen and oxygen atoms are the outer atoms in the Lewis dot structure of HNO2.
Each hydrogen (H) atom requires a total of 2 valence electrons in order to achieve a stable duplet electronic configuration.
The O-H single bond represents 2 valence electrons around the H-atom. This means it already has a complete duplet, and we do not need to make any changes in the Lewis structure obtained so far with regard to the hydrogen atom.
In contrast, an O-atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
An O-H bond and an N-O bond represent a total of 2 + 2 = 4 electrons around this oxygen atom. This denotes it is still deficient in 4 more electrons to complete its octet. So these 4 electrons are placed as 2 lone pairs around this O-atom.
The other oxygen atom only forms an N-O bond which denotes 2 electrons. Thus, it is still deficient in 6 more electrons that are required to complete its octet. Consequently, these 6 electrons are placed as 3 lone pairs around this O-atom, as shown below.
5. Complete the octet of the central atom and make a covalent bond if necessary
Total valence electrons used till step 4 = 3 single bonds + electrons placed around the O-atom forming O-H bond + electrons placed around the other O-atom, shown as dots = 3(2) + 4 + 6 =16 valence electrons.
Total valence electrons – 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 the HNO2 Lewis structure.
In this way, in the above diagram, 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 a terminal O-atom (not the O-H bonded O-atom) 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 2 single bonds + 2 lone pairs, respectively.
So, let’s check the stability of this Lewis structure 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.
Non-bonding electrons = no lone pairs = 0 electrons
Formal charge = 1-0-2/2 =1-0-1 = 1-1= 0
The absence of any formal charges on the bonded atoms makes the HNO2 Lewis structure exceptionally stable.
Another interesting fact is that the following two resonance structures are possible for nitrous acid (HNO2).
Each resonance structure is a way of representing the Lewis structure of a molecule. However, out of the above two HNO2 resonance structures, structure 1 is much more stable than resonance structure 2.
This is because zero formal charges are present on all bonded atoms in structure 1, while in structure 2, a +1 formal charge is present on one O-atom, while a -1 charge is present on the other O-atom; therefore, it is not very stable.
As a result, these two resonance structures are not equivalent. The actual HNO2 structure is a hybrid of the two resonance structures with a greater proportion of structure 1 than structure 2 in this hybrid.
What are the electron and molecular geometry of HNO2?
The electron geometry of HNO2 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 molecule. Thus HNO2 adopts a shape or molecular geometry different from its ideal electron pair geometry, i.e., bent or V-shaped.
Molecular geometry of HNO2
The molecular geometry or shape of nitrous acid (HNO2) 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=OH bonds tilt inwards, forming an inverted V (the alphabet) and consequently adopting a bent shape.
Always keep in mind that the molecular geometry or shape of a molecule 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 HNO2
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 HNO2, the hydroxyl (OH) functional group is considered one region of electron density. The OH group, along with 1 O-atom and a lone pair, makes a total of 3 electron density regions or electron domains around the central N-atom in HNO2. Thus its electron geometry is trigonal planar.
A more straightforward way of determining the shape and geometry of a 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 shape and geometry of a molecule based on the VSEPR concept.
AXN notation for the HNO2 molecule
A in the AXN formula represents the central atom. In HNO2, nitrogen (N) acts as the central atom, so A = N.
X denotes the electron domains bonded to the central atom.1 O-atom and 1 OH group are directly bonded to the central N-atom in HNO2; thus, X=2.
N stands for the lone pairs present on the central atom. As per the Lewis dot structure of HNO2, there is 1 lone pair present on central nitrogen; thus, N=1.
So, the AXN generic formula for HNO2 is AX2N1.
Now, look at the VSEPR chart below to identify where you find AX2N1.
The VSEPR chart confirms that the molecular geometry or shape of a molecule 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 nitrous acid (HNO2) molecule.
Hybridization of HNO2
The central N-atom has sp2 hybridization in HNO2.
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 HNO2.
The other two sp2 hybrid orbitals contain a single electron each which they use for sigma (σ) bond formation on each side of HNO2 by sp2-sp2 overlap (in N=O) and sp2-sp3 orbital overlap (in N-OH).
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 is by using its steric number against the table given below.
The steric number of central N-atom in HNO2 is 3, so it has sp2 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The HNO2 bond angle
The ideal bond angle in a symmetrical trigonal planar molecule is 120°. However, a strong electron repulsive effect in the bent HNO2 molecule decreases the concerned bond angles. O=N-O bond angle becomes 110° while the N-O-H bond angle is even smaller, i.e., approx., 102° in HNO2.
Conversely, as per the different types of bonds present, there are three different bond lengths in the nitrous acid molecule. The N=O bond length is 120 pm, the N-O bond length is 146 pm, and the O-H bond length is 98 pm, as shown below.
Pauling’s electronegativity scale states that a covalent bond is polar if the bonded atoms possess an electronegativity difference between 0.5 to 1.6 units.
An electronegativity difference of 0.4 units exists between the bonded nitrogen (E.N = 3.04) and oxygen (E.N = 3.44) atoms in each of the N-O and N=O bonds in the HNO2 molecule. Thus both N-O and the N=O bonds are slightly polar in the HNO2 molecule.
Similarly, an electronegativity difference of 1.24 units exists between the bonded hydrogen (E.N = 2.20) and oxygen atoms (E.N = 3.44). So the O-H bond is also polar and possesses a specific dipole moment value (symbol µ).
The dipole moments of these polar bonds do not get canceled equally; thus, the HNO2 molecule overall is polar with a non-uniformly distributed electron cloud (net µ > 0).
The Lewis structure for nitrous acid (HNO2) displays a total of 18 valence electrons i.e., 18/2 = 9 electron pairs.
Out of these 9 electron pairs, there are 4 bond pairs and 5 lone pairs of electrons.
The bond pairs are comprised of a N=O bond, an N-O bond, and an O-H bond.
Out of these 5 lone pairs, there are 2 lone pairs on each O-atom, while 1 lone pair is present on the central N-atom.
Is it correct to draw the Lewis structure of HNO2 in both ways as H-O-N=O and O-N=O-H?
The following two forms are known as the resonance structures of HNO2. Each resonance structure is a way of representing the Lewis structure of a molecule.
So yes, HNO2 can be represented both ways H-O-N=O and O-N=O-H. However, the former is much more stable than the latter due to fewer formal charges present.
Does HNO2 have the same shape and electron geometry?
No.The molecular geometry or shape of HNO2 is different from its ideal electron pair geometry. This is due to a lone pair of electrons on the central N-atom in HNO2 that lone pair-bond pair repulsions exist in the molecule in addition to a bond pair-bond pair repulsive effect.
Thus the molecule adopts a distorted/ asymmetrical bent shape as opposed to its trigonal planar electron geometry.
How is the shape of HNO2 different from that of HNO3?
The AXN generic formula for HNO2 is AX2N1, so according to the VSEPR concept, the shape of nitrous acid is bent or V-shaped.
Contrarily, the AXN generic formula for HNO2 is AX3, so its shape is trigonal planar. The bonded atoms lie along the three vertices of an equilateral triangle.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
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