Nitric oxide (NO) Lewis structure, molecular geometry or
shape, electron geometry, bond angle, hybridization, formal
charges, polar or nonpolar
NO is the chemical formula for nitric oxide, aka nitrogen monoxide, an odorless, colorless gas naturally found in the atmosphere.
NO is a highly reactive and unstable diatomic molecule. Its extreme reactivity is accredited to the presence of an unpaired electron in its structure, making it a free radical. Under atmospheric oxygen, NO readily transforms into NO2, which dissolves in water to form nitric acid (HNO3).
In this article, you will find why NO is a free radical as you will learn to draw the Lewis dot structure of NO. Moreover, you will also learn interesting facts about molecular geometry, electron geometry, bond angles, formal charges, hybridization and polarity of nitric oxide (NO).
So, for all this valuable information, continue reading the article!
The Lewis structure of NO comprises only two atoms, i.e., a nitrogen (N) atom and an oxygen (O) atom, double-covalently bonded to each other. The O-atom carries 2 lone pairs of electrons, while 1 lone pair of electrons and an unpaired electron is present on the N-atom in NO.
Drawing the Lewis structure of NO may seem tricky at first, but you can easily learn it by following the uncomplicated steps given below.
Steps for drawing the Lewis dot structure of NO
1. Count the total valence electrons present in NO
In the first step of drawing the Lewis structure of a molecule, you need to count the total valence electrons present in it.
The valence electrons of an elemental atom can be determined from its group number.
NO is composed of only two elemental atoms. Nitrogen (N) is present in Group V A (or 15) of the Periodic Table of Elements, meaning it has 5 valence electrons. In contrast, Oxygen (O) is present in Group VIA (or 16). Each O-atom thus has a total of 6 valence electrons.
Total number of valence electrons in nitrogen = 5
Total number of valence electrons in oxygen = 6
The NO molecule comprises 1 N-atom and 1 O-atom only.
∴ Therefore, the total valence electrons available for drawing the Lewis dot structure of NO = 1(5) + 1(6) = 11 valence electrons.
2. Choose the central atom
By convention, the least electronegative atom out of all those available is chosen as the central atom while drawing the Lewis structure of a molecule.
The least electronegative atom can easily form covalent bonds with other atoms by sharing its electrons.
In nitric oxide (NO), nitrogen (E.N = 3.04) is less electronegative than oxygen (E.N = 3.44), so it is chosen as the central atom.
3. Connect the outer atom with the central atom
Only an O-atom is the outer atom in the Lewis structure drawn to this step. So the O-atom is joined to the central N-atom using a single straight line.
The straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons. This implies that 2 valence electrons are already consumed from the 11 initially available in drawing the NO Lewis structure.
4. Complete the octet of the outer atom
The outer O-atom needs a total of 8 valence electrons in order to achieve a complete octet electronic configuration.
An N-O bond represents 2 valence electrons only. So 8- 2 = 6 valence electrons are placed as 3 lone pairs around the O-atom, as shown below.
5. Put the remaining electrons on the central atom and convert lone pairs into covalent bonds if necessary
Total valence electrons used till step 4 = 1 N-O single bond + electrons placed around the O-atom, shown as dots =1(2) + 6 = 8 valence electrons.
Total valence electrons – electrons used till step 4 = 11 – 8 = 3 valence electrons.
Thus, the remaining 3 electrons are placed as a lone pair and an unpaired electron on the N-atom in the NO Lewis structure.
However, the N-atom still has an incomplete octet in this Lewis structure. It only consists of a total of 5 valence electrons surrounding it.
So to overcome this electron deficiency, 1 lone pair from the terminal O-atom is converted into an extra bond between the N-atom and the O-atom, as shown below.
The central N-atom still has an incomplete octet (only 7 electrons), but we must accept that it is impossible to give every atom a full octet electronic configuration with an odd number of electrons initially available.
So, sufficing at the NO Lewis structure obtained above, we just need to check its stability by applying the formal charge concept, as shown in the next step.
6. Check the stability of Lewis’s 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 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 NO-bonded atoms.
Zero formal charges on both the N-atom and the O-atom ensure no overall charge is present on the NO molecule. Hence, the structure obtained below is the correct Lewis representation of nitric oxide (NO).
However, you must note that it is not a stable structure. NO is a free radical. The presence of an unpaired electron makes it a highly reactive and unstable molecule.
Furthermore, NO can be represented by the following resonance forms.
Each resonance form is a way of representing the Lewis structure of a molecule.
The non-bonded and pi-bonded electrons keep moving from one position to another on the molecule. The actual molecular structure is a weighted average of all the possible resonance forms, known as the resonance hybrid.
Structure I is the most contributing resonance form as it is a neutral molecule with zero formal charges.
Structure II holds a greater contribution than structures III and IV as the -1 formal charge is present on the more electronegative atom (O-atom); also, the N-atom has a complete octet in it.
It is difficult to give definite answers related to the molecular shape, electron geometry, and hybridization of a molecule that has an odd number of valence electrons.
But we have tried our best to do so in our next section by using the two most important resonance forms of NO.
What are the electron and molecular geometry of NO (Nitric oxide)?
The NO molecule possesses an identical electron and molecular geometry or shape, i.e., linear. The central N-atom carries lone pairs and an unpaired electron on it. However, these non-bonded electrons do not disturb the overall geometry and orientation of the molecule.
Molecular geometry of NO
The molecular geometry or shape of NO is linear.
As per resonance form I, the unbonded electrons on the double-covalently bonded N-atom and O-atom are maximally separated. The lone pair-lone pair and lone pair-bond pair electronic repulsions are minimized to zero.
Thus, the molecule occupies a perfectly symmetrical, linear shape. There is no distortion in the shape and geometry of the molecule.
Electron geometry of NO
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.
To avoid any confusion, we can better explain the electron geometry of NO w.r.t resonance form II. The central N-atom consists of 1 lone pair, and it is triple-covalently bonded to the terminal O-atom.
The triple covalent bond is considered as 1 electron density region while determining the electron geometry of a molecule. This makes a total of 2 electron density regions around the central N-atom in NO; thus, its ideal electron pair geometry is linear.
An easy trick to finding a molecule’s electron and molecular geometry is using the AXN method.
AXN is a simple formula representing the number of bonded atoms and lone pairs on the central atom.
It is used to predict the shape and geometry of a molecule using the VSEPR concept.
AXN notation for NO molecule
A in the AXN formula represents the central atom. As per resonance form II, in the NO molecule, a nitrogen (N) atom is present at the center, so A = N.
X denotes the atoms bonded to the central atom. In NO, only an O-atom is bonded to the nitrogen atom. So X = 1 for NO.
N stands for the lone pairs present on the central atom. As per the Lewis structure of NO, the central N-atom carries 1 lone pair of electrons. Thus, N= 1 for NO.
As a result, the AXN generic formula for NO is AX1N1or simply AXN.
Now, you may have a look at the VSEPR chart below.
The VSEPR chart confirms that the molecular geometry or shape of a molecule with an AXN generic formula is identical to its electron pair geometry, i.e., linear, as we already noted down for nitrogen monoxide/nitric oxide (NO).
Hybridization of NO
As per a linear shape, the central N-atom is sp hybridized in NO. In contrast, the O-atom is sp2 hybridized.
The electronic configuration of nitrogen (N) is 1s2 2s2 2p3. During chemical bonding, the 2s atomic orbital of nitrogen hybridizes with one of the three half-filled 2p orbitals to produce 2 sp hybrid orbitals.
Each sp hybrid orbital possess a 50 % s-character and a 50 % p-character. However, the two sp hybrid orbitals of N-atom are not equivalent in NO.
One sp hybrid orbital consist of paired electrons which are situated as a lone pair on nitrogen in NO. Contrarily, the other sp hybrid orbital contains a single unpaired electron. It forms the N-O sigma bond by sp-sp2 overlap with the adjacent O-atom.
The unhybridized p-orbital of nitrogen forms a pi-bond with the O-atom by p-p orbital overlap.
The unpaired electron of nitrogen is accommodated in the second unhybridized p-orbital.
Also, the hybridization present in NO can be determined using resonance form II and the table is given below. The steric number of the N-atom in NO is 2, so it has sp hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The bond angle of NO
The nitrogen and oxygen atoms lie on a perfectly straight line in NO, forming a bond angle of 180°.
The nitrogen-oxygen bond length is experimentally determined to be 116.4 pm. You may note that this bond length is intermediate between the bond lengths of a typical N-O single (136 pm) and N=O double bond (115 pm), thus, another evidence of the electronic delocalization and resonance present within the nitric oxide molecule.
A specific electronegativity difference of 0.4 units exists between a nitrogen (E.N = 3.04) and an oxygen (E.N = 3.44) atom.
Oxygen being strongly electronegative, attracts the N=O bonded electrons toward itself to a greater extent. The N-atom thus gains a partial positive (δ+) charge while the O-atom obtains a partial negative (δ–) charge.
The charged electron cloud stays non-uniformly distributed between the two bonded atoms as the N=O dipole moment stays uncancelled even in its linear shape.
Consequently, NO is overall polar (net µ = 0.15 Debye).
The Lewis dot structure of nitric oxide (NO) displays a total of 11 valence electrons.
Out of the 11 valence electrons, there are 2 bond pairs and 3 lone pairs of electrons, while 1 electron is present as an unpaired electron on the N-atom.
The N-atom and O-atom are double-covalently bonded to each other.
The O-atom carries 2 lone pairs of electrons, while the N-atom carries a lone pair and an unpaired electron.
How many resonance structures are possible for representing the NO molecule?
The following resonance structures are possible for representing the NO molecule.
However, among all the possible resonance forms, structure 1 and structure 2 are the most important. Structure 1 is a neutral molecule with zero or no formal charges on both the bonded atoms.
However, the N-atom has an incomplete octet in structure 1. In contrast, in Structure 2, both atoms have a complete octet while it carries formal charges. It is represented as NO–.
What is the molecular shape and geometry of NO?
The NO molecule possesses a linear shape and electron or molecular geometry.
How is the shape of NO different from that of NO2?
NO has a linear shape, while nitrogen dioxide (NO2) possesses a bent, angular, or V-shape.
In NO, the N-atom is bonded to one other atom only, and there is only one lone pair of electrons on the N-atom. This lone pair is kept at a maximum distance from the N=O bond pair in a planar linear shape.
Contrarily, in NO2, the central N-atom carrying 1 lone pair of electrons is bonded to two other O-atoms. Strong, lone pair-bond pair repulsions make the molecule adopt a bent shape.
Is the shape of HNO3 w.r.t central N-atom the same or different from NO?
The two molecules mentioned above possess entirely different shapes. NO is a linear molecule, while the molecular shape ofHNO3 is trigonal planar w.r.t central N-atom.
In HNO3, the central N-atom is covalently bonded to three O-atoms via double and single covalent bonds.
There is no lone pair of electrons on the N-atom in HNO3, thus no lone pair-lone pair or lone pair-bond pair electronic repulsions exist. Hence, the shape of HNO3 is trigonal planar w.r.t N-atom.
The total number of valence electrons available for drawing the NO Lewis structure is 11, i.e., an odd number.
The nitric oxide (NO) molecule possesses an identical electron and molecular geometry or shape, i.e., linear.
The central N-atom is sp hybridized in NO, while the O-atom is sp2.
The N-atom and O-atom lie on a straight line, forming a mutual bond angle of 180° in NO.
NO is a polar molecule (net µ > 0) as the more electronegative O-atom pulls the N=O shared electron cloud away from the N-atom.
Zero or no formal charges are present on the covalently bonded atoms in the NO molecule in its most contributing resonance form. However, the N-atom has an incomplete octet in this resonance form. The presence of an unpaired electron on the nitrogen atom makes NO a highly reactive, unstable free radical.
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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