Nitrosonium (NO+) ion Lewis structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity
NO+ represents the nitrosonium ion, also known as nitrosyl cation. NO+ is a relatively stable, highly reactive ion. It is an inorganic ion commonly used as a ligand in coordination chemistry.
It is also valuable in spectroscopic studies giving information about chemical bonding and the electronic structure of molecules.
This article is very important for you if you want to learn how to draw the Lewis dot structure of the nitrosonium (NO+) ion, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity, etc.
So, without any further delay, let’s start reading!
The nitrosonium (NO+) ion consists of only a nitrogen (N) atom and an oxygen (O) atom triple-covalently bonded to each other. There is 1 lone pair of electrons on both the N-atom and the O-atom in the NO+ Lewis dot structure.
Follow us through the simple steps given below and draw the Lewis dot structure of NO+ with us.
Steps for drawing the Lewis dot structure of NO+
1. Count the total valence electrons present in NO+
NO+ consists of two distinct elements, i.e., nitrogen and oxygen.
Nitrogen (N) is present in Group V A (or 15) of the Periodic Table of Elements. Thus, it has a total of 5 valence electrons in each atom.
In contrast, oxygen (O) is located in Group VI A (or 16), containing 6 valence electrons in each atom.
Total number of valence electrons in nitrogen = 5
Total number of valence electrons in oxygen = 6
The NO+ ion comprises 1 N-atom and 1 O-atom.
An important point to remember is that the NO+ ion carries a positive (+1) charge, which means 1 valence electron is removed from this Lewis structure.
∴ Therefore, the total valence electrons available for drawing the Lewis dot structure of NO+ = 1(5) + 1(6) = 11 – 1 = 10 valence electrons.
2. Find the least electronegative atom and place it at the center
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 or molecular ion.
The least electronegative atom can easily form covalent bonds with other atoms by sharing its electrons.
Among the two types of atoms present in NO+, nitrogen (E.N = 3.04) is less electronegative than oxygen (E.N = 3.44).
Hence the N-atom is chosen as the central atom in the NO+ Lewis structure while the O-atom is placed next to it, as shown below.
3. Connect the outer atom with the central atom
In this step, the terminal O-atom is joined to the central N-atom using a single straight line.
A straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons.
Two valence electrons used out of the ten initially available means we are now left with 8 valence electrons.
So let’s see in the next steps where we can place these remaining valence electrons.
4. Complete the octet of the outer atom
An O-atom needs a total of 8 valence electrons to gain a full octet configuration.
Therefore, to complete the octet of the O-atom in the NO+ Lewis structure, 3 lone pairs are placed around it, as shown below.
5. Complete the octet of the central atom
Total valence electrons used till step 4 = 1 single bond + electrons placed around the O-atom, shown as dots = 2 + 6 = 8 valence electrons.
Total valence electrons – electrons used till step 4 = 10 – 8 = 2 valence electrons.
Thus these 2 valence electrons are placed as a lone pair on the central N-atom in the NO+ Lewis structure.
However, a problem here is that this N-atom still has only 4 valence electrons surrounding it. Hence 2 lone pairs from the terminal O-atom are converted into additional covalent bonds between the adjacent N and O-atoms, as shown below.
In this way, the central N-atom also has a complete octet with 1 triple bond +1 lone pair surrounding it in NO+ Lewis structure.
As a final step, let’s check the stability of the above Lewis structure by applying the formal charge concept.
6. Check the stability of Lewis’s 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 present on the NO+– bonded atoms.
As per the above calculation, zero or no formal charge is present on the N-atom, while the O-atom carries a +1 formal charge in the NO+ Lewis structure which is also the charge present on the nitrosonium ion overall.
This implies that we have drawn the NO+ Lewis structure correctly.
However, you must note that there is another possibility for drawing the NO+ Lewis structure, as shown below.
Structures I and II are known as two different resonance forms of the nitrosonium ion. These resonance forms are achieved by the delocalization, i.e., movement of pi-bonded electrons and lone pairs from one position to another on the molecular ion, which in turn changes the position of +1 formal charge from oxygen (Structure I) to nitrogen (Structure II).
The actual NO+ structure is an average of the above two resonance forms, known as the resonance hybrid.
Now that we know everything about NO+ Lewis structure, let’s move ahead and discuss its electron and molecular geometry or shape.
What are the electron and molecular geometry of NO+?
The nitrosonium (NO+) ion possesses an identical electron and molecular geometry or shape, i.e., linear. The lone pairs present on the N-atom and the O-atom are equally displaced. Therefore, there is no distortion in the shape and geometry of NO+.
Molecular geometry of NO+
The molecular geometry or shape of the nitrosonium (NO+) ion is linear.
The only two atoms present in NO+, i.e., nitrogen and oxygen, are triple-covalently bonded to each other. Both atoms carry 1 lone pair of electrons each.
These two lone pairs are equally displaced; therefore, the lone pair-bond pair repulsions are minimized, while the lone pair-lone pair repulsions are negligible in NO+. Hence the NO+ ion occupies an ideal linear shape.
Electron geometry of NO+
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule or molecular ion containing a total of 2 electron density regions around the central atom is linear.
In NO+, the central N-atom is directly bonded to an O-atom on one side, and it has a lone pair of electrons on the other side. This makes a total of 2 electron-density regions surrounding this N-atom. As a result, the ideal electron pair geometry of NO+ is linear.
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.
AXN notation for the nitrosonium (NO+) ion
A in the AXN formula represents the central atom. In NO+, a nitrogen (N) atom is present at the center, so A = N.
X denotes the atoms bonded to the central atom. In NO+, an O-atom is directly bonded to the central N-atom, so X= 1.
N stands for the lone pairs present on the central atom. As per the Lewis structure of NO+, the central N-atom has 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 or molecular ion with an AXN generic formula is identical to its electronic geometry, i.e., linear, as we already noted down for the nitrosonium (NO+) ion.
Hybridization of NO+
Both the atoms, i.e., the N-atom and the O-atom, are sp hybridized in NO+.
During chemical bonding, the 2s atomic orbital of nitrogen hybridizes with one of its three half-filled 2p orbitals to produce two sp hybrid orbitals.
Each sp hybrid orbital possesses a 50% s-character and a 50% p-character.
One of these sp hybrid orbitals contains paired electrons which are situated as a lone pair on the N-atom.
Contrarily, the sp hybrid orbital of nitrogen-containing a single electron only forms an sp-sp sigma bond with the adjacent oxygen atom.
On the other hand, the unhybridized p-orbitals of nitrogen and oxygen form the N≡O pi bonds.
Refer to the figure drawn below.
Another shortcut to finding the hybridization present in a molecule or molecular ion is using its steric number against the table 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 NO+ bond angle
The concept of bond angle is irrelevant in NO+ as it comprises only two atoms; however, a bond angle is formed by the intersection of two lines i.e., three atoms A-B-C involved.
However, the N≡O bond length equals 106 pm in the nitrosonium ion.
As per Pauling’s electronegativity scale, a polar covalent bond is formed between two dissimilar atoms with an electronegativity difference between 0.4 and 1.6 units.
In NO+, an electronegativity difference of exactly 0.4 units is present between a nitrogen (E.N = 3.04) and an oxygen (E.N = 3.44) atom.
The more electronegative oxygen atom thus gains a partial negative (δ–) charge by strongly attracting the bonded electrons towards itself, while the adjacent nitrogen atom obtains a partial positive (δ+) charge.
The strong N≡O dipole moment stays uncancelled even in the planar linear shape of the nitrosonium ion.
Therefore, the nitrosonium (NO+) ion is overall polar (net µ > 0) with a non-uniformly charged electron cloud spread over it.
The Lewis dot structure of the nitrosonium (NO+) ion displays a total of 10 valence electrons, i.e., 10/2 = 5 electron pairs.
Out of the 5 electron pairs, there are 3 bond pairs and 2 lone pairs of electrons.
An N-atom is triple covalently bonded to the O-atom in the NO+ Lewis structure.
1 lone pair is present on each of the bonded atoms.
What is the molecular shape of NO+?
The molecular geometry or shape of the nitrosonium (NO+) ion is linear.
Why is the shape of NO+ the same as its electron geometry, although it has lone pairs?
The lone pairs present on nitrogen and oxygen in NO+ are uniformly displaced at equal distances from the center. Lone pair-bond pair and lone pair-lone pair electronic repulsions are minimized to a great extent.
Therefore, there is no distortion present in the shape and geometry of NO+ thus, it is linear.
How is the shape of NO2– different from that of NO+?
In nitrite (NO2–), two oxygen atoms are covalently bonded to the central nitrogen atom. The central N-atom also has a lone pair of electrons.
Hence the shape of NO2– is bent, angular, or V-shaped w.r.t the central N-atom. In contrast, only 1 O-atom is triple covalently bonded to an N-atom in NO+. Therefore, its shape is linear.
How is the shape of CO32- the same or different from that of NO+?
The shape of the carbonate (CO32-) ion is different from that of NO+. In CO32-, three O-atoms are directly bonded to the central C-atom that has no lone pair of electrons.
The total number of valence electrons available for drawing the nitrosonium (NO+) ion Lewis structure is 10.
NO+ possesses an identical electron and molecular geometry or shape, i.e., linear.
NO+ has sp hybridization.
The N≡ O bond length is equal to 106 pm in the nitrosonium ion.
NO+ is a polar cation as the charged electron cloud stays non-uniformly distributed between the triple covalently bonded nitrogen and oxygen atoms.
Zero or no formal charge is present on the N-atom, while the O-atom carries a +1 formal charge which is also the charge present on the NO+ ion overall.
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