Nitronium (NO2+) ion Lewis dot structure, molecular geometry
or shape, electron geometry, bond angles, hybridization,
formal charges, polar vs nonpolar
NO2+ is the chemical formula for nitronium ion, a highly reactive cation that functions as an electrophile in various organic chemistry reactions.
It is involved as a key component in the nitration process, essential for the synthesis of explosives, dyes and pharmaceuticals.
In this article, we will teach you how to draw the Lewis dot structure of the nitronium (NO2+) ion, its molecular geometry or shape, electron geometry, bond angles, hybridization, formal charges, resonance structures, polarity, etc.
But for all that, you just have to do one thing, i.e., continue reading this article till the end.
The Lewis structure of the nitronium (NO2+) ion comprises a nitrogen (N) atom at the center. It is double covalently bonded to two oxygen (O) atoms, one on either side of the cation. There is no lone pair of electrons on the central N-atom, while each terminal O-atom carries 2 lone pairs, respectively.
Follow us through the simple steps given below and draw the Lewis dot structure of NO2+ within no time.
Steps for drawing the Lewis dot structure of NO2+
1. Count the total valence electrons present in NO2+
NO2+ 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 NO2+ ion comprises 1 N-atom and 2 O-atoms.
An important point to remember is that the NO2+ ion also 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 NO2+ = 1(5) + 2(6) = 17 – 1 =16 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 NO2+, 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 NO2+ Lewis structure, while the two O-atoms are placed at the peripheries, as shown below.
3. Connect the outer atoms with the central atom
In this step, the terminal O-atoms are joined to the central N-atom using single straight lines.
A straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons.
2(2) = 4 valence electrons used out of the 16 initially available means we are now left with 12 valence electrons.
So let’s see in the next steps where we can place these remaining valence electrons in the NO2+ Lewis structure.
4. Complete the octet of the outer atoms
An O-atom needs a total of 8 valence electrons to gain a full octet configuration.
An N-O bond represents 2 valence electrons already present around an O-atom.
Therefore, to complete the octets of the two O-atoms in the NO2+ Lewis structure, 3 lone pairs are placed around each, as shown below.
5. Complete the octet of the central atom and convert lone pairs into covalent bonds if necessary
Total valence electrons used till step 4 = 2 single bonds + 2(electrons placed around each O-atom, shown as dots) = 2(2) + 2(6) = 16 valence electrons.
Total valence electrons – electrons used till step 4 = 16 – 16 = 0 valence electrons.
As all 16 valence electrons initially available for drawing the NO2+ Lewis structure are already consumed so there is no lone pair on the central N-atom.
However, a problem here is that the central N-atom still has an incomplete octet with only 2 single bonds, i.e., 2(2) = 4 valence electrons surrounding it.
But don’t worry because we can easily solve this problem by converting lone pairs into extra covalent bonds.
A lone pair from each terminal O-atom is converted into an extra covalent bond between the central N-atom and the respective O-atom, as shown below.
This leads to the formation of two double bonds around the N-atom, completing its electron deficiency.
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 NO2+– bonded atoms.
As per the above calculation, zero or no formal charge is present on the O-atoms, while the central N-atom carries a +1 formal charge which is also the charge present on the nitronium (NO2+) ion overall.
However, you may note that the following resonance structures are possible for representing the nitronium ion. The pi-bonded electrons and lone pairs keep revolving from one position to another on the molecular ion.
Among all the NO2+ resonance forms, Structure I is the most stable as the formal charges present on the individual atoms are as minimized in it as possible.
The actual NO2+ structure is an average of the above structures, known as theresonance hybrid.
In conclusion, we have drawn the correct Lewis structure of the nitronium (NO2+) ion. So let’s move ahead and discuss its electron and molecular geometry or shape.
What are the electron and molecular geometry of NO2+?
The nitronium (NO2+) ion possesses an identical electron and molecular geometry or shape, i.e., linear. There is no lone pair of electrons on the central N-atom, so no distortion is present in the shape and geometry of the molecular ion.
Molecular geometry of NO2+
The molecular geometry or shape of the nitronium (NO2+) ion is linear.
The central N-atom is directly bonded to two O-atoms via double covalent bonds, one on either side. There is no lone pair of electrons on the central N-atom, so no lone pair-lone pair or lone pair-bond pair electronic repulsions are present in the molecular ion.
Hence the NO2+ ion occupies an ideal linear shape in which the bonded atoms lie on a single straight line, as shown below.
Electron geometry of NO2+
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 NO2+, the central N-atom is directly bonded to 2 O-atoms, and it has no lone pair of electrons. This makes a total of 2 electron-density regions surrounding the central N-atom. As a result, the ideal electron pair geometry of NO2+ 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 or molecular ion using the VSEPR concept.
AXN notation for the nitronium (NO2+) ion
A in the AXN formula represents the central atom. In NO2+, a nitrogen (N) atom is present at the center, so A = N.
X denotes the atoms bonded to the central atom. In NO2+, two O-atoms are directly bonded to the central N-atom, so X= 2.
N stands for the lone pairs present on the central atom. As per the Lewis structure of NO2+, the central N-atom has no lone pair of electrons. Thus, N= 0 for NO2+.
As a result, the AXN generic formula for NO2+ is AX2N0or simply AX2.
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 AX2 generic formula is identical to its electron geometry, i.e., linear, as we already noted down for the nitronium (NO2+) ion.
Hybridization of NO2+
The central N-atom is sp hybridized in NO2+.
The electronic configuration of nitrogen is 1s2 2s2 2p3.
Upon excitation, the N-atom loses one of its valence electrons. The electronic configuration of N+ thus becomes 1s2 2s2 2p2.
During chemical bonding in NO2+, one of the two 2s electrons of nitrogen shifts to its empty 2p atomic orbital. Consequently, the half-filled 2s and one of the three 2p orbitals hybridize to produce two sp hybrid orbitals.
Each sp hybrid orbital possesses a 50% s-character and a 50% p-character, and both contain a single unpaired electron only.
Nitrogen uses these sp hybrid orbitals to form the N-O sigma bonds by sp-sp2 orbital overlap on either side of the molecular ion.
In contrast, the unhybridized p-orbitals of nitrogen form the N=O pi bonds by p-p orbital overlap with adjacent O-atoms, as shown 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 NO2+ is 2, so it has sp hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The bond angle of NO2+
The O=N=O bond angle is exactly equal to 180° due to the linear shape of the nitronium (NO2+) 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 NO2+, 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 in each N=O bond.
The more electronegative oxygen atoms thus gain partial negative (δ–) charges by strongly attracting the bonded electrons from each N=O bond, while the central N-atom obtains a partial positive (δ+) charge in NO2+.
However, it is due to the symmetrical linear shape of the nitronium cation that the oppositely-directed N=O dipole moments get canceled equally.
This leads to a uniform electron cloud distribution in NO2+ and an overall non-polar molecular ion (net µ =0).
The Lewis dot structure of the nitronium (NO2+) ion displays a total of 16 valence electrons, i.e., 16/2 = 8 electron pairs.
Out of the 8 electron pairs, there are 4 bond pairs and 4 lone pairs of electrons.
An N-atom at the center is surrounded by two O-atoms, one on each side, via double covalent bonds.
There is no lone pair of electrons on the central N-atom, while each O-atom has 2 lone pairs, respectively.
What is the molecular shape of NO2+?
Nitronium (NO2+) is an AX2-type molecular ion. Its molecular shape is thus identical to its ideal electronic geometry, i.e., linear.
How is the shape of NO2+ different from that of NO2–?
The nitrite (NO2–) ion possesses a bent, angular or V-shape w.r.t the central N-atom. A lone pair of electrons is present on the central N-atom, which leads to strong lone pair-bond pair electronic repulsions, in turn distorting the overall molecular shape.
It is thus different from the shape of the nitronium (NO2+) ion, i.e., linear.
Among NO2+ and NO2–, which bond angle is larger?
The O=N=O bond angle is greater in NO2+ i.e., 180° due to its linear shape.
Contrarily, it is due to the distortion present in NO2– that the O=N-O bond angle reduces to 134° as per the bent shape of the molecular ion.
Which of the following nitrogen sources has a trigonal planar shape?
A) NO2–
B) NO3–
C) NO2+
D) NO+
Option B is the correct answer. The nitrate (NO3–) ion possesses a symmetrical trigonal planar shape. To an N-atom at the center, three O-atoms are covalently bonded. There is no lone pair of electrons on the central N-atom, so no distortion is present in the molecular ion.
The total number of valence electrons available for drawing the nitronium (NO2+) ion Lewis structure is 16.
NO2+ possesses an identical electron and molecular geometry or shape, i.e., linear.
The central N-atom is sp hybridized in NO2+.
The O=N=O bond angle equals 180° in the nitronium ion.
NO2+ is overall non-polar (net µ = 0) as the equal and opposite N=O dipole moments get canceled uniformly overall.
Zero or no formal charges are present on the O-atoms, while the central N-atom carries a formal charge of +1 which is also the charge present on the nitronium (NO2+) ion overall.
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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