Nitrosyl fluoride (NOF) Lewis dot structure, molecular geometry or shape, hybridization, bond angle, resonance structure, polar or nonpolar
NOF is the chemical formula used to represent nitrosyl fluoride, a colorless gas possessing a pungent odor. NOF reacts violently with water, alkalis and some metals. It is used as a fluorinating agent, a rocket propellant, in organic synthesis and in the semiconductor industry.
In this article, we will discuss the Lewis dot structure of NOF, its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity, resonance structure, etc.
The Lewis dot structure of nitrosyl fluoride (NOF) consists of a nitrogen (N) atom at the center. It is double-covalently bonded to an oxygen (O) atom on one side and single-covalently bonded to a fluorine (F) atom on the other side. The central N-atom carries a lone pair of electrons as well.
If you want to draw the Lewis structure of NOF, then follow the simple steps given below.
Steps for drawing the Lewis dot structure of NOF
1. Count the total valence electrons present in NOF
NOF consists of three distinct elements, i.e., nitrogen, oxygen, and fluorine.
Nitrogen (N) is present in Group VA (or 15) of the Periodic Table of elements. It consists of a total of 5 valence electrons in it.
Oxygen (O) is present in Group VI A (or 16), containing 6 valence electrons in each atom.
In contrast, fluorine (F) is a halogen located in Group VII A (or 17), possessing a total of 7 valence electrons.
Total number of valence electrons in nitrogen = 5
Total number of valence electrons in oxygen = 6
Total number of valence electrons in fluorine = 7
The NOF molecule comprises 1 N-atom, 1 O-atom and 1 F-atom.
∴ Therefore, the total valence electrons available for drawing the Lewis dot structure of NOF = 1(5) + 1(6) + 1(7) = 18 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.
Nitrogen (E.N = 3.04) is less electronegative than both oxygen (E.N = 3.44) and fluorine (E.N = 3.98), so it is chosen as the central atom in the NOF Lewis structure.
In contrast, the O-atom and F-atom occupy terminal positions, as shown below.
3. Connect the outer atoms with the central atom
In this step, both the O-atom and the F-atom 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.
In the above structure, there are 2 single bonds, i.e., 2(2) = 4 valence electrons are already consumed out of the 18 initially available. But we still have 14 valence electrons remaining.
Now let’s see where we can place these remaining valence electrons in the NOF Lewis structure.
4. Complete the octet of the outer atoms
Both the outer atoms, i.e., the F-atom and the O-atom in the NOF Lewis structure, need a total of 8 valence electrons to complete their octets.
Therefore, 3 lone pairs of electrons are placed around both atoms, 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 + electrons placed around F-atom + electrons placed around O-atom, shown as dots = 2(2) + 6 + 6 = 16 valence electrons.
Total valence electrons – electrons used till step 4 = 18 – 16 = 2 valence electrons.
The remaining 2 valence electrons are placed as a lone pair on the central N-atom.
But a problem here is that this N-atom still has an incomplete octet with 2 single bonds + 1 lone pair.
However, don’t worry; we can easily solve this issue by converting a lone pair from an outer atom into an additional bond between the central N-atom and the corresponding outer atom.
Out of the two outer atoms available, O-atom is less electronegative than F-atom, so it can easily let go of its lone pairs.
Therefore, a lone pair from the O-atom is converted into an extra bond between N and O, as shown below.
Now let’s check the stability of this 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, 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 NOF-bonded atoms.
For nitrogen atom
Valence electrons of nitrogen = 5
Bonding electrons = 1 single bond + 1 double bond = 2 + 4 = 6 electrons
The absence of any formal charges on all the bonded atoms in the NOF Lewis structure marks its incredible stability.
However, you may note that the following resonance structures are possible for representing the NOF molecule.
Each resonance structure is a way of representing the Lewis structure of a molecule.
The lone pairs and pi-bonded electrons are delocalized, i.e., they keep revolving from one position to another on the molecule. The actual NOF structure is a weighted average of the above resonance forms.
However, Structure I has the maximum contribution in the resonance hybrid as it is the most stable with no formal charges on any of the bonded atoms.
In conclusion, we have obtained the correct Lewis representation of NOF, and now we are good to proceed to the next section of the article.
What are the electron and molecular geometry of NOF?
The molecular geometry or shape of nitrosyl fluoride (NOF) w.r.t the central N-atom is bent, angular, or V-shaped. However, its ideal electron pair geometry is trigonal planar.
The presence of a lone pair of electrons on the central N-atom leads to a strong repulsive effect, thus distorting the overall molecular geometry and shape of NOF.
Molecular geometry of NOF
The molecular geometry or shape of NOF w.r.t the central N-atom is bent, angular or V-shaped.
There is a lone pair of electrons on the central N-atom in NOF, leading to lone pair-bond pair electronic repulsions. As a result, the terminal atoms tilt away from the center to minimize this strong repulsive effect.
Hence, the NOF molecule occupies a non-planar bent or V-shape, as shown below.
Electron geometry of NOF
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule containing a total of 3 electron density regions around the central atom is trigonal planar.
In NOF, the N-atom at the center is surrounded by 2 bond pairs (N=O bond and N-F bond) and a lone pair of electrons, making a total of 3 electron density regions.
Hence, the ideal electron pair geometry of the NOF molecule is trigonal planar.
An easy trick to finding a molecule’s electron and molecular geometry is 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 NOF molecule
A in the AXN formula represents the central atom. In the NOF molecule, a nitrogen (N) atom is present at the center, so A = N.
X denotes the atoms bonded to the central atom. In NOF, 1 O-atom and 1 F-atom are directly bonded to the central N-atom. So X = 2 for NOF.
N stands for the lone pairs present on the central atom. As per the Lewis structure of NOF, the central N-atom carries 1 lone pair of electrons. Thus, N = 1 for NOF.
As a result, the AXN generic formula for NOF is AX2N1.
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 AX2N1 generic formula is bent, angular, or V-shaped, while its electron geometry is trigonal planar, as we already noted down for the NOF molecule.
During chemical bonding in NOF, the 2s atomic orbital of nitrogen hybridizes with two of its three half-filled 2p atomic orbitals to produce three sp2 hybrid orbitals.
Each sp2 hybrid orbital possesses a 33.3 % s-character and 66.7 % p-character. However, all the sp2 hybrid orbitals are not equivalent in NOF.
One sp2 hybrid orbital contains paired electrons which are situated as a lone pair on the central N-atom in NOF. Contrarily, the half-filled sp2 hybrid orbitals of nitrogen form the N=O and N-F sigma (σ) bonds.
Conversely, the unhybridized p-orbital of nitrogen forms the N=O pi (π) bond by p-p orbital overlap with oxygen.
Refer to the figure drawn below.
A shortcut to finding the hybridization present in a molecule is using its steric number against the table shown below.
The steric number of the N-atom in NOF is 3, so it has sp2 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The bond angle of NOF
Due to the non-planar, bent shape of NOF, the O=N-F bond angle is reduced from an ideal value of 120° to 110°.
Pauling’s electronegativity scale states that a covalent chemical bond is polar if an electronegativity difference of 0.5 to 1.6 units exists between the bonded atoms.
The two main covalent chemical bonds present in NOF are N=O and N-F.
The N=O bond is moderately polar as a specific electronegativity difference of 0.40 units exists between a nitrogen (E.N = 3.04) and an oxygen (E.N = 3.44) atom.
In comparison, the N-F bond is strongly polar as per an electronegativity difference of 0.94 units between the nitrogen and the fluorine (E.N = 3.98) atom.
The more electronegative O-atom and F-atom attract the N=O and N-F shared electron clouds away from the central N-atom in NOF. Consequently, the central N-atom gains a partial positive charge (δ+) while both the terminal atoms obtain partial negative charges (δ–).
However, the dipole moment value of the N-F bond is higher than that of the N=O bond.
The unequal N-F and N=O dipole moments do not get canceled in the asymmetric bent shape of NOF. Thus, NOF is a strongly polar molecule overall (net µ = 1.81 Debye).
The Lewis dot structure of nitrosyl fluoride (NOF) displays a total of 18 valenceelectrons, i.e., 18/2 = 9 electron pairs.
Out of the 9 electron pairs, there are 3 bond pairs and 6 lone pairs of electrons.
An N-atom at the center is double-covalently bonded to an O-atom on one side and single-covalently bonded to an F-atom on the other side.
The central N-atom carries a lone pair of electrons while the terminal O-atom and F-atom carry 2 and 3 lone pairs, respectively.
What is the molecular shape of NOF?
The molecular geometry or shape of NOF w.r.t the central N-atom is bent, angular or V-shape.
Is the molecular geometry of NOF the same as its electron geometry?
No, the molecular geometry of NOF is bent, angular, or V-shaped, while its ideal electron pair geometry is trigonal planar.
The presence of a lone pair of electrons on the central N-atom leads to strong lone pair-bond pair electronic repulsions. The strong repulsive effect distorts the overall molecular shape, and it becomes asymmetric bent.
How is the shape of NOF different from that of ClF3?
In NOF, the central N-atom is bonded to two other atoms, and it has a lone pair of electrons. The lone pair-bond pair repulsions distort the molecular geometry of NOF, and it occupies a bent shape.
In ClF3, a chlorine (Cl) atom is present at the center. It is single-covalently bonded to 3 F-atoms at the sides, and the central Cl-atom has 2 lone pairs of electrons as well.
The electron-repulsive effect makes the molecule occupy a T-shape, unlike its ideal electron pair geometry, i.e., trigonal bipyramidal.
How are the shapes of the following fluorine-containing compounds the same or different?
In NOF, the central N-atom is directly bonded to an O-atom and an F-atom, and it has a lone pair of electrons. Lone pair-bond pair repulsions tilt the terminal atoms away from the center and towards each other. Thus, NOF occupies a bent shape.
In OF2, the central O-atom is bonded to 2 F-atoms, and it has 2 lone pairs of electrons. Lone pair-lone pair and lone pair-bond pair electronic repulsions make the molecule occupy a bent shape.
HF is a linear molecule as only two atoms are directly bonded to each other in a planar arrangement.
The shape of NO2F is trigonal planar as the central N-atom is bonded to 1 F-atom and 2 O-atoms, and there is no lone pair of electrons on it. The absence of electronic repulsions makes the molecule occupy a symmetrical trigonal planar shape and molecular geometry.
The total number of valence electrons available for drawing the nitrosyl fluoride (NOF) Lewis structure is 18.
The molecular geometry or shape of NOF w.r.t the central N-atom is bent, angular or V-shaped.
The ideal electron pair geometry of NOF is trigonal planar.
NOF has sp2 hybridization.
The O=N-F bond angle is 110° in nitrosyl fluoride.
NOF is a polar molecule (net µ > 0) as the unequal dipole moments of individually polar N=O and N-F bonds stay uncancelled in the non-planar, asymmetric bent shape.
Zero or no formal charges on the covalently bonded atoms in the NOF molecule ensure the extraordinary stability of the Lewis structure drawn in this article.
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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