Nitrogen trifluoride or NF3 is a nitrogen halide compound that is slightly water-soluble. Its noticeable characteristics include being colorless and carrying a musty or moldy odor.
NF3 has a molar mass of around 71.002 g/mol and a density of 3.003 kg/m3.
One of the main reasons why this chemical compound is an essential topic is because it is a greenhouse gas. It is said to have quite high global warming potential but a comparatively low value of radiative forcing.
Other than being a greenhouse gas and contributing to the climatic change of the planet, NF3 has several applications. It is used to produce chemical fluoride lasers and remove silicon-based compounds during semiconductor manufacturing. Also, nitrogen trifluoride is a component of high-energy fuels.
NF3 is a toxic non-flammable gas that can prove to be harmful if inhaled. It can lead to a condition called methemoglobinemia.
NF3 Lewis Structure
Nitrogen trifluoride is a chemical molecule having one atom of nitrogen and three atoms of fluorine.
In Lewis Structure, we try to find a suitable pictorial representation of a molecule to have an idea of the chemical bonding occurring inside the molecular structure.
Before we proceed, we need to learn about a very important concept: Valence Electrons.
The electrons that are present in the outermost shell of an atomic nucleus are known as valence electrons. The combining or the binding power or capacity is termed as Valency.
As we can see, the above diagram gives us the Periodic table. Nitrogen belongs to group 15 and fluorine, being a halogen, belongs to group 17.
nitrogen has 5 and fluorine has seven valence electrons.
The total number of valence electrons in a NF3 molecule = 5 + 7*3 = 26.
Nitrogen is more electropositive than fluorine and we will keep it in the center. This is because electronegative atoms are usually reluctant to share electrons. Electropositive atoms, on the other hand, are more prone to share and aid in bond formation.
In Lewis Structure, we denote valence electrons by dot notations. Let us sketch a probable Lewis Structure for Nitrogen Trifluoride.
We have put the valence dot structures and made a sketch of the NF3 molecule.
Let us now understand the next important concept of Lewis Structure:
The main group (group 1- 17) elements tend to attain the electronic configuration of the outermost shell of noble gas elements.
Noble gases like Neon, Argon, Xenon have eight electrons in their valence shells. Main group elements present in the same period have this tendency to achieve a valency of eight while forming bonds inside a molecule.
This is known as the octet rule.
Let us check if we have fulfilled the octet rule in our Lewis structure.
The above graphical representation of NF3 shows that all the three surrounding Fluorine atoms and the central Nitrogen atom have achieved octet fulfillment.
Our last step towards getting the required structural sketch is the concept of Formal Charge.
The formal charge can be defined as the electric charge assigned to any given molecular atom assuming that electrons are always shared equally amongst the constituent elements.
We calculate formal charge via the below-mentioned formula:
Formal charge for Nitrogen atom = 5 – 0.5*6 – 2 = 0.
Formal charge for each Fluorine atom = 7 – 0.5*2 – 6 = 0.
As we can see, all the atoms inside the NF3 molecule have the least possible formal charge values.
Therefore, we have attained our most perfect Lewis Structure diagram.
If we indicate single bonds by straight lines, a 2D representation of the Nitrogen Trifluoride molecule will look like this:
NF3 Molecular Geometry
Molecular geometry or molecular shape is an important concept that we need to decipher while we are learning the chemical bonding of any chemical composition.
While Lewis Structure gives us an idea about the internal bond types and valence electron sharing inside a given molecule, it can only explain a two-dimensional geometry. To understand the 3D molecular shape of the molecule, we have an interesting theory: VSEPR theory.
VSEPR or Valence Shell Electron Pair Repulsion model is our go-to theory to predict the molecular geometry of NF3.
Electrons present around the nuclei of constituent atoms of different molecules result in a negatively charged atmosphere. Since these subatomic particles happen to carry negative charges, they also develop a repulsive tendency that needs to be minimized to account for molecular stability.
VSEPR theory talks about the minimization of repulsive forces developed by valence shell electrons and helps us predict the perfect molecular geometry.
For the NF3 molecule, we will try to find out the 3D shape.
While we are doing this, remember we will only take into account the valence electrons around the central nitrogen in the polyatomic halide molecule.
AXnEx is the notation of the VSEPR model.
A: the central atom, here Nitrogen will be A.
X stands for the surrounding atoms for nitrogen, ∴ n = 3.
E stands for lone pairs of the central atom, ∴ n = 1.
The resultant notation will be AX3N1.
Let us look at the VSEPR notation chart.
According to the VSEPR chart, the molecular geometry of nitrogen trifluoride is trigonal bipyramidal. We can also predict electron geometry via electron groups through VSEPR theory.
In NF3, the central nitrogen atom has four electron groups surrounding it: three single bonds( three bonded pairs) and one lone pair.
Therefore, the electron geometry is tetrahedral and the bond angle is around 102.50 degrees.
Note: The bond angle in NF3 is less than that in NH3 owing to the higher electronegativity value of fluorine.
Hybridization, a concept of chemical bonding, is related to orbitals.
While orbits are fixed paths of electrons around the atomic nuclei, orbitals are given by the mathematical probability of the presence of electrons in any region of space.
When atomic orbitals come together to fuse and form hybrid orbitals, we call it the orbital hybridization in chemistry.
We have several types of atomic orbitals like s, p, d, f which have different shapes.
We also have the existence of several hybrid orbitals in a molecule like sp, sp2, sp3, sp3d, etc.
Note, according to Valence Bond Theory, only atomic orbitals of the same atom inside the molecule can combine to form hybridized orbitals in hybridization.
Let us look at the electronic configuration of the central Nitrogen atom of the NF3 molecule:
N: 1s2 2s2 2p3
The one 2s orbital and 2px, 2py, and 2pz orbitals of p orbital fuse to form four hybrid orbitals.
Also, by the concept of steric number,
Steric number = Number of atoms bonded to central atom inside a molecule + Number of lone pair of electrons attached to the central atom
The steric number = 3 + 1 = 4.
The hybridization H value is 4 and the type, therefore, is sp3.
Polarity is the concept of understanding charge separation inside a molecule. Electronegativity talks about the degree to which an element can attain or gain electrons. To comprehend polarity, we require to decipher the electronegativity concept as well.
The Pauling Electronegativity chart is the first and foremost step towards understanding the polar nature of a molecule
As we can see here clearly, Nitrogen and Fluorine have a high difference in electronegativity.
While Nitrogen has an electronegativity value of 3.04, that of fluorine is 3.98 making the difference more than 0.4-0.5.
This indicates that there will be a separation of electric charges. Nitrogen will possess a partial positive charge (δ+) and fluorine will possess a partial negative charge (δ-).
Each N-F bond will therefore have negative and positive poles. This leads to the rise of three polar bonds inside the NF3 molecule.
Now, the polarity of the whole molecule is dependent not only on the nature of bonds but also on the symmetry of the molecule. In NF3 we do not have linear geometry.
Due to the asymmetry resulting from the presence of a lone pair on the central nitrogen atom, we get polarity.
NF3 is a polar molecule.
You must also go through the detailed article written on the polarity of NF3.
Note: NF3 is quite less polar than NH3 (since the net dipole moment is less).
NF3 Molecular Orbital (MO) Diagram
Molecular orbital theory or MOT is a complex and interesting theoretical approach towards understanding the chemical bonding occurring inside a given molecular structure.
In MOT, we do not only discuss the fusion of AOs (atomic orbitals) of the same atom of the same energies but we consider the orbitals of the molecule as a whole.
Here, the valence electrons of all constituent atoms can be shared inside the molecule. They combine to form MOs or molecular orbitals and the total concept is explained graphically via an MO diagram.
In molecular orbital theory, we consider the orbitals containing the inner shell electrons to be non-bonding orbitals.
Other than this, we have bonding and anti-bonding orbitals. Anti-bonding orbitals are higher in energy, less stable, and contain a node perpendicular to the internuclear axis. Bonding orbitals are more stable in nature.
N: 1s2 2s2 2p3
F: 1s2 2s2 2p5
Here, the 1s2 orbitals of the four atoms are the nonbonding orbitals.
Here, we have attached the MO diagram of another halide molecule BF3 for reference.
In this article, we have discussed the chemical bonding nature inside a molecule of Nitrogen Trifluoride. We have explained the Lewis Structure, VSEPR theory, Hybridization, Polarity, and MOT diagram in detail.