Hydrogen fluoride is a colorless liquid or a gaseous compound having the chemical formula HF. It tends to dissolve in water and the colorless aqueous solution is known as hydrofluoric acid.
It has a melting point of -118.50 F and a boiling point of about 670 F.
HF has a molar mass of 20.0064 g/mol and a density of 1.15 g/litre as a gas at 250 C.
HF has a wide range of applications. It acts as a precursor to the halogen fluorine via electrolysis procedure. Other than this, it also acts as the precursor of several metal fluorides like aluminum fluoride and uranium hexafluoride.
Anhydrous HF has catalyzing properties and is hence used in the process of petroleum alkylation to increase the octane number of petroleum. Also, hydrogen fluoride can be used to manufacture herbicides, fluorescent light bulbs, refrigerants, and so on.
Hydrogen fluoride gas is dangerous and aqueous HF acid is corrosive and toxic.
The below reaction provides us the procedure of formation of HF:
CaF2 + H2SO4 ——–> CaSO4 + 2HF (endothermic reaction)
HF Lewis Structure
Chemical bonding is the study of atomic attraction that results in the formation of new products. The science behind the formation of chemical bonds helps us understand several chemical and physical properties exhibited by different molecules and compounds.
To learn the nature of bonding inside a molecule of HF, we will first understand the concept of Lewis Structure so that we can draw the correct diagram.
Lewis Structure gives a 2-dimensional sketch representing the distribution of electrons among atoms inside a molecule.
Here, we will talk about the valence electrons i.e. the electrons in the outermost shell of the atom which take part in bonding.
Let us discuss the step-by-step procedure to draw the Lewis Structure diagram for an HF molecule:
Step 1: We have to first count the total number of valence electrons inside a single hydrogen fluoride molecule.
HF has one atom of hydrogen and one atom of fluorine.
Hydrogen belongs to period 1 and has 1 valence electron whereas Fluorine belongs to period 17(group of halogens) and therefore has 7 valence electrons. We can check the atomic number from the periodic table to confirm the valence electron number.
Total number of valence electrons in an HF molecule = 1 + 7 = 8.
Step 2: Following the general procedure, now that we have found the valence electron number, we have to work out the element which will take the central position in the molecule. Usually, the element having the least electronegativity value acts as the central atom.
Hydrogen fluoride is a diatomic molecule, so here the concept of a central atom does not take place. We can place them accordingly:
Step 3: Lewis Structure is also known as an electron-dot structure since we use dot notations to represent the valence electrons surrounding the atoms.
Let us have a look at the schematic sketch of HF after we have placed the dot electrons:
Step 4: Now, we will check the octet rule.
The octet rule deals with the octet fulfillment of valence shells of elements. It says that the elements of group 1 to group 17 tend to achieve the octet outer shell electron configuration of noble gas elements.
As we can see in HF, we have fluorine which will tend to have eight electrons in its valence shell. As per the above sketch, we have fulfilled the required criterion.
In the case of hydrogen, however, we see that it will have a tendency to have two electrons in its outermost shell as it will achieve [He] configuration.
In the diagram, we can see that hydrogen has got two electrons surrounding its atom and therefore this criterion has also been fulfilled. Let us proceed to the following step.
Step 5: Before we can confirm a certain Lewis Structure diagram, we have to check one last thing and that is the Formal Charge.
The formal charge is defined as the charge we assign to a given atom inside a molecule following the assumption that electrons are shared equally among elements for bond formation.
We calculate formal charge via the following formula:
In HF, the formal charge for H = 1 – 0.5*2 – 0 = 0.
The formal charge for F = 7 – 0.5*2 – 6 = 0.
As we can see, the elements are present in their least possible formal charge values.
We can hence confirm that we have got our correct Lewis Structure diagram for hydrogen fluoride.
We have a single bond between H and F due to sharing of an electron pair.
Having a 2D idea about a molecule is not enough to understand the nature of chemical bonding. Lewis Structure has its own limitations and therefore we will now use a different and modified approach.
We will use Valence Shell Electron Pair Repulsion (VSEPR) theory to decipher the 3- dimensional molecular geometry of the given molecule.
This will help us visualize the structural arrangement inside the molecule in a better manner and also aid in understanding several properties that it can exhibit.
VSEPR theory says that the like-charged electron particles form a negative cloud atmosphere around the nuclei which can cause repulsion. In order to minimize repulsion among the electrons and to maintain the stability of the geometry, the atoms remain farther away from each other.
Let us check the molecular geometry of hydrogen fluoride (HF).
Molecular Geometry of HF
To use the VSEPR model for predicting the 3D molecular shape, we have to first look at the Lewis Structure.
Here we have our Lewis Structure for HF
Now, since we have a heteronuclear diatomic molecule, we are having two electron clouds surrounding the two atoms, one on each, and therefore to reduce the repulsive forces, they will be placed farthest from each other.
The only possible bond angle is 180 degrees and the H and F atoms are forming a straight line with each other as per VSEPR theory.
To confirm whether our prediction went right, we can check with VSEPR notations (AXnEx). We usually do this for polyatomic molecules ( more than two atoms) where there is a central atom playing a role.
In HF, let us consider H to be A.
X= no of surrounding atoms, ∴ n = 1.
E= no of lone pairs on A, ∴ x = 0.
The notation is AX1E0.
If F is considered to A, the notation will be AX1E3.
For a notation of AX1Ex, we will have a linear structure and a bond angle of 180 degrees.
There are several models and theories that are used to explain the mechanism of bond formation. One such model to discuss the mechanism behind covalent bonding is known as orbital hybridization.
Here, we do not talk about electron orbits but rather we deal with orbitals.
Orbital is a terminology of quantum mechanics that explains the wave nature of electrons and gives us a mathematical function of the probability of electron density near atomic nuclei.
Atomic orbitals are of different types like s,p,d,f. The process of hybridization says that these atomic orbitals of an atom inside a molecule combine and fuse to form hybridized orbitals like sp, sp2,sp3,sp3d, and so on. These hybrid orbitals then take part in bond formation.
For HF, let us look at the fluorine atom and its electronic configuration.
Fluorine has an atomic number of 9 and the electronic configuration is 1s22s22p5.
Here, the s orbital and the three p orbitals fuse and form 4 sp3 hybridized orbitals and the hybridization that exists is sp3.
But there are other theories explaining the hybridization in a different manner.
It is said that the 1s orbital of hydrogen overlaps and fuses with the 2p orbital of fluorine in a molecule of HF. According to Molecular Orbital Theory, the 2s orbital of F is non-bonding, and the 2pz orbital of F combines with 1s of H.
Polarity is yet another important topic of chemistry that we are going to discuss in this article.
It is a property exhibited by molecular structures and it deals with the separation of electric charges between the constituent atoms inside the molecule.
Now, how can we term a molecule polar or non-polar?
To find this out, we have to first understand what a polar bond means.
A bond is said to be polar if there is a considerable electronegativity difference (more than 0.4-0.5) between the two atoms which results in the formation of a dipole moment.
If a molecule is not symmetrical then the dipoles formed in its bonds do not get canceled out totally with one end having denser negative charges than the other and it turns out to be polar in nature.
H has an electronegativity value of 2.20 while that of F is 3.98. The difference = 3.98 – 2.20 = 1.78.
In the H-F bond, H will bear a partial 𝛅+ and F will bear a 𝛅- charge. Even though HF is linear, there is only one bond and the net resultant dipole doesn’t cancel out. Thus, we have a polar molecule in HF.
In this article on Hydrogen Fluoride, we have included the topics of Lewis Structure, Hybridization, Molecular Geometry, and Polarity to explain the covalent bonding inside the molecule.