Hydrogen fluoride, sometimes known as HF, is a colourless liquid or gas with the chemical formula HF. Hydrofluoric acid is a colourless aqueous solution that tends to dissolve in water.
It has a melting point of -118.50 degrees Fahrenheit and a boiling point of roughly 670 degrees Fahrenheit.
At 250 degrees Celsius, HF has a molar mass of 20.0064 g/mol and a density of 1.15 g/litre as a gas.
HF can be used in a variety of ways. Through electrolysis, it serves as a precursor to the halogen fluorine. It also serves as a precursor to a number of metal fluorides, including aluminium fluoride and uranium hexafluoride.
Because anhydrous HF has catalytic capabilities, it is utilised in the petroleum alkylation process to raise the octane number. Herbicides, fluorescent light bulbs, refrigerants, and other products can all be made with hydrogen fluoride.
Aqueous hydrogen fluoride acid is caustic and toxic, and hydrogen fluoride gas is harmful.
The following reaction shows how HF is formed:
—––––––––––––––––––––––––––––––––––––––––––– (endothermic reaction)
Lewis Structure in HF
The study of atomic attraction that leads to the development of new products is known as chemical bonding. The chemistry underpinning chemical bond formation aids our understanding of a variety of chemical and physical properties displayed by various molecules and compounds.
To comprehend the nature of bonding within an HF molecule, we must first grasp the notion of Lewis Structure in order to build the correct figure.
Lewis Structure is a two-dimensional diagram that depicts the electron distribution among atoms within a molecule.
The valence electrons, or electrons in the atom’s outermost shell that participate in bonding, will be discussed here.
Let’s go over how to build a Lewis Structure diagram for an HF molecule step by step:
The entire amount of valence electrons inside a single hydrogen fluoride molecule must first be counted.
One atom of hydrogen and one atom of fluorine make up HF.
Fluorine belongs to period 17 (group of halogens) and hence has 7 valence electrons, whereas hydrogen belongs to period 1 and has one valence electron. To confirm the valence electron number, we can look up the atomic number in the periodic table.
In an HF molecule, the total number of valence electrons is 1 + 7 = 8.
Step 2: Using the standard technique, we must determine which element will occupy the central location in the molecule now that we have determined the valence electron number. The centre atom is usually the element with the lowest electronegativity value.
Because hydrogen fluoride is a diatomic compound, the concept of a core atom is absent. We can put them in the following order:
Step 3: Because we use dot notations to represent the valence electrons around the atoms, the Lewis Structure is also known as an electron-dot structure.
After we’ve placed the dot electrons, have a look at the HF schematic sketch:
Step 4: Next, we’ll look at the octet rule.
The octet rule is concerned with the octet completion of element valence shells. It claims that noble gas elements in groups 1 to 17 tend to acquire the octet outer shell electron configuration.
We have fluorine, which has eight electrons in its valence shell, as we can see in HF. We have met the required criterion, as shown in the diagram above.
However, because hydrogen will acquire the [He] configuration, it will have a tendency to have two electrons in its outermost shell.
As can be seen in the diagram, hydrogen has two electrons encircling its atom, indicating that this requirement has also been met. Let us go on to the next phase.
Step 5: Before we can certify a Lewis Structure diagram, one last thing must be checked, and that is the Formal Charge.
The formal charge is the charge that we attribute to a certain atom within a molecule based on the premise that electrons are shared evenly among elements during bond formation.
The following formula is used to determine formal charge:
H has a formal charge of 1 – 0.5*2 – 0 = 0 in HF.
F = 7 – 0.5*2 – 6 = 0 has a formal charge of 0.
The elements are present in their least formal charge values, as can be seen.
As a result, we can validate that our Lewis Structure diagram for hydrogen fluoride is correct.
Because H and F share an electron pair, they form a single bond.
Geometry of Molecular Structure
Understanding the nature of chemical bonding requires more than a 2D understanding of a molecule. Because the Lewis Structure has its own set of constraints, we’ll now take a new, modified approach.
The 3-dimensional molecular geometry of the supplied molecule will be deciphered using Valence Shell Electron Pair Repulsion (VSEPR) theory.
This will assist us in better visualising the structural arrangement within the molecule as well as comprehending numerous of its features.
According to the VSEPR hypothesis, like-charged electron particles form a negative cloud environment around nuclei, causing repulsion. The atoms are kept farther apart to avoid electron repulsion and maintain the geometry’s stability.
Let’s have a look at the chemical structure of hydrogen fluoride (HF).
HF Molecular Geometry
We must first examine the Lewis Structure before using the VSEPR model to estimate the 3D molecule structure.
Our Lewis Structure for HF is shown below.
Now, because we have a heteronuclear diatomic molecule, we have two electron clouds around the two atoms, one on each, and they will be put farthest apart to lessen repulsive interactions.
According to VSEPR theory, the only feasible bond angle is 180 degrees, and the H and F atoms form a straight line with each other.
We can use VSEPR notations to confirm whether our forecast was correct (AXnEx). This is typically done for polyatomic compounds (those with more than two atoms) in which a central atom plays a role.
Let’s pretend that H is A in HF.
X = number of atoms in the immediate vicinity, n = 1.
x = 0; E= number of lone pairings on A.
AX1E0 is the notation.
The notation will be AX1E3 if F is believed to be A.
We’ll use an AX1Ex notation with a linear structure and a bond angle of 180 degrees.
Hybridization with HF
The mechanism of bond creation is explained using a variety of models and ideas. Orbital hybridization is one model for discussing the mechanism behind covalent bonding.
We’re not talking about electron orbits here; instead, we’re talking about orbitals.
Orbital is a quantum mechanics concept that describes the wave behaviour of electrons and provides a mathematical formula for calculating the probability of electron density near atomic nuclei.
Different types of atomic orbitals exist, such as s, p, d, and f. These atomic orbitals of an atom inside a molecule join and fuse to generate hybridised orbitals like sp, sp2,sp3,sp3d, and so on, according to the hybridization process. These hybrid orbitals then participate in the creation of bonds.
Let’s have a look at the fluorine atom and its electrical state in relation to HF.
The electronic configuration of fluorine is 1s22s22p5. Fluorine has an atomic number of 9 and an electronic configuration of 1s22s22p5.
The s orbital and the three p orbitals combine here to generate four sp3 hybridised orbitals, with sp3 hybridization.
However, there are alternative theories that explain hybridization in a different way.
In a molecule of HF, the 1s orbital of hydrogen is said to overlap and fuse with the 2p orbital of fluorine. The 2s orbital of F is non-bonding, and the 2pz orbital of F joins with the 1s of H, according to Molecular Orbital Theory.
Polarity of HF
Another key topic of chemistry that we will examine in this post is polarity.
It’s a feature of molecular structures that deals with the separation of electric charges between the molecule’s constituent atoms.
What does it mean to call a molecule polar or non-polar now?
To figure this out, we must first comprehend the concept of a polar link.
A bond is said to be polar if the electronegativity difference between the two atoms is significant (greater than 0.4-0.5), resulting in the creation of a dipole moment.
If a molecule is not symmetrical, the dipoles created in its bonds do not cancel out completely, and one end has denser negative charges than the other, resulting in polarity.
H has an electronegativity of 2.20, whereas F has a value of 3.98. 3.98 – 2.20 = 1.78 is the difference.
H will have a partial + charge and F will have a – charge in the H-F bond. Despite the fact that HF is linear, it only has one bond, and the net resulting dipole does not cancel out. In HF, we now have a polar molecule.
Conclusion
To describe the covalent bonding inside the molecule, we have incorporated the themes of Lewis Structure, Hybridization, Molecular Geometry, and Polarity in this article on Hydrogen Fluoride.
Good luck with your studies!
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