MO Diagram, Molecular Geometry, Hybridization, Polarity, and F2 Lewis Structure

Fluorine is a pale yellow-colored diatomic gas with a pungent odour. Its chemical formula is F2. The molecular weight of F2 is 37.997 g/mol. It has a boiling temperature of 188 degrees Celsius and a melting point of 219 degrees Celsius.

It is poisonous in nature, causing chemical burns on the skin and being fatal if ingested. It’s extremely reactive, can corrode metals, and is unstable at high temperatures.

F2 is produced commercially by high-temperature electrolysis of molten potassium bifluoride (KHF2); fluorine gas is liberated at the anode, while hydrogen gas is liberated at the cathode.

Fluorine is primarily utilised in the nuclear fuel cycle to prepare uranium fluoride. It’s also how fluorides like SF6, ClF3, and CoF3 are made.

Lewis Structure F2

The Lewis theory of chemical bonding aids in visualising the arrangement of atoms in molecules, as well as how they are linked or bound.

Because the valence electrons in each atom are the ones that participate in bonding, they are the only ones visible in Lewis structures. It should be highlighted, however, that this hypothesis of electrical structure is quite basic and limiting.

Each valence electron is depicted as a dot in a typical Lewis structure, while a covalent connection between two atoms (made by sharing two electrons) is represented as a line.

Chemical bonding causes several atoms to seek eight electrons in their valence shell; this is known as the octet rule, and it is mirrored in the Lewis structure of a molecule.

However, hydrogen is an exception; it desires a duplet rather than an octet since it only has one electron in its K shell and hence requires only one extra to reach the K shell’s maximal capacity.

Noble gases do not need to bind or react with other atoms/molecules because their valence shells are already totally filled. This is why they are referred to as “noble.”

Take a look at the chemical bonding in F2 as represented by the Lewis structure.

Step 1: Determine how many more valence electrons each fluorine atom requires to complete an octet by multiplying the amount of valence electrons in each atom of F2.

Fluorine has an atomic number of 9 and hence has 9 electrons in its neutral atomic form. Its K shell has 2 electrons while its L shell has 7 electrons. Each fluorine atom thus contains seven valence electrons. Each atom need one extra electron to complete the octet.

Step 2: Fluorine is one of the most straightforward cases. It’s not rocket science to figure out that both fluorine atoms can share one pair of electrons while still being content with their own octets!

The lone pairs are the three unshared pairs of electrons on each fluorine atom.

Step 3: Using the information from step 2, create a skeleton for the F2 molecule. Because both fluorine atoms share one pair of electrons, they form a single covalent bond.

Do you have any idea how much energy would be required to break this bond? It has a calorie density of 157 KJ/mol.

This is about half the energy needed to break a single carbon–carbon bond. Fluorine’s low bond energy explains why it is reactive.

Anyway, back to the article’s main point! The fluorine molecule is neutral, which means it has no charge. Each fluorine atom, in reality, has a formal charge of zero. In any case, the formula for the official charge is supplied below for your convenience.

Valence electrons – 0.5*bonding electrons – non-bonding electrons = Formal charge (FC).

In F2, hybridization occurs.

A chemical bond is formed between two atoms in a molecule when their valence orbitals overlap by sharing a pair of electrons.

In terms of orbitals, consider the ground state electronic arrangement of fluorine atom(s) in F2.

F: 1s22s22px22py22pz1

The valence orbitals of a fluorine atom—2s, 2px, 2py, and 2pz—hybridize to generate four identical sp3 orbitals, each with the same energy.

Three of these hybrid orbitals are fully occupied (each with two electrons), while the fourth is half-filled (with one electron) and so has room for another.

The needed electron originates from the second fluorine atom’s half-filled sp3 hybrid orbital. This leads in the production of a single link (also known as a sigma bond) between two fluorine atoms’ half-filled sp3 hybrid orbitals. The lone pairs are the sp3 hybrid orbitals on each atom that are already fully filled.

The following formula can be used to determine the hybridization of the centre atom.

Number of sigma bonds + number of lone pairs = steric number

Steric numberType of hybridization
2sp
3sp2
4sp3
5sp3d
6sp3d2

Molecular Geometry (F2)

When it comes to establishing the geometry of a molecule, the Valence Shell Electron Pair Repulsion (VSEPR) theory is the favoured method.

According to this, the constituent atoms of a molecule organise themselves in such a way that the valence shell electrons on all atoms cause the least amount of repulsion.

The number of bond pairs on the central atom, stoichiometry, and the amount of lone pairs on the central atom all influence how this theory is applied to a molecule. On the basis of VSEPR, the table below lists numerous chemical stoichiometries and the geometries that the molecule adopts.

Instead of following the rules, let us use our common sense in the example of F2. Because there are just two atoms, the two atoms can only arrange themselves in three dimensions in one way: linearly, regardless of how many bond pairs and lone pairs there are.

F2 is a linear molecule as a result!

Diagram of the F2 Molecular Orbital (MO)

All of the constituent atoms in a molecule contribute to the development of molecular orbitals, according to molecular orbital (MO) theory. These MOs are formed by combining the atomic orbitals in a linear fashion. As a result, electrons in a molecule are assigned to molecular orbitals rather than atomic orbitals.

Let’s have a look at the F2 MO diagram. Both F atoms’ 2s orbitals combine to create a low-energy bonding orbital and a high-energy antibonding orbital (as shown below).

2pz orbitals combine to generate one bonding orbital (sigma) and one antibonding orbital (sigma) by end-on overlap. Both F atoms’ 2px and 2py orbitals combine (through sideways overlap) to generate two bonding pi orbitals and two antibonding pi orbitals.

Because of the differences in how the constituent atomic orbitals overlap, the shapes of sigma and pi orbitals vary, as do the forms of sigma* and pi*.

It’s worth noting that eight atomic orbitals combine to generate eight molecule orbitals. Starting with the least energy molecular orbital, fourteen electrons are filled in.

The greatest energy antibonding orbital is the only one that remains empty, as indicated. In the Lewis structure, we can see that there is a single sigma connection between two F atoms.

Let us now use the following formula to calculate the bond order in the F2 molecule using MO theory.

BO = 0.5 * (number of electrons in bonding orbitals – number of electrons in antibonding orbitals)

BO = 0.5 (8 – 6) = 1 for fluorine molecule.

Because of the symmetry and diatomicity of the molecule, the case of F2 is straightforward.

The level of mixing and hence the contribution of individual atomic orbitals to generate a particular molecular orbital in more complicated molecules (polyatomic and asymmetric) is determined by the relative energy alignment of the atomic orbitals.

Polarity F2

Because both constituent atoms are of the same element F, the F2 molecule is symmetric. As a result, there is no difference in electronegativity between the two constituent atoms, making F2 non-polar.

This means that the FF bond’s dipole moment is 0! As a result, the F2 molecule as a whole is non-polar.

Conclusion

The information presented in this article aids in creating a fundamental grasp of F2 structure via chemical bonding.

In fact, the knowledge in this page might be extrapolated to comprehend chemical bonding in additional halogen compounds such as Cl2, Br2, and I2.

Read more: Geometry, Hybridization, and Polarity of N2F2 Lewis Structure

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.

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