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

Benzene is a white liquid chemical molecule with the molecular formula C6H6. It is classified as an aromatic hydrocarbon since its structure is made up entirely of carbon and hydrogen.

Aromatic hydrocarbons are unsaturated hydrocarbons with single and double bonds in their ring configurations.

Benzene is a highly flammable, volatile chemical that smells like gasoline. As a byproduct of oil refining, benzene can be detected in crude oil.

It is frequently utilised in the chemical industry as a solvent or an intermediary in the synthesis of numerous compounds. It’s also used to make colours, plastics, detergents, insecticides, and pharmaceuticals, among other things.

Benzene has a molar mass of 78.11 g/mol. 5.53 °C and 80.1 °C are the melting and boiling points, respectively. Because benzene has a lower density than water due to its lower flashpoint of less than 0°F. As a result, it’s just mildly soluble in water.

Let’s look at the features of benzene, such as its structure, geometry, hybridization, and MO diagram.

Lewis Structure of Benzene

A molecule’s valence shell electrons are demonstrated in the Lewis structure. It depicts the arrangement of individual electrons in a molecule. It’s also known as an electron dot structure since each bond between two atoms is represented by two dots.

Let’s have a look at the Lewis structure of benzene one step at a time.

The total number of valence electrons of each atom present in benzene (C6H6), i.e. carbon and a hydrogen atom, is determined in the first step.

In carbon, there are four valence electrons.

In hydrogen, there are 1 valence electrons.

In carbon, the total number of valence electrons is 6 X 4 = 24.

In hydrogen, the total number of valence electrons is 6 X 1 = 6.

Step 2– Calculate the total number of valence electrons in the benzene molecule, which is as follows:

The total number of valence electrons in C6H6 is equal to the number of valence electrons in six carbon atoms plus the number of valence electrons in six hydrogen atoms.

30 electrons = 24 + 6

Step 3– Determine how many electrons carbon and hydrogen atoms require to complete their octet.

To complete an octet, a carbon atom will need 4 more electrons, whereas a hydrogen atom will need 1 more electron in its outermost shell.

As a result, each carbon atom will establish a single hydrogen atom bond.

All hydrogen atoms are now paired, but each carbon atom will require three additional electrons in its outermost shell.

Step 4– Calculate the number of electrons needed for benzene to achieve a stable structure.

Each carbon atom needs three additional electrons to complete their octet, so benzene will need 18 electrons to establish a stable configuration.

The remaining valence electrons are arranged so that the carbon octet is complete. The electrons involved in bond formation are depicted by a dot in the structure below.

As a result, 2 dot denotes 2 electrons, indicating a single bond, and 4 dot denotes 4 electrons, indicating a double bond.

Step 5– Benzene Lewis structure

As a result, benzene is made up of six carbon atoms arranged in a planar ring structure with alternate single and double bonds, and each carbon atom has a single bond with a hydrogen atom.

Molecular Geometry of Benzene

The Valence Shell Electron Pair Repulsion Theory (VSEPR) is used to calculate a molecule’s molecular geometry.

According to this hypothesis, a molecule’s geometry and structure are determined by minimising repulsion between valence shell electron pairs.

A molecule’s form and geometry can be explained as follows:

The number of electron pairs

Bonding electron pair count

No. of electron pairNo. of Bonding electron pairNo. of Lone pair of electronsGeometryBond angle
220Linear180°
330Trigonal Planar120°
21Bent117°
440Tetrahedral109.6°
31Trigonal pyramidal107°
22Bent104°
550Trigonal Bipyramidal120° + 90°
41Seesaw117° + 90°
32T-shaped90°
23Linear180°
660Octahedral90°

Each carbon atom makes three bonds, two with surrounding carbon atoms and one with the hydrogen atom, as we explained in the Lewis structure. As a result, carbon contains three bonding electron pairs, resulting in trigonal planar geometry.

As a result, benzene’s 3D structure will be as follows.

All C-C-C and H-C-C bond angles of 120° and C-C bond lengths of 139 pm conform to benzene’s trigonal planar shape.

Now we’ll talk about benzene molecule hybridization.

Hybridization of Benzene

Pauling and Slater proposed the concept of hybridization. They claim that blending distinct atomic orbitals with similar energies produces a new set of orbitals known as “Hybrid orbitals.”

The number of hybrid orbitals created should be the same as the number of mixed atomic orbitals.

Valence bond theory will help you understand the concept of hybridization (VBT). When incompletely filled atomic orbitals meet, a chemical bond is established between two atoms, according to this idea.

The benzene hybridization will be explained as follows:

The electrical configuration of carbon is 1s2 2s2 2p2, whereas the ground state of the hydrogen atom is 1s1.

When a carbon atom is excited, its electrical configuration changes to 1s2 2s1 2px1 2py1 2pz1, with four unpaired electrons in its valence shell. These unpaired electrons will be involved in the creation of bonds.

In the case of benzene, not all four carbon orbitals are utilised to create bonds.

Because each carbon atom has three bonds with two other nearby carbon atoms and a hydrogen atom, as well as one link with one of the neighbouring carbon atoms, each carbon atom is sp2 hybridised and lies in a single plane at an angle of 120°.

The unhybridized 2pz-orbital will lay perpendicular to the hybridised orbital’s plane. Two lobes, one above and one below the plane, make up this 2pz orbital. This orbital will form by overlapping sideways with the 2pz orbitals of nearby carbon atoms, resulting in the formation of a -bond.

The delocalization of -electrons is responsible for the benzene molecule’s stability.

Molecular Orbital (MO) Diagram of Benzene

Energy level diagrams illustrate chemical bonding in molecules in a molecular orbital (MO) diagram. In 1928, Robert S. Mulliken and Friedrich Hund proposed them.

The following are the postulates of MO theory:

Electrons in a molecule’s molecular orbitals are taken into account.

The development of molecular orbitals is caused by the mixing of atomic orbitals with identical energies.

Atomic orbitals should be symmetrical and have similar energy.

It is polycentric because electrons in molecular orbitals are impacted by two or more nuclei.

The number of molecular orbitals generated is equal to the number of atomic orbitals added together.

Bonding, Antibonding, and Nonbonding orbitals are the three types of molecular orbitals.

Bonding orbitals have a low energy and hence stabilise the molecule since they are closer to the nuclei, promoting molecule bonding.

Antibonding orbitals have a high energy, therefore they oppose molecular bonding.

Because non-bonding orbitals have the same energy as atomic orbitals, they do not participate in molecular bonding.

The Aufbau principle, Pauli’s exclusion principle, and Hund’s multiplicity rule are used to fill electrons in molecular orbitals.

One can forecast the stability, bond order, and magnetic property of a molecule by drawing molecular orbital diagrams.

Let’s take a look at the benzene MO diagram now.

As indicated in the picture, the p-orbitals on each carbon atom can overlap to produce six molecular orbitals, three bonding orbitals (1–3), and three antibonding orbitals (4–6).

MO 1 has the lowest energy and is the most stable bonding because it contains all six carbon atoms, i.e. it is delocalized. MO 2 and 3 are degenerate, which means they have the same energy.

They are both over 1, implying that they have more energy than 1. Antibonding orbitals 4 and 5 are degenerate antibonding orbitals that lie above bonding orbitals 2 and 3.

The remaining antibonding orbitals 6, which are higher in energy than 4 and 5, are located above 4 and 5. All electrons are present in bonding orbitals and all electrons are coupled, implying that benzene is stable. The closed shell of delocalized electrons is formed as a result of these properties.

Benzene’s Polarity

Benzene is a chemical that is nonpolar. This is owing to its symmetric ring-like form, which maintains a neutral charge distribution on the overall charge.

In nature, symmetric molecules are always non-polar. You should also read the article on the polarity of Benzene for more thorough information.

Conclusion

The Lewis structure, molecular geometry, hybridization, and molecular orbital diagram of the benzene molecule distinguish it from other molecules. Its structure demonstrates benzene characteristics that are distinct from those of other ring configurations.

Because many fragrant oils contain the benzene ring, these substances are referred to as aromatic. The majority of the accessible information on benzene is explained in layman’s terms here.

After that, anyone can easily go through any other benzene idea. This page covers all of benzene’s fundamental properties.

I hope you find the answers to your long-awaited questions at the end of this post. Please do not hesitate to ask any benzene-related queries.

Thank you for taking the time to read this article.

Read more: Honey’s pH: Acidic or Alkaline?

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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