sp2 Hybridisation

The Hybridisation of p orbitals is an important factor in understanding molecular structures and properties. Hybridisation involves electrons in the s and p orbitals interacting chemically to form bonding and nonbonding electron pairs. These orbital arrangements are further classified as sp, sp3, and sp2 Hybridisation. Hybridisation is the overlap of atomic orbitals around a nucleus after the electrons are arranged in an energy level diagram, often called “atomic orbital contributions to molecular orbitals.” The number of AOCM per energy level is determined by the number of s, p, and d orbitals available to each element involved in a bond.

sp2 Hybridisation: 

Some atoms have more than one hydrogen atom attached (such as CH4). The second H has a +1 formal charge when this is the case. An sp2 hybridisation occurs when a pair of p electrons are shared between the two H atoms.

The formal charge on the second hydrogen does not change the chemical bond energy because only one orbital can accept electron pairs. The resulting is termed an sp2 hybridised molecule. 

The sp2 hybrids are important in organic chemistry because they represent molecular structures and interactions, both covalent (water molecules) and ionic (amino acids, steroids, phospholipids, etc.). 

The carbonyl group occupies the lowest unoccupied molecular orbital level. It cannot accept a lone pair from an adjacent atom and form bonds. The lone pair of electrons can only be shared between two atoms or molecules through an sp2-hybridised bond.

The very highest occupied molecular orbital for this molecule is invariantly sp2 hybridised. It means that the electrons in the highest occupied orbital of this tetrahedral molecule are all paired. The only orbitals that can accept the lone pair are the sp2 hybridisation type. The third carbon (not shown) has a lone pair that can be shared with one oxygen atom or with another carbon atom.

sp2 Hybridisation of acenes:

The carbonyl group occupies the lowest unoccupied molecular orbital level. It cannot accept a lone pair from an adjacent atom and form bonds. The lone pair can only be shared between two atoms or molecules through an sp2-hybridised bond.

 This molecule’s highest occupied molecular orbital is invariantly sp2 hybridised, which means that the electrons in the highest occupied orbital of this tetrahedral molecule are all paired. The only orbitals that can accept the lone pair are the sp2 hybridisation type. The third carbon (not shown) has a lone pair that can be shared with one oxygen atom or with another carbon atom.

The shape of sp2 Hybridisation: 

The p orbitals have a linear arrangement with a slight positive lobe and a negative lobe at one end. This arrangement having no heart-like shape is termed planar. The symbol for sp2 Hybridisation represents one plane (XY) of the two planes that make up an sp2 hybridised orbital pair.

Properties of sp2 Hybridisation:

An sp2 hybridised orbitals have a linear shape, which allows for considerable overlap with other orbitals. One of the most common properties of s, p, and d orbitals is their ability to overlap in different configurations called molecular orbital (MO) levels.

sp2 hybridised molecular orbitals have some specific properties that make them very useful in organic chemistry:

• Electron density is concentrated between the nuclei of atoms or groups attached to the p atomic orbitals.

• It allows for considerable overlap with atomic orbitals on adjacent atoms through a hybridisation process, which results in the formation of σ (sigma) and π (pi) bonds. The sigma bonds are the bonds that hold molecules together.

• Hybridised orbitals exhibit resonance, an effect in which one electron pair can move between two equivalent positions without violating their wave function by distributing the pair over more than one position.

• Just as p orbitals are subject to the octet rule, s and p electrons also exhibit similar behaviour, where a maximum of eight electrons can be in any one orbital. As with the d-orbitals, electrons in sp orbitals exhibit spin-orbit coupling. sp2 Hybridisation allows for the existence of an unpaired electron.

Molecular orbitals: 

Describing molecular orbitals is tricky, especially when they are nonbonding levels. A bond has a fixed shape and energy, which cannot be changed after being formed. In contrast, the orbitals described here can be placed in many different possible shapes. These shapes are mathematical constructs. The virtual orbitals themselves are not as easy to explain because they are very strange spaces that defy conventional geometry and the rules of three-dimensional space.

The shape of a molecule’s orbitals is described in terms of their symmetry elements and the Hybridisation of electron pairs in individual orbital domains; these descriptions often look like mathematical formulas. It is because molecules and their various shapes are larger, more complex space molecules, which have unique mathematics.

 Conclusion 

A single bond cannot separate the electrons; therefore, it is impossible to expect one lone electron in the s-orbital. Thus, an sp2 hybridised molecule always has a linear shape. The theoretical orbital will have a spherical shape, and no molecular orbitals are linear in any carbon molecule. The sp hybridised molecular orbital is linear because s-orbitals are strictly nonbonding and do not overlap with other orbitals. There are no lone electrons in its bond during an sp2 hybridised molecule formation.