Hybridization

Hybridization is an important concept in chemistry introduced by Linus Pauling in 1931 to explain the molecular structure of simple molecules, like methane (CH4). We know, in methane, carbon is the central atom surrounded by four hydrogen atoms. The electronic configuration of carbon is [He] 2s2 2p2. Carbon has four electrons in the outermost shell and requires four more to complete its octet. Common sense says carbon in methane would bond with three hydrogen atoms using its three 2p orbitals and with one hydrogen atom using its 2s orbital. As we have earlier studied, all the three p orbitals are perpendicular to each other. So, three C-H bonds should be at right angles to each other and one C-H bond would be in an arbitrary direction. However, this is not the case. In methane, all four C-H bonds are of equal strength and separated by 109°28′, making a tetrahedral geometry. How is this possible? The answer to this is explained by hybridization.

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals, which are suitable for chemical bonding. The total number of atomic orbitals before mixing always equals the number of newly formed hybrid orbitals. In methane, four atomic orbitals (one 2s and three 2p) mix to form four sp3 hybrid orbitals—we will explain this in detail later on.

Remember: Hybridization is a conceptual explanation of molecular geometry; it is not a real process. The real behavior of the atomic/molecular orbital is dictated by the mathematical equations of Quantum Mechanics.

Hybridization allows chemists to explain the geometry of various chemical molecules without dealing with complex math. We have covered below common hybridization types with some examples.

sp3 Hybridization

When an s orbital and three p orbital mix, four sp3 hybrid orbitals are formed. Note: the number of orbital before and after mixing are always equal.

Methane is a classical example in which the atomic orbitals of carbon are sp3 hybridized. Also, a note to consider: Hybridization relates to an individual atom, not the molecule. When we say, methane is sp3 hybridized, we mean the carbon atom in methane is sp3 hybridized, not methane as a whole or four hydrogen atoms.

The electronic configuration of carbon in the ground state is [He] 2s2 2p2. This can be better understood from the below diagram.

Each orbital occupies two electrons. In the ground state, 2s orbital holds two electrons and two 2p orbitals hold the other one electron each, while the remaining 2p orbital remains vacant.

In such a condition, carbon cannot form bonds with four hydrogen atoms. Carbon needs four orbitals filled with a single electron for bonding. So, carbon promotes one of its electrons from 2s to the vacant 2p orbital. This is an excited state of carbon. After excitation, all the outermost orbitals are partially filled and ready for hybridization.

What happens now is the mixing of these four orbitals to form four hybrid orbitals. Since one s and three p orbitals are combined, we name the new hybrid orbitals as sp3.

These hybrid orbitals differ from their antecedents in shape, size, orientation, and energy. The hybrid orbitals of carbon in methane are spread out evenly in space. The angle between each is 109.5°, giving a tetrahedral shape to the methane molecule. The orientation of hybrid orbitals is decided by the energy of the system. The atomic system tries to minimize the total energy, so hybrid orbitals try to distribute themselves into space at the farthest separation from each other. In methane, this is achieved at an angle of 109.5°. The 3D structure of the methane is depicted in the below diagram.

In the above diagram, carbon is at the center of the structure and four hydrogens bonded to four sp3 hybrid orbitals of carbon. The angle between H-C-H is 109.5°. Since there is head-to-head overlap between all four C-H bonds, they are all sigma (σ) bonds.

For sp3 hybridization, we need four bonds. But a bond can be replaced by a lone pair of electrons if a central atom has any. For example, in ammonia (NH3), there are three sigma bonds (three N-H bonds), and nitrogen has a lone pair of electrons. Three of four sp3 orbitals of nitrogen are used for N-H bonding. And the last remaining sp3 orbital is used for holding a lone pair of electrons—see the diagram below.

The bond angle is reduced to 107° in ammonia due to the presence of a lone pair of electrons. A rationale behind this is greater repulsion offered by a lone pair of electrons than the electrons in a sigma bond. The reduced bond angle causes a bend in the tetrahedral structure.

In a water (H-O-H) molecule, the bond angle is further reduced to 104° because oxygen has two lone pairs of electrons.

sp2 Hybridization

In the sp2 hybridization, one s orbital and two p orbitals are hybridized to give three sp2 hybrid orbitals. One p orbital of three remains unhybridized. This unhybridized orbital is utilized for π bonding.

A classical example of sp2 hybridization is ethene (H2C=CH2). In ethene, two hydrogen atoms make two sigma bonds with a carbon atom. There is one C-C σ bond and one C-C π bond. The 2s orbital and two 2p orbitals of carbon hybridize to form three sp2 orbitals. One 2p remains unhybridized, which forms the π bond with 2p of the other carbon in ethene.

All three sigma bonds (2 C-H and 1 C-C) are in a single plane. And the bond angle is 120°. The diagram below illustrates the same.

The other examples of sp2 hybridization are boron trifluoride (BF3) and ozone (O3).

BF3 has three fluorine atoms separated at 120°, giving it a trigonal planar structure. In O3, we have a lone pair of electrons in central O, which reduces the bond angle to 117°.

sp Hybridization

In the sp hybridization, one s and one p orbitals hybridize to form two sp orbitals. An example of sp hybridization is acetylene (H-C≡C-H). In acetylene, both carbon atoms are sp hybridized. One 2s and one 2p orbital undergo hybridization and the remaining two 2p orbitals remain unhybridized. One sp orbital forms a σ bond with hydrogen and the other sp forms a σ bond with the other carbon. The two unhybridized p orbitals form two π bonds with the other carbon atom.

The bond angle is 180°. It is a linear geometry. Hydrogen, carbon, carbon, and hydrogen are all in a straight line.

Other examples of sp hybridization are beryllium chloride (BeCl2) and hydrogen cyanide (HCN).