The hybridization of carbon involves the combination of two pure atomic orbitals to form a new molecular orbital "hybrid" with its own characteristics. The notion of atomic orbital gives a better explanation than the previous concept of orbit, to establish an approximation of where there is a greater probability of finding an electron within an atom.

In other words, an atomic orbital is the representation of quantum mechanics to give an idea of ​​the position of an electron or pair of electrons in a certain area within the atom, where each orbital is defined according to the values ​​of its numbers quantum.

Quantum numbers describe the state of a system (such as that of the electron inside the atom) at a given moment, through the energy belonging to the electron (n), the angular momentum that it describes in its motion (l), the related magnetic moment (m) and the spin of the electron as it travels within the atom (s).

These parameters are unique for each electron in an orbital, so two electrons cannot have exactly the same values ​​of the four quantum numbers and each orbital can be occupied by at most two electrons.

What is carbon hybridization?

To describe the hybridization of carbon, it must be taken into account that the characteristics of each orbital (its shape, energy, size, etc.) depend on the electronic configuration of each atom.

That is, the characteristics of each orbital depend on the arrangement of the electrons in each "shell" or level: from the closest to the nucleus to the outermost, also known as the valence shell.

The electrons at the outermost level are the only ones available to form a bond. Therefore, when a chemical bond is formed between two atoms, the overlap or superposition of two orbitals (one from each atom) is generated and this is closely related to the geometry of the molecules.

As previously mentioned, each orbital can be filled with a maximum of two electrons but the Aufbau Principle must be followed, by means of which the orbitals are filled according to their energy level (from the smallest to the largest), as shown shows below:

In this way, first the 1 s level is filled, then the 2 s, followed by the 2 p and so on, depending on how many electrons the atom or ion has.

Thus, hybridization is a phenomenon corresponding to molecules, since each atom can contribute only pure atomic orbitals (s, p, d, f) and, due to the combination of two or more atomic orbitals, the same amount of hybrid orbitals that allow links between elements.

Main types

Atomic orbitals have different shapes and spatial orientations, increasing in complexity, as shown below:

It is observed that there is only one type of s orbital (spherical shape), three types of p orbital (lobular shape, where each lobe is oriented on a spatial axis), five types of d orbital and seven types of f orbital, where each type of orbital possesses exactly the same energy as those of its kind.

The carbon atom in its ground state has six electrons, whose configuration is 1 s ² 2 s ² 2 p ^2. That is, they should occupy the level 1 s (two electrons), the 2 s (two electrons) and partially the 2p (the two remaining electrons) according to the Aufbau Principle.

This means that the carbon atom only has two unpaired electrons in the 2 p orbital, but thus it is not possible to explain the formation or geometry of the methane (CH ₄) molecule or other more complex ones.

So to form these bonds the hybridization of the s and p orbitals is needed (in the case of carbon), to generate new hybrid orbitals that explain even the double and triple bonds, where the electrons acquire the most stable configuration for the formation of molecules..

Sp hybridization

The sp ³ hybridization consists of the formation of four “hybrid” orbitals from the pure 2s, 2p _x, 2p _y and 2p _z orbitals.

Thus, there is the rearrangement of the electrons at level 2, where there are four electrons available for the formation of four bonds and they are arranged in parallel to have less energy (greater stability).

An example is the ethylene molecule (C ₂ H ₄), whose bonds form 120 ° angles between the atoms and give it a planar trigonal geometry.

In this case, CH and CC single bonds (due to the sp ² orbitals) and a CC double bond (due to the p orbital) are generated to form the most stable molecule.

Sp hybridization

Through sp ² hybridization, three "hybrid" orbitals are generated from the pure 2s orbital and three pure 2p orbitals. Furthermore, a pure p orbital is obtained that participates in the formation of a double bond (called pi: "π").

An example is the ethylene molecule (C ₂ H ₄), whose bonds form 120 ° angles between the atoms and give it a planar trigonal geometry. In this case, CH and CC single bonds (due to the sp ² orbitals) and a CC double bond (due to the p orbital) are generated to form the most stable molecule.