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Hybridization involves the mixing of atomic orbitals to form new, hybrid orbitals with different characteristics than the originals, and helps explain molecular geometries or shapes of molecules.  The primary driving force behind hybridization is the achievement of a more stable and energetically favorable molecular structure. By hybridizing atomic orbitals, the resulting hybrid orbitals are better suited to overlap efficiently with other orbitals, promoting stronger and more directional bonding.



S orbitals are atomic orbitals with a simple, spherical symmetry that are defined by the principal quantum number (n: 0,1,2…etc).  An important characteristic of s orbitals is their ability to hold a maximum of two electrons with opposite spins.  Additionally, when s-orbitals of separate atoms or molecules overlap to form bonds, they do so at an orientation of 180 degrees.

Unlike s orbitals, which are spherical, p orbitals are atomic orbitals that come in sets of three per energy level.  Each is aligned along one of the three perpendicular axes (x, y, and z) in space, therefore all three p orbitals are oriented at 90 degrees to each other.  Additionally, they are characterized by a dumbbell shape.  The orientation of overlap of p-orbitals is more complicated in that depends on whether the orbital is hybridized or not.

An electron domain refers to a region in the space around a nucleus where electrons are likely to be found.  A domain can be a single bond, a double bond, a triple bond, or a lone pair of electrons. These domains influence the overall shape of a molecule because electrons, being negatively charged, repel each other. The VSEPR theory posits that electron pairs will arrange themselves around a central atom in a way that minimizes repulsion and maximizes the distance between them.


Hybrid orbitals are the result of mixing of atomic orbitals from the same atom to create new, hybrid orbitals with different shapes and energies than the original orbitals.  The most common types of hybridization involve the combination of s and p orbitals. For example, in sp3 hybridization, one s orbital and three p orbitals from a single atom combine to form four equivalent sp3 hybrid orbitals. These hybrid orbitals then arrange themselves in a tetrahedral geometry, ideal for forming sigma (σ) bonds. Hybridization helps to rationalize molecular shapes and bond angles, providing a more precise model than would be possible by considering only the individual atomic orbitals.  A general rule of thumb the number of atomic orbitals used to hybridize is equal to the number of new hybrid orbitals formed.


Sigma (σ) bonds describe a type of overlap of orbitals.  These bonds hold atoms together within molecules.  In a sigma bond, there is at least one hybridized orbital involved in the overlap and electron density is concentrated along the internuclear x-axis.  Sigma bonds are typically associated with single bonds in molecules, but also occur in multiple-bonded molecules, such as those with double or triple bonds. In such cases, the first bond formed between two atoms is a sigma bond, while additional bonds involve pi (π) bonds that form from the side-to-side overlap of p orbitals.

Pi (π) bonds are associated with double and triple bonds in molecules.  Unlike sigma (σ) bonds, which result from head-on overlap of atomic orbitals along the internuclear axis, pi bonds arise from the side-to-side overlap of p orbitals. In a pi bond, the electron density is concentrated above and below the plane defined by the two bonded atoms. This region of electron density forms a "pi bond cloud" that is perpendicular to the internuclear axis. The formation of pi bonds involves the overlap of parallel un-hybridized p orbitals, and each pi bond is accompanied by a sigma bond. For example, in a molecule with a double bond, the first bond formed is a sigma bond, while the second bond is a pi bond resulting from the overlap of two parallel p orbitals.


Electron Domains

Hybrid Orbitals

Sigma Bonds

Pi Bonds

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