Hybridization and Hybrid Orbitals

Core Concepts

In this article, you will learn hybridization, how to identify hybridization orbitals, and the way they affect molecular geometry.

Topics Covered in Other Articles

• Electron Orbitals and Orbital Shapes
• Molecular Geometry and Bond Angles
• VSEPR Theory
• What is a Chemical Bond?
• Coordination Number in Chemistry

What is hybridization?

Simply put, hybridization is the way that distinct atomic orbitals combine together to form identical hybrid orbitals which can participate in bonding much more favorably than unhybridized ones.

As you know, every atom has electron orbitals surrounding it. These orbitals are denoted s, p, d, and f, and each have distinct shapes and characteristics. When atoms bond covalently, there is a physical overlap of these orbitals in space that lets the two bonding atoms share their electrons. However, it is not the raw atomic orbitals that participate, but hybrid orbitals formed from the combination of the original ones.

Why?

Consider methane, CH₄. The central carbon atom has four orbitals at its disposal for valence (bonding) electrons to occupy – one spherical s orbital and three orthogonal p orbitals. If these orbitals were responsible for bonding, the p-bonded hydrogens and the s-bonded hydrogen would be very different, occupying different space and forming bonds of different energy to the central carbon. This is not observed. Instead, the four hydrogens in methane are all perfectly identical, splayed apart from each other and equal in energy. What is happening? How can different orbitals form identical bonds? The answer is that the orbitals are not different. Instead of three p orbitals and one s orbital, there are four sp³ hybrid orbitals.

Hybrid orbitals are extremely useful for explaining the characteristics of bonds, as well as predicting the geometry of different molecules. The latter makes use of VSEPR theory (Valence Shell Electron Pair Repulsion). For an in-depth explanation of the energetics and favorability of orbital hybridization, see our article on valence bond theory.

Sigma and Pi bonding

Before we jump into the different types of hybridization, it’s important to clarify something first. There are two main types of covalent bond: sigma (σ) and pi (π). The names seem to refer to s and p orbitals, but that connection is misleading. Sigma bonds are formed by direct, head-on overlap of orbitals. Pi bonds are formed by side-on overlap.

In the ethylene molecule above, the short bond between the two carbon atoms is a sigma bond (head on overlap). The greenish cloud is the pi bond, formed by side-on overlap. This is the key distinction: sigma bonds are formed by hybrid orbitals, and pi bonds are formed by unhybridized p orbitals.

sp³ Hybridization

This is arguably the most common hybridization state one encounters in chemistry. The carbon backbones of organic molecules are sp³ hybridized, as is water. There are four sp³ orbitals, meaning any atom with exactly four sigma bonds/lone pairs is sp³ hybridized.

Orbitals involved

Sp³ hybrid orbitals are composed of one s orbital and three p orbitals. There are no leftover p orbitals.

Geometries

There are three common geometries for this hybridization, corresponding to different numbers of lone pairs. Without lone pairs, the molecule takes a tetrahedral geometry, with the four ligands spread 109.5° apart. With one lone pair, it takes a pyramidal geometry, with an angle of 107.3° between each ligand. With two lone pairs, the molecule is called bent with angles of 104.5°.

Example compounds

sp² Hybridization

Sp² is the next most common hybridization state. Benzene is a famous molecule whose 6 carbon atoms are all sp² hybridized. Sp² is notable because it features an unhybridized p orbital that can participate in pi bonding (the exception to this being boron, which just doesn’t have enough electrons for the p orbital to be filled).

Orbitals involved

One s orbital and two p orbitals combine to form three sp² hybrid orbitals, in addition to one remaining unhybridized p orbital.

Geometries

Sp² almost always forms the trigonal planar geometry with all three ligands 120° apart. It can also form a “bent” geometry, with bond angles of slightly less than 120°. This geometry is uncommon enough that it’s not worth noting the exact value (quotes are used around “bent” because bent generally refers to the sp³ version). A linear geometry can also be formed.

Example compounds

sp Hybridization

This hybridization state is relatively uncommon. CO₂ and nitrile compounds (-CN, also called cyanide) are the most common. Acetylene (C₂H₂) and N₂ are also sp hybridized.

Orbitals involved

Sp hybrid orbitals are composed of one s and one p orbital. Atoms with sp hybridization have two leftover unhybridized p orbitals which participate in pi bonding, most often as triple bonds (although “2+2” bonding is also seen, like in CO₂ or organic molecules called ketenes, for example).

Geometries

Sp-hybridized atoms EXCLUSIVELY forms linear geometries.

Example compounds

sp³d Hybridization

Here begins the more complex hybridizations. With the introduction of sp3d1, we have entered hypervalence. You will notice that between one s, three p, and one d orbital, there are more than 8 electrons involved. Expanded octets, as they are called, are typical of heavier chalcogens (oxygen family) and pnictogens (nitrogen family) beginning with sulfur and phosphorus respectively.

Orbitals involved

One s, three p, and one d orbital result in five sp³d hybrid orbitals, with no unhybridized p orbitals left over.

Geometries

The sp³d hybridization works somewhat counterintuitively. The “saturated” geometry, with no lone pairs, is called trigonal bipyramidal, and looks like if an sp geometry was shoved through the middle of a trigonal planar one. The angles between the equatorial ligands are all 120°, the angle between the axial ligands is 180°, and the angle between the axial and planar ligands is 90°.

Here’s where things get counterintuitive. Instead of replacing the axial ligands with lone pairs, the equatorial ones go first. This results in some funky-looking geometries we’ll look at now.

The next sp³d geometry is called “see-saw” due to its resemblance to the playground structure. The 180° and 90° angles remain constant, but the angle between equatorial ligands becomes slightly smaller than the normal 120°, similarly to the “bent” geometry we saw for sp².

The last major geometry for this hybridization is called T-shaped. Like the name implies, molecules with this geometry look like a letter T, with 90° angles between the equatorial and axial ligands and the same 180° angle between the axial ones.

Technically speaking, there are more possible arrangements of an sp³d structure, including central atoms with >2 lone pairs. However, sp³d hybridization is uncommon, and rarely appears except in single molecules with only sigma bonds. Thus, the geometries of those more obscure arrangements won’t be covered here.

Example compounds

sp³d² Hybridization

This is the last major orbital hybridization one will see. Sp³d² is even less common than Sp³d, and is generally not seen in practice.

Orbitals involved

This orbital hybridization is formed from the combination of one s, three p, and two d orbitals. There are no leftover p orbitals to participate in pi bonding.

Geometries

The main sp³d² geometry one encounters is called octahedral, with bond angles of 90° between equatorial ligands, 90° between axial and equatorial ligands, and a 180° angle between axial ligands as always.

The next sp³d² geometry is called square pyramidal. Contrary to the behavior of sp³d hybridization, when a ligand in an octahedral geometry is swapped for a lone pair, an AXIAL one is lost, not an equatorial, resulting in the shape seen below.

The last and final geometry is called square planar. Like above, the next lone pair substitution occurs at the axial position, leaving the four equatorial ligands. This results in, as the name implies, a geometry with four planar ligands in a square.

Example compounds

Hybridization Chart

Video Tutorial on Hybridization

Please enjoy our animated video tutorial explaining how and why hybridization occurs.

Hybridization Practice Problems

Problem 1

Phosphate pentafluoride, PF₅, has 5 fluoride ligands bound to a central phosphate that has no lone electron pairs. What is the hybridization of the phosphate?

Problem 2

Propinal (C₃H₂O) has the following chemical structure:

What molecular geometries are present in propinal? Note: the oxygen has two lone electron pairs not shown.

Hybridization Practice Problem Solutions

1: sp³d

2: linear (hydrogens and alkyne carbons), trigonal planar (carbonyl carbon and oxygen)