Have you ever dormed with a roommate? You each have your own space, but you are a pair sharing a room. This is how electrons form bonds, their "space" (called orbitals) overlap and that bond is their "shared room". These orbitals sometimes need to hybridize (which we will discuss in detail later) so that their electrons are free to form bonds of equal energies. Imagine you were moving into your new apartment to find someone already in your bed or that you and your roommate have keys to completely different floors! This is why hybridization is important in molecules.

In this article, we will be discussing bond hybridization and how orbitals hybridize themselves to form different types of bonds.

• This article covers bond hybridization.
• First, we will look at the definition of hybridization.
• Next, we will walk through single-bond hybridization.
• Then, we will explain why pi-bonds are important in hybridization.
• Thereafter, we will discuss both double- and triple-bond hybridization.
• Lastly, we will look at the bond angles in different types of hybridized molecules.

Hybridization Definition

There are two theories that describe how bonds are made and what they look like. The first is valence bond theory. It states that two orbitals, each with one electron, overlap to form a bond. When orbitals directly overlap, that is called a σ-bond and a sideways overlap is a π-bond.

However, this theory doesn't perfectly explain all types of bonds, which is why the hybridization theory was created.

These orbitals can be used to create hybridization pi bonds and sigma bonds. The s-, p-, and d-orbitals can all be mixed to create these hybridized orbitals.

Single-bond hybridization

The first type of hybridization is single-bond hybridization or sp³ hybridization

So, why is this hybridization necessary? Let's look at CH₄ (methane) and see why hybridization is better at explaining the bonding than valence bond theory.

This is what carbon's valence (outermost) electrons look like:

In CH₄, carbon makes 4 equal bonds. However, based on the diagram, it doesn't make sense why that is the case. Not only are 2 of the electrons already paired, but these electrons are in a different energy level than the other two. Carbon instead forms 4 sp3 orbitals so that there are 4 electrons ready for bonding at the same energy level.

Now that the orbitals have been hybridized, carbon can make four σ-bonds with hydrogen. CH₄ as well as all sp³ hybridized molecules form the tetrahedral geometry.

Carbon's sp3 orbital and hydrogen's s-orbital overlap to form a σ-bond (single-bond). This geometry is called a tetrahedral and resembles a tripod.

Carbon's sp³ orbitals form four equal σ-bonds (single-bonds) by overlapping with each hydrogen's s-orbital. Each overlapping pair contains 2 electrons, one from each orbital.

Hybridization pi bonds

As mentioned previously, there are two types of bonds: σ- and π-bonds. Π-bonds are caused by the sideways overlap of orbitals. When a molecule forms a double-bond, one of the bonds will be a σ-bond, and the other will be a π-bond. For triple-bonds, two will be a π-bond and the other is a σ-bond.

Π-bonds also come in pairs. Since p-orbitals have two "lobes", if the top one is overlapping, the bottom one will too. However, they are still considered one bond.

Here we can see how the p-orbitals overlap to form the π-bonds. These bonds are present in both double- and triple-bond hybridization, so it's helpful to understand what they look like by themselves.

Double-bond hybridization

The second type of hybridization is double-bond hybridization or sp² hybridization.

Carbon hybridizes 1 2s orbital and 2 2p orbitals to form 3 sp2 orbitals, leaving one 2p orbital unhybridized. Vaia Original

The 2p-orbital is left unhybridized to form the C=C π-bond. Π-bonds can only be formed with orbitals of "p" energy or higher, so it is left untouched. Also, the 2sp² orbitals are lower in energy than the 2p orbital, since the energy level is an average of the s and p energy levels.

Let's see what these bonds look like:

Carbon's sp2 orbitals overlap with hydrogen's s-orbital and the other carbon's sp2 orbital to form single (σ) bonds. The unhybridized carbon p-orbitals overlap to form the other bond in the carbon-carbon double bond (π-bond).

Like before, the carbon hybridized orbitals (here sp² orbitals) overlap with hydrogen's s-orbital to form single bonds. The carbon p-orbitals overlap to form the second bond in the carbon-carbon double bond (π-bond). The π-bond is shown as a dotted line since the electrons in the bond are in the p-orbitals, not the sp² orbitals as shown.

Triple-bond hybridization

Lastly, let's look at triple-bond hybridization (sp-hybridization).

Let's see this bonding in action!

Carbon's sp-orbitals form a single (σ) bond by overlapping with hydrogen's s-orbitals and the other carbon's sp-orbital. The unhybridized p-orbitals form 1 π-bond each to form the second and third bond in the carbon-carbon triple bond. Vaia Original.

As before, carbon's hybridized orbitals overlap with hydrogen's s-orbital and the other carbon's hybridized orbital to form σ-bonds. The unhybridized p-orbitals overlap to form π-bonds (shown by the dotted line).

sp³, sp and sp² Hybridization and bond angles

Each type of hybridization has its own geometry. Electrons repel each other, so each geometry maximizes the distance between orbitals.

First up are single-bond/sp³ hybridized orbitals, which have the tetrahedral geometry:

Sp3/single-bond hybridized orbitals form the tetrahedral geometry. The bonds are 109.5 degrees apart. Vaia Original.

In a tetrahedral, the bond lengths and bond angles are all the same. The bond angle is 109.5°. The bottom three orbitals are all on one plane, with the top orbital sticking upward. The shape is similar to a camera tripod.

Next, double-bond/sp² hybridized orbitals form the trigonal planar geometry:

Sp2/double-bond hybridized orbitals have the trigonal planar geometry. The bond angle is 120 degrees. Vaia Original.

When we label a molecule's geometry, we base it on the center atom's geometry. When there is no main center atom, we label the geometry based on what central atom we choose. Here we consider each carbon to be a center atom, both of these carbons have the trigonal planar geometry.

Trigonal planar geometry is shaped like a triangle, with each element being on the same plane. The bond angle is 120°. In this example, we have two overlapping triangles, with each carbon being at the center of its own triangle. Sp² hybridized molecules will have two trigonal planar shapes within them, with the elements in the double-bond being their own center.

Lastly, we have triple-bond/sp hybridized orbitals, which form the linear geometry:

Sp/triple-bond hybridized orbitals form the linear geometry. The bond angles are 180 degrees. Vaia Original.

Like with the previous example, this geometry is for both elements in the triple-bond. Each carbon has a linear geometry, so it has 180° bond angles between it and what it is bonded to. Linear molecules are, as the name implies, shaped like a straight line.

In summary: