Sigma and pi bonds

🧠 Exam tips: Sigma & pi bonding

✅ Sigma (σ) = head-on overlap along the internuclear axis.
✅ Pi (π) = sideways-overlap of p-orbitals (electron density above and below the axis).
⚠️ A double bond is 1 σ + 1 π; a triple bond is 1 σ + 2 π.
💪 σ bonds are generally stronger than π bonds because the overlap is greater.
🔄 A π bond restricts rotation (so alkenes can show E/Z isomerism).
📝 Use the right language: say “overlap of orbitals” and “shared pair of electrons”, not “orbitals bond together”.
🎯 When asked about strength/length: more overlap → stronger bond; weaker bonds are usually longer.

Covalent bonds can be considered as either sigma (σ) or pi (π) bonds. The difference between these two types of covalent bonds is due to the way in which the atomic orbitals on each atom involved in forming the covalent bond overlap with each other during bond formation.

Sigma (σ) bonds

As an example, consider what happens when two hydrogen atoms combine to form a molecule of hydrogen (H₂). Hydrogen atoms contain a single electron and have an electronic configuration of 1s¹; so when two hydrogen atoms react, a covalent bond forms between the two 1s orbitals which contain the electrons. Recall that a covalent bond simply involves sharing a pair of electrons. These 1s orbitals will overlap with each other and form a covalent bond between the two atoms. When the two 1s orbitals overlap or merge with each other they will form new molecular orbitals, which will contain the 2 electrons present in the hydrogen molecule. This is outlined below:

Bond formation

As the two hydrogen atoms approach each other, the overlap of the 1s orbitals increases. This results in a stronger bond. However, as the distance between the two nuclei decreases, the amount of electrostatic repulsion between the two positively charged nuclei will increase. A point is therefore reached where there is a balance between the attraction of the nuclei for the shared pair of electrons and the repulsion between the two nuclei. The covalent bond formed between the two hydrogen atoms is formed by the head-on overlap of the 1s orbital. Covalent bonds which are formed by this head-on overlap of atomic orbitals are also called sigma (σ) bonds.

Sigma (σ) bonds and p-orbitals

When a hydrogen molecule is formed a covalent bond forms between the 2 hydrogen atoms by the head-on overlap of two 1s-orbitals. This results in the formation of a sigma bond. However sigma bonds can also form when atomic orbitals other than s-orbitals overlap. p-orbitals can also overlap in a head-on manner to form sigma bonds. Now recall that there are three p-orbitals; p_x, p_y and p_z with each orbital arranged along each of the x, y and z-axes as shown below:

The sigma bonding in fluorine

Fluorine is a halogen in group 7 of the periodic table. It has an atomic number of 9 and an electronic configuration of 1s²2s²2p⁵ or 1s²2s²2p_z²2p_y²2p_x¹. There are 5 electrons in the outer valence 2p sub-level, so fluorine is able to make one covalent bond in order to fill its outer energy level and end up with the same electron configuration as the noble gas neon (2s²2p⁶).

Fluorine gas (F₂), like hydrogen gas, consists of a diatomic molecule with a single covalent bond between the fluorine atoms. To completely fill the outer 2p sub-level each fluorine atom will form a covalent bond and so share a pair of electrons by a head-on overlap using one of its p-orbitals which contains the one unpaired electron; that is, it will form a covalent sigma bond by overlapping these two atomic p-orbitals. This is shown in the image below:

Pi(π) bonding

As we have seen above the p-orbitals in an atom can overlap in a head-on manner to form a sigma bond. However there is another way in which the p-orbitals can overlap to form a covalent bond. Instead of the p-orbitals overlapping in a head-on manner, they can overlap in a sideways- manner; this sideways-overlap results in the formation of a new type of covalent bond called a pi (π) bond. This is shown below: Two p-orbitals can overlap in a partial or sideways manner to form a pi bond.

Two p-orbitals can overlap in a partial or sideways manner to form a pi bond.

A pi (π) bond has a smaller area of overlap between the p-orbitals than a sigma bond, which means it is generally a weaker covalent bond. In many molecules with double and even triple bonds you are likely to find that the covalent bonds between the atoms in the molecule will contain a mixture of one sigma bond and one or more pi bonds.

Multiple bonds

The bonding present in molecules with single covalent bonds is likely to be a sigma bond, where there is full head-on overlap of the atomic orbitals. However in molecules which have multiple bonds such as double or triple bonds then the covalent bonds between the atoms in the molecule will be a mixture of sigma and pi(π) bonds. This is shown below:

Formation of multiple bonds in molecules

A possible way in which p-orbitals which contain only one electron could overlap and form sigma (σ) and pi(π) covalent bonds is shown below; in the diagram the P_x orbitals will overlap in a head-on manner to form a sigma bond (σ) while the P_y and P_z will partially overlap to form two pi(π) bonds.

Key Points

⚠️ Common exam mistakes

❌ Saying “orbitals bond together” instead of overlap of orbitals.
❌ Forgetting that every multiple bond contains exactly one σ bond.
❌ Saying a π bond is stronger than a σ bond.
❌ Mixing up head-on overlap (σ) with sideways-overlap (π).
❌ Forgetting that π bonds restrict rotation about a C=C bond.
❌ Describing bond strength without mentioning orbital overlap.

Sigma(σ) and pi(π) bonds