Difficulty: Introductory | Prerequisites: Basic quantum mechanics concepts (wave-particle duality, electron configuration)
This topic is the foundation for everything else in organic chemistry. Before you can understand why molecules form, how they react, or what shapes they take, you need to know how electrons occupy space around atoms and how those electron regions overlap to form bonds. If you have missed earlier material on electron configuration and quantum numbers, revisit those first. Everything from hybridisation to resonance to reaction mechanisms builds directly on what is covered here.
Orbitals are regions in space where electrons are likely to be found, governed by quantum mechanics. When atomic orbitals on different atoms overlap, they form molecular orbitals, which can be bonding (lower energy, stable) or anti-bonding (higher energy, destabilising). Sigma bonds come from head-to-head overlap; pi bonds come from side-to-side overlap and are weaker.
Orbital
A region in space where an electron is likely to be found. In simple terms, think of it as the "zone" an electron lives in around a nucleus, described mathematically by quantum mechanics.
Atomic orbital (AO)
An orbital belonging to a single, isolated atom (e.g. 1s, 2s, 2p). Think of it as the electron's home address before the atom bonds to anything.
Molecular orbital (MO)
An orbital that forms when two or more atomic orbitals overlap during bonding. It belongs to the molecule as a whole, not to one atom. In simple terms, when atoms share space, their orbitals merge into new shared orbitals.
Bonding molecular orbital
A molecular orbital that is lower in energy than the original atomic orbitals. Electrons here stabilise the bond. Think of it as the "glue" holding two atoms together.
Anti-bonding molecular orbital
A molecular orbital that is higher in energy than the original atomic orbitals. Electrons here weaken or destabilise the bond. For every bonding MO that forms, one anti-bonding MO also forms.
Sigma (σ) bond
A bond formed by direct, head-to-head overlap of orbitals along the internuclear axis. This is the strongest type of covalent bond overlap.
Pi (π) bond
A bond formed by side-to-side overlap of parallel p orbitals. Weaker than a sigma bond because the overlap is less direct.
Orbitals are regions in space where electrons are likely to be found
The study of orbitals falls under quantum mechanics
Bonds are made by the overlap of atomic orbitals, which gives rise to molecular orbitals
When two atomic orbitals combine, they always produce two molecular orbitals: one bonding, one anti-bonding
The bonding MO sits at lower energy than the original atomic orbitals, stabilising the system
The anti-bonding MO sits at higher energy, destabilising the system
For every bonding interaction, there is an anti-bonding counterpart
This is a universal rule, not an exception
Sigma bonds involve direct, head-to-head overlap of orbitals. There are three common combinations:
S + S overlap
Two s orbitals overlap head-on
Example: H₂ (each hydrogen contributes a 1s orbital)
S + P overlap
One s orbital overlaps head-on with one p orbital
Example: HCl (hydrogen's 1s overlaps with chlorine's 3p)
P + P overlap (head-to-head)
Two p orbitals overlap end-to-end along the bond axis
Results in a sigma bond (not a pi bond, because the overlap is head-on)
Formed by side-to-side overlap of p orbitals
The overlap occurs above and below (or in front of and behind) the internuclear axis, not along it
Weaker than sigma bonds because the lateral overlap is less extensive
Pi bonds explain a great deal about reactivity in organic chemistry
Double bonds consist of one σ bond + one π bond
Triple bonds consist of one σ bond + two π bonds
Example: ethylene (CH₂=CH₂) has one σ and one π bond between the carbons
P + P (head-to-head) → σ bond
P + P (side-to-side) → π bond
Double bond = 1σ + 1π
Triple bond = 1σ + 2π
Number of unhybridised p orbitals = number of π bonds × 2 (across the molecule)
Sigma and pi bonding directly determines the rigidity and reactivity of organic molecules. The restricted rotation around double bonds (caused by the pi bond locking the geometry) is the reason cis/trans isomers exist, which matters in pharmaceuticals, polymer chemistry, and biological signalling.
Students often think that all p orbital overlaps form pi bonds. They do not. If two p orbitals overlap head-to-head (end-to-end), the result is a sigma bond. Only side-to-side overlap gives a pi bond.
Students sometimes forget that anti-bonding orbitals always form alongside bonding orbitals. You cannot have one without the other.
A common error is thinking pi bonds are "extra" or unimportant because they are weaker. Pi bonds are central to reactivity, colour, and molecular geometry in organic chemistry.
⚠️ You will almost certainly be asked to identify sigma vs pi bonds in a given molecule. Count carefully: every single bond is one σ; every double bond is one σ + one π; every triple bond is one σ + two π.
⚠️ The bonding/anti-bonding concept reappears in MO theory questions, UV-Vis spectroscopy, and later in frontier molecular orbital theory (HOMO/LUMO). Get it right now.
⚠️ "For every bonding MO, there is an anti-bonding MO" is a frequently tested true/false point.
True or False: A pi bond is stronger than a sigma bond.
Fill in the blank: When two atomic orbitals combine, they produce ______ molecular orbital(s).
True or False: Side-to-side overlap of two p orbitals produces a sigma bond.
Fill in the blank: A double bond consists of ______ sigma bond(s) and ______ pi bond(s).
True or False: Anti-bonding molecular orbitals are lower in energy than the original atomic orbitals.
Answers: 1. False 2. Two (one bonding, one anti-bonding) 3. False (it produces a pi bond) 4. One sigma, one pi 5. False (they are higher in energy)
Q: In a molecule of acetylene (C₂H₂), how many sigma bonds and how many pi bonds are present in total?
A: Three sigma bonds (one C–C σ, two C–H σ) and two pi bonds (from the triple bond between the carbons).
Q: Explain why a pi bond is weaker than a sigma bond.
A: A pi bond is formed by side-to-side (lateral) overlap of p orbitals, which produces less orbital overlap than the direct, head-to-head overlap that forms a sigma bond. Less overlap means less electron density between the nuclei, so the bond is weaker.
Q: What is always true about the number of bonding and anti-bonding molecular orbitals formed when atomic orbitals combine?
A: They are always equal. For every bonding molecular orbital formed, exactly one anti-bonding molecular orbital is also formed.
Q: A molecule has 2 double bonds and 1 triple bond. How many pi bonds does it contain?
A: Four pi bonds (one from each double bond = 2, plus two from the triple bond = 2, total = 4).
This material connects directly to hybridisation (sp³, sp², sp), which explains the geometry that results from different combinations of sigma and pi bonds. It also lays the groundwork for resonance, where pi electrons delocalise across aligned p orbitals, and for reaction mechanisms, where pi bonds are often the site of attack by electrophiles.
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