Overview of Sigma and Pi Bonds
This lesson covers the fundamental concepts of sigma (σ) and pi (π) bonds, essential for understanding covalent bonding in organic chemistry.
Key Prerequisites
- A covalent bond forms when two atoms share a pair of electrons
- Bonding occurs when half-filled atomic orbitals overlap to create a bonding orbital
- Electrons in this orbital are attracted to both atomic nuclei, pulling the atoms together
- For a deeper refresher, review the differences between Understanding Ions and Ionic Bonds in Chemistry and covalent bonding.
What is a Sigma (σ) Bond?
- Formed by the direct overlap of two atomic orbitals (e.g., s-s, s-p, or p-p facing each other)
- The bonding orbital is close to both nuclei, resulting in strong attraction
- Single bonds are always sigma bonds
Key Features:
- Strong bond (electrons near nuclei)
- Free rotation around the bond axis is possible
What is a Pi (π) Bond?
- Formed by the sideways overlap of two p orbitals
- Occurs after a sigma bond is already formed between the same atoms
- Creates two regions of electron density: one above and one below the sigma bond
Key Features:
- Weaker than a sigma bond (electrons are further from nuclei)
- Restricted rotation – rotating one atom would twist and potentially break the pi bond
Origin in Orbital Geometry
- Each atom has three p orbitals (px, py, pz) oriented at 90° to each other
- After one p orbital forms a sigma bond, the remaining p orbitals can bend sideways to form a pi bond
- The electronic effects influencing this geometry are explained in Understanding Dipole Moments and Electronic Effects in Chemistry.
Comparing Sigma and Pi Bonds
| Feature | Sigma Bond | Pi Bond | |---------|------------|--------| | Overlap type | Direct (end-on) | Sideways | | Strength | Strong | Weaker | | Electron position | Near nuclei | Further from nuclei | | Rotation | Free | Restricted | | Bond type | Always single bonds | Part of double/triple bonds |
Case Studies: Carbon Bonding
Ethane (C2H6)
- Single C-C bond = 1 sigma bond
- Each carbon has sp3 hybridization (three bonds to H + one half-filled p orbital)
- The remaining p orbitals overlap directly to form the sigma bond
- Understanding the underlying chemical composition is aided by Understanding Chemical Formulas: Types, Ratios, and Structures Explained.
Ethene (C2H4)
- Double C-C bond = 1 sigma + 1 pi bond
- Each carbon has sp2 hybridization (two bonds to H + two half-filled p orbitals)
- Bonding sequence:
- Direct overlap of one p orbital from each carbon → sigma bond
- Sideways overlap of the remaining two p orbitals → pi bond
- The double bond is stronger than a single bond, but not twice as strong (pi bond is weaker than sigma bond)
Reactivity Insight
- Pi bonds are easier to break than sigma bonds
- This makes alkenes (like ethene) more reactive than alkanes (like ethane)
- If a pi bond breaks, the sigma bond remains, keeping the molecule intact
- To learn more about these reactivity patterns, see Interpreting 1D NMR Spectra: Chemical Shifts and Spin-Spin Coupling Explained, which often uses alkene systems as examples.
Summary
- Sigma bond: Direct orbital overlap, strong, free rotation
- Pi bond: Sideways p-orbital overlap, weaker, restricted rotation, has two electron density regions
- In double bonds: one sigma + one pi bond together
- For a broader view of atomic bonding types, explore Understanding Atomic, Molecular Elements and Ionic vs Molecular Compounds.
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hello matt here from chemistrystudent.com in this video we're going to look at sigma and pi bonds
we're going to talk about what sigma and pi bonding actually is how sigma and pi bonds arise and compare
them both by looking at the carbon bonding in ethane and ephene covalent bonding has been covered in a
separate video check the links in the description below before we talk in detail about sigma and
pi bonding there are a few essential ideas you need to be comfortable with a covalent bond is formed when two atoms
share a pair of electrons it is an example of an atomic bond when a covalent bond forms half-filled
atomic orbitals from two different atoms overlap creating a bonding orbital that a pair of electrons can exist in
electrons are negatively charged and the positively charged nuclei of both atoms are attracted to the electron density in
the new in orbital this pulls both atoms together and creates a covalent bond this is drawn as a single line between
the two atoms there are different shaped orbitals electrons can be in depending on their
distance from the nucleus of an atom the orbital shape just refers to the area that an electron pair is likely to
be in at any one time they are essentially constantly moving around within this space
at this level s and p shaped orbitals are the ones most commonly studied other half-filled atomic orbitals can
sometimes also overlap or merge creating another bonding orbital between the two atoms leading to a double bond
the atoms are harder to split apart when double bonded together although the second bond is weaker than the first
single bond recap done let's go as mentioned to form a covalent bond
between two atoms atomic orbitals must overlap or merge to create a bonding orbital
the two atomic orbitals face each other and if the orbitals can overlap easily the bond formed is called a sigma bond
shown with the greek symbol for sigma the bonding orbital is close to the nuclei of both atoms and the nuclei both
have a high level of attraction to the electrons in the orbital making the bond very strong
single bonds are always examples of sigma bonds as atoms will always try and arrange themselves to maximize the
overlap of two orbitals for example in an oxygen molecule each oxygen atom has a half-filled p-shaped
orbital the p-shaped orbitals have lobes that stick out of each side of the nucleus
and the load from each p orbital from each atom can point towards each other and the two can overlap and merge
the bond formed is a sigma bond as the both orbitals overlap directly the bond is also free to rotate as
rotating each nuclei has no impact on the bonding orbital before the oxygen atoms formed a sigma
bond they each had two half-filled p orbitals the sigma bond was formed from the direct overlap of one of these from
each atom this means that each oxygen atom still has a half-filled p orbital and wants to
form another bond now things can get a bit interesting remember that there are three p shaped
orbitals around the nucleus p x p y and p z and their lobes are each pointing in
different directions or 90 degrees to each other if two half-filled p orbitals overlap
fully to form a sigma bond the half-filled orbitals left over cannot directly line up or overlap
they are simply pointing in the wrong direction as the atoms are now quite close
together due to the single sigma bond the other half-filled orbitals can actually bend inwards slightly and
overlap sideways to create a new bonding orbital between the two atoms the electrons in this orbital are
further from the nuclei of both atoms compared to the sigma bond and the bond is therefore weaker
this type of bond is called a pi bond shown with the greek symbol pi due to each orbital having two lobes the
p orbitals overlap sideways creating two areas of electron density above and below the sigma bond already formed
this means the bond is now unable to rotate freely if one atom rotates the pi bond in
orbitals will try to rotate as well meaning they'll become twisted and if twisted too far will break
sigma bonds are covalent bonds that form from the direct overlap of two orbitals from two atoms
electrons in the bond are close to the nuclei of both atoms and the bond is strong
pi bonds are covalent bonds that form from the sideways overlap of two p-shaped orbitals from two atoms
electrons in the bond are further from the nuclei of both atoms and as a result pi bonds are weaker than sigma bonds
at this level the carbon bonding in ethane and ethene is often used to model sigma and pi bonding
in ethane the carbon atoms are bonded by the direct overlap of two orbitals a sigma bond
each carbon atom has three bonds to hydrogen atoms and this leaves one half-filled p-shaped orbital left over
the carbon atoms arrange themselves in such a way that enables the p shaped orbitals to overlap
in ifene the carbon atoms are again bonded by the direct overlap of two orbitals a sigma bond as well as a
sideways overlap of two other orbitals a pi bond this is referred to as a carbon-carbon
double bond each carbon atom has two bonds to hydrogen atoms and this leaves two
half-filled p-shaped orbitals left over to start with the carbon atoms will arrange themselves to have that direct
overlap of two orbitals and the sigma bond will form this will leave two p shaped orbitals
left over that can overlap sideways between the two atoms forming a pi bond there are now essentially two bonds
between the carbon atoms a sigma bond and a pi bond this is called a double bond and is shown as two lines between
the carbon atoms one line represents the sigma bond and the other the pi bond
as the carbon atoms are held by two bonds a carbon-carbon double bond is stronger and harder to break than a
single carbon bond sigma bonds are stronger than pi bonds however meaning that a carbon-carbon
double bond isn't twice as strong as a carbon-carbon single bond it is easier to break a pi bond than a
sigma bond this actually gives ethene a higher reactivity of an ethane and explains why alkenes are more reactive
than alkanes if a pi bond is broken the sigma bond between the atoms still remains
so to summarize covalent bonds are formed from the overlap of two atomic orbitals from two
atoms sigma bonds form from the direct overlap of atomic orbitals from two atoms
electrons in the bond are close to the nuclei of both atoms given a high level of attraction that is hard to break
making sigma bonds strong sigma bonds are free to rotate pi bonds formed from the sideways
overlap or bending of p shaped atomic orbitals from two atoms electrons in the bond are further from
the nuclei of both atoms compared to a sigma bond given a weaker attraction and a bond that is easier to break
as p shaped orbitals have two lobes a pi bond has two areas of electron density one above and one below a sigma bond
already formed between the atoms there is restricted rotation and the pi bond is unable to freely rotate and
twist i hope you found this video useful please check out other relevant videos
in the links given in the description below and visit chemistryshooting.com for free notes and revision materials
A sigma bond is formed by the direct, end-on overlap of atomic orbitals (such as s-s, s-p, or p-p facing each other). In contrast, a pi bond is formed by the sideways overlap of two p orbitals, creating two regions of electron density above and below the sigma bond. This difference in overlap type determines their distinct properties.
Pi bonds are weaker because their electrons are located further from the atomic nuclei due to the sideways orbital overlap, resulting in less effective attraction. This makes pi bonds easier to break than sigma bonds, which is why alkenes (containing double bonds with a pi bond) are more reactive than alkanes (with only sigma bonds). Even if a pi bond breaks, the stronger sigma bond keeps the molecule intact.
No, a single bond is always a sigma bond. In a double bond, there is exactly one sigma bond and one pi bond. A triple bond consists of one sigma bond and two pi bonds. The sigma bond forms first through direct overlap, and then the additional pi bonds form through sideways overlap of unhybridized p orbitals.
Restricted rotation arises because rotating one atom relative to the other in a double bond would twist and potentially break the pi bond, which requires the p orbitals to overlap sideways continuously. This restricted rotation is crucial for molecular geometry and can lead to cis-trans isomerism in alkenes, affecting their properties and reactivity.
In ethene, each carbon atom has sp² hybridization with two p orbitals remaining for bonding. First, one p orbital from each carbon overlaps directly to form a sigma bond. Then, the remaining two unhybridized p orbitals overlap sideways to form a single pi bond. This results in a double bond that is stronger than a single bond but not twice as strong due to the weaker pi bond.
A sigma bond allows free rotation around the bond axis because its electron density is directly between the nuclei, and it is a strong bond due to electrons being near the nuclei. In contrast, a pi bond has restricted rotation because rotating would disrupt the sideways overlap, and it is weaker since its electrons are further from the nuclei. These differences influence molecular stability and behavior.
The type of hybridization determines how many pi bonds can form. In sp³ hybridization (e.g., ethane), all orbitals are used for sigma bonds, so only single bonds exist. In sp² hybridization (e.g., ethene), one p orbital remains unhybridized to form a pi bond, enabling a double bond. In sp hybridization (e.g., ethyne), two unhybridized p orbitals allow two pi bonds, creating a triple bond. This directly influences the bond order and structure.
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