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Sigma and Pi Bond - Difference, Definition, Structure, Properties, Uses, FAQs

Sigma and Pi Bond - Difference, Definition, Structure, Properties, Uses, FAQs

Team Careers360Updated on 02 Jul 2025, 04:45 PM IST

Chemical bonds are very important in the field of chemistry. Chemical bonds are mainly classified into four categories they are; coordinate bonds, covalent bonds, ionic bonds, and hydrogen bonds. The bonds which are formed by sharing of electrons between two atoms are called covalent bonds and it is also known as a molecular bond. Sigma and Pi bonds are two different types of bonds that are formed in a covalent bond. The bond that is formed by the headways overlapping of atomic orbitals is called the sigma bond and is the strongest among the covalent bonds. Sigma bonds are denoted by the letter σ . And the bond which is formed by the lateral overlapping of atoms is called the pi bond is denoted by π.

This Story also Contains

  1. Write the difference between sigma bond and pi bond.
  2. Distinguishing features of sigma and pi bonds:
  3. Sigma Antibonding and pi Antibonding Orbitals

Also read -

Write the difference between sigma bond and pi bond.

Sigma bond and pi bond difference is shown in the below table.

Sigma bond

Pi bond

Sigma bonds are formed by the headways overlapping of atomic orbitals.

Pi bonds are formed by the lateral overlapping of atomic orbitals.

Sigma bonds are denoted by σ .

A pi bond is denoted by π.

Sigma bond is the strongest bond.

The pi bonds are weaker.

It can exist independently and free rotation is possible.

Pi bonds can be accessed only with the help of a sigma bond.

It will determine the shape of the molecule.

Do not have a role in the shape of molecules.

The overlapped orbitals formed by pure orbitals, hybrid orbitals and one hybrid orbital, and one pure orbital.

Overlapping orbitals are only pure orbitals.

The sigma bond that can be formed between two atoms is one.

More than one pi bond can be formed between atoms.

Commonly Asked Questions

Q: How do sigma and pi bonds contribute differently to the strength of carbon-carbon bonds?
A:
Both sigma and pi bonds contribute to the strength of carbon-carbon bonds, but in different ways. Sigma bonds provide the primary strength and stability of the bond through direct overlap of orbitals along the internuclear axis. Pi bonds add additional strength by providing extra electron sharing between atoms, but they are generally weaker than sigma bonds. The overall bond strength increases with the number of bonds (e.g., triple bonds are stronger than double bonds, which are stronger than single bonds), but the incremental increase is less for each additional pi bond.
Q: Why are reactions involving pi bonds often faster than those involving only sigma bonds?
A:
Reactions involving pi bonds are often faster because pi electrons are more accessible and reactive than sigma electrons. The electron density in pi bonds is located above and below the bond axis, making it more exposed and available for interaction with reactants. Additionally, pi bonds are generally weaker than sigma bonds, requiring less energy to break. These factors combined make pi bonds more susceptible to attack by electrophiles or other reactive species, leading to faster reaction rates compared to reactions involving only sigma bonds.
Q: Why are sigma bonds typically not involved in resonance structures while pi bonds are?
A:
Sigma bonds are typically not involved in resonance structures because they are localized between two specific atoms and have cylindrical symmetry around the bond axis. This localization and symmetry make it difficult for sigma electrons to be shared across multiple atoms. Pi bonds, on the other hand, have electron density above and below the bond axis, making their electrons more mobile and able to be delocalized across multiple atoms. This ability to delocalize is fundamental to the concept of resonance, where electron density can be distributed over several atoms or bonds in a molecule.
Q: Why are pi bonds more easily broken than sigma bonds in most chemical reactions?
A:
Pi bonds are more easily broken than sigma bonds in most chemical reactions because they are generally weaker and their electrons are more accessible. The side-by-side overlap of p orbitals in pi bonds results in less effective sharing of electrons compared to the head-on overlap in sigma bonds. Additionally, the electron density in pi bonds is located above and below the bond axis, making it more exposed and available for interaction with reactants. This accessibility, combined with the lower bond strength, makes pi bonds more susceptible to breaking or r
Q: Can you explain how the concept of hybridization relates to the formation of sigma and pi bonds?
A:
Hybridization is closely related to the formation of sigma and pi bonds. Hybrid orbitals, formed by mixing atomic orbitals, are primarily responsible for forming sigma bonds. For example, sp3 hybrid orbitals form four equivalent sigma bonds, sp2 hybrids form three sigma bonds, and sp hybrids form two. The remaining unhybridized p orbitals are then available to form pi bonds through side-by-side overlap. Thus, the hybridization state of an atom determines how many sigma bonds it can form and how many p orbitals are left for potential pi bonding.

Distinguishing features of sigma and pi bonds:

The orbital overlapping of sigma and pi bonds is different, for a sigma bond formation headways overlapping between atomic orbitals will take place. The following images show the formation of a sigma bond between two atoms.

Sigma bond formation between two atoms.

For a pi bond formation, the atomic orbitals are overlapped in a sideways manner. The following image shows the pi bond formation.

Pi bond formation between two atoms

Sigma bond is usually defined for diatomic molecules and is formed by the headways overlapping between atomic orbitals. In the case of a Pi bond which is formed due to the lateral overlapping of orbitals and the bond is weaker than a sigma bond. A single bond always contains only one sigma bond and no pi bond is there eg. alkanes. In the case of multiple bonds like a triple and double bond, it contains pi bonds too. A double bond contains a one sigma bond and one Pi bond eg. alkenes. Similarly, in the case of a triple bond, there is one sigma bond and 2 Pi bonds are there eg. alkynes. The electron cloud in the case of the pi bond is unsymmetrical and in the case of the sigma bond is symmetrical.

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Commonly Asked Questions

Q: What is the fundamental difference between sigma and pi bonds?
A:
Sigma bonds are formed by head-on overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of bonded atoms. Pi bonds, on the other hand, are formed by side-by-side overlap of p orbitals, creating electron density above and below the plane of the nuclei. This difference in orbital overlap leads to distinct properties and behaviors of sigma and pi bonds.
Q: How do sigma and pi bonds differ in their contribution to molecular polarity?
A:
Sigma bonds contribute to molecular polarity through the electronegativity difference between bonded atoms, creating bond dipoles. Pi bonds, while not directly contributing to bond polarity, can affect the overall molecular polarity by influencing electron distribution. In molecules with multiple pi bonds, such as benzene, the delocalized pi electrons can lead to a more uniform electron distribution, potentially reducing overall molecular polarity.
Q: Why are sigma bonds generally stronger than pi bonds?
A:
Sigma bonds are typically stronger than pi bonds because of the greater orbital overlap in sigma bonds. The head-on overlap of atomic orbitals in sigma bonds allows for more effective sharing of electrons between atoms, resulting in a stronger bond. Pi bonds, with their side-by-side orbital overlap, have less direct interaction between atomic nuclei, leading to weaker bond strength.
Q: How do sigma and pi bonds contribute differently to molecular orbital theory?
A:
In molecular orbital theory, sigma bonds form molecular orbitals that are symmetrical around the internuclear axis, while pi bonds form molecular orbitals that have nodes in this axis. Sigma molecular orbitals can be formed from s orbitals or head-on overlap of p orbitals, whereas pi molecular orbitals are always formed from side-by-side overlap of p orbitals. This difference affects the energy levels and electron distribution in molecules, influencing their properties and reactivity.
Q: Why are sigma bonds considered the "backbone" of organic molecules?
A:
Sigma bonds are considered the backbone of organic molecules because they form the primary framework of covalent connections between atoms. These bonds determine the overall shape and structure of molecules, as they allow for rotation and flexibility. While pi bonds contribute to reactivity and certain properties, it's the sigma bonds that hold the molecule together and provide its fundamental skeletal structure.

Sigma Antibonding and pi Antibonding Orbitals

When an electron is present outside the region between two nuclei it is called an antibonding molecular orbital. The orbitals of two atoms overlap when they approach each other and thereby form a bond but in the case of antibonding a bond will not be formed. No overlapping of orbitals takes place; this is the only repulsion. Sigma antibonding is represented by a sigma star that is raised to an asterisk, σ*. Pi antibonding is represented by the pi star that is raised to an asterisk, π*.\

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The Following Figure Shows the Formation of Antibonding Orbitals.

Formation of Pi antibonding orbitals and sigma antibonding orbitals.

When there is a presence of extra electron like in the case of H-2 the extra electron will be present in the antibonding orbital. Thus thereby avoiding the extra electron interacting with other electrons. Antibonding orbitals have more energy compared to the bonding molecular orbitals. Only the bonding molecular orbitals participate in the bond formation. And the energy difference between two bonding molecular orbital and antibonding molecular orbital is high. The figure below shows the bonding and antibonding molecular orbital formation in the hydrogen molecule.

Bonding and antibonding molecular orbital formation.

The bonding molecular orbital is formed by the combination of + and + part or – and – part of atomic orbitals and thereby forming a bond in between the two atoms. While the antibonding molecular orbital is formed by the combination of – and + part of atomic orbitals and thereby no bond formation is taking place. And for an antibonding molecular orbital the probability of finding an electron is low and correspondingly the electron density is low.

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Commonly Asked Questions

Q: How does the presence of pi bonds affect the reactivity of a molecule?
A:
Pi bonds generally increase the reactivity of a molecule because the electrons in pi bonds are more loosely held and more accessible for chemical reactions. This accessibility makes pi bonds common sites for addition reactions, where new atoms or groups can be added to the molecule by breaking the pi bond. The increased electron density above and below the bond also makes pi bonds attractive to electrophiles (electron-seeking species).
Q: Can you explain why double bonds don't rotate freely like single bonds?
A:
Double bonds consist of one sigma bond and one pi bond. While the sigma bond allows rotation, the pi bond restricts it. This is because the pi bond is formed by the side-by-side overlap of p orbitals, which must maintain their parallel alignment to preserve the bond. Rotation would disrupt this alignment, breaking the pi bond. This restriction leads to the phenomenon of geometric isomerism in molecules with double bonds.
Q: How do sigma and pi bonds contribute to the overall bond order?
A:
Bond order represents the number of electron pairs shared between two atoms. Sigma bonds always contribute 1 to the bond order, while pi bonds can contribute additional integers. For example, a single bond (purely sigma) has a bond order of 1, a double bond (one sigma, one pi) has a bond order of 2, and a triple bond (one sigma, two pi) has a bond order of 3. Higher bond orders generally indicate stronger overall bonding between atoms.
Q: What role do pi bonds play in the concept of conjugation?
A:
Pi bonds are crucial in conjugation, which occurs when p orbitals on adjacent atoms overlap to create a delocalized electron system. In conjugated systems, pi electrons can move freely across multiple atoms, leading to increased stability and unique properties. This delocalization is responsible for phenomena like resonance in organic molecules and the vibrant colors of many organic compounds.
Q: How does the presence of sigma and pi bonds affect the hybridization of carbon atoms?
A:
The type and number of bonds a carbon atom forms determine its hybridization state. Carbon atoms forming only sigma bonds (single bonds) are typically sp3 hybridized. Those involved in one pi bond (double bond) are usually sp2 hybridized, while those in two pi bonds (triple bond) are sp hybridized. This progression from sp3 to sp2 to sp hybridization corresponds to an increase in s-character and a decrease in the bond angle between sigma bonds.

Frequently Asked Questions (FAQs)

Q: How does the presence of pi bonds affect the UV-Vis spectroscopy of organic compounds?
A:
Pi bonds significantly affect UV-Vis spectroscopy because they allow for electronic transitions at lower energies compared to sigma bonds. The pi to pi* transition (from bonding to antibonding pi orbitals) often occurs in the UV or visible region of the spectrum. As the number of conjugated pi bonds in a molecule increases, the energy required for these transitions decreases, shifting absorption to longer wavelengths. This is why many highly conjugated organic compounds are colored – they absorb light in the visible spectrum due to their extensive pi bond systems.
Q: How does the presence of pi bonds affect the reactivity of adjacent sigma bonds?
A:
Pi bonds can significantly affect the reactivity of adjacent sigma bonds through electronic effects. The electron-rich nature of pi bonds can make neighboring sigma bonds more polarized, influencing their reactivity. For example, in allylic systems (where a sigma bond is adjacent to a pi bond), the pi electrons can participate in resonance with the sigma bond, making the allylic position more susceptible to certain types of reactions, such as free radical substitution. This interaction between pi and sigma systems is also important in understanding the reactivity of conjugated systems and aromatic compounds.
Q: Can you explain how the concept of bond order relates to sigma and pi bonds?
A:
Bond order represents the number of electron pairs shared between two atoms in a covalent bond. Sigma bonds always contribute 1 to the bond order, while pi bonds can add additional integers. For example, a single bond (purely sigma) has a bond order of 1, a double bond (one sigma, one pi) has a bond order of 2, and a triple bond (one sigma, two pi) has a bond order of 3. Higher bond orders generally indicate stronger overall bonding between atoms and correlate with shorter bond lengths and higher bond energies.
Q: How do sigma and pi bonds contribute differently to the infrared (IR) spectroscopy of molecules?
A:
In IR spectroscopy, both sigma and pi bonds contribute to molecular vibrations, but they typically appear in different regions of the spectrum. Stretching vibrations of sigma bonds usually occur at higher frequencies (wavenumbers) than those of pi bonds. For example, C-H sigma bond stretches appear around 2850-3000 cm^-1, while C=C pi bond stretches appear around 1640-1680 cm^-1. Pi bonds can also participate in bending vibrations that are not possible with single bonds. The presence and position of these vibrations in IR spectra provide valuable information about molecular structure.
Q: Can you explain how sigma and pi bonds affect the acidity of organic compounds?
A:
Both sigma and pi bonds can affect the acidity of organic compounds, but in different ways. Sigma bonds influence acidity through inductive effects, where electronegativity differences along sigma bonds can stabilize or destabilize an acid's conjugate base. Pi bonds, on the other hand, can affect acidity through resonance effects. For example, in carboxylic acids, the pi system of the carbonyl group can delocalize the negative charge of the conjugate base, increasing the acid's strength. In general, pi bonds adjacent to acidic protons often increase acidity by stabilizing the conjugate base through resonance.
Q: How do sigma and pi bonds contribute differently to the heat of hydrogenation?
A:
The heat of hydrogenation is primarily affected by pi bonds. When a pi bond is hydrogenated (converted to a sigma bond), energy is released. This is because pi bonds are generally weaker than sigma bonds, so breaking a pi bond releases less energy than is gained by forming two new sigma bonds (C-H bonds). The more pi bonds a molecule contains, the greater its heat of hydrogenation. Sigma bonds, being already saturated, do not directly contribute to the heat of hydrogenation.
Q: Why do pi bonds make molecules more susceptible to photochemical reactions?
A:
Pi bonds make molecules more susceptible to photochemical reactions because they can absorb light at wavelengths that have enough energy to promote electrons from pi bonding to pi antibonding orbitals. This electronic excitation can lead to various photochemical processes, such as isomerization, cycloaddition, or bond cleavage. The lower energy gap between pi and pi* orbitals compared to sigma orbitals means that pi bonds can often absorb light in the UV or visible range, making these reactions more accessible under normal conditions.
Q: How does the presence of sigma and pi bonds affect the dipole moment of a molecule?
A:
Sigma bonds contribute to the dipole moment of a molecule through the electronegativity difference between bonded atoms. Pi bonds, while not directly creating bond dipoles, can affect the overall electron distribution in a molecule. In molecules with multiple pi bonds, especially in conjugated or aromatic systems, the delocalized pi electrons can lead to a more uniform electron distribution, potentially reducing the overall dipole moment. However, in molecules with polarized pi bonds (like carbonyls), the pi system can significantly enhance the dipole moment.
Q: How do sigma and pi bonds differ in their contribution to molecular orbital diagrams?
A:
In molecular orbital diagrams, sigma bonds form molecular orbitals that are symmetrical around the internuclear axis, while pi bonds form molecular orbitals with a node in this axis. Sigma molecular orbitals are typically lower in energy than pi molecular orbitals. The energy difference between bonding and antibonding orbitals is usually larger for sigma bonds than for pi bonds. This difference in energy levels and orbital shapes influences the electronic properties and reactivity of molecules.
Q: Why are sigma bonds always present in covalent bonding, while pi bonds are not?
A:
Sigma bonds are fundamental to covalent bonding because they represent the primary sharing of electrons between atoms. They form the basic connection needed for atoms to bond covalently. Pi bonds, on the other hand, are additional bonds that form when atoms have extra electrons to share beyond those used in sigma bonding. Not all covalently bonded atoms have the electron configuration or orbital alignment necessary to form pi bonds, making them less universal than sigma bonds.
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