Lewis Dot Structures: Sigma & Pi Bonds - Online Tutor, Practice Problems & Exam Prep

In chemistry, a covalent bond involves the sharing of two valence electrons between atoms, forming a strong connection known as a sigma bond (σ-bond). Sigma bonds are the primary bonds that directly link atoms in a molecule. Pi bonds (π-bonds), on the other hand, are weaker and serve to reinforce the sigma bond, much like a phone cover protects the phone. While a single bond consists of one sigma bond and no pi bonds, a double bond includes one sigma and one pi bond, and a triple bond has one sigma bond with two pi bonds. As the number of pi bonds increases, the overall bond strength increases, but the bond length decreases. This means that a triple bond is stronger and shorter than a double bond, which in turn is stronger and shorter than a single bond. Understanding the relationship between bond types, strength, and length is crucial for grasping molecular structure and stability.

Lewis Dot Structures: Sigma & Pi Bonds

Now recall a single covalent bond involve the sharing of two valence electrons between elements. So let's say you add element A and element A, the bond that connects them. That single bond, each one contributes an electron to do it. So they're each contributing A valence electrons and that's why a chemical bond, a single chemical bond has two valence electrons involved.

Now with these types of chemical bonds we can talk about Sigma versus π bonds. Now a Sigma bond which uses the symbol σ. It's the strongest form of a covalent bond that directly connects elements together. A π bond, which uses the symbol π is the weaker form of a covalent bond that insulates and protects the Sigma bond. And what we need to realize here is that as the number of π bonds increases between elements, the bond strength increases but the bond length decreases.

So let's talk about this whole idea of Sigma and π bonds when looking at single, double and triple bonds. Now if we look at a single bond, a single bond has one Sigma bond which remember is the strongest type of bond which directly connects 2 elements together and it has zero Pi bonds. So here 2 carbon single bonded to each other. We can see that the length of a bond involved when we talk about a double bond, a double bond has still 1 Sigma bond which is still this bond here. But then remember, we set a π bond insulates and protects the signal bond. So it has one Pi bond, which is this red bond here.

Now individually, that π bond is not stronger than a Sigma bond, It's just adding a little bit of extra cushion. Think of it as the cover of your cell phone, your cell phone itself, the solid piece of metal or plastic in your hand. Think of that as your Sigma bond, the phone cover that you have on it. Think of that as the π bond around it. By itself, the phone cover is not stronger than your phone, it's just offering a little bit more protection for your phone itself. Think of it like this when looking at a Sigma bond versus a π bond, the π bond is just offering a little bit of cover or protection for the Sigma bond.

Now a triple bond still has only one Sigma bond, so no matter if you're looking at a single bond, a double bond, or a triple bond, it's always one Sigma bond. Each π bond gets added makes that bond makes the total bond a little bit stronger. So here a Sigma bond has a triple bond, has one Sigma bond and 2π bonds. So we have here on top and here on the bottom. So this Sigma bond in the middle is super protected.

So if we look, we can see here that the single bond, we see that it's the longest and then we see a double bond, it strengthens, is increasing because we added a π bond. But look at the length, the length is a little bit shorter. And then when we get to the triple bond, we're very protected. Our single bond, our Sigma bond is very protected and insulated because we have a π bond on the top and the bottom. But look at the length, it's so much shorter. So just remember, as we increase the strength the covering for the Sigma bond, we have to compromise on the length of the bond itself.

Lewis Dot Structures: Sigma & Pi Bonds Example 1

Here it says which of the following statements best describes the relationship between bond length and bond strength for a series of compounds involving bonds between the same 2 atoms. All right, the greater the bond strength, the longer the bond. Now we know there's an inverse relationship between bond strength and bond length.

If one goes up, the other one has to come down. So the bond would have to get shorter, not longer. This one here is the correct answer. The greater the bond strength and the shorter the bond will get. Here, bond strength and bond length are not related. That's not true. They're inversely proportional.

The relationship between bond length and bond strength depends on other factors. No. The only correct answer here would have to be option B.

How many pi bonds does the following molecule contain? 

• A. 1
• B. 2
• C. 3
• D. 4

How many sigma bonds does the following molecule contain?

• A. 12
• B. 8
• C. 9
• D. 10

Which has greater bond strength between the carbon–carbon bond.

C₂Cl₂  vs.  C₂Cl₆

• A. C₂Cl₂
• B. C₂Cl₆

Draw the total number of sigma and pi bonds of the sulfur trioxide molecule, SO₃.

• A. 4 sigma, 0 pi
• B. 3 sigma, 1 pi
• C. 2 sigma, 2 pi
• D. 1 sigma, 3 pi

Here’s what students ask on this topic:

What is a sigma bond?

A sigma bond (σ bond) is the strongest type of covalent chemical bond where two atomic orbitals directly overlap between the nuclei of two atoms. It is formed by the head-on or axial overlapping of atomic orbitals, which allows for the maximum overlap and therefore a strong bond. In a sigma bond, the electron density is concentrated along the axis connecting the two bonding nuclei.

Sigma bonds can be formed from the overlap of s orbitals with other s orbitals, s orbitals with p orbitals, or p orbitals with p orbitals, as long as the orbitals are oriented to overlap end-to-end. Each single bond in a molecule is a sigma bond, and it allows for free rotation about the bond axis because the electron density is cylindrically symmetrical around the bond axis. In contrast, pi bonds (Π bonds), which can accompany sigma bonds in double and triple bonds, do not allow for this kind of rotation due to their different orbital overlap.

What does a sigma bond look like?

A sigma bond (σ bond) is the strongest type of covalent chemical bond and is formed by the direct overlap of atomic orbitals from two different atoms. To visualize a sigma bond, imagine two spherical s orbitals or two dumbbell-shaped p orbitals from two separate atoms coming together so that their electron clouds overlap directly between the nuclei of the atoms. This head-on overlap allows for a strong bond because the electron density is concentrated along the axis connecting the two nuclei, which maximizes the attraction between the positively charged nuclei and the negatively charged electrons.

In diagrams, sigma bonds are often represented as a single line connecting two atoms, indicating a single pair of shared electrons. For example, in a molecule of hydrogen (H₂), the two hydrogen atoms are connected by a single line, representing a sigma bond formed by the overlap of their 1s orbitals.

What is a pi bond?

A pi bond (Π bond) is a type of covalent bond that forms when two p orbitals overlap in a side-by-side fashion. Unlike sigma bonds (σ bonds), which occur when orbitals overlap end-to-end and allow for free rotation of the bonded atoms, pi bonds restrict rotation because of the parallel overlap. Pi bonds are found in double and triple bonds alongside sigma bonds. In a double bond, there is one sigma bond and one pi bond, while in a triple bond, there is one sigma bond and two pi bonds.

The electron density in a pi bond is located above and below the plane of the nuclei of the bonding atoms, which gives these bonds their characteristic properties, such as the inability to rotate and the increased reactivity compared to sigma bonds. Pi bonds are less strong than sigma bonds because the overlap is not as effective, making them more susceptible to reactions. Understanding pi bonds is crucial in organic chemistry, as they play a key role in the structure and reactivity of molecules.

What does a pi bond look like?

A pi bond (Π bond) is a type of covalent bond that forms when parallel orbitals overlap and electrons are shared between atoms. Unlike sigma bonds, which are formed by end-to-end overlapping and are found in all single bonds, pi bonds are formed from the side-to-side overlap of two p orbitals.

To visualize a pi bond, imagine two adjacent atoms, each with an electron cloud shaped like a dumbbell, which represents a p orbital. A pi bond occurs when these dumbbell-shaped clouds overlap on their sides. This overlap happens above and below the plane of the atoms involved in the bond.

In diagrams, pi bonds are often represented by a line or dash parallel to the sigma bond, indicating the side-by-side connection. In a double bond, for example, you have one sigma bond and one pi bond. The sigma bond is the first bond formed and is stronger and more stable, while the pi bond is the second bond, which is weaker and gives the double bond its additional reactivity compared to a single bond.

What is the difference between a sigma bond and a pi bond?

Sigma (σ) bonds and pi (Π) bonds are types of covalent bonds that differ in the way the atomic orbitals overlap. A sigma bond is the strongest type of covalent bond and is formed by the end-to-end (head-on) overlap of atomic orbitals. Typically, this involves the s orbitals, but p orbitals and hybrid orbitals can also form sigma bonds. A single bond between two atoms consists of one sigma bond.

On the other hand, a pi bond is formed by the side-to-side overlap of two p orbitals, with their electron density located above and below the plane of the nuclei of the bonding atoms. Pi bonds are generally weaker than sigma bonds. In a double bond, there is one sigma bond and one pi bond, while in a triple bond, there is one sigma bond and two pi bonds. The presence of pi bonds causes rigidity in the molecule's structure, restricting rotation around the bond axis, which is not the case with sigma bonds.