Average Bond Order: Videos & Practice Problems

Understanding the relationship between bond order, bond strength, and bond length is crucial in chemistry. Bond order indicates the number of chemical bonds between a pair of atoms, with single, double, and triple bonds having bond orders of one, two, and three, respectively. The bond strength is directly proportional to the bond order—higher bond orders result in stronger bonds. Conversely, bond length is inversely proportional to bond order; as the bond order increases, the bond becomes shorter. Recognizing these relationships helps predict the physical properties of molecules and their reactivity.

Average Bond Order

The average bond order is defined as the average number of chemical bonds between a pair of elements. In this context, a single bond is assigned a bond order of 1, a double bond has a bond order of 2, and a triple bond carries a bond order of 3. Understanding bond order is crucial because it directly influences the properties of the bond. As the average bond order increases, the strength of the bond also increases, while the length of the bond decreases. This relationship indicates that average bond order and bond strength are directly proportional, meaning that stronger bonds tend to be shorter. This concept is essential for predicting the behavior of molecules in chemical reactions and understanding their stability.

Average Bond Order Example 1

To determine the average bond order of the sulfur-oxygen bonds within the sulfite ion (SO₃^2-), we start by analyzing its resonance structures. In one common resonance structure, sulfur is double bonded to one oxygen atom and single bonded to the other two oxygen atoms. This configuration allows us to visualize the distribution of bonds among the surrounding elements.

First, we count the total number of bonds between sulfur and oxygen. In this case, there are four bonds: one double bond (counted as two) and two single bonds (counted as one each). Therefore, the total number of bonds is:

2 (from the double bond) + 1 + 1 (from the two single bonds) = 4 bonds.

Next, we identify the number of surrounding elements, which in this case are the three oxygen atoms. To find the average bond order, we divide the total number of bonds by the number of surrounding elements:

Average bond order = Total bonds / Number of surrounding elements = 4 / 3 = 1.33.

This average bond order of 1.33 indicates that while each oxygen atom is single bonded to sulfur, the presence of resonance allows for the sharing of a pi bond among the three oxygen atoms. Thus, each oxygen effectively has a third of that pi bond, leading to the fractional bond order. This method of calculating average bond order can be applied to other molecules by counting the total bonds and dividing by the number of surrounding atoms involved.

What is the bond order of the phosphate–oxygen bonds within the phosphate ion, PO₄^3–?

• A. 1.25
• B. 2
• C. 2.5
• D. 1

Answer: (A)

Here’s what students ask on this topic:

How can I find the average bond order?

To find the average bond order, you need to consider the number of bonding and antibonding electrons within a molecule's molecular orbitals. Bond order gives you an indication of the stability and strength of a bond between atoms. Here's a simple step-by-step process to calculate it:

1. Determine the Bonding Electrons: Count the total number of electrons in the bonding molecular orbitals. These are the orbitals that stabilize the molecule when they are filled with electrons.
2. Determine the Antibonding Electrons: Count the total number of electrons in the antibonding molecular orbitals. Antibonding orbitals, typically denoted with an asterisk (*), destabilize the molecule when they are filled with electrons.
3. Calculate the Bond Order: Subtract the number of antibonding electrons from the number of bonding electrons, and then divide the result by two. The formula looks like this:

$\mathrm{Bond}\mathrm{Order}=\frac{(\mathrm{Number}\mathrm{of}\mathrm{Bonding}\mathrm{Electrons}-\mathrm{Number}\mathrm{of}\mathrm{Antibonding}\mathrm{Electrons})}{2}$

A higher bond order typically indicates a stronger bond. For example, a bond order of 1 represents a single bond, 2 represents a double bond, and 3 represents a triple bond. If you're dealing with resonance structures, the bond order will be an average of these structures, reflecting the delocalization of electrons.

How do you calculate average bond order?

To calculate the average bond order in a molecule, you first need to understand the concept of bond order itself. Bond order represents the number of chemical bonds between a pair of atoms. For example, a single bond has a bond order of 1, a double bond has a bond order of 2, and a triple bond has a bond order of 3.

For molecules that have resonance structures (different ways of drawing the molecule that show different bond arrangements), the average bond order can be calculated by taking the sum of the bond orders for a given bond in all resonance structures and dividing by the number of resonance structures.

Here's a step-by-step process:

1. Draw all possible resonance structures for the molecule.
2. Count the total number of bonds of a particular type between two atoms across all the resonance structures.
3. Divide this total by the number of resonance structures to get the average bond order.

For example, in the carbonate ion (CO₃^2-), there are three resonance structures, each with one double bond and two single bonds between the carbon and oxygen atoms. The average bond order for the C-O bonds would be calculated as follows:

$(2+1+1)/3=\frac{4}{3},\text{or approximately}1.33$

What is the average bond order?

The average bond order refers to the number of chemical bonds between a pair of atoms, averaged over similar types of bonds or across a molecule. It's a way to describe the strength and stability of a bond in molecular orbital theory. To calculate the average bond order, you divide the total number of bonding electrons by the total number of molecular orbitals involved in bonding.

For example, in the molecular ion O₂⁺, there are a total of 15 electrons. In the molecular orbital diagram, 8 electrons occupy bonding orbitals (2 in σ_2s, 2 in σ*_2s, 2 in σ_2p, and 2 in Π_2p), and 7 electrons occupy antibonding orbitals (3 in Π*_2p and 4 in σ*_2p). The bond order is calculated as (number of bonding electrons - number of antibonding electrons) / 2. So for O₂⁺, it would be (8 - 7) / 2 = 0.5. This indicates a weaker bond compared to O₂, which has a bond order of 2.