Question

How would one expect the bond strength of \(\mathrm{NO}\) to compare with that of \(\mathrm{O}_{2}\) ?

Short answer

The bond strength of O₂ is expected to be greater than that of NO, as O₂ has a bond order of 2 (stronger double bond) while NO has a bond order of 1.5 (relatively weaker bond).

Step-by-step solution

Find the Electronic Configuration of the Atoms

Start by determining the electronic configuration of Nitrogen (N) and Oxygen (O) atoms. Nitrogen has 7 electrons and Oxygen has 8 electrons. Their electronic configurations are: Nitrogen (N): 1s² 2s² 2p³ Oxygen (O): 1s² 2s² 2p⁴

Sketch the Molecular Orbital Diagrams

Draw molecular orbital diagrams for both NO and O₂ to determine the bond orders and the type of bond formed in each molecule. Since they are both diatomic molecules, you can use the molecular orbital diagram of diatomic molecules involving 2s and 2p orbitals.

Determine the Bond Order for Each Molecule

Bond order is calculated as the number of electrons in bonding orbitals minus the number of electrons in antibonding orbitals, divided by 2. For NO and O₂, the molecular orbital diagrams will show the electrons in bonding (σ and π) and antibonding (σ* and π*) orbitals. For NO: Bond order = (8 - 5) / 2 = 1.5 For O₂: Bond order = (8 - 4) / 2 = 2

Compare the Bond Strengths

With higher bond order, the bond is generally stronger. Based on the calculated bond orders, we can compare the bond strengths of the two molecules: NO has a bond order of 1.5, which indicates a relatively weaker bond as compared to O₂. O₂ has a bond order of 2, which indicates a stronger double bond. Thus, we would expect the bond strength of O₂ to be greater than the bond strength of NO.

Key concepts

Molecular Orbital Theory

Understanding the strength of chemical bonds in molecules starts with grasping the molecular orbital theory. This profound concept reshaped the way chemists view how atoms share electrons within a molecule. Simplistically, when two atoms approach each other to form a bond, their atomic orbitals, areas of space where electrons are most likely to be found, overlap to create molecular orbitals. These orbitals are distinct in two forms: bonding and antibonding.
Bonding molecular orbitals are lower in energy than the atomic orbitals from which they form, stabilizing the molecule as electrons populate these spaces. Conversely, antibonding molecular orbitals are higher in energy and can destabilize a molecule if filled with electrons. The arrangement and filling of these orbitals determine the molecule's electronic structure and, ultimately, key properties such as bond strength, magnetism, and reactivity.

• Bonding Orbitals: Stabilize the molecule and lower its energy.
• Antibonding Orbitals: Potentially destabilize the molecule, increasing its energy.

The theory becomes particularly insightful when dealing with diatomic molecules, such as the NO and O₂ examples, by helping predict bonding patterns and molecular stability through orbital diagrams.

Bond Order Calculation

The bond order is a critical index reflecting the stability and strength of a bond between atoms. It is calculated using a simple formula derived from molecular orbital theory. The formula is:
Bond order = (Number of electrons in bonding orbitals - Number of electrons in antibonding orbitals) / 2.
Higher bond orders correspond to stronger, shorter bonds with greater electron sharing between the atoms. For instance, a bond order of 2 often suggests a double bond, while a bond order of 1 indicates a single bond. In the case of diatomic molecules like NO and O₂, the bond order provides a straightforward comparison of their relative bond strengths.

Interpreting Bond Order:

• Bond order of 1: Single bond, longest and weakest type of bond.
• Bond order of 2: Double bond, stronger and shorter than a single bond.
• Bond order of 3: Triple bond, strongest and shortest bond, indicative of a robust electron sharing.

The step-by-step solving strategy demonstrates that an NO molecule has a bond order of 1.5, suggesting a bond strength between that of a single and a double bond, while O₂, with a bond order of 2, is firmly in the double bond category, hence stronger.

Diatomic Molecules

Diatomic molecules, composed of two atoms, are the simplest molecules and serve as a foundation for understanding molecular bonding. These molecules can be homonuclear, comprising the same atoms like O₂, H₂, or N₂, or heteronuclear, consisting of different atoms such as CO and NO. Their behavior is a central focus of molecular orbital theory.
In studying diatomic molecules, such as NO and O₂, we examine the molecular orbital configurations that arise from combinations of the atomic orbitals of each atom. For homonuclear diatomic molecules, where the atoms possess identical electronegativity, orbital mixing is straightforward. In contrast, heteronuclear diatomic molecules like NO display more complex interactions due to differing atomic characteristics. Understanding these interactions is critical for predicting molecular properties.

Key Points:

• Diatomic molecules can be the bedrock for understanding more complex molecular structures.
• The molecular orbital theory applies to both homonuclear and heteronuclear diatomic molecules but with varying complexities.
• Predicting molecular properties and reactivity often starts with evaluating diatomic molecular behavior.

Electronic Configuration

The electronic configuration of an atom is the distribution of electrons in its atomic orbitals. It is foundational for predicting how atoms will bond and react with others. For example, nitrogen (N) with an electronic configuration of 1s² 2s² 2p³, has five valence electrons available for bonding, while oxygen (O) with an electronic configuration of 1s² 2s² 2p⁴, has six valence electrons.
In the formation of diatomic molecules such as NO and O₂, the valence electrons from each atom play a pivotal role. They fill the molecular orbitals based on the principles of molecular orbital theory, significantly impacting the resultant bond order and the molecular properties. Proper configuration and subsequent filling of the orbitals in NO and O₂ result in different bond strengths and are a key factor in various chemical and physical behaviors, such as magnetic properties and bond lengths.

Significance of Electronic Configuration:

• Dictates how atoms bond and the type of bonds formed.
• Valence electrons are the primary participants in bond formation within molecules.
• Directly influences bond order and, consequently, bond strength.

Understanding the electronic configuration of the involved atoms is the first step in determining the molecular structure and properties, such as the bond strength comparison between NO and O₂ as examined in the textbook solution.