Problem 107 - Chemistry The Science in Context

Question

The odd-electron molecule ClO affects the atmospheric chemistry of chlorofluorocarbons as illustrated by the reaction (where the \(^{*}\) indicates an excited-state oxygen atom): $$\mathrm{CF}_{2} \mathrm{Cl}_{2}+\mathrm{O}^{*} \rightarrow \mathrm{ClO}+\mathrm{CF}_{2} \mathrm{Cl}$$ Draw a molecular orbital diagram for ClO. Is the odd electron in a bonding or antibonding orbital?

Step-by-Step Solution

Answer

Answer: The odd electron in the ClO molecule is in an antibonding molecular orbital.

Step 1: Identify Atomic Orbitals

The valence electron configuration of Chlorine (Z = 17) is: $$1s^2 2s^2 2p^6 3s^2 3p^5$$. The valence electron configuration of Oxygen (Z = 8) is: $$1s^2 2s^2 2p^4$$. The atomic orbitals involved in forming molecular orbitals for ClO are 3s and 3p for Chlorine and 2s and 2p for Oxygen.

Step 2: Combine Atomic Orbitals to Form Molecular Orbitals

We will begin by combining the s atomic orbitals. The 3s orbital from Chlorine and the 2s orbital from Oxygen will combine to form σ and σ* molecular orbitals. Next, we will combine the p atomic orbitals. Because p orbitals have three orientations - px, py and pz - we have to consider each combination: 1. Chlorine 3px and Oxygen 2px form σ and σ* molecular orbitals 2. Chlorine 3py and Oxygen 2py form π and π* molecular orbitals 3. Chlorine 3pz and Oxygen 2pz form π and π* molecular orbitals

Step 3: Draw MO Diagram

Now we'll draw the molecular orbitals in an energy level diagram, placing the bonding orbitals lower in energy than the antibonding orbitals: 3s and 2s orbitals: - σ (bonding) - σ* (antibonding) 3p and 2p orbitals: - σ (bonding) - σ* (antibonding) - π (bonding) - π* (antibonding)

Step 4: Fill in Electrons and Determine Odd Electron Location

Add up the valence electrons from both Chlorine and Oxygen: 5 electrons from Chlorine (3p^5) and 4 electrons from Oxygen (2p^4) give a total of 9 electrons. Fill in the electrons to the lowest energy molecular orbitals first following Pauli's exclusion principle and Hund's rule. 1. σ (bonding) - 2 electrons 2. σ (bonding) - 2 electrons 3. π (bonding) - 4 electrons 4. σ* (antibonding) - 1 electron There is an odd electron in the σ* antibonding molecular orbital, which is formed by the combination of Chlorine 3px and Oxygen 2px atomic orbitals. So the odd electron is in an antibonding orbital.

Key Concepts

Odd-Electron Molecule

Have you encountered an odd-electron molecule yet? It's a small but mighty species in the molecular world that features an unpaired electron. Common in reactive intermediates, these molecules are often associated with free radical reactions, which are pivotal in various chemical processes, including some that occur in our own atmosphere.

Take for example Chlorine monoxide (ClO), a byproduct of reactions involving chlorofluorocarbons and excited-state oxygen atoms, which has an odd number of electrons. Because the total number of valence electrons in ClO is 13, an unpaired electron is bound to exist in one of the molecule's orbitals. This is a hallmark characteristic of odd-electron molecules: they have at least one electron that does not pair up with another, often leading to high reactivity and a predilection for participating in chemical reactions.

Molecular Orbital Theory

Understanding the behavior of electrons in a molecule can be like solving a puzzle, and Molecular Orbital (MO) Theory provides the framework to do just that. Rather than considering electrons to reside in isolated atomic orbitals of individual atoms, MO theory proposes that when atoms combine to form molecules, their atomic orbitals overlap to create new orbitals, called molecular orbitals, which extend over the entire molecule.

In these orbitals, electrons are considered to be delocalized across the molecule, which can influence properties such as bond order, magnetism, and stability. The theory elegantly explains why some molecules like oxygen are paramagnetic, and it's essential for predicting the electronic structures of complex molecules.

Antibonding Orbitals

Now, let's delve into the concept of antibonding orbitals, a somewhat mysterious space where bonding's counterpart lies. Within the framework of MO theory, when atomic orbitals combine, they can form both bonding and antibonding molecular orbitals. While bonding orbitals stabilize a molecule through electron sharing, antibonding orbitals do the opposite; they contain electrons that actually increase the energy of the molecule, making it less stable.

These peculiar orbitals are typically higher in energy than the corresponding bonding orbitals and are denoted with an asterisk (For example, \(\sigma^*\) or \(\pi^*\)). When an odd electron populates such an orbital, it indicates that not only is the electron unpaired, but it also contributes to molecular instability, a condition that is often linked to a molecule's reactivity.

Valence Electron Configuration

The valence electron configuration is like the DNA of an atom's chemistry, revealing how it will behave and bond with other atoms. It refers to the arrangement of the outermost electrons, which are primarily responsible for the chemical properties and reactions of an element. For molecules, understanding the valence electron configuration of constituent atoms is pivotal in predicting how molecules will form.

Diving into Chlorine monoxide's valence configuration, it's evident that with 5 valence electrons from Chlorine's 3p orbital and 4 from Oxygen's 2p orbital, we have a total of 9 active players in the molecular game. In the context of MO diagrams, the configuration will determine how these electrons fill the available bonding and antibonding orbitals, hinging on rules such as Pauli's Exclusion Principle and Hund's Rule, with the ultimate goal of achieving the most stable structure possible. Understanding these configurations leads to the prediction of the molecule's structure, bond order, and magnetic properties, vastly demystifying the otherwise complex nature of molecular chemistry.