Question

Arrange the following from lowest to highest ionization energy: \(\mathrm{O}, \mathrm{O}_{2}, \mathrm{O}_{2}^{-}, \mathrm{O}_{2}^{+} .\) Explain your answer.

Short answer

The order from lowest to highest ionization energy for the given species is: \(\mathrm{O} < \mathrm{O}_{2}^{-} < \mathrm{O}_{2} < \mathrm{O}_{2}^{+}\) . This is due to the relative stabilities of these species resulting from their bond strengths and molecular orbital configurations.

Step-by-step solution

1. Analyzing Species Stability

We will start by examining the stability of each species, which depends on factors such as bond strength and molecular orbital configurations. - O: For a single oxygen atom, there are 6 electrons in the valence shell and we know that half-filled or completely filled shells tend to be stable. In this case, O is not at its most stable state. - O2: Diatomic oxygen forms a double bond between the two O atoms, leading to a relatively high bond strength and a stable molecular configuration. - O2- (Oxygen anion): Adding an electron to O2 results in the anion O2-. The additional electron occupies an antibonding molecular orbital, reducing the bond strength between the two O atoms; though, O2- is still more stable than a single O atom. - O2+ (Oxygen cation): Removing an electron from O2 results in the cation O2+. This removal of one electron increases the bond strength between the O atoms, making it even more stable compared to O2 and O2-.

2. Arrange in order of Ionization Energy

Based on the stability analysis, we can arrange the given species in ascending order of ionization energy as follows: - O (least stable) ⟶ O2- ⟶ O2 ⟶ O2+ (most stable) Since species with higher stability have higher ionization energies, the order from lowest to highest ionization energy is: \(\mathrm{O} < \mathrm{O}_{2}^{-} < \mathrm{O}_{2} < \mathrm{O}_{2}^{+}\)

Key concepts

Stability of Molecular Species

The stability of molecular species is a crucial aspect in determining a molecule's chemical behavior, including its ionization energy, which is the energy required to remove an electron. Let’s delve into the specifics of different oxygen species and how their stability varies.

• Oxygen Atom (O): A single oxygen atom, with 6 electrons in its valence shell, is not in its most stable state. Atoms tend to be more stable when they have a complete or half-filled valence shell, which oxygen lacks in its elemental form.
• Diatomic Oxygen (O2): Two oxygen atoms bond together to form a molecule of O2 with a double bond. This strong bond results in a more stable configuration compared to a single atom.
• Oxygen Anion (O2-): Additional electrons weaken the bond strength as they occupy antibonding orbitals. This makes O2- less stable than O2 but generally more stable than a single oxygen atom.
• Oxygen Cation (O2+): By removing an electron from O2, the bond strength between the atoms increases, resulting in a more stable species than neutral O2, similar to achieving a noble gas-like configuration.

As stability increases, the ionization energy required to remove an electron does too. Thus, O2+ is the most stable in terms of ionization, and O is the least stable.

Molecular Orbital Theory

Molecular Orbital Theory (MOT) allows us to understand the electronic structure of molecules and predict their stability and reactivity. It describes how atomic orbitals combine to create molecular orbitals that can either bond or antibond atoms.

• Bonding vs. Antibonding Orbitals: When atomic orbitals overlap constructively, they form bonding molecular orbitals, which increase stability by lowering energy. Conversely, antibonding orbitals are formed when overlap is destructive, increasing energy and weakening stability.
• Oxygen Molecules: In O2, the electrons fill up bonding orbitals before occupying any antibonding ones. O2 has fewer electrons in antibonding orbitals compared to its anion state (O2-) and more in the cation state (O2+), explaining its intermediate stability.
• The Importance for Ionization: Removing an electron from bonding vs. antibonding orbitals directly impacts the energy needed. Losing an antibonding electron, as in O2+ formation, stabilizes the species, increasing its ionization energy.

Through MOT, we see that as the electron configuration shifts towards lower energy (more bonding than antibonding), the species becomes more stable and has higher ionization energies.

Oxygen Species Ionization

Ionization energy is predominantly influenced by the stability of molecular species and their electron configuration. Understanding this concept is vital in chemistry, particularly when analyzing why certain species have higher ionization energies.

• Lowest Ionization Energy - Oxygen Atom (O): The single oxygen atom has the lowest ionization energy. It’s less stable due to its unpaired electrons and high reactivity.
• Progression to O2-: Although O2- is an anion and contains more electrons than O2, these additional electrons are located in antibonding orbitals, weakening bonds and requiring less energy to ionize than neutral O2.
• Neutral Oxygen Molecule (O2): O2, being a stable molecule with equally filled bonding molecular orbitals, requires more energy for ionization compared to O and O2-.
• Highest Ionization Energy - Oxygen Cation (O2+): O2+ is the most stable and has the highest ionization energy. The removal of an electron from O2 allows the molecular orbitals to become more bonding, reducing potential for reactivity and thus demanding more energy to ionize.

From these points, it’s clear that the molecular configuration and bond strength largely dictate ionization energy, with stability being the driving factor in determining order.