Question

The transport of \(\mathrm{O}_{2}\) in the blood is carried out by hemoglobin. Carbon monoxide (CO) can interfere with \(\mathrm{O}_{2}\) transport because hemoglobin has a stronger affinity for CO than for \(\mathrm{O}_{2}\) If \(\mathrm{CO}\) is present, normal uptake of \(\mathrm{O}_{2}\) is prevented, depriving the body of needed \(\mathrm{O}_{2} .\) Using the molecular orbital model, write the electron configurations for \(\mathrm{CO}\) and for \(\mathrm{O}_{2} .\) From your configurations, give two property differences between \(\mathrm{CO}\) and \(\mathrm{O}_{2}\)

Short answer

The electron configurations for CO and O₂ based on the molecular orbital model are [σ(1s)² σ*(1s)² σ(2s)² σ*(2s)² σ(2p)² π(2p)⁴] for CO and [σ(1s)² σ*(1s)² σ(2s)² σ*(2s)² σ(2p)² π(2p)⁴ π*(2p)²] for O₂. Two property differences between CO and O₂ based on these configurations are: 1. CO has a triple bond (bond order = 3) between its atoms, while O₂ has a double bond (bond order = 2) between its atoms, making CO's bond stronger. 2. CO is diamagnetic (all electrons paired), while O₂ is paramagnetic (unpaired electrons in π*(2p) orbitals).

Step-by-step solution

Reminder on atomic orbitals for C and O

To build the molecular orbitals, let's recall the electron configurations of carbon (C) and oxygen (O) in their atomic forms: - Carbon has six electrons: 1s², 2s², and 2p². - Oxygen has eight electrons: 1s², 2s², and 2p⁴.

Construct molecular orbitals for CO

Use molecular orbital theory to combine the atomic orbitals of carbon and oxygen to form molecular orbitals for CO: - The \(1s\) orbital from both C and O will form a sigma (σ) bond and their antibonding counterpart (σ*). - The \(2s\) orbital from both C and O will form another sigma bond (σ) and their antibonding counterpart (σ*). - The \(2p\) orbitals will form three bonding orbitals (σ, π, π) and their antibonding counterparts (σ*, π*, π*). Now, fill in the electrons in the molecular orbitals, using the Aufbau principle (from the lowest to a higher orbital) and remembering that CO has a total of 14 electrons: - σ(1s): 2 e⁻ - σ*(1s): 2 e⁻ - σ(2s): 2 e⁻ - σ*(2s): 2 e⁻ - σ(2p): 2 e⁻ - π(2p): 4 e⁻ (2 electrons in each degenerate orbital) The electron configuration for CO is: [σ(1s)² σ*(1s)² σ(2s)² σ*(2s)² σ(2p)² π(2p)⁴]

Construct molecular orbitals for O₂

Follow the same procedure as for CO, but now for O₂, which contains 16 electrons: - σ(1s): 2 e⁻ - σ*(1s): 2 e⁻ - σ(2s): 2 e⁻ - σ*(2s): 2 e⁻ - σ(2p): 2 e⁻ - π(2p): 4 e⁻ (2 electrons in each degenerate orbital) - π*(2p): 2 e⁻ (1 electron in each degenerate orbital, as they have the same energy) The electron configuration for O₂ is: [σ(1s)² σ*(1s)² σ(2s)² σ*(2s)² σ(2p)² π(2p)⁴ π*(2p)²]

Identify two property differences based on electron configurations

1. Bond order: CO has a bond order of 3, which can be calculated as (number of bonding electrons - number of antibonding electrons) / 2 = (10 - 4) / 2 = 3. This means that CO has a triple bond between its atoms. On the other hand, O₂ has a bond order of 2 (calculated as (10 - 6)/2), which implies that O₂ has a double bond between its atoms. Therefore, CO has a stronger bond between its Carbon and Oxygen atoms than O₂ has between its two Oxygen atoms. 2. Paramagnetism: Electrons in the π*(2p) orbitals in O₂ are unpaired, making O₂ paramagnetic. In contrast, CO has all its electrons paired, so it is diamagnetic. This difference in magnetism is due to their distinct electron configurations.

Key concepts

Electron Configuration

Electron configuration provides the arrangement of electrons within an atom or a molecule. It maps out which molecular orbitals these electrons occupy. For molecules like CO and O₂, we use the molecular orbital theory to determine how atomic orbitals from individual atoms combine to form molecular orbitals.

These molecular orbitals are filled with electrons according to the Aufbau principle—all starting from the lowest energy level. For CO, with a total of 14 electrons, this results in an electron configuration of \[\sigma(1s)^2 \sigma^*(1s)^2 \sigma(2s)^2 \sigma^*(2s)^2 \sigma(2p)^2 \pi(2p)^4\].
While for O₂, which has 16 electrons, the configuration is \[\sigma(1s)^2 \sigma^*(1s)^2 \sigma(2s)^2 \sigma^*(2s)^2 \sigma(2p)^2 \pi(2p)^4 \pi^*(2p)^2\].

• The difference in the electron count affects the properties of these molecules.
• Understanding their electron configurations helps predict how they will interact with other molecules.

Bond Order

Bond order is an indicator of the strength and stability of a bond between two atoms within a molecule. It is defined as half the difference between the number of electrons in bonding and antibonding molecular orbitals.

For CO, the bond order is 3, calculated using the expression \((10 - 4) / 2 = 3\).

• This signifies that CO has a triple bond, indicating a very strong bond between the carbon and oxygen atoms.
• Triple bonds often mean higher bond energy and shorter bond lengths.
• Such a triplet in CO allows it to have a high binding affinity, hence why CO can strongly attach to hemoglobin.

In contrast, O₂ has a bond order of 2, \((10 - 6) / 2 = 2\), meaning it possesses a double bond.

• A double bond suggests moderate bond strength and stability.
• It helps oxygen remain reactive, fulfilling its role in biological processes like cellular respiration.

Paramagnetism

Paramagnetism in molecules is a result of unpaired electrons in their molecular orbitals. A molecule will be paramagnetic if even a single electron is unpaired.

O₂ has unpaired electrons in the \(\pi^*(2p)\) orbitals, reflecting its paramagnetic nature.

• This means that O₂ is attracted to magnetic fields.
• Paramagnetic properties arise because these unpaired electrons generate a magnetic field of their own.

On the other hand, CO is diamagnetic because all its electrons are arranged in pairs; no unpaired electrons exist.

• Diamagnetic substances are repelled by magnetic fields.
• CO's diamagnetism results in a lack of magnetic attraction, attributed to its full pairing of the molecular orbital electrons.
• The absence of unpaired electrons in CO makes it more stable in terms of magnetic interaction.