Question

Molecular nitrogen, carbon monoxide, and cyanide ion are isoelectronic. (a) Draw an MO diagram for each. (b) \(\mathrm{CO}\) and \(\mathrm{CN}^{-}\) are toxic. What property may explain why \(\mathrm{N}_{2}\) isn't?

Short answer

N_2's stability and inert nature make it non-toxic, unlike CO and CN^- which disrupt biological processes.

Step-by-step solution

- Understanding Isoelectronic Species

Isoelectronic species have the same number of electrons. Our first task is to confirm whether molecular nitrogen (N_2), carbon monoxide (CO), and cyanide ion (CN^-) are isoelectronic by counting their total number of electrons.

- Electron Count

N_2: Nitrogen has 7 electrons each, so N_2 has 14 electrons. CO: Carbon has 6 electrons and oxygen has 8 electrons, totaling 14 electrons for CO. CN^-: Carbon has 6 electrons, nitrogen has 7 electrons, and the extra negative charge adds 1 more, summing to 14 electrons. Thus, N_2, CO, and CN^- all have 14 electrons, confirming they are isoelectronic.

- MO Diagram for N_2

For molecular nitrogen, construct the MO diagram: 1. Atomic orbitals overlap to form bonding and antibonding molecular orbitals. 2. Fill the orbitals following Aufbau principle. - \(\text{Total electrons} = 14\) - \(\text{MO Diagram:} 1σ_{2s}^2, 1σ_{2s}^*2, 1σ_{2p}^2, 2σ_{2p}^2, 1π_{2p}^4, 1π_{2p}^*2\)

- MO Diagram for CO and CN^-

Both CO and CN^- have identical MO diagrams due to being isoelectronic with N_2. - The MO diagrams for CO and CN^-: - \(\text{MO Diagram:} 1σ_{2s}^2, 1σ_{2s}^*2, 1σ_{2p}^2, 2σ_{2p}^2, 1π_{2p}^4, 1π_{2p}^*2\)

- Comparison of Properties

A possible reason for N_2 not being toxic, compared to CO and CN^-, is its triple bond stability and inert nature. - N_2 forms very stable diatomic molecules due to a strong triple bond. - CO can bind to hemoglobin, preventing oxygen transport, and CN^- inhibits cellular respiration enzyme, leading to toxicity.

Key concepts

Molecular Orbital Theory

Molecular Orbital (MO) Theory is a fundamental concept in chemistry that helps us understand how atoms combine to form molecules. Unlike the valence bond theory that focuses on bonds localized between atoms, MO theory treats electrons as delocalized over the entire molecule.
This theory explains bonding using molecular orbitals formed by the linear combination of atomic orbitals. These molecular orbitals can be bonding, antibonding, or non-bonding:

• Bonding orbitals: These are lower in energy and stabilize the molecule.
• Antibonding orbitals: These are higher in energy and destabilize the molecule.
• Non-bonding orbitals: These do not contribute to the bond strength or stability.

Electrons fill these molecular orbitals starting from the lowest energy level, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

MO Diagrams

Molecular Orbital (MO) diagrams are visual representations that show the relative energy levels of molecular orbitals formed when atomic orbitals combine. These diagrams help us understand the bonding, antibonding, and non-bonding interactions in molecules.
Let's break down the MO diagram for isoelectronic species like \( \text{N}_2 \), CO, and CN^-:

• First, write down the total number of electrons.
• Place the electrons in molecular orbitals starting from the lowest energy ones.
• The typical MO diagram includes combinations of \( \text{σ} \) and \( \text{π} \) orbitals from \( \text{2s} \) and \( \text{2p} \) atomic orbitals.

For \( \text{N}_2 \), CO, and CN^-, the MO diagram will be:
\[ 1σ_{2s}^2, 1σ_{2s}^*2, 1σ_{2p}^2, 2σ_{2p}^2, 1π_{2p}^4, 1π_{2p}^*2 \] The notation \(1σ_{2s}^2\) means two electrons are in the first bonding sigma orbital from the \(2s\) state.

Molecular Stability

Molecular stability is determined by the arrangement and filling of molecular orbitals. The number of electrons in bonding and antibonding orbitals greatly influences this stability:

• More electrons in bonding orbitals: The molecule is more stable.
• More electrons in antibonding orbitals: The molecule is less stable.

For instance, nitrogen (\text{N}_2) has a very stable molecular structure with a triple bond due to a favorable filling of bonding orbitals. The strong triple bond in \( \text{N}_2 \) makes it inert and less reactive under normal conditions, contributing to its high stability.
On the other hand, both CO and CN^- possess strong bonds, but their interaction with biological molecules leads to toxicity. Understanding molecular stability is key in predicting reactivity and properties of molecules.

Toxicity of Molecules

When comparing the toxicity of molecules like \( \text{N}_2 \), CO, and CN^-, we must consider how these molecules interact with biological systems.

• Carbon Monoxide (CO): CO is toxic because it binds tightly to hemoglobin in the blood, preventing oxygen from being transported to cells, which can lead to suffocation and death.
• Cyanide Ion (CN^-): CN^- is highly toxic as it inhibits cytochrome c oxidase in mitochondria, blocking cellular respiration, and leading to cell death.
• Nitrogen (\text{N}_2): Although also stable, \( \text{N}_2 \) is not toxic because it does not interact negatively with biological systems. Its strong triple bond makes it inert and prevents any significant biological activity.

It’s the specific biochemical interactions of CO and CN^- that make them toxic, in contrast to the harmless nature of \( \text{N}_2 \).

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom or molecule. For molecules, this involves filling molecular orbitals based on their energy levels.
Let's look at the electron configuration for \( \text{N}_2 \), CO, and CN^-:

• \text{N}_2: \[ 1σ_{2s}^2, 1σ_{2s}^*2, 2σ_{2p}^2, 1π_{2p}^4, 3σ_{2p}^2 \]
• CO: \[ 1σ_{2s}^2, 1σ_{2s}^*2, 2σ_{2p}^2, 1π_{2p}^4, 3σ_{2p}^2 \]
• CN^-: \[ 1σ_{2s}^2, 1σ_{2s}^*2, 2σ_{2p}^2, 1π_{2p}^4, 3σ_{2p}^2 \]

Each of these configurations has a total of 14 electrons, reflecting their isoelectronic nature.
These configurations illustrate which orbitals are fully occupied and which contribute to the bonding or antibonding character of the molecules, influencing their stability and reactivity.