During the change of O2 to O2^-, the incoming electron goes to the orbital.
(a) \pi 2p_x
(b) \pi^* 2

Q: Why does the electron go to the anti-bonding orbital instead of a bonding one?

Electrons always fill orbitals from lowest energy to highest energy. In the $O_2$ molecule, all the lower-energy bonding orbitals ($\sigma 2p_z$, $\pi 2p_x$, $\pi 2p_y$) are already fully occupied with two electrons each. The lowest available space is in the anti-bonding $\pi^* 2p$ orbitals.

Q: Does it matter if the electron goes to $\pi^* 2p_x$ or $\pi^* 2p_y$?

No, it does not matter. The $\pi^* 2p_x$ and $\pi^* 2p_y$ orbitals are degenerate, meaning they exist at the exact same energy level. The electron can enter either one with equal probability.

Q: Why is $O_2$ paramagnetic if it has an even number of electrons (16)?

This is a classic triumph of Molecular Orbital Theory! Because the two highest energy orbitals ($\pi^* 2p_x$ and $\pi^* 2p_y$) are degenerate, Hund's Rule states that electrons will fill them singly before pairing up. Thus, the last two electrons remain unpaired, making $O_2$ paramagnetic.

Q: What happens to the bond order when this electron is added?

The bond order is calculated as $\frac{N_b - N_a}{2}$. Since the new electron enters an anti-bonding orbital ($N_a$ increases by 1), the bond order decreases from 2.0 to 1.5. This means the bond in $O_2^-$ is weaker and longer than in $O_2$.

Q: Is $O_2^-$ diamagnetic now that we paired an electron?

No, it is still paramagnetic. While one of the $\pi^*$ orbitals is now fully paired, the other degenerate $\pi^*$ orbital still contains one unpaired electron. As long as there is at least one unpaired electron, the species is paramagnetic.

Question

During the change of $O_2$ to $O_2^-$, the incoming electron goes to the orbital.

Options:

(A) $\pi 2p_x$

(B) $\pi^* 2p_x$

(C) $\pi 2p_y$

(D) $\sigma^* 2p_z$

Sigma Percentile · Accepted Answer

### The Journey of an Electron: From Oxygen to Superoxide

Imagine you are observing a neutral oxygen molecule ($O_2$) floating in space. It is perfectly stable, but suddenly, a stray electron approaches it. The oxygen molecule captures this electron, transforming into a superoxide ion ($O_2^-$). But where exactly does this new electron go? To answer this, we must dive into the elegant world of Molecular Orbital (MO) Theory.

### Analyzing the Setup: The Oxygen Molecule

Before we can place the new electron, we need to understand the current real estate of the oxygen molecule. A neutral $O_2$ molecule has a total of 16 electrons. According to MO theory, these electrons fill the molecular orbitals in increasing order of energy.

The electronic configuration for the valence electrons of $O_2$ is:

$$ \sigma 2s^2 \ \sigma^* 2s^2 \ \sigma 2p_z^2 \ \pi 2p_x^2 = \pi 2p_y^2 \ \pi^* 2p_x^1 = \pi^* 2p_y^1 $$

Notice the last part of this configuration. The Highest Occupied Molecular Orbitals (HOMO) are the anti-bonding $\pi^*$ orbitals. Because $\pi^* 2p_x$ and $\pi^* 2p_y$ are degenerate (they have the exact same energy), Hund's Rule dictates that they must be singly occupied before any pairing occurs. This leaves $O_2$ with two unpaired electrons, which is why liquid oxygen is famously paramagnetic and can be trapped between the poles of a strong magnet!

### The Arrival of the 17th Electron

Now, the 17th electron arrives. Nature always seeks the lowest possible energy state. The bonding orbitals ($\sigma 2p_z$, $\pi 2p_x$, and $\pi 2p_y$) are already completely full, holding two electrons each.

The lowest available energy slots are in the half-filled anti-bonding orbitals: $\pi^* 2p_x$ and $\pi^* 2p_y$. The incoming electron will enter one of these degenerate orbitals and pair up with the electron already residing there.

### The Final Configuration

Once the electron settles in, the new electronic configuration for the superoxide ion ($O_2^-$) becomes:

$$ \sigma 2s^2 \ \sigma^* 2s^2 \ \sigma 2p_z^2 \ \pi 2p_x^2 = \pi 2p_y^2 \ \mathbf{\pi^* 2p_x^2} = \pi^* 2p_y^1 $$

**The 17th electron has successfully entered the $\pi^* 2p_x$ (or equivalently, the $\pi^* 2p_y$) orbital.**

This addition has profound physical consequences. Because an electron was added to an anti-bonding orbital, the overall bond order decreases from 2.0 in $O_2$ to 1.5 in $O_2^-$, making the oxygen-oxygen bond weaker and longer. Furthermore, while $O_2$ had two unpaired electrons, $O_2^-$ now only has one, meaning it is still paramagnetic, but its magnetic moment is reduced.

Understanding where electrons go isn't just an abstract exercise; it directly predicts the physical and chemical behavior of the molecules that make up our universe!

Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: During the change of $O_{2}$ to $O_{2}^{-}$ , the incoming electron goes to the orbital.

Select Answer:

Visualized Solution

$O_{2} and O_{2}^{-}$

\text{Electronic Configuration of } O_2

\text{Adding an Electron}

\text{Configuration of } O_2^-

\text{Magnetic Properties}

The Sigma Insight: Molecular Orbital Theory

The Journey of an Electron

Analyzing the Setup

The Arrival of the 17th Electron

The Final Configuration

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Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: .math-text-wrapper .katex { font-size: 1.05em !important; } .math-text-wrapper .katex-display { margin: 1rem 0 !important; } During the change of O2​ to O2−​, the incoming electron goes to the orbital.