Which of the following species exhibits diamagnetic behaviour?
(a) O2^2-
(b) O2^+
(c) O2
(d) NO

Q: Why can't we just say an even number of electrons means the species is diamagnetic?

That's a very common trap! While odd electron species are always paramagnetic, even electron species can be paramagnetic too. The classic examples are $B_2$ (10 electrons) and $O_2$ (16 electrons), which both have two unpaired electrons due to Hund's rule in degenerate orbitals.

Q: How do I quickly find the number of unpaired electrons without drawing the whole MO diagram?

You can use the '10, 16 rule'. Molecules with exactly 10 or 16 total electrons are paramagnetic with 2 unpaired electrons. Odd electron molecules always have at least 1 unpaired electron. Most other even electron molecules (like 14 or 18) are diamagnetic.

Q: What does the 'star' (*) mean in $\pi^* 2p_x$?

The star denotes an **antibonding** molecular orbital. These orbitals have higher energy than their corresponding bonding orbitals and possess a node between the nuclei, which destabilizes the chemical bond when filled with electrons.

Q: Why do the extra electrons in $O_2^{2-}$ go into the $\pi^*$ orbitals?

Electrons fill orbitals from lowest energy to highest energy (Aufbau principle). In the $O_2$ molecule, the $\pi^* 2p$ orbitals are the lowest available unfilled orbitals (LUMO), so any added electrons must go there to minimize the energy of the system.

Q: Does adding electrons to form $O_2^{2-}$ make the oxygen-oxygen bond stronger or weaker?

It makes the bond **weaker**. The added electrons go into antibonding ($\pi^*$) orbitals. This decreases the bond order from 2 in $O_2$ to 1 in $O_2^{2-}$, making the bond longer and significantly weaker.

Question

Which of the following species exhibits diamagnetic behaviour?

Options:

(A) $O_2^{2-}$

(B) $O_2^+$

(C) $O_2$

(D) $NO$

Sigma Percentile · Accepted Answer

### The Magnetic Mystery

Have you ever wondered why liquid oxygen gets trapped between the poles of a strong magnet, while nitrogen gas just floats away? The answer lies in the invisible dance of electrons within these molecules. This phenomenon is governed by the magnetic properties of the species: **paramagnetism** and **diamagnetism**.

A species is considered **paramagnetic** if it contains one or more unpaired electrons. These unpaired electrons act like tiny bar magnets, causing the substance to be weakly attracted to an external magnetic field. On the flip side, a species is **diamagnetic** if all of its electrons are perfectly paired up. The opposing spins of paired electrons cancel out their magnetic fields, resulting in a weak repulsion from external magnetic fields.

### The Power of Molecular Orbital Theory

To determine whether a molecule's electrons are paired or unpaired, simple Lewis dot structures often fail us. For instance, a Lewis structure of $O_2$ shows all electrons paired, yet experimentally, $O_2$ is paramagnetic! This is where **Molecular Orbital Theory (MOT)** swoops in to save the day.

MOT tells us that atomic orbitals combine to form molecular orbitals, which span the entire molecule. Electrons fill these molecular orbitals following the Aufbau principle, Pauli exclusion principle, and Hund's rule of maximum multiplicity.

### Analyzing the Suspects

Let's put our options under the MOT microscope by counting their total electrons and filling their molecular orbitals.

**1. The Oxygen Molecule ($O_2$)**
Oxygen has an atomic number of 8, so $O_2$ has $8 + 8 = 16$ electrons. The molecular orbital configuration is:

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

Notice the last two electrons. Because the $\pi^* 2p_x$ and $\pi^* 2p_y$ orbitals are degenerate (have the same energy), Hund's rule dictates that they must be singly occupied before pairing occurs. Thus, $O_2$ has **two unpaired electrons** and is paramagnetic.

**2. The $O_2^+$ and $NO$ Ions**
$O_2^+$ has $16 - 1 = 15$ electrons. $NO$ has $7 (from N) + 8 (from O) = 15$ electrons.

Since both have an odd number of electrons, it is mathematically impossible for all electrons to be paired. If we look at their configurations, they both end with a single electron in a $\pi^*$ orbital. Therefore, both are paramagnetic.

### The Diamagnetic Champion

Finally, let's examine the peroxide ion, **$O_2^{2-}$**.

This ion has $16 + 2 = 18$ electrons. Let's write out its full configuration:

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

Those two extra electrons (compared to $O_2$) have entered the previously half-filled $\pi^* 2p_x$ and $\pi^* 2p_y$ orbitals. Now, every single molecular orbital is completely filled. There are absolutely **zero unpaired electrons**.

Because all of its electrons are paired, $O_2^{2-}$ perfectly exhibits **diamagnetic behaviour**. The mystery is solved!

Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: Which of the following species exhibits diamagnetic behaviour?

Select Answer:

Visualized Solution

$D iama g n e t i s m v s P a r ama g n e t i s m$

$M O C o n f i g u r a t i o n o f O_{2}$

$A na l y z in g O_{2}^{+} an d N O$

$M O C o n f i g u r a t i o n o f O_{2}^{2 -}$

$C o n c l u s i o n$

The Sigma Insight: Molecular Orbital Theory

The Magnetic Mystery

The Power of Molecular Orbital Theory

Analyzing the Suspects

The Diamagnetic Champion

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