Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: Number of paramagnetic oxides among the following given oxides is ${\text{Li}}_2\text{O},{\text{Na}}_2{\text{O}}_2,{\text{KO}}_2,\text{HgO}$ and ${\text{K}}_2\text{O}$

Imagine you are Michael Faraday in the mid-19th century. You take a powerful electromagnet and pour liquid oxygen between its poles. To your absolute astonishment, the pale blue liquid doesn't just fall through; it sticks to the magnet! It defies gravity. This groundbreaking experiment proved that oxygen is paramagnetic—it is attracted to magnetic fields.

But there was a massive problem. According to the classical Lewis dot structures that chemists relied on, the ${\text{O}}_2$ molecule had a perfect double bond with all electrons neatly paired up. Paired electrons mean diamagnetism. The theory completely failed to explain reality.

It wasn't until the advent of Molecular Orbital Theory (MOT) that the mystery was solved. MOT revealed that the last two electrons in an ${\text{O}}_2$ molecule don't pair up; they occupy two degenerate ${\pi }^{\ast }$ antibonding orbitals separately, spinning in the same direction. These two unpaired electrons are the source of oxygen's magnetic magic.

In our problem, we are given five solid metal oxides: ${\text{Li}}_2\text{O}$, ${\text{Na}}_2{\text{O}}_2$, ${\text{KO}}_2$, $\text{HgO}$, and ${\text{K}}_2\text{O}$. We need to find out how many of them inherit this paramagnetic property.

The magnetic behavior of these ionic solids is entirely dictated by their anions. The metal cations (${\text{Li}}^{+}$, ${\text{Na}}^{+}$, ${\text{K}}^{+}$, ${\text{Hg}}^{2+}$) all possess stable, fully paired electronic configurations. They are magnetically silent. Our focus must be entirely on the oxygen species.

Magnetically Silent

Let's group the standard oxides first: ${\text{Li}}_2\text{O}$, $\text{HgO}$, and ${\text{K}}_2\text{O}$.

In these compounds, oxygen exists as the simple oxide ion, ${\text{O}}^{2-}$. A neutral oxygen atom has 8 electrons. By gaining two electrons from the metals, it reaches a total of 10 electrons.

This makes the oxide ion isoelectronic with Neon ($1s^{2}2s^{2}2p^6$). Because every single electron is paired in a closed shell, the oxide ion is strictly diamagnetic.

A Full House

Next, we examine Sodium peroxide, ${\text{Na}}_2{\text{O}}_2$. The anion here is the peroxide ion, ${\text{O}}_2^{2-}$.

A neutral ${\text{O}}_2$ molecule has 16 electrons. The peroxide ion has gained two extra electrons, bringing its total to 18 electrons. Let's consult the Molecular Orbital Theory for a diatomic molecule with 18 electrons. The filling order for the highest energy levels is:

Notice the ${\pi }^{\ast }$ antibonding orbitals. In a neutral ${\text{O}}_2$ molecule, these orbitals hold only 2 electrons (one in each). But in the peroxide ion, the two extra electrons completely fill these ${\pi }^{\ast }$ orbitals. With a full house and zero unpaired electrons, Sodium peroxide is also diamagnetic.

The Smoking Gun

Finally, we arrive at Potassium superoxide, ${\text{KO}}_2$. The anion here is the superoxide ion, ${\text{O}}_2^{-}$.

This species has 17 valence electrons (16 from ${\text{O}}_2$ plus 1 extra). Let's fill the molecular orbitals again. The first 16 electrons fill up to the ${\pi }^{\ast }$ level just like in neutral oxygen. But where does that 17th electron go?

It pairs up with one of the electrons in the ${\pi }^{\ast }$ orbitals, leaving exactly one unpaired electron in the other ${\pi }^{\ast }$ orbital. This single, lonely electron is the smoking gun! It gives the superoxide ion a net magnetic moment, making it paramagnetic.

Out of the five oxides provided, only Potassium superoxide (${\text{KO}}_2$) contains an unpaired electron. Therefore, the number of paramagnetic oxides is exactly 1.

As a parting thought, you might wonder why Potassium forms a superoxide while Lithium only forms a normal oxide. It all comes down to Lattice Energy. A large, low-charge-density cation like ${\text{K}}^{+}$ is required to stabilize a large, low-charge-density anion like ${\text{O}}_2^{-}$. Small cations like ${\text{Li}}^{+}$ have such high charge density that they force the complete reduction of oxygen to ${\text{O}}^{2-}$. Keep this lattice energy principle in your arsenal; it's a classic concept that unlocks many inorganic chemistry mysteries!