Question

Compare the relative stability of the following species and indicate theirmagnetic properties; \(O_2\), \(O^+_2\), \(O_2^-\) (superoxide), \(O_2^{2-}\)(peroxide)

Answer 1

Concept:

According to Molecular Orbital (MO) theory:

• Bond order = (Bonding electrons − Antibonding electrons) / 2 
• Higher bond order ⇒ greater stability 
• Presence of unpaired electrons ⇒ paramagnetic 
• Absence of unpaired electrons ⇒ diamagnetic

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Molecular orbital considerations:

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(i) O₂

Bond order = 2 
Two unpaired electrons present in π* orbitals.

Magnetic nature: Paramagnetic 
Stability: Moderate

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(ii) O₂⁺

One electron is removed from antibonding π* orbital.

Bond order = 2.5 
One unpaired electron present.

Magnetic nature: Paramagnetic 
Stability: Highest (maximum bond order)

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(iii) O₂⁻ (Superoxide)

One electron is added to antibonding π* orbital.

Bond order = 1.5 
One unpaired electron present.

Magnetic nature: Paramagnetic 
Stability: Less than O₂

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(iv) O₂²⁻ (Peroxide)

Two electrons are added to antibonding π* orbitals.

Bond order = 1 
All electrons are paired.

Magnetic nature: Diamagnetic 
Stability: Least stable

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Comparative Summary: 

Species	Bond Order	Stability	Magnetic Nature
O2+	2.5	Highest	Paramagnetic
O2	2	High	Paramagnetic
O2−	1.5	Lower	Paramagnetic
O22−	1	Lowest	Diamagnetic

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Final Order of Stability:

O₂⁺ > O₂ > O₂⁻ > O₂²⁻