D Block Elements

Paramagnetism arises from the presence of unpaired electrons in a molecule or ion. Species with one or more unpaired electrons are attracted to an external magnetic field. Diamagnetic species have all electrons paired and are weakly repelled by a magnetic field. A simple way to predict paramagnetism for simple molecules is to count the total number of valence electrons. If the total number is odd, the molecule must have at least one unpaired electron and will be paramagnetic. If the total number is even, the molecule is usually diamagnetic (though exceptions like O₂ exist, requiring Molecular Orbital Theory for explanation). Let's calculate the total valence electrons for each oxide:

• (A) SO₂: Sulfur (Group 16) has 6 valence electrons. Oxygen (Group 16) has 6 valence electrons.
  Total = 6 + 2(6) = 18 electrons (Even). Expected to be diamagnetic.
• (B) NO₂: Nitrogen (Group 15) has 5 valence electrons. Oxygen (Group 16) has 6 valence electrons.
  Total = 5 + 2(6) = 17 electrons (Odd). Expected to be paramagnetic.
• (C) SiO₂: Silicon (Group 14) has 4 valence electrons. Oxygen (Group 16) has 6 valence electrons.
  Total = 4 + 2(6) = 16 electrons (Even). Expected to be diamagnetic. (Note: SiO₂ is a network solid, but the fundamental bonding involves paired electrons).
• (D) CO₂: Carbon (Group 14) has 4 valence electrons. Oxygen (Group 16) has 6 valence electrons.
  Total = 4 + 2(6) = 16 electrons (Even). Expected to be diamagnetic.

Since NO₂ has an odd number of total valence electrons (17), it must contain at least one unpaired electron, making it paramagnetic. The other oxides have an even number of valence electrons and are diamagnetic.