Why is CO₂ a gas while SiO₂ is a solid at room temperature?

The difference in physical state between carbon dioxide (CO₂) and silicon dioxide (SiO₂) at room temperature is due to their molecular structure and bonding.

Carbon dioxide (CO₂):
CO₂ is a simple molecular compound. Each molecule consists of one carbon atom double-bonded to two oxygen atoms in a linear structure (O=C=O). The molecules are held together by weak van der Waals forces (intermolecular forces). These weak forces require very little energy to overcome, so CO₂ exists as a gas at room temperature.

Silicon dioxide (SiO₂):
SiO₂, in contrast, forms a giant covalent (network) structure. Each silicon atom is covalently bonded to four oxygen atoms in a tetrahedral arrangement, and each oxygen connects to two silicon atoms. This three-dimensional network creates an extremely strong and rigid lattice. Breaking the bonds in this network requires a large amount of energy, so SiO₂ exists as a solid at room temperature.

Other options are less relevant:

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  A) Carbon non-metal, silicon semi-metal: This only describes elemental properties and does not explain the difference in physical state.

  
  


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  C) CO₂ forms multiple bonds, silicon does not: While CO₂ has double bonds, the multiple bonding alone does not determine gas vs solid; intermolecular forces are key.

  
  


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  D) Silicon has only sigma bonds: The presence of sigma bonds is part of the network but does not fully explain the difference in state.

  
  

In summary, the physical state difference is primarily due to:

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  CO₂: Molecular compound with weak intermolecular forces → gas.

  
  


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  SiO₂: Network covalent solid with strong covalent bonds → solid.

  
  

Correct Answer: B) CO₂ = molecule, SiO₂ = network structure