why CO2 is gas while SiO2 is solid at room temperature?

Answer 1

The difference in physical states between carbon dioxide (CO2) and silicon dioxide (SiO2) at room temperature can be attributed to the nature of their chemical bonds and the molecular structures.

Carbon dioxide (CO2) is a linear molecule consisting of one carbon atom bonded to two oxygen atoms. The carbon-oxygen double bonds in CO2 are relatively strong, but they are still considered as covalent bonds. In the gas phase at room temperature, the kinetic energy of the CO2 molecules is sufficient to overcome the attractive forces between the molecules, allowing them to move freely and remain in the gaseous state.

On the other hand, silicon dioxide (SiO2) has a three-dimensional network structure in which each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms. These strong covalent bonds form a crystalline structure called quartz or silica. The arrangement of the silicon and oxygen atoms in SiO2 creates a rigid and highly stable network that requires a significant amount of energy to disrupt.

At room temperature, the thermal energy is not sufficient to break the strong covalent bonds in SiO2, causing it to remain in a solid state. SiO2 has a high melting point (around 1,710 degrees Celsius or 3,110 degrees Fahrenheit) due to its strong intermolecular forces, which require substantial energy input to overcome.

In summary, the difference in physical states between CO2 (gas) and SiO2 (solid) at room temperature arises from the molecular structures and the strength of the chemical bonds. CO2 has a linear structure with relatively weaker covalent bonds, allowing it to exist as a gas. In contrast, SiO2 forms a three-dimensional network structure with strong covalent bonds, resulting in a solid state at room temperature.