\(\mathrm{SiO}_{2}\), the main component of glass, is a solid at room temperature, and \(\mathrm{CO}_{2}\) is a gas at room temperature. How do the structures of these compounds explain this fact?

SiO2 mainly forms a network covalent molecular structure, consisting of silicon atoms bonded to four oxygen atoms in a tetrahedral arrangement. The bond formed between silicon and oxygen is a strong covalent bond. CO2, on the other hand, has carbon double bonded to two oxygen atoms in a linear arrangement, forming a molecular structure. The bond formed between carbon and oxygen is also a covalent bond.

Since both compounds have covalent bonds, we should consider the intermolecular forces acting between the molecules to explain their state at room temperature. SiO2 forms an extensive three-dimensional network structure due to the strong covalent bonds between silicon and oxygen atoms. This results in a more rigid and strong arrangement, which has high melting and boiling points, making it a solid at room temperature. On the other hand, CO2 has a linear molecular shape with no net dipole moment, as the polarity of the C=O bonds cancels out due to the symmetrical shape. Thus, CO2 exhibits weak London dispersion forces between the molecules which are influenced by the temporary dipoles that arise from the random motion of electrons.

The extensive network structure with strong covalent bonds in SiO2 makes it robust and rigid, resulting in high melting and boiling points. This causes SiO2 to be a solid at room temperature. On the contrary, CO2 has weaker London dispersion forces between the molecules compared to the covalent bonds in SiO2. The weak forces between CO2 molecules are easily overcome at room temperature, causing CO2 to exist as a gas. In conclusion, the difference in their molecular structures and type of intermolecular forces explain why SiO2 is a solid and CO2 is a gas at room temperature.