Question

\(\mathrm{SiO}_{2}\), 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?

Short answer

In conclusion, SiO2 has a network covalent structure with strong covalent bonds, resulting in a rigid arrangement and high melting and boiling points, making it a solid at room temperature. CO2 has a linear molecular shape with weaker London dispersion forces between the molecules, allowing it to exist as a gas at room temperature. The difference in molecular structures and intermolecular forces explains their distinct states at room temperature.

Step-by-step solution

Determine the Type of Bonding in Each Compound

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.

Analyze the Intermolecular Forces

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.

Compare the Intermolecular Forces and Determine the Physical State

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.

Key concepts

Covalent Bonds

Covalent bonds are the glue that holds many molecules together. They occur when atoms share electrons to achieve stability, forming a robust connection. Each atom in a covalent bond contributes one or more electrons to form the shared pair, creating a bond that can be single, double, or triple, depending on the number of shared electron pairs.
For example, in \( ext{SiO}_2\), silicon atoms form strong covalent bonds with oxygen atoms. This arrangement results in a stable and complex network, contributing to its solid state at room temperature. In \( ext{CO}_2\), each carbon atom forms double covalent bonds with two oxygen atoms, creating a linear molecule.
Covalent bonds give molecules their definitive shapes and contribute significantly to their chemical stability and physical properties. Understanding the type and strength of covalent bonds helps to explain why different substances have varying states of matter at room temperature.

Intermolecular Forces

Intermolecular forces are the forces of attraction or repulsion between neighboring molecules. Although weaker than covalent bonds, they play a crucial role in determining the physical properties of substances, such as melting and boiling points. They are responsible for the state (solid, liquid, or gas) a substance is found in at different temperatures.
In \( ext{SiO}_2\), although the primary covalent bonds dominate, the vast network's strength also contributes to the strong intervening forces. Contrastingly, \( ext{CO}_2\) molecules experience relatively weak intermolecular forces because they are mostly influenced by London dispersion forces due to their non-polarity.
Understanding intermolecular forces allows us to not just explain the state of matter but also anticipate how substances might behave under changing conditions.

Physical States of Matter

The physical state of matter—solid, liquid, or gas—is determined by the way molecules interact with each other. At room temperature, \( ext{SiO}_2\) remains solid while \( ext{CO}_2\) is a gas, and this can be explained by their molecular structures and force interactions.
In solids like \( ext{SiO}_2\), molecules are packed tightly due to strong bonding and intermolecular forces, preventing them from moving freely. The complex network of covalent bonds creates a rigid structure, giving the substance a high melting and boiling point.
Gases like \( ext{CO}_2\) display much weaker forces between molecules, which are free to move. This freedom allows gases to expand and occupy more volume since the weak forces do not bind them in one position.
Recognizing how molecular structure influences these states provides insight into both the physical characteristics and potential applications of materials.

London Dispersion Forces

London dispersion forces are a type of weak intermolecular force. They arise from temporary dipoles that occur when electrons in neighboring molecules interact and create slight attractions. Even though these forces are momentary and weak alone, collectively, they can influence certain properties of matter.
In \( ext{CO}_2\), which lacks a net dipole moment due to its linear structure, London dispersion forces are the primary intermolecular forces. They are generally weaker than other forces, such as hydrogen bonds, making it easier for \( ext{CO}_2\) molecules to separate and exist as a gas at room temperature.
Understanding London dispersion forces is valuable for explaining why some molecules with similar atomic compositions might have different physical properties and states of matter. This highlights the importance of even the weakest molecular forces in determining the macroscopic behavior of substances.