Question

Aluminum chloride \(\left(\mathrm{AlCl}_{3}\right)\) has a melting point of \(463 \mathrm{~K}\left(190^{\circ} \mathrm{C}\right),\) whereas aluminum oxide \(\left(\mathrm{Al}_{2} \mathrm{O}_{3}\right)\) has a melting point of \(2343 \mathrm{~K}\left(2070^{\circ} \mathrm{C}\right) .\) Explain why the melting points of the two compounds are so different.

Short answer

AlCl₃ has weaker intermolecular forces, resulting in a lower melting point compared to the strong ionic bonds in Al₂O₃, which require more energy to break.

Step-by-step solution

Understanding the Compounds

Aluminum chloride (AlCl₃) and aluminum oxide (Al₂O₃) are compounds formed by aluminum and different types of bonding. Begin by examining the structure of these compounds. Aluminum chloride is a molecular compound, existing as AlCl₃ molecules, whereas aluminum oxide is an ionic compound represented as a lattice of Al³⁺ and O²⁻ ions.

Explore the Type of Bonding

Identify the type of bonding present in both compounds. AlCl₃ consists mainly of covalent bonds due to the sharing of electrons between Al and Cl atoms. In contrast, Al₂O₃ is formed by ionic bonds caused by the electrostatic attraction between positively charged Al³⁺ cations and negatively charged O²⁻ anions.

Evaluate the Intermolecular vs. Ionic Forces

Compare the forces holding these compounds together. Aluminum chloride experiences weaker intermolecular forces (van der Waals forces) between its molecules, whereas aluminum oxide is held together by strong ionic bonds between its ions in a lattice structure.

Analyze the Energy Required to Break the Forces

Consider the energy required to overcome these attractions. Breaking the covalent and Van der Waals forces in AlCl₃ requires significantly less energy compared to breaking the strong ionic bonds in Al₂O₃. This greater energy requirement for Al₂O₃ results in its much higher melting point.

Reasoning the Conclusion

Summarize why these melting points differ significantly. The covalent and weak intermolecular forces in aluminum chloride result in a lower melting point, while the robust ionic bonds in aluminum oxide lead to a much higher melting point.

Key concepts

Ionic Compounds

Ionic compounds are made of ions that form a crystal lattice. These ions are held together by strong electrostatic forces called ionic bonds, where oppositely charged ions attract each other.
For example, in aluminum oxide \( \text{(Al}_2\text{O}_3) \), aluminum ions \( (\text{Al}^{3+}) \) have lost three electrons, while oxide ions \( (\text{O}^{2-}) \) have gained two. This charge difference leads to a strong attraction between them.
This structure creates a solid compound with a high melting point since a great deal of energy is needed to break these ionic bonds and separate the ions.

• Strong attractions lead to high stability and melting points.
• Lattice structure is repeating and spacious, contributing to the strength.

This robust nature makes ionic compounds like \( \text{Al}_2\text{O}_3 \) highly resistant to heat.

Covalent Compounds

Covalent compounds are formed when atoms share electrons to achieve stability. This type of bonding is present in aluminum chloride \( (\text{AlCl}_3) \).
In these compounds, atoms overlap their outer electron shells, creating a bond by mutual sharing rather than full transfer.
This results in molecules that are often not as tightly bound as ionic compounds.

• These bonds are called covalent bonds.
• Each bond involves the sharing of electron pairs.

Because of this shared bonding, covalent compounds like \( \text{AlCl}_3 \) exhibit lower melting points.
Covalent compounds tend to have weaker forces holding their molecules together—such as van der Waals forces—which are easier to break than the ionic bonds in compounds like \( \text{Al}_2\text{O}_3 \).
This leads to reduced energy needs for breaking them.

Melting Points

Melting points indicate the temperature at which a solid becomes a liquid. This concept is crucial when comparing ionic and covalent compounds.
Aluminum chloride \( (\text{AlCl}_3) \) has a melting point of 463 K (190°C), significantly lower than that of aluminum oxide \( (\text{Al}_2\text{O}_3) \), which melts at 2343 K (2070°C).
This difference is due to the type of bonding and the forces involved in each compound.

• Ionic compounds often have high melting points due to strong lattice structures.
• Covalent compounds tend to have lower melting points due to weaker intermolecular forces.

The energy required to overcome these bonds determines the melting point, with stronger ionic bonds in \( \text{Al}_2\text{O}_3 \) requiring more energy compared to the covalent and van der Waals forces in \( \text{AlCl}_3 \).
Thus, melting points give insight into the fundamental nature of compounds and their interactions.