Question

Predict which compound in each pair will have the higher melting point, and explain why for each pair. (a) HI or KI (b) \(\mathrm{Na}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{CaF}_{2}\) or \(\mathrm{HF}\) (d) \(\mathrm{SiO}_{2}\) or \(\mathrm{CO}_{2}\)

Short answer

(a) KI will have a higher melting point than HI as ionic compounds have stronger electrostatic forces compared to covalent molecules. (b) \(\mathrm{Na}_{2} \mathrm{O}\) will have a higher melting point than \(\mathrm{H}_{2} \mathrm{O}\) because ionic compounds have stronger attractive forces compared to molecular compounds. (c) \(\mathrm{CaF}_{2}\) will have a higher melting point than \(\mathrm{HF}\) since the ionic bonds in \(\mathrm{CaF}_{2}\) are stronger than the hydrogen bonds in \(\mathrm{HF}\). (d) \(\mathrm{SiO}_{2}\) will have a higher melting point than \(\mathrm{CO}_{2}\) due to its giant covalent structure with strong covalent bonds, while \(\mathrm{CO}_{2}\) has weaker dispersion forces between its molecules.

Step-by-step solution

(a) Compare HI and KI

To compare HI (Hydrogen iodide) and KI (Potassium iodide), let's first identify the bonding types. HI is a covalent molecule, while KI is an ionic compound. Ionic compounds typically have stronger electrostatic forces between their positively and negatively charged ions, leading to higher melting points. Therefore, we can predict that KI will have a higher melting point than HI.

(b) Compare \(\mathrm{Na}_{2} \mathrm{O}\) and \(\mathrm{H}_{2} \mathrm{O}\)

Both \(\mathrm{Na}_{2} \mathrm{O}\) (Sodium oxide) and \(\mathrm{H}_{2} \mathrm{O}\) (Water) contain oxygen, but while \(\mathrm{Na}_{2} \mathrm{O}\) is an ionic compound, \(\mathrm{H}_{2} \mathrm{O}\) is a molecular compound. As previously explained, ionic compounds usually have stronger attractive forces, leading to higher melting points. Thus, we can predict that \(\mathrm{Na}_{2} \mathrm{O}\) will have a higher melting point than \(\mathrm{H}_{2} \mathrm{O}\).

(c) Compare \(\mathrm{CaF}_{2}\) and \(\mathrm{HF}\)

\(\mathrm{CaF}_{2}\) (Calcium fluoride) is an ionic compound, while \(\mathrm{HF}\) (Hydrogen fluoride) is a molecular compound with hydrogen bonding, which is a stronger type of intermolecular force. Despite that, the ionic bonds in \(\mathrm{CaF}_{2}\) are stronger than the hydrogen bonds in \(\mathrm{HF}\). Therefore, we can predict that \(\mathrm{CaF}_{2}\) will have a higher melting point than \(\mathrm{HF}\).

(d) Compare \(\mathrm{SiO}_{2}\) and \(\mathrm{CO}_{2}\)

Both \(\mathrm{SiO}_{2}\) (Silicon dioxide) and \(\mathrm{CO}_{2}\) (Carbon dioxide) are molecular compounds. \(\mathrm{SiO}_{2}\) forms a giant covalent structure, while \(\mathrm{CO}_{2}\) is a simple covalent molecule. The giant covalent structure of \(\mathrm{SiO}_{2}\) has very strong covalent bonds that require a large amount of energy to break. In contrast, the intermolecular forces between \(\mathrm{CO}_{2}\) molecules are weaker dispersion forces. As a result, we can predict that \(\mathrm{SiO}_{2}\) will have a higher melting point than \(\mathrm{CO}_{2}\).

Key concepts

Ionic Compounds

Ionic compounds form when metals and nonmetals exchange electrons, resulting in positively and negatively charged ions. These oppositely charged ions attract each other strongly, creating what we call ionic bonds. These bonds are quite strong, making ionic compounds have high melting and boiling points.

For example, in the comparison of HI (Hydrogen iodide) and KI (Potassium iodide), KI is an ionic compound. The ionic bonds in KI are much stronger than the covalent bonds in HI, leading KI to have a significantly higher melting point. Examples of ionic compounds include table salt (NaCl) and magnesium oxide (MgO). They are usually solid at room temperature due to the strength of the ionic bonds.

Covalent Compounds

Covalent compounds are formed when two nonmetals share electrons. The shared electrons create covalent bonds, which are generally weaker than ionic bonds. As a result, covalent compounds tend to have lower melting and boiling points compared to ionic compounds.

In the exercise, water (H\(_2\)O) is a covalent compound compared to sodium oxide (Na\(_2\)O), which is ionic. Na\(_2\)O has stronger bonding due to its ionic nature, leading to a higher melting point than H\(_2\)O. Carbon dioxide (CO\(_2\)) is also a covalent compound where molecules are held together by weaker dispersion forces, making it easier to change states from solid to liquid or gas.

Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules. They play a critical role in determining the physical properties of substances, including their melting points. These forces are generally weaker than covalent or ionic bonds.

Different types of intermolecular forces include hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Hydrogen bonding is a stronger type of intermolecular force, as seen in HF (Hydrogen fluoride), but still weaker than ionic bonds. As presented in the exercise, despite having hydrogen bonds, HF has a lower melting point than CaF\(_2\) (Calcium fluoride) due to the stronger ionic bonds in CaF\(_2\). Understanding the types and strengths of these forces helps explain why some molecular substances transition different states more readily than others.

Bonding Types

Bonding types refer to the types of connections formed between atoms or ions within a compound. The main types are ionic, covalent, and metallic bonds. Each type of bond varies in strength, influencing the properties of the resulting compounds.

In the exercise, different bonding types are explored through paired compounds to predict melting points. For instance, SiO\(_2\) (Silicon dioxide) has a giant covalent structure, resulting in very high melting points due to the robust covalent bonds forming a lattice. On the other hand, the simple covalent structure of CO\(_2\) results in weaker intermolecular forces, leading to a lower melting point. By understanding these bonding types, students can better anticipate the physical and chemical properties of various compounds.