Question

Indicate which compound in each of the following pairs is more likely to form ion pairs in water: (a) \(\mathrm{NaCl}\) or \(\mathrm{Na}_{2} \mathrm{SO}_{4},\) (b) \(\mathrm{MgCl}_{2}\) or \(\mathrm{MgSO}_{4},\) (c) \(\mathrm{LiBr}\) or \(\mathrm{KBr}\).

Short answer

(a) \(\mathrm{Na}_{2} \mathrm{SO}_{4}\), (b) \(\mathrm{MgSO}_{4}\), (c) \(\mathrm{LiBr}\).

Step-by-step solution

Analyze Sodium Compounds

In pair (a), compare \(\mathrm{NaCl}\) and \(\mathrm{Na}_{2} \mathrm{SO}_{4}\). Ions with higher charges and larger ionic radii tend to form more ion pairs. \(\mathrm{NaCl}\) consists of \(\mathrm{Na}^+\) and \(\mathrm{Cl}^-\), both with a charge of ±1. \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) contains \(\mathrm{Na}^+\) and \(\mathrm{SO}_{4}^{2-}\), where \(\mathrm{SO}_{4}^{2-}\) has a higher charge. Hence, \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) is more likely to form ion pairs in water due to the higher charge of the \(\mathrm{SO}_{4}^{2-}\) ion which increases interionic interactions.

Compare Magnesium Compounds

In pair (b), evaluate \(\mathrm{MgCl}_{2}\) versus \(\mathrm{MgSO}_{4}\). \(\mathrm{MgCl}_{2}\) involves \(\mathrm{Mg}^{2+}\) and \(\mathrm{Cl}^-\) ions, while \(\mathrm{MgSO}_{4}\) contains \(\mathrm{Mg}^{2+}\) and \(\mathrm{SO}_{4}^{2-}\). Both \(\mathrm{Mg}^{2+}\) and \(\mathrm{SO}_{4}^{2-}\) have higher charges compared to \(\mathrm{Cl}^-\), facilitating stronger electrostatic attraction in \(\mathrm{MgSO}_{4}\). Therefore, \(\mathrm{MgSO}_{4}\) is more likely to form ion pairs in water.

Evaluate Lithium and Potassium Bromides

For pair (c), compare \(\mathrm{LiBr}\) and \(\mathrm{KBr}\).Both compounds contain bromide \(\mathrm{Br}^-\), but they differ in the cation, having \(\mathrm{Li}^+\) and \(\mathrm{K}^+\) respectively. \(\mathrm{Li}^+\) ions have a smaller ionic radius compared to \(\mathrm{K}^+\), leading to stronger ion-ion interactions in \(\mathrm{LiBr}\). Hence, \(\mathrm{LiBr}\) is more likely to form ion pairs in water.

Key concepts

Ionic Compounds

Ionic compounds are fascinating chemical substances, made up of ions, that are held together by strong electrostatic forces. These ions are typically formed when atoms gain or lose electrons, leading to positive or negative charges. This process is most commonly seen between metals, which lose electrons and become positively charged cations, and non-metals, which gain electrons and become negatively charged anions. The interaction results in the formation of a solid lattice structure.

Examples of common ionic compounds include sodium chloride (\( \mathrm{NaCl} \)) and magnesium sulfate (\( \mathrm{MgSO}_4 \)). These compounds are characterized by their high melting and boiling points. This is a direct result of the strong attractions between the ions in their crystal lattices. However, when dissolved in water or melted, they readily conduct electricity due to the mobility of the ions. This sets them apart from molecular compounds, which typically do not conduct electricity when dissolved.

Water Solubility

Water, often dubbed the "universal solvent," is especially adept at dissolving ionic compounds. This is due to its polar nature—having a partial positive charge near the hydrogen atoms and a partial negative charge near the oxygen atom. When ionic compounds encounter water, the polar water molecules interact with the charged ions.

For example, in a sodium chloride solution, the water molecules surround the \( \mathrm{Na}^+ \) and \( \mathrm{Cl}^- \) ions, effectively breaking them apart and dispersing them throughout the solution. This process is known as hydration, which significantly contributes to the solubility of ionic compounds in water. It is important to note that not all ionic compounds are equally soluble in water. Factors such as ionic size and charge play a crucial role. Bigger or more heavily charged ions might not dissolve as readily, as they form stronger ion pairs.

Electrostatic Interactions

Electrostatic interactions are the forces that hold the ions together in an ionic compound. These forces are the result of attractions between oppositely charged ions. The strength of these interactions is dependent on two main factors: the magnitude of the charge on the ions and the distance between them.

Stronger interactions occur when the ion charges are larger. For instance, the attraction between \( \mathrm{Mg}^{2+} \) and \( \mathrm{SO}_4^{2-} \) in \( \mathrm{MgSO}_4 \) is stronger than that between \( \mathrm{Na}^+ \) and \( \mathrm{Cl}^- \) in \( \mathrm{NaCl} \), because of the higher charges. The smaller the distance, the stronger the interaction, which is why compounds with smaller ionic radii might have stronger electrostatic interactions. These interactions not only influence the formation of ion pairs but also affect properties like melting points and solubility.

Ionic Radius

The ionic radius refers to the size of an ion in a crystal lattice. It plays a significant role in determining how ions interact in a compound. Smaller ions can get closer together, strengthening the electrostatic forces between them and affecting the overall properties of the compound.

For example, in the case of \( \mathrm{Li}^+ \) versus \( \mathrm{K}^+ \), the \( \mathrm{Li}^+ \) ion is smaller, which means it can get closer to the \( \mathrm{Br}^- \) ion in \( \mathrm{LiBr} \), resulting in stronger interactions compared to \( \mathrm{KBr} \). This enhanced interaction often leads to increased ion pair formation. The radius also influences solubility and the stability of the ionic compound structure. When comparing compounds with similarly charged ions, the one composed of smaller ions typically exhibits stronger interactions.