Question

Determine whether each compound is soluble or insoluble. If the compound is soluble, list the ions present in solution. a. \(\mathrm{AgNO}_{3}\) b. \(\mathrm{Pb}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2}\) c. \(\mathrm{KNO}_{3}\) d. \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{~S}\)

Short answer

\(\mathrm{AgNO}_{3}\), \(\mathrm{Pb}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2}\), \(\mathrm{KNO}_{3}\), and \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}\) are all soluble. Ions present in solution: \(\mathrm{Ag}^{+}\), \(\mathrm{NO}_{3}^{-}\); \(\mathrm{Pb}^{2+}\), \(2\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}\); \(\mathrm{K}^{+}\), \(\mathrm{NO}_{3}^{-}\); \(2\mathrm{NH}_{4}^{+}\), \(\mathrm{S}^{2-}\).

Step-by-step solution

Refer to solubility rules

The solubility of a compound in water is determined by using a set of guidelines commonly known as solubility rules. These rules help to predict whether a compound will dissolve (be soluble) in water or form a precipitate (be insoluble).

Evaluate the solubility of \(\mathrm{AgNO}_{3}\)

According to solubility rules, all nitrates (\mathrm{NO}_{3}^{-}) are soluble in water. Therefore, \(\mathrm{AgNO}_{3}\) is soluble, and will dissociate into \(\mathrm{Ag}^{+}\) and \(\mathrm{NO}_{3}^{-}\) ions.

Evaluate the solubility of \(\mathrm{Pb}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2}\)

Lead (\(\mathrm{Pb}^{2+}\)) acetate is typically soluble because acetates (\(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}\)) are generally soluble in water. Thus, \(\mathrm{Pb}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2}\) is soluble and will dissociate into \(\mathrm{Pb}^{2+}\) and \(2\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}\) ions.

Evaluate the solubility of \(\mathrm{KNO}_{3}\)

Potassium nitrate is soluble because both potassium (\(\mathrm{K}^{+}\)) and nitrate (\(\mathrm{NO}_{3}^{-}\)) are on the list of ions that form soluble compounds regardless of the other ion. It dissociates into \(\mathrm{K}^{+}\) and \(\mathrm{NO}_{3}^{-}\) ions.

Evaluate the solubility of \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}\)

Ammonium sulfide is soluble in water because ammonium (\(\mathrm{NH}_{4}^{+}\)) salts are generally soluble. It will dissociate into \(2\mathrm{NH}_{4}^{+}\) and \(\mathrm{S}^{2-}\) ions.

Key concepts

Chemical Solubility

Chemical solubility refers to the property of a substance—typically a solid—to dissolve in a solvent, creating a homogenous solution. In the context of a chemistry course or while doing homework, students often encounter the need to determine if a given ionic compound is soluble in water. This is a central concept not only for laboratory work but also for understanding reactions that occur in biological systems, industrial processes, and the environment.

To determine if an ionic compound is soluble, scientists use established solubility rules. These rules are based on empirical observations and patterns. For example, nitrates and acetates are generally soluble, while carbonates and phosphates are typically insoluble unless paired with certain cations. To improve student understanding, it is beneficial to provide examples and engage in exercises that apply these rules, just as the problem-solving exercise demonstrates.

Understanding solubility is not just about memorizing the rules—it's about recognizing patterns and predicting the behavior of compounds in different solutions. By practicing with various compounds, like those in the given textbook exercise, students reinforce their ability to predict which compounds dissolve in water and which form
a precipitate.

Ionic Compounds

Ionic compounds consist of positive and negative ions held together by ionic bonds, which are electrical attractions between oppositely charged particles. The properties of ionic compounds, such as their high melting and boiling points, are derived from these strong bonds. When dissolved in water, many ionic compounds dissociate into their constituent ions, influencing their solubility.

The solubility of ionic compounds in water can vary widely. To determine whether a substance will dissolve, students can consult solubility rules. These rules consider the ions that make up the compound. For instance, silver nitrate \(\mathrm{AgNO}_{3}\) is an ionic compound that dissociates completely into silver \(\mathrm{Ag}^{+}\) and nitrate \(\mathrm{NO}_{3}^{-}\) ions in solution, making it highly soluble. It's critical for students to not only memorize the solubility of common ions but also to apply this knowledge practically, as was the case in working through the textbook exercise.

Precipitation Reactions

Precipitation reactions occur when two soluble ionic compounds in solution mix and form an insoluble product, known as the precipitate. This process is the result of the formation of an insoluble ionic compound from the combination of the cations and anions of the reacting compounds.

Understanding precipitation reactions is essential in chemistry for predicting the outcome of mixing different solutions. This is crucial for everything from synthesizing desired products to treating waste water. The solubility rules help us predict which compounds will form a precipitate. As shown in the step-by-step solutions in the textbook, no precipitate forms because the ionic compounds in the exercise are all soluble. To aid comprehension, explaining this process through visualization, like writing out the complete ionic equations, can be particularly helpful. Furthermore, experimenting with actual solutions in a lab setting can dramatically improve a student's grasp of the concepts involved in precipitation reactions.