Question

You know that an unlabeled bottle contains an aqueous solution of one of the following: \(A g N O_{3},\) CaCl \(_{2},\) or \(A l_{2}\left(S O_{4}\right)_{3}\) . A friend suggests that you test a portion of the solution with \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) and then with NaCl solutions. According to your friend's logic, which of these chemical reactions could occur, thus helping you identify the solution in the bottle? (a) Barium sulfate could precipitate. (b) Silver chloride could precipitate. (c) silver sulfate could precipitate. (d) More than one, but not all, of the reactions described in answers a-c could occur. (e) All three reactions described in answers a-c could occur.

Short answer

(d) More than one, but not all, of the reactions described in answers a-c could occur.

Step-by-step solution

Write the possible reactions

For the given aqueous solutions, let's write the possible reactions when adding \(Ba(NO_{3})_{2}\) and NaCl: 1. Aqueous solution of \(AgNO_{3}\): - With \(Ba(NO_{3})_{2}\): \(AgNO_{3} + Ba(NO_{3})_{2} \rightarrow Ag(NO_{3})_{2} + Ba(NO_{3})\) - With NaCl: \(AgNO_{3} + NaCl \rightarrow AgCl\downarrow + NaNO_{3}\) 2. Aqueous solution of \(CaCl_{2}\): - With \(Ba(NO_{3})_{2}\): \(CaCl_{2} + Ba(NO_{3})_{2} \rightarrow Ca(NO_{3})_{2} + BaCl_{2}\) - With NaCl: \(CaCl_{2} + NaCl \rightarrow CaCl_{2} + NaCl\) 3. Aqueous solution of \(Al_{2}(SO_{4})_{3}\): - With \(Ba(NO_{3})_{2}\): \(Al_{2}(SO_{4})_{3}+ 3Ba(NO_{3})_{2} \rightarrow 2Al(NO_{3})_{3} + 3BaSO_{4}\downarrow\) - With NaCl: \(Al_{2}(SO_{4})_{3}+6NaCl \rightarrow 2AlCl_{3}+ 6Na_{2}SO_{4}\)

Determine the reactions that could occur

Analyzing the reactions from Step 1, we see that the following reactions can occur: - Silver chloride could precipitate (In the case of the aqueous solution of \(AgNO_{3}\) with NaCl) - Barium sulfate could precipitate (In the case of the aqueous solution of \(Al_{2}(SO_{4})_{3}\) with \(Ba(NO_{3})_{2}\))

Identify the correct option

Based on our analysis, we found that: (a) Barium sulfate could precipitate. (True) (b) Silver chloride could precipitate. (True) (c) Silver sulfate could precipitate. (False) Thus, the correct option is: (d) More than one, but not all, of the reactions described in answers a-c could occur.

Key concepts

Precipitation Reactions

A precipitation reaction is a type of chemical reaction that occurs when two aqueous solutions of ionic compounds combine to form an insoluble solid, known as a precipitate. These reactions are key in various fields, from industrial chemistry to environmental science.

For instance, when an aqueous solution of silver nitrate (AgNO₃) is mixed with sodium chloride (NaCl), silver chloride (AgCl) precipitate forms:
AgNO₃ + NaCl → AgCl↓ + NaNO₃

Here, the arrow pointing downward (↓) signifies the formation of a precipitate. Learning to predict and balance precipitation reactions is foundational for chemistry students, as it underpins many real-world applications like water purification and metallurgical processes.

What is essential to understand is that not all combinations of ionic compounds in aqueous solutions result in a precipitate. Whether a reaction forms a precipitate depends on the solubility rules, which are guidelines to predict the solubility of different compounds in water.

Solubility Rules

The solubility rules are a set of guidelines that help us predict the solubility of different ionic compounds in water. These rules say that certain ions tend to form soluble compounds and others tend to form insoluble compounds.

For example:

• Most nitrates (NO₃^-) are soluble.
• Most sulfates (SO₄^2-) are soluble, except those of barium (Ba²⁺), lead (Pb²⁺), mercury (Hg²⁺), and calcium (Ca²⁺) to some extent.
• Most chloride salts are soluble (Cl^-), but salts of silver (Ag⁺), mercury (Hg₂²⁺), and lead (Pb²⁺) are exceptions and are generally insoluble.

Understanding these patterns allows chemists to predict outcomes of reactions, such as precipitation. As seen in the example given in the textbook exercise, silver chloride (AgCl) and barium sulfate (BaSO₄) are predicted to precipitate based on these rules, as they are generally insoluble in water.

Ionic Compounds in Aqueous Solutions

When ionic compounds dissolve in water, they separate into their constituent ions – a process known as dissociation. This is crucial for precipitation reactions, since the ions become 'free' to interact with others in the solution.

For example, when calcium chloride (CaCl₂) is dissolved in water, it dissociates into calcium ions (Ca²⁺) and chloride ions (Cl^-):
CaCl₂ → Ca²⁺ + 2Cl^-

The behavior of these ions in solution can be influenced by several factors, including the concentration of ions and the presence of other compounds. In the classroom scenario, the identification of the original ionic compound relies on observing which precipitation reactions occur upon mixing with other ionic solutions. This is a direct application of solubility rules and understanding the nature of ionic interactions in aqueous solutions.