Problem 69 - Chemistry and Chemical Reactivity

Question

In each of the following cases, decide whether a precipitate will form when mixing the indicated reagents, and write a balanced equation for the reaction. (a) \(\mathrm{NaBr}(\mathrm{aq})+\mathrm{AgNO}_{3}(\mathrm{aq})\) (b) \(\mathrm{KCl}(\mathrm{aq})+\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\)

Step-by-Step Solution

Answer

(a) Yes, AgBr precipitates. (b) Yes, PbCl₂ precipitates.

Step 1: Identify the potential reactions

When two aqueous solutions are mixed, possible reactions include the exchange of ions between the compounds. In the given cases: (a) Sodium bromide (NaBr) and silver nitrate (AgNO₃) can potentially react to form silver bromide (AgBr) and sodium nitrate (NaNO₃). (b) Potassium chloride (KCl) and lead(II) nitrate (Pb(NO₃)₂) can potentially form lead(II) chloride (PbCl₂) and potassium nitrate (KNO₃).

Step 2: Check solubility rules

According to the solubility rules: - Silver bromide (AgBr) is insoluble in water. - Lead(II) chloride (PbCl₂) is also insoluble in water. In both reactions, these compounds will form precipitates.

Step 3: Write the balanced chemical equations

For each reaction, write the balanced equation: (a) NaBr(aq) + AgNO₃(aq) → AgBr(s) + NaNO₃(aq) This shows AgBr as the precipitate. (b) 2KCl(aq) + Pb(NO₃)₂(aq) → PbCl₂(s) + 2KNO₃(aq) This shows PbCl₂ as the precipitate.

Step 4: Determine if a precipitate is formed

Based on the balanced equations and solubility rules, both reactions will form a precipitate: (a) AgBr will precipitate. (b) PbCl₂ will precipitate.

Key Concepts

Solubility Rules

Understanding solubility rules is crucial in predicting whether a precipitation reaction will occur when two solutions are mixed. Solubility rules are a set of guidelines that help determine the solubility of various compounds in water. Here's how they work:

• Most nitrate (\( ext{NO}_3^-\)) salts are soluble in water. That means they dissolve without forming a solid precipitate.
• Sodium (\( ext{Na}^+\)), potassium (\( ext{K}^+\)), and ammonium (\( ext{NH}_4^+\)) compounds are almost always soluble.
• Compounds containing silver (\( ext{Ag}^+\)), lead (\( ext{Pb}^{2+}\)), and mercury (\( ext{Hg}_2^{2+}\)) are generally insoluble, with a few exceptions.

In the given exercise, knowing that silver bromide (\( ext{AgBr}\)) is insoluble and will form a solid helps us predict that a precipitation reaction will occur when mixing \( ext{NaBr(aq)}\) and \( ext{AgNO}_3(aq)\). Similarly, lead(II) chloride (\( ext{PbCl}_2\)) is also insoluble, indicating a solid will form in the reaction between \( ext{KCl(aq)}\) and \( ext{Pb(NO}_3)_2(aq)\). These rules are your fundamental tool for anticipating the outcomes of reactions in aqueous solutions.

Balanced Chemical Equations

Writing balanced chemical equations is essential to accurately representing what happens in a chemical reaction. Balancing an equation ensures that the same number of each type of atom appears in both the reactants and the products. Here are the basic steps to achieve this:

• Write the chemical formulas for all reactants and products involved in the reaction.
• Count the number of each type of atom on both sides of the equation.
• Adjust the coefficients (the numbers before each compound) to achieve the same number of atoms of each element on each side.

For instance, if we consider the reaction between \( ext{NaBr(aq)}\) and \( ext{AgNO}_3(aq)}\), we write:\[\text{NaBr(aq) + AgNO}_3\text{(aq) → AgBr(s) + NaNO}_3\text{(aq)}\]Each side of the equation has one sodium (\( ext{Na}\)), one bromine (\( ext{Br}\)), one silver (\( ext{Ag}\)), and one nitrate (\( ext{NO}_3\)) ion, ensuring balance. It's a consistent reflection of mass and charge, demonstrating the conservation of mass in reactions.

Insoluble Compounds

Insoluble compounds are those that do not dissolve well in water. When mixed, these compounds often form a solid, known as a precipitate. Understanding which compounds are insoluble is key to predicting precipitation reactions.Key characteristics of insoluble compounds:

• Typically, insoluble compounds involve ions that form strong ionic bonds, making them less eager to separate into individual ions in water.
• As seen with lead(II) chloride (\( ext{PbCl}_2\)) and silver bromide (\( ext{AgBr}\)), the presence of lead (\( ext{Pb}^{2+}\)) or silver (\( ext{Ag}^+\)) generally implies low solubility, hence the tendency to precipitate.

In the exercise, when \( ext{NaBr(aq)}\) was mixed with \( ext{AgNO}_3(aq)}\), \( ext{AgBr}\) formed as a faint yellow precipitate because it is insoluble. Similarly, mixing \( ext{KCl(aq)}\) with \( ext{Pb(NO}_3)_2(aq)\) resulted in solid white \( ext{PbCl}_2\). Recognizing and predicting the formation of these insoluble compounds is crucial to understanding and navigating various chemical reactions.