Problem 71 - 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

Precipitates form in both cases: AgBr and PbCl₂ are the precipitates.

Step 1: Determine ion combinations

When mixing solutions, the ions from each compound dissolve in water and float around. For each pair of reactants, identify the possible cation-anion combinations that may form precipitates. These combinations are: (a) NaBr (aq) + AgNO₃ (aq): Na⁺, Br⁻, Ag⁺, NO₃⁻ which combine to form NaNO₃ and AgBr. (b) KCl (aq) + Pb(NO₃)₂ (aq): K⁺, Cl⁻, Pb²⁺, NO₃⁻ which combine to form KNO₃ and PbCl₂.

Step 2: Apply solubility rules

Use solubility rules to check if any of the combinations result in a precipitate: (a) According to solubility rules, AgBr is insoluble in water while NaNO₃ is soluble. Thus, AgBr will form as a precipitate. (b) PbCl₂ is insoluble in water, as per solubility rules, while KNO₃ is soluble. Hence, PbCl₂ will form as a precipitate.

Step 3: Write balanced chemical equations

For the reactions that form a precipitate, write down the balanced chemical equation: (a) NaBr (aq) + AgNO₃ (aq) → AgBr (s) + NaNO₃ (aq) (b) 2KCl (aq) + Pb(NO₃)₂ (aq) → PbCl₂ (s) + 2KNO₃ (aq)

Key Concepts

Solubility Rules

Understanding solubility rules is essential when predicting the formation of precipitates in chemical reactions. These rules are guidelines that help us know whether an ionic compound will dissolve in water.
Key Points about Solubility Rules:

• Compounds containing sodium (Na⁺), potassium (K⁺), and nitrate (NO₃⁻) ions are generally soluble in water.
• Most chlorides (Cl⁻), bromides (Br⁻), and iodides (I⁻) are soluble, except when paired with silver (Ag⁺), lead (Pb²⁺), or mercury (Hg₂²⁺).
• Sulfates (SO₄²⁻) are usually soluble, with exceptions like calcium sulfate (CaSO₄), barium sulfate (BaSO₄), and lead sulfate (PbSO₄).
• Carbonates (CO₃²⁻), phosphates (PO₄³⁻), and sulfides (S²⁻) are typically insoluble, except when with alkali metals (like Na⁺, K⁺) or ammonium (NH₄⁺).

These rules provide a quick reference to determine if a compound forms a soluble solution or a solid, known as a precipitate, during reactions. Knowing solubility rules helps predict the outcome of mixing various ionic solutions.

Chemical Equations

Chemical equations are expressions that describe the transformation of reactants into products in a chemical reaction. They provide a concise way to communicate chemical changes.
Start recognizing reactants and products on either side of the equation:

• Reactants are substances you begin with, on the left side of the equation.
• Products are the substances formed, listed on the right side.

A balanced chemical equation adheres to the law of conservation of mass, meaning the number of atoms for each element must be the same on both sides.
Example:

• For AgBr formation: \(\text{NaBr (aq) + AgNO}_3 \text{(aq) → AgBr (s) + NaNO}_3 \text{(aq)}\)
• For PbCl₂ formation: \(2\text{KCl (aq) + Pb(NO}_3\text{)}_2\text{(aq) → PbCl}_2\text{(s) + 2KNO}_3\text{(aq)}\)

Each equation shows reactants converted to products, with states (solid \((s)\), liquid \((l)\), aqueous \((aq)\), gas \((g)\)) in parentheses to indicate the physical state of each compound.

Cation-Anion Combinations

Cation-anion combinations occur when positively charged ions (cations) and negatively charged ions (anions) associate to form compounds. In aqueous solutions, ions dissociate and can recombine when different solutions are mixed.

• Cations within the exercise include Na⁺, K⁺, Ag⁺, and Pb²⁺.
• Anions include Br⁻, Cl⁻, and NO₃⁻.

These ions in separate solutions can mix and form new combinations. Such mixing can lead to products like insoluble compounds, precipitating out of the solution.
In the examples given, the cation of one reactant meets the anion of the other, resulting in either soluble or insoluble compounds. For instance:

• Na⁺ + NO₃⁻ form soluble NaNO₃.
• Ag⁺ + Br⁻ form the insoluble compound AgBr (precipitate).
• K⁺ + NO₃⁻ form soluble KNO₃.
• Pb²⁺ + Cl⁻ form the insoluble compound PbCl₂ (precipitate).

This process shows why knowing ion combinations is critical for predicting the formation of a precipitation.

Insoluble Compounds

Insoluble compounds are those that do not dissolve significantly in water, resulting in the formation of a solid precipitate when solutions combine during a reaction. Understanding why certain compounds are insoluble helps in predicting the outcome of an ionic reaction.
Substances with certain ion pairs, like silver bromide (AgBr) or lead chloride (PbCl₂), are known to be insoluble. Upon interaction:

• AgBr is formed when Ag⁺ and Br⁻ combine, precipitating as a solid because it does not dissolve in water.
• PbCl₂ forms under similar circumstances from Pb²⁺ and Cl⁻, appearing as a solid precipitate.

Not all ionic compounds will dissolve when introduced to water, leading to substances that appear as precipitates, which can be visible as cloudy forms in the mixture. Recognizing which compounds are likely to be insoluble allows for the prediction of a reaction’s visual outcome and the identification of resultant products.