Problem 58 - Chemistry: The Central Science

Question

For each of the following slightly soluble salts, write the net ionic equation, if any, for reaction with acid: (a) MnS, (b) \(\mathrm{PbF}_{2}\), (c) \(\mathrm{AuCl}_{3}\) (d) \(\mathrm{Hg}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), (e) \(\mathrm{CuBr}\).

Step-by-Step Solution

Answer

The net ionic equations for the reaction of slightly soluble salts with acid are: a) \( MnS(s) + 2H⁺(aq) \rightarrow Mn²⁺(aq) + H₂S(g) \) b) \( PbF₂(s) + 2Cl⁻(aq) \rightarrow PbCl₂(s) + 2F⁻(aq) \) c) No Reaction d) \( Hg₂C₂O₄(s) + 2H⁺(aq) \rightarrow Hg₂Cl₂(s) + 2HOC₂O₂(aq) \) e) No Reaction

Step 1: Identify the acid and the ions present in the salt

Generally, acids that are commonly used are hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). For this exercise, we will use hydrochloric acid (HCl) as the acid. Let's identify the ions present in each salt: a) MnS - Mn²⁺ (Manganese ion), S²⁻ (Sulfide ion) b) PbF₂ - Pb²⁺ (Lead ion), F⁻ (Fluoride ion) c) AuCl₃ - Au³⁺ (Gold ion), Cl⁻ (Chloride ion) d) Hg₂C₂O₄ - Hg₂²⁺ (Mercury ion), C₂O₄²⁻ (Oxalate ion) e) CuBr - Cu⁺ (Copper ion), Br⁻ (Bromide ion)

Step 2: Predict the possible products of the reaction

In this step, we will predict the possible products of the reaction between the salt and hydrochloric acid (HCl). Usually, slightly soluble salts will form a solid (precipitate) in the presence of an acid. a) MnS + 2HCl → MnCl₂ + H₂S b) PbF₂ + 2HCl → PbCl₂ + 2HF c) AuCl₃ + 3HCl → No Reaction (AuCl₃ is soluble in HCl and forms chloroauric acid, HAuCl₄) d) Hg₂C₂O₄ + 2HCl → Hg₂Cl₂ + 2HOC₂O₂ e) CuBr + HCl → No Reaction (CuBr is not soluble in HCl)

Step 3: Write the balanced chemical equations

Now, we will write the balanced chemical equations for the reactions that took place: a) MnS(s) + 2HCl(aq) → MnCl₂(aq) + H₂S(g) b) PbF₂(s) + 2HCl(aq) → PbCl₂(s) + 2HF(aq) d) Hg₂C₂O₄(s) + 2HCl(aq) → Hg₂Cl₂(s) + 2HOC₂O₂(aq)

Step 4: Write the net ionic equations by eliminating spectator ions

Finally, we will write the net ionic equations for the reactions by eliminating the spectator ions: a) MnS(s) + 2H⁺(aq) → Mn²⁺(aq) + H₂S(g) b) PbF₂(s) + 2H⁺(aq) + 2Cl⁻(aq) → PbCl₂(s) + 2F⁻(aq) + 2H⁺(aq) d) Hg₂C₂O₄(s) + 2H⁺(aq) → Hg₂Cl₂(s) + 2HOC₂O₂(aq) Note that in (b), the 2H⁺ ions are present on both sides of the equation. So, they are spectator ions and can be eliminated: b) PbF₂(s) + 2Cl⁻(aq) → PbCl₂(s) + 2F⁻(aq)

Key Concepts

Net Ionic Equations

A net ionic equation is a way to express a chemical reaction that focuses only on the ions that participate directly in the formation of the products. In such equations, spectator ions, which do not change during the reaction, are excluded. This simplifies the equation and highlights the chemistry occurring in the solution.
For instance, in the reaction of MnS with HCl, the full chemical equation is:

• MnS(s) + 2HCl(aq) → MnCl₂(aq) + H₂S(g)

To create a net ionic equation, we remove the spectator ions (in this case, Cl⁻ ions). The net ionic equation, which shows only the reacting species, becomes:

• MnS(s) + 2H⁺(aq) → Mn²⁺(aq) + H₂S(g)

Net ionic equations help clarify the essence of the reaction by excluding parts that don’t directly participate. They are particularly useful in understanding precipitation reactions, redox reactions, and acid-base reactions.

Slightly Soluble Salts

Slightly soluble salts are compounds that do not dissolve well in water, meaning only a small amount will break apart into ions. This "limited solubility" often results in the formation of a solid or precipitate in solutions. Understanding the behavior of these salts is crucial when predicting reactions, especially in aqueous environments.
Taking lead(II) fluoride (PbF₂) as an example, this salt is slightly soluble in water. When it reacts with hydrochloric acid (HCl), it forms lead(II) chloride (PbCl₂), which is less soluble than PbF₂ and precipitates out:

• PbF₂(s) + 2HCl(aq) → PbCl₂(s) + 2HF(aq)

Writing the net ionic equation involves removing spectator ions. Here, ions like H⁺ do not appear in the net equation as they do not contribute directly to the formation of the precipitate. Recognition and understanding of slightly soluble salts help in predicting reactions and their products in various chemical settings.

Reaction with Acids

Reactions between slightly soluble salts and acids are fascinating as they typically involve the exchange of ions that form either new acids, dissolved species, or precipitates. Commonly used acids include hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃), with HCl being most frequently used in educational settings for ease of understanding.
These reactions often dissolve the original salt or alter it to produce a different compound. For instance, Hg₂C₂O₄ (mercury(II) oxalate) reacts with HCl as follows:

• Hg₂C₂O₄(s) + 2HCl(aq) → Hg₂Cl₂(s) + 2HOC₂O₂(aq)

This example illustrates the transformation of a slightly soluble salt into another compound (Hg₂Cl₂), which also precipitates. The ability to recognize and predict such transformations provides deep insights into the behavior of ions in solutions.

Spectator Ions

Spectator ions are ions that do not participate directly in the chemical reaction. They appear on both sides of a chemical equation, unchanged, and their role is essentially to "watch" the reaction unfold, hence the term "spectator."
To illustrate, consider the reaction involving PbF₂ and HCl mentioned earlier. The spectator ions are Cl⁻, present on both sides of the equation:

• PbF₂(s) + 2Cl⁻(aq) → PbCl₂(s) + 2F⁻(aq)

In this reaction, the Cl⁻ ions do not partake in the formation of the precipitate (PbCl₂), which means they are spectators. Their presence in ionic reactions does not affect the product outcome directly. Understanding and identifying these ions allows chemists to simplify complex reactions into more manageable net ionic equations, providing clearer insights into the chemical processes at play.