Problem 64 - Chemistry The Central Science

Question

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

Step-by-Step Solution

Answer

Reacting MnS with Strong Acid: \(MnS + 2H^+ \rightarrow Mn^{2+} + H_2S\)

Step 1: Identify given information

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

Step 2: Determine what to find

We identify the unknown quantity in this Chemistry problem and select the appropriate approach.

Step 3: Select the appropriate formula

Based on the given data and unknown, we choose the relevant equation from Chemistry.

Step 4: Substitute values and solve

Substituting known values into the equation, we perform the calculation.

Step 5: State the result

The result is: Reacting MnS with Strong Acid: \(MnS + 2H^+ \rightarrow Mn^{2+} + H_2S\)

Key Concepts

Slightly Soluble Salts

Slightly soluble salts are compounds that only dissolve to a small extent in water. These salts have low solubility, meaning that only a small amount of the salt can dissolve in the solution at any given time. In practical terms, when these salts are added to water, they tend to form a solid precipitate rather than dissolving completely. Examples include substances like calcium fluoride (\( \text{CaF}_2 \) ) and magnesium carbonate (\( \text{MgCO}_3 \)).Understanding slightly soluble salts is crucial because their solubility can be affected by different factors, such as changes in pH or the presence of other ions in the solution. Chemists often deal with these salts to predict and analyze the output of various chemical reactions in aqueous (water-based) environments. When dealing with chemical equations, slightly soluble salts commonly appear on the reactant side, as they are generally present in the system as solid forms.

Strong Acid Reactions

Strong acids play a significant role in chemical reactions, particularly when dealing with ionic compounds. Known for their complete ionization in water, strong acids like hydrochloric acid (\( \text{HCl} \)) and sulfuric acid (\( \text{H}_2\text{SO}_4 \)) are often used to facilitate reactions with slightly soluble salts.During a reaction with a strong acid, the donation of protons (\( \text{H}^+ \)) can change the solubility of the salt and cause some ions previously locked in solid form to dissolve.For example, in the case of \( \text{MnS} \), the addition of a strong acid such as hydrochloric acid results in an ionic reaction where \( \text{MnS} \) reacts with the \( \text{H}^+ \) ions to form \( \text{Mn}^{2+} \) and hydrogen sulfide gas (\( \text{H}_2\text{S} \)).Thus, strong acids are valuable tools in manipulating the solubility and reactivity of slightly soluble salts in various chemical evaluations.

MnS Reaction

The reaction between manganese sulfide (\( \text{MnS} \)) and strong acids is a classic example of how slightly soluble salts interact in solution to form new products.In this reaction, \( \text{MnS} \), which is only slightly soluble in water, reacts with strong acids to form manganese ions (\( \text{Mn}^{2+} \)) and hydrogen sulfide gas (\( \text{H}_2\text{S} \)).The net ionic equation is as follows:\[ \text{MnS} (s) + 2\text{H}^+ (aq) \rightarrow \text{Mn}^{2+} (aq) + \text{H}_2\text{S} (g) \]This equation signifies that the strong acid supplies the \( \text{H}^+ \) ions needed to dissolve the otherwise slightly soluble \( \text{MnS} \), resulting in the formation of sulfur gas and soluble manganese ions.Such reactions are vital in the realm of chemistry for extracting or dissolving specific components from mixtures.

Solubility Rules

Solubility rules are guidelines that help predict the solubility of various ionic compounds in water. These rules assist chemists in determining which salts will dissolve and which will form precipitates.Here are a few basic solubility rules to consider:

• Most nitrate (\( \text{NO}_3^- \)) salts are soluble in water.
• Salts containing alkali metal ions and ammonium (\( \text{NH}_4^+ \)) are generally soluble.
• Chloride (\( \text{Cl}^- \)), bromide (\( \text{Br}^- \)), and iodide (\( \text{I}^- \)) salts are soluble, except those of silver, lead (\( \text{Pb}^{2+} \)), and mercury (\( \text{Hg}_2^{2+} \)).
• Sulfate (\( \text{SO}_4^{2-} \)) salts are soluble, with exceptions like barium sulfate (\( \text{BaSO}_4 \)) and calcium sulfate (\( \text{CaSO}_4 \)).
• Most sulfide (\( \text{S}^{2-} \)), carbonate (\( \text{CO}_3^{2-} \)), and phosphate (\( \text{PO}_4^{3-} \)) salts are insoluble, except those containing alkali metals and \( \text{NH}_4^+ \).

Applying these solubility rules aids in anticipating the outcomes of reactions involving ionic compounds and in understanding which products will dissolve or precipitate in water.