Question

Predict whether aqueous solutions of the following compounds are acidic, basic, or neutral: (a) \(\mathrm{NH}_{4} \mathrm{Br}\), (b) \(\mathrm{FeCl}_{3},(\mathrm{c})\) \(\mathrm{Na}_{2} \mathrm{CO}_{3},\) (d) \(\mathrm{KClO}_{4},\) (e) \(\mathrm{NaHC}_{2} \mathrm{O}_{4}\)

Short answer

(a) \(\mathrm{NH}_{4} \mathrm{Br}\) is acidic as it increases the concentration of hydronium ions (H₃O⁺) when dissolved in water. (b) \(\mathrm{FeCl}_{3}\) is acidic as ferric ions react with water molecules to release hydronium ions. (c) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is basic as carbonate ions react with water to form hydroxide ions. (d) \(\mathrm{KClO}_{4}\) is neutral as neither potassium nor perchlorate ions participate in significant equilibrium reactions with water. (e) \(\mathrm{NaHC}_{2} \mathrm{O}_{4}\) is nearly neutral as hydrogen tartrate ions have a weak acidic character that doesn't significantly influence the solution's acidity or basicity.

Step-by-step solution

(a) Predicting the nature of \(\mathrm{NH}_{4} \mathrm{Br}\) solution

When ammonia (\(\mathrm{NH}_{4}\)), a weak base, is mixed with hydrobromic acid (\(\mathrm{HBr}\)), a strong acid, it forms the salt \(\mathrm{NH}_{4}\mathrm{Br}\). When this compound is dissolved in water, it dissociates into ammonium (\(\mathrm{NH}_{4}^{+}\)) and bromide (\(\mathrm{Br}^-\)) ions. The ammonium ion can further react with water to give hydronium ions and ammonia. This will increase the hydronium ion concentration, making the solution acidic. \[NH_{4}^{+}(aq) + H_2O(l) \rightleftharpoons NH_{3}(aq) + H_{3}O^{+}(aq)\]

(b) Predicting the nature of \(\mathrm{FeCl}_{3}\) solution

When ferric chloride (\(\mathrm{FeCl}_{3}\)) is dissolved in water, it will dissociate into ferric (\(\mathrm{Fe}^{3+}\)) and chloride ions (\(\mathrm{Cl}^-\)). The \(\mathrm{Fe}^{3+}\) ions will react with the water molecules to form hydrated ions, releasing hydronium ions (H₃O⁺), which will increase the H₃O⁺ ion concentration, making the solution acidic. \[Fe^{3+}(aq) + 6H_2O(l) \rightleftharpoons [Fe(H_2O)_{6}]^{3+}(aq) \rightleftharpoons [Fe(H_2O)_{5}(OH)]^{2+}(aq) + H_{3}O^{+}(aq)\]

(c) Predicting the nature of \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) solution

When sodium carbonate (\(\mathrm{Na}_{2}\mathrm{CO}_{3}\)) is dissolved in water, it will dissociate into two sodium ions (\(\mathrm{Na}^{+}\)) and one carbonate ion (\(\mathrm{CO}_{3}^{2-}\)). The sodium ions have no significant interaction with water molecules, but the carbonate ions will react with water molecules to form bicarbonate (\(\mathrm{HCO}_{3}^-\)) and hydroxide (\(\mathrm{OH}^-\)) ions. \[CO_{3}^{2-}(aq) + H_2O(l) \rightleftharpoons HCO_{3}^{-}(aq) + OH^{-}(aq)\] The presence of hydroxide ions makes the solution basic.

(d) Predicting the nature of \(\mathrm{KClO}_{4}\) solution

When potassium perchlorate (\(\mathrm{KClO}_{4}\)) is dissolved in water, it dissociates into potassium (\(\mathrm{K}^{+}\)) and perchlorate (\(\mathrm{ClO}_{4}^-\)) ions. Both potassium and perchlorate ions are weak conjugate ions and don't participate in significant equilibrium reactions with water. Therefore, their concentrations won't affect the H₃O⁺ or OH⁻ concentrations, and the solution will remain neutral.

(e) Predicting the nature of \(\mathrm{NaHC}_{2}\mathrm{O}_{4}\) solution

Sodium hydrogen tartrate (\(\mathrm{NaHC}_{2} \mathrm{O}_{4}\)) is the sodium salt of the weak acid, tartaric acid and can be considered the product of a neutralization reaction between a weak acid and a strong base. When dissolved in water, it dissociates into sodium (\(\mathrm{Na}^{+}\)) and hydrogen tartrate (\(\mathrm{HC}_{2}\mathrm{O}_4^-\)) ions. The sodium ions do not significantly interact with water, but hydrogen tartrate ions can participate in an equilibrium with water molecules: \[HC_{2}O_4^-(aq) + H_2O(l) \rightleftharpoons C_{2}O_{4}^{2-}(aq) + H_3O^{+}(aq)\] However, hydrogen tartrate ions have a weak acidic character and don't significantly influence the acidity or basicity of the solution. Therefore, the solution can be considered nearly neutral. In conclusion, (a) \(\mathrm{NH}_{4} \mathrm{Br}\) is acidic, (b) \(\mathrm{FeCl}_{3}\) is acidic, (c) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is basic, (d) \(\mathrm{KClO}_{4}\) is neutral, and (e) \(\mathrm{NaHC}_{2} \mathrm{O}_{4}\) is nearly neutral.

Key concepts

Acid-Base Reactions

Acid-base reactions are pivotal in understanding the pH of solutions. An acid-base reaction involves the transfer of protons (\( H^+ \)) from an acid to a base. This interaction forms different products, influencing the pH of a solution. For example, when ammonia (\( NH_4^+ \)) from ammonium bromide (\( NH_4Br \)) reacts with water, it forms hydronium ions (\( H_3O^+ \)), making the solution acidic. Key points to remember:

• Acids release hydrogen ions in solution, increasing hydronium ion concentration.
• Bases accept hydrogen ions, often increasing the concentration of hydroxide ions (\( OH^- \)).
• Strong acids like \( HBr \) can lower the pH more significantly than weak acids.

Understanding these reactions is crucial for predicting whether a solution will be acidic, basic, or neutral.

Ion Dissociation

Ion dissociation is the process where ionic compounds separate into ions when dissolved in water. This process is fundamental to understanding solution chemistry and the resultant pH of the solution. When sodium carbonate (\( Na_2CO_3 \)) dissolves, it dissociates into sodium (\( Na^+ \)) and carbonate (\( CO_3^{2-} \)) ions. The carbonate ions interact with water to form bicarbonate and hydroxide ions. This increase in hydroxide ions makes the solution basic. Notable points:

• Ionic compounds dissociate depending on their solubility in water.
• The dissociation process affects the concentration of other reactive species in the solution.
• Dissociation can result in solutions that are either acidic or basic, depending on the ions formed.

Knowing how different ionic compounds dissociate helps predict the behavior of the solution.

Equilibrium Reactions

Equilibrium reactions occur when chemical reactions proceed in both directions at the same rate. In the context of pH, equilibrium reactions often involve weak acids or bases interacting with water. For instance, the dissociation of ammonium ions (\( NH_4^+ \)) in water reaches an equilibrium between ammonia (\( NH_3 \)), ammonium, hydronium (\( H_3O^+ \)) ions. Key aspects:

• Equilibrium reactions balance the reactants and products, affecting pH levels.
• They result in dynamic constancy, where concentrations remain steady over time.
• Shifting equilibrium can change the solution's acidity or basicity.

Grasping these concepts is vital for analyzing how weak acids and bases adjust the pH in aqueous solutions.

Chemical Compounds

Chemical compounds are substances formed from two or more elements. Their behavior in solution relates directly to their ionic and molecular nature, influencing the pH and reactivity of the solution.Understanding chemical compounds' components helps predict solution behavior. For instance, the salt \( FeCl_3 \) in a solution will dissociate to form ferric ions (\( Fe^{3+} \)) and chloride ions (\( Cl^- \)). The ferric ions can interact with water, lowering the pH and making the solution more acidic.Key considerations:

• The identity and strength of component ions affect how they influence pH.
• Knowing the reaction tendencies of each ion provides insight into the compound's impact.
• Strong acids and bases formed from dissociation result in more significant pH changes.

Studying these helps in making educated predictions about solution characteristics.

Solution Chemistry

Solution chemistry deals with how solutes dissolve in solvents and the resulting interactions affecting the solution's properties. Factors include the concentration and nature of solutes and their interactions with the solvent. In aqueous solutions, water acts as the solvent, and the nature of dissolved ions determines solution properties like pH. Potassium perchlorate (\( KClO_4 \)) illustrates a neutral solution because both \( K^+ \) and \( ClO_4^- \) do not significantly interact with water to shift the balance of \( H_3O^+ \) or \( OH^- \) ions.Considerations:

• Solution behavior depends on solute-solvent interactions.
• Equilibrium of ion concentrations influences the resultant pH.
• Understanding solubility and reactivity of solutes aids in predicting solution outcomes.

Gaining a grasp on how substances dissolve and interact helps explain why certain solutions remain neutral, acidic, or basic in nature.