Question

The solubility of copper(II) hydroxide in water can be increased by adding either the base \(\mathrm{NH}_{3}\) or the acid HNO \(_{3}\). Explain. Would added \(\mathrm{NH}_{3}\) or \(\mathrm{HNO}_{3}\) have the same effect on the solubility of silver acetate or silver chloride? Explain.

Short answer

The solubility of copper(II) hydroxide (\(\mathrm{Cu(OH)_{2}}\)) increases with the addition of either ammonia (\(\mathrm{NH}_{3}\)) or nitric acid (\(\mathrm{HNO}_{3}\)). In the presence of ammonia, a soluble complex (\(\mathrm{[Cu(NH_{3})_{4}]^{2+}}\)) is formed, while the addition of nitric acid reacts with the hydroxide ions to form water, decreasing the concentration of hydroxide ions, thus favoring dissolution according to Le Châtelier's principle. The solubility of silver acetate (\(\mathrm{AgC}_{2}\mathrm{H}_{3}\mathrm{O}_{2}\)) is not significantly affected by the presence of ammonia or nitric acid because it is a relatively soluble compound in water. However, the solubility of silver chloride (\(\mathrm{AgCl}\)) increases in the presence of ammonia due to the formation of a soluble complex (\(\mathrm{[Ag(NH_{3})_{2}]^{+}}\)). In contrast, the addition of nitric acid decreases the solubility of silver chloride as it reacts with chloride ions to form HCl, causing silver chloride to precipitate.

Step-by-step solution

Solubility of Copper(II) Hydroxide in the presence of NH3

The solubility of copper(II) hydroxide, \(\mathrm{Cu(OH)_{2}}\), is relatively low in water. When we add NH3 to the solution, it forms a complex with the copper(II) ions and produces the soluble complex \(\mathrm{[Cu(NH_{3})_{4}^{2+}}]\). The overall reaction taking place is: \[\mathrm{Cu(OH)_{2}(s) + 4NH_{3}(aq) \rightleftharpoons [Cu(NH_{3})_{4}]^{2+}(aq) + 2OH^{-}(aq)}\] The formation of this soluble complex increases the solubility of copper(II) hydroxide.

Solubility of Copper(II) Hydroxide in the presence of HNO3

When we add an acid like HNO3, it reacts with the hydroxide ions to form water. The overall reaction is: \[\mathrm{Cu(OH)_{2}(s) + 2HNO_{3}(aq) \rightleftharpoons Cu(NO_{3 })_{2}(aq) + 2H_{2}O(l)}\] This reaction decreases the concentration of hydroxide ions (OH-) in the solution, thus favoring the dissolution of more copper(II) hydroxide in water according to Le Châtelier's principle.

Solubility of Silver Acetate and Silver Chloride in the presence of NH3

Silver acetate, \(\mathrm{AgC}_{2}\mathrm{H}_{3}\mathrm{O}_{2}\), is a relatively soluble ionic compound in water and both \(\mathrm{Ag^{+}}\) and \(\mathrm{C}_{2}\mathrm{H}_{3}\mathrm{O}_{2}^{-}\) ions reach relatively high concentration in water. That being said, neither NH3 nor HNO3 have a significant effect on its solubility. On the other hand, silver chloride, \(\mathrm{AgCl}\), has a low solubility in water. When NH3 is added to the solution, it forms a soluble complex called \(\mathrm{[Ag(NH_{3})_{2}]^{+}}\). The overall reaction is: \[\mathrm{AgCl(s) + 2NH_{3}(aq) \rightleftharpoons [Ag(NH_{3})_{2}]^{+}(aq) + Cl^{-}(aq)}\] Adding NH3 considerably increases the solubility of silver chloride.

Solubility of Silver Acetate and Silver Chloride in the presence of HNO3

In the case of silver acetate, as previously mentioned, adding HNO3 will have no significant effect on its solubility. For silver chloride, the addition of HNO3 will decrease its solubility. This is because the chloride ions will form a strong acid (HCl) when reacting with HNO3, decreasing the concentration of available chloride ions in the solution. \[\mathrm{HNO}_{3}(aq) + Cl^{-}(aq) \rightleftharpoons HCl(aq) + NO_{3}^{-}(aq)\] As the concentration of chloride ions decreases, silver chloride will precipitate out of the solution.

Key concepts

Le Châtelier's principle

Understanding Le Châtelier's principle is crucial when discussing changes to systems in chemical equilibrium. This principle states that if a system at equilibrium experiences a change, such as in concentration, temperature, or pressure, the system will adjust to counteract the imposed change and restore a new balance.

Consider the dissolution of copper(II) hydroxide in water. If hydroxide ions are removed from the solution, for instance by the addition of HNO₃, the equilibrium shifts to replace them, resulting in more copper(II) hydroxide dissolving. Alternatively, adding NH₃ pulls copper ions out of equilibrium by forming a complex ion, effectively increasing its solubility. Le Châtelier's principle explains how the system compensates for these changes to maintain equilibrium, which is fundamental in predicting how a solution's solubility will adjust to various additions.

Complex Ion Formation

Complex ion formation is a key concept when studying solubility in solutions. A complex ion is formed when central metal ions bond with several molecules or ions, called ligands. These ligands possess at least one pair of electrons available to create a coordinate covalent bond with the metal ion.

For instance, when NH₃ is added to a solution of copper(II) hydroxide, it acts as a ligand, binding with the copper ions to form a complex ion: \[\mathrm{[Cu(NH_{3})_{4}]^{2+}}\]. This formation reduces the concentration of free copper ions, shifting the equilibrium toward copper(II) hydroxide dissolving to replace them, hence increasing its overall solubility in water.

Acid-Base Reactions

Acid-base reactions are pivotal in altering the solubility of substances. An acid, by definition, is a substance that can release a proton (or H⁺ ion), while a base can accept a proton. When an acid reacts with a base, a neutralization reaction occurs, producing water and a salt.

This concept elucidates why adding HNO₃, an acid, to copper(II) hydroxide increases its solubility. It neutralizes the OH^- ions, forming water and reducing their concentration. This reaction tips the scales in the equilibrium, resulting in the dissolution of more solid to maintain the ionic equilibrium. Such acid-base reactions are a cornerstone of solubility-related equilibria in chemistry.

Ionic Equilibrium

Ionic equilibrium is a state where the rate of dissociation of ions is equal to the rate of their recombination. It is a subset of chemical equilibrium that focuses specifically on the balance between ions in a solution.

When discussing solubility, the concepts of ionic equilibrium become significant. For silver chloride (\(\mathrm{AgCl}\)), the addition of NH₃ forms a complex ion with silver, significantly changing the ionic concentration of available Ag⁺ ions. This changes the dynamic of how molecules interact and will lead to a shift in the equilibrium to dissolve more AgCl. However, in the case of silver acetate, the already high ionic concentration means that complex ion formation or acid addition does not have a notable impact. The state of ionic equilibrium guides how reactants and products interact in solution and subsequently their solubility behavior.