Question

For which salt in each of the following groups will the solubility depend on \(\mathrm{pH}\) ? a. \(\mathrm{AgF}, \mathrm{AgCl}, \mathrm{AgBr}\) c. \(\mathrm{Sr}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{Sr}\left(\mathrm{NO}_{2}\right)_{2}\) b. \(\mathrm{Pb}(\mathrm{OH})_{2}, \mathrm{PbCl}_{2}\) d. \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{Ni}(\mathrm{CN})_{2}\)

Short answer

The solubility depends on pH for the following salts in each group: - Group a: None - Group b: \(\mathrm{Sr}\left(\mathrm{NO}_{2}\right)_{2}\) - Group c: \(\mathrm{Pb}(\mathrm{OH})_{2}\) - Group d: \(\mathrm{Ni}(\mathrm{CN})_{2}\)

Step-by-step solution

Group a: \(\mathrm{AgF}, \mathrm{AgCl}, \mathrm{AgBr}\)

All the silver salts contain halides as anions (F-, Cl-, Br-). None of these ions have significant acid-base properties, therefore the solubility of these salts is not dependant on the pH.

Group b: \(\mathrm{Sr}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{Sr}\left(\mathrm{NO}_{2}\right)_{2}\)

Both strontium salts contain nitrate and nitrite anions. Nitrate (NO\(_3^-\)) is not considered to have significant acid-base properties, while nitrite (NO\(_2^-\)) can react with acids and bases easily. So, the solubility of \(\mathrm{Sr}\left(\mathrm{NO}_{2}\right)_{2}\) will depend on pH.

Group c: \(\mathrm{Pb}(\mathrm{OH})_{2}, \mathrm{PbCl}_{2}\)

The lead salts contain hydroxide (OH-) and chloride (Cl-) ions as anions. Hydroxide ion (OH-) has strong basic properties, and the solubility of \(\mathrm{Pb}(\mathrm{OH})_{2}\) in acidic solutions will be higher as it can react with H+ ions. The solubility of \(\mathrm{Pb}(\mathrm{OH})_{2}\) depends on pH.

Group d: \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{Ni}(\mathrm{CN})_{2}\)

Both nickel salts contain nitrate and cyanide anions. Nitrate (NO\(_3^-\)) ion doesn't have significant acid-base properties, while cyanide (CN-) ion can react with both acids and bases. Therefore, the solubility of \(\mathrm{Ni}(\mathrm{CN})_{2}\) will depend on pH. In conclusion, the salts for which the solubility depends on pH are: - Group a: None - Group b: \(\mathrm{Sr}\left(\mathrm{NO}_{2}\right)_{2}\) - Group c: \(\mathrm{Pb}(\mathrm{OH})_{2}\) - Group d: \(\mathrm{Ni}(\mathrm{CN})_{2}\)

Key concepts

pH and Solubility

The concept of pH is crucial when studying the solubility of salts, which can significantly be influenced by the acidity or basicity of the solution.

Solubility refers to the maximum amount of a substance that can be dissolved in a solvent at a given temperature. In an aqueous solution, the pH level indicates the concentration of hydrogen ions \( \[H^+\] \) present, which can impact the solubility of certain salts. Acidic environments (lower pH) increase the concentration of \( \[H^+\] \) ions, while basic environments (higher pH) decrease it.

Salts that contain anions with acidic or basic properties will have their solubility affected by the pH of the solution. For example, salts with basic anions such as \(OH^-\) will be more soluble in acidic solutions due to the reaction of hydroxide ions with \( \[H^+\] \) ions, forming water. On the other hand, salts with anions like \(CN^-\), which can bind to \( \[H^+\] \) ions, will have their solubility altered in both acidic and basic solutions. This relationship between pH and solubility is essential when predicting the behavior of salts in different environmental conditions.

Acid-Base Properties

Exploring acid-base properties in the context of solubility requires understanding how anions and cations interact with \( \[H^+\] \) and \( \[OH^-\] \) ions in water.

Anions derived from weak acids or bases can undergo protonation or deprotonation reactions in water, depending on the pH of the solution. Considering the problem we discussed, \(NO_2^-\) ions from \( \[Sr\left(\mathrm{NO}_{2}\right)_{2}\] \) can react with \( \[H^+\] \) ions, thus acting as bases, and changing their solubility as the pH varies. Similarly, \(CN^-\) ions can behave both as bases by reacting with \( \[H^+\] \) ions or as acids in the presence of \( \[OH^-\] \) ions, which explains why the solubility of \( \[Ni(\mathrm{CN})_{2}\] \) is pH-dependent.

Enhanced Solubility in Acidic Conditions

The case of \( \[Pb(\mathrm{OH})_{2}\] \) is a good example where solubility increases in acidic solutions. The hydroxide ions from the salt can readily react with the \( \[H^+\] \) ions, leading to a decrease in \(OH^-\) concentration and thereby shifting the equilibrium to dissolve more \( \[Pb(\mathrm{OH})_{2}\] \) to restore the balance.

Solubility Product Principle

The solubility product principle, represented by the symbol \( K_{sp} \) is a useful tool for understanding the solubility behavior of ionic compounds.

\( K_{sp} \) is the equilibrium constant for the solubility of a slightly soluble ionic compound. The principle states that the product of the ions' concentrations, each raised to the power of its coefficient in the balanced equation, equals a constant value at a given temperature. This constant is specific for each solute and is influenced by the presence of other substances in the solution.

For example, the solubility product of \( \[AgCl\] \) is the product of the concentrations of \( \[Ag^+\] \) and \( \[Cl^-\] \) ions in the solution at equilibrium. If other ions are added that form complexes with \( \[Ag^+\] \) or \( \[Cl^-\] \) ions, this changes the \( K_{sp} \) value, illustrating that solubility is affected by the chemical composition of the solvent. Even though \( K_{sp} \) does not directly factor pH into its equation, the solubility it represents can change with pH alterations because \( \[H^+\] \) or \( \[OH^-\] \) can react with the ions that make up the salt, altering the concentrations of these ions, thus affecting the \( K_{sp} \) equilibrium.