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Chapter 4: Problem 6

We have seen that ions in aqueous solution are stabilized by the attractions between the ions and the water molecules. Why then do some pairs of ions in solution form precipitates? [Section 4.2]

Short Answer

Expert verified
In aqueous solutions, ions are stabilized by their interactions with water molecules through solvation. However, some pairs of ions in solution form precipitates because their attractive forces with each other are stronger than the stabilizing interactions with water molecules. These precipitates are formed when the solubility equilibrium is reached, and the ion product exceeds the solubility product constant (Ksp) for the substance involved. This indicates that the interactions between the ions overcome the stabilizing effect of water molecules.

Step by step solution

01

Introduction to solvation and stabilization

In aqueous solutions, ions (either positively charged cations or negatively charged anions) are surrounded by water molecules. These water molecules are attracted to the ions due to their polar nature. This interaction is called "solvation" and it stabilizes ions in the solution.
02

Formation of precipitates

Despite the stabilization imparted by water molecules, some pairs of ions in the solution might form precipitates. This is because the ions in the solution can interact with each other as well, forming products with very low solubility. Precipitation happens when the attractive forces between the ions in the pair are stronger than the stabilizing interaction with the water molecules.
03

Examples of precipitation reactions

An example of this interaction causing precipitation is the reaction between a soluble silver salt (e.g., silver nitrate, AgNO3) and a soluble chloride salt (e.g., sodium chloride, NaCl), leading to the formation of silver chloride (AgCl), a highly insoluble precipitate: \[ AgNO_3 + NaCl \rightarrow AgCl_{(s)} + NaNO_3 \] In this case, the strong attraction between the Ag+ and Cl- ions is stronger than the interactions between these ions and the water molecules, which leads to the formation of AgCl precipitate.
04

An understanding of solubility equilibrium

It is important to remember that the solubilities of different compounds vary, and this plays a significant role in the formation of the precipitate. Precipitation reactions are governed by the solubility equilibrium and the solubility product constant (Ksp), which is specific for each substance. If the ion product of a solution exceeds the Ksp value for the substance, the excess ions will form insoluble precipitates. This indicates that the attractive forces between the ions in the solution overcome the stabilizing interactions with water molecules.

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Most popular questions from this chapter

Write balanced molecular and net ionic equations for the reactions of (a) hydrochloric acid with nickel, (b) dilute sulfuric acid with iron, (c) hydrobromic acid with magnesium, (d) acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}\), with zinc.

An 8.65-g sample of an unknown group \(2 \mathrm{~A}\) metal hydroxide is dissolved in \(85.0 \mathrm{~mL}\) of water. An acid-base indicator is added and the resulting solution is titrated with \(2.50 \mathrm{M}\) \(\mathrm{HCl}(a q)\) solution. The indicator changes color signaling that the equivalence point has been reached after \(56.9 \mathrm{~mL}\) of the hydrochloric acid solution has been added. (a) What is the molar mass of the metal hydroxide? (b) What is the identity of the metal cation: \(\mathrm{Ca}^{2+}, \mathrm{Sr}^{2+}, \mathrm{Ba}^{2+}\) ?

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Would you expect that anions would be physically closer to the oxygen or to the hydrogens of water molecules that surround it in solution?
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