4.7: Solubility of Ionic Compounds

Solubility is the measure of the maximum amount of solute that can be dissolved in a given quantity of solvent at a given temperature and pressure. Solubility is usually measured in molarity (M) or moles per liter (mol/L). A compound is termed soluble if it dissolves in water.

When soluble salts dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution; this process represents a physical change known as dissociation. Potassium chloride (KCl) is an example of a soluble salt. When solid KCl is added to water, the positive (hydrogen) end of the polar water molecules is attracted to the negative chloride ions, and the negative (oxygen) ends of water are attracted to the positive potassium ions. The water molecules surround individual K⁺ and Cl⁻ ions, reducing the strong forces that bind the ions together and letting them move off into solution as solvated ions.

Another example of a soluble salt is silver nitrate, AgNO₃, which dissolves in water as Ag⁺ and NO₃^- ions. Nitrate, NO₃^-, is a polyatomic ion, and in solution, it stays intact as a single whole unit. Unlike monatomic ions (K⁺, Cl^-, Ag⁺), which contain only one atom, polyatomic ions are a group of atoms that carry a charge (NO₃^-, SO₄^2-, NH₄⁺). They remain such in solution and do not split into individual atoms.

A compound is termed insoluble if it does not dissolve in water. However, in reality, “insoluble” compounds dissolve to some extent, that is, less than 0.01 M.

In the case of insoluble salts, the strong interionic forces that bind the ions in the solid are stronger than the ion-dipole forces between individual ions and water molecules. As a result, the ions stay intact and do not separate. Thus, most of the compound remains undissolved in water. Silver chloride (AgCl) is an example of an insoluble salt. The water molecules cannot overcome the strong interionic forces that bind the Ag⁺ and Cl^- ions together; hence, the solid remains undissolved.

Solubility Rules

The solubility of ionic compounds in water depends on the type of ions (cation and anion) that form the compounds. For example, AgNO₃ is water-soluble, but AgCl is water-insoluble. The solubility of a salt can be predicted by following a set of empirical rules (listed below), developed based on the observations on many ionic compounds.

i) Compounds containing ammonium ions (NH₄⁺) and alkali metal cations are soluble
ii) All nitrates and acetates are always soluble.
iii) Chloride, bromide, and iodide compounds are soluble with the exception of those of silver, lead, and mercury(I)
iv) All sulfate salts are soluble except their salts with silver, lead, mercury(I), barium, strontium, and calcium
v) All carbonates, sulfites, and phosphates are insoluble except their salts with ammonium and alkali metal cations.
vi) Sulfides and hydroxides of all salts are insoluble, with the exception of their salts with alkali metal cations, ammonium ion, and calcium, strontium, and barium ions.
vii) All oxide-containing compounds are insoluble except their compounds with calcium, barium, and alkali metal cations.

This text is adapted from OpenStax Chemistry 2e, Section 11.2: Electrolytes.

The solubility of a solute is its maximum possible concentration at solubility equilibrium in a given amount of solvent. Solubility is affected by temperature and other physical conditions.

Substances that dissolve in water are called water-soluble. A simple water-soluble ionic compound like sodium chloride dissolves in water by breaking up into monatomic ions. Here, it is more favorable for the water molecules and ions to interact in solution than it is for the ions to remain in the ordered solid.

A more complex water-soluble ionic compound like sodium nitrate contains ions that are composed of multiple atoms covalently bound together, or polyatomic ions. When sodium nitrate dissolves, the polyatomic nitrate ions do not split into nitrogen and oxygen. Instead, the ions disperse in solution as intact units.

Substances that do not dissolve in water are water-insoluble. For example, silver chloride is a water-insoluble ionic compound. In this case, it is more favorable for the ions to remain in the ordered solid than to interact with water and disperse in the solution.

The solubility of an ionic compound in water depends on the ion pair that makes up the compound. Chemists have formulated a set of empirical guidelines to predict the solubility of ionic compounds in water. Exceptions to these guidelines are rare.

All nitrates and acetates are soluble. Similarly, all ammonium and non-lithium alkali metal compounds are soluble, as are nearly all lithium salts.

Sulfate compounds are soluble, with the exception of its salts with lead, mercury, and silver – remember the acronym LMS or the phrase Let Me See – and calcium, barium, and strontium – remember the acronym CBS or the phrase Come By Soon.

All chloride, bromide, and iodide salts are soluble, with the exception of their salts with LMS – lead, mercury, and silver – as well as copper(I) and thallium.

Moving to insoluble compounds, sulfides and hydroxides are insoluble, with the exception of their salts with alkali metals and barium. In addition, ammonium sulfide is soluble, and strontium hydroxide is soluble when heated.

Similarly, carbonates and phosphates are insoluble, with the exception of their ammonium and non-lithium alkali metal salts.