Which Common Compounds Are Insoluble in Water? Examples and Solubility Tips

Calcium carbonate and barium sulfate are insoluble in water. Their chemical structures have strong ionic bonds and high lattice energies, so water molecules can’t break them apart easily. Ions with higher charges or larger sizes form stronger bonds, which lowers solubility. For example, Ca²⁺ and CO3²⁻ in calcium carbonate create a lattice that resists dissolving.

Some insoluble compounds can dissolve in certain conditions. Silver chloride is usually insoluble but dissolves in ammonia because ammonia forms a complex with silver ions. This reaction reduces free silver ion concentration, making the salt dissolve. Similarly, calcium carbonate dissolves in acid because the carbonate ions react with hydrogen ions.

Periodic table trends can help predict insolubility. Alkaline earth metal sulfates get less soluble down the group—magnesium sulfate is soluble, but barium sulfate isn’t. Silver salts often have low solubility, except for silver nitrate. These trends aren’t perfect, though, because factors like pH or complexing agents can change things.

In labs, insoluble compounds are useful for precipitation reactions to identify ions, like using barium sulfate to detect sulfate ions. They’re also used in gravimetric analysis, where the mass of a precipitate measures a substance’s amount. Some serve as filtration aids or inert supports in experiments due to their stability.

Calcium carbonate and barium sulfate remain insoluble in water because their lattice energy—the energy holding ions together in the crystal—far exceeds the hydration energy released when ions bond with water. Their structures, featuring divalent cations (Ca²⁺, Ba²⁺) and divalent anions (CO₃²⁻, SO₄²⁻), form strong electrostatic attractions that resist dissociation. Silver chloride, typically insoluble in water, dissolves in ammonia due to a complex ion reaction: AgCl + 2NH₃ → Ag(NH₃)₂⁺ + Cl⁻, where ammonia ligands break the lattice.

Solubility trends show Group 2 sulfates become less soluble down the group—MgSO₄ dissolves readily, while BaSO₄ is highly insoluble. Most metal carbonates are insoluble except for Group 1 and ammonium salts. In labs, insoluble compounds like BaSO₄ are vital for gravimetric analysis to quantify ions by mass, and precipitates like AgCl help identify ions in spot tests. Their low solubility also makes them useful for filtering and purifying solutions, underpinning both qualitative and quantitative chemical analysis.

Many ionic compounds like calcium carbonate (CaCO₃) and barium sulfate (BaSO₄) exhibit very low solubility in water due to their chemical structures and the balance of intermolecular forces. The solubility of an ionic compound depends on the relative strengths of the lattice energy (the energy holding the solid together) and the hydration energy (the energy released when ions are surrounded by water molecules). In compounds like CaCO₃ and BaSO₄, the lattice energy is exceptionally high because of the strong electrostatic attractions between the divalent cations (Ca²⁺ and Ba²⁺) and the large, multiply charged anions (CO₃²⁻ and SO₄²⁻). This high lattice energy makes it difficult for water molecules to overcome these forces and dissolve the compound, resulting in very low solubility.

The chemical structure plays a critical role in determining solubility. The size and charge of the ions are particularly important. Smaller ions with higher charges tend to form compounds with higher lattice energies, making them less soluble. For example, barium sulfate has a very low solubility product constant (Ksp) of about 1.1 × 10⁻¹⁰, meaning only a tiny amount dissolves in water. Similarly, calcium carbonate has a Ksp of around 3.4 × 10⁻⁹, indicating its limited solubility. The large size of the carbonate and sulfate ions also contributes to their low solubility because they cannot be effectively hydrated by water molecules, reducing the hydration energy that would otherwise help dissolve the compound.

Exceptions to this general trend do exist. For instance, silver chloride (AgCl) is typically insoluble in water but dissolves in ammonia due to the formation of a complex ion, Ag(NH₃)₂⁺. This demonstrates how solubility can be influenced by the presence of complexing agents that alter the equilibrium of the dissolution process. Similarly, calcium carbonate can dissolve in acidic solutions because the acid reacts with the carbonate ion to form carbon dioxide and water, effectively removing one of the products of the dissolution equilibrium and driving the reaction forward.

Periodic table trends can provide some predictive power regarding solubility. Generally, compounds formed from elements in the lower right of the periodic table (like barium and sulfate) tend to be less soluble due to the high lattice energies associated with larger, highly charged ions. Conversely, compounds formed from elements in the upper left (like sodium and chloride) are typically more soluble because of their lower lattice energies and higher hydration energies.

Compounds such as calcium carbonate (CaCO3) and barium sulfate (BaSO4) typically remain insoluble in water primarily because of their high lattice energies. Lattice energy refers to the energy required to separate one mole of a solid ionic compound into its gaseous ions. In the case of CaCO3 and BaSO4, the strong electrostatic forces between their cations (calcium and barium) and anions (carbonate and sulfate) are extremely robust. These forces require more energy to break apart than what is released during the process of ion hydration, which is when water molecules surround and interact with the ions. As a result, the compounds have limited dissociation in water, staying mostly in their solid form. For example, BaSO4’s low solubility makes it highly valuable in medical imaging as a barium meal. When ingested, it doesn't dissolve in the digestive fluids of the stomach and intestines, allowing doctors to clearly visualize the gastrointestinal tract using X - rays.

The chemical structure of a compound plays a crucial role in determining its solubility. This is influenced by both lattice energy and ion hydration. Smaller, more highly charged ions tend to form compounds with higher lattice energies, and this often results in reduced solubility. Take aluminum compounds as an example; aluminum ions (Al3+) are small and highly charged, and many aluminum - containing compounds have relatively low solubility in water.

However, there are exceptions to the general rule of insolubility. Silver chloride (AgCl), which is insoluble in water, can dissolve in ammonia. This occurs due to complex formation. When ammonia is present, the reaction AgCl + 2NH3 → Ag(NH3)2+ + Cl– takes place. The formation of the Ag(NH3)2+ complex ion lowers the free energy of the system, promoting the solubility of AgCl.

In laboratory settings, insoluble compounds have numerous practical applications. Calcium carbonate (CaCO3) is commonly used to neutralize acids. When an acid is spilled or needs to be made less acidic, CaCO3 can react with the acid in a neutralization reaction, producing a salt, water, and carbon dioxide. Barium sulfate (BaSO4) is employed in gravimetric analysis, a method used to determine the amount of a particular substance in a sample by precipitating it out and weighing the precipitate. Silver chloride (AgCl) is useful in precipitation titrations, which are used to determine the concentration of a particular ion in a solution.