Calcium carbonate (\(\text{CaCO}_3\)) is one of the most abundant substances on Earth, forming the basis of rocks like limestone, marble, and chalk. It is also the main component of seashells, eggshells, and coral skeletons. The most direct answer is that it is largely insoluble in pure water under normal conditions. While a very small amount does dissolve, the compound is classified as poorly soluble, meaning the solid material remains mostly intact when mixed with water.
Why Calcium Carbonate Does Not Dissolve Easily
Calcium carbonate’s low solubility results from its strong internal structure, which requires significant energy to break apart. This compound is an ionic salt, held together by strong electrostatic forces between positively charged calcium ions (\(\text{Ca}^{2+}\)) and negatively charged carbonate ions (\(\text{CO}_3^{2-}\)). These forces create a tightly packed, stable crystal lattice structure.
For a substance to dissolve, the energy released when water molecules hydrate the ions must be greater than the energy required to break the crystal lattice. In calcium carbonate, the lattice energy is too high. The polar nature of water molecules is not strong enough to effectively pull the ions away from each other and into the solution.
Only a negligible amount of the compound dissociates. At room temperature, only about 0.013 grams of calcium carbonate will dissolve in a liter of pure water. This limited dissociation explains why limestone and marble are durable materials used in construction.
How Acidity Changes the Equation
Although calcium carbonate resists dissolving in pure water, its solubility increases dramatically when the water becomes even slightly acidic. This behavior explains why the compound dissolves in the natural world, such as in the formation of caves and sinkholes. The presence of hydrogen ions (\(\text{H}^+\)) from an acid changes the chemical environment, upsetting the natural equilibrium of the dissolution reaction.
When acid is introduced, the hydrogen ions react directly with the carbonate ion (\(\text{CO}_3^{2-}\)). This reaction first forms bicarbonate (\(\text{HCO}_3^-\)), and then quickly forms carbonic acid (\(\text{H}_2\text{CO}_3\)). Carbonic acid is unstable and decomposes into water (\(\text{H}_2\text{O}\)) and carbon dioxide gas (\(\text{CO}_2\)).
By reacting with the carbonate ions, the acid continuously removes them from the solution, preventing them from re-forming the solid calcium carbonate. This removal forces more of the solid \(\text{CaCO}_3\) to dissolve to replace the lost carbonate ions and re-establish the chemical balance. The calcium carbonate is converted into highly soluble calcium salts, allowing it to dissolve much more readily than in neutral water.
In nature, water is made acidic primarily by dissolved carbon dioxide from the atmosphere. Rainwater absorbs \(\text{CO}_2\), which forms a weak solution of carbonic acid. Even this slight acidity (pH usually around 5.6) is enough over long periods to drive the dissolution of massive amounts of calcium carbonate rock. The result is the formation of soluble calcium bicarbonate (\(\text{Ca}(\text{HCO}_3)_2\)).
Calcium Carbonate in Nature and Daily Life
The unique solubility behavior of calcium carbonate is responsible for a variety of geological formations and common household issues. The slow, acid-driven dissolution process creates karst topography, characterized by underground caves, caverns, and sinkholes. Water that has absorbed carbon dioxide seeps through limestone rock, slowly dissolving the stone over thousands of years.
This process also contributes to hard water, which is common in areas with limestone bedrock. As water dissolves the calcium carbonate, it carries the dissolved calcium ions (\(\text{Ca}^{2+}\)) into the water supply. When this hard water is heated, such as in a kettle or water heater, the conditions reverse, causing the dissolved calcium to precipitate back out as solid calcium carbonate. This white, chalky deposit is known as limescale or boiler scale.
Calcium carbonate is also involved in biology and medicine. It is the main structural component used by many organisms, such as oysters, clams, and corals, to build their protective shells and skeletons. In medicine, its ability to react with acid is utilized in antacid tablets, where it neutralizes excess stomach acid to relieve heartburn and indigestion. It is also a widely used dietary supplement to provide the body with calcium.