Magnesium carbonate (\(\text{MgCO}_3\)) is an inorganic salt that appears as a white, odorless powder, often found naturally in the mineral magnesite. This compound is composed of a magnesium cation (\(\text{Mg}^{2+}\)) and a carbonate anion (\(\text{CO}_3^{2-}\)), held together by strong ionic bonds. \(\text{MgCO}_3\) is widely used in various industrial and medical applications.
Defining Solubility: The Direct Answer for \(\text{MgCO}_3\)
Magnesium carbonate is classified as sparingly soluble or practically insoluble in pure water at room temperature. The term “sparingly soluble” means that only a very small amount of the compound will dissolve to form a solution. Its solubility is measured at approximately \(0.01\) to \(0.02\) grams per \(100\) milliliters of water. For most practical purposes, magnesium carbonate is treated as if it does not dissolve.
Chemists quantify this low solubility using the solubility product constant (\(\text{K}_{\text{sp}}\)), which for \(\text{MgCO}_3\) is about \(3.5 \times 10^{-8}\) at \(25^\circ\text{C}\). This small value indicates that the equilibrium between the solid compound and its dissolved ions heavily favors the undissolved solid. The resulting molar solubility is only about \(1.87 \times 10^{-4}\) moles per liter.
Why \(\text{MgCO}_3\) Resists Dissolving in Pure Water
Magnesium carbonate resists dissolving due to the strength of its internal structure. Solid \(\text{MgCO}_3\) is held together by a crystal lattice formed by the electrostatic attraction between the \(\text{Mg}^{2+}\) and \(\text{CO}_3^{2-}\) ions. The energy required to break this lattice, known as the lattice energy, is very high.
For the compound to dissolve, the energy released when water molecules surround and stabilize the separated ions (hydration energy) must be sufficient to overcome the high lattice energy. In the case of magnesium carbonate, the energy gain from hydrating the ions is not enough to compensate for the significant energy cost of breaking the strong ionic bonds. This energetic imbalance ensures that the vast majority of the \(\text{MgCO}_3\) remains in its solid, undissolved form in neutral water.
The Crucial Exception: Solubility in Acidic Water
The solubility of magnesium carbonate changes significantly in the presence of even weak acids. When \(\text{MgCO}_3\) encounters an acidic environment, a chemical reaction occurs that drives the dissolution process forward. The carbonate ion (\(\text{CO}_3^{2-}\)) reacts readily with hydrogen ions (\(\text{H}^{+}\)) from the acid.
This reaction converts the poorly soluble carbonate ion into the much more soluble bicarbonate ion (\(\text{HCO}_3^{-}\)), often releasing carbon dioxide gas in the process. The removal of carbonate ions from the solution shifts the dissolution equilibrium, causing more solid \(\text{MgCO}_3\) to dissolve to replace the consumed ions. This acidic reaction is why magnesium carbonate is listed as soluble in dilute acids. Natural water that contains dissolved carbon dioxide forms carbonic acid, which increases \(\text{MgCO}_3\) solubility compared to \(\text{CO}_2\)-free water.
Practical Consequences of Low Solubility
The solubility properties of magnesium carbonate have several real-world implications. In medicine, \(\text{MgCO}_3\) is widely used as a non-systemic antacid because it reacts with and neutralizes highly acidic stomach acid. Its low solubility prevents significant absorption into the bloodstream. The reaction releases carbon dioxide, which can contribute to the effervescence often experienced with antacids.
Geologically, the solubility of \(\text{MgCO}_3\) plays a part in the formation of hard water and karst landscapes. Magnesium ions dissolved from the rock contribute to water hardness, along with calcium ions. The high solubility of magnesium bicarbonate formed by acidic groundwater dissolving the mineral is responsible for the erosion of carbonate rock structures and the creation of caves over geological time.