When carbon dioxide (\(\text{CO}_2\)) dissolves in water (\(\text{H}_2\text{O}\)), the vast majority of the gas molecules simply remain as dissolved \(\text{CO}_2\) in the aqueous solution. A small but significant fraction, however, reacts chemically with the water molecules. This reaction creates a new compound, the weak acid known as carbonic acid (\(\text{H}_2\text{CO}_3\)). This chemical process is the underlying cause for the change in water’s acidity and is fundamental to global environmental cycles and biological systems.
The Formation of Carbonic Acid
The initial chemical reaction involves dissolved carbon dioxide combining directly with a water molecule to yield carbonic acid. This process is represented by the reversible chemical equation \(\text{CO}_2 + \text{H}_2\text{O} \leftrightarrow \text{H}_2\text{CO}_3\). The double-arrow symbol indicates that this reaction can proceed in both the forward direction, forming carbonic acid, and the reverse direction, where the acid breaks back down into its original components.
Carbonic acid is characterized as an unstable weak acid. Due to its instability, the equilibrium strongly favors the reactants, meaning that the concentration of carbonic acid molecules is much lower than the concentration of dissolved \(\text{CO}_2\). The rate at which the carbonic acid breaks down is significantly faster than the rate at which it forms in the absence of a catalyst. Despite its low concentration, this compound acts as the precursor for all subsequent changes in the water’s chemistry.
The Subsequent Chemical Equilibrium
Following its formation, the carbonic acid molecule begins a sequential process of dissociation. The first step involves the carbonic acid losing a single hydrogen ion (\(\text{H}^+\)) to form the bicarbonate ion (\(\text{HCO}_3^-\)). This release of \(\text{H}^+\) ions lowers the water’s pH, making the solution more acidic.
The bicarbonate ion can then undergo a second, less frequent dissociation step, where it releases another hydrogen ion to form the carbonate ion (\(\text{CO}_3^{2-}\)). The entire system exists in a dynamic chemical equilibrium involving dissolved \(\text{CO}_2\), carbonic acid, bicarbonate ions, and carbonate ions. The relative concentrations of these species are highly dependent on the water’s pH.
In natural waters, such as the oceans, the vast majority of the dissolved carbon exists in the form of bicarbonate ions, accounting for approximately 92% of the dissolved inorganic carbon. The remainder is split between dissolved \(\text{CO}_2\) and a smaller percentage of carbonate ions. This balance is maintained by the rapid exchange of hydrogen ions between the different forms, acting as a natural buffer system to resist drastic changes in acidity.
Consequences in Natural and Engineered Systems
The chemistry of carbon dioxide dissolving in water has consequences across various natural and human-engineered systems. In the context of the global ocean, the absorption of atmospheric \(\text{CO}_2\) creates a phenomenon known as ocean acidification. As the concentration of dissolved \(\text{CO}_2\) increases, the elevated levels of hydrogen ions bind with existing carbonate ions, reducing their availability.
This reduction in carbonate ions directly impacts calcifying organisms, such as corals, oysters, and pteropods, which rely on carbonate to form calcium carbonate structures for their shells and skeletons. The increased acidity makes it harder for these organisms to build their protective layers and can also cause existing shells to dissolve.
This reversible reaction is also the fundamental principle behind the carbonation of beverages like soda. The drink is bottled under high pressure, which, according to Le Chatelier’s principle, forces the equilibrium to shift toward the formation of dissolved \(\text{CO}_2\) and carbonic acid. When the container is opened, the sudden release of pressure causes the equilibrium to shift rapidly in the opposite direction. This results in the rapid conversion of dissolved \(\text{CO}_2\) back into gas, creating the visible bubbles and the characteristic “fizz” as the gas escapes the liquid.
In the human body, this same chemical system forms the bicarbonate buffer that maintains the blood’s physiological pH within a narrow range, typically between 7.35 and 7.45. The rapid conversion between \(\text{CO}_2\) and bicarbonate is accelerated by the enzyme carbonic anhydrase, which is present in red blood cells. This enzyme speeds up the reaction by a factor of over a million, quickly converting metabolic \(\text{CO}_2\) into bicarbonate for transport and managing the concentration of free hydrogen ions to prevent dangerous shifts in blood acidity.