When a gas dissolves into a liquid, the interaction can be purely physical or a chemical process that fundamentally changes the liquid’s composition. Carbon dioxide ($\text{CO}_2$) is significantly soluble in water ($\text{H}_2\text{O}$), and its dissolution involves a rapid and reversible chemical transformation. This interaction is central to global biogeochemical cycles and has profound implications for the chemistry of natural water bodies, including the world’s oceans. Understanding the specific chemical species that form is the foundation for grasping the resulting change in water acidity.
The Formation of Carbonic Acid
The moment a carbon dioxide molecule enters the water, it begins a chemical reaction with a water molecule to form carbonic acid ($\text{H}_2\text{CO}_3$). This reaction is represented by the equation $\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3$. The double-headed arrow signifies that the reaction is reversible and exists in a state of chemical equilibrium. Carbonic acid is the primary product formed, but it is considered a weak acid because only a small fraction of the dissolved $\text{CO}_2$ actually converts into this compound. The majority of the dissolved gas remains as unreacted aqueous $\text{CO}_2$ molecules. Under typical conditions, carbonic acid is unstable and decomposes quickly.
The Subsequent Chemical Dissociation
The newly formed carbonic acid then undergoes a process called dissociation, which occurs in two sequential stages and is the source of the solution’s increased acidity. In the first stage, the carbonic acid molecule loses a single hydrogen ion ($\text{H}^+$), yielding a bicarbonate ion ($\text{HCO}_3^-$). This reaction is written as $\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-$. The release of this positively charged hydrogen ion into the water alters the solution’s chemistry.
The bicarbonate ion formed in the first stage can then dissociate further in a second, less frequent stage, losing another hydrogen ion to form a carbonate ion ($\text{CO}_3^{2-}$). The three resulting carbon-containing species—carbonic acid ($\text{H}_2\text{CO}_3$), bicarbonate ($\text{HCO}_3^-$), and carbonate ($\text{CO}_3^{2-}$)—exist in a dynamic chemical equilibrium. The relative concentrations of these three forms are controlled by the solution’s overall acidity, or $\text{pH}$.
Measuring the Resulting Acidity
The concentration of the free hydrogen ions ($\text{H}^+$) released during the dissociation of carbonic acid is the standard measure of the water’s acidity. This is quantified using the $\text{pH}$ scale, which is an inverse, logarithmic measure of the hydrogen ion concentration. A lower $\text{pH}$ value corresponds to a higher concentration of hydrogen ions and, consequently, a more acidic solution.
When carbon dioxide dissolves and forms carbonic acid, the resulting increase in hydrogen ions causes the water’s $\text{pH}$ to drop, making the solution more acidic. This phenomenon is easily observed in everyday life with carbonated beverages, such as sparkling water or soda. These drinks are produced by dissolving $\text{CO}_2$ under high pressure, and the slightly tart or biting taste is the direct result of the mild acidity created by the formation of carbonic acid.
Environmental and Biological Significance
The process of $\text{CO}_2$ dissolving in water is not limited to laboratory settings or beverage production; it is a fundamental process driving large-scale environmental changes. The world’s oceans absorb approximately 30% of the carbon dioxide released into the atmosphere from human activities. This massive uptake of $\text{CO}_2$ leads to a shift in the marine carbon chemistry, resulting in the phenomenon known as ocean acidification.
Since the start of the Industrial Revolution, the average $\text{pH}$ of the ocean’s surface waters has fallen by approximately 0.1 $\text{pH}$ units, which represents about a 30% increase in acidity. This change has significant consequences for many marine organisms, particularly those that build shells and skeletons out of calcium carbonate ($\text{CaCO}_3$), such as corals, oysters, and mussels. The increase in hydrogen ions from the dissolved $\text{CO}_2$ not only lowers the $\text{pH}$ but also reacts with the available carbonate ions ($\text{CO}_3^{2-}$), effectively reducing the concentration of the building blocks these organisms need. A reduced concentration of carbonate ions makes it more energetically difficult for calcifying organisms to build and maintain their protective structures, threatening the survival of entire marine ecosystems.