Carbon dioxide (\(\text{CO}_2\)) is technically neither a traditional acid nor a base in its pure gaseous state. Its acidic nature only becomes apparent when it interacts with water. When dissolved, \(\text{CO}_2\) behaves as an acidic anhydride, a compound that reacts with water to produce an acid. This interaction forms carbonic acid, a weak acid that plays a significant role in both natural processes and environmental systems.
Defining Acidity
Acidity and basicity are measured using the \(\text{pH}\) scale, a logarithmic range from 0 to 14. A \(\text{pH}\) value of 7 is considered neutral, such as pure water, while anything below 7 is acidic and anything above 7 is basic. Because the scale is logarithmic, a drop of one \(\text{pH}\) unit represents a tenfold increase in acidity.
The Arrhenius definition states that an acid is a substance that produces hydrogen ions (\(\text{H}^+\)) when dissolved in water. The Brønsted-Lowry definition characterizes an acid as any substance capable of donating a proton (\(\text{H}^+\)) to another substance. Carbon dioxide falls into the category of a non-metal oxide that reacts with water to form an acid, which is the definition of an acidic anhydride.
The Chemical Reaction That Forms Carbonic Acid
The acidic nature of carbon dioxide is revealed through its reaction with water. First, dissolved \(\text{CO}_2\) molecules form carbonic acid (\(\text{H}_2\text{CO}_3\)). This initial reaction is represented as \(\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3\).
Carbonic acid ionizes, giving the solution its acidic properties. The carbonic acid molecule acts as an acid by donating one of its protons (\(\text{H}^+\)). This ionization is shown as \(\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-\), producing a hydrogen ion and a bicarbonate ion. The release of the \(\text{H}^+\) proton lowers the \(\text{pH}\) and confirms the substance’s acidic behavior.
Carbonic acid is classified as a weak acid because its ionization is incomplete. The equilibrium strongly favors the reactants, meaning most \(\text{H}_2\text{CO}_3\) molecules remain intact. \(\text{CO}_2\) is termed the acidic anhydride of \(\text{H}_2\text{CO}_3\) because it is the oxide from which the acid is formed.
Natural Systems Where \(\text{CO}_2\) Acidity Matters
The \(\text{CO}_2\)/carbonic acid/bicarbonate system is a buffer mechanism that maintains the \(\text{pH}\) balance in the human bloodstream. This system keeps blood \(\text{pH}\) within the narrow, slightly alkaline range of 7.35 to 7.45. Carbon dioxide, a waste product of cellular respiration, is converted to carbonic acid inside the red blood cells by the enzyme carbonic anhydrase.
The bicarbonate ion (\(\text{HCO}_3^-\)) serves as the base component of the buffer, ready to neutralize any excess acid. Conversely, if the blood becomes too alkaline, carbonic acid releases more \(\text{H}^+\) ions to lower the \(\text{pH}\) back to the safe range. The respiratory system regulates this balance by adjusting the rate of \(\text{CO}_2\) exhalation, while the kidneys control the concentration of bicarbonate ions.
The absorption of atmospheric \(\text{CO}_2\) by the ocean is causing a phenomenon known as ocean acidification. The oceans have absorbed about a quarter of human-caused \(\text{CO}_2\) emissions. This massive influx of \(\text{CO}_2\) shifts the overall equilibrium, increasing the concentration of \(\text{H}^+\) ions in the seawater.
The average \(\text{pH}\) of the ocean surface has dropped from 8.15 to about 8.05. This increasing acidity has consequences for marine organisms that build shells and skeletons out of calcium carbonate, such as corals and shellfish. The excess hydrogen ions bond with carbonate ions, making them less available for these calcifying organisms, which hinders their ability to maintain their structures and threatens entire food chains.