bikarbonaattihiilidioksidijärjestelmä

Bicarbonate, also known as hydrogen carbonate, refers to the bicarbonate ion HCO3− and its salts with metals such as sodium, potassium, and calcium. It is the conjugate base of carbonic acid (H2CO3) and participates in the equilibrium CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3−, linking carbon dioxide in the air and water to pH balance in solutions.

In nature, bicarbonate appears in minerals and dissolved in natural waters. The dissolution of carbonate rocks in carbon dioxide-rich water forms calcium bicarbonate, contributing to the buffering capacity of rivers and springs. This buffering helps stabilize pH against fluctuations from respiration, photosynthesis, and other processes.

Bicarbonates have widespread practical uses. In cooking, sodium bicarbonate (baking soda) functions as a leavening agent, releasing carbon dioxide when heated or when reacting with acids. In food processing and home cleaning, bicarbonates act as mild alkalies, deodorizers, and gentle abrasives. Medical applications include bicarbonate salts used as antacids and, in some clinical settings, to correct certain acid–base disorders. In personal care, bicarbonates are common in toothpastes and deodorants for their mild abrasive and odor-neutralizing properties.

Physiologically, the bicarbonate buffer system helps regulate blood pH. The kidneys and lungs work with bicarbonate and carbon dioxide to maintain a narrow pH range (about 7.35 to 7.45) essential for enzyme function and cellular processes. CO2 produced by metabolism is transported in the blood partly as dissolved bicarbonate.

Safety and environmental considerations are generally favorable for consumer uses. Concentrated bicarbonate solutions can irritate skin or eyes, and large-scale releases may alter water chemistry and aquatic life, requiring standard wastewater treatment and monitoring. Overall, bicarbonates are widely regarded as safe when used appropriately.