bicarbonat

Bicarbonat refers to salts containing the bicarbonate (hydrogen carbonate) anion, HCO3−. The most familiar example is sodium bicarbonate (NaHCO3), commonly known as baking soda. Other common bicarbonates include potassium bicarbonate (KHCO3) and ammonium bicarbonate (NH4HCO3). In chemistry, bicarbonate is the conjugate base of carbonic acid (H2CO3) and participates in the carbonic acid–bicarbonate system in water.

In aqueous solution, bicarbonate acts as part of the buffering system for carbon dioxide in water. The equilibrium CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3− helps regulate pH in natural waters, with pKa values around 6.3 for H2CO3/HCO3− and around 10.3 for HCO3−/CO3^2−. This buffering capacity is important in biological fluids and environmental contexts.

Occurrence and minerals: Bicarbonate is common in natural waters, where calcium and magnesium bicarbonate contribute to water hardness. The mineral nahcolite is a natural sodium bicarbonate. Soda lakes and evaporite deposits can concentrate bicarbonate salts.

Production and supply: Industrially, bicarbonate can be produced by reacting carbon dioxide with sodium hydroxide to form sodium bicarbonate, which may be crystallized and sold. In the Solvay process, sodium bicarbonate forms as an intermediate toward sodium carbonate (washing soda), and can be isolated under certain conditions.

Uses: Sodium bicarbonate is widely used as a leavening agent in baking, a gentle antacid, a buffering agent in foods and beverages, and a cleaning and deodorizing compound. It is also used in fire suppression powders and as a soil pH adjuster in agriculture.

Safety and environment: Bicarbonates are generally regarded as safe at typical dietary levels, but excessive intake can cause metabolic alkalosis or electrolyte disturbances. In the environment, bicarbonate ions help stabilize pH in water bodies but very high concentrations can affect aquatic life.