Bicarbonate

In inorganic chemistry, bicarbonate (IUPAC-recommended nomenclature: hydrogencarbonate) is an intermediate form in the deprotonation of carbonic acid. It is a polyatomic anion with the chemical formula HCO3-. Bicarbonate serves a crucial biochemical role in the physiological pH buffering system. The term "bicarbonate" was coined in 1814 by the English chemist William Hyde Wollaston. The name lives on as a trivial name. The bicarbonate ion (hydrogencarbonate ion) is an anion with the empirical formula HCO3- and a molecular mass of 61.01 daltons; it consists of one central carbon atom surrounded by three oxygen atoms in a trigonal planar arrangement, with a hydrogen atom attached to one of the oxygens. It is isoelectronic with nitric acid HNO3. The bicarbonate ion carries a negative one formal charge and is an amphiprotic species which has both acidic and basic properties. It is both the conjugate base of carbonic acid H2CO3; and the conjugate acid of CO32−, the carbonate ion, as shown by these equilibrium reactions: CO32− + 2 H2O HCO3- + H2O + OH− H2CO3 + 2 OH− H2CO3 + 2 H2O HCO3- + H3O+ + H2O CO32− + 2 H3O+. A bicarbonate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound. Many bicarbonates are soluble in water at standard temperature and pressure; in particular, sodium bicarbonate contributes to total dissolved solids, a common parameter for assessing water quality. Bicarbonate (HCO3-) is a vital component of the pH buffering system of the human body (maintaining acid–base homeostasis). 70%–75% of CO2 in the body is converted into carbonic acid (H2CO3), which is the conjugate acid of HCO3- and can quickly turn into it. With carbonic acid as the central intermediate species, bicarbonate – in conjunction with water, hydrogen ions, and carbon dioxide – forms this buffering system, which is maintained at the volatile equilibrium required to provide prompt resistance to pH changes in both the acidic and basic directions.