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In inorganic chemistry, bicarbonate (IUPAC-recommended nomenclature: hydrogencarbonate) is an intermediate form in the deprotonation of carbonic acid.
Bicarbonate serves a crucial biochemical role in the physiological pH buffering system.[1]
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Chemical properties
The bicarbonate ion (hydrogen carbonate) 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. The bicarbonate ion carries a negative one formal charge and is the conjugate base of carbonic acid, H2CO3; it is 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, particularly sodium bicarbonate and magnesium bicarbonate; both of these substances contribute to total dissolved solids, a common parameter for assessing water quality.
Biochemical role
Bicarbonate is an alkaline, and a vital component of the pH buffering system[1] of the body (maintaining acid-base homeostasis). 86%-90% of CO2 in the body is converted into carbonic acid (H2CO3), which can quickly turn into bicarbonate (HCO3−).
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[1] required to provide prompt resistance to drastic pH changes in both the acidic and basic directions. This is especially important for protecting tissues of the central nervous system, where pH changes too far outside of the normal range in either direction could prove disastrous. (See acidosis, or alkalosis.)
Bicarbonate also acts to regulate pH in the small intestine. It is released from the pancreas in response to the hormone secretin to neutralize the acid chyme entering the duodenum from the stomach [2]
Other uses
The most common salt of the bicarbonate ion is sodium bicarbonate, NaHCO3, which is used as baking soda. When exposed to an acid such as acetic acid (vinegar), sodium bicarbonate releases carbon dioxide. This is used as a leavening agent in baking.
The flow of bicarbonate ions from rocks weathered by the carbonic acid in rainwater is an important part of the carbon cycle.
Bicarbonate also serves in the digestive system. It raises the internal pH of the stomach, after highly acidic digestive juices have finished in their digestion of food. Ammonium bicarbonate is used in digestive biscuit manufacture.
Diagnostics
In diagnostic medicine, the blood value of bicarbonate is one of several indicators of the state of acid-base physiology in the body.
The parameter Standard bicarbonate concentration (SBCe) is the bicarbonate concentration in the blood at a CO2 of 5.33kPa, full oxygen saturation and 37 degrees Celsius.[3]
Bicarbonate compounds
References
- ^ a b c http://www.biology.arizona.edu/biochemistry/problem_sets/medph/intro.html
Biology.arizona.edu - October 2006. Clinical correlates of pH levels: bicarbonate as a buffer. - ^ Berne & Levy, Principles of Physiology
- ^ Acid Base Balance (page 3)
See also
External links
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- This page was last modified on 6 August 2008, at 20:17.
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