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| Names | |
|---|---|
| IUPAC name Hydrogencarbonate | |
| Systematic IUPAC name Hydroxidodioxidocarbonate(1−)[1] | |
Other names
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| Identifiers | |
| |
3D model (JSmol) | |
| 3903504 | |
| ChEBI | |
| ChEMBL | |
| ChemSpider |
|
| 49249 | |
| KEGG |
|
| UNII | |
| |
| |
| Properties | |
| HCO−3 | |
| Molar mass | 61.0168 g mol−1 |
| logP | −0.82 |
| Acidity (pKa) | 10.3 |
| Basicity (pKb) | 7.7 |
| Conjugate acid | Carbonic acid |
| Conjugate base | Carbonate |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Ininorganic chemistry,bicarbonate (IUPAC-recommended nomenclature:hydrogencarbonate[2]) is an intermediate form in thedeprotonation ofcarbonic acid. It is apolyatomicanion with the chemical formulaHCO−3.
Bicarbonate serves a crucial biochemical role in the physiologicalpHbuffering system.[3]
The term "bicarbonate" wascoined in 1814 by the English chemistWilliam Hyde Wollaston.[4][5] The name lives on as atrivial name.
The bicarbonate ion (hydrogencarbonate ion) is ananion with theempirical formulaHCO−3 and a molecular mass of 61.01daltons; it consists of one central carbonatom surrounded by three oxygen atoms in atrigonal planar arrangement, with a hydrogen atom attached to one of the oxygens. It isisoelectronic withnitric acid (HNO3). The bicarbonate ion carries a negative oneformal charge and is anamphiprotic species which has both acidic and basic properties. It is both theconjugate base ofcarbonic acid (H2CO3); and theconjugate acid ofCO2−3, thecarbonate ion, as shown by theseequilibrium reactions:
A bicarbonate salt forms when apositively charged ion attaches to the negatively charged oxygen atoms of the ion, forming anionic compound. Many bicarbonates aresoluble inwater atstandard temperature and pressure; in particular, sodium bicarbonate contributes tototal dissolved solids, a common parameter for assessingwater quality.[6]
Bicarbonate is a vital component of thepHbuffering system[3] of the human body (maintainingacid–base homeostasis). 70%–75% of CO2 in the body is converted intocarbonic acid (H2CO3), which is theconjugate acid ofHCO−3 and can quickly turn into it.[7]
With carbonic acid as thecentral intermediatespecies, bicarbonate – in conjunction with water,hydrogen ions, andcarbon dioxide – forms this buffering system, which is maintained at the volatile equilibrium[3] required to provide prompt resistance to pH changes in both the acidic andbasic directions. This is especially important for protectingtissues of thecentral nervous system, where pH changes too far outside of the normal range in either direction could prove disastrous (seeacidosis oralkalosis). Recently it has been also demonstrated that cellular bicarbonate metabolism can be regulated by mTORC1 signaling.[8]
Additionally, bicarbonate plays a key role in the digestive system. It raises the internal pH of the stomach, after highly acidic digestive juices have finished in their digestion of food. Bicarbonate also acts to regulate pH in the small intestine. It is released from thepancreas in response to the hormonesecretin to neutralize the acidicchyme entering theduodenum from the stomach.[9]
Bicarbonate is the dominant form ofdissolved inorganic carbon in sea water,[10] and in most fresh waters. As such it is an important sink in thecarbon cycle.
Some plants likeChara utilize carbonate and produce calcium carbonate (CaCO3) as a result of biological metabolism.[11]
In freshwater ecology, strongphotosynthetic activity by freshwater plants in daylight releases gaseousoxygen into the water and at the same time produces bicarbonate ions. These shift the pH upward until in certain circumstances the degree of alkalinity can become toxic to some organisms or can make other chemical constituents such asammonia toxic. In darkness, when no photosynthesis occurs, respiration processes release carbon dioxide, and no new bicarbonate ions are produced, resulting in a rapid fall in pH.[citation needed]
The flow of bicarbonate ions from rocks weathered by the carbonic acid in rainwater is an important part of thecarbon cycle.
The most common salt of the bicarbonate ion issodium bicarbonate,NaHCO3, which is commonly known asbaking soda. When heated or exposed to anacid such asacetic acid (vinegar), sodium bicarbonate releasescarbon dioxide. This is used as aleavening agent inbaking.[12]
Ammonium bicarbonate is used in the manufacturing of some cookies, crackers, and biscuits.[13]
Indiagnostic medicine, theblood value of bicarbonate is one of several indicators of the state ofacid–base physiology in the body. It is measured, along withchloride,potassium, andsodium, to assesselectrolyte levels in anelectrolyte panel test (which hasCurrent Procedural Terminology, CPT, code 80051).[14]
The parameterstandard bicarbonate concentration (SBCe) is the bicarbonate concentration in the blood at aPaCO2 of 40 mmHg (5.33 kPa), full oxygen saturation and 36 °C.[15]
The most common indicator of water quality is the concentration of total dissolved solids (TDS)
