Bicarbonate

In inorganic chemistry, bicarbonate (IUPAC-recommended nomenclature: hydrogencarbonate[2]) is an intermediate form in the deprotonation of carbonic acid. It is a polyatomic anion with the chemical formula HCO
3
.

Bicarbonate serves a crucial biochemical role in the physiological pH buffering system.[3]

The term "bicarbonate" was coined in 1814 by the English chemist William Hyde Wollaston.[4] The prefix "bi" in "bicarbonate" comes from an outdated naming system and is based on the observation that there is twice as much carbonate (CO2−
3
) per sodium ion in sodium bicarbonate (NaHCO3) and other bicarbonates than in sodium carbonate (Na2CO3) and other carbonates.[5] The name lives on as a trivial name.

According to the Wikipedia article IUPAC nomenclature of inorganic chemistry, the prefix bi– is a deprecated way of indicating the presence of a single hydrogen ion. The recommended nomenclature today mandates explicit referencing of the presence of the single hydrogen ion: sodium hydrogen carbonate or sodium hydrogencarbonate. A parallel example is sodium bisulfite (NaHSO3).

Bicarbonate
Skeletal formula of bicarbonate with the explicit hydrogen added
Ball and stick model of bicarbonate
Names
Systematic IUPAC name
Hydroxidodioxidocarbonate(1−)[1]
Other names
Hydrogencarbonate[1]
Identifiers
3D model (JSmol)
3DMet
3903504
ChEBI
ChEMBL
ChemSpider
49249
KEGG
UNII
Properties
HCO
3
Molar mass 61.0168 g mol−1
log P −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 their standard state (at 25 °C [77 °F], 100 kPa).

Chemical properties

The bicarbonate ion (hydrogencarbonate ion) is an anion with the empirical formula HCO
3
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 HNO
3
. 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 H
2
CO
3
; and the conjugate acid of CO2−
3
, the carbonate ion, as shown by these equilibrium reactions:

CO2−
3
+ 2 H2O ⇌ HCO
3
+ H2O + OH ⇌ H2CO3 + 2 OH
H2CO3 + 2 H2O ⇌ HCO
3
+ H3O+ + H2O ⇌ CO2−
3
+ 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.

Physiological role

Riassorbimento bicarbonati e respirazione cellulare
CO2 produced as a waste product of the oxidation of sugars in the mitochondria reacts with water in a reaction catalyzed by carbonic anhydrase to form H2CO3, which is in equilibrium with the cation H+ and anion HCO3. It is then carried to the lung, where the reverse reaction occurs and CO2 gas is released. In the kidney (left), cells (green) lining the proximal tubule conserve bicarbonate by transporting it from the glomerular filtrate in the lumen (yellow) of the nephron back into the blood (red). The exact stoichiometry in the kidney is omitted for simplicity.

Bicarbonate (HCO
3
) is a vital component of the pH buffering system[3] 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 HCO
3
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[3] required to provide prompt resistance to 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 serves much 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 the pancreas in response to the hormone secretin to neutralize the acidic chyme entering the duodenum from the stomach.[6]

Bicarbonate in the environment

Bicarbonate is the dominant form of dissolved inorganic carbon in sea water,[7] and in most fresh waters. As such it is an important sink in the carbon cycle.

In freshwater ecology, strong photosynthetic activity by freshwater plants in daylight releases gaseous oxygen 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 as ammonia 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.

Other uses

The most common salt of the bicarbonate ion is sodium bicarbonate, NaHCO3, which is commonly known as baking soda. When heated or 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.

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. It is measured, along with carbon dioxide, chloride, potassium, and sodium, to assess electrolyte levels in an electrolyte panel test (which has Current Procedural Terminology, CPT, code 80051).

The parameter standard bicarbonate concentration (SBCe) is the bicarbonate concentration in the blood at a PaCO2 of 40 mmHg (5.33 kPa), full oxygen saturation and 36 °C.[8]

Reference ranges for blood tests - by molarity
Reference ranges for blood tests, comparing blood content of bicarbonate (shown in blue at right) with other constituents.

Bicarbonate compounds

See also

References

  1. ^ a b "hydrogencarbonate (CHEBI:17544)". Chemical Entities of Biological Interest (ChEBI). UK: European Institute of Bioinformatics. IUPAC Names. Archived from the original on 2015-06-07.
  2. ^ Nomenclature of Inorganic Chemistry IUPAC Recommendations 2005 (PDF), IUPAC, p. 137, archived (PDF) from the original on 2017-05-18
  3. ^ a b c "Clinical correlates of pH levels: bicarbonate as a buffer". Biology.arizona.edu. October 2006. Archived from the original on 2015-05-31.
  4. ^ William Hyde Wollaston (1814) "A synoptic scale of chemical equivalents," Philosophical Transactions of the Royal Society, 104: 1-22. On page 11, Wollaston coins the term "bicarbonate": "The next question that occurs relates to the composition of this crystallized carbonate of potash, which I am induced to call bi-carbonate of potash, for the purpose of marking more decidedly the distinction between this salt and that which is commonly called a subcarbonate, and in order to refer at once to the double dose of carbonic acid contained in it."
  5. ^ "Classroom Resources - Argonne National Laboratory". www.newton.dep.anl.gov. Archived from the original on 26 February 2015. Retrieved 2 May 2018.
  6. ^ Berne & Levy, Principles of Physiology
  7. ^ "The chemistry of ocean acidification : OCB-OA". www.whoi.edu. Woods Hole Oceanographic Institution. 24 September 2012. Archived from the original on 19 May 2017. Retrieved 17 May 2017.
  8. ^ Acid Base Balance (page 3) Archived 2002-06-13 at the Wayback Machine

External links

Acetazolamide

Acetazolamide, sold under the trade name Diamox among others, is a medication used to treat glaucoma, epilepsy, altitude sickness, periodic paralysis, idiopathic intracranial hypertension (raised brain pressure of unclear cause), and heart failure. It may be used long term for the treatment of open angle glaucoma and short term for acute angle closure glaucoma until surgery can be carried out. It is taken by mouth or injection into a vein.Common side effects include numbness, ringing in the ears, loss of appetite, vomiting, and sleepiness. It is not recommended in those with significant kidney problems, liver problems, or who are allergic to sulfonamides. Acetazolamide is in the diuretic and carbonic anhydrase inhibitor families of medication. It works by decreasing the amount of hydrogen ions and bicarbonate in the body.Acetazolamide came into medical use in 1952. It is on the World Health Organization's List of Essential Medicines, which lists the safest and most effective medicines needed in a health system. Acetazolamide is available as a generic medication. The wholesale cost in the developing world is about US$1.40–16.93 per month. In the United States the wholesale cost is about US$125.34 per month.

Acid–base homeostasis

Acid–base homeostasis is the homeostatic regulation of the pH of the body's extracellular fluid (ECF). The proper balance between the acids and bases (i.e. the pH) in the ECF is crucial for the normal physiology of the body, and cellular metabolism. The pH of the intracellular fluid and the extracellular fluid need to be maintained at a constant level.Many extracellular proteins such as the plasma proteins and membrane proteins of the body's cells are very sensitive for their three dimensional structures to the extracellular pH. Stringent mechanisms therefore exist to maintain the pH within very narrow limits. Outside the acceptable range of pH, proteins are denatured (i.e. their 3-D structure is disrupted), causing enzymes and ion channels (among others) to malfunction.

In humans and many other animals, acid–base homeostasis is maintained by multiple mechanisms involved in three lines of defence:

The first line of defence are the various chemical buffers which minimize pH changes that would otherwise occur in their absence. They do not correct pH deviations, but only serve to reduce the extent of the change that would otherwise occur. These buffers include the bicarbonate buffer system, the phosphate buffer system, and the protein buffer system.The second line of defence of the pH of the ECF consists of controlling of the carbonic acid concentration in the ECF. This is achieved by changes in the rate and depth of breathing (i.e. by hyperventilation or hypoventilation), which blows off or retains carbon dioxide (and thus carbonic acid) in the blood plasma.The third line of defence is the renal system, which can add or remove bicarbonate ions to or from the ECF. The bicarbonate is derived from metabolic carbon dioxide which is enzymatically converted to carbonic acid in the renal tubular cells. The carbonic acid spontaneously dissociates into hydrogen ions and bicarbonate ions. When the pH in the ECF tends to fall (i.e. become more acidic) the hydrogen ions are excreted into the urine, while the bicarbonate ions are secreted into the blood plasma, causing the plasma pH to rise (correcting the initial fall). The converse happens if the pH in the ECF tends to rise: the bicarbonate ions are then excreted into the urine and the hydrogen ions into the blood plasma.Physiological corrective measures make up the second and third lines of defence. This is because they operate by making changes to the buffers, each of which consists of two components: a weak acid and its conjugate base. It is the ratio concentration of the weak acid to its conjugate base that determines the pH of the solution. Thus, by manipulating firstly the concentration of the weak acid, and secondly that of its conjugate base, the pH of the extracellular fluid (ECF) can be adjusted very accurately to the correct value. The bicarbonate buffer, consisting of a mixture of carbonic acid (H2CO3) and a bicarbonate (HCO−3) salt in solution, is the most abundant buffer in the extracellular fluid, and it is also the buffer whose acid to base ratio can be changed very easily and rapidly.An acid–base imbalance is known as acidaemia when the acidity is high, or alkalaemia when the acidity is low.

Ammonium bicarbonate

Ammonium bicarbonate is an inorganic compound with formula (NH4)HCO3, simplified to NH5CO3. The compound has many names, reflecting its long history. Chemically speaking, it is the bicarbonate salt of the ammonium ion. It is a colourless solid that degrades readily to carbon dioxide, water and ammonia.

Band 3 anion transport protein

Band 3 anion transport protein, also known as anion exchanger 1 (AE1) or band 3 or solute carrier family 4 member 1 (SLC4A1), is a protein that is encoded by the SLC4A1 gene in humans.

Band 3 anion transport protein is a phylogenetically-preserved transport protein responsible for mediating the exchange of chloride (Cl−) with bicarbonate (HCO3−) across plasma membranes. Functionally similar members of the AE clade are AE2 and AE3.

Bicarbonate buffer system

The bicarbonate buffer system is an acid-base homeostatic mechanism involving the balance of carbonic acid (H2CO3), bicarbonate ion (HCO−3), and carbon dioxide (CO2) in order to maintain pH in the blood and duodenum, among other tissues, to support proper metabolic function. Catalyzed by carbonic anhydrase, carbon dioxide (CO2) reacts with water (H2O) to form carbonic acid (H2CO3), which in turn rapidly dissociates to form a bicarbonate ion (HCO−3 ) and a hydrogen ion (H+) as shown in the following reaction:

As with any buffer system, the pH is balanced by the presence of both a weak acid (for example, H2CO3) and its conjugate base (for example, HCO−3) so that any excess acid or base introduced to the system is neutralized.

Failure of this system to function properly results in acid-base imbalance, such as acidemia (pH<7.35) and alkalemia (pH>7.45) in the blood.

Bicarbonate transporter protein

In molecular biology, bicarbonate transporter proteins are proteins which transport bicarbonate. Bicarbonate (HCO3 −) transport mechanisms are the principal regulators of pH in animal cells. Such transport also plays a vital role in acid-base movements in the stomach, pancreas, intestine, kidney, reproductive organs and the central nervous system. Functional studies have suggested four different HCO3 − transport modes. Anion exchanger proteins exchange HCO3 − for Cl− in a reversible, electroneutral manner. Na+/HCO3 − co-transport proteins mediate the coupled movement of Na+ and HCO3 − across plasma membranes, often in an electrogenic manner. Na+ driven Cl−/HCO3 − exchange and K+/HCO3 − exchange activities have also been detected in certain cell types, although the molecular identities of the proteins responsible remain to be determined.

Sequence analysis of the two families of HCO3 − transporters that have been cloned to date (the anion exchangers and Na+/HCO3 − co-transporters) reveals that they are homologous. This is not entirely unexpected, given that they both transport HCO3 − and are inhibited by a class of pharmacological agents called disulphonic stilbenes. They share around ~25-30% sequence identity, which is distributed along their entire sequence length, and have similar predicted membrane topologies, suggesting they have ~10 transmembrane (TM) domains.

A conserved domain is found at the C terminus of many bicarbonate transport proteins. It is also found in some plant proteins responsible for boron transport. In these proteins it covers almost the entire length of the sequence.

The Band 3 anion exchange proteins that exchange bicarbonate are the most abundant polypeptide in the red blood cell membrane, comprising 25% of the total membrane protein. The cytoplasmic domain of band 3 functions primarily as an anchoring site for other membrane-associated proteins. Included among the protein ligands of this domain are ankyrin, protein 4.2, protein 4.1, glyceraldehyde-3-phosphate dehydrogenase (GAPDH), phosphofructokinase, aldolase, hemoglobin, hemichromes, and the protein tyrosine kinase (p72syk).

Calcium bicarbonate

Calcium bicarbonate, also called calcium hydrogen carbonate, has a chemical formula Ca(HCO3)2. The term does not refer to a known solid compound; it exists only in aqueous solution containing the calcium (Ca2+), bicarbonate (HCO−3), and carbonate (CO2−3) ions, together with dissolved carbon dioxide (CO2). The relative concentrations of these carbon-containing species depend on the pH; bicarbonate predominates within the range 6.36–10.25 in fresh water.

All waters in contact with the atmosphere absorb carbon dioxide, and as these waters come into contact with rocks and sediments they acquire metal ions, most commonly calcium and magnesium, so most natural waters that come from streams, lakes, and especially wells, can be regarded as dilute solutions of these bicarbonates. These hard waters tend to form carbonate scale in pipes and boilers and they react with soaps to form an undesirable scum.

Attempts to prepare compounds such as solid calcium bicarbonate by evaporating its solution to dryness invariably yield instead the solid calcium carbonate:

Ca(HCO3)2(aq) → CO2(g) + H2O(l) + CaCO3(s).Very few solid bicarbonates other than those of the alkali metals except lithium and ammonium ion are known to exist.

The above reaction is very important to the formation of stalactites, stalagmites, columns, and other speleothems within caves, and for that matter, in the formation of the caves themselves. As water containing carbon dioxide (including extra CO2 acquired from soil organisms) passes through limestone or other calcium carbonate-containing minerals, it dissolves part of the calcium carbonate, hence becomes richer in bicarbonate. As the groundwater enters the cave, the excess carbon dioxide is released from the solution of the bicarbonate, causing the much less soluble calcium carbonate to be deposited.

In the reverse process, dissolved carbon dioxide (CO2) in rainwater (H2O) reacts with limestone calcium carbonate (CaCO3) to form soluble calcium bicarbonate (Ca(HCO3)2). This soluble compound is then washed away with the rainwater. This form of weathering is called carbonation.

In medicine, calcium bicarbonate is sometimes administered intravenously to immediately correct the cardiac depressor effects of hypokalemia by increasing calcium concentration in serum, and at the same time, correcting the acid usually present.

Carbonate

In chemistry, a carbonate is a salt of carbonic acid (H2CO3), characterized by the presence of the carbonate ion, a polyatomic ion with the formula of CO2−3. The name may also refer to a carbonate ester, an organic compound containing the carbonate group C(=O)(O–)2.

The term is also used as a verb, to describe carbonation: the process of raising the concentrations of carbonate and bicarbonate ions in water to produce carbonated water and other carbonated beverages – either by the addition of carbon dioxide gas under pressure, or by dissolving carbonate or bicarbonate salts into the water.

In geology and mineralogy, the term "carbonate" can refer both to carbonate minerals and carbonate rock (which is made of chiefly carbonate minerals), and both are dominated by the carbonate ion, CO2−3. Carbonate minerals are extremely varied and ubiquitous in chemically precipitated sedimentary rock. The most common are calcite or calcium carbonate, CaCO3, the chief constituent of limestone (as well as the main component of mollusc shells and coral skeletons); dolomite, a calcium-magnesium carbonate CaMg(CO3)2; and siderite, or iron(II) carbonate, FeCO3, an important iron ore. Sodium carbonate ("soda" or "natron") and potassium carbonate ("potash") have been used since antiquity for cleaning and preservation, as well as for the manufacture of glass. Carbonates are widely used in industry, e.g. in iron smelting, as a raw material for Portland cement and lime manufacture, in the composition of ceramic glazes, and more.

Carbonic acid

Not to be confused with carbolic acid, an antiquated name for phenol.Carbonic acid is a chemical compound with the chemical formula H2CO3 (equivalently: OC(OH)2). It is also a name sometimes given to solutions of carbon dioxide in water (carbonated water), because such solutions contain small amounts of H2CO3. In physiology, carbonic acid is described as volatile acid or respiratory acid, because it is the only acid excreted as a gas by the lungs. It plays an important role in the bicarbonate buffer system to maintain acid–base homeostasis.

Carbonic acid, which is a weak acid, forms two kinds of salts: the carbonates and the bicarbonates. In geology, carbonic acid causes limestone to dissolve, producing calcium bicarbonate, which leads to many limestone features such as stalactites and stalagmites.

It was long believed that carbonic acid could not exist as a pure compound. However, in 1991 it was reported that NASA scientists had succeeded in making solid H2CO3 samples.

Diabetic ketoacidosis

Diabetic ketoacidosis (DKA) is a potentially life-threatening complication of diabetes mellitus. Signs and symptoms may include vomiting, abdominal pain, deep gasping breathing, increased urination, weakness, confusion and occasionally loss of consciousness. A person's breath may develop a specific "fruity" smell. Onset of symptoms is usually rapid. People without a previous diagnosis of diabetes may develop DKA as the first obvious symptom.DKA happens most often in those with type 1 diabetes but can also occur in those with other types of diabetes under certain circumstances. Triggers may include infection, not taking insulin correctly, stroke and certain medications such as steroids. DKA results from a shortage of insulin; in response, the body switches to burning fatty acids, which produces acidic ketone bodies. DKA is typically diagnosed when testing finds high blood sugar, low blood pH and ketoacids in either the blood or urine.The primary treatment of DKA is with intravenous fluids and insulin. Depending on the severity, insulin may be given intravenously or by injection under the skin. Usually, potassium is also needed to prevent the development of low blood potassium. Throughout treatment, blood sugar and potassium levels should be regularly checked. Antibiotics may be required in those with an underlying infection. In those with severely low blood pH, sodium bicarbonate may be given; however, its use is of unclear benefit and typically not recommended.Rates of DKA vary around the world. In the United Kingdom, about 4% of people with type 1 diabetes develop DKA each year, while in Malaysia the condition affects about 25% of type-1 diabetics a year. DKA was first described in 1886 and, until the introduction of insulin therapy in the 1920s, it was almost universally fatal. The risk of death with adequate and timely treatment is around 1–4%.

Electrogenic sodium bicarbonate cotransporter 1

Electrogenic sodium bicarbonate cotransporter 1 is a membrane transport protein that in humans is encoded by the SLC4A4 gene.

Electrogenic sodium bicarbonate cotransporter 4

Electrogenic sodium bicarbonate cotransporter 4 is a protein that in humans is encoded by the SLC4A5 gene.

Electroneutral sodium bicarbonate exchanger 1

Electroneutral sodium bicarbonate exchanger 1 is a protein that in humans is encoded by the SLC4A8 gene.

Gastric acid

Gastric acid, gastric juice, or stomach acid, is a digestive fluid formed in the stomach and is composed of hydrochloric acid (HCl), potassium chloride (KCl), and sodium chloride (NaCl). The acid plays a key role in digestion of proteins, by activating digestive enzymes, and making ingested proteins unravel so that digestive enzymes break down the long chains of amino acids.

Gastric acid is produced by cells in the lining of the stomach, which are coupled in feedback systems to increase acid production when needed. Other cells in the stomach produce bicarbonate, a base, to buffer the fluid, ensuring that it does not become too acidic. These cells also produce mucus, which forms a viscous physical barrier to prevent gastric acid from damaging the stomach. The pancreas further produces large amounts of bicarbonate and secretes bicarbonate through the pancreatic duct to the duodenum to completely neutralize any gastric acid that passes further down into the digestive tract.

The main constituent of gastric acid is hydrochloric acid which is produced by parietal cells (also called oxyntic cells) in the gastric glands in the stomach. Its secretion is a complex and relatively energetically expensive process. Parietal cells contain an extensive secretory network (called canaliculi) from which the hydrochloric acid is secreted into the lumen of the stomach. The pH of gastric acid is 1.5 to 3.5 in the human stomach lumen, the acidity being maintained by the proton pump H+/K+ ATPase. The parietal cell releases bicarbonate into the bloodstream in the process, which causes a temporary rise of pH in the blood, known as an alkaline tide.

The highly acidic environment in the stomach lumen causes proteins from food to lose their characteristic folded structure (or denature). This exposes the protein's peptide bonds. The gastric chief cells of the stomach secrete enzymes for protein breakdown (inactive pepsinogen, and in infancy rennin). Hydrochloric acid activates pepsinogen into the enzyme pepsin, which then helps digestion by breaking the bonds linking amino acids, a process known as proteolysis. In addition, many microorganisms have their growth inhibited by such an acidic environment, which is helpful to prevent infection.

Metabolic acidosis

Metabolic acidosis is a disorder that occurs when the body produces excessive amounts of acid, such as ketoacids or lactic acid; the kidneys are unable to remove enough acid produced from normal metabolism; or the body loses too much bicarbonate ion (HCO−3). If unchecked, metabolic acidosis can lead to acidemia, which is arterial blood pH lower than 7.37 due to increased production of hydrogen ions by the body or the loss and/or inability of the body to form bicarbonate (HCO−3) in the kidney or gastrointestinal tract. Its causes are diverse, and its consequences can be serious, including coma and death. Together with respiratory acidosis, it is one of the two general causes of acidemia.

Terminology:

Acidosis refers to a process that causes a low pH in blood and tissues.

Acidemia refers specifically to a decrease in pH in circulating blood and is caused by acidosis.In most cases, acidosis occurs first for reasons explained below. Free H+ ions then diffuse into the blood, lowering the pH. Arterial blood gas analysis detects acidemia (pH lower than 7.35). When acidemia is present, acidosis is presumed.

Potassium bicarbonate

Potassium bicarbonate (also known as potassium hydrogen carbonate or potassium acid carbonate) is the inorganic compound with the chemical formula KHCO3. It is a white solid.

Sodium bicarbonate

Sodium bicarbonate (IUPAC name: sodium hydrogen carbonate), commonly known as baking soda, is a chemical compound with the formula NaHCO3. It is a salt composed of a sodium cation (Na+) and a bicarbonate anion (HCO3−). Sodium bicarbonate is a white solid that is crystalline, but often appears as a fine powder. It has a slightly salty, alkaline taste resembling that of washing soda (sodium carbonate). The natural mineral form is nahcolite. It is a component of the mineral natron and is found dissolved in many mineral springs.

Sodium bicarbonate cotransporter 3

Sodium bicarbonate cotransporter 3 is a protein which in humans is encoded by the SLC4A7 gene.

Sodium bicarbonate transporter-like protein 11

Sodium bicarbonate transporter-like protein 11 is a protein that in humans is encoded by the SLC4A11 gene.

Common oxides
Exotic oxides
Polymers
Compounds derived from oxides
Common for blood tests (CPT 82000–84999)
Electrolytes
Acid-base
Iron tests
Hormones
Metabolism
Cardiovascular
Liver function tests
Pancreas
Compounds
Carbon ions
Oxides and related

Languages

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.