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Bicarbonate
Bicarbonate
Bicarbonate
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Bicarbonate
Skeletal formula of bicarbonate with the explicit hydrogen added
Skeletal formula of bicarbonate with the explicit hydrogen added
Ball and stick model of bicarbonate
Ball and stick model of bicarbonate
Names
IUPAC name
Hydrogencarbonate
Systematic IUPAC name
Hydroxidodioxidocarbonate(1−)[1]
Other names
  • Hydrogen carbonate[1]
  • Hydrocarbonate
Identifiers
3D model (JSmol)
3903504
ChEBI
ChEMBL
ChemSpider
49249
KEGG
UNII
  • InChI=1S/CH2O3/c2-1(3)4/h(H2,2,3,4)/p-1 checkY
    Key: BVKZGUZCCUSVTD-UHFFFAOYSA-M checkY
  • OC([O-])=O
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).

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][5] The name lives on as a trivial name.

Chemical properties

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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 (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 CO2−3, the carbonate ion, as shown by these equilibrium reactions:

CO2−3 + 2 H2O ⇌ HCO3 + H2O + HO− ⇌ H2CO3 + 2 HO−
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.[6]

Physiological role

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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 HCO−3. 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 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.[7]

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). 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 the pancreas in response to the hormone secretin to neutralize the acidic chyme entering the duodenum from the stomach.[9]

Bicarbonate in the environment

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Bicarbonate is the dominant form of dissolved inorganic carbon in sea water,[10] and in most fresh waters. As such it is an important sink in the carbon cycle.

Some plants like Chara utilize carbonate and produce calcium carbonate (CaCO3) as a result of biological metabolism.[11]

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.[citation needed]

The flow of bicarbonate ions from rocks weathered by the carbonic acid in rainwater is an important part of the carbon cycle.

Other uses

[edit]

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.[12]

Ammonium bicarbonate is used in the manufacturing of some cookies, crackers, and biscuits.[13]

Diagnostics

[edit]

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 chloride, potassium, and sodium, to assess electrolyte levels in an electrolyte panel test (which has Current Procedural Terminology, CPT, code 80051).[14]

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.[15]

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

Bicarbonate compounds

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See also

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Bicarbonate

View on Grokipedia
from Grokipedia
Bicarbonate, also known as hydrogen carbonate, is a polyatomic anion with the HCO₃⁻ and a molecular weight of 61.017 g/mol. It serves as the conjugate base of (H₂CO₃), formed by the dissociation of carbonic acid into a (H⁺) and the bicarbonate ion. This ion is fundamental in aqueous solutions, where it participates in equilibrium reactions that influence levels. In biological systems, bicarbonate plays a critical role in maintaining acid-base homeostasis, acting as a key component of the bicarbonate buffer system in blood and extracellular fluids. It is generated through the reaction of carbon dioxide (CO₂) with water, catalyzed by the enzyme carbonic anhydrase, to form carbonic acid, which then dissociates into H⁺ and HCO₃⁻. This process allows bicarbonate to neutralize excess acids by combining with H⁺ to produce CO₂, which is subsequently exhaled by the lungs, thereby regulating blood pH within a narrow range of approximately 7.35 to 7.45. Bicarbonate is also a metabolic byproduct transported in the blood to the lungs for elimination as CO₂. Chemically, bicarbonate exhibits amphoteric properties, capable of acting as both an acid and a base; it can donate a proton to form (CO₃²⁻) or accept one to reform . In aqueous environments, it contributes to the of solutions, as seen in natural waters and biological buffers where it resists changes. Bicarbonate ions are prevalent in mammalian tissues and fluids, where they support physiological processes such as respiration, electrolyte balance, and enzyme function. Disruptions in bicarbonate levels, such as in or , can lead to significant health issues, underscoring its essential role in human physiology.

Fundamentals

Definition and Nomenclature

Bicarbonate refers to the anion HCO₃⁻, which serves as the conjugate base of (H₂CO₃) and constitutes an essential component of the buffering system in aqueous environments. This ion carries a of -1 and plays a central role in acid-base equilibria involving dissolution. Its molecular formula is CHO₃, comprising one carbon atom, one , and three oxygen atoms in a 1:1:3 atomic ratio. In , the systematic IUPAC name for the anion is hydrogen carbonate, reflecting its structure as a protonated form of the . The term "bicarbonate" remains in widespread common usage, particularly for naming salts such as (NaHCO₃), where the cation pairs with the HCO₃⁻ anion. This naming distinguishes it from the (CO₃²⁻), which is the fully deprotonated conjugate base of and features a -2 charge without the . The etymology of "bicarbonate" traces back to 1814, when English chemist coined the term from the prefix "bi-" (indicating two) and "," based on an early equivalent-weight system that viewed the ion as containing two equivalents of per base equivalent. This historical convention underscores its intermediate position in the carbonic acid dissociation sequence, though modern favors the more descriptive "hydrogen carbonate."

Historical Discovery

The discovery of bicarbonate as a chemical entity emerged from early investigations into gases and alkalis in the . Scottish chemist played a pivotal role through his experiments on "fixed air," the term he coined for (CO₂), identified in 1754 while studying the heating of magnesia alba (magnesium carbonate). Black demonstrated that fixed air was released during the calcination of carbonates and could be absorbed by caustic alkalis, such as (potassium hydroxide), to form mild alkalis like potassium carbonate; he further observed that these mild alkalis could absorb additional fixed air to produce an , now recognized as (KHCO₃), formed by the absorption of additional fixed air into solutions of mild alkalis such as potash. These findings, detailed in Black's 1756 dissertation and lectures, distinguished fixed air as a unique substance and highlighted its role in bicarbonate formation, marking a key milestone in pneumatic chemistry. In the late , precursors to industrial bicarbonate production appeared with French Nicolas Leblanc's 1791 process for manufacturing from salt, , and , which indirectly facilitated bicarbonate-related chemistry by providing a source of soda ash. The term "bicarbonate" itself was coined in 1814 by English to describe salts containing the HCO₃⁻ ion, such as bicarbonate of , reflecting the "bi-" prefix for compounds with twice the acid content relative to the base compared to ordinary carbonates. By the mid-19th century, commercial production of began in the United States in 1846, when physician Austin Church and entrepreneur John Dwight established a refining operation in New York, packaging the compound—refined from crude soda ash—for use as a , initially distributed by hand. A major advancement came with the , patented in 1861 by Belgian chemist , which synthesized directly and efficiently by reacting ammoniated brine with under controlled conditions, yielding NaHCO₃ as a key intermediate before conversion to . This method, building on earlier experiments like Leblanc's, revolutionized large-scale production by reducing costs and environmental impact compared to prior techniques. Solvay's first plant in Couillet, , began operations in 1863, establishing bicarbonate as an industrially viable compound essential for subsequent chemical applications.

Chemical Properties

Molecular Structure and Bonding

The bicarbonate ion (HCO₃⁻) features a central carbon atom bonded to a hydroxyl group (C–OH single bond), a double-bonded oxygen atom (C=O), and a singly bonded oxygen atom bearing a negative charge (C–O⁻). This arrangement satisfies the octet rule for all atoms, with the carbon contributing four valence electrons and each oxygen six, plus the additional electron from the negative charge. Resonance structures delocalize the double bond and negative charge between the two non-hydroxyl oxygen atoms, resulting in equivalent bond characters for those two C–O linkages rather than distinct single and double bonds. The electron geometry and molecular shape around the central carbon atom are trigonal planar, arising from three electron domains (one and two s) with no lone pairs on carbon. Bond angles are approximately 120°, consistent with sp² hybridization of the carbon atom, where one 2s and two 2p orbitals form three equivalent sp² hybrid orbitals for bonding, and the remaining p orbital participates in pi bonding. Due to , the two equivalent C–O bond lengths are intermediate, typically around 1.26–1.28 , shorter than a pure C–O (1.43 ) but longer than a C=O (1.20 ); the C–OH bond is longer at about 1.36 . In bicarbonate salts such as (NaHCO₃), the bonding is predominantly ionic between the Na⁺ cation and the HCO₃⁻ anion, with the anion's negative charge distributed via stabilizing the lattice. In aqueous solutions, the bicarbonate ion engages in hydrogen bonding, where the O–H group acts as a donor to molecules and the resonant oxygen atoms serve as acceptors, influencing and contributing to the ion's amphoteric behavior. Infrared spectroscopy reveals characteristic absorption bands for the bicarbonate ion, including the C=O stretching at approximately 1650 cm⁻¹, reflecting the partial double-bond character, and the O–H stretching mode around 2500–3000 cm⁻¹, broadened by hydrogen bonding in solution. These bands arise from the vibrational modes of the resonant structure and are useful for identifying the ion in various environments.

Acidity, Basicity, and Equilibrium

The bicarbonate ion (HCO₃⁻) is amphiprotic, meaning it can function as both a Brønsted-Lowry acid and a base due to its ability to donate or accept a proton. As an acid, it dissociates according to the equilibrium: \ceHCO3−<=>H++CO32−\ce{HCO3^- <=> H^+ + CO3^{2-}} with an Ka2K_{a2} corresponding to pKa2=10.33\mathrm{p}K_{a2} = 10.33 at 25°C and zero . This step represents the second dissociation of the system, where bicarbonate loses a proton to form the carbonate ion. As a base, bicarbonate accepts a proton to form : \ceHCO3−+H+<=>H2CO3\ce{HCO3^- + H^+ <=> H2CO3} This reaction is the reverse of the first dissociation of , which occurs via: \ceH2CO3<=>H++HCO3−\ce{H2CO3 <=> H^+ + HCO3^-} with pKa1=6.35\mathrm{p}K_{a1} = 6.35 under the same conditions. The overall process ties into the hydration of dissolved : \ceCO2(aq)+H2O<=>H2CO3<=>H++HCO3−\ce{CO2 (aq) + H2O <=> H2CO3 <=> H^+ + HCO3^-} where the hydration step is relatively slow without , but the apparent for the combined process reflects the effective pKa1\mathrm{p}K_{a1} value. These equilibria position bicarbonate as a key intermediate in the buffer system, with the pKa values indicating moderate acidity for the first step and weaker acidity for the second. The buffering capacity of the bicarbonate system arises from these proton transfer reactions, which resist changes by shifting equilibria upon addition of or base. For the H₂CO₃/HCO₃⁻ pair, the is described by the Henderson-Hasselbalch : pH=pKa1+log⁡10([\ceHCO3−][\ceH2CO3])\mathrm{pH} = \mathrm{p}K_{a1} + \log_{10} \left( \frac{[\ce{HCO3^-}]}{[\ce{H2CO3}]} \right) Optimal buffering occurs near = pK_{a1} (≈6.35), where equal concentrations of the conjugate and base maximize resistance to shifts. Similarly, for the HCO₃⁻/CO₃²⁻ pair, buffering is effective around = pK_{a2} (≈10.33). The system's capacity depends on total concentration and is influenced by the rapid equilibration of CO₂(aq) with H₂CO₃, though true H₂CO₃ constitutes only a small (≈0.3%) of the "" pool. Bicarbonate stability and solubility in aqueous solutions vary with pH, temperature, and pressure, primarily through shifts in the above equilibria. At low pH (below ≈6), protonation drives decomposition: \ceHCO3−+H+−>H2CO3−>CO2(aq)+H2O\ce{HCO3^- + H^+ -> H2CO3 -> CO2 (aq) + H2O} releasing CO₂ gas, which reduces bicarbonate concentration and solubility. Higher temperatures decrease CO₂ solubility (e.g., Henry's constant increases by ≈4% per °C from 0–30°C), promoting decomposition and lowering effective bicarbonate stability in open systems. Conversely, elevated pressure enhances CO₂ dissolution (solubility proportional to partial pressure via Henry's law), stabilizing bicarbonate at given pH by favoring the hydration equilibrium. These dependencies are critical in contexts like geochemical systems, where pH and pressure control speciation. In addition to acid-induced decomposition in solution, solid bicarbonates undergo thermal decomposition upon heating, independent of pH. All metal bicarbonates decompose to the corresponding carbonate, CO₂, and H₂O, following the general reaction: \ce2MHCO3−>M2CO3+CO2+H2O\ce{2MHCO3 -> M2CO3 + CO2 + H2O} (where M is a metal cation, typically alkali or alkaline earth metals). For example, sodium bicarbonate decomposes as: \ce2NaHCO3−>Na2CO3+CO2+H2O\ce{2NaHCO3 -> Na2CO3 + CO2 + H2O} This reaction typically initiates at temperatures around 80–100°C for sodium bicarbonate and varies for other metals, with alkali metal bicarbonates decomposing around 300°C.

Biological Roles

In Human Physiology

In human physiology, bicarbonate plays a central role in maintaining acid-base , particularly through the bicarbonate-carbonic acid buffer system in , which stabilizes between 7.35 and 7.45. This system is essential for counteracting the daily production of acids from and respiration, where (CO₂) generated in tissues diffuses into red cells and reacts with to form , which dissociates into bicarbonate (HCO₃⁻) and ions (H⁺). The H⁺ ions are buffered by , while bicarbonate is transported out of the red cells into plasma via the , an anion exchange where chloride ions (Cl⁻) enter the cells to maintain electroneutrality. This mechanism facilitates the of approximately 70-80% of CO₂ from tissues to the lungs as bicarbonate, preventing excessive acidification of and enabling efficient . The kidneys further regulate acid-base balance by controlling bicarbonate levels through and , ensuring long-term pH stability. In the proximal tubules, nearly all filtered bicarbonate is reabsorbed via carbonic anhydrase-mediated conversion to CO₂ and water, which then re-enters tubular cells to reform bicarbonate for return to the blood. In response to , distal tubules increase bicarbonate generation and H⁺ , often as titratable acids or , while in , bicarbonate rises to lower plasma levels. This renal compensation corrects by elevating bicarbonate to normalize or addresses by reducing it, with full adjustments taking hours to days depending on the severity. Bicarbonate also contributes to digestion by neutralizing acidic chyme in the duodenum, creating an optimal environment for enzymatic activity. Stimulated by in response to low duodenal , the secretes a bicarbonate-rich fluid into the , raising the from around 2-3 to 6-7, which protects the intestinal mucosa and activates pancreatic enzymes like and . This process prevents autodigestion of the gut lining and facilitates breakdown and absorption. Normal plasma bicarbonate concentrations range from 22 to 28 mEq/L, reflecting balanced production, , and excretion. Disruptions occur in conditions like , where elevates CO₂ levels, prompting renal compensation that increases plasma bicarbonate to 24-30 mEq/L or higher to restore ; conversely, in , bicarbonate falls below 22 mEq/L due to acid overload or loss, such as in or .

In Other Organisms

In , bicarbonate plays a crucial role in carbon concentrating mechanisms (CCMs) that enhance by accumulating inorganic carbon around the enzyme , particularly in CO₂-limited environments. These mechanisms involve active uptake and transport of HCO₃⁻ into cells and chloroplasts, where it serves as a stable carbon source for CO₂ fixation via carbonic anhydrase-mediated conversion. In terrestrial , such as those engineered with cyanobacterial bicarbonate transporters, HCO₃⁻ accumulation in chloroplasts can boost photosynthetic rates by up to 30% under ambient CO₂ conditions. Aquatic plants and algae, including microalgae and seagrasses, rely heavily on HCO₃⁻ uptake for CO₂ fixation due to the low diffusivity of CO₂ in water. Chloroplasts in species like Chlamydomonas and Elodea simultaneously transport HCO₃⁻ and CO₂, enabling active accumulation of inorganic carbon even at low external concentrations. This process is integral to biophysical CCMs, where HCO₃⁻ is dehydrated to CO₂ in a localized microenvironment, suppressing photorespiration and supporting growth in CO₂-poor aquatic habitats. In marine organisms, bicarbonate is essential for calcification processes in corals and , where it provides the carbon source for forming (CaCO₃) skeletons and shells. Corals actively pump HCO₃⁻ into the subcalicoblastic space via anion transporters like SLC4γ, elevating and carbonate saturation to drive . , such as oysters and mussels, similarly utilize HCO₃⁻ in extrapallial fluid for shell formation, but ocean acidification reduces ion availability by shifting equilibria toward CO₂, impairing rates by 20-50% in some species. Microbial communities across diverse taxa employ bicarbonate in key metabolic pathways, including and adaptations in extremophilic environments. In anaerobic methanogenic like Methanobacterium, HCO₃⁻ acts as a substrate for acetate-dependent methanogenesis, where elevated concentrations enhance microbial diversity and shift community composition toward hydrogenotrophic pathways. Certain in soda lakes, such as haloalkaliphilic Thioalkalivibrio species, adapt to high HCO₃⁻ levels (tens to hundreds of millimolar) by evolving specialized transporters and enzymes for carbon fixation, enabling oxidation and cycling in pH 9-11 conditions. Among non-marine animals, and utilize bicarbonate for acid-base in their respective fluids. Freshwater and marine gills transport HCO₃⁻ via ionocytes and , accumulating it in plasma to counteract during , with net HCO₃⁻ influx balancing H⁺ efflux for stability. In , buffering relies on HCO₃⁻/CO₂ equilibrium, with concentrations around 10-20 mM contributing up to 50% of total buffer capacity in locusts, decreasing with to maintain near neutrality during metabolic fluctuations.

Environmental Occurrence

In Natural Waters and Cycles

Bicarbonate (HCO₃⁻) serves as the dominant anion in many natural waters, particularly in rivers and , where it typically constitutes over 50% of the total anionic composition due to its formation through geochemical processes. In global river systems, average concentrations range from approximately 100 mg/L, reflecting inputs from rock dissolution across diverse watersheds. often exhibits higher levels, with mean concentrations around 250 mg/L, influenced by longer residence times and interaction with carbonate-rich aquifers. These concentrations arise primarily from the weathering of and rocks, where atmospheric CO₂ dissolves in water to form , which reacts with minerals to release bicarbonate ions via reactions such as CaSiO3+2CO2+3H2O→Ca2++2HCO3−+H4SiO4\text{CaSiO}_3 + 2\text{CO}_2 + 3\text{H}_2\text{O} \rightarrow \text{Ca}^{2+} + 2\text{HCO}_3^- + \text{H}_4\text{SiO}_4 In ocean chemistry, bicarbonate dominates the dissolved inorganic carbon pool, comprising about 90% of the total at concentrations of roughly 2.3 mM in surface seawater. This form contributes significantly to the ocean's total alkalinity, which averages around 2.4 meq/L, providing a key measure of the water's capacity to neutralize acids. Bicarbonate's presence helps maintain the marine pH buffer system, stabilizing conditions despite fluctuations in CO₂ inputs. Bicarbonate plays a crucial role in regulating and water hardness in natural aquatic systems through its buffering capacity, where it resists acidification by accepting protons and forms part of the carbonate equilibrium. In freshwater, it contributes to temporary hardness by associating with calcium and magnesium ions from dissolved , influencing the suitability of for various uses. Measurements of , largely driven by bicarbonate, serve as indicators of environmental stress; elevated levels can signal from or agricultural runoff, while changes may reflect dynamics where nutrient enrichment alters buffering. On a global scale, rivers deliver approximately 0.6 Gt of carbon annually to the in the form of bicarbonate, representing a major flux in the hydrological cycle that sustains marine carbon reservoirs. exacerbates this process by accelerating the dissolution of carbonate rocks, thereby increasing bicarbonate release and altering local water chemistry in affected regions.

Geological and Atmospheric Presence

Bicarbonate is present in the primarily as constituent minerals in deposits formed through the evaporation of ancient alkaline lakes. (NaHCO₃) occurs abundantly in the layers of the Parachute Creek Member within the Eocene Green River Formation of the Piceance Basin, northwestern , where it crystallized under conditions of high and low during episodic lake regressions. (Na₂CO₃·NaHCO₃·2H₂O), a hydrated , forms in similar evaporitic environments, notably in the Wilkins Peak Member of the Green River Formation and in modern and ancient soda lakes worldwide, resulting from the concentration of sodium-rich brines. Silicate weathering represents a fundamental geological process linking bicarbonate to the long-term carbon cycle, acting as a primary regulator of atmospheric CO₂ levels over millions of years. In this reaction, carbonic acid derived from atmospheric CO₂ reacts with silicate minerals, exemplified by the simplified equation: CaSiO3+2CO2+3H2O→Ca2++2HCO3−+H4SiO4\text{CaSiO}_3 + 2\text{CO}_2 + 3\text{H}_2\text{O} \rightarrow \text{Ca}^{2+} + 2\text{HCO}_3^- + \text{H}_4\text{SiO}_4 This weathering produces bicarbonate ions that are transported to oceans, where they facilitate carbonate precipitation and burial, effectively sequestering carbon and stabilizing Earth's climate on geological timescales. Atmospheric connections to bicarbonate arise from volcanic activity and cloud processes, integrating gas-phase emissions with geological sinks. Volcanic eruptions release substantial CO₂, which serves as a key precursor for bicarbonate formation through subsequent hydration and reaction in aqueous environments. In the atmosphere, CO₂ hydration within droplets can generate bicarbonate ions, particularly when interacting with alkaline dust particles like , contributing to the formation of hygroscopic aerosols that influence microphysics and radiative properties. Fossil records preserved in sediments offer insights into historical bicarbonate dynamics, reflecting variations in tied to global climate shifts. During the Paleocene-Eocene Thermal Maximum approximately 56 million years ago, widespread dissolution in deep-sea sediments indicated a transient increase in ocean acidity and reconfiguration of the system, including elevated bicarbonate levels from massive carbon inputs, which amplified warming feedbacks.

Applications and Uses

Medical and Diagnostic Uses

Sodium bicarbonate is administered intravenously to treat severe , particularly in severe cases of (DKA) with arterial less than 7.0 despite initial fluid and insulin therapy, often accompanied by serum bicarbonate levels below 5-10 mEq/L and elevated blood glucose. This therapy is recommended for cases with refractory acidosis, hemodynamic instability, or , as it helps correct the acid-base imbalance by increasing plasma bicarbonate concentration. However, routine use in mild to moderate DKA is not supported due to potential risks like , and it is reserved for severe presentations. Oral supplementation is a standard intervention for managing in (CKD), where it counters acid buildup by raising serum bicarbonate levels and slowing disease progression. Clinical trials have demonstrated that this therapy improves kidney function, reduces the risk of end-stage renal disease, and enhances overall outcomes in patients with CKD stages 3-5. Dosing typically aims to maintain serum bicarbonate above 22 mEq/L, with long-term use showing benefits in preserving renal function without significant adverse effects in most cases. As an antacid, neutralizes excess gastric to alleviate symptoms of and acid indigestion, providing rapid symptomatic relief by raising the in the . It is commonly formulated in effervescent tablets or powders combined with , which reacts to release gas, aiding dissolution and promoting burping to relieve . This mechanism not only buffers acid but also enhances patient compliance through the fizzy effervescence, though prolonged use may lead to if not monitored. In diagnostics, arterial blood gas (ABG) analysis quantifies bicarbonate concentration ([HCO₃⁻]) to assess acid-base disorders, with normal values ranging from 22-28 mEq/L; deviations indicate metabolic acidosis (low [HCO₃⁻]) or alkalosis (high [HCO₃⁻]). Bicarbonate levels from ABG are integral to calculating the anion gap, typically using the formula AG = Na⁺ - (Cl⁻ + HCO₃⁻), where an elevated gap (>12 mEq/L) suggests high-anion-gap metabolic acidosis, such as in lactic acidosis or DKA. This measurement helps differentiate causes of acid-base imbalances and guides therapeutic decisions. Recent advancements include the use of in , which maintain a neutral and reduce issues compared to lactate-buffered alternatives, leading to lower rates of and improved peritoneal membrane function. Studies from 2020-2025, including trials as of 2024-2025, have explored sodium bicarbonate's role in managing associated with severe , where low bicarbonate levels predict poor prognosis; preliminary investigations suggest adjuvant intravenous administration for correcting acidemia or nebulized for potential respiratory benefits in cases such as inhibiting viral entry or improving ARDS, with low-certainty evidence indicating possible acceleration of recovery, though larger trials are needed to confirm efficacy and safety. A 2025 randomized trial (BICARICU-2) investigated bicarbonate infusion in severe ( ≤7.20) with , showing potential benefits in recovery but no reduction in day-90 mortality; further studies are ongoing.

Industrial and Household Applications

Sodium bicarbonate, commonly known as , serves as a in by reacting with acidic components such as or to release gas, which causes to rise. It is approved as a under the code E500 in the and is widely used in processed foods like biscuits and cakes for this purpose. In household cleaning, acts as a mild in toothpastes and scouring powders, effectively removing stains without damaging surfaces due to its low abrasiveness. It also neutralizes odors in refrigerators and carpets by absorbing acidic volatile compounds and releasing . Industrially, is a key component in dry chemical fire extinguishers, where it decomposes under heat to produce , smothering flames in Class B and E fires involving flammable liquids and electrical equipment. In , it is added to adjust levels in systems, neutralizing acidity to prevent in pipes and maintain around 30-100 mg/L as needed. For , functions as a buffer, stabilizing the alkaline conditions required for reactive dyes to fix onto fabrics like , typically at pH 10-11. Beyond these, is used as an in pharmaceutical formulations to aid in tablet disintegration and control during manufacturing. In agriculture, is applied as an organic on crops such as grapes and berries to combat powdery mildew by disrupting fungal spore germination on contact. For environmental remediation, neutralizes by raising and precipitating like iron and aluminum, reducing toxicity in affected waterways.

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