Carbonic acid

Not to be confused with carbolic acid, an antiquated name for phenol.
Carbonic acid
Structural formula
Ball-and-stick model
Names
Preferred IUPAC name
Carbonic acid[1]
Other names
Carbon dioxide solution
Dihydrogen carbonate
Hydrogen bicarbonate
Acid of air
Aerial acid
Hydroxymethanoic acid
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.133.015
EC Number
  • 610-295-3
KEGG
Properties
CH2O3
Molar mass 62.024 g·mol−1
Density 1.668 g/cm3
Only stable in solution
Acidity (pKa) 3.6 (pKa1 for H2CO3 only), 6.3 (pKa1 including CO2(aq)), 10.32 (pKa2)
Conjugate base Bicarbonate
Related compounds
Related compounds
Acetone
Urea
Carbonyl fluoride
Hazards
NFPA 704
Flammability code 0: Will not burn. E.g. waterHealth code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g. sodium chlorideReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
0
1
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

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.[2] 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.[3]

Chemical equilibrium

When carbon dioxide dissolves in water it exists in chemical equilibrium with carbonic acid:[4]

The hydration equilibrium constant at 25 °C is called Kh, which in the case of carbonic acid is [H2CO3]/[CO2] ≈ 1.7×10−3 in pure water[5] and ≈ 1.2×10−3 in seawater.[6] Hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO2 molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2 + H2O → H2CO3) and 23 s−1 for the reverse reaction (H2CO3 → CO2 + H2O). The addition of two molecules of water to CO2 would give orthocarbonic acid, C(OH)4, which exists only in minute amounts in aqueous solution.

Addition of base to an excess of carbonic acid gives bicarbonate (hydrogen carbonate). With excess base, carbonic acid reacts to give carbonate salts.

Role of carbonic acid in blood

Bicarbonate is an intermediate in the transport of CO2 out of the body by respiratory gas exchange. The hydration reaction of CO2 is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which increases the reaction rate, producing bicarbonate (HCO3) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO2 and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.[7] A 2016 theoretical report suggests that carbonic acid may play a pivotal role in protonating various nitrogen bases in blood serum.[8]

Role of carbonic acid in ocean chemistry

The oceans of the world have absorbed almost half of the CO2 emitted by humans from the burning of fossil fuels.[9]  It has been estimated that the extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about −0.1 unit from pre-industrial levels. This is known as ocean acidification, even though the ocean remains basic.[10]

Acidity of carbonic acid

Carbonic acid is a carboxylic acid with a hydroxyl group as the substituent. It is also a polyprotic acid — specifically it is diprotic, meaning that it has two protons that may dissociate from the parent molecule. Thus, there are two dissociation constants, first of which is for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3:

Ka1 = 2.5×10−4;[4] pKa1 = 3.6 at 25 °C.

Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution, carbonic acid exists in equilibrium with carbon dioxide, and the concentration of H2CO3 is much lower than the concentration of CO2. In many analyses, H2CO3 includes dissolved CO2 (referred to as CO2(aq)), H2CO3* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows:[4]

H2CO3* ⇌ HCO3 + H+
Ka(app) = 4.47×10−7; pK(app) = 6.35 at 25 °C and ionic strength = 0.0.

Whereas this apparent pKa is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO2-containing solutions. A similar situation applies to sulfurous acid (H2SO3), which exists in equilibrium with substantial amounts of unhydrated sulfur dioxide.

The second constant is for the dissociation of the bicarbonate ion into the carbonate ion CO32−:

Ka2 = 4.69×10−11; pKa2 = 10.329 at 25 °C and ionic strength = 0.0.

The three acidity constants are defined as follows:

pH and composition of carbonic acid solutions

At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2 solution) is completely determined by the partial pressure of carbon dioxide above the solution. To calculate this composition, account must be taken of

  • the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution:
CO2(gas) ⇌ CO2(dissolved) with where kH = 29.76 atm/(mol/L) (Henry constant) at 25 °C
  • the hydration equilibrium between dissolved CO2 and H2CO3 with constant (see above)
  • the first dissociation equilibrium of carbonic acid (see Ka1 above)
  • the second dissociation equilibrium of carbonic acid (see Ka2 above)
  • the dissociation equilibrium of water:
  • the charge neutrality condition

Taken at face value, the above are 6 equations for the 6 unknowns [CO2]aq, [H2CO3], [H+], [OH], [HCO3] and [CO32−]. Note, however, that the first 2 equations express [CO2]aq and [H2CO3] as simple linear functions of , reducing the problem to the latter 4 equations with 4 unknowns. Either way, this demonstrates that the composition of the solution is fully determined by . When isolating [HCO3] in the first dissociation equilibrium, [HCO32−] in the second dissociation equilibrium and [OH] in the dissociation equilibrium of water, then substituting all three in the charge neutrality condition, a cubic equation in [H+] is obtained, whose numerical solution yields the values for the pH and the concentrations of the different species in the following table. (Note that the second dissociation equilibrium can be neglected for this particular problem, reducing the cubic equation to a simple square root; see remarks below the table.)


(atm)
pH [CO2]
(mol/L)
[H2CO3]
(mol/L)
[HCO3]
(mol/L)
[CO32−]
(mol/L)
1.0 × 10−8 7.00 3.36 × 10−10 5.71 × 10−13 1.42 × 1009 7.90 × 10−13
1.0 × 10−7 6.94 3.36 × 1009 5.71 × 10−12 5.90 × 1009 1.90 × 10−12
1.0 × 10−6 6.81 3.36 × 1008 5.71 × 10−11 9.16 × 1008 3.30 × 10−11
1.0 × 10−5 6.42 3.36 × 1007 5.71 × 10−10 3.78 × 1007 4.53 × 10−11
1.0 × 10−4 5.92 3.36 × 1006 5.71 × 1009 1.19 × 1006 5.57 × 10−11
3.5 × 10−4 5.65 1.18 × 1005 2.00 × 1008 2.23 × 1006 5.60 × 10−11
1.0 × 10−3 5.42 3.36 × 1005 5.71 × 1008 3.78 × 1006 5.61 × 10−11
1.0 × 10−2 4.92 3.36 × 1004 5.71 × 1007 1.19 × 1005 5.61 × 10−11
1.0 × 10−1 4.42 3.36 × 1003 5.71 × 1006 3.78 × 1005 5.61 × 10−11
1.0 × 10+0 3.92 3.36 × 1002 5.71 × 1005 1.20 × 1004 5.61 × 10−11
2.5 × 10+0 3.72 8.40 × 1002 1.43 × 1004 1.89 × 1004 5.61 × 10−11
1.0 × 10+1 3.42 3.36 × 1001 5.71 × 1004 3.78 × 1004 5.61 × 10−11
  • In the total range of pressure, the pH is always much lower than pKa2 (= 10.3) so that the CO32− concentration is always negligible with respect to HCO3 concentration. In fact, CO32− plays no quantitative role in the present calculation (see remark below).
  • For vanishing , the pH is close to the one of pure water (pH = 7), and the dissolved carbon is essentially in the HCO3 form.
  • For normal atmospheric conditions ( atm), we get a slightly acidic solution (pH = 5.7), and the dissolved carbon is now essentially in the CO2 and HCO3 forms.
  • For a CO2 pressure typical for bottled carbonated drinks ( ~ 2.5 atm), we get a relatively acidic medium (pH = 3.7) with a high concentration of dissolved CO2. These features contribute to the sour and sparkling taste of these drinks.
  • Between 2.5 and 10 atm, the pH crosses the pKa1 value (3.60), giving [H2CO3] > [HCO3] at high pressures.
  • A plot of the equilibrium concentrations of these different forms of dissolved inorganic carbon (and which species is dominant) as a function of the pH of the solution is known as a Bjerrum plot.
Remark

As noted above, [CO32−] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H+]:

Pure carbonic acid

It was long believed that carbonic acid could not exist as a pure compound and would immediately decompose into its more stable components, water and carbon dioxide. However, in 1991 scientists at NASA's Goddard Space Flight Center (USA) succeeded in making solid H2CO3 samples.[3] They did so by exposing a frozen mixture of water and carbon dioxide to high-energy proton radiation, and then warming to remove the excess water. The carbonic acid that remained was characterized by infrared spectroscopy. The fact that the carbonic acid was prepared by irradiating a solid H2O + CO2 mixture, or even by irradiation of dry ice alone, has given rise to suggestions that H2CO3 might be found in outer space or on Mars, where frozen ices of H2O and CO2 are found, as well as cosmic rays.[11][12] It was announced in 1993 that solid carbonic acid had been created by a cryogenic reaction of potassium bicarbonate and HCl dissolved in methanol.[13][14] Later work showed that in fact the methyl ester had been formed, but other methods were successful.[12] Theoretical calculations showed that a single molecule of water can catalyze the decomposition of a gas-phase carbonic acid molecule to carbon dioxide and water. In the absence of water, the dissociation of gaseous carbonic acid has been predicted to be very slow, with a half-life of 180,000 years.[11] This only applies if the molecules are few and far between, because it has also been predicted that gas-phase carbonic acid will catalyze its own decomposition by forming dimers, which then break apart into two molecules each of water and carbon dioxide.[15]

For a while it was thought that there were two polymorphs of solid carbonic acid, called α and β. The polymorph denoted β-carbonic acid was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuum, which causes protonation of the bicarbonate, followed by removal of the solvent. The previously suggested α-carbonic acid, which was prepared by the same technique using methanol rather than water as a solvent, was later shown to be a monomethyl ester CH3OCOOH.[12]

See also

References

  1. ^ "Front Matter". Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. pp. P001–P004. doi:10.1039/9781849733069-FP001. ISBN 978-0-85404-182-4.
  2. ^ Acid-Base Physiology 2.1 – Acid-Base Balance by Kerry Brandis.
  3. ^ a b M. H. Moore; R. K. Khanna (1990). "Infrared and mass spectral studies of proton irradiated H2O + CO2 ice: Evidence for carbonic acid". Spectrochimica Acta Part A. 47 (2): 255–262. Bibcode:1991AcSpA..47..255M. doi:10.1016/0584-8539(91)80097-3.
  4. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 310. ISBN 978-0-08-037941-8.
  5. ^ Housecroft and Sharpe, Inorganic Chemistry, 2nd ed, Prentice-Pearson-Hall 2005, p. 368.
  6. ^ Soli, A. L.; R. H. Byrne (2002). "CO2 system hydration and dehydration kinetics and the equilibrium CO2/H2CO3 ratio in aqueous NaCl solution". Marine Chemistry. 78 (2–3): 65–73. doi:10.1016/S0304-4203(02)00010-5.
  7. ^ "excretion". Encyclopædia Britannica. Encyclopædia Britannica Ultimate Reference Suite. Chicago: Encyclopædia Britannica, 2010.
  8. ^ "Reaction Mechanism for Direct Proton Transfer from Carbonic Acid to a Strong Base in Aqueous Solution I: Acid and Base Coordinate and Charge Dynamics", S. Daschakraborty, P. M. Kiefer, Y. Miller, Y. Motro, D. Pines, E. Pines, and J. T. Hynes. J. Phys. Chem. B (2016), 120, 2271.
  9. ^ Sabine, C. L.; et al. (2004). "The Oceanic Sink for Anthropogenic CO2". Science. 305 (5682): 367–371. Bibcode:2004Sci...305..367S. doi:10.1126/science.1097403. hdl:10261/52596. PMID 15256665. Archived from " the original on July 6, 2008.
  10. ^ National Research Council. "Summary". Ocean Acidification: A National Strategy to Meet the Challenges of a Changing Ocean. Washington, DC: The National Academies Press, 2010. 1. Print.
  11. ^ a b Loerting, T.; Tautermann, C.; Kroemer, R. T.; Kohl, I.; Hallbrucker, E.; Mayer, A.; Liedl, K. R. (2000). "On the Surprising Kinetic Stability of Carbonic Acid". Angew. Chem. Int. Ed. 39 (5): 891–895. doi:10.1002/(SICI)1521-3773(20000303)39:5<891::AID-ANIE891>3.0.CO;2-E. PMID 10760883.
  12. ^ a b c Reisenauer, H. P.; Wagner, J. P.; Schreiner, P. R. (2014). "Gas-Phase Preparation of Carbonic Acid and Its Monomethyl Ester". Angew. Chem. Int. Ed. 53 (44): 11766–11771. doi:10.1002/anie.201406969. PMID 25196920.
  13. ^ Hage, W.; Hallbrucker, A.; Mayer, E. (1993). "Carbonic Acid: Synthesis by Protonation of Bicarbonate and FTIR Spectroscopic Characterization Via a New Cryogenic Technique". J. Am. Chem. Soc. 115 (18): 8427–8431. doi:10.1021/ja00071a061.
  14. ^ "Press release: International First: Gas-phase Carbonic Acid Isolated". Technische Universität Wien. 11 January 2011. Archived from the original on 9 August 2017. Retrieved 9 August 2017.
  15. ^ de Marothy, S. A. (2013). "Autocatalytic decomposition of carbonic acid". Int. J. Quantum Chem. 113 (20): 2306–2311. doi:10.1002/qua.24452.

Further reading

External links

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.

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 HCO−3.

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 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. 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 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.

Bjerrum plot

A Bjerrum plot (named after Niels Bjerrum) is a graph of the concentrations of the different species of a polyprotic acid in a solution, as functions of the solution's pH, when the solution is at equilibrium. Due to the many orders of magnitude spanned by the concentrations, they are commonly plotted on a logarithmic scale. Sometimes the ratios of the concentrations are plotted rather than the actual concentrations. Occasionally H+ and OH− are also plotted.

Most often, the carbonate system is plotted, where the polyprotic acid is carbonic acid (a diprotic acid), and the different species are carbonic acid, carbon dioxide, bicarbonate, and carbonate. In acidic conditions, the dominant form is CO2; in basic (alkalinic) conditions, the dominant form is CO32−; and in between, the dominant form is HCO3−. At every pH, the concentration of carbonic acid is assumed to be negligible compared to the concentration of CO2, and so is often omitted from Bjerrum plots. These plots are typically used in ocean chemistry to track the response of an ocean to changes in both pH and of inputs in carbonate and CO2.The Bjerrum plots for other polyprotic acids, including silicic, boric, sulfuric and phosphoric acids, can also be constructed.

Carbon dioxide

Carbon dioxide (chemical formula CO2) is a colorless gas with a density about 60% higher than that of dry air. Carbon dioxide consists of a carbon atom covalently double bonded to two oxygen atoms. It occurs naturally in Earth's atmosphere as a trace gas. The current concentration is about 0.04% (410 ppm) by volume, having risen from pre-industrial levels of 280 ppm. Natural sources include volcanoes, hot springs and geysers, and it is freed from carbonate rocks by dissolution in water and acids. Because carbon dioxide is soluble in water, it occurs naturally in groundwater, rivers and lakes, ice caps, glaciers and seawater. It is present in deposits of petroleum and natural gas. Carbon dioxide is odorless at normally encountered concentrations, but at high concentrations, it has a sharp and acidic odor.As the source of available carbon in the carbon cycle, atmospheric carbon dioxide is the primary carbon source for life on Earth and its concentration in Earth's pre-industrial atmosphere since late in the Precambrian has been regulated by photosynthetic organisms and geological phenomena. Plants, algae and cyanobacteria use light energy to photosynthesize carbohydrate from carbon dioxide and water, with oxygen produced as a waste product.CO2 is produced by all aerobic organisms when they metabolize carbohydrates and lipids to produce energy by respiration. It is returned to water via the gills of fish and to the air via the lungs of air-breathing land animals, including humans. Carbon dioxide is produced during the processes of decay of organic materials and the fermentation of sugars in bread, beer and wine making. It is produced by combustion of wood and other organic materials and fossil fuels such as coal, peat, petroleum and natural gas. It is an unwanted byproduct in many large scale oxidation processes, for example, in the production of acrylic acid (over 5 million tons/year).It is a versatile industrial material, used, for example, as an inert gas in welding and fire extinguishers, as a pressurizing gas in air guns and oil recovery, as a chemical feedstock and as a supercritical fluid solvent in decaffeination of coffee and supercritical drying. It is added to drinking water and carbonated beverages including beer and sparkling wine to add effervescence. The frozen solid form of CO2, known as dry ice is used as a refrigerant and as an abrasive in dry-ice blasting.

Carbon dioxide is the most significant long-lived greenhouse gas in Earth's atmosphere. Since the Industrial Revolution anthropogenic emissions – primarily from use of fossil fuels and deforestation – have rapidly increased its concentration in the atmosphere, leading to global warming. Carbon dioxide also causes ocean acidification because it dissolves in water to form carbonic acid.

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.

Carbonate ester

A carbonate ester (organic carbonate or organocarbonate) is an ester of carbonic acid. This functional group consists of a carbonyl group flanked by two alkoxy groups. The general structure of these carbonates is R1O(C=O)OR2 and they are related to esters R1O(C=O)R and ethers R1OR2 and also to the inorganic carbonates.

Monomers of polycarbonate (e.g. Lexan) are linked by carbonate groups. These polycarbonates are used in eyeglass lenses, compact discs, and bulletproof glass. Small carbonate esters like dimethyl carbonate, ethylene carbonate, propylene carbonate are used as solvents. Dimethyl carbonate is also a mild methylating agent.

Carbonation

Carbonation is the chemical reaction of carbon dioxide to give carbonates, bicarbonates, and carbonic acid. In chemistry, the term is sometimes used in place of carboxylation, which refers to the formation of carboxylic acids.

In inorganic chemistry and geology, carbonation is common. Metal hydroxides (MOH) and metal oxides (M'O) react with CO2 to give bicarbonates and carbonates:

MOH + CO2 → M(HCO3)

M'O + CO2 → M'CO3In reinforced concrete construction, the chemical reaction between carbon dioxide in the air and calcium hydroxide and hydrated calcium silicate in the concrete is known as neutralisation.

Dimethyl carbonate

Dimethyl carbonate (DMC) is an organic compound with the formula OC(OCH3)2. It is a colourless, flammable liquid. It is classified as a carbonate ester. This compound has found use as a methylating agent and more recently as a solvent that is exempt from the restrictions placed on most volatile organic compounds (VOCs) in the US. Dimethyl carbonate is often considered to be a green reagent.

Effervescence

Effervescence is the escape of gas from an aqueous solution and the foaming or fizzing that results from that release. The word effervescence is derived from the Latin verb fervere (to boil), preceded by the adverb ex. It has the same linguistic root as the word fermentation.

Effervescence can also be observed when opening a bottle of champagne, beer or carbonated beverages such as soft drinks. The visible bubbles are produced by the escape from solution of the dissolved gas (which itself is not visible while dissolved in the liquid).

Although CO2 is most common for beverages, nitrogen gas is sometimes deliberately added to certain beers. The smaller bubble size creates a smoother beer head. Due to the poor solubility of nitrogen in beer, kegs or widgets are used for this.

In the laboratory, a common example of effervescence is seen if hydrochloric acid is added to a block of limestone. If a few pieces of marble or an antacid tablet are put in hydrochloric acid in a test tube fitted with a bung, effervescence of carbon dioxide can be witnessed.

This process is generally represented by the following reaction, where a pressurized dilute solution of carbonic acid in water releases gaseous carbon dioxide at decompression:

In simple terms, it is the result of the chemical reaction occurring in the liquid which produces a gaseous product.

Keto acid

Keto acids or ketoacids (also called oxo acids or oxoacids) are organic compounds that contain a carboxylic acid group and a ketone group. In several cases, the keto group is hydrated. The alpha-keto acids are especially important in biology as they are involved in the Krebs citric acid cycle and in glycolysis.Common types of keto acids include:

Alpha-keto acids, Alpha-ketoacids, or 2-oxoacids, such as pyruvic acid, have the keto group adjacent to the carboxylic acid. One important alpha-keto acid is oxaloacetic acid, a component of the Krebs cycle. Another is alpha-ketoglutarate, a 5-carbon ketoacid derived from glutamic acid. Alpha-ketoglutarate participates in cell signaling by functioning as a coenzyme, and is commonly used in transamination reactions. Alpha-keto acids possesses extensive chemistry as acylation agents.

Beta-keto acids, Beta-ketoacids, or 3-oxoacids, such as acetoacetic acid, have the ketone group at the second carbon from the carboxylic acid. They can be formed by the Claisen condensation.

Gamma-keto acids, Gamma-ketoacids, or 4-oxoacids, such as levulinic acid, have the ketone group at the third carbon from the carboxylic acid.Keto acids appear in a wide variety of anabolic pathways in metabolism, across living organisms. For instance, in plants (specifically, in hemlock, pitcher plants, and fool's parsley), 5-oxo-octanoic acid is converted in enzymatic and non-enzymatic steps into the cyclic class of coniine alkaloids.When ingested sugars and carbohydrate levels are low, stored fats and proteins are the primary source of energy production. Glucogenic amino acids from proteins are converted to glucose. Ketogenic amino acids can be deaminated to produce alpha keto acids and ketone bodies.

Alpha keto acids are used primarily as energy for liver cells and in fatty acid synthesis, also in the liver.

One Piece (season 20)

The twentieth season of the One Piece anime series was produced by Toei Animation, and directed by Tatsuya Nagamine. The season began broadcasting in Japan on Fuji Television on July 7, 2019. Like the rest of the series, it follows the adventures of Monkey D. Luffy and his Straw Hat Pirates. The story arc, called "Wano Country", adapts material from the rest of the 90th volume onwards. It deals with the alliance between the pirates, samurai, and minks to liberate Wano Country from the shogun, who has allied with the Beast Pirates led by Kaido. Episodes 895–896 contain an original story arc, "Carbonic Acid King" which ties into the movie One Piece: Stampede. Episode 907 is a special of the second one-shot, which is a predecessor to the series.

The opening theme music for this season is "Over The Top" (オーバーザトップ, Ōbā Za Toppu) by Hiroshi Kitadani.

Phosgene

Phosgene is the organic chemical compound with the formula COCl2. It is a colorless gas; in low concentrations, its odor resembles freshly cut hay or grass. Phosgene is a valued industrial building block, especially for the production of urethanes and polycarbonate plastics. However, it is very poisonous and was used as a chemical weapon during World War I where it was responsible for 85,000 deaths. In addition to its industrial production, small amounts occur from the breakdown and the combustion of organochlorine compounds.

Propylene carbonate

Propylene carbonate (often abbreviated PC) is an organic compound with the formula C4H6O3. It is a cyclic carbonate ester derived from propylene glycol. This colorless and odorless liquid is useful as a polar, aprotic solvent. Propylene carbonate is chiral, but is used exclusively as the racemic mixture in most contexts.

Solutional cave

A solutional cave or karst cave is a cave usually formed in the soluble rock limestone. It is the most frequently occurring type of cave. It can also form in other rocks, including chalk, dolomite, marble, salt beds, and gypsum.

Speleogenesis

Speleogenesis is the origin and development of caves, the primary process that determines essential features of the hydrogeology of karst and guides its evolution. It often deals with the development of caves through limestone, caused by the presence of water with carbon dioxide dissolved within it, producing carbonic acid which permits the dissociation of the calcium carbonate in the limestone.

Svante Arrhenius

Svante August Arrhenius (; 19 February 1859 – 2 October 1927) was a Swedish scientist. Originally a physicist, but often referred to as a chemist, Arrhenius was one of the founders of the science of physical chemistry. He received the Nobel Prize for Chemistry in 1903, becoming the first Swedish Nobel laureate. In 1905, he became director of the Nobel Institute, where he remained until his death.Arrhenius was the first to use basic principles of physical chemistry to estimate the extent to which increases in atmospheric carbon dioxide are responsible for the Earth's increasing surface temperature. In the 1960s, David Keeling demonstrated that human-caused carbon dioxide emissions are large enough to cause global warming.Arrhenius' contributions to science are memorialized by the Arrhenius equation, Arrhenius acid, lunar crater Arrhenius, Martian crater Arrhenius, the mountain of Arrheniusfjellet, and the Arrhenius Labs at Stockholm University.

Total inorganic carbon

The total inorganic carbon (CT, or TIC) or dissolved inorganic carbon (DIC) is the sum of inorganic carbon species in a solution. The inorganic carbon species include carbon dioxide, carbonic acid, bicarbonate anion, and carbonate. It is customary to express carbon dioxide and carbonic acid simultaneously as CO2* . CT is a key parameter when making measurements related to the pH of natural aqueous systems, and carbon dioxide flux estimates.

CT = [CO2*] + [HCO3−] + [CO32−]where,

CT is the total inorganic carbon

[CO2*] is the sum of carbon dioxide and carbonic acid concentrations ( [CO2*] = [CO2] + [H2CO3])

[HCO3−] is the bicarbonate concentration

[CO32−] is the carbonate concentrationEach of these species are related by the following pH-driven chemical equilibria:

CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3− ⇌ 2H+ + CO32−The concentrations of the different species of DIC (and which species is dominant) depends on the pH of the solution, as shown by a Bjerrum plot.

Total inorganic carbon is typically measured by the acidification of the sample which drives the equilibria to CO2. This gas is then sparged from solution and trapped, and the quantity trapped is then measured, usually by infrared spectroscopy.

Weerman degradation

Weerman degradation, also named Weerman reaction, is a name reaction in organic chemistry. It is named after Rudolf Adrian Weerman, who discovered it in 1910. In general, it is an organic reaction in carbohydrate chemistry in which amides are degraded by sodium hypochlorite, forming an aldehyde with one less carbon. Some have regarded it as an extension of the Hofmann rearrangement.

Hydrogen compounds
Common oxides
Exotic oxides
Polymers
Compounds derived from oxides
Compounds
Carbon ions
Oxides and related

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