Carbonic acid

Names
IUPAC name Carbonic acid[1]
Other names Hydroxyformic acid Hydroxymethanoic acid Dihydroxycarbonyl
Identifiers
CAS Number	463-79-6 Y
3D model (JSmol)	Interactive image
ChEBI	CHEBI:28976 Y
ChEMBL	ChEMBL1161632 Y
ChemSpider	747 Y
DrugBank	DB14531
ECHA InfoCard	100.133.015
EC Number	610-295-3
Gmelin Reference	25554
KEGG	C01353 Y
PubChem CID	767
CompTox Dashboard (EPA)	DTXSID9043801
show InChI InChI=1S/CH2O3/c2-1(3)4/h(H2,2,3,4) Y Key: BVKZGUZCCUSVTD-UHFFFAOYSA-N Y InChI=1/H2O3/c2-1(3)4/h(H2,2,3,4) Key: BVKZGUZCCUSVTD-UHFFFAOYAU
show SMILES O=C(O)O
Properties
Chemical formula	H2CO3
Melting point	−80 °C (−112 °F; 193 K) (decomposes)
Conjugate base	Bicarbonate, Carbonate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Y  (what is YN ?)
Infobox references

In chemistry, carbonic acid is a dibasic acid with the chemical formula H₂CO₃. The pure compound decomposes at temperatures greater than ca. −80 °C.^[2]

In biochemistry, the name "carbonic acid" is often applied to aqueous solutions of carbon dioxide, which play an important role in the bicarbonate buffer system, used to maintain acid–base homeostasis.^[3]

Chemical equilibria[]

Equilibrium constant values[]

Bjerrum plot for carbonate speciation in seawater (ionic strength 0.7 mol dm⁻³)

In aqueous solution carbonic acid behaves as a dibasic acid. The Bjerrum plot shows typical equilibrium concentrations, in solution, in seawater, of carbon dioxide and the various species derived from it, as a function of pH.^[4] The acidification of natural waters is caused by the increasing concentration of carbon dioxide in the atmosphere, which is believed to be caused by the burning of increasing amounts of coal and hydrocarbons.^[5][6]

Expected change refers to predicted effect of continued ocean acidification.^[7] It has been estimated that the increase in dissolved carbon dioxide has caused the ocean's average surface pH to decrease by about 0.1 from pre-industrial levels.

The stability constants database contains 136 entries with values for the overall protonation constants, β₁ and β₂, of the carbonate ion. In the following expressions [H⁺] represents the concentration, at equilibrium, of the chemical species H⁺, etc.

The value of log β₁ decreases with increasing ionic strength, ${\displaystyle {\ce {I}}}$. At 25 °C:

${\displaystyle {\ce {CO3^{2-}{+ H+}<=> HCO3^-}}}$ : ${\displaystyle \beta _{1}={\frac {[{\text{HCO}}_{3}^{-}]}{[{\text{H}}^{+}][{\text{CO}}_{3}^{2-}]}}}$

${\displaystyle log\beta _{1}=0.54I^{2}-0.96I+9.93}$ (selected data from SC-database)

The value of log β₂ also decreases with increasing ionic strength.

${\displaystyle {\ce {CO3^{2-}{+ 2H+}<=> H2CO3}}}$ : ${\displaystyle \beta _{2}={\frac {[{\text{H}}_{2}{\text{CO}}_{3}]}{[{\text{H}}^{+}]^{2}[{\text{CO}}_{3}^{2-}]}}}$

${\displaystyle log\beta _{2}=-2.5I^{2}-0.043I+16.07}$

At ${\displaystyle {\ce {I}}}$=0 and 25 °C the pK values of the stepwise dissociation constants are

pK₁ = logβ₂ - logβ₁ = 6.77.

pK₂ = logβ₁ = 9.93.

When pH = pK the two chemical species in equilibrium with each other have the same concentration.

Note 1: There are apparently conflicting values in the literature for pK_a. Pines et al. cite a value for "pK_app" of 6.35, consistent with the value 6.77, mentioned above.^[8] They also give a value for "pK_a" of 3.49 and state that

pK_a = pK_app − log K_D (eqn. 5)

where K_D=[CO₂]/[H₂CO₃]. (eqn. 3) The situation arises from the way that the dissociation constants are named and defined, which is clearly stated in the text of the Pines paper, but not in the abstract.

Note 2: The numbering of dissociation constants is the reverse of the numbering of the numbering of association constants, so pK₂ (dissociation)= log β₁ (association). The value of the stepwise constant for the equilibrium

${\displaystyle {\ce {HCO3- <=> CO3^{2-}{+ H+}}}}$

is given by

pK₁(dissociation)₁ = log β₂ − log β₁ (association)

In non-biological solutions[]

The hydration equilibrium constant at 25 °C is called K_h, which in the case of carbonic acid is [H₂CO₃]/[CO₂] ≈ 1.7×10⁻³ in pure water^[9] and ≈ 1.2×10⁻³ in seawater.^[10] 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 and 23 s⁻¹ for the reverse reaction.

In nature, limestone may react with rainwater, forming a solution of calcium bicarbonate; evaporation of such a solution will result re-formation of solid calcium carbonate. These processes occur in the formation of stalactites and stalagmites.

In biological solutions[]

When the enzyme carbonic anhydrase is also present in the solution the following reaction takes precedence.^[11]

${\displaystyle {\ce {HCO3^- {+}H^+ <=> CO2 {+}H2O}}}$

When the amount of carbon dioxide created by the forward reaction exceeds its solubility, gas is evolved and a third equilibrium

${\displaystyle {\ce {CO_2 (soln) <=> CO_2 (g)}}}$

must also be taken into consideration. The equilibrium constant for this reaction is defined by Henry's law. The two reactions can be combined for the equilibrium in solution.

${\displaystyle {\ce {HCO3^- {+}H^{+}<=> CO_2(soln) {+}H2O}}}$ :    ${\displaystyle {\ce {K_3 = {\frac {[H^+][HCO_3^{-}]}{[CO_2(soln)]}}}}}$

When Henry's law is used to calculate the value of the term in the denominator care is needed with regard to dimensionality.

In physiology, carbon dioxide excreted by the lungs may be called volatile acid or respiratory acid.

Use of the term carbonic acid[]

Speciation for a monoprotic acid, AH as a function of pH.

Strictly speaking the term "carbonic acid" refers to the chemical compound with the formula ${\displaystyle {\ce {H2CO3}}}$.

Since pK_a1 has a value of ca. 6.8 , at equilibrium carbonic acid will be almost 50% dissociated in the extracellular fluid (cytosol) which has a pH of ca.7.2. Note that dissolved carbon dioxide in extracellular fluid is often called as "carbonic acid" in biochemistry literature, for historical reasons. The reaction in which it is produced

HCO₃⁻ + H⁺ ⇌ CO₂ + H₂O

is fast in biological systems. Carbon dioxide can be described as the anhydride of carbonic acid.

Pure carbonic acid[]

Carbonic acid, H₂CO₃, is stable at ambient temperatures in strictly anhydrous conditions. It decomposes to form carbon dioxide in the presence of any water molecules.^[12]

Carbonic acid forms as a by-product of CO₂/H₂O irradiation, in addition to carbon monoxide and radical species (HCO and CO₃).^[2] Another route to form carbonic acid is protonation of bicarbonates (HCO₃⁻) with aqueous HCl or HBr. This has to be done at cryogenic conditions to avoid immediate decomposition of H₂CO₃ to CO₂ and H₂O.^[13] Amorphous H₂CO₃ forms above 120 K, and crystallization takes place above 200 K to give "β-H₂CO₃", as determined by infrared spectroscopy. The spectrum of β-H₂CO₃ agrees very well with the by-product after CO₂/H₂O irradiation.^[2] β-H₂CO₃ sublimes at 230–260 K largely without decomposition. Matrix-isolation infared spectroscopy allows for the recording of single molecules of H₂CO₃.^[14]

The fact that the carbonic acid may form by irradiating a solid H₂O + CO₂ mixture or even by proton-implantation of dry ice alone^[15] 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.^[12] The surprising stability of sublimed H₂CO₃ up to rather high-temperatures of 260 K even allows for gas-phase H₂CO₃, e.g., above the pole caps of Mars.^[14] Ab initio calculations showed that a single molecule of water catalyzes 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 is predicted to be very slow, with a half-life in the gas-phase of 180,000 years at 300 K.^[12] This only applies if the molecules are few and far apart, 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.^[16]

Solid "α-carbonic acid" was claimed to be generated by a cryogenic reaction of potassium bicarbonate and a solution of HCl in methanol.^[17][18] This claim was disputed in a PhD thesis submitted in January 2014.^[19] Instead, isotope labeling experiments point to the involvement of carbonic acid monomethyl ester (CAME). Furthermore, the sublimed solid was suggested to contain CAME monomers and dimers, not H₂CO₃ monomers and dimers as previously claimed.^[20] Subsequent matrix-isolation infrared spectra confirmed that CAME rather than carbonic acid is found in the gas-phase above "α-carbonic acid".^[21] The assignment as CAME is further corroborated by matrix-isolation of the substance prepared in gas-phase by pyrolysis.^[22]

Despite its complicated history, carbonic acid may still appear as distinct polymorphs. Carbonic acid forms upon oxidization of CO with OH-radicals.^[23] It is not clear whether carbonic acid prepared in this way needs to be considered as γ-H₂CO₃. The structures of β-H₂CO₃ and γ-H₂CO₃ have not been characterized crystallographically.

At high pressure[]

Although molecules of H₂CO₃ do not constitute a significant portion of the dissolved carbon in aqueous "carbonic acid" under ambient conditions, significant amounts of molecular H₂CO₃ can exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres), such as can occur in planetary interiors.^[24][25]

Carbonic acid should be stabilized under pressures of 0.6–1.6 GPa at 100 K, and 0.75–1.75 GPa at 300 K. These pressures are attained in the cores of large icy satellites such as Ganymede, Callisto, and Titan, where water and carbon dioxide are present. Pure carbonic acid, being denser, would then sink under the ice layers and separate them from the rocky cores of these moons.^[26]

References[]

Further reading[]

•  "Climate and Carbonic Acid" in Popular Science Monthly Volume 59, July 1901
• Welch, M. J.; Lifton, J. F.; Seck, J. A. (1969). "Tracer studies with radioactive oxygen-15. Exchange between carbon dioxide and water". J. Phys. Chem. 73 (335): 3351. doi:10.1021/j100844a033.
• Jolly, W. L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 978-0-07-112651-9.
• Moore, M. H.; Khanna, R. (1991). "Infrared and Mass Spectral Studies of Proton Irradiated H2O+Co2 Ice: Evidence for Carbonic Acid Ice: Evidence for Carbonic Acid". Spectrochimica Acta. 47A (2): 255–262. Bibcode:1991AcSpA..47..255M. doi:10.1016/0584-8539(91)80097-3.
• W. Hage, K. R. Liedl; Liedl, E.; Hallbrucker, A; Mayer, E (1998). "Carbonic Acid in the Gas Phase and Its Astrophysical Relevance". Science. 279 (5355): 1332–1335. Bibcode:1998Sci...279.1332H. doi:10.1126/science.279.5355.1332. PMID 9478889.
• Hage, W.; Hallbrucker, A.; Mayer, E. (1995). "A Polymorph of Carbonic Acid and Its Possible Astrophysical Relevance". J. Chem. Soc. Faraday Trans. 91 (17): 2823–2826. Bibcode:1995JCSFT..91.2823H. doi:10.1039/ft9959102823.

External links[]

• Carbonic acid/bicarbonate/carbonate equilibrium in water: pH of solutions, buffer capacity, titration and species distribution vs. pH computed with a free spreadsheet
• How to calculate concentration of Carbonic Acid in Water