In chemistry, carbonic acid is an organic compound with the chemical formula H2CO3. The molecule rapidly converts to water and carbon dioxide in the presence of water. However, in the absence of water, it is (contrary to popular belief) quite stable at room temperature.[4][5] The interconversion of carbon dioxide and carbonic acid is related to the breathing cycle of animals and the acidification of natural waters.[3]
In chemistry, the term "carbonic acid" strictly refers to the chemical compound with the formula H 2CO 3. Some biochemistry literature effaces the distinction between carbonic acid and carbon dioxide dissolved in extracellular fluid.
In physiology, carbon dioxide excreted by the lungs may be called volatile acid or respiratory acid.
According to neutron diffraction of dideuterated carbonic acid (D 2CO 3) in a hybrid clamped cell (Russian alloy/copper-beryllium) at 1.85 GPa, the molecules are planar and form dimers joined by pairs of hydrogen bonds. All three C-O bonds are nearly equidistant at 1.34Å, intermediate between typical C-O and C=O distances (respectively 1.43 and 1.23Å). The unusual C-O bond lengths are attributed to delocalized π bonding in the molecule's center and extraordinarily strong hydrogen bonds. The same effects also induce a very short O—O separation (2.13Å), through the 136° O-H-O angle imposed by the doubly hydrogen-bonded 8-membered rings.[3] Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid, where the distances exceed 2.4Å.[10]
The hydrationequilibrium constant at 25 °C is [H 2CO 3]/[CO2] ≈ 1.7×10−3 in pure water[11] and ≈ 1.2×10−3 in seawater.[12] Hence the majority of carbon dioxide at geophysical or biological air-water interfaces does not convert to carbonic acid, remaining dissolved CO2 gas. However, the uncatalyzed equilibrium is reached quite slowly: the rate constants are 0.039 s−1 for hydration and 23 s−1 for dehydration.
In biological solutions
In the presence of the enzyme carbonic anhydrase, equilibrium is instead reached rapidly, and the following reaction takes precedence:[13]
When the created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium
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:
When Henry's law is used to calculate the denominator care is needed with regard to units since Henry's Law constant can be commonly expressed with 8 different dimensionalities.[14]
Significant amounts of molecular H 2CO 3 exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres) in planetary interiors.[15][16] Pressures of 0.6–1.6 GPa at 100 K, and 0.75–1.75 GPa at 300 K 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, is expected to have sunk under the ice layers and separate them from the rocky cores of these moons.[17]
Relationship to bicarbonate and carbonate
Bjerrum plot of speciation for a hypothetical monoprotic acid: AH concentration as a function of the difference between pK and pH
where brackets indicate the concentration of specie. At 25 °C, these equilibria empirically satisfy[18]
Note that log(β1) decreases with increasing I, as does log(β2). In a solution absent other ions (e.g. I = 0), these curves imply the following stepwise dissociation constants:
Direct values for these constants in the literature include pK1 = 6.35 and pK2 - pK1 = 3.49.[19]
To interpret these numbers, note that two chemical species in an acid equilibrium are equiconcentrated when pK = pH. In particular, the extracellular fluid (cytosol) in biological systems exhibits pH ≈ 7.2, so that carbonic acid will be almost 50%-dissociated at equilibrium.
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.[7][8] As human industrialization has increased the proportion of carbon dioxide in Earth's atmosphere, the proportion of carbon dioxide dissolved in sea- and freshwater as carbonic acid is also expected to increase. This rise in dissolved acid is also expected to acidify those waters, generating a decrease in pH.[20][21] It has been estimated that the increase in dissolved carbon dioxide has already caused the ocean's average surface pH to decrease by about 0.1 from pre-industrial levels.
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 ed.). McGraw-Hill. ISBN978-0-07-112651-9.
^ abWinkel, Katrin; Hage, Wolfgang; Loerting, Thomas; Price, Sarah L.; Mayer, Erwin (2007). "Carbonic Acid: From Polyamorphism to Polymorphism". Journal of the American Chemical Society. 129 (45): 13863–71. doi:10.1021/ja073594f. PMID17944463.
^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.
![{\displaystyle {\begin{aligned}{\ce {HCO3^-{}+ H+{}<=> CO2(soln){}+ H2O))&&K_{3}={\frac {[{\ce {H+))][{\ce {HCO3^-))]}{[{\ce {CO2(soln)))]))\end{aligned))}](https://wikimedia.org/api/rest_v1/media/math/render/svg/d5b661a568f432e0b2827ba6de39431e96ea1f69)
![{\displaystyle {\begin{aligned}{\ce {CO3^{2-}{}+ H+{}<=> HCO3^-))&&\beta _{1}={\frac {[{\ce {HCO3^-))]}{[{\ce {H+))][{\ce {CO3^{2-))}]))\\{\ce {CO3^{2-}{}+ 2H+{}<=> H2CO3))&&\beta _{2}={\frac {[{\ce {H2CO3))]}{[{\ce {H+))]^{2}[{\ce {CO3^{2-))}]))\end{aligned))}](https://wikimedia.org/api/rest_v1/media/math/render/svg/09f8d396f1143f199cf1136f1ad60c7ce873b848)
