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.[1]

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.

The long-held belief that carbonic acid could not exist as a pure compound has reportedly been recently disproved by the preparation of the pure substance in both solid and gas form by University of Innsbruck researchers.[2]

Chemical equilibriumEdit

When carbon dioxide dissolves in water it exists in chemical equilibrium producing carbonic acid:[3]

CO2 + H2O 15px H2CO3

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[4] and ≈ 1.2×10−3 in seawater.[5] 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). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. 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 Edit

Carbonic acid is an intermediate step in the transport of CO2 out of the body via 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 both increases the reaction rate and dissociates a hydrogen ion (H+) from the resulting carbonic acid, leaving 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.[6]

Role of carbonic acid in ocean chemistryEdit

The oceans of the world have absorbed almost half of the CO2 emitted by humans from the burning of fossil fuels.[7]  The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels.[8] This process is known as ocean acidification.[9]

Acidity of carbonic acidEdit

Carbonic acid is one of the polyprotic acids: It is diprotic - it has two protons, which may dissociate from the parent molecule. Thus, there are two dissociation constants, the first one for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3:

H2CO3 15px HCO3 + H+
Ka1 = 2.5×10−4;[3] 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:[3]

H2CO3* 15px HCO3 + H+
Ka(app) = 4.6×10−7; pK(app) = 6.3 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−:

HCO3 15px CO32− + H+
Ka2 = 4.69×10−11; pKa2 = 10.329 at 25 °C and ionic strength = 0.0

The three acidity constants are defined as follows: $ K_{a1}=\frac{[H^+][HCO_3^-]}{[H_2CO_3]} \qquad K_a{(app)}=\frac{[H^+][HCO_3^-]}{[H_2CO_3]+[CO_2(aq)]} \qquad K_{a2}=\frac{[H^+][CO_3^{2-}]}{[HCO_3^-]} $

pH and composition of carbonic acid solutionsEdit

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 $ \scriptstyle p_{CO_2} $ of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2CO3, HCO3 and CO32−) as well as of the hydration equilibrium between dissolved CO2 and H2CO3 with constant $ \scriptstyle K_h=\frac{[H_2CO_3]}{[CO_2]} $ (see above) and of the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution:

CO2(gas) 15px CO2(dissolved) with $ \scriptstyle \frac{[CO_2]}{p_{CO_2}}=\frac{1}{k_\mathrm{H}} $ where kH=29.76 atm/(mol/L) at 25 °C (Henry constant)

The corresponding equilibrium equations together with the $ \scriptstyle[H^+][OH^-]=10^{-14} $ relation and the charge neutrality condition $ \scriptstyle[H^+]=[OH^-]+[HCO_3^-]+2[CO_3^{2-}] $ result in six equations for the six unknowns [CO2], [H2CO3], [H+], [OH], [HCO3] and [CO32−], showing that the composition of the solution is fully determined by $ \scriptstyle p_{CO_2} $. The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

$ \scriptstyle p_{CO_2} $
pH [CO2]
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
  • We see that in the total range of pressure, the pH is always largely lower than pKa2 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 $ \scriptstyle p_{CO_2} $, the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO3 form.
  • For normal atmospheric conditions ($ \scriptstyle p_{CO_2}=3.5\times 10^{-4} $ atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO2 form. From this pressure on, [OH] becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H+ and HCO3.
  • For a CO2 pressure typical of the one in soda drink bottles ($ \scriptstyle p_{CO_2} $ ~ 2.5 atm), we get a relatively acid 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 a dominant H2CO3 concentration (with respect to 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.


As noted above, [CO32−] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H+]:
$ \scriptstyle[H^+] \simeq \left( 10^{-14}+\frac {K_hK_{a1}}{k_\mathrm{H}} p_{CO_2}\right)^{1/2} $

Spectroscopic studies of carbonic acidEdit

Theoretical calculations show that the presence of even a single molecule of water causes carbonic acid to revert 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 of 180,000 years.[10] This may only apply to isolated carbonic acid molecules however, as it has been predicted to catalyze its own decomposition[11]

It has long been recognized that pure carbonic acid cannot be obtained at room temperatures (about 20 °C or about 70 °F). It can be generated by exposing a frozen mixture of water and carbon dioxide to high-energy 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 may suggest that H2CO3 might be found in outer space, where frozen ices of H2O and CO2 are common, as are cosmic rays and ultraviolet light, to help them react.[10] The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuo, which causes protonation of bicarbonate, followed by removal of the solvent. The previously suggested alpha-carbonic acid, which was prepared by the same technique using methanol rather than water as a solvent was shown to be a monomethyl ester CH3OCOOH.[12]

See alsoEdit

References Edit

  1. Acid-Base Physiology 2.1 - Acid-Base Balance by Kerry Brandis
  2. "International first: Gas-phase carbonic acid isolated",, Jan 11, 2011
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  4. Housecroft and Sharpe, Inorganic Chemistry, 2nd ed, Prentice-Pearson-Hall 2005, p.368.
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  6. "excretion." Encyclopædia Britannica. Encyclopædia Britannica Ultimate Reference Suite. Chicago: Encyclopædia Britannica, 2010.
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  9. 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.
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Further readingEdit

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External links Edit

Template:Hydrogen compounds

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