Carbonic acid

Carbonic acid

Names
Other names Carbon dioxide solution, dihydrogen carbonate (IUPAC)
Identifiers
CAS Number	463-79-6 N
3D model (JSmol)	Interactive image
ChemSpider	747
ECHA InfoCard	100.133.015
CompTox Dashboard (EPA)	DTXSID9043801
InChI InChI=1/CH2O3/c2-1(3)4/h(H2,2,3,4) Key: BVKZGUZCCUSVTD-UHFFFAOYAU
SMILES O=C(O)O
Properties
Chemical formula	H2CO3
Molar mass	62.03 g/mol
Density	1.0 g/cm3 (dilute soln.)
Melting point	n/a
Solubility in water	Exists only in solution
Acidity (pKa)	see acidity of carbonic acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Y verify (what is YN ?) Infobox references

Carbonic acid (ancient name acid of air or aerial acid) has the formula H₂CO₃. It is also a name sometimes given to solutions of carbon dioxide in water, which contain small amounts of H₂CO₃. The salts of carbonic acids are called bicarbonates (or hydrogen carbonates) and carbonates. It is a weak acid. Carbonic acid should never be confused with carbolic acid, an antiquated name for phenol.

Carbon dioxide dissolved in water is in equilibrium with carbonic acid:

CO₂ + H₂O ⇌ H₂CO₃

The hydration equilibrium constant at 25°C is K_h= 1.70×10⁻³: hence, the majority of the carbon dioxide is not converted into carbonic acid and stays 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.

Role of carbonic acid in blood

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 disassociates 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.

Carbonic acid also plays a very important role as a buffer in mammalian blood. The equilibrium between carbon dioxide and carbonic acid is very important for controlling the acidity of body fluids, and the carbonic anhydrase increases the reaction rate by a factor of nearly a billion to keep the fluids at a stable pH.

Role of carbonic acid in anthropogenic climate change

The oceans of the world have absorbed almost half of the CO₂ emitted by humans from the burning of fossil fuels.  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 process is known as ocean acidification.  Depending on the rate and magnitude of future emissions, the ocean's pH could drop by as much as 0.35 units by the mid-21st century.

Acidity of carbonic acid

Carbonic acid is diprotic: it has two hydrogen atoms 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 ⇌ HCO₃⁻ + H⁺

K_a1 = 4.3×10⁻⁷ ; pK_a1 = 6.37 at 25 °C.

and the second for the dissociation of the bicarbonate ion into the carbonate ion CO₃²⁻:

HCO₃⁻ ⇌ CO₃²⁻ + H⁺

K_a2 = 5.61×10⁻¹¹ ; pK_a2 = 10.25 at 25 °C.

Care must be taken when quoting and using the first dissociation constant of carbonic acid. The value given above is correct for the H₂CO₃ molecule, and shows that it is a stronger acid than acetic acid or formic acid. This is expected from the influence of the electronegative oxygen substituent. However, in aqueous solution carbonic acid only exists in equilibrium with carbon dioxide, and the concentration of H₂CO₃ there is much lower than the CO₂ concentration, thus reducing the measured acidity. The equation may be rewritten as follows (c.f. sulfurous acid):

CO₂ + H₂O ⇌ HCO₃⁻ + H⁺

K_a = 4.30×10⁻⁷; pK_a = 6.36.

This figure is quoted as the dissociation constant of carbonic acid, although this is ambiguous: it might better be referred to as the acidity constant of carbon dioxide, as it is particularly useful for calculating the pH of CO₂ solutions.

pH and composition of a carbonic acid solution

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 $\scriptstyle K_{h}={\frac {[H_{2}CO_{3}]}{[CO_{2}]}}$ (see above) and of the following equilibrium between the dissolved CO₂ and the gaseous CO₂ above the solution:

CO₂(gas) ↔ CO₂(dissolved) with $\scriptstyle {\frac {[CO_{2}]}{p_{CO_{2}}}}={\frac {1}{k_{\mathrm {H} }}}$ where k_H=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 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:

$\scriptstyle P_{CO_{2}}$ (atm)	pH	[CO2] (mol/L)	[H2CO3] (mol/L)	[HCO3−] (mol/L)	[CO32−] (mol/L)
10−8	7.00	3.36 × 10−10	5.71 × 10−13	1.42 × 10−9	7.90 × 10−13
10−6	6.81	3.36 × 10−8	5.71 × 10−11	9.16 × 10−8	3.30 × 10−11
10−4	5.92	3.36 × 10−6	5.71 × 10−9	1.19 × 10−6	5.57 × 10−11
3.5 × 10−4	5.65	1.18 × 10−5	2.00 × 10−8	2.23 × 10−6	5.60 × 10−11
10−3	5.42	3.36 × 10−5	5.71 × 10−8	3.78 × 10−6	5.61 × 10−11
10−2	4.92	3.36 × 10−4	5.71 × 10−7	1.19 × 10−5	5.61 × 10−11
10−1	4.42	3.36 × 10−3	5.71 × 10−6	3.78 × 10−5	5.61 × 10−11
1	3.92	3.36 × 10−2	5.71 × 10−5	1.20 × 10−4	5.61 × 10−11
2.5	3.72	8.40 × 10−2	1.43 × 10−4	1.89 × 10−4	5.61 × 10−11
10	3.42	0.336	5.71 × 10−4	3.78 × 10−4	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.

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

$\scriptstyle [H^{+}]\simeq \left(10^{-14}+{\frac {K_{h}K_{a1}}{k_{\mathrm {H} }}}p_{CO_{2}}\right)^{1/2}$

Instability of carbonic acid

It has long been recognized that it is impossible to obtain pure carbonic acid at room temperatures (about 20 °C or about 70 °F). However, in 1991 scientists at NASA's Goddard Space Flight Center (USA) succeeded in making the first pure H₂CO₃ samples. They did so 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 has given rise to suggestions 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. The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by a cryotechnique at the University of Innsbruck: alternating layers of glassy aqueous solutions of bicarbonate and acid were heated in vacuo, which causes protonation of bicarbonate, and the solvent was subsequently removed. A second polymorph (denoted alpha-carbonic acid) was prepared by the same technique at the University of Innsbruck using methanol rather than water as a solvent.

It has since been shown, by theoretical calculations, that the presence of even a single molecule of water causes carbonic acid to revert to carbon dioxide and water fairly quickly. Pure carbonic acid is predicted to be stable in the gas phase, in the absence of water, with a calculated half-life of 180,000 years.

There is a hypothetical acid orthocarbonic acid which is even more hydrated, being H₄CO₄.

See also

• Carbonated water
• Ocean acidification
• Carbon dioxide

References

• Welch, M. J.; Lipton, 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.{{cite journal}}: CS1 maint: date and year (link)
• Jolly, W. L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 0-07-112651-1.
• 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: 255–262. doi:10.1016/0584-8539(91)80097-3.{{cite journal}}: CS1 maint: date and year (link)
• Loerting, T.; Tautermann, C.; Kroemer, R.T.; Kohl, I.; Mayer, E.; Hallbrucker, A.; Liedl, K. R. (2001). "On the Surprising Kinetic Stability of Carbonic Acid". Angew. Chem. Int. Ed. 39: 891–895. doi:10.1002/(SICI)1521-3773(20000303)39:5<891::AID-ANIE891>3.0.CO;2-E.{{cite journal}}: CS1 maint: date and year (link)
• W. Hage, K. R. Liedl; Mayer, E. (1998). "Carbonic Acid in the Gas Phase and Its Astrophysical Relevance". Science. 279: 1332–1335. doi:10.1126/science.279.5355.1332. PMID 9478889.{{cite journal}}: CS1 maint: date and year (link)
• 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: 8427–8431. doi:10.1021/ja00071a061.{{cite journal}}: CS1 maint: date and year (link)
• Hage, W.; Hallbrucker, A.; Mayer, E. (1995). "A Polymorph of Carbonic Acid and Its Possible Astrophysical Relevance". J. Chem. Soc. Farad. Trans. 91: 2823–2826. doi:10.1039/ft9959102823.{{cite journal}}: CS1 maint: date and year (link)

External links

• Ask a Scientist: Carbonic Acid Decomposition
• Why was the existence of carbonic acid unfairly doubted for so long?