Carbonic acid

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 was disproved, in 1993, by the preparation of the pure substance.^[2] This was independently confirmed in 2011.^[3]

Chemical equilibrium

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

CO₂ + H₂O ⇌ H₂CO₃

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^[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). 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

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.^[7] Recent theoretical report suggests that carbonic acid may have 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 hypothesised 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 theory is known as ocean acidification.^[10]

Acidity of carbonic acid

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 ⇌ HCO₃⁻ + H⁺

K_a1 = 2.5×10⁻⁴;^[4] pK_a1 = 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⁺

K_a(app) = 4.47×10⁻⁷; pK(app) = 6.35 at 25 °C and ionic strength = 0.0

Whereas this apparent pK_a 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⁻ ⇌ CO₃²⁻ + H⁺

K_a2 = 4.69×10⁻¹¹; pK_a2 = 10.329 at 25 °C and ionic strength = 0.0

The three acidity constants are defined as follows: $K_{a1}=\frac{[\text{H}^+][\text{HCO}_3^-]}{[\text{H}_2\text{CO}_3]} \qquad K_a\text{(app)}=\frac{[\text{H}^+][\text{HCO}_3^-]}{[\text{H}_2\text{CO}_3]+[\text{CO}_2\text{(aq)}]} \qquad K_{a2}=\frac{[\text{H}^+][\text{CO}_3^{2-}]}{[\text{HCO}_3^-]}$

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 $p_{\text{CO}_2}$ 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 $\textstyle K_h=\frac{[\text{H}_2\text{CO}_3]}{[\text{CO}_2]}$ (see above) and of the following equilibrium between the dissolved CO₂ and the gaseous CO₂ above the solution:

CO₂(gas) ⇌ CO₂(dissolved) with

\textstyle \frac{[\text{CO}_2]}{p_{\text{CO}_2}}=\frac{1}{k_\text{H}},

where k_H = 29.76 atm/(mol/L) (Henry constant) at 25 °C.

The corresponding equilibrium equations together with the $[\text{H}^+][\text{OH}^-] = 10^{-14}$ relation and the charge neutrality condition $[\text{H}^+]=[\text{OH}^-]+[\text{HCO}_3^-]+2[\text{CO}_3^{2-}]$ 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 $p_{\text{CO}_2}$ . The equation obtained for [H⁺] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

$p_{\text{CO}_2}$ (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 × 10⁻⁰⁹	7.90 × 10⁻¹³
1.0 × 10⁻⁷	6.94	3.36 × 10⁻⁰⁹	5.71 × 10⁻¹²	5.90 × 10⁻⁰⁹	1.90 × 10⁻¹²
1.0 × 10⁻⁶	6.81	3.36 × 10⁻⁰⁸	5.71 × 10⁻¹¹	9.16 × 10⁻⁰⁸	3.30 × 10⁻¹¹
1.0 × 10⁻⁵	6.42	3.36 × 10⁻⁰⁷	5.71 × 10⁻¹⁰	3.78 × 10⁻⁰⁷	4.53 × 10⁻¹¹
1.0 × 10⁻⁴	5.92	3.36 × 10⁻⁰⁶	5.71 × 10⁻⁰⁹	1.19 × 10⁻⁰⁶	5.57 × 10⁻¹¹
3.5 × 10⁻⁴	5.65	1.18 × 10⁻⁰⁵	2.00 × 10⁻⁰⁸	2.23 × 10⁻⁰⁶	5.60 × 10⁻¹¹
1.0 × 10⁻³	5.42	3.36 × 10⁻⁰⁵	5.71 × 10⁻⁰⁸	3.78 × 10⁻⁰⁶	5.61 × 10⁻¹¹
1.0 × 10⁻²	4.92	3.36 × 10⁻⁰⁴	5.71 × 10⁻⁰⁷	1.19 × 10⁻⁰⁵	5.61 × 10⁻¹¹
1.0 × 10⁻¹	4.42	3.36 × 10⁻⁰³	5.71 × 10⁻⁰⁶	3.78 × 10⁻⁰⁵	5.61 × 10⁻¹¹
1.0 × 10⁺⁰	3.92	3.36 × 10⁻⁰²	5.71 × 10⁻⁰⁵	1.20 × 10⁻⁰⁴	5.61 × 10⁻¹¹
2.5 × 10⁺⁰	3.72	8.40 × 10⁻⁰²	1.43 × 10⁻⁰⁴	1.89 × 10⁻⁰⁴	5.61 × 10⁻¹¹
1.0 × 10⁺¹	3.42	3.36 × 10⁻⁰¹	5.71 × 10⁻⁰⁴	3.78 × 10⁻⁰⁴	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 $p_{\text{CO}_2}$ , 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 ( $p_{\text{CO}_2}=3.5\times 10^{-4}$ 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 carbonated drink bottles ( $p_{\text{CO}_2}$ ~ 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.

Remark

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

[\text{H}^+] \simeq \left( 10^{-14} + \frac{K_hK_{a1}}{k_\text{H}} p_{\text{CO}_2}\right)^{1/2}.

Spectroscopic studies of carbonic acid

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.^[11] This may only apply to isolated carbonic acid molecules, however, as it has been predicted to catalyze its own decomposition.^[12]

The longstanding belief was that pure carbonic acid cannot be obtained at room temperatures (about 20 °C, or about 70 °F) due to its catalytic destruction by water.^[11] It has been discovered that 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 remains is 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.^[11] The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuum, 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.^[13]

References

↑ Acid-Base Physiology 2.1 - Acid-Base Balance by Kerry Brandis
↑ 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.
↑ "International first: Gas-phase carbonic acid isolated", phys.org, Jan 11, 2011
1 2 3 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 310. ISBN 0-08-037941-9.
↑ Housecroft and Sharpe, Inorganic Chemistry, 2nd ed, Prentice-Pearson-Hall 2005, p.368.
↑ 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.
↑ "excretion." Encyclopædia Britannica. Encyclopædia Britannica Ultimate Reference Suite. Chicago: Encyclopædia Britannica, 2010.
↑ "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.
↑ Sabine, C.L.; et al. (2004). "The Oceanic Sink for Anthropogenic CO₂". Science 305 (5682): 367–371. doi:10.1126/science.1097403. PMID 15256665. Archived from " the original on July 6, 2008.
↑ 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.
1 2 3 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.
↑ de Marothy, S. A. (2013). "Autocatalytic decomposition of carbonic acid". Int. J. Quantum Chem 113 (20): 2306–2311. doi:10.1002/qua.24452. Retrieved 19 January 2015.
↑ 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.

External links

Hydrogen compounds

Oxocarbons

Common oxides	CO 2 CO

Exotic oxides	CO 3 CO 4 CO 5 CO 6 C 2O C 2O 2 C 2O 3 C 2O 4 C 2O 4 C 3O 2 C 3O 3 C 3O 6 C 4O 2 C 4O 4 C 4O 6 C 5O 2 C 5O 5 C 6O 6 C 6O 6 C 8O 8 C 9O 9 C 10O 8 C 10O 10 C 12O 6 C 12O 9 C 12O 12

Polymers	Graphite oxide C₃O₂ CO CO₂

Compounds derived from oxides	Metal carbonyls Carbonic acid Bicarbonates Carbonates Dicarbonates Tricarbonates

Authority control	GND: 4192064-8

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