- Not to be confused with carbolic acid, an antiquated name for phenol.
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. It plays an important role in the bicarbonate buffer system to maintain acid–base homeostasis.
|Preferred IUPAC name
Carbon dioxide solution
Acid of air
3D model (JSmol)
|Molar mass||g·mol−1 62.024|
|Only stable in solution|
|Acidity (pKa)||3.6 (pKa1 for H2CO3 only), 6.3 (pKa1 including CO2(aq)), 10.32 (pKa2)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
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 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 and ≈ 1.2×10−3 in seawater. 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). 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.
Role of carbonic acid in bloodEdit
Bicarbonate is an intermediate in the transport of CO2 out of the body by 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 increases the reaction rate, producing 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. A 2016 theoretical report suggests that carbonic acid may play a pivotal role in protonating various nitrogen bases in blood serum.
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. It has been estimated 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 is known as ocean acidification, even though the ocean remains basic.
Acidity of carbonic acidEdit
Carbonic acid is a carboxylic acid with a hydroxyl group as the substituent. It is also a polyprotic acid — specifically it is diprotic, meaning that it has two protons that may dissociate from the parent molecule. Thus, there are two dissociation constants, first of which is for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3−:
- Ka1 = 2.5×10−4; 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:
- H2CO3* ⇌ HCO3− + H+
- Ka(app) = 4.47×10−7; pK(app) = 6.35 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.
- Ka2 = 4.69×10−11; pKa2 = 10.329 at 25 °C and ionic strength = 0.0.
The three acidity constants are defined as follows:
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 of carbon dioxide above the solution. To calculate this composition, account must be taken of
- the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution:
- CO2(gas) ⇌ CO2(dissolved) with where kH = 29.76 atm/(mol/L) (Henry constant) at 25 °C
- the hydration equilibrium between dissolved CO2 and H2CO3 with constant (see above)
- the first dissociation equilibrium of carbonic acid (see Ka1 above)
- the second dissociation equilibrium of carbonic acid (see Ka2 above)
- the dissociation equilibrium of water:
- the charge neutrality condition
Taken at face value, the above are 6 equations for the 6 unknowns [CO2]aq, [H2CO3], [H+], [OH−], [HCO3−] and [CO32−]. Note, however, that the first 2 equations express [CO2]aq and [H2CO3] as simple linear functions of , reducing the problem to the latter 4 equations with 4 unknowns. Either way, this demonstrates that the composition of the solution is fully determined by . When isolating [HCO3−] in the first dissociation equilibrium, [HCO32−] in the second dissociation equilibrium and [OH−] in the dissociation equilibrium of water, then substituting all three in the charge neutrality condition, a cubic equation in [H+] is obtained, whose numerical solution yields the values for the pH and the concentrations of the different species in the following table. (Note that the second dissociation equilibrium can be neglected for this particular problem, reducing the cubic equation to a simple square root; see remarks below the table.)
|10−8||7.00||3.36 × 10−10||5.71 × 10−13||1.42 × 10−9||7.90 × 10−13|
|10−7||6.94||3.36 × 10−9||5.71 × 10−12||5.90 × 10−9||1.90 × 10−12|
|10−6||6.81||3.36 × 10−8||5.71 × 10−11||9.16 × 10−8||3.30 × 10−11|
|10−5||6.42||3.36 × 10−7||5.71 × 10−10||3.78 × 10−7||4.53 × 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|
|10 0||3.92||3.36 × 10−2||5.71 × 10−5||1.20 × 10−4||5.61 × 10−11|
|2.5 × 100||3.72||8.40 × 10−2||1.43 × 10−4||1.89 × 10−4||5.61 × 10−11|
|10 1||3.42||3.36 × 10−1||5.71 × 10−4||3.78 × 10−4||5.61 × 10−11|
- In the total range of pressure, the pH is always much lower than pKa2 (= 10.3) 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 , 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 ( atm), we get a slightly acidic solution (pH = 5.7), and the dissolved carbon is now essentially in the CO2 and HCO3− forms.
- For a CO2 pressure typical for bottled carbonated drinks ( ~ 2.5 atm), we get a relatively acidic 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 [H2CO3] > [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+]:
Pure carbonic acidEdit
It was long believed that carbonic acid could not exist as a pure compound. However, in 1991 scientists at NASA's Goddard Space Flight Center (USA) succeeded in making solid H2CO3 samples. They did so by exposing a frozen mixture of water and carbon dioxide to high-energy proton 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, or even by irradiation of dry ice alone, 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. It was announced in 1993 that solid carbonic acid had been created by a cryogenic reaction of potassium bicarbonate and HCl dissolved in methanol. Later work showed that in fact the methyl ester had been formed, but other methods were successful. Theoretical calculations showed that a single molecule of water can catalyze 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 has been predicted to be very slow, with a half-life of 180,000 years. This only applies if the molecules are few and far between, 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.
For a while it was thought that there were two polymorphs of solid carbonic acid, called α and β. The polymorph denoted β-carbonic acid was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuum, which causes protonation of the bicarbonate, followed by removal of the solvent. The previously suggested α-carbonic acid, which was prepared by the same technique using methanol rather than water as a solvent, was later shown to be a monomethyl ester CH3OCOOH.
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