# Carbonic acid

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.[2] It plays an important role in the bicarbonate buffer system to maintain acid–base homeostasis.

Carbonic acid
Names
Preferred IUPAC name
Carbonic acid[1]
Other names
Carbon dioxide solution
Dihydrogen carbonate
Hydrogen bicarbonate
Acid of air
Aerial acid
Hydroxymethanoic acid
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.133.015
EC Number
• 610-295-3
KEGG
Properties
CH2O3
Molar mass 62.024 g·mol−1
Density 1.668 g/cm3
Only stable in solution
Acidity (pKa) 3.6 (pKa1 for H2CO3 only), 6.3 (pKa1 including CO2(aq)), 10.32 (pKa2)
Conjugate base Bicarbonate
Related compounds
Related compounds
Acetone
Urea
Carbonyl fluoride
Hazards
NFPA 704 (fire diamond)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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.

It was long believed that carbonic acid could not exist as a pure compound. However, in 1991, it was reported that NASA scientists had succeeded in making solid H2CO3 samples.[3]

## Chemical equilibrium

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

${\displaystyle {\ce {CO2 + H2O <=> H2CO3}}}$

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[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). 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

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. The equilibration plays an important role as a buffer in mammalian blood.[7] A 2016 theoretical report suggests that carbonic acid may play a 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 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.[10]

## Acidity of carbonic acid

Carbonic acid is a carboxylic acid with a hydroxyl group as the substituent. It is also a polyprotic acid: it is specifically diprotic and so has two protons that may dissociate from the parent molecule. Thus, there are two dissociation constants, the first of which is for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3:

${\displaystyle {\ce {H2CO3 <=> HCO3^- + H+}}}$
Ka1 = 2.5×10−4;[4] pKa1 = 3.6 at 25 °C.

Care must be taken when the first dissociation constant of carbonic acid is quoted and used. In aqueous solutions, 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+
Ka(app) = 4.47×10−7; pK(app) = 6.35 at 25 °C and ionic strength = 0.0.

Whereas the apparent pKa is quoted as the dissociation constant of carbonic acid, it is ambiguous, and 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−:

${\displaystyle {\ce {HCO3^- <=> CO3^{2-}{}+ H+}}}$
Ka2 = 4.69×10−11; pKa2 = 10.329 at 25 °C and ionic strength = 0.0.

The three acidity constants are defined as follows:

${\displaystyle K_{a1}={\frac {[{\text{H}}^{+}][{\text{HCO}}_{3}^{-}]}{[{\text{H}}_{2}{\text{CO}}_{3}]}},}$
${\displaystyle K_{a}{\text{(app)}}={\frac {[{\text{H}}^{+}][{\text{HCO}}_{3}^{-}]}{[{\text{H}}_{2}{\text{CO}}_{3}]+[{\text{CO}}_{2}{\text{(aq)}}]}},}$
${\displaystyle K_{a2}={\frac {[{\text{H}}^{+}][{\text{CO}}_{3}^{2-}]}{[{\text{HCO}}_{3}^{-}]}}.}$

### pH and composition of carbonic acid solutions

The composition of a carbonic acid solution is fully determined by the partial pressure ${\displaystyle p_{{\text{CO}}_{2}}}$  of carbon dioxide above the solution. To calculate the 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 ${\displaystyle \textstyle {\frac {[{\text{CO}}_{2}]_{aq}}{p_{{\text{CO}}_{2}}}}={\frac {1}{k_{\text{H}}}},}$  where kH = 29.76 atm/(mol/L) (Henry constant) at 25 °C
${\displaystyle \Leftrightarrow [{\text{CO}}_{2}]_{aq}=p_{{\text{CO}}_{2}}/k_{\text{H}}}$
• the hydration equilibrium between dissolved CO2 and H2CO3 with constant ${\displaystyle K_{h}={\tfrac {[{\text{H}}_{2}{\text{CO}}_{3}]_{aq}}{[{\text{CO}}_{2}]_{aq}}}}$  (see above)
${\displaystyle \Leftrightarrow [{\text{H}}_{2}{\text{CO}}_{3}]_{aq}=K_{h}[{\text{CO}}_{2}]_{aq}={\frac {K_{h}}{k_{\text{H}}}}p_{{\text{CO}}_{2}}}$
• the first dissociation equilibrium of carbonic acid (see Ka1 above)
${\displaystyle \Leftrightarrow [{\text{H}}^{+}][{\text{HCO}}_{3}^{-}]=K_{a1}[{\text{H}}_{2}{\text{CO}}_{3}]={\frac {K_{h}K_{a1}}{k_{\text{H}}}}p_{{\text{CO}}_{2}}}$
• the second dissociation equilibrium of carbonic acid (see Ka2 above)
• the dissociation equilibrium of water: ${\displaystyle [{\text{H}}^{+}][{\text{OH}}^{-}]=10^{-14}}$
• the charge neutrality condition ${\displaystyle [{\text{H}}^{+}]=[{\text{OH}}^{-}]+[{\text{HCO}}_{3}^{-}]+2[{\text{CO}}_{3}^{2-}]}$

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. (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.)

${\displaystyle p_{{\text{CO}}_{2}}}$
(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 × 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
• 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 ${\displaystyle p_{{\text{CO}}_{2}}}$ , 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 (${\displaystyle p_{{\text{CO}}_{2}}=3.5\times 10^{-4}}$  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 (${\displaystyle p_{{\text{CO}}_{2}}}$  ~ 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.
Remark

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

${\displaystyle [{\text{H}}^{+}]\simeq {\sqrt {10^{-14}+{\frac {K_{h}K_{a1}}{k_{\text{H}}}}p_{{\text{CO}}_{2}}}}}$

## Pure carbonic acid

Carbonic acid forms as a by-product of CO2/H2O irradiation, in addition to carbon monoxide and radical species (HCO and CO3).[3] Another route to form carbonic acid is protonation of bicarbonates (HCO3) with aqueous HCl or HBr. This has to be done at cryogenic conditions to avoid immediate decomposition of H2CO3 to CO2 and H2O.[11] Amorphous H2CO3 forms above 120 K, and crystallization takes place above 200 K to give "β-H2CO3", as determined by infrared spectroscopy. The spectrum of β-H2CO3 agrees very well with the by-product after CO2/H2O irradiation.[3] β-H2CO3 sublimes at 230 - 260 K largely without decomposition. Matrix-isolation infared spectroscopy allows for the recording of single molecules of H2CO3.[12]

The fact that the carbonic acid may form by irradiating a solid H2O + CO2 mixture or even by proton-implantation of dry ice alone[13] 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.[14][15] The surprising stability of sublimed H2CO3 up to rather high-temperatures of 260 K even allows for gas-phase H2CO3, e.g., above the pole caps of Mars.[12] Ab initio calculations showed that a single molecule of water catalyzes 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 is predicted to be very slow, with a half-life in the gas-phase of 180,000 years at 300 K.[14] This only applies if the molecules are few and far apart, 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.[16]

Solid "α-carbonic acid" was claimed to be generated by a cryogenic reaction of potassium bicarbonate and a solution of HCl in methanol.[17][18] This claim was disputed in a PhD thesis submitted in January 2014.[19] Instead, isotope labeling experiments point to the involvement of carbonic acid monomethyl ester (CAME). Furthermore, the sublimed solid was suggested to contain CAME monomers and dimers, not H2CO3 monomers and dimers as previously claimed.[20] Subsequent matrix-isolation infrared spectra confirmed that CAME rather than carbonic acid is found in the gas-phase above "α-carbonic acid".[21] The assignment as CAME is further corroborated by matrix-isolation of the substance prepared in gas-phase by pyrolysis.[15]

Despite its complicated history, carbonic acid may still appear as distinct polymorphs. Carbonic acid forms upon oxidization of CO with OH-radicals.[22] It is not clear whether carbonic acid prepared in this way needs to be considered as γ-H2CO3. The structures of β-H2CO3 and γ-H2CO3 have not been characterized crystallographically.

## References

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