Carbonic acid

Chemical equilibrium

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

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 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.

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 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 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 chemistry

The oceans of the world have absorbed almost half of the CO₂ emitted by humans from the burning of fossil fuels.  It has been theorized 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, even though the ocean remains basic.

Acidity of carbonic acid

Carbonic acid is a polyprotic acid — specifically it is diprotic meaning 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; 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:

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 a 1 = [ H + ] [ HCO 3 − ] [ H 2 CO 3 ] K a (app) = [ H + ] [ HCO 3 − ] [ H 2 CO 3 ] + [ CO 2 (aq) ] K a 2 = [ H + ] [ CO 3 2 − ] [ 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 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 K h = [ 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 [ CO 2 ] p CO 2 = 1 k H , where k_H = 29.76 atm/(mol/L) (Henry constant) at 25 °C.

The corresponding equilibrium equations together with the [ H + ] [ OH − ] = 10 − 14 relation and the charge neutrality condition [ 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 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:

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 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 CO 2 = 3.5 × 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 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⁺]:

[ H + ] ≃ ( 10 − 14 + K h K a 1 k H p CO 2 ) 1 / 2 .

Pure carbonic acid

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 beta-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 alpha-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.