Carbonic acid

Bjerrum plot for carbonate speciation in seawater (ionic strength 0.7 mol dm⁻³)

In aqueous solution carbonic acid behaves as a dibasic acid. The Bjerrum plot shows typical equilibrium concentrations, in solution, in seawater, of carbon dioxide and the various species derived from it, as a function of pH.[8][9] The acidification of natural waters is caused by the increasing concentration of carbon dioxide in the atmosphere, which is caused by the burning of increasing amounts of coal and hydrocarbons.[10][11]

Expected change refers to predicted effect of continued ocean acidification.[12] It has been estimated that the increase in dissolved carbon dioxide has caused the ocean's average surface pH to decrease by about 0.1 from pre-industrial levels.

The stability constants database contains 136 entries with values for the overall protonation constants, β₁ and β₂, of the carbonate ion. In the following expressions [H⁺] represents the concentration, at equilibrium, of the chemical species H⁺, etc.

The value of log β₁ decreases with increasing ionic strength, ${\displaystyle {\ce {I}}}$. At 25 °C:

${\displaystyle {\ce {CO3^{2-}{+ H+}<=> HCO3^-}}}$ : ${\displaystyle \beta _{1}={\frac {[{\text{HCO}}_{3}^{-}]}{[{\text{H}}^{+}][{\text{CO}}_{3}^{2-}]}}}$

${\displaystyle \log \beta _{1}=0.54I^{2}-0.96I+9.93}$ (selected data from SC-database)

The value of log β₂ also decreases with increasing ionic strength.

${\displaystyle {\ce {CO3^{2-}{+ 2H+}<=> H2CO3}}}$ : ${\displaystyle \beta _{2}={\frac {[{\text{H}}_{2}{\text{CO}}_{3}]}{[{\text{H}}^{+}]^{2}[{\text{CO}}_{3}^{2-}]}}}$

${\displaystyle \log \beta _{2}=-2.5I^{2}-0.043I+16.07}$

At ${\displaystyle {\ce {I}}}$=0 and 25 °C the pK values of the stepwise dissociation constants are

pK₁ = logβ₂ - logβ₁ = 6.77.

pK₂ = logβ₁ = 9.93.

When pH = pK the two chemical species in equilibrium with each other have the same concentration.

Note 1: There are apparently conflicting values in the literature for pK_a. Pines et al. cite a value for "pK_app" of 6.35, consistent with the value 6.77, mentioned above.[13] They also give a value for "pK_a" of 3.49 and state that

pK_a = pK_app − log K_D (eqn. 5)

where K_D=[CO₂]/[H₂CO₃]. (eqn. 3) The situation arises from the way that the dissociation constants are named and defined, which is clearly stated in the text of the Pines paper, but not in the abstract.

Note 2: The numbering of dissociation constants is the reverse of the numbering of the numbering of association constants, so pK₂ (dissociation)= log β₁ (association). The value of the stepwise constant for the equilibrium

${\displaystyle {\ce {HCO3- <=> CO3^{2-}{+ H+}}}}$

is given by

pK₁(dissociation)₁ = log β₂ − log β₁ (association)

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[14] and ≈ 1.2×10⁻³ in seawater.[15] 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 and 23 s⁻¹ for the reverse reaction.

In the beverage industry, sparkling or "fizzy water" is usually referred to as carbonated water. It is made by dissolving carbon dioxide under a small positive pressure in water. Many soft drinks treated in the same way make them effervescent.

Significant amounts of molecular H₂CO₃ exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres), such as can occur in planetary interiors.[16][17] Pressures of 0.6–1.6 GPa at 100 K, and 0.75–1.75 GPa at 300 K are attained in the cores of large icy satellites such as Ganymede, Callisto, and Titan, where water and carbon dioxide are present. Pure carbonic acid, being denser, would then sink under the ice layers and separate them from the rocky cores of these moons.[18]

When the enzyme carbonic anhydrase is also present in the solution the following reaction takes precedence.[19]

${\displaystyle {\ce {HCO3^- {+}H^+ <=> CO2 {+}H2O}}}$

When the amount of carbon dioxide created by the forward reaction exceeds its solubility, gas is evolved and a third equilibrium

${\displaystyle {\ce {CO_2 (soln) <=> CO_2 (g)}}}$

must also be taken into consideration. The equilibrium constant for this reaction is defined by Henry's law. The two reactions can be combined for the equilibrium in solution.

${\displaystyle {\ce {HCO3^- {+}H^{+}<=> CO_2(soln) {+}H2O}}}$ :    ${\displaystyle {\ce {K_3 = {\frac {[H^+][HCO_3^{-}]}{[CO_2(soln)]}}}}}$

When Henry's law is used to calculate the value of the term in the denominator care is needed with regard to dimensionality.

In physiology, carbon dioxide excreted by the lungs may be called volatile acid or respiratory acid.