]} \\
\ce{CO3^{2-}{} + 2H+{} <=> H2CO3} && \beta_2 = \frac{[\ce{H2CO3}]}{[\ce{H+}]^2[\ce{CO3^{2-]}
\end{align}</math> where brackets indicate the concentration of species. At 25 °C, these equilibria empirically satisfy<math display="block">\begin{alignat}{6}
\log(\beta_1) =&& 0&.54&I^2 - 0&.96&I +&& 9&.93 \\
\log(\beta_2) =&& -2&.5&I^2 - 0&.043&I +&& 16&.07
\end{alignat}</math> decreases with increasing , as does . In a solution absent other ions (e.g. ), these curves imply the following stepwise dissociation constants:<math display="block">\begin{alignat}{3}
p\text{K}_1 &= \log(\beta_2) - \log(\beta_1) &= 6.77 \\
p\text{K}_2 &= \log(\beta_1) &= 9.93
\end{alignat}</math> Direct values for these constants in the literature include and .
To interpret these numbers, note that two chemical species in an acid equilibrium are equiconcentrated when . In particular, the extracellular fluid (cytosol) in biological systems exhibits , so that carbonic acid will be almost 50%-dissociated at equilibrium.
Ocean acidification
thumb|class=skin-invert-image|upright=1.4|Carbonate speciation in seawater (ionic strength 0.7 mol/dm<sup>3</sup>). The expected change shown is due to the [[Greenhouse gas emissions|current anthropogenic increase in atmospheric carbon dioxide concentration.|left]]
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. As human industrialization has increased the proportion of carbon dioxide in Earth's atmosphere, the proportion of carbon dioxide dissolved in sea- and freshwater as carbonic acid is also expected to increase. This rise in dissolved acid is also expected to acidify those waters, generating a decrease in pH. It has been estimated that the increase in dissolved carbon dioxide has already caused the ocean's average surface pH to decrease by about 0.1 from historical pre-industrial levels.