In chemistry, the term "carbonic acid" strictly refers to the chemical compound with the formula H
₂CO
₃. Some biochemistry literature effaces the distinction between carbonic acid and carbon dioxide dissolved in extracellular fluid.

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

At ambient temperatures, pure carbonic acid is a stable gas.^[6] There are two main methods to produce anhydrous carbonic acid: reaction of hydrogen chloride and potassium bicarbonate at 100 K in methanol and proton irradiation of pure solid carbon dioxide.^[3] Chemically, it behaves as a diprotic Brønsted acid.^[8][9]

Carbonic acid monomers exhibit three conformational isomers: cis–cis, cis–trans, and trans–trans.^[10]

At low temperatures and atmospheric pressure, solid carbonic acid is amorphous and lacks Bragg peaks in X-ray diffraction.^[11] But at high pressure, carbonic acid crystallizes, and modern analytical spectroscopy can measure its geometry.

According to neutron diffraction of dideuterated carbonic acid (D
₂CO
₃) in a hybrid clamped cell (Russian alloy/copper-beryllium) at 1.85 GPa, the molecules are planar and form dimers joined by pairs of hydrogen bonds. All three C-O bonds are nearly equidistant at 1.34 Å, intermediate between typical C-O and C=O distances (respectively 1.43 and 1.23 Å). The unusual C-O bond lengths are attributed to delocalized π bonding in the molecule's center and extraordinarily strong hydrogen bonds. The same effects also induce a very short O—O separation (2.13 Å), through the 136° O-H-O angle imposed by the doubly hydrogen-bonded 8-membered rings.^[4] Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid, where the distances exceed 2.4 Å.^[11]

In even a slight presence of water, carbonic acid dehydrates to carbon dioxide and water, which then catalyzes further decomposition.^[6] For this reason, carbon dioxide can be considered the carbonic acid anhydride.

The hydration equilibrium constant at 25 °C is [H
₂CO
₃]/[CO₂] ≈ 1.7×10⁻³ in pure water^[12] and ≈ 1.2×10⁻³ in seawater.^[13] Hence the majority of carbon dioxide at geophysical or biological air-water interfaces does not convert to carbonic acid, remaining dissolved CO₂ gas. However, the uncatalyzed equilibrium is reached quite slowly: the rate constants are 0.039 s⁻¹ for hydration and 23 s⁻¹ for dehydration.

In the presence of the enzyme carbonic anhydrase, equilibrium is instead reached rapidly, and the following reaction takes precedence:^[14] $${\displaystyle {\ce {HCO3^- {+}H^+ <=> CO2 {+}H2O}}}$$ 

When the created carbon dioxide exceeds its solubility, gas evolves 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: $${\displaystyle {\begin{aligned}{\ce {HCO3^{-}{}+ H+{}<=> CO2(soln){}+ H2O}}&&K_{3}={\frac {[{\ce {H+}}][{\ce {HCO3^-}}]}{[{\ce {CO2(soln)}}]}}\end{aligned}}}$$  When Henry's law is used to calculate the denominator care is needed with regard to units since Henry's law constant can be commonly expressed with 8 different dimensionalities.^[15]

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 the same way effervesce.

Significant amounts of molecular H
₂CO
₃ exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres) 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, is expected to have sunk under the ice layers and separate them from the rocky cores of these moons.^[18]

Carbonic acid is the formal Brønsted–Lowry conjugate acid of the bicarbonate anion, stable in alkaline solution. The protonation constants have been measured to great precision, but depend on overall ionic strength I. The two equilibria most easily measured are as follows: $${\displaystyle {\begin{aligned}{\ce {CO3^{2-}{}+ H+{}<=> HCO3^-}}&&\beta _{1}={\frac {[{\ce {HCO3^-}}]}{[{\ce {H+}}][{\ce {CO3^{2-}}}]}}\\{\ce {CO3^{2-}{}+ 2H+{}<=> H2CO3}}&&\beta _{2}={\frac {[{\ce {H2CO3}}]}{[{\ce {H+}}]^{2}[{\ce {CO3^{2-}}}]}}\end{aligned}}}$$  where brackets indicate the concentration of specie. At 25 °C, these equilibria empirically satisfy^[19]$${\displaystyle {\begin{alignedat}{6}\log(\beta _{1})=&&0&.54&I^{2}-0&.96&I+&&9&.93\\\log(\beta _{2})=&&-2&.5&I^{2}-0&.043&I+&&16&.07\end{alignedat}}}$$ log(β₁) decreases with increasing I, as does log(β₂). In a solution absent other ions (e.g. I = 0), these curves imply the following stepwise dissociation constants:$${\displaystyle {\begin{alignedat}{3}p{\text{K}}_{1}&=\log(\beta _{2})-\log(\beta _{1})&=6.77\\p{\text{K}}_{2}&=\log(\beta _{1})&=9.93\end{alignedat}}}$$  Direct values for these constants in the literature include pK₁ = 6.35 and pK₂ - pK₁ = 3.49.^[20]

To interpret these numbers, note that two chemical species in an acid equilibrium are equiconcentrated when pK = pH. In particular, the extracellular fluid (cytosol) in biological systems exhibits pH ≈ 7.2, so that carbonic acid will be almost 50%-dissociated at equilibrium.

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] 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.^[21][22] 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 pre-industrial levels.

•   "Climate and Carbonic Acid" in Popular Science Monthly Volume 59, July 1901
• Welch, M. J.; Lifton, J. F.; Seck, J. A. (1969). "Tracer studies with radioactive oxygen-15. Exchange between carbon dioxide and water". J. Phys. Chem. 73 (335): 3351. doi:10.1021/j100844a033.
• Jolly, W. L. (1991). Modern Inorganic Chemistry (2nd ed.). McGraw-Hill. ISBN 978-0-07-112651-9.
• Moore, M. H.; Khanna, R. (1991). "Infrared and Mass Spectral Studies of Proton Irradiated H₂O+CO₂ Ice: Evidence for Carbonic Acid Ice: Evidence for Carbonic Acid". Spectrochimica Acta. 47A (2): 255–262. Bibcode:1991AcSpA..47..255M. doi:10.1016/0584-8539(91)80097-3.
• W. Hage, K. R. Liedl; Liedl, E.; Hallbrucker, A; Mayer, E (1998). "Carbonic Acid in the Gas Phase and Its Astrophysical Relevance". Science. 279 (5355): 1332–5. Bibcode:1998Sci...279.1332H. doi:10.1126/science.279.5355.1332. PMID 9478889.
• Hage, W.; Hallbrucker, A.; Mayer, E. (1995). "A Polymorph of Carbonic Acid and Its Possible Astrophysical Relevance". J. Chem. Soc. Faraday Trans. 91 (17): 2823–6. Bibcode:1995JCSFT..91.2823H. doi:10.1039/ft9959102823.

• Carbonic acid/bicarbonate/carbonate equilibrium in water: pH of solutions, buffer capacity, titration, and species distribution vs. pH, computed with a free spreadsheet
• How to calculate concentration of carbonic acid in water