Carbonic acid

Carbonic acid is a chemical compound with the formula H2CO3. It behaves as a diprotic Brønsted acid, but in the presence of water, it rapidly breaks down into water and carbon dioxide. In its pure, anhydrous form, however, it is surprisingly stable. The reversible conversion between carbon dioxide and carbonic acid plays a key role in the respiration of all aerobic organisms and in the acidification of natural waters.

The history of the name is tangled. In 1781, a group of chemists led by Lavoisier named carbon dioxide "carbonic acid" because it was produced by burning charcoal and was considered an acid based on its ability to turn litmus red and taste sour. Later, as the definition of an acid changed, carbon dioxide was no longer accepted as one, and the name "carbonic acid" was reassigned to H2CO3. For many years, H2CO3 was thought to exist only in solution and not as an independent molecule.

In the 1960s, researchers succeeded in making adducts of carbonic acid. At -78 °C, they prepared an adduct with diethyl ether. Another adduct, with dimethyl ether, was synthesized at -30 °C from dimethyl ether and sodium carbonate and confirmed by chemical, thermochemical, IR, and NMR methods; this pure solid decomposes at 5 °C. In 1987, gaseous H2CO3 was produced by heating ammonium bicarbonate and confirmed by IR and collisional activation mass spectrometry. In 1991, solid carbonic acid was made by irradiating a frozen 1:1 mixture of water and carbon dioxide at 20 K, a form later called α‑H2CO3. Another route involved reacting hydrogen bromide and potassium bicarbonate in water, then sublimating the water away at 200 K, yielding a form called β‑H2CO3, confirmed by FTIR. However, α‑H2CO3 was later argued to actually be a monomethyl ester, (CH3)O(COOH).

The reason for carbonic acid’s surprising kinetic stability—and why it took so long to isolate it as a solid—is explained by transition state theory. Pure H2CO3 is metastable: at 300 K, its theoretical half-life is about 105 years, but in the presence of just two water molecules, the half-life drops to roughly one minute. Water catalyzes its decomposition.

Anhydrous carbonic acid can be made by reacting hydrogen chloride and potassium bicarbonate at 100 K in methanol, or by proton irradiation of pure solid carbon dioxide. At room temperature, it is predicted to be a kinetically stable gas. Its monomers exist as three conformational isomers: cis–cis, cis–trans, and trans–trans. At low temperature and atmospheric pressure, solid carbonic acid is amorphous and shows no Bragg peaks in X-ray diffraction. Under high pressure, it crystallizes. Neutron diffraction of dideuterated carbonic acid (D2CO3) at 1.85 GPa in a hybrid clamped cell reveals planar molecules that form dimers joined by pairs of hydrogen bonds. The three C–O bonds are nearly equidistant at 1.34 Å, intermediate between typical single and double bond lengths (1.43 and 1.23 Å). This is attributed to delocalized π bonding and extraordinarily strong hydrogen bonds, which also produce a very short O–O separation of 2.13 Å through the 136° O–H–O angle in the doubly hydrogen-bonded eight-membered rings. For comparison, strong intramolecular hydrogen bonds in oxalic acid have O–O distances exceeding 2.4 Å. Trimers and higher polymers are predicted to be even more stable, leading to suggestions that solid carbonic acid could be used for carbon sequestration.

In aqueous solution, even a trace of water causes carbonic acid to dehydrate to carbon dioxide and water, which then catalyzes further decomposition. At 25 °C, the hydration equilibrium constant [H2CO3]/[CO2] is about 1.7×10⁻³ in pure water and 1.2×10⁻³ in seawater, meaning most dissolved carbon dioxide at air-water interfaces does not convert to carbonic acid. The uncatalyzed reaction is slow: the rate constant for hydration is 0.039 s⁻¹ and for dehydration is 23 s⁻¹. In biological systems, the enzyme carbonic anhydrase accelerates the equilibrium, allowing the rapid interconversion of carbon dioxide, bicarbonate, and protons.