Carbonic acid

In the early days of chemistry, CO₂ was known by several names. Led by Lavoisier, a group of chemists systematized the nomenclature of chemicals then known. CO₂ was named "carbonic acid" in 1781, because it was produced by the combustion of charcoal (charbon in French), and it was considered an acid. Back then, an acid was something that, when dissolved in water, turned litmus red and had a distinctive acidic taste. Later, chemists reconceptualized "acid", and no longer accepted CO₂ to be an acid, and "carbonic acid" became the systematic name of H₂CO₃ instead. It was thought that H₂CO₃ could not exist independently, and could only exist in a solution.

In the 1960s, experimenters produced adducts of H₂CO₃. The diethyl ether adduct H₂CO₃·O(CH₂CH₃)₂ was prepared at -78 °C. The dimethyl ether adduct H₂CO₃·O(CH₃)₂ was synthesized at -30 °C using dimethyl ether and sodium carbonate, and confirmed by chemical, thermochemical, IR, and NMR methods. The pure H₂CO₃·O(CH₃)₂ is a solid that decomposes at 5 °C.

In 1987, gaseous H₂CO₃ was prepared by thermolysis of NH₄HCO₃, and confirmed by IR and collisional activation mass spectrometry. In 1991, solid H₂CO₃ was prepared by solid irradiation of a 1:1 H₂O + CO₂ ice mixture at 20 K. The result of this preparation route is termed α−H₂CO₃. Another synthesis route was to begin by reacting HBr and KHCO₃ in an aqueous solution, then sublimate away water at 200 K. Confirmed by FTIR spectroscopy. The result of this preparation route is termed β−H₂CO₃. However, α−H₂CO₃ is later argued to be (CH₃)O(COOH), a monomethyl ester.

This then leaves the question of why H₂CO₃ has "surprising kinetic stability", and why it took so long for solid-state H₂CO₃ to be isolated. An explanation using transition state theory is that pure H₂CO₃ is in fact rather metastable. At 300 K, its half-life is theoretically predicted to be ~10⁵ years, but in the presence of 2 water molecules, it decreases to ~1 minute. In short, water catalyzes the decomposition of H₂CO₃.

According to quantum chemical calculations, at room temperature (300 K), pure carbonic acid is expected to be a kinetically stable gas. 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. Chemically, it behaves as a diprotic Brønsted acid.

Carbonic acid monomers exhibit three conformational isomers: cis–cis, cis–trans, and trans–trans.

At low temperature and atmospheric pressure, solid carbonic acid is amorphous and lacks Bragg peaks in X-ray diffraction. But at high pressure, carbonic acid crystallizes, and modern analytical spectroscopy can measure its geometry, which under certain conditions have space group Cmc2₁.

According to neutron diffraction of dideuterated carbonic acid (D
₂CO
₃) in a hybrid clamped cell (Ni–Cr–Al 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. The crystal has space group P2₁/C. Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid, where the distances exceed 2.4 Å. Trimers and higher polymers are predicted to be even more stable. Based on the stability, it was suggested that solid carbonic acid can be used to sequestrate carbon.

In the presence of even a slight amount of water, carbonic acid dehydrates to carbon dioxide and water, which then catalyzes further decomposition.

The hydration equilibrium constant at 25 °C is [H
₂CO
₃]/[CO₂] ≈ 1.7×10⁻³ in pure water and ≈ 1.2×10⁻³ in seawater. 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 biological solutions

In the presence of the enzyme carbonic anhydrase, equilibrium is instead reached rapidly, and the following reaction takes precedence:

When the carbon dioxide concentration exceeds its solubility, gas evolves, and a third equilibrium must also be considered. The equilibrium constant for this reaction is defined by Henry's law.

The two reactions can be combined for the equilibrium in solution: 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.

In water pH control

In wastewater treatment and agricultural irrigation, carbonic acid is used to acidify the water, similar to sulfuric acid and sulfurous acid produced by sulfur burners.

Under high CO₂ partial pressure

In the beverage industry, sparkling or "fizzy water" is usually referred to as carbonated water. It is made by dissolving carbon dioxide in water under a small positive pressure. Many soft drinks are treated the same way and 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. 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 than ice, is expected to have sunk beneath the ice layers and separated them from the rocky cores of these moons.

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: where brackets indicate the concentration of species. At 25 °C, these equilibria empirically satisfylog(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: Direct values for these constants in the literature include pK₁ = 6.35 and pK₂ - pK₁ = 3.49.

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.

Ocean acidification

A Bjerrum plot (also known as a Sillén diagram) shows typical equilibrium concentrations of carbon dioxide and its derived species in solution (or seawater) 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.