Carbonic acid

Carbonic acid is a chemical compound with the chemical formula Template:Chem2. The molecule rapidly converts to water and carbon dioxide in the presence of water. In constrast to early-twentieth century textbooks, numerous studies conducted since ca. 1990 identified Template:Chem2 as a real molecule with a distinct Raman spectrum and with a first-order life-time of ca. 20 ms at 37 °C. Solid anhydrous carbonic acid has also been isolated.

The interconversion of carbon dioxide and carbonic acid is related to the breathing process of all aerobic organisms and to the acidification of natural waters.

History

During the early days of chemistry, Template:Chem2 received multiple names. Led by Lavoisier, a group of chemists systematized the nomenclature of chemicals then known. Template:Chem2 was named "carbonic acid" in 1781, because it was produced by combustion of charcoal (charbon in French), and it was considered an acid. Back then, an acid was something that, when dissolved in water, produces red in a litmus test and a distinctive acidic taste. Later, chemists reconceptualized "acid", and no longer accepted Template:Chem2 to be an acid, and "carbonic acid" became the systematic name of Template:Chem2 instead. It was thought that it Template:Chem2 could not exist independently, and could only exist in a solution.

In the 1960s, experimenters produced adducts of Template:Chem2. The diethyl ether adduct Template:Chem2 was prepared at -78 °C. The dimethyl ether adduct Template:Chem2 was synthesized at -30 °C using dimethyl ether and sodium carbonate, and confirmed by chemical, thermochemical, IR, and NMR methods. The pure Template:Chem2 is a solid that decomposes at 5 °C.

In 1987, gaseous Template:Chem2 was prepared by thermolysis of [[Ammonium bicarbonate|Template:Chem2]], and confirmed by IR and collisional activation mass spectrometry. In 1991, solid Template:Chem2 was prepared by solid irradiation of a 1:1 Template:Chem2 ice mixture at 20 K. The result of this preparation route is termed Template:Chem2. Another synthesis route was to begin by reacting [[Hydrogen bromide|Template:Chem2]] and [[Potassium bicarbonate|Template:Chem2]] in an aqueous solution, then sublimate away water at 200 K. Confirmed by FTIR spectroscopy. The result of this preparation route is termed Template:Chem2. However, Template:Chem2 is later argued to be Template:Chem2, a monomethyl ester.

This then leaves the question of why Template:Chem2 has "surprising kinetic stability", and why it took so long for solid-state Template:Chem2 to be isolated. An explanation using transition state theory is that pure Template:Chem2 is in fact rather metastable. At 300 K, its half-life is theoretically predicted to be ~10⁵ years, but the half-life in the presence of 2 water molecules decreases to ~1 minute. In short, water catalyzes the decomposition of Template:Chem2.

Anhydrous carbonic acid

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 (Template:Chem) 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 aqueous solution

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 Template:Awrap 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 Template: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: $$\mathrm{H}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{-}\text{+}\mathrm{H}{\phantom{A}}^{+}\stackrel{-\rightharpoonup }{\leftharpoondown -}\mathrm{C}\mathrm{O}{\phantom{A}}_2\text{+}\mathrm{H}{\phantom{A}}_2\mathrm{O}$$

When the created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium $$\mathrm{C}\mathrm{O}{\phantom{A}}_2(\mathrm{s}\mathrm{o}\mathrm{l}\mathrm{n})\stackrel{-\rightharpoonup }{\leftharpoondown -}\mathrm{C}\mathrm{O}{\phantom{A}}_2(\mathrm{g})$$ 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: $$\begin{array}{ccc}\mathrm{H}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{-}+\mathrm{H}{\phantom{A}}^{+}\stackrel{-\rightharpoonup }{\leftharpoondown -}\mathrm{C}\mathrm{O}{\phantom{A}}_2(\mathrm{s}\mathrm{o}\mathrm{l}\mathrm{n})+\mathrm{H}{\phantom{A}}_2\mathrm{O}& & K_3=\frac{[\mathrm{H}{\phantom{A}}^{+}][\mathrm{H}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{-}]}{[\mathrm{C}\mathrm{O}{\phantom{A}}_2(\mathrm{s}\mathrm{o}\mathrm{l}\mathrm{n})]}\end{array}$$ 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 agriculture 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 under a small positive pressure in water. Many soft drinks treated the same way effervesce.

Significant amounts of molecular Template:Chem 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 the ice, is expected to have sunk beneath the ice layers and to separate them from the rocky cores of these moons.

Relationship to bicarbonate and carbonate

Bjerrum plot of speciation for a hypothetical monoprotic acid: AH concentration as a function of the difference between Template:Math and Template:Math

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 Template:Mvar. The two equilibria most easily measured are as follows: $$\begin{array}{ccc}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{2-}+\mathrm{H}{\phantom{A}}^{+}\stackrel{-\rightharpoonup }{\leftharpoondown -}\mathrm{H}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{-}& & {\beta }_1=\frac{[\mathrm{H}\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{-}]}{[\mathrm{H}{\phantom{A}}^{+}][\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{2-}]}\\ \mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{2-}+2\mathrm{H}{\phantom{A}}^{+}\stackrel{-\rightharpoonup }{\leftharpoondown -}\mathrm{H}{\phantom{A}}_2\mathrm{C}\mathrm{O}{\phantom{A}}_3& & {\beta }_2=\frac{[\mathrm{H}{\phantom{A}}_2\mathrm{C}\mathrm{O}{\phantom{A}}_3]}{[\mathrm{H}{\phantom{A}}^{+}{]}^2[\mathrm{C}\mathrm{O}{\phantom{A}}_3{\phantom{A}}^{2-}]}\end{array}$$ where brackets indicate the concentration of species. At 25 °C, these equilibria empirically satisfy$$\begin{array}{cccccccccc}\log ({\beta }_1)=& & 0& .54& I^2-0& .96& I+& & 9& .93\\ \log ({\beta }_2)=& & -2& .5& I^2-0& .043& I+& & 16& .07\end{array}$$Template:Math decreases with increasing Template:Mvar, as does Template:Math. In a solution absent other ions (e.g. Template:Math), these curves imply the following stepwise dissociation constants:$$\begin{array}{ccc}p{\text{K}}_1& =\log ({\beta }_2)-\log ({\beta }_1)& =6.77\\ p{\text{K}}_2& =\log ({\beta }_1)& =9.93\end{array}$$ Direct values for these constants in the literature include Template:Math and Template:Math.

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

Ocean acidification

Carbonate speciation in seawater (ionic strength 0.7 mol/dm³). The expected change shown is due to the current anthropogenic increase in atmospheric carbon dioxide concentration.

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.

Further reading

• Template:WsPSM2
• Template:Cite journal
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External links

• 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