Carbonic acid
Chemical compound
- 6 min read
- 4 chapters
Related on Lably
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, CO2 received multiple 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 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 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 the half-life in the presence of 2 water molecules decreases to ~1 minute. In short, water catalyzes the decomposition of H2CO3.
Sign in to keep reading
3 more chapters are free with a Lably account.
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 (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 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 [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 created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium 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: 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 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 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
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
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.