Carbonic acid

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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. However, in the absence of water, it is quite stable at room temperature.<ref>Template:Greenwood&Earnshaw2nd</ref><ref name="lo">Template:Cite journal</ref> The interconversion of carbon dioxide and carbonic acid is related to the breathing cycle of animals and the acidification of natural waters.<ref name=neutron/>

In biochemistry and physiology, the name "carbonic acid" is sometimes applied to aqueous solutions of carbon dioxide. These chemical species play an important role in the bicarbonate buffer system, used to maintain acid–base homeostasis.<ref>Acid-Base Physiology 2.1 – Acid-Base Balance by Kerry Brandis.</ref>

Terminology in biochemical literature

In chemistry, the term "carbonic acid" strictly refers to the chemical compound with the formula Template:Chem. 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.

Anhydrous carbonic acid

At ambient temperatures, pure carbonic acid is a stable gas.<ref name="lo" /> 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.<ref name="sublime" /> Chemically, it behaves as a diprotic Brønsted acid.<ref name=peroxide /><ref name="Andersen" />

Carbonic acid monomers exhibit three conformational isomers: cis–cis, cis–trans, and trans–trans.<ref>Template:Cite journal</ref>

At low temperatures and atmospheric pressure, solid carbonic acid is amorphous and lacks Bragg peaks in X-ray diffraction.<ref name="wi">Template:Cite journal</ref> But at high pressure, carbonic acid crystallizes, and modern analytical spectroscopy can measure its geometry.

According to neutron diffraction of dideuterated carbonic acid (Template:Chem) 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.<ref name="neutron" /> Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid, where the distances exceed 2.4 Å.<ref name="wi" />

In aqueous solution

In even a slight presence of water, carbonic acid dehydrates to carbon dioxide and water, which then catalyzes further decomposition.<ref name="lo" /> For this reason, carbon dioxide can be considered the carbonic acid anhydride.

The hydration equilibrium constant at 25 °C is Template:Awrap in pure water<ref name="HS">Template:Cite book</ref> and ≈ 1.2×10⁻³ in seawater.<ref name="SB">Template:Cite journal</ref> 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:<ref name="Lindskog_1997">Template:Cite journal</ref> <chem display=block>HCO3^- {+} H^+ <=> CO2 {+} H2O</chem>

When the created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium <chem display=block>CO_2 (soln) <=> CO_2 (g)</chem> 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: <math chem="" display="block">\begin{align} \ce{HCO3^{-}{} + H+{} <=> CO2(soln){} + H2O} && K_3 = \frac{[\ce{H+}][\ce{HCO3^-}]}{[\ce{CO2(soln)}]} \end{align}</math> 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.<ref>Template:Cite journal</ref>

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.<ref>Template:Cite web</ref>

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.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> 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.<ref name="Saleh-Scirep">Template:Cite journal</ref>

Relationship to bicarbonate and carbonate

File:Weak acid speciation.svg

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: <math chem display="block">\begin{align} \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{align}</math> where brackets indicate the concentration of species. At 25 °C, these equilibria empirically satisfy<ref>IUPAC (2006). "Stability constants" (database). </ref><math display="block">\begin{alignat}{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{alignat}</math>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:<math display="block">\begin{alignat}{3} p\text{K}_1 &= \log(\beta_2) - \log(\beta_1) &= 6.77 \\ p\text{K}_2 &= \log(\beta_1) &= 9.93 \end{alignat}</math> Direct values for these constants in the literature include Template:Math and Template:Math.<ref>Template:Cite journal</ref>

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

File:Carbonate system of seawater.svg

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.<ref name="peroxide">Template:Cite journal</ref><ref name="Andersen">Template:Cite journal</ref> 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.<ref name="cald03">Template:Cite journal</ref><ref name="sabine">Template:Cite journal</ref> 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.

References

External links

Template:Inorganic compounds of carbon Template:Carbonates Template:Hydrogen compounds Template:Oxides of carbon

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