Carbonic acid
From Freepedia
| Carbonic acid | |
|---|---|
| Image:Carbonic acid.png | |
| Other names | Carbon dioxide solution |
| Molecular formula | H2CO3 |
| SMILES | C(=O)(O)O |
| Molar mass | 62.03 g/mol |
| CAS number | 463-79-6 |
| Density | 1.0 g/cm3 (dilute solution) |
| Solubility (water) | exists only in solution |
| Acidity (pKa) | 3.60 (see text) 10.25 |
| Disclaimer and references | |
Carbonic acid is a carbon-containing acid with the formula H2CO3. It is also a name sometimes given to solutions of carbon dioxide in water, which contain small amounts of H2CO3. The salts of carbonic acids are called bicarbonates (or hydrogencarbonates) and carbonates.
Carbon dioxide dissolved in water is in equilibrium with carbonic acid:
- CO2 + H2O → H2CO3
The equilibrium constant at 25°C is 1.70×10−3: hence, the majority of the carbon dioxide is not converted into carbonic acid and stays as CO2 molecules. In the absence of a catalyst, the equlibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2 + H2O → H2CO3) and 23 s−1 for the reverse reaction (H2CO3 → CO2 + H2O).
The equilibrium between carbon dioxide and carbonic acid is important for controlling the acidity of body fluids, and almost all living organisms have an enzyme, carbonic anhydrase, which catalyzes the conversion between the two compounds, increasing the reaction rate by a factor of nearly 109.
Contents |
Acidity of carbonic acid
Carbonic acid has two acidic hydrogens and so two dissociation constants:
- H2CO3 ⇌ HCO3− + H+
- Ka1 = 2.5×10−4 L/mol; pKa1 = 3.60.
- H2CO3 ⇌ HCO3− + H+
- HCO3− ⇌ CO32− + H+
- Ka2 = 5.61×10−11 L/mol; pKa2 = 10.25.
- HCO3− ⇌ CO32− + H+
Care must be taken when quoting and using the first dissociation constant of carbonic acid. The value quoted above is correct for the H2CO3 molecule, and shows that it is a stronger acid than acetic acid or formic acid: this might be expected from the influence of the electronegative oxygen substituent. However, carbonic acid only ever exists in solution in equilibrium with carbon dioxide, and so the concentration of H2CO3 is much lower than the concentration of CO2, reducing the measured acidity. The equation may be rewritten as follows (c.f. sulfurous acid):
- CO2 + H2O ⇌ HCO3− + H+
- Ka = 4.30×10−7 L/mol; pKa = 6.36.
- CO2 + H2O ⇌ HCO3− + H+
This figure is often quoted as the dissociation constant of carbonic acid, although this is ambiguous: it might better be referred to as the acidity constant of carbon dioxide, as it is particularly useful for calculating the pH of CO2 solutions.
Instability of carbonic acid
It is usually not possible to obtain pure hydrogen bicarbonate as the presence of even a single molecule of water causes the carbonic acid to revert to carbon dioxide and water fairly quickly. However, pure carbonic acid has been found to be quite stable in the absence of water, with a calculated half-life of 180,000 years. There is a hypothetical acid orthocarbonic acid which is even more hydrated, being H4CO4.
Carbonic acid and rain water
A solution of carbon dioxide in water in equilibrium with the atmosphere (0.033% CO2) has a pH of 5.6. Rain water is normally not quite saturated in CO2, and has a pH of around 6 in the absence of atmospheric pollutants. This effect is separate from the phenomenon of acid rain, where industrial pollutants such as sulfur dioxide dissolve in rain water and lower its pH drastically. However, the acidity of rain water has important geological consequences for carbonate rocks such as chalk and limestone. An equilibrium is established between the calcium carbonate of the rock and calcium bicarbonate in solution:
- CaCO3 + CO2 + H2O ⇌ Ca(HCO3)2
This can erode underground caverns around fault lines which water runs down. As the calcium-rich water evaporates, the calcium carbonate precipitates, often as stalactites and stalagmites. Water drawn from chalk aquifers contains dissolved calcium bicarbonate, and is described as "hard".
See also
References
- Welch, M. J.; Lipton, J. F.; Seck, J. A. (1969). J. Phys. Chem. 73:3351.
- {{{Author|}}}{{|{{{3}}}}}}|show1| (1991)}}{{{{{Year|}}}}}}|show1|.}} {{|{{{3}}}}}}|show1|[{{{URL}}}}} Modern Inorganic Chemistry (2nd Edn.){{|{{{3}}}}}}|show1|]}}{{|{{{3}}}}}}|show1|, {{{Pages}}}}}{{|{{{3}}}}}}|Show1|, New York:MgGraw-Hill}}. {{{ID|}}}
External links
- Ask a Scientist: Carbonic Acid Decomposition
- Why was the existence of carbonic acid unfairly doubted for so long?
