Carbonate
Systematic name Carbonate
Properties
Molecular formula	CO2− 3
Molar mass	60.01 g mol-1
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

In chemistry, a carbonate is a salt of carbonic acid, characterized by the presence of the carbonate ion, CO2−
3. The name may also mean an ester of carbonic acid, an organic compound containing the carbonate group C(=O)(O–)₂.

The term is also used as a verb, to describe carbonation: the process of raising the concentrations of carbonate and bicarbonate ions in water to produce carbonated water and other carbonated beverages — either by the addition of carbon dioxide gas under pressure, or by dissolving carbonate or bicarbonate salts into the water.

In geology and mineralogy, the term "carbonate" can refer both to carbonate minerals and carbonate rock (which is made of chiefly carbonate minerals), and both are dominated by the carbonate ion, CO2−
3. Carbonate minerals are extremely varied and ubiquitous in chemically-precipitated sedimentary rock. The most common are calcite or calcium carbonate, CaCO₃, the chief constituent of limestone (as well as the main component of mollusc shells and coral skeletons); dolomite, a calcium-magnesium carbonate CaMg(CO₃)₂; and siderite, or iron(II) carbonate, FeCO₃, an important iron ore. Sodium carbonate ("soda" or "natron") and potassium carbonate ("potash") have been used since antiquity for cleaning and preservation, as well as for the manufacture of glass. Carbonates are widely used in industry, e.g. in iron smelting, as a raw material for Portland cement and lime manufacture, in the composition of ceramic glazes, and more.

Structure and bonding[link]

The carbonate ion is the simplest oxocarbon anion. It consists of one carbon atom surrounded by three oxygen atoms, in a trigonal planar arrangement, with D_3h molecular symmetry. It has a molecular mass of 60.01 daltons and carries a negative two formal charge. It is the conjugate base of the hydrogen carbonate (bicarbonate) ion, HCO₃⁻, which is the conjugate base of H₂CO₃, carbonic acid.

The Lewis structure of the carbonate ion has two (long) single bonds to negative oxygen atoms, and one short double bond to a neutral oxygen

This structure is incompatible with the observed symmetry of the ion, which implies that the three bonds are equally long and that the three oxygen atoms are equivalent. As in the case of the isoelectronic nitrate ion, the symmetry can be achieved by a resonance between three structures:

This resonance can be summarized by a model with fractional bonds and delocalized charges:

Chemical properties[link]

Metal carbonates generally decompose on heating, liberating carbon dioxide from the long term carbon cycle to the short term carbon cycle and leaving behind an oxide of the metal. This process is called calcination, after calx, the Latin name of quicklime or calcium oxide, CaO, which is obtained by roasting limestone in a lime kiln.

A carbonate salt forms when a positively charged ion, M⁺, attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound:

2 M⁺ + CO2−
3 → M₂CO₃

M²⁺ + CO2−
3 → MCO₃

2 M³⁺ + 3 CO2−
3 → M₂(CO₃)₃

Most carbonate salts are insoluble in water at standard temperature and pressure, with solubility constants of less than 1×10⁻⁸. Exceptions include lithium, sodium, potassium and ammonium carbonates, as well as many uranium carbonates.

In aqueous solution, carbonate, bicarbonate, carbon dioxide, and carbonic acid exist together in a dynamic equilibrium. In strongly basic conditions, the carbonate ion predominates, while in weakly basic conditions, the bicarbonate ion is prevalent. In more acid conditions, aqueous carbon dioxide, CO₂(aq), is the main form, which, with water, H₂O, is in equilibrium with carbonic acid - the equilibrium lies strongly towards carbon dioxide. Thus sodium carbonate is basic, sodium bicarbonate is weakly basic, while carbon dioxide itself is a weak acid. Note that although the carbonate salts of most metals are insoluble in water, the same is not true of the bicarbonate salts. This equilibrium between carbonate, bicarbonate, carbon dioxide and carbonic acid in water can, under changing temperature or pressure conditions, and in the presence of metal ions with insoluble carbonates, result in formation of insoluble compounds. This is responsible for the buildup of scale inside pipes caused by hard water.

Carbonated water is formed by dissolving CO₂ in water under pressure. When the partial pressure of CO₂ is reduced, for example when a can of soda is opened, the equilibrium for each of the forms of carbonate (carbonate, bicarbonate, carbon dioxide, and carbonic acid) shifts until the concentration of CO₂ in the solution is equal to the solubility of CO₂ at that temperature and pressure. In living systems an enzyme, carbonic anhydrase, speeds the interconversion of CO₂ and carbonic acid.

Organic carbonates[link]

In organic chemistry a carbonate can also refer to a functional group within a larger molecule that contains a carbon atom bound to three oxygen atoms, one of which is double bonded. These compounds are also known as organocarbonates or carbonate esters, and have the general formula ROCOOR′, or RR′CO₃. Important organocarbonates include dimethyl carbonate, the cyclic compounds ethylene carbonate and propylene carbonate, and the toxic gas-producing solid triphosgene.

Biological significance[link]

It works as a buffer in the blood as follows: when pH is too low, the concentration of hydrogen ions is too high, so one exhales CO₂. This will cause the equation to shift left, essentially decreasing the concentration of H⁺ ions, causing a more basic pH.

When pH is too high, the concentration of hydrogen ions in the blood is too low, so the kidneys excrete bicarbonate (HCO₃⁻). This causes the equation to shift right, essentially increasing the concentration of hydrogen ions, causing a more acidic pH.

There are 3 important reversible reactions that control the above pH balance.^[1]

1. H₂CO₃(aq) H⁺(aq) + HCO₃^-(aq)

2. H₂CO₃(aq) CO₂(aq) + H₂O(l)

3. CO₂(aq) CO₂(g)

Exhaled CO₂(g) depletes CO₂(aq) which in turn consumes H₂CO₃ causing the aforementioned shift left in the first reaction by Le Chatelier's principle. By the same principle when the pH is too high, the kidneys excrete bicarbonate (HCO₃^-) into urine as urea via the Urea Cycle (aka the Krebs-Henseleit Ornithine Cycle). By removing the bicarbonate more H⁺ is generated from carbonic acid (H₂CO₃) which come from CO₂(g) produced by cellular respiration.

Carbonate salts[link]

• Carbonate overview:

Presence outside Earth[link]

It is generally thought that the presence of carbonates in rock is strong evidence for the presence of liquid water. Recent observations of the Planetary nebula NGC 6302 shows evidence for carbonates in space,^[2] where aqueous alteration similar to that on Earth is unlikely. Other minerals have been proposed which would fit the observations.

Significant carbonate deposits have not been found on Mars via remote sensing or in situ missions, even though Martian meteorites contain small amounts. Groundwater may have existed at both Gusev^[3] and Meridiani Planum.^[4]

See also[link]

• Oxalate
• Sodium percarbonate
• Peroxocarbonate

References[link]

External links[link]

• Carbonate/bicarbonate/carbonic acid equilibrium in water: pH of solutions, buffer capacity, titration and species distribution vs. pH computed with a free spreadsheet