The chemical formula is: HCO3- + H+ = H20 + CO2
In chemistry, an oxocarbon anion
is a negative ion consisting solely of carbon and oxygen atoms, and therefore having the general formula Cx
for some integers x
, and n
The most common oxocarbon anions are carbonate, CO3
2−, and oxalate, C2
2−. There is however a large number of stable anions in this class, including several ones that have research or industrial use. There are also many unstable anions, like CO2
− and CO4
−, that have a fleeting existence during some chemical reactions; and many hypothetical species, like CO4
4−, that have been the subject of theoretical studies but have yet to be observed.
Stable oxocarbon anions form salts with a large variety of cations. Unstable anions may persist in very rarefied gaseous state, such as in interstellar clouds. Most oxocarbon anions have corresponding moieties in organic chemistry, whose compounds are usually esters. Thus, for example, the oxalate moiety [–O–(C=O–)2
–O–] occurs in the ester dimethyl oxalate H3
In many oxocarbon anions each of the extra electrons responsible for the negative electric charges behaves as if it were distributed over several atoms. Some of the electron pairs responsible for the covalent bonds also behave as if they were delocalized. These phenomena are often explained as a resonance between two or more conventional molecular structures that differ on the location of those charges and bonds. The carbonate ion, for example, is considered to have an "average" of three different structures
so that each oxygen has the same negative charge equivalent to 2/3 of one electron, and each C–O bond has the same average valence of 4/3. This model accounts for the observed threefold symmetry of the anion.
Similarly, in a deprotonated carboxyl group –, each oxygen is often assumed to have a charge of −1/2 and each C–O bond to have valence 3/2, so the two oxygens are equivalent. The croconate anion also has fivefold symmetry, that can be explained as the superposition of five states leading to a charge of −2/5 on each oxygen. These resonances are believed to contribute to the stability of the anions.
An oxocarbon anion Cx
can be seen as the result of removing all protons from a corresponding acid Cx
. Carbonate CO3
2−, for example, can be seen as the anion of carbonic acid H2
. Sometimes the "acid" is actually an alcohol or other species; this is the case, for example, of acetylenediolate C2
2− that would yield acetylenediol C2
. However, the anion is often more stable than the acid (as is the case for carbonate); and sometimes the acid is unknown or is expected to be extremely unstable (as is the case of methanetetracarboxylate C(COO−)4
Every oxocarbon anion Cx
can be matched in principle to the electrically neutral (or oxidized) variant Cx
, an oxocarbon (oxide of carbon) with the same composition and structure except for the negative charge. As a rule, however, these neutral oxocarbons are less stable than the corresponding anions. Thus, for example, the stable carbonate anion corresponds to the extremely unstable neutral carbon trioxide CO3
; oxalate C2
2− correspond to the even less stable 1,2-dioxetanedione C2
; and the stable croconate anion C5
2− corresponds to the neutral cyclopentanepentone C5
, which has been detected only in trace amounts.
Conversely, some oxocarbon anions can be reduced to yield other anions with the same structural formula but greater negative charge. Thus rhodizonate C6
2− can be reduced to the tetrahydroxybenzoquinone (THBQ) anion C6
4− and then to benzenehexolate C6
An oxocarbon anion Cx
can also be associated with the anhydride of the corresponding acid. The latter would be another oxocarbon with formula Cx
; namely, the acid minus n
/2 water molecules H2
O. The standard example is the connection between carbonate CO3
2− and carbon dioxide CO2
. The correspondence is not always well-defined since there may be several ways of performing this formal dehydration, including joining two or more anions to make an oligomer or polymer. Unlike neutralization, this formal dehydration sometimes yields fairly stable oxocarbons, such as mellitic anhydride C12
from mellitate C12
6− via mellitic acid C12
For each oxocarbon anion Cx
there are in principle n
−1 partially hydrogenated anions with formulas Hk
(, where k
ranges from 1 to n
−1. These anions are generally indicated by the prefixes "hydrogen"-, "dihydrogen"-, "trihydrogen"-, etc. Some of them, however, have special names: hydrogencarbonate is commonly called bicarbonate, and hydrogenoxalate is known as binoxalate.
The hydrogenated anions may be stable even if the fully protonated acid is not (as is the case of bicarbonate).
The carbide anions, such as acetylide C2
2− and methanide C4−, could be seen as extreme cases of oxocarbon anions Cx
, with y
equal to zero. The same could be said of oxygen-only anions such as oxide O2−, superoxide, O2
−, peroxide, O2
2−, and ozonide O3
Here is an incomplete list of the known or conjectured oxocarbon anions
Several other oxocarbon anions have been detected in trace amounts, such as , a singly ionized version of rhodizonate.
In inorganic chemistry, bicarbonate
(IUPAC-recommended nomenclature: hydrogen carbonate
) is an intermediate form in the deprotonation of carbonic acid. It is an anion with the chemical formula HCO3
Bicarbonate serves a crucial biochemical role in the physiological pH buffering system.
The bicarbonate ion (hydrogen carbonate ion) is an anion with the empirical formula HCO3
− and a molecular mass of 61.01 daltons; it consists of one central carbon atom surrounded by three oxygen atoms in a trigonal planar arrangement, with a hydrogen atom attached to one of the oxygens. It is isoelectronic with nitric acid . The bicarbonate ion carries a negative one formal charge and is the conjugate base of carbonic acid ; it is the conjugate acid of , the carbonate ion, as shown by these equilibrium reactions.
A bicarbonate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound. Many bicarbonates are soluble in water at standard temperature and pressure, in particular sodium bicarbonate contributes to total dissolved solids, a common parameter for assessing water quality.]
Bicarbonate is alkaline, and a vital component of the pH buffering system of the human body (maintaining acid-base homeostasis). 70-75% of CO2
in the body is converted into carbonic acid (H2
), which can quickly turn into bicarbonate (HCO3
With carbonic acid as the central intermediate species, bicarbonate – in conjunction with water, hydrogen ions, and carbon dioxide – forms this buffering system, which is maintained at the volatile equilibrium required to provide prompt resistance to drastic pH changes in both the acidic and basic directions. This is especially important for protecting tissues of the central nervous system, where pH changes too far outside of the normal range in either direction could prove disastrous (see acidosis or alkalosis).
Bicarbonate also acts to regulate pH in the small intestine. It is released from the pancreas in response to the hormone secretin to neutralize the acidic chyme entering the duodenum from the stomach.
In freshwater ecology, strong photosynthetic activity by freshwater plants in daylight releases gaseous oxygen into the water and at the same time produces bicarbonate ions. These shift the pH upward until in certain circumstances the degree of alkalinity can become toxic to some organisms or can make other chemical constituents such as ammonia toxic. In darkness, when no photosynthesis occurs, respiration processes release carbon dioxide, and no new bicarbonate ions are produced, resulting in a rapid fall in pH.
The most common salt of the bicarbonate ion is sodium bicarbonate, NaHCO3
, which is commonly known as baking soda. When heated or exposed to an acid such as acetic acid (vinegar), sodium bicarbonate releases carbon dioxide. This is used as a leavening agent in baking.
The flow of bicarbonate ions from rocks weathered by the carbonic acid in rainwater is an important part of the carbon cycle.
Bicarbonate also serves much in the digestive system. It raises the internal pH of the stomach, after highly acidic digestive juices have finished in their digestion of food. Ammonium bicarbonate is used in digestive biscuit manufacture.
In diagnostic medicine, the blood value of bicarbonate is one of several indicators of the state of acid-base physiology in the body.
The parameter standard bicarbonate concentration
) is the bicarbonate concentration in the blood at a 2
P of 40 mmHg (5.33 kPa), full oxygen saturation and 36 °C.
noco/acba/cong/tumr, sysi/epon, urte
proc/itvp, drug (G4B), blte, urte
noco (d)/cong/tumr, sysi/epon
proc, drug (A10/H1/H2/H3/H5)
noco/cong/tumr, sysi/epon, injr
proc, drug (C1A/1B/1C/1D), blte
anat (t, g, p)/phys/devp/enzy
proc, drug (A2A/2B/3/4/5/6/7/14/16), blte
Calcium hydrogen carbonate
, also called calcium hydrogen carbonate
, has a chemical formula Ca(HCO3
. The term does not refer to a known solid compound; it exists only in aqueous solution containing the calcium (Ca2+), bicarbonate (HCO3
–), and carbonate (CO3
2–) ions, together with dissolved carbon dioxide (CO2
). The relative concentrations of these carbon-containing species depend on the pH; bicarbonate predominates within the range 6.36-10.25 in fresh water.
All waters in contact with the atmosphere absorb carbon dioxide, and as these waters come into contact with rocks and sediments they acquire metal ions, most commonly calcium and magnesium, so most natural waters that come from streams, lakes, and especially wells, can be regarded as dilute solutions of these bicarbonates. These hard waters tend to form carbonate scale in pipes and boilers and they react with soaps to form an undesirable scum.
Attempts to prepare compounds such as calcium bicarbonate by evaporating its solution to dryness invariably yield the solid carbonate instead: Ca(HCO3
(aq) → CO2
(g) + H2
O(l) + 3
CaCO(s). Very few solid bicarbonates other than those of the alkali metals and ammonium ion are known to exist.
The above reaction is very important to the formation of stalactites, stalagmites, columns, and other speleothems within caves and, for that matter, in the formation of the caves themselves. As water containing carbon dioxide (including extra CO2
acquired from soil organisms) passes through limestone or other calcium carbonate containing minerals, it dissolves part of the calcium carbonate and hence becomes richer in bicarbonate. As the groundwater enters the cave, the excess carbon dioxide is released from the solution of the bicarbonate, causing the much less soluble calcium carbonate to be deposited.
Dissolved carbon dioxide (CO2
) in rainwater (H2
O) reacts with limestone, calcium carbonate (CaCO3
) to form soluble calcium bicarbonate (Ca(HCO3
). This soluble compound is then washed away with the rainwater. This is form of weathering is called 'Carbonation'.
175–176 °C (dec.)
180–187 °C (dec.)
or butynedioic acid
is an organic compound (a dicarboxylic acid) with the formula C4
H. It is a crystalline solid that is soluble in diethyl ether.
The removal of two protons yields the acetylenedicarboxylate dianion C4
2−, which consists only of carbon and oxygen, making it an oxocarbon anion. Partial ionization yields the monovalent hydrogenacetylenedicarboxylate anion HC4
The acid was first described in 1877 by Polish chemist Ernest Bandrowski. It can be obtained by treating α,β-dibromosuccinic acid with potassium hydroxide KOH in methanol or ethanol. The reaction yields potassium bromide and potassium acetylenedicarboxylate. The salts are separated and the latter is treated with sulfuric acid.
Acetylenedicarboxylic acid is used in the synthesis of dimethyl acetylenedicarboxylate, an important laboratory reagent. Both the acid and the monobasic salt potassium hydrogenacetylenedicarboxylate KC4
are commonly traded as laboratory chemicals.
Carbon dioxide solution; Dihydrogen carbonate; acid of air; Aerial acid; Hydroxymethanoic acid
is the chemical compound with the formula H2
). It is also a name sometimes given to solutions of carbon dioxide in water (carbonated water), because such solutions contain small amounts of H2
. Carbonic acid, which is a weak acid, forms two kinds of salts, the carbonates and the bicarbonates.
When carbon dioxide dissolves in water it exists in chemical equilibrium producing carbonic acid:
The hydration equilibrium constant at 25°C is called Kh
, which in the case of carbonic acid is [H2
] ≈ 1.7×10−3 in pure water and ≈ 1.2×10−3 in seawater. Hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO2
molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2
O → H2
) and 23 s−1 for the reverse reaction (H2
O). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. The addition of two equivalents of water to CO2
would give orthocarbonic acid
, which exists only in minute amounts in aqueous solution.
Addition of base to an excess of carbonic acid gives bicarbonate. With excess base, carbonic acid reacts to give carbonate salts.
Carbonic acid is an intermediate step in the transport of CO2
out of the body via respiratory gas exchange. The hydration reaction of CO2
is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which both increases the reaction rate and dissociates a hydrogen ion (H+) from the resulting carbonic acid, leaving bicarbonate (HCO3
-) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO2
and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.
The oceans of the world have absorbed almost half of the CO2
emitted by humans from the burning of fossil fuels. The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels. This process is known as ocean acidification.
Carbonic acid is one of the polyprotic acids: It is diprotic - it has two protons, which may dissociate from the parent molecule. Thus, there are two dissociation constants, the first one for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3
Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution, carbonic acid exists in equilibrium with carbon dioxide, and the concentration of H2
is much lower than the concentration of CO2
. In many analyses, H2
includes dissolved CO2
(referred to as CO2
* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows:
Whereas this apparent pKa
is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO2
-containing solutions. A similar situation applies to sulfurous acid (H2
), which exists in equilibrium with substantial amounts of unhydrated sulfur dioxide.
The second constant is for the dissociation of the bicarbonate ion into the carbonate ion CO3
At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2
solution) is completely determined by the partial pressure
of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2
− and CO3
2−) as well as of the hydration equilibrium between dissolved CO2
(see above) and of the following equilibrium between the dissolved CO2
and the gaseous CO2
above the solution:
The corresponding equilibrium equations together with the
relation and the charge neutrality condition
result in six equations for the six unknowns [CO2
], [H+], [OH−], [HCO3
−] and [CO3
2−], showing that the composition of the solution is fully determined by
. The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:
Theoretical calculations show that the presence of even a single molecule of water causes carbonic acid to revert to carbon dioxide and water. In the absence of water, the dissociation of gaseous carbonic acid is predicted to be very slow, with a half-life of 180,000 years.
It has long been recognized that pure carbonic acid cannot be obtained at room temperatures (about 20 °C or about 70 °F). It can be generated by exposing a frozen mixture of water and carbon dioxide to high-energy radiation, and then warming to remove the excess water. The carbonic acid that remained was characterized by infrared spectroscopy. The fact that the carbonic acid was prepared by irradiating a solid H2
O + CO2
mixture may suggest that H2
might be found in outer space, where frozen ices of H2
O and CO2
are common, as are cosmic rays and ultraviolet light, to help them react. The same carbonic acid polymorph (denoted beta
-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuo
, which causes protonation of bicarbonate, followed by removal of the solvent. Alpha
-carbonic acid was prepared by the same technique using methanol rather than water as a solvent.
An oxocarbon or oxide of carbon is an inorganic compound consisting only of carbon and oxygen.
The simplest and most common oxocarbons are carbon monoxide (CO) and carbon dioxide (CO2). Many other stable or metastable oxides of carbon are known, but they are rarely encountered, such as carbon suboxide (C3O2 or O=C=C=C=O) and mellitic anhydride (C12O9).
While textbooks will often list only the first three, and rarely the fourth, a large number of other oxides are known today, most of them synthesized since the 1960s. Some of these new oxides are stable at room temperature. Some are metastable or stable only at very low temperatures, but decompose to simpler oxocarbons when warmed. Many are inherently unstable and can be observed only momentarily as intermediates in chemical reactions or are so reactive that they can exist only in the gas phase or under matrix isolation conditions.
The inventory of oxocarbons appears to be steadily growing. The existence of graphene oxide and of other stable polymeric carbon oxides with unbounded molecular structures suggests that many more remain to be discovered.
Carbon dioxide (CO2) occurs widely in nature, and was incidentally manufactured by humans since pre-historical times, by the combustion of carbon-containing substances and fermentation of foods such as beer and bread. It was gradually recognized as a chemical substance, formerly called spiritus sylvestre ("forest spirit") or "fixed air", by various chemists in the 17th and 18th centuries.
Carbon monoxide may occur in combustion, too, and was used (though not recognized) since antiquity for the smelting of iron from its ores. Like the dioxide, it was described and studied in the West by various alchemists and chemists since the Middle Ages. Its true composition was discovered by William Cruikshank in 1800.
Carbon suboxide was discovered by Brodie in 1873, by passing electric current through carbon dioxide.
The fourth "classical" oxide, mellitic anhydride (C12O9), was apparently obtained by Liebig and Wöhler in 1830 in their study of mellite ("honeystone"), but was characterized only in 1913, by Meyer and Steiner.
Brodie also discovered in 1859 a fifth compound called graphite oxide, consisting of carbon and oxygen in ratios varying between 2:1 and 3:1; but the nature and molecular structure of this substance remained unknown until a few years ago, when it was renamed graphene oxide and became a topic of research in nanotechnology.
Notable examples of unstable or metastable oxides that were detected only in extreme situations are dicarbon monoxide radical (:C=C=O), carbon trioxide (CO3), carbon tetroxide (), carbon pentoxide (), carbon hexoxide () and 1,2-dioxetanedione (C2O4). Some of these reactive carbon oxides were detected within molecular clouds in the interstellar medium by rotational spectroscopy.
Many hypothetical oxocarbons have been studied by theoretical methods but have yet to be detected. Examples include oxalic anhydride (C2O3 or O=(C2O)=O), ethylene dione (C2O2 or O=C=C=O) and other linear or cyclic polymers of carbon monoxide (-CO-)n (polyketones), and linear or cyclic polymers of carbon dioxide (-CO2-)n, such as the dimer 1,3-dioxetanedione (C2O4) and the trimer 1,3,5-trioxanetrione (C3O6).
Normally carbon is tetravalent while oxygen is divalent, and in most oxocarbons (as in most other carbon compounds) each carbon atom may be bound to four other atoms, while oxygen may be bound to at most two. Moreover, while carbon can connect to other carbons to form arbitrarily large chains or networks, chains of three or more oxygens are rarely if ever observed. Thus the known electrically neutral oxocarbons generally consist of one or more carbon skeletons (including cyclic and aromatic structures) connected and terminated by oxide (-O-, =O) or peroxide (-O-O-) groups.
Carbon atoms with unsatisfied bonds are found in some oxides, such as the diradical C2O or :C=C=O; but these compounds are generally too reactive to be isolated in bulk. Loss or gain of electrons can result in monovalent negative oxygen (-), trivalent positive oxygen (≡), or trivalent negative carbon (≡). The last two are found in carbon monoxide, −C≡O+. Negative oxygen occurs in most oxocarbon anions.
One family of carbon oxides has the general formula CnO2, or O=(C=)nO — namely, a linear chain of carbon atoms, capped by oxygen atoms at both ends. The first members are
Some higher members of this family have been detected in trace amounts in low-pressure gas phase and/or cryogenic matrix experiments, specifically for n = 7:p.97 and n = 17, 19, and 21.:p.95
Another family of oxocarbons are the linear carbon monoxides CnO. The first member, ordinary carbon monoxide CO, seems to be the only one that is stable in the pure state at room temperature. Photolysis of the linear carbon dioxides in a cryogenic matrix leads to loss of CO, resulting in detectable amounts of even-numbered monoxides such as C2O, C4O, and C6O. The members up to n=9 have also been obtained by electrical discharge on gaseous C3O2 diluted in argon. The first three members have been detected in interstellar space.
When n is even, the molecules are believed to be in the triplet (cumulene-like) state, with the atoms connected by double bonds and an unfilled orbital in the first carbon — as in :C=C=O, :C=C=C=C=O, and, in general, :(C=)n=O. When n is odd, the triplet structure is believed to resonate with a singlet (acetylene-type) polar state with a negative charge on the carbon end and a positive one on the oxygen end, as in −C≡C-C≡O+, −C≡C-C≡C-C≡O+, and, in general, −(C≡C-)(n-1)/2C≡O+. Carbon monoxide itself follows this pattern: its predominant form is believed to be −C≡O+.
Another family of oxocarbons that has attracted special attention are the cyclic radialene-type oxocarbons CnOn or (CO)n. They can be regarded as cyclic polymers of carbon monoxide, or n-fold ketones of n-carbon cycloalkanes. Carbon monoxide itself (CO) can be regarded as the first member. Theoretical studies indicate that ethylene dione (C2O2 or O=C=C=O) and cyclopropanetrione C3O3 do not exist. The next three members — 4O4C, 5O5C, and 6O6C — are theoretically possible, but are expected to be quite unstable, and so far they have been synthesized only in trace amounts.
On the other hand, the anions of these oxocarbons are quite stable, and some of them have been known since the 19th century. They are
The cyclic oxide C6O6 also forms the stable anions of tetrahydroxy-1,4-benzoquinone (C6O64−) and benzenehexol (C6O66−), The aromaticity of these anions has been studied using theoretical methods.
Many new stable or metastable oxides have been synthesized since the 1960s, such as:
Many relatives of these oxides have been investigated theoretically, and some are expected to be stable, such as other carbonate and oxalate esters of tetrahydroxy-1,2-benzoquinone and of the rhodizonic, croconic, squaric, and deltic acids.
Carbon suboxide spontaneously polymerizes at room temperature into a carbon-oxygen polymer, with 3:2 carbon:oxygen atomic ratio. The polymer is believed to be a linear chain of fused six-membered lactone rings, with a continuous carbon backbone of alternating single and double bonds. Physical measurements indicate that the mean number of units per molecule is about 5–6, depending on the formation temperature.
Carbon monoxide compressed to 5 GPa in a diamond anvil cell yields a somewhat similar reddish polymer with a slightly higher oxygen content, which is metastable at room conditions. It is believed that CO disproportionates in the cell to a mixture of CO2 and C3O2; the latter forms a polymer similar to the one described above (but with a more irregular structure), that traps some of the CO2 in its matrix.
Another carbon-oxygen polymer, with C:O ratio 5:1 or higher, is the classical graphite oxide and its single-sheet version graphene oxide.
−111.3 °C, 161.9 K
6.8 °C, 280.0 K
, or tricarbon dioxide
, is an oxide of carbon with chemical formula C3
or O=C=C=C=O. Its four cumulative double bonds make it a cumulene. It is one of the stable members of the series of linear oxocarbons O=Cn
=O, which also includes carbon dioxide (CO2
) and pentacarbon dioxide (C5
The substance was discovered in 1873 by Benjamin Brodie by submitting carbon monoxide to an electric current. He claimed that the product was part of a series of "oxycarbons" with formulas Cx+1
, namely C, C2
, ..., and to have identified the last two; however only C3
is known. In 1891 Marcellin Berthelot observed that heating pure carbon monoxide at about 550 °C created small amounts of carbon dioxide but no trace of carbon, and assumed that a carbon-rich oxide was created instead, which he named "sub-oxide". He assumed it was the same product obtained by electric discharge and proposed the formula C2
O. Otto Diels later stated that the more organic names dicarbonyl methane and dioxallene were also correct.
It is commonly described as an oily liquid or gas at room temperature with an extremely noxious odor.
It is synthesized by warming a dry mixture of phosphorus pentoxide (P4
) and malonic acid or the esters of malonic acid. Therefore, it can be also considered as the anhydride of malonic anhydride, i.e. the "second anhydride" of malonic acid. Malonic anhydride (not to be confused with maleic anhydride) is a real molecule.
Several other ways for synthesis and reactions of carbon suboxide can be found in a review from 1930 by Reyerson.
Carbon suboxide polymerizes spontaneously to a red, yellow, or black solid. The structure is postulated to be poly(α-pyronic), similar to the structure in 2-pyrone (α-pyrone). In 1969, it was hypothesized that the color of Martian surface was caused by this compound; this was disproved by the Viking Mars probes.
Carbon suboxide is used in the preparation of malonates; and as an auxiliary to improve the dye affinity of furs.
A chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using a single line of chemical element symbols, numbers, and sometimes also other symbols, such as parentheses, dashes, brackets, and plus (+) and minus (−) signs. These are limited to a single typographic line of symbols, which may include subscripts and superscripts. A chemical formula is not a chemical name, and it contains no words. Although a chemical formula may imply certain simple chemical structures, it is not the same as a full chemical structural formula. Chemical formulas are more limiting than chemical names and structural formulas.
The simplest types of chemical formulas are called empirical formulas, which use only letters and numbers indicating atomic proportional ratios (the numerical proportions of atoms of one type to those of other types). Molecular formulas indicate the simple numbers of each type of atom in a molecule of a molecular substance, and are thus sometimes the same as empirical formulas (for molecules that only have one atom of a particular type), and at other times require larger numbers than do empirical formulas. An example of the difference is the empirical formula for glucose, which is CH2O, while its molecular formula requires all numbers to be increased by a factor of six, giving C6H12O6.