Carbon and its Inorganic Compounds
Created | Updated Jul 3, 2016
Carbon is second only to hydrogen in the number of compounds that it forms. Thus it has been estimated that there are of the order of 22 million organic (carbon-based) compounds1. These compounds also contain hydrogen but are not the main focus of this Entry. In this Entry the inorganic chemistry of carbon is discussed.
The number of the beast
Here is wisdom. Let him that hath understanding count the number of the beast: for it is the number of a man; and his number is six hundred threescore and six.
- Rev 13:18
Consideration of the atomic structure of carbon, symbolised in the box above, reveals that it contains 6 protons, 6 neutrons and 6 electrons; ie can be represented by the number 666 - the 'number of the beast'. This led a correspondent to the Daily Mail (John Whitehead, Letters, 12th March, 2003) to suggest that:
This seems to be a perfect description of science, which is meddling with the genetic code despite having no coherent philosophy of being.
Some people also make a link to the role of carbon dioxide in the potentially devastating consequences of global warming.
Carbon makes up 0.08% of the Earth's crust, occurring in its elementary state as graphite, diamond and coal. It also occurs in the lithosphere as mineral carbonates, such as limestone (CaCO3), magnesite and dolomite. The atmosphere contains approximately 0.03% carbon dioxide by volume from which a huge variety of biological compounds are derived. For example, living tissue consists of organic molecules such as proteins. The adult human body contains about 10kg of carbon, which is sufficient to make 9,000 lead pencils.
The element principally exists in two isotopic forms, 12C (~99%) and 13C (~1%). Despite its low abundance, 13C is an important isotope. The magnetic properties of its nucleus enable it to act as an effective natural marker for the characterisation of carbon-containing compounds in 13C nuclear magnetic resonance (NMR) spectroscopy.
Another isotope also exists, 14C, which is radioactive. This isotope is incorporated only into living organisms. It is formed in the atmosphere by the action of cosmic rays on 14N nuclei and ends up as 14CO2. Since the level of cosmic rays is pretty much constant, this labelled carbon dioxide maintains a steady concentration in the atmosphere and is taken up by plants during photosynthesis, thereby incorporating 14C atoms in their structures. Because of this, all plants and animals have a minute but measurable quantity of this isotope (1.2 x 10-10% of total carbon). The nuclei are radioactive with a half-life of 5730 years. Since the 14C content of an organism is continually regenerated during its life but stops in death, the radioactive decay of the element can be used to find its age. Radio-carbon dating therefore finds applications in archaeology for the dating of human remains, fabrics and leathers and other organic materials to a practical limit of 50,000 years.
Bonding in carbon compounds
The chemistry of carbon is almost totally covalent in nature though some Ionic species are known. The C4+ is not observed but C4- may exist in compounds which are called carbides formed with highly electropositive metals.
An important feature of the chemistry of carbon and which forms the basis of organic chemistry, is the ability of the element to catenate. This is its ability to form bonds to itself leading to highly stable chains or rings. This is chiefly due to the high strength of the C-C single bond (356 kJ mol-1). Along with the other first row non-metals, carbon can form C=C double bonds and C≡C triple bonds through sideways overlap of atomic p orbitals and these can be incorporated into the carbon-carbon bonded frameworks of its compounds, giving rigidity since these bonds can't rotate. The combined effects of catenation and multiple bond formation also allows the formation of planar cyclic aromatic2 compounds such as benzene. Other elements such as sulphur and silicon can also catenate, but far less effectively due to much weaker bonds (Si-Si 226 kJ mol-1). Multiple bonding for silicon is also highly unstable due to poor π overlap of its p orbitals. Because of these properties of carbon, it is not difficult to see why nature uses carbon as the element of choice to give structure to its biomolecules.
Allotropes of Carbon
There are several known allotropes (natural structural forms of the element) of carbon, which include graphite, diamond, linear acetylenic carbon (LAC) and the fullerenes.
Graphite is the most thermodynamically stable allotrope of carbon. The structure consist of sheets of sp2 hybridised carbon atoms in interlocking hexagons (see figure 1). The remaining p orbital of each carbon atom engages in π bonding, and thus each sheet may be regarded as a fused system of benzene rings. The delocalised π electron of each atom accounts for the ability of graphite to conduct electricity. Sheets of these hexagonal arrays of atoms stack on top of each other, the sheets being about 3.35 Angstroms apart. This separation is about twice the covalent radius of carbon (1.54 Angstroms). The sheets are held together by relatively weak van der Waals forces. The wide separation accounts for the low density of graphite (2.2g cm-3) compared to diamond (3.51g cm-3). This structure means that graphite is quite soft and is able to cleave easily between the layers. The layers are also able to slide over each other quite easily. Thus graphite has lubricating properties, in contrast to the abrasive properties of diamond. Thus graphite finds application as the 'lead' in pencils and as a component of penetrating oils.
Graphite occurs naturally in two forms α-graphite and β-graphite. Both have identical material properties except for their crystalline structure. The alpha type exhibits a hexagonal crystalline structure, while the beta type has a rhombohedral structure. The hexagonal structure occurs almost exclusively in synthetic materials.
Diagram of the structure of a layer of graphite.
Each carbon-carbon bond has partial double bond
character as the p electron of each atom is delocalised
over the whole sheet
Intercalation compounds of graphite
Despite being more stable than diamond, it is more reactive and, even under mild conditions, can undergo a number of unusual reactions due to the reactivity of the delocalised π electrons. Thus it is possible for atoms, ions or molecules to penetrate into the graphite lattice in such a way that only the hexagonal layer sheets are displaced perpendicular to the layer planes. This produces a swelling of the graphite crystal lattice and forms a group of compounds known as intercalation compounds. The first reported material of this type was the potassium intercalation compound C8K, discovered in 1926.
The intercalation compounds of graphite fall into two main groups according to the nature of the chemical bonding:
Group 1. Covalent Bonding Between C-Atoms and Intercalation Atoms (Non-Conducting)
(a) Graphite Oxide - produced by oxidising graphite with a mixture of perchloric acid, conc. sulphuric acid and conc. nitric acid.
Ideally, graphite oxide should have the formula C2O but, in practice, the oxygen atoms are associated with hydrogen atoms to form ether or carboxyl groups. Thus the C:O ratio rarely approaches 2:1. The probable structure involves a buckling of the hexagonal sheets.
(b) Carbon Monofluoride - whereas soot, wood charcoal and so on will burn in fluorine gas at room temperature, graphite is stable up to about 400°C. Between 420 and 460°C graphite and fluorine will react to form the grey monofluoride (CF)n. In practice, this compound is non-stoichiometric and has a composition range from CF0.68 to CF0.995. It is a white compound. The layer spacing is about 8 Angstroms and, again, the hexagonal sheets are buckled. If the temperature is increased to 700°C, a mixture of fluorocarbons is obtained.
(c) Tetracarbon monofluoride (TCMF) - If a fluorine/hydrogen fluoride mixture is passed over graphite at room temperature, the blackish-blue compound TCMF (C4F) is obtained. In this intercalation compound, the rings are not puckered and fluorine atoms are arranged in two layers, one above and one below each carbon plane. The distance between the graphite layers is 5.5 Angstroms.
(d) Graphite bisulphate - In the presence of strong oxidising agents, sulphuric acid reacts with graphite to form graphite bisulphate, a blue solution of formula (Cx)+HSO4-.2H2SO4. Here, the bisulphate ions are intercalated between giant cations formed by the hexagonal-structured sheets of carbon atoms.
Group 2. Ionic Bonding Between C-Atoms and Intercalation Atoms (Electrically Conducting)
The majority of graphitic compounds can be classified under this heading, which includes intercalation compounds of alkali metals and graphite, alkaline earth metals and graphite, graphitic salts, halides and their derivatives.
(a) Alkali metal - These compounds are fairly easily formed. Graphite reacts directly with molten alkali metals to form intercalation compounds of type C8M. Lithium and sodium are the least reactive metals of the group in the formation of these compounds, which are highly coloured (reddish) and extremely reactive.
When heated in vacuo they lose alkali metal to form, eventually, the blue C24M compound. In fact, this is only the first stage in the decomposition, further stages yielding compounds of formula C12nM, where n=2 for the first stage, 3 for the 2nd stage, etc.
Structurally, the metal atoms in C8M are triangularly packed between the graphite layers.
Uses of Group 2 Compounds are fairly limited, although the intercalation compound of lithium has found application in batteries.
Diamond is marginally less thermodynamically stable than graphite by just under 3 kJ mol-1 but is much less reactive. In the structure, each carbon atom is sp3 hybridised and forms four bonds to four other carbon atoms arranged tetrahedrally around it. Because of this more rigid structure diamond is a much harder substance than graphite. Diamond can be made synthetically from graphite by heating to about 2,730°C (3,000K) under a pressure of about 125,000 atmospheres using a metal catalyst such as chromium, iron or platinum. Synthetic diamonds have applications in machining tools.
Melting diamond involves breaking the strong covalent bonds which extend in all directions, hence the melting point is extraordinarily high - about 3,600°C (3,873K).
Otherwise known as buckminsterfullerenes or 'bucky balls', fullerenes are a class of carbon allotrope that was only recently discovered (c. 1990). They are produced by hitting graphite with powerful lasers or an electrical arc.
The fullerenes have very interesting structures. Consider a portion of a graphite-like sheet but with one of the hexagonal rings being replaced by a five-membered ring, bordered by five interlocking six-membered rings. In this instance we would have a curved rather than a flat sheet. If one now places several five membered rings in the structure bridged by six-membered rings the opposing edges of the sheet will meet to give a spherical or near spherical molecule, resembling a football. Indeed, a football is a perfect representation of C60 (Buckminsterfullerene).
Fullerenes come in a variety of forms with different numbers of carbon atoms, the most commonly studied being the perfectly spherical C60. The π electrons in the structure are not fully delocalised however but act chemically more like isolated alkenes. Because they have a cage-like structure, formation of a fullerene in the presence of other atoms, ions or compounds can lead to their encapsulation. Example include gases like helium and metals ranging from potassium to uranium.
By altering the conditions under which bucky balls are formed, it is possible to form structures similar to graphite sheets rolled into tubes, called carbon nanotubes. These promise a revolution in materials science and electronics. Thus, it is possible to insert quite large molecules, including biomolecules such as cytochrome C and beta-lactamase into nanotubes, giving the potential for developing biomedical devices such as biosensors. Bundles of carbon nanotubes are calculated to have a tensile strength of between 50 and 100 times that of steel, and to conduct electricity as well as copper. Hence, other potential application include molecular wires and hydrogen storage devices for hydrogen powered vehicles.
Carbon will form binary compounds, called carbides, with most elements. These generally fall into three classes; ionic (or salt-like), covalent and interstitial.
1. Ionic (salt-like) carbides
Ionic carbides are those formed with highly electropositive metals ie the alkali and alkaline earth metals and also the lanthanides, and metals like aluminium.
The ionic carbides are easily hydrolysed by water and can be classified according to the aliphatic hydrocarbon which they yield. The ionic carbides are formed mainly by elements from Groups 1, 2 and 3 of the Periodic Table.
For example Be2C; Al4C3
These are presumably ionised, containing metal cations and anionic carbon in the form C4- and are exceedingly reactive with water, giving methane. For example:
Al4C3 + 12 H2O → 4 Al(OH)3 + 3 CH4
Both these carbides are made by direct combination of the elements at about 1800K, and are much harder materials than the acetylides/ethynides (see below).
1(b) Acetylides (Ethynides)
This is the large group of acetylides, which have the carbon atoms in pairs, forming the anion [C≡C]2-.
These are formed by the elements in Groups 1 to 6 of the Periodic Table; especially Groups 1, 2 and 3, e.g:
Na2C2, Cu2C2, BeC2, CaC2, Al2C6.
These compounds are also exceedingly reactive with water, giving ethyne (acetylene), for example
CaC2 + 2 H2O → Ca(OH)2 + HC≡CH
This reaction was utilized in the acetylene (carbide) lamps, once used in mines, vintage cars and for bicycle lamps.
Calcium carbide, therefore, is an important raw material for the manufacture of many organic compounds, eg ethanal, and ethanoic acid, where ethyne is an intermediate.
Calcium carbide itself is made by heating coke and calcium oxide (lime) in an electric furnace. These reactions provide us with one of the few methods available for making organic compounds directly from carbon:CaC2 + 2H2O > C2H2 + Ca(OH)2
The only common one is Mg2C3, which has the anion [C=C=C]2-.
This yields propyne on hydrolysis:
Mg2C3 + 4H2O > 2Mg(OH)2 + CH3-C-C≡CH
2. Covalent carbides
These are of two kinds, (a) volatile and (b) non-volatile (macromolecules).
These are small molecules, formed especially by hydrogen and elements from Groups 4 to 6; eg CH4, (CN)2, CO, CS2, CCl4 etc.
2(b) Non-volatile macromolecules
Covalent carbides are formed with elements such as boron and silicon. These materials are extremely hard and inert. Silicon carbide (carborundum), SiC, has a diamond structure in which tetrahedral carbon atoms are bonded to four silicon atoms which in turn are tetrahedral and each bonded to four carbons.
In boron carbide (B12C3) the carbon atoms are arranged in groups of 3 linearly, the boron atoms forming a continuous structure. Boron atoms are arranged in icosahedral groups of 12. The two groups are packed together in a NaCl type array. B12C3 is even harder than SiC.
3. Interstitial carbides
Interstitial carbides are materials where carbon atoms occupy the spaces between metal atoms (interstitial spaces) in metal lattices.
They are formed by transition metals, particularly chromium, molybdenum and tungsten, and the carbon atoms occupy octahedral sites in the close-packed metal lattice.
Compounds of this interstitial type have certain properties (such as opacity, conductivity, metallic lustre, variable composition) not unlike those of metal alloys, ie the presence of carbon in the interstitial sites does not affect the electrical conductivity of the metal.
If the radius of the metal atoms is greater than 1.3 Angstroms, then carbon atoms in interstitial sites do not distort the metal lattice. In fact, carbon atoms stabilise the lattice and thus increase the hardness and melting point. Thus materials such as tungsten carbide (WC), are extremely hard inert materials with very high melting points. Tungsten carbide finds frequent use in machining tools.
Five oxides of carbon are known, ie CO, CO2, C3O2, (C5O2, C12O9)
Carbon Monoxide, CO
Carbon monoxide, CO, is a colourless, odourless gas that is formed through incomplete combustion due to lack of oxygen. It is flammable and will form carbon dioxide if burnt in air. The combustion is considerably exothermic, and so carbon monoxide is an important fuel.
Carbon monoxide is a good reducing agent and can reduce many metal oxides to the metal, eg:
Fe2O3 + 3CO → 2Fe + 3CO2
This reaction occurs in the blast furnace.
Carbon monoxide is made on industrial scales as a feedstock for other chemical processes and a mixture with hydrogen called synthesis gas is important for the production of methanol.
The nature of the bonding between the oxygen and carbon in this molecule makes it able to donate a lone pair of electrons to a central transition metal atom, thus forming a coordinate bond. The products are known as carbonyl compounds. For this reason, carbon monoxide is also highly toxic: it binds irreversibly to the central Fe ions in haemoglobin, the mammalian oxygen transport protein.
Carbon Dioxide, CO2
Carbon dioxide is the most stable oxide of carbon at ambient temperatures. It is a colourless gas produced during respiration in animals and plants, the carbon dioxide being the product of the oxidation of glucose in mitochondria. Carbon dioxide is also present in volcanic gas.
If carbon dioxide is cooled below -78°C it solidifies, in which form it is known as 'dry ice' or 'cardice'. Above this temperature it sublimes3.
Liquid carbon dioxide is shipped in high pressure cylinders, and is used in fire extinguishers.
An important use of carbon dioxide is its use as a supercritical fluid4 for the removal of caffeine from coffee (de-caffeination).
Carbon Suboxide, C3O2
This is prepared by dehydration of malonic acid using phosphorus pentoxide (P2O5).
Its structure can be represented as O=C=C=C=O in equilibrium with zwitterionic species containing C≡O and C≡C.
It is a foul-smelling, lachrymatory gas, being stable at -78°C. At 25°C the compound becomes unstable and polymerises to form highly coloured (yellow, red or brown) solid products. Under the influence of ultraviolet light it photolyses, decomposing to form ketene, C2O, a very reactive molecule.
Since carbon suboxide is the acid anhydride of malonic acid, it reacts slowly with water to produce that acid.
Carbon suboxide is widely used in the laboratory as a source of atomic carbon and may account for the CO and C emissions seen in the comae of comets such as Halley's Comet.
When carbon dioxide dissolves in water, it is in equilibrium with a small quantity of carbonic acid, H2CO3.
CO2 + H2O ↔ HCO3- + H+↔ H2CO3
Carbon forms a wide range of carbonate compounds containing the CO32- or HCO3- ions. Most carbonates, apart from those of the alkali metals, are insoluble in water and precipitate from aqueous solution as anyone living in a hard water area will know well. The most common carbonate is calcium carbonate, or limestone. Carbonates will react with acids to give carbon dioxide and water in the reverse of the carbonic acid formation reaction, ie
Na2CO3 + 2 HCl → CO2 + H2O + 2 NaCl
The simplest sulphur compound of carbon is carbon disulphide, CS2. It is a pale yellow toxic liquid and is structurally analogous to carbon dioxide. In contrast to carbon dioxide however, it is much more reactive and is flammable. It is sometimes used as a solvent for some specific chemical reaction where there isn't an alternative. It is produced via the reaction of methane and elemental sulphur over silica (SiO2) or alumina (Al2O3) at around 1,000°C.
CH4 + 4 S → CS2 + 2 H2S
CS2 will undergo reaction with a variety of nucleophiles. It will reaction with the ion HS- producing trithiocarbonate CS32-, with alcohols in the presence of a base to from compounds called xanthates and with amines to give compounds called thiocarbamates (see figure 2).
CS2 + S2-→ CS32-
CS2 + NaOH + ROH → (RO-CS2)Na + H2O
CS2 + R2NH → R2N-CS2H
The -CS2 motif of xanthates and thiocarbamates are analogous to carboxylates and can form metal complexes in the same way. Carbon disulphide can also form complexes with metals; the first example of this was discovered by the group of the Nobel prize winner Geoffrey Wilkinson. CS2 was added to a solution of the platinum complex [Pt(P(C6H5)3)3] to give [Pt(CS2 )(P(C6H5)3)2] in which the CS2 ligand is bent (figure 3).
Carbon does not form a free sulphur analogue of carbon monoxide, as CS is an unstable radical (containing unpaired electrons) which will react with just about anything. It does however exist as a ligand in a variety of metal complexes. The analogous selenium and tellurium compounds have also been observed. The carbon dioxide analogues S=C=Se and S=C=Te are also known.
Halides and Oxohalides
Carbon form tetrahalides with all of the halogens from fluorines to iodine. The lightest of these is tetrafluoromethane, CF4, and is an extremely stable gas due to very strong C-F bonds. It is commonly used in the etching of electronic circuit boards in industry and as a coolant.
Carbon will also form compounds analogous to alkanes but in which every hydrogen atom is replaced by fluorine. These are highly inert substances, an important example being the polymer polytetrafluoroethene (PTFE) otherwise known as Teflon. Tetrachloromethane (carbon tetrachloride) CCl4, is a toxic liquid that is most often used as a solvent but was also used to make freons, better known as chlorofluorocarbons or CFCs (such as CFCl3, CF2Cl2, CF3Cl). These non-toxic, non-flammable and highly chemically unreactive compounds are excellent refrigerants and were widely used in the 1960s and 1970s but have since been outlawed because of their role in stratospheric ozone depletion. These have been temporarily replaced by hydrochlorofluorocarbons (HCFCs) (which have hydrogen atoms in them to make them more reactive and break down harmlessly before being able to reach the ozone layer) and HFCs (which don't contain chlorine) but these have their own problems in being potent greenhouse gases.
Carbon tetrabromide, CBr4, and carbon tetraiodide, CI4 are much less thermally stable than the lighter tetrahalides. They are both solids, CBr4 being pale yellow and CI4 being bright red and crystalline. There is also a series of mixed halides containing bromine called halons. These are non-flammable gases and are used as fire suppressants.
Carbon also forms a series of highly reactive trigonal planar oxohalide compounds (X2C=O). The homohalides F2C=O, Cl2C=O and Br2C=O are known and are gases at room temperature, though I2C=O has not been made and is probably too unstable to exist. Carbonyl chloride (Cl2C=O), otherwise known as phosgene, is a highly toxic gas that was briefly used as an ineffective chemical weapon during the first world war. It is formed when carbon monoxide and chlorine react together in sunlight. It is prepared by performing the reaction at 150°C in the presence of active charcoal as catalyst. It is now a major industrial chemical for the production of other bulk chemicals.
Mixed halides are also known, for example, FClC=O and FBrC=O, and though the diiodide is not known, the compound FIC=O does occur.
The oxohalides are reactive compounds and will react with nucleophiles (compounds with non-bonding lone pairs of electrons) such as water, hydroxide and amines. For example, phosgene reacts in the presence of ammonia to give urea as the major product.
Cl2C=O + 2 NH3→ (H2N)2C=O + 2 HCl
In addition to the vast array of nitrogen-containing organic compounds, carbon also forms a wide range of inorganic compounds with C-N bonds. The most important of these compounds are the cyanides, cyanates and thiocyanates.
An important compound is cyanogen, (CN)2. It is a linear molecule with the structure N≡C-C≡N and is a highly toxic and flammable gas and produces one of the hottest known flames. This is made industrially by overall gaseous reaction of hydrogen cyanide with oxygen using NO2 as a catalyst.
2 HCN + NO2→ (CN)2 + NO + H2O
NO + ½O2→ NO2
If pure, cyanogen is very stable but trace impurities lead it to polymerise under heating to around 500°C. The C≡N fragment has been compared in its reactivity with the halogens. For example, it can dissociate into CN radicals like halogen X2 molecules can and will also undergo a reaction with some metal complexes called oxidative addition. Here the C-C (or X-X) bond is broken by donation of two electron by the metal and two M-C bonds form raising the oxidation state of the metal by two units, hence the reaction is oxidative. For example
((C6H5)3P)4Pd0 + N≡C-C≡N → ((C6H5)3P)2PdII(C≡N)2 + 2 P(C6H5)3
Also like the halogen, cyanogen can undergo disproportionation5 in basic solutions.
(CN)2 + 2 OH-→ CN- + OCN- + H2O
Another important carbon-nitrogen compound is hydrogen cyanide, HCN. It is also an extremely poisonous compound and is only just a gas at room temperature having a boiling point of 25.6°C. The relatively high boiling point for such a small molecule is due to the presence of hydrogen bonding. HCN is thought to have been an important compound in the formation of biological molecules in the early primordial atmosphere of Earth. Under mild pressure and in the presence of water and ammonia it will spontaneously form adenine, one of the bases that make up DNA.
It is made industrially by the reaction of methane, ammonia and oxygen over a platinum metal catalyst.
2 CH4 + 2 NH3 + O2→ 2 HCN + 6 H2O
It has many industrial uses, for example, it is used in the production of feedstocks for nylon production. It is also used for the production of metal cyanides. HCN is weakly acidic and will partly dissociate in water to give H+ and CN- ions and can therefore be neutralised by alkali metal hydroxides to give the metal hydroxide. The metal hydroxide can then be crystallised from solution.
HCN + NaOH → NaCN + H2O
The cyanides of highly electropositive metals, eg the alkali metals, are very water soluble, however those of silver(I), mercury(I) and lead(II) are very insoluble. The CN- ion will readily form complexes with many metals. It form complexes with metallic silver and gold and was used in the extraction of these metals from their lower grade ores.