Alkanes occur both on Earth and in the solar system, however only the first hundred or so, and even then mostly only in traces. The light hydrocarbons, especially methane and ethane for example, have been detected both in the tail of the comet Hyakutake and in some meteorites such as carbonaceous chondrites. They also form an important portion of the atmospheres of the outer gas planets Jupiter, Saturn, Uranus and Neptune. On Titan, the satellite of Saturn, it is believed that there were once large oceans of these and longer chain alkanes: smaller seas of liquid ethane are thought still to exist there.

Traces of methane (about 0.0001% or 1 ppm) occur in the Earth's atmosphere, produced primarily by forms of Archaea. The content in the oceans is negligible due to the low solubility in water: however, at high pressures and low temperatures, methane can co-crystallize with water to form a solid methane hydrate. Although they cannot be commercially exploited at the present time, the calorific value of the known methane hydrate fields exceeds the energy content of all the natural gas and oil deposits put together—methane extracted from methane hydrate is considered therefore a candidate for future fuels.


Today, the most important commercial sources for alkanes are clearly natural gas and oil, which are the only organic compounds to occur as minerals in nature. Natural gas contains primarily methane and ethane, with some propane and butane: oil is a mixture of liquid alkanes and other hydrocarbons. Both were formed when dead marine animals and plants (zooplankton and phytoplankton) sank to the bottom of ancient seas and were covered with sediments in an anoxic environment and converted over many millions of years at high temperatures and high pressure to their current form. Natural gas resulted thereby for example from the following reaction:

C₆H₁₂O₆ → 3CH₄ + 3CO₂

Alkanes are both important raw materials of the chemical industry and the most important fuels of the world economy.

The starting materials for the processing are always natural gas and crude oil. The latter is separated in an oil refinery by fractional distillation and processed into many different products, for example gasoline. The different "fractions" of crude oil have different boiling points and can be isolated and separated quite easily: within the individual fractions the boiling points lie closely together.

The domain of usage of a certain alkane can be determined quite well according to the number of carbon atoms, although the following demarcation is idealized and not perfect. The first four alkanes are used mainly for heating and cooking purposes, and in some countries for electricity generation. Methane and ethane are the main components of natural gas; they are normally stored as gases under pressure. It is however easier to transport them as liquids: this requires both compression and cooling of the gas.

Propane and butane can be liquefied at fairly low pressures, and are well known as liquified petroleum gas (LPG). Propane, for example, is used in the propane gas burner, butane in disposable cigarette lighters (where the pressure is a mere 2 bar). The two alkanes are used as propellants in aerosol sprays.

From pentane to octane the alkanes are highly volatile liquids. They are used as fuels in internal combustion engines, as they vaporise easily on entry into the combustion chamber without forming droplets which would impair the unifomity of the combustion. Branched-chain alkanes are preferred, as they are much less prone to premature ignition which causes knocking than their straight-chain homologues. This propensity to premature ignition is measured by the octane rating of the fuel, where 2,2,4-trimethylpentane (isooctane) has an arbitrary value of 100 and heptane has a value of zero. Apart from their use as fuels, the middle alkanes are also good solvents for nonpolar substances.

Alkanes from nonane to, for instance, hexadecane (an alkane with sixteen carbon atoms) are liquids of higher viscosity, less and less suitable for use in gasoline. They form instead the major part of diesel and aviation fuel. Diesel fuels are characterised by their cetane number, cetane being an old name for hexadecane. However, the higher melting points of these alkanes can cause problems at low temperatures and in polar regions, where the fuel becomes too thick to flow correctly.

Alkanes from hexadecane upwards form the most important components of fuel oil and lubricating oil. In latter function they work at the same time as anti-corrosive agents, as their hydrophobic nature means that water cannot reach the metal surface. Many solid alkanes find use as paraffin wax, for example in candles. This should not be confused however with true wax, which consists primarily of esters.

Alkanes with a chain length of approximately 35 or more carbon atoms are found in bitumen, used for example in road surfacing. However, the higher alkanes have little value and are usually split into lower alkanes by cracking.

The molecular structure of the alkanes directly affects their physical and chemical characteristics. It is derived from the electron configuration of carbon, which has four valence electrons. The carbon atoms in alkanes are always sp³ hybridised, that is to say that the valence electrons are said to be in four equivalent orbitals derived from the combination of the 2s orbital and the three 2p orbitals. These orbitals, which have identical energies, are arranged spatially in the form of a tetrahedron, the angle of cos⁻¹(−⅓) ≈ 109.47° between them.

An alkane molecule has only C–H and C–C single bonds. The former result from the overlap of a sp³-orbital of carbon with the 1s-orbital of a hydrogen; the latter by the overlap of two sp³-orbitals on different carbon atoms. The bond lengths amount to 1.09×10⁻¹⁰ m for a C–H bond and 1.54×10⁻¹⁰ m for a C–C bond.


The spatial arrangement of the bonds is similar to that of the four sp³-orbitals—they are tetrahedrally arranged, with an angle of 109.47° between them. Structural formulae which represent the bonds as being at right angles to one another, while both common and useful, do not correspond with the reality.

The structural formula and the bond angles are not usually sufficient to completely describe the geometry of a molecule. There is a further degree of freedom for each carbon–carbon bond: the torsion angle between the atoms or groups bound to the atoms at each end of the bond. The spatial arrangement described by the torsion angles of the molecule is known as its conformation.


Ethane forms the simplest case for studying the conformation of alkanes, as there is only one C–C bond. If one looks down the axis of the C–C bond, then one will see the so-called Newman projection. The hydrogen atoms on both the front and rear carbon atoms have an angle of 120° between them, resulting from the projection of the base of the tetrahedron onto a flat plane. However, the torsion angle between a given hydrogen atom attached to the front carbon and a given hydrogen atom attached to the rear carbon can vary freely between 0° and 360°. This is a consequence of the free rotation about a carbon–carbon single bond. Despite this apparent freedom, only two limiting conformations are important: eclipsed conformation and staggered conformation.

The two conformations, also known as rotamers, differ in energy: The staggered conformation is 12.6 kJ/mol lower in energy (more stable) than the eclipsed conformation.

This difference in energy between the two conformations, known as the torsion energy, is low compared to the thermal energy of an ethane molecule at ambient temperature. There is constant rotation about the C-C bond, albeit with short "pauses" at each staggered conformation. The time taken for an ethane molecule to pass from one staggered conformation to the next, equivalent to the rotation of one CH₃-group by 120° relative to the other, is of the order of 10⁻¹¹ seconds.


The situation with respect to the two C-C bonds in propane is qualitatively similar to that of ethane: it is more complex, however, for butane and higher alkanes.

If one takes the central C-C bond of butane as the reference axis, each of the two central carbon atoms is bound to two hydrogen atoms and a methyl group. Four different conformations can be defined by the torsion angle between the two methyl groups and, as in the case of ethane, each has its characteristic energy.

The difference in energy between the fully eclipsed conformation and the antiperiplanar conformation is about 19 kJ/mol, and is therefore still relatively small at ambient temperature.

The case of higher alkanes is similar: the antiperiplanar conformation is always the most favoured around each carbon-carbon bond. For this reason, alkanes are usually shown in a zigzag arrangement in diagrams or in models. The actual structure will always differ somewhat from these idealised forms, as the differences in energy between the conformations are small compared to the thermal energy of the molecules: alkane molecules have no fixed structural form, whatever the models may suggest.

The molecular structure, particularly the surface area of the molecule, determines the boiling point of the alkane: the smaller the surface, the lower the boiling point, as the van der Waals forces between the molecules are weaker. A reduction of the surface area can be achieved by chain-branching or by a circular structure. This means in practice that alkanes with higher number of carbon atoms usually have higher boiling points than those with lower numbers of carbon atoms, and that branched-chain alkanes and cycloalkanes have lower boiling points than their straight-chain homologues. Under standard conditions, from CH₄ to C₄H₁₀ alkanes are gaseous; from C₅H₁₂ to C₁₇H₃₆ they are liquids; and after C₁₈H₃₈ they are solids. The boiling point increases between 20 and 30 °C per CH₂-group.

The melting points of the alkanes also rise with the increase in the number of carbon atoms (with only one exception, propane). However the melting points rise more slowly than the boiling points, in particular for the higher alkanes. In addition, the melting points of alkanes with an odd number of carbon atoms increase faster than the melting points of alkanes with an even number of carbon atoms (see figure): the cause of this phenomenon is the higher packing density of the alkanes with an even number of carbon atoms. • The melting points of branched-chain alkanes can be either higher or lower than those of the corresponding straight-chain alkanes, depending on the efficiency of molecular packing: this is particularly true for isoalkanes (2-methyl isomers), which often have melting points higher than those of their normal analogues.

Alkanes do not conduct electricity, nor are they substantially polarized by an electric field. For this reason they do not form hydrogen bonds and are insoluble in polar solvents such as water. Since the hydrogen bonds between individual water molecules are aligned away from an alkane molecule, the coexistence of an alkane and water leads to an increase in molecular order (a reduction in entropy). As there is no significant bonding between water molecules and alkane molecules, the second law of thermodynamics suggests that this reduction in entropy should be minimised by minimising the contact between alkane and water: alkanes are said to be hydrophobic in that they repel water.

Their solubility in nonpolar solvents is relatively good, a property which is called lipophilicity. Different alkanes are, for example, miscible in all proportions among themselves.

The density of the alkanes usually increases with increasing number of carbon atoms, but remains less than that of water. Hence, alkanes form the upper layer in an alkane-water mixture.