Methane (US: METH-ayn, UK: MEE-thayn) is a chemical compound that has the chemical formula CH4 (one carbon atom bonded to four hydrogen atoms). It is a group-14 hydride, the simplest alkane, and the main constituent of natural gas. The abundance of methane on Earth makes it an economically attractive fuel, although capturing and storing it is difficult because it is a gas at standard temperature and pressure. In the Earth's atmosphere methane is transparent to visible light but absorbs infrared radiation, acting as a greenhouse gas. Methane is an organic hydrocarbon, and among the simplest of organic compounds. Naturally occurring methane is found both below ground and under the seafloor and is formed by both geological and biological processes. The largest reservoir of methane is under the seafloor in the form of methane clathrates. When methane reaches the surface and the atmosphere, it is known as atmospheric methane. Methane has also been detected on other planets, including Mars, which has implications for astrobiology research.
Properties and bonding Methane is a tetrahedral molecule with four equivalent C–H bonds. Its electronic structure is described by two bonding molecular orbitals (MOs) resulting from the overlap of the valence orbitals on C and H. The lowest-energy MO is the result of the overlap of the 2s orbital on carbon with the in-phase combination of the 1s orbitals on the four hydrogen atoms. Above this energy level is a triply degenerate MO that overlaps with various linear combinations of the 1s orbitals on hydrogen atom. The resulting "three-over-one" bonding scheme is consistent with photoelectron spectroscopic measurements. Methane is an odorless, colourless and transparent gas at standard temperature and pressure. It does absorb visible light, especially at the red end of the spectrum, due to overtone bands, but the effect is only noticeable if the light path is very long. This is what gives Uranus and Neptune their blue or bluish-green colors, as light passes through their atmospheres containing methane and is then scattered back out. The familiar smell of natural gas as used in homes is achieved by the addition of an odorant, usually blends containing tert-butylthiol, as a safety measure. Methane has a boiling point of −161.5 °C at a pressure of one atmosphere. As a gas, it is flammable over a range of concentrations (5.4%–17%) in air at standard pressure. Solid methane exists in several modifications, of which nine are known. Cooling methane at normal pressure results in the formation of methane I. This substance crystallizes in the cubic system (space group Fm3m). The positions of the hydrogen atoms are not fixed in methane I, i.e. methane molecules may rotate freely. Therefore, it is a plastic crystal.
Chemical reactions The primary chemical reactions of methane are combustion, steam reforming to syngas, and halogenation. In general, methane reactions are difficult to control.
Acid–base reactions Like other hydrocarbons, methane is an extremely weak acid. Its pKa in DMSO is estimated to be 56. It cannot be deprotonated in solution, but the conjugate base is known in forms such as methyllithium. A variety of positive ions derived from methane have been observed, mostly as unstable species in low-pressure gas mixtures. These include methenium or methyl cation CH3+, methane cation CH4+, and methanium or protonated methane CH5+. Some of these have been detected in outer space. Methanium can also be produced as diluted solutions from methane with superacids. Cations with higher charge, such as CH62+ and CH73+, have been studied theoretically and conjectured to be stable. Despite the strength of its C–H bonds, there is intense interest in catalysts that facilitate C–H bond activation in methane (and other lower numbered alkanes).
Oxidation
Methane's heat of combustion is 55.5 MJ/kg. Combustion of methane is a multiple step reaction summarized as follows:
CH4 + 2 O2 → CO2 + 2 H2O ΔcH = −891 kJ/mol, at standard conditions (liquid water is formed at standard conditions, ΔcH = −802 kJ/mol if water vapor formation is considered) Peters four-step chemistry is a systematically reduced four-step chemistry that explains the burning of methane. Partial oxidation (combustion) of methane to methanol (CH3OH) in effected by some enzymes. The enzymes methane monooxygenase produces methanol from methane. One group of bacteria catalyze methane oxidation with nitrite as the oxidant in the absence of oxygen, giving rise to the so-called anaerobic oxidation of methane. Industrial scale oxidation has never achieved commercial success, even with an insufficient supply of oxygen. Some homogeneously catalyzed systems and heterogeneous systems have been demonstrated, but no such processes has proven economical. These generally operate by generating protected products which are shielded from overoxidation. Examples include the Catalytica system, copper zeolites, and iron zeolites stabilizing the alpha-oxygen active site.
Methane radical reactions Given appropriate conditions, methane reacts with halogen radicals as follows:
•X + CH4 → HX + •CH3 •CH3 + X2 → CH3X + •X where X is a halogen: fluorine (F), chlorine (Cl), bromine (Br), or iodine (I). This mechanism for this process is called free radical halogenation. It is initiated when UV light or some other radical initiator (like peroxides) produces a halogen atom. A two-step chain reaction ensues in which the halogen atom abstracts a hydrogen atom from a methane molecule, resulting in the formation of a hydrogen halide molecule and a methyl radical (•CH3). The methyl radical then reacts with a molecule of the halogen to form a molecule of the halomethane, with a new halogen atom as byproduct. Similar reactions can occur on the halogenated product, leading to replacement of additional hydrogen atoms by halogen atoms with dihalomethane, trihalomethane, and ultimately, tetrahalomethane structures, depending upon reaction conditions and the halogen-to-methane ratio. This reaction is commonly used with chlorine to produce dichloromethane and chloroform via chloromethane. Carbon tetrachloride can be made with excess chlorine.
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