Ethane

Properties

At standard temperature and pressure, ethane is a colorless, odorless gas. It has a boiling point of −88.5 °C (−127.4 °F; 184.6 K) and melting point of −182.8 °C (−296.9 °F; 90.4 K). Ethane is only very sparingly soluble in water.

The bond parameters of ethane have been measured to high precision by microwave spectroscopy and electron diffraction: r_C−C = 1.528(3) Å, r_C−H = 1.088(5) Å, and ∠CCH = 111.6(5)° by microwave and r_C−C = 1.524(3) Å, r_C−H = 1.089(5) Å, and ∠CCH = 111.9(5)° by electron diffraction (the numbers in parentheses represents the uncertainties in the final digits).

Rotational barrier

Ethane adopts an staggered conformation, in which the methyl groups are rotated relative to each other so that the hydrogen atoms are interleaved when viewed along the C−C axis. The rotational barrier, the energy barrier for the methyl groups to twist about the C−C bond so as to pass through the eclipsed conformation, is 12 kJ/mol (2.9 kcal/mol). As early as 1890–1891, chemists proposed that ethane molecules preferred the staggered conformation.

Rotating a molecular substructure about a twistable bond usually requires energy. Ethane gives a classic, simple example of such a rotational barrier, sometimes called the "ethane barrier". Among the earliest experimental evidence of this barrier (see diagram at left) was obtained by modelling the entropy of ethane. The three hydrogens at each end are free to pinwheel about the central carbon–carbon bond when provided with sufficient energy to overcome the barrier. The physical origin of the barrier is still not completely settled, although the overlap (exchange) repulsion between the hydrogen atoms on opposing ends of the molecule is perhaps the strongest candidate, with the stabilizing effect of hyperconjugation on the staggered conformation contributing to the phenomenon. Theoretical methods that use an appropriate starting point (orthogonal orbitals) find that hyperconjugation is the most important factor in the origin of the ethane rotation barrier.

Solid phases

Ethane exists in several solid phases. On cooling at normal pressure, it freezes into phase I, a body-centered cubic structure (space group Im3m). In this form, the positions of the hydrogen atoms are not fixed and the molecules rotate readily around the C−C axis, making it a plastic crystal. In the very narrow temperature range between 89.78 K (−183.37 °C) and 89.68 K (−183.47 °C) it exhibits a metastable state (phase II) that is orthorhombic. In phase II the molecules are orientationally ordered, meaning that the methyl groups do not rotate freely. When cooled further (below 89.68 K), it forms the low-temperature monoclinic phase, phase III (space group P2₁/n), which is also orientationally ordered. At room temperature and pressures above 2.46 gigapascals (24,600 bar), it forms phase IV, a tetragonal crystal (space group P4₂/mnm).

Atmospheric

Ethane occurs as a trace gas in the Earth's atmosphere, currently having a concentration at sea level of 0.5 ppb. Global ethane quantities have varied over time, likely due to flaring at natural gas fields. Global ethane emission rates declined from 1984 to 2010, though increased shale gas production at the Bakken formation in the U.S. has arrested the decline by half.

Although ethane is a greenhouse gas, it is much less abundant than methane, has a lifetime of only a few months compared to over a decade, and is less efficient at absorbing radiation relative to mass. In fact, ethane's global warming potential largely results from its conversion in the atmosphere to methane.

In Earth's atmosphere, hydroxyl radicals convert ethane to methanol vapor with a half-life of around three months.

Extraterrestrial

Ethane has been detected as a trace component in the atmospheres of all four giant planets, and in the atmosphere of Saturn's moon Titan. Atmospheric ethane results from the Sun's photochemical action on methane gas, which is also present in the atmospheres of these bodies. The conversion from ethane to methane results from ultraviolet photons of wavelengths shorter than 160 nm that photo-dissociate the methane molecule into a methyl radical and a hydrogen atom. When two methyl radicals recombine, the result is ethane:

CH₄ → CH₃• + •H

CH₃• + •CH₃ → C₂H₆

It is suspected that ethane produced on Titan by photochemical conversion from methane rains back onto the moon's surface, and over time has accumulated into hydrocarbon seas covering much of the moon's polar regions. In mid-2005, the Cassini orbiter discovered Ontario Lacus in Titan's south polar regions. Further analysis of infrared spectroscopic data presented in July 2008 provided additional evidence for the presence of liquid ethane in Ontario Lacus. Several significantly larger hydrocarbon lakes, Ligeia Mare and Kraken Mare being the two largest, were discovered near Titan's north pole using radar data gathered by Cassini. These lakes are believed to be filled primarily by a mixture of liquid ethane and methane. Surface temperatures of Titan are roughly 89–94 K, which is close to the melting points of pure ethane (90.4 K) and pure methane (90.7 K). Ethane and methane form a eutectic mixture at molar fractions of 36% ethane and 64% methane having a melting point of 72 K, which is well below the melting points of either component alone, which may allow these lakes to remain liquid even below 90 K.

In 1996, ethane was detected in Comet Hyakutake, and it has since been detected in some other comets. The existence of ethane in these distant solar system bodies may implicate ethane as a primordial component of the solar nebula from which the sun and planets are believed to have formed.

In 2006, Dale Cruikshank of NASA/Ames Research Center (a New Horizons co-investigator) and his colleagues announced the spectroscopic discovery of ethane on Pluto's surface.

Chemistry

The reactions of ethane involve chiefly free radical reactions. Ethane can react with the halogens, especially chlorine and bromine, by free-radical halogenation. This reaction proceeds through the propagation of the ethyl radical:

Cl₂ → 2 Cl•

C₂H₆ + Cl• → C₂H₅• + HCl

C₂H₅• + Cl₂ → C₂H₅Cl + Cl•

Cl• + C₂H₆ → C₂H₅• + HCl

The combustion of ethane releases 1559.7 kJ/mol, or 51.9 kJ/g, of heat, and produces carbon dioxide and water according to the chemical equation:

2 C₂H₆ + 7 O₂ → 4 CO₂ + 6 H₂O + 3120 kJ

Combustion may also occur without an excess of oxygen, yielding carbon monoxide, acetaldehyde, methane, methanol, and ethanol. At higher temperatures, especially in the range 600–900 °C (1,112–1,652 °F), ethylene is a significant product:

2 C₂H₆ + O₂ → 2 C₂H₄ + 2 H₂O

Such oxidative dehydrogenation reactions are relevant to the production of ethylene.

Production

After methane, ethane is the second-largest component of natural gas. Natural gas from different gas fields varies in ethane content from less than 1% to more than 6% by volume. Prior to the 1960s, ethane and larger molecules were typically not separated from the methane component of natural gas, but simply burnt along with the methane as a fuel. Today, ethane is an important petrochemical feedstock and is separated from the other components of natural gas in most well-developed gas fields. Ethane can also be separated from petroleum gas, a mixture of gaseous hydrocarbons produced as a byproduct of petroleum refining.

Ethane is most efficiently separated from methane by liquefying it at cryogenic temperatures. Various refrigeration strategies exist: the most economical process presently in wide use employs a turboexpander, and can recover more than 90% of the ethane in natural gas. In this process, chilled gas is expanded through a turbine, reducing the temperature to approximately −100 °C (−148 °F). At this low temperature, gaseous methane can be separated from the liquefied ethane and heavier hydrocarbons by distillation. Further distillation then separates ethane from the propane and heavier hydrocarbons.

Usage

The chief use of ethane is the production of ethylene (ethene) by steam cracking. Steam cracking of ethane is fairly selective for ethylene, while the steam cracking of heavier hydrocarbons yields a product mixture poorer in ethylene and richer in heavier alkenes (olefins), such as propene (propylene) and butadiene, and in aromatic hydrocarbons.

Ethane has been investigated as a feedstock for other commodity chemicals. Oxidative chlorination of ethane has long appeared to be a potentially more economical route to vinyl chloride than ethylene chlorination. Many patent exist on this theme, but poor selectivity for vinyl chloride and corrosive reaction conditions have discouraged the commercialization of most of them. Presently, Ineos operates a 1000 t/a (tonnes per annum) ethane-to-vinyl chloride pilot plant at Wilhelmshaven in Germany.

SABIC operates a 34,000 t/a plant at Yanbu to produce acetic acid by ethane oxidation. The economic viability of this process may rely on the low cost of ethane near Saudi oil fields, and it may not be competitive with methanol carbonylation elsewhere in the world.

Ethane can be used as a refrigerant in cryogenic refrigeration systems.

In the laboratory

On a much smaller scale, in scientific research, liquid ethane is used to vitrify water-rich samples for cryo-electron microscopy. A thin film of water quickly immersed in liquid ethane at −150 °C (−238 °F) or colder freezes too quickly for water to crystallize. Slower freezing methods can generate cubic ice crystals, which can disrupt soft structures by damaging the samples and reduce image quality by scattering the electron beam before it can reach the detector.

Health and safety

At room temperature, ethane is an extremely flammable gas. When mixed with air at 3.0–12.5% by volume, it forms an explosive mixture.

Ethane is not a carcinogen.