Ethane

Last updated
Ethane
Skeletal formula of ethane with all implicit hydrogens shown Ethan Skelett.svg
Skeletal formula of ethane with all implicit hydrogens shown
Skeletal formula of ethane with all implicit carbons shown, and all explicit hydrogens added Ethan Lewis.svg
Skeletal formula of ethane with all implicit carbons shown, and all explicit hydrogens added
Ball and stick model of ethane Ethane-A-3D-balls.png
Ball and stick model of ethane
Spacefill model of ethane Ethane-3D-vdW.png
Spacefill model of ethane
Names
Preferred IUPAC name
Ethane [1]
Systematic IUPAC name
Dicarbane (never recommended [2] )
Identifiers
3D model (JSmol)
1730716
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.000.741 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 200-814-8
212
MeSH Ethane
PubChem CID
RTECS number
  • KH3800000
UNII
UN number 1035
  • InChI=1S/C2H6/c1-2/h1-2H3 Yes check.svgY
    Key: OTMSDBZUPAUEDD-UHFFFAOYSA-N Yes check.svgY
  • CC
Properties
C2H6
Molar mass 30.070 g·mol−1
AppearanceColorless gas
Odor Odorless
Density
  • 1.3562 kg/m3 (gas at 0 °C) [3]

544.0 kg/m3 (liquid at -88,5 °C)
206 kg/m3 (at critical point 305.322 K)

Contents

Melting point −182.8 °C; −296.9 °F; 90.4 K
Boiling point −88.5 °C; −127.4 °F; 184.6 K
Critical point (T, P)305.32 K (32.17 °C; 89.91 °F) 48.714 bars (4,871.4 kPa)
56.8 mg L−1 [4]
Vapor pressure 3.8453 MPa (at 21.1 °C)
19 nmol Pa−1 kg−1
Acidity (pKa)50
Basicity (pKb)−36
Conjugate acid Ethanium
-37.37·10−6 cm3/mol
Thermochemistry
52.49 J K−1 mol−1
−84 kJ mol−1
−1561.0–−1560.4 kJ mol−1
Hazards
GHS labelling:
GHS-pictogram-flamme.svg
Danger
H220, H280
P210, P410+P403
NFPA 704 (fire diamond)
1
4
0
SA
Flash point −135 °C (−211 °F; 138 K)
472 °C (882 °F; 745 K)
Explosive limits 2.9–13%
Safety data sheet (SDS) inchem.org
Related compounds
Related alkanes
Related compounds
Supplementary data page
Ethane (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Yes check.svgY  verify  (what is  Yes check.svgYX mark.svgN ?)

Ethane ( US: /ˈɛθeɪn/ ETH-ayn, UK: /ˈiːθeɪn/ EE-thayn) is an organic chemical compound with chemical formula C
2H
6. At standard temperature and pressure, ethane is a colorless, odorless gas. Like many hydrocarbons, ethane is isolated on an industrial scale from natural gas and as a petrochemical by-product of petroleum refining. Its chief use is as feedstock for ethylene production.

Related compounds may be formed by replacing a hydrogen atom with another functional group; the ethane moiety is called an ethyl group. For example, an ethyl group linked to a hydroxyl group yields ethanol, the alcohol in beverages.

History

Ethane was first synthesised in 1834 by Michael Faraday, applying electrolysis of a potassium acetate solution. He mistook the hydrocarbon product of this reaction for methane and did not investigate it further. [5]

During the period 1847–1849, in an effort to vindicate the radical theory of organic chemistry, Hermann Kolbe and Edward Frankland produced ethane by the reductions of propionitrile (ethyl cyanide) [6] and ethyl iodide [7] with potassium metal, and, as did Faraday, by the electrolysis of aqueous acetates. They mistook the product of these reactions for the methyl radical (CH3), of which ethane (C2H6) is a dimer.

This error was corrected in 1864 by Carl Schorlemmer, who showed that the product of all these reactions was in fact ethane. [8] Ethane was discovered dissolved in Pennsylvanian light crude oil by Edmund Ronalds in 1864. [9] [10]

Properties

At standard temperature and pressure, ethane is a colorless, odorless gas. It has a boiling point of −88.5 °C (−127.3 °F) and melting point of −182.8 °C (−297.0 °F). Solid ethane exists in several modifications. [11] On cooling under normal pressure, the first modification to appear is a plastic crystal, crystallizing in the cubic system. In this form, the positions of the hydrogen atoms are not fixed; the molecules may rotate freely around the long axis. Cooling this ethane below ca. 89.9 K (−183.2 °C; −297.8 °F) changes it to monoclinic metastable ethane II (space group P 21/n). [12] Ethane is only very sparingly soluble in water.

Molecular geometry of ethane based on rotational spectroscopy. Ethane-staggered-CRC-MW-dimensions-2D.png
Molecular geometry of ethane based on rotational spectroscopy.

The bond parameters of ethane have been measured to high precision by microwave spectroscopy and electron diffraction: rC−C = 1.528(3) Å, rC−H = 1.088(5) Å, and ∠CCH = 111.6(5)° by microwave and rC−C = 1.524(3) Å, rC−H = 1.089(5) Å, and ∠CCH = 111.9(5)° by electron diffraction (the numbers in parentheses represents the uncertainties in the final digits). [13]

Atmospheric and extraterrestrial

A photograph of Titan's northern latitudes. The dark features are hydrocarbon lakes containing ethane Titan North Pole Lakes PIA08630.jpg
A photograph of Titan's northern latitudes. The dark features are hydrocarbon lakes containing ethane

Ethane occurs as a trace gas in the Earth's atmosphere, currently having a concentration at sea level of 0.5 ppb, [14] though its preindustrial concentration is likely to have been only around 0.25 part per billion since a significant proportion of the ethane in today's atmosphere may have originated as fossil fuels. Global ethane quantities have varied over time, likely due to flaring at natural gas fields. [15] Global ethane emission rates declined from 1984 to 2010, [15] though increased shale gas production at the Bakken Formation in the U.S. has arrested the decline by half. [16] [17]

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, [18] and is also 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. [19] It has been detected as a trace component in the atmospheres of all four giant planets, and in the atmosphere of Saturn's moon Titan. [20]

Atmospheric ethane results from the Sun's photochemical action on methane gas, also present in these atmospheres: ultraviolet photons of shorter wavelengths than 160 nm can 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

In Earth's atmosphere, hydroxyl radicals convert ethane to methanol vapor with a half-life of around three months. [18]

It is suspected that ethane produced in this fashion on Titan rains back onto the moon's surface, and over time has accumulated into hydrocarbon seas covering much of the moon's polar regions. In December 2007 the Cassini probe found at least one lake at Titan's south pole, now called Ontario Lacus because of the lake's similar area to Lake Ontario on Earth (approximately 20,000 km2). Further analysis of infrared spectroscopic data presented in July 2008 [21] 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.

In 1996, ethane was detected in Comet Hyakutake, [22] 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. [23]

Chemistry

Ethane can be viewed as two methyl groups joined, that is, a dimer of methyl groups. In the laboratory, ethane may be conveniently synthesised by Kolbe electrolysis. In this technique, an aqueous solution of an acetate salt is electrolysed. At the anode, acetate is oxidized to produce carbon dioxide and methyl radicals, and the highly reactive methyl radicals combine to produce ethane:

CH3COO− → CH3• + CO2 + e−
CH3• + •CH3 → C2H6

Synthesis by oxidation of acetic anhydride by peroxides, is conceptually similar.

The chemistry of ethane involves 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:

C2H5• + Cl2 → C2H5Cl + Cl•
Cl• + C2H6 → C2H5• + HCl

Because halogenated ethanes can undergo further free radical halogenation, this process results in a mixture of several halogenated products. In the chemical industry, more selective chemical reactions are used for the production of any particular two-carbon haloalkane.

Combustion

The complete 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, forming a mix of amorphous carbon and carbon monoxide.

2 C2H6 + 3 O2 → 4 C + 6 H2O + energy
2 C2H6 + 5 O2 → 4 CO + 6 H2O + energy
2 C2H6 + 4 O2 → 2 C + 2 CO + 6 H2O + energy etc.

Combustion occurs by a complex series of free-radical reactions. Computer simulations of the chemical kinetics of ethane combustion have included hundreds of reactions. An important series of reaction in ethane combustion is the combination of an ethyl radical with oxygen, and the subsequent breakup of the resulting peroxide into ethoxy and hydroxyl radicals.

C2H5• + O2 → C2H5OO•
C2H5OO• + HR → C2H5OOH + •R
C2H5OOH → C2H5O• + •OH

The principal carbon-containing products of incomplete ethane combustion are single-carbon compounds such as carbon monoxide and formaldehyde. One important route by which the carbon–carbon bond in ethane is broken, to yield these single-carbon products, is the decomposition of the ethoxy radical into a methyl radical and formaldehyde, which can in turn undergo further oxidation.

C2H5O• → CH3• + CH2O

Some minor products in the incomplete combustion of ethane include 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. It arises through reactions such as this:

C2H5• + O2 → C2H4 + •OOH

Similar reactions (with agents other than oxygen as the hydrogen abstractor) are involved in the production of ethylene from ethane in steam cracking.

Barrier

Ethane (shown in Newman projection) barrier to rotation about the carbon-carbon bond. The curve is potential energy as a function of rotational angle. Energy barrier is 12 kJ/mol or about 2.9 kcal/mol. Ethane conformations and relative energies.svg
Ethane (shown in Newman projection) barrier to rotation about the carbon-carbon bond. The curve is potential energy as a function of rotational angle. Energy barrier is 12 kJ/mol or about 2.9 kcal/mol.

Rotating a molecular substructure about a twistable bond usually requires energy. The minimum energy to produce a 360° bond rotation is called the rotational barrier.

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. [25] 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, [26] although the overlap (exchange) repulsion [27] 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. [28] 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. [29] [30]

As far back as 1890–1891, chemists suggested that ethane molecules preferred the staggered conformation with the two ends of the molecule askew from each other. [31] [32] [33] [34]

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. When diluted with steam and briefly heated to very high temperatures (900 °C or more), heavy hydrocarbons break down into lighter hydrocarbons, and saturated hydrocarbons become unsaturated. Ethane is favored for ethylene production because the 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.

Experimentally, ethane is under investigation 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 processes for producing this reaction have been patented, but poor selectivity for vinyl chloride and corrosive reaction conditions (specifically, a reaction mixture containing hydrochloric acid at temperatures greater than 500 °C) 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.

Similarly, the Saudi Arabian firm SABIC has announced construction of a 30,000 t/a plant to produce acetic acid by ethane oxidation at Yanbu. 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. 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 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.

Some additional precautions are necessary where ethane is stored as a cryogenic liquid. Direct contact with liquid ethane can result in severe frostbite. Until they warm to room temperature, the vapors from liquid ethane are heavier than air and can flow along the floor or ground, gathering in low places; if the vapors encounter an ignition source, the chemical reaction can flash back to the source of ethane from which they evaporated.

Ethane can displace oxygen and become an asphyxiation hazard. Ethane poses no known acute or chronic toxicological risk. It is not a carcinogen. [35]

See also

Related Research Articles

<span class="mw-page-title-main">Alkane</span> Type of saturated hydrocarbon compound

In organic chemistry, an alkane, or paraffin, is an acyclic saturated hydrocarbon. In other words, an alkane consists of hydrogen and carbon atoms arranged in a tree structure in which all the carbon–carbon bonds are single. Alkanes have the general chemical formula CnH2n+2. The alkanes range in complexity from the simplest case of methane, where n = 1, to arbitrarily large and complex molecules, like pentacontane or 6-ethyl-2-methyl-5-(1-methylethyl) octane, an isomer of tetradecane.

<span class="mw-page-title-main">Alkene</span> Hydrocarbon compound containing one or more C=C bonds

In organic chemistry, an alkene is a hydrocarbon containing a carbon–carbon double bond.

<span class="mw-page-title-main">Ethylene</span> Hydrocarbon compound (H2C=CH2)

Ethylene is a hydrocarbon which has the formula C2H4 or H2C=CH2. It is a colourless, flammable gas with a faint "sweet and musky" odour when pure. It is the simplest alkene.

<span class="mw-page-title-main">Ethanol</span> Organic compound (CH3–CH2–OH)

Ethanol is an organic compound. It is an alcohol with the chemical formula C2H6O. Its formula can be also written as CH3−CH2−OH or C2H5OH. Ethanol is a volatile, flammable, colorless liquid with a characteristic wine-like odor and pungent taste. It is a psychoactive recreational drug, the active ingredient in alcoholic drinks.

<span class="mw-page-title-main">Hydrocarbon</span> Organic compound consisting entirely of hydrogen and carbon

In organic chemistry, a hydrocarbon is an organic compound consisting entirely of hydrogen and carbon. Hydrocarbons are examples of group 14 hydrides. Hydrocarbons are generally colourless and hydrophobic, and their odors are usually weak or exemplified by the odors of gasoline and lighter fluid. They occur in a diverse range of molecular structures and phases: they can be gases, liquids, low melting solids or polymers.

<span class="mw-page-title-main">Methanol</span> Organic compound (CH3OH); simplest alcohol

Methanol (also called methyl alcohol and wood spirit, amongst other names) is an organic chemical and the simplest aliphatic alcohol, with the formula CH3OH (a methyl group linked to a hydroxyl group, often abbreviated as MeOH). It is a light, volatile, colourless, flammable liquid with a distinctive alcoholic odour similar to that of ethanol (potable alcohol). A polar solvent, methanol acquired the name wood alcohol because it was once produced chiefly by the destructive distillation of wood. Today, methanol is mainly produced industrially by hydrogenation of carbon monoxide.

In organic chemistry, a methyl group is an alkyl derived from methane, containing one carbon atom bonded to three hydrogen atoms, having chemical formula CH3. In formulas, the group is often abbreviated as Me. This hydrocarbon group occurs in many organic compounds. It is a very stable group in most molecules. While the methyl group is usually part of a larger molecule, bounded to the rest of the molecule by a single covalent bond, it can be found on its own in any of three forms: methanide anion, methylium cation or methyl radical. The anion has eight valence electrons, the radical seven and the cation six. All three forms are highly reactive and rarely observed.

<span class="mw-page-title-main">Methyl radical</span> Chemical compound

Methyl is an organic compound with the chemical formula CH•
3. It is a metastable colourless gas, which is mainly produced in situ as a precursor to other hydrocarbons in the petroleum cracking industry. It can act as either a strong oxidant or a strong reductant, and is quite corrosive to metals.

<span class="mw-page-title-main">Cracking (chemistry)</span> Process whereby complex organic molecules are broken down into simpler molecules

In petrochemistry, petroleum geology and organic chemistry, cracking is the process whereby complex organic molecules such as kerogens or long-chain hydrocarbons are broken down into simpler molecules such as light hydrocarbons, by the breaking of carbon-carbon bonds in the precursors. The rate of cracking and the end products are strongly dependent on the temperature and presence of catalysts. Cracking is the breakdown of a large alkane into smaller, more useful alkenes. Simply put, hydrocarbon cracking is the process of breaking a long chain of hydrocarbons into short ones. This process requires high temperatures.

The Fischer–Tropsch process is a collection of chemical reactions that converts a mixture of carbon monoxide and hydrogen, known as syngas, into liquid hydrocarbons. These reactions occur in the presence of metal catalysts, typically at temperatures of 150–300 °C (302–572 °F) and pressures of one to several tens of atmospheres. The process was first developed by Franz Fischer and Hans Tropsch at the Kaiser Wilhelm Institute for Coal Research in Mülheim an der Ruhr, Germany, in 1925.

<span class="mw-page-title-main">Conformational isomerism</span> Different molecular structures formed only by rotation about single bonds

In chemistry, conformational isomerism is a form of stereoisomerism in which the isomers can be interconverted just by rotations about formally single bonds. While any two arrangements of atoms in a molecule that differ by rotation about single bonds can be referred to as different conformations, conformations that correspond to local minima on the potential energy surface are specifically called conformational isomers or conformers. Conformations that correspond to local maxima on the energy surface are the transition states between the local-minimum conformational isomers. Rotations about single bonds involve overcoming a rotational energy barrier to interconvert one conformer to another. If the energy barrier is low, there is free rotation and a sample of the compound exists as a rapidly equilibrating mixture of multiple conformers; if the energy barrier is high enough then there is restricted rotation, a molecule may exist for a relatively long time period as a stable rotational isomer or rotamer. When the time scale for interconversion is long enough for isolation of individual rotamers, the isomers are termed atropisomers. The ring-flip of substituted cyclohexanes constitutes another common form of conformational isomerism.

<span class="mw-page-title-main">Eclipsed conformation</span> Molecular form in which substituents on two adjacent atoms are closest together

In chemistry an eclipsed conformation is a conformation in which two substituents X and Y on adjacent atoms A, B are in closest proximity, implying that the torsion angle X–A–B–Y is 0°. Such a conformation can exist in any open chain, single chemical bond connecting two sp3-hybridised atoms, and it is normally a conformational energy maximum. This maximum is often explained by steric hindrance, but its origins sometimes actually lie in hyperconjugation.

<span class="mw-page-title-main">Hyperconjugation</span> Concept in organic chemistry

In organic chemistry, hyperconjugation refers to the delocalization of electrons with the participation of bonds of primarily σ-character. Usually, hyperconjugation involves the interaction of the electrons in a sigma (σ) orbital with an adjacent unpopulated non-bonding p or antibonding σ* or π* orbitals to give a pair of extended molecular orbitals. However, sometimes, low-lying antibonding σ* orbitals may also interact with filled orbitals of lone pair character (n) in what is termed negative hyperconjugation. Increased electron delocalization associated with hyperconjugation increases the stability of the system. In particular, the new orbital with bonding character is stabilized, resulting in an overall stabilization of the molecule. Only electrons in bonds that are in the β position can have this sort of direct stabilizing effect — donating from a sigma bond on an atom to an orbital in another atom directly attached to it. However, extended versions of hyperconjugation can be important as well. The Baker–Nathan effect, sometimes used synonymously for hyperconjugation, is a specific application of it to certain chemical reactions or types of structures.

<span class="mw-page-title-main">Methane</span> Saturated hydrocarbon with formula CH4

Methane ( MEH-thayn, MEE-thayn) is a chemical compound with 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 relative abundance of methane on Earth makes it an economically attractive fuel, although capturing and storing it poses technical challenges due to its gaseous state under normal conditions for temperature and pressure.

The oxidative coupling of methane (OCM) is a chemical reaction discovered in the 1980s for the direct conversion of natural gas, primarily consisting of methane, into value-added chemicals. The OCM process has not been commercially practiced.

<span class="mw-page-title-main">Ethenium</span>

In chemistry, ethenium, protonated ethylene or ethyl cation is a positive ion with the formula C
2H+
5. It can be viewed as a molecule of ethylene with one added proton, or a molecule of ethane minus one hydride ion. It is a carbocation; more specifically, a nonclassical carbocation.

Radical theory is an obsolete scientific theory in chemistry describing the structure of organic compounds. The theory was pioneered by Justus von Liebig, Friedrich Wöhler and Auguste Laurent around 1830 and is not related to the modern understanding of free radicals. In this theory, organic compounds were thought to exist as combinations of radicals that could be exchanged in chemical reactions just as chemical elements could be interchanged in inorganic compounds.

Methane functionalization is the process of converting methane in its gaseous state to another molecule with a functional group, typically methanol or acetic acid, through the use of transition metal catalysts.

1,2-Difluoroethane is a saturated hydrofluorocarbon containing an atom of fluorine attached to each of two carbons atoms. The formula can be written CH2FCH2F. It is an isomer of 1,1-difluoroethane. It has a HFC name of HFC-152 with no letter suffix. When cooled to cryogenic temperatures it can have different conformers, gauche and trans. In the liquid form these are about equally abundant and easily interconvert. As a gas it is mostly the gauche form.

<span class="mw-page-title-main">Steam cracking</span> Petrochemical process to break down saturated hydrocarbons in smaller molecules

Steam cracking is a petrochemical process in which saturated hydrocarbons are broken down into smaller, often unsaturated, hydrocarbons. It is the principal industrial method for producing the lighter alkenes, including ethene and propene. Steam cracker units are facilities in which a feedstock such as naphtha, liquefied petroleum gas (LPG), ethane, propane or butane is thermally cracked through the use of steam in steam cracking furnaces to produce lighter hydrocarbons. The propane dehydrogenation process may be accomplished through different commercial technologies. The main differences between each of them concerns the catalyst employed, design of the reactor and strategies to achieve higher conversion rates.

References

  1. ↑ International Union of Pure and Applied Chemistry (2014). Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013. The Royal Society of Chemistry. p. 133. doi:10.1039/9781849733069. ISBN   978-0-85404-182-4. The saturated unbranched acyclic hydrocarbons C2H6, C3H8, and C4H10 have the retained names ethane, propane, and butane, respectively.
  2. ↑ IUPAC 2014, p. 4. "Similarly, the retained names 'ethane', 'propane', and 'butane' were never replaced by systematic names 'dicarbane', 'tricarbane', and 'tetracarbane' as recommended for analogues of silane, 'disilane'; phosphane, 'triphosphane'; and sulfane, 'tetrasulfane'."
  3. ↑ "Ethane – Compound Summary". PubChem Compound. USA: National Center for Biotechnology Information. 16 September 2004. Retrieved 7 December 2011.
  4. ↑ Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. p. 8.88. ISBN   0-8493-0486-5.
  5. ↑ Faraday, Michael (1834). "Experimental researches in electricity: Seventh series". Philosophical Transactions. 124: 77–122. Bibcode:1834RSPT..124...77F. doi:10.1098/rstl.1834.0008. S2CID   116224057.
  6. ↑ Kolbe, Hermann; Frankland, Edward (1849). "On the products of the action of potassium on cyanide of ethyl". Journal of the Chemical Society. 1: 60–74. doi:10.1039/QJ8490100060.
  7. ↑ Frankland, Edward (1850). "On the isolation of the organic radicals". Journal of the Chemical Society. 2 (3): 263–296. doi:10.1039/QJ8500200263.
  8. ↑ Schorlemmer, Carl (1864). "Ueber die Identität des Aethylwasserstoffs und des Methyls". Annalen der Chemie und Pharmacie. 132 (2): 234–238. doi:10.1002/jlac.18641320217.
  9. ↑ Roscoe, H.E.; Schorlemmer, C. (1881). Treatise on Chemistry. Vol. 3. Macmillan. pp. 144–145.
  10. ↑ Watts, H. (1868). Dictionary of Chemistry. Vol. 4. p. 385.
  11. ↑ Van Nes, G.J.H.; Vos, A. (1978). "Single-crystal structures and electron density distributions of ethane, ethylene and acetylene. I. Single-crystal X-ray structure determinations of two modifications of ethane" (PDF). Acta Crystallographica Section B. 34 (6): 1947. doi:10.1107/S0567740878007037. S2CID   55183235.
  12. ↑ "Ethane as a solid" . Retrieved 2019-12-10.
  13. ↑ Harmony, Marlin D. (1990-11-15). "The equilibrium carbon–carbon single‐bond length in ethane". The Journal of Chemical Physics. 93 (10): 7522–7523. Bibcode:1990JChPh..93.7522H. doi:10.1063/1.459380. ISSN   0021-9606.
  14. ↑ Trace gases (archived). Atmosphere.mpg.de. Retrieved on 2011-12-08.
  15. 1 2 Simpson, Isobel J.; Sulbaek Andersen, Mads P.; Meinardi, Simone; Bruhwiler, Lori; Blake, Nicola J.; Helmig, Detlev; Rowland, F. Sherwood; Blake, Donald R. (2012). "Long-term decline of global atmospheric ethane concentrations and implications for methane". Nature. 488 (7412): 490–494. Bibcode:2012Natur.488..490S. doi:10.1038/nature11342. PMID   22914166. S2CID   4373714.
  16. ↑ Kort, E. A.; Smith, M. L.; Murray, L. T.; Gvakharia, A.; Brandt, A. R.; Peischl, J.; Ryerson, T. B.; Sweeney, C.; Travis, K. (2016). "Fugitive emissions from the Bakken shale illustrate role of shale production in global ethane shift". Geophysical Research Letters. 43 (9): 4617–4623. Bibcode:2016GeoRL..43.4617K. doi: 10.1002/2016GL068703 .
  17. ↑ "One oil field a key culprit in global ethane gas increase". University of Michigan. April 26, 2016.
  18. 1 2 Aydin, Kamil Murat; Williams, M.B. and Saltzman, E.S.; ‘Feasibility of reconstructing paleoatmospheric records of selected alkanes, methyl halides, and sulfur gases from Greenland ice cores’; Journal of Geophysical Research; volume 112, D07312
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