Ethane
 Molecular geometry of ethane based on rotational spectroscopy.
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| Names | |||
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| Preferred IUPAC name
Ethane[1] | |||
| Systematic IUPAC name
Dicarbane (never recommended[2]) | |||
| Identifiers | |||
3D model (
JSmol ) |
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| 1730716 | |||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider | |||
ECHA InfoCard
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100.000.741 | ||
| EC Number |
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| 212 | |||
| MeSH | Ethane | ||
PubChem CID
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RTECS number
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| UNII | |||
| UN number | 1035 | ||
CompTox Dashboard (EPA)
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| Properties | |||
| C2H6 | |||
| Molar mass | 30.070 g·mol−1 | ||
| Appearance | Colorless gas | ||
| Odor | Odorless | ||
| Density |
544.0 kg/m3 (liquid at -88,5 °C) | ||
| 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) | ||
Henry's law
constant (kH) |
19 nmol Pa−1 kg−1 | ||
| Acidity (pKa) | 50 | ||
| Basicity (pKb) | −36 | ||
Conjugate acid
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Ethanium | ||
| -37.37·10−6 cm3/mol | |||
| Thermochemistry | |||
Heat capacity (C)
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52.14± 0.39 J K−1 mol−1 at 298 Kelvin[5] | ||
Std enthalpy of (ΔfH⦵298)formation |
−84 kJ mol−1 | ||
Std enthalpy of (ΔcH⦵298)combustion |
−1561.0–−1560.4 kJ mol−1 | ||
| Hazards | |||
| GHS labelling: | |||
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| Danger | |||
| H220, H280 | |||
| P210, P410+P403 | |||
| NFPA 704 (fire diamond) | |||
| Flash point | −135 °C (−211 °F; 138 K) | ||
| 472 °C (882 °F; 745 K) | |||
Explosive limits
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2.9–13% | ||
| Safety data sheet (SDS) | inchem.org | ||
| Related compounds | |||
Related alkanes
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Related compounds
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| 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).
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Ethane (
Related compounds may be formed by replacing a hydrogen atom with another
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.[6] The process is now called Kolbe electrolysis:
During the period 1847–1849, in an effort to vindicate the
This error was corrected in 1864 by
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.[12] 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).[13] 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: 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).[14]
Rotating a molecular substructure about a twistable bond usually requires energy. The minimum energy to produce a 360° bond rotation is called the
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.[16] 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,[17] although the overlap (exchange) repulsion[18] 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.[19] 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.[20][21]
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.[22][23][24][25]
Atmospheric and extraterrestrial
Ethane occurs as a trace gas in the
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,[30] 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.[31] 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.[32]
Atmospheric ethane results from the Sun's
- 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.[30]
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
In 1996, ethane was detected in
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.[35]
Chemistry
The chemistry of ethane involves chiefly
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:
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 + H2O
Such oxidative dehydrogenation reactions are relevant to the production of ethylene.[36]
Production
After
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
Experimentally, ethane is under investigation as a feedstock for other commodity chemicals.
Similarly, the
Ethane can be used as a refrigerant in cryogenic refrigeration systems. On a much smaller scale, in scientific research, liquid ethane is used to
MAN Energy Solutions currently manufactures two-stroke dual fuel engines (B&W ME-GIE) which can run on both Marine diesel oil and ethane.
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
See also
- Biogas: carbon-neutral alternative to natural gas
- Biorefining
- Biodegradable plastic
- Drop-in bioplastic
References
- ISBN 978-0-85404-182-4.
The saturated unbranched acyclic hydrocarbons C2H6, C3H8, and C4H10 have the retained names ethane, propane, and butane, respectively.
- ^ 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'."
- ^ "Ethane – Compound Summary". PubChem Compound. US: National Center for Biotechnology Information. 16 September 2004. Retrieved 7 December 2011.
- ISBN 0-8493-0486-5.
- ^ "Ethane". webbook.nist.gov. National Institute of Standards and Technology. Retrieved 2024-05-16.
- S2CID 116224057.
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- ^ Roscoe, H.E.; Schorlemmer, C. (1881). Treatise on Chemistry. Vol. 3. Macmillan. pp. 144–145.
- ^ Watts, H. (1868). Dictionary of Chemistry. Vol. 4. p. 385.
- S2CID 55183235.
- ^ "Ethane as a solid". Retrieved 2019-12-10.
- ISSN 0021-9606.
- ISBN 9780840054449.
- .
- .
- .
- PMID 15065281.
- S2CID 205017635.
- PMID 12370897.
- .
- .
- .
- .
- ^ "Trace gases (archived)". Atmosphere.mpg.de. Archived from the original on 2008-12-22. Retrieved 2011-12-08.
- ^ S2CID 4373714.
- hdl:2027.42/142509.
- ^ "One oil field a key culprit in global ethane gas increase". University of Michigan. April 26, 2016.
- ^ .
- doi:10.1002/asl.804.
- ^ Brown, Bob; et al. (2008). "NASA Confirms Liquid Lake on Saturn Moon". NASA Jet Propulsion Laboratory. Archived from the original on 2011-06-05. Retrieved 2008-07-30.
- S2CID 4398324.
- S2CID 27362518.
- ^ Stern, A. (November 1, 2006). "Making Old Horizons New". The PI's Perspective. Johns Hopkins University Applied Physics Laboratory. Archived from the original on August 28, 2008. Retrieved 2007-02-12.
- S2CID 231946397.
- ISBN 9780123750891.
External links
