Ethylene - Thermophysical Properties

Ethylene, C2H 4 , is a highly flammable, colorless and noncorrosive gas with a sweet odor. It is easily ignited and a flame can easily flash back to the source of the leak. Under prolonged exposure to fire or heat the containers may rupture violently and rocket. Can cause explosion. Vapors arising from the boiling liquid are lighter than air. Ethylene is not toxic, but is a simple asphyxiant.

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Ethylene is used as an anesthetic, a refrigerant, and to make other chemicals as polymers and plastics.

Ethylene is produced in petrochemical processes, as steam cracking where hydrocarbons and steam are heated to 750&#;950 °C. This process converts large hydrocarbons into smaller ones and introduces unsaturated products.

The phase diagram of ethylene is shown below the table.

Chemical, physical and thermal properties of ethylene:
Values are given for gas phase at 25 oC /77 oF / 298 K and 1 bara, if not other phase, temperature or pressure given.

For full table with Imperial Units - rotate the screen!

Follow the links below to get values for the listed properties of ethylene at varying pressure and temperature :

See also more about atmospheric pressure , and STP - Standard Temperature and Pressure & NTP - Normal Temperature and Pressure ,
as well as Thermophysical properties of: Acetone , Acetylene , Air , Ammonia , Argon , Benzene , Butane , Carbon dioxide , Carbon monoxide , Ethane , Ethanol , Helium , Hydrogen , Hydrogen sulfide , Methane , Methanol , Nitrogen , Oxygen , Pentane , Propane , Toluene , Water and Heavy water, D2O .

Ethylene is a gas at standard conditions. However, at low temperature and/or high pressures the gas becomes a liquid or a solid.

The ethylene phase diagram shows the phase behavior with changes in temperature and pressure. The curve between the critical point and the triple point shows the ethylene boiling point with changes in pressure. It also shows the saturation pressure with changes in temperature.

At the critical point there is no change of state when pressure is increased or if heat is added.

The triple point of a substance is the temperature and pressure at which the three phases (gas, liquid, and solid) of that substance coexist in thermodynamic equilibrium.

Hydrocarbon compound (H&#;C=CH&#;)

"Ethene" redirects here. Not to be confused with ethane or ethyne

Chemical compound

Ethylene (IUPAC name: ethene) 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.[7] It is the simplest alkene (a hydrocarbon with carbon&#;carbon double bonds).

Ethylene is widely used in the chemical industry, and its worldwide production (over 150 million tonnes in [8]) exceeds that of any other organic compound.[9][10] Much of this production goes toward creating polythene, which is a widely used plastic containing polymer chains of ethylene units in various chain lengths. Production emits greenhouse gases, including methane from feedstock production and carbon dioxide from any non-sustainable energy used.

Ethylene is also an important natural plant hormone and is used in agriculture to induce ripening of fruits.[11] The hydrate of ethylene is ethanol.

Structure and properties

This hydrocarbon has four hydrogen atoms bound to a pair of carbon atoms that are connected by a double bond. All six atoms that comprise ethylene are coplanar. The H-C-H angle is 117.4°, close to the 120° for ideal sp² hybridized carbon. The molecule is also relatively weak: rotation about the C-C bond is a very low energy process that requires breaking the π-bond by supplying heat at 50 °C.[citation needed]

The π-bond in the ethylene molecule is responsible for its useful reactivity. The double bond is a region of high electron density, thus it is susceptible to attack by electrophiles. Many reactions of ethylene are catalyzed by transition metals, which bind transiently to the ethylene using both the π and π* orbitals.[citation needed]

Being a simple molecule, ethylene is spectroscopically simple. Its UV-vis spectrum is still used as a test of theoretical methods.[12]

Uses

Major industrial reactions of ethylene include in order of scale: 1) polymerization, 2) oxidation, 3) halogenation and hydrohalogenation, 4) alkylation, 5) hydration, 6) oligomerization, and 7) hydroformylation. In the United States and Europe, approximately 90% of ethylene is used to produce ethylene oxide, ethylene dichloride, ethylbenzene and polyethylene.[13] Most of the reactions with ethylene are electrophilic addition.[citation needed]

Polymerization

Polyethylene production uses more than half of the world's ethylene supply. Polyethylene, also called polyethene and polythene, is the world's most widely used plastic. It is primarily used to make films in packaging, carrier bags and trash liners. Linear alpha-olefins, produced by oligomerization (formation of short-chain molecules) are used as precursors, detergents, plasticisers, synthetic lubricants, additives, and also as co-monomers in the production of polyethylenes.[13]

Oxidation

Ethylene is oxidized to produce ethylene oxide, a key raw material in the production of surfactants and detergents by ethoxylation. Ethylene oxide is also hydrolyzed to produce ethylene glycol, widely used as an automotive antifreeze as well as higher molecular weight glycols, glycol ethers, and polyethylene terephthalate.[14][15]

Ethylene oxidation in the presence of a palladium catalyst can form acetaldehyde. This conversion remains a major industrial process (10M kg/y).[16] The process proceeds via the initial complexation of ethylene to a Pd(II) center.[citation needed]

Halogenation and hydrohalogenation

Major intermediates from the halogenation and hydrohalogenation of ethylene include ethylene dichloride, ethyl chloride, and ethylene dibromide. The addition of chlorine entails "oxychlorination", i.e. chlorine itself is not used. Some products derived from this group are polyvinyl chloride, trichloroethylene, perchloroethylene, methyl chloroform, polyvinylidene chloride and copolymers, and ethyl bromide.[17]

Alkylation

Major chemical intermediates from the alkylation with ethylene is ethylbenzene, precursor to styrene. Styrene is used principally in polystyrene for packaging and insulation, as well as in styrene-butadiene rubber for tires and footwear. On a smaller scale, ethyltoluene, ethylanilines, 1,4-hexadiene, and aluminium alkyls. Products of these intermediates include polystyrene, unsaturated polyesters and ethylene-propylene terpolymers.[17]

Oxo reaction

The hydroformylation (oxo reaction) of ethylene results in propionaldehyde, a precursor to propionic acid and n-propyl alcohol.[17]

Hydration

Ethylene has long represented the major nonfermentative precursor to ethanol. The original method entailed its conversion to diethyl sulfate, followed by hydrolysis. The main method practiced since the mid-s is the direct hydration of ethylene catalyzed by solid acid catalysts:[18]

C2H4 + H2O &#; CH3CH2OH

Dimerization to butenes

Ethylene is dimerized by hydrovinylation to give n-butenes using processes licensed by Lummus or IFP. The Lummus process produces mixed n-butenes (primarily 2-butenes) while the IFP process produces 1-butene. 1-Butene is used as a comonomer in the production of certain kinds of polyethylene.[19]

Fruit and flowering

Ethylene is a hormone that affects the ripening and flowering of many plants. It is widely used to control freshness in horticulture and fruits.[20] The scrubbing of naturally occurring ethylene delays ripening.[21] Adsorption of ethylene by nets coated in titanium dioxide gel has also been shown to be effective.[22]

Niche uses

An example of a niche use is as an anesthetic agent (in an 85% ethylene/15% oxygen ratio).[23] Another use is as a welding gas.[13][24] It is also used as a refrigerant gas for low temperature applications under the name R-.[25]

Production

Global ethylene production was 107 million tonnes in ,[9] 109 million tonnes in ,[26] 138 million tonnes in , and 141 million tonnes in .[27] By , ethylene was produced by at least 117 companies in 32 countries. To meet the ever-increasing demand for ethylene, sharp increases in production facilities are added globally, particularly in the Mideast and in China.[28] Production emits greenhouse gas, namely significant amounts of carbon dioxide.[29]

Industrial process

Ethylene is produced by several methods in the petrochemical industry. A primary method is steam cracking (SC) where hydrocarbons and steam are heated to 750&#;950 °C. This process converts large hydrocarbons into smaller ones and introduces unsaturation. When ethane is the feedstock, ethylene is the product. Ethylene is separated from the resulting mixture by repeated compression and distillation.[17] In Europe and Asia, ethylene is obtained mainly from cracking naphtha, gasoil and condensates with the coproduction of propylene, C4 olefins and aromatics (pyrolysis gasoline).[30] Other technologies employed for the production of ethylene include Fischer-Tropsch synthesis and methanol-to-olefins (MTO).[31]

Laboratory synthesis

Although of great value industrially, ethylene is rarely synthesized in the laboratory and is ordinarily purchased.[32] It can be produced via dehydration of ethanol with sulfuric acid or in the gas phase with aluminium oxide or activated alumina.[33]

Biosynthesis

Ethylene is produced from methionine in nature. The immediate precursor is 1-aminocyclopropane-1-carboxylic acid.[34]

Ligand

Ethylene is a fundamental ligand in transition metal alkene complexes. One of the first organometallic compounds, Zeise's salt is a complex of ethylene. Useful reagents containing ethylene include Pt(PPh3)2(C2H4) and Rh2Cl2(C2H4)4. The Rh-catalysed hydroformylation of ethylene is conducted on an industrial scale to provide propionaldehyde.[36]

History

Some geologists and scholars believe that the famous Greek Oracle at Delphi (the Pythia) went into her trance-like state as an effect of ethylene rising from ground faults.[37]

Ethylene appears to have been discovered by Johann Joachim Becher, who obtained it by heating ethanol with sulfuric acid;[38] he mentioned the gas in his Physica Subterranea ().[39] Joseph Priestley also mentions the gas in his Experiments and observations relating to the various branches of natural philosophy: with a continuation of the observations on air (), where he reports that Jan Ingenhousz saw ethylene synthesized in the same way by a Mr. Enée in Amsterdam in and that Ingenhousz subsequently produced the gas himself.[40] The properties of ethylene were studied in by four Dutch chemists, Johann Rudolph Deimann, Adrien Paets van Troostwyck, Anthoni Lauwerenburgh and Nicolas Bondt, who found that it differed from hydrogen gas and that it contained both carbon and hydrogen.[41] This group also discovered that ethylene could be combined with chlorine to produce the Dutch oil, 1,2-dichloroethane; this discovery gave ethylene the name used for it at that time, olefiant gas (oil-making gas.)[42] The term olefiant gas is in turn the etymological origin of the modern word "olefin", the class of hydrocarbons in which ethylene is the first member.[citation needed]

In the mid-19th century, the suffix -ene (an Ancient Greek root added to the end of female names meaning "daughter of") was widely used to refer to a molecule or part thereof that contained one fewer hydrogen atoms than the molecule being modified. Thus, ethylene (C
2H
4) was the "daughter of ethyl" (C
2H
5). The name ethylene was used in this sense as early as .[43]

In , the German chemist August Wilhelm von Hofmann proposed a system of hydrocarbon nomenclature in which the suffixes -ane, -ene, -ine, -one, and -une were used to denote the hydrocarbons with 0, 2, 4, 6, and 8 fewer hydrogens than their parent alkane.[44] In this system, ethylene became ethene. Hofmann's system eventually became the basis for the Geneva nomenclature approved by the International Congress of Chemists in , which remains at the core of the IUPAC nomenclature. However, by that time, the name ethylene was deeply entrenched, and it remains in wide use today, especially in the chemical industry.

Following experimentation by Luckhardt, Crocker, and Carter at the University of Chicago,[45] ethylene was used as an anesthetic.[46][7] It remained in use through the s use even while chloroform was being phased out. Its pungent odor and its explosive nature limit its use today.[47]

Nomenclature

The IUPAC nomenclature rules made an exception for retaining the non-systematic name ethylene;[48] however, this decision was reversed in the rules,[49] and it remains unchanged in the newest recommendations,[50] so the IUPAC name is now ethene. In the IUPAC system, the name ethylene is reserved for the divalent group -CH2CH2-. Hence, names like ethylene oxide and ethylene dibromide are permitted, but the use of the name ethylene for the two-carbon alkene is not. Nevertheless, use of the name ethylene for H2C=CH2 (and propylene for H2C=CHCH3) is still prevalent among chemists in North America.[51]

Greenhouse gas emissions

"A key factor affecting petrochemicals life-cycle emissions is the methane intensity of feedstocks, especially in the production segment."[52] Emissions from cracking of naptha and natural gas (common in the US as gas is cheap there) depend a lot on the source of energy (for example gas burnt to provide high temperatures[53]) but that from naptha is certainly more per kg of feedstock.[54] Both steam cracking and production from natural gas via ethane are estimated to emit 1.8 to 2kg of CO2 per kg ethylene produced,[55] totalling over 260 million tonnes a year.[56] This is more than all other manufactured chemicals except cement and ammonia.[57] According to a report using renewable or nuclear energy could cut emissions by almost half.[54]

Safety

Like all hydrocarbons, ethylene is a combustible asphyxiant. It is listed as an IARC group 3 agent, since there is no current evidence that it causes cancer in humans.[58]

See also

• RediRipe, an ethylene detector for fruits.

References

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