Xylene encompasses three isomers of dimethylbenzene. The isomers are distinguished by the designations ortho- (o-), meta- (m-), and para- (p-), which specify to which carbon atoms (of the benzene ring) the two methyl groups are attached. Counting the carbon atoms from one of the ring carbons bonded to a methyl group, and counting towards the second ring carbon bonded to a methyl group, the o- isomer has the IUPAC name of 1,2-dimethylbenzene, the m- isomer has the IUPAC name of 1,3-dimethylbenzene, and the p- isomer has the IUPAC name of 1,4-dimethylbenzene. The mixture is a slightly greasy, colourless liquid commonly encountered as a solvent. Several million tons are produced annually.[1]
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Xylenes represent about 0.5–1% of crude oil, depending on the source (hence xylenes are found in small amounts in gasoline and airplane fuels). It is mainly produced from reformate. It is also obtained from coal carbonisation derived from coke ovens. It is produced by dehydrocyclodimerization and by methylating of toluene and benzene.[1][2]
Via the Isomar process, the ratio of isomers can be shifted to favor p-xylene, which is most valued. This conversion is catalysed by zeolites.[1]
Some chemical and physical properties differ from isomer to isomer. The melting point ranges from −47.87 °C (−54.17 °F) (m-xylene) to 13.26 °C (55.87 °F) (p-xylene). The boiling point for each isomer is around 140 °C (284.00 °F). The density of each is around 0.87 g/mL (7.26 lb/U.S. gallon or 8.72 lb/imp gallon) and thus is less dense than water. Xylene in air can be smelled at 0.08 to 3.7 parts of xylene per million parts of air (ppm) and can begin to be tasted in water at 0.53 to 1.8 ppm.
| Xylene Isomers | ||||
|---|---|---|---|---|
| General | ||||
| Common name | Xylenes | o-Xylene | m-Xylene | p-Xylene |
| Systematic name | Dimethylbenzenes | 1,2-Dimethylbenzene | 1,3-Dimethylbenzene | 1,4-Dimethylbenzene |
| Other names | Xylols | o-Xylol; Orthoxylene |
m-Xylol; Metaxylene |
p-Xylol; Paraxylene |
| Molecular formula | C8H10, C6H4(CH3)2 or C6H4C2H6 | |||
| SMILES | Cc1c(C)cccc1 | Cc1cc(C)ccc1 | Cc1ccc(C)cc1 | |
| Molar mass | 106.16 g/mol | |||
| Appearance | clear, colorless liquid | |||
| CAS number | [1330-20-7] | [95-47-6] | [108-38-3] | [106-42-3] |
| Properties | ||||
| Density and phase | 0.864 g/mL, liquid | 0.88 g/mL, liquid | 0.86 g/mL, liquid | 0.86 g/mL, liquid |
| Solubility in water | practically insoluble | |||
| Soluble in non-polar solvents such as aromatic hydrocarbons | ||||
| Melting point | −47.4 °C (−53.3 °F; 226 K) | −25 °C (−13 °F; 248 K) | −48 °C (−54.4 °F; 225 K) | 13 °C (55.4 °F; 286 K) |
| Boiling point | 138.5 °C (281.3 °F; 412 K) | 144 °C (291.2 °F; 417 K) | 139 °C (282.2 °F; 412 K) | 138 °C (280.4 °F; 411 K) |
| Viscosity | 0.812 cP at 20 °C (68 °F) | 0.62 cP at 20 °C (68 °F) | 0.34 cP at 30 °C (86 °F) | |
| Hazards | ||||
| MSDS | Xylenes | o-Xylene | m-Xylene | p-Xylene |
| EU Classification | Harmful (Xn) | |||
| NFPA 704 | ||||
| Flash point | 30 °C (86 °F) | 17 °C (63 °F) | 25 °C (77 °F) | 25 °C (77 °F) |
| R/S statement | R10, R20/21, R38: (S2), S25 | |||
| RTECS number | ZE2450000 | ZE2275000 | ZE2625000 | |
| Supplementary data page | ||||
| Structure & properties | n, εr, etc. | |||
| Thermodynamic data | Phase behaviour Solid, liquid, gas |
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| Spectral data | UV, IR, NMR, MS | |||
| Related compounds | ||||
| Related aromatic hydrocarbons |
toluene, mesitylene, benzene, ethylbenzene | |||
| Related compounds | xylenols - types of phenols | |||
| Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Infobox disclaimer and references |
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Xylenes form azeotropes with water and a variety of alcohols. With water the azeotrope consists of 60% xylenes and boils at 90 °C. As with many alkylbenzene compounds, xylenes form complexes with various halocarbons. The complexes of different isomers often have dramatically different properties from each other.[3]
p-Xylene is the principal precursor to terephthalic acid and dimethyl terephthalate, both monomers used in the production of polyethylene terephthalate (PET) plastic bottles and polyester clothing. 98% of p-xylene production, and half of all xylene, is consumed in this way.[4][5] o-Xylene is an important precursor to phthalic anhydride. The demand for isophthalic acid is relatively modest so m-xylene is rarely sought (and hence the utility of its conversion to the o- and p-isomers).
Xylene is used as a solvent. In this application, the mixture of isomers is often referred to as xylenes or xylol. Solvent xylene often contains a small percentage of ethylbenzene. Like the individual isomers, the mixture is colorless, sweet-smelling, and highly flammable. Areas of application include printing, rubber, and leather industries. It is a common component of ink, rubber, adhesive,[6] and leather industries. In thinning paints and varnishes, it can be substituted for toluene where slower drying is desired. Similarly it is a cleaning agent, e.g., for steel, silicon wafers, and chips. In dentistry, xylene can be used to dissolve gutta percha, a material used for endodontics (root canal treatments).
It is used in the laboratory to make baths with dry ice to cool reaction vessels, and as a solvent to remove synthetic immersion oil from the microscope objective in light microscopy.[7] In histology, xylene is also used for clearing the tissues following dehydration in preparation for paraffin wax infiltration. It is also used after sections have been stained to make them hydrophobic so that a coverslip may be applied with a resin in solvent.
Although conversion to terephthalic acid is the dominant chemical conversion, xylenes are precursors to other chemical compounds. For instance chlorination of both methyl groups gives the corresponding xylene dichlorides (bis(chloromethyl)benzenes).
Xylenes are not highly toxic as indicated by the high values of the LD50, which range from 200 to 4000 mg/kg for animals. The principal mechanism of detoxification is oxidation to methylbenzoic acid and hydroxylation to hydroxylene.[1]
Chisholm, Hugh, ed (1911). "xylene". Encyclopædia Britannica (11th ed.). Cambridge University Press.