Difference between revisions of "Benzene" - New World Encyclopedia

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==Structure==
 
==Structure==
:''Main article: [[aromaticity]]''
 
  
The formula of benzene (C<sub>6</sub>H<sub>6</sub>) mystified scientists who could not figure out its structure. [[Friedrich August Kekulé von Stradonitz]] was the first to deduce the ring structure of benzene. An often-repeated story claims that after years of studying carbon bonding, benzene and related molecules, he dreamt one night of the [[Ouroboros]], a [[snake]] eating its own tail, and that upon waking he was inspired to deduce the ring structure of benzene. However, the story first appeared in the ''Berichte der Durstigen Chemischen Gesellschaft'' (Journal of the Thirsty Chemical Society), a [[parody]] of the ''Berichte der Deutschen Chemischen Gesellschaft'', which appeared annually in the late-19th century on the occasion of the congress of German chemists; as such, it is probably to be treated with circumspection.
+
Scientists who were familiar with the chemical formula of benzene (C<sub>6</sub>H<sub>6</sub>) were mystified about its molecular structure. They knew that each molecule of benzene contained six carbon atoms, but the substance did not behave as though each molecule was an open-ended chain. [[Friedrich August Kekulé von Stradonitz]]* is usually credited with being the first to deduce the ring structure of benzene, in 1865. It is reported that after he had spent a long time mentally wrestling over the matter, he had a dream of a [[snake]] swallowing its own tail. This image inspired him to think of a ring structure for benzene.
  
While his (more formal) claims were well-publicized and accepted, by the early-1920s Kekulé's biographer came to the conclusion that Kekulé's understanding of the tetravalent nature of carbon bonding depended on the previous research of [[Archibald Scott Couper]] (1831-1892); furthermore, [[Josef Loschmidt]] (1821-1895) had earlier posited a cyclic structure for benzene as early as 1861. The cyclic nature of benzene was finally confirmed by the eminent crystallographer [[Kathleen Lonsdale]].
+
When Kekulé made his formal claims, they were well-publicized and accepted. It now appears that a lesser known scientist, [[Josef Loschmidt]]* (1821-1895), had posited a cyclic structure for benzene in a booklet published in 1861. Whether Kekulé actually had the dream or whether he borrowed from Loschmidt's publication are matters of debate. It is entirely possible that the two scientists thought of benzene's ring structure independently&mdash;a type of occurrence that is not unusual in science. The cyclic nature of benzene was finally confirmed by the crystallographer [[Kathleen Lonsdale]]*.
  
Benzene presents a special problem in that, to account for all the bonds, there must be alternating [[double bond|double]] carbon bonds:
+
Benzene presents a special problem in that, to account for all the bonds, there must be alternating single and double [[covalent bond]]s between carbon atoms, which may be represented as:
  
 
[[image:benz1.png|Benzene with alternating double bonds]]
 
[[image:benz1.png|Benzene with alternating double bonds]]
  
Using [[X-ray diffraction]], researchers discovered that all of the carbon-carbon bonds in benzene are of the same length of 140 [[picometre]]s (pm). The C-C [[bond length]]s are greater than a double bond (134pm) but shorter than a single bond (147pm). <!--147pm is the sp2-sp2 single bond length without conjugation —>  This intermediate distance is explained by electron [[delocalized|delocalization]]: the electrons for C-C bonding are distributed equally between each of the six carbon atoms. One representation is that the structure exists as a superposition of so-called [[resonance structure]]s, rather than either form individually. This delocalisation of electrons is known as [[aromaticity]], and gives benzene great stability. This enhanced stability is the fundamental property of aromatic molecules that differentiates them from molecules that are non-aromatic. To reflect the delocalised nature of the bonding, benzene is often depicted with a circle inside a hexagonal arrangement of carbon atoms:
+
Using the technique known as [[X-ray diffraction]]*, researchers discovered that all the carbon-carbon (C-C) bonds in benzene have the same length (140 picometers (pm)). The length of each C-C bond is greater than that of a double bond (134 pm) but shorter than a single bond (147 pm). The bond length of 140 pm, which is intermediate in length, is explained by the concept of "electron [[delocalized|delocalization]]*": the electrons for C-C bonding are distributed equally among the six carbon atoms. [One representation is that the structure exists as a superposition of two "[[resonance structure]]*s," rather than either form individually.]
 +
 
 +
This delocalization of electrons is known as ''[[aromaticity]]*'', which gives benzene great stability. This enhanced stability is a fundamental property of a class of molecules called "aromatic molecules," differentiating them from molecules that are not aromatic. To reflect the delocalized nature of the bonding, benzene is often depicted with a circle inside a hexagonal arrangement of carbon atoms (which are not labeled):
  
 
[[image:benz4.png|Benzene structure with a circle inside the hexagon]]
 
[[image:benz4.png|Benzene structure with a circle inside the hexagon]]
 
As is common in organic chemistry, the carbon atoms in the diagram above have been left unlabeled.
 
 
Benzene occurs sufficiently often as a component of organic molecules that there is a [[Unicode]] symbol with the code 232C to represent it: <font size="20">⌬</font>.
 
Many fonts do not have this Unicode character, so a browser may not be able to display it correctly.''
 
  
 
==Substituted benzene derivatives==
 
==Substituted benzene derivatives==

Revision as of 15:57, 21 September 2006

Benzene
Benzene
General
Systematic name Benzene
Other names Benzol
Molecular formula C6H6
SMILES c1ccccc1
C1=CC=CC=C1
InChI InChI=1/C6H6
/c1-2-4-6-5-3-1/h1-6H
Molar mass 78.11 g/mol
Appearance Colorless liquid
CAS number [71-43-2]
Properties
Density and phase 0.8786 g/cm³, liquid
Solubility in water 1.79 g/L (25 °C)
Melting point 5.5 °C (278.6 K)
Boiling point 80.1 °C (353.2 K)
Viscosity 0.652 cP at 20 °C
Structure
Molecular shape Planar
Symmetry group D6h
Dipole moment 0 D
Hazards
MSDS External MSDS
EU classification Flammable (F)
Carc. Cat. 1
Muta. Cat. 2
Toxic (T)
NFPA 704

NFPA 704.svg

3
2
0
 
R-phrases R45, R46, R11, R36/38,
R48/23/24/25, R65
S-phrases S53, S45
Flash point −11 °C
Autoignition temperature 561 °C
RTECS number CY1400000
Related compounds
Related
hydrocarbons
cyclohexane
naphthalene
Related compounds toluene
borazine
Except where noted otherwise, data are given for
materials in their standard state (at 25°C, 100 kPa)
Infobox disclaimer and references

Benzene (also known as benzol or [6]-annulene) is a colorless, flammable, sweet-smelling liquid. It is a natural constituent of crude oil but is usually synthesized from other compounds present in petroleum. Chemically, it is classified as an aromatic hydrocarbon, which is a group of organic compounds. Its chemical formula is C6H6. It is carcinogenic but is a valuable solvent used in the laboratory and industry. It is an important precursor in the production of a wide range of materials, including drugs, plastics, synthetic rubber, and dyes.

History

Benzene has been the subject of studies by many famous scientists, including Michael Faraday and Linus Pauling. In 1825, Faraday reported its isolation from oil gas and gave it the name bicarburet of hydrogen. In 1833, Eilhard Mitscherlich produced it by the distillation of benzoic acid (from gum benzoin) and lime (calcium oxide). Mitscherlich named the compound benzin. In 1845, Charles Mansfield, working under August Wilhelm von Hofmann, isolated benzene from coal tar. Four years later, Mansfield began the first industrial-scale production of benzene, based on the coal-tar method.

Structure

Scientists who were familiar with the chemical formula of benzene (C6H6) were mystified about its molecular structure. They knew that each molecule of benzene contained six carbon atoms, but the substance did not behave as though each molecule was an open-ended chain. Friedrich August Kekulé von Stradonitz is usually credited with being the first to deduce the ring structure of benzene, in 1865. It is reported that after he had spent a long time mentally wrestling over the matter, he had a dream of a snake swallowing its own tail. This image inspired him to think of a ring structure for benzene.

When Kekulé made his formal claims, they were well-publicized and accepted. It now appears that a lesser known scientist, Josef Loschmidt (1821-1895), had posited a cyclic structure for benzene in a booklet published in 1861. Whether Kekulé actually had the dream or whether he borrowed from Loschmidt's publication are matters of debate. It is entirely possible that the two scientists thought of benzene's ring structure independently—a type of occurrence that is not unusual in science. The cyclic nature of benzene was finally confirmed by the crystallographer Kathleen Lonsdale.

Benzene presents a special problem in that, to account for all the bonds, there must be alternating single and double covalent bonds between carbon atoms, which may be represented as:

Benzene with alternating double bonds

Using the technique known as X-ray diffraction, researchers discovered that all the carbon-carbon (C-C) bonds in benzene have the same length (140 picometers (pm)). The length of each C-C bond is greater than that of a double bond (134 pm) but shorter than a single bond (147 pm). The bond length of 140 pm, which is intermediate in length, is explained by the concept of "electron delocalization": the electrons for C-C bonding are distributed equally among the six carbon atoms. [One representation is that the structure exists as a superposition of two "resonance structures," rather than either form individually.]

This delocalization of electrons is known as aromaticity, which gives benzene great stability. This enhanced stability is a fundamental property of a class of molecules called "aromatic molecules," differentiating them from molecules that are not aromatic. To reflect the delocalized nature of the bonding, benzene is often depicted with a circle inside a hexagonal arrangement of carbon atoms (which are not labeled):

Benzene structure with a circle inside the hexagon

Substituted benzene derivatives

Main article: Aromatic hydrocarbons

Many important chemicals are derived from benzene, wherein with one or more of the hydrogen atoms is replaced with another functional group. Examples of simple benzene derivatives are phenol, toluene, and aniline, abbreviated PhOH,PhMe, and PhNH2, respectively. Linking benzene rings gives biphenyl, C6H5-C6H5. Further loss of hydrogen gives "fused" aromatic hydrocarbons, such naphthalene and anthracene. The limit of the fusion process is the hydrogen-free material graphite.

In heterocycles, carbon atoms in the benzene ring are replaced with other elements. The most important derivatives are the rings containing nitrogen. Replacing one CH with N gives the compound pyridine, C5H5N. Although benzene and pyridine are structurally related, benzene cannot be converted into pyridine. Replacement of a second CH bond with N gives, depending on the location of the second N, pyridazine, pyrimidine, and pyrazine.

Production

Trace amounts of benzene may result whenever carbon-rich materials undergo incomplete combustion. It is produced in volcanoes and forest fires, and is also a component of cigarette smoke.

Up until World War II, most benzene was produced as a byproduct of coke production in the steel industry. However, in the 1950s, increased demand for benzene, especially from the growing plastics industry, necessitated the production of benzene from petroleum. Today, most benzene comes from the petrochemical industry, with only a small fraction being produced from coal.

Three chemical processes contribute equally to industrial benzene production: catalytic reforming, toluene hydrodealkylation, and steam cracking.

Catalytic reforming

In catalytic reforming, a mixture of hydrocarbons with boiling points between 60-200 °C is blended with hydrogen gas and then exposed to a bifunctional platinum chloride or rhenium chloride catalyst at 500-525 °C and pressures ranging from 8-50 atm. Under these conditions, aliphatic hydrocarbons form rings and lose hydrogen to become aromatic hydrocarbons. The aromatic products of the reaction are then separated from the reaction mixture by extraction with any one of a number of solvents, including diethylene glycol or sulfolane, and benzene is then separated from the other aromatics by distillation.

Toluene hydrodealkylation

Toluene hydrodealkylation converts toluene to benzene. In this process, toluene is mixed with hydrogen, then passed over a chromium, molybdenum, or platinum oxide catalyst at 500-600 °C and 40-60 atm pressure. Sometimes, higher temperatures are used instead of a catalyst. Under these conditions, toluene undergoes dealkylation according to the chemical equation:

C6H5CH3 + H2 → C6H6 + CH4

Typical reaction yields exceed 95%. Sometimes, xylene and heavier aromatics are used in place of toluene, with similar efficiency.

Toluene disproportionation

Where a chemical complex has similar demands for both benzene and xylene, then toluene disproportionation (TDP) may be an attractive alternative. Broadly speaking 2 toluene molecules are reacted and the methyl groups rearranged from one toluene molecule to the other, yielding one benzene molecule and one xylene molecule.

Given that demand for para-xylene (p-xylene) substantially exceeds demand for other xylene isomers, a refinement of the TDP process called Selective TDP (STDP) may be used. In this process, the xylene stream exiting the TDP unit is approximately 90% paraxylene.

Steam cracking

Steam cracking is the process for producing ethylene and other olefins from aliphatic hydrocarbons. Depending on the feedstock used to produce the olefins, steam cracking can produce a benzene-rich liquid byproduct called pyrolysis gasoline. Pyrolysis gasoline can be blended with other hydrocarbons as a gasoline additive, or distilled to separate it into its components, including benzene.

Uses

Early uses

In the 19th and early-20th centuries, benzene was used as an after-shave lotion because of its pleasant smell. Prior to the 1920s, benzene was frequently used as an industrial solvent, especially for degreasing metal. As its toxicity became obvious, benzene has been supplanted by other solvents.

In 1903, Lugwig Roselius popularized the use of benzene to decaffeinate coffee. This discovery lead to the production of Sanka, -ka for kaffein, This process was later discontinued.

As a petrol additive, benzene increases the octane rating and reduces knocking. Consequently, petrol often contained several percent benzene before the 1950s, when tetraethyl lead replaced it as the most widely-used antiknock additive. However, with the global phaseout of leaded petrol, benzene has made a comeback as a gasoline additive in some nations. In the United States, concern over its negative health effects and the possibility of benzene entering the groundwater have led to stringent regulation of petrol's benzene content, with values around 1% typical. European petrol specifications now contain the same 1% limit on benzene content.

Current uses of benzene

Today benzene is mainly used as an intermediate to make other chemicals. Its most widely-produced derivatives include styrene, which is used to make polymers and plastics, phenol for resins and adhesives (via cumene), and cyclohexane, which is used in the manufacture of Nylon. Smaller amounts of benzene are used to make some types of rubbers, lubricants, dyes, detergents, drugs, explosives and pesticides.

In laboratory research, toluene is now often used as a substitute for benzene. The solvent-properties of the two are similar but toluene is less toxic and has a wider liquid range.

Benzene has been used as a basic research tool in a variety of experiments including analysis of a two-dimensional gas.

Reactions of benzene

Electrophilic aromatic substitution of benzene
  • Electrophilic aromatic substitution is a general method of derivatizing benzene. Benzene is sufficiently nucleophilic that it undergoes substitution by acylium ions or alkyl carbocations to afford give substituted derivatives.
    • The Friedel-Crafts acylation is a specific example of electrophilic aromatic substitution. The reaction involves the acylation of benzene (or many other aromatic rings) with an acyl chloride using a strong Lewis acid catalyst such as aluminium chloride..
Friedel-Crafts alkylation of benzene with methyl chloride
    • Like the Friedel-Crafts acylation, the Friedel-Crafts alkylation involves the alkylation of benzene (and many other aromatic rings) usng an alkyl halide in the presence of a strong Lewis acid catalyst.
    • sulfonation.
    • Nitration: Benzene undergoes nitration with nitronioum ions (NO2+) as the electrophile. Thus, warming benzene with a combination of concentrated sulphuric and nitric acid gives nitrobenzene.
  • Hydrogenation: Benzene and derivatives convert to cyclohexane and derivatives when treated with hydrogen at high hydrogen pressures.
  • Benzene is an excellent ligand in the organometallic chemistry of low-valent metals. Important examples include the sandwich and half-sandwich complexes respectively Cr(C6H6)2 and [RuCl2(C6H6)]2.

Health effects

Benzene exposure has serious health effects. Breathing high levels of benzene can result in death, while low levels can cause drowsiness, dizziness, rapid heart rate, headaches, tremors, confusion, and unconsciousness. Eating or drinking foods containing high levels of benzene can cause vomiting, irritation of the stomach, dizziness, sleepiness, convulsions, rapid heart rate, and death.

The major effect of benzene from chronic (long-term) exposure is to the blood. Benzene damages the bone marrow and can cause a decrease in red blood cells, leading to anemia. It can also cause excessive bleeding and depress the immune system, increasing the chance of infection.

Some women who breathed high levels of benzene for many months had irregular menstrual periods and a decrease in the size of their ovaries. It is not known whether benzene exposure affects the developing fetus in pregnant women or fertility in men.

Animal studies have shown low birth weights, delayed bone formation, and bone marrow damage when pregnant animals breathed benzene.

The US Department of Health and Human Services (DHHS) classifies benzene as a human carcinogen. Long-term exposure to high levels of benzene in the air can cause leukemia, a potentially fatal cancer of the blood-forming organs. In particular, Acute Myeloid Leukemia (AML) may be caused by benzene.

Several tests can show if you have been exposed to benzene. There is a test for measuring benzene in the breath; this test must be done shortly after exposure. Benzene can also be measured in the blood; however, since benzene disappears rapidly from the blood, measurements are accurate only for recent exposures.

In the body, benzene is metabolized. Certain metabolites can be measured in the urine. However, this test must be done shortly after exposure and is not a reliable indicator of how much benzene you have been exposed to, since the same metabolites may be present in urine from other sources.

The US Environmental Protection Agency has set the maximum permissible level of benzene in drinking water at 0.005 milligrams per liter (0.005 mg/L). The EPA requires that spills or accidental releases into the environment of 10 pounds (4.5 kg) or more of benzene be reported to the EPA.

The US Occupational Safety and Health Administration (OSHA) has set a permissible exposure limit of 1 part of benzene per million parts of air (1 ppm) in the workplace during an 8-hour workday, 40-hour workweek.

In March 2006, the official Food Standards Agency in Britain conducted a survey of 150 brands of soft drinks. It found that four contained benzene levels above World Health Organization limits. The affected batches were removed from sale.

In recent history there have been many examples of the harmful health effects of benzene and its derivatives. Toxic Oil Syndrome caused localised immune-suppression in Madrid in 1981 from people ingesting benzene-contaminated olive-oil. Chronic Fatigue Syndrome has also been highly correlated with people who eat "denatured" food that use solvents to remove fat or contain benzoic acid.

Workers in various industries that make or use benzene may be at risk for being exposed to high levels of this carcinogenic chemical. Industries that involve the use of benzene include the rubber industry, oil refineries, chemical plants, shoe manufacturers, and gasoline related industries. In 1987, OSHA estimated that about 237,000 workers in the United States were potentially exposed to benzene, and it is not known if this number has substantially changed since then.

Water and soil contamination are important pathways of concern for transmission of benzene contact. In the U.S. alone there are approximately 100,000 different sites which have benzene soil or groundwater contamination. In 2005, the water supply to the city of Harbin in China with a population of almost nine million people, was cut off because of a major benzene exposure. Benzene leaked into the Songhua River, which supplies drinking water to the city, after an explosion at a China National Petroleum Corporation (CNPC) factory in the city of Jilin on 13 November.

Main article: benzene in soft drinks

References
ISBN links support NWE through referral fees

  • Archibald Scott Couper, On a New Chemical Theory, Philosophical Magazine 16, 104-116 (1858)
  • Josef Loschmidt, Chemische Studien I, Carl Gerold's Sohn, Vienna (1861),
  • Josef Loschmidt, Chemische Studien I, Aldrich Chemical Co, Milwaukee (1989), catalog no. Z-18576-0, and (1913) catalog no. Z-18577-9
  • Kathleen Lonsdale, "The Structure of the Benzene Ring in Hexamethylbenzene," Proceedings of the Royal Society 123A: 494 (1929).
  • Kathleen Lonsdale, "An X-Ray Analysis of the Structure of Hexachlorobenzene, Using the Fourier Method," Proceedings of the Royal Society 133A: 536 (1931).
  • "FDA: Too Much Benzene In Some Drinks", CBS News, May 19, 2006, retrieved July 11, 2006

External links

Template:ChemicalSources


Simple:Benzene

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