Pyridine

Pyridine
IUPAC name Pyridine
Other names Azabenzene Azine py
Identifiers
CAS number	110-86-1 Y
PubChem	1049
ChemSpider	1020 Y
UNII	NH9L3PP67S Y
EC number	203-809-9
KEGG	C00747 Y
ChEBI	CHEBI:16227 N
ChEMBL	CHEMBL266158 Y
Jmol-3D images	Image 1
SMILES c1ccncc1
InChI InChI=1S/C5H5N/c1-2-4-6-5-3-1/h1-5H Y Key: JUJWROOIHBZHMG-UHFFFAOYSA-N Y InChI=1/C5H5N/c1-2-4-6-5-3-1/h1-5H Key: JUJWROOIHBZHMG-UHFFFAOYAY
Properties
Molecular formula	C5H5N
Molar mass	79.1 g mol−1
Appearance	colourless liquid
Density	0.9819 g/cm3, liquid
Melting point	-41.6 °C, 232 K, -43 °F
Boiling point	115.2 °C, 388 K, 239 °F
Solubility in water	Miscible
Vapor pressure	18 mmHg
Acidity (pKa)	5.25[1][2]
Refractive index (nD)	1.5093
Viscosity	0.88 cP
Dipole moment	2.2 D[3]
Hazards[4]
MSDS	External MSDS
EU classification	Flammable (F) Harmful (Xn)
R-phrases	R20 R21 R22 R34 R36 R38
NFPA 704	3 3 0
Flash point	21 °C
Threshold Limit Value	5 ppm (TWA)
Related compounds
Related amines	Picoline Quinoline
Related compounds	Aniline Pyrimidine Piperidine
Supplementary data page
Structure and properties	n, εr, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
N (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Pyridine is a basic heterocyclic organic compound with the chemical formula C₅H₅N. It is structurally related to benzene, with one C-H group replaced by a nitrogen atom. The pyridine ring occurs in many important compounds, including azines and the vitamins niacin and pyridoxal.

Pyridine was discovered in 1849 by the Scottish chemist Thomas Anderson as one of the constituents of bone oil. Two years later, Anderson isolated pure pyridine through fractional distillation of the oil. It is a colorless, highly flammable, weakly alkaline, water-soluble liquid with a distinctive, unpleasant fish-like odor.

Pyridine is used as a precursor to agrochemicals and pharmaceuticals and is also an important solvent and reagent. Pyridine is added to ethanol to make it unsuitable for drinking. It is used in the synthesis of sulfapyridine (a drug against bacterial and viral infections), antihistaminic drugs tripelennamine and mepyramine, as well as water repellents, bactericides and herbicides. Some chemical compounds, although not synthesized from pyridine, contain its ring structure. They include B vitamins niacin and pyridoxal, an anti-tuberculosis drug isoniazid, nicotine and other nitrogen-containing plant products.^[5] Historically, pyridine was produced from coal tar and as a by-product of the coal gasification. However, increased demand for pyridine resulted in the development of more economical methods of synthesis from acetaldehyde and ammonia, and more than 20,000 tonnes per year are manufactured worldwide.

History

Thomas Anderson

Impure pyridine was undoubtedly prepared by early alchemists by heating animal bones and other organic matter,^[6] but the earliest documented reference is attributed to the Scottish scientist Thomas Anderson.^[7][8] In 1849, Anderson examined the contents of the oil obtained through high-temperature heating of animal bones.^[8] Among other substances, he separated from the oil a colorless liquid with unpleasant odor, from which he isolated pure pyridine two years later. He described it as highly soluble in water, readily soluble in concentrated acids and salts upon heating, and only slightly soluble in oils. Owing to its flammability, Anderson named the new substance pyridine, after Greek: πῦρ (τὸ) (pyr) meaning fire. The suffix -idine was added in compliance with the chemical nomenclature, as in toluidine, to indicate a carbon cycle containing a nitrogen atom.^[9]

The chemical structure of pyridine was determined decades after its discovery. Wilhelm Körner (1869)^[10] and James Dewar (1871)^[11] independently suggested that, in analogy between quinoline and naphthalene, the structure of pyridine is derived from benzene by substituting one C-H unit with a nitrogen atom.^[12][13] The suggestion by Körner and Dewar was later confirmed in an experiment where pyridine was reduced to piperidine with sodium alcohol. In 1876, William Ramsay combined acetylene and hydrogen cyanide into pyridine in a red-hot iron-tube furnace. This was the first synthesis of a hetero-aromatic compound.^[14][15]

The contemporary methods of pyridine production had a low yield, and the increasing demand for the new compound urged to search for more efficient routes. A breakthrough came in 1924 when the Russian chemist Aleksei Chichibabin invented a pyridine synthesis reaction which was based on inexpensive reagents.^[16] This method is still used for the industrial production of pyridine.^[17]

Occurrence

Pyridine is not abundant in nature, except for the leaves and roots of belladonna (Atropa belladonna)^[18] and in marshmallow (Althaea officinalis).^[19] Pyridine derivatives, however, are often part of biomolecules such as the eponymous pyridine nucleotides and alkaloids. In daily life, trace amounts of pyridine are components of the volatile organic compounds that are produced in roasting and canning processes, e.g. in fried chicken,^[20] sukiyaki,^[21] fried bacon,^[22] Beaufort cheese,^[23] coffee aroma,^[24] black tea^[25] and sunflower honey.^[26] The smoke of tobacco^[27][28] and of marijuana^[15] contains pyridine, as do vaginal secretions.

Nomenclature

The systematic name of pyridine, within the Hantzsch–Widman nomenclature recommended by the IUPAC, is azine. However, systematic names for simple compounds are used very rarely, instead heterocyclic nomenclature follows historically established common names. IUPAC discourages the use of azine in favor of pyridine.^[29] The numbering of the ring atoms in pyridine starts at the nitrogen (see infobox). An allocation of positions by letter of the Greek alphabet (α-γ) and the substitution pattern nomenclature common for homoaromatic systems (ortho, meta, para) are used sometimes. Here α (ortho), β (meta) and γ (para) refer to the 2, 3 and 4 position, respectively. The systematic name for the pyridine derivatives is pyridinyl, wherein the position of the substituted atom is preceded by a number. However, here again the historical name pyridyl is encouraged by the IUPAC and used instead of the systematic name.^[30] The cationic derivative formed by the addition of an electrophile to the nitrogen atom is called pyridinium.

• 

  4-bromopyridine

• 

  2,2′-bipyridine

• 

  pyridine-2,6-dicarboxylic acid (dipicolinic acid)

• 

  General form of the pyridinium cation

Production

Historically, pyridine was extracted from coal tar or obtained as a by-product of coal gasification. The process was labor consuming and inefficient: coal tar contains only about 0.1% pyridine,^[31] and therefore a multi-stage purification was required, which further reduced the output. Nowadays, most pyridine is produced synthetically using various name reactions, and the major ones are discussed below.^[17]

Chichibabin synthesis

The Chichibabin pyridine synthesis was reported in 1924 and is still in use industrially.^[16] In its general form, the reaction can be described as a condensation reaction of aldehydes, ketones, α,β-Unsaturated carbonyl compounds, or any combination of the above, in ammonia or ammonia derivatives.^[32] In particular, unsubstituted pyridine is produced from formaldehyde and acetaldehyde, which are inexpensive and widely available. First, acrolein is formed in a Knoevenagel condensation from the acetaldehyde and formaldehyde. It is then condensed with acetaldehyde and ammonia into dihydropyridine, and then oxidized with a solid-state catalyst to pyridine. This process is carried out in a gas phase at 400–450 °C. The product consists of a mixture of pyridine, simple methylated pyridines (picoline) and lutidine; its composition depends on the catalyst used and can be adapted to the needs of the manufacturer. The catalyst is usually a transition metal salt such as cadmium(II) fluoride or manganese(II) fluoride, but cobalt and thallium compounds can also be used. The recovered pyridine is separated from by-products in a multistage process.^[17]

Formation of acrolein from acetaldehyde and formaldehyde

Condensation of pyridine from acrolein and acetaldehyde

Practical application of the traditional Chichibabin pyridine synthesis are limited by its consistently low yield, typically about 20%. This low yield, together with the high prevalence of byproducts, render unmodified forms of Chichibabin's method unpopular.^[32]

Dealkylation of alkylpyridines

Pyridine can be prepared by dealkylation of alkylated pyridines, which are obtained as by-products in the syntheses of other pyridines. The oxidative dealkylation is carried out either using air over vanadium(V) oxide catalyst,^[33] by vapor-dealkylation on nickel-based catalyst,^[34][35] or hydrodealkylation with a silver or platinum based catalyst.^[36] Yields of pyridine up to be 93% can be achieved with the nickel-based catalyst.^[17]

Hantzsch pyridine synthesis

The first major synthesis of pyridine derivatives was described in 1881 by Arthur Rudolf Hantzsch.^[37] The Hantzsch pyridine synthesis typically uses a 2:1:1 mixture of a β-keto acid (often acetoacetate), an aldehyde (often formaldehyde), and ammonia or its salt as the nitrogen donor. First, a double hydrogenated pyridine is obtained, which is then oxidized to the corresponding pyridine derivative. Emil Knoevenagel showed that unsymmetrically substituted pyridine derivatives can be produced with this process.^[38]

Hantzsch pyridine synthesis with acetoacetate, formaldehyde and ammonium acetate, and iron(III) chloride as the catalyst.

Bönnemann cyclization

Bönnemann cyclization

The trimerization of a part of a nitrile molecule and two parts of acetylene into pyridine is called Bönnemann cyclization. This modification of the Reppe synthesis can be activated either by heat or light. While the thermal activation requires high pressures and temperatures, the photoinduced cycloaddition proceeds at ambient conditions with CoCp₂(cod) (Cp = cyclopentadienyl, cod = 1,5-cyclooctadiene) as a catalyst, and can be performed even in water.^[39] A series of pyridine derivatives can be produced in this way. When using acetonitrile as the nitrile, 2-methylpyridine is obtained, which can be dealkylated to pyridine.

Other methods

The Kröhnke pyridine synthesis involves the condensation of 1,5-diketones with ammonium acetate in acetic acid followed by oxidation.

The Ciamician-Dennstedt rearrangement entails the ring-expansion of pyrrole with dichlorocarbene to 3-chloropyridine.^[40]

In the Gattermann-Skita synthesis,^[41] a malonate ester salt reacts with dichloromethylamine.^[42]

Biosynthesis

Several pyridine derivatives play important roles in biological systems. While its biosynthesis is not fully understood, nicotinic acid (vitamin B₃) occurs in some bacteria, fungi and mammals. Mammals synthesize nicotinic acid through oxidation of the amino acid tryptophan, where an intermediate product, aniline, creates a pyridine derivative, kynurenine. On the contrary, the bacteria Mycobacterium tuberculosis and Escherichia coli produce nicotinic acid by condensation of glyceraldehyde 3-phosphate and aspartic acid.^[43]

Properties

Physical properties

Crystal structure of pyridine

Pyridine is a colorless liquid that boils at 115.2 °C and freezes at −41.6 °C. Its density, 0.9819 g/cm³, is close to that of water, and its refractive index is 1.5093 at a wavelength of 589 nm and a temperature of 20 °C.^[44] Addition of up to 40 mol% of water to pyridine gradually lowers its melting point from −41.6 °C to −65.0 °C.^[45] The molecular electric dipole moment is 2.2 Debye.^[3] Pyridine is diamagnetic and has a diamagnetic susceptibility of −48.7×10⁻⁶ cm³·mol⁻¹.^[46] The standard enthalpy of formation is 100.2 kJ·mol⁻¹ in the liquid phase^[47] and 140.4 kJ·mol⁻¹ in the gas phase. At 25 °C pyridine has a viscosity^[48] of 0.88 mPa/s and thermal conductivity of 0.166 W·m⁻¹·K⁻¹.^[49][50] The enthalpy of vaporization is 35.09 kJ·mol⁻¹ at the boiling point and normal pressure.^[51] The enthalpy of fusion is 8.28 kJ·mol⁻¹ at the melting point.^[52]

Pyridine crystallizes in an orthorhombic crystal system with space group Pna2₁ and lattice parameters a = 1752, b = 897, c = 1135 pm, and 16 formula units per unit cell (measured at 153 K). For comparison, benzene crystal is also orthorhombic, with space group Pbca, a = 729.2 pm, b = 947.1 pm, c = 674.2 pm (at 78 K), but the number of molecules per cell is only 4.^[53] This difference is partly related to the lower symmetry of the individual pyridine molecule (C_2v vs. D_6h for benzene). A trihydrate (pyridine·3H₂O) is known; it also crystallizes in an orthorhombic system in the space group Pbca, lattice parameters a = 1244, b = 1783, c = 679 pm and eight formula units per unit cell (measured at 223 K).^[45]

The critical parameters of pyridine are pressure 6.70 MPa, temperature 620 K and volume 229 cm³·mol⁻¹.^[54] In the temperature range 340–426 °C its vapor pressure p can be described with the Antoine equation

$\log_{10} p = A-\frac{B}{C+T}$

where T is temperature, A = 4.16272, B = 1371.358 K and C = -58.496 K.^[55]

The optical absorption spectrum of pyridine in hexane contains three bands at the wavelengths of 195 nm (π → π*transition, molar absorptivity ε = 7500 L·(mol·cm)⁻¹), 251 nm (π → π*transition, ε = 2000 L·(mol·cm)⁻¹) and 270 nm (n → π*transition, ε = 450 L·(mol·cm)⁻¹).^[56] The ¹H nuclear magnetic resonance (NMR) spectrum of pyridine contains three signals with the integral intensity ratio of 2:1:2 which correspond to the three chemically different protons in the molecule. These signals originate from the α-protons (chemical shift 8.5 ppm), γ-proton (7.5 ppm) and β-protons (7.1 ppm). The carbon analog of pyridine, benzene, has only one proton signal at 7.27 ppm. The larger chemical shifts of the α- and γ-protons in comparison to benzene result from the lower electron density in the α- and γ-positions, which can be derived from the resonance structures. The situation is rather similar for the ¹³C NMR spectra of pyridine and benzene: pyridine shows a triplet at δ (α-C) = 150 ppm, δ (β-C) = 124 ppm and δ (γ-C) = 136 ppm, whereas benzene has a single line at 129 ppm. All shifts are quoted for the solvent-free substances.^[57] Pyridine is conventionally detected by the gas chromatography and Mass spectrometry methods.^[15]

Chemical properties

Pyridine is miscible with water and virtually all organic solvents.^[44] It is weakly basic, and with hydrochloric acid it forms a crystalline hydrochloride salt which melts at 145–147 °C.^[58] Most chemical properties of pyridine are typical of a heteroaromatic compound. In organic reactions, pyridine behaves both as a tertiary amine, undergoing protonation, alkylation, acylation, and N-oxidation at the nitrogen atom, and as an aromatic compound, undergoing nucleophilic substitutions.

Because of the electronegative nitrogen in the pyridine ring, the molecule is relatively electron deficient. It therefore enters less readily electrophilic aromatic substitution reactions, which are characteristic of benzene derivatives. However, unlike benzene and its derivatives, pyridine is more prone to nucleophilic substitution and metalation of the ring by strong organometallic bases.^[59][60] The reactivity of pyridine can be distinguished for three chemical groups. With electrophiles, electrophilic substitution takes place where pyridine expresses aromatic properties. With nucleophiles, pyridine reacts via its 2nd and 4th carbon atoms and thus behaves similar to imines and carbonyls. The reaction with many Lewis acids results in the addition to the nitrogen atom of pyridine, which is similar to the reactivity of tertiary amines. The ability of pyridine and its derivatives to oxidize, forming amine oxides (N-oxides), is also a feature of tertiary amines.^[61]

1,10-phenanthroline

The nitrogen center of pyridine features a basic lone pair of electrons. Because this lone pair is not part of the aromatic ring, pyridine is a base, having chemical properties similar to those of tertiary amines. The pKa of the conjugate acid is 5.25. Pyridine is protonated by reaction with acids and forms a positively charged aromatic polyatomic ion called pyridinium. The bond lengths and bond angles in pyridine and pyridinium are almost identical.^[62] The pyridinium cation is isoelectronic with benzene. Pyridinium p-toluenesulfonate (PPTS) is an illustrative pyridinium salt; it is produced by treating pyridine with p-toluenesulfonic acid.

Pyridine can act as Lewis base, donating its pair of electron to a Lewis acid as in the sulfur trioxide pyridine complex.

Pyridine itself is a relatively weak ligand in forming complexes with transition metal ions. For example, it forms a 1:1 complexes with nickel(II), Ni²⁺, and copper(II), Cu²⁺, with logK₁ values of ca. 1.9 and 2.6, respectively.^[63] The infrared spectra of pyridine complexes have been discussed in detail.^[64] Picolinic acid, which is a substituted derivative of pyridine, forms strong complexes due to the chelate effect. 2,2'-bipyridine and 1,10-phenanthroline, which can also be viewed as substituted derivatives of pyridine, also form strong complexes, such as in Ferroin which can be used as an redox indicator in the quantitative analysis of iron.^[65]

The η⁶ coordination mode, as occurs in η⁶ benzene complexes, is only observed in sterically encumbered derivatives that block the nitrogen center.^[66]

Molecular properties

Pyridine with its free electron pair

Pyridine has a conjugated system of six π-electrons that are delocalized over the ring. The molecule is planar and thus follows the Hückel criteria for aromatic systems. In contrast to benzene, the electron density is not evenly distributed over the ring, reflecting the negative inductive effect of the nitrogen atom. For this reason, pyridine has a dipole moment and a weaker resonant stabilization than benzene (resonance energy 117 kJ·mol⁻¹ in pyridine vs. 150 kJ·mol⁻¹ in benzene).^[67] The electron localization in pyridine is also reflected in the shorter C-N ring bond (137 pm for the C-N bond in pyridine vs. 139 pm for C-C bond in benzene),^[68] whereas the carbon-carbon bonds in the pyridine ring have the same 139 pm length as in benzene. These bond lengths lie between the values for the single and double bonds and are typical of aromatic compounds.

All the ring atoms in the pyridine molecule are sp²-hybridized. The nitrogen atom "donates" its three hybridized electrons to the ring system, and its extra electron pair lies in the molecule plane, projecting outwards, in the plane of the ring. This lone pair does not contribute to the aromatic system but importantly infuences the chemical properties of pyridine, as it easily supports bond formation via an electrophilic attack. However, because of the separation of the lone pair from the aromatic system of the ring affects, the nitrogen atom can not exhibit a positive mesomeric effect.

Many analogues of pyridine are known where N is replaced by other heteroatoms (see Figure). Substitution of one CH in pyridine with a second N gives rise to the "diaza" heterocycles (C₄H₄N₂), with the names pyridazine, pyrimidine, and pyrazine.

Bond lengths and angles of benzene, pyridine, phosphorine, arsabenzene, stilbabenzene, and bismabenzene	Electron orbitals in pyridine
Resonance structures of pyridine	Electron orbitals in protonated pyridine

Reactions

Many reactions that are characteristic of benzene proceed with pyridine either at more complicated conditions or/and with low yield. Owing to the decreased electron density in the aromatic system, electrophilic substitutions are suppressed in pyridine and its derivatives in favor of addition of nucleophiles at the electron-rich nitrogen atom. The nucleophilic addition at the nitrogen atom leads to a further deactivation of the aromatic properties and hindering of the electrophilic substitution. On the other hand, free-radical and nucleophilic substitutions occur more readily in pyridine than in benzene.^[3][59]

Electrophilic substitutions

Many electrophilic substitutions on pyridine either do not proceed or proceed only partially; however, the heteroaromatic character can be activated by electron-donating functionalization. Common alkylations and acylations, such as Friedel–Crafts alkylation or acylation, usually fail for pyridine because they only lead to the addition at the nitrogen atom. Substitutions usually occur at the 3-position which is the electron-richest carbon atom in the ring and is therefore more susceptible to an electrophilic addition.

Structure of pyridine-N oxide

Substitutions to pyridine at the 2- or 4-position result in an energetically unfavorable σ complex. They can be promoted, however, using clever experimental techniques, such as conducting electrophilic substitution on the pyridine-N-oxide followed by deoxygenation of the nitrogen atom. Addition of oxygen reduces electron density on the nitrogen atom and promotes substitution at the 2- and 4-carbons. The oxygen atom can then be removed via several routes, most commonly with compounds of trivalent phosphorus or divalent sulfur which are easily oxidized. Triphenylphosphine is a frequently used reagent, which is oxidized in this reaction to triphenylphosphine oxide. The following paragraphs describe representative electrophilic substitution reactions of pyridine.^[59]

Direct nitration of pyridine requires harsh conditions and has very low yields. The 3-nitropyridine can be obtained instead by reacting pyridine with dinitrogen pentoxide in presence of sodium.^[69][70][71] Pyridine derivatives where the nitrogen atom is screened sterically and/or electronically can be obtained by nitridation with nitronium tetrafluoroborate (NO₂BF₄). In this way, 3-nitropyridine can be obtained via the synthesis of 2,6-dibromopyridine followed by removal of the bromine atoms.^[72][73] Direct sulfonation of pyridine is even more difficult than direct nitridation. However, pyridine-3-sulfonic acid can be obtained at acceptable yield by boiling pyridine in an excess of sulfuric acid at 320 °C.^[74] Reaction with the SO₃ group also facilitates addition of sulfur to the nitrogen atom, especially in the presence of a mercury(II) sulfate catalyst.^[59][75]

In contrast to the nitration and sulfonation, the direct bromination and chlorination of pyridine proceed well. The reaction of pyridine with molecular bromine in sulfuric acid at 130 °C readily produced 3-bromopyridine. The yield is lower for 3-chloropyridine upon chlorination with molecular chlorine in the presence of aluminium chloride at 100 °C. Both 2-bromopyridine and 2-chloropyridine can be produced by direct reaction with halogen with a palladium(II) chloride catalyst.^[76]

Nucleophilic substitutions

In contrast to benzene, pyridine efficiently supports several nucleophilic substitutions, and is regarded as a good nucleophile (donor number 33.1). The reason for this is relatively lower electron density of the carbon atoms of the ring. These reactions include substitutions with elimination of a hydride ion and elimination-additions with formation of an intermediate arine configuration, and usually proceed at 2- or 4-position.^[59][60]

Many nucleophilic substitutions occur easier not with bare pyridine, but with pyridine modified with bromine, chlorine, fluorine or sulfonic acid fragments which then become a leaving group. So fluorine is the best leaving group for the substitution with organolithium compounds. The nucleophilic attack compounds may be alkoxides, thiolates, amines, and ammonia (at elevated pressures).^[77]

The hydride ion is generally a poor leaving group and occurs only in a few heterocyclic reactions. They include the Chichibabin reaction which yields pyridine derivatives aminated at the 2-position. Here sodium amide is used as the nucleophile yielding 2-aminopyridine. The hydride ion released in this reaction combines with a proton of an available amino group forming a hydrogen molecule.^[60][78]

Analogous to benzene, nucleophilic substitutions to pyridine can result in the formation of heteroarine intermediates. For this purpose, pyridine derivatives can be eliminated with good leaving groups using strong bases such as sodium and potassium tert-butoxide. The subsequent addition of a nucleophile to the triple bond has low selectivity and the result is a mixture of the two possible adducts.^[59]

Radical reactions

Pyridine supports a series of radical reactions, which is used in its dimerization to bipyridines. Radical dimerization of pyridine with elemental sodium or Raney nickel selectively yields 4,4'-bipyridine^[79] or 2,2'-bipyridine,^[80] which are important precursor reagents in the chemical industry. One of the name reactions involving free radicals is the Minisci reaction. It can produce 2-tert-butylpyridine upon reacting pyridine with pivalic acid, silver nitrate and ammonium in sulfuric acid with a yield of 97%.^[59]

Reactions on the nitrogen atom

Additions of various Lewis acids to pyridine

Lewis acids easily add to the nitrogen atom of pyridine forming pyridinium salts. The reaction with alkylhalides leads to alkylation of the nitrogen atom. This creates a positive charge in the ring that increases the reactivity of pyridine to both oxidation and reduction. The Zincke reaction is used for the selective introduction of radicals in pyridinium compounds (it has no relation to the chemical element zinc).

Hydrogenation and reduction

Reduction of pyridine to piperidine with Raney nickel

Hydrogen-saturated piperidine is obtained through reaction with hydrogen gas in the presence of Raney nickel.^[81] This reaction releases 193.8 kJ·mol⁻¹ of energy,^[82] which is slightly less than the energy of the hydrogenation of benzene (205.3 kJ·mol⁻¹).^[82]

Partially hydrogenated derivatives are obtained under milder conditions. For example, reduction with lithium aluminium hydride yields a mixture of 1,4-dihydropyridine, 1,2-dihydropyridine and 2,5-dihydropyridine.^[83] Selective synthesis of 1,4-dihydropyridine is achieved in the presence of organometallic complexes of magnesium and zinc,^[84] and (Δ3,4)-tetrahydropyridine is obtained by electrochemical reduction of pyridine.^[85]

Applications

Use of pyridine in the chemical industry, VEB Berlin-Chemie, 1959.

Pyridine is an important raw material in of the chemical industry, with the 1989 production of 26,000 tonnes in world-wide.^[17] Among major 25 production sites for pyridine eleven are located in Europe (as of 1999).^[15] The major producers of pyridine include Evonik Industries, Rütgers Chemicals, Imperial Chemical Industries and Koei Chemical.^[17] The pyridine production has significantly increased in the early 2000s, with an annual production capacity of 30,000 tonnes in mainland China alone.^[86] The US-Chinese joint venture Vertellus is currently the world leader in pyridine production.^[87]

Pyridine is used as polar, basic, low-reactive solvent, for example in Knoevenagel condensations.^[15] It is especially suitable for the dehalogenation, where it acts as the base of the elimination reaction and bonds the resulting hydrogen halide to form a pyridinium salt. In esterifications and acylations pyridine activates the carboxylic acid halides or anhydrides. Even more active in these reactions are the pyridine derivatives 4-dimethylaminopyridine (DMAP) and 4-(1-pyrrolidinyl) pyridine. Pyridine is also used as a base in condensation reactions.^[88]

Elimination reaction with pyridine to form pyridinium

Pyridinium chlorochromate was developed by Elias James Corey and William Suggs in 1975 and is used to oxidize primary alcohols to aldehydes and secondary alcohols to ketones.^[89] It is obtained by adding pyridine to a solution of chromic acid and concentrated hydrochloric acid:

C₅H₅N + HCl + CrO₃ → [C₅H₅NH][CrO₃Cl]

The carcinogenicity of the side-product chromyl chloride (CrO₂Cl₂) urged to look for alternative routes, such as treating chromium(VI) oxide with pyridinium chloride:^[90]

[C₅H₅NH⁺]Cl⁻ + CrO₃ → [C₅H₅NH][CrO₃Cl]

The Cornforth reagent (pyridinium dichromate, PDC),^[91] pyridinium chlorochromate (PCC), the Collins reagent (complex of chromium(VI) oxide with pyridine in dichloromethane)^[92][93] and the Sarret reagent (complex of chromium(VI) oxide with pyridine in pyridine) are similar chromium-based pyridine compounds, which are also used for oxidation, namely conversion of primary and secondary alcohols to ketones. The Collins and Sarret reagents are both difficult and dangerous to prepare, they are hygroscopic and can inflame during preparation. For this reason, the use of PCC and PDC was preferred. Those reagents were rather popular in the 1970s–1980s, but because of their toxicity and confirmed carcinogenic status, they are rarely used nowadays.^[94]

Oxidation of an alcohol to aldehyde with the Collins reagent.

structure of the Crabtree's catalyst

When a pyridine ligand is part of a metal complex, it can be easily replaced by a stronger chelating Lewis base. This property is exploited in catalysis of polymerization^[95][96] and hydrogenation reactions, using, for example, the Crabtree's catalyst.^[97] The pyridine ligand replaced during the reaction is restored after its completion.

In the pharmaceutical industry pyridine serves as a building block for making a variety of drugs, insecticides and herbicides. It was and is used in large quantities in the production of herbicides diquat and paraquat, which contain bipyridine fragments. The first synthesis step of insecticide chlorpyrifos consists of the chlorination of pyridine. Pyridine is also the starting compound for the preparation of pyrithione-based fungicides.^[15] Cetylpyridinium and laurylpyridinium, which can be produced from pyridine with a Zincke reaction, are used as antiseptic in oral and dental care products.^[3]

Synthesis of paraquat^[98]

In addition to pyridines, piperidine derivatives are also important synthetic building blocks. A common synthesis of piperidine is the reduction of pyridine with a nickel, cobalt or ruthenium-based catalyst at elevated temperatures.^[99]

Pyridine is used as a solvent in the manufacture of dyes and rubber.^[100] It is also used in the textile industry to improve network capacity of cotton.^[3] Pyridine is added to ethanol to make it unsuitable for drinking.^[3] In low doses, pyridine is added to foods to give them a bitter flavor, and such usage is approved by the US Food and Drug Administration.^[15] The detection threshold for pyridine in solutions is about 1–3 mmole·L⁻¹ (79–237 mg·L⁻¹).^[101] As a base, pyridine can be used as the Karl Fischer reagent, but it is usually replaced by alternatives with a more pleasant odor, such as imidazole.^[102]

Pyridine is widely used as a ligand in coordination chemistry. Also important are its chelating derivatives 2,2'-bipyridine, consisting of two pyridine molecules joined by a single bond, and terpyridine, a molecule of three pyridine rings linked together. Pyridine is easily attacked by alkylating agents to give N-alkylpyridinium salts. One example is cetylpyridinium chloride, a cationic surfactant that is a widely used disinfection and antiseptic agent. Pyridinium salts can be obtained in the Zincke reaction. Useful adducts of pyridine include pyridine-borane, C₅H₅NBH₃ (melting point  10–11 °C), a mild reducing agent with improved stability relative to NaBH₄ in protic solvents and improved solubility in aprotic organic solvents. Pyridine-sulfur trioxide, C₅H₅NSO₃ (melting point 175 °C) is a sulfonation agent used to convert alcohols to sulfonates, which in turn undergo C-O bond scission upon reduction with hydride agents.

Hazards

Pyridine has a flash point (the lowest temperature at which it can vaporize to form an ignitable mixture in air) of only 17 °C and is therefore highly flammable. Its ignition temperature is 550 °C, and mixtures of 1.7–10.6 vol% of pyridine with air are explosive. The thermal modification of pyridine starts above 490 °C, resulting in bipyridine (mainly 2,2'-bipyridine and to a lesser extent 2,3'-bipyridine and 2,4'-bipyridine), nitrogen oxides and carbon monoxide.^[50] Pyridine easily dissolves in water and harms both animals and plants in aquatic systems.^[103] The permitted maximum allowable concentration of pyridine was 15–30 parts per million (ppm, or 15–30 mg·m⁻³ in air) in most countries in the 1990s,^[15] but was reduced to 5 ppm in the 2000s.^[104] For comparison, indoor air contaminated with tobacco smoke may contain up to 16 µg·m⁻³, and one cigarette contains 21–32 µg of pyridine.^[15]

Health issues

Metabolism of pyridine

Pyridine is harmful if inhaled, swallowed or absorbed through the skin.^[105] Effects of an acute pyridine intoxication include dizziness, headache, lack of coordination, nausea, salivation and loss of appetite. They may progress into abdominal pain, pulmonary congestion and unconsciousness.^[106] One person died after accidental ingestion of half a cup of pyridine.^[15] The lowest known lethal dose (LD_Lo) for the ingestion of pyridine in humans is 500 mg·kg⁻¹. In high doses pyridine has a narcotic effect and its vapor concentrations of above 3600 ppm pose health risk.^[17] The LD₅₀ in rats (oral) is 891 mg·kg⁻¹. Pyridine is flammable.

Evaluations as a possible carcinogenic agent showed there is inadequate evidence in humans for the carcinogenicity of pyridine, albeit there is limited evidence of carcinogenic effects on animals.^[106] Available data indicate that "exposure to pyridine in drinking-water led to reduction of sperm motility at all dose levels in mice and increased estrous cycle length at the highest dose level in rats".^[106]

Pyridine might also have minor neurotoxic, genotoxic and clastogenic effects.^{[15][50][107]} Exposure to pyridine would normally lead to its inhalation and absorption in the lungs and gastrointestinal tract, where it either remains unchanged or is metabolized. The major products of pyridine metabolism are N-methylpyryliumhydroxide, which is formed by N-methyltransferases^{[disambiguation needed ]} (e.g. pyridine N-methyltransferase), as well as pyridine-N oxide, and 2-, 3- and 4-hydroxypyridine, which are generated by the action of monooxygenase. In humans, pyridine is metabolized only into N-methylpyryliumhydroxide.^[50][107] Pyridine is readily degraded by bacteria to ammonia and carbon dioxide.^[108] The unsubstituted pyridine ring degrades more rapidly than picoline, lutidine, chloropyridine, or aminopyridine^{[disambiguation needed ]},^[109] and a number of pyridine degraders have been shown to overproduce riboflavin in the presence of pyridine.^[110]

Minor amounts of pyridine are released into environment from some industrial processes such as steel manufacture,^[111] processing of oil shale, coal gasification, coking plants and incinerators.^[15] The atmosphere at oil shale processing plants can contain pyridine concentrations of up to 13 µg·m⁻³,^[112] and 53 µg·m⁻³ levels were measured in the groundwater in the vicinity of a coal gasification plant.^[113] According to a study by the US National Institute for Occupational Safety and Health, about 43,000 Americans work in contact with pyridine.^[114]

References

1. ^ Linnell, Robert (1960). Journal of Organic Chemistry 25 (2): 290. doi:10.1021/jo01072a623. 
2. ^ Pearson, Ralph G.; Williams, Forrest V. (1953). Journal of the American Chemical Society 75 (13): 3073. doi:10.1021/ja01109a008. 
3. ^ ^{a b c d e f} RÖMPP Online – Version 3.5. Stuttgart: Georg Thieme. 2009. 
4. ^ Pyridine MSDS
5. ^ Pyridine Encyclopedia Britannica on-line
6. ^ A. Weissberger (Ed.), A. Klingberg (ed.), R.A. Barnes, F. Brody, P.R. Ruby: Pyridine and its Derivatives, Volume 1, 1960, Interscience Pub., New York
7. ^ Anderson, T. Transactions of Royal Society of Edinburg, 16 (1849) 123
8. ^ ^{a b} Von Anderson, Th. (1849). "Producte der trocknen Destillation thierischer Materien". Annalen der Chemie und Pharmacie 70: 32. doi:10.1002/jlac.18490700105. 
9. ^ Anderson, Th. (1851). "Ueber die Producte der trocknen Destillation thierischer Materien". Annalen der Chemie und Pharmacie 80: 44. doi:10.1002/jlac.18510800104. 
10. ^ Körner, W. Giorn. academ. Palermo, vol. 5 (1869)
11. ^ Dewar, J. Chem. News, 23 (1871) 38
12. ^ Albert Ladenburg Lectures on the history of the development of chemistry since the time of Lavoisier., pp. 283–287
13. ^ Raj K. Bansal Heterocyclic Chemistry, (1999) ISBN 8122412122, p. 216
14. ^ "A. Henninger, aus Paris. 12. April 1877". Berichte der deutschen chemischen Gesellschaft 10: 727. 1877. doi:10.1002/cber.187701001202. 
15. ^ ^{a b c d e f g h i j k l} Pyridine, IARC Monogrpahs Vol. 77, OSHA, Washington D.C., 1985
16. ^ ^{a b} Tscihtschibabin, A.E. (1924). "Über Kondensation der Aldehyde mit Ammoniak zu Pyridinebasen". Journal für praktische Chemie 107: 122. doi:10.1002/prac.19241070110. http://gallica.bnf.fr/ark:/12148/bpt6k90877m/f132.chemindefer. 
17. ^ ^{a b c d e f g} S. Shimizu, N. Watanabe, T. Kataoka, T. Shoji, N. Abe, S. Morishita, H. Ichimura Pyridine and Pyridine Derivatives, in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a22_399
18. ^ G. A. Burdock (ed.) Fenaroli's Handbook of Flavor Ingredients, Vol. II, 3rd Edition, CRC Press, Boca Raton, 1995, ISBN 0-8493-2710-5
19. ^ A. Täufel, W. Ternes, L. Tunger, M. Zobel: Lebensmittel-Lexikon, 4th ed., p. 450, Behr, 2005, ISBN 3-89947-165-2
20. ^ Tang, Jian; Jin, Qi Zhang; Shen, Guo Hui; Ho, Chi Tang; Chang, Stephen S. (1983). "Isolation and identification of volatile compounds from fried chicken". Journal of Agricultural and Food Chemistry 31 (6): 1287. doi:10.1021/jf00120a035. 
21. ^ Shibamoto, Takayuki; Kamiya, Yoko; Mihara, Satoru (1981). "Isolation and identification of volatile compounds in cooked meat: sukiyaki". Journal of Agricultural and Food Chemistry 29: 57. doi:10.1021/jf00103a015. 
22. ^ Ho, Chi Tang; Lee, Ken N.; Jin, Qi Zhang (1983). "Isolation and identification of volatile flavor compounds in fried bacon". Journal of Agricultural and Food Chemistry 31 (2): 336. doi:10.1021/jf00116a038. 
23. ^ Dumont, Jean Pierre; Adda, Jacques (1978). "Occurrence of sesquiterpene in mountain cheese volatiles". Journal of Agricultural and Food Chemistry 26 (2): 364. doi:10.1021/jf60216a037. 
24. ^ Aeschbacher, HU; Wolleb, U; Löliger, J; Spadone, JC; Liardon, R (1989). "Contribution of coffee aroma constituents to the mutagenicity of coffee". Food and chemical toxicology 27 (4): 227–32. PMID 2659457. 
25. ^ Vitzthum, Otto G.; Werkhoff, Peter.; Hubert, Peter. (1975). "New volatile constituents of black tea flavored". Journal of Agricultural and Food Chemistry 23 (5): 999. doi:10.1021/jf60201a032. 
26. ^ A. Täufel, W. Ternes, L. Tunger, M. Zobel: Lebensmittel-Lexikon, 4th ed., p. 226, Behr, 2005, ISBN 3-89947-165-2
27. ^ Curvall, Margareta; Enzell, Curt R.; Pettersson, Bertil (1984). "An evaluation of the utility of four in vitro short term tests for predicting the cytotoxicity of individual compounds derived from tobacco smoke". Cell Biology and Toxicology 1 (1): 173. doi:10.1007/BF00125573. PMID 6400922. 
28. ^ Schmeltz, Irwin.; Hoffmann, Dietrich. (1977). "Nitrogen-containing compounds in tobacco and tobacco smoke". Chemical Reviews 77 (3): 295. doi:10.1021/cr60307a001. 
29. ^ Powell, W. H. (1983). "Revision of the extended Hantzsch-Widman system of nomenclature for hetero mono-cycles". Pure and Applied Chemistry 55 (2): 409–416. doi:10.1351/pac198855020409. http://www.iupac.org/publications/pac/1983/pdf/5502x0409.pdf. 
30. ^ D. Hellwinkel: Die systematische Nomenklatur der Organischen Chemie, 4th ed., p. 45, Springer, Berlin, 1998, ISBN 3-540-63221-2
31. ^ A. Gossauer: Struktur und Reaktivität der Biomoleküle, 2006, p. 488, Wiley-VCH Weinheim, ISBN 3-906390-29-2
32. ^ ^{a b} Frank, R.L.; Seven, R. P. (1949). "Pyridines. IV. A Study of the Chichibabin Synthesis". Journal of the American Chemical Society 71 (8): 2629–2635. doi:10.1021/ja01176a008. 
33. ^ ICI DE-AS Patent No 1,917,037, (1968)
34. ^ Nippon Kayaku, Japanese patent No. 7039545 (1967)
35. ^ Koei Chemicals, Patent BE 758 201 (1969)
36. ^ F. Mensch Erdöl Kohle Erdgas Petrochemie, 1969, 2, pp. 67–71
37. ^ Hantzsch, A. (1881). "Condensationsprodukte aus Aldehydammoniak und ketonartigen Verbindungen". Berichte der deutschen chemischen Gesellschaft 14 (2): 1637. doi:10.1002/cber.18810140214. 
38. ^ Knoevenagel, E.; Fries, A. (1898). "Synthesen in der Pyridinreihe. Ueber eine Erweiterung der Hantzsch'schen Dihydropyridinsynthese". Berichte der deutschen chemischen Gesellschaft 31: 761. doi:10.1002/cber.189803101157. 
39. ^ A. Behr: Angewandte homogene Katalyse, 2008, p. 722, Wiley-VCH, Weinheim, ISBN 3-527-31666-3
40. ^ Ciamician-Dennstedt Rearrangement, G. L. Ciamician, M. Dennstedt, Ber. 14, 1153 (1881); A. H. Corwin, Heterocyclic Compounds 1, 309 (1950); H. S. Mosher, ibid. 475; P. S. Skell, R. S. Sandler, J. Am. Chem. Soc. 80, 2024 (1958); E. Vogel, Angew. Chem. 72, 8 (1960).
41. ^ L. Gattermann, A. Skita (1916). "Eine Synthese von Pyridin-Derivaten". Ber. 49 (1): 494–501. doi:10.1002/cber.19160490155. 
42. ^ Gattermann-Skita, Institute of Chemistry, Skopje, Macedonia
43. ^ Tarr, J. B.; Arditti, J. (1982). "Niacin Biosynthesis in Seedlings of Zea mays". Plant Physiology 69 (3): 553. doi:10.1104/pp.69.3.553. PMC 426252. PMID 16662247. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=426252. 
44. ^ ^{a b} Lide, p. 3-448
45. ^ ^{a b} Mootz, D. (1981). "Crystal structures of pyridine and pyridine trihydrate". The Journal of Chemical Physics 75 (3): 1517. doi:10.1063/1.442204. 
46. ^ Lide, p. 3-673
47. ^ Lide, p. 5-28
48. ^ Lide, p. 6-211
49. ^ Lide, p. 6-221
50. ^ ^{a b c d} Record of Pyridine in the GESTIS Substance Database from the IFA
51. ^ V. Majer, V. Svoboda: Enthalpies of Vaporization of Organic Compounds: A Critical Review and Data Compilation, Blackwell Scientific Publications, Oxford, 1985, ISBN 0-632-01529-2
52. ^ Domalski, Eugene S.; Hearing, Elizabeth D. (1996). "Heat Capacities and Entropies of Organic Compounds in the Condensed Phase". Journal of Physical and Chemical Reference Data 25: 1. doi:10.1063/1.555985. 
53. ^ Cox, E. (1958). "Crystal Structure of Benzene". Reviews of Modern Physics 30: 159. Bibcode 1958RvMP...30..159C. doi:10.1103/RevModPhys.30.159. 
54. ^ Lide, p. 6–67
55. ^ McCullough, J. P.; Douslin, D. R.; Messerly, J. F.; Hossenlopp, I. A.; Kincheloe, T. C.; Waddington, Guy (1957). "Pyridine: Experimental and Calculated Chemical Thermodynamic Properties between 0 and 1500°K.; a Revised Vibrational Assignment". Journal of the American Chemical Society 79 (16): 4289. doi:10.1021/ja01573a014. 
56. ^ Joule, p. 14
57. ^ Joule, p. 16
58. ^ Pyridine hydrochloride MSDS, Alfa Aesar, 26 June 2010
59. ^ ^{a b c d e f g} Joule, pp. 125–141
60. ^ ^{a b c} D. T. Davies Aromatic heterocyclic chemistry, Oxford University Press, 1992, ISBN 0198556608
61. ^ R. Milcent, F. Chau: Chimie organique hétérocyclique: Structures fondamentales, pp. 241–282, EDP Sciences, 2002, ISBN 2-86883-583-X
62. ^ T. M. Krygowski, H. Szatyowicz, and J. E. Zachara (2005). "How H-bonding Modifies Molecular Structure and π-Electron Delocalization in the Ring of Pyridine/Pyridinium Derivatives Involved in H-Bond Complexation^†". J. Org. Chem. 70 (22): 8859. doi:10.1021/jo051354h. PMID 16238319. 
63. ^ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
64. ^ Nakamoto, K. (1997). Infrared and Raman spectra of Inorganic and Coordination compounds. Part A (5th ed.). Wiley. ISBN 0471 16394 5. . Part B. ISBN 0471 16392 9.  Part B, p24
65. ^ Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M.J.K.; Denney, R. C.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, p. 418, ISBN 0-582-22628-7 
66. ^ C. Elschenbroich Organometallchemie, 6th ed., pp. 524–525, Vieweg+Teubner, 2008, ISBN 3-8351-0167-6
67. ^ Joule, p. 7
68. ^ C. Elschenbroich Organometallchemie, 6th ed., p. 218, Vieweg+Teubner, 2008, ISBN 3-8351-0167-6
69. ^ Joule, p. 129
70. ^ Bakke, Jan M.; Hegbom, Ingrid; Verne, Hans Peter; Weidlein, Johann; Schnckel, Hansgeorg; Paulsen, Gudrun B.; Nielsen, Ruby I.; Olsen, Carl E. et al. (1994). "Dinitrogen Pentoxide-Sulfur Dioxide, a New nitrate ion system". Acta Chemica Scandinavica 48: 181. doi:10.3891/acta.chem.scand.48-0181. 
71. ^ Ono, Noboru; Murashima, Takashi; Nishi, Keiji; Nakamoto, Ken-Ichi; Kato, Atsushi; Tamai, Ryuji; Uno, Hidemitsu (2002). "Preparation of Novel Heteroisoindoles from nitropyridines and Nitropyridones". Heterocycles 58: 301. doi:10.3987/COM-02-S(M)22. 
72. ^ Duffy, Joseph L.; Laali, Kenneth K. (1991). "Aprotic Nitration (NO₂⁺BF₄⁻) of 2-Halo- and 2,6-Dihalopyridines and Transfer-Nitration Chemistry of Their N-Nitropyridinium Cations". The Journal of Organic Chemistry 56 (9): 3006. doi:10.1021/jo00009a015. 
73. ^ Joule, p. 126
74. ^ Gabriel, S. (1882). "Note on nicotinic acid from pyridine". Berichte der deutschen chemischen Gesellschaft 15: 834. doi:10.1002/cber.188201501180. 
75. ^ Möller, Ernst Friedrich; Birkofer, Leonhard (1942). "Konstitutionsspezifität der Nicotinsäure als Wuchsstoff bei Proteus vulgaris und Streptobacterium plantarum". Berichte der deutschen chemischen Gesellschaft (A and B Series) 75 (9): 1108. doi:10.1002/cber.19420750912. 
76. ^ Joule, p. 130
77. ^ Joule, p. 133
78. ^ Shreve, R. Norris; Riechers, E. H.; Rubenkoenig, Harry; Goodman, A. H. (1940). "Amination in the Heterocyclic Series by Sodium amide". Industrial & Engineering Chemistry 32 (2): 173. doi:10.1021/ie50362a008. 
79. ^ Badger, G; Sasse, W (1963). The action of metal catalysts on pyridines. 2. pp. 179. doi:10.1016/S0065-2725(08)60749-7. 
80. ^ Sasse, W. H. F. (1966). "2,2'-bipyridine". Organic Syntheses 46: 5–8. http://www.orgsyn.org/orgsyn/pdfs/CV5P0102.pdf. 
81. ^ Burrows, George H.; King, Louis A. (1935). "The Free Energy Change that Accompanies Hydrogenation of pyridines to piperidines". Journal of the American Chemical Society 57 (10): 1789. doi:10.1021/ja01313a011. 
82. ^ ^{a b} J.D. Cox, G. Pilcher: Thermochemistry of Organic and Organometallic Compounds, Academic Press, New York, 1970, p. 1–636, ISBN 0-12-194350-X
83. ^ Tanner, Dennis D.; Yang, Chi Ming (1993). "On the structure and mechanism of formation of the Lansbury reagent, lithium tetrakis (N-dihydropyridyl) aluminate". The Journal of Organic Chemistry 58 (7): 1840. doi:10.1021/jo00059a041. 
84. ^ De Koning, A (1980). "Specific and selective reduction of aromatic nitrogen heterocycles with the bis-pyridine complexes of bis (1,4-dihydro-1-pyridyl) zinc and bis (1,4-dihydro-1-pyridyl) magnesium". Journal of Organometallic Chemistry 199 (2): 153. doi:10.1016/S0022-328X(00)83849-8. 
85. ^ M. Ferles Collection of Czechoslovak Chemical Communications, 1959, 24, pp. 1029–1033
86. ^ Pyridine’s Development in China, AgroChemEx 2010
87. ^ About Vertellus
88. ^ Sherman, A. R. “Pyridine” in e-EROS (Encyclopedia of Reagents for Organic Synthesis) (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289X.rp280 2001.
89. ^ Corey, E.J.; Suggs, W. (1975). "Pyridinium Chlorochromate. An Efficient Reagent for Oxidation of Primary and Secondary Alcohols to Carbonyl Compounds". Tetrahedron Lett. 16 (31): 2647–2650. doi:10.1016/S0040-4039(00)75204-X. 
90. ^ Agarwal, S; Tiwari, H. P.; Sharma, J. P. (1990). "Pyridinium Chlorochromate: an Improved Method for its Synthesis and use of Anhydrous acetic acid as catalyst for oxidation reactions". Tetrahedron 46 (12): 4417–4420. doi:10.1016/S0040-4020(01)86776-4. 
91. ^ Cornforth, R.H.; Cornforth, J.W.; Popjak, G. (1962). "Preparation of R-and S-mevalonolactones". Tetrahedron 18 (12): 1351–4. doi:10.1016/S0040-4020(01)99289-0. 
92. ^ J. C. Collins, W. W. Hess and F. J. Frank (1968). "Dipyridine-chromium(VI) oxide oxidation of alcohols in dichloromethane". Tetrahedron Lett. 9 (30): 3363–3366. doi:10.1016/S0040-4039(00)89494-0. 
93. ^ J. C. Collins, W.W. Hess (1988), "Aldehydes from Primary Alcohols by Oxidation with Chromium Trioxide: Heptanal", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv6p0644 ; Coll. Vol. 6: 644 
94. ^ G. Tojo and M. Fernâandez (2006). Oxidation of alcohols to aldehydes and ketones: a guide to current common practice. New York: Springer. pp. 28, 29, 86. ISBN 0387236074. http://books.google.com/books?id=O6USLyDIBOUC&pg=PA86. 
95. ^ C. H. Bamford, C. F. H Tipper: Comprehensive Chemical Kinetics: Non-radical Polymerisation, Elsevier, Amsterdam, 1980, ISBN 0-444-41252-2
96. ^ A. V. Hopper: Recent Developments in Polymer Research, Nova Science Publisher, 2007, ISBN 1-60021-346-4
97. ^ Crabtree, Robert (1979). "Iridium compounds in catalysis". Accounts of Chemical Research 12 (9): 331. doi:10.1021/ar50141a005. 
98. ^ Environmental and health criteria for paraquat and diquat, World Health Organization, Geneva, 1984
99. ^ K. Eller, E. Henkes, R. Rossbacher, H. Hoke: Amines, Aliphatic, in: Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH Weinheim
100. ^ C. E. Terry, R. P. Ryan, S. S. Leffingwell: Toxicology Desk Reference: The Toxic Exposure & Medical Monitoring Index: The Toxic Exposure and Medical Monitoring Index, 5th ed., p. 1062, Taylor & Francis, ISBN 1-56032-795-2
101. ^ A. Täufel, W. Ternes, L. Tunger, M. Zobel: Lebensmittel-Lexikon, 4th ed., p. 218, Behr, 2005, ISBN 3-89947-165-2
102. ^ Wasserbestimmung mit Karl-Fischer-Titration, Jena University
103. ^ Database of the Environmental Protection Agency (EPA)
104. ^ Pyridine MSDS, Alfa Aesar, 3 June 2010
105. ^ Aylward, G, (2008), "SI Chemical Data 6th Ed.", ISBN 978 0 470 81638 7 (pbk.)
106. ^ ^{a b c} International Agency for Research on Cancer (IARC) (2000-08-22). "Pyridine Summary & Evaluation". IARC Summaries & Evaluations. IPCS INCHEM. http://www.inchem.org/documents/iarc/vol77/77-16.html. Retrieved 2007-01-17. 
107. ^ ^{a b} N. Bonnard, M.T. Brondeau, S. Miraval, F. Pillière, J.C. Protois, O. Schneider: Pyridine, Fiche Toxicologique, INRS (in French)
108. ^ Sims, G.K. and O'Loughlin, E.J. (1989). "Degradation of pyridines in the environment". CRC Critical Reviews in Environmental Control 19 (4): 309–340. doi:10.1080/10643388909388372. 
109. ^ Sims, G. K. and L.E. Sommers (1986). "Biodegradation of pyridine derivatives in soil suspensions". Environmental Toxicology and Chemistry 5 (6): 503–509. doi:10.1002/etc.5620050601. 
110. ^ Sims, G. K. and E.J. O'Loughlin (1992). "Riboflavin production during growth of Micrococcus luteus on pyridine". Applied and Environmental Microbiology 58 (10): 3423–3425. PMC 183117. PMID 16348793. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=183117. 
111. ^ Junk, G.A.; Ford, C.S. (1980). "A review of organic emissions from selected combustion processes". Chemosphere 9 (4): 187. doi:10.1016/0045-6535(80)90079-X. 
112. ^ Hawthorne, Steven B.; Sievers, Robert E. (1984). "Emissions of organic air pollutants from shale oil wastewaters". Environmental Science & Technology 18 (6): 483. doi:10.1021/es00124a016. 
113. ^ Stuermer, Daniel H.; Ng, Douglas J.; Morris, Clarence J. (1982). "Organic contaminants in groundwater near to underground coal gasification site in northeastern Wyoming". Environmental Science & Technology 16 (9): 582. doi:10.1021/es00103a009. 
114. ^ National Occupational Exposure Survey 1981–83, Cincinnati, OH, Department of Health and Human Services, Public Health Service, Centers for Disease Control, National Institute for Occuptional Safety and Health.

Bibliography

• Joule, J. A. and Mills, K. Heterocyclic Chemistry, 5th ed., Blackwell Publishing, Chichester, 2010, ISBN 1-4051-3300-7
• Lide, D. R. (ed.) Handbook of Chemistry and Physics, 90th edition, CRC Press, Boca Raton, 2009, ISBN 978-1-4200-9084-0

External links

• Synthesis and propierties of pyridines at chemsynthesis.com
• International Chemical Safety Card 0323
• NIOSH Pocket Guide to Chemical Hazards
• Synthesis of pyridines (overview of recent methods)