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Halogenation
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In chemistry, halogenation is a chemical reaction which introduces one or more halogens into a chemical compound. Halide-containing compounds are pervasive, making this type of transformation important, e.g. in the production of polymers, drugs.[1] This kind of conversion is in fact so common that a comprehensive overview is challenging. This article mainly deals with halogenation using elemental halogens (F2, Cl2, Br2, I2). Halides are also commonly introduced using halide salts and hydrogen halide acids. Many specialized reagents exist for introducing halogens into diverse substrates, e.g. thionyl chloride.
Organic chemistry
[edit]Several pathways exist for the halogenation of organic compounds, including free radical halogenation, ketone halogenation, electrophilic halogenation, and halogen addition reaction. The nature of the substrate determines the pathway. The facility of halogenation is influenced by the halogen. Fluorine and chlorine are more electrophilic and are more aggressive halogenating agents. Bromine is a weaker halogenating agent than both fluorine and chlorine, while iodine is the least reactive of them all. The facility of dehydrohalogenation follows the reverse trend: iodine is most easily removed from organic compounds, and organofluorine compounds are highly stable.
Free radical halogenation
[edit]Halogenation of saturated hydrocarbons is a substitution reaction. The reaction typically involves free radical pathways. The regiochemistry of the halogenation of alkanes is largely determined by the relative weakness of the C–H bonds. This trend is reflected by the faster reaction at tertiary and secondary positions.
Free radical chlorination is used for the industrial production of some solvents:[2]
- CH4 + Cl2 → CH3Cl + HCl
Naturally occurring organobromine compounds are usually produced by free radical pathway catalyzed by the enzyme bromoperoxidase. The reaction requires bromide in combination with oxygen as an oxidant. The oceans are estimated to release 1–2 million tons of bromoform and 56,000 tons[which?] of bromomethane annually.[3]
The iodoform reaction, which involves degradation of methyl ketones, proceeds by the free radical iodination.
Fluorination
[edit]Because of its extreme reactivity, fluorine (F2) represents a special category with respect to halogenation. Most organic compounds, saturated or otherwise, burn upon contact with F2, ultimately yielding carbon tetrafluoride. By contrast, the heavier halogens are far less reactive toward saturated hydrocarbons.
Highly specialised conditions and apparatus are required for fluorinations with elemental fluorine. Commonly, fluorination reagents are employed instead of F2. Such reagents include cobalt trifluoride, chlorine trifluoride, and iodine pentafluoride.[4]
The method electrochemical fluorination is used commercially for the production of perfluorinated compounds. It generates small amounts of elemental fluorine in situ from hydrogen fluoride. The method avoids the hazards of handling fluorine gas. Many commercially important organic compounds are fluorinated using this technology.
Addition of halogens to alkenes and alkynes
[edit]Unsaturated compounds, especially alkenes and alkynes, add halogens:
- R−CH=CH−R' + X2 → R−CHX−CHX−R'
In oxychlorination, the combination of hydrogen chloride and oxygen serves as the equivalent of chlorine, as illustrated by this route to 1,2-dichloroethane:
- 4 HCl + 2 CH2=CH2 + O2 → 2 Cl−CH2−CH2−Cl + 2 H2O
The addition of halogens to alkenes proceeds via intermediate halonium ions. In special cases, such intermediates have been isolated.[5]
Bromination is more selective than chlorination because the reaction is less exothermic. Illustrative of the bromination of an alkene is the route to the anesthetic halothane from trichloroethylene:[6]
Iodination and bromination can be effected by the addition of iodine and bromine to alkenes. The reaction, which conveniently proceeds with the discharge of the color of I2 and Br2, is the basis of the analytical method. The iodine number and bromine number are measures of the degree of unsaturation for fats and other organic compounds.
Halogenation of aromatic compounds
[edit]Aromatic compounds are subject to electrophilic halogenation:
- R−C6H5 + X2 → HX + R−C6H4−X
This kind of reaction typically works well for chlorine and bromine. Often a Lewis acidic catalyst is used, such as ferric chloride.[7] Many detailed procedures are available.[8][9] Because fluorine is so reactive, other methods, such as the Balz–Schiemann reaction, are used to prepare fluorinated aromatic compounds.
Other halogenation methods
[edit]In the Hunsdiecker reaction, carboxylic acids are converted to organic halide, whose carbon chain is shortened by one carbon atom with respect to the carbon chain of the particular carboxylic acid. The carboxylic acid is first converted to its silver salt, which is then oxidized with halogen:
- R−COO−Ag+ + Br2 → R−Br + CO2 + Ag+Br−
- CH3−COO−Ag+ + Br2 → CH3−Br + CO2 + Ag+Br−
Many organometallic compounds react with halogens to give the organic halide:
- RM + X2 → RX + MX
- CH3CH2CH2CH2Li + Cl2 → CH3CH2CH2CH2Cl + LiCl
Inorganic chemistry
[edit]All elements aside from argon, neon, and helium form fluorides by direct reaction with fluorine. Chlorine is slightly more selective, but still reacts with most metals and heavier nonmetals. Following the usual trend, bromine is less reactive and iodine least of all. Of the many reactions possible, illustrative is the formation of gold(III) chloride by the chlorination of gold. The chlorination of metals is usually not very important industrially since the chlorides are more easily made from the oxides and hydrogen chloride. Where chlorination of inorganic compounds is practiced on a relatively large scale is for the production of phosphorus trichloride and disulfur dichloride.[10]
See also
[edit]
Wikiquote has quotations related to Halogenation.
- Dehalogenation
- Haloalkane (Alkyl halide)
- Halogenoarene (Aryl halide)
- Free radical halogenation
- Haloketone
- Electrophilic substitution
References
[edit]- ^ Hudlicky, Milos; Hudlicky, Tomas (1983). "Formation of Carbon-Halogen Bonds". In S. Patai; Z. Rappoport (eds.). Halides, Pseudo-Halides and Azides: Part 2 (1983). PATAI's Chemistry of Functional Groups. pp. 1021–1172. doi:10.1002/9780470771723.ch3. ISBN 9780470771723.
- ^ Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a06_233.pub2. ISBN 978-3-527-30673-2.
- ^ Gribble, Gordon W. (1999). "The diversity of naturally occurring organobromine compounds". Chemical Society Reviews. 28 (5): 335–346. doi:10.1039/a900201d.
- ^ Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a11_307. ISBN 3-527-30673-0.
- ^ T. Mori; R. Rathore (1998). "X-Ray structure of bridged 2,2′-bi(adamant-2-ylidene) chloronium cation and comparison of its reactivity with a singly bonded chloroarenium cation". Chem. Commun. (8): 927–928. doi:10.1039/a709063c.
- ^ Synthesis of Essential Drugs, Ruben Vardanyan, Victor Hruby; Elsevier 2005 ISBN 0-444-52166-6
- ^ Beck, Uwe; Löser, Eckhard (2011). "Chlorinated Benzenes and Other Nucleus-Chlorinated Aromatic Hydrocarbons". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.o06_o03. ISBN 978-3527306732.
- ^ Organic chemistry by Jonathan Clayden, Nick Grieves, Stuart Warren, Oxford University Press
- ^ Edward R. Atkinson, Donald M. Murphy, and James E. Lufkin (1951). "dl-4,4′,6,6′-Tetrachlorodiphenic Acid". Organic Syntheses. 31: 96. doi:10.15227/orgsyn.031.0096
{{cite journal}}: CS1 maint: multiple names: authors list (link). - ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
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Halogenation
View on Grokipediafrom Grokipedia
Halogenation is a fundamental class of chemical reactions involving the incorporation of one or more halogen atoms—fluorine, chlorine, bromine, or iodine—into molecules, prominently in organic substrates but also in inorganic contexts, typically through substitution of other atoms or addition across multiple bonds.[1] These processes are essential for synthesizing halogenated compounds, which are versatile building blocks in chemical synthesis due to the reactivity of the carbon-halogen bond.[2]
The primary types of halogenation reactions vary by mechanism and substrate. Free-radical halogenation occurs with alkanes under light or heat, replacing a hydrogen with a halogen via a chain mechanism involving initiation, propagation, and termination steps, though it often yields mixtures due to varying hydrogen reactivity.[3] Electrophilic addition halogenation targets alkenes and alkynes, where halogens like bromine or chlorine add across the double bond to form vicinal dihalides, proceeding through a cyclic halonium ion intermediate.[4] In aromatic systems, electrophilic aromatic substitution introduces a halogen to benzene rings, requiring a Lewis acid catalyst such as FeBr₃ for bromination to generate the electrophilic halogen species and stabilize the sigma complex intermediate.[5] Additionally, alpha-halogenation of carbonyl compounds under acidic or basic conditions selectively functionalizes the carbon adjacent to the carbonyl group, exploiting enol or enolate reactivity. Inorganic halogenation includes reactions like the formation of metal halides or interhalogen compounds, covered in detail in subsequent sections.
The concept of halogenation dates back to the late 18th century, with the discovery of chlorine by Carl Wilhelm Scheele in 1774 and its properties explored by Humphry Davy, who coined the term "halogen" in 1811 for chlorine, bromine, and iodine. These reactions underpin much of organic synthesis, facilitating the creation of pharmaceuticals, agrochemicals, and advanced materials by providing precursors for cross-coupling reactions like Suzuki or Heck couplings.[2] Halogenation's versatility stems from the distinct properties of each halogen—chlorine's abundance for industrial scale-up, bromine's moderate reactivity, iodine's mild conditions, and fluorine's unique electron-withdrawing effects—making it indispensable despite challenges like regioselectivity and over-halogenation.
Handling halogens requires stringent safety measures due to their toxicity and corrosivity; fluorine gas is extremely reactive and corrosive, attacking glass and most metals, while chlorine and bromine cause severe respiratory irritation and burns upon exposure, and iodine vapors can lead to thyroid disruption. All halogens produce hazardous fumes and must be managed in well-ventilated fume hoods with appropriate protective equipment to mitigate inhalation and contact risks.[21]
with kJ/mol, followed by the methyl radical reacting with chlorine:
which is highly exothermic at kJ/mol. These steps propagate the chain efficiently, though the overall reaction is exothermic.[29] Selectivity differs markedly between halogens due to variations in the hydrogen abstraction step's energetics and bond dissociation energy (BDE) differences for C-H bonds (primary ~410 kJ/mol, secondary ~397 kJ/mol, tertiary ~381 kJ/mol). Chlorination shows low selectivity, with relative reactivities per hydrogen of 1:3.8:5 for primary:secondary:tertiary positions, as the early transition state for the exothermic abstraction by Cl• is less sensitive to radical stability. Bromination, however, exhibits high selectivity at 1:80:1600, because the endothermic abstraction by Br• has a late transition state that closely resembles the more stable tertiary radical, amplifying BDE differences./09:_Free_Radical_Substitution_Reaction_of_Alkanes/9.04:_Chlorination_vs_Bromination) Reactions are conducted in the gas phase or at elevated temperatures (300–500°C) to promote radical formation and mixing, often with controlled halogen ratios to minimize side reactions like polyhalogenation, where further substitution on RX occurs due to weakened C-H bonds alpha to the halogen.[30] A key example is the industrial chlorination of methane to chloromethane (CH₃Cl), performed at 400–500°C with UV or thermal initiation, yielding a mixture separable by distillation; this process produces precursors for dichloromethane, chloroform, and other chlorinated solvents used in chemical manufacturing.[31]
This process occurs continuously in industrial settings, with the reaction mixture serving as a solvent to control the exothermic nature.[54] Phosphorus halides exemplify the influence of reaction conditions on product formation. Phosphorus trichloride is obtained by direct chlorination, but exposure to excess chlorine converts it to phosphorus pentachloride:
Similarly, for bromides, low temperatures are employed during the reaction of red phosphorus with bromine to favor phosphorus tribromide and prevent over-halogenation to phosphorus pentabromide:
Excess halogen promotes higher oxidation states, such as +5 for phosphorus, while temperature control avoids mixtures of tri- and penta-halides.[55][56] Sulfur fluorides demonstrate halogenation leading to high coordination numbers. Sulfur reacts vigorously with fluorine to produce sulfur hexafluoride:
or, for the elemental form:
This reaction requires careful handling due to the extreme reactivity of fluorine and occurs under controlled conditions to achieve complete fluorination. Silicon, another nonmetal, forms silicon tetrachloride by direct combination with chlorine:
This synthesis typically involves heating silicon or ferrosilicon with chlorine gas.[57][58] The resulting compounds feature covalent bonding, often with expanded octets for central atoms in higher oxidation states. In SF6, sulfur exhibits hypervalency, accommodating 12 valence electrons around the central atom through d-orbital involvement or three-center four-electron bonds, exceeding the octet rule. Phosphorus pentachloride similarly displays an expanded octet in its trigonal bipyramidal structure. Oxygen, however, typically does not form stable binary halogen compounds under standard halogenation conditions, with rare exceptions like oxygen difluoride being unstable and atypical. These binary nonmetal halides contrast with interhalogen compounds, which involve reactions between different halogens.
Introduction
Definition and Scope
Halogenation is a chemical reaction that introduces one or more halogen atoms into a molecule, typically through processes such as substitution, addition, or direct combination.[1] The halogens involved are the Group 17 elements of the periodic table: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At).[6] The scope of halogenation encompasses organic chemistry, where it often forms carbon-halogen (C-X) bonds essential for synthesis; inorganic chemistry, such as the direct reaction of halogens with metals to produce metal halides; and biochemical contexts, where enzymes facilitate halogen incorporation into natural products like antibiotics.[7][6][8] This distinguishes halogenation from dehalogenation, which involves the removal of halogen atoms from molecules.[8] Key types include substitution, where a halogen replaces hydrogen or another group; addition, which occurs across unsaturated bonds; and oxidative halogenation, utilizing oxidants to generate halogenating agents from halides.[7][9] Halogenation relies on the atomic structure of halogens, which possess seven valence electrons in their outer shell, and their high electron affinity, driving the acquisition of an additional electron to achieve a stable octet configuration.[6] Reactivity trends among halogens decrease from fluorine to iodine due to increasing atomic size and weakening electronegativity.[6]Historical Background
The discovery of the halogens began in the late 18th century with the isolation of chlorine by Swedish chemist Carl Wilhelm Scheele in 1774, who produced the greenish-yellow gas by reacting hydrochloric acid with manganese(IV) oxide.[10] Scheele described its properties, including its bleaching action and corrosiveness, though he initially viewed it as a compound rather than an element. In 1810, British chemist Humphry Davy confirmed chlorine as a distinct element and named it from the Greek word chloros, meaning pale green. The term "halogens" for the group was coined in 1826 by Swedish chemist Jöns Jacob Berzelius, recognizing their shared ability to form salts.[11][10] Iodine was isolated in 1811 by French chemist Bernard Courtois during the extraction of potassium nitrate from seaweed ash for gunpowder production, where he observed violet vapors upon adding sulfuric acid; the element was characterized and named by Joseph Louis Gay-Lussac and others shortly thereafter.[12] Bromine followed in 1826, discovered by French chemist Antoine-Jérôme Balard in the bittern (residual salt liquor) from Montpellier salt marshes, where he isolated the reddish-brown liquid through chlorine displacement and ether extraction.[13] Fluorine, the most reactive halogen, was isolated in 1886 by French chemist Henri Moissan through electrolysis of potassium hydrogen fluoride in anhydrous hydrogen fluoride.[14] Astatine, the heaviest halogen, was first synthesized in 1940 by Dale R. Corson, Kenneth R. MacKenzie, and Emilio Segrè via alpha-particle bombardment of bismuth.[15] In the mid-19th century, notable advances in halogenation techniques emerged through the work of key chemists. Toward the end of the century, Belgian chemist Frédéric Swarts developed fluorination methods in the 1890s, introducing halogen exchange reactions using antimony trifluoride (SbF3) to convert chlorides and bromides into fluorides, such as the first synthesis of dichlorodifluoromethane (Freon-12), enabling access to previously challenging organofluorine compounds.[16] The 20th century marked a shift from elemental isolations to mechanistic understanding and industrial-scale halogenation. In the 1930s, the free radical mechanism for chlorination of alkanes was elucidated, with researchers like F.O. Rice demonstrating chain reactions initiated by light or heat, allowing selective substitution in hydrocarbons.[17] Fluorination saw breakthroughs during the 1940s Manhattan Project, where large-scale production of elemental fluorine was achieved for uranium enrichment, spurring safe handling techniques and synthetic routes that expanded organofluorine chemistry beyond laboratory scales.[18] Post-World War II, halogenation evolved into a cornerstone of organic synthesis, with controlled methods enabling the production of pharmaceuticals, agrochemicals, and polymers, transitioning from qualitative observations to precise, scalable processes driven by mechanistic insights.General Principles
Properties of Halogens
The halogens, comprising fluorine, chlorine, bromine, and iodine, exhibit distinct physical properties that vary systematically down Group 17 of the periodic table. Fluorine and chlorine exist as pale yellow and greenish-yellow diatomic gases at room temperature, respectively, while bromine is a volatile red-brown liquid, and iodine forms shiny black-violet crystals that sublime to a violet gas. This progression from gaseous to solid states correlates with increasing atomic size and molecular mass, leading to stronger intermolecular dispersion forces; boiling points rise from 85 K for F₂ to 457 K for I₂. Atomic radii also increase down the group due to additional electron shells, with covalent radii of 64 pm for fluorine, 99 pm for chlorine, 114 pm for bromine, and 133 pm for iodine, influencing steric effects in molecular interactions. Electronegativities follow the Pauling scale, decreasing from 3.98 for fluorine to 2.66 for iodine, reflecting the tighter hold of valence electrons by smaller, more compact atoms at the top of the group.[19][20] Chemically, halogens are highly reactive nonmetals whose oxidizing power diminishes down the group, with fluorine being the strongest oxidant capable of reacting with nearly all elements, while iodine shows the least reactivity toward common substrates. This trend stems from decreasing electronegativity and electron affinity, where chlorine exhibits the highest first electron affinity at -349 kJ/mol, slightly surpassing fluorine's -328 kJ/mol due to less electron-electron repulsion in its larger orbitals; bromine and iodine follow at -324 kJ/mol and -295 kJ/mol, respectively. Bond dissociation energies of the X-X bonds deviate from the expected trend, with F-F being anomalously weak at 159 kJ/mol owing to lone-pair repulsions in the compact fluorine atoms, compared to 243 kJ/mol for Cl-Cl, 193 kJ/mol for Br-Br, and 151 kJ/mol for I-I; this weakness facilitates homolytic cleavage in reactions. The increasing atomic size down the group enhances steric hindrance, moderating reactivity in processes involving close molecular approaches.| Halogen | Covalent Radius (pm) | Electronegativity (Pauling) | First Electron Affinity (kJ/mol) | X-X Bond Energy (kJ/mol) | Boiling Point (K) |
|---|---|---|---|---|---|
| F₂ | 64 | 3.98 | -328 | 159 | 85 |
| Cl₂ | 99 | 3.16 | -349 | 243 | 239 |
| Br₂ | 114 | 2.96 | -324 | 193 | 332 |
| I₂ | 133 | 2.66 | -295 | 151 | 457 |
Common Mechanisms
Halogenation reactions proceed through several common mechanisms depending on the substrate, halogen involved, and reaction conditions. The free radical mechanism is prevalent in the substitution of alkanes with halogens under light or heat. It consists of three stages: initiation, propagation, and termination. In the initiation step, homolytic cleavage of the halogen molecule (X₂) generates halogen radicals: , where X is Cl, Br, or I.[22] Propagation involves two key steps that sustain the chain: the halogen radical abstracts a hydrogen from the substrate (RH), forming HX and an organic radical (R•) (), followed by the organic radical reacting with X₂ to regenerate X• and produce the halogenated product (RX) ().[23] Termination occurs when radicals combine, such as two X• forming X₂ () or R• with X• forming RX, halting the chain.[22] The chain length, defined as the number of product molecules per initiation event, can reach thousands, enhancing efficiency, while selectivity favors tertiary over primary hydrogens, with bromine showing higher selectivity (relative rate ~1600:82:1 for 3°:2°:1°) than chlorine (~5:4:1) due to the stability of the transition state in propagation.[23] The electrophilic mechanism typically involves the halogen acting as an electrophile, often forming a halonium ion intermediate in additions to unsaturated systems. For example, in the addition of Br₂ to an alkene, the π electrons of the double bond attack the polarized Br δ+, forming a three-membered bromonium ion cyclic intermediate, with the bromide ion as a counterion.[24] This is followed by backside attack of the bromide on the bromonium ion, leading to anti addition stereochemistry and trans-1,2-dibromide product, as the nucleophilic attack occurs from the opposite face to avoid steric repulsion.[24] In aromatic substitutions, the mechanism similarly features electrophilic attack by X⁺ (generated via Lewis acid catalysis) to form a sigma complex (arenium ion), followed by deprotonation, though without a true halonium ring.[25] Nucleophilic mechanisms in halogenation occur when a halide ion acts as a nucleophile attacking an electron-deficient carbon center, commonly in substitution reactions of alkyl halides. These follow SN2 or SN1 pathways: in SN2, the halide (e.g., I⁻) inverts configuration at a primary or secondary carbon via backside attack on the carbon bearing a leaving group like Cl, as in the Finkelstein reaction converting chlorides to iodides.[26] For SN1, ionization forms a carbocation intermediate at tertiary or benzylic centers, followed by front-side or back-side capture by the halide, leading to racemization.[26] This mechanism is less common for initial halogen introduction but facilitates exchange of halogens. An oxidative mechanism is observed in the alpha-halogenation of carbonyl compounds, where the halogen serves as an oxidant. Under acid catalysis, the enol tautomer of the carbonyl attacks the electrophilic halogen (e.g., Br₂), forming the alpha-bromo carbonyl and HBr; the rate-determining step is enol formation, independent of halogen concentration.[27] The halogen's oxidizing role stabilizes the enol's nucleophilic attack, and multiple halogenations can occur due to increased acidity of remaining alpha hydrogens.[27] Several factors influence these mechanisms, including solvent effects and catalysts. Polar aprotic solvents enhance nucleophilic attacks in SN2 by solvating cations without hindering anions, while protic solvents stabilize ions in SN1.[26] In electrophilic additions, inert nonpolar solvents like CCl₄ prevent side reactions by avoiding nucleophilic interference with the halogen.[28] Catalysts such as Lewis acids (e.g., FeBr₃ for aromatic bromination) polarize X₂ to generate X⁺, lowering the activation energy, while acid catalysts (e.g., H⁺ for enol formation) promote tautomerization in oxidative halogenation.[25]Organic Halogenation
Free Radical Halogenation
Free radical halogenation involves the substitution of a hydrogen atom in an aliphatic hydrocarbon (RH) with a halogen atom (X), producing an alkyl halide (RX) and hydrogen halide (HX), where X is typically chlorine or bromine. This reaction requires initiation by ultraviolet light or heat to generate halogen radicals./09:_Free_Radical_Substitution_Reaction_of_Alkanes/9.01:_Free_Radical_Halogenation_of_Alkanes) The process follows a chain mechanism with initiation, propagation, and termination steps. In the key propagation steps for chlorination of methane, the chlorine radical abstracts a hydrogen atom:with kJ/mol, followed by the methyl radical reacting with chlorine:
which is highly exothermic at kJ/mol. These steps propagate the chain efficiently, though the overall reaction is exothermic.[29] Selectivity differs markedly between halogens due to variations in the hydrogen abstraction step's energetics and bond dissociation energy (BDE) differences for C-H bonds (primary ~410 kJ/mol, secondary ~397 kJ/mol, tertiary ~381 kJ/mol). Chlorination shows low selectivity, with relative reactivities per hydrogen of 1:3.8:5 for primary:secondary:tertiary positions, as the early transition state for the exothermic abstraction by Cl• is less sensitive to radical stability. Bromination, however, exhibits high selectivity at 1:80:1600, because the endothermic abstraction by Br• has a late transition state that closely resembles the more stable tertiary radical, amplifying BDE differences./09:_Free_Radical_Substitution_Reaction_of_Alkanes/9.04:_Chlorination_vs_Bromination) Reactions are conducted in the gas phase or at elevated temperatures (300–500°C) to promote radical formation and mixing, often with controlled halogen ratios to minimize side reactions like polyhalogenation, where further substitution on RX occurs due to weakened C-H bonds alpha to the halogen.[30] A key example is the industrial chlorination of methane to chloromethane (CH₃Cl), performed at 400–500°C with UV or thermal initiation, yielding a mixture separable by distillation; this process produces precursors for dichloromethane, chloroform, and other chlorinated solvents used in chemical manufacturing.[31]
Addition to Alkenes and Alkynes
Halogenation of alkenes involves the electrophilic addition of molecular halogens (X₂, where X = Cl, Br, or I) across the carbon-carbon double bond, yielding vicinal dihalides as the primary products.[32] The reaction proceeds via a two-step mechanism: first, the π-electrons of the alkene attack the electrophilic halogen, forming a cyclic halonium ion intermediate; second, a halide ion (X⁻) acts as a nucleophile, attacking the more substituted carbon of the halonium ion from the opposite face. This process can be represented by the general equation: where the product is a 1,2-dihalide.[32] For bromine addition to cyclohexene, the reaction exemplifies the stereospecificity, producing trans-1,2-dibromocyclohexane due to the anti addition facilitated by the bridged bromonium ion intermediate.[33] The stereochemistry of halogen addition to alkenes is predominantly trans, arising from the backside attack on the halonium ion, which prevents syn addition and leads to racemic mixtures or meso compounds depending on the alkene's symmetry.[34] In unsymmetric alkenes, regioselectivity is observed particularly when the reaction occurs in nucleophilic solvents; for instance, in water, halohydrins form instead of dihalides, with the halogen attaching to the less substituted carbon and the hydroxyl group to the more substituted carbon, following a Markovnikov-like orientation for the nucleophile.[32] An example is the addition of Br₂ in H₂O to propene, yielding 1-bromo-2-propanol as the major product.[33] Reaction conditions typically involve non-polar solvents such as dichloromethane or carbon tetrachloride to favor dihalide formation and prevent side reactions with protic solvents; chlorine and bromine additions are rapid and irreversible under these conditions, whereas iodine addition is reversible and equilibrium-limited, rendering it less practical for synthesis.[32] For alkynes, halogenation occurs stepwise due to the presence of two π-bonds, initially forming vinyl dihalides and, with excess halogen, tetrahalides.[35] The mechanism mirrors that of alkenes, involving a halonium ion intermediate for each addition, with the first step yielding a trans-vinyl dihalide.[35] Acetylene provides a representative example: These additions are typically conducted in non-aqueous media to control the extent of halogenation, and the process exhibits anti stereochemistry in the vinyl halide intermediate.[36]Electrophilic Aromatic Halogenation
Electrophilic aromatic halogenation is a substitution reaction in which a hydrogen atom on an aromatic ring is replaced by a halogen atom, typically chlorine, bromine, or iodine, under the influence of a Lewis acid catalyst. The general reaction is represented as Ar-H + X₂ → Ar-X + HX, where Ar denotes an aromatic system and X is the halogen. This process requires a Lewis acid such as FeX₃ (e.g., FeBr₃ for bromination) to generate the electrophilic halogen species, as halogens alone are not sufficiently reactive toward the electron-rich aromatic ring. Unlike aliphatic halogenation, which often proceeds via free radical mechanisms, aromatic halogenation preserves the aromatic π-system through substitution rather than addition.[37] The mechanism involves three key steps: generation of the electrophile, formation of the sigma complex, and rearomatization. First, the halogen molecule (X₂) coordinates with the Lewis acid (MXₙ), polarizing the X-X bond to produce an electrophilic species such as [X-MXₙ]⁺ and X⁻ (approximating X⁺). The aromatic ring then attacks this electrophile, forming a resonance-stabilized sigma complex (also known as the Wheland intermediate or arenium ion), which is the rate-determining step due to the loss of aromaticity. Finally, a base (typically X⁻) abstracts a proton from the sigma complex, restoring aromaticity and yielding the halogenated product. This mechanism was first proposed by G. W. Wheland in 1942. A representative example is the chlorination of benzene: C₆H₆ + Cl₂ (FeCl₃) → C₆H₅Cl + HCl.[38][39] Substituent groups on the aromatic ring exert directing effects that influence both the rate and regioselectivity of halogenation. Electron-donating groups, such as -OH, act as strong activators and ortho-para directors by increasing electron density at the ortho and para positions through resonance, stabilizing the sigma complex at those sites. In contrast, electron-withdrawing groups like -NO₂ are deactivators and meta directors, as they destabilize the sigma complex at ortho and para positions via inductive withdrawal of electrons, making the meta position relatively more favorable. Halogens themselves are unique: they are ortho-para directors due to resonance donation from their lone pairs but overall deactivators because of strong inductive electron withdrawal, resulting in slower reactions compared to unsubstituted benzene.[40][41] Polyhalogenation, where multiple hydrogens are replaced, can occur if the ring becomes activated or conditions are not controlled, but it is typically minimized to favor monohalogenation. This is achieved by using an excess of the aromatic substrate relative to the halogen, diluting the electrophile concentration, or employing mild conditions such as limited catalyst amounts and lower temperatures to prevent over-substitution.[37]Other Organic Methods
Nucleophilic substitution plays a central role in the alpha-halogenation of carbonyl compounds, where the enol or enolate tautomer acts as the nucleophile toward electrophilic halogens. In acid-catalyzed conditions, the enol form of an aldehyde or ketone reacts with X (X = Cl, Br, or I) to form the alpha-halo carbonyl, with the reaction proceeding via addition-elimination at the enol double bond; for instance, acetophenone undergoes bromination at the alpha position in acetic acid to yield phenacyl bromide in high yield. Base-catalyzed variants generate the enolate ion, which attacks X directly, enabling polyhalogenation under controlled conditions, as seen in the preparation of alpha, alpha-dibromo ketones from methyl ketones. This method's regioselectivity stems from the thermodynamic stability of the enol or enolate at the alpha carbon, making it essential for synthesizing intermediates in pharmaceuticals and natural products.[42][43] The Finkelstein reaction provides an efficient halogen exchange route for converting alkyl chlorides or bromides to iodides through nucleophilic substitution. Typically, an alkyl chloride (RCl) is treated with sodium iodide (NaI) in acetone, where the SN2 mechanism favors iodide as the nucleophile due to its polarizability and the insolubility of NaCl, driving the equilibrium forward; benzyl chloride, for example, yields benzyl iodide in over 90% yield under reflux. This reaction is particularly valuable for primary alkyl halides, avoiding elimination side products common in other methods, and serves as a key step in preparing iodides for cross-coupling reactions like the Heck or Sonogashira. Conversion of alcohols to alkyl halides often employs thionyl chloride (SOCl) for chlorides or phosphorus tribromide (PBr) for bromides, transforming the poor leaving group OH into a suitable halide. With SOCl, the alcohol oxygen nucleophilically attacks sulfur to form a chlorosulfite intermediate (RO-SOCl), which decomposes via chloride attack, yielding RCl, SO, and HCl; secondary alcohols like 2-propanol produce isopropyl chloride with retention of configuration in non-polar solvents due to an internal return (SNi) mechanism. PBr similarly coordinates to the alcohol, forming a phosphonium intermediate that undergoes bromide displacement with inversion for primary and secondary substrates, as in the transformation of 1-octanol to 1-bromooctane, producing HOPBr as byproduct. These reagents minimize carbocation rearrangements compared to HX acids, enhancing stereochemical control in synthesis.[44] Oxidative halogenation utilizes hypohalites like sodium hypochlorite (NaOCl) to generate electrophilic halogen species in situ for selective introduction into organic substrates. In aqueous or biphasic media, NaOCl oxidizes enolizable carbonyls or activated aromatics, with the hypochlorite acting as both oxidant and halogen source; for example, cyclohexanone undergoes alpha-chlorination to 2-chlorocyclohexanone in moderate yields under mild conditions. This approach is advantageous for large-scale processes due to the availability and low cost of bleach solutions, though it requires pH control to prevent over-oxidation. Hypohalites also facilitate oxidative cleavage in haloform-like reactions, converting methyl ketones to carboxylic acids and haloform.[45] Contemporary advancements incorporate phase-transfer catalysis (PTC) and microwave irradiation to optimize halogenation efficiency and sustainability. PTC employs quaternary ammonium salts to shuttle halide ions across immiscible phases, accelerating nucleophilic substitutions like alpha-halogenation of ketones with KBr and Oxone, achieving up to 95% yield in water-organic systems without phase incompatibility issues. Chiral PTC variants enable asymmetric alpha-fluorination or chlorination of carbonyls, with enantioselectivities exceeding 90% ee using cinchona alkaloid catalysts. Microwave assistance further expedites these processes by rapid heating, as in the bromination of acetophenone with N-bromosuccinimide, completing in 5 minutes at 100°C with 98% yield, reducing energy consumption and solvent volume compared to conventional heating. These techniques align with green chemistry principles, minimizing waste in industrial applications.[46][47]Inorganic Halogenation
Halogenation of Metals
Halogenation of metals refers to the direct reaction between elemental metals and diatomic halogens (X₂, where X = F, Cl, Br, or I) to produce metal halides. These reactions generally follow the stoichiometry 2M + X₂ → 2MX, though variations occur depending on the metal's oxidation state and the halogen involved. For highly electropositive metals such as alkali metals, the process is highly exothermic, often proceeding vigorously upon ignition and yielding ionic halides like sodium chloride from the reaction of sodium with chlorine gas: 2Na + Cl₂ → 2NaCl.[48][48] The nature of the resulting metal halide—ionic or covalent—depends on the metal's position in the periodic table. Electropositive metals, including alkali metals (e.g., Na, K) and alkaline earth metals (e.g., Ca, Mg), form predominantly ionic halides with the general formulas MX or MX₂, where the metal achieves its group oxidation state. For instance, calcium reacts with chlorine to yield calcium chloride (CaCl₂), an ionic compound. In contrast, transition metals often produce covalent halides due to their intermediate electronegativity and variable oxidation states. Titanium, for example, reacts directly with chlorine to form titanium(IV) chloride (TiCl₄), a covalent liquid: Ti + 2Cl₂ → TiCl₄. This distinction arises because transition metal halides frequently involve d-orbital participation, leading to more covalent bonding compared to the predominantly electrostatic interactions in halides of s-block metals.[49][50] Reaction conditions vary with metal reactivity. Highly reactive alkali metals ignite spontaneously in halogen gases, requiring controlled environments to manage the exothermic release of energy. Less reactive metals, such as iron, necessitate elevated temperatures; for example, iron reacts with chlorine at around 300°C to form iron(III) chloride: 2Fe + 3Cl₂ → 2FeCl₃. Specific examples include the preparation of silver(I) fluoride by passing fluorine gas over silver metal: 2Ag + F₂ → 2AgF, a reaction that proceeds under controlled conditions due to fluorine's high reactivity. Industrially, magnesium is chlorinated to produce magnesium chloride (MgCl₂) for electrolytic magnesium production, often involving direct combination in high-temperature processes. Transition metals like copper exhibit variable stoichiometry based on oxidation states; copper(II) chloride forms upon heating copper with chlorine at 400–500°C: Cu + Cl₂ → CuCl₂, reflecting copper's ability to adopt +1 or +2 states in halides.[51][52][53]Halogenation of Nonmetals
Halogenation of nonmetals involves the direct reaction of elemental nonmetals, excluding carbon, with diatomic halogens to form binary covalent compounds of the general form NX_y, where N is the nonmetal and X is the halogen. These reactions are typically exothermic and proceed via free radical or direct combination mechanisms, depending on the elements involved. For instance, white phosphorus reacts with chlorine to yield phosphorus trichloride:This process occurs continuously in industrial settings, with the reaction mixture serving as a solvent to control the exothermic nature.[54] Phosphorus halides exemplify the influence of reaction conditions on product formation. Phosphorus trichloride is obtained by direct chlorination, but exposure to excess chlorine converts it to phosphorus pentachloride:
Similarly, for bromides, low temperatures are employed during the reaction of red phosphorus with bromine to favor phosphorus tribromide and prevent over-halogenation to phosphorus pentabromide:
Excess halogen promotes higher oxidation states, such as +5 for phosphorus, while temperature control avoids mixtures of tri- and penta-halides.[55][56] Sulfur fluorides demonstrate halogenation leading to high coordination numbers. Sulfur reacts vigorously with fluorine to produce sulfur hexafluoride:
or, for the elemental form:
This reaction requires careful handling due to the extreme reactivity of fluorine and occurs under controlled conditions to achieve complete fluorination. Silicon, another nonmetal, forms silicon tetrachloride by direct combination with chlorine:
This synthesis typically involves heating silicon or ferrosilicon with chlorine gas.[57][58] The resulting compounds feature covalent bonding, often with expanded octets for central atoms in higher oxidation states. In SF6, sulfur exhibits hypervalency, accommodating 12 valence electrons around the central atom through d-orbital involvement or three-center four-electron bonds, exceeding the octet rule. Phosphorus pentachloride similarly displays an expanded octet in its trigonal bipyramidal structure. Oxygen, however, typically does not form stable binary halogen compounds under standard halogenation conditions, with rare exceptions like oxygen difluoride being unstable and atypical. These binary nonmetal halides contrast with interhalogen compounds, which involve reactions between different halogens.