Chloral
Chloral, also known as trichloroacetaldehyde, is an organic compound with the molecular formula C₂HCl₃O and a molecular weight of 147.39 g/mol.[1] It is a colorless, oily, hygroscopic liquid with a pungent odor, a boiling point of 97.8°C, a melting point of -57.5°C, and a density of 1.51 g/cm³ at 20°C.[1] Soluble in water, ethanol, chloroform, and diethyl ether, chloral readily reacts with water to form chloral hydrate (C₂H₃Cl₃O₂), a geminal diol that exists in equilibrium with the aldehyde.[1] This reactivity also leads to polymerization under exposure to light or sulfuric acid, forming metachloral.[1] First synthesized in 1832 by German chemist Justus von Liebig via chlorination of ethanol, chloral marked an early achievement in organic synthesis.[2] Commercial production began in the 1940s, primarily through chlorination of acetaldehyde or ethanol using catalysts like antimony trichloride, with global output peaking at around 40,000 tonnes annually in the United States by 1963. Its industrial significance grew during World War II due to demand for insecticides, but production declined sharply after the 1972 U.S. ban on DDT, reducing U.S. output to negligible levels by the 1990s. The primary use of chloral has been as an intermediate in the manufacture of pesticides, including dichlorodiphenyltrichloroethane (DDT), methoxychlor, naled, trichlorfon, and dichlorvos, accounting for about 40% of U.S. consumption in 1975.[2] It also serves in producing trichloroacetic acid (a herbicide and protein precipitant), polyurethane foams, and pharmaceuticals such as chloral hydrate (a historical sedative), chloral betaine, α-chloralose (a bird repellent), and triclofos sodium. Additionally, chloral induces swelling in starch granules for industrial applications and occurs as a disinfection byproduct in chlorinated drinking water, with concentrations typically ranging from 0.01 to 28 µg/L in U.S. supplies.[1] Chloral is toxic by ingestion and inhalation, causing irritation to the skin, eyes, and respiratory tract, with an oral LD50 of 480 mg/kg in rats (for the hydrate form).[1] It has low bioconcentration potential and is classified by the International Agency for Research on Cancer (IARC) as Group 2A (probably carcinogenic to humans) as of 2014.[3] Occupational exposure limits are limited, with Russia setting a short-term workplace threshold of 5 mg/m³ and an ambient air standard of 0.01 mg/m³; the U.S. EPA has proposed a drinking water limit of 0.04 mg/L. Today, chloral is produced by a small number of companies in countries like China and India, reflecting its reduced but ongoing role in specialized chemical synthesis.[2]Properties
Molecular structure
Chloral has the molecular formula C₂HCl₃O and is systematically named 2,2,2-trichloroacetaldehyde or trichloroethanal.[1] The molecular structure consists of an aldehyde functional group (-CHO) in which the carbonyl carbon is double-bonded to an oxygen atom and single-bonded to a hydrogen atom and to a trichloromethyl group (CCl₃-), represented as Cl₃C–CHO.[1] This arrangement positions the three chlorine atoms on the α-carbon adjacent to the carbonyl, creating a linear chain with two carbons total.[1] The C–Cl bonds exhibit significant polarity due to the electronegativity difference between carbon and chlorine (electronegativities of 2.55 and 3.16, respectively), rendering the trichloromethyl group electron-withdrawing and enhancing the electrophilic character of the carbonyl carbon.[1] The overall molecule is polar, with a topological polar surface area of 17.1 Ų, primarily from the polar aldehyde and C–Cl bonds.[1] Chloral possesses no stable isomers, existing solely in this constitutional form as a chlorinated derivative of acetaldehyde (CH₃CHO), where the three methyl hydrogens are replaced by chlorines.[1]Physical properties
Chloral appears as a colorless oily liquid at room temperature.[4] It has a melting point of −57.5 °C and a boiling point of 97.8 °C at 760 mmHg.[5] The density is 1.51 g/cm³ at 20 °C.[4] Chloral is miscible with organic solvents such as ethanol, ether, and chloroform.[4] It reacts with water to form chloral hydrate, which is highly soluble; the anhydrous form has limited intrinsic solubility, attributable to the polar carbonyl group contrasted with the hydrophobic trichloromethyl moiety.[4] Chloral exhibits a pungent, irritating odor.[4] It is volatile, with a vapor pressure of 35 mmHg at 20 °C, and remains stable under ambient conditions when maintained in anhydrous form.[4] In infrared spectroscopy, the carbonyl group shows absorption around 1740 cm⁻¹.[6] The aldehyde proton in ¹H NMR spectroscopy appears at approximately 9.4 ppm.History
Discovery
Chloral was first synthesized in 1832 by the German chemist Justus von Liebig at the University of Giessen through the chlorination of anhydrous ethanol with dry chlorine gas.[2] This reaction produced chloral as an intermediate, which Liebig isolated during experiments aimed at understanding the halogenation of alcohols.[7] Liebig detailed the process in a publication in the inaugural issue of Annalen der Pharmacie, noting that the chlorination proceeded at varying temperatures to yield the compound.[8] Liebig characterized chloral as a colorless, oily, hygroscopic liquid with a pungent, irritating odor, which was volatile and tended to form a hydrate upon exposure to moisture. Initially, he identified it as a distinct substance related to chlorinated alcohols but distinct from chloroform, though early samples were sometimes confounded with the latter due to overlapping reaction products.[9] This discovery occurred amid the burgeoning field of organic chemistry in the early 19th century, where chemists like Liebig were pioneering systematic studies of halogenated organic compounds to elucidate reaction mechanisms and synthetic pathways.[10] Liebig's work on chloral exemplified the era's emphasis on empirical experimentation and structural elucidation, contributing to foundational advances in synthetic organic chemistry.[11]Development and early recognition
Following the initial synthesis of chloral in 1832 by Justus von Liebig through the chlorination of ethanol, mid-19th-century chemists advanced its purification and characterization. In 1847, Adolph Staedeler reported detailed studies on chloral, including methods for obtaining purer forms by distillation and reaction analysis, which helped distinguish it from impurities in early preparations.[12] These efforts built on Liebig's work and facilitated a clearer understanding of chloral as trichloroacetaldehyde, separate from related chlorinated compounds.[13] A pivotal advancement came in 1869 when German pharmacologist Oscar Liebreich explored the hydrate form of chloral (chloral hydrate) for potential medical applications. Liebreich's experiments on animals demonstrated that subcutaneous or oral administration induced deep sleep without the analgesic effects of chloroform, attributing the hypnotic action initially to in vivo decomposition into chloroform.[14] His monograph, Das Chloralhydrat: Ein neues Hypnoticum und Anaestheticum, published that year, detailed these findings and advocated its use as a sedative for anxiety and insomnia.[15] This work marked the shift from chloral as a chemical curiosity to a recognized pharmaceutical agent, with early human trials confirming its efficacy in inducing sleep.[16] By the early 1870s, chloral hydrate gained widespread medical adoption in Germany as the first synthetic hypnotic, appearing in pharmacopoeias and clinical practice for sedation. Pharmacological studies published around 1870 further documented its rapid onset and short duration, solidifying its role in treating sleep disorders and nervous conditions.[17] However, early scientific debates arose over its mechanism, stemming from confusion with chloroform due to structural similarities and shared chlorine content; Liebreich's hypothesis of metabolic conversion was later challenged in the 1870s. Notably, in 1875, physiologist Claude Bernard demonstrated through experiments that no chloroform is produced in the body upon administration of chloral hydrate. The exact mechanism of action, involving reduction to the active metabolite trichloroethanol independent of chloroform formation, was elucidated later in 1948.[18]Production
Industrial methods
The primary industrial process for chloral production involves the chlorination of ethanol or acetaldehyde under acidic conditions to achieve stepwise substitution of hydrogen atoms with chlorine.[2] This method ensures commercial viability by controlling side reactions through a gradual temperature increase, typically starting at 0 °C and ramping up to 90 °C, which optimizes yield and minimizes over-chlorination.[2] The reaction proceeds in hydrochloric acid solution, where antimony trichloride may serve as an additional catalyst to enhance reaction rates.[19] A simplified overview of the process using ethanol as feedstock is:\ce{CH3CH2OH + 4Cl2 -> Cl3CCHO + 5HCl}
This stepwise chlorination first forms intermediates like chloral alcoholate before yielding trichloroacetaldehyde (chloral).[2] Hydrochloric acid acts as both solvent and catalyst, facilitating the substitution while the temperature ramp—often in stages from 0–30 °C, 50–60 °C, to 80–90 °C—prevents excessive heat that could lead to decomposition or unwanted byproducts.[20] Following chlorination, the mixture undergoes distillation to isolate chloral, separating it from byproducts such as chloroform formed via minor haloform reaction pathways.[19] This process was developed in the late 19th century to meet growing pharmaceutical demand for chloral hydrate, the hydrated form of chloral used as a sedative.[2] Industrial-scale operations, often batch or semicontinuous, emerged around the 1870s–1890s as chloral hydrate gained medical prominence, with production scaled via large reactors handling chlorine gas feeds efficiently.[19] Today, global output remains limited and closely tied to chloral hydrate needs in pharmaceuticals and niche applications, with major producers in regions like China, Europe, and Japan maintaining capacities on the order of thousands of tonnes annually.[2] Environmental considerations in chloral production center on the handling of chlorine gas, a hazardous reactant sourced from the chlor-alkali industry, and the generation of hydrochloric acid as a coproduct, which requires neutralization or recycling to manage waste streams.[2] Modern facilities incorporate scrubbers and distillation recovery to mitigate emissions, though the process's reliance on chlorine contributes to overall chlorine consumption in the chemical sector.[19]