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Calcium hypochlorite
View on Wikipediafrom Wikipedia
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| Identifiers | |
3D model (JSmol)
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| ChEBI | |
| ChEMBL | |
| ChemSpider | |
| ECHA InfoCard | 100.029.007 |
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| UNII | |
| UN number | 1748 2208 |
CompTox Dashboard (EPA)
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| Properties | |
| Ca(OCl)2 | |
| Molar mass | 142.98 g·mol−1 |
| Appearance | white/gray powder |
| Density | 2.35 g/cm3 (20 °C) |
| Melting point | 100 °C (212 °F; 373 K) |
| Boiling point | 175 °C (347 °F; 448 K) decomposes |
| 21 g/(100 mL) at 25 °C | |
| Solubility | reacts in alcohol |
| Hazards | |
| GHS labelling: | |
   
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| Danger | |
| H272, H302, H314, H400 | |
| P210, P220, P221, P260, P264, P270, P273, P280, P301+P312, P301+P330+P331, P303+P361+P353, P304+P340, P305+P351+P338, P310, P321, P330, P363, P370+P378, P391, P405, P501 | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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850 mg/kg (oral, rat) |
| Safety data sheet (SDS) | ICSC 0638 |
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Calcium hypochlorite is an inorganic compound with chemical formula Ca(ClO)2, also written as Ca(OCl)2. It is a white solid, although commercial samples appear yellow. It strongly smells of chlorine, owing to its slow decomposition in moist air. This compound is relatively stable as a solid and solution and has greater available chlorine than sodium hypochlorite.[1] "Pure" samples have 99.2% active chlorine. Given common industrial purity, an active chlorine content of 65-70% is typical.[2] It is the main active ingredient of commercial products called bleaching powder,[a] used for water treatment and as a bleaching agent.[3]
History
[edit]Charles Tennant and Charles Macintosh developed an industrial process in the late 18th century for the manufacture of chloride of lime, patenting it in 1799.[4] Tennant's process is essentially still used today,[4][3] and became of military importance during World War I, because calcium hypochlorite was the active ingredient in trench disinfectant.[4]
Uses
[edit]Sanitation
[edit]Calcium hypochlorite is commonly used to sanitize public swimming pools and disinfect drinking water. Generally the commercial substances are sold with a purity of 65% to 73% with other chemicals present, such as calcium chloride and calcium carbonate, resulting from the manufacturing process. In solution, calcium hypochlorite could be used as a general purpose sanitizer,[5] but due to calcium residue (making the water harder), sodium hypochlorite (bleach) is usually preferred.
Organic chemistry
[edit]Calcium hypochlorite is a general oxidizing agent and therefore finds some use in organic chemistry.[6] For instance the compound is used to cleave glycols, α-hydroxy carboxylic acids and keto acids to yield fragmented aldehydes or carboxylic acids.[7] Calcium hypochlorite can also be used in the haloform reaction to manufacture chloroform.[8] Calcium hypochlorite can be used to oxidize thiol and sulfide byproducts in organic synthesis and thereby reduce their odour and make them safe to dispose of.[9] The reagent used in organic chemistry is similar to the sanitizer at ~70% purity.[10]
Production
[edit]Calcium hypochlorite is produced industrially by reaction of moist slaked calcium hydroxide with chlorine gas. The one-step reaction is shown below:[3]
- 2 Cl2 + 2 Ca(OH)2 → CaCl2 + Ca(OCl)2 + 2 H2O
Industrial setups allow for the reaction to be conducted in stages to give various compositions, each producing different ratios of calcium hypochlorite, unconverted lime, and calcium chloride.[3] In one process, the chloride-rich first stage water is discarded, while the solid precipitate is dissolved in a mixture of water and lye for another round of chlorination to reach the target purity.[2] Commercial calcium hypochlorite consists of anhydrous Ca(OCl)2, dibasic calcium hypochlorite Ca3(OCl)2(OH)4 (also written as Ca(OCl)2·2Ca(OH)2), and dibasic calcium chloride Ca3Cl2(OH)4 (also written as CaCl2·2Ca(OH)2).[11][12]
Reactions
[edit]Calcium hypochlorite reacts rapidly with acids producing calcium chloride, chlorine gas, and water:[citation needed]
- Ca(ClO)2 + 4 HCl → CaCl2 + 2 Cl2 + 2 H2O
Safety
[edit]It is a strong oxidizing agent, as it contains a hypochlorite ion at the valence +1 (redox state: Cl+1).[citation needed]
Calcium hypochlorite should not be stored wet and hot, or near any acid, organic materials, or metals. The unhydrated form is safer to handle.[citation needed]
See also
[edit]References
[edit]- ^ also chlorine powder, chloride of lime, chlorinated lime, "dry chlorine"
- ^ Gerald F. Connell. "Key operating strategies for chlorine disinfection operating systems" (PDF). Retrieved 19 October 2014.
- ^ a b "Calcium Hypochlorite - 3V Tech". www.3v-tech.com.
- ^ a b c d Vogt, H.; Balej, J; Bennett, J. E.; Wintzer, P.; Sheikh, S. A.; Gallone, P.; Vasudevan, S.; Pelin, K. (2010). "Chlorine Oxides and Chlorine Oxygen Acids". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a06_483.pub2. ISBN 978-3527306732. S2CID 96905077.
- ^ a b c "Calcium hypochlorite". Chemistry World.
- ^ Chemical Products Synopsis: Calcium Hypochlorite (Technical report). Asbuiy Park, NJ: Mannsvile Chemical Products. 1987.
- ^ Nwaukwa, Stephen; Keehn, Philip (1982). "The oxidation of aldehydes to acids with calcium hypochlorite [Ca(ClO)2]". Tetrahedron Letters. 23 (31): 3131–3134. doi:10.1016/S0040-4039(00)88577-9.
- ^ Nwaukwa, Stephen; Keehn, Philip (1982). "Oxidative cleavage of α-diols, α-diones, α-hydroxy-ketones and α-hydroxy- and α-keto acids with calcium hypochlorite [Ca(ClO)2]". Tetrahedron Letters. 23 (31): 3135–3138. doi:10.1016/S0040-4039(00)88578-0.
- ^ Cohen, Julius (1900). Practical Organic Chemistry for Advanced Students. New York: Macmillan & Co. p. 63.
- ^ National Research Council (1995). Prudent Practices in the Laboratory: Handling and Disposal of Chemicals. Washington, DC: The National Academies Press. p. 161. doi:10.17226/4911. ISBN 978-0-309-05229-0.
- ^ "8.41799 Calcium hypochlorite for synthesis". Sigma-Aldrich.
Assay (iodometric): 67.0 - 75.0 %
- ^ W.L Smith, Inorganic Bleaches, Production of Hypochlorite in Handbook of Detergents,Part F, (2009) Ed. U Zoller and Paul Sosis, CRC Press, ISBN 978-0-8247-0349-3
- ^ Aleksandrova, M.M.; Dmitriev, G.A.; Avojan, R.L. (1968). "The probable model of the crystal structure of the twobase calcium hypochlorite". Armyanskii Khimicheskii Zhurnal. 21: 380-386.
{{cite journal}}: CS1 maint: multiple names: authors list (link)
External links
[edit]Calcium hypochlorite
View on Grokipediafrom Grokipedia
History
Early discovery
The initial observations of hypochlorite compounds emerged in the late 18th century amid investigations into chlorine's chemical properties. French chemist Claude-Louis Berthollet, who had been studying chlorine since the mid-1780s, first synthesized a hypochlorite solution in 1785 by passing chlorine gas over a solution of potash (potassium carbonate), producing what became known as Javel water or potassium hypochlorite.[7] This marked the earliest recognition of hypochlorites as bleaching agents, with Berthollet publishing detailed accounts of chlorine's reactions with alkalies around 1789, laying the groundwork for subsequent hypochlorite derivatives.[8] Building on Berthollet's findings, early synthesis attempts focused on reacting chlorine gas with calcium-based compounds to produce a more stable solid form. Scottish chemist Charles Tennant conducted experiments in the late 1790s, initially proposing in 1798 a liquid solution of calcium hypochlorite obtained by treating slaked lime (calcium hydroxide) with chlorine gas, as an alternative to liquid bleaches.[9] This approach addressed the instability of earlier hypochlorite solutions and highlighted calcium's role in forming a dry, transportable bleaching powder. Tennant's work culminated in a pivotal 1799 patent for the industrial preparation of "chloride of lime," a mixture primarily consisting of calcium hypochlorite, achieved by absorbing chlorine into dry slaked lime.[10] This patent not only confirmed calcium hypochlorite as a distinct compound through empirical testing of its bleaching efficacy but also established the foundational chemical process—chlorination of calcium hydroxide—that would define its identity.[9] These early experiments underscored the compound's potential beyond liquid forms, setting the stage for broader applications.Commercial and military development
The transition of calcium hypochlorite from a laboratory compound to a cornerstone of industrial chemistry occurred in the late 18th century through the efforts of Scottish chemists Charles Tennant and Charles Macintosh. Macintosh contributed significantly to refining the dry production process originally pioneered by Tennant, enabling the creation of bleaching powder—a stable, powdered form of calcium hypochlorite mixed with calcium chloride and hydroxide—by reacting chlorine gas with dry slaked lime. This innovation addressed the limitations of earlier liquid bleach solutions, which were unstable and difficult to transport. Tennant secured a patent for the process in 1799, marking the birth of commercial-scale manufacturing.[9][11] Commercialization accelerated with the establishment of the St. Rollox chemical works near Glasgow in 1800 by Tennant, Macintosh, and partners, which rapidly scaled to produce approximately 10,000 tons of bleaching powder annually within five years, making it the world's largest chemical facility at the time. This output revolutionized textile bleaching, reducing processing times from months to days and fueling the Industrial Revolution's textile boom. By the mid-19th century, Tennant's firm had expanded globally, with bleaching powder exports supporting industries in paper production and sanitation. Advancements in purity followed, elevating the available chlorine content from around 25-30% in early batches to 35-40% by the late 1800s through optimized reaction controls and raw material quality, though high-purity forms exceeding 65% emerged later. The company's growth culminated in its 1890 merger into the United Alkali Company, further consolidating production.[9][11] Calcium hypochlorite's military significance peaked during World War I, where it served as a vital disinfectant in the trenches to combat waterborne diseases like dysentery and typhoid amid unsanitary conditions. Soldiers used bleaching powder to treat drinking water and clean equipment, while its formulation into eusol—an antiseptic solution combining equal parts chloride of lime (calcium hypochlorite) and boric acid—became standard for irrigating infected wounds, reducing sepsis rates.[9] Post-war, civilian applications surged, with calcium hypochlorite adopted for municipal water treatment and household disinfection, expanding from wartime stockpiles into everyday products. The 20th century brought key refinements, including stabilized granular calcium hypochlorite with 65-70% available chlorine purity, developed to minimize decomposition and enhance safety during storage and handling. These forms, often produced via the sodium process involving sodium hypochlorite and lime, replaced dusty powders with easier-to-use tablets and briquettes for swimming pools and sanitation, reflecting lessons from military logistics.[9]Structure and properties
Molecular structure
Calcium hypochlorite is an inorganic compound with the chemical formula .[1] It exists primarily as an ionic solid, composed of one calcium dication () electrostatically bound to two hypochlorite anions ().[1] The hypochlorite anion features a polar covalent bond between the chlorine and oxygen atoms, where the oxygen carries a partial negative charge and possesses three lone pairs of electrons, while the chlorine atom has three lone pairs.[1] In commercial preparations, calcium hypochlorite is frequently encountered as the dihydrate, , which adopts a tetragonal crystal system with flat, square plate-like morphology.[12] The anhydrous form, however, crystallizes in an orthorhombic lattice belonging to the Ccce space group (No. 68), forming one-dimensional ribbons along the a-axis.[13] Structural analyses, including those derived from density functional theory and consistent with experimental diffraction patterns, reveal calcium ions coordinated to oxygen atoms from the hypochlorite ligands, with Ca-O bond distances ranging approximately from 2.34 Å to 2.38 Å and Cl-O bonds around 1.71 Å.[13] In comparison to sodium hypochlorite (), which forms highly soluble ionic solutions typically handled as liquids, the molecular structure of calcium hypochlorite results in a less soluble solid due to the divalent calcium cation strengthening the ionic lattice and reducing hydration tendencies.[14] This structural difference enables calcium hypochlorite to be distributed and stored as stable granules or tablets, contrasting with the aqueous nature of its sodium counterpart.[14]Physical and chemical properties
Calcium hypochlorite is typically a white or grayish crystalline powder, available in granular, pellet, or tablet forms, and pure samples are odorless, though commercial products often emit a chlorine-like odor due to trace decomposition or impurities.[1][15] It is hygroscopic and deliquescent, readily absorbing moisture from the air, which can cause clumping in storage.[15] Key physical properties include a density of 2.35 g/cm³ at 25°C and no true melting point, as the compound decomposes above 100°C, releasing oxygen and chlorine gases.[1][16] Its solubility in water is approximately 21 g/100 mL at 20–25°C, though it undergoes gradual decomposition in solution.[1][16] Chemically, calcium hypochlorite is a strong oxidizer, with commercial grades containing 65–70% available chlorine, quantified via iodometric titration.[15][17] Aqueous solutions are strongly basic, exhibiting a pH of 11–12 depending on concentration, due to hydrolysis of the hypochlorite ions.[18] The oxidizing power stems from the hypochlorite moieties in its ionic structure.[1]Production
Industrial processes
The primary industrial process for calcium hypochlorite production is the calcium method, which involves the direct chlorination of slaked lime (calcium hydroxide) slurry with chlorine gas in a continuous reactor system. In this process, chlorine gas is absorbed into the aqueous suspension of Ca(OH)₂, typically within absorption towers or agitated reactors, where the exothermic reaction proceeds according to the overall equation: 2 Ca(OH)₂ + 2 Cl₂ → Ca(OCl)₂ + CaCl₂ + 2 H₂O. The reaction mixture is maintained at controlled temperatures, often below 35°C, to minimize decomposition and ensure efficient conversion, with the hypochlorite product crystallizing out as a dihydrate.[19] Process flow typically includes preparation of the lime slurry, chlorination in multi-stage reactors for complete absorption, followed by solid-liquid separation via filtration or centrifugation to isolate the crude calcium hypochlorite crystals from the calcium chloride-rich mother liquor.[20] Purification steps are essential to remove calcium chloride impurities, which can lower product stability and available chlorine content; these involve washing the crystals with water or dilute hypochlorite solution and sometimes recrystallization to achieve desired purity levels.[21] A common variation is the sodium method, which incorporates sodium hydroxide into the slurry to convert calcium chloride to soluble sodium chloride, yielding a higher-purity product with reduced impurities, though it requires additional caustic input. The sodium method is the most common process used domestically in the United States.[5] For producing anhydrous calcium hypochlorite, a hot drying process is employed post-crystallization, where the dihydrate is heated to remove water, contrasting with the standard cold crystallization that retains hydration for stability in hydrated grades.[22] Global production is concentrated in China, the leading manufacturer and top exporter, followed by India; in the United States, manufacturing occurs at facilities in Tennessee and West Virginia, co-located with chlor-alkali plants. As of 2018, total worldwide capacity was estimated at approximately 400,000 metric tons per year, though more recent production estimates suggest around 1.26 million metric tons globally as of 2024.[23][24] Economic viability hinges on raw material costs, with chlorine (derived from chlor-alkali processes) and lime accounting for over 60% of expenses, alongside energy for drying; product pricing varies by purity grades, such as 65% available chlorine (standard for water treatment) versus 70% (premium for bleaching), typically ranging from $800 to $1,500 per metric ton.[25][26]Laboratory preparation
Calcium hypochlorite can be prepared in the laboratory by bubbling chlorine gas through a suspension of calcium hydroxide in water. The reaction proceeds as follows: To perform this synthesis, slaked lime (Ca(OH)₂) is first prepared by adding water to quicklime and allowing hydration, then suspended in distilled water to form a slurry (typically 10-20% solids by weight). Chlorine gas, generated from hydrochloric acid and potassium permanganate or manganese dioxide, is slowly introduced via a gas delivery tube while stirring the suspension vigorously to ensure even distribution and prevent localized overheating. The reaction is exothermic and should be conducted at temperatures below 30°C to minimize side products like calcium chlorate. The mixture turns milky as the white precipitate of calcium hypochlorite forms, along with some calcium chloride as a byproduct if excess chlorine is used. After 1-2 hours of chlorination, the reaction is stopped, and the precipitate is filtered using a Buchner funnel, washed with cold water to remove soluble impurities, and dried under vacuum or in a desiccator at low temperature (around 40-50°C) to avoid decomposition. Analytical verification involves iodometric titration to determine available chlorine content, typically aiming for 65-70%.[27][28] An alternative laboratory method employs double decomposition between sodium hypochlorite and calcium chloride solutions. Aqueous solutions of sodium hypochlorite (prepared from commercial bleach, ~10-15% NaOCl) and calcium chloride (saturated, ~40% CaCl₂) are mixed in a 2:1 molar ratio: The solutions are cooled to 5-10°C to promote precipitation of the less soluble calcium hypochlorite, stirred for 30-60 minutes, then filtered, washed with ice-cold water, and dried similarly to the previous method. This approach avoids direct handling of chlorine gas but requires careful pH control (around 10-11) to prevent hypochlorite decomposition. The precipitate is again verified by titration for purity. Yields for both methods typically range from 80-90%, depending on reactant quality and temperature control, with the chlorine-lime method often yielding slightly higher due to fewer impurities.[22][29] Safety precautions are essential given the hazardous nature of chlorine gas and the oxidizing properties of hypochlorite. All procedures must be conducted in a well-ventilated fume hood to avoid inhalation of toxic chlorine fumes, which can cause severe respiratory irritation. Protective equipment includes chemical-resistant gloves, safety goggles, face shields, and lab coats; avoid skin contact as calcium hypochlorite causes burns. Chlorine generation should use a gas trap to scrub excess gas with sodium hydroxide solution. The product must be stored in airtight, opaque containers away from acids, organics, and heat to prevent explosive decomposition or chlorine release. Small-scale experiments (under 100 g) are recommended to minimize risks.[2] Historical laboratory methods from the 19th century, pioneered by chemists like Charles Tennant around 1799-1800, closely mirrored the standard chlorine-lime procedure but on a smaller scale using hand-generated chlorine from salt and acid, often for bleaching experiments; these early syntheses produced impure forms but established the core reaction for educational and research use.[9]Reactions
Hydrolysis and aqueous behavior
When dissolved in water, calcium hypochlorite dissociates into calcium ions and hypochlorite ions (ClO⁻), which subsequently undergo hydrolysis according to the equilibrium reaction: This process generates hydroxide ions, rendering the resulting solution alkaline with a pH typically above 10. The hypochlorous acid (HOCl) formed is a weak acid and serves as the primary active species for disinfection, while the equilibrium shifts based on solution conditions.[30][31] The compound exhibits high solubility in water, approximately 21 g per 100 mL at 25°C, and the dissolution process is exothermic, releasing heat that can accelerate minor decomposition reactions. During dissolution, partial decomposition occurs, evolving small amounts of chlorine gas (Cl₂) and oxygen (O₂), which contributes to the overall heat generation and potential for localized temperature increases in concentrated solutions. Aqueous stability is limited, with solutions prone to degradation over time; for instance, available chlorine content can decrease by up to 40–50% within a month at room temperature, and decomposition accelerates significantly at 40°C, potentially leading to 14% loss over 50 days in stabilized formulations.[1][32][16][33][34] Speciation of hypochlorite in solution is highly pH-dependent, governed by the acid-base equilibrium HOCl ⇌ H⁺ + ClO⁻ (pKₐ ≈ 7.5). At pH values greater than 9, the hypochlorite ion (ClO⁻) predominates, enhancing solution stability but reducing the proportion of the more biocidal HOCl. Under neutral conditions (pH ≈ 7), a roughly equal mixture of HOCl and ClO⁻ exists, and further pH reduction can promote chlorine gas evolution via HOCl + HCl → Cl₂ + H₂O, though this is more pronounced in slightly acidic environments.[30] Detection of hypochlorite in aqueous solutions commonly employs UV absorbance spectroscopy, where ClO⁻ exhibits a characteristic peak at approximately 292 nm, allowing quantification via Beer's law, or iodometric titration, in which hypochlorite oxidizes iodide to iodine, which is then titrated with thiosulfate using starch as an indicator. These methods provide sensitive and selective analysis, with detection limits in the micromolar range suitable for environmental and water treatment monitoring.[35]Oxidation and acid reactions
Calcium hypochlorite reacts vigorously with acids to produce chlorine gas, calcium chloride, and water, as exemplified by the reaction with hydrochloric acid:This process involves the protonation of hypochlorite ions, leading to rapid chlorine evolution that poses significant hazards due to the toxicity and pressure buildup of the gas.[27] As a strong oxidizing agent, calcium hypochlorite facilitates the conversion of primary alcohols to aldehydes or carboxylic acids and secondary alcohols to ketones, typically under mild conditions in aqueous or phase-transfer media. It also enables epoxidation of alkenes, forming oxiranes through addition across the double bond. In these oxidations, the hypochlorite ion (ClO⁻) serves as an electrophile, often generating hypochlorous acid (HOCl) in situ, which attacks nucleophilic sites on substrates to transfer oxygen or chlorine equivalents.[36] In the haloform reaction, calcium hypochlorite oxidizes methyl ketones (CH₃COR) under basic conditions, yielding a carboxylic acid salt, haloform (typically CHCl₃), and chloride ions, with the stoichiometry:
This involves sequential chlorination of the methyl group followed by cleavage.[37] Upon heating, calcium hypochlorite undergoes thermal decomposition via disproportionation:
The chlorate product can further decompose to release oxygen and calcium chloride, contributing to the compound's instability at elevated temperatures above 100 °C.[38]