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Calcium oxide
Calcium oxide
Calcium oxide
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Calcium oxide
Calcium oxide
Calcium oxide
Ionic crystal structure of calcium oxide
  Ca2+   O2-
Powder sample of white calcium oxide
Names
IUPAC name
Calcium oxide
Other names
  • Lime
  • Quicklime
  • Burnt lime
  • Unslaked lime
  • Free lime (building)
  • Caustic lime
  • Pebble lime
  • Calcia
  • Oxide of calcium
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.013.763 Edit this at Wikidata
EC Number
  • 215-138-9
E number E529 (acidity regulators, ...)
485425
KEGG
RTECS number
  • EW3100000
UNII
UN number 1910
  • InChI=1S/Ca.O
    Key: ODINCKMPIJJUCX-UHFFFAOYSA-N
  • InChI=1/Ca.O/rCaO/c1-2
    Key: ODINCKMPIJJUCX-BFMVISLHAU
  • O=[Ca]
Properties
CaO
Molar mass 56.0774 g/mol
Appearance White to pale yellow/brown powder
Odor Odorless
Density 3.34 g/cm3[1]
Melting point 2,613 °C (4,735 °F; 2,886 K)[1]
Boiling point 2,850 °C (5,160 °F; 3,120 K) (100 hPa)[2]
Reacts to form calcium hydroxide
Solubility in Methanol Insoluble (also in diethyl ether, octanol)
Acidity (pKa) 12.8
−15.0×10−6 cm3/mol
Structure
Cubic, cF8
Thermochemistry
40 J·mol−1·K−1[3]
−635 kJ·mol−1[3]
Pharmacology
QP53AX18 (WHO)
Hazards
GHS labelling:
GHS05: CorrosiveGHS07: Exclamation mark
Danger
H302, H314, H315, H335
P260, P261, P264, P270, P271, P280, P301+P312, P301+P330+P331, P302+P352, P303+P361+P353, P304+P340, P305+P351+P338, P310, P312, P321, P330, P332+P313, P362, P363, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
3
0
2
W
Flash point Non-flammable[4]
Lethal dose or concentration (LD, LC):
>2000 mg/kg oral, female rat [5]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 5 mg/m3[4]
REL (Recommended)
 TWA 2 mg/m3[4]
IDLH (Immediate danger)
 25 mg/m3[4]
Safety data sheet (SDS) ICSC 0409
Related compounds
Other anions
Other cations
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Calcium oxide (formula: CaO), commonly known as quicklime or burnt lime, is a widely used chemical compound. It is a white, caustic, alkaline, crystalline solid at room temperature. The broadly used term lime connotes calcium-containing inorganic compounds, in which carbonates, oxides, and hydroxides of calcium, silicon, magnesium, aluminium, and iron predominate. By contrast, quicklime specifically applies to the single compound calcium oxide. Calcium oxide that survives processing without reacting in building products, such as cement, is called free lime.[6]

Quicklime is relatively inexpensive. Both it and the chemical derivative calcium hydroxide (of which quicklime is the base anhydride) are important commodity chemicals.

Preparation

[edit]

Calcium oxide is usually made by the thermal decomposition of materials, such as limestone or seashells, that contain calcium carbonate (CaCO3; mineral calcite) in a lime kiln. This is accomplished by heating the material to above 825 °C (1,517 °F),[7][8] a process called calcination or lime-burning, to liberate a molecule of carbon dioxide (CO2), leaving quicklime behind. This is also one of the few chemical reactions known in prehistoric times.[9]

CaCO3(s) → CaO(s) + CO2(g)

The quicklime is not stable and, when cooled, will spontaneously react with CO2 from the air until, after enough time, it will be completely converted back to calcium carbonate unless slaked with water to set as lime plaster or lime mortar.

Annual worldwide production of quicklime is around 283 million tonnes. China is by far the world's largest producer, with a total of around 170 million tonnes per year. The United States is the next largest, with around 20 million tonnes per year.[10]

Hydroxyapatite's free CaO content rises with increased calcination temperatures and longer times. It also pinpoints particular temperature cutoffs and durations that impact the production of CaO, offering information on how calcination parameters impact the composition of the material.

Uses

[edit]
A demonstration of slaking of quicklime as a strongly exothermic reaction. Drops of water are added to pieces of quicklime. After a while, a pronounced exothermic reaction occurs ("slaking of lime"). The temperature can reach up to some 300 °C (572 °F).
CaO (s) + H2O (l) ⇌ Ca(OH)2 (aq) (ΔHr = −63.7 kJ/mol of CaO)
As it hydrates, an exothermic reaction results and the solid puffs up. The hydrate can be reconverted to quicklime by removing the water by heating it to redness to reverse the hydration reaction. One litre of water combines with approximately 3.1 kilograms (6.8 lb) of quicklime to give calcium hydroxide plus 3.54 MJ of energy. This process can be used to provide a convenient portable source of heat, as for on-the-spot food warming in a self-heating can, cooking, and heating water without open flames. Several companies sell cooking kits using this heating method.[12]
  • It is a food additive used as an acidity regulator, a flour treatment agent and a leavener.[13] It has E number E529.
  • Light: When quicklime is heated to 2,400 °C (4,350 °F), it emits an intense glow. This form of illumination is known as a limelight, and was used broadly in theatrical productions before the invention of electric lighting.[14]
  • Cement: Calcium oxide is a key ingredient for the process of making cement.
  • As a cheap and widely available alkali.[15]
  • Petroleum industry: Water detection pastes contain a mix of calcium oxide and phenolphthalein. Should this paste come into contact with water in a fuel storage tank, the CaO reacts with the water to form calcium hydroxide. Calcium hydroxide has a high enough pH to turn the phenolphthalein a vivid purplish-pink color, thus indicating the presence of water.
  • Chemical pulping: Calcium oxide is used to make calcium hydroxide, which is used to regenerate sodium hydroxide from sodium carbonate in the chemical recovery at kraft pulp mills.
  • Plaster: There is archeological evidence that Pre-Pottery Neolithic B humans used limestone-based plaster for flooring and other uses.[16][17][18] Such Lime-ash floor remained in use until the late nineteenth century.
  • Chemical or power production: Solid sprays or slurries of calcium oxide can be used to remove sulfur dioxide from exhaust streams in a process called flue-gas desulfurization.
  • Carbon capture and storage: Calcium oxide can be used to capture carbon dioxide from flue gases in a process called calcium looping.
  • Mining: Compressed lime cartridges exploit the exothermic properties of quicklime to break rock. A shot hole is drilled into the rock in the usual way and a sealed cartridge of quicklime is placed within and tamped. A quantity of water is then injected into the cartridge and the resulting release of steam, together with the greater volume of the residual hydrated solid, breaks the rock apart. The method does not work if the rock is particularly hard.[19][20][21]
  • Disposal of corpses: Historically, it was mistakenly thought that quicklime was efficacious in accelerating the decomposition of corpses. The application of quicklime can, in fact, promote preservation. Quicklime can aid in eradicating the stench of decomposition, which may have led people to the erroneous conclusion.[22]
  • It has been determined that the durability of ancient Roman concrete is attributed in part to the use of quicklime as an ingredient. Combined with hot mixing, the quicklime creates macro-sized lime clasts with a characteristically brittle nano-particle architecture. As cracks form in the concrete, they preferentially pass through the structurally weaker lime clasts, fracturing them. When water enters these cracks it creates a calcium-saturated solution which can recrystallize as calcium carbonate, quickly filling the crack.[23]
  • The thermochemical heat storage mechanism is greatly impacted by the sintering of CaO and CaCO3. It demonstrates that the storage materials become less reactive and denser at increasing temperatures. It also pinpoints particular sintering processes and variables influencing the efficiency of these materials in heat storage.

Weapon

[edit]

Quicklime is also thought to have been a component of Greek fire. Upon contact with water, quicklime would increase its temperature above 150 °C (302 °F) and ignite the fuel.[24]

David Hume, in his History of England, recounts that early in the reign of Henry III, the English Navy destroyed an invading French fleet by blinding the enemy fleet with quicklime.[25] Quicklime may have been used in medieval naval warfare – up to the use of "lime-mortars" to throw it at the enemy ships.[26]

Substitutes

[edit]

Limestone is a substitute for lime in many applications, which include agriculture, fluxing, and sulfur removal. Limestone, which contains less reactive material, is slower to react and may have other disadvantages compared with lime, depending on the application; however, limestone is considerably less expensive than lime. Calcined gypsum is an alternative material in industrial plasters and mortars. Cement, cement kiln dust, fly ash, and lime kiln dust are potential substitutes for some construction uses of lime. Magnesium hydroxide is a substitute for lime in pH control, and magnesium oxide is a substitute for dolomitic lime as a flux in steelmaking.[27]

Safety

[edit]

Because of vigorous reaction of quicklime with water, quicklime causes severe irritation when inhaled or placed in contact with moist skin or eyes. Inhalation may cause coughing, sneezing, and labored breathing. It may then evolve into burns with perforation of the nasal septum, abdominal pain, nausea and vomiting. Although quicklime is not considered a fire hazard, its reaction with water can release enough heat to ignite combustible materials.[28][better source needed]

Mineral

[edit]

Calcium oxide is also a separate mineral species (with the unit formula CaO), named 'Lime'.[29][30] It has an isometric crystal system, and can form a solid solution series with monteponite. The crystal is brittle, pyrometamorphic, and is unstable in moist air, quickly turning into portlandite (Ca(OH)2).[31]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Calcium oxide

View on Grokipedia
from Grokipedia
Calcium oxide, commonly known as quicklime or lime, is an with the chemical formula and a molecular weight of 56.077 g/mol. It appears as an odorless, white or gray-white solid, often in the form of hard lumps or powder, and is highly caustic and alkaline, reacting vigorously with to form and generate significant . This compound is a strong irritant to , eyes, and mucous membranes, necessitating careful handling in industrial and settings. Calcium oxide is primarily produced industrially through the of (, CaCO₃) in rotary kilns at temperatures between 900°C and 1,200°C, following the reaction CaCO₃ → CaO + CO₂. This process, known as lime burning, has been utilized for and remains the dominant method due to the abundance of limestone deposits worldwide. The resulting quicklime is then processed into various forms, including hydrated lime (Ca(OH)₂), for diverse applications. Key physical properties include a high of approximately 2,613°C, a around 2,850°C, and a of 3.34 g/cm³, making it suitable for high-temperature processes. Chemically, it exhibits basic properties with a of about 12.5 in aqueous suspension and is sparingly soluble in (about 1.65 g/L at 20°C), but it readily absorbs and from the air, reverting to over time. In industry, calcium oxide serves as a fundamental material in construction, where it is a key ingredient in and for binding aggregates. It is also essential in as a flux to remove impurities like phosphorus and sulfur, and in for softening, adjustment, and sludge conditioning. Additional applications include as a soil amendment and , pulp and paper production for bleaching, and chemical manufacturing for processes like the Solvay soda production. Historically, it has been used in incendiary devices, and in modern contexts for due to its exothermic .

Properties

Physical properties

Calcium oxide is an odorless, white or gray-white crystalline solid, typically appearing as hard lumps, fine powder, or granules in its pure form. It exhibits a caustic and alkaline nature due to its basic properties, with a density of 3.34 g/cm³ at 20 °C. This compound demonstrates high thermal stability, with a melting point of 2613 °C and a boiling point of 2,850 °C, at which it tends to sublime rather than boil in a conventional sense. Its specific heat capacity is approximately 0.75 J/g·K for the solid phase near room temperature. Regarding solubility, calcium oxide is practically insoluble in water—reacting instead to form with significant evolution—but it dissolves in while remaining insoluble in and acetone. In commercial quicklime products, particle morphology varies by processing, often featuring irregular crystalline granules or lumps with sizes ranging from fine powder (less than 0.1 mm) to larger aggregates up to 150 mm, influencing handling and reactivity.

Chemical properties

Calcium oxide, with the chemical formula CaO, has a of 56.077 g/mol. It is an ionic consisting of calcium cations (Ca²⁺) and anions (O²⁻), where calcium exhibits an of +2 and the ion serves as a strong Lewis base. The of calcium oxide is face-centered cubic, adopting the rock salt (NaCl) structure with Fm³m, in which each Ca²⁺ is octahedrally coordinated to six O²⁻ and vice versa. This ionic lattice contributes to its high and stability under standard conditions. Calcium oxide is a strong base owing to the basicity of the ion, which readily accepts protons; an aqueous suspension exhibits a of approximately 12.8. Key chemical reactions of calcium oxide highlight its reactivity. It undergoes slaking with in an : CaO+H2O→Ca(OH)2\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2 releasing significant heat, up to 65 kJ/mol. With , it forms : CaO+CO2→CaCO3\text{CaO} + \text{CO}_2 \rightarrow \text{CaCO}_3 a reaction driven by the affinity of the ion for CO₂. Calcium oxide also reacts vigorously with acids, such as , in a neutralization reaction: CaO+2HCl→CaCl2+H2O\text{CaO} + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} demonstrating its utility as a basic reagent. Calcium oxide is hygroscopic, readily absorbing atmospheric moisture and CO₂ to form surface layers of and , which can cause crumbling. It remains stable under ambient conditions but decomposes into elemental calcium and oxygen at extremely high temperatures above 2500 K, as dictated by its thermodynamics.

Production

Industrial production

Calcium oxide, commonly known as quicklime, is primarily produced on an industrial scale through the of (, CaCO₃) in large rotary kilns at temperatures ranging from 900°C to 1,200°C, following the reaction CaCO₃ → CaO + . This decomposes the carbonate, releasing carbon dioxide gas and leaving behind the oxide, with rotary kilns favored for their efficiency in handling high volumes of material in continuous operation. The primary raw material is high-purity limestone, often sourced from quarries, though marble can also serve as a suitable alternative due to its similar calcium carbonate composition; impurities such as silica or clay in the feedstock can lead to reduced product purity and affect downstream applications by forming silicates or other contaminants. Various kiln types are employed depending on scale and desired product characteristics, including shaft kilns for smaller or intermittent operations and advanced flash calcination processes that rapidly heat fine limestone particles to produce highly reactive quicklime with greater surface area. Energy consumption for these processes typically ranges from 5 to 7 GJ per ton of lime produced, primarily from fuels like natural gas or coal used to sustain the high temperatures required. Global production of lime, including quicklime, reached 424 million metric tons in 2023 and an estimated 420 million metric tons in 2024, underscoring its status as a major commodity chemical essential to industries like and ; leading producers include (310 million tons), (17 million tons), the (16 million tons), and several European countries. A significant by-product of calcination is CO₂, with the lime industry accounting for about 1% of global industrial CO₂ emissions, predominantly from the inherent process reaction; mitigation efforts include technologies integrated into operations to reduce these emissions. Produced lime is classified into quality grades based on reactivity, with quicklime offering high reactivity for applications requiring rapid slaking and dead-burned lime, achieved through prolonged high-temperature exposure, exhibiting lower reactivity and suited for materials due to its denser structure.

Laboratory preparation

Calcium oxide is commonly prepared in the laboratory through the of , a method analogous to but scaled down from industrial processes. Analytical-grade (CaCO₃) is heated in a at temperatures between 900 and 1000 °C for 1–2 hours, leading to the reaction: CaCO3(s)→CaO(s)+CO2(g)\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)} This produces a white, powdery residue of calcium oxide with yields typically approaching the theoretical value of 56% based on the ratio, often achieving 90–95% practical recovery when using pure and monitoring weight loss to confirm CO₂ evolution. The equipment includes or crucibles to hold the sample and a for controlled high-temperature heating, ensuring even decomposition without contamination. An alternative route involves the of , which is heated to 500–600 °C in a similar furnace setup. The reaction proceeds as: Ca(OH)2(s)→CaO(s)+H2O(g)\mathrm{Ca(OH)_2(s) \rightarrow CaO(s) + H_2O(g)} This method yields high-purity calcium oxide, with decomposition initiating around 512 °C and completing efficiently at higher temperatures, providing yields of 80–90% depending on the hydroxide's initial purity. It is particularly useful when starting from slaked lime, and the produced can be vented to prevent of the product. Calcium oxide can also be synthesized by direct oxidation of calcium metal in an oxygen atmosphere. Finely divided calcium is ignited or heated in a dry oxygen stream within a controlled inert setup, such as a , following: 2Ca(s)+O2(g)→2CaO(s)2\mathrm{Ca(s) + O_2(g) \rightarrow 2CaO(s)} This requires careful temperature control (around 300–500 °C) to avoid excessive , yielding nearly stoichiometric calcium oxide with purity exceeding 95% if high-purity metal is used. precautions include using an inert atmosphere initially to prevent premature ignition and handling the reactive metal with in a glove box. Other methods utilize the calcination of calcium salts like oxalate or nitrate for specialized applications. Calcium oxalate (CaC₂O₄) is decomposed at 600–800 °C to form CaO, CO, and CO₂, while calcium nitrate (Ca(NO₃)₂) decomposes around 600 °C to CaO, NO₂, and O₂; both require analytical-grade precursors and yield 85–95% CaO after purification steps like washing to remove impurities. Throughout these preparations, purity is ensured by using reagent-grade starting materials and post-reaction treatments such as sieving or acid washing, with overall yields ranging from 80–95%. Equipment like muffle furnaces and crucibles is standard, but safety notes emphasize wearing heat-resistant gloves, using fume hoods for gaseous byproducts, and avoiding direct contact with hot samples to prevent burns or reactions with moisture.

Applications

Construction and materials

Calcium oxide, commonly known as quicklime, plays a pivotal role in the production of , where it reacts with silica and alumina from clay or at high temperatures to form clinker, the primary component of cement. This clinker is then ground with to produce the final cement powder used extensively in for building foundations, bridges, and . In lime mortar applications, quicklime is slaked with to form , which is mixed with to create a breathable binder suitable for historic restorations and modern eco-friendly buildings, allowing moisture vapor to pass through walls and prevent damage from trapped . The high reactivity of calcium oxide enables rapid setting and strong adhesion in these mortars, contributing to their durability and flexibility compared to rigid cement-based alternatives. In steel manufacturing, calcium oxide serves as a flux in basic oxygen furnaces (BOF), where it combines with impurities such as and to form , which is then removed to purify the molten . This process, essential for producing high-quality used in beams, , and structural components, consumes significant quantities of quicklime to maintain the basicity of the and facilitate efficient impurity removal. Calcium oxide is also employed in water treatment processes within construction to soften hard water used for concrete mixing, where it reacts with calcium and magnesium ions to precipitate them as insoluble carbonates, preventing adverse effects on cement hydration and concrete strength. In glass and ceramics production, calcium oxide acts as a flux to lower the melting point of silica, enabling the formation of stable soda-lime glass through the lime-soda process, which combines limestone (source of CaO), soda ash, and sand for windows, bottles, and containers used in building materials. Similarly, in ceramics, it stabilizes glazes and bodies by adjusting viscosity and promoting vitrification at lower temperatures, enhancing the quality of tiles and sanitary ware for construction interiors. The construction industry is the largest end-user of calcium oxide, accounting for over half of global demand as of 2025. The material's advantages, such as its high reactivity for quick setting and its historical application in —where unslaked quicklime created self-healing properties—continue to influence modern sustainable construction practices.

Chemical manufacturing

Calcium oxide serves as a versatile and in numerous industrial chemical processes, leveraging its strong basicity and reactivity to facilitate the synthesis of key compounds. In the production of , a critical intermediate for and other chemicals, calcium oxide reacts with carbon at high temperatures. The reaction proceeds as follows: CaO+3C→CaC2+CO\mathrm{CaO + 3C \rightarrow CaC_2 + CO} This typically occurs in furnaces at approximately 2,000°C, with lime derived from calcined providing the calcium source and or serving as the carbon input. The resulting is then hydrolyzed to produce gas, which is essential for and . In the pulp and paper industry, calcium oxide plays a pivotal role in the kraft process's chemical recovery cycle, specifically during the causticizing step. Here, quicklime (CaO) reacts with sodium carbonate in green liquor to regenerate sodium hydroxide, the active cooking agent for wood pulping: CaO+Na2CO3→CaCO3+2NaOH\mathrm{CaO + Na_2CO_3 \rightarrow CaCO_3 + 2NaOH} This reaction precipitates calcium carbonate mud, which is subsequently calcined in a lime kiln to recover CaO, enabling a closed-loop system that minimizes waste. The process efficiency depends on the slaking of CaO to calcium hydroxide prior to reaction, ensuring high conversion rates of sodium carbonate to hydroxide. Calcium oxide is widely employed as a heterogeneous base catalyst in via the of vegetable oils or animal fats with or . Typically, 1-2% CaO loading by weight of the oil is used, promoting the conversion of triglycerides to fatty acid methyl esters () and at temperatures around 60-65°C. Yields exceeding 90% are achievable under optimized conditions, such as a 6:1 -to-oil ratio and 2-3 hours reaction time, with the catalyst's activity stemming from its Lewis basic sites. Waste-derived CaO, like from eggshells, enhances while maintaining efficacy. In petroleum refining, calcium oxide functions as a desulfurization agent to remove sulfur compounds from fuels and as a drying agent to eliminate moisture from hydrocarbons. It neutralizes acidic impurities and absorbs water in processes like caustic washing or additive formulations, with specialized grades like PetroCal O tailored for these applications. Its role supports compliance with low-sulfur fuel standards by facilitating the conversion of sulfur oxides or mercaptans. For , calcium oxide acts as a agent to remove from reaction mixtures and as a base in esterification reactions, promoting the formation of esters from acids and alcohols. In variants, it catalyzes the exchange of alkyl groups, often at mild conditions, while its dehydrating properties prevent side reactions. These applications benefit from CaO's mild basicity, avoiding over-alkalinity issues seen with homogeneous catalysts. High-grade calcium oxide, with purity exceeding 95%, is essential for catalytic applications to ensure minimal impurities that could deactivate sites or contaminate products. Commercial lime for these uses typically contains 90-95% free CaO when properly stored, but catalytic processes demand refined grades approaching 99% to optimize activity and selectivity.

Agriculture and environmental uses

Calcium oxide, commonly known as quicklime, plays a vital role in agriculture by neutralizing acidic soils through liming, which raises soil pH and enhances nutrient availability for crops. This process counters soil acidity caused by factors such as acid rain, fertilizer application, and crop nutrient uptake, thereby improving plant growth and yield. Application rates typically range from 1 to 5 tons per hectare, determined by soil tests that assess buffering capacity, target pH, and crop needs; quicklime's high neutralizing value (150-175% calcium carbonate equivalent) allows for lower quantities compared to agricultural limestone. In animal feed, calcium oxide serves as a supplemental source of calcium for and , supporting health, formation, and overall metabolic functions. It is (GRAS) when incorporated into formulations at low levels, such as 1% of in rations to improve feed digestibility and reduce acidity in by-products like . Quicklime is often slaked to for safer integration into feed blocks or supplements containing , minerals, and non-protein nitrogen. Calcium oxide is widely employed in and for its ability to precipitate and , aiding in pollution control. In removal, lime raises to form insoluble , effectively reducing levels in effluents from sources like agricultural runoff or . For such as lead, , and , quicklime precipitates them as hydroxides at concentrations exceeding 1000 mg/L, enabling efficient separation from industrial wastewater. Additionally, in (FGD) systems at power plants, calcium oxide reacts with to form , mitigating precursors: CaO+SO2→CaSO3\text{CaO} + \text{SO}_2 \rightarrow \text{CaSO}_3 This dry or wet scrubbing process captures up to 90% of SO₂ emissions. In sugar refining, particularly for beet juice clarification, calcium oxide is added to raw after heating, where it neutralizes organic acids and facilitates impurity removal through subsequent . The addition of CO₂ converts excess lime to precipitates, which trap proteins, colorants, and other non-sugars, producing clear juice for and . This step enhances purity and yield, with quicklime preferred for its rapid reactivity in industrial-scale operations. For odor control in , calcium oxide neutralizes and reduces volatile compounds like by elevating above 12, creating an alkaline environment inhospitable to odor-producing . In stabilization, quicklime addition minimizes off-gassing of amines and during processing and storage, improving handling safety and land application viability. This application also aids reduction, aligning with regulatory standards for residuals. The use of calcium oxide in and environmental applications promotes by enhancing and reducing reliance on synthetic fertilizers, as improved optimizes natural cycling and microbial activity. These sectors represent a significant portion of global calcium oxide production, underscoring its role in eco-friendly practices that lower chemical inputs and mitigate environmental .

Historical and miscellaneous uses

Calcium oxide, known historically as quicklime, has been utilized since ancient times in for its binding properties when slaked to form . Archaeological evidence indicates its use in around 2600 BCE, where lime-based mortars helped bond blocks, contributing to the enduring structures of . The Romans advanced this application by combining quicklime with , a , to create hydraulic that set underwater and exhibited superior durability, as described in Vitruvius's engineering texts and evidenced in structures like the Pantheon. In military contexts, quicklime is speculated to have served as an incendiary component in Byzantine , a petroleum-based weapon deployed from the onward, where its with intensified flames and caused burns upon contact. During the medieval period, lime mortars incorporating quicklime were widely employed in European fortifications, such as castle walls and cathedrals, providing flexible yet strong bonding that accommodated structural shifts over time. For applications requiring adjusted reactivity, substitutes like have been used in place of calcium oxide, particularly in fluxes where higher refractoriness is needed, and dolomitic lime (containing both calcium and magnesium oxides) offers slower slaking for controlled adjustments in treatment. provides similar basicity but with greater thermal stability, making it preferable in high-temperature processes. Among miscellaneous uses, quicklime plays a key role in tanning through depilation, where it swells and loosens hair from hides via alkaline , a process dating back to ancient civilizations and still employed today. Slaked lime, derived from calcium oxide, is essential in , the traditional Mesoamerican processing of corn, where it enhances nutritional by breaking down cell walls and adding calcium. Recent research since 2010 has explored calcium oxide nanoparticles for catalytic applications, particularly in via , where their high surface area enables efficient conversion of triglycerides with yields exceeding 95% under mild conditions. These also show promise in CO2 capture and , leveraging their basic sites for enhanced reactivity. Culturally, quicklime held symbolic significance in as a of transformation, representing the stage where base materials were purified into higher forms, akin to spiritual enlightenment in medieval esoteric texts. This of rebirth underscored its role in alchemical pursuits of the .

Safety and environmental considerations

Health and safety hazards

Calcium oxide, commonly known as quicklime, is highly corrosive and acts as a severe irritant to the skin, eyes, and . Upon contact with moist tissues, it undergoes exothermic hydration, generating significant heat that can cause chemical burns and tissue damage. This reactivity makes it particularly hazardous in handling, as even brief exposure can lead to painful blisters, ulceration, or perforation of affected areas. Primary exposure routes include of dust particles, which can irritate the upper and, in cases of impure calcium oxide containing silica, pose a risk of —a chronic lung disease characterized by scarring of lung tissue. Skin contact often results in , with symptoms ranging from redness and itching to severe burns, while eye exposure can cause immediate pain, swelling, and potential vision impairment. is less common but can lead to gastrointestinal irritation and burns in the mouth and . Toxicity assessments indicate low acute oral toxicity, with an LD50 greater than 2000 mg/kg in rats, suggesting it is not highly poisonous when swallowed in small amounts. The International Agency for Research on Cancer (IARC) does not classify calcium oxide as a . However, chronic inhalation exposure has been associated with lung issues, including and nasal septum ulceration, due to prolonged irritation of respiratory passages. To mitigate risks, (PPE) such as chemical-resistant gloves, safety goggles, protective clothing, and respirators approved for particulate matter is essential during handling. Storage should occur in cool, dry, well-ventilated areas to prevent moisture absorption and spontaneous reactions, and workplaces must adhere to the OSHA (PEL) of 5 mg/m³ as an 8-hour time-weighted average for calcium oxide dust. like local exhaust ventilation further reduce airborne concentrations. In case of exposure, immediate is critical: for skin or , flush thoroughly with water for at least 15-20 minutes while removing contaminated clothing, and seek medical attention; for , move the affected person to fresh air and provide oxygen if breathing is difficult; if ingested, do not induce vomiting but rinse the mouth and contact poison control. Historical incidents in lime kilns frequently involved severe burns from hot calcium oxide or dust exposure, contributing to worker injuries before modern safeguards. Under REACH regulations, calcium oxide is classified as causing skin (Category 2), serious eye damage (Category 1), and specific target organ from single exposure ( , Category 3), mandating strict handling protocols across member states.

Environmental impact

The production of calcium oxide through calcination of limestone releases approximately 0.78 tons of CO₂ per ton of CaO, primarily due to the decomposition of calcium carbonate, contributing significantly to global and . Airborne lime from production and handling processes can degrade local air quality by forming fine particulates that scatter and deposit on ecosystems, while alkaline runoff from sites elevates water levels, potentially causing oxygen depletion and disrupting aquatic habitats. Limestone quarrying for calcium oxide feedstock results in and loss, leading to declines as vegetation is cleared and structures are altered, with long-term effects on local and populations. To mitigate these impacts, (CCS) technologies are being integrated into lime kilns, capturing up to 99% of process CO₂ for geological sequestration, while waste materials such as eggshells—calcined to produce CaO—reduces reliance on virgin and diverts organic waste from landfills. Regulatory measures like the (EU ETS) impose caps on CO₂ emissions from lime production facilities, requiring allowances for verified outputs and incentivizing reductions through trading, while life-cycle assessments indicate a of approximately 1.0 kg CO₂ equivalent per kg of calcium oxide, encompassing extraction, production, and transport stages. On a positive note, calcium oxide's application in SO₂ scrubbing processes neutralizes emissions from industrial sources, thereby reducing formation and associated ecosystem damage.

Natural occurrence

Associated minerals

Calcium oxide, mineralogically known as lime (IMA symbol: Lm), is an extremely rare naturally occurring mineral due to its high reactivity, which causes it to rapidly form hydrated or carbonated compounds in the presence of or atmospheric CO₂. It has been documented in trace amounts as a product of high-temperature processes, such as coal-seam fires where decomposes, and in altered xenoliths within volcanic or sublimates. Additionally, pure CaO has been identified in calcium-aluminum-rich inclusions within meteorites like Allende. A related mineral is (Ca(OH)₂), the hydrated form of calcium oxide, which occurs more frequently in natural environments such as altered ultramafic rocks and combustion metamorphosed limestones, though it differs from pure CaO by incorporating water of hydration. This hydration reflects the instability of anhydrous CaO under typical terrestrial conditions. When naturally occurring quicklime is found, it is typically impure, often containing significant magnesia (MgO) in dolomitic varieties derived from magnesium-bearing limestones, along with minor traces of silica (SiO₂) and alumina (Al₂O₃) from the parent rock composition. Over 99% of calcium oxide in existence is anthropogenic, produced via industrial , underscoring the negligible role of natural sources. The identification of calcium oxide in potential natural samples relies on techniques like X-ray diffraction (XRD), which detects its through characteristic peaks at 2θ values of approximately 32°, 37°, 54°, 64°, and 67°; such analyses often reveal it in industrial residues or slags that may be erroneously classified as minerals.

Geological formation

Calcium oxide, commonly known as lime, primarily forms in geological settings through the of precursor minerals such as (, CaCO₃), where heat and pressure drive decarbonation reactions that release . This process occurs during contact or regional , typically at temperatures exceeding 700–900°C under low-pressure conditions, decomposing into CaO and CO₂. Similarly, in meteoritic contexts, CaO appears as a minor oxide phase within calcium-aluminum-rich inclusions (CAIs) in chondritic meteorites, representing early condensates from the solar nebula formed at high temperatures over 4.5 billion years ago. The geological stability of calcium oxide is limited by its rapid hydration in the presence of atmospheric moisture, forming (Ca(OH)₂) and preventing the accumulation of large natural deposits. Exploration for potential CaO sources focuses on geophysical surveys targeting precursors, such as seismic and magnetic methods to map layers, given the rarity of pure lime, which constitutes less than 0.01% of as a distinct phase.

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