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Calcination
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Calcination is thermal treatment of a solid chemical compound (e.g. mixed carbonate ores) whereby the compound is raised to high temperature without melting under restricted supply of ambient oxygen (i.e. gaseous O2 fraction of air), generally for the purpose of removing impurities or volatile substances and/or to incur thermal decomposition.[1]
The root of the word calcination refers to its most prominent use, which is to remove carbon from limestone (calcium carbonate) through combustion to yield calcium oxide (quicklime). This calcination reaction is CaCO3(s) → CaO(s) + CO2(g). Calcium oxide is a crucial ingredient in modern cement, and is also used as a chemical flux in smelting. Industrial calcination generally emits carbon dioxide (CO2).
A calciner is a steel cylinder that rotates inside a heated furnace and performs indirect high-temperature processing (550–1150 °C, or 1000–2100 °F) within a controlled atmosphere.[2]
Etymology
[edit]The process of calcination derives its name from the Latin calcinare 'to burn lime'[3] due to its most common application, the decomposition of calcium carbonate (limestone) to calcium oxide (lime) and carbon dioxide, in order to create cement. The product of calcination is usually referred to in general as "calcine", regardless of the actual minerals undergoing thermal treatment.
Industrial processes
[edit]
Calcination is carried out in furnaces or reactors (sometimes referred to as kilns or calciners) of various designs including shaft furnaces, rotary kilns, multiple hearth furnaces, and fluidized bed reactors.
Examples of calcination processes include the following:
- decomposition of carbonate ores, as in the calcination of limestone to drive off carbon dioxide;
- decomposition of hydrated minerals, as in the calcination of bauxite and gypsum, carbonate ore to remove water of crystallization as water vapor;
- decomposition of volatile matter contained in raw petroleum coke;
- heat treatment to effect phase transformations, as in conversion of anatase to rutile or devitrification of glass materials;
- removal of ammonium ions in the synthesis of zeolites;
- defluorination of uranyl fluoride to create uranium dioxide and hydrofluoric acid gas;
- heat treatment of anthracite through electrically fired calcining furnace or gas calcination which results in development of graphitic structure.
Reactions
[edit]Calcination reactions usually take place at or above the thermal decomposition temperature (for decomposition and volatilization reactions) or the transition temperature (for phase transitions). This temperature is usually defined as the temperature at which the standard Gibbs free energy for a particular calcination reaction is equal to zero.
Limestone calcination
[edit]In limestone calcination, a decomposition process that occurs at 900 to 1050 °C, the chemical reaction is
CaCO3(s) → CaO(s) + CO2(g)
Today,[when?] this reaction largely occurs in a cement kiln.
The standard Gibbs free energy of reaction in [J/mol] is approximated as ΔG°r ≈ 177,100 J/mol − 158 J/(mol*K) * T.[4] The standard free energy of reaction is 0 in this case when the temperature, T, is equal to 1121 K, or 848 °C.
Oxidation
[edit]In some cases, calcination of a metal results in oxidation of the metal to produce a metal oxide. In his essay "Formal response to the question, why Tin and Lead increase in weight when they are calcined" (1630), Jean Rey notes that "having placed two pounds six ounces of fine English tin in an iron vessel and heated it strongly on an open furnace for the space of six hours with continual agitation and without adding anything to it, he recovered two pounds thirteen ounces of a white calx". He claimed "That this increase in weight comes from the air, which in the vessel has been rendered denser, heavier, and in some measure adhesive, by the vehement and long-continued heat of the furnace: which air mixes with the calx (frequent agitation aiding) and becomes attached to its most minute particles: not otherwise than water makes heavier sand which you throw into it and agitate, by moistening it and adhering to the smallest of its grains", presumably the metal gained weight as it was being oxidized.[5]
At room temperature, tin is quite resistant to the impact of air or water, as a thin oxide film forms on the surface of the metal. In air, tin starts to oxidize at a temperature of over 150 °C: Sn + O2 → SnO2.[6]
Antoine Lavoisier explored this experiment with similar results time later.[7]
Alchemy
[edit]In alchemy, calcination was believed to be one of the 12 vital processes required for the transformation of a substance.
Alchemists distinguished two kinds of calcination, actual and potential. Actual calcination is that brought about by actual fire, from wood, coals, or other fuel, raised to a certain temperature. Potential calcination is that brought about by potential fire, such as corrosive chemicals; for example, gold was calcined in a reverberatory furnace with mercury and salammoniac; silver with common salt and alkali salt; copper with salt and sulfur; iron with salammoniac and vinegar; tin with antimony; lead with sulfur; and mercury with nitric acid.[8]
There was also philosophical calcination, which was said to occur when horns, hooves, etc., were hung over boiling water, or other liquor, until they had lost their mucilage, and were easily reducible into powder.[8]
According to the obsolete phlogiston theory, the 'calx' was the true elemental substance that was left after phlogiston was driven out of it in the process of combustion.[9]
References
[edit]- ^ "Calcination". The IUPAC Compendium of Chemical Terminology. 2014. doi:10.1351/goldbook.C00773.
- ^ "High-Temperature Processing with Calciners".
- ^ Mosby's Medical, Nursing and Allied Health Dictionary, Fourth Edition, Mosby-Year Book Inc., 1994, p. 243
- ^ Gilchrist, J.D. (1989). Extraction Metallurgy (3rd ed.). Oxford: Pergamon Press. p. 145. ISBN 978-0-08-036612-8.
- ^ Rey, Jean (1953). Essays of Jean Rey, doctor of medicine, on an enquiry into the cause wherefore tin and lead increase in weight on calcination (1630). E. & S. Livingstone for the Alembic Club. OCLC 154124030.
- ^ "Tin: its oxidation states and reactions with it".
- ^ "Lavoisier tin calcination".
- ^ a b
This article incorporates text from a publication now in the public domain: Chambers, Ephraim, ed. (1728). "Calcination". Cyclopædia, or an Universal Dictionary of Arts and Sciences (1st ed.). James and John Knapton, et al.
- ^ Daintith, John, ed. (2008). "Phlogiston theory". A Dictionary of Chemistry (6th ed.). Oxford University Press. doi:10.1093/acref/9780199204632.001.0001. ISBN 9780191726569 – via Oxford Reference.
In the early 18th century Georg Stahl renamed the substance phlogiston (from the Greek for 'burned') and extended the theory to include the calcination (and corrosion) of metals. Thus, metals were thought to be composed of calx (a powdery residue) and phlogiston; when a metal was heated, phlogiston was set free and the calx remained. The process could be reversed by heating the metal over charcoal (a substance believed to be rich in phlogiston, because combustion almost totally consumed it). The calx would absorb the phlogiston released by the burning charcoal and become metallic again.
Calcination
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Calcination is a pyrometallurgical process involving the thermal treatment of solid materials at high temperatures, typically above 900°C, to drive off volatile components such as carbon dioxide, chemically bound water, or organic matter without causing fusion or melting of the substance.[1] This endothermic reaction often results in chemical decomposition, phase transitions, or the purification of the material, producing a more stable oxide or residue known as calcine.[2] The process is conducted in controlled atmospheres, such as in the absence of oxygen to prevent oxidation, and is distinguished from roasting by its focus on decomposition rather than sulfation.[1]
Historically, calcination originated from the Latin term calcinare, referring to the heating of limestone (calcium carbonate) to produce quicklime (calcium oxide) by releasing CO₂, a reaction exemplified by CaCO₃(s) → CaO(s) + CO₂(g) at temperatures exceeding 900°C.[3] Over time, the term has broadened beyond lime production to encompass similar thermal treatments in various chemical and metallurgical contexts, evolving with industrial advancements in the 19th and 20th centuries to include ore beneficiation and material synthesis.[1]
In modern applications, calcination plays a critical role in industries such as cement manufacturing, where limestone is decomposed into lime for clinker production, and in alumina extraction via the Bayer process, converting aluminum hydroxide to alumina at 1300–1500°C through reactions like 2Al(OH)₃(s) → Al₂O₃(s) + 3H₂O(g).[3][1] It is also essential in pigment and catalyst production, enhancing material properties like purity, porosity, and reactivity by removing impurities and volatiles.[2] Common equipment includes rotary kilns for continuous processing, which allow precise control of temperature profiles and retention times—often 6 hours to several days depending on the material—and fluidized bed calciners for uniform heating of fine particles.[3][2]
Key process parameters, including temperature (typically 800–1340°C), particle size, and partial pressure of gases like CO₂, significantly influence reaction efficiency and product quality, with higher temperatures yielding harder, less reactive limes suitable for refractory uses.[3] Environmental considerations are prominent, as calcination emits substantial CO₂, contributing to efforts in carbon capture and sustainable kiln designs.[3]
This reaction reaches equilibrium around 400°C under standard conditions, with the forward decomposition favored at higher temperatures due to the Le Chatelier principle shifting the balance away from the gaseous CO₂ product.[33] Similarly, hydroxide dehydration, such as that of aluminum hydroxide (Al(OH)₃, often in gibbsite form), yields alumina (Al₂O₃) and water vapor:
This stepwise process typically initiates between 200–500°C, involving intermediate boehmite or transition aluminas before stabilizing as α-Al₂O₃ at 1100–1300°C. The kinetics of these decompositions follow the Arrhenius equation, , where the rate constant depends exponentially on temperature , with activation energy quantifying the energy barrier for bond rupture. For calcium carbonate (CaCO₃) decomposition—a benchmark for carbonate calcination— is approximately 200 kJ/mol, reflecting the stability of the CO₃²⁻ ion and the need for substantial thermal activation.[34] Key influencing factors include particle size, which affects heat transfer and surface area exposure (smaller particles accelerate decomposition due to shorter diffusion paths for volatiles), and the surrounding atmosphere, where inert gases like N₂ promote faster rates by minimizing reverse carbonation compared to CO₂-rich environments.[35] During calcination, these reactions induce phase changes from the precursor solid (e.g., hydrated or carbonated mineral) to a more stable anhydrous oxide, often preserving nanoscale porosity in the product to maintain reactivity. Careful temperature control is essential to avoid sintering, where excessive heat causes particle coalescence and pore collapse, which would reduce the oxide's surface area and catalytic potential; thus, calcination protocols typically limit dwell times and temperatures to below the sintering threshold for the target oxide.[31]
This process strengthens pellet structure by forming a dense oxide layer, often conducted at 1100–1200°C with limited air to balance oxidation and prevent over-sintering. These reactions highlight the redox nature, where metals increase in oxidation state, driven by oxygen as the oxidizing agent.[37] Thermodynamically, these oxidation processes are governed by the Gibbs free energy change (ΔG), which must be negative for spontaneity: ΔG = ΔH - TΔS < 0 at elevated temperatures. For metal oxide oxidations, the enthalpy (ΔH) is typically negative due to strong metal-oxygen bonds formed, while entropy (ΔS) contributions from gas evolution further favor the reaction at high temperatures, as illustrated in Ellingham diagrams where oxidation lines slope downward. The partial pressure of oxygen plays a critical role; lower pO₂ in limited-air calcination shifts the equilibrium toward partial oxidation, preventing excessive exothermicity and allowing precise control over reaction extent.[38] Unlike pure thermal decomposition reactions, which are predominantly endothermic and focus on volatile release without oxygen involvement, oxidation-inclusive calcination incorporates exothermic contributions from redox steps, potentially reducing overall energy requirements and promoting sintering of particles into cohesive structures. This exothermicity can elevate local temperatures, accelerating diffusion and oxide layer formation, though careful air control is essential to avoid agglomeration.[37]
Fundamentals
Definition and Process Overview
Calcination is a thermal treatment process applied to solid materials, involving heating at high temperatures—typically in the range of 800–1400°C—within a controlled atmosphere, often with limited oxygen or air, to induce chemical changes such as decomposition, phase transformations, or the removal of volatile components, all without causing the material to melt.[4][5] This process is fundamental in materials science and chemical engineering for altering the physical and chemical properties of substances like ores, minerals, and precursors, enhancing attributes such as stability, reactivity, or purity.[6] The key steps in calcination include preheating the solid feedstock in a furnace or kiln to gradually elevate the temperature, maintaining the heat under controlled conditions to drive the desired reactions—such as the thermal decomposition of carbonates or hydroxides into oxides—while closely monitoring to prevent fusion or excessive sintering, and finally cooling the product to stabilize it.[2] Resulting products often include stable oxides; for instance, the calcination of calcium carbonate follows the general reaction: where limestone yields quicklime and carbon dioxide gas, though detailed mechanisms are governed by reaction kinetics and thermodynamics.[7] Calcination is distinct from related thermal processes: unlike roasting, which typically occurs in an oxidizing atmosphere with ample air to convert sulfides into oxides or sulfates, calcination emphasizes decomposition in a limited-oxygen environment to avoid such oxidation.[8] In contrast to sintering, which primarily bonds particles through diffusion at high temperatures without significant chemical decomposition, calcination focuses on inducing compositional changes like dehydration or decarbonation while avoiding particle fusion.[4] Common equipment includes rotary kilns, which rotate to ensure uniform heating and material movement at temperatures up to 1400°C; fluidized bed reactors, utilizing gas flow to suspend fine particles for efficient heat transfer at 800–1200°C; and shaft kilns, employing vertical gravity-fed designs for continuous operation in the 900–1300°C range.[9][10]Etymology
The term "calcination" derives from the Medieval Latin verb calcināre, meaning "to heat" or "to burn to lime," which itself stems from the Late Latin calcīna (lime) and the classical Latin noun calx (lime, limestone, or burnt stone).[11] This root reflects the process's ancient association with heating limestone to produce lime, a practice central to early construction and metallurgy.[12] The word entered English in the 14th century as calcinen or similar forms, borrowed via Old French calciner (to calcine), with the earliest recorded use of "calcination" appearing around 1386 in the works of Geoffrey Chaucer, where it denoted the fiery purification of substances.[13] Initially tied to alchemical contexts of reducing metals or minerals to a powdery "calx" through intense heat, the term's meaning evolved to encompass broader thermal treatments for decomposition. Etymologically linked to "calcium," the chemical element isolated from lime compounds and named in 1808 by Humphry Davy from Latin calx, the word also connects to "quicklime" (unslaked lime, or calx viva), underscoring its origins in lime production.[14] The Latin calx traces further to Ancient Greek kháliks (χάλιξ), meaning "pebble" or "small stone," highlighting the linguistic progression from natural stone to processed lime.[15] Over time, "calcination" shifted semantically from alchemical "purgation by fire" to its modern denotation of controlled thermal decomposition in chemistry.[16]Historical Development
Ancient and Medieval Uses
The practice of calcination originated in the Near East as early as 12,000 BCE, where early communities produced quicklime by heating limestone in simple hearths or kilns to create binders for plasters and mortars. Archaeological evidence from sites like Çatalhöyük in Turkey reveals lime plasters used in wall coatings and flooring as early as 6000 BCE, indicating controlled calcination processes for construction and symbolic purposes, such as modeling human skulls. Similarly, at Aşıklı Höyük in central Anatolia, dated to circa 7000 BCE, lime-based plasters containing calcined materials demonstrate the technology's role in stabilizing mud-brick structures and creating durable surfaces.[17][18] By the time of ancient Egypt, around 2600 BCE, calcined materials including gypsum and lime were integral to monumental architecture. Gypsum mortar served as a key component for binding massive limestone blocks in the pyramids at Giza, while lime plasters provided weather-resistant coatings.[19] This application extended the utility of calcination beyond basic shelter to large-scale engineering feats, enhancing structural integrity in arid environments. In the Roman era, from the 1st century BCE onward, calcination produced lime for opus caementicium, a hydraulic concrete mixed with pozzolanic additives like volcanic ash, enabling the construction of enduring infrastructure such as aqueducts, harbors, and the Pantheon. Archaeological remnants of Roman lime kilns and mortar samples from sites like Sagalassos confirm on-site calcination, which facilitated imperial expansion by supporting vast building projects that conveyed water over long distances and housed public works.[20][17] Medieval advancements refined calcination techniques, with improved kiln designs emerging in both Europe and the Islamic world to boost efficiency and output for diverse applications. In the Islamic realm, early structures like the Great Mosque of Kairouan in Tunisia (built 670 CE and expanded in the 9th century) incorporated calcined lime mortars for enhanced durability and water resistance, while Umayyad-period kilns in the Near East, dated to the 7th–8th centuries, featured advanced draft systems for consistent high-temperature burning. European developments paralleled this, as seen in Norman castles from the 11th century, where calcined lime supported rapid fortification efforts. These innovations extended calcination to glassmaking, where quicklime acted as a stabilizer in alkali-lime recipes prevalent in 9th–10th century Abbasid Iraq, and to pottery, aiding in the production of glazed wares through lime-based fluxes that improved firing and surface quality.[17][21]Alchemical Significance
In alchemy, practiced from approximately 300 BCE to 1700 CE, calcination constituted the inaugural phase of the Magnum Opus, or Great Work, embodying the nigredo stage of blackening that signified decomposition and the dissolution of base elements through intense heat, ultimately yielding a calx or ash as the purified residue. This process not only dismantled physical substances but also symbolized the alchemist's inner confrontation with chaos and ego, facilitating transmutation toward higher states of perfection.[22] The eighth-century polymath Jabir ibn Hayyan, often regarded as a foundational figure in Islamic alchemy, detailed calcination as an essential technique for metal purification, involving high-temperature roasting to separate and identify constituent principles such as sulfur and mercury, thereby enabling the refinement of base materials. In the sixteenth century, Paracelsus extended calcination's application to medicinal alchemy, employing fire in sealed vessels to incinerate impurities from metals like silver (Luna) and gold (Sol), extracting noble essences for therapeutic elixirs that promoted bodily purification and vitality.[22][23][24] Alchemical practitioners executed calcination through repeated heating of metals or minerals in specialized self-regulating furnaces known as athanors, which sustained gentle, constant warmth to volatilize volatile components while preserving essential principles, frequently in conjunction with philosophical mercury to isolate pure quintessences. Symbolically, fire in calcination served as the supreme purifying agent, mirroring the alchemist's spiritual journey toward enlightenment by incinerating material dross and unveiling latent divinity.[25][22]Transition to Modern Chemistry
In the late 18th century, Antoine Lavoisier revolutionized the understanding of calcination by reinterpreting it through rigorous experimentation, particularly his studies on the calcination of mercury conducted between 1774 and 1778. Observing that mercury gained weight upon forming its calx (oxide) when heated in air, and that the "fixed air" (dephlogisticated air, later identified as oxygen) diminished in volume, Lavoisier concluded that calcination involved the fixation of oxygen by the metal rather than the release of phlogiston as proposed by earlier theories. This empirical framework shifted calcination from a mystical alchemical process to a quantifiable chemical reaction, laying the groundwork for oxygen-based combustion theory.[26] The 19th century saw accelerated advancements in calcination studies, driven by the Industrial Revolution's demand for efficient thermal processes in materials production. Chemists began conducting precise volumetric and gravimetric analyses of decomposition reactions, with Joseph Louis Gay-Lussac contributing key observations in the 1810s and 1830s on the thermal decomposition of carbonates, such as limestone (calcium carbonate), where he noted that the presence of water vapor or carbon dioxide could influence the rate and completeness of calcination to form oxides. These investigations emphasized controlled heating conditions and gas evolution, providing foundational data for scaling chemical processes beyond laboratory settings.[27][28] By the early 20th century, calcination had been fully incorporated into modern thermodynamics and chemical kinetics, enabling predictive modeling of its underlying mechanisms. The introduction of activation energy by Svante Arrhenius in 1889 offered a quantitative tool to describe the minimum energy required for bond breaking in thermal decompositions, such as those yielding metal oxides from carbonates or hydroxides; this concept was rapidly applied to calcination kinetics, revealing how temperature thresholds govern reaction feasibility and equilibrium. Such integrations transformed calcination into a cornerstone of physical chemistry, with standardized equations for heat transfer and phase changes.[29] A pivotal milestone occurred with the development of the periodic table in 1869, where Dmitri Mendeleev explicitly linked calcination-derived oxide formulas to elemental classification, using properties like oxide valency and stability (e.g., MO, M₂O₃ patterns) to predict undiscovered elements and their behaviors under thermal treatment. This contextualized calcination as essential for revealing periodic trends in oxide formation across groups, solidifying its role in systematic inorganic chemistry.[30]Chemical Principles
Thermal Decomposition Reactions
Thermal decomposition reactions in calcination represent a fundamental non-oxidative process where inorganic compounds, such as carbonates and hydroxides, undergo endothermic breakdown upon heating, driven by the absorption of thermal energy to sever chemical bonds and liberate volatile gases like CO₂ or H₂O. This mechanism proceeds via a topotactic decomposition pathway, particularly for carbonates, where the crystal lattice of the precursor facilitates the nucleation and growth of the resulting oxide phase while releasing the gas in a controlled manner.[31] The endothermic nature ensures that sufficient heat input is required to overcome the lattice energy, preventing spontaneous reaction at ambient conditions and allowing precise control over the transformation.[32] A primary example is the decomposition of magnesium carbonate, where magnesite (MgCO₃) converts to magnesium oxide (MgO) and carbon dioxide:This reaction reaches equilibrium around 400°C under standard conditions, with the forward decomposition favored at higher temperatures due to the Le Chatelier principle shifting the balance away from the gaseous CO₂ product.[33] Similarly, hydroxide dehydration, such as that of aluminum hydroxide (Al(OH)₃, often in gibbsite form), yields alumina (Al₂O₃) and water vapor:
This stepwise process typically initiates between 200–500°C, involving intermediate boehmite or transition aluminas before stabilizing as α-Al₂O₃ at 1100–1300°C. The kinetics of these decompositions follow the Arrhenius equation, , where the rate constant depends exponentially on temperature , with activation energy quantifying the energy barrier for bond rupture. For calcium carbonate (CaCO₃) decomposition—a benchmark for carbonate calcination— is approximately 200 kJ/mol, reflecting the stability of the CO₃²⁻ ion and the need for substantial thermal activation.[34] Key influencing factors include particle size, which affects heat transfer and surface area exposure (smaller particles accelerate decomposition due to shorter diffusion paths for volatiles), and the surrounding atmosphere, where inert gases like N₂ promote faster rates by minimizing reverse carbonation compared to CO₂-rich environments.[35] During calcination, these reactions induce phase changes from the precursor solid (e.g., hydrated or carbonated mineral) to a more stable anhydrous oxide, often preserving nanoscale porosity in the product to maintain reactivity. Careful temperature control is essential to avoid sintering, where excessive heat causes particle coalescence and pore collapse, which would reduce the oxide's surface area and catalytic potential; thus, calcination protocols typically limit dwell times and temperatures to below the sintering threshold for the target oxide.[31]
Oxidation and Redox Processes
In the context of calcination processes involving oxygen, the procedure entails heating ores or metal compounds in a controlled atmosphere with limited air supply to facilitate the oxidation of impurities while avoiding complete combustion or fusion or sulfation. This variant contrasts with full roasting, which employs excess air to ensure thorough oxidation and volatilization of sulfur as SO₂. The limited oxygen environment promotes selective redox reactions that convert lower-valence metal species into stable oxides, enhancing the material's purity for subsequent processing without excessive energy input or unwanted side products.[36] Key redox reactions in such calcination include the oxidation of lower-valence iron oxides, exemplified in iron ore preparation for pelletizing, where magnetite undergoes oxidation to hematite:This process strengthens pellet structure by forming a dense oxide layer, often conducted at 1100–1200°C with limited air to balance oxidation and prevent over-sintering. These reactions highlight the redox nature, where metals increase in oxidation state, driven by oxygen as the oxidizing agent.[37] Thermodynamically, these oxidation processes are governed by the Gibbs free energy change (ΔG), which must be negative for spontaneity: ΔG = ΔH - TΔS < 0 at elevated temperatures. For metal oxide oxidations, the enthalpy (ΔH) is typically negative due to strong metal-oxygen bonds formed, while entropy (ΔS) contributions from gas evolution further favor the reaction at high temperatures, as illustrated in Ellingham diagrams where oxidation lines slope downward. The partial pressure of oxygen plays a critical role; lower pO₂ in limited-air calcination shifts the equilibrium toward partial oxidation, preventing excessive exothermicity and allowing precise control over reaction extent.[38] Unlike pure thermal decomposition reactions, which are predominantly endothermic and focus on volatile release without oxygen involvement, oxidation-inclusive calcination incorporates exothermic contributions from redox steps, potentially reducing overall energy requirements and promoting sintering of particles into cohesive structures. This exothermicity can elevate local temperatures, accelerating diffusion and oxide layer formation, though careful air control is essential to avoid agglomeration.[37]