Calcination
Calcination refers to 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), making it a major contributor to climate change.
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]
Industrial processes
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. 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.
Reactions
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
In limestone calcination, a decomposition process that occurs at 900 to 1050 °C, the chemical reaction is
- CaCO3(s) → CaO(s) + CO2(g)
Today, 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
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
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 sal ammoniac 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]
References
- ^ "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.