Industrial processes. Examples of calcination processes include the following:

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2 Industrial processes The process of calcination derives its name from the Latin calcinare (to burn lime) 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 minerals, as in the calcination of limestone to drive off carbon dioxide; - decomposition of hydrated minerals, as in the calcination of bauxite and gypsum, to remove crystalline water 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.

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8 Aluminothermic reactions are exothermic chemical reactions using aluminium as the reducing agent at high temperature. The process is industrially useful for production of alloys of iron. The most prominent example is the thermite reaction between iron oxides and aluminium to produce iron itself: Fe 2 O Al 2 Fe + Al 2 O 3 This specific reaction is however not relevant to the most important application of aluminothermic reactions, the production of ferroalloys. For the production of iron, a cheaper reducing agent, coke, is used instead via the carbothermic reaction.

9 An aluminothermic reaction using iron(iii) oxide. The sparks flying outwards are globules of molten iron trailing smoke in their wake.

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14 Elemental aluminium cannot be produced by the electrolysis of an aqueous aluminium salt because hydronium ions readily oxidize elemental aluminium. Although a molten aluminium salt could be used instead, aluminiu oxide has a melting point of 2072 C so electrolysing it is impractical. In the Hall Héroult process, alumina, Al 2 O 3, is dissolved in molten synthetic cryolite, Na 3 AlF 6, to lower its melting point for easier electrolysis.

15 Pure cryolite has a melting point of 1009 ± 1 C. With a small percentage of alumina dissolved in it, its melting point drops to about 1000 C. Besides having a relatively low melting point, cryolite is used as an electrolyte because among other things it also solves alumina well, conducts electricity, dissociates electrolytically at higher voltage than alumina and has a lighter density than aluminum at the temperatures required by the electrolysis.

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20 Liquation is a metallurgical method for separating metals from an ore or alloy. The material must be heated until one of the metals starts to melt and drain away from the other and can be collected. This method was largely used to remove lead containing silver from copper, but it can also be used to remove antimony minerals from ore, and refine tin.

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22 Poling is a metallurgical method employed in the purification of copper which contains copper oxide as an impurity and also in the purification of tin ("Sn") which contains tin oxide (stannic oxide or "SnO 2 ") as an impurity. The impure metal, usually in the form of molten blister copper, is placed in an anode furnace for two stages of refining. In the first stage, sulfur and iron are removed by gently blowing air through the molten metal to form iron oxides and sulfur dioxide. The iron oxides are skimmed or poured off the top of the copper and the gaseous sulfur dioxide exits the furnace via the off-gas system. Once the first oxidation stage is complete, the second stage (reduction or poling) begins. This involves using a reducing agent, normally natural gas or diesel (but ammonia, liquid petroleum gas, and naphtha can also be used), to react with the oxygen in the copper oxide to form copper metal. In the past, freshly cut ("green") trees were used as wooden poles. The sap in these poles acted as the reducing agent. The heat of the copper makes the pole emit a gas that reduces the cuprous oxide to copper. It was the use of these poles gave rise to the term "poling". Care must be taken to avoid removing too much of the oxygen from the anode copper, as this will cause other impurities to change from their oxide to metallic states and they will remain in solid solution in the copper, reduce its conductivity and change its physical properties.

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