What is generally referred to as the ‘lime production process‘ actually involves the burning and processing of by-products obtained from different types of minerals available in nature.
For the sake of explanatory clarity, the chemical diagram below is not so distinctly divided into the minerals that are used industrially, which represent a more or less complex mixture of the types mentioned.
CaCO3 MgCO3SiO2 Fe2O3 Al2O3
This mix of possible chemical compositions of the starting mineral obviously leads to a greater complexity of the “real” material behaviour both during its work and in terms of chemical and physical characteristics of the finished products of the productive process.
|CaCO3||+||Kcal||↔||CaO||+||CO2||(+ 760 kcal/kg)|
|MgCO3||+||Kcal||↔||MgO||+||CO2 ↑||(+ 723 kcal/kg)|
|CaCO3 MgCO3||+||Kcal||↔||CaO MgO||+||2CO2 ↑||(+ 723 kcal/kg)|
|CaO||+||H2O||↔||Ca(OH)2||+||Kcal||(- 273 kcal/kg)|
|CaO MgO||+||H2O||↔||Ca(OH)2 MgO||+||Kcal|
|CaO MgO||+||2H2O||↔||Ca(OH)2 Mg(OH)2||+||Kcal||(- 211 kcal/kg)|
The first group of reactions summarises the decarbonation process of the three main types of carbonates existing in nature.
They represent the reaction that, industrially, takes place in the limestone kiln when, starting from the mineral coming from the quarry and providing a prefixed quantity of heat, the dissociation of the carbonate (CaCO3) in an oxide (CaO) and in carbon dioxide (CO2) is obtained.
The reaction is endothermic and therefore requires a caloric input obtained by using a fuel.
It takes 760 kcal/kg to decarbonate CaCO3 and 723 kcal/kg for MgCO3.
Note: In the following, unless otherwise noted, we will refer to the case of calcium carbonate and/or lime oxide for simplicity of exposition, but in the case of the other types of components mentioned, the mechanisms by which the reactions take place are quite similar.
The dissociation of limestone takes place in five consecutive steps:
The main factors affecting the speed of decarbonation are:
The term hydration or slaking of lime refers to the process by which lime oxide (CaO) is transformed into lime hydroxide (Ca(OH)2).
From a chemical point of view, there is only one type of reaction, but from an industrial point of view, there are two different types of process.
One is the hydration process itself, in which the oxide is reacted with the stoichiometric amount of water (32% H2O) to obtain a hydroxide in the form of a powder containing a maximum of 1.5% free water.
The other is the slaking process where the lime oxide is made to react with a quantity of water much greater than the stoichiometric content, obtaining a lime hydroxide in suspension with very variable concentration values depending on the use required.
The chemical nature of the hydration reaction is very elementary, but the kinetics of the reaction, linked to crystallisation and agglomeration, is much more complex and depends not only on the physical and chemical characteristics of the oxide to be extinguished, but also on the way in which this simple reaction is carried out.
At temperatures below 350 °C, the oxide reacts completely with water, giving rise to an exothermic reaction with the development of 276 kcal/kg. CaO, at higher temperatures the reaction occurs in the opposite direction resulting in the separation of the reaction water.
Magnesium oxide is not very reactive with water and under normal conditions only 25% of it reacts.
In order to achieve complete hydration of MgO, it is necessary for the reaction to take place at a temperature above 100° C using pressurised equipment.
The main factors affecting hydration are:
In the industrial context, the hydration reaction is carried out using hydrators, in the case of the production of powdery hydroxide on a dry basis (stoichiometric water) and rotating drum slackers if the production of lime hydroxide in slurry form (excess water) is desired.
Both systems are schematically speaking mixers in which mechanical agitation produces very close contact between the lime oxide and the reaction water.
The recarbonation reaction is the last reaction in the life cycle of lime, it controls the process that transforms lime back into carbonate by reabsorbing CO2 from the environment and giving back to lime the chemical and physical characteristics that the original limestone had.
This is the fundamental reaction that allows lime, once applied to a wall, to set and harden.
Man has learnt to control this reaction for his use in the production of PCC (Precipitated Calcium Carbonate), a process in which by controlling the parameters influencing recarbonation it is possible to modify the grain size and morphology of the calcium carbonate produced.
At ambient conditions the recarbonation of lime is very modest, but from 290 °C up to 600 °C the speed of recarbonation increases rapidly, exponentially increasing the affinity of CaO with CO2.
The speed of this reaction is also significantly increased by both the specific surface area of the oxide and the speed with which the oxide is mixed with carbon dioxide. It can be seen that magnesium oxide has significantly longer recarbonation times than calcium oxide.
The presence of water, both in the form of moisture and steam, enables rapid recarbonation to take place even at ambient temperature and pressure, since water acts as a catalyst in the reaction.
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