Food & Beverages
The inorganic chemical with the formula CaSO4 is calcium sulphate, often known as calcium sulphate, and its associated hydrates. It functions as a desiccant when it is present as -anhydrite (the anhydrous form). One of these hydrates is more well known as plaster of Paris, whereas the other is found naturally as the mineral gypsum. In industry, it is put to a lot of different uses. All forms are insoluble in water and are white solids.  Permanent hardness in water is caused by calcium sulphate.
Hydration conditions and crystal formations
There are three stages of hydration for the chemical, which correspond to various mineral crystal structures and crystallographic structures:
Anhydrous condition of CaSO4 (anhydrite). The structure is similar to zirconium orthosilicate (zircon), with Ca2+ having an 8-coordinate, SO24 having a tetrahedral shape, and O having a 3-coordinate.
Gypsum with selenite (minerals): dihydrate of CaSO42H2O.
Plaster of Paris is another name for the mineral hemihydrate CaSO41/2H2O (bassanite). There are two specific hemihydrates that can sometimes be distinguished: – and -hemihydrate.
Stucco and plaster of Paris are the two principal products made using calcium sulphate. These uses take advantage of the fact that calcium sulphate that has been powdered and calcined creates a moldable paste upon hydration and hardens as crystalline calcium sulphate dihydrate. It is also useful that calcium sulphate is weakly soluble in water and does not readily dissolve in contact with water after it has solidified.
Hydration and dehydration reactions
Gypsum can be heated carefully to transform it into the mineral known as bassanite or plaster of Paris, which is partially dehydrated. CaSO4(nH2O) is the formula for this substance, where 0.5–0.8 is the value of n.  Its structure must reach temperatures between 100 and 150 °C (212–302 °F) to drive off the water. Ambient humidity affects the time and temperature data. In commercial calcination, temperatures as high as 170 °C (338 °F) are employed, but at these temperatures, -anhydrite starts to develop. When heat is applied to gypsum at this point (the heat of hydration), it usually evaporates water (as water vapour) rather than raising the mineral’s temperature, which climbs gradually until the water is gone and then rises more quickly. The equation for the partial dehydration is:
CaSO4 · 2 H2O → CaSO4 · 1/2 H2O + 11/2 H2O↑
The endothermic quality of this reaction is essential to how well drywall performs in providing fire resistance to homes and other structures. In a fire, the area beneath a layer of drywall will stay relatively cool as water is vaporised from the gypsum, preventing (or significantly delaying) damage to the framing and subsequent structural collapse (due to combustion of wood components or loss of steel strength at high temperatures). However, calcium sulphate will release oxygen and function as an oxidising agent at higher temperatures. Aluminothermy makes use of this quality. Calcined gypsum has an unusual property: when combined with water at normal (ambient) temperatures, it quickly reverts chemically to the preferred dihydrate form, while physically “setting” to form a rigid and relatively strong gypsum crystal lattice. This is in contrast to most minerals, which when rehydrated simply form liquid or semi-liquid pastes, or remain powdery.
CaSO4 1 / 2 H2O CaSO4 2 H2O + 3 / 2 H2O
The ease with which gypsum may be formed into different shapes, such as sheets (for drywall), sticks (for blackboard chalk), and moulds, is due to this exothermic process (to immobilise broken bones, or for metal casting). It has been utilised as a cement for bone healing when combined with polymers. As an alternative to adobe, cast earth is strengthened by the addition of small amounts of calcined gypsum (which loses its strength when wet). The porosity of the hemihydrate can be adjusted by altering the circumstances of dehydration, giving rise to the so-called – and -hemihydrates (which are more or less chemically identical).
The almost water-free form, known as -anhydrite (CaSO4nH2O, where n = 0 to 0.05), is generated when heating to 180 °C (356 °F). Some commercial desiccants take advantage of the fact that anhydrite slowly reacts with water to revert to the dihydrate form. The totally anhydrous form, often known as “natural” anhydrite, is created when a substance is heated over 250 °C. Even across geological timeframes, natural anhydrite does not react with water unless it is extremely finely powdered.
Since their virtually identical crystal structures feature “channels” that can hold varying amounts of water or other tiny molecules like methanol, the hemihydrate and -anhydrite’s varied composition and ease of interconversion are caused by this.
Products like tofu use calcium sulphate hydrates as a coagulant.
According to the FDA, it is allowed in cheese, cheese-related products, cereal flours, bakery goods, frozen desserts, artificial sweeteners for jam and preserves, condiment vegetables, condiment tomatoes, and some sweets. It is identified as E516 in the E number series and is also referred to as a firming agent, flour treatment agent, sequestrant, and leavening agent by the UN’s FAO.
The use of calcium sulphate in dentistry is not new. It has been utilised in directed bone tissue regeneration as a barrier and as a graft material and graft binder (or extender). It is a biocompatible substance that totally dissolves after implantation. It induces a minimal host reaction and produces a calcium-rich environment near the site of implantation.
The numerous crystalline forms of calcium sulphate dissolve in water in an exothermic process that produces heat (reduction in enthalpy: H 0). As a direct result, in order to continue, the dissolving process needs to expel this heat, which is a reaction byproduct. Calcium sulphate will dissolve more readily if the system is chilled because the dissolving equilibrium will move towards the right in accordance with the Le Chatelier principle. As a result, calcium sulphate becomes more and more soluble as temperature drops. Le Chatelier’s principle states that if the system’s temperature is increased, the reaction heat cannot evaporate and the equilibrium will regress to the left. As temperature rises, calcium sulfate’s solubility declines. Retrograde solubility is the name for this counterintuitive solubility behaviour. It is less frequent than for the majority of salts whose solubility increases with temperature and whose dissolving process is endothermic (i.e., the reaction consumes heat: increase in Enthalpy: H > 0). Because its dissolution reaction is equally exothermic and generates heat, another calcium compound, calcium hydroxide (Ca(OH)2, portlandite), also exhibits retrograde solubility for the same thermodynamic purpose. Therefore, rather than raising the temperature of the solution, it is required to cool it down close to its freezing point in order to dissolve the greatest quantity of calcium sulphate or calcium hydroxide in water.
The precipitation of calcium carbonate, whose solubility also decreases when CO2 degasses from hot water or can escape out of the system, and calcium sulphate are both caused by the retrograde solubility of calcium sulphate. Calcium sulphate precipitates in the hotter zone of heating systems and contributes to the formation of scale in boilers.