Calcium sulfate

With judicious heating, gypsum converts to the partially dehydrated mineral called bassanite or plaster of Paris. This material has the formula CaSO₄·(nH₂O), where 0.5 ≤ n ≤ 0.8.[8] Temperatures between 100 and 150 °C (212–302 °F) are required to drive off the water within its structure. The details of the temperature and time depend on ambient humidity. Temperatures as high as 170 °C (338 °F) are used in industrial calcination, but at these temperatures γ-anhydrite begins to form. The heat energy delivered to the gypsum at this time (the heat of hydration) tends to go into driving off water (as water vapor) rather than increasing the temperature of the mineral, which rises slowly until the water is gone, then increases more rapidly. The equation for the partial dehydration is:

CaSO₄ · 2 H₂O   →   CaSO₄ · 1/2 H₂O + 1+1/2 H₂O↑

The endothermic property of this reaction is relevant to the performance of drywall, conferring fire resistance to residential and other structures. In a fire, the structure behind a sheet of drywall will remain relatively cool as water is lost from the gypsum, thus preventing (or substantially retarding) damage to the framing (through combustion of wood members or loss of strength of steel at high temperatures) and consequent structural collapse. But at higher temperatures, calcium sulfate will release oxygen and act as an oxidizing agent. This property is used in aluminothermy. In contrast to most minerals, which when rehydrated simply form liquid or semi-liquid pastes, or remain powdery, calcined gypsum has an unusual property: when mixed 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:

CaSO₄ · 1/2 H₂O + 1+1/2 H₂O   →   CaSO₄ · 2 H₂O

This reaction is exothermic and is responsible for the ease with which gypsum can be cast into various shapes including sheets (for drywall), sticks (for blackboard chalk), and molds (to immobilize broken bones, or for metal casting). Mixed with polymers, it has been used as a bone repair cement. Small amounts of calcined gypsum are added to earth to create strong structures directly from cast earth, an alternative to adobe (which loses its strength when wet). The conditions of dehydration can be changed to adjust the porosity of the hemihydrate, resulting in the so-called α- and β-hemihydrates (which are more or less chemically identical).

On heating to 180 °C (356 °F), the nearly water-free form, called γ-anhydrite (CaSO₄·nH₂O where n = 0 to 0.05) is produced. γ-Anhydrite reacts slowly with water to return to the dihydrate state, a property exploited in some commercial desiccants. On heating above 250 °C, the completely anhydrous form called β-anhydrite or "natural" anhydrite is formed. Natural anhydrite does not react with water, even over geological timescales, unless very finely ground.

The variable composition of the hemihydrate and γ-anhydrite, and their easy inter-conversion, is due to their nearly identical crystal structures containing "channels" that can accommodate variable amounts of water, or other small molecules such as methanol.

The calcium sulfate hydrates are used as a coagulant in products such as tofu.[9]

For the FDA, it is permitted in Cheese and Related Cheese Products; Cereal Flours; Bakery Products; Frozen Desserts; Artificial Sweeteners for Jelly & Preserves; Condiment Vegetables; and Condiment Tomatoes and some candies.[10]

It is known in the E number series as E516, and the UN's FAO knows it as a firming agent, a flour treatment agent, a sequestrant, and a leavening agent.[10]

Calcium sulfate has a long history of use in dentistry.[11] It has been used in bone regeneration as a graft material and graft binder (or extender) and as a barrier in guided bone tissue regeneration. It is a biocompatible material and is completely resorbed following implantation.[12] It does not evoke a significant host response and creates a calcium-rich milieu in the area of implantation.[13]

Drierite

When sold at the anhydrous state as a desiccant with a color-indicating agent under the name Drierite, it appears blue (anhydrous) or pink (hydrated) due to impregnation with cobalt(II) chloride, which functions as a moisture indicator.

Up to the 1970s, commercial quantities of sulfuric acid were produced in Whitehaven (Cumbria, UK) from anhydrous calcium sulfate. Upon being mixed with shale or marl, and roasted, the sulfate liberates sulfur dioxide gas, a precursor in sulfuric acid production, the reaction also produces calcium silicate, a mineral phase essential in cement clinker production.[14]

2 CaSO₄ + 2 SiO₂ → 2 CaSiO₃ + 2 SO₂ + O₂ [15]

The plant made sulfuric acid by the “Anhydrite Process” , in which cement clinker itself was a by-product. In this process, anhydrite (calcium sulfate) replaces limestone in a cement rawmix, and under reducing conditions, sulfur dioxide is evolved instead of carbon dioxide. The sulfur dioxide is converted to sulfuric acid by the Contact Process using a vanadium pentoxide catalyst. [16]

CaSO₄ + 2 C → CaS + 2CO₂

3 CaSO₄ + CaS + 2 SiO₂ → 2 Ca₂SiO₄ (belite) + 4 SO₂

3 CaSO₄ + CaS → 4 CaO + 4 SO₂

Ca₂SiO₄ + CaO → Ca₃OSiO₄ (alite)

2 SO₂ + O₂ → 2 SO₃ (in the presence of the catalyst vanadium pentoxide)

SO₃ + H2O → H₂SO₄ [17]

Because of its use in an expanding niche market, the Whitehaven plant continued to expand in a manner not shared by the other Anhydrite Process plants. The anhydrite mine opened on 11/1/1955, and the acid plant started on 14/11/1955. For a while in the early 1970s, it became the largest sulfuric acid plant in the UK, making about 13% of national production, and it was by far the largest Anhydrite Process plant ever built. [18]