Know The Importance Of This Chemical Compound In The Food Industry As Well As Its Other Uses

About this Compound – Calcium Sulfate

Calcium sulfate which is also known as calcium sulphate and related hydrates is an inorganic compound having the formula CaSO4.¹https://en.wikipedia.org/wiki/Calcium_sulfate It is utilized as a desiccant in the form of an anhydrite (the anhydrous form).

One hydrate is more well recognized as Plaster of Paris, while another naturally occurs as that of the mineral gypsum. It has several applications in the industry. All states are white solids which are water-insoluble. Water becomes permanently hardened due to calcium sulfate.

CaSO4 (calcium sulfate) is found naturally in calcium salt. It is most well recognized in its dihydrate state, CaSO42H2O, as gypsum, a white or colorless powder.

The sulfate is used as a soil conditioner since it is uncalcined gypsum. Calcined gypsum is used in the production of lath, wallboard, tile, and a variety of plasters.

When gypsum is warmed to around 120 °C (250 °F), it sheds three-quarters of the water and transforms into the hemihydrate CaSO4.H2O, plaster of Paris.

Plaster of Paris, when combined with water, maybe sculpted into forms before hardening by recrystallizing it to a dihydrate state. Groundwater may include calcium sulfate, which causes hardness or stiffness that cannot be eliminated by boiling.

Watch the video below to know about calcium sulphate:

Crystallographic Structures and Hydration States of Calcium Sulfate

The molecule has three hydration states that correlate to various crystallographic structures and minerals:

• CaSO4 (anhydrite): In anhydrous form; the structure is similar to the structure of zirconium orthosilicate (zircon): Ca2+ has an 8-coordinate, SO2-4 has a tetrahedral structure, and O has a 3-coordinate.
• CaSO4-2H2O: Dihydrate {gypsum and selenite (mineral)}.
• CaSO4.H2O: Hemihydrate, often known as the Plaster of Paris. Certain hemihydrates are occasionally distinguished: α-hemihydrate as well as β-hemihydrate.

Medical Subject Headings (MeSH) Entry Terms of Calcium Sulfate

• Alabaster
• Anhydrous Sulfate of Lime
• Artificial Dental Stone
• Calcium Sulfate
• Calcium Sulfate (1:1), Dihydrate
• Calcium Sulfate (1:1), Hemihydrate
• Calcium Sulfate (2:1)
• Calcium Sulfate Dihydrate
• Calcium Sulfate, Anhydrous
• Calcium Sulfate, Dihydrate
• Calcium Sulfate, Hemihydrate
• Calcium Sulphate
• Dental Gypsum
• Dental Stone, Artificial
• Drierite
• Gypsite
• Gypsum
• Gypsum, Dental

Physical and Chemical Properties of Calcium Sulfate

Experimental Properties of Calcium Sulphate

Physical Description

• Calcium sulfate appears as odorless, white powder or colorless, crystalline solid.²https://www.britannica.com/science/calcium-sulfate Crystals sometimes have a blue, gray or reddish tinge or can be brick red. Density: 2.96 g cm-3.
• Dry Powder; Dry Powder, Other Solid; Dry Powder, Pellets or Large Crystals; Liquid; Liquid, Other Solid; NKRA; Other Solid; Pellets or Large Crystals; Pellets or Large Crystals, Liquid; Pellets or Large Crystals, Water or Solvent Wet Solid; Water or Solvent Wet Solid; Water or Solvent Wet Solid, Other Solid
• Fine, white to slightly yellowish-white odorless powder
• Odorless, white powder or colorless, crystalline solid. [Note: May have blue, gray, or reddish tinge
• White hygroscopic powder or crystalline powder.

Production and Occurrence of Calcium Sulphate

The main sources of calcium sulfate are naturally occurring gypsum and anhydrite, which occur at many locations worldwide as evaporites. These may be extracted by open-cast quarrying or by deep mining. World production of natural gypsum is around 127 million tons per annum.

In addition to natural sources, calcium sulfate is produced as a by-product in a number of processes:

In flue-gas desulfurization, exhaust gasses from fossil-fuel power stations and other processes (e.g., cement manufacture) are scrubbed to reduce their sulfur oxide content, by injecting finely ground limestone: SO2 + 0.5 O2 + CaCO3 → CaSO4 + CO2

Related sulfur-trapping methods use lime and some produce an impure calcium sulfite, which oxidizes on storage to calcium sulfate.

In the production of phosphoric acid from phosphate rock, calcium phosphate is treated with sulfuric acid, and calcium sulfate precipitates. The product, called phosphogypsum is often contaminated with impurities making its use uneconomic.

In the production of hydrogen fluoride, calcium fluoride is treated with sulfuric acid, precipitating calcium sulfate.

In the refining of zinc, solutions of zinc sulfate are treated with hydrated lime to co-precipitate heavy metals such as barium.

Calcium sulfate can also be recovered and reused from scrap drywall at construction sites.

These precipitation processes tend to concentrate radioactive elements in the calcium sulfate product. This issue is particular with the phosphate by-product since phosphate ores naturally contain uranium and its decay products such as radium-226, lead-210, and polonium-210.

Extraction of uranium from phosphorus ores can be economical on its own depending on prices in the uranium market or the separation of uranium can be mandated by environmental legislation and its sale is used to recover part of the cost of the process.

Calcium sulfate is also a common component of fouling deposits in industrial heat exchangers because its solubility decreases with increasing temperature (see the specific section on the retrograde solubility below).

The Retrograde Solubility of Calcium Sulfate

The dissolution of the different crystalline phases of calcium sulfate in water is exothermic and releases heat (decrease in Enthalpy: ΔH < 0). As an immediate consequence, to proceed, the dissolution reaction needs to evacuate this heat that can be considered as a product of the reaction.

If the system is cooled, the dissolution equilibrium will evolve towards the right according to the Le Chatelier principle and calcium sulfate will dissolve more easily. Thus, the solubility of calcium sulfate increases as the temperature decreases and vice versa.

If the temperature of the system is raised, the reaction heat cannot dissipate and the equilibrium will regress towards the left according to Le Chatelier principle. The solubility of calcium sulfate decreases as temperature increases.

This counter-intuitive solubility behavior is called retrograde solubility. It is less common than for most of the salts whose dissolution reaction is endothermic (i.e., the reaction consumes heat: increase in Enthalpy: ΔH > 0) and whose solubility increases with temperature.

Another calcium compound, calcium hydroxide (Ca(OH)2, portlandite) also exhibits a retrograde solubility for the same thermodynamic reason: because its dissolution reaction is also exothermic and releases heat.

So, to dissolve the maximum amount of calcium sulfate or calcium hydroxide in water, it is necessary to cool the solution down close to its freezing point instead of increasing its temperature.

Temperature dependence of the solubility of calcium sulfate (3 phases) in pure water.

The retrograde solubility of calcium sulfate is also responsible for its precipitation in the hottest zone of heating systems and for its contribution to the formation of scale in boilers along with the precipitation of calcium carbonate whose solubility also decreases when CO2 degasses from hot water or can escape out of the system.

The Uses of Calcium Sulfate

Calcium sulfate is mostly used to make stucco and Plaster of Paris. These applications take advantage of the fact that powdered and calcined calcium sulfate produces a molded paste after hydration and solidifies as crystalline calcium sulfate dihydrate.

It is also helpful that calcium sulfate is weakly water soluble and does not rapidly dissolve after solidification when in contact with water.

Uses of calcium sulfate in hydration and dehydration reactions

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.

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 the 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 the 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.

Uses of calcium Sulfate in the food industry

Calcium sulfate hydrates are used as a coagulant in products such as tofu.³https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-sulfate 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.

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.

Use of calcium sulfate in dentistry

Calcium sulfate has a long history of use in dentistry. 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. It does not evoke a significant host response and creates a calcium-rich milieu in the area of implantation.

Therapeutic Uses of Calcium Sulfate

The use of calcium sulfate (plaster of Paris) has been advocated to repair bony defects because of its unique capability of stimulating osteogenesis. Plaster of Paris can be used as a bony alloplastic and it can be analyzed histologically.

Sinus roentgenograms and technetium Tc 99m medronate bone scanning further support the use of plaster of Paris as an alloplastic and assess its osteogenic capacity when implanted in the frontal sinus of dogs; complete bone regeneration was demonstrated in six dogs within four to six months.

The use of plaster of Paris for bone reconstruction in the head and neck can be applied in surgery. The experience with plaster of Paris to date although limited shows it to be safe and highly encouraging as an effective bone allograft.

Other Uses of calcium sulfate

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.

2 CaSO4 + 2 SiO2 → 2 CaSiO3 + 2 SO2 + O2

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 raw mix, 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.

• CaSO4 + 2 Cas → 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

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.

Toxicological Information of Calcium Sulfate – Side Effects of Calcium Sulphate

Exposure Routes: Inhalation, skin, and/or eye contact

Symptoms include:

1. Irritation of the eyes, skin, and upper respiratory system
2. Conjunctivitis
3. Rhinitis
4. Epistaxis (nosebleed)
5. Inhalation Symptoms 
6. Cough

• Eye Symptoms: Redness
• Ingestion Symptoms: Abdominal pain

Toxicological studies indicate the following:

• Gypsum: Gypsum dust has an irritant action on mucous membranes of the respiratory tract & eyes, and there have been reports of conjunctivitis, chronic rhinitis, laryngitis, pharyngitis, impaired sense of smell & taste, bleeding from the nose, & reactions of tracheal & bronchial membranes in exposed workers.
• Plaster of Paris: Because it hardens quickly after absorbing moisture, its ingestion may result in obstruction, particularly at the pylorus. To delay "setting," drink glycerin or gelatin solutions, or large volumes of water. Surgical relief may be necessary.
• Calcium sulfate caused no lung disease in calcium sulfate miners.

All right, guys, that is it for now for calcium sulfate. I hope Healthsoothe answered any questions you had concerning calcium sulphate. 

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Additional resources and citations

• 1
  https://en.wikipedia.org/wiki/Calcium_sulfate
• 2
  https://www.britannica.com/science/calcium-sulfate
• 3
  https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-sulfate

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