Chemical elements
  Calcium
    Isotopes
    Energy
    Preparation
    Physical Properties
    Chemical Properties
      Calcium Hydride
      Calcium Subfluoride
      Calcium Fluoride
      Calcium Subchloride
      Calcium Chloride
      Calcium Bromide
      Calcium Subiodide
      Calcium Iodide
      Calcium Periodides
      Calcium Halides
      Calcium Perhalides
      Calcium Oxychloride
      Calcium Hypochlorite
      Bleaching Powder
      Calcium Chlorite
      Calcium Chlorate
      Calcium Perchlorate
      Calcium Oxybromide
      Calcium Hypobromite
      Brome Bleaching Powder
      Calcium Bromate
      Calcium Oxyiodide
      Calcium Hypoiodite
      Iodine Bleaching Powder
      Calcium Iodate
      Calcium Periodate
      Calcium Manganites
      Calcium Manganate
      Calcium Permanganate
      Calcium Oxide
      Caustic lime
      Calcium Suboxide
      Calcium Hydroxide
      Lime
      Calcium Peroxide
      Calcium Peroxyhydrates
      Calcium Tetroxide
      Calcium Monosulphide
      Calcium Hydrosulphide
      Calcium Polysulphides
      Calcium Hydroxyhydrosulphide
      Calcium Oxysulphides
      Calcium Thiosulphate
      Calcium Hyposulphite
      Calcium Sulphite
      Calcium Bisulphite
      Calcium Sulphite
      Calcium Dithionate
      Calcium Trithionate
      Calcium Sulphate
      Acid Calcium Sulphates
      Calcium Pyrosulphate
      Calcium Selenide
      Calcium Selenite
      Calcium Selenate
      Calcium Telluride
      Calcium Tellurite
      Calcium Tellurate
      Calcium Chromite
      Calcium Chromate
      Calcium Dichromate
      Calcium Tetrachromate
      Basic Calcium Chromate
      Calcium Chlorochromate
      Calcium Molybdate
      Calcium Tungstate
      Calcium Uranate
      Calcium Peruranate
      Calcium Nitride
      Calcium Azide
      Calcium Hexammoniate
      Calcium Ammonium
      Calcium Amide
      Calcium Imide
      Calcium Hydroxylamite
      Calcium Imidosulphonate
      Calcium Hyponitrite
      Calcium Nitrohydroxylaminate
      Calcium Nitrite
      Calcium Nitrate
      Basic Calcium Nitrates
      Calcium Phosphide
      Calcium Dihydrohypophosphite
      Calcium Hydrophosphite
      Neutral Calcium Phosphite
      Calcium Dihydrophosphite
      Acid Calcium Phosphite
      Neutral Calcium Hypophosphate
      Acid Calcium Hypophosphate
      Calcium Orthophosphates
      Calcium Pyro- Meta-phosphates
      Calcium Ultraphosphates
      Calcium Selenophosphate
      Basic Calcium Phosphates
      Phosphatic Fertilisers
      Calcium Arsenide
      Calcium Arsenites
      Calcium Arsenates
      Calcium Pyroarsenate
      Calcium Thioarsenites
      Calcium Thio-oxyarsenate
      Calcium Antimonide
      Calcium Antimonate
      Calcium Orthovanadate
      Calcium Pyrovanadate
      Calcium Metavanadate
      Calcium Pervanadate
      Calcium Pyro- Meta- niobates
      Calcium Pyro- Meta-tantalate
      Calcium Potassium Pertantalate
      Calcium Carbide
      Calcium Formate
      Calcium Acetate
      Calcium Oxalate
      Calcium Carbonate
      Calcium Bicarbonate
      Calcium Trithiocarbonate
      Calcium Perthiocarbonate
      Calcium Cyanide
      Calcium Oxycyanide
      Calcium Cyanamide
      Calcium Cyanate
      Calcium Cyanurate
      Calcium Thiocyanate
      Calcium Silicide
      Calcium Monosilicide
      Calcium Silicalcyanide
      Monocalcium Silicate
      Calcium Meta-silicate
      Calcium Orthosilicate
      Dicalcium Silicate
      Tricalcium Silicate
      Acid Calcium Silicate
      Calcium Fluosilicate
      Calcium Aluminates
      Monocalcium Aluminate
      Tricalcium Aluminate
      Pentacalcium Aluminate
      Calcium Stannate
      Calcium Chlorostannate
      Calcium Silicostannate
      Calcium Orthoplumbate
      Calcium Metaplumbate
      Acid Calcium Plumbate
      Calcium Metatitanate
      Calcium Fluotitanate
      Calcium Silicotitanate
      Calcium Zirconate
      Calcium Silicozirconate
      Calcium Boride
      Calcium Borates
      Calcium Silicoborate
      Calcium Borostannate
      Calcium Perborate
      Calcium Ferrate
    Glass
    Cement
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Calcium Sulphate, CaSO4






Calcium Sulphate, CaSO4, occurs naturally in two forms - gypsum and anhydrite. The former, crystallising with 2 molecules of water, is more frequently met than the latter, which, as its name indicates, is anhydrous. There are several varieties of gypsum: rock or massive gypsum, density 2.30 – 2.33, composed of minute crystals; large monoclinic crystals, often twinned, of selenite, which can be split into sheets and is the standard for the second degree of hardness in Mohs' scale; fibrous gypsum, or satin-spar; a compact, crystalline, slightly translucent form of rock gypsum, used for purposes of ornament and known as alabaster; and, finally, gypsite, formed by the evaporation of gypsiferous waters as a soft, incoherent, earthy deposit. Gypsum is often found interstratified with limestone and may have originated in the interaction of the limestone with soluble sulphates, with sulphuric acid resulting from the decomposition of sulphides, or with sulphurous vapours and solutions of volcanic origin. It is found in most countries. Anhydrite occurs as thin seams in rock salt, as large masses coated with gypsum produced by hydration of the anhydrite, or as a replacement of limestone. Local uplifts, and fracture or crumpling of rocks, are often caused by the hydration of subterranean beds of anhydrite.

Calcium sulphate is also a constituent of certain minerals: for example, glauberite, CaNa2(SO4)2; syngenite, CaK2(SO4)2.H2O; polyhalite, Ca2K2Mg(SO4)4.2H2O; penta-salt, Ca5K2(SO4)6.H2O; and krugite, Ca4K2Mg(SO4)6.2H2O. The conditions for the formation of these minerals in oceanic salt deposits have been studied by van't Hoff and his colleagues.

Calcium sulphate is largely responsible for the permanent hardness in natural waters.


Calcium Sulphate Dihydrate, CaSO4

2H2O.Calcium Sulphate Dihydrate, CaSO4.2H2O, is formed by the precipitation of soluble salts of calcium by alkali sulphates, or by neutralisation of the carbonate or oxide by dilute sulphuric acid. For technical purposes, however, natural gypsum is almost exclusively employed. In Great Britain alone, between 1903 and 1916, the average annual output was 250,000 tons.

In addition to the forms already mentioned, a hemihydrate, 2CaSO4.H2O, and certain allotropic modifications of anhydrite can be prepared. Although these are of considerable commercial importance and have been the object of much careful study, the literature on the subject is full of contradictions as to the conditions of their formation and the relations between them.

Calcium sulphate may also be precipitated in a gelatinous form by the action of sulphuric acid on a methyl alcoholic solution of calcium oxide, or by the addition of alcohol to an equal volume of a saturated solution of calcium sulphate. It contains more water than gypsum, but no combined alcohol. It is also more soluble than gypsum.

Calcium Sulphate Hemihydrate 2CaSO4

H2O.Calcium Sulphate Hemihydrate or Plaster of Paris, 2CaSO4.H2O. - When gypsum is heated between 100° and 200° C., a product is obtained corresponding to the formula 2CaSO4.H2O. This possesses the property of hydrating again and setting to a hard mass, which expands during solidification, thus giving a sharp impression of the mould in which it is placed. In reality there is a total contraction, made up of a preliminary contraction, followed by a smaller expansion. The reaction takes place with considerable heat evolution. From the fact that deposits of gypsum at Montmartre were first used for the production of this compound on an extensive scale, it received the name of plaster of Paris.

In the manufacture of plaster of Paris the maximum temperature employed in England is 110°-120° C., and in Germany 130° C. In America 180°-200° C. is often reached, although it may also be discharged at 120° C. The transition temperature from gypsum to the half-hydrate is 107° C., but this does not exclude the possibility of carrying out the dehydration at a still lower temperature, provided the vapour pressure is reduced by a current of dry air, or by the use of an open crucible. In England the gypsum is heated in an oven in which it is piled in lumps upon arches under which the fuel is burnt. In America the gypsum is often finely ground before being heated in pots or kettles 8-14 feet in diameter and 6-10 feet deep, and made of boiler plates with a convex iron or steel bottom. The gypsum is stirred. The agitation of the powder by the escaping water vapour is described as " boiling." This method produces a good quality of plaster, but it is slow and expensive and is being gradually replaced by the use of rotary heaters.

The hemihydrate forms rhombic crystals, which, through repeated twinning, appear hexagonal. The density is 2.63. On exposure to air, moisture is absorbed with gradual formation of gypsum. The hemihydrate has been found as a crystalline deposit in boilers working at two atmospheres pressure.

Davis stated that when freshly prepared crystals of gypsum are heated for three or four hours with water at 100° C., a half-hydrate, which sets very slowly, is obtained. This appears to be identical with a half-hydrate obtained by Le Chatelier in a similar way at 130°-150° C., and Davis thought that it might be a second modification. Probably it is merely in a less finely divided state than ordinary plaster.

The Setting of Plaster of Paris

Calcium sulphate is only very slightly soluble in water, but it has a great tendency to form supersaturated solutions. Le Chatelier attributed to this phenomenon the setting of plaster. He supposed that the partially dehydrated calcium sulphate dissolves, forming a solution which is supersaturated with respect to gypsum. From this solution is deposited hydrated sulphate as interlocking needles which offer great frictional resistance to disruption. The breaking of set plaster is not due to fracture of the crystals themselves, as in the case of metals, but to the overcoming of the forces of cohesion at the crystal surfaces. According to Davis, the plaster on setting first forms orthorhombic crystals which slowly change to monoclinic, the two changes corresponding with the observed contraction and expansion respectively. Recent views, however, require the existence of an intermediate colloidal state from which crystallisation takes place.

The rate of setting depends on the time and temperature of calcination of the gypsum. It is also greatly influenced by the presence of small quantities of foreign substances, some of which retard whilst others accelerate it. Rohland and others showed that those substances which increase the solubility of calcium sulphate also increase the rate of setting of plaster of Paris, and vice versa. Sulphates, chlorides, nitrates, bromides, and iodides of potassium, sodium, and ammonium, potassium hydroxides, sulphuric, hydrochloric, and nitric acids, sulphates of zinc, iron, and copper, calcium oxide, soap, potassium dichromate, ammonium fluoride, aluminium sulphate and chloride, lithium chloride and magnesium chloride accelerate the rate of setting. Boric acid, borax, sodium and potassium silicates, sodium, calcium, and magnesium carbonates, glycerol, and alcohol retard it.

Haddon, on rather insufficient grounds, appears to consider that Rohland's theory is not entirely justifiable.

Emulsoids, such as gum acacia, glue, etc., retard the setting in consequence of adsorption, and this fact is made use of by artists and modellers.

Anhydrous Calcium Sulphate

Natural anhydrite crystallises in the rhombic system and has a density of 2.8 – 3.0 and hardness 3.3 - 5. It gradually hydrates to gypsum in the presence of water, but does not set unless ground to an extremely fine powder. The hydration is influenced by foreign substances in the same way as the setting of plaster of Paris.

It may be prepared artificially by fusing calcium chloride with potassium sulphate and dissolving out the potassium chloride with water.

Two modifications of anhydrite are obtained by heating the hemi-hydrate. The first one, formed under 200° C., takes up water very quickly, probably forming the hemihydrate almost instantaneously and passing more slowly to gypsum. This variety differs from the natural form and is known as soluble anhydrite. It forms triclinic needles of density 2.45. Van't Hoff and his co-workers obtained it by dehydrating precipitated calcium sulphate in vacuo at 60°-90° C. over phosphorus pentoxide or sulphuric acid, or by shaking plaster of Paris with twenty times its weight of water at 100° C.

By heating soluble anhydrite to bright redness another compound is formed which hydrates very slowly and does not harden, and is probably identical with natural anhydrite. Hoppe-Seyler apparently obtained it by heating gypsum with a concentrated solution of sodium chloride in a sealed tube at 125°-130° C. A temperature a little above 30° C. is sufficient however.

At high temperatures an anhydrous form, capable of both hydration and setting, is obtained. It sets more slowly than plaster of Paris, but forms a very hard and highly resistant plaster which finds technical application under the name of Estrich gypsum or flooring plaster. The rate of setting can also be accelerated or retarded by suitable catalysts, but the same substances do not necessarily produce the same effect as in plaster of Paris.

At high temperatures dead-burnt gypsum, which will neither hydrate nor set, is also produced, and there are two rival theories as to the nature of Estrich gypsum and dead-burnt gypsum. According to van't Hoff, the rate of setting continuously falls as the temperature of burning is increased. If the temperature is raised so high that the setting and hydrating power disappear, the crystalline structure is lost and the plaster is then dead burnt. From the work on soluble and insoluble anhydrite, Estrich gypsum would thus appear to be a mixture of the two. The proportions, and, therefore, the setting power, depend on the temperature of calcination - dead-burnt gypsum being obtained when the temperature is high enough to transform all the soluble into insoluble anhydrite.

Rohland, on the other hand, considered that Estrich gypsum is formed at a higher temperature than dead-burnt gypsum. Hence arose the belief, denied, however, by van't Hoff, that the former might be a basic sulphate. Von Glasenapp favoured the latter view, and stated that Estrich gypsum may be produced at a temperature, namely 900°-1000° C., or even 1300° C., much higher than is usually considered suitable, and a considerable amount of calcium oxide may be present before the product is incapable of hardening. He observed glassy and crystalline portions; the first he regarded as a solid solution of lime in neutral sulphate, and the second as pure anhydrite.

Grengg studied the dehydration products of gypsum under the microscope, and also concluded that the formation of dead-burnt gypsum takes place first. He regarded it as completely dehydrated calcium sulphate, probably identical with natural anhydrite. He also regarded calcium oxide as one of the constituents of Estrich gypsum, and, like von Glasenapp, found a combination of a glassy and crystalline structure, but supposed both to be solid solutions of calcium oxide in the sulphate - the proportions and concentrations of these varying with the time and temperature of heating. Observations on the temperature of formation of the oxide seem to cast doubt on the basic sulphate theory. When calcium sulphate is heated to constant weight at 1000° C. only 0.21 per cent, of oxide is formed, at 1300° C. 3.0 per cent., and at 1375° C. 98.67 per cent., the salt then melting. The rate becomes much more rapid above 1380° C.

It is possible that the supposed modifications of anhydrite owe their differences largely to differences merely in the fineness of their particles.

When calcium sulphate is strongly heated there is a transition point at 1193° C. The α-form stable above the transition point is completely miscible with α-strontium sulphate, and is, therefore, isomorphous with it and with α-barium sulphate.

The melting-point of calcium sulphate is 1450° C. According to Calcagni and Mancini, pure calcium sulphate decomposes so rapidly at 1000° C. that the melting-point cannot be determined directly, but extrapolation from the freezing-point curve of sodium and calcium sulphates gave 1375° C.

Gallo suggested a scheme for the estimation of the different modifications of calcium sulphate in plaster of Paris which depends on the amount of water given up or absorbed under different conditions.

The Vapour Pressure and Solubility Relations between the different Modifications of Calcium Sulphate

The equilibrium relations between gypsum, the half-hydrate, soluble anhydrite, and natural anhydrite, between 0° and 100° C., have been studied by van't Hoff and his co-workers. According to the phase rule, the system calcium sulphate and water has one degree of freedom when three phases are present - for example, solid, solution, and vapour, or two solids and vapour; but reaches an invariant point when four phases appear - for example, two solids, solution, and vapour. Van 't Hoff followed the change of vapour pressure with composition and obtained the results recorded in the following table: -

Temp.,° C.Vapour Pressure of Water or Saturated Solution of Dihydrate (I.). mm
Half-hydrate (II.). mmSol. Anhyd. (III.). mmNat. Anhyd. (IV.). mm
04.571.171.522.06
56.511.842.343.17
109.142.783.554.79
1512.74.215.297.12
2017.46.247.7710.5
3031.512.716.121.6
4054.926.331.842.5
5092.050.059.979.7
6014991.4108143
65187122143175 (at 63.5° C.)
75289210245
90526446513
95634565588 (at 93° C.)
105906888
1101073971 (at 107° C.)


In the presence of concentrated solutions of other salts, such as sodium or magnesium chlorides, these transition points are lowered. For example, a saturated sodium chloride solution lowers the temperature of equilibrium between gypsum and natural anhydrite to 36° C.

This explains the production of natural anhydrite from gypsum by heating with concentrated sodium chloride solution a little above 30° C. It also explains the existence of anhydrite in contact with beds of rock salt. The inter-relations of gypsum, natural anhydrite, soluble anhydrite, and the half-hydrate, account for the fact that only the two former are found in nature, since the two latter are always unstable with respect to both the others.

The energies of transformation of the different modifications into gypsum are given by van't Hoff as follows: -

From half-hydrate 737-6.88t small cal. (t in ° C.).
From soluble anhydrite 602-6.48t small cal. (t in ° C.).
From natural anhydrite 435-6.84t small cal. (t in ° C.).

The solubilities of gypsum and soluble and natural anhydrite were studied by Melcher by means of electrical conductivity determinations. His results are combined with some of other investigators in the following table: -

FromSolubility in Milli-equivalents per Litre at Temperatures, ° C.
0183040506575100156200218
Gypsum25.929.530.730.830.028.327.123.3
Sol. Anhyd.22.86.42.3
Nat. Anhyd9.22.70.90.7


Solubility of calcium sulphate
Solubility of calcium sulphate
The solubility curves for the three compounds are given in Fig. The transition points as found by van't Hoff are represented by the points where the solubility curves for the two anhydrites meet that for gypsum. Soluble anhydrite is evidently a meta-stable form. No determinations, other than qualitative ones, have been made for plaster of Paris, but presumably the solubility curve for this substance would show higher values than soluble anhydrite at the same temperature, and would cut the curve for gypsum at about 107° C. Janecke confirmed van't Hoff's observation that gypsum melts at 107° C., giving the hemihydrate.

The solubility curve for gypsum shows a maximum at about 40° C., the solubility being 30.826 milli-equivalents per litre. Earlier investigators found a slightly lower temperature, 38° C.

An observation of Hulett's explains the difficulty found in arriving at concordant solubility values. By shaking powdered gypsum with water he obtained a supersaturated solution which required seventeen days to reach a state of final equilibrium. The degree of supersaturation depends on the size of the particles.

There is apparently a considerable difference between the rates of solubility at the different surfaces of a selenite crystal, although the actual solubilities are naturally the same in each case.

Gypsum is almost insoluble in alcohol but slightly soluble in glycerol, 0.0072 gram-molecule per litre, the solubility increasing with temperature.

The solubility in water is greatly increased by the presence of acids. The following values have been obtained: -

Cc. containing 1 mgm.- equivalent of Acid.Increase in Solubility at 20° C. in mgm. per mgm.-equivalent of Acid.
HCl.HNO3.CH2ClCOOH.CH3COOH.
0.57.619.27
111.5113.10.410.31
215.7520.40.17
1023.023.4


The solubility in different salt solutions has been studied by a number of investigators.

Calcium sulphate gives a brilliant green cathode-ray fluorescence if a trace of manganese be present, and bismuth make it orange-red.

Uses of Calcium Sulphate

In the form on either plaster of Paris or Estrich gypsum, calcium sulphate forms the principal ingredients of many of the cements and wall-plasters on the market - for example, Keene's, Martin's, and Parian cement. A retarder, such as glue or blood, or an accelerator, alum, sodium sulphate, etc., is mixed with it. Stucco is a hard plaster which can be polished, and is obtained by treating plaster of Paris with a lukewarm solution of size, fish-glue, or gum-arabic, or by mixing with freshly slaked lime and dipping the casting in a strong solution of magnesium sulphate, sodium silicate, or alum. Fibrous plaster for temporary buildings is made from plaster toughened by tow, asbestos, or slag wool, and backed by coarse canvas. Gypsum boards and tiles are also made. The tensile strength of plaster is affected by the conditions of gauging, or wetting, and by the added constituents. Gypsum is often put in Portland cement in small quantity to act as a retarder. It is sometimes used in the natural form as building-stone.

There seems to be some doubt as to whether or not calcium sulphate is of any value as a fertiliser. It is probably more advantageous to employ it along with basic rather than with acid fertilisers. It has been suggested that it increases the availability of potash in the soil, but this has not, apparently, been confirmed.

Various patents were taken out in Germany during the war for the manufacture of sulphuric acid from calcium sulphate. The methods proposed consist chiefly in the reduction by coal or hydrocarbons, and the removal and oxidation of sulphuretted hydrogen, or in the decomposition of the sulphate by silicates.

The production of ammonium sulphate from gypsum for use as a fertiliser, has also been suggested.

Calcium sulphate is an ingredient of certain painters' colours, and is employed in paper manufacture for weighting the paper. Plaster of Paris is used in the making of moulds for many purposes, for example, for rubber stamps, pottery, terra cotta, and special foundry castings, also for making statuary, relief maps and models, surgical casts, relief decorations on walls and ceilings, for bedding plate-glass for polishing, and for many other purposes.
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