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Calcium Carbonate, CaCO3
Calcium Carbonate, CaCO3 is one of the most frequently occurring compounds of calcium in both inorganic and organic nature. It is often found in large masses almost chemically pure. Chalk, limestone, marble, calcspar, coral, and the shells of molluscs are the principal sources. It is also found as dolomite, in which some of the calcium is replaced by magnesium, and as the double calcium lead carbonate, plumbocalcite.
Two crystalline forms undoubtedly exist. The most widely distributed is the hexagonal form, calcite, calcspar, or Iceland spar, showing marked double refraction. Aragonite in rhombic crystals is much rarer. Other forms have been described - conchite, ktypeite, or vaterite, and lublinite. According to more recent investigations, however, these latter are identical with one or other of the two well- known forms and a third form, μ-calcium carbonate, consisting of microscopic hexagonal plates of density 2.54, which is obtained by precipitation at 60° C., but is always contaminated with the other two. Another form, α-calcium carbonate, only stable at very high temperatures and with optical properties apparently identical with those of calcite, has also been mentioned. The apparently amorphous varieties probably consist of minute crystals of calcite.
Marble and limestone consist of small interlocking calcite crystals.
Calcite and aragonite form one of the earliest recognised cases of polymorphism. Both are anhydrous calcium carbonate, but, besides the difference in crystalline form already noted, there is a difference in density, which is 2.715 for calcite and 2.934 for aragonite. Calcite is isomorphous with sodium nitrate, and aragonite with strontianite. There is an extensive literature of the mutual relations of calcite and aragonite, the former being the stable variety under ordinary conditions. When calcium carbonate is formed by precipitation, it is very probable that, in accordance with Ostwald's rule, one of the unstable forms, μ-calcium carbonate or aragonite, is first obtained, and that this changes with more or less rapidity into calcite. In some cases, owing to the high temperature or to the catalysing effect of the salts present, the velocity of transformation may be so great that ealeite appears to be the first precipitated.
In the dry state aragonite is not changed into calcite with appreciable rapidity until a temperature above 400° C. is reached. At 470° C. the change takes place in a few minutes. If under any conditions aragonite can be the stable form, it must be at very low temperatures. The transition temperature has been given as - 43°±5° C., rising 1° C. per 40 atmospheres increase of pressure. It can only be in the coldest regions of the earth, therefore, that natural aragonite is the stable form. The following are probably the main conditions under which aragonite forms in nature: (a) through organic agencies, for example shells; (b) by deposition from hot springs; (c) when an isomorphous carbonate is present to serve as a nucleus; (d) by chemical precipitation in certain saline waters in the presence of sulphate. The preservation of aragonite in any but recent geological formations may depend on its having been kept dry, or on the presence of impurities, which are more frequent in aragonite than in calcite and which reduce the rate of transformation by water by reducing the solubility below that of calcite. In older rocks, fossil shells of such a type that they were presumably aragonite originally, now consist of calcite. The change from aragonite to calcite is accompanied by a 10 per cent, increase in volume. The heat of transformation is so small, however, that there appears to be some doubt as to whether it is positive or negative. It is possible that it may be positive at ordinary temperatures and negative at higher ones.
μ-Calcium carbonate is too unstable to be found as a mineral, and α-calcium carbonate is only stable above 970° C., the temperature of transition from calcite.
A number of colour tests have been suggested for distinguishing aragonite from calcite in minerals. These depend on the different reactivities of the two modifications towards different salt solutions.
Calcium carbonate may be prepared by precipitation of a soluble calcium salt with an alkali carbonate, or by neutralisation of calcium hydroxide solution with carbon dioxide. The latter is not an instantaneous process but may take a considerable time before it is complete. It is probably governed by the rate of hydration of carbon dioxide.
By fusion of the precipitated carbonate with a mixture of sodium and potassium chlorides, calcite in a visibly crystalline form may be obtained.
The specific heat of calcium carbonate at ordinary temperatures is 0.2027. Regnault gave the following values for the specific heats of the different mineral forms at 99° C.: -
Iceland spar – 0.20858
Aragonite – 0.20850
White marble – 0.21585
Grey marble – 0.20989
White chalk – 0.21485
Calcite is phosphorescent under the cathode rays, and some specimens also when heated, or on insolation.
On heating calcium carbonate it decomposes into calcium oxide and carbon dioxide. The reaction is reversible at high temperatures, and for every temperature there is a corresponding equilibrium pressure of carbon dioxide. This behaviour is to be expected if the system consists of two solid phases, calcium oxide and calcium carbonate, and a vapour phase, carbon dioxide. At low temperatures dissociation and re-absorption proceed very slowly, and depend very much on the condition of the solid phases. Jolibois and Bouvier found that the reaction is not completely reversible when pure calcium carbonate alone forms the starting-point. A mixture of calcium oxide and calcium carbonate must be used. Probably the difficulty is of purely mechanical origin. These authors regarded it as an indication that a solid solution of calcium oxide in calcium carbonate is formed.
The following values have been given for the dissociation pressures of calcium carbonate at different temperatures: -
The temperature of dissociation at atmospheric pressure, 760 mm., according to Riesenfeld is 908°±5° C. Hedvall found 913°-923° C., and Johnston 898° C. The latter obtained the following equation for the pressure-temperature curve: -
log10p = - 9840/T + 1.1log10T – 0.0012T + 8.882, where p is the pressure in mm. mercury and T the absolute temperature.
Since aragonite is the unstable form, it should have a higher dissociation pressure than calcite, but at the temperature at which rapid transformation into calcite begins, the pressures are only of the order of 0.01 mm. and differences cannot be determined with accuracy.
The molecular heat of formation of precipitated calcium carbonate from calcium oxide and carbon dioxide is 43.3 Cal., of calcspar 42.0 Cal., of aragonite 42.6 Cal., and of the carbonate at the temperature of decomposition under a pressure of 1 atmosphere 34.76 Cal.
Marble is generally said to be of igneotis origin, and Hall claimed to have obtained marble by the fusion of chalk, but the experiments were repeated by later investigators without success. Recently a substance of the characteristic appearance of marble has been obtained by heating anhydrous calcium chloride with hydrated sodium carbonate, or hydrated chloride with anhydrous carbonate, for eight hours at 300°-305° C. under 24 atmospheres pressure; or by saturating with carbon dioxide, at room temperature, a paste of precipitated chalk with saturated sodium chloride solution, and heating for eight hours at 300°-305° C. under a pressure of 27 atmospheres. Le Chatelier supposed that he had obtained evidence of fusion of calcium carbonate, with or without pressure, at 1020° C., but this was questioned by Joannis, and also by Boeke, who found that it partially dissociated but did not melt even when heated to 1400°-1500° C., under a pressure of carbon dioxide of 30 atmospheres, although the original fine powder became crystalline. In later experiments, however, the latter succeeded in fusing Iceland spar at 1289° C. in carbon dioxide at 110 atmospheres pressure, but the work of Smith and Adams seems to indicate that the melting-point of calcium carbonate lies a little above 1340° C. under a pressure of about 1050 atmospheres. An eutectic mixture containing 50 per cent, of calcium oxide fuses at 1240°±1° C. under about 40 atmospheres pressure.
Calcium carbonate is only very slightly soluble in water, but the solubility is greatly influenced by the concentration of carbon dioxide in solution, and this in its turn depends on the concentration or partial pressure of the gas in the vapour space above the solution. This influence is explained by the hydrolysis of calcium carbonate by water, forming the more soluble calcium hydroxide when the partial pressure of carbon dioxide is low, and by the formation of calcium bicarbonate, which is also more soluble than the carbonate, when the partial pressure of carbon dioxide is high. After prolonged boiling of a suspension of calcium carbonate, the substance passing into solution is practically all calcium hydroxide. By passing a very rapid stream of carbon dioxide through saturated lime-water at 15° C., a concentration of 2.29 grm. of calcium carbonate may be reached, but this only persists for a short time.
A study of the solubility of calcium carbonate is, therefore, a study of the equilibrium conditions in the ternary system, calcium oxide: carbon dioxide: water, and for any one temperature a pressure-concentration curve, consisting of three branches, might be drawn, as shown in Fig., which gives a general, but not an accurate, indication of the course of the curve - the low-pressure part of the curve being on a much larger scale than the high-pressure part.
The branch AB represents the slightly increasing concentration of the solution in contact with solid calcium hydroxide for partial pressures up to about 10-14 atmosphere, at which point, represented by B on the diagram, the solid phase, calcium carbonate, appears. The solubility then falls with increasing pressure to a minimum value at about 3.8×10-7 atmosphere, as calculated by Johnston and Williamson from the solubility product. The solubility rises again, rapidly at first and then more gradually, until, at about 15 atmospheres pressure, marked by point C, a fresh solid phase, calcium bicarbonate, should appear. The line CD shows a slightly decreasing concentration of bicarbonate with further increase of pressure.
It is probable that even at the minimum concentration of calcium ions in solution, which is the maximum of carbonate ions, more than 50 per cent, of the negative ions still consist of hydroxyl and bicarbonate ions.
The solubility of calcium carbonate under normal atmospheric conditions, at a partial pressure of carbon dioxide of 3×10-4 atmospheres, is 63 mgm. per litre. The failure of some investigators to state the partial pressures under which their experiments were conducted makes difficult any comparison between the determinations made. It is also possible that in some cases equilibrium was not reached.
The solubility of calcium carbonate decreases with rise of temperature. The following values have been obtained for the solubility of calcite at different temperatures under a partial pressure of 3.2×10-4 atmospheres: -
The solubility of aragonite is greater than that of calcite, in accordance with the general relationship between stable and unstable forms. The ratio of solubilities has an average value of 1.07, but there is a tendency for this to decrease as the temperature falls, which suggests that aragonite becomes the stable form at low temperatures.
A convenient method of expressing the solubility is by the solubility product constant, [Ca••]×[CO3''], from which the saturation concentration for any given partial pressure of carbon dioxide can then be calculated. Certain assumptions must be made as to the amount of free and hydrated carbon dioxide in solution, the degree of ionisation of the different electrolytes, and so on. Values for calcite which are of the same order of magnitude have been calculated by different authors and are recorded in the table on the following page.
The last value was obtained by recalculation from the data of McCoy and Smith, assuming that only a small part of the dissolved carbon dioxide is combined with water, and that, consequently, carbonic acid is a stronger acid than has previously been supposed.
Johnston's formula for the relation between the solubility product of calcium carbonate and the partial pressure of carbon dioxide is,
where "n" is the proportion of total carbon dioxide in the solution as carbonic acid, H2CO3, "r" is the ratio k1/k2, that is the ratio
"K" is the solubility product constant, and "cP" is the molar concentration of the dissolved gas, "P" being the partial pressure expressed in atmospheres, and c=a/22.4 where "a" is the absorption coefficient of carbon dioxide.
Using Wells' data, Johnston and Williamson also calculated the solubility product for different temperatures: -
The changes of solubility caused by slight variations in the carbon dioxide content of the atmosphere are of great importance geologically, for many natural waters are saturated with calcium carbonate, and alternate solution and deposition may bring about the transference of large quantities of it. The presence of other salts has also an important influence on the solubility. Ammonium salts, probably through the formation of a complex calcium-ammonium ion, can increase it to a considerable extent. The solubility is, in general, increased by the presence of chlorides, nitrates, and sulphates, but decreased by alkali carbonates and alkaline earth salts. It is also increased in water containing humus, especially when alkali salts are also present. In chloride solutions the solubility passes through a maximum. Calcium carbonate is soluble in acids, with evolution of carbon dioxide.
A colloidal solution of calcium carbonate can be obtained by passing carbon dioxide into a suspension of calcium oxide in methyl alcohol. Ultimately a solid jelly, soluble in methyl alcohol, is obtained. It is miscible with benzene, chloroform, toluene, and ether, but is precipitated by acetone and carbon disulphide.
Two crystalline forms undoubtedly exist. The most widely distributed is the hexagonal form, calcite, calcspar, or Iceland spar, showing marked double refraction. Aragonite in rhombic crystals is much rarer. Other forms have been described - conchite, ktypeite, or vaterite, and lublinite. According to more recent investigations, however, these latter are identical with one or other of the two well- known forms and a third form, μ-calcium carbonate, consisting of microscopic hexagonal plates of density 2.54, which is obtained by precipitation at 60° C., but is always contaminated with the other two. Another form, α-calcium carbonate, only stable at very high temperatures and with optical properties apparently identical with those of calcite, has also been mentioned. The apparently amorphous varieties probably consist of minute crystals of calcite.
Marble and limestone consist of small interlocking calcite crystals.
Calcite and aragonite form one of the earliest recognised cases of polymorphism. Both are anhydrous calcium carbonate, but, besides the difference in crystalline form already noted, there is a difference in density, which is 2.715 for calcite and 2.934 for aragonite. Calcite is isomorphous with sodium nitrate, and aragonite with strontianite. There is an extensive literature of the mutual relations of calcite and aragonite, the former being the stable variety under ordinary conditions. When calcium carbonate is formed by precipitation, it is very probable that, in accordance with Ostwald's rule, one of the unstable forms, μ-calcium carbonate or aragonite, is first obtained, and that this changes with more or less rapidity into calcite. In some cases, owing to the high temperature or to the catalysing effect of the salts present, the velocity of transformation may be so great that ealeite appears to be the first precipitated.
In the dry state aragonite is not changed into calcite with appreciable rapidity until a temperature above 400° C. is reached. At 470° C. the change takes place in a few minutes. If under any conditions aragonite can be the stable form, it must be at very low temperatures. The transition temperature has been given as - 43°±5° C., rising 1° C. per 40 atmospheres increase of pressure. It can only be in the coldest regions of the earth, therefore, that natural aragonite is the stable form. The following are probably the main conditions under which aragonite forms in nature: (a) through organic agencies, for example shells; (b) by deposition from hot springs; (c) when an isomorphous carbonate is present to serve as a nucleus; (d) by chemical precipitation in certain saline waters in the presence of sulphate. The preservation of aragonite in any but recent geological formations may depend on its having been kept dry, or on the presence of impurities, which are more frequent in aragonite than in calcite and which reduce the rate of transformation by water by reducing the solubility below that of calcite. In older rocks, fossil shells of such a type that they were presumably aragonite originally, now consist of calcite. The change from aragonite to calcite is accompanied by a 10 per cent, increase in volume. The heat of transformation is so small, however, that there appears to be some doubt as to whether it is positive or negative. It is possible that it may be positive at ordinary temperatures and negative at higher ones.
μ-Calcium carbonate is too unstable to be found as a mineral, and α-calcium carbonate is only stable above 970° C., the temperature of transition from calcite.
A number of colour tests have been suggested for distinguishing aragonite from calcite in minerals. These depend on the different reactivities of the two modifications towards different salt solutions.
Calcium carbonate may be prepared by precipitation of a soluble calcium salt with an alkali carbonate, or by neutralisation of calcium hydroxide solution with carbon dioxide. The latter is not an instantaneous process but may take a considerable time before it is complete. It is probably governed by the rate of hydration of carbon dioxide.
By fusion of the precipitated carbonate with a mixture of sodium and potassium chlorides, calcite in a visibly crystalline form may be obtained.
The specific heat of calcium carbonate at ordinary temperatures is 0.2027. Regnault gave the following values for the specific heats of the different mineral forms at 99° C.: -
Iceland spar – 0.20858
Aragonite – 0.20850
White marble – 0.21585
Grey marble – 0.20989
White chalk – 0.21485
Calcite is phosphorescent under the cathode rays, and some specimens also when heated, or on insolation.
On heating calcium carbonate it decomposes into calcium oxide and carbon dioxide. The reaction is reversible at high temperatures, and for every temperature there is a corresponding equilibrium pressure of carbon dioxide. This behaviour is to be expected if the system consists of two solid phases, calcium oxide and calcium carbonate, and a vapour phase, carbon dioxide. At low temperatures dissociation and re-absorption proceed very slowly, and depend very much on the condition of the solid phases. Jolibois and Bouvier found that the reaction is not completely reversible when pure calcium carbonate alone forms the starting-point. A mixture of calcium oxide and calcium carbonate must be used. Probably the difficulty is of purely mechanical origin. These authors regarded it as an indication that a solid solution of calcium oxide in calcium carbonate is formed.
The following values have been given for the dissociation pressures of calcium carbonate at different temperatures: -
| Temperature, ° C | 500 | 550 | 600 | 650 | 700 | 750 | 800 | 850 | 900 | 950 | 1000 |
| Pressure, mm. Mercury. | 0.11 | 0.57 | 2.35 | 8.20 | 25.3 | 68 | 168 | 373 | 773 | 1490 | 2710 |
The temperature of dissociation at atmospheric pressure, 760 mm., according to Riesenfeld is 908°±5° C. Hedvall found 913°-923° C., and Johnston 898° C. The latter obtained the following equation for the pressure-temperature curve: -
log10p = - 9840/T + 1.1log10T – 0.0012T + 8.882, where p is the pressure in mm. mercury and T the absolute temperature.
Since aragonite is the unstable form, it should have a higher dissociation pressure than calcite, but at the temperature at which rapid transformation into calcite begins, the pressures are only of the order of 0.01 mm. and differences cannot be determined with accuracy.
The molecular heat of formation of precipitated calcium carbonate from calcium oxide and carbon dioxide is 43.3 Cal., of calcspar 42.0 Cal., of aragonite 42.6 Cal., and of the carbonate at the temperature of decomposition under a pressure of 1 atmosphere 34.76 Cal.
Marble is generally said to be of igneotis origin, and Hall claimed to have obtained marble by the fusion of chalk, but the experiments were repeated by later investigators without success. Recently a substance of the characteristic appearance of marble has been obtained by heating anhydrous calcium chloride with hydrated sodium carbonate, or hydrated chloride with anhydrous carbonate, for eight hours at 300°-305° C. under 24 atmospheres pressure; or by saturating with carbon dioxide, at room temperature, a paste of precipitated chalk with saturated sodium chloride solution, and heating for eight hours at 300°-305° C. under a pressure of 27 atmospheres. Le Chatelier supposed that he had obtained evidence of fusion of calcium carbonate, with or without pressure, at 1020° C., but this was questioned by Joannis, and also by Boeke, who found that it partially dissociated but did not melt even when heated to 1400°-1500° C., under a pressure of carbon dioxide of 30 atmospheres, although the original fine powder became crystalline. In later experiments, however, the latter succeeded in fusing Iceland spar at 1289° C. in carbon dioxide at 110 atmospheres pressure, but the work of Smith and Adams seems to indicate that the melting-point of calcium carbonate lies a little above 1340° C. under a pressure of about 1050 atmospheres. An eutectic mixture containing 50 per cent, of calcium oxide fuses at 1240°±1° C. under about 40 atmospheres pressure.
Calcium carbonate is only very slightly soluble in water, but the solubility is greatly influenced by the concentration of carbon dioxide in solution, and this in its turn depends on the concentration or partial pressure of the gas in the vapour space above the solution. This influence is explained by the hydrolysis of calcium carbonate by water, forming the more soluble calcium hydroxide when the partial pressure of carbon dioxide is low, and by the formation of calcium bicarbonate, which is also more soluble than the carbonate, when the partial pressure of carbon dioxide is high. After prolonged boiling of a suspension of calcium carbonate, the substance passing into solution is practically all calcium hydroxide. By passing a very rapid stream of carbon dioxide through saturated lime-water at 15° C., a concentration of 2.29 grm. of calcium carbonate may be reached, but this only persists for a short time.
 |
| Concentration of CaO vs Pressure of CO2 |
The branch AB represents the slightly increasing concentration of the solution in contact with solid calcium hydroxide for partial pressures up to about 10-14 atmosphere, at which point, represented by B on the diagram, the solid phase, calcium carbonate, appears. The solubility then falls with increasing pressure to a minimum value at about 3.8×10-7 atmosphere, as calculated by Johnston and Williamson from the solubility product. The solubility rises again, rapidly at first and then more gradually, until, at about 15 atmospheres pressure, marked by point C, a fresh solid phase, calcium bicarbonate, should appear. The line CD shows a slightly decreasing concentration of bicarbonate with further increase of pressure.
It is probable that even at the minimum concentration of calcium ions in solution, which is the maximum of carbonate ions, more than 50 per cent, of the negative ions still consist of hydroxyl and bicarbonate ions.
The solubility of calcium carbonate under normal atmospheric conditions, at a partial pressure of carbon dioxide of 3×10-4 atmospheres, is 63 mgm. per litre. The failure of some investigators to state the partial pressures under which their experiments were conducted makes difficult any comparison between the determinations made. It is also possible that in some cases equilibrium was not reached.
The solubility of calcium carbonate decreases with rise of temperature. The following values have been obtained for the solubility of calcite at different temperatures under a partial pressure of 3.2×10-4 atmospheres: -
| Temperature, ° C. | 1 | 21 | 23 | 30 |
| Mgm. CaCO3 per litre. | 82 | 60 | 57 | 55 |
The solubility of aragonite is greater than that of calcite, in accordance with the general relationship between stable and unstable forms. The ratio of solubilities has an average value of 1.07, but there is a tendency for this to decrease as the temperature falls, which suggests that aragonite becomes the stable form at low temperatures.
A convenient method of expressing the solubility is by the solubility product constant, [Ca••]×[CO3''], from which the saturation concentration for any given partial pressure of carbon dioxide can then be calculated. Certain assumptions must be made as to the amount of free and hydrated carbon dioxide in solution, the degree of ionisation of the different electrolytes, and so on. Values for calcite which are of the same order of magnitude have been calculated by different authors and are recorded in the table on the following page.
The last value was obtained by recalculation from the data of McCoy and Smith, assuming that only a small part of the dissolved carbon dioxide is combined with water, and that, consequently, carbonic acid is a stronger acid than has previously been supposed.
Johnston's formula for the relation between the solubility product of calcium carbonate and the partial pressure of carbon dioxide is,
where "n" is the proportion of total carbon dioxide in the solution as carbonic acid, H2CO3, "r" is the ratio k1/k2, that is the ratio
"K" is the solubility product constant, and "cP" is the molar concentration of the dissolved gas, "P" being the partial pressure expressed in atmospheres, and c=a/22.4 where "a" is the absorption coefficient of carbon dioxide.
Using Wells' data, Johnston and Williamson also calculated the solubility product for different temperatures: -
| Temp., ° C. | 0 | 5 | 10 | 15 | 20 | 25 | 30 |
| K×109 | 12.2 | 11.4 | 10.6 | 9.9 | 9.3 | 8.7 | 8.1 |
The changes of solubility caused by slight variations in the carbon dioxide content of the atmosphere are of great importance geologically, for many natural waters are saturated with calcium carbonate, and alternate solution and deposition may bring about the transference of large quantities of it. The presence of other salts has also an important influence on the solubility. Ammonium salts, probably through the formation of a complex calcium-ammonium ion, can increase it to a considerable extent. The solubility is, in general, increased by the presence of chlorides, nitrates, and sulphates, but decreased by alkali carbonates and alkaline earth salts. It is also increased in water containing humus, especially when alkali salts are also present. In chloride solutions the solubility passes through a maximum. Calcium carbonate is soluble in acids, with evolution of carbon dioxide.
A colloidal solution of calcium carbonate can be obtained by passing carbon dioxide into a suspension of calcium oxide in methyl alcohol. Ultimately a solid jelly, soluble in methyl alcohol, is obtained. It is miscible with benzene, chloroform, toluene, and ether, but is precipitated by acetone and carbon disulphide.
The Hydrates of Calcium Carbonate
By exposure of calcium saccharate to the air, or to an atmosphere of carbon dioxide at 0° C., or by injecting carbon dioxide into an ice-cold sugar-lime solution through a capillary tube, crystals of the hexahydrate, CaCO3.6H2O, formerly supposed to be the.pentahydrate, are obtained. They may also be prepared by slow precipitation at 0° C. In this case it is advisable to add a little potassium hydroxide, and the precipitate should be filtered, washed with alcohol and ether, and dried by suction as rapidly as possible. The hexahydrate is thus obtained in colourless monoclinic crystals, which rapidly pass into calcite and water even when kept at 0° C. An attempt has been made by Mackenzie to determine the transition temperature by means of the dilatometer, and it was found that below 5° C. the hydrate is fairly stable, but that, when the change has set in, lowering the temperature does not stop it. By boiling with absolute alcohol Pelouze obtained a trihydrate. According to Johnston and his co-workers only the hexahydrate exists, but Tschirwinsky described two mineral forms, trihydrocalcite, CaCO3.3H2O, and pentahydrocalcite, CaCO3.5H2O.
Double Salts of Calcium Carbonate
The best known, and most widely distributed, naturally occurring double carbonate of calcium is dolomite, or bitter spar, which may contain varying proportions of calcium to magnesium, and is, therefore, to be regarded as a solid solution rather than as a compound. It forms yellow or white, massive, crystalline aggregates of density 2.8 – 2.95. It is harder than limestone, and less easily soluble in acids. It is sometimes used for building-stone. On heating to 400° C. it becomes capable of setting with water like cement. Dolomite may be formed artificially by the action of ammonium carbonate on calcium and magnesium chlorides in an atmosphere of carbon dioxide. Plumbo-calcite has already been mentioned.
By the action of sodium carbonate solution on calcium hydroxide, pirssonite, Na2Ca(CO3)2.2H2O, gaylussite, Na2Ca(CO3)2.5H2O, or the anhydrous salt, Na2Ca(CO3)2, may be formed, depending on the conditions of temperature and concentration. Three double calcium potassium salts, K2Ca(CO3)2, K6Ca(CO3)4, and K6Ca2(CO3)5.6H2O, have also been obtained. The anhydrous double salts, Na2Ca(CO3)2 and K2Ca(CO3)2, may be prepared by fusion.
A compound of calcium chloride with calcium carbonate has been described.
By fusing together calcium and lithium carbonates in a current of carbon dioxide, a compound, probably of composition CaCO3.2Li2CO3, is formed. It is readily decomposed by water.
By the action of sodium carbonate solution on calcium hydroxide, pirssonite, Na2Ca(CO3)2.2H2O, gaylussite, Na2Ca(CO3)2.5H2O, or the anhydrous salt, Na2Ca(CO3)2, may be formed, depending on the conditions of temperature and concentration. Three double calcium potassium salts, K2Ca(CO3)2, K6Ca(CO3)4, and K6Ca2(CO3)5.6H2O, have also been obtained. The anhydrous double salts, Na2Ca(CO3)2 and K2Ca(CO3)2, may be prepared by fusion.
A compound of calcium chloride with calcium carbonate has been described.
By fusing together calcium and lithium carbonates in a current of carbon dioxide, a compound, probably of composition CaCO3.2Li2CO3, is formed. It is readily decomposed by water.
Basic Calcium Carbonate
By heating quicklime in carbon dioxide Raoult obtained two compounds, to which he gave the formulae 2CaO.CO2 and 3CaO.2CO2, and possibly also the compounds 4CaO.3CO2 and 5CaO.2CO2. The two former are hydrated and set like cement when treated with water, giving CaCO3.Ca(OH)2 and 2CaCO3.Ca(OH)2. The vapour pressure curve of calcium carbonate at high temperatures, however, hardly favours the view that a basic carbonate may be formed, nor yet that solid solutions of oxide in carbonate are possible.
Uses of Calcium Carbonate
The chief use of calcium carbonate is for the production of quicklime. It can also be used wherever a cheap mild alkali is required.
Calcium carbonate, as well as lime, has a beneficial action on the soil apart from the neutralisation of the soil acids. It is more suitable for the lighter soils, but can also increase to some extent the permeability of heavy soils. It increases the availability of nitrogen.
Calcium carbonate, as well as lime, has a beneficial action on the soil apart from the neutralisation of the soil acids. It is more suitable for the lighter soils, but can also increase to some extent the permeability of heavy soils. It increases the availability of nitrogen.
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