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Ferrous sulphate, FeSO4
The usual form of Ferrous sulphate is the heptahydrate, FeSO4.7H2O, which is obtained when iron is dissolved in dilute sulphuric acid and the solution allowed to crystallise. The crystals belong to the monoclinic system, and owing to their green colour the salt has long been known as green vitriol. The density of the pure salt at 14.8° C. is 1.8987, and at 15° С. 1.899 ± 0.001, and it is stable in contact with a saturated solution of ferrous sulphate up to 56.6° C. which is the transition point of the hepta- to the tetra-hydrate under these conditions. It occurs native as the mineral melanterite or copperas, which may be crystalline, but is more usually massive. Mixed crystals of the heptahydrates of ferrous and cupric sulphates occur as pisanite, (Fe, Cu)SO4.7H2O, in the form of blue transparent crystals. Both these minerals are isomorphous with Boothite, CuSO4.7H2O.
Drainage water from coal mines is frequently charged with ferrous sulphate consequent upon the oxidation of pyrites in the coal, and upon evaporation yields the impure salt. The salt is readily obtained in a pure state by dissolving electrolytic iron in dilute sulphuric acid and allowing to crystallise, preferably in an inert atmosphere. Commercially ferrous sulphate is obtained by exposing heaps of pyrites to the oxidising action of the air. Ferrous sulphate and free sulphuric acid drain off into tanks, the acid being neutralised with scrap iron: -
FeS2 + 7O + H2O = FeSO4 + H2SO4,
H2SO4 + Fe = FeSO4 + H2.
The salt obtained in this way is not pure, but contains small quantities of ferric sulphate and the sulphates of metals such as manganese naturally occurring in the pyrites. Copper sulphate is removed by allowing the liquors to remain a sufficient length of time in contact with the scrap iron, the copper being precipitated out: -
CuSO4 + Fe = FeSO4 + Cu.
Ferric sulphate is removed by re-crystallisation, but small quantities of the other salts remain.
When exposed to the air the ordinary commercial salt gradually oxidises, yielding a basic salt. It is easily dried, however, by powdering and repeatedly pressing between the folds of filter paper. It is then quite stable in air at 15° C. and neither oxidises, effloresces, nor deliquesces. When heated in chloroform vapour a mixture of ferrous and ferric chlorides is produced.
The salt melts at 64° C. When heated in vacuo, ferrous sulphate heptahydrate undergoes dehydration, six molecules of water being removed at 140° C., complete dehydration being effected at a slightly higher temperature. When heated in a tube open at both ends, ferrous sulphate begins to decompose at 150° C., yielding ferric sulphate, Fe2(SO4)3, and at higher temperatures, in the presence of air, ferric oxide results. Its coefficient of expansion per degree C. rise in temperature is 7.2×10-5.
Its specific heat is as follows: -
and its molecular specific heat between 22° and 45° C. is 96.2 calories. The heat of hydration of the heptahydrate is 1912 calories, and its heat of solution is -4323 calories at 13.5° C.
Crystals of the heptahydrate possess the same vapour tension at 44.01° C. as magnesium sulphate, MgSO4.7H2O. Below this temperature their dissociation pressure is greater, and above it is less, than that of the magnesium salt. In the case of zinc sulphate, ZnSO4.7H2O, the equilibrium temperature between the two salts is 16.4° C.
Examination of ferrous sulphate crystals by X-ray methods indicates that the seven molecules of water are not symmetrically disposed or equivalent in their structural relations to the other constituents.
The heptahydrate is dimorphous, a rhombic variety occurring in nature as tauriscite in a more or less impure condition. It has not been isolated in a pure condition in the laboratory, but Rammelsberg obtained it in association with rhombic magnesium sulphate.
Experiment shows that crystals of magnesium sulphate heptahydrate, MgSO4.7H2O, will hold up to 19 per cent, of ferrous sulphate, FeSO4.7H2O, in solid solution. Now, the pure magnesium salt is rhombic, and has a density of 1.677, whilst the usual form of ferrous sulphate is monoclinic, density 1.898. The mixed crystals, however, are rhombic. Determination of their density shows that it is not the mean value as calculated from the densities of the constituent salts, but is slightly less. In other words, it appears that the magnesium sulphate is mixed with a ferrous sulphate of density 1.875, rhombic in form, and too unstable to exist alone when quite pure under ordinary conditions. Similarly crystals of ferrous sulphate heptahydrate can hold 54 per cent, of magnesium sulphate in solid solution. The mixed crystals are, in this case, monoclinic, and their density is greater than the mean calculated from the densities of the constituents. It would thus appear that in this case the ferrous sulphate is associated with a monoclinic variety of magnesium sulphate heptahydrate, of density 1.691, but which is too unstable to exist alone in ordinary circumstances. Hence the two heptahydrates are isodimorphous.
The hexahydrate, FeSO4.6H2O, results on passing hydrogen chloride into a saturated solution of ferrous sulphate. Ferrous chloride separates first, and, concentrating the mother liquor, tabular crystals of ferrous sulphate hexahydrate separate out.
When a crystal of copper sulphate, CuSO4.5H2O, is introduced into a supersaturated solution of ferrous sulphate, triclinic crystals of the pentahydrate, FeSO4.5H2O, separate out, isomorphous with the copper salt, and of density 1.89.
The pentahydrate may also be obtained by allowing an acidified solution of ferrous sulphate to concentrate in vacuo. The heptahydrate crystallises first, next the pentahydrate, and finally the tetrahydrate, FeSO4.4H2O, the relative proportions of these hydrates depending upon the amount of free sulphuric acid in solution.
The tetrahydrate is also formed when the heptahydrate is kept for several days over concentrated sulphuric acid. It is monoclinic, and isomorphous with the corresponding hydrate of manganous sulphate, MnSO4.4H2O. Its limits of stability in contact with a saturated solution of ferrous sulphate are 56.6° C. to 64.4° C. Its heat of solution at 13.5° C. is 1599 calories, and its molecular heat is 63.587 calories.
The trihydrate, FeSO4.3H2O, and dihydrate, FeSO4.2H2O, have also been obtained - the former, by solution of the heptahydrate in concentrated hydrochloric acid; the latter by separation from a concentrated solution of ferrous sulphate on addition of sulphuric acid in small quantities at a time. The existence of this latter hydrate is clearly indicated by a break in the time-dehydration curve at 100° C.
The monohydrate, FeSO4.H2O, occurs in nature as the mineral ferro-palladite, in Chili, and results when the heptahydrate is heated to 140° C. in vacuo, or is allowed to effloresce for prolonged periods in dry air, and by passing air, dried over sulphuric acid, over the heptahydrate at 100° C. It also results by bringing the heptahydrate into contact with sulphuric acid of concentration not less than 12.5 normal, and by addition of concentrated sulphuric acid to saturated aqueous solutions of ferrous sulphate. This is conveniently effected by dissolving 400 grams of the pure heptahydrate in 200 c.c. of 50 per cent, sulphuric acid on the water-bath. Almost immediately after solution is complete, the monohydrate separates out as a white, crystalline powder. It is dried with alcohol and ether, and finally over sulphuric acid. The salt is permanent in air, and is not hygroscopic. It is useful, therefore, for standardising permanganate solutions for volumetric analysis. It clings most tenaciously to its combined water, but may be made to part with it to a small extent by heating to 100° C. in a current of air dried over phosphorus pentoxide. The suggestion has therefore been made that the salt is a hemihydrol of the formula
It is stable in contact with ferrous sulphate solution at temperatures above 64.4° C., which point is the transition temperature in these circumstances between the mono- and tetra-hydrates. Its heat of solution in water at 13.5° C. is 7538 calories. It absorbs ammonia vapour, yielding the pentammoniate, FeSO4.5NH3.H2O.
The anhydrous salt, FeSO4, is extremely difficult to prepare in a state of purity. It results, more or less contaminated with a basic salt, when any hydrate is heated in vacuo to a temperature somewhat above 140° C. It is stated also to result when the heptahydrate is dissolved in concentrated sulphuric acid. It then separates as microscopic prisms. It is white in appearance, insoluble both in alcohol and in concentrated sulphuric acid. Its heat of solution in water is 14,901 calories at 13.5° C. Density 2.841. At dull red heat it decomposes, yielding first a basic sulphate and finally ferric oxide. Thus: -
2FeSO4 = Fe2O3 + SO3 + SO2.
Ferrous sulphate readily dissolves in water, and if acidified with dilute sulphuric acid the solution is fairly stable in the cold. Exposure to light accelerates its rate of oxidation. The solubility of ferrous sulphate in water has been determined at different temperatures by many investigators, the most recent and reliable work being that of Fraenckel, who gives the following data, shown graphically in fig. At 25° C. 100 grams of saturated solution contain 22.98 grams (0.1503 mol.) of FeSO4.
The density of a solution of ferrous sulphate, saturated and in contact with crystals of the salt, at 8.9° C. is 1.1949.
The rate of oxidation of ferrous sulphate solution upon exposure to air is proportional to the partial pressure of the oxygen. Hence it is reduced by addition of concentrated solutions of inert soluble salts, such as chlorides and sulphates of sodium, potassium, and magnesium, owing to their presence causing a decrease in the solubility of the oxygen. The oxidation depends upon the un-ionised portion of the dissolved salt.
Ferrous sulphate undergoes hydrolysis when its solution is boiled with potassium iodide and iodate. Thus: -
3FeSO4 + 5KI + KIO3 + 3H2O = 3Fe(HO)2 + 3K2SO4 + 3I2.
The excess of iodate then oxidises the ferrous hydroxide to the ferric condition.
6Fe(OH)2 + KIO3 + 3H2O = KI + 6Fe(OH)3.
Drainage water from coal mines is frequently charged with ferrous sulphate consequent upon the oxidation of pyrites in the coal, and upon evaporation yields the impure salt. The salt is readily obtained in a pure state by dissolving electrolytic iron in dilute sulphuric acid and allowing to crystallise, preferably in an inert atmosphere. Commercially ferrous sulphate is obtained by exposing heaps of pyrites to the oxidising action of the air. Ferrous sulphate and free sulphuric acid drain off into tanks, the acid being neutralised with scrap iron: -
FeS2 + 7O + H2O = FeSO4 + H2SO4,
H2SO4 + Fe = FeSO4 + H2.
The salt obtained in this way is not pure, but contains small quantities of ferric sulphate and the sulphates of metals such as manganese naturally occurring in the pyrites. Copper sulphate is removed by allowing the liquors to remain a sufficient length of time in contact with the scrap iron, the copper being precipitated out: -
CuSO4 + Fe = FeSO4 + Cu.
Ferric sulphate is removed by re-crystallisation, but small quantities of the other salts remain.
When exposed to the air the ordinary commercial salt gradually oxidises, yielding a basic salt. It is easily dried, however, by powdering and repeatedly pressing between the folds of filter paper. It is then quite stable in air at 15° C. and neither oxidises, effloresces, nor deliquesces. When heated in chloroform vapour a mixture of ferrous and ferric chlorides is produced.
The salt melts at 64° C. When heated in vacuo, ferrous sulphate heptahydrate undergoes dehydration, six molecules of water being removed at 140° C., complete dehydration being effected at a slightly higher temperature. When heated in a tube open at both ends, ferrous sulphate begins to decompose at 150° C., yielding ferric sulphate, Fe2(SO4)3, and at higher temperatures, in the presence of air, ferric oxide results. Its coefficient of expansion per degree C. rise in temperature is 7.2×10-5.
Its specific heat is as follows: -
| Temperature Range. °C. | Specific Heat of FeSO4.7H2O. |
| -78.4 to +22 | 0.292±0.002 |
| -190.0 to +22 | 0.234±0.002 |
| -190.0 to -78.4 | 0.182±0.004 |
and its molecular specific heat between 22° and 45° C. is 96.2 calories. The heat of hydration of the heptahydrate is 1912 calories, and its heat of solution is -4323 calories at 13.5° C.
Crystals of the heptahydrate possess the same vapour tension at 44.01° C. as magnesium sulphate, MgSO4.7H2O. Below this temperature their dissociation pressure is greater, and above it is less, than that of the magnesium salt. In the case of zinc sulphate, ZnSO4.7H2O, the equilibrium temperature between the two salts is 16.4° C.
Examination of ferrous sulphate crystals by X-ray methods indicates that the seven molecules of water are not symmetrically disposed or equivalent in their structural relations to the other constituents.
The heptahydrate is dimorphous, a rhombic variety occurring in nature as tauriscite in a more or less impure condition. It has not been isolated in a pure condition in the laboratory, but Rammelsberg obtained it in association with rhombic magnesium sulphate.
Experiment shows that crystals of magnesium sulphate heptahydrate, MgSO4.7H2O, will hold up to 19 per cent, of ferrous sulphate, FeSO4.7H2O, in solid solution. Now, the pure magnesium salt is rhombic, and has a density of 1.677, whilst the usual form of ferrous sulphate is monoclinic, density 1.898. The mixed crystals, however, are rhombic. Determination of their density shows that it is not the mean value as calculated from the densities of the constituent salts, but is slightly less. In other words, it appears that the magnesium sulphate is mixed with a ferrous sulphate of density 1.875, rhombic in form, and too unstable to exist alone when quite pure under ordinary conditions. Similarly crystals of ferrous sulphate heptahydrate can hold 54 per cent, of magnesium sulphate in solid solution. The mixed crystals are, in this case, monoclinic, and their density is greater than the mean calculated from the densities of the constituents. It would thus appear that in this case the ferrous sulphate is associated with a monoclinic variety of magnesium sulphate heptahydrate, of density 1.691, but which is too unstable to exist alone in ordinary circumstances. Hence the two heptahydrates are isodimorphous.
The hexahydrate, FeSO4.6H2O, results on passing hydrogen chloride into a saturated solution of ferrous sulphate. Ferrous chloride separates first, and, concentrating the mother liquor, tabular crystals of ferrous sulphate hexahydrate separate out.
When a crystal of copper sulphate, CuSO4.5H2O, is introduced into a supersaturated solution of ferrous sulphate, triclinic crystals of the pentahydrate, FeSO4.5H2O, separate out, isomorphous with the copper salt, and of density 1.89.
The pentahydrate may also be obtained by allowing an acidified solution of ferrous sulphate to concentrate in vacuo. The heptahydrate crystallises first, next the pentahydrate, and finally the tetrahydrate, FeSO4.4H2O, the relative proportions of these hydrates depending upon the amount of free sulphuric acid in solution.
The tetrahydrate is also formed when the heptahydrate is kept for several days over concentrated sulphuric acid. It is monoclinic, and isomorphous with the corresponding hydrate of manganous sulphate, MnSO4.4H2O. Its limits of stability in contact with a saturated solution of ferrous sulphate are 56.6° C. to 64.4° C. Its heat of solution at 13.5° C. is 1599 calories, and its molecular heat is 63.587 calories.
The trihydrate, FeSO4.3H2O, and dihydrate, FeSO4.2H2O, have also been obtained - the former, by solution of the heptahydrate in concentrated hydrochloric acid; the latter by separation from a concentrated solution of ferrous sulphate on addition of sulphuric acid in small quantities at a time. The existence of this latter hydrate is clearly indicated by a break in the time-dehydration curve at 100° C.
The monohydrate, FeSO4.H2O, occurs in nature as the mineral ferro-palladite, in Chili, and results when the heptahydrate is heated to 140° C. in vacuo, or is allowed to effloresce for prolonged periods in dry air, and by passing air, dried over sulphuric acid, over the heptahydrate at 100° C. It also results by bringing the heptahydrate into contact with sulphuric acid of concentration not less than 12.5 normal, and by addition of concentrated sulphuric acid to saturated aqueous solutions of ferrous sulphate. This is conveniently effected by dissolving 400 grams of the pure heptahydrate in 200 c.c. of 50 per cent, sulphuric acid on the water-bath. Almost immediately after solution is complete, the monohydrate separates out as a white, crystalline powder. It is dried with alcohol and ether, and finally over sulphuric acid. The salt is permanent in air, and is not hygroscopic. It is useful, therefore, for standardising permanganate solutions for volumetric analysis. It clings most tenaciously to its combined water, but may be made to part with it to a small extent by heating to 100° C. in a current of air dried over phosphorus pentoxide. The suggestion has therefore been made that the salt is a hemihydrol of the formula
It is stable in contact with ferrous sulphate solution at temperatures above 64.4° C., which point is the transition temperature in these circumstances between the mono- and tetra-hydrates. Its heat of solution in water at 13.5° C. is 7538 calories. It absorbs ammonia vapour, yielding the pentammoniate, FeSO4.5NH3.H2O.
The anhydrous salt, FeSO4, is extremely difficult to prepare in a state of purity. It results, more or less contaminated with a basic salt, when any hydrate is heated in vacuo to a temperature somewhat above 140° C. It is stated also to result when the heptahydrate is dissolved in concentrated sulphuric acid. It then separates as microscopic prisms. It is white in appearance, insoluble both in alcohol and in concentrated sulphuric acid. Its heat of solution in water is 14,901 calories at 13.5° C. Density 2.841. At dull red heat it decomposes, yielding first a basic sulphate and finally ferric oxide. Thus: -
2FeSO4 = Fe2O3 + SO3 + SO2.
 |
| Solubility of ferrous sulphate in water |
The density of a solution of ferrous sulphate, saturated and in contact with crystals of the salt, at 8.9° C. is 1.1949.
The rate of oxidation of ferrous sulphate solution upon exposure to air is proportional to the partial pressure of the oxygen. Hence it is reduced by addition of concentrated solutions of inert soluble salts, such as chlorides and sulphates of sodium, potassium, and magnesium, owing to their presence causing a decrease in the solubility of the oxygen. The oxidation depends upon the un-ionised portion of the dissolved salt.
Ferrous sulphate undergoes hydrolysis when its solution is boiled with potassium iodide and iodate. Thus: -
3FeSO4 + 5KI + KIO3 + 3H2O = 3Fe(HO)2 + 3K2SO4 + 3I2.
The excess of iodate then oxidises the ferrous hydroxide to the ferric condition.
6Fe(OH)2 + KIO3 + 3H2O = KI + 6Fe(OH)3.
Ferrous Sulphate as a Reducing Agent
Ferrous sulphate is frequently used as a mild reducing agent. Thus, auric chloride is reduced to the metal in aqueous solution - a reaction made use of in photographic toning: -
AuCl3 + 3FeSO4 = Au + FeCl3 + Fe2(SO4)3.
Silver salts are similarly reduced to the metal thus: -
Ag2SO4 + 2FeSO4 = 2Ag + Fe2(SO4)3
and
3AgNO3 + 3FeSO4 = 3Ag + Fe2(SO4)3 + Fe(NO3)3.
These reactions with silver are particularly interesting inasmuch as Landolt finds that they are accompanied by a loss in weight greater than that attributable to experimental error, and suggests that the atoms lose a small portion of their mass in the reaction, the detached particles passing through the walls of the containing vessel.
Under the influence of light, potassium ferricyanide is reduced to ferrocyanide by ferrous sulphate in alkaline solution.
Potassium permanganate is instantaneously decolorised by ferrous sulphate in acid solution, being reduced to manganous sulphate. Thus: -
10FeSO4 + 2KMnO4 + 8H2SO4 = 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O.
Similarly potassium bichromate is reduced to chromic sulphate in accordance with the equation: -
6FeSO4 + K2Cr2O7 + 7H2SO4 = K2SO4 + Cr2(SO4)3 + 3Fe2(SO4)3 + 7H2O.
Aqueous solutions of ferrous sulphate readily absorb nitric oxide, the extent of absorption depending upon the concentration of the iron, the temperature, and the pressure. The limit of absorption is reached when one molecule of NO is present to each atom of iron and the brown solution undoubtedly contains the compound FeSO4.NO, probably more or less combined with the solvent. The addition of small quantities of sulphuric acid to the solution of ferrous sulphate tends to diminish the absorption of nitric oxide, the equilibrium represented by the equation
FeSO4 + NO ⇔ FeSO4.NO
being pushed towards the left. Further increase of the acid, however, assists absorption, a maximum being reached in the presence of 82 per cent, of acid. Under these conditions the solution is cherry-red in colour. The colour is not due to the formation of ferrous nitroso sulphonate,
as assumed by Raschig. The complex salt FeSO4.NO may be isolated by adding a concentrated aqueous solution of nitric oxide in ferrous sulphate to ice-cold sulphuric acid in an atmosphere of nitric oxide; the compound crystallises in small red leaflets, but is very unstable. A second complex salt of the formula FeSO4NO.FeSO4.13H2O has also been obtained by addition of ethyl alcohol to aqueous ferrous sulphate in an atmosphere of nitric oxide. It crystallises in small rectangular plates brown in colour and which slowly decompose on exposure to air.
On addition of nitric acid to a solution of ferrous sulphate acidified with sulphuric acid, nitric oxide is formed, a portion of the ferrous salt being oxidised to ferric. Thus: -
6FeSO4 + 3H2SO4 + 2HNO3 = 3Fe2(SO4)3 + 4H2O + 2NO.
This is the basis of the "ring test" for nitric acid or nitrates, which usually consists in pouring a cold solution of ferrous sulphate gently down the sides of an inclined test-tube on to a layer of concentrated sulphuric acid, containing the nitrate. Since ferric sulphate yields a red compound, possibly Fe2(SO4)3.4NO, with nitric oxide in the presence of concentrated sulphuric acid, the colour of the ring formed at the junction of the two liquids will depend upon whether the nitric oxide compound is formed in the concentrated acid or in the aqueous layer, being brown in the latter, but ruddy in the former.
The dark colour of the nitric oxide derivative may be utilised in the volumetric estimation of nitric acid, as its appearance depends upon the presence of excess ferrous sulphate in solution. Further addition of nitric acid oxidises this remaining ferrous sulphate and, when oxidation is complete, the colour disappears. The discharge of the colour is found to be sufficiently definite to enable the end point to be determined with reasonable accuracy.
The formation of this brown compound can advantageously be utilised as a method of detecting ferrous salts in the presence of other metals that would obscure the more usual ferricyanide reaction. The solution to be tested is mixed with an equal volume of concentrated sulphuric acid, and a crystal of potassium nitrate added. The last-named becomes surrounded with red-brown streaks of the nitroso compound.
A mixed solution of ferrous sulphate and catechol in the presence of an alkali readily absorbs oxygen, giving a deep red colour. The reaction is exceedingly sensitive, and is recommended as a delicate test for oxygen.
Ferrous sulphate has been used as a dressing for crops, but apparently it is beneficial only when the soil contains an excess of lime, which is thereby converted into gypsum. It is this last-named salt which really benefits the soil. Ferrous salts in general are toxic and are usually regarded as one cause of sterility of badly aerated soils.
The double sulphate with sodium, FeSO4.Na2SO4.4H2O, has been prepared by crystallisation of the mixed solutions above 45° C.
Double Sulphates of the Type (M'', N'')SO4.xH2O, or M''SO4.N''SO4.xH2O.
Ferrous sulphate yields, with the sulphates of certain divalent metals, mixed salts of the general type (M'', N'')SO4.H2O. They are prepared by mixing solutions of the constituent sulphates, and then adding concentrated sulphuric acid, when the mixed sulphates separate out on cooling.
AuCl3 + 3FeSO4 = Au + FeCl3 + Fe2(SO4)3.
Silver salts are similarly reduced to the metal thus: -
Ag2SO4 + 2FeSO4 = 2Ag + Fe2(SO4)3
and
3AgNO3 + 3FeSO4 = 3Ag + Fe2(SO4)3 + Fe(NO3)3.
These reactions with silver are particularly interesting inasmuch as Landolt finds that they are accompanied by a loss in weight greater than that attributable to experimental error, and suggests that the atoms lose a small portion of their mass in the reaction, the detached particles passing through the walls of the containing vessel.
Under the influence of light, potassium ferricyanide is reduced to ferrocyanide by ferrous sulphate in alkaline solution.
Potassium permanganate is instantaneously decolorised by ferrous sulphate in acid solution, being reduced to manganous sulphate. Thus: -
10FeSO4 + 2KMnO4 + 8H2SO4 = 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O.
Similarly potassium bichromate is reduced to chromic sulphate in accordance with the equation: -
6FeSO4 + K2Cr2O7 + 7H2SO4 = K2SO4 + Cr2(SO4)3 + 3Fe2(SO4)3 + 7H2O.
Aqueous solutions of ferrous sulphate readily absorb nitric oxide, the extent of absorption depending upon the concentration of the iron, the temperature, and the pressure. The limit of absorption is reached when one molecule of NO is present to each atom of iron and the brown solution undoubtedly contains the compound FeSO4.NO, probably more or less combined with the solvent. The addition of small quantities of sulphuric acid to the solution of ferrous sulphate tends to diminish the absorption of nitric oxide, the equilibrium represented by the equation
FeSO4 + NO ⇔ FeSO4.NO
being pushed towards the left. Further increase of the acid, however, assists absorption, a maximum being reached in the presence of 82 per cent, of acid. Under these conditions the solution is cherry-red in colour. The colour is not due to the formation of ferrous nitroso sulphonate,
as assumed by Raschig. The complex salt FeSO4.NO may be isolated by adding a concentrated aqueous solution of nitric oxide in ferrous sulphate to ice-cold sulphuric acid in an atmosphere of nitric oxide; the compound crystallises in small red leaflets, but is very unstable. A second complex salt of the formula FeSO4NO.FeSO4.13H2O has also been obtained by addition of ethyl alcohol to aqueous ferrous sulphate in an atmosphere of nitric oxide. It crystallises in small rectangular plates brown in colour and which slowly decompose on exposure to air.
On addition of nitric acid to a solution of ferrous sulphate acidified with sulphuric acid, nitric oxide is formed, a portion of the ferrous salt being oxidised to ferric. Thus: -
6FeSO4 + 3H2SO4 + 2HNO3 = 3Fe2(SO4)3 + 4H2O + 2NO.
This is the basis of the "ring test" for nitric acid or nitrates, which usually consists in pouring a cold solution of ferrous sulphate gently down the sides of an inclined test-tube on to a layer of concentrated sulphuric acid, containing the nitrate. Since ferric sulphate yields a red compound, possibly Fe2(SO4)3.4NO, with nitric oxide in the presence of concentrated sulphuric acid, the colour of the ring formed at the junction of the two liquids will depend upon whether the nitric oxide compound is formed in the concentrated acid or in the aqueous layer, being brown in the latter, but ruddy in the former.
The dark colour of the nitric oxide derivative may be utilised in the volumetric estimation of nitric acid, as its appearance depends upon the presence of excess ferrous sulphate in solution. Further addition of nitric acid oxidises this remaining ferrous sulphate and, when oxidation is complete, the colour disappears. The discharge of the colour is found to be sufficiently definite to enable the end point to be determined with reasonable accuracy.
The formation of this brown compound can advantageously be utilised as a method of detecting ferrous salts in the presence of other metals that would obscure the more usual ferricyanide reaction. The solution to be tested is mixed with an equal volume of concentrated sulphuric acid, and a crystal of potassium nitrate added. The last-named becomes surrounded with red-brown streaks of the nitroso compound.
A mixed solution of ferrous sulphate and catechol in the presence of an alkali readily absorbs oxygen, giving a deep red colour. The reaction is exceedingly sensitive, and is recommended as a delicate test for oxygen.
Ferrous sulphate has been used as a dressing for crops, but apparently it is beneficial only when the soil contains an excess of lime, which is thereby converted into gypsum. It is this last-named salt which really benefits the soil. Ferrous salts in general are toxic and are usually regarded as one cause of sterility of badly aerated soils.
The double sulphate with sodium, FeSO4.Na2SO4.4H2O, has been prepared by crystallisation of the mixed solutions above 45° C.
Double Sulphates of the Type (M'', N'')SO4.xH2O, or M''SO4.N''SO4.xH2O.
Ferrous sulphate yields, with the sulphates of certain divalent metals, mixed salts of the general type (M'', N'')SO4.H2O. They are prepared by mixing solutions of the constituent sulphates, and then adding concentrated sulphuric acid, when the mixed sulphates separate out on cooling.
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