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Sulphuric Acid, H2SO4

Sulphuric Acid, H2SO4, the commonest derivative of sulphur trioxide and the most important of all acids from a technical and commercial aspect, has been known from early times, although its production on a large scale and at a low price dates from the success of the "lead chamber" process of manufacture, which revolutionised chemical industry in the early part of the nineteenth century.


Rivers and springs which have their sources in volcanic regions frequently contain appreciable quantities of sulphuric acid, the amount rising sometimes to 0.5 per cent.; such waters are found for example in Texas. The free acid is also found in the ground water on the moors near Danzig. Certain molluscs, for example Dolium galea, secrete a considerable amount of free sulphuric acid. Traces are present in the air of towns to the extent of a few thousandths of one per cent., but its action on finely divided salt (from sea spray) suspended in the air gives rise to traces of sodium sulphate; the latter is the probable cause of the almost immediate crystallisation of supersaturated solutions of sodium sulphate when exposed to the air.

Sulphates occur abundantly in nature, the chief being those of calcium, barium, strontium, magnesium, aluminium, iron, zinc, copper, sodium and potassium.

Early History

Although the destructive distillation of partially dehydrated sulphate such as alum was undoubtedly practised very early, even perhaps in the tenth century, the first definite details of the process emanated from Basil Valentine to wards the end of the fifteenth century, who described the distillation of green vitriol with silica; the name oil of vitriol is derived from this process. Basil Valentine also obtained an acid product by the combustion of a mixture of sulphur and saltpetre, the identity of this "sulphuric acid" with "oil of vitriol" being first proved by A. Libavius in 1595.

Glauber was the first to separate sulphur from the acid, by heating its salts with coal and acidifying the aqueous extract of the product. On account of this relation between sulphur and the acid, the French chemists de Morveau, Lavoisier, Berthollet and Fourcroy proposed the name "sulphuric acid," which has been retained.

The first experimental investigation indicating the quantitative composition of the acid was made by Gay-Lussac in 1807.


Sulphuric acid is commonly produced in the oxidation of sulphur compounds less rich in oxygen:
  1. Finely divided sulphur in a moist atmosphere undergoes very slow oxidation, some sulphuric acid being produced; the oxidation can rapidly be effected by nitric acid or by aqueous solutions of chlorine or bromine.
  2. Sulphur dioxide in a moist condition also readily oxidises to sulphuric acid (see sulphur dioxide and trioxide), the change being aided by the presence of nitric oxide or by contact with finely divided platinum. With aqueous solutions, free oxygen can be replaced by various oxidising agents such as nitric acid, the halogens, or hydrogen peroxide.
  3. Oxidation of the salts richer in sulphur or poorer in oxygen, for example the polythionates, thiosulphates, hydrosulphites and sulphides, together with their corresponding acids, also produces sulphates as the final stable product.


The world's annual output of sulphuric acid is approximately 10 million tons, and this huge amount is supplied almost entirely by the "lead chamber" and "contact" processes. The dry distillation of green vitriol as a technical operation has now been abandoned. In Great Britain and Northern Ireland the output for the year 1928 in terms of 100 per cent, sulphuric acid and including oleum was 928,000 tons.

Physical Properties of Sulphuric Acid

The "pure" product obtained by evaporation or distillation whether of a dilute or a fuming acid contains as a maximum approximately 98.4 per cent, of H2SO4, the actual figure varying slightly with the pressure. At extremely low pressures evaporation may yield almost pure sulphuric acid. The commercial acid commonly has a concentration of 94 to 97 per cent.

As stated earlier, absolute sulphuric acid of 100 per cent, strength can be obtained as a crystalline solid by fractional crystallisation of the ordinary concentrated acid or by adding the requisite quantity of sulphur trioxide to it; the solid has a melting-point of 10.5° C. and a latent heat of fusion of 22.82 calories per gram.

Absolute 100 per cent, sulphuric acid when heated begins to boil at 290° C., but the temperature rapidly rises to a constant boiling-point of 338° C., when 98.4 per cent, acid distils. This behaviour is due to dissociation of the pure acid, which at temperatures above 30° C. is made evident by its fumes, and which can be demonstrated by partial separation of the dissociation products by making use of their unequal rates of diffusion. The vapour of sulphuric acid is largely dissociated into sulphur trioxide and water; acid above 98.4 per cent, concentration loses trioxide more rapidly than water, whilst acid below this strength parts preferentially with water, the tendency in both cases being for the remaining liquid to attain the concentration of 98.4 per cent., which is that of the acid mixture of minimum vapour pressure.

Vapour density or vapour pressure determinations indicate very considerable dissociation of sulphuric acid vapour into steam and sulphur trioxide at 450° C., and at higher temperatures the dissociation is still further increased, attaining about 50 per cent, at 623° c. under atmospheric pressure. At these higher temperatures the sulphur trioxide itself suffers partial dissociation:

H2SO4H2O + SO3,
2SO3 ⇔ 2SO2 + O2.

Near the boiling-point of the acid the vapour density gives indications of a partial association of the sulphuric acid molecules into (H2SO4)x, where x is probably 2, in addition to the partial dissociation already mentioned.

Pure sulphuric acid at the ordinary temperature is an odourless, oily liquid; at 0° C. it has a specific gravity of 1.854, at 15° C. 1.8372, and at 24° C., referred to water at the same temperature, 1.834. The pure acid is a very poor conductor of electricity; k = 0.01 at 25° C. It has a specific heat of 0.335 between 5° and 22° C., 0.355 between 22° and 80° C., the value increasing with rise of temperature to 0.370 at 160° to 170° C. Other physical characteristics include a refractive index, nD, of 1.42922, and a specific inductive capacity of 41.6 at the ordinary temperature.

The gases oxygen, nitrogen and carbon dioxide are dissolved by concentrated sulphuric acid almost to the same extent as by water, but the diluted acid is a poorer solvent in this respect, the minimum solvent power for oxygen and nitrogen occurring with acid mixed with water in the proportions which give maximum contraction on mixing. This is not the case, however, with hydrogen chloride; the quantity of this gas absorbed by aqueous solutions of sulphuric acid diminishes with increasing concentration of the latter to a minimum at about 89 per cent, sulphuric acid.

Sulphuric acid is a good solvent for most solids, many organic substances in particular dissolving in it readily; the molecular weights of substances dissolved in pure sulphuric acid are generally lower than the commonly accepted values, probably due to the formation of dissociable additive compounds between solute and solvent.

Behaviour towards Water

Water is eagerly absorbed by sulphuric acid, and the concentrated acid is commonly applied as a dehydrating agent for inert gases and organic compounds; with the latter the effect is so marked that compounds containing hydrogen and oxygen, for example the carbohydrates and tartaric acid, are readily carbonised by the acid. The preparation of carbon monoxide from formic acid or oxalic acid is also due to this dehydrating action, but it is remarkable that sulphuric acid of 100 per cent, concentration acts on oxalic acid very slowly in comparison with the slightly diluted acid, the addition of 0.1 per cent, of water increasing the velocity of reaction to seventeen times that observed with the pure acid.

On mixing with water much heat is liberated, 17.7 Calories being liberated per gram-molecule H2SO4 if a large excess of water is used; over half this quantity of heat is liberated during the addition of the first two molecular proportions of water. It is uncertain to what extent the evolution of heat is due respectively to hydrate formation and to mere dilution.

As the ordinary heat of formation of sulphuric acid from its elements is 193.0 Calories, the heat of formation from the elements in the presence of a large quantity of water is 210.7 Calories.

In contrast to its behaviour with water, sulphuric acid with ice forms freezing mixtures, the temperature falling from 0° C. even to -24° C., the result being still more satisfactory if the acid is previously diluted with a bi- to ter-molecular proportion of water. The actual absorption of heat is due to the enforced melting of a relatively large quantity of ice.

A mixture of water with sulphuric acid occupies a smaller volume than the two unmixed constituents, considerable contraction being observable, the maximum effect occurring with 67 to 69 per cent, of acid by weight. With gradual increase in concentration the acid does not exhibit a continuous increase in density, but attains a maximum specific gravity of 1.8442 at 15° C. at a concentration of 97.5 per cent., so that the specific gravity is of little value in testing the acid concentration between 95 and 100 per cent. The following table gives an indication of this abnormal behaviour, the specific gravities being given for the acid at 15° C. compared with water at the same temperature.

Specific Gravity.Per cent. HSO4.

Between 66 and 81 per cent, the relation between the concentration of sulphuric acid and the specific gravity is represented very closely by the formula

x = 86 S-69.00,

where x is the concentration and S the specific gravity at 15° C. referred to water at the same temperature. The concentration of stronger acid can be determined by diluting with a known proportion of water so that the weakened acid falls somewhere within the stated limits of- concentration, then finding the specific gravity.

Another interesting physical method for determining the percentage of sulphuric acid in the concentrated acid is to measure the contraction on mixing with water. If in a 300 c.c. graduated flask 200 c.c. of the acid under investigation are added to 100 c.c. of water, then after restoring the temperature to 15° C. a contraction is evident; its amount can be measured by the addition of an inert liquid (for example vaseline oil) from a burette to the mixture until the 300 c.c. mark of the flask is reached by the level of the liquid. The contraction with acid of various strengths is as follows:

Percentage of Sulphuric Acid.Contraction, c.c.

The aqueous vapour pressure of diluted sulphuric acid has been determined over a large range of concentration and temperature. The results given in the following table for temperatures up to 90° C. are due to Sorel, and those from 100° to 200° C. to Briggs.

Vapour pressure of aqueous sulphuric acid (in mm. Hg)

Concentration H2SO4 (per cent.). \ Temperature, ° C102030405060708090
444.48.515.528.148.3. . .. . .. . .. . .

Concentration H2SO4 (per cent.). \ Temperature, ° C100120140160180200
77.5120.271.0168.5. . .. . .. . .
81.818.532.575.7207.5. . .. . .
91.22.. .. . .10.025.062149

Hydrates of Sulphuric Acid

The avidity of sulphuric acid for water naturally suggests the possible formation of definite compounds or hydrates, and the existence of such hydrates has been demonstrated in various ways.

Freezing point SO3-H2O
Freezing point Curve for SO3-H2O.
The freezing-point curve for mixtures of two substances is usually regarded as affording definite evidence of the occurrence or otherwise of chemical combination between the constituents, and with mixtures of sulphuric acid and water very definite results are obtained. With sulphur trioxide and water as the two components, maxima on the freezing-point curve (fig.) occur at proportions corresponding with the compositions of the following compounds: SO3, H2S2O7, H2SO4, H2SO4. H2O, H2SO4.2H2O, H2SO4.4H2O, H2O. Thus the existence of a mono-, a di- and a tetra-hydrate of sulphuric acid in the solid condition is proved, but the possibility of other hydrates, for instance H2SO4. 3H2O, 3H2SO4.H2O and H2SO4.12H2O, under other conditions, is not excluded. The final portion of the curve towards 100 per cent. SO3 is uncertain owing to the nature of sulphur trioxide, and is therefore not included in the figure.

Other indications of the existence of hydrates are found in the curve representing the electrical conductivity of solutions of various concentrations, minima occurringat the compositions SO3, H2SO4, H2SO4. H2O, H2O; also in the refractive index curve 2 a maximum is found at the proportion H2SO4.H2O, whilst sharp breaks occur at H2SO4, H2SO4.2H2O and H2SO4.4H2O. The viscosity curve 3 also indicates a series of compounds SO3, H2S2O7, H2SO4, H2SO4.H2O, H2O, whilst the thermal expansion of sulphuric acid of various concentrations gives indications of the hydrates H2SO4.H2O and H2SO4.2H2O.Particularly clear evidence of hydrate formation is available from determinations of the molecular weight of various mixtures of sulphuric acid and water, using acetic acid as cryoscopic solvent, the results demonstrating the occurrence of H2SO4.H2O and H2SO4.2H2O, and also the probability of a higher hydrate, H2SO4.3H2O, and a lower one, 2H2SO4.H2O.

The relation between specific gravity and concentration gives no definite confirmation of the presence of hydrates in diluted sulphuric acid.

Sulphuric acid monohydrate, H2SO4.H2O, crystallises from acid of the corresponding concentration in hexagonal prisms which melt at +9° C.; the dihydrate, H2SO4.2H2O, melts at -37° C., and the tetrahydrate, H2SO4.4H2O, at -24.5° C. These hydrates are relatively unstable and in the molten condition are dissociated to a considerable extent, behaving as ordinary dilute sulphuric acid. As is indicated in fig., the melting-points of these hydrates are lowered very considerably by the presence of a slight excess of either constituent.

It is worthy of note that the physical properties of mixtures of sulphuric acid and ethyl ether also indicate the formation of additive compounds, for example H2SO4.(C2H5)2O, in which the ether may be regarded as functioning in the same manner as does the molecule of water in the hydrate H2SO4.H2O. Ethyl alcohol, which is analogous in many ways to water, behaves in a different manner with sulphuric acid, giving ethyl hydrogen sulphate, (C2H5)HSO4, and other products.

Ionisation of Sulphuric Acid

In aqueous solution sulphuric acid undergoes electrolytic dissociation with formation of H and HSO4' ions, the latter dissociating in part still further into H and SO4'' ions, especially if the solution is dilute. Examination by the various available physico-chemical methods shows that, regarded from the point of view of the first stage of dissociation or, in other words, considering the acid as monobasic, sulphuric acid is comparable in strength (as distinct from concentration) with hydrochloric and nitric acids; the second stage, HSO4' ⇔ H + SO4'', however, has much less tendency to proceed to completion, and, in consequence of this, sulphuric acid, acting as a dibasic acid, is considerably inferior in strength to nitric acid and the halogen hydracids, excluding hydrofluoric acid. Its ability to displace these acids completely from their salts is dependent on its lower volatility.


On electrolysis of acid below 20 per cent, concentration the products are mainly hydrogen and oxygen (with ozone), but with more concentrated acid the conversion of HSO4' ions at the anode into free HSO4 groups is succeeded by the coupling of the latter to form molecules of perdisulphuric acid, H2S2O8, which changes more or less rapidly according to the concentration into permonosulphuric acid, H2SO5; this last product, again, can undergo further change, giving rise to hydrogen peroxide.

Sulphuric acid of 98 per cent, strength undergoes decomposition on electrolysis. At 50° c. hydrogen, hydrogen sulphide and sulphur are liberated at the cathode; at higher temperatures the cathodic products are sulphur dioxide and sulphur, until at 300° C. only sulphur is liberated. Up to 200° c. oxygen alone is evolved at the anode, but with rise in temperature sulphur dioxide also appears, resulting from oxidation of sulphur which has diffused from the cathode chamber; above 280° c. the two gases leave the anode in quantities in accordance with Faraday's Law.

Chemical Properties of Sulphuric Acid

Concentrated sulphuric acid possesses marked oxidising power, especially in the presence of certain metallic salts, such as those of mercury and copper. Gaseous hydrogen begins to undergo appreciable oxidation by sulphuric acid at 160° C., but nascent hydrogen or hydrogen in contact with finely divided platinum can effect the reduction of the acid even at the ordinary temperature. If a mixture of hydrogen and sulphuric acid vapour is passed over silica heated at 700° to 900° c., quantitative reduction to hydrogen sulphide occurs. The acid slowly oxidises carbon to carbon dioxide, this process occurring in the Kjeldahl process for the estimation of nitrogen in organic substances. The reaction with carbon monoxide proceeds according to the equation

H2SO4 + CO = CO2 + SO2 + H2O,

as long as the acid concentration does not fall below 91 per cent.; certain catalysts, e.g. mercury, silver, selenium, palladium and iridium, but not platinum or osmium, accelerate this reaction. Phosphorus is oxidised to phosphoric acid by sulphuric acid, and sulphur to sulphur dioxide, in both cases the assistance of heat being necessary.

Hydrogen sulphide reacts with concentrated solutions of sulphuric acid yielding sulphur and sulphurous acid, thus:

H2S + H2SO4 = H2S>O3 + H2O + S.

At the ordinary temperature the concentration of the acid must not be much less than 25N for separation of sulphur to occur. It has been suggested, therefore, that it is really pyrosulphuric acid, H2S2O7, which reacts with the hydrogen sulphide, this acid being present in small equilibrium quantity in the concentrated acid solution. The first stage of the reaction would thus be:

H2S2O7 + H2SH2S3O6 + H2O.

The trithionie acid would then decompose, ultimately yielding sulphurous acid and sulphur.

Hydrobromic acid reduces sulphuric acid to sulphur dioxide, as also does hydriodic acid, but with the latter in high concentration the reduction goes further, producing sulphur and hydrogen sulphide. For this reason when metallic bromides and iodides are treated with sulphuric acid, the halogen element is liberated, the temperature at which this becomes evident depending on the concentration of the acid; thus, with the potassium salts, the following observations have been made:

Concentration of acid, per cent. H2SO42530354050
Temperature, ° C at which I2 detected10060504536
Concentration of acid, per cent. H2SO460708090. . .
Temperature, ° C at which Br2 detected1691004638. . .

Stannous chloride also causes reduction to hydrogen sulphide.

Sulphuric acid is oxidised to permonosulphuric, perdisulphuric and fluorosulphonie acids when an ice-cold aqueous solution is treated with fluorine. In addition to the foregoing products, ozone and a very unstable compound believed to be a tetroxide (SO4 or S2O8) are also produced. A similar oxidation occurs when fluorine is passed into cold solutions of alkali sulphates or hydrogen sulphates.

Sodium and potassium attack the pure acid even in the cold, the action with hot acid being explosive. Other metals, for example iron, zinc, magnesium and manganese, which liberate hydrogen readily from hydrochloric acid, behave similarly with dilute sulphuric acid, but with the concentrated acid the reaction is more sluggish, generally requiring to be aided by heat, and, as any hydrogen which might be produced becomes oxidised by the acid, the gaseous product is sulphur dioxide. Silver, lead, mercury and copper are attacked only by the hot concentrated acid; by-reactions frequently occur, such as the formation of some metallic sulphide due to reduction of sulphate, and in the case of copper not only are cupric sulphate and cuprous sulphide commonly obtained, but cuprous sulphate may also be found in solution; indeed, the reaction, if carried out at about 200° c., provides a means of preparing the last-named salt. Platinum is attacked very appreciably by sulphuric acid at 250° c., the effect being diminished by the presence of reducing agents such as carbon, sulphur, sulphur dioxide or arsenious oxide.

The interaction of a mixture of nitric oxide and nitrogen dioxide with sulphuric acid has already been referred to nitric oxide alone is only slightly absorbed by the pure acid and not appreciably by the somewhat diluted acid such as is used in the nitrometer.

Mention has been made already of the action of phosphorus penta- chloride on sulphuric acid.

Applications of Sulphuric Acid

In addition to various purposes already mentioned, sulphuric acid is applied on an immense scale in the inorganic chemical industry for the manufacture of acids such as hydrochloric, nitric, hydrofluoric and phosphoric acids, and directly or indirectly in the production of many other chemical substances such as phosphorus, chlorine, bromine, iodine, sodium carbonate, hydrogen peroxide and, of course, sulphates. The "superphosphate of lime" industry, by which ordinary insoluble calcium phosphate is converted into a form which can act as a fertiliser, supplying phosphorus in a condition suitable for absorption by the roots of plants, consumes immense quantities of sulphuric acid, as also does the production of ammonium sulphate in coke ovens and gas works. Much sulphuric acid is also used in electric batteries of various types, accumulator cells in particular requiring a dilute acid of a high degree of purity. Metallurgical industry also calls for large quantities of sulphuric acid, especially for use in " pickling " metals, i.e. cleansing metallic surfaces from rust and dirt; thus, in South Wales, steel plate is almost exclusively pickled with sulphuric acid before tinning.

The organic chemical industry likewise consumes much sulphuric acid for a wide variety of purposes. The acid is not only used for the preparation of sulphonic acids, a process of particular importance in the coal tar dye industry, but is also applied as an oxidising agent, for example in the production of phthalic anhydride from naphthalene, which is also an important step in the manufacture of various dyes such as indigo and eosin; in applying the sulphuric acid for sulphonation or for oxidation purposes, traces of mercury compounds frequently exert a marked catalytic influence. Sulphuric acid is also used in the purification of various kinds of oils, especially mineral oils, and directly or indirectly in the hydrolytic decomposition of various animal and vegetable oils and fats for the production of glycerine and fatty -acids, the latter being further convertible into soap and candles. The conversion of paper into "vegetable parchment," of viscose into artificial silk, and of starch into glucose for various purposes, are also processes involving sulphuric acid. Together with nitric acid, sulphuric acid is applied to the production of many organic nitro-compounds, for example nitrobenzene, dinitrobenzene, picric acid and the various nitrotoluenes, which are of importance for the manufacture of dyes and explosives, and of certain organic nitrates, in particular the misnamed nitrocellulose and nitroglycerine, which should more correctly be termed cellulose nitrate and glyceryl nitrate.

The use of weak solutions of sulphuric acid (1 to 2 per cent.) as a weed spray has been suggested.

Molecular Weight and Constitution of Sulphuric Acid

Allowing for the effect of electrolytic dissociation, the behaviour of sulphuric acid in aqueous solution is consistent with the formula H2SO4. In the vapour state also, allowance being made for gaseous dissociation, the molecular weight agrees fairly well with this formula, except that at temperatures in the neighbourhood of the boiling-point there is distinct evidence of partial association into double molecules. In the liquid state it is probable that the pure acid is associated into larger molecules, the low vapour pressure and the surface tension supplying evidence in favour of a higher molecular weight than H2SO4, which receives confirmation from the ready formation of double sulphates. It may be assumed that formation of the larger molecules occurs by the oxygen atoms acting as links, as expressed in the bimolecular formula

which is analogous with the formula already proposed for dimeric sulphur trioxide.

In accordance with the dualistic views of earlier days, the composition of sulphuric acid was at one time expressed by the formula H2O.SO3, but to-day the formula SO2(OH)2 is generally adopted, commonly with the assumption of sexavalent sulphur, so that the structural formula is

The presence of the two hydroxyl groups in the molecule is shown by the action of chlorine and of phosphorus pentachloride, which respectively produce chlorosulphonic acid, HSO3Cl, and pyrosulphuryl chloride, S2O5Cl2. Moreover, both these products with water yield sulphuric and hydrochloric acids. Since the alkali alkylsulphates, for example C2H5O.SO2.OK, do not exist in two isomeric forms, it is evident that the two hydroxyl groups are symmetrically placed in the molecule. That the hydroxyl groups are directly attached to the sulphur atom may be deduced from such transformations as the following: Concentrated sulphuric acid reacts directly with benzene on heating to produce benzenesulphonic acid; this with phosphorus pentachloride yields the corresponding chloride, which on reduction gives thiophenol,


The latter compound on oxidation again yields the sulphonic acid. Again, sulphuryl chloride reacts with benzene in the presence of aluminium chloride to give diphenylsulphone, (C6H5)2SO2, a compound which can also be obtained by oxidation of diphenyl sulphide, (C6H5)2S. On the assumption that the phenyl radicals remain attached to the sulphur atom throughout these, transformations, it is obvious that the hydroxyl radicals in sulphuric acid must also be directly attached to sulphur.

The somewhat remarkable fact that many sulphates when heated retain one molecule of water of crystallisation (called " water of constitution ") very persistently, whilst the other molecules are eliminated with relative ease, this phenomenon being observable with copper, manganese, ferrous, nickel, cobalt, magnesium and zinc sulphates, has led to a suggestion that these monohydrates, of the general formula M2SO4. H2O, are actually to be regarded as acid salts derived from a sulphuric acid of the constitution O=S(OH)4, which corresponds with the composition of the monohydrate of sulphuric acid.

Detection and Estimation of Sulphuric Acid

The usual dry test for a sulphate is reduction on charcoal in the presence of sodium or potassium carbonate; alkali sulphide in the fused product can easily be detected by moistening on a clean silver coin or by the application of other suitable tests.

An insoluble sulphate may be detected by the formation of turpeth mineral on the addition of a 10 per cent, solution of mercuric nitrate in dilute nitric acid (1 in 100); the small yellow tetragonal crystals of the basic sulphate may be identified microscopically. The reaction takes place immediately in the cold with calcium and mercurous sulphates, less readily with strontium and lead sulphates, whilst with barium sulphate boiling is necessary.

For solutions containing sulphuric acid or a sulphate the reagent commonly applied is barium chloride, both when the test is to be qualitative and when quantitative. Precipitation is effected by the gradual addition of barium chloride to the boiling solution containing a little hydrochloric acid, but for the production of pure barium sulphate, and therefore in order to ensure accuracy, certain precautions must be observed. Nitrates, perchlorates, phosphates, tervalent metals and large quantities of salts of the alkali metals (particularly potassium) and of the alkaline earth metals are to be avoided, as they cause the precipitated barium sulphate to be rendered impure by occlusion of otherwise soluble substances. Such impurities may be accounted for partly by reaction between the sulphion and the intermediate ions of any ternary electrolytes present; thus, in the presence of excess of barium chloride the reaction

SO4'' + 2BaCl = Ba2Cl2SO4

tends to increase the weight of the precipitate, whilst in the presence of potassium sulphate the opposite effect results from the reaction

2KSO4' + Ba•• = BaK2(SO4)2.

Another source of error in this method of estimating sulphuric acid appears to be in the formation of the complex ion [Ba5(SO4)6]'', the potassium salt of which has been isolated. It is essential that the barium chloride be added slowly, and the precipitate should be collected from the hot solution and washed with hot aqueous acetic acid.

Various modifications by which the reaction with barium chloride may be subjected to volumetric treatment have been suggested. Thus the precipitant may be added in excess and suitably back-titrated. Other reagents which may be applied to the volumetric determination of sulphuric acid and sulphates are barium chromate and benzidine. In the case of the former, the solution of sulphate is precipitated by a solution of barium chromate in hydrochloric acid; on subsequent neutralisation of the filtrate, a quantity of chromic acid, equivalent to the barium sulphate which has been precipitated, remains in solution and may be estimated iodometrically. Benzidine, on the other hand, is an organic base which forms a very sparingly soluble sulphate; the solution of mineral sulphate is treated with a solution of benzidine hydrochloride and the precipitated benzidine sulphate removed by filtration; when subsequently suspended in pure water the benzidine sulphate undergoes hydrolysis to a sufficient extent to permit titration of the sulphuric acid with standard alkali. Lead nitrate may also be used in the presence of alcohol to titrate sulphuric acid, a few drops of potassium iodide being added as indicator. Concordant but slightly low results are obtained.

Another volumetric process which has been recommended consists in reducing the sulphuric acid or sulphate (excepting barium sulphate) by heating with a mixture of hydriodic acid, phosphorus and phosphoric acid, absorbing the resulting hydrogen sulphide in a solution of zinc acetate, and estimating the precipitated zinc sulphide iodometrically.

The importance of the determination of sulphuric acid and sulphates lies very largely in the fact that sulphur in various forms of combination, and also free sulphur, is frequently determined quantitatively by primary conversion into sulphuric acid or sulphate, followed by actual estimation in the latter form. For example, a convenient volumetric process applicable to the Carius method for determining sulphur in organic compounds is to neutralise the solution of sulphuric acid after removal of the excess of nitric acid by evaporation, and then add silver nitrate, by which the sulphate is converted into silver sulphate. This is separated from the excess of silver nitrate by means of its insolubility in alcohol, and is then estimated by dissolving in dilute nitric acid and titrating the silver with standard thiocyanate solution.

For solutions containing sulphuric acid only, direct titration with standard alkali, and measurement of the specific gravity, are possible as methods of estimation, provided that the process in either case is, if necessary, preceded by suitable dilution. Thermometric methods have also been suggested, depending on the rise in temperature when the acid is mixed with water, or when titrated with barium chloride solution. The water content of the concentrated acid may be determined by similar titration with oleum which has been standardised thermometrically by 80 per cent, sulphuric acid.

An electrometric method for determining soluble sulphates consists in precipitation of the latter by the addition of a measured excess of a standard solution of lead nitrate in the presence of alcohol, removing the lead sulphate by filtration, washing it with alcohol, and titrating the unchanged lead nitrate in the filtrate electrometrically with standard ferrocyanide solution.

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