Chemical elements
  Sulphur
    Isotopes
    Energy
    Extraction
    Refining
    Applications
    Allotropy
    Crystalline
    Amorphous Sulphur
    Colloidal Sulphur
    Physical Properties
    Chemical Properties
    Detection
    Estimation
    Compounds
      Hydrogen Sulphide
        Metallic Sulphides
        Detection and Estimation
      Metal Polysulphides
      Hydrogen Polysulphides
      Hydrogen Pentasulphide
      Hydrogen Trisulphide
      Hydrogen Disulphide
      Sulphur Monofluoride
      Sulphur Tetrafluoride
      Sulphur Hexafluoride
      Sulphur Monochloride
      Sulphur Dichloride
      Sulphur Tetrachloride
      Sulphur Monobromide
      Thionyl Fluoride
      Sulphuryl Fluoride
      Fluorosulphonic Acid
      Thionyl Chloride
      Sulphuryl Chloride
      Sulphur Oxytetrachloride
      Pyrosulphuryl Chloride
      Chlorosulphonic Acid
      Thionyl Bromide
      Sodium Sulphoxylate
      Sulphur Dioxide
      Sulphurous Acid
      Sulphites
      Sulphur Trioxide
      Pyrosulphuric Acid
      Pyrosulphates
      Sulphuric Acid
      Persulphuric Anhydride
      Persulphuric Acid or Perdisulphuric Acid
      Perdisulphates
      Permonosulphuric Acid
      Amidopermonosulphuric Acid
      Thiosulphuric Acid
      Thiosulphates
      Polythionic Acids
      Dithionic Acid
      Trithionic Acid
      Trithionates
      Tetrathionic Acid
      Tetrathionates
      Pentathionic Acid
      Pentathionates
      Wackenroders Solution
      Hexathionic Acid
      Polythionic Acids
      Sulphur Sesquioxide
      Hydrosulphurous Acid
      Hydrosulphites
      Nitrogen Sulphide
      Nitrogen Persulphide
      Nitrogen Pentasulphide
      Sulphammonium
      Hexasulphamide
      Nitrogen Chlorosulphide
      Trithiazyl Chloride
      Thiotrithiazyl Chloride
      Dithiotetrathiazyl Chloride
      Nitrogen Bromosulphide
      Thiotrithiazyl Bromide
      Thiotrithiazyl Iodide
      Thiotrithiazyl Nitrate
      Thiotrithiazyl Hydrogen Sulphate
      Thiotrithiazyl Thiocyanate
      Thionylamide
      Sulphamide
      Imidodisulphamide
      Sulphimide
      Sulphonic Acids
      Amidosulphonic Acid
      Imidosulphonic Acid
      Nitrilosulphonic Acid
      Hydroxylamine-monosulphonic Acid
      Nitrososulphonic Acid
      Hydroxylamine-disulphonic Acid
      Hydroxylamine-isodisulphonic Acid
      Hydroxylamine-trisulphonic Acid
      Dihydroxylamidosulphonic Acid
      Sulphazinic Acid
      Sulphazotinic Acid
      Dehydrosulphazotinic Acid
      Nitrosulphonic Acid
      Nitrosulphonyl Chloride
      Nitrosulphonic Anhydride
      Nitrosulphuric Acid
      Nitrosodisulphonic Acid
      Sulphonitronic Acid
      Sulphates of Hydroxylamine
      Hydroxylamine Dithionate
      Hydrazine Dithionate
      Hydrazine Amidosulphonate
      Carbon Subsulphide
      Carbon Monosulphide
      Carbon Disulphide
      Thioformaldehyde
      Thiocarbonic Acid
      Ammonium thiocarbonate
      Thiolcarbonic Acid
      Xanthic Acid
      Perthiocarbonic Acid
      Sodium perthiocarbonate
      Carbonyl Sulphide
      Thiocarbonyl Chloride
      Thiocarbonyl Tetrachloride or
      Carbon Hexachlorosulphide
      Trichloromethyl Disulphide
      Thiocarbonyl Sulphochloride
      Carbon Bromosulphide
      Amino-derivatives of Thiocarbonic Acid
      Dithiocarbamic Acid
      Thiocarbamide
      Azidodithiocarbonic Acid
      Thiocyanogen
      Cyanogen Monosulphide
      Cyanogen Trisulphide
      Sulphur Thiocyanate
      Disulphur Dithiocyanate
      Thiocyanic Acid
      Thiocyanates
      Dithiocyanic Acid
      Trithiocyanuric Acid
      Perthiocyanic Acid
      Perthiocyanogen
      Sulphates

Hydrogen Sulphide, H2S





Occurrence of Hydrogen Sulphide

Hydrogen Sulphide, also known as Sulphuretted Hydrogen or Hydrosulphuric Acid, H2S, is present in volcanic gases, probably owing to the action of steam on sulphides or sulphur at a high temperature in the earth. It is also found in the waters of certain spas, as for example Harrogate and Strathpeffer in this country and Aix-les-Bains on the Continent; the hydrogen sulphide of such "sulphurous" waters has probably been formed at least in part by the biochemical reduction of mineral sulphates. Not many bacteria able to reduce sulphates are known, however, and those that are known are non-sporogenic and strictly anaerobic. The reducing action appears to be associated with the oxidation of organic matter, whence the necessary energy is derived. Decomposing organic matter, especially of animal origin, also frequently gives rise to hydrogen sulphide, due to the decomposition of the albuminoid substances under the influence of micro-organisms.


History

The fact that sulphur is soluble in aqueous solutions of alkaline substances was known to the alchemists, who realised that it could again be liberated by acidifying, but although the simultaneous liberation of a "sulphurous air" is mentioned in their records, the gas was not submitted to careful examination. During the phlogistic period the combustibility of the gas was discovered as well as its precipitant action on solutions of metallic salts. In 1777, Scheele obtained it by the action of acids on calcium polysulphide and also on manganese sulphide and ferrous sulphide; he observed the solubility of the gas in water and its oxidation to free sulphur by atmospheric air and other oxidising agents. On account of the phlogistic views prevalent at the time, however, Scheele and his contemporaries failed to recognise the real nature of the gas, which received such names as "liver of sulphur air," "hepatic air." The gas was first recognised as an oxygen-free acid by Berthollet in 1796.

Formation and Preparation of Hydrogen Sulphide

From its Elements

(a) Above 200° C. hydrogen and sulphur interact with appreciable velocity, forming hydrogen sulphide. Below 350° C. the combination proceeds slowly until one or other of the reagents is entirely consumed, but above this temperature, although the reaction is naturally more rapid, the final product is an equilibrium mixture, the change being representable thus:

H2 + SH2S.

The higher the temperature the lower the proportion of hydrogen sulphide in the equilibrium mixture.

A careful examination of the velocity of the reaction under varied conditions indicates that it is proportional to the pressure of the hydrogen and to the square root of the pressure of the sulphur vapour; this result is interpreted as due to the reaction occurring in stages, the first being a slow, reversible change, S8 ⇔ 4S2, followed by a very rapid dissociation, S2 ⇔ 2S, the combination of molecular hydrogen and atomic sulphur then proceeding with measurable velocity. Platinum black or red phosphorus accelerates the combination, as also does exposure to ultra-violet light.

The formation of hydrogen sulphide in this manner can be realised experimentally by passing hydrogen or even purified coal gas into sulphur boiling in a flask, when the issuing gas will contain appreciable quantities of hydrogen sulphide. Under pressures of 5 to 10 atmospheres the reaction proceeds at 250° to 300° C., giving a satisfactory yield if the liquid sulphur is subjected to vigorous agitation.

(b) Sulphur can also be reduced to hydrogen sulphide at the ordinary temperature if "nascent" hydrogen is used; thus, powdered sulphur yields some hydrogen sulphide if treated with aluminium, tin, iron or zinc and hydrochloric acid, the result being improved by the additional presence of acetic acid or alcohol, which will increase the solubility of the sulphur. The reduction can also be effected electrolytically by having powdered sulphur in contact with a platinum cathode immersed in dilute acid, e.g. hydrochloric acid.

When heated with hydrogen iodide or concentrated aqueous hydriodic acid, sulphur is reduced to hydrogen sulphide, but once more the reaction is incomplete, leading only to an equilibrium mixture:

2HI + SH2S + I2

Other processes which are known to produce hydrogen sulphide from sulphur include treatment with steam or water (e.g. in a sealed tube at 200° C. or higher temperature), heating with organic matter (e.g. a mixture of vaseline with sulphur in the proportions 7: 3 gives rise to very pure hydrogei) sulphide when heated), and the reducing action of certain anaerobic micro-organisms in the presence of water at the ordinary temperature, as well as ordinary alcoholic fermentation.

From Sulphur-Oxygen Compounds

Sulphites in aqueous solution are easily reduced to hydrogen sulphide by nascent hydrogen, produced, for example, by the interaction of zinc and dilute sulphuric acid:

H2SO3 + 6H = 3H2O + H2S.

The reduction of sulphites and of sulphates can also be effected by certain bacteria, the presence of hydrogen sulphide in some mineral waters probably being due to reduction of calcium sulphate in this way. The first product is probably the corresponding sulphide, which subsequently undergoes hydrolytic decomposition.

Hydrogen sulphide is also formed when sulphur dioxide is carried in a blast of steam through red-hot coke; the actual reduction is probably effected by hydrogen, since this gas is known to reduce sulphur dioxide at a dull red heat.

From Sulphides

Almost all the methods commonly employed for the production of sulphuretted hydrogen belong to this class.

Some sulphides, e.g. aluminium sulphide, are decomposed in the cold by water, with liberation of hydrogen sulphide:

Al2S3 + 6H2O = 2Al(OH)3 + 3H2S.

The sulphides of the alkali metals and of the alkaline earth metals are readily decomposed by weak acids, even by carbonic acid. By the use of these sulphides and dilute sulphuric acid it is possible to obtain hydrogen sulphide in a high degree of purity. The hydrosulphides of the alkaline earth metals are convenient sources of hydrogen sulphide, yielding the gas on treatment with carbon dioxide or even by merely heating. The carbon dioxide method is applied on a technical scale in the extraction of sulphur from alkali waste.

When the gas is intended for ordinary laboratory purposes, ferrous sulphide and hydrochloric acid are the reagents commonly employed for preparing hydrogen sulphide. The reaction occurs readily at the ordinary temperature, but as the ferrous sulphide is produced by heating together iron and sulphur, it commonly contains at least traces of metallic iron which cause the evolved gas to be contaminated with hydrogen, in addition to impurities such as arsine due to the presence of impurities in the iron. In order to avoid these impurities the use of precipitated ferrous sulphide has been suggested, but such a process would have various disadvantages in addition to increased cost. Manganese sulphide and zinc sulphide have also been recommended in place of ferrous sulphide.

Hydrogen Sulphide Generator
Improved Hydrogen Sulphide Generator
The chief defects of the ordinary laboratory forms of hydrogen sulphide generators result from the comparative slowness of the reaction between the acid and the sulphide, especially when the concentration of the former is low. The apparatus described below (fig.) is more suitable for meeting continuous heavy demands, providing a rapid evolution of the gas from minimum quantities of acid, which undergoes complete neutralisation. The reaction is brought about at a temperature in the neighbourhood of 100° C. by surrounding the sulphide container with steam; a short air condenser at A ensures practically complete condensation of the outgoing steam, so that little attention is necessary. The acid holders contain commercial hydrochloric acid diluted with an equal volume of water, and when the tap B is first opened, if the sulphide is thoroughly heated reaction takes place with almost explosive violence and each drop of acid is soon completely exhausted. A too sudden entry of acid is prevented by the insertion at D of a piece of capillary tubing, 3 cm. long and 1 mm. bore. In order to facilitate the displacement of air from the reservoir K by the incoming spent acid, the tube L has a number of holes blown in it.

For the production of hydrogen sulphide free from uncombined hydrogen, the mineral stibnite is frequently used; this reacts with concentrated hydrochloric acid, but the reaction is slow unless aided by warming:

Sb2S3 + 6HCl = 2SbCl3 + 3H2S.

Purification of Hydrogen Sulphide

The commonest impurities in hydrogen sulphide are free hydrogen and arsine. The latter, which is due to the presence of arsenic in the reagents, can be removed by chemical means, for example by passing over heated potassium polysulphide ("liver of sulphur") at 350° C., the arsenic being converted into potassium thioarsenate. Other methods of treatment for the removal of arsenic include the action of solid iodine at the ordinary temperature, by which the arsine is converted into arsenic iodide and hydrogen iodide, whilst the hydrogen sulphide passes on almost unaltered and can be freed from hydrogen iodide by washing with water. Mere passage of the gas through a glass tube packed with glass fragments at a dull red heat is also effective in causing decomposition of the arsine.

Hydrogen sulphide is generally dried by passage through anhydro,us calcium chloride, but as this may lead to the introduction of small quantities of hydrogen chloride, calcium chloride is not so satisfactory a drying agent as phosphoric oxide.

Moissan, in 1903, applied a very elegant method to the removal of all impurities from hydrogen sulphide. The gas was dried by slow passage through two or three glass tubes at -50° to -70° C., which procedure is as effective as the chemical drying agents commonly applied. The hydrogen sulphide was then collected as a solid in a glass tube immersed in liquid air, and any air or free hydrogen present was drawn away by exhausting with a pump. On allowing the mass to warm gradually, it melted, and then attained a state of ebullition, the pure gas being collected when the boiling-point was constant.

Physical Properties of Hydrogen Sulphide

Hydrogen sulphide is a colourless gas with the unpleasant odour which is commonly associated with a bad egg, the smell of which is actually due largely to this gas. It is 1.189 times as dense as air and one litre at N.T.P. weighs 1.539 grams.

In the formation of hydrogen sulphide from its elements heat is evolved, gaseous hydrogen sulphide when referred to hydrogen gas and solid sulphur being exothermic to the extent of 2.73 Calories per gram- molecule. The specific heat at constant pressure is 0.2423, the ratio Cp/Cv at 20° C. having the value 1.315. The heat capacity decreases with increasing temperature.

Water dissolves hydrogen sulphide fairly readily at the ordinary temperature, the solubility following Henry's Law. Solution is accompanied by an evolution of 4.56 Calories per gram-molecule of gas dissolved.

The following table gives the volume of gas, corrected to N.T.P., which can be absorbed by one cubic centimetre of water under a hydrogen sulphide pressure of 760 mm.:

°C.Absorption Coefficient.°C.Absorption Coefficient.°C.Absorption Coefficient.
04.621252.257701.010
53.935302.014800.906
103.362401.642900.835
152.913501.3761000.800
202.554601.176


It will be seen that the solubility decreases rapidly as the temperature is raised, and this is still more evident in the following table, in which the solubility (volume of gas, corrected to N.T.P., absorbed by one volume of water at 760 mm. pressure) at various temperatures is given (the figures are somewhat less accurate than the foregoing). For the purpose of comparison analogous figures are given for alcohol, in which hydrogen sulphide is more soluble:

°C.Solubility.
In Water.In Alcohol.
04.3717.89
53.9714.78
103.5911.99
153.239.54
202.917.42
252.615.62
302.33. . .
352.08. . .
401.86. . .


The gas was first liquefied in 1823 by M. Faraday, who, twenty-two years later, successfully reduced the temperature by means of a mixture of solid carbon dioxide and ether to such a degree that the substance solidified. In Faraday's method, using an inverted U-tube, liquid hydrogen sulphide is obtained by placing in one limb materials such as hydrogen polysulphide or charcoal saturated with hydrogen sulphide, from which the necessary hydrogen sulphide is liberated by gently warming; alternatively, the material enclosed in the one arm of the tube may be ferrous sulphide and concentrated hydrochloric acid, from which the hydrogen sulphide is obtained by subsequent interaction. If the empty arm of the inverted U-tube is then sufficiently cooled, the gas liquefies there under its own pressure. At the present day, however, on account of the ease with which low temperatures are obtainable, the liquefaction and solidification of hydrogen sulphide present no difficulties, passage of the gas into a tube cooled externally by liquid air being sufficient to produce the solid.

On elimination of sulphur from hydrogen sulphide an equal volume of hydrogen remains, which of course is evidence for the composition represented by the formula H2S. Vapour density determinations also show that the gas consists of single H2S molecules, a value 34.085 having been obtained for the molecular weight, after making due corrections to the vapour density result. In the liquid state also there appears to be no association of the molecules and the molecular formula is still H2S.

Liquid hydrogen sulphide is a colourless, very mobile fluid, which resembles the dry gas in being neutral to litmus. Its vapour pressure increases rapidly with rise of temperature, finally attaining a critical pressure of 88.9 atmospheres at the critical temperature, 100.4° C.

Temperature18.2° 50°100° C
Vapour pressure (atmospheres)10.2516.9535.5688.7


At its ordinary boiling-point, -60.2° C. (760 mm.), the density of the liquid is 0.964. The solidified substance forms a crystalline, snowlike mass, m.pt. -83° C., of greater density than the liquid. The liquid is more highly refractive than water (nD = 1.384 at 18.5° C.); the surface tension at -60° C. is 25.43 dynes per cm., and the dielectric constant 10.2 (air = 1).

In the liquid state, as in the gaseous condition, hydrogen sulphide consists of simple H2S molecules. Like water, it is practically a nonconductor of electricity, its specific conductivity at about -80° C. being 1×10-11 mho; it is a good solvent for many substances; it dissolves sulphur without combining with it; hydrogen chloride and hydrogen bromide yield solutions in hydrogen sulphide which, in contrast to their aqueous solutions, are non-conducting. Ammonium chloride also yields a non-conducting solution, but alkylamine hydrochlorides, however, yield solutions which do conduct electricity, as also do certain organic ammonium salts, certain alkaloids and a number of oxygen compounds. The behaviour of such solutions is not analogous to that of aqueous solutions of ordinary electrolytes, since the conductivity increases enormously when the concentration is increased, and is, therefore, probably due to the formation, between solute and solvent, of some additive compound of conducting power. The halides of phosphorus, arsenic and antimony dissolve in liquid hydrogen sulphide, and their solutions are electrically conducting, the conductivity considerably increasing with the atomic weight of the element.

Iodine dissolves in liquid hydrogen sulphide without appreciable chemical change, the product being a deep red solution; bromine, however, is attacked vigorously, with formation of sulphur bromide. Hydrogen chloride also dissolves to a considerable extent without reaction, as also do some of the halides, although there is a tendency (e.g. with the chlorides of Hg, Hg••, Ag and Cu) to "thiohydrolysis," sulphide and hydrosulphide being formed. With the halides, in general, solubility and tendency to reaction increase as the basicity of the positive radical decreases. Liquid hydrogen sulphide reacts vigorously with dry liquid sulphur dioxide, and more slowly with selenium dioxide, forming selenium, sulphur and water. Oxidising agents cause separation of sulphur. Commercial iron and copper, even after prolonged immersion, are not affected by liquid hydrogen sulphide, which may therefore be stored in containers made of these metals.

Chemical Properties of Hydrogen Sulphide

As has already been indicated, hydrogen sulphide exhibits dissociation when heated above a certain temperature. The effect becomes appreciable near 400° C., and with rise in temperature the equilibrium mixture steadily contains less hydrogen sulphide until when near 1350° C. 50 per cent., and at 1700° C. approximately 75 per cent., is in the form of the free elements. A silent electric discharge through the gas hastens the decomposition, as also does the radiation from radium or radium emanation. Heating the gas by a series of electric sparks naturally induces a similar dissociation, and, as the products diffuse from the path of the sparks into a cooler region, the sulphur gradually deposits in the solid condition, so that the gas undergoes a slow but finally complete decomposition, leaving an equal volume of hydrogen:

H2S H2 + S.

Contact with a heated platinum filament also promotes the decomposition.

The gas will not support the combustion of substances which are commonly termed combustible, but will itself burn readily in air of in oxygen, with a blue flame. In the presence of a relatively plentiful supply of air or oxygen, a condition easily obtained by burning the gas at a very fine jet, the products of combustion are water and sulphur dioxide:

2H2S + 3O2 = 2H2O + 2SO2,

but unless care is taken to ensure an excess of oxygen, the combustion is generally incomplete, the foregoing products being accompanied by sulphur, which is produced according to the equation:

2H2S + O2 = 2H2O + 2S.

A mixture of hydrogen sulphide and oxygen explodes on the application of a flame, no free sulphur being formed if the proportion of oxygen is equal to or in excess of that required by the first of the two foregoing equations. In a closed vessel, an undried mixture of hydrogen sulphide with a termolecular proportion of oxygen will explode on being heated to 250° C. The heat of combustion of hydrogen sulphide to water and gaseous sulphur dioxide is 13.67 Cals. per gram-molecule.

The bluish-violet layer observed against the glass surface of a flask when the latter is depressed on the upper part of a Bunsen flame appears to be due to the presence of hydrogen sulphide in the coal gas.

As might be expected, slow flameless combustion of hydrogen sulphide is possible, and the process can be accelerated by contact with certain heated solids, e.g. charcoal, iron oxide, pumice or finely divided! platinum or palladium. It is interesting that a catalytic effect of this type was applied to the recovery of sulphur from the " alkali waste of the Leblanc soda process, so that an otherwise troublesome and unpleasant waste product was not only deprived of its nuisance-creating characteristics, but concurrently made to yield sulphur of a good quality. In the Chance-Claus process the hydrogen sulphide mixed with air was passed through a heated porous mass of iron oxide (bog iron ore) supported on fragments of firebrick. A process applicable to the removal of hydrogen sulphide from crude coal gas employs activated charcoal in a similar manner; in this case the gases can be passed through the material at a high velocity and emerge from the filters completely free from hydrogen sulphide, whilst pure sulphur results from the oxidation.

Dry hydrogen sulphide is stable in the air at ordinary temperatures, but when moist it undergoes oxidation, the rate being especially appreciable if the temperature is raised a little, e.g. to 70°-80° C. Aqueous solutions of the gas undergo a similar atmospheric oxidation, causing the well-known deposit of sulphur in " hydrogen sulphide water," and at the same time acquire traces of sulphuric acid. Exposure to light facilitates oxidation of both gas and solution. In solution oxidation is accelerated by the presence of finely divided nickel.

For the preservation of hydrogen sulphide solutions a covering layer of paraffin oil or the addition of such substances as sugar, glycerol or salicylic acid has been suggested; in the case of the latter substances it is possible that their 44 negative catalytic effect " may be due to their rendering inactive traces of some otherwise powerful catalyst.

On account of its tendency to oxidation, hydrogen sulphide is frequently employed as a reducing agent in organic chemistry. In the case of solutions in N-hydrochloric acid the oxidation is catalytically accelerated by a mixture of manganese and iron.

With oxidising agents other than free oxygen, hydrogen sulphide yields sulphur or sulphuric acid, according to the conditions. Concentrated sulphuric acid oxidises the gas to free sulphur, being itself reduced to sulphur dioxide; sulphur dioxide also effects oxidation to sulphur; hydrogen peroxide solution causes gradual oxidation to sulphur, but in the presence of alkalis oxidation proceeds further, to the formation of sulphate.

If the gas is passed over sodium peroxide, a violent reaction occurs both in the presence and the absence of air. If the peroxide is previously heated, the reaction is accompanied by flame, and if excess of air is present a loud detonation is produced. In the absence of air the products consist of sulphide and polysulphide of sodium, together with a small amount of thiosulphate and sulphate; in the presence of air very little sulphide is formed, sodium sulphate and free sulphur being obtained.

Ozone causes partial conversion of aqueous hydrogen sulphide to sulphuric acid, sulphur being simultaneously produced. With nitric acid the oxidation is very vigorous and with the fuming acid may be explosive, the product being sulphuric acid; hydrogen sulphide will burn in nitric acid vapour. The fact that carbon dioxide appears to liberate sulphur from hydrogen sulphide at a red heat may not be due to direct oxidation but to previous thermal dissociation into sulphur and hydrogen, the latter subsequently being converted into water, whilst the carbon dioxide is reduced to monoxide. Hydrogen sulphide reacts with nitric oxide to form sulphur, nitrogen and water, especially in the presence of catalysts such as silica gel (alone or impregnated with ferric oxide) or glass wool:

2H2S + 2NO = 2S + N2 + 2H2O;

no trace of nitrous oxide or ammonium sulphide is formed. The reaction, however, appears to be more complex than is indicated in the equation, and it has been suggested that the nitric oxide first forms condensed molecules, N2O2, which are adsorbed by the catalyst and bring about the oxidation. With nitrous acid solution, hydrogen sulphide produces nitric and nitrous oxides when the nitrous acid is in excess, but ammonia and hydroxylamine when the hydrogen sulphide is in excess; free sulphur and even a little sulphuric acid are produced.

Aqueous solutions of alkali chromates yield with hydrogen sulphide a precipitate of chromium hydroxide contaminated with sulphur, whilst alkali sulphide, polysulphide, thiosulphate and colloidal sulphur remain in the solution.

Chlorine and bromine liberate sulphur from gaseous hydrogen sulphide, the reaction being capable of going further in aqueous solution, because under these conditions the sulphur may be oxidised to sulphuric acid. With iodine, appreciable reaction occurs only in aqueous solution, and even then the chemical change may not be complete, ceasing when the hydriodic acid attains a concentration of approximately 25 per cent, in the solution. Fluorine attacks gaseous hydrogen sulphide so violently as to cause spontaneous inflammation.

H2S + Cl2 = 2HCl + S,
S + 3Cl2 + 4H2O = H2SO4 + 6HCl,
H2S + I2 ⇔ 2HI + S.

Under considerably increased pressure, or when strongly cooled, hydrogen sulphide and water can combine to form a crystalline compound of which the composition is probably H2S.6H2O; if the temperature and pressure are allowed to revert to the normal conditions, the crystals at once dissociate into the constituent substances. An additive compound of methyl ether and hydrogen sulphide, (CH3)2O.H2S, melting at -148.5° C., is also capable of existence at low temperatures; although the nature of this compound may be allied to that of the additive compound with water, it appears more probable that the methyl ether compound is an oxonium salt.

Hydrogen sulphide appears to be able to play, in a feeble manner, a role analogous to that of water in compounds containing so-called

"water of crystallisation"; thus, aluminium bromide in a fused condition or dissolved in carbon disulphide absorbs gaseous hydrogen sulphide with formation of a colourless crystalline compound, AlBr3. H2S, which melts at 83° C. and is decomposed by water with liberation of hydrogen sulphide. Boron trichloride reacts with liquid hydrogen sulphide, forming white crystals of composition BCl3.12H2S.

Hydrogen Sulphide as an Acid

In the anhydrous condition, whether as gas or liquid, hydrogen sulphide has no acidic properties. When moist or in aqueous solution, however, it behaves as a feeble acid, whence the occasional name " hydrosulphuric acid." For the characteristic test with lead acetate paper or for the reddening of litmus paper, the presence of at least a little moisture is essential.

With ammonia, the gas combines to form ammonium sulphide or ammonium hydrosulphide, according to the relative quantities of the reagents. The alkaloids, which may be regarded as organic derivatives of ammonia, also combine with hydrogen sulphide, forming crystalline salts; such salts of cinchonine, quinine, strychnine, brucine and nicotine have been known for many years. When heated in the gas, the alkali metals yield the acid sulphides:

2K + 2H2S = 2KSH + H2,

which can also be obtained by the action of excess of the gas on solutions of the hydroxides, or of the metals in absolute alcohol.

Tin, when heated in the gas, undergoes conversion into stannous sulphide without the gaseous volume being permanently altered thereby:

Sn + H2S = SnS + H2.

Many other metals yield sulphides if treated with gaseous hydrogen sulphide under suitable conditions; thus mercury, silver and copper fail to react with dry hydrogen sulphide, but if the gas be moist, and especially if oxygen or air also be present, the metals react readily, with formation of the corresponding sulphide, whilst in the presence of oxygen the hydrogen is oxidised to water; the reaction for copper is represented by the equation:

4Cu + 2H2S + O2 = 2Cu2S + 2H2O.

Boron and silicon likewise decompose hydrogen sulphide at high temperatures, liberating hydrogen and forming sulphides.

The oxides of many metals, if heated in a stream of the gas, become converted into the corresponding sulphides, sometimes with simultaneous formation of sulphur dioxide if the reaction becomes vigorous. With peroxides and dioxides the heat of the reaction may be so great as to cause the mass to become incandescent.

Metallic hydroxides, on account of their more decidedly alkaline character, generally react more readily than the corresponding oxides, and a mixture of soda-lime, if submitted to the action of a current of hydrogen sulphide mixed with air, becomes white hot.

On account of the feeble acidity of hydrogen sulphide, the alkali carbonates in aqueous solution are decomposed only as far as the hydrogen carbonates, an equilibrium being attained:

Na2CO3 + H2S = NaSH + NaHCO3,
NaHCO3 + H2SNaSH + H2CO3.

In agreement with the relative reactivities of the alkali hydroxides and carbonates towards hydrogen sulphide, it has been observed that the blackening of basic lead carbonate by the gas is due to the conversion into sulphide of the lead hydroxide only and not of the carbonate. An aqueous solution of hydrogen sulphide is only feebly acidic, the dissolved substance being only partly ionised; the ions present consist almost entirely of H and SH', from the dissociation:

H2SH + SH',

and very few S' ions are present as a result of further dissociation of the hydrosulphide ion.

Physiological Action of Hydrogen Sulphide

If breathed into the lungs hydrogen sulphide has an exceedingly poisonous action, air containing as little as 0.1 per cent, being capable in time of producing a fatal effect. Its action is believed to be due to the formation of sodium sulphide in the blood, which then affects the nerve centres. In considerably more dilute condition than 0.1 per cent, it will produce sickness and headache. Poisoning by absorption of the gas through the skin or mucous membrane is apparently not possible. On account of the harmful action of the gas, many devices have been suggested for the precipitation of sulphides in analytical processes without liberating the hydrogen sulphide in the gaseous condition, for example, the use of an alkali sulphide.
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