Chemical Properties of Sulphur

Chemical Properties of Sulphur are related with electronegativity: sulphur is less electronegative than oxygen and the halogens. It combines directly with most of the elements, notable exceptions being nitrogen, gold, platinum and beryllium.

When heated in oxygen or in air it inflames near 260° C. with formation of sulphur dioxide and a little sulphur trioxide. As would be expected, the atomic heat of combustion of octahedral sulphur (71,080 calories per gram-atom) is less than that of prismatic sulphur (71,720 calories) if considered at the ordinary temperature, the difference being equal to the heat of transformation of the one form into the other. According to Mondain-Monval, however, the transition S_α → S_β absorbs only 87 calories per gram-atom.

At 200° C. sulphur can undergo slow oxidation, manifested by a distinct phosphorescence; oxidation can also occur even at the ordinary temperature, especially with finely divided sulphur in a moist condition. " Flowers of sulphur," when stirred with water, usually imparts a feeble acid reaction to the liquid. To this slow oxidation probably is to be attributed any beneficial effects resulting from the customary introduction of lumps of sulphur into dogs' drinking troughs.

Sulphur is soluble in solutions of the sulphides of the alkali metals, including ammonium, with the formation of yellow solutions of poly-sulphides. The alkali carbonates and the hydroxides of the alkali and alkaline earth metals, in aqueous solution, also dissolve sulphur, producing sulphides or polysulphides together with thiosulphates and sulphites. In all probability the ideal equation for hydroxides is:

6MOH + (n + 2 )S = 2M₂S_n + M₂S₂O₃ + 3H₂O.

The main polysulphide products are trisulphide, tetrasulphide and pentasulphide, but on account of the tendency of the polysulphides to decompose, yielding thiosulphate and hydrogen sulphide, the quantity of thiosulphate usually exceeds that indicated in the equation. Loss of sulphur by the thiosulphate yields sulphite. The so-called "lime-sulphur" washes, used as insecticides in agricultural work, are obtained by treating sulphur with milk of lime in this way.

Many metals combine readily with sulphur; for example copper in the form of foil or wire and the vapour of boiling sulphur react so vigorously, forming cuprous sulphide, that the metal becomes white-hot; sodium burns brilliantly on molten sulphur. The reaction between aluminium and sulphur, started by means of magnesium and barium peroxide, takes place with explosive violence; also, magnesium and sulphur combine explosively if ignited by means of a mixture of sulphur and potassium chlorate. Many investigations have been made on the fusion curves of mixtures of metals with varying amounts of sulphur, and the formation of compounds has been revealed in many cases; antimony and sulphur under these conditions yield only Sb₂S₃, whilst lead and sulphur form only PbS.

Most oxidising agents affect sulphur, the vigour of the action varying with the conditions and with the nature of the agent. A mixture of sulphur and potassium chlorate is highly explosive and will detonate violently on slight shock. Nitric acid oxidises sulphur quantitatively to sulphuric acid. When heated with sulphur, most metallic oxides are converted into sulphide and sulphate; mercury, lead, bismuth and cadmium oxides are common examples:

e.g. 4PbO + 4S = 3PbS + PbSO₄.

In some cases, such as with the oxides of silver and copper, the sulphate can undergo further reduction by sulphur with formation of sulphide and sulphur dioxide; the oxides of zinc, tin and iron are not greatly attacked by sulphur, whilst chromium trioxide reacts so violently as to cause inflammation of the sulphur.

Sulphates of the alkali and alkaline earth metals, when heated with sulphur, are converted into sulphide, polysulphide and thiosulphate, with simultaneous formation of sulphur dioxide; many other sulphates, e.g. those of copper, mercury, silver and lead, yield only sulphide.7 Other salts of the metals behave in a similar manner, undergoing transformation into sulphides, the change being effected more readily with the salts of the heavy metals, many of which indeed react slowly with sulphur even at 100° C. in the presence of water. At 150° to 200° C. mercuric, stannic and ferric salts in aqueous solution are quantitatively reduced by sulphur; mercurous, cupric, bismuth and lead salts are slowly but quantitatively precipitated as sulphides. Nitrates, permanganates and iodates cause oxidation of the sulphur to sulphuric acid.

With sufficiently powerful reducing agents, especially hydriodic acid, sulphur is reducible to hydrogen sulphide, and at higher temperatures a similar reduction can be effected by organic matter generally.

Liquid ammonia dissolves sulphur with formation of an additive compound, S(NH₃)_x, decomposition slowly occurring with subsequent formation of hydrogen sulphide and nitrogen sulphide. Aqueous solutions of ammonia resemble solutions of the ordinary alkalis in their action on sulphur, but are less rapid in their effect.

As has already been mentioned when discussing the allotropy of the element, especially in the gaseous condition, sulphur can exist as molecular aggregates of variable magnitude, so that the molecular weight is not constant and in the gaseous condition ranges actually from S₈ to S. In the molten condition it is probable that at least two types of molecules exist, namely S₈ and S₆, but in solution at the ordinary temperature sulphur appears to be present almost entirely as S₈ molecules, even with such a variety of solvents as bromoform, iodine, anthraquinone, white phosphorus, antimony trichloride, carbon disulphide, carbon tetrachloride, naphthalene, toluene, dimethylaniline and rubber. The various investigations carried out with solutions of sulphur in these liquids include cryoscopic and ebullioscopic measurements, and measurements of vapour pressure, and from the results obtained the evidence for the existence of octa-atomic sulphur in the dissolved condition is overwhelming.