August 31, 1911] 



NATURE 



297 



as a typically active solvent. The ordinary organic solvents 

 exhibit intermediate degrees of activity. 



l'"or the purpose of illustrating the effect of solvents on a 

 dissolved substance one may conveniently take a coloured 

 substance in a series of colourless solvents. If the sub- 

 Stance is unaffected by the solvent, we might reasonably 

 expect the colour of the solution to be the same as the 

 colour of the vapour of the substance at equal concentra- 

 tion. Iodine, for instance, gives rise to the familiar violet 

 vapour. Its solution in carbon disulphide has a colour 

 practically similar, but its solution in alcohol or water is 

 of a brown tint quite different from the other. In the in- 

 different hydrocarbons and in chloroform the colour is like 

 that in carbon disulphide, in methyl or ethyl alcohol it is 

 brown. We conclude therefore roughly that iodine dis- 

 solved in saturated hydrocarbons, in chloroform, carbon 

 tetrachloride and carbon disulphide is little affected by the 

 solvent, whereas in water and the alcohols it is greatly 

 affected, probably by way of combination, since in all the 

 solvents two atoms of iodine seem to be associated in the 

 molecule. That combination between the iodine and the 

 active solvents has really occurred receives confirmation from 

 the behaviour of iodine in dilute solution in glacial acetic 

 acid. If the colour of this solution is observed in the cold 

 it is seen to be brown, resembling in colour the aqueous 

 solution. If the solution be now heated to the boiling- 

 point, the colour changes to pink, which may be taken to 

 indicate that the compound of iodine and acetic acid which 

 is stable at the ordinary temperature becomes to a large 

 extent dissociated at ioo°. 



Now, as I have said, a general theory of solution must be 

 applicable to all classes of solution, and herein lies the 

 importance of van 't Hoff's osmotic pressure theory. It 

 applies equally to mixtures of gases, to mixtures of inert 

 liquids, and to mixtures such as those of sulphuric acid 

 and water ; and it has the further advantage that so long as 

 the solutions considered are dilute there are simple relations 

 connecting the osmotic pressure with other easily measurable 

 properties of the solutions. It has been unfortunately the 

 custom to oppose the osmotic pressure theory of solution to 

 the hydrate, or more generally the solvate, theory, in which 

 combination between solute and solvent is assumed. The 

 solvate theory is, in the first place, not a general theory, 

 and in the second place it is perfectly compatible with the 

 osmotic pressure theory. It is in fact with regard to a 

 general theory of solutions on the same plane as the elec- 

 trolytic dissociation theory of Arrhenius. This theory of 

 ionisation applies to a certain class of solutions, those, 

 namely, which conduct electricity, and is a welcome and 

 necessary adjunct in accounting for the numerical values 

 of the osmotic pressure found in such solutions. Similarly 

 the hydrate, or more generally the solvate, theory is 

 applicable only to those solutions in which combination 

 between solvent and solute occurs, and will no doubt in 

 time afford valuable information with regard to the osmotic 

 pressure, especially of concentrated solutions in which the 

 afrtnit\ between solvent and solute is most evident. It can 

 tell us nothing about solutions in which one, or both, com- 

 ponents is inactive, just as the electrolytic dissociation theory 

 can tell us nothing about solutions which do not conduct 

 electricity. 



The great practical advantage bequeathed to chemists by 

 the genius of van 't Hoff is the assimilation of substances 

 in dilute solution to substances in the gaseous state. Here 

 all substances obey the same physical laws, and a secure 

 basis is offered for calculation connecting measurable 

 physical magnitudes, irrespective of the chemical nature of 

 tin- substances and of the solvents in which they are dis- 

 solved, provided only that the solutions are non-electrolytes. 

 If the solutions are electrolytes, the dissociation theory of 

 Arrhenius, developed independently of the osmotic pressure 

 theory of van 't Hoff, gives the necessary complement, and 

 for aqueous solutions offers a simple basis for calculation. 

 Van 't Hoff has given to science the numerically definable 

 conception of osmotic pressure ; Arrhenius has contributed 

 the numerically definable conception of coefficient of activity 

 of electrolytes in aqueous solution, or what is now called 

 the degree of ionisation. 



Of late there has been a tendency in some thermo- 

 (lynamical quarters to belittle the importance of the con- 

 ception of osmotic pressure. It is quite true that from the 



NO. 2183, VOL. 87] 



mathematical thermodynamical point of view it may be 

 relegated to a second place, and even dispensed with alto- 

 gether, for it is thermodynamically related to other magni- 

 tudes which can be substituted for it. But it may be 

 questioned if without the conception the cultivators of the 

 thermodynamic method would ever have arrived at the 

 results obtained by van 't Hoff through osmotic pressure. 

 Van 't Hoff was only an amateur of thermodynamics, but 

 the results achieved by him in that field are of lasting 

 importance, and his work and the conception of osmotic 

 pressure have given a great stimulus to the cultivation of 

 thermodynamics to chemistry. 



And here we trench on a question on which a certain 

 confusion of thought often exists. To the investigator it 

 is open to choose that one of several equivalent methods or 

 conceptions which best suits his personal idiosyncrasy. To 

 the teacher such a choice is not open. He must choose the 

 method or conception which is most clearly intelligible to 

 students, and is at the same time least likely to lead to 

 misconception. Osmotic pressure is a conception which the 

 chemical student of mediocre mathematical attainments can 

 grasp, and it is not difficult to teach the general elementary 

 theory of dilute solutions by means of it and of reversible 

 cycles without liability to radical error or misconception. 

 I should' be sorry on the other hand to try to teach the 

 theory of solutions to ordinary chemical students by means 

 of any thermodynamic function. The two methods are 

 thermodynamicaliy equivalent, and the second is mathe- 

 matically more elegant and in a way simpler, but it affords 

 less opportunity than the first for the student to subnjit his 

 methods to any practical check or test, and in nine cases 

 out of ten would lead to error and confusion. The difficulty 

 of the student is not the mathematical one ; with the excel- 

 lent teaching of mathematics now afforded to students of 

 physics and chemistry the mathematical difficulty has 

 practically disappeared — the difficulty lies in critically 

 scrutinising the conditions under which each equation used 

 is applicable. 



Of the mechanism of osmotic pressure we still know 

 nothing, but with the practical measurement of osmotic 

 pressure great advances have been made in recent years. 

 In particular the admirable work of Morse and Frazer is of 

 the first importance in establishing for solutions up to 

 normal concentration the relationship between osmotic 

 pressure and composition, and its variation with the tem- 

 perature. Much may be anticipated from the continuation 

 of these accurate and valuable researches, the experimental 

 difficulties of which are enormous. 



We are indebted to America not only for these researches, 

 and for the voluminous material of H. C. Jones and his 

 collaborators dealing with hydrates in solution, but also to 

 A. A. Noyes and his school for accurate experimental work 

 and for systematic treatment of solutions on the theoretical 

 side. They, and also van Laar, have shown how solutions 

 not coming within the ordinary range of dilute solutions 

 to which van 't Hoff's simple law is applicable, may in some 

 cases at least be made amenable to mathematical treatment. 

 Van 't Hoff chose one simplification of the general theory 

 by considering only very dilute solutions, for which very 

 simple laws hold good, 'just as they do for dilute gases. 

 Even a single gas in the concentrated or compressed form 

 diverges widely from the simple gas laws ; much more then 

 may concentrated solutions diverge from the simple osmotic 

 pressure law. The other simplification is to consider solu- 

 tions of which the components are miscible in all propor- 

 tions and are without action on each other ; and this method 

 has been developed with marked success from the point of 

 view of osmotic pressure and other colligative properties. 



The outstanding practical problem in the domain of 

 electrolytic solutions is to show why the strong electrolytes 

 are not subservient to the same laws as govern weak 

 electrolytes. If we apply the general mass-action law of 

 chemistry to the electrically active and inactive parts of a 

 dissolved substance (the ions and un-ionised moIecules)_ as 

 deduced from the conductivities by the rule of Arrhenius, 

 we find that for a binary substance a certain formula con- 

 necting concentration and ionisation should be followed, 

 a' formula which we know by the name of Ostwald's 

 dilution law. This law seems to be strictly applicable to 

 solutions of feeble electrolytes, but to solutions of strong 

 electrolytes it is altogether without application. Wherein 



