ON SOLUBILITY. 453 



C (ii)) is applicable also to magnesium carbonate, but only in a limited 

 number of rases to the action of dibasic acids on the' solubility of 

 their neutral salts. The third 1 ™ treated in an extended form of the 

 question of the equivalent precipitation of various chlorides by hydrogen 

 cldoride and is a resume of previous publications in the ' Compt. rend.' 

 1889 A \ furtlier P a P er I49 b y th e same author contained an interesting 

 brief historical review of the subject under investigation, and the 

 collected results of an extensive study of the influence of hydrogen 

 chloride on the solubility of many metallic chlorides — namely, those of 

 tin, copper, cobalt, potassium, lead, mercury, and platinum. 



Graphs representing solubility of potassium chloride in water at various 

 temperatures, alone and in presence of sodium chloride, were found by 

 Etard l51 ^ to converge towards a temperature of 913°, and the 



composition of the mixture at the ' limit of solubility temperature '* 



738° C— was calculated 152 to be 16'7 per. cent. NaCl and 83"3 per cent, 

 KC1. These weights contain approximately equal quantities of metal 

 (K and Na) and non-metal (CI). 



]889 Nernst 144 contributed an important theoretical discussion of 



the results obtained by Engel and the general nature of the influ- 

 ence exerted by one soluble substance upon another in solution, the 

 explanation offered being strictly comparable with the gas laws and the 

 observations of Horstmannf on dissociated vapours. The conclusions 

 he arrived at were supported qualitatively by the precipitation of 

 potassium chlorate from its saturated solution by potassium chloride 

 and by_ sodium chlorate, and approximately quantitatively by results of 

 determining the solubility of silver acetate as influenced by sodium 

 acetate and by silver nitrate. He was the first to advance the theory 

 that at a given temperature the solubility of a sparingly soluble electro- 

 lyte in water or in aqueous solutions of other electrolytes, is dependent 

 upon a constant which he called the ' solubility product.' 



The study of mutual salt influence carried out by Eiidorff was 

 extended by Meyerhoffer 157 to conditions where double salt formation 

 takes place. This author examined the change of solubility at the 

 transformation temperature of double salt, and came to the conclusion 

 that each individual substance has a definite solubility temperature- 

 curve, and at the point of the intersection of such curves the composition 

 of each saturated solution must be identical for each curve. 

 1890 Noyes 164 examined the application of Nernst 's formula for 



calculating the solubility of mixed salts, and in some cases with 

 dilute solutions; the results he found were in approximate agreement 

 with the values obtained experimentally. So strong was this author's 

 faith in the above-mentioned hypothesis that he concluded the discre- 

 pancy between his calculated values and his experimentally determined 

 values for the solubility was due to the degree of dissociation hot being 

 accurately determined by conductivity measurements. He proceeded 

 to calculate the dissociation constant (K) for thallium nitrate from his 

 solubility results, and found a smaller and a fairly constant value. This 



* The temperature at which the proportion of water has beoome reduced to zero, 

 owing to the increased quantity of salt, 

 t Ber., 1881, 14, 1242. 



