164. Scientific Intelligence. 



with solutions as weak as 0*07 per cent H 2 SO^, or 0'0125 H 2 S0 4 

 to 100 H„0. The molecular depression shown, even in this ex- 

 treme region, instead of being constant as it should be according 

 to the theory of osmotic pressure, varies between 2*95° and 2'1°. 

 The molecular depression in the various curves ranges from 0'01° 

 to 2*95° according to the nature of the solvent and that of the 

 dissolved substance, these numbers referring only to solutions 

 containing not more than one foreign molecule to 100 solvent 

 molecules. In no two cases out of the seven investigated does it 

 possess the same value. In every case an increase in the strength 

 of the solution beyond this proportion entails an abnormally low 

 freezing point. According to all existing physical theories of 

 solution the freezing points of such solutions should be abnor- 

 mally high." — J. Chem. Soc, lvii, 331, May, 1890. g. f. b. 



2. On the Molecular- Mass of Iodine, of Phosphorus, and of 

 Sulphur in Solution. — Beckmakx has applied his method of de- 

 termining the molecular mass of a substance by means of the ele- 

 vation which it produces in the boiling point of a solvent, to the 

 cases of iodine, of sulphur and of phosphorus, dissolved in carbon 

 disulphide, and to that of iodine in ether. The molecular eleva- 

 tion of the boiling point in the case of cai'bon disulphide using 

 100 grams was 23 - 75° and referred to 100 c. c. was 19'43°. In the 

 case of ether the molecular elevation was, for 100 grams 21 "05°, 

 and for 100 c. c. 30*21°. As a result it appeared that the molecu- 

 lar mass of iodine, in solution both in ether and in carbon disul- 

 phide, is 254 very nearly ; corresponding to the formula I„. The 

 molecular mass of phosphorus in carbon disulphide is 124, cor- 

 responding to the formula P 4 ; and that of sulphur in the same 

 solvent is 256, corresponding to S 6 . — Zeitschr. Physik. Chem., v, 76, 

 February, 1890. g. f. b. 



3. On the conditions of Equilibrium betioeen Electrolytes. — 

 Arrhe^ius has compared the experimental conditions of equilib- 

 rium in solutions with those pointed out by theory. In the case 

 of an acid and one of its salts, if x represents the fraction of dis- 

 sociated acid and d the amount of the salt, V being the volume 

 in liters containing one gram molecule of the acid and n mole- 

 cules of the salt, then (nd-\-x)x=~KV(l — x), K being a con- 

 stant of dissociation determinable from the conductivity of the 

 acid. This equation is found to hold in the case of acetic and 

 formic acids and their sodium salts. Since for feeble acids x is 

 small, and may be neglected in comparison with nd or 1, and 

 since d is practically independent of the dilution, it follows that 

 the degree of dissociation, that is, the strength of a feeble acid 

 when a salt is present in the same solution, is proportional to the 

 amount of the salt. Between a feeble acid such as acetic, and a 

 salt such as sodium chloride, the equilibrium is that between four 

 substances, not only the two mentioned, but also the sodium 

 acetate and hydrogen chloride resulting from their reaction. If 

 the fractional dissociation of these substances in the order named 

 be expressed by d 1 d i c? 2 d z , and one molecule of acetic acid on 



