H. Gruener — Iodometric Determination of Nitrates. 47 



The process seems reliable enough, then, for estimating 

 nitrates in small quantities not exceeding an equivalent of 0*04 

 grin, or 0'05 grm. of potassium nitrate. Every experiment per- 

 formed with quantities not larger than 005 grm. has been given 

 here, except four in which an imperfection in the apparatus 

 showed an obvious mechanical loss. The errors in these ex- 

 periments lie between the extremes of 0016 grm. + and 0008 

 grm. — on the nitrate, with an average error of 0*00016 grm. If 

 series 3 alone is taken, in which the dilution of the phosphoric 

 acid was regulated, we have as extreme errors 0*0012 grm. + and 

 0*0006 grm. — with a mean error of 0*0002 grm. +. With quan- 

 tities of nitrate above 0'05 grm. the process does not seem a safe 

 one, inasmuch as with a moderate amount of water present 

 some nitric acid distils over undecom posed, and with little 

 water present other complications as seen above arise. 



The method so far as it is applicable may be summed up as 

 follows. The nitrate not to exceed in amount 0*05 grm. of potas- 

 sium nitrate, is introduced into a retort, together with ten times 

 its weight of potassium iodide, and 17 to 20 cm 3 of phosphoric 

 acid, of specific gravity 1 43. All water used should be recently 

 boiled. Carbon dioxide is passed from a receiver carefully set 

 up. The neck of the retort passes into a receiver containing 

 a known amount of decinormal arsenious oxide, alkaline with a 

 good excess of hydrogen sodium carbonate and diluted to a 

 convenient bulk. To this flask is attached for additional safety 

 a simple trap containing water. The solution in the retort is 

 boiled until it is clear that no more iodine remains, when the 

 receiver, after proper washing and addition of the liquid in the 

 trap, is titrated with iodine to find the amount of arsenious 

 oxide still left. This gives the measure of the iodine evolved 

 and consequently of the nitrate present, according to the 

 equation 



2HN0 3 + 6HI = 4H 2 + 2NO + 3I-I 



The Decomposition of Nitrates by Antimonious Chloride. 

 — The failure mentioned above, in attempting to use arsenious 

 oxide to register the action of nitric acid, led to the trial of 

 antimonious chloride as a substitute, inasmuch as this substance 

 is easily oxidizable and less volatile than arsenious chloride in 

 presence of hydrochloric acid. This latter fact is of great 

 importance in the decomposition of a nitrate. The point to be 

 tested was whether the complete decomposition of nitrates by 

 the action of antimonious chloride in hydrochloric acid solution, 

 and the absorption of the nascent oxygen to form antimonic 

 chloride would be secured, so that the antimonious chloride left 

 at the end as compared with the amount taken should give the 

 measure of the nitrate used, according to the equation : 



