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PHARMACEUTICAL MEETING. 
or q. s.; water, 9 fl. oz. Dissolve the hypophosphite in 6 oz. of the water, and 
the acid in the remainder with the aid of heat; mix the solutions, pour the 
mixture on a white paper filter, and when the liquid has passed, add distilled 
water carefully, till it measures 10 fl. oz., and evaporate this to 8| fl. oz. The 
solution thus prepared is stated to contain about 10 per cent, of real acid 
(HPHoOo). 
Having frequent occasion to prepare this syrup of hypophosphite of iron, I at 
first resorted to the process given by Mr. Parrish, and I made a quantity of the 
hypophosphorous acid by his method as just described. But upon carefully 
testing the acid I obtained, I found it was not so pure as theory would indicate 
it should be ; in fact, it contained an appreciable quantity both of oxalic acid 
and of lime. The free hypophosphorous acid had dissolved a portion of the 
oxalate of lime. This impurity seemed to me to be of some importance, and, as 
it could not be got rid of, it constituted an objection to the process. It was 
easy to produce a purer acid by resorting to hypophosphite of baryta, and exactly 
decomposing its solution with sulphuric acid. But as the syrup of hypophos¬ 
phite of iron is only occasionally demanded, and as, from its proneness to oxida¬ 
tion, it deteriorates greatly by long keeping, it appeared desirable to possess a 
more expeditious method for its preparation than that which involves the elimi¬ 
nation of the acid, and the precipitation of the iron as carbonate. I consequently 
abandoned the use of the acid, and resorted to a process of double decomposi¬ 
tion. 
When hypophosphite of lime and sulphate of iron are brought into contact, 
sulphate of lime is precipitated, and ferrous hypophosphite remaius in solution. 
Ca(P H 2 0 2 ) 2 + Fe S 0 4 =z Fe(P H, 0 2 ) 2 + Ca S 0 4 . 
Using these materials in their atomic proportions, the iron salt is obtained 
contaminated only by a small quantity of sulphate of lime, and the amount of 
the latter may be reduced to a minimum by employing very little water as a 
solvent. 
It became necessary, however, for the trustworthiness of this method, to 
ascertain whether commercial hypophosphite of lime is sufficiently pure and defi¬ 
nite to effect always an exact decomposition of the sulphate of iron. I there¬ 
fore analysed several samples of the salt, and found that the amount of real 
hypophosphite (Ca (PH 2 0 2 ) 2 ) varied from 92 to 94 per cent., the remainder 
consisting chiefly of water, with a little carbonate and phosphate of lime. 
This variation in composition I regard as too slight to be of any material im¬ 
portance. In calculating the quantity of the salts necessary for the reaction, the 
lime salt may be regarded as containing 90 per cent, of real calcic hypophos¬ 
phite, a slight excess of the latter being preferable to any surplus of the sulphate 
of iron. The proportions to be then employed for 320 grs. of ferrous hypophos¬ 
phite, Fe(PH,0 3 ) 2 , are 480 grs. of crystallized sulphate of iron and 326 grs. 
of commercial hypophosphite of lime. If the two salts are triturated with 2| 
oz. of water, the resulting paste pressed out, and the filtered liquid mixed with 
seven times its volume of simple syrup, the product contains 2 grs. of hypophos¬ 
phite of iron in each fluid drachm. 
But the syrup so produced does not keep long unless atmospheric air is 
thoroughly excluded. After a few hours’ exposure, a precipitate begins to form 
at the surface, and gradually passes downwards. To prevent this, it is neces¬ 
sary to introduce a free acid, and phosphoric or citric acid is best suited for the 
purpose. The former is preferable, because it is more in chemical accordance 
with the other constituents of the syrup. I have tried the use of free hypophos¬ 
phorous acid, but it does not answer well, being a very bad solvent for the inso¬ 
luble ferric hypophosphite which forms after a time in the syrup. The process 
I follow, therefore, stands as follows :— 
