T H E J O U R N A L OF I N D U S T R I A L A N D ENGINEERING C H E M I S T R Y
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Extraction 1 .2 3. 4 5 6 7 8 9.. 10
........................ ........................ ....................... ........................ ........................ ........................ ........................ ........................ ...................... ........................
Alkaloidal Material
Ether Used Gal. 160 140 145 160 150 150 150 150 150 150
Removed 02.
14.22 9.41 8.62 5 7.66 .80 5.80 5.44 2.89 5.33 4.41
METHOD %-To 300 lbs. Lloyd’s reagent (assay 2.36 per cent) 40 lbs. CaO were added and the mixture extracted for about 3 hrs. with t h e following volumes of 80 per cent alcohol: 80 Per cent Alcohol Used
Gal.
Alkaloidal Material Removed 02.
45.44 17.61 13.68 8.92 7.42
An assay of the extracted Lloyd’s reagent gave 0.18 per cent alkaloid, showing t h a t 7.6 per cent of the alkaloidal material still remained in it. T h e following is of interest in connection with the alcohol extraction of the alkaloidal material from Lloyd’s reagent. A mixture of 10 g. of Lloyd’s reagent with 1.3 g. of slaked lime was extracted for 14 hrs. with 50 cc. of alcohol of the following strengths: Strength of Alcohol Per cent 100 90 80 70 60
50
T h e volume of follows:
80 per cent alcohol was varied as
Alcohol Used cc. 20 30 40 ..
50 80 100
Alkaloidal Material Removed Per cent 0.00 25.42 41.25 41.25 33.82 23.50
Alkaloidal Material Removed Per cent 26.8 32.6 35.0 41.25 43.6 47.0
PURIFICATION OF THE CRUDE ATROPINE
The alkaloidal material was extracted from the Lloyd’s reagent with 95 per cent alcohol, using lime t o obtain the proper alkalinity. The extractions were acidulated with acetic acid and the solution concentrated first t o 12 per cent, and then under diminished pressure t o 2 per cent of its original volume. This procedure was sufficient t o convert all the hyoscyamine into atropine. After neutralization with ammonia, the solution was allowed t o stand over night and filtered. A test portion of the filtrate was shaken with ether. If a n emulsion resulted, t h e solution was ‘diluted about one-fourth and returned t o the vacuum still. Distilling the neutral liquid, and again filtering, usually prevented the troublesome emulsion with ether. Ammonia was added until the solution was alkaline and the atropine alkaloid extracted with ether. After evaporation of the ether, t h e alkaloid was carefully dried a t about 3 5 ” C. The dried alkaloid was dissolved in ethyl alcohol in t h e proportion of one ounce of alkaloid t o two fluid ounces of solvent, and the solution almost neutralized
Vol. 13, No. 3
with sulfuric acid, using cochineal as indicator. After filtering i t was evaporated on the water bath t o a thin sirup, and t o this sirup, while still warm, acetone was added almost t o the point of precipitation of t h e atropine sulfate. On cooling, the atropine sulfate crystallized. If not sufficiently pure the crystals were dissolved in alcohol and recrystallized as outlined above. The acetone was evaporated from the mother liquor, and the alcoholic solution of atropine sulfate poured into a large volume of water. From this the alkaloid was extracted with ether, and if not of sufficient purity t h e process already outlined was repeated. AN INVESTIGATION OF THE U. S. P. ASSAY FOR PHOSPHORIC ACID AND SOLUBLE PHOSPHATES‘ By A. E. Steam, H. V. Farr and N. P. Knowlton MALLINCKRODT CHEMICAL WORKS, ST. LOUIS, MISSOURI
I n routine analysis of samples of phosphoric acid in this laboratory, it was noted t h a t , although aliquots from the same solution when assayed according t o the directions given in the U. S. Pharmacopeia2 gave concordant checks, i t was difficult t o obtain check results when two different samples were weighed out and made u p t o volume, unless the size of t h e sample happened t o be nearly the same in both cases. Briefly, the method is t o transform the acid t o t h e disodium salt b y neutralizing with NaOH t o a phenolphthalein end-point, precipitate with a n excess of standard silver nitrate solution, bring the solution t o neutrality t o litmus with ZnO, and determine the excess AgN03. Calculations were made of the error introduced by the actual volume occupied by the precipitate. This error, assuming a specific gravity of I for t h e precipit a t e (the value is given as between 7 and 8 ) , adsorption of water t o the extent of one mole per mole of phosphate, the presence of the equivalent of 50 cc. 0.1N salt, and an equal volume occupied by the excess of ZnO added, was shown t o have a maximum possible value of 0.5 per cent, and more reasonable assumptions reduced this error t o 0 . 0 8 per cent on a 90 per cent sample. This small error by no means explained t h e large discrepancies of 5 t o I O per cent met with a t times. A few preliminary experiments seemed t o indicate t h a t the results were influenced very markedly by the size of t h e sample. T h e larger the am0 sample, the lower were t h e results obtained. I n the filtrates, after the Ag3P04 had been filtered off, a test with ammonium molybdate showed the presence of significant quantities of phosphate which t h e silver, though present in considerable excess, had failed t o carry down. The importance of this particular method may be realized when we recall t h a t i t forms the basis for t h e assay not only of phosphoric acid and t h e alkali phosphates, but also of many hypophosphites, such a s those of Ca, Na, K, Mn, “4, etc. The method is not confined t o t h e U. S. P., but is found in t h e N. F. and even in the “New and Non-official Remedies.” 1 Presented before the Division of Medicinal Products Chemistry, a t the 60th Meeting of the American Chemical Society, Chicago, Ill., S e p tember 6 to 10, 1920. 2 Ninth Revision, p. 21.
Mar., 1921
T H E J O U R N A L OF I N D U S T R I A L A N D ENGINEERING C H E M I S T R Y
It was therefore well worth while t o investigate the cause of these phenomena, confirm them, and determine, if possible, a reliable procedure, or a t least a size of sample which would give fair results in routine work. R O U T I N E ANALYSIS O F STOCK P H O S P H O R I C ACID
The effect of varying the size of sample was first studied in the case of a stock solution of the acid taken from the laboratory shelf. A solution was made u p of approximately I O t o 11 g. per liter, and analyses of different aliquots were made. T h e procedure was as nearly as possible t h a t of rapid routine work. Room temperature was considered sufficiently constant (though on warm days the solution was cooled t o the temperature a t which it was made up).
75
i
I
I
5
I
I
IO
221
could give was added. The solution was made u p a t z s 0 C. and maintained within,^.^' of t h a t temperature; t h e volumes of the standard solutions were in all cases corrected for even slight changes in temperature, and great care was exercised in making the solution neutral t o litmus, after precipitation with AgNO3. I t was not sufficient t o take a drop and touch it t o litmus paper, or even t o float pieces of red and blue litmus on t h e surface. The acidity was regulate& more by means of dilute NaOH t h a n b y ZnO, though this was added. It was manipulated until pieces of red and blue litmus, after vigorous shaking in the solution, kept their respective colors side by side.
-75
Cc. Solution fn Sample
FIG. 1-SHOWING
EFFECT OF SIZE OF SAMPLE OF Hap04 ON PER CENT OF Hap04 OBTAINED. U. S. P. METHOD
Table I gives the results so obtained. These d a t a are also plotted in Fig. I , the number of cc. of acid solution in t h e sample being plotted as abscissae, while the per cents of HsP04 found are plotted as ordinates. The phosphate in t h e filtrate was determined by precipitation as phosphomolybdate and titration of t h e precipitate with standard alkali. This method is sufficiently accurate for t h e small quantities in t h e filtrate, though i t was found unsatisfactory as a check method on the total H3P04. The method actually used as reference was the standard pyrophosphate method, results of which are also given for comparison. TABLG I-ANALYSIS OF
FIG.2-S€iOWING
EFFECT 08 SIZE OF SAMPLE OF NazHPO4 OF P E R CENT OF NazHPOd OBTAINED. U. S. P. METHOD
The results of this series of experiments are given in Table 11, and the d a t a plotted in Fig. 2. I n Column 5 are found data upon which t h e size of the points in Fig. z is based, this size representing the error introduced into t h e position of the point by a n error of one drop (0.0445 cc. in the buret used) in titration. I n a certain sense the size of a point represents its accuracy from a manipulative point of view. This does not mean, however, t h a t the points can be considered significant only t o t h e extent of their respective areas, TABLE 11-ANALYSES OF NAzHP04
c";
Tntal
---I_
.... .. ..
*.
..
*.
*. ..
ANALYSIS O F DISODIUM P H O S P H A T E
I n order t o ascertain whether t h e rapidity and inaccuracies of routine work were responsible for t h e discrepancies shown in Table I, a sample of c. P. sodium phosphate was recrystallized twice, centrifuged, and dried t o constant weight a t IIO', where i t was completely exsiccated. Practically the same procedure was followed as in t h e case of t h e phosphoric acid, except t h a t all t h e accuracy which time and precautions
OF 10.158
Reaction Sample to Cc. Litmus 1 Neutral 2 Neutral 3 Neutral 5 Neutral 7 Neutral 10 Neutral 12.5 Neutral 15 Neutral 15 Verysl. acid 17.5 Neutral 20 Neutral 20 S1. acid 20 Very SI. alk. 22.5 Neutral 22.5 Neutral 22.5 Neutral 25 Neutral 25 Neutral 25 Neutral 25 Verysl. alk. 15 cc. gave 0.1195 15 cc. gave 0.1195
G. MADEUP
TO
53
zu
OF 23.362 G.PHOSPHORIC ACID MADS UP TO 2 LITERS
A SOLUTION
0.1 N HsPOI Total HaPo4 AgNOa Hap04 Recovered HsPO4 Present (PyroSample Consumed Found from Filtrate Found phosphate) cc. Cc. Per cent Per cent Per cent Per cent 1 3.20 89.56 89.56 86.7 3 9.58 89.37 89.37 5 15.75 88.15 88.15 7 21.98 87.87 87.87 *. 30.97 86.66 trace 86.66 .. lo 12 37.20 86.75 0.1 86.85 15 43.90 81.90 3.7 85.6 17.5 49.95 78.29 8.35 86.63 10 CC. gave 0.115 1 g. MgzPzO7 or 86.76 per cent Hap04 0.1150 86.69 per cent Av.. 8 6 . 7 Each value given represents the average of a t least two titrations which in the case of samples u p t o 10 cc. checked t o the drop, and in practically all other cases checked within one drop. T h e same is true of data presented in the other tables.
SOLUTION
ONE LITER hp.l
9"$
2j !i
-a S %
28
p
2.21 103.01 2.10 4.41 102.78 1.05 6.51 101.15 0.70 10.76 100.31 0.42 15.05 100.22 0.30 21.36 99.59 0.21 (4 det.) 26.77 99.82 0.16 31.97 99 34 0 13 31.47 98:78 0113 37.26 98.22 0.10 42.53 99.12 0.10 40.86 95.23 0.10 42.94 100.07 0.10 46.11 95.52 0.10 46 15 95.59 0 10 4.5113 93.49 0:lO 49.85 92.95 0.10 50.00 93.23 0.10 50.80 94.72 0.10 53.50 99.75 0.10 g. MgzPzO7 or 100.06 per g. MgzPzO7 or 100.06 per
....
.. .. .. .. ..
$R H
0
103.01 102.78 101.15 100.31 100.22 99.59
... ...
100.06
... ... ... ... ...
99.82 99.34 98.8+ 98.22 99.12 si'ppt. 95.23+ 100.07 95.52 95.59 1:92 95.27 2.45 95.40 3.50 96.73 2.15 96.87 ... 99.75 cent NazHP04 cent NazHPO4 Av., 100.06 sl:bpt.
...
..
... ... ......
..
...
..
..
... ... ... ... ...
for a n error of two drops is more likely on the portion of the curve where t h e points are small than t h a t of half a drop where they are large. I n Fig. 2 t h e same tendency as in Fig. I will b e
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‘222
noted, though t o a less marked degree. The column representing “Reaction” is given because this condition was found t o be a source of error which outweighed all others. It will be noted (in the case of one of t h e 20-cc. samples) t h a t very slight acidity gives low results, while very slight alkalinity gives seemingly high results, a distinctly alkaline solution precipitating X g 2 0 , and the back titration showing too little AgN03. For this reason phenolphthalein cannot be used as a n indicator. Berthelotl suggests neutrality t o phenol-phthalein for complete precipitation of phosphate with silver; and if the result looked for is merely t h e complete precipitation of the PO4 ion this is all right, but if t h e excess of t h e silver is t o be determined i t will not work. The solubility of AgZO is o . 0 0 0 1 0 8 mole per liter a t 25’. Silver oxide in solution is practically completely hydrated and is a comparatively strong base, so t h a t we have the following equilibrium: A Ag,O HzO zAgOH 28g’ zOH-
+
-
e
+
At z j O we have in solution in one liter 0.0001 mole 9 g z 0 or 0 . 0 0 0 2 g. ions each of Ag+ and OH-, so t h a t its solubility product is of the order of (O.OOOZ)~=
1.6 X
10-l~.
I n ordinary assays we have an excess of 2 0 cc. 0 . I N AgNOs per I O O cc., or 0 . 0 2 equivalent per liter. For a red reaction to phenolphthalein t h e OH- concentration is approximately IO-^. Thus (o.oz)X ~
(O.OOOOI)~= 4
X
10-l~.
This value exceeds by some thirty times t h e solubility product of silver hydroxide, so t h a t a considerable quantity of Ag2O would come down. On t h e other hand, AgsPOd is soluble in acid solutions forming the acid phosphates. DISCUSSION
Referring t o the standard pyrophosphate values, it will be noted t h a t in both series of experiments very small samples gave abnormally high values; the sample recommended in t h e U. S. P. gave fair t o good results, and large samples gave very low results, though i n these cases i t is difficult t o obtain checking duplicates, and different determinations even on the same size of sample vary widely. Several factors affect the first-mentioned case. The inaccuracy of measuring such small volumes seems t o the writer t o play an important part, as does also the inaccuracy of titration where one drop of 0 . I N solution corresponds t o over 2 per cent in t h e result. It must be borne in mind, however, t h a t the titration is a back titration, and t h a t a n overtitration would cause the error t o throw t h e results low instead of high, so t h a t the high results seem difficult of explanation a t the present time. The fact t h a t results go from much too high gradually t o much too low as the sample is increased shows t h a t i t must cross the line of a true result at a certain point or with a certain size of sample. I t is this size, whether accidentally or not, which happens t o be given i n t h e U. S. P. The results obtained at this point 1
Ann. clztm. phys., 171 25 (19021, 160.
Vol. 13, No. 3
are probably due t o the accidental compensation of a number of errors a t t h a t particular concentration. At first thought, from a theoretical point of view, neglecting t h e possible influence of any mechanical occlusion of either N a 2 H P 0 4or AgN03 in the precipitate, the method should work over the entire range studied. The idea of mechanical occlusion, which was formerly overworked in many cases where discrepancies were thought t o have been noted, has given place largely t o the idea of higher order compounds,l or t o t h a t of intermediate compounds in t h e case of polyvalent substances.2 Some light on the present case can be obtained from the studies of two men, Berthelot and Y . Osaka,3 who state: According t o Berthelot, in the action of sodium phosphate on silver nitrate, the precipitate of trisilver phosphate retains a certain amount of the disilver phosphate, and, under certain conditions, of a silver-sodium phosphate. The quantities of these substances vary with the composition of the supernatant liquor with which it is in equilibrium. This equilibrium has been studied by Osaka. Upon precipitating N a 2 H P 0 4with three equivalents of AgN08, Berthelot found by analysis of t h e precipit a t e and of t h e filtrate t h a t equilibrium conditions were expressed very satisfactorily by the equation: 3AgN03 NaQHP04= zhTaN03 0.78AgN03 0.2HN03 o 6 AgaP04 0.4AgHzP04
+
+ +
+
+
These conditions are fairly represented by t h e larger samples in Tables I and 11, where the excess of silver is qiiite small. The solubility of Ag3P04 in a solution of AgN03 is very small itself if t h e solution is not acid. Much acid tends t o form t h e acid salts AgzHPOd and AgHzP04. I n t h e present study we are interested in two equilibria. We have always a n excess of AgNOp present and we remove t h e ” 0 3 , so t h a t our conditions after t h e precipitation of silver phosphate are: AgHZP04 f AgN03 A AgzHP04 HNOa
-
+
+
NaOH (or ZnO)
dr
AgzHP04 f
c.
Ag3P04
+
NaN03 “03
(I)
+ HOH
+
NaOH (or ZnO)
(3)
11
NaNO3 f HOH I n t h e case of Equation I we have a salt highly acid t o litmus, and NaOH is added, shifting t h e equilibrium far t o t h e right and transforming most of t h e salt into t h e nearly alkaline Ag2HP04. I n pure water this is said t o hydrolyze readily into AgsPOs and &Pod. This hydrolysis would be greatly retarded by t h e presence of the neutral salts which are i n solut i o ~ and ~ , ~ of t h e trace of H N 0 3 . I n t h e case of Equation 2 we have a condition with only a very small change in Hf-ion concentration for relatively large additions of alkali. I n other words, the conditions exG. McP Smith, J . Am. Chem. S O L 39 , (1917), 1152. Berthelot, LOG.cit. s Mem. COX Sci., Kyoto I m p . Unio., 1 (1904-51, 188. 4 Compare Treadwell-Hall, “Analytical Chemistry,” 2, 587. 1
e
Mar., 1921
T B E J O U R A T AL 0 F I X D U S T R I A L A iV D E N G I N E E R I N G C H i 3 M I S T R Y
pressed by Equations I and z seem t o represent conditions which not only give a pronounced “buffer” action, but which seem t o be practically neutral t o litmus over a wide range. Thus there would seem t o be a good deal of uncertainty as t o when t h e equilibrium was completed t o the right. On t h e other hand, t h e addition of a just sufficient quantity of NaOH should theoretically shift t h e equilibrium completely t o the right and precipitate all of t h e phosphate, since under no conditions would NaN03 be hydrolyzed t o the same extent as Ag2HP04. With this in mind, calculations were made as t o the theoretical amount of HX03 which would be liberated from a 25-cc. sample of t h e sodium phosphate solution (Table 11). This was found t o be equivalent t o 3 . 5 8 cc. of 0.5 N alkali. Two samples were taken and into one were put 3 . 5 0 cc. 0 . 5 N alkali, while into the other were p u t 3.60 cc. 0 . 5 N alkali, and t h e phosphate was precipitated. The results are given i n Table 111. 0.5 N Alkali Sample Added cc. cc. 25 3.50
25
0.5 N AgNOa Required Cc. 53.27
TABLEI11 NazHP04 Found Per cent 99.88
3.60 53.67 100.07 Compare Table 11, Lines 13 and 20.
NazHPO4 NazHP04 by Pyrophosphate in Filtrate (Molybdate) Per cent S1. but distinct yel... low coloration 100.06 Nocolor
These figures seem a very significant confirmation of t h e above-mentioned buffer action when viewed in the light of three facts: ( I ) Referring to Table 11, it will be noted that when the solution containing a 20-cc. sample was made very slightly on the alkaline side of neutrality to litmus, a result of 100.07 per cent was obtained, and the same treatment of a 25-cc. sample yielded 99.75 per cent, while the ordinary determinations on the same sized samples were running up to some 8 per cent lower. (2) When the first point a t which the solution reacts neutral is taken as the point of complete precipitation, increasingly low results are obtained with increase of sample since the smaller the eacess of AgN03, the greater the concentration of acid salts formed in the original mixture, and the wider the range of the buffer action. ( 3 ) With increasingly large samples it becomes increasingly difficult to determine this “first point of neutrality,” so that, although fair checks are often obtained on duplicates run side by side and treated almost exactly the same, it will be noted in Table I1 that it is increasingly difficult t o get two sets of determinations with like-sized samples to check. In other w-ords, it is harder t o strike the same point on the wider buffer range every time. I
This fact, coupled with t h e striking results in Table 111, suggested t h a t reliable results might be obtained if we should have present in solution just enough sodium hydroxide t o take up all t h e nitric acid as fast as it mas liberated in the precipitation of the Ag3P04. Such a method would seem t o have several advantages over neutralization with zinc oxide. Thus all of the nitric acid which can be liberated is taken care of by a n excess of sodium hydroxide already present, and the acid salts AgH2P04 and Ag2HP04 are prevented from forming by the distinct alkalinity of the solutions during the precipitation. So long as there is any phosphate present no silver will be precipitated as oxide, owing t o the greater insolubility of the phosphate. The solubility of Ag20
223
a t zoo is o.oooog mole per liter, while t h a t of Ag3P04 is O . O O O O I ; mole per liter. Thus no Ag20 will be precipitated until t h e concentration of the P O 4 ion is reduced t o 0.O O O O I ~ mole per liter even in a n alkaline solution, provided there is a n excess of AgN03. I n case there is not an excess it can be easily shown t h a t a t equilibrium there will be only 0.003 times as much PO4--- in solution as there is OH.’ An apparent disadvantage may suggest itself in overtitration of the phosphoric acid and consequent excessive addition of alkali in taking care of the nitric acid t o be liberated. As a matter of fact, the hydrolysis of the NasHP04, which is distinctly alkaline t o phenolphthalein, takes care of this. Indeed, the hydrolysis occurs t o such a n extent t h a t if t h e titration be made on a plain acid solution a t room temperature the excess alkali will be significantly deficient for neutralization of the nitric acid liberated during the precipitation. This hydrolysis is inhibited by having the solution ice cold and by the presence of a neutral salt. Treadwell-Hall recommends sodium chloride. This, of course, cannot be used here, but sodium nitrate can be introduced. The difference in titration of phosphoric acid with and without these precautions is shown by the following figures: 0.1 N NaOH Required cc. Ice cold, with NaCl.. .............................. 20.75 Room temperature, without NaC1.. . . . . . . . . . . . . . . . . .20.45 17 .OO Ice cold, with N a C l . . .............................. Room temperature, without NaCl.. . . . . . . . . . . . . . . . . . 16.70 Ice cold, with NaN031. ........................... 47.20 Room temperature, without NaN03. 46.35 1 The NaN03 was tested and found to react neutral.
................
A small overtitration affects results much less t h a n i t might a t first seem. I n the first place, the error throws the results high and tends t o compensate for the slight error introduced by t h e volume occupied by the precipitate. One drop overtitration of 0 . I N NaOH, even if the alkali were completely precipitated as AgZO, would amount t o 0.045/3o.o, or 0 . 1 5 per cent, since a representative sample consumes 3 0 cc. of 0 . I A7 AgNO8. I n Table IV are given the results of a TABLEIV 0.1 N Alkali Added Amount of Na?HPO4 Found1 Run Cr. Overtitration Per cent 1....................... 10.75 0.0 100.49 2. ...................... 10.85 0.1 100.75 0.2 101.12 3 ....................... 10.95 4 . . . . . . . . . . . . . . . . . . . . . . .11.05 101.43 0.3 0.4 5 ....................... 11.15 101.74 6 ....................... 11.25 102.05 0.5 7 ....................... 11.35 0.6 102.24 8 ....................... 11.45 0.7 102.49 9 ....................... 11.75 1.0 103.48
series of determinations upon the effect of overtitration. The solution of specially purified NazHPOl was used so t h a t there was no question about the endpoint. The assay of somewhat over I O O per cent in the first figure is probably due t o t h e fact t h a t t h e solution, which had assayed 100.06 by the pyrophosphate 1
For example: I n same way Cubing (1) Squaring (2)
AgZ X (0H)z Ag3 X ( P O P ) Age X ( O H p Age X (OH)6
= 1.6 X 10-16
(1)
5 X 10-20 4.1 X 10-45 25 X 10-40
(2)
= = =
(3)
(4)
Dividing (3) by (4) z= 1/6 X 10 5 or POa = 770 (0H)s (PO41 Taking any definite value for PO;, such as 0.0001, and solving for (OH} we obtain the relation: OH = 300 PO4 (5)
T H E JOURNAL OF INDUSTRIAL A N D ENGINEERING CHEMISTRY
224
method and 100.07 by this method, had been standing with a loose cork stopper and had become slightly more concentrated, since this experiment was made some time after the rest of the work had been finished. The theoretical amount of alkali t o be added is I O . 7 5 cc. 0.I N t o a I 5-cc. charge. I n these titrations t h e average number of cc. of 0.1N AgN03 consumed was 33. The error per one cc. overtitration, then, should be 1/33 X 100,or 3 per cent. Glancing a t t h e results of Runs I and 9, we find 103.48 t o 100.49, giving a n error of 3 per cent. There is in practice, however, very little danger of overtitration with alkali, on account of the aforementioned hydrolysis. PROCEDURE SUGGESTED FOR THE A S S A Y O F PHOSPHORIC
I n the light of the foregoing the following procedure is suggested for the assay of phosphoric acid. Weigh out about I O g. of t h e acid and make up t o 1000 cc. Introduce I O cc. into a 100-cc. standard flask. Add 3 t o 5 g. C. P. NaN03, cool in an ice bath, and titrate with NaOH, using phenolphthalein as indicator. Take the number of cc. of standard alkali required, divide it by two, and add this quantity in excess t o the sample. Add 50 cc. 0 .I N AgN03, make u p t o the mark, mix, filter through a dry filter, rejecting the first 2 0 cc. of the filtrate. To 50 cc. of this filtrate add 5 cc. concentrated, H N 0 8 , and titrate with 0.I N sulfocyanate. Table V gives the results of a series of determinations made according t o the above procedure. The method has also been used by the writer as well as by others on a number of samples t o be tested, and its use was attended with apparent satisiaction. On various occasions i t has been checked up very favorably with pyrophosphate determinations on t h e same sample. TABLEV-ANALYSIS
OF A SOLUTION OF 11.447 G. OF PHOSPEORIC MADEUP TO O N E LITER
ACID
HaPOa HaPOi HSPOA 0.1 N Calc. from Calc-. from Pyro-AgNOa NaOH AgNOa phosphate Sample Added Consumed Titration Consumed Method cc. cc. cc. Cc. Per cent Per cent Per cent 1 2.08 3.10 3.20 89.11 91.38 91.3 2 4.25 6.37 6.40 91.04 91.40 5 10.50 15.75 16.00 89.96 91.38 7 14.60 21.90 22.35 89.60 91.20 31.50 31.90 90.00 91.10 10 21.00 31.80 89.15 90.81 31.20 10’ 20.80 38.10 38.40 90.68 91.38 12 25.40 47.85 47 85 91.10 91.10 15 31.90 10 cc. gave 0,1188 g. MgzPz0, or 91.38 per cent Hap00 0.1186 91.23 ver cent A v . , 91.3 A t room temperature.
0.1 N NaOH Titer
Cc. Aliquot Taken FIG.J-UPPER TION
ACID
NaOH
.. .. .. .... .. ..
Column 5 furnishes a n answer t o the question as t o whether i t would not be just as well t o calculate the strength of the acid from the NaOH titration as t o go on through the entire procedure. The results in this column vary over a range of z per cent, while in t h e next column the variation is confined t o only 0.3 per cent. It will be noted t h a t , in the case of one of t h e determinations on a IO-cc. sample, significantly low results were obtained by not titrating the solution ice cold. A word as t o the use of zinc oxide as a neutralizing agent may be in place. The solubility of zinc oxide at room temperature is o.oooo5 mole per liter. The solubility product (Zn) X (OH)2 = o.00005 X
Vol. 13, No. 3
PER CENT HsPOi O B T A I N E D AS A FUNCO F SIZE O F SAMPLE. MODIFIED u. s. P . M E T H O D . LOWERCURVB: R E S U L T S OF NaOH TITRATION OF &PO4 SHOWINQ VARIABLE CURVE: SHOWINQ
( O . O O O I ) ~ = 5 . 2 X 10-13. A representative sample of phosphoric acid (0.I 9.) liberates 0.001 mole of H N 0 3 in I O O cc. The zinc oxide going into solution t o neutralize this acquires a concentration of 0 . 0 0 1 / 2 X I O or 0.005. The maximum OH- concentration a t the final point of neutrality then should be j.2
X
10-1a
or IO-^.
0.005
Obviously, this is a sufficiently alkaline solution. Yet a few facts may be pointed out in this connection. The concentration of t h e OH ion in a saturated solution of ZnO, assuming complete hydration with solution, is 0.0001,and yet even after boiling a suspension of ZnO i t requires some hours for i t slowly t o t u r n red litmus blue. This indicates t h a t , even though hydrat i p is probably complete with solution, its rate is extremely slow in a solution approaching neutrality; thus the neutralization process in t h e HaPo4 assay would be very slow even with t h e last traces of free HNOa, but when precipitation takes place in acid solution we have not free H N 0 3 finally, but t h e acid phosphates of silver, which i t is doubtful if the zinc oxide would ever neutralize. Conditions approaching good working conditions might possibly be obtained by introducing the zinc oxide first and then adding t h e AgN03 very slowly, thus neutralizing the free acid as i t is formed, if not formed too rapidly. This procedure is not practicable, however, and t h e one outlined above has all the advantages of this last suggestion. SUGGESTED
P R O C E D U R E FOR A S S A Y OF
Na2HP04
AND
Na2HP04.1zHzO
As a n outgrowth of the above-suggested procedure for t h e assay of phosphoric acid, a simple modification for t h e assay of t h e salts has been tried with success. I n the case of t h e exsiccated salt which will run 99 t o IOO per cent Na2HP04 there will be, for every molecule of salt, one molecule of HNO, liberated b y the silver precipitation. On the assumption of a g g per cent or a I O O per cent material i t is easy t o calculate the quantity of alkali necessary t o neutralize this acid, as follows: Weigh out 1 5 g. of dried sample and make up to
Mar., 1921
T H E J O U R N A L OF I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S I ' R Y
cc. To I O cc. of this solution add, for every g. of the salt, I O . 50 cc. 0.x N alkali. (This is one drop overtitration for a gg per cent salt and one drop undertitration for a I O O per cent salt.) Add 50 ec. 0.I N AgN03, make up t o volume in a 100-cc. standard flask, mix, filter through a dry filter, and titrate 50 cc. of the filtrate with 0.I N sulfocyanate. The crystals NazHP04.IzHz0 might be treated t h e same way, adding I O . 50 cc. alkali per 0 . 3 7 8 g. of t h e salt. Here, however, the method falls down, as any appreciable efflorescence would cause the calculkted amount of alkali t o be too small and the results would be low. It would be better t o exsiccate the salt a t xxo0, determine the moisture, and then analyze the dried salt as above. 1000
0.15
SUMMARY
I-The U. S. P. method for the assay of phosphoric acid is incapable of yielding true results except a t one specific concentration, namely. 6 . 2 mg. per cc. The error varies from about + 3 per cent a t a concentration of 0 . 6 2 mg. per cc. to -8 per cent a t a concentration of 11 mg. per cc. 2-This is probably due t o the formation of acid phosphates of silver which are slightly soluble, the amount formed increasing rapidly as the phosphate concentration is increased and the excess of silver nitrate is simultaneously decreased. 3-The fair results obtained a t the specific concentration given in the U. S. P. is probably due t o the accidental compensation of a number of errors a t t h a t particular concentration. 4-By modifying the method t o the extent of transforming the acid t o the tri-sodium salt, results are obtained which coincide with the results yielded by the pyrophosphate method, and which are independent of the Concentration of the phosphate. The Corn Products Refining Company is said to have concluded negotiations for the taking over of plants in England, Germany, and France, after negotiations extending over several months. The company, capitalized a t $80,000,000, intends to manufacture and distribute its goods in Europe, and has worked out a plan on a large scale in order to overcome the high duties collected on business transactions between the United States and European nations. The German plants are to be located a t Hallem, Steutz, Grafenhainichen, and Nierstein. Reports of the glass industryin West Virginia show that a majority of the glasshouses in the state have shut down, only the largest ones which have big contracts going ahead with the completion of their orders. Plants which are in operation have their forces cut down sometimes as much as 75 per cent. Belgian glass is said to be selling a t two dollars per box less than the American price. Lever Brothers Company, of Cambridge, Mass., the American auxiliary of the British company of the same name, has increased its capitalization t o $150,000,000, preliminary to taking over the American Linseed Company. During the past 12 mo. the company has taken over by purchase wholly or in part the capital stock of a majority of the British oil mills and refineries and has consummated a gigantic combination of the industry.
225
NEW METHOD FOR THE DETERMINATION OF POTASSIUM IN SILICATES' By Jerome J. Morgan COLUMBIA UNIVERSITY,NSw YORK,N. Y. Received December 9, 1920
The usual methods of determiniqg potassium in silicates, particularly in fused residues, have been found t o present numerous difficulties. I n some work2 on the volatilization of potassium oxide from natural silicates, a combination of the J. Lawrence Smith method and the perchloric acid method was used in determining the amount of potassium in the residues from an experiment, whenever t h a t residue was in such form t h a t i t could be removed from the platinum boat in which the experiment was made, and ground t o a powder. In many cases the residue had fused, and t o remove i t from the boat i t was necessary t o dissolve it with hydrofluoric acid. I n such cases a combination of the hydrofluoric acid method of Krishnayya3 and the perchloric acid method was employed. While it was possible by either of these methods t o obtain results in duplicate t h a t agreed fairly well, the results by one method did not agree with those by the other, and both methods were tedious and timeconsuming. It was therefore evident t h a t if much work was t o be done on the volatilization of potassium salts from silicate mixtures, i t would be necessary t o find a more rapid and more reliable method for the determination of potassium in silicates. I n order that the method might be applicable t o residues which had been melted and solidified in a platinum boat, i t was necessary t o decompose the silicate with hydrofluoric acid, and hence i t was decided t o get rid of the fluorine by evaporating with perchloric acid instead of with sulfuric acid. This was found t o work very satisfactorily. Only in mixtures containing considerable calcium is i t necessary t o repeat the evaporation after taking up with hot water and adding more perchloric acid t o transform the fluorides completely t o perchlorates. This procedure offers an additional advantage, since the perchlorates of all of the bases, except potassium, commonly found in silicates, are soluble in the alcohol wash used in the regular perchlorate method.4 Hence the residue after evaporation can be treated a t once with alcohol wash, and the insoluble potassium perchlorate transferred from the dish in which the silicate is dissolved directly t o the Gooch crucible in which i t is weighed. The only common substances likely t o interfere in the analysis are ammonium and sulfur compounds. The ammonium compounds can be removed by preliminary heating, and some experiments, which will be mentioned later, seem t o show that any interference of the sulfur compounds can be estimated and a correction applied. The method has been used on a large number of mixtures of feldspar and glauconite with sodium chloride, calcium chloride, calcium carbonate, and limestone, 1 Part of a thesis submitted in partial fulfilment of the requirement for the degree of Doctor of Philosophy in the Faculty of Pure Science, Columbia University, New York, N. Y. 9 Jackson and Morgan, T H I S JOURNAL, 13 (1921), 110. 3 Chem. News, 107 (1913), 100. 4 Scholl, J . Am. Chem. SOC., 36 (1911), 2085.
