Analyses of Sodium Meta-, Pyro-, and Orthophosphates With Some Annotations on Methods ARTHUR B. GERBER
AND
FRANCIS T. MILES, Monsanto Chemical Company, Anniston, Ala.
I
N A previous paper (6) the authors described the deter-
mination of meta-, pyro-, and orthophosphoric acids in mixtures by means of a series of three acidimetric titrations in which the acids are considered, with respect to one atom of phosphorus, as mono-, di-, and tribasic. The end points were defined in terms of colorimetric pH. The present paper outlines the application of these p H titrations to the assay of sodium phosphates and compares the method with a number of other modes of attack, in an endeavor to rationalize some of the larger discrepancies of analytical data shown by various test procedures. I n using the colorimetric p H titrations the same considerations (6) which applied to the polyphosphoric acids apply to the polyphosphates. The method determines the amount of phosphorus pentoxide associated with one, two, and three molecules of sodium oxide or vater, which are reported as meta-, pyro-, and orthophosphate, respectively, but it does not distinguish between a polyphosphate and a mixture of meta- and pyrophosphate of equivalent proportions.
cent by theory, with an average deviation from mean of 0.02 per cent phosphorus pentoxide in 10 titrations. In addition to the items described in the earlier aper (6), sulfuric or nitric acid, roughly 0.5 N , is necessary. Either acid at 0.1409 N is useful in some modifications of the general method. The acid is standardized by titrating a measured amount with the 0.1409 N sodium hydroxide using phenolphthalein indicator. The distilled water used in the titrations and in preparing the standard solutions should be vigorously aerated to remove carbon dioxide. Not more than 0.05 cc. of the 0.1409 N alkali should be required to bring 100 cc. of the aerated water to pH 8.8.
Procedure The procedure given below is a general one for water-soluble sodium or potassium phosphates. Because carbonate may be present in the more alkaline phosphates, treatment to remove interference from this source is necessary. The removal is accomplished by aeration of a n acidified solution of the sample; boiling is inadmissible because of rapid hydration t o orthophosphate. PREPARATION OF SOLUTION.Transfer a weighed sample containing about 1.6 grams of phosphorus pentoxide to a 100-cc. volumetric flask. Dissolve the sample in an excess of 0.5 N sulfuric or nitric acid, measured from a buret, together with sufEcient water to give a volume of 50 to 70 cc. An excess of 3 to 5 cc., as shown by a drop of 0.4 per cent bromocresol green indicator, is suitable. Fit the flask with a 2-hole rubber stopper with one short tube for connection to an aspirator and one capillary tube drawn out to a very fine tip reaching nearly to the bottom of the flask. Pull air through the acidified solution under reduced pressure (less than 100 mm. of mercury) for 5 minutes to remove carbon dioxide. After dilution to the mark and mixing, transfer three 25-cc. aliquots, containing about 0.4 gram of phosphorus pentoxide, to each of three Almquist titrating flasks. BROMOCRESOL GREENTITRATION.Titrate one aliquot with standard alkali to the monobasic end point as described in the prior paper (6). METHYLRED TITRATION. To the above titrated solution at the monobasic end point, add silver nitrate in excess and titrate to the tribasic end point with a standard alkali as described in the prior paper. Add water as desired to improve sedimentation. At the end point, a red color does not return upon further addition of silver nitrate but the addition of one drop of 0.1409 N acid should bring back a red or orange-red color which is not removed by vigorous agitation. The red color will sometimes slowly reappear upon standing, owing to hydration of meta- or pyrophosphate of the sample, but acidity from this source obviously should not be included in the titration. THYMOL BLUETITRATION. To another aliquot add 20 grams of sodium nitrate and titrate with standard alkali to the dibasic end point as described in the prior paper, substituting oleo red B indicator if required. A sodium nitrate blank, found by adjusting 100 cc. of water to pH 8.8, adding 20 grams of sodium nitrate, and titrating t o pH 8.8, must always be subtracted from this titration. METHYLRED TITRATIONFOR TOTALPHOSPHORUS. To the third aliquot add 7 cc. of nitric acid, specific gravity 1.42, and a few glass beads and dilute to 100 cc. Boil for 15 minutes to convert to orthophosphate. Cool, add one drop of bromocresol green indicator, and nearly neutralize with clarified 50 per cent sodium hydroxide solution. Cool again, add 0.5 cc. of 0.4 per cent bromocresol green solution, and accurately adjust with 0.1409 N alkali to pH 4.6 a t 100-cc. volume. Add 25 cc. of 0.85 N neutral silver nitrate solution and 0.5 cc. of 0.2 er cent methyl red indicator, then titrate with 0.1409 N alkali asiefore until the red color of the indicator is just discharged, the solution then becoming greenish yellow in color.
Determination by pH Titrations The description of the analytical procedure for sodium phosphates is restricted, in so far as practicable, to an outline, since the theory, technique, limitations, and accuracy as well as many of the operations and solutions are the same as given in the earlier paper (6). The procedure has been in constant use in this laboratory for more than four years and has proved to be of much value in research and control work. Reference should have been made in the earlier paper to prior workers who employed titration with alkali to the tribasic end point using excess silver nitrate, in particular to Balareff (3) who projected its application to mixtures of the three acids. The titration has been recurrently described without coming into general use. As indicator, lacmoid was used by Balareff, phenolphthalein in the presence of sodium acetate by Wilkie (27), methyl red by Moerck and Hughes (17) and by Sanfourche and Focet (22), and bromothymol blue by Simmich (23). The potential high accuracy of this titration is not realized, however, when used, as customary, in conjunction with the less satisfactory methyl orange and phenolphthalein titrations to the mono- and dibasic end points. The accuracy of these end points is greatly enhanced by definition in terms of colorimetric p H and by recognizing the differences between the equivalence points of the pyroand orthophosphates.
Apparatus and Solutions The method depends upon titrations with 0.1409 N sodium hydroxide solution (1 cc. = 0.01 gram of phosphorus pentoxide) changed from 0.1408 N (6) to agree with the 1940 atomic weight of hosphorus. The solution must be essentially carbonate-free an$ protected at all times against atmospheric carbon dioxide. The solution is Standardized against benzoic acid, using standard (sample 39) and procedure of the National Bureau of Standards. When so standardized, the mean phosphorus pentoxide content of Sorensen grade primary potassium orthophosphate, dried a t 105" C. and titrated directly to the silver nitrate-methyl red end point, was found to be 52.17 per cent, against 52.16 per 406
ANALYTICAL EDITION
June 15, 1941
ACID EQUIVALENT TITRATION. Transfer the same amount of 0.5 N acid as used in preparing the sample solution to a 100-cc. volumetric flask and dilute to the mark. Titrate a 25-cc. aliquot with 0.1409 N sodium hydroxide using phenolphthalein indicator. CALCULATIONS. Using the nomograph (6, Figure 3), determine the pH values to be used as end points for the phosphate mixture under test. Correct the titration volumes using these end points. Using these corrected volumes and with the titrations in order of their description designated as BCG, MR, TB, MRT, and AE, calculate the composition.
+ +
MRT x 0.005 = grams of total PZOS( 0 p m) (TB - BCG) x 0.01 = grams of P,Or (0 f p ) [ MR - (TB - BCG) ] x 0.01 = grams of PzOs ( 0 ) CAE - BCG) X 0.004367 = grams of alkalinity in terms of NazO [ M R - (AE- - BCG)] x 0.801269 = gramsbf replaceable hydrogen in terms of HzO Example. A 2.8-gram sample of a mixture of anhydrous salts (equal parts of sodium trimetaphosphate, sodium tripolyphosphate, tetrasodium pyro hosphate, and primary sodium orthophosphate) was dissolvea and treated with 20 cc. of approximately 0.5 N acid. Aliquots, each representing 0.7 gram of sample, gave titration volumes, corrected to the appropriate end points, as below. All figures refer to 0.1409 N alkali. cc
Substituting these values: 83.80 X 0.005 = 0.4190 gram of total P2O5 ( 0 f p 4- m) (29.36 - 2.89) X 0.01 = 0.2647 gram of Pz05 (0 4- p ) [36.65 - (29.36 - 2.89)] X 0.01 = 0.1018 gram of Pzo5 (0) (18.90 - 2.89) X 0.004367 = 0.0699 gram of alkalinity as NazO [36.65 - (18.90 - 2.89)] X 0.001269 = 0.0262 gram of water of
constitution
By subtractions: Meta-PgO5 = 0.1543 gram or 22.04 per cent Pyro-Pz05 = 0.1629 gram or 23.27 per cent Ortho-PZOs= 0.1018 gram or 14.54 per cent COMPOSITION AS OXIDES. The foregoing calculations give values for total phosphorus pentoxide and water of constitution. The composition in terms of oxides is completed, in the absence of any free acidity, by use of the equation Total NazO in grams = grams of alkalinity f (grams of total P& X 0.4367) where 0.4367 is a conversion factor to obtain that untitrated sodium oxide which is bound to the (total) phosphorus pentoxide in equimolecular proportions. In the above titration example, the oxide components would then be: 36.1 per cent sodium oxide 3 . 7 per cent water of constitution 59.9 per cent phosphorus pentoxide
407
This calculation is sometimes useful in the examination of saltr but is, of course, applicable only to pure sodium phosphates.
COMPOSITION AS SALTS.Although the method cannot identify mixtures in terms of salts of specific basicities, it provides groupings or limits of variable latitude within which the salt components must fall. Compositions in terms of salts, calculated empirically, are often convenient forms of expression in control work, but the equivocal nature of such terms should be understood. In the titration example above, an empirical composition would be: 31.7 per cent sodium metaphosphate
4 2 . 8 per cent tetrasodium pyrophosphate
0 . 6 per cent sodium acid pyrophosphate 2 4 . 6 per cent primary sodium orthophosphate
As is often the case, other salt compositions are possible. The presence of the polyphosphate as such is not revealed by the results, but the equivalent quantities of meta- and pyrophosphate are represented.
Modifications The general procedure described above may be simplified if the predominating composition is known. Commercial pyrophosphates may usually be titrated directly to p H 4.2 and 9.1 without other readings, orthophosphates directly to p H 4.6 and 8.4. If metaphosphate is known to be absent, the determination of the total phosphorus content after hydration is unnecessary. Because the acid salts commonly contain no carbonate, these may often be titrated to the appropriate end points without the acidification and aeration step. Substitution of potentiometric means (6) for the methyl red indicator is sometimes convenient. The general procedure is amenable to considerable rearrangement. The methyl red titration can be made, if preferred, directly on a separate aliquot or it can follow the thymol blue titration in the manner described in the earlier paper. When performed after the thymol blue titration, the sodium nitrate blank which is subtracted from the first titration should be added back to the methyl red titration because the p H a t the methyl red end point is lower than that of the preceding end point. When little or no orthophosphate is present, as in commercial sodium pyrophosphate, and the methyl red titration follows the titration with oleo red B indicator, it is advisable t o add a measured amount, say 1 cc., of 0.1409 N acid to the solution before adding the silver nitrate to ensure approach to the end point from the acid side. Aside from such rearrangements, the titration procedure is rather inflexible, since the p H values used as end points are based on specific conditions of phosphate concentration, salt effects. and pH-color calibration.
TABLE I. SCHEMATIC OUTLIXES OF ANALYSIS Employed b y Britske a n d Dragunow (6) Kiehl and Coats ( 1 1 )
Travers and Chu ($4)
Wurzschmitt and Schuhknecht (28) Madorsky and Clark (14)
Pyrophosphate Titrimetric measurement of acidity produced on adding excess zinc sulfate under controlled conditions 1. Precipitated with zinc acetate under controlled conditions of acidity (pH 3.3) and concentration, then determined gravimetrically 1. Precipitated with zinc sulfate under controlled conditions of acidity ( p H 3.7-4.7) and NH4C1 concentration, then determined gravimetrically Titrimetric measurement of acid produced on adding excess zinc ammonium iodide. Modified Britske titration Determined gravimetrically b y modification of Travers a n d Chu method (acidity a t p H 2.7-2.8)
Orthophosphate
.......................... 2. Determined gravimetrically in filtrate
from pyro b y precipitation with magnesia mixture in presence of ammonium chloride 2 . Determined gravimetrically in filtrate from pyro by neutralizing t o p H 7 t o precipitate zinc ammonium orthophosphate Separated by extracting phosphomolybd a t e complex with ethyl acetate, then determined gravimetrically Determination of total phosphorus permits calculation b y difference. Difference includes polyphosphates
Metaphosphate
....................... Known total phosphorus content permitted calculation b y difference 3. Determined gravimetrically in filtrate
from ortho after boiling with acid t o hydrate t o ortho form
Determination of total phosphorus permits calculation b y difference Determined gravimetrically by precipitation with barium chloride a t controlled acidity, p H 2.2-2.3
408
Vol. 13, No. 6
INDUSTRIAL AND ENGINEERING CHEMISTRY
With obvious changes in factors, the general procedure may be used for potassium phosphates. It is not applicable to the ammonium phosphates where ammonia acts as a buffer in the alkaline range nor to the calcium phosphates where insoluble salts are formed.
TABLE11. PYROPHOSPH-~TE BY DIRECT TITRATIOK NMP~O; Gram
Taken Added G mm
T';a4PZOl Found Gram
1.0000
Sone
(1.000)
0.1000
0.100s 0.1012 0.1128
H
0.6000 0.5000 0,5000
Sone 0.9 (XaPOda 0.9 metaglass 0 . 9 Graham'ssalt 0 . 1 0 IJaH~POa 0 . 4 3 NaHzPOk 0.50 hazHP0k 0 . 4 3 NaHzPOI
1
0.5000
0.56 NasPOk
J
None
0.86 KaHIPO,
K
hTone
1.OO KazHPOk
L M
None None 0.1000 0.1000
0 . 8 6 NaHd'O4 1.OO NazHPOk 0.86 NaHzPOI 1.00 NazHP04 0.50 NazHPO4 0.50 meta glass
Test
Methods and Discrepancies Aside from the acidimetric or neutralization methods, exemplified by the procedure described above, many workers have employed selective precipitations for the estimation of phosphate mixtures. Reactions of phosphate with metallic ions, particularly zinc, are frequently used. A number of these methods are briefly outlined in Table I. Limitations and difficulties of such methods are known, especially to workers closely identified with the field, but apparently have not been accorded general recognition and acceptance. A n enumeration of some of these difficulties and discrepancies is believed to have use in appraising the worth of existing methods, experimental data, and concepts of composition. The enumeration is restricted to phases which may involve discrepancies of a t least several per cent or the identity of phosphates. It is hoped that a wider perception of these difficulties of analysis may lead to much-needed improvements in the estimation and identification of the several phosphates. PYROPHOSPHATE BY DIRECTTITRATION. The titrimetric method of Britske and Dragunow (5) has been employed in the trade for the assay of commercial pyrophosphates and is sometimes useful even in the presence of sodium silicate, carbonate, and soap. The precision of the titration is improved by using, as is now common practice, a glass or quinhydrone electrode with end point of, say, pH 3.8. Britske and Dragunow observed that the accuracy of the titration is somewhat impaired in the presence of much metaphosphate. Inaccuracy on this account, however, seems open to some doubt according to tests in this laboratory, where sodium trimetaphosphate and a sodium metaphosphate glass were employed as sources of metaphosphate. Sodium trimetaphosphate as used in this paper is that water-soluble polymer (16), solutions of which show no precipitate .ivith silver or barium cations. The metaphosphate glass was prepared after Pascal (19) by the fusion and quick cooling of the trimetaphosphate, since this gives a product of higher metaphosphate content than the usual Graham's salt. The potentiometric titration was performed as follows: Dissolve the sample, usually 1 gram, in 50 cc. of water, add 0.2 N hydrochloric or sulfuric acid in slight excess (pH less than 3.8), then dilute to 100- to 110-cc. volume. Csing a glass or quinhydrone electrode titration assembly, adjust to pH 3.8 with 0.1 N sodium hydroxide solution. Add 40 cc. of 0.5 M zinc sulfate solution (equivalent by theory to 2.6 grams of tetrasodium pyroghosphate) which has been adjusted with sulfuric acid to pH 3.8. low y titrate the acid, liberated by the precipitation of the normal zinc pyrophosphate, with 0.1 N alkali while stirring vigorously until pH 3.8 is reached and maintained for at least 5 minutes. The alkali consumed is a measure of the pyrophosphate. Because the reaction as measured is not strictly stoichiometric, the standard alkali is calibrated against pure tetrasodium pyrophosphate under like test conditions.
The test results indicate that metaphosphates, a t least the In the presence of the metaphosphate glass some apparent inaccuracy is found (Table 11, Test C) but this may be due to somewhat imperfect composition of the metaphosphate (Graham's Yalt, Test D, illustrates imperfect composition) or, as is more probable, to slight hydration during test. I n the presence of orthophosphate the Britske and Dragunow titration is subject to some inaccuracy (which increases with the orthophosphate content) because of the salt effect (1.3) of the zinc sulfate reagent upon the hydrogen-ion cont w o varieties employed, cause no particular difficulty.
A B
C
D
E F G
N
0 P Q
0.1000 0.1000 0.1000 0.9000
0 1000
Kone
+
........... .... ........... . . . .
........... ............. .. .. .. ............
0.9058
Titrated a t 20' C. Titrated a t 20' C. 0.55+ At 32' C., unstable point O.54-t .4t 32' C., unstable point O.25+ .It 32' C.. unstable point 0.2S-t At 32' C., unstable point 0.0206n At 2 5 O C., increased 0 . 0 2 0 P A t 25' C., increased 0.12160 At 25' C., increased 0.1223" A t 25' C., increased 0.1196" At 2 5 O C., increased 0.5272
0.5288
10 7101
R
None
0 4943
S
None
0.42680
a
Calibration, 4 tests; 74.62, 74.58, 74.46, 74.68 cc. of 0.1 N alkali
0,1546
\ 0.5157 1.0 KaaPs01o
Remarks
end end end end
vol. vol. vol. vol. vol.
First reading, unstable a d point Final end point: solid phase contains 0.562 gram of 0.579 gram PZOSpresent Stable end point Increased volume, no ppt. a t end point
Initial dilution to 250 cc. instead of 100-110 cc.
centration. In addition, large amounts of orthophosphate may interfere through precipitation. The interference is accompanied by impermanent end points. The amount of orthophosphate which can be tolerated is greatly influenced by concentration and temperature. Under the abovedescribed titration conditions, orthophosphate constituting one-half of the sample (0.25 gram of orthophosphate as PzO5) is tolerated but only in cool solution (Tests F to K). When orthophosphate predominates in the sample, a greater initial dilution, say to 250 instead of 100 to 110 cc., will permit completion of the titration without precipitation of orthophosphate though still subject to the salt error (Tests L to 0). Wurzschmitt and Schuhknecht (28) reported that the Britske and Dragunow titration was inaccurate in mixtures containing primary sodium orthophosphate, although the same inaccuracy did not occur in mixtures containing equivalent quantities of secondary or tertiary orthophosphate. The primary orthophosphate was said to result in formation of a precipitate which vitiated the pyrophosphate values; polymeric forms of orthophosphate were suggested in explanation. This exceptional phenomenon cannot, however, be confirmed in this laboratory. Under the titration conditions described above, no difference in behavior could be discerned (Tests E to 0) whether primary or secondary and tertiary orthophosphate were present. Even without resort to refinements of procedure and calibration, the Britske and Dragunow titration is seen to be useful in the analysis of mechanical mixtures of pure sodium meta-, pyro-, and orthophosphates. The method, however, is not used in the trade as a general one for sodium phosphates, because of the ability (IO)of sodium meta- and pyrophosphate melts to form complexes or polyphosphates which are further dependent upon the heat treatment of the melts. When these complexes exist, the usual reactions of meta- and pyrophosphate with metallic ions are altered. For example. when the Britske and Dragunow method is applied to sodium tripolyphosphate, Na6P3010, titrations are obtained corre-
ANALYTICAL EDITION
June 15, 1941
sponding to 42 to 71 per cent of pyrophosphate, depending upon test conditions. I n t,he smaller volume (Test Q), 97 per cent of the phosphorus has been precipitated a t the end point. In the larger volume (Test S), no precipitation occurs although considerable acidity is developed. Obviously the meta- and pyrophosphates have been altered by heat treatment; pyrophosphate titrations are inaccurate in t'he presen ce of polyph osphate. The tripolyphosphate referred to in this paper is that fused and tempered phosphate obtained according to the description of Andress and \Tust ( 1 ) . It was prepared in the laboratory from a mechanical mixture containing 72.3 per cent of tetrasodium pyrophosphate and 27.7 per cent of sodium metaphosphate, corresponding t o equimolecular proportions. An aqueous solution of the powdery polyphosphate showed positive in both cases when tested by the qualitative methods described by Wurzschmitt and Schuhknecht (68) for hexametaphosphate and for trimetaphosphate. It coagulated albumen as does metaphosphate. However, it gave negative tests when the precipitation methods of Madorsky and Clark (14) for metaphosphate and for pyrophosphate radicals were applied. The tripolyphosphate showed the following values when tested by t'lie colorimetric pH titration method: 19 . O per 38.9 per 0 . 0 per 0.034 per 4 2 . 1 per (16.9 per
cent metaphosphate as PzOs cent pyrophosphate as PzOa cent orthophosphate as PzOs cent water of constitution cent sodium oxide aent alkalinity)
Empiric composition as salt's: 7 2 . 3 per cent tetrasodium pyrophosphate 0 . 4 per cent disodium pyrophosphate 2 7 . 3 per cent sodium metaphosphate
The agreement of the above empiric composition with t'he known starting materials indicates that the corresponding tripolyphosphoric acid dissociates in substantially the same manner [3 strongly and 2 weakly dissociated hydrogen atom,< (20, 26) but with none of the feeble dissociation shown by orthophosphate] as does an equivalent mixture of meta- ant1 pyrophosphoric acids. SEPARATIOXS WITH ZIM ION. Some workers have determined the pyrophosphate radical by separating the precipitate formed with zinc ion in the presence of ammonium chloride and under conditions of controlled acidity, variously defined between p H 2.7 and 4.7. In addition, the separation is also influenced by the concentration and distribution of the several phosphoric acid radicals, the amount of zinc reagent, temperature, and time allowed for precipitation. So far as known to the authors, this multiplicity of variables has not to date been well enough defined to warrant the application of the zinc separation as a general method for sodium phosphate mixtures. This does not necessarily imply that the variables cannot be suitably defined. In this paper a few of the possible sources of inaccuracy are mentioned but no attempt is made to delineate satisfactory conditions of separation or of application. The separation of the pyrophosphate radical seems to be particularly disturbed by the ability of sodium meta- and pyrophosphate melts to bind (10) many metallic ions as water-soluble complexes. A number of iiidust'rial applications of such melts are predicated upon this and related phenomena, variously described ( 1 , 7) as the inhibition or retardation of numerous precipitations, the pept'ization of metal salts, and the stabilization of conditions of supersaturation. Although attention has been directed chiefly to reactions with calcium ion, analogous behavior may occur with other multivalent cations including zinc. This is illustrated by applications of the zinc sulfate precipitat'ion method to known mixtures.
409
In these tests, the known phosphates were dissolved in 200 cc. of water, the solution was adjusted with dilute hydrochloric acid
to pH 3.0 to 3.3, 20 cc. of 0.5 M zinc sulfate solution were added, and the acidity was readjusted with dilute sodium hydroxide solution until pH 3.3 a t 25" C. was maintained for 10 minutes with final volume of 250 cc. Aliquots of the solution, clarified by filtration, were analyzed for phosphorus content by the molybdate method. The difference between the known total phosphorus and the phosphorus found in the liquid hase gives the phosphorus in the solid phase. This latter vayue accordingly does not involve either the composition of the precipitate or its washing technique. Tests were also made in the presence of 25 grams of ammonium chloride, since Travers and Chu recommend its addition (10 per cont) to prevent the precipitation of zinc metaphosphate. The typical results in Table 111,where the per cent recoveiy of pyrophosphate is given in the right-hand column, shov that the presence of pho~pliateshaving metallic ion binding power may lead to gross errors when the pyrophosphate radical is sought by the reaction with zinc cation. I n vieu of the many variables,, the test conditions used above give a very limited portrayal of the possibilities of the reaction but they seem sufficient LO indicate that the determination of pyrophosphate in this manner !Tarrants more than casual investigation and description. In some limited applications, of course, the method employing the zinc pyrophosphate precipitation is appropriate. A specific instance is wemingly afforded in the work reported by Kiehl and Coats 111)in their study of the hydration of sodium metaphosphate in alkaline solution. From the mode of preparation and the reported water of crystallization, the metaphosphate under study appears to have been the trimetaphosphate, althmgh erroneously designated as the monometaphosphate IZS noted by Sylen (18). The trimetaphosphate has no metallic ion binding power (26) nor does it form a precipitate ivith zinc ion (15) under the test conditions as may the glassy metaphosphate. From the manner of preparation, the presence of polyphosphates or complexes x i t h their attendant difficulties may be excluded. Moreover, the difficult condition (28) of small amounts of pyrophosphate was not encountered because the end products of the hydration mere equimolecular quantities of pyro- and orthophosphate. Although the separation as employed appears to be suitable for the particular problem at hand, its application as a general method mas not investigated nor implied and would be manifestly inappropriate. On the other hand, 'l'ravers and Chu (24) reported commercial glacial metaphosphoric acid to be a mixture of the meta and ortho forms and found 18.6 per cent of sodium oxide. The ortho form was deduced from acidimetric titrations
KITH ZINC Iox TABLE111. SEP.$RATIOSS ---
-
Take,,-p----.-
Other phosphate M u PZOi
Pi05
106.8 106 8 53.4 None None 53.4 106.8 106.8 106.8 106. h Yon? Sone 53.4 106.8 106.8 106 8 106.8 53.4 53 4
Sone None None ? 4 7 . 0 Graham's salt $347.0 Graham's salt !47.0 Graham's salt ,347.0 Graham's salt 347 0 Graham's salt 1 3 8 . 4 Graham's salt 1 3 8 , 4 Graham's salt "88 5 Xia6PaOlo 288.5 NasPs0l.i 288 5 NanPaOlo 288 5 NasPs013 288.5 SasPaO10
106 S 106 8 53.4
115.4 115.4 347.0 312 6
NaiP:O1, Na6Pa010
(SaPOa): Sa2HPO..
Mg.
347.0 347.0 400.4
453,s 453 8 245.2 245.2 288.5 258 5 341.9 :395 2 ,393 2 222.2
'21,2 400 4 :36G 0
NEIaCI Grams Sone 25 Sone None 25 None Xone 25 Sone 25 None 25 Zonr Son? 25
104.7 91.4 51.8 Sone None Xone 96.3 None 127.7 None Kone Sone None 111.2 Yone
Sonc 25
119..i
None
53.0
Sone
-
--Found Solid PZOSin solids phase PnOa in pyro M g . PnOs
,-
Total
NruP20J f g . PnOs
None
52.8
98.0 85.6 97.0
... ...
0.0 90.2 0 0 119.5 0.0
...
0:o 104.1 0.0 111.9 0.0 99.2 98.9
410
INDUSTRIAL AND ENGINEERING CHEMISTRY
using methyl orange and phenolphthalein indicators which showed the presence of a multibasic acid; pyrophosphoric acid was eliminated because of the absence of precipitate with zinc ion, thus leading to the assumption that the acid was the orthophosphoric form. As shown in Table 111, however, the absence of a precipitate may not be conclusive evidence of the absence of pyrophosphate. Moreover, the possibility of polyphosphates or complexes was not eliminated. When tested in this laboratory by pH titrations, a fresh sample of a commercial c. P. grade of metaphosphoric acid (glassy sticks with little white opaque coating) showed 2 5 . 5 ner cent metaDhosnhate as P&r
36.8 per cent pyrophosbhate as PlOs 9 . 2 per cent orthophosphate as PgOr 1 7 . 2 per cent base a8 sodium oxide 11 1 per cent water of constitution
The free acid content was equivalent to 32.1 per cent phosphorus pentoxide. I n a supplementary test the orthophosphate content was found to be 9.6 per cent by a modification of the molybdate precipitation method of Hitchens and McCauley (8). That the glacial acid is more than the mixture of meta and ortho forms, which Travers and Chu report, is further indicated by its behavior in the Britske and Dragunow titration where the alkali consumed corresponds to 28.7 per cent pyrophosphate. Although the titration is not accurate for this type of material, the large alkali consumption is a proof that more than meta- and orthophosphoric acids are present. In view of the interference introduced by the presence of polyphosphates and polymers, the question naturally occurs as to the frequency with which this may be encountered. Because metaphosphates are seldom, indeed rarely, free from complexes having ainc ion binding power, the presence of metaphosphate (or of polyphosphate) in any phosphate mixture may be regarded as a potential source of more or less difficulty in the determination of pyrophosphate using zinc ion. METAPHOBPHATE. The direct gravimetric determination of metaphosphate using barium chloride as precipitant has heen employed by Holt and Myers (9) and by Madorsky and Clark (14). The test, however, is not a generic one because of differences in behavior of the polymeric forms. Thus, the determination surely would not include trimetaphosphate, nhich forms no precipitate with barium ion (16) under the test conditions. The full significance and manner of application of the barium precipitation seem incompletely developed at this time. Holt and Myers reported polymeric forms of pyrophosphoric acid, (1) a simple form obtained by displacement from lead pyrophosphate and (2) a polymer of formula between (H4P207)d and (H,Pa0,)5 which was obtained by dehydration of orthophosphoric acid. The latter form, however, is a mixture of acids in a condition of equilibrium as had been adduced by Berthelot and Andre (4) and recently described by Lum, Malowan, and Durgin ( I S ) . An equilibrium constant has been determined (6) by the use of pH titrations. For the examination of the metaphosphate reactions, Graham’s salt is frequently used as a source of the sodium salt. When prepared by the usual fusion of primary sodium orthophosphate a t red heat, the conversion to metaphosphate is more or less incomplete because of the difficulty with which the last bit of water is expelled. As a result, solutions of the melt are somewhat acid in reaction. The residual mater of constitution can be readily estimated by quickly dissolving a portion of the chilled melt in cold \Cater, adding an excess of silver nitrate reagent, and titrating with standard sodium hydroxide solution using methyl red indicator. From their examination of Graham’s salt, Treadwell and
Vol. 13, No. 6
Leutwyler (95)concluded that the metaphosphate is always attended by an appreciable amount of orthophosphate (10 per cent KaH2POI in an example) which is present as a reaction product; a theoretical explanation involving splitting of molecules was given. The orthophosphate was said to be best determined from its buffer action by measuring the caustic consumption from pH 4.3 to 9.5. Although this titration is an excellent tool in the examination of sodium metaphosphate for composition, it is not conclusive as a measure of orthophosphate unless the absence of pyro- and polyphosphates or complexes has been fully assured, for which purpose the electrometric titration curves as used seem inadequate. When Graham’s salt was prepared in this laboratory by heating primary sodium phosphate (25 grams) in an oven a t 700” C. for an hour, an analysis of the chilled melt by the pH titration method showed the presence of 3.8 per cent pyrophosphate as P205 but no orthophosphate, while the water of constitution was 0.46 per cent. Moreover (Table 11,Test D), Graham’s salt (from the above preparation) behaves in the Britske and Dragunow titration as though it contained some pyro- or polyphosphate; such behavior would not be expected if the buffering salt were orthophosphate. From this and many similar observations, the conclusion is reached that, contrary to Treadwell and Leutwyler, orthophosphate is not a reaction product of the dehydration of primary sodium phosphate a t red heat, even though important amounts of buffer salt are present. Salih (21) reported that a hexametaphosphoric acid, obtained by displacement from the lead salt, showed 4 strongly and 2 weakly dissociated hydrogen atoms. Such dissociation, if true, would invalidate the colorimetric pH titration method. However, Treadwell and Leutwyler contest (25) Salih’s work and find metaphosphoric acids to be strongly dissociated (26), an observation in agreement with that of the authors and of Rudy and Schloesser (20). The value CpH 9.1) given by Salih for the hydrogen-ion concentration of a 0.1 N solution of the sodium salt suggests that the acid under examination may have inchded phosphoric acid of a higher degree of hydration than the meta stage. Wursschmitt and Schuhknecht (28) recently described qualitative tests for the detection of ortho-, pyro-, hexameta-, and trimetaphosphate. Although the tests are said to be applicable to phosphates in the presence of one another, this phase seems insufficiently developed and no known mixtures were included in the tabulated results showing test performance. Unfortunately, many reactions, even though suitable for the identification or estimation of phosphates occurring singly, are inadequate when two or more phosphates are present. An example of such inadequacy is afforded by the qualitative test where the presence of phosphate in the filtrate from a precipitation with barium chloride in a slightly alkalinized solution of the sample is said to indicate the presence of trimetaphosphate. I n parallel tests where sodium orthophosphate, Graham’s salt, and a 1 to 5 mixture of the two were treated as directed by Wurzschmitt and Schuhknecht, the filtrate corresponding to the 1 to 5 mixture showed a prompt and strong test for phosphate. On standing, all three of the treated filtrates showed the yellow phosphomolybdate precipitate; the least precipitate corresponded to the orthophosphate sample while by far the most precipitate corresponded to the 1 to 5 mixture. Although the strong test shown by the 1 to 5 mixture would indicate, according to directions, the presence of much trimetaphosphate, such a conclusion obviously is not justified in this case. Instead, the prompt positive test seems only to indicate that the precipitation of orthophosphate with barium ion may be retarded when Graham’s salt is present, thus allowing orthophosphate
June 15, 1941
ANALYTICAL EDITION
to appear in the filtrate. Analogous interference in the precipitation of pyrophosphate radical n i t h zinc ion has been noted in Table 111. I n addition to being thus exposed to false positive or false negative results in the presence of phosphates having metallic ion binding power, the qualitative tests of Kurzschmitt and Schuhknecht are confined t o the four phosphate radicals with no cognizance being taken of complexes or polyphosphates. Because of these limitations, employment of the tests in a general way leaves much to be desired. Thus commercial sodium metaphosphates were identified (28, Table I) only as the hexametaphosphatc with occasional traces of orthophosphate although very considerable amounts of buffering salt have always been found when such phosphates have been tested acidimetrically. A domestic commercial purified sodium metaphosphate reagent, which the manufacturer affirms to be typical, was recently found by pH titrations to have an empiric salt composition of 11.9 per cent sodium metaphosphate, 86.8 per cent sodium acid pyrophosphate, and 1.1per cent tetrasodium pyrophosphate. The metaphosphate content of 11.9 per cent was essentially the water-insoluble variety. NOMENCLATURE.Determinations of the phosphoric acids by various modes of attack introduce an ambiguity in nomenclature. I n the acidimetric method, meta-, pyro-, and orthophosphoric acids are regarded as phosphorus pentoxide in three stages of hydration or union with water (1, 2, and 3 moles of water per mole of phosphorus pentoxide) which can be estimated by means of a 3-step neutralization. The test result shows phosphorus pentoside hydrated to some particular degree which is expressed in terms of the three stages. On the other hand, if the determination is done by the separation and estimation of the acid radicals-for example, Pz07 using zinc ion or (POa), using barium ion-the phosphoric acids cannot be considered as a ternary group except perhaps as they may be regarded as ortho-, poly-, and metaphosphoric acids, in which case pyrophosphoric acid is deemed (16) to be a member (bipolyphosphoric acid) of the series of polyphosphoric acids. I n any event, consideration cannot be confined to acid radicals of the ortho-, pyro-, and metaphosphates but must be extended (IS, 26) to include the more complex combinations. Failure to recognize this situation has been common. As a result much work that has appeared regarding composition and transformations of the phosphates is of little value unless the methods of analysis are first examined. Until the number and identity of the acid radicals of the polyphosphates-and also of the polymers of metaphosphoric acid-are better defined, the practice of determining any phosphate radical by difference is rarely admissible. Because of these differences in the manner of attack, the value for pyrophosphate, as an example, when determined by the acidimetric method may a t times involve complexes or polyphosphate and thus be higher than the pyrophosphate found by precipitations which would reflect only the pyrophosphate radical, Pz07. Such ambiguity, involving definition of terms, has been a not infrequent source of discrepancies in phosphate data since workers have variously employed acidimetry, precipitations, and a t times, jumbled combinations of the two. Some workers have employed x-ray diffraction methods in the examination of sodium phosphates for the purpose of identifying the chemical compounds. Thus Andress and Wust reported (1) differentiation of the more distinctly crystalline compounds but found ( 2 ) the x-ray method unsuitable for the examination of certain vitreous phosphate complexes which are essentially amorphous. llthough the pH titration method cannot perform the function of distinguishing complexes or polyphosphates from
411
an equivalent mixture of meta- and pyrophosphates, it has been found useful in practice when applied to a wide variety of phosphates encountered in manufacture and research. In these applications the method serves to follow molecular dehydrations in heat treatments or molecular hydrations in contact with water, to determine the base-acid ratio or the composition in terms of the oxides, and to provide limits within which the salt composition must fall. The acidimetric method appears to be applicable to the assay of phosphoric acids and their sodium salts excepting only those compounds, usually high in the meta form, which hydrate under the test conditions. Such hydration errors are small or negligible except in the case of those acids, practically acid anhydrides, containing 88 to 90 per cent of phosphorus pentoxide and higher, which are not brought into aqueous solution without considerable hydration. The error in this case has been readily discerned because the sum of the phosphorus pentoxide and the combined water, calculated from the titrations, then exceeds the weight of sample.
Summary Acidimetric titrations, previously used for strong phosphoric acids, are applied to the analysis of sodium phosphates where they provide a useful classification in terms of meta-, pyro-, and orthophosphate. Polyphosphates, if present, titrate as mixtures of meta- and pyrophosphate. Sources of discrepancies, when using various methods of phosphate analysis, are enumerated. I n those methods where the estimation of phosphate radicals is accomplished by precipitations using metallic ions, consideration cannot be confined, as has been common, to acid radicals of the meta-, pyro-, and orthophosphates but must be extended to embrace the phosphate complexes or polyphosphates. I n the presence of polyphosphate, analysis by acidimetry and by precipitation introduces a potential source of confusion in terminology. The determination of phosphate radicals, using metal salts as precipitants, may be subject t o interference because of the metallic ion binding power of sodium meta- and pyrophosphate melts. Thus the titrimetric or gravimetric estimation of the pyrophosphate radical using zinc ion is exposed to inaccuracy in the presence of polyphosphate. Other causes of substantial discrepancies in phosphate identity or i n analytical data are considered.
Literature Cited Andress, K. R., and Wiist, K., 2. anorg. allgem. Chem., 237, 113-31 (1938). [bid., 241, 196-204 (1939). Balareff, D.. Ibid., 99, 184-6 (1917). Berthelot, M., and Andrt., G., Compt. rend., 123, 776-82 (1896). Britske, E. V.. and Dragunow, S. S., J . Chem. Ind. (Moscow) 4, 49-51 (1927). Gerber, A. B., and Miles, F. T., IND.ENG. CHEM.,Anal. Ed., 10, 519-24 (1938). Hatch, G . B., and Rice, O., IND. ENG.CHEM.,31,51-7 (1939). Hitchens. R. M., and McCauley. M. S.. J. Am. Pharm. Assoc., 25, 990-2 (1936). Holt, A., and Myers, J. E., J. Chcm. Soc., Trans., 99, 384-91 (1911), Huber, H., Angew. Chem., 50, 323-6 (1937). Kiehl, S. J., and Coats, H. P., J . Am. Chem. Soc., 49, 2180-93 (1927). Kolthoff, I. M., and Furman, N. H., “Volumetric Analysis”, Vol. 11, p. 143, New York, John Wiley & Sons, 1929. Lum, J. H., Malowan, J. E., and Durgin, C. B., Chem. Met. Ew., 44, 721-5 (1937). Madorsky, S. L., and Clark, K . G., IND.ENG.CHEM.,32, 24.4-8 (1940). Mellor. J. W., “Treatise on Inorganic and Theoretical Chemistry”, Vol. VIII, pp. 986-7, London and New York, Longmans, Green and Co., 1928. ~I
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(16) Ibid., p. 990. (17) Moerck, F. X., and Hughes, E. J.. A m . J . Pharm., 94, 65&5 (1922). (18) Nylen, P., 2. anorg. aZZgem. Chern., 229, 30-5 (1936). (19) Pascal, P., Compt. Tend., 177, 1298-1300 (1923). (20) Rudy, H., and Schloesser, H., Ber.. 73B, 484-92 (1940). (21) Salih, Mme. R., Bull. SOC. chim., [5] 3, 1391-6 (1936). (22) Sanfourche, A.. and Focet, B.. Ibid.. (41, 53, 963-9 (1933).
(23) (24) (25) (26) (27)
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Simmich, H . , Angew. Chent., 48, 566 (1935). Travers, A,, and Chu, Y. K.. HeZv. Chim. Acta, 16, 913-17 (1933:. Treadwell, W. D., and Leutnyler, F., Ibid., 20, 931-6 (1937). I b i d . , 21, 1450-9 (1938) Wilkie. J. M., J. SOC.Chem. I n d . , 28, 68-9, 464, 980 (1909);
29, 794-9 (1910). (28) Wurzschmitt, B., and Schuhknecht, IT..Angew. Chem., 52, 711-8 (1939).
Determination of Succinic Acid in Plant Tissues GEORGE W. PUCHER AND HUBERT BRADFORD VICKERY Connecticut .4gricultural Experiment Station, New Haven, Conn.
Succinic acid is extracted from plant tissue with ether, freed from contamination with other substances by oxidation, converted into its anhydride, and condensed in toluene solution with p-toluidine to the insoluble crystalline succinyl-p-toluide. The properties of this substance are such as to permit of substantially quantitative isolation, and of identification by means of
S
UCCINIC acid is now known to be a metabolite involved
in the respiration system of certain animal tissues ( 1 , 2 , 8, l4,16),and current speculation suggests (4,21) that it may play an analogous part in plants as well. Accordingly, the recognition and, if possible, the accurate determination of succinic acid in plants become a matter of concern to the further development of our understanding of the physiological functions of the organic acids that form so important a part of most plant tissues. Early methods of determining succinic acid, most of which depend on the precipitation of an insoluble inorganic salt, were thoroughly reviewed in 1909 by von der Heide and Steiner (11) who recommended the use of the silver salt a5 being the most satisfactory. More recent modifications of the silver salt method, designed for application to the analysis of animal tissues, have been described by several investigators (6, 7,9,16). During the past few years, however, attention has been directed chiefly to methods that depend on the oxidation of succinic acid in the presence of an enzyme that occurs in muscle tissue (2, 12, 22). These methods are especially suited to the estimation of the minute amounts of succinic acid that are encountered in the small samples of animal tissue employed for the study of the details of metabolism. For work with plant tissues, none of these methods is completely satisfactory. Our knowledge of the qualitative composition of the mixture of organic acids present in most plants is very limited and salt precipitation methods are specific only under the most carefully controlled conditions. The enzyme method, although probably specific, requires the use of the IF'arburg manometric apparatus and of a highly specialized technique, and its applicability to the conditions encountered in the plant field is still to be established. It seemed desirable, therefore, to develop a simple method which would permit the isolation of a characteristic derivative of succinic acid in order to provide both qualitative and quantitative evidence of the tissue composition.
the melting point and crystalline form. An empirical solubility correction and a conversion factor are provided that lead to average recoveries of the order of 99 per cent, and single determinations can be made within * 5 per cent over the range from 1 to 20 mg. of succinic acid. The only known interfering substance is a-ketoglutaric acid.
Auwers, in 1896 (S), observed that anhydrides of organic acids condense with certain aromatic amines to give insoluble crystalline derivatives. Succinic acid is readily converted into its anhydride when heated with acetyl chloride (6),and the condensation product of succinic anhydride with p-toluidine mentioned by Auwers is a well-crystallized substance that is almost insoluble in toluene and possesses excellent properties for quantitative isolation and for identification. The reactions involved are: CHZ-COOH CH3COClI CHZ-COOH Succinic acid Molecular weight
CH2-CO
1
CHpCO/ Succinic anhydride
' 0
+ H*N.C$Hd.CHI p-Toluidine
118.1
CHz-CO-NH-CeH4.CHa
--+ AH2-COOH Succinyl-p-toluide Molecular weight 207.2 The theoretical yield from 1 mg. of succinic acid is 1.75 mg. of toluide. Although the use of amines of even larger molecular weight would appear to be desirable, none of those tested gave condensation products with as favorable properties as the p-toluide.
Preparation of Organic Acid Fraction Fresh plant tissue is prepared for analysis by being dried in it ventilated oven at 80' C. and is then ground t o a powder. Of this, 1.0 gram is accurately weighed and is extracted with ether its described by Pucher, Wakeman, and Vickery (17). The ether extract is treated with 25 ml. of water and the ether is evaporated. Troublesome frothin is sometimes encountered if alkali is added before evaporation. %he solution is diluted with 25 ml. of water,
