Oxidation of Hydrazine IX. Mono- and Di-delectronation of Hydrazine by Permanganate in Hydrochloric Acid Solution A. G. HOUPT,K. W. SHERK,AND A. W. BROWIVE Baker Laboratory of Chemistry, Cornel1 University, Ithaca, N. Y.
T
HE action of various oxidizing agents upon hydrazine
less for practical purposes,” and, with the comment that “neither he (Medri) nor Browne and Shetterly succeeded in clearing up this question,” suggested (6) the substitution of hydrochloric for sulfuric acid as a modification of the Petersen method. Kolthoff stated that the titration of a hydrochloric acid solution a t boiling gave very good results, and that the oxidation corresponds quantitatively to the equation N2Hd 20 +Nt 2Ha0
has engaged the attention of numerous investigators, chiefly analytical chemists, whose principal motivation in the work has been the formulation of reliable nitrometric or volumetric methods for the determination of either hydrazine or some oxidizing agent. One of the main objects of the investigations recorded in the current series of articles (4) has been the elucidation of the complex mechanism involved in the oxidation of hydrazine (9). Incidentally, in view of the circumstance that ammonia and hydronitric acid are frequently formed as products of the incomplete oxidation of hydrazine, a warning to those who would devise analytical methods involving the use of hydrazine has been sounded to the effect that “positive proof of the quantitative conversion of the hydrazine to nitrogen and water should be obtained , , , , before reliance is placed upon any method which under ordinary conditions is subject to the error just mentioned” (8).
+
Further (6) that “the color change may be detected sharply, and therefore hydrazine sulfate shduld be a good primary substance for the standardization of permanganate.” In comparing the results obtained by the Petersen-Kolthoff method in this laboratory with those obtained by the bromate method of Kurtenacker and Wagner (7) and by titration of hydrazine against permanganate in alkaline solution, as suggested by Sabanejeff (IS), it was found by the authors that a given amount of hydrazine will reduce in alkaline solution nearly double the amount of permanganate reduced in hydrochloric acid solution. This is in qualitative accord with the potentiometric titrations carried out by Stelling (14). The results of numerous preliminary experiments indicate that one atom of oxygen, only, is required for the oxidation of one molecule of hydrazine, instead of two atoms, as stated by Kolthoff. Determinations of the ammonia in the residual solutions pointed toward the conclusion that about one-third of the nitrogen of the hydrazine is converted to ammonia, in accordance with the tentative equation 3N2H4 3 0 +2Nv 2NHa 3Hz0
..
U75 1250
1225 1200 1175
+
0 if50
+
+
In order to subject the reaction to careful scrutiny, an extended series of titrations has been carried out, under widely varying conditions, as well as under the particular conditions that prevailed in the work of Kolthoff,
r“ sf725 0
i=
p & J
EXPERIMENTAL
1075
MATERIALS USED.The solutions of hydrazine sulfate and hydrochloride (Kahlbaum) were standardized by the bromate method (7). The permanganate solutions were titrated against standard solutions of both oxalic acid and sodium thiosulfate. All burets and pipets used in the work were carefully calibrated, FACTORS STUDIED. Among the factors found to influence the course of the reaction were (1) amount of hydrochloric acid added, (2) initial concentration of the hydrazine salt, (3) concentration of manganous ion, (4)temperature, and (5) mode of bringing the reagents together.
f 050
1025
lm a975
FIGURE 1. INFLUENCE OF MANGANOUS IONON RATIOOF HYDRAZINE TO OXYGEN
+
THE
PROCEDURE. The analytical procedure was in general briefly as follows: The solution obtained by adding a specified volume of 4 N hydrochloric acid to a measured volume of the standard hydrazine sulfate or hydrochloride solution was heated to the boiling point in a 250-cc. Erlenmeyer flask and immediately thereafter was titrated with ermanganate, added drop by drop,
The reaction between potassium permanganate and hydrazine sulfate in sulfuric acid solution has been studied by Petersen (9),Roberto and Roncali (I@, Medri (8), Browne and Shetterly (I), Raschig ( I I ) , Cuy, Rosenberg, and Bray (3), and Purgotti (IO), with results construed in some cases as favorable and in others as adverse to the use of this reaction in the determination of either hydrazine or permanganate. Kolthoff (6) characterized the Petersen method as “worth-
until the pink color persistezfor several minutes. In studying the influence of the five factors enumerated above, the procedure just outlined was of necessit modified in certain cases. (1). In experiments 1 to 36 (Table the volume of 4 N hydrochloric acid was varied over a range of 6 to 40 cc. for 50 cc.
fi
54
ANALYTICAL EDITION
January 15, 1935
of the hydrazine solution. ( 2 ) In experiments 42 to 60 (Table 11) 10 cc. of 4 N hydrochloric acid were added to different volumes of the hydrazine solution, and the total volume was made up t o 52 cc. by addition of appropriate amounts of water. (3) In experiments 61 t o 81 (Table 111) varying amounts of manganous sulfate (0.1 N solution) were added to the hydrazine solution with 10 cc. of 4 N hydrochloric acid. (4).During the titration the temperature of the solution was permitted t o fall gradually
from the boiling point to about 55" C., exce t i n experiments 4,5, and 18, where it was maintained at 90" over an electric hot plate. ( 5 ) The permanganate solution was added drop by drop,
cf:
55
sulfate solution mixed with either sulfuric or hydrochloric acid were introduced dropwise beneath the surface of a boiling solution of potassium permanganate prepared by dissolving 10 grams of the solid in 90 cc. of water. Boiling was continued for 10 minutes after all the hydrazine had been introduced. After the solution had cooled, the excess of permanganate was reduced by adding 10 per cent hydrogen per1.525
with a sufficient interval between drops t o permit complete discharge of the color. It was found inadvisable to introduce more than 25 cc. of the oxidant in 7 minutes, since any appreciable local excess of the permanganate favors the liberation of chlorine. The effect of a large excess of permanganate was investigated by dropping known amounts of hydrazine, both with and without hydrochloric acid, into boiling permanganate solution. Ammonia was determined in the residual solutions by distillation from alkaline solution into a measured excess of standard hydrochloric acid. A series of experiments was performed in order to determine the ratio of hydrazine oxidized to nitrogen evolved as gas. A water-jacketed nitrometer of the Lunge t e was used. The permanganate solution was slowly droppegnto the hydrazine salt solution, to which 10 cc. of 4 N hydrochloric acid had been added. The temperature of the reacting solution was maintained a t 100' C. by means of a water bath, and the evolution flask was subjected to continuous mechanical agitation. Control experiments with water and permanganate solution were performed under com arable conditions in order t o evaluate the correction to be appied for gases dissolved in the reagents.
I500 1475
1450 I425
1400
2 2 1.375
-.
2
~ " 1 30 5
3 & 1325
TABLE I. INFLUENCE OF HYDROCHLORIC ACIDUPON TITRATION OF HYDRAZINE SALTWITH PERMANGANATE
1300
UP1
(0.6gri N )
KMIlO4
4 N HCl (0.08715N ) Cc. Av. cc.
EXPT.
1275
RATIOOB NH:
N n H 4 TO
Av. cc.
1.250
HYDRAZINI SULFATE (0.1628 GRAM OF NzH4.HzSO&)
A.
1-3 4 5 6-9 10-14 15-17 18 19-21 22-24
5 5 5 8 10
20 20 30 40 B.
RATIOO F E Q U I V . TO N z H 4 TO 0 NHa
27.35 27.32 27.37 28.78 29.31 29.89
1.050 1.051 1.049 0.997 0.980 0.961
29:iz 29.77
0 986 0.964
:
16.60 16.59
1.466 1.467
1e:io
1:493 1.533 1.603 1.608 1.532 1.569
15.88 15.30 15.25 15.89 15.51
1.225
10 20 30
29.61 29.96 29.50
0.975 0.965 0.979
15.82 15.51 15.80
1.548 1.578 1.550
In experiments 37 to 41 the nitrogen evolved by the titration of hydrazine sulfate with permanganate in hydrochloric acid solution was measured as described above. For 0.1628 gram of NzH4.H80dthe corrected volumes of nitrogen evolved were 18.38, 18.03, 18.44, 17.82, and 18.23 cc. These correspond to an average of 8.11 X moles of nitrogen which gives an average value of 1.543 for the ratio of hydrazine to nitrogen. TABLE11. INFLUENCE OF INITIAL CONCENTRATION OF HYDRAZINE UPON ITS REACTION WITH PERMANGANATE IN HYDROCHLORIC ACIDSOLUTION EXPT. 42,43 44, 45 46, 47 48,49 50,51 52,53 54,55 56,57 58, 59 60
NnH4
KMnOa (0.0900N )
RATIOO F NZH4 TO 0
Moles X 10-4 1.981 3.954 5.990 7.964 10.325 11.90 15.90 19.87 23.90 31.72
An. C C . 4.62 9.17 13.62 17.82 22.97 26.32 34.63 42.62 50.16 65.69
0.954 0.958 0.977 0.993 0.999 1.005 1.021 1.036 1.058 1.073
Ammonia was invariably formed as a product of the oxidation of hydrazine by permanganate under the conditions prevailing in the experiments recorded in Tables I, 11, and 111. In order to investigate the reaction under conditions more favorable to complete oxidation of the hydrazine to nitrogen and water, measured volumes of standard hydrazine
1 10
1 20 b n + T Moles/cc.r
1 30
1
1
IOd
FIGURE 2. INFLUENCE OF MANGANOUS IONON THE RATIOOF HYDRAZINE TO AMMONIA
PYDRAZINE HYDROCHLORIDI (0.1SZZ QRAM OF NzH44HCl)
25-28 29-32 33-36
I
1200
oxide drop by drop until the color was discharged. After the solution had been boiled to remove any excess of peroxide, it was made alkaline, and the ammonia was distilled into standard acid. TABLE111. INFLUENCE OF MANGANOUS IONUPON REACTION BETWBEN HYDRAZINE SALTAND PERMANGANATE IN HYDROCHLORIC ACIDSOLUTION TTnl
nu
KMnO4
EXPT. A.
IMn++l (0.0900N ) Mbles/cc.' x 10-6 cc. '
cc
.
HYDRAZINB HYDROCHLORIDE (0.1060 GRAM OF NzH49HCl)
61 62 63 64 65 66 67 68
1.37 3.42 6.15 11.0 15.0 23.0 30.5 33.8 B.
69 70 71 72 73 74 75 76 77 .. 78 79 80 81
(0.05682N ) RATIOOF RATIOOF EQUIV. TO NzH1 TO N z H 4 TO 0 NHa NHs
22.69 22.29 21.84 20.42 20.02 19.44 19.11 19.07
0.979 0.997 1.017 1.088 1.109 1.143 1.163 1.166
11.67 12.27 12.32 12.84
1.510 1.435 1.429 1.372
ii:i3 13.57 13.73
1:iio 1.298 1.282
HYDRAZINE SULFATB (0.1844 GRAM OF N z H c H ~ 8 0 4 )
1.35 1.58 3.21 3.28 5.95 6.27 12.3 13.0 19.5 19.7 32.1 32.5 37.7
22.61 22.55 22.09 21.94 21.39 21.29 20.54 20.34 19.37 19.25 18.55 18.40 18.29
1.015 1.017 1.039 1.046 1.072 1.078 1.117 1.129 1.177 1.193 1.237 1.247 1.254
12.00 12.12 12.45 12.58 12.78 12.78 13.37 13.58 13.83 13.99 14.58 14.67
...
1.515 1.499 1.481 1.444 1.422 1.422 1.360 1.338 1.314 1.298 1.246 1.238
..
In sulfuric acid solution about 14 moles of ammonia were obtained from 19 of hydrazine. Control experiments with
INDUSTRIAL AND ENGINEERING CHEMISTRY
56
weighed amounts of ammonium sulfate in place of hydrazine sulfate showed that 99 per cent of the ammonia could be recovered. In hydrochloric acid solution, however, no ammonia was obtained. Controls with ammonium chloride yielded only 10 per cent of the ammonia taken. These r e sults show that the ammonia was oxidized by free chlorine, but not to any appreciable extent by the permanganate.
Vol. 7, No. 1
Values of the ratio of hydrazine to ammonia as experimentally determined agree closely with the values as calculated from the empirical equation R=- 3 4
2
7
in which R is the ratio between moles of hydrazine and moles of ammonia, and R' the ratio between moles of hydrazine and atoms of oxygen. TABLE IV. COMPARISON OF OBSERVED AND CALCULATED OF THE RATIO OF MOLES OF N2HnTO MOLES OF NHs VALUES EXPTB.
0.955
asx,
t
R OBEERVED R CALCULATED 1.466 1.432 1.493 1.505 1.533 1.531 1.605 1.564 1.532 1.521 1.569 1.561 1.548 1.540 1.578 1.560 1.550 1.531
OP io
N&
Mdas
20
X
Ire
30
FIGURE3. INFLUENCEOF HYDRAZINE CONCENTRATION ON TBE RATIO OF HYDRAZINE TO OXYGEN '
No'evidence of the formation of hydronitric acid during these experiments or during the titrations was obtained from tests made under various conditions during the course of the work. DISCUSSION OF RESULTS From Table I it is clear that the reaction between either hydrazine sulfate or hydrazine hydrochloride and potassium permanganate in hydrochloric acid solution does not result in complete oxidation of the hydrazine to nitrogen and water as stated by Kolthoff. Under the prevailing conditions the ratios of moles of hydrazine oxidized to (a) atoms of oxygen required, (b) moles of ammonia formed, and (c) moles of nitrogen liberated are in fair agreement with the equation 3N2Hi
R' 1.050 0.997 0.980 0.961 0.986 0.964 0.975 0.965 0.979
1-5 6-9 10-14 15-18 19-21 22-24 25-28 29-32 33-36
+ 3 0 +2"s
+ 2N2 + 3HzO
which may be considered to express the resultant of two concurrent reactions of the complex delectronator (4) potassium permanganate: (1) complete di-delectronation of hydrazine bylpermanganate ion N2H4 2 0 +Nz 2HzO
+
+
and (2) incomplete mono-delectronation of hydrazine by manganic ion (3)
+
2NzH4 2Mn++++NI
+ 2"s
+ 2Mn++ f 2H+
or, in terms of oxygen used 2NzH4
+ 0 +Nz + 2"s
+ H20
Under the conditions prevailing in experiments 1 to 41 approximately one-third of the hydrazine undergoes didelectronation and two-thirds mono-delectronation. In experiments 6 to 9 the results approach exact conformity with this ratio. On the other hand, by increasing the concentration of manganous ion (and thereby favoring increase in concentration of manganic ion) either directly by adding 0.1 N manganous sulfate (Table 111,Figures 1 and 2), or indirectly by increasing the concentration of hydrazine sulfate (Table 11, Figure 3), the relative amount of mono-delectronation is increased, as evidenced by increase in the ratio of hydrazine to oxygen and decrease in the ratio of hydrazine t o ammonia.
Values of R observed and calculated for experiments 61 to 80 are in similar accord. I n experiments 4, 5, and 18 the effect of variations in temperature upon the ratios between hydrazine and oxygen and hydrazine and ammonia was found to be negligible. In the last of these experiments, however, a liberation of chlorine began after the hydrazine had been completely oxidized, and for this reason a satisfactory end point could not be obtained. In view chiefly of the influence of the concentration of hydrazine upon the course of its reaction with permanganate, the Petersen method as modified by Kolthoff cannot be recommended for practical use either in the determination of hydrazine or in the standardization of permanganate solutions. Under carefully controlled conditions consistent but not necessarily correct results may be obtained. SUMMARY The oxidation of hydrazine by permanganate in hydrochloric acid solution has been shown to involve two concurrent reactions : (1) complete di-delectronation of hydrazine by permanganate ions N2Hd
+ 2 0 +N2 + 2HzO
and (2) incomplete mono-delectronation of hydrazine by manganic ions, formed by interaction of manganous and permanganate ions 2NzH4
+ 0 +Na + 2"s
+ Hz0
The relative amount of hydrazine oxidized in these two ways depends on various factors, the influence of which has been investigated. The reaction between hydrazine sulfate and potassium permanganate in hydrochloric acid solution is therefore adjudged unsatisfactory for use either in the determination of hydrazine or in the standardization of permanganate solutions. LITERATURE CITED (1) Browne and Shetterly, J. Am. Chem. SOC.,31, 221 (1909). (2) Ibid., 31, 783-99 (1909). (3) Guy, Rosenberg, and Bray, Ibid., 46, 1796 (1924). (4) Kirk and Browne, Ibid., 50, 337 (1928). (This article, No. VI11 in the current series, contains references to all earlier articles of the series.) ( 5 ) Kolthoff, I. M., Ibid., 46, 2009-16 (1924). (6) Kolthoff, I. M., and Menzel, H., "Volumetric Analysis," tr. by N. H. Furman, p. 301, New York, John Wiley & Sons, 1929. (7) Kurtenacker and Wagner, 2. anorg. allgem. Chem., 120, 261 ( 1 922).
January 15, 1935
ANALYTICAL EDITION
(8) Medri, Gazz. chim. ital., 36, I , 373 (1906); Chem. Zentr., 1906, 11, 459. (9) Petersen,2.anorg. Chem., 5 , l (1893). (10) Purgotti, Ann. ist. super. agrar. Portici, [31, 3, 47 (1929). (11) Raschig, "Schwefel- und Stickstoffstudien,"P. 188, Leiprig, Verlag Chemie, 1924. (12) Roberto and Roncali, Ind. chim., 6, 178 (1904) ; Chem Zentr., 1904, 11, 616.
57
(13) Sabanejeff, 2. anorg. Chem., 20, 21 (1899). (14) Stelling, Svensk Kem. Tid., 45, 3 (1933). RECEIVED October 19, 1934. Baaed upon a part of the thesie presented to the Faculty of the Graduate School of Cornell University by A. G. Houpt in partial fulfilment of the requirements for the degree of Master of Chemistry. The work of Doctor Sherk has been supported through the personal generosity of Haymo V. Pfister of the Pfister Chemical Company.
Standardization of Potassium Dichromate
vv
HOBART H. WILLARDAND PHILENA YOUNG, University of Michigan, Ann Arbor, Mich.
HEN p o t a s s i u n i dithe reaction between arsenious The oxidaiion value of solid potassium dichromate is used as a acid and potassium dichromate chromate or of solutions of the reagent may be in t h e p r e s e n c e of osmium volumetric oxidizing determined very accurately against the primary tetroxide, which appears to be a agent, its o x i d a t i o n v a l u e is standard, arsenious oxide, by treatment of checked most satisfactorily in catalyst for arsenic and similar arsenious acid in a sulfuric acid solution with one of two ways. The c. P. salt materials, would be sufficiently may be recrystallized a number rapid for a direct titration. A less than its equivalent of dichromate, and titraof times from water and the final procedure for comparing potastion of the excess reducing agent, in the presence product after drying used as a sium dichromate against such a of osmium tetroxide as catalyst and o-phenprimary standard, or solutions of primary standard is of imporanthroline ferrous complex as indicator, either the approximate normality detance because the present methwith ceric sulfate or with potassium permansired may be standardized, often ods of checking the normality of by a procedure similar t o that dichromate solutions are too inganate. The titration may also be made potenfollowed in an actual analysis. direct-such as t i t r a t i o n with tiometrically with potassium bromate in a hydroThere are objections to both of ferrous sulfate, which has been chloric acid solution. S i x samples of reagent checked against permanganate these methods. Suitable subquality and chemically pure potassium dichrostandardized a g a i n s t B u r e a u stances of definitely known conmate made by different Jirms showed practically tent against which to standardize of S t a n d a r d s s o d i u m oxalate. This particular method, a dichromate solution are not 100 per cent purity. always available, and e i t h e r which is widely used, has been p r o c e d u r e f o r preparing shown to be undesirable as the - - a standard solution may be time-consuming. Precise and titration with ferrous sulfate is influenced by acidity, volume more direct methods for testing the purity of a supply of of solution, and concentration of dichromate (1, 3). potassium dichromate or for the standardization of solutions, In the titration of arsenious acid with potassium dichromate preferably against primary reduction standards, are needed. in a sulfuric acid solution containing two or three drops of There has been considerable doubt whether the salt as pur- 0.01 M osmium tetroxide as catalyst and two drops of chased is sufficiently pure for a primary standard. The 0.025 M o-phenanthroline ferrous complex as indicator, present study shows that such standardizations may be fairly concordant results could be obtained at 50" C. with made accurately and rapidly &gainst arsenious oxide. solutions containing from 40 to 60 cc. of 10 N sulfuric acid per 100 cc. of solution a t the beginning of a titration. HowEXPERIMENTAL METHODS AND RESULTS ever, as much as 0.10 cc. of 0.1 N dichromate in excess was REAGENTS AND SOLUTIONS. Bureau of Standards arsenious ordinarily used, an indication that the reaction was too slow oxide was dissolved in sodium hydroxide, and sufficient sulfuric acid added to react with the latter, followed by 10 grams of for a visual end point. sodium bicarbonate. TABLEI. POTENTIOMETRIC TITRATION OF ARSENIOUS ACID The potassium dichromate was reagent quality material as WITH POTASSIUM DICHROMATE received. (Osmium tetroxide a6 catalyst) The ceric sulfate solutions were from large supplies made by NORMALITY OF dissolving ceric ammonium sulfate in 0.5 M sulfuric acid. They 0.01 M Os04 10 N &SO4 KzCraOr had been standardized against Bureau of Standards sodium Drops CC. oxalate (4). 4 30 0.1003 4 40 0.1002 A 0.025 M solution of o-phenanthroline ferrous complex 4 50 0.1004 (CJf8N2.H20)aFe,was prepared by dissolving the correct 4 60 0.1004 amount of o-Dhenanthroline in a 0.025 M aaueous solution of 0 0.00985 40 9 40 0.1006 ferrous sulfate'. 12 40 0.1005 The 0.01 M solution of the catalyst was made by dissolving 1 gram of osmium tetroxide (sometimes called perosmic acid) in The potentiometric method was tried because it permitted 400 cc. of 0.1 N sulfuric acid. the use of a higher temperature with the possibility thereby REACTION BETWEEN ARSENIOUSOXIDE AND POTASSIUM of increasing the velocity of the reaction. Twenty-cubic DICHROMATE. It has been stated recently by Gleu (2) that centimeter portions of 0.1009 N sodium arsenite were diluted the titration of arsenious acid either with ceric sulfate or with water and the volume of 10 N sulfuric acid specified with potassium permanganate, using o-phenanthroline ferrous in Table I to 100 cc. A few drops of 0.01 M osmium tetroxide complex as oxidation-reduction indicator, is satisfactory mere added and the arsenic was titrated potentiometrically a t room temperature in a sulfuric acid solution, provided a t 70" to 80" C. with 0.1004 N potassium dichromate. that a very small amount of osmium tetroxide is present as The results obtained are shown in Table I. The reaction is catalyst. His data are inadequate, however, to show the quantitative, but the titration much too slow to be of any precision to be expected. The question arises as t o whether practical value. With no catalyst too much of the dichro-
