Stability of Concent Hydrogen Peroxide J
WALTEK C. SCHUJPB Massachusetts I n s t i t u t e of Technology, Cambridge, Mass.
Stability characteristics of concentrated hydrogen peroxide solutions have been studied a t 50" to 100" C. over the range 70 to 90 weight 70hydrogen peroxide, in both the absence and presence of stabilizers; 50' C. proved satisfactory for the gas evolution method used to determine rate of decomposition of peroxide samples. Previous reports of high stability for very pure hydrogen peroxide at concentrations of 85 to 90y' were codwmed. The temperature coefficient of the decomposition rate of unstabilized peroxide solutions was determined for the intervals 50' to 60' C. and 60" to 10" C. as 2.2 * 0.1, which is believed t o be approximately applicable over any 10" range between 80' and 100' 6. Methods of preparing the surface of
Pyrex containers for use with concentrated hydrogen peroxide were developed which make possible obtaining reproducible measurements of the decomposition rate. Pure aluminum and pure cast tin were found to he suitable materials of construction. Vigorous agitation was without appreciable effect on the measured rate of decomposition of 85% unstabilized hydrogen peroxide. The importance of pH on rate of decomposition was emphasized by measurements with and without added catalytic ions. Data were secured on the catalytic efyect of known concentrations of ferric or cupric ion in 85 t o 90% hydrogen peroxide a t 50' C. and on the required concentrations of Stabilizers to restrain such catalytic action.
A
hydrogen peroxide, are consistent with the present author's data. They point out that no attainable pressure has any noticcablr effect upon the rate of decomposition, and that probably the only known material which itself actually increases the stability of hydrogen peroxide is the hydrogen ion, as indicated by the equation: HOOH Hr OOH-, for which Joyner ( 8 ) gave the They point out also equilibrium constant, IC = 2.4 X that the effect of added stabilizers is to nullify the influence of positive catalysts which may be present in the solution. A list of representative metal ions added in small proportions was reported to have the effects shown in Table I1 upon the decomposition of 90% hydrogen peroxide. Table I1 indicates that aluminum and tin are without effect, and that iron, copper, and chromium are active decomposition catalysts. I n the present investigation, i t was considered advisable to study the stability of the purest attainable unstabilized hydrogen peroxide, more particularly in the higher ranges of concentration (70 to 90 weight %), over a range of temperature and in containers of various materials of construction. With data for the unstabilized peroxide established, the influence o i added stabi-
PPLICATION of the decomposition of concentrated solutions of hydrogen peroxide, with or without added fuels, as a source of power in propelling various types of missiles, aircraft, or submerged craft during the past war brought with i t problems of storage of large quantities of this material under varying climatic conditions. Because it was well known that the decomposition process was directly dependent, among other factors, upon the presence of catalytic substances dissolved or suspended in the liquid and upon the nature of the m+lls of the containing vessel, the need for the addition of stabilizers to offset the effect of such catalytic influences was recognized and as a result of considerable experimental work carried out during the past decade, effective stabilizers have been devised which permit storage for many months with little change in the peroxide concentration. The stability of pure hydrogen peroxide and the accelerating effect upon its decomposition caused by catalytic metal ions and the deactivation of catalysts by stabilizers have been recognized and described (10). As early as 1894, Wolffenstein ( $ 2 ) pointed out that a solution of hydrogen peroxide, free from alkali or traces of heavy metal compounds, possesses a considerable degree of stability and that it may be distilled and concentrated without appreciable loss by decomposition. Quantitative measurements of the stability of concentrated hydrogen peroxide solutions could be significant only in the event that the peroxide samples employed were scrupulously freed from contaminants, both dissolved or suspended. The careful investigations of Maass and his associates ( 7 ) yielded the most reliable data which have been available up to recent times. Regnault and Le Noir de Carlan ( 9 ) also prepared hydrogen peroxide of nearly 100% concentration by fractional distillation of commercial "100 to 130 volume" material in Pyrex apparatus, following concentrative evaporation to approximately 60% HzOz. They commented upon the stability of solutions of the pure peroxide. Very recently Shanley and Greenspan ($1) of the Buffalo Electro-Chemical Company have published the results of measurements carried out concurrently with those presented in this paper. The values of Table 'I, given by these authors for 90%
-+-
TABLE I. DECOWO~ITJOS O F HYDROGEN PEROXIDE Temperature,
30 66 100
140
Decompositinti
C
'7 per year
:.p
per week 2% per 24 hours Rapid decomposition \;rich boiling
OF ADDED IONS TABLE 11. EFFECT
Added Inn None
A1 Sn
Zn Fe
cu cu Cr
992
Amount Added. Mg./L.
7 Osiginal .kcti& o Lost in 24 Hours a t 100' C.
...
2
10 10
10
2 2
10
1.0
15
0.01 0.1 0.1
96
24
85
May 1949
INDUSTRIAL AND ENGINEERING CHEMISTRY
h e r s was t o be determined, in both the absence and presence of known amounts of catalytic substances. As the attempts to prepare the purest possible aqueous solutions of hydrogen peroxide progressed, it became clear that the decomposition of the pure peroxide of itself is exceedingly slow over the entire range of concentration studied (up to more than 90%) and that when sufficiently high purity was secured, very little further reduction in the decomposition rate by addition of stabilizers could be effected, and then only by addition of rather minute concentrations of the stabilizer (of the order of a few parts per million). On practical grounds it was realized that the likelihood of contamination of stored peroxide solutions could not be completely eliminated, and therefore a knowledge of the restraining power of known concentrations of stabilizers present in the peroxide upon known concentrations of dissolved catalytic impurities (or known amounts of undissolved impurities) was necessary. I n addition t o the temperature and concentration of the s o h tion and the nature of the containing vessel, other factors that bear upon the observed decomposition rate were considered of importance, including the pH of the solution, the influence of the surface-to-volume ratio of the containing vessel, and the conceivable influence of agitation of the solution. Furthermore, an adequate knowledge of the mechanism of the stabilization processes in the presence of known impurities is of primary interest and much thought and effort have been given to the attempt t o secure a completely satisfactory explanation of the phenomena here involved. [The results of certain studies of the mechanism of catalysis of the decomposition of hydrogen peroxide solutions, carried out by allied groups of investigators, have recently been published ( I , IO).]
,EXPERIMENTAL PREPARATION O F SOLUTIONS
Most of the experimental work was carried out with peroxide samples obtained by the fractional distillation of 90% unstabilized peroxide (from Buffalo Electro-Chemical Company, socalled Becco). The quality of the distillate thus obtained could not be further improved by subjecting it to fractional freezing. The high purity of the product was indicated both by its very low decomposition rate and by its low specific conductivity. Some measurements were carried out with the 90% Becco peroxide and with 70% Du Pont peroxide (stabilized), as received; but for the greater part of the work described, redistillation of the commercial material preceded all sampling for these stability measurements. CONTAINING VESSELS
As it has long been realized that the rate of decomposition of peroxide solutions is greatly dependent upon the nature and the condition of the surface of the container, a series of tests was made with metal, glass, and quartz vessels to determine their suitability for the construction of storage containers, and considerable attention was given to the matter of cleaning the contact surface so as to obtain a minimum of “wall effect.” Pyrex, fresh from stock, and subjected to a careful cleaning as described below, was used for most of the routine measurements. However, pure aluminum (99.6%or better, with minimum copper content), the surface of which had been anodized in an oxalic acid electrolyte, and the pores further sealed by soaking in boiling water, was found to give equally satisfactory results. (The aluminum may also be treated by degreasing by a short soaking in dilute caustic alkali solution, followed by a prolonged pickling in 4 t o 5% sulfuric acid, and rinsing until the surface assumes a uniform, dull appearance.) The use of clear fused quartz in place of Pyrex did not lead t o noticeable improvement and the varieties of stainless steel tested showed behavior inferior t o that of Pyrex or pure aluminum, as far as the decomposition of the peroxide contained therein was concerned. [Reports have been
993
received of satisfactory performance by certain varieties of stainless steels (303, 304, 309, etc.) as containers for concentrated hydrogen peroxide, after the contact surfaces of the containers had been conditioned by a suitable pickling process. The reported decomposition rates in the case of stainless steel containers, however, are not so low as those observed in the case of pure aluminum or of Pyrex vessels.] Although aluminum vessels of sufficiently high purity and with contact surfacm properly conditioned are found to be well adapted to the storage of concentrated hydrogen peroxide, and are commonly so used, in the present work most of the quantitative measurements were made on solutions contained in Pyrex flasks, because of greater convenience in observing the solutions during the period of experiment and the ability to obtain reproducibility in memurements when the glass surfaces were subjected to a standardized cleaning process. Preparatory cleaning of the glass surfaces before use of the flasks in stability measurements was a subject of much concern a t the outset, because reproducibility of measurements was known to be greatly dependent upon the effectivenessof the preliminary treatment. For example, cleaning solution composed of concentrated sulfuric acid containing chromic acid cannot be used for this purpose, as the minute quantity of chromic acid left adhering t o the walls after thorough washing is still able noticeably t o catalyze the rate of decomposition of the peroxide. After a number of trials with various cleansing procedures, some of which were more elaborate than that described below, the following method was adopted for routine treatment of the glass apparatus. The Pyrex flasks (250-ml. volumetric, with the neck replaced by the inner part of a 14/35 standard-taper ground joint) used in determining the rate of decomposition of peroxide solutions were taken from wholly new stock and were not used again unless with identical contents. The flasks were allowed to stand for about 4 days filled with fuming nitric acid (or else heated with the acid for 1 day a t 60” C.). They were then rinsed thoroughly with distilled water and filled with concentrated, distilled peroxide and again allowed to stand for a matter of days or weeks (or else heated a t 60”C. for 1day) until just prior to use in measurements, Each flask was then emptied and filled with conductivity water and heated by immersion in a bath of boiling water for about 8 hours. When emptied and freed as far as possible from adhering water by shaking, the flask was ready for use, without the inner walls being allowed t o dry completely. When flasks were allowed to stand, they were kept closed by means of small, inverted beakers which themselves had been cleaned by a similar process. All flasks and beakers were protected as far as possible by suitable coverings from contamination from atmospheric fumes and dust. When flasks, thus cleaned, were filled with a sample of 85 to 90 % unstabilized hydrogen peroxide and warmed in a thermostat a t 50” C., in the great majority of cases decomposition proceeded in a slow and regular manner without the appearance of gas “streamers” originating from minute points on the wall of thevessel. If the flask had not been well cleaned or, occasionally, in spite of careful cleaning, such streams of minute bubbles formed a t points on the wall’s surface-if this occurred in any considerable volume, the results of such determinations would be valueless. An occasional, very minute stream could be neglected in cases where the decomposition rate to be measured was not very small. By adhering strictly to the routine method of procedure described above, it was found possible to obtain a satisfactory degree of reproducibility in the measurements of decomposition rate which are presented below in detail. PRELIMINARY TESTS ON SAMPLES
Relative Purity. A comparison of the relative purity of various batches of redistilled, concentrated hydrogen peroxide was made possible by considering the results of such measurements as the following: (1) the specific conductivity of the sample at 25” C.; (2) the specific conductivity of the water obtained by complete decomposition of the sample, in the presence of platinum; (3) the p H of the mixture resulting from the dilution of the sample tenfold with conductivity water, as measured on a
INDUSTRIAL AND ENGINEERING CHEMISTRY
994 = a
Vol. 41, No. 5
of 2 X l o p 6ohm-' em.-' a t 25" C.. and Calvert ( 8 ) previously had determined the specific conductivity of a 4.5% solution as 2.89 X ohm-' cm.-'
5
i
86
I n the present work the specific conductivity of some of the purest sampIes of concentrated, unstabiE4 lized, redistilled hydrogen peroxide was determined as z a function of its concentration, by progressively dilut0" 4 iiig the sample with triply distilled conductivity water, u2 the specific conductivity of which was about 5 x 2?i lo-' ohm-' cm.-' at 25" C. An alternative method, 9 0 20 30 40 5 60 ao ," 3o suggested as a check upon the purity of hydrogen perC 3 F C E N T R A T O N OF H2O2 I 4 1 F C S i E N ' oxide samples, in which the concentrated peroxide &ras gradually decomposed at room temperature by use ~i~~~~ 1. Specific Conductivity of Hydrogen Peroxide Solutions of sheet platinum in a Pyrex or fused clear auarta at 25" C. flask (sealed to a reflux condenser) and the conductivity of the water obtained thereby measured, was not found to be so satisfactory as the dilution method; the Beckman pH meter; (4) qualitative spectroscopic analysis of time required for the decomposition t o become complete was the solid residue obtained by the decomposition of 1 liter oi the so long that inevitable leaching of materials from the glass or sample and evaporation to dryness: and ( 5 ) the rate of decomquartz walls of the vessel occurred, resulting in high conductivity position a t 50" C., by the method described below. (In some readings. cases, in addition to the pH measurements, the free acid content of the sample was determined by tenfold diiution of the sample Thus, a sample of redistilled Becco peroxide, of 98.5 weight with distilled w&ter, followed by titration with 0.01 S sodium % hydrogen peroxide, exhibiting a specific conductivity of 8.2 x hydroxide, using a mixture of methyl red and methylene blue as ohm-' cm.-' a t 25' C., on dilution gave the results indicated indicator.) in Figure 1. The measurements Rere made in a Pyrex cell containing electrodes fashioned of cast pure tin, the plane faces of Nitrogen. I n some of the batches of redistilled peroxide the which were about 25 mm. in diameter, spaced about 4 mm. apart. content of nitrogen, present as ammonla or ammonium ion, was The cell constant was 0.08037, as determined vith standard determined. potassium chloride solution, containing 0.7476 gram per 1000 A sample (about 70 ml.) of distillate was decomposed (with grams of water. The specific conductivity of this soIution at the aid of a piece of sheet platinum), a 25-ml. sample of the 25" C. was taken as 0.001410 ohm-' cm.-l ( 6 ) . resulting water was diluted with an equal volume of ammoniaThe alternating current measurements were carried out at free water, and 2 drops of 10% copper sulfate solution and 1 ml. 1000 cycles with the aid of a General Radio Company impedance of 50% sodium hydroxide were added. After the precipitate bridge, Type 650-A. The cell was placed inside a small oil had settled for 24 hours, a 1-ml. pipetful of the supernatant bath, which in turn n-as immersed in a constant temperature liquid, diluted to 50 ml. 6 a s nessler&d, and the color-was comwater bath a t 25.00' * 0.02" C. pared with tubes containin a series of standards. The content The curve is symmetrical and the maximum value, 5.8 X of ammoniacal nitrogen &us determined in 90% redistilled 10-6 ohm-' cm.-l, occurs at 50% hydrogen peroxide. Although peroxide was found not to exceed the lowest detectable value, the extrapolation of the curve to zero concentration gives a 0.5 p.p.m. conductivity value somewhat higher than that of pure water itself, the deviation therefrom is not large and the reliability of Other determinations, carried out on the 90% unstabilized the data as a whole is indicated. peroxide, as received, included nitrate nitrogen, sulfate, and phosphate. These three impurities were found to be of t,he order The author believes that these data will not be subject to of 1to 2 p.p.m. each. serious revision if a possible closer approach to 100% puritj of the peroxide should be attained a t some future date. Judged by Nitrate nitrogen, determined by the micro-Kjeldahl technique, these results, the value reported by Cuthbertson and lIaass for for the 90% gave 1.8 p.p.m. of nitrogen, expressed as "3, unst,abilized sample. 100% hydrogen peroxide would appear to require revision to not Suljute, 1.5 p.p.m., was determined by decomposition and conover 0.5 X 10-6 ohm-' cm-l centration of 2000 grams of 90% hydrogen peroxide to a final A summary of other data obtained in preliminary tests upon volume of 300 ml., followed by turbidimetric determination of representative samples of 90% Becco unstabilized and 70% Du the sulfate as barium sulfate. Phosphate was determined colorimetrically, by the method of Pont (stabilized) hydrogen peroxide is given in Table 111. Fiske and Subbarow (4) applied to the decomposed peroxide. A value of approximately 2 p.p.m. of phosphate was obtained MEASUREMENT OF STABILITY O F with the aid of a Klett-Summerson colorimeter for the 90% HYDROGEN PEROXIDE SOLUTIONS unstabilized material and 45 to 50 p.p.m. of phosphate (expressed as PnO,----) for a 70% stabilized sample. For some purposes the stability of concentrated hydrogen peroxide solutions has been evaluated by analysis of samples of The variation in pH of redistilled peroxide with concentration the solution taken a t intervals over known periods of time; the was examined over the range 0 to 85% hydrogen peroxide. As loss in percentage of hydrogen peroxide contained is determined pH measurements were usually carried out at a tenfold dilution by titration with an acid solution of potassium permanganate, (conductivity water being added), this range corresponded acwhich reacts as follows: tually to 0. to 8.5y0hydrogen peroxide in the sample as measured. The pH measured in this way changed but slightly: An increase 2MnO; 5H202 6 H + +2Mn++ i8H2O -t 502 was observed from 5.35 for the most concentrated (84%) t,o 5.67 for 0%. pH measurement,s on undiluted samples n-ere not taken, Long-range tests at room temperature are conveniently carried as a rule, as their actual significance is doubtful., hIillivolt readout in this way-for example, a sample of 90% peroxide contained ings were recorded in such cases, however, as there is reason to in a drum may be found 6 months later to show a content of 89% believe that the potentials thus registered with the help of the hydrogen peroxide. However, such measurements are subject to glass electrode bear a direct relationship to the change in such errors as losses of water (and some peroxide) by evaporation -log CH over a considerable range. and gain of water content through absorption from the surroundSpecific Conductivity of Hydrogen Peroxide Solutions. The ing moist air. For laboratory testing the gas evolution method, specific conductivity of pure hydrogen peroxide was measured by carried out a t such a temperature as will provide a satisfactory Cuthbertson and Maass (3) who reported an upper limiting value
2
c
+
f
+
May 1949
INDUSTRIAL AND ENGINEERING CHEMISTRY TABLE 111. SUMMARY OF DATA 90% Becco (as Recd.) 90.1 5.0
70% Du Pont (as R,epd.), Redistilled Stabilized Beceoa 70.9 95.7
Concentration, % HzOz 4.8b 4.7 p H (diluted tenfold) Specific conductivity, (ohm-1) (cm.-1) 9 3 X 10-6 3 . 2 X 1 0 - 9 1 . 1 X 10-6 Decomposition rate a t 50° C . , 0.0027 0.00034 0.0010 Yo per hour Although this sample was not derived from the same hatch of peroxide as t h a t shown in the first column the d a t a given indicate in a general way the degree of change brought abodt on distillation. b If readings were taken on the undiluted peroxide, a n "apparent pH" of 1.4 was obtained. The significance of p H readings in such concentrated peroxide solutions is very doubtful. For comparative purposes the use of "apparent pH" values, or better, of the millivolt readings indicated by t h e p H meter, may be justifiable, regardless of the lack of theoretical significance of such "pH" readings. Q
volume of oxygen in a convenient period of time, has been preferred. Gas Evolution Method. In the attempt to evaluate the rate of decomposition of concentrated, unstabilized peroxide, samples were secured of special "cuts" obtained in the distillation of concentrated peroxide solutions, taken under conditions that would lead t o the purest obtainable product. The samples were stored in aluminum drums, the contact surfaces of which had become "conditioned" by prolonged standing filled with concentrated
peroxide. When such selected materjal was subjected to still further purification b y redistillation in Pyrex apparatus, in a few cases with the addition of a minimum quantity of sodium hydroxide solution t o hold back any possible volatile acidic impurities, a high degree of purity was attained. With commercial peroxide samples of still higher purity no alkali was found t o be required prior to redistillation. From such samples the measurements of specific conductivity as a function of concentration, described above, were obtained. The decomposition rate of these purest samples of unstahilized distilled peroxide, measured as described below, was found not to exceed 0.00270 per hour a t 50' C., a minimum value of about 0.0008% being observed. (Addition of 0.5 p.p.m. of sodium stannate trihydrate as a stabilizer to such samples would reduce the rate a t 50" C. from about 0.002 to about 0.0004~0 per hour.) Although the rate of decomposition of hydrogen peroxide solutions was studied a t temperatures ranging from 50" to 100" C., in order t o establish the temperature coefficient of the reaction, most of the rate measurements were carried out at 50' C., inasmuch as below this temperature the rate is inconveniently slow if the sample is of reasonable purity, whereas at 100" C. i t was difficult t o obtain satisfactory reproducibility of measurements, (For more rapid evaluation in routine testingof thestabilityof samples of concentrated commercial peroxide, an accelerated 100" C. test extending for 16 or 24 hours nevertheless has been much
Leads to
Figure 2. 1.
2.
3. 4. 5. 6.
7. 8. 9. 10.
995
111
I
Apparatus for Gas Evolution Method
Wooden box Glass wool Pyrex jar, 12-inch diameter 10-ml. microburet with water jacket Pyrex capillary tubing 10/30 standard-taper ground-glass joints Manometer (water) 250-ml. flask with 14/35 standard-taper ground-glass joints Mercury leveling bulb Steel rod for support of leveling bulb
11. Rubber tubing 12. Stirrer motor
15E. Overflow
15d. Rubber stopper with vent
16.
Perforated copper plate for support of flash
17. Thermometer, 0' to loo0 C.
996
IN D U S T R I A E A N D E N G IN E E R, IN G C H E M I S T R Y
Vol. 41, No. 5
utilized. In this test a sample of several grams of peroxide solution, contained in a small bulb attached to a capillary tube, is maintained a t constant temperature for the desired period, and the change in composition is determined by titration with potassium permanganate.) The rate of decomposition was determined by a gas evolution method.
Samples of 70% stabilized ( P u Pont) peroxide in 10 months were reported to show a maximum decline of 0.6y0in hydrogen peroxide content, or 0.85% of the original concentration. I n laboratory tests on similar samples of stabilized 70% peroxide, the calculated decomposition rate a t 25" C. was found to bt, 0.00032 weight % per hour or 2.3% in 10 monthr. The 1one1 reported decomposition rate in the long-range test a t room temperature thus illustrates the influence of the surface-volume ratio on the observed decomposition rate of hydrogen peroxide.
The apparatus used consisted of 250-ml. Pyrex flasks, the necks of which were drawn down and sealed to the inner parts of 14/35 standard-taper ground joints, leading through capillarier extending from the thermostat bath to gas microburets. The burets contained mercury with a thin layer of water above i t to ensure saturation of the oxygen with water vapor a t the time of measurement. Six thermostat baths were used, in which the temperature of the water was controlled by conventional means at 50" * 0.05" C. for all routine measurements. The pressure of the confined gas and water vapor was brought to atmospheric by adjusting the mercury leveling tube until the levels indicated by a small auxiliary water manometer, teed into the line, were equal. Figure 2 shows the general arrangement of the apparatus employed.
Temperature Coefficient of Rate of Decomposition. The temperature coefficient, a, of the rate of decomposition of unstabilized hydrogen peroyide solutions, represented by the expression
By a series of trials it was found that rate measurements carried out during the first few hours were likely to show some lack of reproducibility, but if the apparatus was left to stand overnight, measurements throughout the following day were satisfactorily constant and reproducible. This procedure was therefore adopted as standard in all later work. In the 16-hour period of standing such conditions as attainment of equilibrium and relief of any large degree of supersaturation of the solutions with oxygen were established. I n the 50' C. measurements of stability, from the observed volume of oxygen (saturated with water vapor) set free from a known weight of the peroxide sample, the composition of which waa known by analysis (titration with standard permanganate solution in acid solution) of a small sample taken a t the start of the run, and from the observed temperature and pressure of the moist gas as collected, the number of moles of oxygen evolved could be calculated, and thus the per cent of decomposition of the hydrogen peroxide originally present in the sample could be found. The rate of decomposition, unless otherwise specifically mentioned in this paper, is expressed as the per cent of the quantity of pure hydrogen peroxide originally present which is decomposed per hour a t the temperature indicated. As the rates were, in general, small, the concentration of hydrogen peroxide in the sample did not alter significantly during the course of the measurements. Effect of Surface-to-Volume Ratio. I n experiments conducted a t 50' C. for the purpose of determining the effect of a change in the surface-to-volume ratio of the containers, the samples were contained, respectively, in the 250-ml. volumetric Pyrex flasks mentioned above, or in Pyrex test tubes containing several Pyrex rods with rounded ends, to increase the surface of contact. The tubes were connected by standard-taper ground joints to the iapillaries leading to the gas burets. I n this way, a change in surface-volume ratio from 1 to 10 cm. was achieved. This tenfold increase in the surface-volume ratio was observed to result in an approximately tenfold increase in the rate of decomposition, and this relationship was found to hold fairly constant over the entire range of concentration. From the results obtained in the study of the effect of the surface-volume ratio upon the rate of decomposition of hydrogen peroxide, i t becomes evident that for storage purposes over protracted periods of time, the larger the container-drums or tanks-the better, as it then becomes possible to attain notably low values of the surface-volume ratio In this connection it is interesting to compare the decomposition rates observed in 50" C. tests in the laboratory, using small glass flasks as the containing vessels, with the rates observed in large scale storage tests a t room temperature over a period of, say, 10 months.
( Y L )
k r = kT,a u a s determined for the 10" intervals 50" to BO" C. and 60" to 70' C., as well as over the interval 50' to 100" C. Conccntrations of the purest available hydrogen peroxide, 15 to 85% by weight, were employed. Over the ranges studied the value a = 2.2 * 0.1 was obtained-that is, for a 10" rise in temperature thr rate of decomposition of the peroxide was more than doubled; and for the rise from 50 to 100' C. the rate was increased in the order of fiftyfold. The measurement3 at 100" C. were found less reproducible than those carried out a t the lower temperatures, and as a result of these comparative measurements, 50" C. wab chosen as the most satisfactory temperature at which to conduct routine tests of the stability of hydrogen peroxide solutions. Effect of Agitation. I n view of the fact that escape of oxygeri brings about a certain degree of stirring of the liquid, it seemed unlikely that agitation would be found to influence materially the rate of decomposition of concentrated hyrogen peroxide, as has indeed been reported previously by other investigators. However, it was decided to check this observation by independent measurements. A rectangular thermostat bath was constructed in such a way that a small reaction flask, connected by a ground joint to a long, flexible, capillary glass tube leading to a microburet, could be subjected to shaking back and forth in the bath. The liquid contents were thus vigorously agitated-much more than would be experienced in any normal course of events in the transportation of peroxide containers. Measurements thus carried out were compared with results obtained when the reaction flask was not shaken, and the conclusion was reached that vigorous agitation had no appreciable effect upon the rate of decomposition of redistilled peroxide a t 50" C., as shown in Table IV. The fact that the rates of decomposition of the shaken samples were so nearly equal to those of the same samples, when left undisturbed for many hours, rules out the likelihood of results being low by reason of supersaturation of the liquid with dissolved oxygen. If any important amount of supersaturation had occurred in the static runs, which likewise extended over many hours, the results of the runs carried out on the same samples Over a similar period of time, with constant, vigorous shaking, mould inevitably have shown definitely higher readings than thosf' observed. The concordance of the rates on shaken and unshaken samples might also be interpreted as indicating that the degree of supersaturation attained a constant value in earh case. O
TABLEIV.
EFFECT OB AGITATION ON RATEOF DECOMPOSITION 85% HYDROGEN PEROXIDE AT 50' C.
OF
Sample No.
Rate of Decomposition (% per Hour! Unshaken Shaken 0.0069 0,0064 0,0064 0.0025 0.0025
0.0073 0.0068 0.0074 0.00ao
0.0023
INDUSTRIAL AND ENGINEERING CHEMISTRY
May 1949
RESULTS AND CONCLUSIONS STABILITY MEASUREMENTS ON REDISTILLED PEROXIDE
The expression employed throughout this work for the rate of decomposition is the rate of loss of active oxygen by decomposition of the hydrogen peroxide, expressed as a weight percentage of the total amount originally present-for example, weight per cent of hydrogen peroxide decomposed per hour at 50" C. For the purest obtainable, redistilled peroxide, when this decomposition rate was plotted as ordinates against hydrogen peroxide concentration as abscissas, a very flat curve with a slope of nearly zero was obtained, rising slightly in the direction of increased dilution. The presence of a maximum in this cur9e is doubtful, although some of the earlier data obtained indicated a possible flat maximum in the vicinity of 15 weight % hydrogen peroxide.
01
10
0
20
30 40 50 60 GONCENTRITION IWT PERCENT HzOd
70
80
I t should be borne in mind when considering the effect of addition of any ion upon the stability'of hydrogen peroxide that we are not able t o observe the effect of any single ion in a given experiment, but only the net effect of the cation and anion together. Thus the effect of variation in pH by addition of sulfuric acid cannot be ascribed at once to the hydrogen ion, unless there is reason t o believe that the effect of the sulfate ion simultaneously added is negligible in comparison with that of the hydrogen ion; or if it is known that, for example, equivalent amounts of hydrogen ion supplied by a variety of acids have essentially identical effects upon the rate of decomposition of hydrogen peroxide under like conditions.
A comparison of the effects of addition of sulfuric acid and of sodium sulfate to concentrated hydrogen peroxide is of interest. In such a test, small increments of sulfuric acid were added to 85% peroxide and were observed to result, first, in a shar decrease in the rate of decomposition, followed by a rise a t s t i i higher acidities. Sodium sulfate, however, when added in large proportions (about 1000 p.p.m. of sulfate) had little catalytic effect when added t o 85% peroxide, and caused only a very slight increase in the decomposition rate. This increase could not be ascribed merely to a change in the pH, which showed a rise only from about 4.0 to 4.9 (tenfold dilution) after addition of the salt.
90
Alternative Methods of Representing Data on Decomposition of Hydrogen Peroxide 7 composition per hour a t 50' C.
Figure 3. A
997
0: c",. of oxygen evolved per gram of solution per hour a t 60" C
If the decomposition data are alternatively presented in terms of cubic centimeters of oxygen evolved per hour per gram of hydrogen peroxide solution, plotted from 0 to 90% hydrogen peroxide, such curves may show a maximum, as illustrated in Figure 3, in which data from a given run are plotted in both ways. The data resented in this paper are in general expressed as the per cent o?the peroxide originally present which is decomposed per hour. The rate of decomposition of purified 85 t o 90% aqueous hydrogen peroxide solutions, in the absence of stabilizers and freed as completely as possible from dissolved or suspended catalysts, measured in carefully cleaned Pyrex or pure aluminum containers, is known to be very small-at 50" C. it is certainly well under 0,001 weight yo per hour. As the ideal condition of complete purity and elimination of any catalytic effects (including those due t o the walls of the container) is approached more closely, the minimum decomposition rates may be expected to decline still further. Up t o the present time the lowest rate for unstabiiized hydrogen peroxide observed by the author was 0.0008% per hour for a 90% solution a t 50" C.
Although in general reagent quality chemicals were employed in the addition of catalysts, stabilizers, or solutions of acids or salts to the peroxide solutions under examination, addition of any considerable quantity of such substances involves the probability of contamination of the peroxide by possible catalytic impurities in the reagents themselves. As a rule, however, such contamination would be of negligible character, because of both the small concentrations ( a few parts per million) of the added reagents employed and their known purity. For this reason, the fact that added acid results first in a decrease in the decomposition rate, and later in an increase, is probably not to be ascribed t o the accumulation in the peroxide of catalytic materials contained in the acid. Although the influence of varying p H on the stability of peroxide solutions was found to be independent of the presence of stabilizing substances, in the presence of certain catalytic ions, such as ferric or qhromic ions, the decomposition rate was remarkedly sensitive t o pH changes over a very narrow range of
ooa
EFFECT OF pH UPON STABILITY OF HYDROGEN PEROXIDE
Employing 90% unstabilized hydrogen peroxide, t o samples of which sulfuric acid was added dropwise in increasing amounts so that the pH of the resulting solutions, when diluted tenfold with conductivity water, was varied from 5.0 to 1.5, the decomposition rates were measured at 50" C. The results of these determinations, as seen in Figure 4, indicate that a n optimum stability is attained at a p H of about 4 (at tenfold dilution) and that the rate increases regularly when the pH is either raised or lowered from this value. I n connection with these measurements, it is perhaps worth while t o note the effect on the pH of the solution brought about by dilution of concentrated peroxide with the purest available conductivity water. A solution containing 8.36 weight yo of hydrogen peroxide, obtained by tenfold dilution of 83.6% hydrogen peroxide, gave a reading of 5.35 on the Beckman pH meter, while the dilution water alone, saturated with air, gave a pH of 5.67.
Figure 4. Effect of Change in pH upon Stability of 90% Unstabilized Hydrogen Peroxide
INDUSTRIAL AND ENGINEERING CHEMISTRY
998
rx
2 cc
L c d 0
t3 Y
2 a
PH
Figure 5. Rate of Decomposition a t 50" C. of 8570 Hydrogen Peroxide Containing Ferric Ion and with Varying pH 0. 0.10 p.p.m. Fe+++ A.
0.05 p.p.m. Fe +
+
Vol. 41, No. 5
effective stabilization the pH of the peroxide solution should be adjusted, after addition of the stabilizer, to the optimum value. Mechanism of Stabilization. While the facts concerning the stabilization of concentrated hydrogen peroxide have been rathcr fully revealed through the work of various investigators, the mechanisms whereby the results are attained have not beer1 established beyond question in all cases, although reasonable ~vorlcing hypotheses are not lacking by which a means of int,erpretatiori of the observed phenomena is afforded. STABILIZATION B Y S O D I U X S T A N K A T E (NazSn03.3f120). concentrat'ed hydrogen peroxide, up to 70% by- weight, of American manufacture (Du Poiit Company) was provided for this investigation, containing as one stabilizing agent about 70 p.p.m. of sodium stannate. [Pioneering 'work on stabilization of hydrogen peroxide by sodium stannate and sodium pyrophosphate is described in t'he patent literature (8, 11-20) as well as in some foreign patents.] It has been generally assumed that the tin compound acts as a stabilizer by virt'ue of its ability to become hydrolyzeti, forming hydrous stannic oxide, SnOa.zHnO,which, in the form of a colloidal dispersion, adsorbs catalyt,ic ions from the peroxido solution and thus inactivates them.
+
.I4
pH, between about pH 3 and pH 4 (at tenfold dilution). Figure 5 shows this phenomenon in the case of solutions containing two different concentrations of ferric sulfate. An interpretat,ion of these results niay be found in the fact that ferric ion undergoes hydrolysis, forming colloidal hydrous ferric oxide (or basic ferric salts). A similar effect is shown by chromic ion, a sharp peak in the decomposition rate being observed at approximately the same pH as that seen in Figure 5 for ferric ion.
12
/' p
i
I
I
l
I
1 I
l
I I
l
1
I
0
The critical pH region referred to (about pII 3 to 4 a t tenfold dilution) is also the region in which hydrogen peroxide solutions show lowest decomposition rates, either in the absence of added catalysts or when adequate amount,s of stabilizers are present. If sufficient sodium stannate is present, for example, the peak shown in Figure 5 mag be entirely eliminated and a minimum decomposition rate secured at, about pH 3.5 * 0.5. The behavior just discussed is also shown by cupric ion, but to a much less pronounced degree (see Figure 6). Apparently, the sharp maxima seen in the case of ferric and chromic ions arc due to the amphoteric character of the oxides of these metals. E F F E C T O F ADDITION OF STABILIZERS
When a stabilizer such as sodium stannate or sodium pyrophosphate is added to commercial unstabilized concentrated hydrogen peroxide, addition of but a fraction of a part per million suffices to cause a sharp drop in the rate of decomposition. Further addition, of the order of 4 t o 10 p.p.m. of these stabilizers, suffices to bring the rate of decomposition to a minimum value (of the order of 0.0004% per hour at 50" '2.); still further addition, even up to 100 p.p.m., causes a slow but steady rise in the rate of decomposition. '17'it,hin certain limits, addition of the proper stabilizer in the proper quantity will nullify or restrain the catalyzing influence of certain impurities upon the decomposition rate of hydrogen peroxide, notably compounds of iron or of copper. However, if the amount of such active impurities present in the peroxide solution exceeds a certain proportion-on the order of 10 t'o 20 p.p.In.the catalysis caused by such contamination cannot be held in check by even relatively large quantities of stabilizer. The addition of such a stabilizer as sodium stannate, or sodium pyrophosphate, has the effect of increasing the p H of the solution, due t,o hydrolysis of these salts, and if this increase in pH is allowed to rise too far, marked instability of the peroxide will be brought about. This effect is responsible for the steady rise in decomposition rate that attended an incrcase in the stabilizer content beyond an optimum value. It is obvious t,hat for most
10
23
30
40
53
60
PH
I A F T E ? I IO DILUTIOII
Figure 6. Rate of Decomposition at 50" C. of 90% Hydrogen Peroxide Containing 0.1 P.P.M. Cupric Ion and with Varying pH Dotted curveindicates poorly defined maximum
I n corroboration of this hypothesis, it was possible to show that the hydrous oxide alone, when separately prepared (by acidification of a stannate solution and prolonged washing of the resulting precipitated oxide) is capable of stabilizing the peroxide solutions quite as well as a quantity of stannat,e solution equivalent t o the same amount of stannic oxide, as indicatcd by the data in Table V. Such results lend credibility to the hypothesis that stahilization by means of sodium stannate results from a kind of scavenging action on the part of colloidal stannic oxide, whereby dis-
TABLEv. RATE O F DECOMPOSITION O F 84% HYDROGEN PEROXIDE STABILIZED B Y TIN COMPOUNDS Unstabilired HaOn 0.0044
0.0034
(Per cent per hour a t 60' C . ) Stabilized with Stabilized w i t h 0.15 P.P..\I. SnOz.aHsO Equivalent to N a z S n 0 d N ~ O 0.15 P.P.M. NanSnOa.3HzO 0.0008
0.00066 0.00055
..oo12
0012
g 0010
Y
.c
I-
8
0
0
a
0028
x
0010
in
t
= 0008
2E 0008
ir
2
L
w
z
z
headdition of furthei quantities of stannate, which provides colloidal hydrous sta,nnic DECOUPOSITIOK O F 85% HYDROGEK oxide as the effective stabilizer. PEROXIDE BY HYDROUS FERRIC OXIDE AND FERRIC SULFATE M i z e d Catalysts. Another interesting feature of the catalytic (Rate after addition of catalvst. % uer hour a t 50' C.) Rate before effect of added metal ions upon the decomposition of hydrogen 0.36 P.P.M. of Ferric 0.70 P.P.31. of Ferric Catalyst Iron Iron peroxide is t'he promoting action of srnall quantities of one such Added, Added a8 .Added as Added a$ Added a.q 9 /ET. ion upon the catalytic effect of a second ion. Thus, the addition sulfate atc600 C. Fe203.zHzO Fe2(SO&j.61120 oxide of a fractional part per million of cupric ion to a solution of .... 0.0018 0,032 ... 0,0039 .... 0.0323 ,... hydrogen peroxide containing 1 or more p.p.m. of ferric ion re0.0018 .... .... 0.0422 sults in a marked increase in the decomposition rate over and 0.0032 .. . , .... 0 ; 6450 above that caused by the same amount of cupric ion when present alone. This promotion appears to be highly specific and wholly This interpretation is strengtheried by the observation that the lackingin the case of many combinations tried. catalytic effects of addition of ferric compounds, either as hyRepresentative data, showing the promoting efi'ect of small drated ferric sulfate [approximately Fe2(SOc)3.6HsO]or as additions of cupric ion to 90% hydrogen peroxide containing hydrous ferric oxide, freshly prepared and very thoroughly either ferric or chromic ion, follon-. washed, are not very different in the two cases, as the data in Table VI11 are intended to indicate. It appears, therefore, that through hydrolysis of the very dilute ferric sulfate solut,ion, the DECOMPOSITION IZATES O F 90% H Y D R O G E S PEROXIDE AT 50" c. primary effect of contamination by ferric salts is the formation of WITH M I X E DCATALYSTS insoluble hydrous ferric oside, or basic ferric sulfate, which is an Rate of Rate of act.ive decomposition catalyst. Hence, the effect is closely similar Addition, Decomposition, Addition, Decomposition. to that obtained when the hydrous oside i s added directly to the %/Hr. P.P.M. %/Hr. P.P.M. Kone 0.0008 Xone 0.0008 peroxide. 0 . 0 2 cu++ 0.002 0.02 Fe++' 0.0016 Zrnportance of Colloidal Charactei of Certain Decomposition 0 . 1 0 cu+* 0.010 0.10 Fe-+' 0,0075 Catalysts. Much of the experimental evidence concerning the Rate of Addition, Decomposition. catalytic behavior of metal ions, such as ferric or cupric ions, P.P.AI. %/=. upon the decomposition of hydrogen peroxide, as well as the 0.10 Cu+' + 0.02 Fe+" 0.024 stabilization of the peroxide against such contamination by means 0.10 F e + + + + 0.02 C u + + 0.040 of sodium stannate, may be interpreted on the basis of the assumption that both catalysis and stabilization in these cases are fundamentally colloidal phenomena. Thus the fact that the These data indicate the greater promoting effect of addition of catalytic &t,ivit,y of ferric salts added to concentrated peroxide a small amount of cupric ion to a larger amount of ferric ion, as shows a sharp peak over a rather narrow range of apparent pH compared to addition of a small amount of ferric ion to a larger of the solution, was accounted for on the basis of the known amount of cupric ion. tendency of ferric salts to be hydrolyzed appreciably a t other than low pH's, forming basic salts, or, a t still higher pH Rate of Rate of values, hydrous ferric oxide. The known colloidal characterAddition, Decomposition. Addition, Decomposition %/Hr. P.P.M. %/Hr. P.P.M. istics of t'hese slightly soluble iron compounds lead to the None 0.0008 None 0,0008 belief that t'he catalytic action is largely to be attributed t o 0.02 c u + + 0.002 0.02 C r + + + 0.0008 the presence of such colloidal material ; the decomposition, there0.010 0.10 C r + + * 0.003 0.10 C u t + fore, would be predorriinaritly a heterogeneous process. With Rate of Addition, Decomposition. still higher p H values, bhe hydrous ferric oxide would be conP.P.M. %/Hr. verted largely into soluble ferrate, resulting in a sharp lowering O.lOCu+++ 0 . 0 2 C r + + + 0.01 of the rate of decomposition: O.lOCr++++O.O2Cu++ 0.02
TABLEVIII.
RATE O F
,
May 1949
INDUSTRIAL AND ENGINEERING CHEMISTRY
I t is indicated in the above table that a small addition of cupric ion markedly increases the catalytic effect of a larger amount of chromic ion, whereas the addition of a small amount of chromic ion t o a larger amount of cupric ion has little or no promoting action. The combination of chromic and ferric ions in a similar series of tests showed no promoting action of small additions of either ion on the catalytic effect of larger amounts of the other ion. S T A B I L I Z A T I OBNY S O D I U M PYROPHOSPHATE, (NaaPZO7 tOH20). Although pyrophosphate ion is not so generally effective as sodium stannate in protecting concentrated hydrogen peroxide against contamination by certain impurities-in particular, by cupric salts-nevertheless, it is effective in its protective action against ferric ion, with which stabIe complex ions are formed, such as
-
[Fe(OH)PzOr]-- or [Fe(OH)2P2O7]--when an excess of pyrophosphate is added t o the solution containing the ferric ion. The intermediate precipitate of white ferric pyrophosphate is an active decomposition catalyst. If a very small quantity of pyrophosphate or stannate stabilizer is added t o a concentrated peroxide solution-say, 0.1 or 0.2 p.p.m.-subsequent addition of a catalyst, such as 0.1 p.p.m. of a ferric compound, results in an increase in the rate of decomposition which is greater than that observed if the stabilizer had been omitted. This result is dependent upon use of an amount of stabilizer insufficient t o inactivate the amount of catalyst present in the solution; for if adequate proportions of stabilizer are provided, this abnormally large increase in decomposition rate will not be observed, but rather a decrease. The interpretation of these results in the case of the pyrophosphate stabilizer would appear to be connected with the fact that addition of a soluble pyrophosphate to a solution of a ferric salt results, first, in the precipitation of insoluble ferric pyrophosphate, or a basic ferric pyrophosphate, which is a pronounced catalyst for the decomposition of peroxide. If the proportions of pyrophosphate and ferric ion are such that this precipitate is not redissolved (forming complex anions), the decomposition rate will therefore show a pronounced rise. As increasing proportions of pyrophosphate are subsequently added and solution of the precipitate occurs, the rate of decomposition shows a progressive decrease t o a minimum value. As the proportion of pyrophosphate is still further increased, a rise in the decomposition rate again is observed, due t o the decrease in acidity of the solution accompanying hydrolysis of the pyrophosphate. To illustrate the fact that against certain (though not all) catalytic impurities which may be present in commercial concentrated hydrogen peroxide yrophosphate provides protection comparable to that afforded gy sodium stannate, the following data may be offered. With a batch of 90% peroxide, which showed an initial decomposition rate of 0.007 weight yoper hour a t 50 C. addition of 3.5 pap.m. of either sodium pyrophosphate or sodium stannate resulted in a decrease in rate to 0.0004% per hour; and a mixture of 1.73 p.p.m. of the pyrophosphate with 1.73 p.p.m. of sodium stannate gave practically the same results. O
~
In connection with the relative effectiveness of pyrophosphate and stannate stabilization, it should be borne in mind that, whereas pyrophosphate ion is hydrolyzed, slowly a t room temperature and more rapidly a t elevated temperatures, to form orthophosphate ( a less effective stabilizer), sodium stannate behaves in an opposite manner, and solutions properly stabilized thereby will gradually improve in stability with age, as hydrolysis of the stannate yields hydrous stannic oxide, a known effective stabilizer. Hence a comparison of these two stabilizers will be expected to yield results dependent upon the temperature and the duration of the tests. In a series of measurements undertaken to ascertain the optimum proportions of pyrophosphate and ferric ions, contamina-
1001
tion by ferric ion was varied from 0.1 to nearly 2.0 p.p.m. and the amounts of pyrophosphate noted which would reduce the catalytic effect of the ferric ion to the point where the original decomposition rate of the unstabilized peroxide was equaledin this case, about 0.003 to 0.004% per hour ( a t 50" C.). Table I X shows that the ratio of pyrophosphate to ferric ion rose steadily. Doubtless a t least some of this rise was due to the fact that, as the proportion of pyrophosphate is increased, the apparent pH of the solution increases, with a consequent falling off in stability.
OPTIMUMPROPORTIONS OF PYROPHOSPHATE FOR STABILIZATION AGAINST FERRIC ION(50 O C.)
TABLEI X .
Ferric Ion Added a8 Sulfate, P.P.M. 0.08 0.16
0.24
0.52
0.48 0.96 1.08 1.92
Average PlOi---- Required, Added as NarP~07.- Pzo7 ----/Fe IOHaO, P.P.M. Ratio 0.30 3.8 0.60 3.8 1.10 1.60 2.50 7.40 10.00 17.20
+
4.6
5.0 5.2 7.7 9.3 9.0
In the case of larger proportions of ferric ion, it was observed that if the pH ( a t tenfold dilution) was held between 3.7 and 4.5, 35 p.p.m. of pyrophosphate would be able t o restrain up to 4 p.p.m. of ferric ion, corresponding to a ratio of approximately 9.0. These ratios are definitely smaller than those required in the case of sodium stannate; but lower decomposition rates were found to be generally attainable with sodium stannate, when added to commercial, unstabilized concentrated peroxide samples, containing the average mixture of impurities, than when sodium pyrophosphate was similarly used. The effective use of sodium pyrophosphate jointly with sodium stannate has suggested the possibility that the pyrophosphate ion, as a multivalent anion, may serve t o maintain the dispersion of the negatively charged stannic oxide hydrosol formed by hydrolysis of the stannate, as well as to tie up certain catalytic ions, such as ferric ion, by forming with them inactive complexes. In the former function, the tetravalent P 2 0 7 - - - would be more effective than the triply charged PO1---, or the doubly charged HPOI--. It is obviously of some importance t o establish whether or not the pyrophosphate ion has an intrinsic stabilizing effect upon hydrogen peroxide, or whether other types of phosphates would serve equally well in this regard. In comparative tests carried out with mono-, di-, and trisodium phosphates, as well as the pyrophosphate, it was found that the pyrophosphate was definitely superior t o the others in bringing about a reduction in the rate of decomposition of commercial unstabilized 85% peroxide a t 50" C.; it alone brought down the decomposition rate from an original value of about 0.008 to approximately 0.001% per hour, a result that was attained by addition of but 0.1 p.p.m. pyrophosphate (added as sodium DvroDhosohate decahvdrate) . In*combarison, ove; 3 p.p:m. of HzP04- or HPOb-- (added as NaH2P04.12Hz0 or Na2HPOd.lOHzO) showed no decline in the decomposition rate of the peroxide (0.008% per hour at 50" C.) while up t o 6 p.p.m. of PO1--- (added as Na8P04.12H20) caused an increase in the rate to 0.011% per hour. I n part these results may be accounted for by the change in p H brought about by the hydrolysis of the salts, but the indication is still clear that sodium pyrophosphate is superior to the other sodium phosphates in its stabilizing action. OTHERSTABILIZING AGENTS. A number of compounds, of both organic and inorganic character, were tested as possible substitutes for those commonly employed, but no inorganic compound was found which proved superior to the stannatephosphate combinations. Inorganic Compounds. Sodium aluminate, NaA102, was found
9002
'
INDUSTRIAL AND ENGINEERING CHEMISTRY
t o be a fairly effective stabilizer. Thus a sample of commercial, unstabilized peroxide showing a a initial decomposition rate of about 0.006% per hour a t 50" C., on addition of about 5 p.p.m. of sodium alumhate gave a reading of 0.0018% per hour. Ilowever, under similar conditions, sodium stannate or sodium pyrophosphate would have shown a definitely lower value-of the .order of 0.0005% per hour for the stannate. Aluminum phosphate, AIPO,, and pyrophosphate, A14(P207)a, 'were similarly found to be about as effective as the aluminate, Hydrous alumina itself, A1203.zH20, prepared by precipitation a n d thorough mashing, showed st,abilizing action of similar magnitude. Various other inorganic compounds, such as sodium silicate or fluosilicate and antimony trifluoride, were found to be of little if any value as stabilizers. An unsuccessful attempt was also made t o prepare pure stannic pyrophosphate, SnPzO,, in the belief that this substance might offer promise as a source of atannic oxide and the pyrophosphate ion, simultaneously. Zinc ion (added as the hydrated sulfate) has also been shown t o have a certain stabilizing action. A 90% solution of hydrogen peroxide containing 0.5 p,p.m. of cupric ion was observed t o decompose at 50" C. a t the low rate of 0.0016% per hour when 280 p.p.m. of zinc ion had been added. Organic Compounds. Among organic st'abjlizers, none was found which was not subject to oxidation by the concentrated peroxide solutions-for example, alizarin was observed to be a n effective stabilizer immediately after addition to concentrated peroxide, but this action was lost after several days a t 50" C. It is believed that, unless the required period of storage is known 60 be only a few months, the use of purely organic stabilizers may result in significant losses by decomposition over extended periods of time, On the other hand, a serious objection to the use of inorganic stabilizers has been the fact that, when used as a propellant, peroxide so stabilized leaves a solid residue which may a c t as a "poison" upon the solid decomposition catalyst employed. Hydrogen peroxide of German manufacture (85%), used during ithe past war as a propellant, was commonly stabilized by means of $-hydroxyquinoline (8-quinolino1, oxine), often in the form of the pyrophosphate, or mixed with sodium pyrophosphate. This stabilizer is slowly oxidized by the peroxide, SO that in about 6 months an appreciable part of it,s protective action has disappeared. Horever, if the time of storage contemplated is only a few months, this type of stabilizer offers certain advantages over the purely inorganic type. Other similar compounds were suggested by German workers in this field, such as hydroxy derivatives of acridine, I n test,s carried out by the author, using 8hydroxyquinoline and acridine in a parallel series of measurements, acridine was found t o be oxidized, like 8-hydroxyquiriolirie, but a t a slower rate. I n the belief that the action of some stabilizers, such as hydrous stannic oxide, is due to their colloidal behavior, the effect of gelatinous organic materials, such as agar-agar, gum karaya, etc., was tried, but without success. Although the &hydroxyquinoline type of stabilizer, used jointly with sodium pyrophosphate, may be considered less permanent than st,annate or stannate-phosphate types of stabilizers in its protective action against the catalytic decomposition of concentrated hydrogen peroxide, it nevertheless has shown itself superior to the latter type of stabilizer in providing protection against such contaminants as ferric ion, the tolerance for which was found to run a t least as high as 15 to 20 p.p.m. without resulting in excessively high decomposition rates a t 50" C. Further experimentation with the objective of developing organic stabilizers resistant t o the powerfully oxidizing action of 90% peroxide is therefore considered very desirable, and at the present time experimentation is in progress in which 4,5-dihydroxyacridine is being synthesized and will be compared as a stabilizer with 8-hydroxyquinoline, in combination with sodium pyrophosphate.
Vol. 41, No. 5
SUMMARY The stability characteristics of concentrated hydrogen peroxide solutions have been studied at 50 ' to 100' C. over the range 70 to 90 weight % hydrogen peroxide, in both the absence and presence of stabilizers, such as sodium stannate, sodium pyrophosphate, or a mixture of 8-hydroxyquinoline (8-quinolinol) and sodium pyrophosphate. Much of the work was carried out at 50" C., which proved most satisfactory for the gas evolution method used in determining the rate of decomposition of the peroxide samples. Previous reports of high stability for very pure hydrogen peroxide a t concentrations of 85 t o 9070, in the absence of catalytic substances and of "wall effects" from the container, were confirmed. Unstabilized 90% hydrogen peroxide has been observed to show a decomposition rate at 50" C. of well under 0.001% per hour; in the presence of a few tenths of 1 p.p.m. of stabilizer this value may decline to 0.0004% per hour or less. The temperature coefficient, a, of the decomposition rate of unstabilized peroxide solutions, in the expression
was determined for the 10" interval 50" to 60" C. and 60" to 70" C. as 2.2 * 0.1; this value is also believed to be approximately applicable over any 10" range between 50" and 100" C. Methods of preparing the surface of Pyrex containers for use with concentrated hydrogen peroxide were developed which made possible obtaining reproducible measurements of the decomposition rates. Pure aluminum (99.60/, or better), preferably with its surface conditioned by anodizing or allowing to stand in contact with concentrated peroxide over a considerable period of time, as well as pure cast tin, was found to be a suitable and practical material of construction for peroxide containers. Tests of the purity of concentrated peroxide are described, as well as specific conductivity measurements upon redistilled peroxide of all concentrations from 0 to 99 weight %. A tenfold increase in the surface-volume ratio WBS observed to result in a tenfold increase in the rate of decomposition at 50' C. This relationship held fairly constant over the entire range of concentration. Vigorous agitation was shown to be without appreciable effect upon the measured rate of decomposition of S57? unstabilized hydrogen peroxide. The known importance of pH upon the rate of decomposition was emphasized by measurements carried out with and without added catalytic ions, such as ferric and cupric ions (added as sulfates). I n the case of ferric ion the effect is particularly marked in the range of p H 3 to 4 (measured a t tenfold dilution), wherein a sharp maximum in the rate is observed. This maximum is also observed with chromic ion, but is less marked in the case of catalysis by cupric ion, and absent in the case of aluminum ion. Data have been secured upon the catalytic effect of known concentrations of ferric ion or cupric ion in 85 to 90% hydrogen peroxide a t 50" C., arid upon the required concentrations of stabilizers t o restrain such catalytic action. Previous reports have been corroborated that gross contamination by active catalytic ions, such as ferric or cupric, cannot be successfully held in check by any concentration of known stabilizers. The fact that the presence of an insufficient quantity of stabilizer may in fact produce a greater rate of decomposition, in the presence of a given concentration of catalytic ions, than in the absence of the stabilizer, was noted and an explanation of this phenomenon is offered. Discussion is presented of the probable mechanisms of stabilization of Concentrated hydrogen peroxide by sodium stannate and sodium pyrophosphate. The possible use of alternative stabilizers, both inorganic and organic, is considered.
May 1949
INDUSTRIAL AND ENGINEERING CHEMISTRY
LITERATURE CITED Broughton and Wentworth, J . Am. Chem. SOC.,69, 741 (1947). Calvert, Ann. Physik, (4) 1, 483 (1900). Cuthbertson and Maass, J . Am. Chem. Soc., 52, 489 (1930). Fiske and Subbarow, J . Biol. Chem., 66, 375 (1925). International Critical Tables, Vol. VI, p. 230, New York, McGraw-Hill Book Co., 1929. Joyner, Z . anorg. C h e m . , 77, 103 (1912). Maass and Hatcher, J . Am. Chem. SOC.,42, 2548 (1920). Maohu. "Das Wasserstoff Peroxid und die Perverbinduneen." -Vienna, Julius Springer, 1937. Regnault, H., and Le Noir de Carlan, R., Congr. chim. ind., Cornpt. rend. 18me Congr., Nancy, 1938, 766. Reichert, J. S., Chem. Eng. News, 21, 480 (1943). Reichert, J. S.,et al. (to E. I. du Pont de Nemours & Go.), U. S. Patent 1,958,204 (1932). Ibid., 2,004,809. Ibid., 2,008,726. Ibid., 2,021,834.
1003
(15) Ibid., 2,027,838. (16) Ibid., 2,027,839. (17) Ibis., 2,091,178, (18) Ibid., 2,224,835. (19) -I b i d . . 2.316.487. (20) I b i d . : 2:347;434 (1941). CHEM.,39, 1536 (1947). (21) Shanley and Greenspan, IND.ENG. (22) Wolffenstein, Ber., 27, 3307 (1894). > - - ,
RECEIVED June 28, 1948. A portion of the work was carried out under N a v y Bureau of Ordnance Contract NOrd 9107, Task C ; the rest was under Contract Njori-78; T.O. XIX, Office of Naval Research, from which permission to publish was received. General supervision was by H. C. Hottel, G. C. Williams, T. K. Sherwood. and C. N. Satterfield, Department, of Chemical Engineering. Experimental work was carried out by the following group in the Department of Chemistry under Projects D I C 6351 and 6552 of Division of Industrial Cooperation: R. C. Young, C. D. Turner. R. Knodel, Mrs. P. H. Blackall, Misses E. M. Bickford, H. 34. Stickley, and C. L. Tucker.
Low Temperature Carbonization
of Alberta Subbituminous Coal RESULTS OF PRELIMINARY TESTS OF THE NEW STANSFIELD RETORT J. GREGORY AND A. MCCULLOCH Research Council of Alberta, Alberta, Canada has been the object of many OTENTIALLY the An account is given of the construction and operation investigat,ions; some of these largest home market for of the new Stansfield retort for the low temperature carhave been developed on the Canadian coal is in the Provbonization of Alberta subbituminous coal. The principles industrial scale. The subinces of Ontario and Quebec of operation of the retort have been established on a basis bituminous coal may be carwith their relatively high of the yields and compositions of the low-temperature bonized at a low temperadensity of population and carbonization products of an Alberta subbituminous coal. ture and the char so protheir comparatively heavy The necessary data have been obtained from stepwise carduced briqueted with a suitaindustrial c o n c e n t r a t i o n . bonization of the coal, in stages of approximately 100" C. ble binder, or under certain This market, however, is more from 320" to 610" C., in a Gray-King low temperature circumstances, briqueting carbonization assay apparatus. Details are given of the easily and economically supmay precede carbonization. plied with coal imported from results of preliminary tests of a retort built on a pilot As early as 1921, Bone (8) the coal fields of Pennsylplant scale. The yields and compositions of the chars pointed out that if immature vania and Ohio. The greater obtained under varying conditions of oneration of the coals are carbonized up to cost of transporting coal from retort are described. temperatures of the order of Alberta is the principal rea400' C., a large proportion of son why coal producers there the oxygen present is evolved as steam and as oxides of carbon find it difficult to compete successfully with coal produced in (principally carbon dioxide) and the calorific value of the solid the United States. Nonetheless, with the aid of a Dominion material thus substantially increased. The carbonization of coals Government subsidy on transport, some coal mined in Alberta is at temperatures as low as 400" C., however, is not economically marketed in Eastern Canada. Apart from the cost of transport, practicable, and in most investigations temperatures not exceedcoal quality is an important factor in the situation. Generally ing about 700' C. have been employed. I n Canada, considerable speaking, the bituminous coals of Alberta are of high quality experimental work of both a small and large scale nature was and, after preparation for the market, admirably suited to most carried out by Stansfield and others ('7) for the Canadian Lignite industrial purposes. On the other hand, although possessing Utilization Board. Various types of retorts were designed and properties such as cleanliness in handling, low ash content, and eventually one was adopted for the construction of a battery of the production of only small quantities of smoke on combustion, six retorts, each having a daily production capacity of 16 tons of which make them attractive solid domestic fuels for use locally, char. The plant, however, was found t u be unsatisfactory in the subbituminous coals have other properties which are discontinuous operation. Meanwhile, American investigators had advantageous when these coals are considered for use in distant developed the Hood-Ode11 system for the carbonization of the markets. Their calorific values are low due to their high moisture lignites of North Dakota (6). The Hood-Ode11 plant (Figure 1) and oxygen contents, and since certain of them slack readily, operated somewhat o n the principle of a simple gas producer. considerable quantities of finely-divided material may be proThe lignite was dried in the upper section of the plant before its duced in transit and during storage before use. carbonization and partial combustion in the lower section. The T o devise an economical method of carbonizing subbituminous air for combustion was admitted through tuyhres in the lower coals or lignites to produce a solid residue of enhanced heating section and the heat developed sufficed for the drying and carvalue and possessing good storage and weathering properties,
P
