Stability of Dilute Alkaline Solutions of Hydrogen Peroxide

Stabilitv of Dilute Alkaline Solutions J -

of Hydrogen Peroxide W. D. NICOLL AND A. F. SMITH

Chemical Depar tment , Experimental S ta t ion , E. I . du P o n t de Nemours & Co., tnc., Wi lmington, Del.

YDROGEN peroxide is an important agent for bleaching such cellulosic materials as cotton cloth and ground-

wood pulp and has also been used as a super bleach for sulfite- and kraft-type wood pulps. Technical problems related to these uses have supplied background for a study of the chemistry of hydrogen peroxide in alkaline media, with special emphasis on the decomposition of hydrogen peroxide to give oxygen and water.

Pure solutions of hydrogen peroxide at low pH are stable, but in the presence of traces of catalysts and with increase in alkalinity there is an increasing tendency for the peroxide t o decompose. In the pH range 10.0 to 12.0, where most bleaching of cellulose is done, special precautions must therefore be taken in order to make the process practical. One objective of the present study was to discover whether this tendency on the part of hydrogen peroxide to decompose at high pH is caused by the hydroxyl ion itself or by catalytic reactions induced by the presence of metal impurities. A study was made of the action of catalysts and inhibitors for decomposition of hydrogen peroxide.

MATERIALS AND TESTS

Hydrogen Peroxide. The hydrogen peroxide used in this study was an especially purified product containing no added stabilizer. It was prepared in the laboratories of the Electro- chemicals Department, Du Pont Co., by treating a freshly prepared 3570 solution of hydrogen peroxide containing no colloidal stabilizer with ion exchange resins to remove unwanted cations and anions. The unstabilized solution was kept a t low temperature (40" F.) in a borosilicate glass container and was found to maintain a nearly constant composition for several months.

Deionized water was obtained by passing ordinary distilled water slowly through a 12-inch bed of Amberlite Mono- bed resin (Rohm & Haas). This treatment increased the electrical resistance of the distilled water from approximately $00,000 to over 12,000,000 ohms. Ordinary tap water was also used for comparison with distilled and deionized water. A trace of copper ion was found to be an impurity in the distilled water used in the study, but was easily removed by treatment with the ion exchange resin. Tap water contains compounds which both catalyze and inhibit the decomposition of hydrogen peroxide. The effect of such water was found to vary with the season of the year.

Caustic Soda. The caustic soda used as an alkali was Merck's reagent grade C.P. sodium hydroxide. Spectroscopic analyses and stability tests indicated that this grade of caustic soda was relatively free of heavy metal ions that significantly affect the stability of peroxide solutions.

Sodium silicate was especially purified for use in this study. Initially, attempts were made to prepare a pure silicate by reaction of a reagent grade silicic acid with C.P. caustic soda, but the resulting product had a strongly catalytic action on hydrogen peroxide. Better results were obtained by precipitating silicic acid from a solution of ordinary commercial sodium silicate and extracting several times with dilute hydro- chloric acid to remove unwanted metal ions. The silicic acid was then washed with deionized water and redissolved in a solu-

Water.

Sodium Silicate,

tion of C.P. caustic soda made up in deionized water. Spectro- graphic analysis of the product showed only trace quantities of metal impurities and this ti'as confirmed by a stability test on hydrogen peroxide.

Because of the sensitivity of hydrogen peroxide solutions to decomposition by traces of various impurities, precautions were taken throughout the study to reduce the effects of impurities accidentally introduced into the peroxide solutions from the walls of glassware or other apparatus employed in the work. Before use, all glassware was thoroughly cleaned by washing successively with 5Y0 caustic soda, water, 5y0 hydro- chloric acid at about 60" C., 5% nitric acid a t 60" C., and finally several times with deionized water. The rinsed glassware was allowed to drain dry and was used within a day or two after cleaning.

Peroxide Determinations. Throughout the study concentrations of hydrogen peroxide in aqueous solutions or bleaching baths were determined by titrating the iodine liberated from an acidified solution of potassium iodide using 0.1N sodium thio- sulfate and ammonium molybdate as catalyst. Check determinations were occasionally made using potassium permanga- nate.

The stabilities reported for peroxide solutions or bleaching baths were determined a t constant temperature (50°, 80°, and 95" C. were used a t various times) by folloKing the procedure described below.

The simplest method used for determining peroxide stability was to place the solution under test in a 1 X 12 inch borosilicate glass test tube closed with a one-hole rubber stopper carrying a 15-inch length of 6-mm. glass tubing to act as air condenser, and to suspend the tube in a constant temperature bath for varying lengths of time. Aliquots were withdrawn from the solution and analyzed for residual peroxide as described. Stability tests m-ere frequently run for various lengths of time, but the period adopted as standard for reporting stability was 1 hour. The stability number is the percentage of initial peroxide remaining.

Glassware.

Stability Tests.

MECHANISM OF PEROXIDE DECOLMPOSITION

In acid media and especially with concentrated solutions of hydrogen peroxide, the mechanism most often proposed for the decomposition reaction producing oxygen and water has involved the formation of free radicals initiated by the presence of various metal ions. This subject has been reviewed by Weiss and Humphrey (f6-17), by Barb and others (7 , 8), and most recently by Uri (14). Other mechanisms which have been used to explain the decomposition of hydrogen peroxide have involved oxidation- reduction reactions [Abel (1-3), Anderson (4-6), Evans, George, and Uri (9), and Pierron (11-13)]. Several mechanisms have included the idea of the formation of unstable metal ion complexes [Abel ( 1 4 , Pierron (11-13), and Glassner (IO)]. Still another possible mechanism used to explain some of the results obtained in this study involves the decomposition of alkaline hydrogen peroxide in the presence of colloidal materials, particularly metal hydroxides which may show catalytic properties based on specific compositions and large reactive surfaces. I n practice, different mechanisms may be expected to operate alone or in combination, depending on the experimental environment.

2548

December 1955 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y 2549

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RESULTS AND DISCUSSION

Effects Produced by Hydroxyl Ion. The role of the hydroxyl ion in the decomposition of hydrogen peroxide has been investigated through stability tests on solutions made up carefully to contain only hydrogen peroxide, deionized water, and C.P. caustic soda. The amount of alkali was varied to produce a range of pH values of about 10 to 12.5, as this was the range most useful in bleaching cellulose. Typical results, shown as solid lines in Figure 1, indicate that in the system under study very little variation in stability occurs with change in alkali concentration. It appears unlikely therefore that the hydroxyl ion is itself an initiator of peroxide decomposition.

Other data in Figure 1 show the effects of increasing pH on the stabilities of peroxide solutions containing distilled and tap waters instead of deionized water. With both of these waters, the effects of increahg alkalinity were very different from that observed with deionized water. Use of distilled water caused the stability of the hydrogen peroxide to decrease continuously with increase in alkali content of the solution and suggested that such water contained a strongly catalytic impurity. Spectro- graphic analyses of the distilled water used in this work indicated the presence of copper ion as the most probable impurity causing decomposition of the peroxide. The concentration of copper ion a as estimated to be approximateIy 0.05 p.p.m. and the catalytic activity of this amount of copper was confirmed by separate tests, in Thich trace quantities of copper were added to peroxide solutions in deionized water. Data in Figure 2 indicate that copper sulfate a t a concentration of 1 X 10-6M (0.06 p.p.m. of copper) hBd about the same effect on peroxide stability as distilled water (cf. Figure 1). The effects of adding traces of ferrous sulfate r e r e less than those produced by copper sulfate, but iron was found to have definite catalytic activity in the system studied.

An interesting feature illustrated by the data for distilled water in Figure 1 and also by the data for copper and iron added in Figure 2 is the reversal in stability of peroxide which occurs in some systems at high concentrations of alkali. Such reversal in the rate of decomposition of peroxide with increase in alkalinity has been observed only in the presence of very low

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Background information for use in commercial problems arising when cellulose is bleached with hydrogen peroxide

concentrations of heavy metal ions and not with higher concentrations of catalytic metals.

Before an attempt is made to explain the reversal in rate of decomposition of hydrogen peroxide in highly alkaline solutions, i t would be well to state the causes for the increase in catalytic activity observed for various metals in the intermediate pH range 10 to 11.5. Traces of heavy metal ions such as copper or iron are believed t o form unstable peroxides or complex per- ions with resulting decomposition of the peroxide; but in addition, colloidal hydroxides are formed as the alkali content of the solution increases and these hydroxides are believed t o be more active catalysts for peroxide decomposition than are the complex peroxides or per- ions. With a large excess of alkali, however, the heavy metal hydroxides will redissolve, in accordance with the well known tendency of such hydroxides to disperse in more concentrated solutions of strong bafies. Higher concentrations of heavy metal ions produced an increase in catalytic activity at both low and intermediate concentrations of alkali and require a large excess of alkali to redissolve the active heavy metal hydroxides. Under the latter condition the rate of decomposition may remain high and no reversal in activity of the heavy metal catalysts can be expected. The point of minimum stability in such solutions was found t o vary with the concentration of catalytic ion, and with increasing amounts of catalysts the minimum shifted in the direction of higher alkalinities (Figure 2).

From Figure 1 it is apparent that the effect of increasing alkalinity on the stability of hydrogen peroxide solutions wade up with tap water was more like that obtained with deionized water than with distilled water. The effect of tap water undoubtedly

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Figure 2. Stability of hydrogen peroxide solution Initial concentration of hydrogen peroxide, 0.15 %. Temp., 50' C. With increase in alkali and in presence of traces of

cupric or ferrous sulfate

2550 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y VOl. 47, No. 12

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Initial concentration of hydrogen peroxide, 0.15%. Temp., 50D C. Insolutions made alkaline with sodium silicate and

C.P. caustic soda

stability of hydrogen peroxide

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can be attributed to the presence of magnesium and calcium salts in relatively high concentration (50 to 100 p.p.m.). Salts of magnesium and calcium have long been known to be stabilizers for hydrogen peroxide and in tap water were apparently present in sufficient concentration to offset the activities of catalytic impurities.

The effects of increasing alkalinity on peroxide stability were also determined for a series of solutions in deionized water containing sodium silicate in addition to caustic soda (Figure 3). It is apparent that sodium silicate, although carefully purified,

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still contained catalytic materials. When distilled water was substituted for deionized water, the total concentration of catalyst supplied by the sodium silicate and the water was apparently increased to an extent g-hich precluded a reversal in stability a t very high alkalinities (Figure 3 ) . On the basis of these and other experiments with sodium silicate, it seems certain that sodium silicate cannot by itself be considered a stabilizer for peroside. Sodium silicate, however, is capable of forming complexes with a large number of metal ions, and these complexes have definite effects on the catalytic processes by which hydrogen peroxide decomposes in alkaline solution.

Effect of Temperature. An increase in temperature reduces the stability of a peroxide solution, regardless of whether deionized water, C.P. chemicals, or other conditions are employed. A comparison of the effects produced by alkali a t 50" and 80" C. is shown in Figure 4. Although temperature had a marked effect on the absolute stability, it had little influence on the change in stability with increase in alkalinity.

The effect of various additives on the catalytic decomposition of alkaline hydrogen peroxide was studied by the addition of selected ions (ferric, nickel, cobalt, cupric, and silver) to solutions of hydrogen peroxide made up with deionized water and pure caustic soda both without and with added sodium silicate. The experiments covered the pH range 10 to about 12.5 and the data obtained are summarized in Figures 5, 6, and 7.

The effects produced by different ions are highly specific and it is not surprising that it is difficult t o explain the decomposition of alkaline hydrogen peroxide in terms of a single mechanism of catalytic activity. Among the effects observed are the following:

1. hlthough the catalytic activities of most of the heavy metal ions increase with increase in alkalinity, this is not the case with nickel ion and the effect is small for cobalt ion (see Figure 6).

The activities of some catalytic metal ions do not increase regularly with Concentration. An example is ferric ion a t concentrations of 10-6 and I O - b J f within the pH range 10.5 to 11.4 (Figure 5). Here feriic ion acts as a stabilizer a t low or moderate concentrations, but is a catalyst a t higher concentrations. The stabilizing effect of ferric ion was also observed in solutions

Effects of Added Metal Ions.

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Figure 5. Effect of pH and added ferric chloride on stability of hydrogen peroxide

In presence of deionized water and C.P. caustic soda


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Figure 6. Effects of pH and added nickel or cobal sulfates on stability of hydrogen peroxide

Initial concentration of hydrogen peroxide, 0.1s 70. 50' C.

Temp., In presence of deionized water and C.P. caustic soda

containing sodium silicate along with caustic soda and i t was observed that the maximum concentration a t which ferric ion produces a stabilizing effect is higher in the presence of sodium silicate than when caustic soda is the sole alkali. .4 possible explanation for the low degree of catalytic activity

or even stabilizing effect shown by iron salts in low concentration is that the peroxides or complex per- ions formed with iron are in themselves fairly stable compositions. Ferric hydroxide, on the other hand, appears to be a strong catalyst for peroxide decomposition and is probably responsible for the increase in catalytic activity at higher concentration and at higher pH. The stabilizing effect which sodium silicate shows with iron salts may possibly be due to the fact that sodium silicate prevents the formation of ferric hydroxide. Some of these relationships are il- lustrabed by data in Table I.

The catalytic activity of copper ion (Figure 7 ) combines some of the features discussed under items 1 and 2 above. In Figure 7 increase in catalyst added is plotted against peroxide found, with p H constant. The data show that at a p H of 11.6 copper sulfate exhibited a very high catalytic activity at a concen-kration of lO-7M and that the catalytic

3.

Table I. Effect of Soluble Iron on Peroxide Stability

Solution (Heated Rapidly and Filtered) a t 80' C. Fe, P.P.M.

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cation is present in a peroxide solution, the effect may be more complex than expected from effects with single ions. This is illustrated by data in Table 11, which shows the effects of adding both copper and iron salts to an alkaline peroxide bath. The combined effect of these ions is greater than the sum of the separate effects and this finding is consistent with the idea that one ion may act as a promoter for the catalytic activity of another ion. Explanations for the behavior include the idea that extremely unstable complex peroxides containing more than one metal may form. I t is possible that a mixture of the colloidal hydroxides of two or more metals may have unusual potency for surface catalysis.

Table 11. Promoter Action of Iron(I1) and Copper(I1)

NO. Reagents

Stability, % HzOz

after 1 Hr. a t 50° C. pH

1 5 x 10 -eM CUSOI + NaOH + HzOz 75 10 62 97 10 50 2 10-jM FeC13 + NaOH +.HzOz

100 10.55 3 5 X 10-5M FeCh + NaOH + HzOz 69 10.72 4 5 X 10-eM CuSOa + lo-5M FeC13 +

U.0. 5 5 j; '?d;-eM CUSOI + 5 X 1O-SM FeC13 + 48 10.82

NaOH + HzOz

Effects Produced by Anions. h number of tests indicated that different anions influence the catalytic decomposition of peroxide in alkaline solutions in different ways. For example, equivalent concentrations of ferric ion introduced as the sulfate had greater catalytic effects than ferric ion added as the chloride. To investigate this point further, the catalytic activities toward

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should be less active catalytically than other hydroxides. stability of hydrogen peroxide

Catalytic Effects with More Than One Cation. J4'hen more than one catalytic deionized water and C.P. caustic soda

Figure 7.

Initial Concentration of hydrogen peroxide, 0.15%.

Effects of pH and added cupric sulfate or silver nitrate on

In presence of Temp., 50' C.

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and sulfate may influence peroxide decomposi- ".>NO M a NaOH tion through reaction with alkaline earth metals.

I n some cases insoluble DreciDitates are formed

in solution. Other anions such as carbonate

Vol. 47, No. 12

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which would tend to prevent the alkaline earth cations from acting as stabilizers for peroxide.

The effect of cyanide ion appears to be very complex. In a solution containing only sodium and hydroxyl ions, along with hydrogen peroxide, cyanide ion seems to be inert, but in the presence of catalytic cations such as

I I copper, cyanide forms complexes with the 10-6 10-5 IO- 4 10-3

LOG MOLAR GONG. MOS04 cation and thereby helps to inhibit peroxide decomposition.

Mechanism of Inhibitor Action. Among the mechanisms that may be cited to explain the actions of various inhibitors in a peroxide bleaching solution, the following appear worth con- sideration.

The inhibitor forms a stable complex with peroxide and so opposes the effects of catalytic ions which form unstable peroxides. It is probable that magnesium and calcium ion o1ve much of their stabilizing activity to the forlnation of such stable peroxides.

The inhibitor may form a stable complex with catalytically active materials such as heavy metal cations and thereby prevent the latter from combining wit,h peroxide as unstable compounds or forming catalytic hydroxides. An excellent example of a stabilizer capable of forming complexes with catalysts is dimethylglyoxime, which forms complexes with cobalt, nickel, silver, copper, and ferric iron (Table IV). Other compounds acting similarly include ethylenediamine, ethvlenediaminetetraacetic

Hydroxyl 90 10.87 acid, sodium citrate, sodium tart,rate? sodium pyrophosphate, Acetate 87 10.86 and sodium silicate. Cyanide ion may also be included in this Bicarbonate 77 10.50 Tetraborate 89 9.37 list, but cyanide is not sufficiently selective in its action toward

Dimethyl- Bromide 87 10.90 Perborate 95 10.90 Pyrophosphate 100 10.83 glyoxime is especially valuable, in that it preferentially forms Dihydrogen phosphate 97 8.15 Iodide 5 11.85 complexes with the heavy metal catalytic cations.

Sodium silicate iA one of the complexing agents that react with Chloride 92 10.60 Citrate 97 10.80 Nitrite alkaline earth metals and this may be the basis for the fact that Nitrate 90 10.82 Tartrate 96 10.80 Sulfate 79 10.92 Carbonate 59 11.21 C,y.an ide 88 10.80 Silicate 91 10.75 Table IV. Dimethylglyoxime as Stabilizer for Alkaline

Figure 8. Stabilizing effects of magnesium sulfate on peroxide solutions containing caustic soda or sodium silicate

Initial concentration of hydrogen peroxide, 0.15%. Deionized water. 10 -5M cupric sulfate. pH 10.5. Temp., 50° C.

hydrogen peroxide of some 18 different negative ions, as sodium salts, were compared. I n these tests the concentration of nega- t,ive ions was that provided by 0.01M concentration of t,he salts. The pH of these solutions was adjusted to a constant value of 10.5 by addition of caustic soda.

1.

2.

Table 111. Stability of Alkaline Peroxide in Presence of Various Anions

(0.15% &OS. Deionized H20 X 0.01N O.P. oaustio soda) % H s O ~ Re-

Anion Added as Sodium Salt, maining after 1 0.01M Hr. at 50' C. Final pH

cations and may combine with alkaline earth metals.

82 10.82

Hydrogen Peroxide (0.15% H20z. O.OO5.V NaOH. pH 10.5 =t 0.2)

Stability, '% of Original H20z after 1 Hr. a t 50' C.

Without Wi1.h 0.01% Table I11 presents the data from these tests. They confirm

of sulfate ion than chloride ion and also show that several other the fact that hydrogen peroxide is less stable in the presence Catalyst DMG DMG

ions, such as iodide, carbonate, and nitrite, affect the decompo- 10-5M CuSOn 83 92 sition of hydrogen peroxide more strongly than hydroxyl ian. 10-JM cos04 Zero 100 On the other hand, stabilizing effects were observed for phosphate, 10-6M AgNOs Zero 69

None, control 93 98 10-4M FeCls 81 100

1 0 - 6 3 4 NiSOh 75 98

pyrophosphate, citrate, tartrate, and borate ions. Acetate, bromide, -chloride, and nitrate ions appeared to be relatively inert toward hydrogen peroxide. Silicate ion also is inert, in the sense that its presence along with caustic soda does not increase or decrease the stability of peroxide a t pH 10.5.

Some of the foregoing anions may actually react with hydrogen peroxide. This is the case with iodide and nitrite ions, as they presumably could consume peroxide through direct oxida- tive reactions. However, it is likely that most of the anions influence the rate of peroxide decomposition through interaction with cations to increase or decrease the activity of the latter. Some anions form complexes with heavy metal cations and these undoubtedly influence the rate of peroxide decomposition. For example, silicate, citrate, and tartrate ions form complexes with

Table V. Effects of Sodium Silicate on Peroxide Stability without and with Magnesium

of Peroxide Remaining after 1 Hour a t 50" C. N a O H +

NaOH + NaSiOz + NaOH NaOH + M g M g

Catalyst PH alone NazSiOa (10-4M) (10-SAM) 80 80 79 98 50 45 74 96

12 34 17 58 92

FeCla (IO-'M) iy:: CuSOi (10-6M) 10.5 60 2 95 92

11.4 0 0 50 85 Cos01 (10-6M) 10.5 46 90 80 97

11.5 46 HO 78 96 12.2 46 38 76 95


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Figure 9. Effect of centrifuging on peroxide stability Initial Concentration of hydrogen peroxide, 0.15 70. Temp., 50' C. Solutions containing deionized water, caustic soda,

cupric sulfate, and magnesium eulfate

sodium silicate is not by itself a very good stabilizer €or peroxide. However, if the concentration of alkaline earth metal greatly exceeds that of catalytic ions. the presence of sodium silicate is favorable, as it prevents the formation of insoluble alkaline earth hydroxides. This is illustrated by data in Table T', u hich com- pare the effects of magnesium with and without sodium silicate Figure 8 shows that the benefits resulting from the inclusion of magnesium are dependent on having present a considerable excess of magnesium over copper as a catalytic ion. The amount of the excess required to ensure good stability varies with the catalytic ion and in this study it was found to be between 5- and 100- fold.

3. The inhibitor may be a colloid physically capable of adsorb- ing or absorbing catalytic ions and preventing their initiating the decomposition of peroxide. The action of sodium silicate in a bleach has often been attributed to its colloidal DroDerties, and because of this the effect

amounts of a colloidal precipitate may result in the absorption of some catalytic materials which would otherwise initiate the decomposition of peroxide. Such effects may possibly be obtained in the case of the hydroxides of alkaline earth metals.

An illustration of this is given in Figure 9 which summarizes the effects of centrifuging an alkaline solution of peroxide containing copper sulfate as added catalyst and magnesium hydibx- ide as stabilizer, Precipitates were formed and the stabilities determined before and after centrifuging a t high speed. The data show that the removal of precipitated hydroxides did result in increased stability. However, in other tests using similar solutions containing added sodium silicate, centrifuging did not produce any change in the rate of decomposition of hydrogen peroxide.

Additional data on the possible importance of colloidal precipitates in solutions of hydrogen peroxide were obtained through the addition of sodium chloride t o an alkaline peroxide solution containing ferric hydroxide. The sodium chloride had a strong i'salting out'' effect on the ferric hydroxide and increased the amount of precipitated hydroxide removed by centrifuging. This resulted in a net gain in stability over a control to which no salt had been added, but in the uncentrifuged solution the addition of yodium chloride reduced the stability of the peroxide. The results confirm the idea thnt ferric hydroxide is an active catalyst for the decomposition of hvdrogen peroxide.

Effects of Order of Addition of Catalytic and Inhibitor Ions. To explore further the effects of added catalysts and inhibitors, the order of addition of ferric chloride as catalyst, magnesium sulfate as inhibitor, and both caustic soda and sodium silicate as alkalies was varied in qolutions containing hydrogen peroxide

Table VI. Variation in Peroxide Stability with Order of Adding Reagents

St;bi&t&

after 1 Hr. Reagents at 50'C. pH NO.

1 2 3 4 5 6 7 8 9

10

FeCla, MgSOp, Na silicate NaOH H z 0 z FeCla Na silicate hfgSO4' NaOH: HzOz HzOz 'FeCla MgdOr Na sheate NaOH FeClj M g S b r NaOH Na silicaie HzOe MgS04, Na siiioate, $aOH, FeCla: HzOz H202 MgSOr Na silicate NaOH, FeCla FeCIi, HzOz, N a silicat,e, NaOH, MgSOa hlgSO4, NaOH, Na sil!cate, FeCla, HzOz NaOH bfgSOa, Na silicate, FeCla, HzOz FeCla, 'Na silicate, NaOH, HzOz (no hlg)

77 68 62 41 41 35 30 19 17 41

11.88 12.02 11.93 12.02 12.02 12.04 12.02 12.02 12.02 12.04

Concentrations. l0-4M FcCla, 10-4M MgSOa, 0.2% SiOp, 0.05.V NaOH, 0.5% HzOz.

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the benefits obtained from sodium silicate are dependent on its buffering action and its ability to form complexes with metal ions. However, it cannot be denied that the presence of large

Figure 10.

Initial concentration of hydrogen peroxide, 0.15 %.

Effect of magnesium sulfate on catalytic action of silver nitrate in decomposition of hydrogen peroxide

Distilled water. C.P. caustic soda. Temp., 50° C.

2554 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y

Determination of the resulting stabilities showed that the order of addition of the different reagents had considerable influence on the stability of the peroxide, as shown by data in Table VI. I n these experiments the solutions were made alkaline to a pH of 12 and the concentrations of both catalyst and inhibitor salts were l0-4M. The percentages of peroxide remaining after 1 hour a t 50" C. varied irom a high of 77 to a low of 17.

The higher stabilities were obtained when sodium silicate was available to react with both magnesium and iron before the addition of sodium hydroxide. The lower stabilities resulted when the order of addition of reagents permitted the formation of insoluble magnesium hydroxide before addition of iron. Inter- mediate stabilities were obtained when the order of addition was such as to permit the formation of some ferric hydroxide. From the results, it would appear that the chief function of sodium silicate in these solutions was to maintain the metal ions, iron and magnesium, in a soluble state and prevent the formation of the insoluble ferric hydroxide. The last item in the table confirms other findings in showing that the catalytic activity of ferric chloride in the presence of silicate is rather low for concentrations below lO-3M.

Magnesium Hydroxide as a Promoter of Catalytic Activity. The complex effects which the presence of two or more cations may produce in a peroxide solution are shown further by the anomalous action of magnesium, which in the form of the hy- droside was found to promote the catalytic activity of silver ion (Figure 10). Magnesium hydroxide apparently acted as a support for finely divided catalytic silver oxide, so that in this instance the colloidal properties of the hydroxide augmented rather than inhibited decomposition of the peroxide.

Effect of Sodium Citrate in a Peroxide Solution. Sodium citrate is known to be an effective complexing agent for a number of metal ions. Glassner (IO) has shown that in an alkaline solution of hydrogen peroxide and in the presence of copper ion, the effect of sodium citrate is to maintain the copper ion in solution. He states, however, that copper is an active catalyst only when the ratio of caustic soda to citrate in the solution is greater than 1 to 1 . He showed further that brown cupric peroxide was formed only when the ratio exceeded this value and that the rate of decomposition of hydrogen peroxide was directly dependent on the concentration of unstable copper peroxide. These findings were cmfirmed in the present study by data showing that the complex formed between copper and citrate has little catalytic activity, and it would appear that copper must be in the form of the hydroxide before it is able to complex with peroxide. If magnesium ion is present along with citrate, the effect is similar to that produced by magnesium silicate and the result is an inhibition of the catalytic activity of copper. Citrate, however, differs from silicate ion in that the citrate anion by itself seems to inhibit the activity of copper, whereas the silicate anion promotes the catalytic activity of copper.

Effects of Cotton Cloth on Peroxide Stability. In view of the complex relationships shown to operate in alkaline solutions between catalysts and inhibitor for the decomposition of hydrogen peroxide, it is not surprising that the introduction of cellulose should exert a strong and sometimes unpredictable effect on the decomposition rate of the peroxide. Because of this, experiments

Vol. 47, No. 12

under actual bleaching conditions are not reported in this paper. Findings reported on the decomposition of hydrogen peroxide have in no way altered the generally accepted ideas regarding the benefits to be derived from the presence of calcium and magnesium salts in the bleach or the presence of sodium silicate as part of the alkali.

SUMMARY

Knowledge of the decomposition of dilute alkaline solutions of hydrogen peroxide has been extended through quantitative studies on the processes of decomposition and stabilization of hydrogen peroxide. Among the findings, one of the more important is that the tendency of solutions of hydrogen peroxide to decompose a t a faster rate with increase in alkalinity is not due to increase in the hydroxyl ion concentration alone, but is caused by increase in the catalytic activity of certain cations present in the solution. Relatively stable solutions of hydrogen peroxide can be produced by deactivating catalytic ions through formation of soluble or insoluble complexes with such materials as dimethylglyoxime or sodium silicate or by physical removal of the offending ion. The catalytic activities for different cations are specific and subject to modification by such influences as concentration, pH, and presence of other cations or anions.

ACKNOWLEDGRIEXT

The authors wish to acknowledge the assistance of the n'iagara Falls Laboratory, Electrochemicals Department, Du Pont Co., and especially to thank J. H. Young for advice and supplying deionized solutions of hydrogen peroxide and samples of un- bleached wood pulps; R. K. Iler, G. B. Alexander, and W. 11, Heston, Jr., Grasselli Chemicals Department, for advice and assistance in the preparation of purified solutions of sodium silicate; ITr. B. Askew and Leah Stenzel, PhyPical and Analytical Division, Chemical Department, for analytical work in gases and various bleached celluloses.

LITERATURE CITED

Abel, E., Xonatsh. , 79, 457-68 (1948).

Abel, E., 2. anorg. u. allgem. Chem., 263, 229-32 (1950). Anderson, V. S., Acta Chem. Scund., 2, 1-13 (1948). Ihzd., 4 , 914-21 (1950). Ibtd., pp. 1538-40.

Ibid. , 81, 685-8.955-64, 965-9 (1950).

Barb, W. G., and others, Nature, 163, 692 (1949). Barb, W. G., and others Trans. Faradau Soc., 47, 462-500

Evans, 11. G., George, Philip, and Uri, S., Ihid., 45, 230 (1949). Glassner, Abraham, J. Chem. Soc., 1950, p. 904. Pierron, Paul, Bull. S O C . chim. France, 16 ( 5 ) , 754-8 (1949).

Pierron, Paul, Compt. rend., 222, 1107-9 (19461. Uri, N., Chem. Reus., 50, 375 (1952). Weiss. J., Ann. Repts. P r o p . Chem., 44, 65-S (1947). Weiss. J., Ezperientiu, 7, 136-6 (1953). Weiss, J., and Humphrey, C . W., Sature , 163, 691 (1949).

(1951).

Ibid., 17 ( 5 ) , 291-3 (1950).

ACCEPTED July 20, 19.55. RECEIVED for review December 10, 1964. Contribution 341, Chemical Department, Experimental Station, E. I. du Pont de Nemours Q Co.

Documents

Stability of Dilute Alkaline Solutions of Hydrogen Peroxide