T H E THERMAL DECOMPOSITION O F HYDROGES PEROXIDE' BY F.
o. R I C E
ASD ORLASD M. REIFF*
The results of numerous investigations on the thermal decomposition of hydrogen peroxide indicate that the reaction takes place mainly in homogeneous solution and is only slightly catalysed by the walls of the reaction vessel. However practically all previous work deals not with pure aqueous solutions, but with solutions containing alkalies, salts, catalytic metals and preserving agents. In spite of the fact that hydrogen peroxide is a simple compound whose rate of decomposition can easily be measured, the values for the velocity constants, when calculated according to an unimolecular formula are characterised by large discrepancies although obtained under apparently identical conditions. For example, C l a y t ~ n in , ~ an extremely careful study of the thermal decomposition of dilute aqueous solutions of hydrogen peroxide, concluded that the purity of the Lvater is an important factor in the decomposition whereas Lemoine' concluded that the state of the surface of the containing vessel is the most important factor. In this paper we present the results of an investigation in which we paid special attention to the following points; ( I ) we ensured the absence of traces of inhibitors by preparing our own hydrogen peroxide according to a method5 already published; ( 2 ) we removed all colloidal matter and suspended dust particles as described6 in a previous article; (3) before making a velocity determination we melted the surface of the vessel to remove the frittering of the surface of glass and quartz that occurs on standing. Our results indicate that the ordinary decomposition of hydrogen peroxide is due mainly to suspended dust particles but is also caused to some extent by the surface of the vessel; any homogeneous decomposition of hydrogen peroxide is negligibly slow compared to the heterogeneous decomposition. The rate of decomposition of hydrogen peroxide was measured chiefly by the volume of gas evolved, but since this method is open to several criticisms the final results were always checked by experiments in which portions of the solution were withdrawn at measured time intervals, weighed and titrated with permanganate in the usual manner. The experiments were conducted usually in a Pyrex vessel surrounded by an outer jacket fitted with a reflux condenser so that the inner vessel could be kept at a given temperature by the vapor of a suitable liquid; usually the experiments were conducted at the temperature of boiling benzene. Contribution from the Chemical Department of Johns Hopkins University. Abstracted from the thesis of Orland 11.Reiff presented in partial fulfillment of the requirements for the degree of Doctor of Philosophy of the Johns Hopkine University. 3Clayton: Trans. Faraday Soc., 11, 164 (191j). Lenioine: J. Chim. phys., 12, I ( 1 9 1 4 ) , Kilpatrick, Reiff and Rice: J. .h Chem. Sol. , 48, 3019 (1926). Rice: J . Am. Chern. SOC.,48, 2099 ( 1 9 2 6 ) .
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I n previous work but little attention has been given to the quality of the hydrogen peroxide used; ordinary commercial solutions contain various substances as impurities, most commonly phosphoric, sulfuric and hydrochloric acids; the most common preservatives are uric acid, barbituric acid, acetanilide and quinine sulfate. Even the very best brands of hydrogen peroxide which are guaranteed free from added inhibitor, usually are supplied in paraffin bottles so that there is the possibility that the paraffin may supply a minute
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FIQ.I Decomposition of Merck’s Perhydrol, D.R.P. 216263,0.674M hydrogen peroxide. Temp. 802°C.
trace of organic preservative either through an actual impurity in the paraffin or by slight oxidation by the hydrogen peroxide. I n order to clear up the effect of inhibitors as far as possible we made a preliminary study of the rate of decomposition of various commercial products and also the rate of decomposition of our own pure hydrogen peroxide to which various inhibitors were added. Fig. I gives the curve showing the decomposition of solutions of Merck’s perhydrol. This curve is characterised by the occurrence of two induction periods. The type of curve was found to be the same regardless of the concentration; however the breaks in the curve are more pronounced in solutions of higher concentration. Exactly the same type of curve was obtained by either the gasometric or titration method, leaving beyond doubt the accuracy of the experimental procedure. It is interesting to note that Baker’s analysed hydrogen peroxide and Merckis German product S. 7456, while giving an initial period of inhibition do not give the second inflexion in the curve. Fig. 2 , Curve I , shows the type of curve obtained in the thermal decomposition of our own pure hydrogen peroxide to which was added any of
F. 0.RICE AND ORLASD M. REIFF
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the following inhibitors: barbituric acid, uric acid, benzamide, acetanilde, tannic acid or quinine sulfate. In these experiments pure 30% hydrogen peroxide prepared by us was added to ordinary distilled water and then the desired amount of inhibitor added. These curves which were all similar, are characterised by an init'ial fast period followed by a single inhibition. Acetanilide and tannic acid were found to be the least active preservatives, being
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FIG.z Decomposition of 0.6 bl hydrogen peroxide containing o.oo3C;'ouric acid. Temperature 80.zcC. Same solution restored t o the original strength by addition of pure 70yc hydrogen peroxide.
less than one third as effective as barbituric acid, benzamide and quinine sulfate. Fig. 2 , Curve z was obtained in the following way; the rate measurement shown in Fig. 2 , Curve I , was made and then the experiment was repeated with the same solution restored t o its original strength by the addition of the required amount of a 70% solution of pure hydrogen peroxide. This experiment was repeated with several other preservatives but in all cases there was complete absence of any inhibition indicating that during the course of any experiment the inhibitor is destroyed. The thermal decomposition of pure aqueous solutions of hydrogen peroxide free from both organic and inorganic inhibitors appears not to have been previously investigat'ed. Curve 2 , Fig. 3, shows the decomposition of a solution of hydrogen peroxide, made by dilution of the pure concentrated product with water made from a silica still. The curve follows a straight line for practically the entire decomposition and represents the minimum velocity obtained with dusty but alkaline free solutions. If however tap water is used not only is the rate of decomposition faster but the curve is no longer a straight line but approximates the curve for an uniniolecular decomposition (Fig. 3 , Curve I ) . The higher rate is due partly to the higher dust content
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but chiefly to the alkalinity of the tap water. We have found that when the peroxide solutions contain impurities such as chlorides, alkaline substances, inhibitors, etc the curves obtained are never straight lines and often approximate curves for unimolecular decompositions. Our next step was to reduce the dust content of the peroxide solution by the method described by Martin.' Unfortunately this method isnot too wellsuited
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FIQ.3 Decomposition of 0.6 R.1 hydrogen peroxide at 80.2"C. Curve I. Pure 70% hydrogen peroxide diluted with t a p water. Curve 2. Pure 70% hydrogen peroxide diluted with distilled water from a silica still.
for the removal of dust from peroxide solutions because the peroxide decomposes very slightly during distillation which causes ebullition to some extent; however we found that we could greatly diminish the dust content and by conducting experiments with this peroxide in a vessel of freshly fused silica we obtained very low rates of decomposition. I n one experiment 4 jcc. of a 2YC solution evolved oxygen a t the rate of 3cc. per hour a t 8o°C. Our whole experience indicates that complete removal of dust and conducting the experiment in a vessel with smooth walls would lead to a negligibly small rate of decomposition, We measured the rate of decomposition of a solution of hydrogen peroxide (free from inhibitors but containing dust) at 60" and 80' C and calculated for K80/k75 the value 3. Lemoine found for his solutions a value between z and 3. The absence of inhibition periods in the curves for the decomposition of pure solutions of hydrogen peroxide led us to investigate the reaction described by Bray and Caulkins* on the catalysis of hydrogen peroxide by the iodineMartin: J. Phys. Chem., 24,478 (1920). Bray and Caulkins: J. Am. Chem. SOC.,43, 1262 (1921).
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iodic acid couple. We were able to obtain results approximately similar to those of Bray and Caulkins by using various commercial samples of hydrogen peroxide; the shape of the curve depends however very much on the regularity of stirring, the rate of evolution of oxygen being very sensitive to this factor. The use of commercial solutions properly purified from inhibitors by distillation, or the use of our own pure hydrogen peroxide gave curves without a periodicity whereas the addition of inhibitors again caused the phenomenon of periodicity. I t seems probable therefore that the decomposition of hydrogen peroxide by the iodine-iodic acid couple is also a heterogeneous reaction similar to the thermal decomposition of hydrogen peroxide; the experimental difficulties however prevented us from removing the dust from these solutions and subjecting the matter to a rigorous test. summary
The decomposition of hydrogen peroxide in the presence of preservatives is characterised by an initial period of inhibition after which the material decomposes at a rate somewhat similar to an unimolecular reaction. 2. When hydrogen peroxide is prepared free from traces of organic inhibitors, chlorides, etc., its decomposition is approximately a zero order reaction. 3. When solutions of hydrogen peroxide are prepared free from suspended matter and in vessels with smooth walls, the rate of decomposition is exceedingly slow. The ordinary decomposition of hydrogen peroxide takes place on the surface of dust particles and on the walls of the vessel and an inhibitor evidently acts by poisoning the surface. 4. The decomposition of hydrogen peroxide by the iodine-iodic acid couple also appears to be a heterogeneous reaction occurring on the surface of dust particles. I.
Baltzmore, M d .
