going beyond. going further
edited by N. H. ETTINGER George Washington High Schwl B~stletonAve. and Verree Road Philadelphia, PA 19116
The Feasibility of Using Hydrogen Peroxide Decomposition Studies for High School Chemistry Gillian E. Carter 41 Oatley Park Ave., Oatley, Sydney, New South Wales, Australia 2223
This article seeks to highlight the difficulties that occur when teachers attempt to devise new experiments and how such seemingly useless results can be turned into a productive student proiect. The bleaching action in washing clothes involves the thermal decomposition of peroxide. The possibility of using this reaction to illustrate the principles of kinetics was considered as it had the advantage of being related to the students' environment. Washing is generallycarried out in tap water under alkaline conditions, and, thus, the method of determining the initial rates of this process was planned to be one in which the concentration of hydrogen peroxide, hydroxide ions, and ions in tap water were varied. A graph of the concentration of peroxide was then to he plotted against time. The unexpected difficulties outlined below suggest the experiment to be unsuitable for classroom use but ootentiallyborthwhile as a long-term project. T o benefit.from this project, the students should be aware of the components of washing powders (peroxohorates and peroxocarhonates). The need to reduce the number of variables will become quickl? apparent and u,ill eventually lead to the use of hydrugen pert,xide solurion. Rate cunstants f ~ rhe r thermal drrompusirion of hvdrogen peroxide were obtained in a period of 20 min when I mL of '"20 vol" hvdrucen oeroxidt: were added to 250 ml. tan water or 0.001 M sodium hydroxide at 75% (the molarit; of the hydrogen peroxide solution was ao~roximatelv 0.005). At .. inter\,ali 25-ml. portions were removed and ;he reaction quenched with 1 0 m l . ofU.1 A1 (tt'sulfuricacid.Thequrnched iamples were titratrd with 0.002 A1 potassium permanganate. I t mas inirially expected that the order of thr reacrinn rould be deduced. At this point, howwer, results obtained under apparently identical conditioni were found to be variable. This led to an extensive literature search. Phragmen in 1919 reported variations for reactions carried out in alkaline conditions (1).hut these have not ~ been~ generally mentioned by later akhors. I t is the study of the variables that provides an oooortunitv for detailed investigation as a senior high schooi student project. The literature search and my experimental results suggest that the main reason for such variable results were the presence of dust particles and the nature of the surface of the vessel. ~
Effect of Ions Present In Tap Water Thz rate of decomposition of hydrogen peroxide in tap
water has been reported t o be 50 times as fast as that of distilled water (21.l'here has been a great deal of discussion in rhe literatureabout which ion (or inns) is (are) responsible for catalyzing the decomposition of hydrogen peroxide. The most popular choices are the iron ions in the form of either Fez+ or Fe3+ (3). Mechanisms for the catalytic activity of both of these ions have been proposed (4). Reproducible results were obtained when EDTA was added to the tao water (.jJ;howe\er, the rate of decomposition in neutral and alkaline solutions becomes too slow to be easilv studird. The students can vary the concentration of the ions in tap water bv diluting the tao water with different orooortions of deionized water. singg glass vessels, I discovered that the rate of decomposition of hydrogen peroxide was proportional to the concentration of ions in tap water to a concentration of 50%tap water to distilled water. Increasing the proportion of tap water beyond 50% had no significant effect on the result. With polyethylene vessels, the rates were reduced, and it now appeared that the rate was proportional to the conce~~tratim uf ions in tap water over the range from distilled water to undiluted tap water. Effect ot pH Storing tap water to ensure that the concentration of ions remains constant can cause a decrease in rate because of the decrease in pH if absorption of carbon dioxide is allowed to take place (6). Preliminary boiling for several minutes to removed carbon dioxide can overco&e this problem. Pierron (7) found that the rate increased to a maximum and then decreased as the alkali concrntration increased. It has lwrn claimed that 'potassium hydroxide decomposes hydrogen peroxide more energetically than sodium hydroxide" (8). Polyethylene vessels give lower rates than glass vessels for decomposition in alkaline solution, an effect also noted with tap water. Effect of~ ~ u s~i ~ The possibility that dust might cause variable results was considered in the 1920's (9). Williams (10) took elaborate precautions to try to exclude dust from the reaction but found it impossible. He was, however, able to show that dust catalyzed the reaction. Pierron (7) found that the rate of decomposition was lower in a freshly prepared solution of hydrogen peroxide than in one that had been standing for some time. Under school conditions, the exclusion of dust would not be practicable. The group,of students must, thereVolume 63 Number 2
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The Rate Constant for the Decomposition of Hydrogen Peroxide in Glass Vessels: The Eflect ot Various Surface Treatments Nature of Treatment
Rate constant X lo's-'
HF followed by immediate use NaOH followed by immediate use HNO, followed by immediate use H3P04 followed by immediate use HF then exposed to air water for one or more days HF then exposed to air for one
+
or more days HF lhen evacuated or exposed to air free water for one or more days Previously used flask which had been HF treated only before first use HF followed by immediate use (one week before above measurernews) HF followed by immediate use (one week after above measurements
+
2.4 0.2. 2.7 i 0.1" 6 3= 5.3d
+
3.7 i 0.4*
8 & la 2.3
+ O.Zo
2.Sd
3.3
O.lb
fore. nlan the nroiect so that the kinetic measurements are completed within"a period of two days to avoid the accumulation of dust particles in the hydrogen peroxide. Effect of the Nature of the Vessel's Surface I t has heen claimed that thr reaction is catalyzed by the wall ofthereaction vessel, that thecatalvtiuactivitv deoends upon the chemical constitution of the "essel (11); and that the reaction is homogeneous in quartz. (12) and polyethylene (13) vessels. Williams (10) found that the reaction proceeds more rapidly on surfaces subject to cleaning treatments than on surfaces such as freshly formed glass. Adsorption of hydrogen peroxide on the surface wall and dust particles was assumed to cause the faster reaction. The literature contains a number of conflicting observations on the prediction of an increase or decrease in the rate of decomposition of gaseous hydrogen peroxide with different surface treatments. Pretreatments that have been reported include coating the glass with wax (10, l l ) , washing it with ammonia (141, hydrofluoric acid (151, other inorganic and chromic acids (14, 161, and prewashing with lead(I1) oxide (15). My results for hydrogen peroxide solution are given in the table. The experimental procedure was as previously described but reactions were carried out in 0.01 M sodium hydroxide and followed for a period of 1h. The same source of hvdroeen was used throuehout the exneriment. " - neroxide . " Rate constants were calculated in each case using results obtained after the first 10 min of reaction when first-order behavior had been established. My observations using hydrofluoric acid and inorganic acids appear to contradict the literature results for gaseous hydrogen peroxide (16, 17). However, my procedures were carried out in alkaline solution so any acid used in surface treatment would be removed prior to reaction, thus creating a different situation. I found that washing the vessel thoroughly with a solution of hydrofluoric acid or hot concentrated sodium hydroxide provided the most reproducible results. All the results can be explained by the catalytic action of oxveen adsorbed on the surface of the vessel. The " greater the amount of adsorption, the greater the rate constant. Treating with hydrofluoric acid or sodium hydroxide would remove the surface layer with its adsorbed oxygen. Journal of Chemical Education
Reaction Order Thermal decomposition of aqueous hydrogen peroxide occurs bv a free radical mechanism in which the slow nrocess is the initial decomposition of hydrogen peroxide into bydroxyl radicals (20)
1.7 f 0.2'
Average of nine readings. 'Average of two readings. =Averageolmrsereadings. lmividual values were 2.1.4.55, and 10.1 X 10-'r-'. The seoond valve was sner further treamsnt with nitric acid and me mird aner treating with hof concentrated acid. reading. .Average of five readings
160
I t has been suggested that a final cleaning with concentrated hydrogen peroxide be employed in the preparation of glassware used for the study of the decomposition of gaseous hydrogen peroxide (18). Complete oxidation of the surface may be as effective in giving reproducible results as the complete removal of oxygen. The simple treatments discussed in this paper should be preferable to complex procedures such as coating the vessel with boric acid and then fusing in a furnace a t 500°C overnight prior to use (19).
Although this suggests that the reaction is first-order with respect to hydrogen peroxide, the rate constant has been described by the same author as second-order (21). Baysul (22) established the rate law Rate u [H202][Fe2+]
for the Fe"-catalyzed decomposition of aqueous hydrogen ~eroxide.Hv followina the decom~ositionfor a Deriod of 70% bf the reaction and applying statistical techniques, I ohserved that the decomposition carried out in tap water (pH = 7) is first-order with respect to the hydrogen peroxide concentration. However, the possibility of fractional-order kinetics has not been ruled out. The literature contains many theories on the mechanism of hvdroeen decomnosition in an alkaline envirou. - neroxide . ment. The experimental curve I obtained when the concentration of hvdroeen neroxide was dotted aeainst time is consistent with tKe thkory of ~ o r a b i i l s k aand ~olodziejezak (23). It is assumed that the reaction is between Hz02 and 02H-. If the initial concentriltion of undissociated hydrogen peroxide (a) is less than that of 0&, (b) the reaction is given by
where r is the quantity of hydrogen peroxide decomposed. Hawever, if a > b, then there is an autocatalytic reaction due to OHions (initial concentration e ) given by
The initial rate is proportional to the hydroxide concentration, and when the hvdropen ~ e r o x i d econcentration becomes sufficiet~tlylow (ahout ;en minutes for the euperimentalconditionsdesrribedj, thereartion is first-order with respect to hydrogen peroxide. Safety Precautions Glass beads should he used t o assist in even boiling to rid distilled water of dissolved carbon dioxide. Ammonia, hydrofluoric acid, nitric acid, and chromic acid are all hazardous. Skin contact andlor inhalation must be avoided and safety goggles should be worn. A fume hood is recommended. Skin contact with sodium hydroxide, sulfuric acid and hydrogen peroxide should also be avoided. Acknowledgment The author wishes to express her gratitude for helpful comments by Professor Frazer of the University of East Anglia and Drs. Marshall, Steadman, and Wellington of . University College, Swansea.
Literature Clted Phragmen. G. Medd. K . Vetenshapsaked. Nobel-lnsl. 1919.22,l. Clayton. W. Tmns.Fomdoy Soc. 1915,2,1M. Haber. F.: Weiss. J.Pmc. Roy. Soc. A 1934,147,332 119341. Gardiner, W. C., Jr. "Fates and Mechanisms of Chemical Reactions"; W. A. Benjamin: London. 1969. p 9. I51 Koubek, E.: Haggeft. M. L.; Battsglia, C. S.; Ibhe-Kauss, K. M.; Pyun, H. Y.; Edwards, J. 0.J.Amar. Chem. Soc. 1963.85.2263. (61 Mongeot,A.; Aubertot, V. Compt. rendaoc, b i d . , 1928.98,9(15. (71 Pierron, P. Compt. Rend. l916,72,1107. (8) Teletov, J. S.; Griatan, D. N. Zopiski Kharkou, Solghokhos. Inat. 1933.1-2.22. (91 Rice. F. 0.: Kilpetrick. M. L. J . Phys. Chem. 1927.31.1507. (10) Williamr. B. H. T m n s . Fomdoy Soe. 1928.24.245. (111 Maelw-Hughes. E. A."The Kineticsof F a a d o m i n Solution"; Clarendon: Oxford, 1917;p 33. I11 I21 (3) I41
(12) Robertson, A. C. J . Amar Chem. Soc., 1931,53,382. 113) Duke,F.R.: Ha-, T. W. J . Phy~.Chom.,1961.63.304. (14) Linnoft. J. W. In "Seventh Symposium on Cambustian"; Butternorth London. ,469
(161 Hoace, D. E.: Waloh. A. D. In "Fifth Symposium on Combustion": Reinhold: New York. 1955. 1161 Giu&rc,P.A. Con J.Re8.. 1947,25, 135. (17) Cheeney, D. E.; Walsh, A. D. Fuel, 1956.35.258. (18) Schubb, W. C.; Satterfield, C. N.: Wentworth. R. L. "Hydrogen Peroxide": Reinhold: New York. 1955. (191 McLane,C. K. J . Chem.Phys. 1349.17.379. (20) Laidtor, K. J. "RcadionKinetics: Vol. 2. Reactions in So1ution";Peqamon: London,
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