<<

PROPERTIES OF PURE PEROXIDE DEPOSITED BY THE COMMITTEE < » N (Srafcuate Studies.

FS^QILL UNIVERSITY LIBRARY

THESIS.

THE PROPERTIES OF PURE .

by

Paul G. Hiebert.

Thesis presented in part fulfillment of the requirements for the degree of Master of Science.

May, 1922. McGill University, This investigation was suggested by Dr. 0. Haass under whose personal supervision the work was conducted. I wish to acknowlege also the kind assistance of Dr. W. H. Hatcher in the preparation of the hydrogen peroxide, and in the many helpful suggestions.

P.G.H. Of the wiany interesting properties of hydrogen peroxide,that which led most directly to the undertaking of this investigation, was the discovery of the fact that when dry hydrogen is passed into anhydrous hydrogen peroxide is liberated, even at low temperatures. This was unexpected, as it was known that in dilute solutions hydrochloric does not react with hydrogen peroxide, and is, in fact, recommended as a stabilizer to inhibit the spontaneous decomposition of hydrogen peroxide when kept in bottles. It was thought interesting therefore, to determine the concen­ trations at which chlorine would be given ofdin the first mentioned reaction, and the concentrations at which would act as stabilizer. Moreover it was thought that at the same time this might give an indication of the oxidizing power of hydrogen peroxide solutions of various strengths, so that a means might be obtained of standardising these solutions, as it were, by means of hydrochloric acid. The first experiments revealed the remarkable fact that even low concentrations of hydrochloric acid would bring about the decomposition of fairly dilute solutions of hydrogen peroxide, without however liberating chlorine. The investigation was therefore extended to cover a wider field its object being not only to determine the critical concentrations at which chlorine would be liberated, but also to throw some light if possible upon the mechanism of the reactions, both when chlorine 2 is given off and when decomposition occurs v/ithout the liberation of chlorine. The question involved is whether the decomposition of hydro­ gen peroxide in the presence of hydrochloric acid is due to either chlorine- or hydrogen-ion alone, or whether it is due to the hydrogen chloride . And furthermore,whether there"is in faci a critical concentration at which chlorine is given off, or whether there are secondary reactions continually taking place which allow the escape of chlorine only when certain concentrations have been reached or certain velocities attained. It was to find a solution to such questions that this study was undertaken. In the course of the investigation it was thought necessary to determine the react ion .velocities in order that the influence of different concentrations of hydrochloric acid might be compared. It was hoped that these would assist to some extent in disentangling the reaction. Consequently the tabulation of the reaction velocities has become an object of considerable importance in this reaearch as it developed. 3 II

In zhe analysis of the mixture; hydrogen peroxide, hydro­ and chlorine, which was expected to be obtained from the reaction of hydrogen peroxide and hydrochloric acid, it was first necessary to develop some simple and quick method whereby the course of the reaction might be followed. Volumetric methods appeared to be best suited to these requirements, but their develop ment was a matter of some difficulty owing to the presence of chlo­ rine interfering with the standard methods. Not only was it thought probable that the presence of chlorine would seriously ham­ per the accurate determination of hydrogen peroxide by means of standard potassium permanganate solution, but it would also inter­ fere with the determination of the hydrochloric acid present owing to its action upon the alkalies as well as upon the indicators. An attempt was therefore made to remove the chlorine from solution by bubbling air through the mixture after first diluting with . This was found to be impractical at room temperatures. The loss of chlorine from solution amounted to only 30$ of the total chlorine present after the bubbles had been drawn through for five hours at the rate of 102 per minute. A further attempt was made to take advantage of the bleaching action of chlorine upon dyes and indicators by standardising a so­ lution of a dye against a known quantity of chlorine. A dozen dif­ ferent dyes were tried, methyl orange being found, to be the most sensitive to the bleaching action. The method as worked out at the concentrations employed was found to be extremely sensitive when only hydrochloric acid was present, it being easily possible to estimate the amount of hydrochloric acid present to within 0.0000I gr. On the addition of hydrogen peroxide however, the blea­ ching action of the chlorine was so inhibited as to render the method useless for exact quantitative work. But it was still pos­ sible by means of comparison solutions to detect chlorine in quantities as small as 0.0004 gr., in the presence of hydrogen peroxide, and this qualitative test was adopted for detecting chlo­ rine during the course of the reaction. Attempts to remove the hydrochloric acid and part of the free chlorine from solution by precipitating as before titrating the hydrogen peroxide with permanganate failed to yield accurate results within 1.5$. A method of estimating the hydrogen peroxide as peroxide of one of the alkaline earths was devised. This method depends for its success on the insolubility of barium peroxide in neutral or slight- ly alkaline solution. To the mixture containing the hydrogen per­ oxide, hydrochloric acid, and chlorine, was added a slight excess of and the whole made alkaline with caustic soda. The precipitate of barium peroxide was filtered off, washed, and redissolved in sulphuric acid. The hydrogen peroxide thus regener­ ated and free from chlorine was then titrated with standard per­ manganate. Low results were obtained and it was found that the filtrate gave a further reaction with permanganate. The limit of accuracy from the titration of precipitate and filtrate together was 1.7$.

In order to obviate the possibility of soluble peroxides of sodium being formed, barium hydroxide was used instead of caustic 5 soda, but no better results were obtained. Numerous attempts were made to obtain greater accuracy. The use of phosphoric acid to dissolve the precipitate made an im­ provement over the use of sulphuric acid, but it was found that in every case a small quantity of the peroxide passed through the filter paper, even when the reagents were cooled in ice and the whole kept at a low temperature throughout. Precipitation of the peroxide as strontium peroxide or as calcium peroxide failed to improve the method. A feature of these attempts was the inconsistency of the re­ sults obtained, the limit of accuracy being 0.7$. But results with an error as great as 3$ were met with,even when the analysis was carried out with greatest precautions as to temperature and fil­ tration. The method was finally abandoned as untrustv/orthy, but not before one of the causes of the inconsistencies was discovered; viz., the varying of barium peroxide when freshly pre­ cipitated. To verify this interesting fact a sample of pure barium per­ oxide was prepared from specially purified hydrogen peroxide. To the hydrogen peroxide and barium chloride in molecular proportions, was added slowly and with constant stirring enough to pre­ cipitate one half of the maximum yield of barium peroxide. The precipitate was collected and washed until all traces of had disappeared. It was then shaken thoroughly with water and o filtered at 20 C. Measured volumes of the filtrate were titrated with standard permanganate. The filtrate gave a turbidity 0f 6

barium sulphate with sulphuric acid. The solubility of barium peroxide in the sample taken was found to be 0.186$ at 20°C. On being allowed to stand over night the solubility had decreased to 0.006$ at the same temperature, but on being thoroughly shaken the solubility increased slightly. This variation in the solubility of barium peroxide trhen freshly precipitated is in agreement with what is known regarding barium sulphate which also exhibits a varying solubility when the precpitated particles are finely divided, depending on the size i) of the division. Signs of decomposition of the barium peroxide were noticed after being allowed to stand over night, and this fact together with the varying solubility led to the abandonment of the method for exact analytical purposes. The method finally adopted war.-the direct titration of the hydrogen peroxide in the presence of manganese sulphate, phos­ phoric acid and sulphuric acid. By largely diluting the sample to be analysed, and by adding proper quantities of the above mentioned titrating solution, it was found possible to estimate the hydrogen peroxide to within 0.5$> the error being consistently low. Since however it was later discovered during the course of the reaction by means of the methyl orange test, that no chlorine was held in solution, the accuracy of the analysis for hydrogen peroxide was easily within 0.1$.

The experiments were carried out in a thermostat at 25°C. A 50 cc. flask was used for the reactions, and this was weighed I) Hulett, Zeit. "phys "5£eim77"377""i§01""" " — —- —— I 2 ffiiel 7

before and after the addition of about 10 gr. or hydrogen per­ oxide solution. It was then cooled to -I0°C. with a freezing mix­ ture of and ice, and while still in the freezing mixture the hydrochloric acid was added. A sample was then removed with a pipette and analysed, after which the flask was quickly warmed to the temperature of the thermostat, allowing the reaction to take place after the proper temperature had been reached. The time at which the reaction flasks was placed in the thermostat, and the time at which the samples were removed for analysis was carefully noted. When it was found that the reaction was proceeding with considerable velocity, the samples as removed were placed in a cooled weighing bottle and kept in a freezing mixture for several minutes. They were then weighed quickly and diluted with water before being analysed.

The method of adding the hydrochloric acid so as not to dilute the peroxide solution below the required concentration, involved the use of a specially built apparatus whereby pure and dry hydro­ gen chloride gas might conveniently be added in definite amounts. (Fig. I )

The hydrogen chloride was generated in the usual manner by dropping sulphuric acid upon a mixture of and hy­ drochloric acid, and was dried by passing through two sulphuric acid wash bottles. By raising or lowering the reservoir containing the large bulb could be filled to the mark with hydro­ chloric acid gas, and this in turn could be driven over through the capillary tube into the hydrogen peroxide solution. Each ti ime 8 after the bulb had been emptied, air was blown through the cal­ cium chloride tube so as to wash out the last traces of hydro­ gen chloride in the capillary tube. The bulb of about 100 cc. capacity was carefully calibrated by titrating the amount of hydrochloric acid driven over after each filling. Prom the weight of the acid so obtained the exact volume of the bulb could be calculated after correcting the volume of gas for prevailing temperature and pressure. The apparatus was made of glass throughout to prevent the possibility of leakages; all connections were glass sealed and the stop-cocks were made air tight with special vacuum grease. The whole was placed in a chamber free from draughts so that a con*. stant temperature might be maintained. In order to obtain greater accuracy in introducing the hydro­ chloric acid into the peroxide solution, special pipettes were made of required sizes. These were calibrated by means of mercury, and the solutions of acid, made up to a definite strength, were introduced by means of the pipette. The pipette was then washed and the washings titrated with standard alkali solution. The weight of the hydrochloric acid thus employed could be ascertained with great exactness, and this method proved especially valuable when working with a high concentration of peroxide and a low con­ centration of hydrochloric acid.

Where the concentrations of both peroxide and acid were low it was possible to weigh the acid solution after it had been added to the peroxide before any signs of reaction could be noticed, providing the mixture was first well cooled in a freezing bath ?

> F«cl and weighed before iile reac"ticR temperature had been reached. In all cases however it was essential that the reaction flask be kept well cooled while the hydrochloric acid v/as added slowly. For reasons mentioned later, a new method v/as developed for measuring the velocity of the reaction, besides the one referred to in which samples were removed and analysed. Not only could the course of the reaction be followed by the new method, but at the same time the reaction products could be collected and each con­ stituent accounted for when the reaction had proceeded as far as was deemed necessary. The apparatus for collecting the reaction products is shown;

(Fig. II)% The rate at which the hydrogen peroxide was decomposing could be determined by measuring the rate at which the resulting from the decomposition was given off. At the same time any water vapour or chlorine carried over with oxygen v/as caught in bulbs containing sulphuric acid and potassium hydroxide solution respectively. Later, however, a solution of potassium iodide was substituted for potassium hydroxide, since it was found that very little of the which sometimes accompanies the de­ composition of hydrogen peroxide, passed through the sulphuric acid bulb into the potassium iodide. The amount of chlorine thus liber­ ated could be directly estimated by titrating the iodine liberated with standard thiosulphate. The bulbs and the reaction flask were filled separately and sealed together after they had been filled. Finally the reaction flask was sealed so as to avoid all possibility of leakages.

The method of manipulating the apparatus for measuring the 10 rate of reaction was as follows; after adjusting the levelling reservoirs so that the pressure in the apparatus was that of the atmosphere, the stop-cock leading into the large oxygen reservoir v/as closed and at the same moment the oxygen was admitted to the gas burette. The number of minutes since the beginning of the reaction when the apparatus was first placed in the thermostat was noted. A convenient quantity of oxygen, generally about 30 cc. was collected, and the time taken to collect this amount was care­ fully measured with a stop-watch. The tap leading into the burette was again closed and the oxygen issuing from the flask again.ad­ mitted to the large reservoir. The oxygen in the burette was mea­ sured over water at atmospheric pressure and room temperature, and its equivalent in terms of hydrogen peroxide calculated. By lowering the large reservoir and opening the two-way stop-cock the gas burette could be again filled with water for another rea­ ding, while at the same time the oxygen in the burette was being drawn into the reservoir where it could be stored and later mea­ sured.

After-the reaction had proceeded as far as considered neces­ sary, the apparatus was cut apart and the residue and products analysed. The results agreed in every case with the calculated quantities within the limits of experimental error*

From the weight of oxygen collected per minute the of hydrogen peroxide decomposing per minute was obtained. A graph was drawn in which this number of grams of hydrogen per­ oxide decomposing per minute is plotted against the number of minutes since the beginning of the reaction. The latter numb >er II was counted from the time when the apparatus was placed in the thermostat to the moment when the oxygen is admitted to the burette, plus one half of the number of minutes necessary to collect a de­ finite volume of it. The error due to the change of rate of de­ composition while the oxygen was being collected, was thus constant for all the points obtained so that a smooth curve resulted. It was necessary, however, to collect as nearly as possible the same quantity of oxygen for each reading. From the curve so obtained the amount of peroxide that had decomposed could be computed by adding up the areas under the curve corresponding to a definite interval of time. From this the per­ centage of peroxide remaining in the flask could be calculated with considerable accuracy. Since the initial weight of peroxide, (y), is known, and the weight of peroxide which has decomposed, (z), is obtained from the graph, the percentage of hydrogen peroxide remaining in the flask is given by the formula; (y-s)IOO (y-z)4, 0.53(y-z)+C where C is the combined weight of the water and hydrochloric acid at the beginning of the reaction, and 0.?3(y-z) the weight of water formed from the hydrogen peroxide that has decomposed. This formula assumes that the loss in weight due to the oxi­ dation of hydrochloric acid and to the removal of water carried away with oxygen is negligible. This was found to be the case when the sulphuric acid bulb was weighed before and after the reaction. 12

Moreover, the calculated percentage agreed v/ell with the percen­ tage obtained upon analysis when the apparatus was taken apart. From the percentages the molecular concentrations corres­ ponding to them could be obtained from the formula; IOOOdp — - M 100x34 I) where M is the molecular concentration, d, the , and p, the percentage. This formula assumes that for low concentrations of hydrochloric acid the density is unaffected. Where the concen­ tration of hydrochloric acid was high compared to that of the per­ oxide a correction was made for the density of the hydrochloric ac id. The velocity constant, (k), for a monomolecular reaction is given by the expression; a log --«- = kt (a-x)

A series of values for k were worked out for every reaction in order that the change of velocity might be studied. These values will be dealt with under experimental data.

The hydrogen peroxide itself was purified from comercial 3$ solution by distillation in vacuo after being concentrated by means of the sulphuric acid concentrator described by 0. Maass.2^

Fifty litres of the commercial hydrogen peroxide were concentrated and distilled, the yield being about half a litre of pure concen­ trate, in samples ranging from 50$ to 85$ in strength.

1)0. Haass & VST. H. Hatcher T A p"q"""V"""77 — — ——— — mer, J.A.C.S., Dec. 1920. 2) Ibid. Every precaution was taken to prevent the decomposition of the hydrogen peroxide due to traces of foreign impurities. All containers, pipettes, or apparatus used to handle the peroxide were first allowed to stand in cleaning solution, followed by repeated washings with distilled v/ater.

14

Ill

The first experiment was carried out at the following concen­ trations; hydrogen peroxide. 50.2$f hydrochloric acid 9.27$. Reaction took place immediately upon removing the flask from the freezing mixture and placing it in the thermostat. Chlorine was given off in considerable quantities. Samples were removed with a pipette at convenient intervals for analysis, and the reaction curve so obtained was plotted. (Fig. Ill) The velocity constants as obtained from this curve approxi­ mated most closely to a monomolecular reaction but were not con­ vincing owing to their variation especially at the start, as the following table shows;

20 minutes .oo3l9 25 •1 .00325 75 •1 .00221 100 •• .00201 125 '1 .00202 200 •• .00149 275 •• .00149 400 •• .00118 550 •• .00109 800 1 .00102 1000 •• .00091 1200 •1 .00090 I?

The liberation of chlorine indicated that the concentration of hydrochloric acid for a 50$ solution of peroxide was higher than the critical concentration at which the minimum amount of chlorine would be given off. (See p-3o), A lower concentration of hydrochloric acid was tried ac­ cordingly, the concentrations in the second experiment being', hydrochloric acid 4.5$j hydrogen peroxide 51.1$. Ho chlorine was given off in this reaction, nor did the samples removed from the reaction flask give the test for chlorine with methyl orange. Moreover the velocity constants as obtained from the reaction curve indicated clearly that a monomolecular reaction had taken place, as is shov/n by the table of values for the constant k. Time. k 20 minutes .000335 40 " .000385 60 " .000418 80 " .000402 100 M .000398 150 " .000404 200 •• .000400 300 » .000381 500 " .000356 700 " .000329 The rate of decomposition is seen to be approximately one tenth as fast as in the first experiment where double the concen­ tration of hydrochloric was employed.

For the third experiment the concentrations were; hydrogen peroxide. 5l*9$» hydrochloric acid 1.74$. 16

Ho chlorine was liberated and the rate of decomposition was slow as the series of constants will indicate. Time k .0000120 20 minutes 40 ti .0000237 it 60 .0000382 70 » 100 11 .0000313 200 it .0000406 400 11 .0000505 700 11 .0000487 1000 11 2000 11 .0000502 2500 11 .0000442 3000 11 .0000518 .0000413 .0000431 Since it was suspected that the decomposition here might be due to a reaction between the peroxide and the glass of the flask, -the decomposition a blank experiment was carried out in order to determineAof a 50$ solution of hydrogen peroxide in glass vessels at 25 C. A gradual decomposition was observed. This decomposition was however slow in comparison to that taking place when the concen­ tration of hydrochloric acid was 1.74$. Moreover the samples of peroxide removed for analysis gave the us.al slight acid reaction with phenolphthalein, and not an alkaline reaction as would have been expected had appreciable amounts of the glass been attacked. These experiments then, revealed what had been expected, viz., that at strong concentrations hydrochloric acid is at­ tacked by hydrogen peroxide with the liberation of chlorine. But they also revealed the unexpected fact that at lower concentrations hydrochloric acid will bring about the decomposition of hydrogen peroxide without the liberation of chlorine. Moreover they showed that the rate of decomposition was related to the amount of hydro- 17 chloric acid present, and that this decomposition obeys the law for a monomolecular reaction, A series of values for the constant (k), were obtained in each case, and it is to be observed that in the case of the high concentrations of hydrochloric acid where chlorine is given off, these constants decrease in value, the de­ crease being greatest at the beginning of the reaction where most chlorine is given off. They then approach regularity for long in­ tervals. In those reactions whel*e no chlorine is given off, the value for k is regular throughout.

The next question to be considered was whether the decompo­ sition was due to hydrogen-ion, chlorine-ion, or the hydrochloric acid molecule. A run was accordingly made in which nitric was sub­ stituted for hydrochloric acid. The concentration of hydrogen-ion was made equal to that in the second experiment, that is, equi­ valent to a concentration of 4.5$ hydrochloric acid. The effect of the on 50$ hydrogen peroxide sor lution was the reverse of that corresponding to hydrochloric acid. It acted as a stabilizer. Samples removed at intervals for six hours showed no signs of decomposition whatsoever. It was obvious then, that the decomposition was not due to hydrogen-ion, but was apparently due to chlorine-ion or to the hydrogen chloride molecule. To decide which was the catalyst, was now added to the solution containing nitric acid by an amount suf­ ficient to give a hydrochloric acid concentration of 2.26$, the calculation being based on an equal distribution of the be- "7 18 {/ tween nitric and hydrochloric since the two acids are of the same strength. Decomposition at once occurred on the addition of potassium chloride. The velocity constant of this decomposition was .00014, a value which might have been anticipated since the concentration of hydrochloric acid lies between those of the se­ cond and third experiments for which the constants have been given. The concentration of the peroxide in this case was 40.85$« To distinguish between the effect of the hydrochloric acid molecule and the effect of the chlorine-ion, a new experiment was run in which potassium chloride was added to a neutral solution of hydrogen peroxide. The concentrations obtained were; hydrogen peroxide 47.1$, the concentration of chlorine-ion equivalent to 1.63$ of hydrochloric acid. The effect of the potassium chloride was neither to acceler­ ate nor to retard the decomposition. The rate remained the same as that of the natural decomposition of hydrogen peroxide in glass vessels. It was evident therefore that the decomposition is brought about by the hydrochloric acid molecule. The results of all these experiments described up to this point may be summed up as follows; hydrogen chloride cause the decomposition of hydrogen peroxide, the rate depending upon the hydrogen chloride concentration, roughly, in the order that doubling the concentration multiplies the rate of decompo-> sition by ten. The rate of decomposition corresponds to a mono­ molecular reaction, but when the concentration of hydrochloric acid reaches a certain critical value the reaction is no longer 19 a monomolecular one, chlorine being liberated. The reactior vhere- fore> is a series of consecutive reactions in which the mo...mole­ cular reaction is the slowest. The work was therefore continued with the original intention of finding out the critcal concentration at which chlorine would be given off, with the added interest that this might throw some light upon the nature of the consecutive reactions. So far the concentrations employed were all in the neighborhood of fifty percent. It remained to be determined to what extent the hydro­ chloric acid would affect the decomposition when the concentration of peroxide was higher or lov/er, and to what extent chlorine would be given off. Before beginning the third series of experiments whereby the amount of chlorine liberated could be determined and the reaction velocities measured, it was thought necessary to determine the effect of dry chlorine itself upon a fairly concentrated solution of hydrogen peroxide. Various tests upon the residue remaining in the reaction flask had in every case failed to show the presence of chlorine in solution. It appeared in every case, including those experiments in which chlorine had escaped as gas, that only hydrochloric acid remained.

An experiment was accordingly carried out to determine the effect of chlorine upon hydrogen peroxide. This was done by cooling a 45$ solution of the peroxide to -I7°C, then passing dry chlorine through it by means of a capillary tube until saturated. On being warmed to the temperature of the thermostat decomposition occurred 20

and the rate of this decomposition v/as measured in the usual man­ ner. The values obtained for the constant,k, were regular, as the table shows; Time k 100 minutes .000102 200 " .000127 300 H .000130 500 M .000125 700 • .000126 1000 ,f .000120 For hydrogen peroxide in the neighborhood of 50$, these values for k correspond to a concentration of hydrochloric acid between 1.74$ and 4.5$. The constants for the latter concentrations have been given and are on the average; .00004, and .00040 respec­ tively. Analysis of the reaction mixture after the reaction had proceeded for thirty minutes showed the presence of an acid cor~ responding to 3.6$ hydrochloric acid when titrated with standard alkali. In addition, the sample gave no test for chlorine after the reaction had proceeded for thirty minutes. It would appear then, that free chlorine i3 reduced by hydro­ gen peroxide to hydrochloric acid, and that the decomposition of the peroxide then proceeds as usual. The series of experiments was then begun in v/hich the velocity of the reaction is measured by the rate at which oxygen is given off, and the chlorine liberated is caught in the bulbs containing potassium iodide solution. A blank experiment was first carried out with the apparatus described, whereby the hydrogen peroxide was decomposed by the catalytic action of .

Traces of ozone or hydrogen peroxide vapour were carried over with 21 the oxygen and liberated iodine from the solution of potassium iodide, as evidenced by a faint yellow tint. The amount liberated after a long interval of time was exceedingly small so that this was not a source of error. After each run the residue was analysed so that all the pro­ ducts of the reaction could be accounted for, knowing the initial concentrat . Nine runs were made with the new apparatus, in the last eight of which the velocity of the reaction was measured. In the first of these a general balancing of the original constituents with the final products took place. The data so obtained are given in the following table: Concentration HO 42.5$. " HOI II.™

H Weight 202 in flask before reaction 5. 411 gr. » « « * after " 0.7937gr. " " recovered as oxygen 4.7005"

Total H202 recovered 5.4942"

Weight HC1 in flask before reaction; 1.665 gr. " " M after M 1.600? gr. M recovered as Cl2 in KI 0.0880 " Total HC1 recovered- 1.6885 gr.

Percent HC1 oxidised to Cl2 5*¥/o

In this experiment the hydrochloric acid remaining in the flask was titrated with silver nitrate by Mohr's method as well 22 as with standard solution, as a check upon the analysis. In those experiments in which the chlorine liberated and the reaction velocity were measured, it was considered unnecessary to check up the oxygen with the equivalent of hydrogen peroxide. In every case however, the mixtures involving chlorine and hydro­ chloric acid were carefully analysed and a check obtained as in the preceeding table. In the following tables the complete data as obtained during the course of one of these runs is given; Experiment A. a) Data obtained before placing in thermostat;

Concentration H202 , 75.64$. " HC1 3.41$. b) Data obtained during the reaction; Temperature of gas burette; 2I°C. Barometer; 764.5mm. Time at which apparatus was placed in thermostat; I2,05i pm (This time was taken as the zero minute.) In the first column is given the time at which the oxygen collection started. In the second column is given the number of minutes required to collect a definite volume of oxygen at the prevailing temperature and pressure. This volume is given in the third column. In the fourth column headed "Interval", is given the time in minutes from the zero minute to the time given in the first column. This interval plus one half of the number of minutes

23 required to collect a definite volume of oxygen is given under the column headed "Average Time." The last column headed"H202" repre­ sents the number of grams of hydrogen peroxide decomposing per minute, and is calculated from the number of grams of oxygen col­ lected per minute.

Time. Minutes. Volume. Interval, Av. Time. H„09 2 2. 12,07 pm 5.66 42.5 cc, 1.5 min. 4.38 min. .0208 gr. 6.86 11 9.0 12.43 .0171 12,14* 11 12,221 7.85 17.0 20.97 .0150 8.75 it 27.0 31.37 .0134 I2,j2i 9.55 11 37.0 41.77 .0123 12,42* 10.30 11 48.0 53.15 .0114 12,53* 11.28 11 62.0 67.64 .0104 11 1,07* 13.00 81.5 88.0 .0090 14.83 11 112.5 119.4 .0079 1,27 16.60 11 138.5 146.8 .0071 1,57 25.05 54.0 176.5 189. e .0060 2,24 22.44 44.0 211.5 •222.5 .0056 3,01 26.16 42.5 288.5 301.6 .0045 36.15 49.0 355.0 373.0 .0037 3,36 35.50 42.5 437.0 454.0 .0033 4,54 48.30 ti 647.0 671.O .0024 6,00 86.20 II 1260.0 1303.0 .0014 !24.0 M 1795-0 1857.0 7,22 .0009 10,52 From the data of the last two columns a graph is plotted as shown in Figure IV. It is from this graph that the values for the velocity constant, k, were obtained as described under experimen­ tal methods. In practice this graph was plotted upon a much larger scale so that the areas might be obtained with greater accuracy. c) Data obtained from graph. The column headed "x" represents the weight in grams of hydro- gen peroxide which has decomposed after a definite interval from the beginning of the reaction. This interval is given in the first column and is headed "Time". The percentages, and the molecular 24 concentrations, (a-x), are calculated from the formulae already given. The value for a, the initial molecular concentration of the hydrogen peroxide is calculated frpm the percentage of hydro­ gen peroxide originally introduced into the flask.

Time. X Percentage. (a-x) k 10 min. .2028gr. 74.51 28.55 .000943 20 .3690 73.56 28.10 .000819 30 .5105 72.61 27.60 .000806 40 • 6395 71.99 27.35 .000728 .7595 71.32 26.95 .OOO690 50 .9760 ^•28 26.30 .000645 70 1.1660 68.84 25.75 .000603 90 1.4180 67.25 25.05 .000552 120 1.6400 66.07 24.45 .000512 1.9550 63.82 23.50 .000421 150 2.2300 61.91 22.55 .000448 200 2.4650 60.39 21.75 .000425 250 2.6680 58.91 21.25 .000394 300 2.8470 57.69 20.75 .000370 3.0090 55.24 19.65 .000383 350 3.4490 53.27 18.85 .000316 400 3.8990 49.83 17.36 .000282 450 4. 5770 44.26 15.10 .000239 600 5.3070 38.31 12.70 .000301 800 1200 1800 It will be seen from the above table that there is a sudden drop in the value of the constants at the beginning of the reaction indicating that chlorine had been liberated. The next set of data verifies this. d) Data obtained by analysis after the reaction; Weight HC1 recovered in reaction flask .3541 sr Weight HC1 oxidised to CI .0380 gr! Total HOI recovered .4021 gr! Weight HC1 originally in flask .3980 gr.

HC1 oxidised to Cl2 9.75^* 25 The complete set of tables and graphs as given above was worked out for each of the eight reactions. In the following tables are given only the original concentrations, the amount of hydrochloric acid oxidised, and the velocity constants for the stated intervals. Experiment B.

Cone. H202; 73.96$. Cone. HC1; 2.02$.

HC1 oxidized to Cl2; .01$. Time. k. 20 min. .000231 40 .000213 60 .000247 100 .000228 140 .000225 200 .000213 300 .000194 400 .000188 500 .000176 700 .OOOI69 000 .000145

From the fact that a very small amount of chlorine was given off, the concentrations in this experiment were regarded as lying upon the critical concentration curve. (Fig. V). Experiment C.

Concentration H202. 46.1$. " HC1 ; 6.2I#. HC1 oxidized to CI ; none. Time. k. 40 min. .000705 80 .000665 120 .000677 160 .000559 200 .000528 280 .000486 400 .000460 600 .000414 800 .000372 1000 .000324 1200 .000345 1400 .000327

Experiment D.

Concentration H202 ; 32.9$. " HCl ; 10.35$. HCl oxidized to CI; .08$.

Time, k. 20 min. .00115 40 .00123 60 .00119 1080 .0009.001172 150 .00104 200 .00099 250 .00095 300 .00092 Experiment E.

Concentrat ion H202 ; 30.4$. II HCl ; 12.70$ HCl oxidized to CI ; 2.4$.. Time. k. 20 min. ,00624 40 „00320 60 .00292 80 .00269 100 .00252 120 .00256 140 .00243 160 .00233 180 .00229 200 .00223 240 .00212 260 .00214 280 .00200

Experiment F. Concentration HO ; 18.61$. " HCl ; 19.26$ HCl oxidized to CI; 14.1$ Time. k. 10 min. .00919 20 .01032 40 .00765 60 .00647 100 .00592 150 .00559 200 .00577

Experiment G.

Concentration H202 ; 10.95$. " HCl ; 1.8.84$ HCl oxidized to CI; I.$$ Time. k.

10 min. .00587 20 .00598 40 .00578 60 .00550 80 .00525 100 .00521 120 .00493 150 .00485 28

From the fact that there is no sudden drop in the series of values for the constant, k, and from the fact that the percentage of hydrochloric acid oxidized was small, it was assumed that the concentrations employed in this experiment were very near to the critical concentrations. Experiment H.

Concentration H202; 13.8$. » HCl; 13.25$. HCl oxidised to CI; none. Time. k. 10 rain. .00147 20 .00145 30 . 00144 40 .00142 80 .00147 120 .00139 150 .00128 200 .00717 250 .00107 300 .00118 Summary of Results

HCl ox'd. Cone. Ho0o. HCl. k k k k * d (20 min.) (40 min.) (60 min.) (Av'ge)

75.64$ 3.41$ .00082 .00073 .00067 .00042 9.7$. 73.96$ 2.02$ .00023 .00021 .00024 .00020 0.01$. Ct, 5L90$ 1.74$ .000012 .000023 .000035 .000040 5I.IO$ 4.52$ .00033 .00038 .00041 .00040 50.20$ 9.27$ .00319 .00350 .00275 .00218 Considerable. 46.10$ 6.21$ .00070 .00070 .00068 .00041 None. ct* 42.50$ 11.50$ 5.3$. 32.90$ 10.35$ .00115 .00123 .00119 .00095 0.08$. ct, 30.40$ 12.70$ .00624 .00320 .00292 .00209 2.4$. 18.61$ 19.26$ .01032 .00765 .00647 .00560 14.0$ * 10.95$ 18.84$ .00598 .00578 .00550 .00500 1.9$' Ct, 13.80$ 13*25$ .00145 .00142 .00140 .00140 None.

ro 6 7 If 7 40 Percentage HCl.

Fig. V

£ > 30 IV

A general summary of the significant results obtained is set forth in the accompanying table. From this table it will be seen that in those experiments in which chlorine is given off there is a marked drop in the value for the velocity constant within an in­ terval of sixty minutes, whereas in those in which there is little or no chlorine given off, the value for k is regular. The final ex­ periment, (H), of the series was carried out at concentrations well below the critical, and it will be seen that the variation of the value for k is slight, and thafc no hydrochloric acid was oxidized. It will be apparent also that the values for k vary greatly with the concentration of hydrochloric acid, as given roughly by the empirical formula; k^C, where C is the concentration of hydrochloric acid. It follows from this that if a certain concen­ tration of hydrochloric acid will bring about an appreciable de­ composition of hydrogen peroxide, then one half of that concer- intie %an\*- t***n« tration will bring about only one fifth of the decompositionA Small amounts of hydrochloric acid would therefore not appreciably affect the rate of decomposition of dilute solutions of hydro­ gen peroxide.

Four points have been selected from the table as lying upon the critical concentration curve. These points, (indicated in the table by Ct.,) were regarded as the concentrations at which the minimum amount of chlorine was given off.The curve obtained from them is given in Figure V. 31

It is apparent that below the critical concentrations the peroxide decomposes as a monomolecular reaction. The hydrochloric acid appears to act as a catalyser and remains behind unchanged. At concentrations above the critical, the reaction is obviously no longer a menomolecular one since chlorine is given off. The explanation is suggested that the following reactions take place;

2HC1 + H202 -»C12 • 2H20 ™(l), in which the hydrogen peroxide acts as an oxidizing agent, followed by a reaction in which it acts as a reducing agent;

2C1 • H202 —> 2HC1 • 02 -- (2)f so that the total reaction is;

H202 —>H20 • -|02 — (3). Reaction 2, may be regarded as considerably slower than reaction I, to account for the liberation of chlorine when the velocity of the first reaction is increased by a rise in con­ centration of hydrochloric acid. As however the velocity of the reactions is a monomolecular one, equations I, and 2, are not altogether appealing, moreover they do not explain the mechanism of the oxidation and reduction. An hypothesis is therefor tentatively suggested, but before doing so it might be well to recapitulate a number of previously mentioned facts brought out by this investigation. The hydrogen chloride acts as a molecule, the reaction being a monomolecular one, (page 18). Chlorine is acted upon by hydro­ gen peroxide, hydrochloric acid being formed, and here also the 32 reaction is monomolecular, (page 20). Hence reaction I, may be written;

H20o * 2HC1 —» H202.2HC1 ----- i 22 2 (A) followed by;

H202.2HC1 —> 2H20 • Cl2 .-- — ii Reaction 2, may be written;

H202 * Cl2 —> H202.C12 ----- i followed by; (B)

H202.C12 —> 2HC1 + 02 ----- ii where the breaking up of the complexes gives the monomolecular nature to the reaction in each case. How the assumption of the formation of these complexes is based on the formula, H-C^O-H, assigned to hydrogen peroxide due to its physical and chemical properties, and on the well known tendency of oxygen towards oxonium compound formation. This tendency is far more pronounced in the case of the hydrides as compared to the free , thus accounting for the re­ latively greater velocity of A, as compared to B. The evidence of 2) the simple oxonium compounds investigated so far, ' tends to show also that each of the in the H202, if the above formula is correct, will tend to hold in molecular complexity one halogen hydride, or one halogen .

I) 0. Maass and W. H. Hatcher. 2) Mcintosh and Maass, J.A.C.S., I912, andJ.A.C.S., I913. 33

From the point of view of the energy changes taking place during the above reactions, the results are in agreement with the experimental facts. A molecular complex can only break down in such a way that heat is liberated, and from the heats of formation and solution it can be shown that heat will be liberated in both A ii, and B ii. Turning to a consideration of the action of hydrobromic acid on hydrogen peroxide, Maass and Hatcher observed that from a dilute solution of hydrobromic acid, bromine is not liberated by a dilute hydrogen peroxide solution, but that here the critical concentrations are much lower than in the case of hydrochloric

acid. The complex, (H202.Br2), breaks up into HBr and 02 with the liberation of 10 large calories as compared to 32 large

calories in the case of the complex, (H202.C12).

In the case of(H202.I2), for a similar reaction to take place, heat would not be given out but 20 Cal. absorbed. Hence, although reaction A ii, takes place very readily, reaction B ii, is impossible, and hence HI, no matter how dilute, is broken up $iviV»3 I and this is in agreement with the experimental facts. In the case of hydrofluoric acid it can be calculated that

a complex, (H202.2HF), reacting according to A ii, is impossible, apart from the fact that HF will probably not form oxonium com­ pounds. Hence hydrofluoric acid will not cause the decomposition of hydrogen peroxide at any concentration.