PART I. THE INHIBITION BY DIOXANE OF THE ALKYLATION OF PHENOL WITH TERTIARY- BUTYL CHLORID E. PART II. THE SOLUBILITY OF HYDROGEN CHLORIDE IN PHENOL-DIOXANE MIXTURES Thesis for the Dogma o! M. S. MICHIGAN STATE COLLEGE .Iahn Junfior Bordeaux I949 This is to certify that the thesis entitled Part I. The Inhibition by Dioxane of the Alkylation of ‘ Phenol with tert-Butyl Chloride. Part II. The Solubility of Hydrogen Chloride in Phenol—Dioxane Mixtures. presented by John Junior Bordeaux has been accepted towards fulfillment of the requirements for H.§.__ degree in Jhemiatry Major profesgor ‘ Date W PART I. THE INHIBITION BY DIOXANE OF THE ALKYLATION OF PHEIOL WITH TERTIARY-BUTTE CHLORIDE. PART II. THE SOLUBILITY OF HYDROGHT CHLORIDE IN PHE‘TOIFDIOXMIE MIXTURES By John Junior Bordeaux A THESIS Submitted to the School of Graduate Studies of Michigan State College of Agriculture and Applied Science in partial fulfillment of the requirements for the degree of MASTER OF SCI EIJCE Department of Chemistry 1949 CHEMISTRY W?" 775‘}? @737 ACKI-IGNLEDGMENT The writer wishes to express his sincere appreciation to Doctor Harold Hart for his aid and guidance through the course of this inves- tigation. *IIIMIKI‘IHHIIIIII ******** *lll**** **** *1! * 218-387 PART PART C ONT HITS I INTRODUCTION.......................................... HISTORICAL............................................ FJZPERIMENTAL............-.............................. I. Materials.................................... II. Apparatu8.................................... III. methOd Of Measuring Rateooeooooooeooeoooooooo RESIILTSOOOOOIO......OOOOOOOOOCOO......OOOOOOOOOO...... I. mtaOCOOOO.........OOOOOOO......OOOOOCCOOOOOO II. Treatment Of Dataooooooooooooooooooeooooooooo DISCUSSIONoooooooooooooOoooooooooooooooooooooooooooooo SUEUE'AARYOO......OOOOOOOOOOOOOOOOO......OOOOOOOOOOOOOOCO BIBLIOGEMPHYOO......O0.0.000.........OOOOOCOOO00...... II II‘ITRODUCTIOIqOOOOOOOOO0.00.00.........OOOOOOOOOOOOO.... m£RI}'1EI\ITAI’OOOOOOOO...0............OOOOOCOOOOOOOOOOOO I. IdateriaISOOOOOOOOO............OOOOOOOIOOOOOO‘ II. Measurement Procedure" .............. ..... ... RESULTS AND CONCHISIONS............................... I. mtaOOOOOOO......OOOOOOOOOOOOOOOOOOOOOOOOOOOQ II. DiSCuSSion Of ROSUItSoooecocoa-000.0...one... BIBLIOGWIIY...0....00.0.00.................O......... Page co-q 11 11 ll 20 24 25 26 27 27 27 29 29 29 32 PART I THE INHIBITION BY DIOXANE OF THE ALKYLATICN OF PHENOL WITH TERTIARY-RUTH CHLORIDE INTRODUCTION The uncatalyzed alkylation of phenol in the para position using tertiary-alkyl halides was first reported by van Alphen (1) with fur- ther work by Bennett and Reynolds (2) and later by Simons and Hart (3). The kinetics and mechanism of this reaction was recently studied by these latter investigators (4) whose work included a study of the ef- fect of various diluents on the rate of alkylation. They observed that the alkylation was inhibited completely by 1,4-dioxane (herein- after referred to as dioxane) when this diluent was used. When para- xylene was used in comparable amounts, the observed decrease was only that calculated for the diminution in phenol concentration. It was postulated that the inhibitory effect of dioxane was due to the form- tion of an oxonium compound between dioxane and phenol. It was the purpose of this investigation to study the inhibitory property of dioxane in the alkylation of phenol with tertiary butyl chloride; to determine the stoichiometry and, if possible, the mechan- ism of the inhibition. -1- HISTORICAL Since oxonium compounds probably play an important role in the mechanism of the inhibitory property of dioxane in the alkylation of phenol, it is well to briefly review the literature of such compounds. The formation of addition compounds of simple ethers with hydro- gen halides was discovered by Friedel in 1875 (5); these compounds and others similar to them were called coconium salts. Two structures for these oxonium salts have been proposed (6), CK /C1 /°\ H [(CH3)20HHH]+ 01" I II Structure I gives rise to a tetrattalent oxygen atom; II shows the for- mation of a salt-like compound which is analogous to the ammonium salts. The anion is not attached by a covalent bond. The Ypyrones were the first definitely established compounds which formed ozonium salts. They formed crystalline compounds with acids whose structures might be can- pared to those formed by such nitrogen containing compounds as the pyri- dines : I. /CH’C{ I" F + av“ 0:0 0-H c1“ -s of / I. “Hm .I II V pyrone Salt pyridine salt Cineole has been reported to form.similar addition compounds with acidic substances indicating thefibasic character of the oxygen atom in the ring. Bellucci and Grassi (7) show evidence of the formation of oxonium.compounds of cineole with phenol, certain substituted phenols, oc- and P-naphthols. Their data shows the formation of addition com- pounds containing one molecule of cineole to one molecule of phenol, indicating the following formula: F’ + 3 H2 l H2 H2 H30- --CH3 H2 L H From the work of Morgan and Pettet (8), the analogies between axon- r. ium.and ammonium.compounds are further strengthened. They confirm the work of Bellucci and Grassi and also show the formation of an addition compound of o-cresol with both mp5-xylidine and cineole. NHZ H3 H3 mp5-xylidine The study of the formation of salts wdth dioxane has been of some interest and according to‘Whitmore (9), it will form.oxonium compounds in a manner similar to these compounds previously discussed. Examina- tion of the structure of dioxane leads one to believe that it might form a compound containing one molecule of dioxane to two molecules of the acid since it has two oxygens available for combination: <3} Faworski (10) and Smeets (11) have shown this to be true with sulfuric and perchloric acids respectively. Recently Beer (12) has prepared dioxane di phosphate: mm as oPo(0H)2 His discussion of the compound stated that it was soluble in certain organic solvents, indicating that the bonds formed by the addition of the acid were of a covalent nature. Intramolecular hydrogenbonding was studied by Batuev (13) using infra-red adsorption spectra, and in this connection, the effect of dioxane on the chelate ring formed in o-chlorophenol: -4- It is pointed out in this paper that the addition of an equivalmt amount of dioxane to a solution of o-chlorophenol in carbongtetra- chloride tends to disrupt the formation of the chelate ring, probably due to the attraction for the hydrogen of the phenol by the oxygen atoms of the dioxane: 032- Hz Cl CH H2 Bartlett and Dauben (14) have evidence to support the formation of oxonium compounds between diozane and phenol. They discussed the increase in the acid strengthof hydrogen chloride in dioxane brought about by phenols and alcohols. The formation of a complex between the phenol and the hydrogen chloride in which there was a hydrogen-bond formed by the hydrogen of the hydroxyl group and the chlorine was postulated: They pointed out that the basic strmgth of dioxane was sufficient to assure them the hydrogen chloride was entirely combined with the sol- vent, as dioxanonium chloride: 7" Hznc _I++ H 201 The formation of these two compounds offers an explanation of the greater dissociation of hydrogen chloride because it can be supposed that there is a pull on both of the atoms of the hydrogen chloride. The same basic characteristic of dioxane may cause it to be combined with phenol in mixtures of these compounds. -6- EXPEIMBITAL I. materials The phenol was J. T. Baker Chemical Company material, which was distilled with benzene to remove the water. The fraction boiling from 177 to 184° was taken. This was in turn fractionated, retaining the sample from.lBO to 182°. This was then.stored under nitrogen.until ready for use. The tertiary butyl chloride was prepared from Eastman Kodak Comp pany tertiary butyl alcohol by shaking portions of it with concentrated hydrochloric acid in a separatory funnel and separating the organic layer from the water layer (15). The tertiary butyl chloride thus pre- pared was distilled keeping the fraction boiling between 49 and 53.5°. This was further fractionated to obtain a sample which boiled at 51 to 52°, at atmospheric pressure. The dioxane, an batman Kodak product, was purified by the method described by Fieser (16). 200 m1. of dioxane was refluxed with 2.7 ml. of concentrated hydrochloric acid and 20:m1. of water for twelve hours, the solution.was cooled and saturated with potassium.hydroxide pellets. Therlayers were separated, and the dioxane was treated with fresh potas- sium.hydroxide pellets, decanted and refluxed with sodium for twelve hours. The dioxane was then distilled from the sodium and stored over fresh sodium. The entire purification was carried out under an.atmes- phere of nitrogen and the purified dioxane was stored in a nitrogen atmosphere. II. Apparatus The apparatus used is shown in figure 1. It consisted of a 125 ml. reaction flask, A, having a gas inlet tube, B, extending to the bottom. In the early work, the flask was sealed at C to a condenser, D, fitted at the top with a 14/35 male ground glass joint, 3. Later the flask was fitted with a 24/40 female ground glass joint at C and the condenser with the same sized male joint. The condenser was attach- ed to a tube, F, with athrtvmy stopcock, G, whose outlets were fitted with 14/35 male ground glass joints, H, to which sodium bicarbonate tubes, I, were connected. The sodium bicarbonate tubes were connected to trap vessels, J, which were kept in a dry ice-acetone bath. The system was stirred with a magnetic stirring device and a continuous stream of dried nitrogen gas was bubbled through the mixture at B. The reaction flask, A, was submerged in a constant temperature bath main- tained at 50.0 1. o.1° during the experiments. III. Method of MeasuringRate The rate of the reaction was followed by converting the evolved hydrogen chloride to sodium chloride and determining the amount of sod- ium chloride formed at definite time intervals. ‘Weigh‘ed; amounts of dioxane and phenol were thoroughly mixed in the reaction flask. In the early stages of this work, the tertiary butyl chloride was added to the reaction mixture by pipetting it down through the condenser. It was later thought that this method of addi- tion might not be satisfactory because of the possibility that tertiary butyl chloride‘might be lost by volatilization during the process of -8- nLerlllmyILwJ c 18 h Fit. :- l'E' tus soaring- zip; 130:; l”! “'0 4 I I; pipetting. This proved to be incorrect since the results were dupli- cated by a different method. The other method of adding the tertiary butyl chloride was to seal a weighed amount in a thin walled ampoule and break it in the re- action flask with the system closed. This method proved highly suc- cessful since, by using a weighed amount of the tertiary butyl chloride, it was possible to calculate the exact number of moles that were added. In the first method of addition, it was necessary to assume that the total amount of the halide in the pipette was added to the solution. The reaction was followed by the hydrogen chloride given off dur- ing the reaction. The hydrogen chloride, when passed through the sodium.bioarbonate tube, reacted to form.sodium.chloride, water and car- bon dioxide. The contents of the tube were then dissolved,:made acid with nitric acid, and the amount of chloride determined by the Volhard method (17). Since it was believed that small amounts of tertiary butyl chloride might distil over during the reaction, a trap cooled by a dry ice-acetone bath.was employed. The trap was rinsed with absolute alco- hol several times; this solution was then diluted with.water and acidi- fied with nitric acid and the chloride determined as before. The total chloride carried over in any one experiment was always less than 10 per cent of the original tertiary butyl chloride used. The values in column 4 of Table I represent values which have been corrected for the small amount of tertiary butyl chloride carried out of the reaction flask by the nitrogen and evolved hydrogen chloride. -10- REULTS 6 I. Data A sample of the data taken during experiment number six is shown in Table I. Space prevents recording the detailed data for each run, but the essential data for all experiments is summarized in Table II. The experimentally determined first order rate constants (k) and the half-times (tlze) are given in columns (6) and (7) respectively. II. Treatment of Data The data obtained ms in the form of moles of hydrogen chloride evolved (column (3), Table I). Since this was equivalent to the moles of tertiary butyl chloride that had reacted, it was subtracted from the number of moles added at the start of the reaction to give the num- ber of moles of tertiary butyl chloride remaining (column (4), Table I). The volumes of the reactants were assumed to be additive and by divid- ing the number of moles of tertiary butyl chloride remaining by the total volume of solution the concentration at the end of each time in- terval was obtained (column (5), Table I). Since the reaction was de- termined to be first order with respect to the halide (4), the minus log of the concentration was plotted against the time (see figure 2). From this graph, the half-time of the reaction was determined (column (7), Table II). The rate constant, which is the slope of the line was calculated (column (6), Table II). The best line was drawn thraigh the experimental points using the method of least squares (18). -11.. (1) (2) Sampl e Time No. 0 1 10 (min.) per Time Interval O 20 4O 60 80 110 150 200 260 320 380 (3) (4) Moles of Moles of HCl EvOlved t-Butyl Chloride Unreacted 0.000000 0.011261 0.000283 0.010978 0.000561 0.010417 0.000606 0.009811 0.001145 0.008666 0.000982 0.007684 0.001425 0.006259 0.001302 0.004955 0.001124 0.003833 0.000823 0.003010 0.001022 0.001988 Table I -12.- Sample of Data from Ekperiment-#6. (5) Concentration of t-Butyl Chloride in moles/ liter 0.4941 0.4815 0.4569 0.4303 0.3801 0.3370 0.2745 0.2173 0.1681 0.1320 0.0872 (1) (2) t-Bu Cl Exp Moles 1 2 10 0.009872 0.009189 0.009745 0.010034 0.009983 0.011261 0.012144 0.010823 0.011748 0.011149 (3) Phenol Moles 0.2172 0.2090 0.2163 0.2184 0.2179 0.2180 0.2073 0.2185 0.2188 0.2167 Table II Summary of Rate Data (4) Dioxane Moles 0.00000 0.00000 0.00615 0.01174 0.01190 0.02315 0.02362 0.03476 0.03538 0.05774 (5) Molarity of (a) Pheno1 10.61 10.58 10.25 10.00 h 9.99 9.56 9.40 9.15 9.09 8.40 (6) k 0.011017 0.010580 0.005258 0.003679 0.004071 0.001940 0.002010 0.000699 0.000688 0.000080 (7) té. (min) 53.3 55.0 79.8’ 104.0 108.1 / 174.2 164.0 / 429.3 435.0 ’ 8663.5(b) (a) These molarities were calculated on the basis of dioxane acting only as a diluent. (b) This value was estimated using the rate constant k. I I I I l l l l l I l .- J 3. ." 1.5 —- 4 -' 1.5 — 0" ‘ lutr- . '— -L08 0 - 4 .9J— . . 5 -‘ 0' .7r— 0 ——1 0 ° 0 C ‘ .5 '— 0 ° /' e '. 0 ° + 1.") ' .0 120 400 43fi9 jhwv time in minutes Plf'. L.) ‘. a r: ,. . , ._ .13 .2“, _, a, , 4‘ 1. . minus if; c 99. 01.0 101 oxy011.Lnts 1 and 1C. (1:1 0 11) 3": L1) 0: U9 The expression in which the time, t, necessary to complete any fraction of the reaction is proportional to n' the total order of the reaction is shown by the following equation (19): t ac 1/an""1 (where a is the initial concentration of reactants. Thus from any two different experiments it is seen that tl/tz : (ea/a1)n"1 Since the reaction is first order with respect to the halide, the half- time is indepmdent of the tertiary halide concentration. Knowing that n' is the total order of the reaction, n'-1 may be taken as the order of the reaction with respect to phenol. 11 may be used to represent n'-1. Now using M, the molarities, to express the concentrations, a, the follow- ing equation is obtained: tie/t; .-.- (won (1) where tg; is the half-time for the reaction with no dioxane. present, tg'é. is the half-time for the particular'reaction, Mo is the molarity of the phenol in the no dioxane experiment, and M is the molarity being studied. The reference values are 1%? -_- 53.5, and 11° :- 10.61 at 50°. It is possible, using equation (1), to calculate the half-time t%_ which one would expect under varying assumptions concerning the effect of dioxane on the reaction. These values are given in Table III, and their significance is discussed below. . The values for 1‘31; (column (3), Table III) were calculated on the assumption that dioxane acts only to decrease the phenol concentration by diluting the solution. As an example, data from experiment eight is taken: -15- Table III (1)_ (2) (3) (4) (5) (6) Ekp. t1. ti? tlb t1c Moles of Dioxane 2 2 '5 '§ Per Mole of Phenol 1 53.3 2 55.0 3 79.8 65.6 78.1 92.9 0.02843 4 104.0 76.0 106.1 148.9 0.05394 5 108.1 81.9 108.0 153.7v’ 0.05461 6 174.2 97.0 195.6 417.1 0.10601 7 164.0 110.2 227.6 518.3/ 0.11408 8 429.3 129.6 367.7 1286.7 0.15908 9 435.0 134.8 396.8 1403.4 J 0.16170 10 8663.5(a) 216.4 3089.4 0.26645 (a) See footnote (b) page 13. a '1‘: 10.61 6 .3 9.15 a , . t3; = 129.6 minutes I“, H M b 5 O} The value of six was used for 11 because Hart and Simons (4) pointed out that the apparent order of the reaction with respect to phenol was fifth to sixth. b (column (4), Table III) were calculated by assum- The values for t: ing that one mole of dioxane would combined with one mole of phenol, and thus decrease the phenol concentration. Using the same experiment as before for the example, t; was found to be: II 2 : , t__1_ :1 367.7 minutes 53.3 7.69" 2 The t}; values (column (5), Table III) were obtained in the same way byassuming that each mole of dioxane combined with two moles of phenol. Again experiment eight will be used as an example. 0 t; 10.61 6 III 2 = ‘ , t3}: 1286.7 minutes 53.3 "—6.24 2 A plot of these ti]? values against the moles of dioxane per mole of phenol is seen in figure 3. These curves may be compared with the curve forthe experimmtal half-times (column (2), Table III). The same type of calculations could be carried out using five as the 11 value. For comparison with the value of tarp: calculated in I, the same concentration will be used, but the order of the reaction will be taken as fifth order: ha 10 61 5 N “F" -.-. ' , tgf = 111.8 minutes 53.3 "9.I5 2 1300 200 [00 mol ratio ---—— t1 —---—- t1“ 2 2 1, i1 _ e - ‘L- Li tr! ”’9'. 1‘1 Q tp vs. moi ratio of dioxzne to phenol. (Table III) It is readily seen that this last value calculated is less than the corresponding value when n is taken as equal to six, thus magnifying the inhibitory effect. It is unlikely that n could be greater than six. -19- DISCUSSION The kinetics and mechanism of the uncatalyzed alkylation of phenol with tertiary butyl chloride has been investigated by Hart and Simons (4). As a result of their studies, in which they found that tertiary butyl chloride could be used to alkylate phenol quantitatively at convenient temperatures and with a measurable rate, they proposed a mechanism for the reaction. This mechanism is shown schematically below: ——H°"' The solid lines are the bonds before the reaction takes place and the dotted lines show what occurs in the reaction and the new bonds formed. The reaction showed an apparent order with respect to phenol of about five or six. They pointed out that the remainder of the phenol mole- cules involvedtin the transition state may be attached to l and 3 in a linear fashion, or that several phmol molecules may be involved in a cluster, solvating both the halogen and the displaced hydrogen. It is important that the bonds be made and broken simltaneously. It is also necessary that phenol act both as an acid (molecule 1) and as a base (molecule 3). Thus, this mechanism has been said to involve the "ampho- teric medium effect" (3). They also point out that the addition of dioxane in an amount of about 21.7 mol % inhibited the reaction almost completely. In the present work it was the purpose to show to what extent dioxane affected the reaction and to draw some conclusions as to why it -20.. affected the reaction as it did. It was previously shown (see Historical) that dioxane'was basic in nature and would form oxonium compounds with acidic substances such as phenol. It is believed that the formation of an oxonium compound is the factor causing the inhibition. As was pointed out by Hart and Simons (4), the ability of the reaction to proceed at a reasonablerrate was dependent upon a large excess of phenol being present. Any decrease in the phenol concentration caused a large decrease in the rate. Therefore, the addition of dioxane would be expected to decrease the rate. The results show this to be correct, but the decrease is great- er than would be expected if dioxane ware acting only as a diluent. The experimentally determined halfatime for experiment six was one hundred seventy-four minutes; the calculated half-time on the basis of dioxane acting as a diluent for the same experiment was ninety-seven.minutes. By calculating the half-time of reaction based on the assumption that one mole of dioxane would combined with one mole of phenol, the half-time was found to be one hundred ninety-five minutes. Comparing this latter value with the experimentally determined halfetime, it can be seen that they are of about the same order of magnitude. The results for experi- ments one through seven also coincide in the manner mentioned above. This is illustrated in Figure 3, where it is seen that the curve repre- senting the experimentally determined half-times follows closely the curve calculated on the basis of a 1:1 oxonium.comp1ex between phenol and dioxane. This leads one to the belief that a mechanism.similar to the one shown in equation I prevails, but that the effective concentration of phenol is decreased by the formation of complexes such as: -21- f2 \ 3 ~© H2 The possibility that one mole of dioxane might be combining with two moles of phenol cannot be entirely discarded since the experimental results for experiments eight through ten show a decrease in the rate greater than that calculated for the combination of one mole of dioxane to one mole of phenol. This point is best seen by comparing the curves in figure 3. The experimental curve begins to approach the curve of the theoretical values calculated for the combination of one mole of dioxane with two moles of phenol. It might be argued that the apparent inhibition shown on the addi- tion of dioxane to the reaction might be due to the increased solubility of hydrogen chloride in dioxane-phenol mixtures rather than the forma- tion of an oxonium compound between phenol and dioxane. This is improb- able, however, since during the reaction, the volume of nitrogen passing through the system was at least eighty times as great as the volume of hydrogen chloride given off. To determine whether or not the rate of the nitrogen flowing through the system had any effect on the rate of reaction, experiments were nun in which mol ratios of dioxane to phenol .'were about 0.15 and the amounts of nitrogen passed through for the same length of time were eighty and forty times as great as the amount of hydrogen chloride evolved. The rate constants for the two experiments checked with each other within the limits of experimental error. These rate constants were also of the same order of magnitude as the rate -22- constants for experiments six and seven shown in Table 11. Further evidence against the possibility that solubility of hydrogen chloride in the reaction mixture might be the governing factor was shown hy Hart and Simons (4). Two experiments were run in which the mol ratio of dioxane t0 phenol was about 0.30; in the first experiment, the reaction was merely carried out with dioxane present in the reaction mixture. The results showed that no alkylation of the phenol was tak- ing place. In case two, Hart and Simone first saturated the reactants with hydrogen chloride and then carried out the reaction as before, The results from this experiment were the same as those for case one, that is, there was complete inhibition of the reaction. Since there was no evolution of hydrogen chloride in either experi- ment, it was concluded that the increased solubility of hydrogen chloride in the reaction mixture was not the cause for an apparent inhibition. --23-- SUM-ABBY The kinetics of the alkylation of phenol with tertiary butyl chloride in the presence of varying amounts of 1,4-dioxane has been studied at 50.0° c. The investigation has shown that dioxane inhibits the reaction to an extent greater than would be expected if it were merely acting as a diluent. Evidence has been presented to show that the mechanism of the inhibition may involve a 1:1 oxonium compound between dioxane and pheiol. -24- 2. 3. 4. 5. 6. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. BIBLIOGRAPHY J. van Alphen, Rec. trav. chim., 49, 287 (1927). G. M. Bennett and F. M. Reynolds, J. Chem. Soc., 131 (1935). J. H. Simons and H. Hart, J. Am. Chem. Soc., pg, 1309 (1944). H. Hart and J. H. Simons, ibid., 11, 345 (1949). C. Friadel, 00111th rende, 81,152 (1875); Jo Chm. 8°C., 28, 1245 (1875). P. Karrer, “Organic Chemistry", Elsevier Pub. Co., Inc., New York, N. Y., 1947’ P. 114' 893. Bellucci and Grassi, Gazz. chim. ital., :3, II, 712 (1913); C. A., 8,1758 (1914). Morgan and Pettet, J. Soc. Chem. Ind., 52, 22 (1935) F. C. Whitmore, "Organic Chemistry", D. van Nostrand Co., Inc., New York, N. Y., 1937, p. 374-5, 654, 900. A. Faworski, J. Russ. Phys. Chem. Soc., 38, 741 (1906); J. Chem. Soc., 92, 274 (1907). c. Smeets, Natuurw. Tijdsche., 19, 12 (1937); c. 4., 31, 1815 (1937). "' "' E. 3001‘, J. m. Chane 800., _6-6', 303 (1944). M. I. Batuev, Compt. rend. acad. sci., U. R. S. 8., 4_9_, 277-81 (1943), C. 4., 38,6191 (1944). P. D. Bartlett and H. J. Dauben, J. Am. Chen. Soc., 6__2_, 1339 (1940). A. H. Blatt, "Organic Synthesis", John Wiley and Sons, Inc., New York, N. Y., 1941, C011. Vol. I, p. 144. L. F. 'Fieser, "Experiments in Organic Chemistry, D. C. Heath and 000, New York’ N. YO. 1941, p. 368. H. H. Willard and N. H. Human, ”Elanentary Quantitative Analy- sis”, D. van Nostrand Co., Inc., New York, N. Y., 1940, p. 185. A. G. Worthing and J. Geffner, "Treatment of Experimental Data", John Wiley and Sons, Inc., New York, N. Y., 1943, p. 240. S. Glasstone, "Textbook of Physical Chemistry", D. van Nostrand Co., Inc., 1940, p. 1046-1047. PART II THE SOLUBILITY OF HYDROGEN CHLORIDE IN PHENOL-DI 0x455: MIXTURES INTRODUCTION In the course of studying the inhibitory power of dioxane in the alkylation of phenol with tertiary butyl chloride, it was believed that the solubility of hydrogen chloride in.mixtures of phenol and dioxane might be of considerable value in trying to show the mechanism of the reaction. A search of the literature failed to yield any critical data on the solubility of hydrogen chloride in such a mixture. However, Bart- lett and Dauben (1) present evidence to show the formation of a dihy- drochloride of dioxane which they call dioxanonium chloride: F— 44' I H- -H 201' L Hz They further postulate that hydrogen chloride'will form a complel'with phenol which is due to the electrondonor property of the chloride por- tion of hydrogen chloride and the acceptor property of the hydroxyl hydrogen of the phenol. 11 Q. Cl-H immune: TAL I. Materials The dioxane was C. P. grade and was purified by refluxing over sodium for two hours and then distilling from sodium, collecting the portion coming over from 100 to 102°. The index of refraction for 209 1.4224, this roughly purified dioxane was 2C)D 1.4228 as compared to the value found in the literature (2). The dioxane was stored over sodium until ready for use. The phmol was Baker Chemical Company material, which was distill- ed with benzene to renove any water present. The portion boiling at 180 to 183° was collected and kept well covered under an atmosphere of nitrogen. Anhydrous hydrogm chloridewas prepared by dropping concentrated sulfuric acid on sodium chloride. The gas evolved was bubbled through concentrated sulfuric acid to dry it. The apparatus consisted of a fifty milliliter erlenmeyer flask fitted with a 24/40 fenale ground glass connection. The stopper, a 24/40 male ground glass connection, was sealed to four inches of capil- lary tubing (3 ) . II . Measurement Procedure In making the solubility measuranents, the calculated amount of phenol for the particular mol fraction was placed in the previously weighed reaction flask and the weight of phmol determined. The calcu- lated amount of dioxane was added and weighed. From these data the mol fractions of each constituent were calculated. The anhydrous -27- hydrogen chloride was bubbled into the mixture for two hours at room temperature. A magnetic stirrer was placed in the flask, the stop- per inserted, and the flask placed in a 50° constant temperature bath. It was noted that when the flask was first placed in the bath there was a rapid evolution of hydrogen chloride from the solution, showing it to be saturated. The capillary tubing was used on the top of the stopper to prevent a rapid escape from the flask of any gas given off and to prevent the back diffusion of air into the system, thus keeping the reaction under an atmosphere of hydrogen chloride. When the mixture had come to constant temperature (about thirty minutes) samples were removed by pipette and titrated to a phenolphthalin’e end- point with standard sodium hydroxide. Thirty minutes after the first sample, a second sample was taken, this was repeated until constant values were obtained. The mixture was stirred constantly throughout the process . -28... RESULTS AI‘TD CONCLUSIONS I e mta Solubility at 50° 0 M01 fraction g of HCl per of Dioxane 100g of Solvent 1.000 14.33 0.804 10.59 0.608 6.86 0.379 4.08 0.196 2.53 0.000 1.01a a. This value was obtained from other data (4). The solubility in grams of hydrogen chloride per hundred grams of solvent mixture was plotted against the mol fraction of dioxane in the solvent mixture (see figure 1). II. Discussion of Results The results show that dioxane is definitely a good base and that there is undoubtedly an oxonium.compound formed between the dioxane and the hydrogen chloride as Bartlett and Dauben postulate (1). They show that there is an increase in acidity of hydrogen chloride-dioxane solutions on adding phenol. This might be attributed to the concerted pull on both the chlorine and the hydrogen by phenol and dioxane re- spectively. This is shown below: H2 2 Q-H. o..c]_..H. .00 H 2 -29- The curve in figure 1 shows that as an oxygen base, dioxane is considerably stronger than phenol. -30- IJt- Il— 1 1‘ l L l 4L’ 1 1%! n I ‘ {F 04 o.‘ mol fraction l t I _ " C a I ‘ t‘ t. ’ j I ' l i ‘ J l i 5 'i a ‘ “ - t. y I V - .A\ , ,‘L ‘ ‘ I "r .L . k l J.‘ k .L k k. r r I LJ J J :d". U V L; . g. 1 I -~ I .. .’ unrn .L‘ ‘ J ( ~01- 2. 3. 4. BIBLIOGRAPHY P. D. Bartlett and H. J. Dauben, J. Am. Chen. Soc., _6_2_, 1339 (1940). Beilsteins Handbuch Der Organishen Chemie, Julius Springer, Berlin, 1934, Vol. XIX, p. 3. J. Reilly and W. N. Rae, "Physico-Chemical Methods", D. van Nostrand Co., Inc., 1939, Vol. I, p. 592. H. Hart, unpublished results. -32- SEP 5 '51‘ MAP 5 1E.“ JUL 2135; Hamisl‘sy uneasy _ 7 T547 218487 : 3727 Bordeaux 1 . 1 5 i gm 5.3”.- .‘Aflxc-‘J I'L—lFJ-MLM_ .._i_—=_n- _ 312930