August 1948
INDUSTRIAL AND ENGINEERING CHEMISTRY
t h a t temperature changes have less effect on'the solvent power of the sebacates than on t h a t of the phthalates. The flexural strengths of t h e sebacato series are shown in Figure 13. The behavior is regular in that t h e strengths at 70" C. tend t o decrease to a lesser extent than in the phthalate series, indicating lower susceptibility t o temperature differences. The same is true of t h e modulus of elasticity shown in Figure 14. The behavior here is regular and the differenre between plasticizers is marked, as noted with the phthalates. The hardness values of the phthalates series are shown in Figure 15. The data reported are Rockwell R scale (A.S.T.M. D785-44T). The methyl ester has the highest Rockwellnumber and the 2-ethyl hexyl ester the lowest. The sebacate series are also regular, as shown in Figure 16. Little difference in hardness r ~ s u l t sabove a flow temperature of 150" C., but in the softer
1485
flows there is considerable difference. .ha group, the sebacate series have lower hardness than 1 IIP phthalates, in agreement with their lower degree of solvation. LITERATURE ClTED (1) (2) (3) (4) (5) (6) (7) (8)
SOC.Testing Materials, Standards on Plastics, May 1946. Doolittle, A. K., IND.ENO.CHEM.,3 6 , 2 3 9 4 4 (1944). Ibid., 38, 5 3 5 4 0 (1946). Fordyce rtnd Meyer, Ibid., 32, 1053-60 (1940). Frith, E. M., Trans. Faraday Soc., 41,17-27 (1945). Ibid., 41, 90-101 (1945). *Jones,H., Trans. Inst. R u b b e r l n d . , 21, 298-322 (1946). Spurlin et al., J.Polymer Sci., 1 , 63-74 (1946). h i .
RECEIVED May 9, 1947. Prrsented before the Division of Paint, Varnish, and Plastics Chemistry a t the 111th Meeting of thr AVnRIcaN CiiE\rIcAL SOCIETY, Atlantic City, N. J.
Photochemical Gaseous Phase Chlorination of Isobutane ROBERT W. TAFT, J K . ~ , AND GEORGE W. STRATTON l'niversity of Kansas, Lawrence, Kan.
T h i s investigation showed the gaseous phase photochemical chlorination of isobutane (in the presence of a liquid phase) at room temperature and pressures to be a complex series of reactions that occur, to a relatively large degree, at the surface of the reactor. The extent of these reactions is shown to be markedly dependent on experimental conditions; this suggests that chlorination to a desired stage-monochloride, dichloride, trichloride, etc. -may become a reality in future experiments. The effects of each of the following factors on the chlorination reaction were studied : temperature, ratio of concentration of reactants, light intensity, presence of water vapor, absolute rgtes of flow, and surface area and construction of the reactor. The relative rates of substitution of hydrogen atoms are, primary to tertiary, as 1.00 to 5.5 for these reactions; the relative rates do not appear to vary with the ahdve experimental condition6 within the range of the experiments conducted.
LTHOUGH Haw, McBee, and others (6) have made extensive investigations on the effect of experimental conditions on the relative rates of substitution of the hydrogen atoms in various positions within a given paraffin hydrocarbon molecule, very little has been done in investigating the manner in which experimental conditions can be used t o control the extent of the chlorination reactions, other than by the ratio of the concentration of the reactants. The present study on the photochemical gaseous phase chlorination of isobutane at room temperatures and atmospheric pressures shows that the experimental conditions under which these chlorinations are conduct,ed have a pronounced effect on the extent of the reaction. Some of the reactions of the polychloride derivatives obtained by this chlorination are given in (21). The results of this investigation show- that the chlorination of iaobutane proceeds in a series of steps as shown in Figure 1; certain of these steps are much slower than others and result in negligible quantities of the products designated by asterisks. 1 Present address, Chemisi,ry Department, Ohio State University, Columbus, Ohio.
In order t o study the effect of experimental conditions on this series of reactions, these chlorinations were conducted so that comparison between two of the runs would show the effect of a single variable; the other variable factors were nearly identical. Thus, by intercomparison of 17 chlorination experiments, the effects of the following factors on the chlorination results were studied: temperature, ratio of concentration of reactants, light intensity, presence of water vapor, absolute rates of flow, and surface area and construction of the reactor. MATERIALS AND EXPERIMENTAL TECHNIQLE
Mathieson C.P. grade isobutane and ordinary liquefied chlorine obtained from the Mathieson Chemical Corporation were used in all the experiments. TheType of chlorinator used was developed by a series of preliminary experiments; the two different size reactors are shown in Figure 2. I n Figure 2, a cold finger condenser, 1, runs the length of the reactor; this is a n important detail not only because a large percentage of the reaction takes place on the surface of this condenser, but also because i t is used t o regulate the temperature of the reaction. Without such a temperature control and the means of removing heat generated by the exothermic chlorination reaction, the reaction (under conditions employed in these experiments : ratio of hydrocarbon t o chlorine of 1t o 1, 1 t o 2, or 1 t o 3, without inert diluent and temperatures between 30' t o 70 O C.) becomes so vigorous that fire and carbonization occur at rates of flow in excess of 100 ml. per minute, with the result t h a t yields under these conditions for a 12-hour run are never greater than 400 grams. After t h e rates of flow of the reactants were set, regulation of the temperature of the condenser, 3, was secured by adjusting the rate of water flow through the condenser so t h a t the heat liberated by the reaction maintained the temperature constant within a few degrees at a n y desired temperature. Temperature readings were obtained b y means of a thermometer immersed centrally near the outlet of the condenser, 3. Although there is a temperature gradient between the center of the condenser and t h e walls of the reactor it is not large, except at the outer surface near the chlorine inlets where the temperature may be as much as 10" C. higher than t h a t of the condenser. The temperature of the condenser water was taken simultaneously with rate of flow readings for the inlet gases, and the average value is t h e one given in the present, data. The outer jacket of a Victor Meyer volatilizing chamber, 2, was cut t o desired length. T h e mixing of the reactants and the reaction took place between t h e inner wall of this
INDUSTRIAL AND ENGINEERING CHEMISTRY
1486 CH3
H,c-+-cH, + ci2 H
2-methylpropone
,
4 CH3 H3C+-GH3
GI H&-$-CH3
CI
H
1- chloro- 2-methylpropone I
r'
t
7%
CI.&IC-F-CH,
p s CIH&-C-CH,CI
CI H2C-C-CH3
H I,l-dichloro-2methylpropone
1,3-dichloro-2methylpropone
I,2-dichloro-2methylpropone
+
*
2- chloro-2-methylpropone y
k
3
b
-
-+-----I FH3 FH2CI CI,HC-$XH2CI CI H2C-F-CH&I CI H&-C-CH&I c H * H dl
w
THJ
CItHC- C - C h C!
I,l,3-trichloro-2- 1,3-dichloro-2- 1,2,3-trichloroI,l,2-trichloromethylpropone chloromethytpropane 2-methylpropone 2-methylpropane
Id
+ FH2CI CI H,C-C-CH&I
I
ClzHC-~-CH&I
CI
CI
I, 2,3- t ric h loro-2chloromethylpropune
r
7 H,C I Gig C-C-CH&I CI
1,1,2,3-tetrochloro2- chloromethyl propane
FH3 C13G-$-CHS
H3
*
I , 1,233- te trochloro-2-methylpropone
7%
C12HC-$-CHC12 CI 1,1,2,3,3-pentochloro2-methylpropane
CI I , 1,1,2- te trachloro-2-methyl propane
i FHa
C13G-$-CH&I
CI 1,1,1,2,3-pentachloro-2-methylpropane
Figure 1
jacket and the outer surface of the condenser; the product was collected in a receiver at the bulb end of this jacket. Inlet 4 is a 2-mm. capillary stopcock through xvhich the isobutane is introduced. Chlorine is introduced a t three (two for small reactor) inlets ( 5 , 6 , 7 ) along the reactor to prevent building up too great a concentration of chlorine a t one spot. I t isimportant t h a t these tubes be of no greater diameter than 1 mm. because the incoming chlorine gas must enter at a high velocity to ensure uniform mixing with the isobutane to prevent firing and carbonizing. I n the numerous trials t h a t were made, no fires or carbonization took place in a reactor of this description with the condenser temperature below 70" C., using rates of flow of isobutane up to (300 ml. per minute and chlorine flow up to 2000 ml. per minute (total for all chlorine inlets). The importance of rapid mixing of the hydrocarbon and chlorinein reactors of other types has been previously noted by Hass and others ( 4 ) . Figure 3 shows the chlorination assemblp. The gases were passed over calcium chloride, 1, as they left the stolage tanks. The flowmeters, 2, were special chlorine-resistant rotameters (Fischerand Porter,Hatboro, Pa., series 700 master enclosed rotameter). T h e metered chlorine gas was passed through a pair of gas bubbling towers, 3, before entering the reactor, 4. The tower nearest the reactor was partially filled with concentrated sulfuric acid (saturated with chlorine) for most of the chlorination runs. On a few of the chlorination runs, water was substituted for the, sulfuric acid t o determine the effect of saturatlng the reactants with water vapor. Isobutane was not passed through concentrated sulfuric acid since i t reacts with it to give tarlike products. A 1 t o 1 mixture of sealing wax and paraffin was found t o be a n excellent material for preventing t h e leakage of chlorine around the ground-glass joints of the towers.
To obtain fairly constant rates of flow each of the inlets was supplied by a separate storage cylinder. Because of the unavailability of chlorine-resistant needle valves, the rates of flow were regulated b y means of the valves on the storage cylinders. It was not possible to acquire absolutely constant rates of flow; therefore the flows were regulated frequently enough to keep variations of the rates of flow never greater than
Vol. 40, No. 8
7'70 throughout the chlorination run. Flowmeter readings nere taken every 0.5 to 1.5 hours and the average rates of flow for the runs are the ones reported. The chlorination reaction r as activated by the use of a series of ordinary tungsten filament electric light bulbs (5, Figure 3). These bulbs were placed alternately above and below the reactor along its length, about 3 to 4 inches from the outer surface. The number of lamps used and their intensities were varied in different experiments. The chlorinated products condensed on the wall of the chlorinator and were collected in a receiver, 6. The hydrogen chloride generated by the reaction passcd through the system and was led outdoors. The outgoing gases were passed through a dry ice-carbon tetrachloride-chloroform vapor trap, 7, and large quantities of 2-chloro-2-methplpropane and 1-chloro-2-methylpropane Tere collected in this trap. The surface of the reactor was not given special treatment before each of the chlorination runs because it was fully activated by previous chlorination. After each chlorination run, the products of the reaction were allon-ed to drain from the surface of the reactor, and the reactor was allowed to stand in the presence of the residual vapors until the next chlorination run was made (never over 24 hours) without the introduction of air or foreign matter. The chlorination product contains appreciable quantities of hydrogen chloride in solution; i t was found that the most desirable way t o remove it was by distillation, rather than by chemical means. Shaking the chlorinated product with water or a dilute sodium hydroxide solution was not satisfactory as it produced evidences of hydrolysis. Preliminary experiments bhowed t h a t the chlorination products of isobutane could be conveniently and efficiently resolved into the folloning fractions: (4) b.p. = 48-56" C.; (B) b.p. = 64-74°C.; (C) b.p. = 98-110°C.; (D) b.p.5, = 57-74'C.i (E) b.p.50= 77-87" C.; (F) b.p.50 = 98-110" C.; and (G) residue. These seven fractions ere collected for all the chlorination runs made. Only fractions (D) and (G) contain more than one compound; these two fractions contain two and three compounds, respectively, 15-hichcannot be conveniently separated by the first fractionation. On completion of the first fractionation of all the chlorination runs, approximately 2 liters of a representative sample of each of the fractions were carefully refractionated and fractionation curves were plotted ( 1 9 ) . From the results of the first fractionation and the composition of these fractions as given in Table I, the composition of the original chlorination product for each of the runs was calculated on the assumption t h a t the fractions having the same boiling range have the same composition. This assumption should be nearly correct for those fractions containing only one compound, but for those containing two or more compounds (D and G) the results are only approximate.
0
Figure 2
, August 1948
INDUSTRIAL AND ENGINEERING CHEMISTRY
TABLE I. COMPOSITION OB'FRACTIONS AS GIVENBY FRACTIONATION CURVES Weight %
Fraction and Compound B.P. = 48' t o 56O C. Forerun 2-Chloro-2-methylpropane 1-Chloro-2-methylpropane B. B.P. = 64O t o 74O C. 2-Chloro-Z-methylpropane ' 1-Chloro-2-methylpropane 1,2-Dichloro-2-methylpropane C. B.P. = 9 8 O to 110' C. 1-Chloro-2-methylpropane 1 ,l-Dichloro-2-methylpropane 1 2-Diohloro-2-methylpropane 1:3-Dichloro-2-methylpropane D. B.P.ao = 57O t o 74' C.
A.
2.5
80.0 18.0
4.4 91.5 4.1 5.1 8.45 82.0
4.7
(1:l ratio (1:Z and 1 : 3 runsb) ratio runs b) 5.2 3.8
1,2-Dichlor0-2-methylpropane 1 3-Diohloro-2-methylpropane 43.5 1: 1,2-Trichloro-2-methylpropan 30.3 1,2,3-Tri~hloro-2-methylpropane 21.0 E. B.P.60 = 77' t o 87' C. L 1 2-Trichioro-2-methylpropane 1'2'3-Trichloro-2-methylpropane 1:1~2,3-Tetrachloro-2-methylpropane F. B.P.60 = looo t o 110' C. 1,2,3-Trichloro-2-methylpropane 1,1,2,3-'l&tre~hloro-2-methylpropane 1,2,3-Tr1chloro-2-chloromethylpropane G. Residue 1 1 2 3-Tetrachloro-2-methylpropane 1'2'3~Trichloro-2-chloromethylpropane l~l~1,2,3-Pentachloro-2-methylpropane 1,1,2,3-Tetrachloro-2-~hlorornethylpropane Others (assumed mol. weight t o average 225)
38.4 35.4 22.4 3.5 83.4 13.1 3.3 80.0 17.0
2.9 29.0 2.4 34.0 32 0
This figureis the result of a very liberal interpretation of fractionation curve. b The two series of figures are based on an rtppropriate interpretation of fractionation curve.
All fractionations were carried out in columns of the type described by Taft and VanderWerf ( 2 2 ) . Any chlorinated compound boiling above 110" C. at atmospheric pressure was fractionated under reduced pressure. A constant ressure was obtained by use of a Newman type manostat (197. All fractionations were conducted a t reflux ratios of 25 or more to l. The weight of the distilling pot was taken before and after the atmospheric and vacuum part of each fractionation, so that the exact fractionation losses were known for these parts. The fractionation losses for each of the fractions were estimated from these results on the basis of the volatility of the fraction, and these relatively small (total fractionation losses were never greater than 4% by weight of the products of a run) fractionation losses were added to the weight of the distillate collected. The results of the chlorination runs are shown in Table 11. Room temperature for these chl2rinations was between 26 and 33" C., and pressure was between 733 and 748 mm. of mercury. Those experiments conducted in the large chlorinator are designated by L after the experiment number; those conducted in the small chlorinator by S. Experiment 15 was carried out by saturating the incoming gases with water vapor, indicated by W after the experiment number. All other experiments reported were conducted under anhydrous conditions. Four 200-watt lamps were used for activation in experiments 1, 2, 3, 10, 11, 12, 16 17, 18, 19, and 20. For experliments 6, 7, 8, and 9, three 200-watt lamps were used. Four 25-watt lamps served as the source of activation for experiments 4 and 5. The experiments marked f" with a n -asterisk are the average results of two trials
.
Y
1487
(conducted under the same conditions) which were in agreement within 10%. As mean values, the results given in these experiments should, therefore, be accurate to 5 or 6%, except for the results on 1,3-dichloro-2-methylpropane,l,l,Ztrichloro-2-methylpropane, 1,2 3- trichloro-2-chloromethylpropane, and 1,1,2,3-tetrachloro-2-chloromethylpropane,which may be in error t o a somewhat greater degree, as previously mentioned under fractionation procedure. The experiments without a n asterisk are trials which were not checked by repetition; the experimental error of these results may be as much as lOy0or greater. EVIDENCE FOR IDENTITY O F PRODUCTS
2-Chloro-2-methylpropane. The physical constants of the fraction, b.p.740= 50.8" t o 50.9" C., d:5 = 0.8374, n%5= 1.3829, are in good agreement with those reported in the literature, b.p.765 = 51.6" C. ( I @ , die = 0.83683 (I6),d:5 = 0.847 (IS), ab8 = 1.3868 (IS) for 2-chloro-2-methylpropane. 1-Chloro-2-methylpropane. A comparison of the physical constants obtained for the fraction b.p.r40 = 68.2" to 68.3" C., d j 5 = 0.8884, n:5 = 1.4046 with those given by Underwood and Gale (16) and Timmermans and Martin (IS) for l-chloro-2methylpropane, b.p.760 = 68.8' C., d:' = 0.8829, nb5 = 1.4010, identifies this fraction. 1,2-Dichloro-2-methylpropane. The fraction b.p.,,, = 56.5 O to 56.8" C. gave the following physical constants: b.p.T37 = 106.1' C., d f 5 = 1.0887, n:' = 1.4347. Hersh and Nelson (8) obtained 1,2-dichloro-2-methylpropaneby a means other than chlorination and gave the following physical constants for this dichloride: b.p.7, = 106.5' C., d!' = 1.093, 4,' = 1.4370. The close agreement of the two sets of constants indicates t h a t this frastion is 1,2-dichloro-2-methylpropane and that it contains very little l,l-dichloro-2-methylpropane. The small amount of forerun for this fraction probably contains some 1,l-dichloro-2methylpropane, b.p.7, = 103' to 105' C.,':d = 1.0111 ( 2 4 ) . 1,3-Dichloro-2-methylpropane.The fraction b.p.,, = 60.7" to 61.4" C. is a dichloride by analysis (55.7% chlorine by determination, 55.9% chlorine by theory) and its physical constant, b.p.,,, = 134.3' to 134.7' C., d i 6 = 1.1325, n:' = 1.4488, correspond to those given in the literature for 1,3-dichloro-2-methylpropane, b.p.700= 136.0" C. (6),dzo = 1.131 (6). 1,1,2-Trichloro-2-methylpropene.The fraction b.L.50 = 67.5 to 68.1" C., b.p.737 = 144.0' C., d t 6 = 1.2677, n2,6 = 1.4638 is a trichloride by analysis (65.9% chlorine by determination, 66.0% chlorine by theory) and its physical constants identify it as 1,1,2-trichloro-2-methylpropane.Rogers and Nelson (17) give the following constants for this trichloride: b.p.,, = 144.5'
CI
7
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ym
Figure 3
. 1488
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 40, No. 8
~
.IT. rates of flow. ml./niin ClHlo inlet, Clz inlet 1 C!z inlet 2 C12 inlet 3 Total moles CiHio pasari1 Total moles C1z passer1 Ratio C!z:CaHio Yipld of chlorides, ai-aiiis
880
847 247 260 250 303 315 303 315 342 308 10.8 17.0 9.90 1 0 . 8 18.0 1 0 . 2 1.00 1.06 1.03 701) l00a 670 890
20.8 24.2 29.6 43.1 2.2 0.6 22.9 7.4 7.9 6.0 4.0 3.4 9 , 2 10.1 1 2 3-Trichloro-2-methylpropane 7 1 ~ 8~ . 8 1 ~ 1 ~ 2 . 3 - T e t r a c h l o r o - 2 - ~ u e t h ~ l ~ ~ o ~2 a. 1 1'2'3~Triohloro-2-chloromethylpro~a1~~~ .. l : l : 1 , 2 , 3 - P c n t a c h l o r o - 2 - m e t h y l ~ ~ o ~ ,~,~ ~ ~ ~, ~, 1,1,2,3-Tetra.ohloro-2-chlorom~thylpropane ., .. Others Assuming niol. wt. 3 210 0.7 1.5 Assunling mol. wt. = 22.5 CtHlo reacted, 010 ,jL$:?8;:6 Ch reacted, % $ 1 3 , ~7 7 . 8
845 732 725 270 375 370 302 352 35.5 315 10.0 0 : O l 9:37 1 0 . 5 8.93 9 . 3 7 1.05 0 . 9 9 1.00 604 ,521 .is5
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:i.x
O:4 i:a 2 : 8 318 8 9 . 4 96.0 9 6 . 5 97.6 9 6 . 6 9Zi2 89.3 9 1 . 4
to 145.4' C., ti:' = 1.2712,n;? = 1.4666; Tish(-hriiko'i ($4) 190.2"10 IHO.8", d:r' = 1.4344, T L =~ 1.4934; those given b y Rogers and Yelson are: h.p.7wi= 190.6" to 391.3", d:5 = 1.4393, constants are: b.p.,a = 145" i o 146" C:., di0 = 1.2588. The VI,"," = 1.4963. fact t h a t this trichloride gives t,1 -dic-hloro-2-methyl-l-propenr The I rctification curve for 1,1,2,3-tetrachloro-2-mcthylprop~111 on treatment with alcoholic pota,sh is in agreement with the diou-rd an appreciable temperature rise over t h r entire fractioti assigned structure (,%'I). 1,2,3-Trichloro-2-methylpropan1 . L t oixiptiriwn of the physthis mav be caused by the presence of small amounts o f 1,1,3trichloro-2-methylpropane and 1,3-~irhloro-2-rhloromt~th vlical properties of the fraction b.p.,,, = 81 .OD to 81 .3 C., b.p i 3 7 = propane. 162.2"to 162.6"C.,dZ6 = 1.3012, = 1.4736 uitti those It'1,2,3-'~richloro-2-chloroniethy~pr~paiie. The fraction b.p.lo = ported in the literature for 1,2,:3-triehlor oisobutaric identified 86.7"to 87.7" C.,b.p.737 = 209.8" t o 210.5' C.,d i 5 = 1.5036, this fraction, u hich Itnaly\i\ tlc indicatc 10 be a tricnhloride ni5 = 1.5082,is a tetrachloride by analysis (molecular weight (molecular neight by Erecziiig int depression of benzene is 161,by theory is 161.5). Rogei s arid Selron ( 1 7 ) give the follo~~-- by fleezing point depression of benzene 196, by theory 196.0) and on the basis of a comparison with the physical constants ing constants for this tiichloride: b p . 7 ~= 162.0"to 163.1"C., given by Rogers and Nelson ( I T ) , and Kleinfeller (IO) is 1,2,3d:' = 1.3020,ng = 1.4765. 1,2,3-Trichloro-2-i~iethplproprtneis trichloro-2-chloromethylpropane. Kleinfeller gives the boiling V I at ( I 1 ) and PorgorshelPki the trichloride obtairietf by (16) by the chlorination of 2-1 lpropcne, although Mouneypoint of this compound as 87" C. at 9 I I I ~ I . The physical constante reported by Rogers and Nelsoa are: b.p.,M = 206" to rat, on the basis of an empirical iule, erroricwusly reported it to b(5 210aC.,d25 = 1.481,nt6 = 1.508. 1,1,2-trichloro-2-1iiethvlpropan~~I'orgorshelski cox rertlp identi1,1,1,2,3-Pentachloro-2-methylplopane. A small aniourit ot a fied this compound on the 1)iisih of it. hydrolysis piodurt, but white crystalline solid was collected between the fracstions t ) . p l o = t h k correction has not heen made in some of the pertinent 88.7" to 87.7" C. and h.p.,o = 97.9' to 98.9" C.; it atpprwed literature. Porgorshelski gavt: thc boiling point as 163.5' to 164" C. at 772 mni. and Mouncjmit gave thr constants, h.p.7,, = 158" to 162" C.,di6 = 1.295. It is likrly that this fraction contains a small amount of 1,1,1t2-tetr.achlor0-2-methylprtipane, as an impurity, for this crystalline iubstancc subliiiirs at a temBoiling.Points Tndex of perature several degrees higher than the boiling poilit or 1,2,3-RefracPressure, t richloro-2-methylpropane. Porrnnla Temp., C. mm. 11 :' tion, n'T; amouiil I,1,1 , 2 - T e t r a c h l o r o - 2 - i n ~ t h ~ l ~ r ~ ~Ar ~ va t~~ rt ~ysmttll ~ 50.8 t o 5 0 . 9 740 0.8374 1.382Y (CHz)&Cl 740 68.2 to 68.3 0,8884 1,4046 \CH3)2C(H)CHnCl of a white crystalline solid ckpohited on the condenser in frac737 1 ,0887 1 4327 (C€I~)zC(CI)CHzCl 106.1 tionating the tailings of the 1,2,3-trichloro-2-niethylpropan~~ 137 56.5 to 56.8 737 1 3 4 . 3 t o 134.7 1 132.5 CRsC(H) (CHzC1)z 1 4488 fraction. The crystalline nature of this product and its boiling 6 0 . 7 to 6 1 . 4 50 144.0 7 4638 737 1,2677 range, b.p.60 = approxirnntely 82" to 84" C., indicate that a very (CHa)zC(CI)CHClz 6 7 . 6 t o 68.1 50 small amount of 1,l,lJ2-tetra chloro-2-inethylpropane is formed 1 4736 1.3012 lgf.3 t o 162.6 CH-I1C(Cl)(CHzC1)z 781 0I.I in the gaseous phase photochemiral chlorination of isobutane. Approx. 82 to 84 50 (CHa)zC(CI)CCls 190.2 t o 190.8 737 1,4344 1.4934 CHJ~(CI)CHCIZ Rogers and Nelson (17) reported the following constants for this substance: b.p,VbO= 174" (siiblinica), b . p . ~ , ,=~ 192" C.,1n.p. = kHzC1 105.1 t o 105.3 50 1 5036 1 5082 (CH2Cl)aCCI 209 8 t o 210.5 737 178.6"t o 179.8"C. 121 t o 122 50 86.7 t o 8 7 . 7 10 1,1,Z,3-Tetrachloro-2-nlethylpropa1le. The fraction b .p. "0 = 737 Melting point = 73.5O 0. 212.8 to 213.4 C HsC (Cl) C C13 104.3" to 105.3"C. is atetrachloride by analysis (71.5% chlorine bHzC1 90 t o 93 10 by determination, 72.4% chlorine by theory). Its physical 1,5686 1.5165 (CH2Cl)zC(Cl)CHCh 2 2 6 . 3 to 226.4 737 constants are in agreement with those reported for 1,1,2,3a Physical constants given in this table are more accurate t h a n those pretetrachloro-2-methylpropane by Rogers and Nelson (17). The viously published by Taft and Stratton (20). physical constants determined for this fraction are b.p.73, = r)"
INDUSTRIAL AND ENGINEERING CHEMISTRY
August 1948
largely in the range b.p.,, = 90 O to 93 O C. On'recrystallization from alcohol, the following constants were obtained: b.p.,,, = 212.8' to 213.4' C., m.p. = 73.5" C. This fraction has the same boiling range as crystalline 1,1,1,2,3-pentachlor0-2-methylpropane sold by Halogen Chemicals (S),b.p. = 209" to 213'; they prepare it by the addition of chlorine to 1,1,3-trichloro-2-methyl-lpropene (87). Jacob (9) reported the melting point of 1,1,1,2,3-pentachloro2-methylpropane as 59.5" C., obtained also by the addition of chlorine to 1,1,3-trichloro-2-methyl-l-propene. A p p a r e n t l y Jacob was not able to remove the last traces of olefin, for his melting point is noticeably low. 1 , 1 , 2 , 3 - Tetrachloro 2 chloromethylpropane. The fraction b.p.,, = 97.9' t o 98.9 O C. is a pentachloride by analysis (molecular weight by freezing point depression of benzene 230, molecular weight by theory 231.5). The following physical properties were obtained for this fraction: b.p.,,, = 226.3' to 226.4' C., b.p.$ = 95.1" to 95.6" C., d i 5 = 1.5686, n%6 = 1.5165. The boiling point a t 12 mm., 100" C., is in good agreeineiit with the boiling point reported by Kleinfeller (IO)for 1,1,2,3-tetrachloro2-chloromethylpropane, b.p.,z = 99 O to 101O C.
- -
DISCUSSION OF RESULTS
EFFECT
OF
TEMPERATURE. Comparisons of experiments
1 to 3, 10 to 12, and 18 to 20 show that increasing the condenser
temperature markedly increases the yield of monochlorides but decreases the yield of polychlorides. These same experiments show that the percentage of isobutane reacting, for the three Clz:CaHlo ratios used, increases in the temperature interval, 18' to 35", t o a small and somewhat indefinite extent but increases between 35 ' t o 54' very definitely in every case. The percentage of chlorine reacting was greatest at 20 O, was a minimum ai 35", and in every case was definitely greater again a t 54" c.
EFFECTOF RATIOOF CoNcmixurIoN OF REACTANTS. Comparisons of experiments 1 and 10, 2 and 11, and 3 and 12 (chlorinations run in the large reactor) show that increasing the ratio of Clz:CaHlodecreases the yield of mono- and dichlorides, and results in large increases in the yields of tri- and tetrachlorides: Increasing the C1z:C4Hlcratio to 1:3 shows, in general, the same effect on the results as experiments 18to 20 indicate. Chlorination runs in the small reactor (comparisons of experiIiients 6 t o 9) show that an increase in the Clz:C4Hlo ratio decreases the yields of monochlorides, and increases the yields of trichlorides, as above, but the yields of dichlorides are increased and the yields of tetrachlorides are decreased. These results are contrary to those obtained with the large chlorinator. The inconsistency of results shows the importance of the contribution made by the construction of the reactor. EFFECTOF LIGHTINTENSITY. Comparison of experiments 2 and 4 shows that light intensity has very little, if any, effect on the results of a chlorination as long as there is sufficient light to activate the reaction. The data show a slight increase in the yields of polychloride with an increase in light intensity resulting from the substitution of four 200-watt lamps for four 25-watt lamps, but the effect is so small as to be within the experimental error of the data.
If the chlorinator is wrapped SO that no light can enter, the yield of chlorinated products collected in the receiver is decreased 80 to 90%; this shows that light is essential for activation. The minimum light intensity for normal chlorination was not determined but experiment 4 shows that 25-watt lamps are as efficient activators as 200-watt lamps. ]l!FFECT O F n-ATER VAPOR. A comparison of experiment 15 with experiment 12 shows that the presence of water vapor greatly decreases the yield of monochlorides but] increases the yield of polychlorides markedly. However, water vapor apparently ib
1489
not a catalyst increasing the rate of chlorination, for the percentage of isobutane reacting is very much lower in the presence of water vapor than under dry conditions. EFFECTOF ABSOLUTE RATESOF FLOW.The absolute rates of flow of the reacting gases have pronounced effects on the results of a chlorination (comparison of experiments 11, 16, and 17). Reducing the rates of flow of the incoming gases while maintaining the same relative rates of flow greatly decreases the yield of monochlorides and increases the yields of polychlorides, as well as the percentage of isobutane and chlorine reacting. EFFECTO F SURFACE AREAAKD COKSTRUCTION O F REACTOR. The results obtained in the large chlorinator differ from those obtained in the small chlorinator, when the relative rates of flows of the reacting gases are changed. Further, comparison< of experiments 3 and 7 and 8 and 12 show, under identical conditions except for the absolute rates of flow, that the small chlorinator with the larger vapor space gives lower yields of monochlorides and greater yields of polychlorides than the large chlorinator. This result indicates conclusively that the volume of the vapor space of the reactor has a greater effect on the results than does surface area. RELATIVE RATEb OF SussmTaTIo&. A comparison of experiments 1 to 3 shows that lowering the temperature below 35" C. for the 1:l runs markedly increases the yield of dichlorides and decreases the yield of monochlorides. The decrease is appreciable in the yield of 1-chloro-2-methylpropane but only slight in the case of 2-ohloro-2-methylpropane. Likewise, a comparison of experiments 1 and 10 and 2 and 11 shows that increasing the Clz:CrHlo ratio markedly decreases the yield of l-chloro-2niethylpropane but decreases only to a slight extent the yield of 2-chloro-2-methylpropane. In part, this must be due to the longer residence of the 1-chloro-2-methylpropane on the chlorination surface because of its lower vapor pressure. However, the order of magnitude of the results indicates that it must also, in part, be due to the greater rate of reactivity of the 1-chloro isomer. This conclusion is similar t o the observation of Vaughan and Rust (B) t h a t isopropyl chloride is less reactive than n-propyl chloride. Further, it is the result that would be expected as 2chloro-2-methylpropane has available only primary hydrogen atoms for further substitution whereas 1-chloro-2-meth lpropane contains a tertiary hydrogen atom which may be suxstituted. Hass, McBee, and Weber (6)have shown t h a t tertiary hydrogen * atoms are much more reactive than primary hydrogen atoms. These investigators found t h a t at 300' C., with the reaction in the vapor phase, hydrogen atoms are always substituted at the relative rates, primary to tertiary, as 1.00 to 4.43. Using the data of experiments 3 and 4 in which the yields of polychlorides are the smallest of any of the runs, the relative rates of substitution are, primary t o tertiary, as 1.00 t o 5.5 for 50' C. for the reaction of gaseous isobutane and chlorine i n the presence of a liquid phase; this indicates t h a t under these conditions the tertiary hydrogen atom is even more relatively reactive than under the conditions at which Hass conducted his experiments. This order of rates of substitution corresponds t o more rapid substitution of the least acid hydrogen of the molecule. The entire chlorination reaction is so complicated-for example, by the formation of the same compound from two isomers, Figure 1, in unknown amounts-that i t is not possible t o follow the course of the reaction quantitatively on the basis of these experiments alone. Therefore, it is not possible t o give quantitatively the relative rates of substitutions of the hydrogen atoms in methyl and 1- and 2- substituted methyl groups. T o do this i t would be necessary t o conduct two separate experiments, similar to those described here starting with the two monochloride isomers instead of with isobutane. However, the results indicate qualitatively t h a t the relative rates of substitution are in the order CHa)CHZCl>CHCl2, and are not of the order of the pronounced difference between the relative rate of substitution of a primary and tertiary hydrogen atom. Thus, after the hydrogen atom on the tertiary carbon has been substituted, the yield of the various polychlorides may be estimated by a consideration of the statistics of chance The above order of rate of substitution corresponds to more rapid substitution of the hydrogen of the greatest electron densitythat is, the least acid hydrogen of the molecule.
.
1490
.
INDUSTRIAL AND ENGINEERING CHEMISTRY
The relative rates of substitution apparently arc. not noticeably effected by any of the factors which were investigated. These factors affected only the extent of the reaction.
OTHER OBSERVATIONS.The chlorination reactions involved no skeleton chain rearrangements. There were no noticeable quantities of olefins formed by the chlorination. It should be possible, then, to recycle unreacted hydrocarbon and chlorine t o obtain high ultimate yields. However, the polychlorides were found to decompose on atmospheric distillation giving small amounts of olefins. CONCLUSIONS
Although there is evidence that the chlorination reaction carried out under similar conditions of pressure, temperature, and light, once initiated a t the surface, extends into the vapor phase ( 7 ) , the results discussed above can best be explained by proposing t h a t a large proportion of all substitutions of a chlorine atom for a hydrogen atom in a molecule of a hydrocarbon or chloride derivative occur a t the reaction surface under conditions similar t,o those employed in these chlorination experiments: reactants in gaseous phase; room temperature and pressure; reaction activated by light; a liquid phase present; and a flow type reaction vessel. Accordingly, the increasing yields of monochlorides with increasing temperature can be explained by the increase of vapor pressure of the monochlorides with temperature, a factor which results in decreasing the time of residence on the surface of the reactor. (This result is contrary to the well accepted fact t h a t the rate of a single phase chemical reaction is always increased by an increase in temperature. Therefore, chlorination reactions are not simple one-phase reactions.) Once the monochloride is evaporated from the liquid surface into the vapor space of the chlorinator by the flow of outgoing gases (unreacted isobutane and chlorine and hydrogen chloride formed by the react,ion) relatively little further substitution occurs. Just vhich chlorides are suscept,ible to removal from the reaction surface under the Conditions employed in theso experiments in amounts sufficient to affect the results appreciably is not completely answered by the data. However, there is evidence that 1,2-dichloro-2-methyIpropane (and l,l-dichloro-2-methylpropane)as well as the monochlorides evaporate from the surface a t 50" t,o 60" C. in considerable quantities, and are, therefore, not subject t'o further substitution in quantity (18). The results observed on the effect of temperature on the percentage of chlorine and isobutane reacting are apparently due to two opposing factors: the decrease of residence of chlorinating substances on the reaction surface with increasing temperature; and the increase in the rate of chlorination with temperature. The gases passing over the liquid chlorination surface vaporize a quantity of monochlorides. Because this quantity will depend upon the velocity of the gases, it would be expected t'hat, as the velocity of the gases a t the inlets as well as through the reactor is increased, t.he time of surface residence would decrease. The observation that the yield of monochloride increases with increasing rates of flow can then be explained as resulting from a decreased extent of reaction with decreased time of surface residence. The decreased yields of monochlorides and increased yields of polychlorides caused by the presence of water vapor can be explained: by the decrease in the partial pressures of monochlorides (a Dalton law effect); and by the decTease in velocity of the hydrogen chloride (formed by the chlorinatioil) through the reactor caused by its reaction with water vapor to form hydrochloric acid. Both these factors increase the time of residence of the monochlorides on the Chlorinating surface and thus account for the reaction proceeding to a greater extent. Further experimental evidence for this conclusion was obtained as follows: When the condenser of the small chlorinator
Vol. 40, No. 8
was heated to and maintained at 100' C. by passing steam through it, and the chlorination run started under the same conditions as experiment 8 (except, of course, for the condenser temperature), no product was obtained regardless of hoiv long the gases were passed through the reactor and the intensit,y of the light sources used. If the steam source was removed, and the condenser temperature allowed t o fall to 50" t o 70" C., a liquid surface was observed to form over the entire length of the chlorinator and chlorination started and continued at its normal rate. Once this liquid surface had formed, the temperature could be raised above t h a t required for its formation (50° t o 70 C.), but the reaction continued normally, apparently carried on by the presence of polpchlorides formed on the surface. Groll, Hearne, Rust, and Vaughan ( 2 ) made almost identical observations on the importance of a liquid phase in the low temperature chlorination of olefins. The fact that chlorination will take place only when a liquid phase is present, indicates coiiclusively t h a t the reaction is initiated at the surface. If the temperature of the surface is high enough to prevent the condensation of the monochloride (the first products of chlorination) or the surface has been treated with inhibitors (such as oxygen) which break the chain reaction before the formation of monochloride, no reaction t,alies place between isobutane and chlorine. Once a liquid layer has formed, further reaction is initiated at this liquid phase. Although, this observation does not require that all the Chlorination reactions occur a t the surface, t,ogether ait,h the results of the chlorination experiments, it does indicate strongly that a relatively large proportion of the reactions do occur a t the surface. The mechanism of these chlorination reactions must be that of the chain type as shown by the inhibitory action of oxygen or isobutene and the necessity that act,ivating light be present in order to produce reaction. The photochemical chlorination of an alkane in the liquid phase has been shown by Brown, Xharasch, and Chao (1) to proceed by t'he transitory existence of t.he alkyl free radical of the alkane (in this case, isobutyl) according to the following chain mechanism: Clz
+ hv
= 2C1.
+ CJIto = CaHv + HCI C ~ H Q+. Clz = CdHgCl + C1. C1.
The fact t,hat a liquid phase must be present in order to obtain reaction betiTeen isobutane and chlorine suggesk t'hat a large proportion of the reaction takes place (the isobutyl radical may be replaced by any mono- or polychloride derivative) a t the surface betwe& the reactants (which are in greater concentration a t the surface than in t,he gas phase) and free radicals formed at this surface. The remainder of the Chlorination process, although initiated a t the surface, propagates into the vapor phase. Once formed, react,ion products either: drain from thc surface, unaffected by further reaction; react further a t the surface; evaporate from the surface or remain in the vapor phase and collect in the receiver without further reaction; or evaporate from the surface or remain in the vapor phase and react' further in this phase. The third possibility must predominate over the fourth possibility according to the data. The results of this investigation show t h a t the extent' of the chlorination reaction can be controlled t o a very appreciable degree by the experimental conditions which were investigated. Using the observations made in these experiment's, it may be possible in future experiments to regulate the reaction so that it may be stopped, to a large degree, a t any desired stage by carrying the conditions found favorable for the formation of the dcsired product to their limit. For example, if it is desired t o produce only monochlorides, the data indicate that very high percentage yields of monochlorides would be obt,ained by running the reaction with a C1,:C4H,, ratio of 1:2 or 1:3 with the dry reactants entering a reactor with a small vapor space but relatively large surface for reaction, a t high velocities; the condenser temperature should be a t 60" to 70" C. and there should be sufficiently high light intensity t,o activate the reaction. Unreacted gases could be recycled to obtain high ultimate yields as olefins are not produced under these conditions. A maximum yield of dichlorides should be obtained by conducting the chlorination a t 15" to 20" C. with a Clz:C4Hm ratio of approximately 1.5:1, and vit,h moist reactants entering a small
INDUSTRIAL A N D ENGINEERING CHEMISTRY
August 1948
reactor with a large vapor space at moderate velocities. There should be sufficiently high light intensity to activate the reaction, and~recyclingof unreacted gases should be possible. Similar predictions for the conditions favoring maximum production of trichlorides and of tetrachlorides may also be derived from the data. ACKNOWLEDGMENT
The authors wish to express their thanks to A. L. Henne for his helpful criticisms and suggestions. LITERATURE CITED
Brown, H. C., Kharasch, J. S., and Chao, T. H., J . Am. Chem’ Soc., 62, 3435 (1940).
Groll, H. P. A., Hearne, G., Rust, F. F., and Vaughan, W. E., IND. ENQ.CHEM., 31,1239-44 (1939). Halogen Chemicals, 616 King Street, Columbia, S. C. Hass, H. B., McBee, E. E., and Weber, Paul, IND. ENG.CHEM., 27, 1190 (1935).
Hass, H. B., McBee, E. E., and Weber, Paul, Ibid., 28, 333 (1936).
Hass, H. B., and Weber, Paul, Ber., 67,9745 (1934). Henne, A. L., private communication, Ohio State Univ., Columbus. Hersh, J. M., and Nelson, R. E., J . Am. Chem. SOC.,58, 1631
1491
Jacob, R . , Bull. soc. chim., 7, 581-6 (1940). Kleinfeller, H., Bor., 62B, 1582-90 (1929). Mouneyrat, Ann. chim. ( 7 ) ,20, 534 (1900). Newman, M. S., IND. ENG.CHEM.,ANAL.ED.,1 2 , 2 7 4 (1940). Norris, J. F., and Olmsted, A . W., “Organic Syntheses,” Collective Vol. I, p. 138, New York, John Wiley & Sons (1932). (14) Oeconomides,Bull. Soc. Chem., 35,498 (1881). (15) Perkin, W. H . , J . prakt. Chem. ( 2 ) , 3 1 , 4 9 3 (1885). (16) Porgorshelski, Z., J . Phys. C h m . (U.S.S.R.), 36, 1129-84
(9) (10) (11) (12) (13)
(1904); Chem.Zentr., 75 ( l ) , 668 (1905). (17) Rogers, A . D., and Nelson, R . E., J . Am. Chem. Soc., 58, 1027 (1936). (18) Taft, R. W., Jr., M. A. thesis, p. 33, Univ. of Kansas, 1946. (19) Ibid., pp. 15-21. (20) Taft, R. W., Jr., and Stratton, G. W., Trans. Kansas Acad. Sci., 48, 319 (1945). (21) Ibid.,50, 225 (1947). (22) Taft, R. W., Jr., and VanderWerf, C. A., J . Chem. Education, 23, 8 2 (1946). (23) Timmermans, J., and Martin, J.Chim. Phys., 23,778 (1926). (24) Tishchenko, D. V., J.Gen. Chem. (U.S.S.R.), 8 , 1232-45 (1938). 125) Underwood, W., and Gale, J. C., J . Am. Chem. Soc., 56, 2119 (1934). (26) Vaughan, W. E., andRust, F. F.,Org. Chem., 5,449-71 (1940). (27) Whaley, B. R., private comniunication, Halogen Chemicals.
a.
RBCEIVED March 6, 1947.
(1936).
HYDROTROPIC SOLUBILITIES Solubilities i n 4 0 Per Cent Sodium Xylenes ulfo na te HAROLD SIMMONS BOOTH AND HOWARD E. EVERSON Western Reserve University, Cleveland 6 , Ohio
A
study was made of the solubility of a variety of materials in aqueous 40% sodium xylenesulfonate solutions at 25.0” C. At the same time, and under as nearly like conditions as possible, the solubility of these materials was determined in distilled water. This afforded a means of indicating the solvent power of the aqueous 40% sodium xylenesulfonate solutions. When solubility of the solute in water is measurable, its solubility in aqueous 40% sodium xylenesulfonate is generally considerably greater than in water.
I
N A limited sense, hydrotropic solvents are aqueous solutions of salts that cause greater solubility of insoluble or slightly water-soluble substances than does pure water a t the same temperature. Neuberg ( I S ) reported this phenomenon in 1916 and made a rather extensive study of a number of these materials. Because these salts increase the solubility of many slightly water-soluble or insoluble materials, many industrial applications are possible. These can be divided roughly ifito four classes: 1. Crystallization media, as in the use of aqueous calcium cymene sulfonate solution for the purification of benzoic acid, sulfanilic acid, and salicylic acid (5). 2. Selective solvents in extraction of one material from another, as in separating lignin from cellulose in the McKee pulp process ( 3 , 6 , 7 ) . In a similar process, Lau (4)made a study of the extraction of lignin from bamboo pulp and gave the advantages of this method as: production of a higher yield of bamboo pulp compared with that from other processes; repeated use of the solution, with simple and complete recovery of the chemicals employed; no evil smelling gas and no difficulty in disposing of waste liquor; and no chemical other than the hydrotropic solution needed. A selective solvent would be advantageous in the separation of materials with similar boiling points, such as aniline and dimethyl-
aniline ( 3 ) . After extraction of one material from another, the solute generally can be removed, to a considerable degree, from the hydrotropic solution by dilution with water, giving solvent and solute layers that can be separated easily. After this second separation, the aqueous hydrotropic solution can be reconcentrated and recycled through the separation process. 3. Reaction media for certain processes that might otherwise require volatile and flammable solvents. This is advantageous from the standpoint of elimination of explosion hazards, toxic vapors, and loss of solvent by evaporation. A very good example of this is a process for the synthesis of etlylenediamine (8). 4. Finallv, solvents in the field of electrochemistry. Certain water-insoluble materials can be dissolved and electrolyzed with comparative ease. Here again, solvent hazards are eliminated. I n addition, lower voltages can be employed as many of the solutions of hydrotropic solvents are good conductors. This has been shown by the process of electroreduction of aromatic nitro compounds in an aqueous solution of hydrotropic salt (9, IO). Electro-oxidation can be carried out, as shown by the work of McKee and Heard ( I I ) , in the electrolytic oxidation of benzaldehyde t o benzoic acid in hydrotropic solution. Further applications of hydrotropic solutions in industry are discussed in an excellent review by McKee (6). The phenomenon of hydrotrophyis onewhichmany investigators have tried to explain. Makara ( l a ) presents a summary of the various theories. Of these, perhaps t h a t of Bancroft (1) gives the best explanation on the basis of a mixed solvent theory according to which the hydrotropic salt dissolves in a solvent, such as water, and in solution the salt itself then acts as a solvent. These theories can be summarized in the statement that “like dissolves like.” The object of this investigation was to study the hydrotropic solvent action of aqueous sodium xylenesulfonate solutions. Of the various known hydrotropic solvents, sodium xylenesulfonate was selected for this study because of its cheapness, availability, and stability of its aqueous solutions; and because of the
