[CONTRIBUTION FROM
THE
CHEMICAL LABORATORY OF THE UNIVERSITY OF CALIFORNIA ]
THE THERMAL REARRANGEMENT OF N-CHLOROACETANILIDE I N AQUEOUS SOLUTION A. R. OLSON
AND
J. C. HORNEL*
Received April 99, 1938
During the fifty years which have elapsed since Bender’ discovered N-chloroacetanilide and its rearrangement to chloroacetanilide under the catalytic influence of hydrochloric acid, many papers have been published on this subject. These articles, for the purposes of the present discussion, can be catalogued as follows: (1) investigations in aqueous solutions; (2) investigations in non-aqueous solutions; (3) investigations in the dry state. Each of these classifications can be divided into (a) thermal and (b) photochemical investigations, and still again into those concerned primarily with the theory of the reaction, and those consisting mainly of tables of rates of reaction. In still another group of papers, the authors have accepted the mechanism of the N-chloroacetanilide rearrangement as proved, and have used the conclusions with respect to this reaction to bolster their arguments for or against a mechanism of some other reaction, such as halogenation, or in support of some general theory of reaction. We shall concern ourselves in this paper largely with the thermal mechanism of the “rearrangement” in aqueous solution. It cannot be too strongly emphasized that any remodelling of the mechanism in aqueous solutions makes it imperative to reexamine the superstructure based upon this foundation. The early work in this field was strongly influenced by Armstrong’s2 views on substitution reactions. He contended that something more than a mere interchange of positions of radicals took place when N-chloroacetanilide rearranged. Late9 when he found that hydrochloric acid is a specific catalyst for this rearrangement he stated : “It appears legitimate to assume that it (the isomeric change) is dependent on the combination of chloroamine with hydrogen chloride. A condition of extreme instability is thus engendered, and probably the first consequent change is one in which an atom of chlorine attached to the nitrogen atom escapes from the
* Commonwealth Fund Fellow. BENDER, Ber., 19, 2272 (1886). ARMSTRONG, Brit. Assoc. Reports, 1899, p. 685. a ARMSTRONQ, J . Chem. Soc., 77, 1047 (1900). 76 2
REARRANGEMENT OF N-CHLOROACETANILIDE
77
molecule together with an atom of hydrogen from the nucleus; a chlorine atom then slips into the nucleus in place of the latter, whilst the atom of hydrogen introduced in the molecule of hydrochloric acid takes the place of the chlorine atom of the chloroamine, the ortho or para derivative being formed according to the conditions prevailing at the moment of change.” Blanksma4 found that the rate of the reaction was monomolecular, and so he concluded that the reaction was truly intramolecular. Ortons adopted this same viewpoint when he suggested that Armstrong’s pentavalent nitrogen compounds might rearrange by way of a quinonoid path. Later Orton and Jones6 mention a “strong prima facie case” for the quinquavalent nitrogen intermediate, but in another article they’ definitely abandoned this point of view. In the meantime, Acree and Johnsons thought that they had found definite proof of this mechanism when they obtained the same products from N-bromoacetanilide and hydrochloric acid, as from N-chloroacetanilide and hydrobromic acid. The idea of a true intramolecular rearrangement persists in articles by Bell.9 He however also considers another path, and concludes that the solvent composition would have a large influence on the relative amounts going by the two paths. Thus, “In aqueous hydrochloric acid the rearrangement of N-chloroacetanilide takes place chiefly” by the other path. Olson, Porter, Long, and Halford’” using radioactive hydrochloric acid as a catalyst, showed that a pentavalent nitrogen intermediate, in which the two chlorines are equivalent, is impossible. They concluded, however, that under their experimental conditions forty per cent. of the N-chloroacetanilide rearranged without coming into radioactive equilibrium with the chloride ion in solution i f the concentration of chloride ion remained constant during the experiment. This result, they stated, was in agreement with the assumption that this fraction rearranged intramolecularly. The genesis of another mechanism is found in Chattaway and Orton’sll discovery that N-chloroacetanilide and hydrochloric acid produce chlorine and acetanilide. The possibility that hypochlorous acid is an intermediate in the chlorine formation is seen from theirI2 experiments on substituted BLANKSMA, Rec. trav. chim., 21, 366 (1902); 22,290 (1903). ORTON,Proc. Roy. SOC.,(London), 71,153 (1902). 6 ORTON AND J O N E S , ~ Chem. . Soc., 96, 1457 (1909). 7 ORTON AND JONES,Proc. Chem. Sac., 26,233 (1909). * ACREEAND JOHNSON, Am. Chem. J . , 37,410 (1906);38,258 (1907). 9 BELL,J . Chem. Sac., 1936, 1154. 1 0 OLSON, PORTER, LONG,AND HALFORD, J . Am. Chem. Soc., 68,2467 (1936). 11 CEATTAWAY AND ORTON, J . Chem. SOC.,76, 1045 (1899). 12 CHATTAWAY AND ORTON,ibid., 77, 134 (1900). 4
6
78
A. R. OLSON A N D J. C. HORNEL
nitrogen chlorides with acids on which hypochlorous acid has no action. Later Orton and Jones6 stated that no evidence had been found which required the intermediate formation of hypochlorous acid during the reaction. In the same article, Orton and Jones stated that the equilibrium, ArNHAc
+ Clz
ArNClAc
+ HC1
always was rapidly established, and then slowly displaced by the formation of chloroanilide. The solvent composition had a large influence on the equilibrium-in glacial acetic acid, hydrochloric acid was not detectable. About the same time, they7 made very definite statements in favor of the chlorine mechanism. Again the following year, they13mentioned the two mechanisms and adhered to the one based upon the direct action of chlorine on the anilide or on a “dynamic isomeride” of it (“quinonoid”). They showed also that the velocity of the reaction was increased by the addition of acetanilide. Barnes and Porter1*found that in glacial acetic acid, the maximum rate of transformation of N-chloroacetanilide obtained by the addition of j3-acetnaphthalide was greater than that which could be obtained by the addition of acetanilide. They stated that this was in conflict with the predictions based on the theory of Orton and Jones but they offered no satisfactory substitute for that theory. Olson, Porter, Long, and Halford’O showed that sixty per cent. of the chlorine came into radioactive equilibrium with the chloride ion in solution before reacting to form chloroacetanilide. This was in harmony with the idea that the mechanism involved a small steady-state concentration of chlorine, but definitely excluded a real equilibrium between chlorine and the reactants. The conclusion that chlorine, if present, must be at a steady-state concentration can be got from the experiments of Orton and King,16who determined the relative amounts of C- and N-chlorination of acetanilide by chlorine. Orton, Soper, and WilliamP rate determinations of C- and N-chlorination of acetanilide by chlorine lead to the same decision. Olson, Halford, and Hornel,’7 correcting for the chloride ion which was produced during the experiment, concluded that all of the N-chloroacetanilide which reacted to form chloroacetanilide came into radioactive equilibrium with the catalyzing chloride ions, and it therefore had passed through the chlorine stage. The production and disappearance of chlorine as steps in the series of reactions involved in this rearrangement thus appears to be well established. The rate of the =arrangement usually ORTONAND JONES,Brit. Assoc. Reports, 1910,p. 85. BARNES AND PORTER, J . Am. Chem. Soc., 62,2973 (1930). 16 ORTON AND KING,J . Chem. SOC.,99,1369 (1911). 16 ORTON,SOPER, AND WILLIAMS,ibid., 1928, 998. 17 OLSON,HALFORD, AND HORNEL, J . Am. Chem. SOC.,69, 1613 (1937).
1%
14
REARRANGEMENT OF N-CHLOROACETANILIDE
79
is followed by determining the loss in the power of solutions to oxidize iodide ion. The production of chlorine involves no such loss. The rate work on this problem therefore involves at least two assumptions-first, that the steady-state concentration of chlorine has been attained, and second that no side reactions occur unless proper allowances have been made for them. Blanksma4 proved that the rate of the rearrangement as followed by the determination of loss of oxidizing power, involves the concentration of N-chloroacetanilide to the first power, and that of hydrochloric acid to the second power. Acree and Johnsons showed that the appearance of the second power of hydrochloric acid concentration in the rate law might be due to the fact that the reaction is catalyzed by hydrogen and chloride ions. Orton and Jones? found that as the concentration of acetic acid was increased, the second power of hydrochloric acid in the rate law changed gradually to the first power. The change was rapid at about sixty-five per cent. acetic acid. Harned and Seltzl*showed that the rate of rearrangement in aqueous hydrochloric acid was proportional to the product of the activities of hydrogen and chloride ions, rather than the concentration product, but that this simple rate law was not obeyed if salts, such as sodium chloride, were added. Soperlg determined the hydrolysis constant of N-chloroacetanilide. Combining this with the hydrolysis constant of chlorine, and the rates of C- and N-chlorination of acetanilide, he was unable to show that the production of chlorine was the slow step in the Orton series of reactions. He emphasized the necessity for adding enough acetanilide to obtain the maximum rate of rearrangement in order to suppress the N-chlorination of chloroacetanilides. The activity rate law proposed by Harned and Seltzl*was checked by Soper and Pryde20after they had corrected for the activity of N-chloroacetanilide and for a side reaction involving hydrolysis of the N-chloroacetanilide. They stated that their results were not in harmony either with a concentration theory or with the Brgnsted theory. Belton2Istudied the rate of rearrangement in aqueous hydrochloric acid solution to which various amounts of (1) sodium chloride, or (2) perchloric acid had been added. In the second series, he found that the results could be expressed by a concentration theory rather than by an activity theory. In the first series, no simple relationship could be obtained using the concentration, the activity, or the Bronsted theory. Dawson and Milletz2not only corroborated Harned and Seltz’s’* conHARNEDAND SELTZ, ibid., 44, 1475 (1922). SOPER,J . Chem. SOC.,127, 98 (1925); J . Phys. Chem., 31, 1192 (1927). 10 SOPER AND PRYDE, J . Chem. SOC.,1937,2761. 21 BELTON, ibid., 1930, 116. ** DAWSON AND MILLET, i b i d . , 1933, 1920. 1s
19
80
A. R. OLSON AND J. C. HORNEL
clusions that the addition of salts invalidated the simple activity law for catalysis by hydrogen and chloride ions, but decided that the rate could best be represented by a law which involved the concentration of nonionized hydrochloric acid. The side reactions mentioned by Soper and PrydeZawere found by Percival and LaMer23to be insignificant since there was a complete absence of time drifts in their velocity constants. The existence of a side reaction which produces chloride ions, was first noticed by Blank~ma.~It was investigated more fully by Orton and Gray2*who thought that it was a reversible hydrolysis of N-chloroacetanilide, and a subsequent reduction of the hypochlorous acid by aniline. The additions of nitric, sulfuric and perchloric acids, in the order named, were effective in producing chloride ions. Soper and PrydeZafound a similar hydrolysis. In the experiments of Olson, Halford, and H0rne1~~ a very considerable fraction of the N-chloroacetanilide was used up in producing chloride ion. Until the influence which this side reaction exerts on loss of oxidizing power has been established, we obviously cannot use this reaction as a basis of deciding for or against the various general theories of reaction rates. The relative amounts of ortho and para chloroacetanilides formed under various conditions were investigated by Orton and Jones26who-found that in the chlorination of acetanilide by chlorine in glacial acetic acid, twice as much para as ortho chloroacetanilide was formed. Dilution of the solvent by water did not affect this ratio. On the other hand, in the chlorination by bleaching powder, or in the rearrangement, the two isomers were formed in equal amounts. These statements were revised by Orton and Bradfield2' who decided that when the acid was diluted to fifty per cent. the para-ortho ratio dropped to three halves for the chlorination of acetanilide by chlorine. Olson, Halford, and H ~ r n erecovered l~~ only fifty per cent. of the chlorine used up in the direct chlorination of acetanilide by chlorine, as p-chloroacetanilide. They therefore assumed that para and ortho chloroacetanilides are formed in equal amounts. Their solvent was aqueous ethyl alcohol. Olson, Halford, and Horne126concluded: (1) that the total loss in oxidizing power could be accounted for by measuring the production of chloroacetanilides and chloride ion, if it were assumed that two equivalents of oxidizing power were lost for each mole of chloride ion or chloroacetanilide formed. PERCIVAL AND LA MER, J . Am. Chem. SOC.,68,2413 (1936). ORTONAND GRAY,Bm't. Assoc. Reports, 1913, p . 136. 26 OLSON,HALFORD, AND HORNEL, J . Am. Chem. SOC.,69,1613 (1937). za ORTON AND JONES,J . Chem. SOC., 96,1058 (1909). p7 ORTON AND BRADFIELD, ibid., 1927, 986. 23 Z4
REARRANGEMENT OF N-CHLOROACETANILIDE
81
The present investigation is an extension of the preceding paper. We do not find it necessary to alter any conclusions stated there, but we have found that the unknown substance X, to which reference was made as being produced along with the chloride ion, is itself an oxidizing agent which contributes to the iodine titre. We have therefore amplified the experimental procedure so as to determine the concentration of this substance at various times. The N-chloroacetanilide was prepared by the method of Barnes and Porter.28 The temperature was 40" throughout this investigation. For most runs a solvent containing twenty per cent. alcohol was used. It was prepared as follows: to two hundred cubic centimeters of absolute ethyl alcohol, measured at 20°, enough water and sulfuric acid were added, so that a t 40" the solution would occupy one liter and have the required hydrogen-ion concentration. The thermostat was placed in a dark corner of the laboratory, but the reaction was not otherwise shielded from light. In making a run, the solvent was brought to temperature, and the weighed, finely divided N-chloroacetanilide was added. After solution which usually required about two minutes, a weighed amount of sodium chloride was added. Zero time was taken as the time of addition of sodium chloride. The total oxidizing power was determined in the usual way by pipetting five or ten cubic centimeters of the reaction mixture into excess potassium iodide solution, and then titrating the liberated iodine with sodium thiosulfate solution. The new oxidizing agent which we will designate as X, was determined by pipetting ten cubic centimeters of the reaction mixture into an equal volume of solvent which was saturated with sodium chloride. After standing for ten minutes at 40",excess potassium iodide solution was added and the iodine was titrated. From the experimental work to be presented, it will be evident that under these conditions the N-chloroacetanilide was completely destroyed, but X was reduced by only a negligible amount. The chloride ion was determined by chilling rapidly sixty-five cubic centimeters of the reacting solution to about 5", and then extracting twice with equal volumes of cold benzene. Chlorine and acetanilide are both preferentially soluble in benzene, and so about fifteen minutes was permitted t o elapse before separating from the aqueous phase the first time, thus permitting the chlorine and acetanilide to react, and the resulting hydrochloric acid to be reextracted by the water. After the second extraction the water layer was filtered through ordinary dry filter paper. To fifty cubic centimeters of this solution, excess silver nitrate was added, and the silver chloride was determined gravimetrically. In Fig. I and Fig. 11, we have plotted the experimental results for two runs. Other runs using intermediate initial chloride ion concentrations 28
BARNESAND PORTER, J . A m . Chem. Soc., 62, 1721 (1930).
82
A. R. OLSON AND J. C. HORNEL
gave intermediate results. Fig. I11 is for a run where the initial N-chloroacetanilide concentration was halved; Fig. IV for a run where the hydrogenion concentration was halved, and finally Fig. V for a run using twentyfive per cent. alcohol. Various experiments also were performed on the addition of acetanilide and ortho and para chloroacetanilides. Only the acetanilide had a definite accelerating effect on the overall oxidizing rate, which was particularly marked in those solutions which were low in
0
FIG. I. Initial concentration of N-chloroacetanilide is 0.04 moles/liter. Initial concentration of chloride ion is 0.04 moles/liter. Initial concentration of hydrogen ion is 1.43 moles/liter. 20% alcohol. Curve No. 1 represents the total oxidizing power. Curve No. 2 represents the increase in chloride ion. Curve No. 3 represents the concentration of X. Curve No. 4 represents the N-chloroacetanilide plus chlorine. Curve No. 5 represents the production of ortho and parachloroacetanilides.
chloride ion. The effects on the production of X and chloride ion were small. It will be noticed that the curve for X in Fig. I1 divides into two branches at t = 200 minutes. At this time one hundred cubic centimeters of solution was withdrawn from our reaction flask and put into an equal volume of solvent saturated with sodium chloride. At definite intervals samples were withdrawn from the mixture to determine the rate a t which X disappears under those conditions. The rate was so slow that we have
0
50
IO0
200
150
250
350
300
400
450
500
Time, minutes
FIG.11. Initial concentration of N-chloroacetanilide is 0.04 moles/liter. Initial concentration of chloride ion is 0.005 moles/liter. Initial concentration of hydrogen ion is 1.43 moles/liter. 20% alcohol. Curve No. 1 represents the total oxidizing power. Curve No. 2 represents the increase in chloride ion. Curve No. 3 represents the concentration of X. Curve No. 4 represents the N-chloroacetonilide plus chlorine. Curve No. 5 represents the production of ortho and parachloroacetanilides.
‘I I
0
25
50
I
75
I
Time, minutes FIG,111. Initial concentration of N-chloroacetanilide is 0.02 moles/liter. Initial concentration of chloride ion is 0.04 moles/liter. Initial concentration of hydrogen ion is 1.43 moles/liter. 20% alcohol. Curve No. 1 represents the total oxidizing power. Curve No. 2 represents the increase in chloride ion. Curve No. 3 represents the concentration of X. Curve No. 4 represents the N-chloroacetanilide plus chlorine. 83
8
G
Q-4
E 2
2
0 0
20
40
60
80
Time, minutes FIG,IV. Initial concentration of N-chloroacetanilide is 0.04 moles/liter. Initial concentration of chloride ion is 0.04 moles/liter. Initial concentration of hydrogen ion is 0.717 moles/liter. 20% alcohol. Curve No. 1 represents the total oxidizing power. Curve No. 2 represents the increase in chloride ion. Curve No. 3 represents the concentration of X. Curve No. 4 represents the N-chloroacetanilide plus chlorine.
FIG.V. Initial concentration of N-chloroacetanilide is 0.04 moles/liter. Initial concentration of chloride ion is 0.04 moles/liter. Initial concentration of hydrogen ion is 1.43 moles/liter. 25% alcohol. Curve No. 1 represents the total oxidizing power. Curve No. 2 represents the increase in chloride ion. Curve No. 3 represents the concentration of X. Curve No. 4 represents the N-chloroacetanilide plus chlorine. a4
REARRANGEMENT OF N-CHLOROACETANILIDE
85
neglected to make a correction for the ten minutes interval which was normally present in all our determinations. From the curves in the above figures, we see that the initial rate of production of X as well as that of chloride ion, (a) is independent of the concentration of chloride ion, i.e. the initial slope of the X curve in Fig. I is the same as the slope of the corresponding curve in Fig. 11; (b) is directly proportional to the initial concentration of N-chloroacetanilide, cf. Fig. I and Fig. 111; (c) is directly proportional to the hydrogen-ion concentration, and (d) is somewhat lowered by an increase in the alcohol concentration. We see furthermore that the initial slope of the chloride ion curve is twice that of the X curve. We must conclude therefore that the initial production of chloride ion and of X are not independent processes. Furthermore when a sufficient time has elapsed so that the oxidizing power of the solution is due solely to X, the chloride ion concentration continues to increase. In fact a t this state of the reaction, the increase in the chloride ion concentration is equal to the decrease in the concentration of X. We therefore are led to postulate that the side reaction which produces chloride ions, involves first a unimolecular activation of a molecule of N-chloroacetanilide, possibly to a quinonoid structure, followed by a rapid reaction with two more N-chloroacetanilide molecules, producing X and two chloride ions and losing four equivalents of oxidizing power; X then slowly decomposes to produce Y, which is not an oxidizing agent, and a chloride ion. Keeping the hydrogen-ion and alcohol concentrations constant, the simplest mechanism which occurs to us in which cognizance is taken of these experimental data is: NCI
+ C1- + H+
P,
ki k4
CH
+ Clz
ks
CCI
+ C1- + H+
NC1’ 2NC1
+
lk6
x + 2c1lkG
Y
+ c1-
where we have designated N-chloroacetanilide by NC1, acetanilide by CH, and chloroacetanilide by CC1. From what has been said, k6 is so fast that the production of X is measured by kz, and so we can neglect ks; ks can be got from Fig. I1 where it is shown to be small. Likewise k4 from the
86
A. R. OLSON AND J. C. HORNEL
work of Soper and his collaborators as well as from the radioactive work is known to be small compared to ka, and therefore we can ignore it here. In addition it can be seen that the concentrations of acetanilide and chlorine are equal unless other chlorinatable substances are present. The rate of the reaction as usually followed gives the change in total oxidizing power, Le.
-
(”(y)+
y )+ ’XI). dt
When account is taken of
the corresponding change of acetanilide, we see that the titration by the
+(‘l). dt
(”-) conservation law also is equal to dt
Therefore when we
substract and total increase in chloride ion, whether it be present as such or temporarily combined as chlorine, from the loss in oxidizing power, we get a curve representing the production of ortho and para chloroacetanilides. We thus obtain curves No. 5 in Fig. I and Fig. 11. The induction period which is so marked in Fig. I1 must also be present in Fig. I. Its absence means merely that we haven’t accurate enough data near the beginning of the reaction. The form of this curve is in complete accord with Orton’s chlorine intermediate mechanism, the concentration of chlorine building up from nothing at zero time to its maximum steadystate value, and then gradually decreasing. It is no doubt this steady state which various authors had in mind when they discussed an “equilibrium.” The influence which factors have on the steady-state concentration gives at least a partial explanation of a number of phenomena which have been observed by various investigators. Thus an increase in the concentration of N-chloroacetanilide, chloride ion or hydrogen ion increases the rate of formation of chlorine. These factors apparently have no effect on the rate of disappearance of chlorine, and so the steady state concentration should be increased. An increase in the concentration of substances like alcohol, increases the activity of ions, but decreases the activity of the N-chloroacetanilide, Le., makes it more soluble. The change in the rate of formation will depend on which effect is predominant. On the other hand the rate of disappearance is decreased because of the decreased activities of chlorine and acetanilide, probably resulting in an increased concentration of chlorine. The addition of chlorinatable substances like acetanilide has no effect on the rate of production but a marked effect on the rate of disappearance of chlorine, resulting in a decrease of its steadystate concentration. The effect on the observed rate of disappearance of total oxidizing power is much more complex. Thus the rate will be proportional to the first power of the activity of N-chloroacetanilide and of hydrogen ions if the concentration of chlorine is kept essentially constant. The rate also will
REARRANGEMENT OF N-CHLOROACETANILIDE
87
be proportional to the activity of chloride ions, if the above condition is fulfilled, and if, in addition, the production of X is negligible. Where one of the above reactants is present initially to only a small amount, the effect of depletion of this constituent due to the steady-state condition of chlorine and acetanilide becomes very noticeable. Thus in the experiment shown in Fig. 11, the building up of the steady-state concentration of chlorine and acetanilide decreases the concentration of chloride ion to less than one-half of its initial value in spite of the production of chloride ion by the X reaction. This effectively checks the further disappearance of N-chloroacetanilide by this path. The rate is increased by the regeneration of chloride ion as chloroacetanilide is formed, and as chloride ion is produced by the other path. This accounts for the induction period in the overall rate in Fig. 11,for the production of chlorine involves no loss in oxidizing power. The addition of acetanilide or acetnaphthalide increases the rate by decreasing the steady state concentration of chlorine, thus increasing the concentration of chloride ion. This explains why the addition of acetanilide is especially effective when the concentration of chloride ion is small. The small decrease in the production of X on the addition of acetanilide may be due entirely to the increased amount of N-chloroacetanilide which disappears by the chlorine path. No doubt effects analogous to those described could be found if the concentrations of the other reactants were decreased. The extensive tables of reaction rates published by Rivett25 and by FonteinSomight prove to be valuable sources of information in this connection. Throughout this article we have been interested in mapping the course of the reactions, and in the qualitative effect of various factors and not in a determination of absolute rate constants. However we frequently were aided in the approach to our conclusions by a more quantitative procedure which we now sketch. From the curve for total oxidizing power and the curve for X we can obtain an approximate curve for N-chloroacetanilide as shown in Fig. I and Fig. 11. It differs from the true curve by the chlorine concentration. We can fit the total chloride ion concentration by the empirical expression (C1-) = a
+ be-a* = F (t)
where a is the initial concentration of chloride ion, b is the initial slope, c is a constant characteristic of each run, and t is the time, which must not be taken too large. From this we could obtain the true concentration of chloride ion if at the various times we knew the amount of chloride ion ID 80
RIVETT,2. physik. Chem., 82, 201 (1913); 86, 113 (1913). FONTEIN, Rec. trav. chim., 47,635 (1928).
88
A. R. OLSON AND J. C. HORNEL
which is combined as chlorine. Lacking this information, we can substitute the approximate expression in d(NC1) -dt
-
-kl(NC1)[F(t)]
- k2(NC1).
This can be integrated directly, and so first approximations to kl and h can be obtained. From this we get kl = 0.60 moles/min. and k p = 3.6 X 10" moles/min. Likewise an approximate value of k3 can be got from the chloroanilide curve by assuming that the concentration of chlorine is stationary at some value of t. Thus in Fig. II., when t = 30 min., d (CCl/dt = 5.56 X moles/min., (NC1) = 0.0339, (C1-)t,t,l = 0.0086. The condition for constant chlorine concentration permits us to write k1
(0.0339
- Z) (0.0086 - X)
= 5.56 X
lo-' = k39,
where z = (Clz) = (acetanilide). Solving, we find 2 = 0.0066, k3/kl = 1.275. TABLE t
30 60
90
(C1-h
0.0086
.0119 ,0147
(NCU
X
0.0339 .0271 .OB7
0.0066
.0077 .008
1.275 1.37 1.32
0.0020
.0042 .0067
We have tabulated the results of calculations at several times. If it were necessary, these approximate values could be used to obtain the next higher approximations, etc., until a self-consistent set of data were obtained. SUMMARY
It has been shown that in aqueous solutions N-chloroacetanilide disappears by at least two paths. In one path hydrogen ion and chloride ion react with N-chloroacetanilide to build up a steady-state, as distinguished from an equilibrium, concentration of chlorine and acetanilide. These substances then react to form ortho and para chloroacetanilides, regenerating hydrogen and chloride ions. The addition of chlorinatable substances like acetanilide or acetnaphthalide increases the rate of this reaction by decreasing the amount of chloride ion fixed as chlorine. This path thus conforms to the mechanism proposed by Orton and Jones. In the second path which is catalysed only by hydrogen ion three molecules of N-chloroacetanilide condense to form a new compound which we
REARRAIWEMENT OF N-CHLOROACETANILIDE
89
have called X and two chloride ions. X, in acid solution, oxidises iodide ion instantaneously, bromide ion measureably fast and chloride ion only very slowly. X furthermore decomposes slowly, losing its oxidising power to form another chloride ion and some new compound. At 40°, using 20 per cent. aqueous alcohol as solvent, about 70 per cent. of the N-chloroacetanilide disappeared by the first path when the initial concentration of chloride ion was 0.04 molar. When the initial concentration of chloride ion was reduced to 0.005 molar, the amount going by that path was reduced to about 25 per cent. The second reaction apparently is prominent enough to warrant a very full investigation of it before any general conclusions based upon the quantitative study of this rearrangement can be accepted.
