The Current State of Positive Halogenating Agents
Ernst Berliner Bryn Mawr College Bryn Mawr, Pennsylvania
The possibility that positive halogen compounds might he the substituting agents in some electrophilic aromatic halogenations has been the subject of speculation and investigation for a long time. After it was firmly established that the substituting agents in usual aromatic substitutions are electrophilic, it was natural to look for positive halogens, which ought to he the most active halogenating species; and the search was intensified after the isolation and study of the nitronium ion by Ingold and his school (1). It must he stated a t the outset that in uncatalyzed halogenation by molecular chlorine or bromine, the real substituting agents have repeatedly been shown to be the free molecules of chlorine or bromine (2, 5). They are not precursors of species that might be derived from them by dissociation, such as Cl+ or Br+, or by hydrolysis, such as HOCI, HOBr, H20CI*, or H20Br+. That they might be was first suggested by Francis (4), but the proof that they are not is fairly straightforward. A species such as Brf could, for instance, be formed Br-. Its conthrough the dissociation B h s Br+ centration would be inversely proportional to the bromide ion concentration,
+
[Br+l
=
K.,[Brpll[Br-I
(1)
and if Br+ were the effective hrominating agent, the rate of bromination would also be inversely proportional to the bromide ion concentration. Such a dependence has never been observed, and one may conclude that in uncatalyzed halogenation by Br* or CI2the negative ion is always part of the transition state. The proof that HOBr or HOCI are not involved is similar. The situation may well be different in catalyzed halogenations by Clz or Br2. It was suggested as early as 1925 that the function of the catalyst is to polarize the halogen and to help in the formation of an ionic organic intermediate (5). Later, this often came to be represented as an equilibrium involving halogen cations, for instance, Brr
+ FeBr, F? (Brt)(FeBr4-)
(2)
This formulation has always been an attractive hypothesis, but no more. I t has never been proved, kinetically or otherwise, and not much work has been done on these systems. If these ions are formed as depicted, i t is not known if they are separate ions, if they are part of an ion pair, or if the Lewis acid functions by merely polarizing the halogen molecule (6). In fact, it is not even established if the catalyst and halogen are involved in a pre-equilibrium, as often written, or if the catalyst functions after the aromatic compound and the halogen have reacted to form an intermediate, possibly of the type HArBr-Br (7). In the present discussion we will not be concerned 124 / Journol of Chemicol Education
with the positive halogens that might be formed from halogens and a catalyst. Instead, attention will be focused on the evidence for positive halogens derived from other sources. Almost all of this evidence is of a kinetic nature, and therefore manifests all of the elegance as well as the frustrations (and sometimes ambiguities) which proofs that are based only on kinetics often carry with them. For the sake of compactness, positive halogen compounds derived from N-halo-compounds will be omitted from the discussion. By the middle 1950's good evidence had been accumulated which indicated that halogen compounds with a unit positive charge might be involved in some halogenation reactions. Brominations
In 1939 Shilov and Kaniaev published an important paper on the bromination of sodium m-anisolesulfonate with hypobromous acid (8). In slightly alkaline solution the rate law for bromination is -d[HOBrl/dt
=
L[ArH][HOBrI
(3)
where ArH stands for the aromatic compound, and the reaction is very slow. However, the addition of mineral acids (HN08 or HzS04)increases the rate considerably, and the rate was then found to be directly proportional to the hydrogen ion concentration: -d[HOBr]/dt
=
kr2[ArHI[HOBrl [H+]
(4)
The term [HOBr] [H+]was taken to beameasureof the formation of positive bromine by H+
+ HOBr = Br+ + Hs0
(5)
or, as it seems more likely today, by Hi
+ HOBr = HIOBrf
(6)
When the catalyzing acids were HBr or HCI, their anions also played a part, and new substituting agents seemed to be formed, such as BrPby the reaction, HOBr
+ H + + Br-
=
Bn
+ H2O
(7)
or BrC1, and the rates of reaction of the various species could be compared. The rate of reaction of H20Br+ could not be determined directly because its concentration was not known with certainty, but the authors made a minimum estimate and arrived a t the following relative reactivities of the various species: u ~ ~ + n
>110,000
BrCl 43,000
Brz 80
HOBr 0.12
This is a significant order of reactivity, which confirms a suggestion which had been put forth by Ingold several years earlier: the halogen carrier, halogen-X, will be
the more powerful, the more electronegative the "carrier" X is (9). The above order is in agreement with this postulate. In particular, the sequence Brz > HOBr had already been noted by Francis (4), and the order CL > HOCl by Soper and Smith (lo), and both have been repeatedly confirmed in subsequent years. I n the above discussion both of the formulations, Br+ and HzOBr+, have been used to represent positive bromine. I t will now be necessary to digress briefly and to inquire if there are two distinct species--or if the latter is merely a solvated Br+and if the two can be distinguished. The formulas C1+, Br+, I+, are meant to represent the halogen cations, the halinium ions (I), not covalently bound by any nucleophiles. However, if they exist in water, they would certainly be hydrated, l i e sodium ions, and are perhaps better represented as X+aq (11). Unlike sodium ions, however, these cations have only six electrons in the outermost shell and should, by Hund's rule, have two unpaired electrons, i.e., :~1:. This prediction can be tested experimentally, because the cations should be paramagnetic with a magnetic moment of 2.8 Bohr magnetons, corresponding to two unpaired electrons. There has been a considerable amount of work on solutions of iodine in concentrated sulfuric acid and oleum, mainly by Symons and his students, and all evidence points to the following reaction of iodine in 65y0 oleum (11) :
The blue solution contains a species which has one positive charge per iodine atom, which is not protonated and which is indeed paramagnetic with a magnetic moment of 1.5 B.M. A possible reason for this somewhat smaller value has been offered (11). It has been known for many years that iodine-oleum solutions contain a powerful iodinating agent, which will iodinate quite unreactive aromatic compounds, such as phthalic anhydride (I@, but some of these solutions, depending on the SO3content, appear to contain positive iodine of higher oxidation states, such as I+3or I+=. NO evidence has been found for the existence of C1+ or Br+ in oleurn solutions (11). The species HzQX+,called by Ingold the hypohalous acidium ions (I), are the conjugate acids of the hypohalous acids, or protonated hypohalous acids. They are not "solvated" halogen cations, because they cont,ain a covalent halogen-oxygen bond and do not contain anpaired electrons. They represent distinct species, and a case has been made by Millen (IS) that they are separated from the ions X + ( q ) by an energy barrier, and are only slowly interconvertible. The reasoning is that the halogen cations, being in triplet states, cannot react with water without prior reorganization of the electrons into a state in which all electrons are paired and one orbital is vacant. This process requires energy, and hence there will be an activation energy for the reactionX+ HzO = HZOX+. In aqueous solutions it will usually not be possible to distinguish these two species kinetically, because they differ only by one molecule of water. The important equation (4) has been confirmed by others, and for other systems. In 1949 Wilson and Soper studied the bromination of benzene and of ouitroani:iole by Brz and by HOBr in water (14). They
+
found that bromination by HOBr is slow, but that HOBr reacts faster tban Brz on addition of acids. The faster reaction can therefore not be due to a conversion of HOBr to Brz, and was ascribed to HzOBr+. Derbyshire and Waters also confirmed the equation of Shilov and Kaniaev for benzoic acid and benzylsulfonic acid, two compounds which are not attacked by bromine water (16). Benzylsulfonic acid does not react with HOBr a t pH 5.8, but immediately reacted a t pH 2. Molecular bromine was estimated to be about 2000 times as fast as HOBr against sodium panisate a t pH 7-8, when the concentration of H20Br+ is very insignificant. In 1954 Branch and Jones extended the reaction of HOBr to the bromination of substituted aromatic ethers in 75% aqueous acetic acid (16). Again, the active species was more powerful than either BrZ, BrCl, or HOBr, and, on addition of acid, the rate was linearly related to the hydrogen ion concentration. In aqueous acetic acid additional brominating species must be considered, and while the authors preferred H20Br+, the protonated acylhypobromite, CHaCOOBrH+, could not be ruled out as the active species, although the unprotonated acylhypobromite, CH&OOBr, could. Since that time, both CHaCOOC1and CH3COOBr,as well as their protonated forms, have been found to be distinct substituting agents in some reactions (17, 18). Their reactivity is quite large, surprisingly, and contrary to expectation, is greater tban that of the respective halogens. An explanation which is based on a cyclic trausition state has been offered for the order CHaCOOC1> Clz (17). It can be taken for certain that the kinetic term [HOBr] [H+] represents a substituting agent different from, and more powerful than, molecular bromine. Further proof was provided in a series of important investigations initiated by de la Mare and co-workers in 1956 (19). Working witb HOBr and HCIOl in aqueous dioxane, in which the kinetic equation is identical witb that found in water and aqueous acetic acid, the authors extended the reaction to a study of substituent effects. They found that HzOBr+ behaves quite differently from Br2, and more like the nitronium ion. A set of data, typical of many others, is shown in Table 1. The numbers are partial rate factors, and the data for bromination by bromine are from the work of H. C. Brown and co-workers (to). With Brz, toluene reacts about 600 times faster than benzene, but only 35 times as fast with HzOBrf, and the distribution of isomers is less discriminating with the latter reagent than with the former. Such differencesmust be due to different substituting agents, and it is in keeping with present theories to draw the conclusion that HzOBr+ is a more powerful electrophile than Br2 (as bad been found earlier by comparisons of rates) and hence a much less selective one.
Table 1
Volume 43, Number
3, March 1966
/
125
I n 1957 de la Mare and Harvey showed that HzOBr+ reacts with benzene and hexadeuterobenzene with identical rates (21). This is in agreement with a reaction in which the rate-determining step involves the attack of H20Br+on the aromatic compound.
-d[HOCl]/dl = k[H+] [Cl-] [HOCI]
Chlorinations
At the same time, very similar observations were made for chlorination by hypochlorous acid. For moderately reactive compounds, such as sodium benzylsulfonate (22)-which does not react with HOCl alone-benzene, or toluene @S), the rate law is -d[HOCl]/dt
=
k[ArHl [HOCl][H+l
(9)
which is consistent with a reaction involving either H,OCI+ or C1+, and no distinction can be made on kinetic grounds. But de la Mare and co-workers made the following additional observations (24) : if the aromatic compound is more reactive than those to which eqn. (9) applies, and within a critical concentration range, the rate law becomes -d[HOCI] = k[HOCl]
+ k'[HOCII[Ht]
(10)
and the aromatic compound (e.g. methyl ptolyl ether, anisole) does not appear in the rate equation and therefore does not take part in the rate-limiting step of the reaction. Consequently, compounds of such different reactivities as phenol, anisole, or p-dimethoxybenzene, react at identical rates, which is also the same as the rate of addition of the same reagent to certain olefins (25). The situation is therefore analogous to zeroth-order nitration, in which the slow step involves only the formation of the substituting agent ( I ) . The lack of dependence of the rate on the concentration of ArH was explained by the authors by the following series of steps:
+
-
HOCl H+ S HaOCI+ H.OCl+ s C1+ HaO Cl+ + ArH prodncts
+
fat slow fast
The slow step is suggested to be the fission of H20CI+to C1+, which then reacts with ArH as fast as it is formed. I t was also shown that under the conditions for which eqn. (9) applies, the chlorinat,ion of benzene has no isotope effect (26). If the aromatic substrate is still more reactive, or above a critical concentration, a third term appears in the rate equation which then becomes:
The aromatic compound (methyl nz-tolyl ether, phenol, anisole) is now so reactive that it cannot wait for bondfission to occur and react,swith H20C1+before the latter dissociates. The first term was suggested to represent bond fission of the unprotonated HOCl to C1+ and OH-. Because of the novelty of the interpretation, the authors made an admirably thorough study of alternative ways in which the kinetics could he explained. The most important alternative might he a rate-determining reaction of HOCl with chloride ion as in eqn. (12), which can always be presumed to he present in small amounts, and the formation of chlorine, which would then he the substituting agent. HOCl
Soper and Smith had, indeed, shown in 1926 that the chlorination of phenol by HOCl has a rate which is independent of phenol in the presence of small amounts of KC1 and catablished the rate law for this reaction as eqn. (13) (10) :
+ H + + CI-
=
Ch
126 / lournd o f Chemical Education
+ Hz0
(12)
(13)
which is the rate of formation of chlorine by eqn. (12). In order to avoid the incursion of this reaction, de la Mare et al. conducted their reactions in the presence of an excess of silver perchlorate, which will reduce the chloride ion concentration. Furthermore, the AgCIOI concentration was varied 20-fold with no change in the rate of chlorination. A rate-determining reaction involving soluble chloride can therefore be ruled out. Another possibility which the authors considered is a reaction through chlorine monoxide, CL0, which could be formed according to eqn. (14) : CIO-
+ HOCl + H+ = CI.0 + H1O
(14)
This reaction was avoided by keeping the acidity sufficiently high to make the concentration of C10negligible, and by keeping the HOG1 concentration low (0.001 M). At higher concentrations of HOCl the kinetics showed deviations in the direction of a reaction second-order in HOCI, as demanded by eqn. (14). Another important alternative is a rate-limiting proton transfer according to eqn. (15). HOC1 + H + = HIOClt
(15)
Proton transfers to oxygen are generally considered to be very fast and practically instantaneous, and the estimate has been made, based on the rate of the acidcatalyzed reaction of t-butanol with HOC1, that proton transfer to HOCl must he many times faster than zerothorder chlorination (27). A rate-determining proton transfer would also be difficult to reconcile with the third term in eqn. (11). If the proton transfer is slower than the rate of chlorination of the compounds for which the second term holds, it must also be slower than that of the more reactive compounds to which the third term applies. The third term should then not be present. Finally, Swain and Ketley showed in 1955 that the chlorination of methyl p-tolyl ether by HOCl and HCIOn is about twice as fast in D 2 0 as in Hz0 (28). This has always been taken to imply a situation in which a protonated species is formed in a fast preequilibrium (i.e. H + HOCl), followed by a ratedetermining reaction of the protonated substrate. If HOCI behaves like other substrates to which this criterion has been applied, these results rule out a ratelimiting proton transfer. Other explanations of the kinetics, which have from time to time been suggested, will he mentioned later. I n halogenations by HOCl and HOBr, it has re peatedly been noted that a t high acid concentrations the rates deviate in the direction of Hammett's acidity function Ho. De la Mare and Hilton conducted brominations with HOBr in very concentrated acids, up to 6 M HClO,, in order to see whether a t these high acid concentrations the reaction would follow the Ho function, or the function Jo, which is characteristic of the ionization of triarylcarbinols in acid solution (23, SO). A decisive result might therefore have distinguished
+
between H20Br+ and Br+. The results were not clearcut but favored the former possibility. lodinations
The positive halogen species so far discussed are derived from HOBr and HOCl as the bulk reagents. No kinetic work has been done on iodination with HOI, which is very unstable, but kinetic evidence early pointed to the possibility of the formation of a positive iodinating agent when molecular iodine is the reagent. I n 1947 Painter and Soper studied the iodination of phenol by molecular iodine in water and in the presence of buffers (Sl), and found that the kinetics of the uncatalyzed reaction were in agreement with two possible reaction paths as shown in eqns. (16) and (17), (16) -d[I*l/dl = klCsHsOH] [HOI] -d[Id/dt = ~'[CBH~O-I [H201+] (17) where in the second case no distinction can be made between H201+ and I+(aq). The ambiguity of eqns. (16) and (17) arises because the proton can be on either the phenoxide ion or the hypoiodous acid, and both reactions are chemically reasonable. A distinction between the two iodinating agents is possible when the aromatic substrate is changed to aniline (32). The iodination of aniline has the same kinetic characteristics, and two analogous possibilities arise, as in eqns. (18) and (19).
The first of these can, however, be safely eliminated on chemical grounds, because the anilinium ion should be meta-directing, whereas only ortho/para-substitution occurs in the iodination of aniline. This leaves H201+ as a possible substituting agent in a reaction which has the same kinetic form as the equation of Shilov, eqn. ( 9 , and which can be represented by the following steps (scheme A) : I2
+ HnOFIK,HIOI++ I-
must be formed in the reaction. The concentration of HzOI+ is equal to Kl[12]/[I-1, which becomes KIKr [(1~-)1/[1-]2 if eqn. (20) is taken into account. For reaction of HzOI+ the observed rate constant then be~ ,is found. The above comes lcobs = ~ I K I K ~ / [ I - ]&S interpretation of the kinetics is valid if either the attack of HzOI+ on ArH or the last step in scheme A is ratelimiting (k,b. = klk2KlK2/k-,[I-]z), and also if the reaction occurs in one step, but an alternative interpretation of the kinetics is possible, which will be presented shortly. The same kinetic characteristics as found for the iodination of phenol and aniline have repeatedly been observed in the iodination of other reactive aromatic compounds. The inverse dependence on the square of the iodide ion concentration has been noted in the iodination of a series of aniline derivatives (94) and of imidazole (55). A thorough study of the iodination of 2,4dichlorophenol, in which a stirred flow technique was employed (56), and a careful investigation of the iodination of acetyltyrosine, which made use of computer programming (37), have revealed the same dependence on iodide ion. All these reactions are in agreement with the same interpretation. Ridd and Choguill studied the reverse of the iodination reaction, i.e., the protodeiodination of p-iodoaniline as in eqn. (21) (98). The reaction is independent
of the iodide ion concentration, and from its further characteristics, they concluded that, in spite of differences in experimcntal conditions, the forward and reverse reactions have the same transition state, which involves aniline and I+, and that by the principle of microscopic reversibility, iodination and deiodination involve the same reaction paths. The kinetics of iodination is actually somewhat more complex, because catalysis by the buffer salts is also observed, in addition to an uncatalyzed reaction. The total reaction for the iodination of aniline in a buffer can, for instance, be represented by
Scheme A
The important kinetic characteristic of these and subsequent iodination reactions is a dependence of the rate on the inverse square of the iodide ion concentration, which arises as follows. Reactions are usually conducted in the presence of an excess of iodine ion, in order to minimize the effect of iodide ion formed during the reaction. I n the presence of iodide ion, molecular iodine is almost completely in the form of the triiodide ion, I2+ I- Fl ls(20) Kn
and the total titratable iodine can be set equaI to 13-. If the substituting agent were molecular iodine, and its attack on ArH were rate-controlling, the observed rate constant should be inversely proportional to the iodide ion concentration, ie., kobs = kKz/[I-] (95). Since it is inversely proportional to [I-12, another iodide ion
This implies a reaction catalyzed by the buffer base, but in buffered reactions at constant pH it is often difficult to decide if catalysis is by the acid or the base component of the buffer. Hence the alternative reaction (23) is also possible. The first terms are chemically equivalent, but in the second term
the general acid HA takes the place of the hydrogen ions in the formation of the effective iodinating species. Shilov and co-workers have repeatedly found that the iodination of some aniline derivatives in the absence of bufierv involves two molecules of amine (54), and the same observation was made by Ridd in the case of imidazole (35). The second molecule of amine could be visualized as taking the place of the buffer, which would point toward a basic catalysis. The view that iodination reactions are subject to base- rather than acid-catalysis has been stressed particularly as a result of some more recent investigations (57, 59). Some authors have described the catalysis in terms of the Volume 43, Number 3, March 1966
/
127
formation of new substituting agents, such as CHr COOI, IHIP04,or CH3COOBrand BrH2POI(in the case of bromination by HOBr), which is also a possibility, although not proved. The role of the catalyst in any of these reactions cannot be considered established, but the incursion of buffer catalysis does not alter the main features of iodination. The iodination of 2,4dichloropbenol, p-chloroaniline, and anisole by iodine monochloride in water-and in the presence of an excess of hydrogen and chloride ionshas kinetic characteristics which are amenable to the same interpretation as the iodination by iodine (40). The kinetics are in agreement with a reaction involving HzOI+, formed by hydrolysis of IC1 according to (24). ICI + HIO 7 3 HxOIf + CI(24) Reevaluation and Reinterpretation
By the middle 1950's there was thus kinetic evidence for the existence, and participation in halogenations, of H20Clt, H20Br+, H201+, or C1+, Br+, and I+. The last ten years, however, have brought forth criticism of some of the interpretations, and a subsequent reevaluation and reinterpretation of some of the evidence. The impetus came from two sides: the calculation of equilibrium constants for the formation of the positive species by Bell and Gelles (dl), and the discovery of an isotope effect in aromatic iodination by Grovenstein and 1Glby (4t). In 1951 Bell and Gelles calculated equilibrium constants for the eouilihria involved in the formation of halogen cations and hypohalous acidium ions, and drew conclusions about the possible existence of these species as kinetic intermediates (41). The calculations for the free cations involvcd the typical Habcr-Born cycle (25). The free energies for some of the processes (a, h, c)
are well known, but estimates had to be made for the solvation energies (d, e) and for ionic radii, and these were obtained by analogy. Equilibrium constants for (25) were calculated from the free energies and are shown in Table 2. By similar calculations and assumptions, the equilibrium constants for the reaction (26) were also obtained and are also listed. There were a. larger number
+ HzO 73 H20X++ X-
XB
(26)
of assumptions and extrapolations in the calculation for the hypohalous acidium ions than for the cations (the free energies of formation of the oxygen-halogen bonds and of solvation of the hypohalous acids). The authors were able to determine experimentally the ionization Table 2
Equilibrium Constants for Positive Haloaen Com~ounds( U )
128
/
Journal of Chemical Education
constant of Iz to H201+, and obtained a value of 1.2 X lo-", in good agreement with the calculated value. Bell and Gelles concluded that the equilibrium constants of all cations, X+(aq), is too small for the species to exist in appreciable concentrations in water, and that the concentration of all species H,OX+ is greater, although only H201+ exists in concentrations sufficient to be a kinetic intermediate, particularly if the iodide ion concentration is low. The treatment and conclusions of Bell and Gelles have been criticized by several authors. De la Mare and Ridd have drawn attention to the possibility that since the cations have open sextets, analogies to stable species, such as sodium ions, may be wrong (45). Arotsky and Symons have recently pointed out that in the cations some ligand field stabilization must occur, and they arrived a t figures for the cations which are uniformly greater by loz0(Cl+:10-40, Br+:10-30, 1+:10-20) (11). These higher figures are derived from a value for the ligand field stabilization, which was estimated from the spectra of iodine in oleum. The authors agree that the participation of C1+ and Br+ as kinetic intermediates is still unlikely, but that I + might exist in aqueous solutions in more than insignificant amounts. Bell and Gelles' data for the hypohalous acidium ions have been criticized more severely. Kinetic evidence points overwhelmingly to the existence of H20Cl+and H20Br+, although according to their calculations these should not exist in water in kinetically detectable quantities. Hine ( 4 4 , and Arotsky and Symons (11), made estimates of the constants, (which are based on the known aciditics of the hypohalous acids and their differences), and arrived a t values greatly a t variance with those of Bell and Gelles. I t is probably fair to conclude that the uncertainties in the latter's figures are very large (44). This problem has not been settled, but the situation is not unlike that of carhonium ions, where in the late 1930's data were presented for carbonium ion stabilities which seemed to preclude their existence as kinetic intermediates in solution, although in subsequent years these figures were successively reduced downwards to quite acceptable values. Apart from thermodynamic considerations, it is not certain what the minimum concentration of a species must be to permit its detection as a kinetic intermediate. However, the calculated and later revised figures for Cl+ and Br+ still appear to be too small, whereas it seems to be certain that the hypohalous acidium ions can exist as such intermediates. Consequently, attention is again focused on the chlorinium ion, Cl+, in zeroth-order chlorination by acidified HOCI. If Cl+(aq) cannot be present in detectable concentrations, the problem is to find a ratedetermining step which involves H + and HOC1, but which does not involve Cl+ and the aromatic compound. The hint from the kinetics is that there are likely to be two charged substituting agents; one of these is probably H20C1+and the other is a species presumably derived from it in a slow step. One alternative suggestion was recently put forward by Icupinskaya and Shilov (46). In a series of papers which make use of the data of Bell and Gelles, Shilov has expressed doubt about the participation of any
form of positive halogen. I n the case of zeroth-order chlorination, he assumes that the active chlorinating agent is invariably molecular chlorine, formed from HOCI, H,+ and Cl- according to eqn. (12). The authors carried out chlorinations of anisole by HOCI in solntions containing H + and small amounts of KCI, and a plot of koba against Cl- resulted in a straight line. If the rate of reaction is determined by the formation of chlorine according to (121, then kobs [HOCI] = k[H+] [Cl-] [HOCI], and /cobs = k[H+] [Cl-1, where k is the reverse of the rate constant of hydrolysis of chlorine. A line constructed from the latter expression lies close to the line representing the dependence of the rate on co-workers worked with an excess of silver ions to minimize the concentration of Cl-, Shilov stated that dissolved AgCl actually acts like chloride ion and hence is a source of Cl- for the formation of CIS. I n support of this view, he showed that small amounts of dissolved AgCl could replace ICCl, with no appreciable change in rate constants. Since dissolved AgCl is probably completely dissociated (46), this coincidence might be expected. But Shilov added Cl- deliberately, either in the form of ICCI or dissolved AgCl, aud, as mentioned earlier, it is well known that this leads to the formation of Clz. If the reaction in the presence of AgC104were due to dissolved AgC1, the rate should have varied with the concentration of silver ions, which was not the case in the experiments of de la Mare. Shilov's conclusion that the hypochlorous acidium ion plays no part in zeroth-order chlorination, and that the reaction actually involves molecular chlorine, can be tested experimentally. The active species in reaction of acidified HOCI, under conditions to which eqn. (9) applies, leads to partial rate factors which are entirely different from those obtained with molecular chlorine. The data for toluene are shown in Table 3 (47), and although the solvents are not the same, a change from water to acetic acid is not likely to have such a pronounced effect on the partial rate factors (48). According to present theories, differences of this type must be ascribed to different substituting agents, and chlorine cannot be the active species in these reactions. Because of the striking differencein isomer distribution in the chlorination of a series of compounds with molecular chlorine and with &butylhypochlorite in acid solutions, Harvey and Norman have similarly concluded that some form of positive chlorine must be involved in the reactions with the latter reagent (49). At present, the partial rate factors are not known for the zeroth-order chlorination, either under the conditions of the experiments of de la Mare or those of Shilov, but such data could tell if CL were involved. Shilov did not choose the experimental conditions under which zeroth-order chlorinations had been observed, and in his case chlorine may very well have been the chlorinating agent. Table 3
Arotsky and Symons have recently proposed that the rate-limiting step in zeroth-order chlorination might be a reaction involving solid AgCl (11) : AgCl(s)
+ H20Cl+ = AgCL+ + Hg0
(27)
The species AgC12+is the suggested chlorinating agent. I t is not certain if zeroth-order chlorinations were always conducted in the presence of solid AgCl, but such is the assumption in this proposal, and if AgC12+ is formed from precipitated AgCl, the lack of dependence of the rate on dissolved silver salts would be accounted for. According to the authors, there is conductometric and other evidence for the existence of the ion AgIz+ (501, aa well aa for AgBr2+ and AgClz+ (11). Although the ions AgX2- are well known, nothing is known about the stability and structure of their positive counterparts, but they are visualized as complexes between Ag+ and Xz. Another suggestion involves the rate-determining isomerization (51) : H \+
/
H
0-Cl
-
+
H-Cl-OH
(28)
+
but the species H-Cl-OH has not been considered to be a very favorable substituting agent for chlorination. A rattxontrolliog formation of the ion Cl(Hz0)2+has also been proposed (58). None of the suggestions can be considered proved or disproved, and there are undoubtedly others which could account for the kinetics. It would seem that one is still waiting for the crucial experiment which will reconcile the kinetic evidence with the thermodynamic data, and there the problem rests. Isotope Effects
Although the species H201+, and even I+(aq), now seem feasible iodinating agents, the kinetic evidence for their existence is also consistent with a different interpretation. I n 1958 Grovenstein and Kilby made the important observation that the iodination of 2,4,6-trideuterophenol proceeds about four times as slowly as the iodination of phenol (48). This is a snfficiently large isotope effect to deduce that the breaking of the C-H bond is part of the rate-determining step. Aromatic substitution is now usually discussed in terms of a stepwise, not a synchronous, reaction (53),and the above experiment is in agreement with a reaction sequence in which the proton loss is the slow step. Grovenstein and ICilhy pointed out that while these results are not in disagreement with a reaction involving H201+ [or I+(aq)], they are also in agreement with a mechanism of iodination of phenol by molecular iodine, in which 1% and ArH are at, or near, equilibrium, This is represented by Scheme B below, where for simplicity the ionization of phenol to phenoxide ion, as well as a proton acceptor in the second step, have been omitted. ArH
+ I2.kk3 ArHI+ + 1~3
Scheme B
I
k2
&E? H~OCI+,H ~ O
ArHI+ 8-50 CI?. HOAC
ArI
+ H+
If the complexing of iodine is taken into account, reaction (20), the observed rate constant k,b, becomes Volume 43,
Number 3, March 7 966
/
129
k,kXz/k-l[I-]2 and exhibits the same inverse dependence on the square of the iodide ion concentration that is characteristic of the formation of H201+. I n the case of phenol, the intermediate may be particularly stable, because of the formation of a neutral species by a loss of a proton, and probably has the structure
Compounds of this type have been isolated in the iodination and bromination of various substituted phenols. However, the stability of the intermediate is not a necessary condition for the occurrence of an isotope effect. An isotope effectwill happen if kz occurs in the rate equation, and this depends on the ratio of k-, or, in this caseof k-,[I-], to k2,i.e., on how theintermediate partitions itself between products and reactants. If the condition holds that kz >> k-l[I-], almost none of the intermediate reverts, and the first step (given hy k ~ )becomes ratecontrolling. But if k-l[I-] >> kz, almost all of the intermediate goes back to reactants, and kz, the rate constant for C-H bond-breaking, will enter into the over-all reaction rate. It was subsequently found that isotope effects are the rule in iodination, although not in hromination or chlorination. Possihle reasons for this have heen discussed elsewhere (55). Reasonably large isotope effects ( k & ~ = 3-4) have been observed in the iodination by Iz of aniline (54), methylaniline (54), anisole (54), 3phenolsulfonic acid (65), p-nitropheuol (641, imidazol (561, and of a whole series, but not all, NJTdimethyIaniline derivatives (54) in water, and of anisole by ICl in water and in acetic acid (57). The almost invariable participation of C-H bond-breaking in the rate-determining step of iodination has been considered to he in agreement with the role of the base, in buffer catalysis, as assisting the brcaking of that bond. But hecause even those aniline derivatives which are not iodinated with an isotope effect show base-catalysis (54,39), and because in the iodination of imidazole the uncatalyzed and self-catalyzed reactions have identical isotope effects (56), it has also been suggested that the basecatalysis may be due to the formation of an active iodinating agent, rather than the breaking of the C H bond (56). This question has not yet been settled. Grovenstein and co-workers have also pointed out that mechanism A, where the iodinating agent is H201+, and scheme B, where it is 12,cannot be distinguished on kinetic grounds, as has since heen repeatedly reaffirmed. I n scheme A, the inverse square dependence on iodide ion concentration arises from the equilibrium involving the formation of H201+,and in scheme B from that involving the formation of ArHI+. Kinetics can only give information about the composition of the transition state. I n this case, it tells us that the transition state consists of the aromatic compound and an iodide cation, and that the iodide anion has departed before the transition state has been reached. I n that sense, the positive cation is the substituting agent. But it does not tell us at what stage the iodide anion has dropped off 130
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and hence whether it comes from HzOI+ or from Iz. The isotope effect further tells us that in the ratedetermining transition state the C-H bond is broken, which is also in agreement with both schemes, although only a necessity for scheme B. Scheme B is sometimes referred to as the quinoid intermediate mechanism, but this is not the distinction between the two. If scheme A were to hold, it would involve the same intermediate, the usual u-complex, which is generally assumed to he involved in aromatic suhstitution, except that in the case of phenols it has become stabilized through the loss of a proton. Although under the conditions under which a full isotope effect is observed a distinction between the two mechanisms cannot be made on kinetic grounds, it should be possible to distinguish between them at the lower extremes of iodide ion concentration, This arises as follows: In scheme B, the isotope effect depends on the ratio k-l[I-]/k2, If the iodide ion concentration is gradually lowered, the hack reaction of the intermediate, on which the occurrence of an isotope effect d e pends, will become less important. The observed isotope effect will decrease, and eventually it will disappear altogether, when the first step becomes ratecontrolling. This can be seen more clearly if the intermediate is assumed to he formed in a small stationary concentration, and if the steady-state treatment is applied, eqn. (29). In the limit, when L [ I - ]
approaches zero, kobsreduces to k1K2/[1-], the first step determines the rate, and no kinetic isotope effect can be observed. I n scheme A, the isotope effect is independent of the iodide ion concentration, because the iodide ion is not a t all concerned with the substitution step. A lowering of the iodide ion concentration will merely increase the rate without affecting the isotope effect. This prediction has been tested in several instances, hut the results are not clear-cut. There are also experimental difficulties, because rates become very fast as the iodide ion concentration is lowered. I n the iodination of imidazole hy Is, the isotope effect is the same at zero initial iodide ion concentration as it is at a higher concentration (56). I n the iodination of anisole by ICI, where by the same reasoning the isot.ope effect should depend on the chloride ion concentration, the isotope effect persists from a chloride ion concentre tion of 0.9 M to a reaction conducted with no initial chloride ion present (57). However, in a very careful study, Grovenstein and Aprahamian have found that in the iodination of p-nitrophenol by 12, the isotope effect decreases from 5.4 to 2.3 as the iodide ion concentration is lowered from 2.2 X to 1.5 X lO"M, in agreement with calculated values according to scheme B (54). Similarly, Vainshtein, Tomilenko, and Shilov have observed that the isotope effect in the iodination of aniline is 3.6 when the iodide ion concentration is 0.125 M, but that it decreases to 2.4 with no initial iodide (58). Both groups of workers concluded that in these two reactions molecular iodine must be the substituting agent. Experiments of this kind are only decisive if they are positive. i.e., if they show a depression of the isotope
effect with decrease in halide ion concentration, as in the cases of p-nitrophenol and aniline. It can be argued that in the cases of imidazole and anisole, the individual rate constants are such that even a t the lowest ohtainable halide ion concentration either the halide ion formed during the reaction or that present from the hydrolysis of iodine is always sufficientlylarge to fulfill the condition that L [ I - ] >> kz (69). One further, and somewhat similar test, can be applied to the problem. At very low iodide ion concentration, the reaction should cease to be strictly proportional to the inverse square of the iodide ion concentration, but should, according to (29), approach the inverse first power if scheme B were operative, whereas the dependence on iodide ion should not change if scheme A represented the mechanism. I n the iodination of pnitrophenol, deviations in accordance with eqn. (29) were indeed encountered (64).
intervention of H201+ can therefore not be ruled out. It should be noted that in solutions of IC1 the concentration of HZOI+is much greater than in solutions of Iz, because the hydrolysis constant of IC1 is many times larger than that of I, (65) (approximately lo-' compared to 10-13). Shilov and co-workers also reconsidered the implication of the kinetic term [HOBr] [H+] and the possible involvement of H20Br+ in reactions of acidified HOBr (66). They calculated approximate collision numbers for reactions of H20Br+,which were based on the data of Bell and Gelles, and arrived a t the conclusion that the calculated numbers of collision were many times less than the actual rates of bromination (and also of iodination) of typical aromatic substrates. This is, of course, the same as saying that the concentration of H20Br+ is too small for it to be involved in hromination reactions. The authors proposed the stepwise mechanism (31), by which a complex between HOBr
Kinetic Role of XZand HOX
H
These experiments make i t reasonably certain that a t least in the iodination of p-nitrophenol and of aniline the substituting agent is molecular iodine (scheme B), and not a positive species derived from it. Because all iodination reactions have strikingly similar kinetic characteristics, one may perhaps draw the cautious generalization that a11 wueous iodinations involve molecular iodine as the actual substituting entity. Such a conclusion, however, would need further suhstantiation, and it is desirable to extend the iodination reaction to such further substrates as will show the deviations a t low iodide ion concentrations which are typical of scheme B. The possibility should also he explored that the substituting agent, Iz or H201+,may change with the substrate and the iodide ion concentration. Solutions of I, in water will always contain some H,OI+, and there have been repeated reports that the iodination of some substrates-notably heterocyclic compounds such as histidine (60) and substituted pyrroles (61), hut also substituted tertiary aniline derivatives (62) and tyrosine (63)-can be represented by two terms of different iodide ion dependence as in eqn. (30). The first term probably represents hoba =
A/[I-I
+ B/[I-1'
(30)
iodination by molecular iodine, and the second may represent reaction of HOI or H,OI+. I n these reactions the second term gains importance a t low iodide ion concentration (Of), as it should if H201+ were to become the substituting ent,ity. This effect a t low iodide ion concentration is the opposite of that observed in the iodination of p-nitrophenol. The possibility that the substituting agent may depend on the reactivity of the substrate and the concentration of iodide ion was also suggested by Ridd (55), and a similar conclusion was recently drawn by Batts and Gold (641, who based it on the results of a study of the chloride ion-catalyzed protodeiodination of iodo-2,4,6-trimethoxyhene, the reverse of iodination with ICl. Scheme B may also account for iodination with ICI, which would make molecular IC1 the iodinating entity, hut here the evidence is inconclusive. I n the one test which has been performed, the isotope effect did not vary with the chloride ion concentration (67), and the
HOBr
HxO---H+--~--B~
Br
\
Art
/
+ 2Hp0
-
products (31)
H
and ArH is formed first, and then reacts with a proton to form a protonated complex, and ultimately the u-complex. If either the protonation or the subsequent step were rate-determining, the sequence would be in accord with the observed kinetics. The kinetics would also be satisfied by a termolecular reaction between HOBr, H+, and ArH, but termolecular reactions, apart from their intrinsic rarity, are difficult to differentiate mechanistically from a reaction involving two fast successive steps, and the distinction between them becomes blurred. Inasmuch as Shilov's conclusions are based on data for H20Br+ which are now considered too low, or a t least rather uncertain, the search for an alternative substituting agent seems much less urgent, and a reaction involving H20Br+does not seem a t all unlikely. However, Shilov's scheme cannot he ruled out on kinetic grounds. Although the two-step sequence involving HOBr and H + is mechanistically different from the simultaneous substitution by H20Br+,the combination of the initially formed complex with H + may be considered to constitute a substituting agent of different electrophilicity than HOBr or Br,, and the distinct substituent effect of HOBr in acid solutions could thus be accounted for. I n a more recent paper, Shilov suggests that the reaction of HOCl in the presence of acid proceeds by a mechanism which is analogous to the bromination scheme outlincd above (31) (67). This is in agreement with the kinetics of chlorination according to (9), but not of the chlorination by HOCl of those compounds whose concentration does not appear in the rate equation (10). I t was shown in the introduction that bromination by molecular bromine does not depend on the bromide ion concentration, after allowance is made for the Volume 43, Number 3, March 1966
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formation of the tribromide ion, and that this is in accord with Brz being the substituting entity. Lately, there have been a few reports of brominations where such a dependence was observed, for instance, in the bromination of sodium p-anisolesulfonate (68), of 1 , s dimethylnaphthalene (60), and of 2,3-dihydro-4,7dimethyl-1,4diazepinium perchlorate (70). Such a dependence is not likely due to the formation of a new substituting agent, but rather to an increase in the term L [ B r - ] and the participation of C-H bondbreaking in the rate-controlling step of a sequence analogous to scheme B. This does not usually happen in bromination, but may be caused by special structural features in the aromatic substrate, as is also the case in azo coupling (71). The correctness of this latter view has been demonstrated for sodium p-anisolesulfonate by Bourns (68), who showed that the small observed isotope effect depends on the bromide ion concentration, and for 1,5dimethylnaphthalene by the demonstration of base-catalysis, which is usually not found in brominations (69). IVIolecular bromine is therefore involved in these reactions, as well as in some bromodecarboxylations and bromodesulfonations, which also show an inverse dependence on the bromide ion concentration (72). I n spite of the ambiguities which are encountered in the interpretation of the work on positive halogen compounds, it is gratifying that the work has led to some important preparative improvementasin halogenations. Substances that are very difficult to halogenate under normal conditions, including heterocyclic compounds (7S), are quite easily, sometimes instantaneously, substituted by solutions containing the halogen and silver sulfate in concentrated sulfuric acid. This method was introduced by Derbyshire and Waters for bromination and iodination (74),but it can also be used for chlorination (75). I n an extensive earlier study, the combination of IZand AgCIO&in organic solvents had been used (76). The function of the silver salt must certainly be to keep the halide ion concentration at a minimum, which would favor cation formation. Not much is known about the halogenating species in these solutions. They may be positive species, perhaps solvated by HPSOlt or such ion pairs as X+HS04-, or, as has recently been suggested (11,50), the ions AgX2+. It has been reported that in the absence of an aromatic compound the halogenation solutions do not deposit any silver halide, even after being kept at 100' for twelve hours, but that they start to deposit silver halide immediately after addition of the aromatic suhstrate (74). This may mean that as halogenation proceeds, a very unfavorable equilibrium, involving possibly a positive halogen compound, is shifted to completion by precipitation of silver halide (74),or that the silver salt functions only after the halogen and the aromatic compound have reacted to form an intermediate. More work is needed to elucidate t,he mechanism of these reactions. Conclusions
The following conclusions may be drawn from the current evidence: in the uncatalyzed halogenations by C1, and Br2 the substituting agents are the undissociated molecules, and this is probably also true in iodinations by It in aqueous solutions. It is reasonably 132
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certain that the species H201+, which may be formed from 1%) and which was formerly held responsible for iodinations, is not involved in some cases, and may not be involved in any. More work is needed to confirm this generalization, and to ascertain if in some situations both I? and H201+may participate. The kinetic terms [HOCI] [H+] and [HOBr] [H+], which are observed in halogenations by acidified hypochlorous and hypobromous acid, may be taken to represent the distinct species H20C1+ and H,OBr+. A mechanism involving a reaction of the hypohalous acid with the aromatic compound, followed by addition of a proton, is possible, except in those chlorinations by HOCI in which the rate is independent of the concentration of aromatic substrate. The intervention of Cl+(up) in zeroth-order chlorination by acidified hypochlorous acid is supported by all kinetic evidence but may be questioned on the basis of its low equilibrium concentration. This problem remains to be resolved. Literature Cited
INGOLD GOLD, C. K., "Structure and Mechanism in Organic Chemistry," Cornell University Press, Ithaca, New York, 1953, Chap. VI. (2) DE LA M ~ R EP. , B. D., A N D RIDD,J. H., "Aromatic Substitution," Buttorwort,hs Scientific Publiestions, London, 1959, Chap. 9. (3) BERLINER, E., AND BECKETT,M. C., J . Am. Chem. SoC., 79, 1425 (1957). (4) FRANCIS, A. W., J . Am. C h a . Soc., 47, 2340 (1925). (5) ~ ~ E E R W EH., I N2. , Ang. Chem., 38, 815 (1925); PFEIFFER, P., AND WIZINGER, I t . , Ann., 461, 132 (1928). (6) OLAH,G. A., KIJHN, S. J., FLOOD, S. H., N D HARDIE, B. A,, J. Am. Chem. Soc., 86, 1039, 1044 (1964). (7) ~ ~ O B E R T S O NP. , W., DE LA I\~ARE, P. B. D., A N D JOHNSTON, W. T. G., J. Chem. Soe., 276 (1943); KEEFER,R. M., OTTENBERG, A., AND ANDREWS, L. J., J . Am. Chem. Soc., 7 8 , 255 (1956). (8) SHILOV,E. A,, A N D KANIAEV,N . P., C . R. Amd. Sn'. U.R.S.S., 24, 890 (1939). (9) INGOLD, C. K., SMITH, E. W., ANI) VASS,C. C. N., J. Chem. Soc., 1245 (1927). (10) SOPER, F. G., AND SMITH, G. F., J. Chem. Soe., 1582 (1926). (11) AnoTsn~,J., AND SYMONS, M. C. I?., Quart. Rev., 16, 282 (1962). (12) ALLEN,C. F. H., CRESSMAN, H. W. J., A N D JOHKSON, H. B., in "Organic Syntheses," John Wiley & Sons, Ine., New York, 1955, Coll. Vol. 3, p. 796. (13) MILLEN, D. J., quoted in ref. 1, p. 293. (14) WILSON, W. J., AND SOPER, F. G., J . Chem. Soe., 3376 (1949). (15) DERBYSHIRE, D. H., AND WATERS,W. A., J . Chem. Soe., 564 (1950). (16) BRANCH, S. J., AND JONES,B., J. Chem. Sac., 2317 (1954).
See also the similar observations by BRADFIELD, A. E., DAVIES,G. I., AND LONG,E., J . Chem. Soc., 1389 (1949). (17) DE LA MARE,P. B. D., HILTON,I. C., AND VERNON, C. A,, J. Chem. Soc.. 4039 (1960): DE LA MARE. P. B. D.. HILTON, I. C., AND V A ~ M A ,s.', J. Chem. Soe., 4044 (1960): See also STANLEY, G., AND SHORTER, J., J . Chem. Soc., 246, 256 (1958). (18)
DE LA
h{.ARE, P. B. D.,
AND
MAXWELL, J. L., J . Chem. Soe.,
4829 (1962).
B. D., AND HARVEY, J. T., J . Chem. Soe., DE I,.& MARE,P. B. D., AND H:
