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COORDINATION NUMBER CHANGES DURING OXIDATION-REDUCTION REACTIONS OF OXYANIONS. THE KINETICS OF THE ANILINE NITROSATION AND OF THE GLYCOL TELLURATE REACTION BYJOHN 0. EDWARDS, JOHN R. ABBOTT,HERBERT R. ELLISON AND JUNE NYBERG Metcalf Chemical Laboratories of Brown University, Providence 18, R. I . Received October
,$f
1968
The rate of oxyanion reactions with reducing agents parallel the rates of oxygen exchange with water and are dependent on the acidity, No rate correlation with oxidation potential is obvious, however. These data are consistent with the postulation of a deoxygenation reaction occurring before or during the rate step. Cases where variability in coordination number of the central atom are known will be discussed as will the role of protons. Similar behavior appears where oxyanions are oxidized and when they act as catalysts for peroxide oxidations. The kinetics of two reactions have been investigated. One reaction is the aniline nitrosation. Previously unrecognized features of the kinetics will be point,ed out; e:g., the rate increases with buffer concentration and the order in nitrite varies between one and two. A mechanism involving the nitrosonium ion is postulated. The second is the formation of gl col complexes by tellurate ion (H6TeOa-). The data can be interpreted as the rate-determining combination of a glycof molecule and dehydrated forms of tellurate anions. Complexing constants plus forward and reverse rate constants for several glycols will be presented.
Introduction One of the more significant problems which faces chemists working on reaction mechanisms is whether an equilibrium combination of two reactants to produce a complex forms an intermediate on the path to the transition state or is just a storehouse for reactants. This problem has been discussed before and no general solution is available although it has been solved in a few specific cases. It is our intention to first show how this problem is important in the consideration of oxyanion reaction mechanisms, particularly those involving hydrogen peroxide, and then to discuss the work on two kinetics problems that have been studied at Brown. Consider an oxyanion such as nitrite ion; the experimental data at hand on the structure show it is V-shaped. Nitrous acid is usually taken therefore as HONO, also V-shaped. I n strongly acidic media, it probably exists as the nitrosonium ion NO+. Yet there are other possibilities which must be considered. Nitrous acid could have a nitrogen-hydrogen bond or it could be N(OH)a, although all of us rather automatically expect it to have the formula "02 with an oxygen-hydrogen bond. Irrcspective of what the main structure is, small amounts of other structures are possible, and may be the important ones in reaction mechanisms. Thus the problem which one faces with oxyanions is this: what is the coordination number of the central atom in the transition state and to what species is the central atom bound? It is the proposition here that the coordination number of the central atom often changes and that the substrate species often joins in a chemical bond with the central atom during or prior to the transition state of the oxidation-reduction reaction. The rates a t which oxyanions oxidize particles like the halide ions or thiosulfate ion cover a vast range. Perchlorate, nitrate and selenate are very slow oxidants, while iodate, nitrite and selenite are rapid only in acid solution. Periodate and hypochlorite are rapid a t higher pH. What data are available indicate that the rates of oxidation parallel the rates of oxygen exchange with solvent wat>er.l Although there are not enough data available to see (1) J. 0. Edwards, J. Chem. Educ., 31, 270 (1054).
whether a linear free energy relation between rates of oxidation and rates of oxygen exchange can be made, there is certainly considerable correlation. On the other hand, there does not seem to be any obvious relation between rates of oxidation and the oxidation potentials of the oxyanions, thus something more than the oxidizing power is involved. Coordination Number Variability.-The central atoms of oxyanions have been found to show considerable variability in coordination number (c.n. hereafter) in their binding t o oxygen. Carbon can be found in c.n. two as in carbon dioxide or c.n. three as in the normal carbonates; in the orthocarbonates, the c.n. is four. Tetravalent sulfur is found both as sulfur dioxide with c.n. of two and as the sulfites with c.n. of three. Periodate ion can exist either as a tetrahedron or an octahedron, as can the molybdates, tungstates and probably the perrhenates and perosmates. Borates can be found in trigonal plane or tetrahedral configurations. Cationic intermediates such as NO2+, NO+ and Clf have been recognized in the chemistry of nitrates, nitrites and hypochlorites, respectively. There is therefore reason to believe that oxyanions in general can undergo hydration-dehydration reactions to give species with different numbers of oxygen in the coordination sphere of the central atom. There is also reason to believe that many reducing particles can replace oxygen in the coordination sphere since these particles are electronrich and therefore are bases in the Lewis sense. I n fact, there is a quantitative correlation between reducing power and strength of Lewis base.3 Dependence of pH.-There is another parallel between rates of oxidation and rates of water exchange which bears mentioning, and that is the strong dependence of rate on the hydrogen ion concentration. In many cases, the order in hydrogen ion concentration is two which strongly suggests a (2) T h e use of the term "cationic intermediate" for species such as NOa+ and C1+ is reasonable; however it emphasizes a factor, the charge, which is probably far less important than the availability of a binding site on the central atom. Indeed, one finds species such as HCrOa+ being postulated merely t o be consistent with the terminology, whereas CrOx has the necessary open orbital without the need of a proton t o supply a positive charge. (3) J. 0. Edwards, J . Am. Chem. Soc., '76, 1540 (1954).
J. 0. EDWARDS, J. R. ABBOTT,H. R. ELLISON AND J. NYBERG
360
preliminary reaction of the type X0,-
+ 2H+
XO,,+
+ H20
In the case of oxygen isotope exchange, the reaction can be accomplished merely by such an equilibrium. When an agent like iodide ion is being oxidized, the reaction XOn-,+
+ I-
X0,J
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evidence that this peroxide does exist at low temperatureP8although it is not stable at room temperature. This reaction could proceed through the steps
+ + +
+ + +
"03 H + +NOe+ H20 NO*+ Hi02 +HOONOz H + HOON02 +0 2 H+ NO2NOzHZ02 +NO1Hz0
+
+
As hydrogen peroxide is both a reducing agent and an oxidizing agent, cyclic decompositions of peroxide are also possible with other oxyanions. The kinetics of the oxidation of peroxide by iodate ion and by the hypohalites are also consistent with unstable peroxyacids as intermediate^.^ The forR = k[ClOs-] [Br-] [a+]' mation and breakdown of unstable peroxychroand the postulated intermediate is mates is well known even to students in qualitative 0 analysis courses. Br-Cl< The oxidations of iodide ion and thiosulfate ion 0 by hydrogen peroxide are not fast, however they can be made so merely by the addition of a small which can then decompose to give Br+ and ClOa-. Strong evidence for the existence of intermediate amount of molybdate. In view of the fact that complexes may be found in the values of the rate substituted peroxides such as peroxyacetic acid and constants for such reactions; in some cases the peroxymonosulfuric acid oxidize rapidly, it seems rates are so large that equilibrium constants for the probable that the molybdate ion acts through the intermediates must be equal to or greater than one6 rapid reversible formation of a peroxymolybdate in order that the data be consistent with the colli- HMoOb-. It is known that replacements on sion theory. While these data do not in any sense molybdate are rapid. The postulated mechanism prove what type of intermediate complex is in- is volved, the type postulated has the advantages of HMoOi- + Hz02 HOOMoOa- + HzO being consistent with known chemical species and HOOMoOa- + I- +HOI + Mood' of providing a smooth path for transfer of two The fact that monosubstituted peroxides are electrons from reductant to oxidant. The oxidation of sulfite by bromate presents a rapid oxidizing agents causes another behavior mechanistic case where an oxyanion being oxidized pattern which is strongly indicative of intermediate appears to undergo a change in c.n. before elec- formation. Nitrite, lo selenite" and phenylborontron transfer.B The results of both kinetic and iso- atela ions all can be oxidized by peroxide, and in tope tracer studies support a mechanism of the each case one of the observed rate laws has an order of two in oxyanion concentration. Taking type the case of selenite as example, the mechanism HzSOs )r SO2 + HzO HSeOs- + HOOH J_ HOOSeOa- + HIO H 0 SO2 + HBrOs J_ o>-O--Br(o +Products HOOSeO2- + SeOa' +HOSeOs- + SeOa0 where selenite ion acts as a nucleophile in a diswith direct oxygen atom transfer from the bromate placement on the peroxyselenite seems reasonable. to sulfur dioxide. The existence of an unstable peroxynitrite plus the For those who wish to see a more detailed isotope tracer work of Taube and Anbarlo in the discussion of some of the problems and complexi- oxidation of nitrite by peroxide support such a ties in oxyanion reaction mechanisms, the article by mechanism. Further support for the postulation Taube' should be consulted. of monosubstituted peroxides as intermediates is Interactions with Peroxide.-Some of the clear- found in the fact that oxyanions such as molybest examples of intermediate formation prior to the date, nitrite, selenite and phenylboronate ions are rate-determining step occur when hydrogen per- substitution labile, and act as catalysts, while oxide is present. Because of the interesting oxyanions such as sulfate, perchlorate, phosphate chemical properties of peroxide, it can react in and selenate which are substitution inert have not many ways; four distinctly different patterns of been observed to act as catalyst for peroxide reacbehavior have been noted which are pertinent t o tions. the present discussion. The fourth type of behavior also is strongly When the oxyanion has a high oxidation poten- indicative of the formation of an intermediate tial, oxidation of hydrogen peroxide to oxygen can peroxyanion. When sulfite ion in acid solution is be observed. Nitric acid is known to decompose oxidized by peroxide, it has been found that two of peroxide and one possible mechanism involves the the four oxygens present in the product sulfate came formation of peroxynitric acid "Oh. There is (8) R. Sohwarz, Z . Cham., 266, 3 (1948). can occur. The species XO,-l may be formed in the transition state or it may be an intermediate on the road to the transition state. In the chlorate oxidation of bromide the observed rate law4is
Q ~ O T ~ .
(4) J. Hirade, J . Chem. SOC.J u ~ Q10, ~ ,97 (1935). (5) J. 0. Edwards, Chem. Rms., 60, 455 (1952). (6) (a) J. Halperin and H. Taube, J . A m . Chem. SOC.,74, 380 (1952); (b) F. 9. Williamaon and E. L. King, ibid., 79, L397 (1957). (7) H. Taube, Reo. Chsm. Proor., 17, 25 (1958).
(9) M. C. R. Symons. J . Chum. Soc., 5596 (1955). (IO) M.Anbar and H. Taube, J . Am. Chem. SOC.,76, 6243 (1954). (11) F. J. Hughea and D. 8. Martin, THISJOURNAL, 69,410 (1955). (12) H. G. Kuivila and A. G. Armour, J . Am. Chem. Soc., 79,5659 (1957).
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from the hydrogen peroxide.13 The mechanism presumably is the rapid formation of a peroxysulfite ion followed by an intramolecular 1,2-shift of an OH+ from oxygen to sulfur. Under certain conditions a similar behavior has been noted for the reaction of nitrite and peroxide.1° Diazotization of Aniline.-Nitrosation, the process by which an NO+ group becomes attached to a molecule with a spare electron pair, is a reaction of considerable importance. The kinetics of this type of reaction have been studied often; however there is little agreement from one laboratory to the next on the mechanism. Some of the difficulties which have arisen are discussed in the papers by Anbar and Taube,lo Hughes, Ingold and Ridd,14 Li and REPlCTlON TIME IN MINUTES. Ritter,16 Dusenberry and Powell,lBSeel" and BunFig. 1.-Fraction of aniline reacted as a function of time; ton, Llewellyn and Stedman.18 Suffice it to say [C6HbNH2]= 5 X lo-' M , [NOz-] = 4 X IO-* M and here that in some cases not even the order in re- concentration of phthalate buffer = 5 X IO-* M. actants is agreed upon by two different investigators. In hopes of clarifying the situation somewhat, we decided in the fall of 1955 to reinvestigate the nitrosation of aniline. This reaction was chosen for several reasons: (a) the simplicity in analytical technique needed, (b) the presence of untested assumptions in the previous work, and (c) the lack of correlation of the kinetics of this reaction with isotope tracer studies, Some weeks after submission of the abstract for this meeting, the series of papers by Hughes, Ingold and RiddlS came across the Atlantic. Where their results and ours are in agreement, only a brief resume will be given here. There are some points of disagreement yet on which further work seems necessary. Our study has been concerned with the mechanism of the aniline diazotization in buffered solutions between pH five and six. The rates, all at Oo, have been followed by coupling of the product diazonium ion with alkaline phenol and measurement of the color developed. The precision of the results obtained by this technique may be seen in Fig. 1, which also shows that the rate of reaction 0.05 0.10 0.15 0.20 is not dependent on the aniline concentration. [CH&OO-]. This observation, plus the fact that different aniFig. 2.-Corrected rate as a function of acetate concentralines diazotize at the same rate, had been known. tion. Different symbols indicate runs on different days. Our results of rate versus p H indicate that the Conditions comparable to those of Fig. 1. reaction is second order in hydrogen ion concentration,20as had been observed previously. The order in nitrite was found to be between one and two; also the rate was found to depend on the nature of the buffer, phthalate having a greater effect than acetate. From these results, it is apparent that only a thorough study of the rate over a range of nitrite concentrations a t dif(13) J. Halperin and H. Taube, J. Am. Chem. Soc.. 74, 380 (1952). (14) E. D. Hughes, C. I
