Dec., 1939
HALOGEN
HYDRATION AND MECHANISM OF %ME HALOGEN REACTIONS
The question of whether or not these peaks may be assigned to N-€3-K bonding between the secondary pyrrole nitrogens and the tertiary nitrogens in these compounds arises from a consideration of the structures involved, It must be stated at once that infrared spectroscopy can only give evidence for intramolecular bonding when a set of absorption frequencies can be compared with those exhibited by like acceptors and donors in different molecules. On the other hand, intermolecular bonding between molecules of like or unlike species may be demonstrated conclusively by observing the effect of concentration changes in solution upon "bonded' or "unbonded" peaks. This criterion of course becomes inapplicable in testing for intramolecular bonding. Since the absorption peaks in question differ markedly from the pyrrole N-H peak a t 2.85 p,12 one is inclined to suggest that N-H-N bonding exists in these molecules. However, this is unproved as yet, since the character of the substituted pyrrole rings may have been changed in the iormation of these compounds. This seems pos(12) A. M. Buswell, J R. Downing and W. H. Rodebush TRIS 61, 3252 (1939)
JOURNAL,
3513
sible in light of the recent work on unbonded N-H freq~encies.'~Hence we can only state a t present that the absorption peaks occur at frequencies which might be expectedI2 if N-H-N bonding were existent in the porphin and dipyrrylmethene nuclei. The authors wish to express their gratitude to Professors -4. M. Buswell and W. H. Rodebush for use of the infrared apparatus and for timely suggestions.
summary 1. The infrared absorption of two dipyrrylmethenes, the copper complex of one of them, etioporphyrin I, and etiohemin I chloride has been measured. Replacement of acidic hydrogens with metal atoms eliminated nitrogen-hydrogen absorption in the dipyrrylmethenes and in etioporphyrin I. 2. The possibility of M-H-N bonding in these compounds is discussed. (la) A. M. Buswell, 0.W.McMillan, Wall, ibid., 61, 2809 (1939). URBANA,ILLINOIS
W.H. Rodebush and F. T.
RECENEDSEPTEMBER 30, 1939
[CONTRIBUTION FROM THE RESEARCH LABORATORY, GENERAL ELECTRIC COMPANY 1
The Hydroiysis of Bromine. The Hydration of the Halogens. The Mechanism of Certain Halogen Reactions BY HERMAN A. LIEBHAFSKY Corrections.--Jones and Baeckstrom' and Liebhafskyz have independently established 5.8 (lo-$) as the value of K,,the equilibrium constant of the bromine hydrolysis, a t 25'. The temperature coeflicients of this reaction (c$ Ref. 2, Table 11, last column) are probably more accurately known than is .Kr a t any temperature except 25'. For these reasons, it: is especially desirable to correct a computational error that Professors R. H. Gillette and H. A. g of the Division of Chemistry, University alifornia a t Davis, have kindly c d e d to my attention. Through some a t 35', the result for Series mischance, 11.9(10 5, instead of 11,3(10-9), the result for Series 7 (cf. Ref. 2, Table 11),was used in deriving an incorrect equation for & as a function of temperature. A table and a conclusion in Ref. 2 consequently should be changed to read as follows. (1) Joncs and Baeck.strom, Tars JOURNAL, 56, 1517 (1934). (2) Liebldsky, ibid., M, 1600 (1934).
TABLE IV (Ref.2)
x
SUMMARY OF BBST VALUES FOR KI lo-' AT TEMPERATURES 00 100 25' 30'
DIFFERENT 350
Liebhafsky 0.69 1.76 5 8 8.2 11 3 Computed .88 1.95 (5.8)' 8.14 (11 3)" Selectedb .70 1.78 5 8 8.3 11 3 Value assumed in deriving the equation log Ki = 0.6876 2660.3/T from which computed values were obtained. Only these finally selected values for K1 are expressed in activities. (Similar changes are to be made in Ref. 3, p. 95, Table I.)
*
Summary (Ref. 2) 4. The temperature variation of the KI for bromine is anomalous in the neighborhood of 10' ; but the departure of d(1og Kl)/d(l/T) from constancy is less pronounced thm that observed by Jakowkin in the case of chlorine. Also, the lime "a. Measurements a t 0" by Jones and Hartmann" has been omitted above the first row of data in Table 111.
3514
HERMAN A. LIEBHAFSKY
Hydration of the Halogens.-The temperature coefficients of halogen equilibria involving aqueType of measurement
Equilibrium Equilibrium Equilibrium Kinetic
(distribution between air and H20) (hydrolysis of the halogen) (distribution between CClr and H,O) (see text below)
ous solutions are generally a n o m a l o ~ s . ~When the corresponding equilibrium constants are calculated in the usual manner-that is, by using (X,), the concentration of free halogen in the aqueous solution-the log Kx2 vs. 1/T plot approaches a limiting straight line at higher temperatures but bends more and more sharply to the 1/T axis as 0' is approached. Figure 1 illustrates this behavior. While D', the distribution ratio of iodine between air and carbon tetrachloride, shows the orthodox temperature dependence, D,the distribution ratio between air and water, behaves anomalo~sly.~This anomaly has found a reasonably satisfactory quantitative explanation in the assumption that the degree of hydration of the halogen is changing in the temperature region where the anomaly occurs; as 0' is approached, the equilibrium U
+ (It - Hz0 H (HI + (U)= (Xd U)
(1)
shifts so as to increase (H), the concentration of the more hydrated form. (The formula of U is 12.uH20; that of H, I2.hH20.) The correction of the above error tends to strengthen this explanation and has a bearing on the important problem of correlating kinetic and equilibrium data (cf. Ref. 3, p. 111). The data in Table 2d, Ref. 3, for the bromine hydrolysis should be : For bromine from reaction 12 (cf. Table I) 35 0 2 5 . 0 10.0 lo0 X KBr,(measd.) 11.3 5.8 1.76 d (d = 0 assumed) 0.045 .lo9 (H)/(U) d = log KBr, (calcd.) - log ICBr, (rneasd.) log K,,, (calcd.) = 0.6876 - 2660.3/T ABII, = -14300 cal. d.
T , "C.
0.0 .69 .lo6 .276
The difference between d for 0' from kinetic (0.21) and equilibrium data (new value, 0.11) is now cut virtually in half (cf. Ref. 3, p. 95, last paragraph). (3) Liebhafsky, Chcm. Rm.,17, 89 (1936). (4) The data for aqueous solutions have been reproduced from Ref. 3, where the original sources are cited. D' has been obtained for each temperature by dividing the vapor pressure of iodine in mm. of mercury [Baxtcr nod Or-, Tmei JOVUHAL, W , 1061 (1916)l by the lolubilltr io mole fractionr of iodloo for that solvmt [Hildah n d mad Jcnta, ibid , 48, 4188 (lPPO]],
Vol. 61
The calculated values3of the heats of hydration (Equilibrium 1) now are: Chlorine
Bromine
Iodine
- 9,700 cal. - 9,900 cal. - 10,700 cal.
-13,100 cal. - 14,300cal.
-11,800cal
-25,000 cal
-22,800 cal
The new value for bromine is in good agreement with the other value derived from equilibriuul data. 200 100
50
b" 8 s"
20
X
5
10 5
2 3.0
3.2
3.4 1 / T X 10'.
3.6
3.8
Fig. 1.-Distribution ratios for iodine between air and water (D, anomalous temperature coefficient), and normal tembetween air and carbon tetrachloride (D', perature coefficient). The plot is semi-logarithmic.
The specific rates at which hydrogen peroxide is oxidized by each of the three halogens, X2, are the kinetic data from which heats of hydration were deduced. For bromine and iodine, these large (and, perhaps, unreasonably large) heats are obtained because the downward curvatures in the 1/T plots are more pronounced for the kinetic than for the equilibrium measurements. These unexplained minor inconsistencies should not be permitted to obscure the more important fact that for bromine and iodine both types of measurements reveal the existence of similar temperature coefficient anomalies. This fact indicates that the assumed mechanisms, which are of the type Rapid equilibria U H20 HXO H+ X- 4-uHzO (2a) H HXO H+ X(it - 1)HtO (2b) Rate-determining step HsOs HXO ----t HIO 0: H+ 4- X- (3)
+
+
+
+ + + + + +
are in the main correct, for mechanisms of t h i a
Dec., 1939
HALOGEN HYDRATION AND MECHANISM OF
type will reflect a change in the ratio (H)/(U) because such a change necessarily causes a change in the ratio (HXO)/'(X2). The reaction between chlorine and hydrogen peroxide,j however, does not appear to proceed mainly by such a mechanism; for the kinetic data give a linear log k vs. l / T plot from 0 to 23" (The temperature coefficient of the chlorine hydrolysis equilibrium is nearly twice as great at 0' as a t the higher temperature.") This conclusion is strengthened greatly by the experimental fact that. hypochlorous acid when mixed directly with hydrogen peroxide reacted a t approximately one-millionth the rate to be expected from kinetic measurements on the chlorine-hydrogen peroxide r e a ~ t i o n . ~The difficulty has been thoroughly discussed by Makower, who presented this conclusion as a third alternative explanation; the evidence from other reactionsa seems convincing enough to warrant rejection of the other two. To summarize: The rate of the chlorine-hydrogen peroxide reaction, unlike the rates of the corresponding bromine and iodine reactions, is unaffected by the shift from G to H which occurs when the temperature of an aqueous Xz solution is lowered. The mechanism of this reaction must therefore differ from the mechanisms of the other two. The Reaction between Arsenious Acid and Iodine.-Roebuck's classical investigationsa of the reaction
+
H ~ A s O ~1 3 -
+ H20 = H~AsO,+ 2H+ 4-31-
(4)
SOMEHALOGEN REACTIONS
3515
senious acid in the absence of iodine, this verification was sought in a study of the temperature coefficient of the forward reaction. The following measurements were made in 1032 in the Chemical Laboratory of the University of California a t Berkeley. Since the rate law of the forward reaction has been established, it sufficed to perform the experiments under conditions such that no concentration except that of total iodine changed greatly during the course of an experiment. Sodium iodide, sodium meta-arsenite and perchloric acid solutions were mixed, diluted and placed in a thermostat; the reaction was begun by adding sodium triiodide solution that was also a t thermostat temperature. The rate was measured by rapidly titrating samples of the reaction mixture withdrawn from time to time, with 0.02 N sodium thiosulfate. The detailed data for Experiment 5 are given in Fig. 2, where the thiosulfate titer is plotted on semilogarithmic paper against the time. The initial
I
loted
50 100 150 Time in minutes. 2-Detailed results for expt. 5 on the rate of reaction 4. The plot is semi-logarithmic c)
was the first complete experimental demonstra- Fig tion of the relation obtaining between chemical kinetics and chemical equilibrium. The mechaconcentrations for all experiments were nearly nism for the forward reaction identical, except that the concentration of arRapid equilibrium: 1,- Jr I2 + I(5) senious acid was cut in half for the experiments at Rapid equilibrium: It + H20 H I 0 + H+ + Ithe two highest temperatures. Under the con(6) centration conditions chosen, the reaction is apKate-determining step: H I 0 + HsAsOa + proximately first order; the deviations from HsAs04 H+ I- ( 7 ) straight lines in the log 2'12) DS. t plots were satisdoes not appear to have been questioned seriously factorily removed by correcting for the changes although the work of Makower and Liebhafskyg in (H+), (I-) and (H&O,) that occurred as the left little doubt that further verification was dereaction proceeded. (The change in (H +) caused sirable. Since hypoiodous acid is so unstable by mixing meta-arsenite and perchloric acid was that it can scarcely be treated directly with arof course allowed for.) In other words, the rate (5) Makower, THISJOURNAL, 66, 1315 (1934) law (6) Jakowkin, 2 physik. Chem., 29, 613 (1899).
+
+
(7) Bray, Chem R e v , 10, 175 (1932). ( 8 ) (a) Roebuck, J Phys. Chem., 6, 365 (1902), and (b) 9, 727 (1905); (c) Liebhdsky, d i d , 88, 1048 (1931); references t o other work will be found in Ref. $c. (9) Mskowcr pad Liobhafsky, Xwnr. Fnrodoy Soc I PB, 687 (19W.
was obeyed satisfactorily. The temperatures and results are summarized in Table I.
3516
Vol. 61
HERMAN A. LIEBHAFSKY
Comparison of the last two columns in Table I pendent of the ratio (H)/(U), Perhaps the shows that here, as in the chlorine-hydrogen per- simplest mechanism that can be devised to meet oxide reaction] the logarithm of the specific rate this requirement is is a linear function of 1/T to within the experiU HsAsOs A H+ I(9) mental error. Roebuck's& value for k; a t 0' in Rapid equilibria B H + I- (10) H HsAsOa sulfuric acid solution is 9.4(10-4) (units: moles, Rate-determining liters, minutes); this result, when corrected for HsASOr H+ Isteps of iden(Ila) tical specsc the decreased acid concentration due to the presrate B +HsAsOa HC I(llb) ence of HS04-, is in good agreement with the Table I value.1° Roebuck gives several values The device of having A and B decompose a t the for the temperature coefficient of ki between 0 same (or nearly the same) specific rate is saved and 10'; his highest value, 3.5," corresponds to a from being an obvious artifice by the fact that heat of activation Q = $19,200 cal.; his lowest, these two intermediate products differ only in the 3.17,8bto +17,710 cal., which is in fair agreement degree of hydration. An analogous mechanism would eliminate the conflict between equilibrium with 16,970 cal., the value from Table I. and kinetic data for the chlorine-hydrogen perTABLEI oxide reaction. TEMPERATURE COEFFICIENT DATAFOR TXE ARSENIOUS A detailed discussion of the arguments for and ACID-IODINE REACTION this mechanism is scarcely warranted. against IO& X k: 104 X k: Experiment T,O K . (meas.) (CalCdJc' Accurate measurements of the equilibrium con6.6 3 273.1 6.7 stant for the iodine hydrolysis over the tempera4 273.1 6.5 6.6 ture range in question would be more valuable 11.2 5 277.7 10.9 than such a discussion ; but these data, especially 21.6 21.1 6 283.5 for the lower temperatures] will be difficult to ob34 7 288.1 36 tain. The existence of accurate hydrolysis equi8 293.1 58 57 librium measurements for the analogous chlorinea Calculated values from the equation log k: = 10.411 hydrogen peroxide reaction to some extent com(3710.0/T). pensates this difficulty. The magnitude of Q and its constancy (or near It may be pointed out, finally, that the equiconstancy) both suggest the necessity of revising librium constant of a halogen reaction could, if the mechanism of the arsenious acid-iodine reac- heat capacity terms are ignored, be a linear function. The heat of dissociation of triiodide ion tion of 1/T. This linearity will result if the (f4320 cal.)sc and the heat of hydrolysis of iodine forward (reverse) reaction has a mechanism like (+22,700 cd.)" must both be included in Q if the that outlined for the arsenious acid, while the reaccepted mechanism is correct; on this basis the verse (forward) reaction proceeds by a mechaheat of activation for the rate-determining step nism whose temperature coefficient is normal. between hypoiodous and arsenious acids would The triiodide equilibrium (Reaction 5 ) appears have a negative heat of activation of 10,000 cal. to be one case of this kind, and Equilibrium 4 anWhen Roebuck's work was reexamined some years other. Further experimental work to test these ago," it was not known that a constant Q was in ideas is highly desirable. conflict with most of the data for aqueous halogen Light Absorption by Iodine Solutions.-The solutionsS; as it now appears, the previous con- General Electric recording spectrophotometer12 clusionll that the rate-determining step has a was used to establish whether a characteristic negative temperature coefficient must be rejected] change in the absorption of visible light accomand the mechanism for the arsenious acid-iodine panied the change in the hydration of iodine bereaction revised. tween 0 and 75'. For purposes of comparison, The revised mechanism must meet the require- carbon tetrachloride solutions were also investiment that the rate a t all temperatures be inde- gated a t different temperatures. The ease with (10) (a) The dissociation constant at 2 5 O of HSOa- in concentrawhich iodine volatilizes from aqueous solutions is tion units for the above reaction mixtures is near 0.03: Noyes and well known; this diiliculty was minimized by Sherrill, THISJOURNAL, 48, 186 (1926). (b) The corresponding heat of dissociation is small; Hamer, ibid., 66, 860 (1934). fling the cell (5.15 cm. long) with the warm or
t
+ +
__
+ + + + + + + +
+
R
(11) (a) Ref. 8c; (b) cf. also, Moelwyn-Hughes, "Kinetics of ~ O inMSolution,'' Clarendon Press, Oxford, 1933, p. 207.
(12) Michaelson and Liebhafsky, Gex. &c.
&.,
89, 445 (1936).
HALOGEN HYDRATION AND
Dec., 1939
MECHANISM OF SOME
HALOGEN REACTIONS
3517
ment toward longer wave lengths of Curve IIa, Fig. 3, indicates that the complete dehydration of iodine is beginning a t 7G0. - _._. .
100,
I
0L-L 4000
L-J_-,-I-L
I.._l-il-_.
-
.~
J
5000 6000 7000 Wave length in hgstroms. Fig. 3.-Effect of temperaturechange on the absorption of visible light by iodine solutions: Curve
Solvent
Moles h/l.
T , 'C.
I Ia I1 IIa I11 IIIa IV IVa
CCla
None None 6X 6 X 10-6 9 x 10-6 9 x 10-6 9 . 5 x 10-6 9 . 5 x 10-6
25 50 24 76 0.3 24 25 51
CC14
HzO Ha0 HzO Hz0 CClr
CCle
~
cold solution, tracing its transmission curve (less than three minutes is required), and permitting the covered cell to stand until its contents reached room temperature, when the transmission was again measured, Some of the curves obtained are shown in Fig. 3; that the effect of changing temperature on the water-filled cell is very small is obvious from the uppermost curves in Fig. 4. Figure 1 shows clearly that aqueous iodine solutions show anomalous temperature coefficients while carbon tetrachloride solutions do not; Fig. 3 shows that changing temperature affects light absorption by both solutions almost identically when the change in the absorption by carbon tetrachloride itself is allowed for. It follows, therefore, that the change in hydration presumably responsiue for the curvature in Fig. 1cannot be detected on the spectrophotometer, and that IJ is itself hydrated. The noticeable displace-
0 4000
I
L..L- J - 1
I .- 1
5000 6000 7000 Wave length in h g s t r o m s . Fig. 4.-Absorption of visible light by gaseous and dissolved iodine: Curve
Solvent
Moles Ia/l.
I Ia
HzO Hz0 Air Air HnO Air CCla
None None None 1 . 6 x 10-6 5.36 X Unknown 2.39 x 10-4
I1 I11
IT'
v
VI
T, OC.
25 50 25 25 25 25 2
