March, 19131
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XOTES
DzO ISOTOPE EFFECTS I N THE CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY METAL IONS BY J. F. HARROD AND J. HALPERS Department of Chemistry, The University of British Columbia, Vancouic r Canada Received December 2 f j 1960
With a view to gaining further insight into the role of t'he solvent in the catalytic activation of molecular hydrogen in solution,' t'he rates of reaction of hydrogen with a number of metal ions and complexes, in HzO and DzO, have been compared. Our results are summarized in Table I. TABLE
Metal ion
Fig. 1.
Rate law
1 Ternl,., OC.
kH20/ kD20
(10.1)
Ref.a
k[&] [Cu2+] 110 1.20 k[Hzl [Ag+12 50 1 23 Ag +* 75 1.26 k[Hzl &'+I Hg2 Hgz2+ k[Hz1[Hga2+l iJ 1.33 e Cu(0Ac)z k[H2][CU(OAC)~] 100 0.93 PdCLk[Hz] [PdCl,'-] 80 0.90 0 RhC1c3k[Hz] [RhC1e3-] 80 1.00 hLln04k[Hz][Mn04-] 50 0.99 Ag+ MnOd- k[H~][Bg+][hInO4-] 40 0.93 a Earlier measurements in HaO are described in the quoted reference. The same procedures were used in the present measurements and the results in HaO agree well with the earlier ones. While at higher temperatures, the reaction of Hz with Ag+ follows a rate law and mechanism siniilar to those for Cuz-, the predominant contribution under the conditions of this comparison is from the "termolecular" path which is believed to involve homolytic splitting of Hz according to equation 3 (ref. d, this Table). E. Peters and J. Halpern, J . Phys. Chem., 5 9 , 793 (1955) ; J. Halpern, E. R. Macgregor and E. Peters, ibid., 60, 1455 (1956). d A. H. Rebster and J. Halpern, ibid., 60, 280 (1956); 61, 1239, 1245 (1957). e G. J. Korinek and J. Halpern, ibid., 60, 285 (1956). E. Peters and J. Halpern, Can. J . Chenz., 33, 356 (1955). 0 J. Halpern, J. F. Harrod and P. E. Potter, ibid., 37, 1446 (1959). J. F. Harrod and J. Halpern, ibid., 37, 1933 (1959). A. H. Webster and J. Halpern, Trans. Faradall Soc., 33, 51 (1957). cu2+
+
of the proton pairs from one another. Thus it is a t first sight surprising that Finegold mould not have been forced into our interpretation by t'he complexity of his spectrum. Unfortunately, in diethyl sulfite, the chemical shift' seems accidentally to be so small that the additional lines are vanishingly weak, so that' the two possible kinds of nonequivalence are not readily distinguishable from one another. From his figure it would appear that the relevant chemical shift is in the vicinit'y of 0.05 p.p.m. If we assume the spiii-spin coupling between the non-equivalent protons to be 10 c.p.s., as it is in our compound, it) is easy to predict12 that the satellite lines should have intensities of only 1-2%; of t'he central components.1 3 Similar phenomena are observed in many other molecules.s,4 In diethyl acetal, for example, we have observed a complex met'hylene multiplet n-hich has been analyzed to give a chemical shift difference of 0.152 0.005 p.p.m., an internal spin coupling of 9.2 0.3 c.p.s., and couplings of about 6.7 and 7.2 C.P.S. between the non-equivalent methylene protons and the methyl group.'? It is significant that' the spectrum of ethylal, which differs from acet'al only in the substitution of a hydrogen atom for the central methyl group, contains a perfectly normal methylene quartet. l 4 This circumst'ance is completely in accord with the arguments preseiitcd here, since the ethylal molecule possesses too much symmetry to allow nonequivalence of thr protoiis in the same methylene group. We wish to thank ;he Sntioiial Science I-oundatioii for support of this work ( 1 2 ) J.
S.n ' a w l i . "1'rocerdin:rs of the IV International Meeting on
1H;Y. Perganion Press, London, in press. (13) I h . Finegold fpiivate communication) has agreed t o the probable correctness of the interpretation proposed here. He has stated that i n many other phosphorus compounds he had studied ( e . q . . (EtO)zP(O)Me. (F,t,O)zPMe. (ETO)zP(S)Cl, (EtO)zP(O)H a n d ( E t 0 ) z P i O ) S J I e all methylene protons appeared t o be equivalent a n d t h a t these obseryations caused him t o choose his previously published explanation for his observations on (ETO)zP(S)Me rather t h a n t h e one proposed here. (IS) C . S. Johnson. Jr.. J . P. Fackler, Jr.. J. S. Waugh a n d F. A . Cotton, unpublished work.
C F
'
+
J
It was hoped, in particular, that these measurements would provide a criterion for distinguishing those mechanisms2 in which Hz is split het'erolytically, transferring a proton to a water molecule, e.g. CU'+ -t HP
+ HzO +CuH+ + HaO"
(1)
from those in which a ligand other than water is believed to serve as the proton acceptor, e.g. Cil(OAc)z
+ Ha +CuH+ + HOAC+ OAC-
(2)
or in which H2is split homolytically without proton traii.qfer 2,4g+
+ Ha +2AgH+
3)
The results in Table I fail, however, to provide any clear-cut indication of such mechanist,ic differences. All the aquo ions included in the comparison (Cu2+, Ag+, Hg2+ and Hg2?+) shorn modest reductions in rate, ranging from 20 to 30%, on passing from HzO to D20, while the rates for complexes containing ligands other than water in (1) J. Halpern, J. P h y s . Chem., 68, 398 (1859): Aduances in Catalys i s , 9, 302 (1957); 11, 301 (1959). (2) Refs. r d Table I.
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their inner coordination shells are nearly equal in the two solvents. This pattern, including the similarity of the isotope effects for Cu2+and Ag+, suggests that differences in the coordinating properties of HzO and D2O (for which other evidence has been ciited3), rather than specific participation of water in the reaction, are largely responsible for the isotope effect in the case of the aquo ions. This is of interest in view of the many other studies* of H20-D20 isotope effects aimed a t elucidating the role of solvent in the mechanisms of various reactions of metal ions. Grateful acknowledgement is made to the Kational R,esearch Council of Canada and to the donors of the Petroleum Research Fund administered by the American Chemical Society for support of this work. (3) J. Halpern and A. C. Harkness, J . Chem. Phys., 31, 1147
The solution (0.5 ml.) was electrolyzed in millicoulometric cell. The reference electrode was a layer of mercury on the bottom of the cell connetted by a plastic tuhc to the mercury resewor. The anodic mercury w m moved slowly up and down at. one minute intervals by movement of the mercury reservoir. Owing to this agitation of the mercury layer, to the large diameter of the cathode capillary, and to the narrowness of the electrolytic cell, the solution was eontinuously but gently stirred during the electrolysis. Before electrolysis the solutions were freed from dissolved oxygen by bubbling hydrogen through them, and during the electrolysis hydrogen gas wa8 passed very slowly over the surface of the electrolyte. The hydrogen was partially saturated with ammonia by treatment with ammonium chlorideammonia buffer of the same concentration as the electrolyte solution. Polarographic curves were recorded before and after electrolysis
.
Results The effect of electrolysis on the diffusion current of Co(I1) ions is shown in Fig. 1. The solutions
(1959). ( 4 ) J. Hudis and R. W. Dodson. J . A m , Chem. Sac., 7 8 , 911 (195G); F. B Baker and T. 1%‘. Newton, J . Phys. Chem., 61, 381 (1957); A . Znirkel and IT. Taube. J . A m . Chem. Soc., 81, 1288 (1959).
POLAROGRAPHIC AND COULOMETRIC 1KVESTIC;ATIOKS OK T H E REDUCTION RATE OF COBALT(I1) IN T H E PRESENCE OF CYSTINE BY EMILIAN B. WERONSKI Department of C‘hemtstry. Unzverstty of Warszawa, Warszawa, Poland Recezved June 88,1060
Enhancement of the diffusion current has been found in “Brdicka’s electrolyte” and other cobalt solutions in the presence of some proteins and amino acids containing sulfhydryl or disulfide groups. This phenomenon has been attributed to the catalytic reduction of hydrogen,’ but the mechanism has not been elucidated.2 In the case of copper(I1) or bismuth(II1) ions in 1 N sulfuric acid, the author observed a similar “catalytic currelit” in the presence of very small concentrations of hydrocarbons, e.g., benzene, toluene, the xylenrs, n-pentane, n-hexane and cyclohe~ane.~ Since the lrttter “catalytic current” is due to an increase in the reduction of copper(I1) or bismuth(II1) ions, not to catalytic reduction of hydrogen, the following experiments were carried out to re-exmiine the current of cobalt ions in the presence of cystine. Experimental The solution studied was a modified Brdicka’s electrolyte: 0.0025 Jf cobalt(I1) chloride and 0.00001 M cystine in a 0.1 A i ainmonium chloride-0.1 M ammonium hydroxide buffer . Polarographic measurements, which followed standard practice, were made with a Cambridge Polarograph (No. C 466573). The capillary had an m-value of 2.06 mg. sec.-1 and a drop time of 4 5 sec. in twice distilled wster a t a h2ight of the mercurj reservoir of 60 cm. (closed circuit) a t 18
.
(1) R Brdickrr, Coll. Czechoslou. Chem. Communs., 5, 148 (1933). (2) R. Brdicka, abad., 11, 614 (1939); Chem. &%sly,34, 59 (1940); Klumpar, Coll. 17zechoslov. Chem. Communs., 13, 11 (1948); A. G.
Strombeia, Zhu7. F Z Z . Khzm.,20, 409 (1946); J. Heyrovsky, Coll. Czechoslov. Chem Communs., 9, 273 (1937); 111. von Stackelberg and 11. Fassbender, 2. Eleklrochem , 6 2 , 834 (1958). (3) Some of the experimental results submitted in partial fulfillrnent of the reiiuircinents for the degree of Doctor. E B. Weronski, Department of ~Cherriistrp,Cni\ersity \$ arsza\ca. 1957.
I
-1
1
-1.4
1
-1.8
V.
Fig. 1,-Curves 1 and 3 arc polarograms of Brdicka’s electrolyte before electrolysis; S denotes the sensitivity of the galvanometer. Curve 2 is the po!arogram of aliquots of Brdicka’s electrolyte after t - l . a s min. of electrolysis at -1.35 v., or after L1.6 min. of electrolysis u t -1.65 v.
were electrolyzed a t either - 1.35 v. (where the diffusion current of cobalt is not distorted) or at -1.65 v. (where the enhancement of current is maximal) until the diffusion current measured at -1.35 v. decreased by about 30%. illiquots electrolyzed a t -1.35 v. for t--1.35 minutes, or a t -1.65 v. for t--1.65 minutes gave identical polarograms (curve 2). We found that the amounts of electricity in-
