1752
JAMES
E. BOYLEAND G. M. HARRIS
in total volume due to the drop in mercury level in the capillary is negligible) the results might not be too reliable due to the increased solubility of nitrogen with increasing pressure. Figure 4 shows the first-order plot for the process, log [ P , / ( P m Pt) ] versus time yielding a straight line. The induction period and the scatter of points in this region probably can be explained by considering the time
Yol.
so
required to build up nitrogen saturation in the system and possible thermal equilibration effects. The value of k d a t 30.5’ was calculated to be 1.31 X 1 0 7 and is included in Fig. 1. Using all the available data for the decomposition of Azo I in solution the equation which best fits the experimental results is: = 1.58 x l O I 5 exp(-30.8 kcal./RT).
[CONTRIBUTION FROM DEPARTMENT OF CHEMISTRY, UXIVERSITY OF BUFFALO 1
Kinetics of the Exchange Reaction between Carbon-14-labeled Carbonate and
Carbonato-bis-(trimethy1enediamine)-Cobalt(II1) Complex in Aqueous Solution. Effect of Steric Hindrance in a Ligand Substitution Process BY JAMES E. BOYLEIAXD G. Ad. H.%RRIS RECEIVED AUGL-?T29, 1957 The study of carbonate ligand exchange reactions of the carbonato-diamine-cobalt(II1) series of complex ions has been extended t o the bis-trimethylenediamine (tnn) analog. The effects of changes in reactant concentration, p H and temperature have been determined. Comparisons are made between the behavior of the tnz complex and those previously investigated; viz., the pentammine ( N g ) , tetrammine ( 5 4 ) and bis-ethylenediamine (en2) types. The rates and mechanisms of the reactions which promote carbonate exchange are very similar for N, and ens, b u t large variations from these are observed for the i T 5 and tn2 forms. The N5 behavior can be reasonably attributed to hydrogen-bonding factors; for the tnz, steric hindrance appears t o play the predominant role.
Previous publications in this series have reported on the nature of the exchange reaction in aqueous solution between free (uncomplexed) carbonate ion and the corresponding ligand group in a number of carbonato-cobalt(II1) complex ions. Included have been the tetrammine (N4),2a,bthe pentammine (N5) and the bis-ethylenediamine analogs. The mechanisms of the exchanges conform to a definite pattern of similarity, but some differences do occur both in the relative rates of the reactions and in details of the mechanisms. The present paper reports an extension of these. studies to the bistrimethylenediamine (tny) complex, which differs from the en2 compound only in the insertion of one more CH2group in each of the diamine-cobalt chelate rings. Evidence accrues from this study concerning the effect of stereochemical factors on the rate and mechanism of the ligand replacement reaction. Experimental Preparation, Purification and Analysis of Materials.-Carbonato-bis-(ethylenediamine)-cobalt(II1) chloride was prepared as described p r e ~ i o u s l y . ~The corresponding tn2 salt was prepared by an extension of the procedure of Bailar and Ur~rk.5,6 (CotnpC1,)Cl was first obtaii:ed. then aquated by allowing it t o stand in dilute aqueous solution a t room temperature for 45 minutes. Excess of AgZCOa mas dispersed in the solution and kept in suspension by rapid stirring for three hours. T h e mixture was filtered and the filtrate retreated with silver carbonate for one hour and refiltered. Slight excess of BaC1, was added t o the bright red filtrate containing (CotnzCOa)zC03. After standing overnight, the BaC03 was filtered off, and the filtrate evaporated to dryness in the open air at 70°, causing (1) Work done a s part of Ph.D. requirement of t h e University of Ruffalo, 195G. Complete report available from University Microfilms, Ann Arbor, Michigan. (2) (a) G >Harris I. and D. R. Stranks, Trans. Faraday Soc., 48, 137 (1952); ( b ) D. R. Stranks, ihid.,48, 011 (1952). (3) D. R . Stranks, ibid., 61, 505 (1955). (4) J. S. Holden and G. hl. Harris, Tms J O U R N A L , 77, 1934 (1955). (5) J. C. Bailar and J. B. Work, ibid., 68, 232 (1946). (6) Experimental quantities of trimethylenediaminr wprp obtained 3s R gift from !he .4wrrjran Cymsinicj C!v71!i~t?y
precipitation of the cycess BaC12 as BaC08. T h e cobaltic complex was extracted from the residue with water, recrystallized twice from water, and air-dried at room texnperature. T h e purified dry solid (CotntC03)Cl,H%Oremained completely stable at room temperature indefintely. I t s complete analysis was carried out as follows: water of crystallization was determined by drying in air a t 110”, cobalt by the method of Hillebran?,? and by direct ignition of the complex to CoaO4 a t 700 , carbonate by a modification of the Fresenius vacuum baryta method,8 chloridp by standard gravimetric p r o c e d ~ r e ,and ~ nitrogen by a modified Kjeldah1 method.1° Anal. Calcd.: HzO, 5.62; Co, 18.38; COS, 1S.71; C1, 11.06; iT, 17.47. Found: HZO,5.75; Co, 1S.26; COa, 18.73; C1, 11.11; N, 17.54. Carbon-14-labeled sodium carbonate was prepared by reaction of barium carbonate-C-14 with lead chloride and subsequent absorption of the active dioxide in carbonatefree sodium hydr0xide.I’ Apparatus and Technique.-The kinetic runs were made in a manner very similar to t h a t of the previous studies.2-‘ A typical run was commenced by niixiny rapidly a weighed portion of the anhydrous complex salt, 15.00 ml. of 0.0831 M sodium borate, 1.00 ml. of 0.3153 h1 inactive sodiuin carbonate, 7.00 ml. of Cos-free water. and after temperature equilibrium was established, 2.00 ml. of 0.0321 M active sodium carbonate. A zero-time sample was removed and the free carbonate precipitated within 90 seconds. Subsequent samples were withdrawn a t convenient time iiiten als, a stream of CO?-free nitrogen being passed through tlie reaction flask during sampling. Assay was effectd h y eiitlwindow counting of the iinifornily-sprend BaC08 samples in nickel-plated planchets. Half-times of exchange mere obtained from the conyeiitional log (1 - f ) vs. time plots,’? using least-squares aiialy(7) W. F. Hillebrand, e i ai., “Applied Inorganic Analysis,” 2 n d Ed., John Wiley and Sons, New York, N . Y., 1953, p. 418. ( 8 ) W. J. Johnson, R o y . A w l . Cizem. Inst. J . PTOC..17, 327 (1950). (9) I. M. Rolthoff and E. B. Sandell, “Textbook of Quantitative Inorganic Analysis,” 3rd Ed., T h e Macmillan Co., New York, S. Y.. 1952, p. 307. (10) A. Eisner and E. C. Wagner, I i i d . E n g . Chein., Anal. Ed., 6 , 473 (1934). (11) N. Zwiebel, J. ’Turkevirh and W. Miller, THISJ O U R N A L , 71. 376 (1949). (12) Friedlander and Kennedy, “ 5 u c l e s r m t l Radiochemistry,” Jghn Wiley and Sqns, I c r , Vrw x 7 w k , PIT. V I ?9.X p. ,‘{I(;,
EFFECTOF STERIC HINDRANCE IN LIGAND SUBSTITUTION PROCESSES
Feb. 20, 195s
TABLE I DATAOF RUN No. 6 Temperature = 50.5" (carbonate) = 0.01518 pH 9.50 (complex) = 0.03505 AI Half-time = 156 min. (borate) = 0.05 AI
M
Spec. activity
(At)
Time, min.
Wt. of BaCOa. mg.
cts./min.'/mg.
0 60 120 180 240 300 350
4.94 4.21 4.73 4.86 4.72 4.15 4.75 4.19
785 6T5 570 507 440 400 375 262
03
(1
-f)
1.000 0.790 .589
.468 .340 .264 .218 000
.
783
and can in fact be abolished by a minor re-interpretation of the data (see Discussion). B. Stability of Aqueous Solutions of Co tnzC03C1.-Solutions of this tnz salt were found t o be unaffected by heating a t 50" for as long as three days in the presence of BaC12, NHdOH, or both. I t s absorption spectrum remained constant,l6 as did its p H and electrical conductivity. It is clear that not even the minor changes observed with the enz analog4take place here. C. Kinetics of the Exchange Reaction.-A series of runs was made in which concentrations of the carbonato complex (a), total free carbonate (b), and hydrogen ion (h) were varied. In each (Carbonate), M X lo2. 2 3
sis of the data to achieve best fit. The fraction of reaction i s f = (A0 - A t ) / ( A o - A , ) , where the A's are specific activities at t = 0, t = t and t = a, respectively. The final activity A , was determined experimentally by allowing exchange to proceed for a minimum of 8 half-times. Data of a typical run are presented in Table I. p H measurements were made on a Beckman Model G pH-Meter, and were corrected for sodium ion concentrationla and temperature."
1 I
I
I
4 0
8
Results
A. Reinvestigation of the pH-Dependence of the Exchange of Carbonate Ion with CoenzC03+ Ion.-Holden and Harris4 studied the pH-dependence of the Coen&Os+/carbonate exchange in absence of borate buffer. It appeared from other data that use of the buffer might have introduced a negative catalytic effect, but no test of the premise was made. It was considered worthwhile to repeat their pH-dependence study in presence of 0.05 M borate. Figure 1 records the results obtained and 2 4 6 8 (H+), M X 1Olo or (complex), M X 10'. Fig. 2.-Variation of exchange rate of tnt complex with reactant concentration : A, (H+) variable; (complex) = 3.51 X 10-2 M; (carbonate) = 1.52 X M; 0, (complex) variable; (H+) = 7.59 X 10+0 M; (carbonate) = 1.52 X 10-2 M; 0, (carbonate) variable; (H+) = 7.59 X 10-10; (complex) = 3.51 X 10-2 M; (borate buffer) = 0.05 M; temperature, 50.5' in all cases.
0
1
0
1
2
3
4
5
6
(H+), M X 10'0. Fig. 1.-Variation of exchange half-time of en2 complex u i t h concentration of H +: 0, present investigation, (borate) = 0.05 M; A, Holden and Harris, no borate; (complex) = 3.51 X 10-2 M ; (carbonate) = 1.52 X lo-* M ; temperature, 25.00'.
compares them with the earlier work. It is apparent that the buffer has no significant influence on the exchange rate. The 0.8 inhibition factor surmised by Holden and Harris is therefore spurious, (13) H.F.Walton, "Principles and Methods of Chemical Analysis," McGraw-Hill Book Co., New York, N. Y., 1951. (14) "International Critical Tables," Vol. I, McGraw-Hill Book Co., New York, N. Y..1926,p. 81.
run, pH control was achieved by a suitably adjusted 0.05 M sodium borate-boric acid The half-times were determined as described above, and the various rates of exchange calculated according to the standard r e l a t i ~ n R~ ~=. ~In~ 2/t1/, x ab/ (a b). The results so obtained are presented in Fig. 2 . Some experiments were performed in which
+
(15) The spectrum was determined between 3000 and 6500 A. on a Beckman D.U. spectrophotometer. Two peaks are observed, which compare with those for the en, complex as follows: ( E = (log Ie/I)/Cl = molarextinctioncoefficient): ens: 5200,E = 127.4; 3600,E 119.8; tnr: 5200, E 108.5; 3600, E 127.7. (16) Several runs were made in which all factors remained constant except the borate Concentration, and no rate di5erences were observed showing t h a t borate has no influence on the exchange in the concentration used. Likewise duplicate control experiments in light and dark or in glass-bead-packed and unpacked vessels ruled out possible photochemical or surface effects. (17) This relation ignores isotope effects. T h e equilibrium isotope effect i n t h e tn; system is about 7% (see forthcoming publication), and only very small errors in t h e rate constant values will be introduced by neglecting i t (see G. M. Harris, Trans. Faraday Soc., 47, 716
-
(1951)).
-
-
784
JAMFS
I
E. BOYLEAND G. M. HARRIS
6
0.5
the ionic strength was varied over a wide range by addition of sodium chloride, as shown in Fig. 3. The temperature dependence of the exchange was studied under three different sets of concentration conditions, and the results were analyzed by means of the usual Arrhenius plots (see Fig. 4). The activation energies were identical in all three cases, 24.8 f 0.9 kcal. /mole.
so
GIL-Y
1
1
I
1.o 1.5 2.0 Ionic strength. Fig. 3.--T.’ariation of exchange rate of complex with ionic strength: (Complex) = 3.51 X hf; (carbonate) = 1.52 X If; (borate) = 0.05 M; pH 9.12; temperature, 50.5O. 0.0
VOl.
5 41.8
1/T 2.0 X lo3, 2.2 degree-’. 2.4 2.6
2.8
dependence of the exchange: 0, Fig. 4.-Temperature Series 1 (complex) = 1 7 5 X -11; (carbonate) = 1 52 10-2 0 , series 2 (complex) = 3 51 (carbollate) = 1.52 x 1 0 - 2 ~f,. A, series 3 (complex) = 3.51 x 10-2 JI; (carbonate) = 4.16 x 1 0 - 2 JF; (borate) = 10-10 in all series. 0 05 ,“.I; (H +) = 7.59
x
lv
x
respectively, the curves of Fig. 2 have been plotted. Excellent agreement with the experimental points is seen to be obtained.18 The variation of R with ionic strength is slight Discussion (Fig. 3 ) , a small decrease occurring a t high pvalues. Figure 2 indicates that the tnz complex exchange Under the conditions of the study, about l / 3 of R reaction is essentially independent of pH. It is is contributed by the second-order ion-ion reaction, first-order in complex concentration and appears and 2,’8 by the first-order aquation. This corto be part zero-order and part first-order in total responds closely to the situation discussed in defree carbonate concentration (HC03COS=). tail previously for the CoN4C03+species,2band the A mechanism which satisfies these conditions and tn2 results may be dealt with similarly. develops logically from the concepts deduced in I t remains to compare the carbonate exchange the previous studies of the series*-4is (L = tnz) behavior of the four analogous carbonato-cobalt(111) salts for which data are now a t hand. The C O L C O ~ + HzO _r C O L C O ~ . I I ~ O + (1) relevant experimental results appear in Table ITA. CoLC03.H20C + H C * 0 3 - z
+
+
COLC*O~~H~ +OHCO3+ + HzO I _ CoL(H20)2+’ + COS” CoL(Hz0)2+++ + C*O,‘ e COLC*O~.€I~O ++H?O COLC*O~.HZO eCOLCO~’+ H2O CoLCO,.H?O+
+
(2)
TABLE IIA
(2)
SUMMARY O F EXPERIMENTAL RATEDATA k’, k”, E’, E”. 1. mole-1 1. mole-1 Temp., cal./deg./ min. -1 min.-l ‘C. mole 7 . 8 x 106 8.7 x 1 0 6 U 0 15.8 1 9 . 9 1 . 3 x 10-2 9 . 0 x 106 20.3 16.7 20.6 1 . 6 X 10-2 0 . 0 x lo6 25 ..b . .b 3.5 x 1 0 - 2 Q o x 5 0 . 5 24.8 24.8
(3’) (1’)
The rate law corresponding to this mechanism is’’
Complex
s, N4 en2 tnn
Ref.
3 2a 4
Present work
a Includes equil. const. Kj. Units are l.z min.? for this k ‘ . * Kot adequately studied for this compound. I h e s not include K s . Units are min.-’ for this k ” .
I n this, the small k’s are forward-reaction rate constants and the large K’s equilibrium constants. Kd refers to the reaction HCO3-
Using “best-fit” values of
and
H+
+ CO3‘
TABLE IIR RATEDATARECALCULATED TO 25’ r\’S
(4)
k’
Complex
k”
9.3 x 1.7 x 6.0 X 3.3 x
1.0x 108
0.021 ,016 tnz ,0012 See text below. N4
en2
-
koa
106 105 106
0.07 .2
10-5
..
t?
. I
(18) Since Kr 0.7 X 10-11 a t 50’ (Harned and Sctioles, Trrrs 63, 1706 (1941)), h / ( h Kt) is nearly constant in t h e p H range of t h e study. Hence the near independence of pH of t h e rates of exchange observed. JOURNAL,
+
EFFECTOF STERIC HINDRANCE IN LIGAND SUBSTITUTION PROCESSES
Feb. 30, 193s
Table IIB shows the same data computed to the common temperature 25 ’, utilizing the known activation energy information. The results for the tetrammine have been recalculated on the basis of the mechanism suggested for the en2 ~ o m p l e x which ,~ consists of reactions I, 2 and 1’ above (L = N4 or en2),together with
+H+ + H2O
CoLC03.H20f CnLHCO?.H20-+
CoLHC03 H,O++ ( 5 )
CoL(H20)2+++
+ HC*O3- _r
CoL(H?O)?++-
+ HCOi-
(6)
+ +
COLHC*O~.HQO++ HzO (6’) C O L H C * O ~ . H ~ OJJ + + C O L C * O ~ . H ~ O +H + ( 5 ’ )
The corresponding rate law is
so that for N4 and en2, k’ is the same as for tn2 and = k&jK1(H20)/(K1(H20) 1). This law not only fits the N4 data much better than that originally proposed,2n but also eliminates the unexplained arbitrary intercept found necessary to make the older expression conform to the PH dependence data. For the en2 complex, k’ and k” have been altered slightly from those given in the earlier publication4 in order to fit the p H variation data more satisfactorily, while eradicating the spurious 0.8 borate inhibition factor mentioned above (see Experimental). The mechanism for the pentammine3 consists of reactions 5 , 5’) 6 and 6’ (reading N5for L.H20),together with
+
k”
CoN5HC03++f HC*03CoNsHC*Oa+’ f HC0’-
The rate law is R
=
+” K1) +
kJSbhab (h-
ksKdH20)ah
(7)
(c)
so that for N5, k’ = k7K5 and k” = k&,(HzO). Of the various values for k’, the “direct” exchange rate constants given in Table IIB, only those for the N4,en2 and tnz compounds are strictly comparable. For the first two, the values differ but little, as would be expected, since in going from N4 to enz, one merely replaces ammonia by NH2R (neglecting chelation effects). However, the value for tn2 is anomalously low, since tn2 differs little from en2 in their most significant chemical properties.Ig The only possible explanation must be a steric hindrance which can reduce the magnitudes of k z or K1 or both, in the case of tnz. There is evidence] based on the lack of an equilibrium isotope effect, that K1 is large for the en2 c ~ m p o u n dand ,~ this probably is also true for N4.20 A much smaller K1 for the tnz complex would justify the observations] and comparison of molecular models of the
785
enz and tnz forms does in fact support a spacial interpretation] since it is sterically much more difficult to carry out reaction 1 with the tn2 than with the en2 compound. One must explain the large discrepancy in the Ns behavior as regards “direct” exchange. This has been attempted previously in terms of ligand electronegativity variations. A more reliable approach now appears to be to consider hydrogenbonding factors. In the N5 compound, only NH3 groups are cis to the COS,while the exchanging species in aqueous solutions of the Nq, en2 and tnz forms have H 2 0 cis to COS. I t is well known that the 0-H . . . . 0 hydrogen bond is much stronger than N-H . . . . 0.21 Hence the coo group in CoLCO3.HZ0+(L = N4, en2, tnz) retains considerable bidentate character and has a rather smaller tendency to accept a further proton than does the N5 type. The latter readily becomes protonated,22 and it is the bicarbonato form that undergoes rapid “direct” exchange,23 accounting for the p H dependence of this process, in which it contrasts with the N4, en2 and tn2 complexes. Turning now to the k” values, the rate constants for the “aquation” mechanism of exchange, only those for N,, N4 and en2 are comparable (k” = k6’ & ( H Z O ) ) . ~The ~ KO’s should be in the approximate ratio 1:3:10 for N5, N4 and en2, respect i ~ e l y . Since ~ ~ Kg for N.5 is 2.5 X lo6 a t 2 5 ° , 2 2the various k6’s are as given in Table IIB, corresponding to K , values of 1.7 X lo5 and 1.5 X l o 4 for N4 and en2, respectively. The tn2 reaction is not only much slower but proceeds by a pH-independent mechanism. Again, a steric explanation is reasonable-the models show that hydrogen bonding is so favored in the compound Cotn3C03.H20 + that reaction 5 has a much smaller tendency to occur than i t has for en2 and Kb is in consequence smaller by perhaps a few more orders of magnitude. The main path for aquation of the tnz complex is thus by direct replacement of water for C03=, according to reaction 3. The latter reaction is also sterically difficult, thereby contributing, along with the small K1,to the smallness of k ” . Finally, the activation energy data, though incomplete, offer some further support of the arguments advanced above. For N5, E’ = E7 AH5. The temperature coefficient of reaction 5 for N5 suggests AH5 - 10 kcal./mole.22 Thus, E , 26 kcal./mole, somewhat higher than E 2 for N4 (17
+
-
(21) C. A. Coulson, Research, 10, 149 (1957). (22) A. B. Lamb and K . Mysels, THISJOURN?LL, 6 1 , 468 (1945), report Ka = 2.5 X 106 at 25’ for Ns. (23) Using Lamb and Mysels’ Ks value, it is seen that k , = 40 1. mole-’ min. -1, about 200 times as great as kn for the Nd species, (24) Assuming K1 at least as great as the order of unity (see ref. 20) for Nd and em. (25) The following rate constants of aquation have been observed for chloro-cobaltic complexes: Co(NHdrCl.HnO++ k = 1.3 x 10-4 min.-l at 20° (J. N. Brbnsted, Z. phvsik. Chem., 122, 383 (1926)); k 2.0 X 10-8 m h - 1 at 30’ (data of J. P. Co enr Cl.HrO++ Mathieu, Bitll. SOL. chim., 151 3 , 2121 (1936), re-interpreted in the manner of Brbnsted, above); Co(NH:)aCl++ - k = 4.0 X 10-4 (19) For example, their basicities (G.Bredig, Z . physik. Chem., 13, 309 (1894)). m h - 1 at 35” (R.G. Pearson, C. R. Boston and F. Basolo, J . Phys. (20) An equilibrium isotope effect has been observed for the tetramChcm., 69, 304 (1955)). Brbnsted’s temperature dependence study mine (D.R . Stranks and G.M. Harris, J . Phys. Chem., 6 6 , 906 (1952)) for the corresponding nitrato salt ( i b i d . ) indicates E.,$ -25 kcal./ though this has been questioned (E. Saito and B. Lazard, J . Inorg. mole for this type of reaction. Thus, it can be estimated that at Nitcleur Chem.. 1 . 218 (1955)). 25’ the k values are roughly in the ratio 1 : 3 : 10 for the Na, Nd and en2 . . I n anv ~.case. reaction 1 need not be complete for Nd-with K I only of the order of unity, K I ( H ~ O ) / K I - types of chloro-cobaltic ions, and one would expect a similar ratio for (HsO) -I-1 E 1. the bicarhonato salts of reaction 6.
-
-
-
RICHARD C. JARNAGIN
786
kcal./mole). However, reaction 2 is rather more complicated than (7), since it involves both carbonate replacement and H + transfer, precluding direct comparisons between them. E” = EO AH6 for both Nd and N6, making E6 30 kcal./ mole for both ammines, in acceptable agreement with Bronsted’s data for the corresponding reaction of the nitrato salt.25 For tn2, E’ lies between E2
-
[CONTRIBUTION NO.
+
1463 FROM
THE
AND JUI
H. ~ V A N G
Vol.
+
so
and E? AHl, depending on how small K1 is. Similarly, E” lies between Ea and Ea A€I; for tnz. A H l is not known but is probably quite small, so that E2 S E3 S 25 kcal./mole in this case. Support of this research by the U. S. Atomic Energy Commission through Contract AT(30-1)-15% is gratefully acknowledged. BUFFALO, N. Y .
+
DEPARTMENT O F CHE>lISTRY, YALE
UNIVERSITY]
Investigation of the Catalytic Mechanism of Catalase and Other Ferric Compounds with Doubly O18-labeled Hydrogen Peroxide
c. JARNAGIN’
B Y RICHARD
AND JUI
€1. 1v.YSNG
RECEIVED AUGUST 19, 1957 A method for determining the origin of molecular oxygen in the catalytic decomposition of hydrogen peroxide is developed and applied t o t h e catalytic decomposition of hydrogen peroxide by catalase, polyamine-Fe(II1) chelates and ferric hydroxide. I n all cases studied, the oxygen molecule originates from the intact 0-0 bond in hydrogen peroxide.
One method of approach to the study of catalysis is via the synthesis and investigation of relatively simple model molecules of known structure. By this means Wang2 discovered the remarkably high catalytic efficiency of triethylenetetramine-Fe(II1) chelate for the decomposition of hydrogen peroxide. The postulated mechanism of catalysis involves the metathetical displacement of a bound OH- by OOH-, followed by the displacement of another bound OH- by the second oxygen atom of the same OOH- as illustrated below.
riched water vapor. If the isotopic atom-fraction of OIs in the resulting HzOz is X , the mole-fraction of HOl80IxHshould be approximately equal to X 2 . (Here it is assumed that the isotope-effect due to the difference in zero point energies of light and heavy 0-atoms in I - 1 2 0 2 is negligible.) When this HzOz is catalytically decomposed, the mole-fraction of 018018 in the oxygen gaq liberated should also be X 2 no matter whether the decomposition takes place through mechanism (A) or mechanism (B), because the isotopic distribution in this HrO2 is in accordance with natural probabilities. However, when \Le/OH I 013 each mole of this 0IX-enrichedhydrogen peroxide is +OHmixed with q moles of ordinary hydrogen peroxide, I I1 the isotopic atom-fraction of OI8 and the tnoleSince the 0-0bond in compound I1 is under strain, fraction of H01a018Hin the mixed H202 become 4) and [X’ q(0.0020)2]/ it becomes more reactive and reacts subsequently [X q ( @ . O O Z O ) ] / ( l q ) , respectively, where 0.0020 is the isotopic with a second OOH- to produce 0 2 and regenerate (1 atom-fraction of Olx in ordinary HzO?. This last the catalyst. There are two alternative mechanisms for this isotopic distribution is unnatural, for had this last reaction step to take place. These are depicted H202 been made directly from enriched water vapor with atom-fraction of Olx equal to [X below as mechanisms ( 4 ) and (B), respectively. q(0.0020)]/(1 q ) , then the mole-fraction of H018018Hin the resulting Hz02 would be [X q(0.0020)]2/(1 a)?. Now if this 018-enriched H202 with unnatural isotopic distribution is catalytically decomposed by, say, triethylenetetramine-Fe(III), the mole-fraction of 018018 in the liberated 0 2 will be different for mechanisms (A) It is of particular interest to find out whether the and (B), respectively. hlechanism (B) involves a catalytic decomposition of hydrogen peroxide by reshuffling of the 0-atoms, ;.e., restoration to the triethylenetetramine-Fe(II1) chelate takes place natural distribution of the isotopic species, and hence predicts a mole-fraction of [X q(0.0020)]2/ through mechanism (A) or (B). 4)’ for O’eO1bin the liberated 0 2 gas. But a In general even for cases in which the detailed (1 mechanism is unknown, it is still of interest to know mechanism such as (-A), which produces all the 0 2 whether the two 0-atoms in each 0 2 liberated came from the intact 0-0bonds, leaves the original isofrom the same HzO2 molecule or from two separate topic distribution undisturbed, and hence predicts the mole-fraction of 0 * 8 0 1 8 to be [X? q(O.OOZO)?]/ H202 molecules. (1 4 ) . Principle of the Method If X = 0.10 and q = 1.0, we have [ X 2 q Suppose O18-enrichedhydrogen peroxide is made (0.0020)2]/(1 q ) = 0.0050 and [X q(0.0020)]2/ directly by an electric discharge through OlX-en- (1 q ) ? = 0.00%. Consequently a mass-spectrometric analysis of the 0 2 produced by the cata(1) National Science Foundation Predoctoral Fellow, 1947-1958. lytic decomposition of hydrogen peroxide could (2) J. H. Wang, TaIs J O U R N A L , 7 7 , 4715 (1955).
[/: \OOH]+
-[>:.
