NOTES
1892
VOl. 65
NOTES THE RATE OF ELECTRON TRANSFER BETWEEN THE TRIS(ETHYLENEDIAMINE)-COBALT(I1) AND COBBLT(II1) IONS
to the chloride with concentrated hydrochloric acid, and crystallized by the addition of alcohol. The chloride was twice recrystallized from aqueous alcohol containing a little hydrochloricoacid and air dried. The ure material gave [Q]D 154 , ( I dm. tube.). Anal. 8alcd. for [CoenslC13.H20: C, 19.82; H, 7.21; N,23.12. Found: C, 19.82; If, 7.19; N,22.90.
+
BY FRANCIS P. DWYER AND ALANM. SARGESON
12 5
Biological Inorgmic Chemistry Scdion, John Curlin School of Medical Rescorch, Australian NotionaE University, Canberra, A ustrolia Rece-iaed December 18, 1980
The rate of the slow electron transfer reaction between tht: tris-(ethylenediamine)-cobalt(I1) and -cobalt(III) complex cations has been measured by Dhe radio-isotope method.’ Since the oxidized species in the reaction is optically stable, the reduced form optically Iabile, and the optical density of the mixed species in solution is not great, the rate of electron transfer can be measured conveniently by the racemization method, as suggested by Busch.2 It will be evident that electron transfer proceeds through two reactions
7‘ 1.00 d
.-s I
2 2 n-
X
g 0.75
B~OW
d-CoenpJ+4-l-Coeuaa+-< GCoenla+
+ d-Coeritz+
(.i)
+ d-Coen*Z+
(B)
BIOW
d-Coen3a
+
-t-d-Coens2 +
+
d-Coen3sf
fmt
d-Coena2
+
1~’
I-Cocntz+
0.50
i
0.0
0.2
0.4 0.6 0.8 “free”(en), ill. Fig. 1.-The dependenw of the rate const:int K on ethylenediamine and cobalt(I1) coiiceritration for [Co(III)] = 0.0551 dl, fi = 0.45, ?’ = 98”, 0 , 0.030!) M Co(I1); 0 ; 0.0103 ,?f Co(I1); 0 , 0.0206 M CoiIIj.
Each act of (electron transfer according to (A) kads to inversion of thc Co(II1) specics, but electron transftx produces no rotational chauge in (13) arid cannot 1)c measured. It follows that the The wtction misturc, prqxtrcd f r o m t;oliit,ions of thr apracemization rat,e is equivalent to twice the elcc- proprkrte concent.rat,ion of piire d-[CocilJjC13.11?(?, i2n:il:tr. tron transfer rate, (A), since the racemization cobalt( 11) chloride &hydrate arid potassium :.hlorid(~,WR: frecd from t,racre of osygcii by continuotie flow of “oYvgc.n rate is twice t,he rate of invcrsioii. frre” nit,t,oyc,n. Thc nitrogen first W ~ Epassod tlirough The dissociation of the tris-(ethS.lciic!~li~~IiiiIie)-anolhcr solution of the same composition, maintairicvl :It thct cobalt(I1) ion in aqueous solution can be linijted same tc~nipcmturr,in ordcr to a.void loss of wntw :tritl ctliyil’iire ethylenediamine (15 M , O X Y ~ O I I froe) by using an excess of the free base lmt, as a result, enetliamiiw. ~ i a :&ed s sul)scqucnt,l:+~ l o st:i.rt thca rcartion forinill:: the a, competing reaction is introduced, \\hich also cohalt(I1) comploz in sill!. Constmt rc?:tcf ion r:ttcs w r ( ! leads tto racemization, but without electroil trans- obt.niricd wit.hiii 15-20 minutes at, (thehe compounds can he separatlxl oiily with difiicsiiltyij. The assignments For cis 1H :LiFI and IH : 3H given Mixture of trans 1H/3H and trans 1H/4H, h.p. 79" in Table I1 ar':: drrived from suggest,ioiismade b y i Musher." Tliesc hare been clecliiccd from the rpH current chemical shifts foulid i l l the reso~iaiice spectra of the c i s 113:2€1, IH:311 and 111:4€1 isomers. From tiles(. :iskigiiincvit\ it (mibe seeii that proximity t o hydrogw ntoriis through 1)onds on the fliioriiie :ttoms in a
{
le,3e
'
'
veitigation of' c i s atid trans perfluorodewlin l'homns6 has field doublet,of t h c low trmperatiire l 9 l ? spwtrum of perfluororyc.lo1icrallr. t o tht. axial fiuorine ar oms. Thc conc~lusionsof Thomas are based on the nssumptioil that ?piti--spin interactioii hrtween trans fluorine atonis on adj:tci>nt ( w \ ) o i i atoms is greater than t'hat' hetween cis fluorine atoms in a cyclohexane ring. Thr. rerulta of Tiers and Thomas and the suggestions made h y Musher nll hare been used R. P. Smith and J. C . Tatlow, J . Chem. Soc., 2505 (1957). ( 5 ) J. I. blusher, private communication. ( 6 ) J. Homer a n d L. F. T h o m a s , Proc. Chrm. a o r . , I 3 0 i i . > ( i l ) (4)
-le,5e
i
-
8
ln,2a
9
2,2r
5,5'
I 2e,2a,,?e,3a _ oe,,ia,fic.,t?:i
',
\
f '1
1 2 3
1 :
5 6 7
1815 2010 2085 2255 5860 590.5
2 100 J
i
6120 6165 Atoms marked with a primc are trans to the nearest hydrogen atom via the bonds in an intcrconvrrting molecule. The boiling points quoted in the paper by P. ,Johnrock, \V.K. Musgrave, J. Feeney and I,. H. Rutcliffc, Chem. ond Ind., 1314 (1959), require interchanging.
',
8 9
'j
in obtaining the modified set of assignments for thq dihydrodecafluorocyclohexane given in Table 11.
NOTES
1896
D I - ~ ~ T I ~ I ~ ~ ~ O X Z E O TAS R OAP EMEA4XS S OF RESOLTJTTON
Vol. 65 (2.57) dt
BY CI IFFORD J. .MCGINN Depurtment of C'hemaslry Le M o y n z College, Sgracuse, N . Y. Recczwd February 97, 1961
This is the rrport of an attempt to resolve a racemic mixture by means of the fractional distillation of the mixture of its djastereoazeotropes. If each of the enantiomers dl and II forms an azeotrope with a third optically artive substance, dz, the two diastereoazrotropes may dit'er sufficiently in boiIing point to permit resolution by fractional distillation, particuiady if the ternary system dlli-dz should i t d f not contain a ternary azeotrope. If the system dots form such a tertiary azeotrope the possible rrwlution will thewby be restricted, the more so the more nearly equal are the concentrations of dl and 1, in the azeotrope. The specific resolution attempted was that of d,l-2-octanol ( = dl and Z1) u4ng d-limonene as the third component ( = d?). Several different solutions, varying jn mole frsctioii of d, from 0.1 to 0.9, were fractionally distilled a t several different temperatures a i d pressures. The fractions were analyzed for the optically active component d-limonene by mea115 (If infrared, and its contribution to the opticd rotation mas subtracted from the observed rotntion. l o positive evidence that resolution had ocwrred rouid be found. The fractionating columri consistrd of a three-necked round-bottom flask inside a hmting mantle which was electrically controilcd. Noiintcd on this was a three-foot colun~n,p:wktd with Raschig rings, wound with nichrome >{ire,and jacketed with asbestos one inch thick. It3 temperatiire was controlled independently of the pot. On t o p of this was a distilling head containing a cold finger. The apparatus as described was equivaleiit to 20 theoretical plates. To determine why resolution had not occurred the dl-ll pair was resolved by conventional methods and the physical properties of the following sys tems were studied d,. 11, dl-Zi, dl-dz, ll-d2 anti dl-iI-d?. The actual systems were formed from d-2-octanol (dl), Z-2-octanol (11) and &limonene (d2). The physical propcrt ies measured were the boiling point, refractir-e index, density, and where applicable the temperature change on mixing. The above physical properties were determined simply to shorn the degree of non-ideality of the systems studird. The reriilts are shown in Table I.
Boiling point diagram (schem:ttic).
differences in the boiling points were expected to be small the actual boiling points were not taken. The recorded values are those take11 with a Beckmann thermometer which was adjusted to the anticipated ranges for the systems studied. The very fact that the racemic d,-ll solution is nzeotropic, and hence non-ideal, is noteworthy. The expectation has been expressed (1) that such mixtures should be strictly ideal, whereas these data show the d.1-2-octanol mixture to deviate from strict additivity in all the properties measured. There seems to be no theoretical reason, however, for evpecting strict ideality and strict addit,ivity. The deviation in the properties will probably always be small, but varying in degree from system to system. In the present considerations we are concerned with the boiling point diagram of the ternary system dl-ll-dt. The binary system dl-11 would be expected to be azeotropic, as observed, and of courie symmetrical. The rest of the ternary diagram, however, must be asymmetric, as it seems to he from consideration of the boiling points: 1.49' for the dl-& azeotrope, and 0.96' for the l r d s azeotrope. The d2 content of these two azeotropes should not be the same. While the results showed the compositions to be different, the differences unfortunately were within the limits of experimental error. The ohserved minimum ternary azeotropic boiling point (9.62') is consistent with the existence of the three hinwy minimurn azeotropes of the system. The cmnposition of the ternary azeotrope is expected to he somewhat on the ZI sidc of the system. The boiling point diagram, then, mth arrows indicating the course of the compositioii of the liquid residue during boiling, is of the following type d t mix. (schematic). According to these relations, some ~ L D (27.0') (27.60) 1.4'232 ..... resolution by fractional distillation should be 1.4232 . . . . . possible. The actual failure of the resolution then 1 ,4246 - 0.618 means that either the asymmetry of the system is 1.4711 . . . . . not sufficient for the fractionating technique used, or that the analytical procedure was not suffciently 1.4466 -1.380 1.447% . . . . . sensitive to detect differences. It should be pointed out that the 2-octanol wnc: rwovered as sodium 1.4460 -1.085 sec-octgl phthalate. It had no observable rotaThe systenis did,, ll-d2 and d,-ll-dz were all tion. 55.070 in $-limonene (the approximate composiThe author iq at present investigating the availtion of the ternary azeotrope dl-l;-d2). Since the .zbilitv of fractionating columns with an efficiency sufficLnt to effect a pGssih1e resolution. (l! 11. ?rlausar, Be?. 90, 299 (1958).
Oct., 1961
1897
hTOTFS
Acknowledgment.-This work is being supported by the American Chemicd Socicty Yetroleum Ikwarch Fund.
TIIE REACTION BETWEEN URANIUM (IV) AND HYDROGEN PEROXIDE'
decomposition over the whole range of reactant concentrations studied. Similar determinations in 0.5 M HCI04 or in solutions containing up to 0.006 M Na2S04,0.006 hl TICl, 0.008 M Co(I1) or 0.009 U(VI), or a t a temperature of 2.4' gave consumption ratios in the range of those listed ilk Table I.
BY F. B. BAKERA N D T. W. NEWTON
TABLEI
Uniterdtv of Califorriia Loa Alamos Scientifi Laboratory. Loa Alamor, New Mezico Received Ala7ch 20 1961
The purpose of this note is to describe some experiments which were done to help understand rhe reaction betneen U(ITr) arid Hz02in aqueous HClOd solutions. The pertinent oxidation potentials are such that the net reaction to be expected is tJ+4
+ HtOs = UOL" + 2II+
(1)
The results to be given here show that the reaction does not proceed directly by this simple reaction but that it is in part, a t least, a chain reaction and that it is accompanied by a small amount of HzOs decomposition. Experimental Solutions of U(VI), IJ(IV), HClO,, NaCIO, and LiCI04 were prepared as before.? Other U(1V) solutions were made by the reduction of U(V1) with metallic Zn or Pb. U(1V) made in all three wayb showed essentially the samc rate behavior. The H?02 was from two murces, 3OW0 solution from the Mallinckrodt Chemical Co. and 90Yo sollition from the Becco Chemical Division of the Food Machinerv and Chemical Corp. Both of these solutions were furthrr purified by distillation under reducrd pressure and gave iitially the samc kinetic results. The second fraction from the 30% solution had a disagr~eahieodor and gave slightly higher rates; so subsequent runs were made with diluted 90% material. Cu(C10,)2 was made by fuming Cu(YOq)awith HClO, and then crystallizing from concentrated HClO,. The other perchlorate salts were from the G . F. Smith Chemical Go and were used without further purification, except for Co( C104)2which was recrystallized. The U(1V) solutions wrre analyzed by titration using standard Ce(1V) sulfate. H20t was analyzed as an oxidising agent against standard Fe(I1) and as a reducing agent against standard C'e(1V) with the same results. Mixtures of U(1V) and HZ02 were itnalyzed for the difference in their concentrations by adding aliquots to eycess standard Fe(I1) and back titrating with standard Ce(1V). Thc rate rnns were madr spertrophotometrically by following the absorhsncr at 6475 A. where C(1V) absorbs relatively strongly The apparatus and general procedure have been dwcribed prPviorislv.2
MOLESOF HtO, CONSUMEDFOR EACRMOLE OF U(IV) REACTED (CONSUMPTION RATIO) Conditions: 2M HC104,23-25', solutions deaeratrd with A r -Initial
P$'l'd; 3.04 3.08 3.03 1.54 2.76 1.35 2.69 1.35 1.48 1.48 1.48 1.48
values---
[HtOzl/[U(IV)]
0.54
.55 .55 1.1 1.1 2.5 2.5 5.0 8.3
10,s 18.8 18.0
Consumption ratio
Approximate rate constant, M - 1 min. - 1
1.024 0.997 1.017
48
1.014 1.010 1. O X $ 1 1. ( J U
57 68 95 111 114 ..
.o:n
1.035 1.029 1.019 1.029
49 47
.. ..
..
Reaction Rates.-Preliminary runs made a t 25" in air-saturated 2 M HClO, showed good reproducibility and were in accord with a rate law which was first order in U(IV) and in H202. &st values for the rate constants for these runs were determined by the use of a non-linear least squares program* wbich minimized the sum of the differences between the observed and calculated absorbancr values. The average serond-order rate constant was found to be 56.8 with a standard deviation of 1.85 and a maximum deviation of 2.9 M-l inin.-'. The second-order rate law fit the experimenQa1data quite well; the average difference between t,he calculated and observed absorbance values was only 0.002. The HC10, used in some of these runs was prepared by vacuum distillation, but no significant differences were observed. Similarly, the water used in some of the runs had been distilled but once, giving no detect.able effect. Alt,hough oxygen in the air reacts with V(IV) a t an insignificant rateSat room temperature in 2 AI Results HCIO,, it has been found that it has a significsnt Stoichiometry .-The stoichiometry of the re- effect on the kinetics of the U(IV)-H202 reaction. action was examined closely since it is known3 that Four additional runs were made in which the soluthe reaction between Fe(I1) and H20?is quantita- tions were deaerated with argon; these gave avertive when Fe(1I) ic. in excess but that a large age second-order rate constants ranging between amount of extra E120?is consumed when HzOzis in 38 and 44 M - I min.-'. Two additional runs using excess. The rt,oichionietry of reaction 1 was stud- air-saturated solutions gave 53 and 54 M-' min. ied by analyzing known mixtures after reaction in agreement with the previous series. Two runs $ahrated solutions gave apparent either for [H?O?]or for the [l?(IV)]-[H,O,] differ- made wing 0% ence. The reliults of some typical determinations of second-order rate constants of 56 and 61 M-' the number of moles of H2O2 consumed for each min.-l. The data from the solutions which had mole of U ( W ) which reacted (consumption ratio) (4) We are indebted to R . H. Moore and Ivan Cherry of the Theoare given in Table T. The data show that, unlike the retical Division of this Laboratory for computations involving the Computer: the former for modifying existing least square case of Fc(II), thcre is a small amount of H202 IBM-704 programs for our purpose (see R. 13. Moore and R. K. Zeigler. LA(1) This work was done under the auspices of the U. S. Atomic Energy Cornmiasion. (2) T. W. Newton, J . Phys. Cham.. 63, 1491 (1959). (3) J. H. Baxendale, in "Advances in Catalysis." Vol. IV. Academic Press Inc., New York, pr'. Y . , 1952, p. 46 ff.
8367) and the latter for slope calculations. (5) J. Halpern and 3. G . Smith, Con. J . Chem., S4, 1419 (IRRC,), report -d[U(IV)]/df ~ [ U ( I V ) I [ O I ] / [ H ~and ] show that in 0 . i .If HClO' a%30° and 0.96 atm. of 0 , the half-tiiiie for the oxidatirrii of U(IV) is about 130 min.
-
I50
0.01 M U(VI), 0.004 M Hg(I1) and 0.003 A I Pb(1I) were without significant effect. The rate u a s roughly tripled by 0.005 J I hg(1) and by 0.01 d1 Mn(I1) while 0.0002 AI Fe(I1) increased the rate about sixfold. The ions of Cu(1I) and Co(I1) nere found to inhibit the reaction. The decreases in rate observed with these ions are summarized in Table 11. It is to be noted that small amounts of Co(I1) cause less inhibition when Ro = 2 than when RO= 0.5, but that this effect is reversed at higher Co(I1) r omen tra tions.
,
'
f 100 P
I
=.
1
50
TABLE
been deaerated with argon fit a second-order rate law much less well than those from the air saturated solutions mentioned above. Apparent second-order rate constants were computed from the slopes of the absorbance versus time functionsJ; these rate couqtants were found to decrease uriiformly during the course of the run, having dropped about 20yGby the time the reaction was 7.5%) complete. It thus appears that oxygen in thf. solution.: increases the rate and allows better adherence to a second-order rate law. Consumption ratio determinations showed that there is a small amount of induced oxygen reaction during the U(IV) -H?O? reaction. The data obtained in all the runs made at 23' in deaerated 2 M IIClO, are summarized in Fig. 1. In order to make the data more readily comparable, the observed rates have been divided by tht, colicentration product [U(IV)] [HZ021 to give ic, thc apparent second-order rate coiistant, and plotted against R, the [H202]/[U(IT.')]ratio. The individual runs are shown as separate lines. Although there is considerable scatter among the various runs, it is clear that, the apparent second-order ratc, constants are larger in the high H202 region. At constant, R values there is no correlation betw-cen the apparent rate constants and the reactant concentration:t. l o r example, when R = 5, the H202concentrations for the runs hhown from top to bottom 111 Vig. 1 \?-ere 3.9, 3.9, 5 0. 2.3, I.!), 5 0, 2.5, 2.5. 3.6, 1.4, 1.9 and 4.5 mniolei per litel. Experimc~nt~ done a t 25' nith R,. the initial value of R, equal to 0.5 showed that decwasirig the HCIOl concmiration from 2 to 0.5 (ionic strength held a t 2 by the use of LiClO, or NnC104) mcieascd the rate by a factor between 2 arid 3. Some rates were measured a t 2.4, 9 8 and 34.2' ti weil its a t 25'. The temperature cocfficieiits of ond-order rate coii+tiits m:ide it ate the over-all activatioii energy fcir the reaction under various conditions. 111 2 11 HCIOa w t t , Bo = 0.57 the activatioii cncrg.rr was found to he 17 kc:tl 'mole and with Xo = 5 the corrcspondnig valuc was 1 6 kea1 mole. 'In 0.5 ;II HC104 with I? = 0 5 , the iictivatiori criergy ma< found t o ti^ 19 kral 'mole. Duc to the knon II watter of the data these activatioii energies arc uneertaiii by a t least 1 kcal.jmole and it is not PO:sible to att:icli m y significance to the differences among the values The catalytic effect of several cations W : L ~invcstigatcd: in 1 Jf HClOd solutions with R,, := 0.5, 1
3
11
EFFECT OF Cu( 11) A V D Co( 11) O N THE RATEOF THE U( 1V)HzO2 REACTION Conditions: 2 j 0 , 2M HClOa. The tabulated values are the fractions of the corresponding uninhibited rates. Inhibitor
R~
CU(I1) Co(I1) CO(I1)
0 5 0 5 2 0
-----Inhibitor 10-4
085
concn -f% 10-3
io-?
10-1
073 83 05
060 56 50
0 33 18
The effect of Ro on the rates becomes smaller as the Co(I1) concentration j s increased; however, in 0.1 X solutions the rate is about 35% greater in solutions with the higher Ro. Small amount? of sulfate did not affect the reaction rate but 1.4 x 10-3 11.1 HC1 in 2 M HC104 iiicreased the rate by a factor of about 1.7. That the reaction is not catalyzed by glass surface w a q shown in an experiment in which the surfare area was tripled by the use of a coiled Pyrex rod placed in the reaction vessel. For these runs the reaction was followed by removing aliquots a t defiute times and quenching them 111 excess Ce-
(IT-). Photochemical effects are unimportant since the rate was not increased significantly when the light from a 100 W tungsten lamp was focussed onto the reaction vessel. The possibility of the formation of peroxide cbomplexes has been considered. Although U(V1) does not form peroxide complexes in acid solution6 the existence of Pu(IV)-peroxide complexes' makes U(T\')-peroxide complexes plausible A qtudy of such complexes is difficult because of :he rapidity of reaction 1. Preliminary experiment5 were made in 0.5 ;I1 IIC'104 at 2.4' 111which absortxuire. ohwved at 647S -1 xere extrapolated t o the t m e ot mixir~g At a con\fant U(IV) concentration of 2.6 >( 10 - I -If the extrapolated absorbance.s were found t o be L: nearly linear function of the H,02 roncent ration. A HzO? concentration of 4 G X 10 B I gax(' an extrapolated absorbance n hi(,h IT as 95' of its value in the absence of H a 0 2 This iiidiwtes that some complexing occurs but the effect i.s tim mlnll to make even LL qualitatix e eitiniatc I t I t s t'xteiit Discussion I n their study of the reaction hctwec-ri L(1Y) and dissolved oxygen, Halpern and Sniithj found that the reaction proceeds i n two stages. the firht a chain process leading to the net reaction
+ + 2H10 = UO?-' + H201 + 2H+ ( 2 ) 0 2
-___I
(6) J Corpel, BzlZZ. BOC chzm. Francs, 752 (1953). (7) R. E Conniek and W. H &IcVey, J Ani Chem %e., 71, 1334 (1949).
NOTES
Oct., 1961
1899
TETRACHLOROPHTHALIC ANHYDRIDEand the second reaction I , which they found to be instantaneous. The chain carriers which were AZAHYDROCARBON COMPLEXES proposed for reaction 2 are U(V) and HOa. BYMIHIRCHOWDHURX~ Our experiments indicate that reaction 1 also is a & Technology, 98, Upper Circdar Rand chain process and that its rate is conveniently Uniaersitv College of ScienceCalcutta-9, India measurable.* The discrepancy in the reported Received March 81, 1961 rates of (1) is probably due to the fact that Halpern It was reported in a previous pap(+ that there and Smith used higher reactant and lower hydrogen was close correspondence between Anla, of T.C.P.A. ion concentration. The principd evidence that reaction 1 proceeds complexes of hydrocarbons and those of cwrrespondby a chain is the fact that it is strongly inhibited by ing azahydrocarbons. This suggest,ad that the Co(I1) and Cu(I1). This conclusion is supported elect.ron involved in the spectral jump came from by the lack of reproducibility which was encoun- the ?r-orbital of the azahydrocarbon. This does not tered in the rate measurements. It is interesting however necessarily mean tha,t t’he stability of the to note that the ions iMn(I1) and Co(I1) which had complex is also due to 7r-r interaction alone. In no effect on reaction 25 were found in the present order to see whether the stabilities also are close to work to catalyze and inhibit reaction 1, respec- each other, we have determined the equilibrium tively. Also the ions of Fe(II), Ag(I), Cu(1I) and constants of T.C.P.A. complexes of three azahydroC1- show the opposite effect in the two reactions. carbons by following a modified Benesi-Hildebrand These results imply that a t least one of the radicals relation. The data are summarized in Table I involved in reaction 1 is different from those in- and the plots are given in Fig. 1. It has not been possible to extend the study to other am-compounds volved in reaction 2 . It is plausible that the chain cariiers in reaction because of incipient precipitation. 1 are U(V) and HO and that in the absence of TABLE I catalysts or inhibitors the most important reaction< EQUILIBRIUM CONSTANTS O F CHARGE-TRANSFER COMPLEXES Inverse of a re Concn. of concn.
+ Hz02 = U(V) + HO TI(’$) + HZ02 U(V1) + HO U( [V) + HO = U(V) + H20 TJ( V) + HO = U(V1) + H20 U(1rV)
(3) (4)
Donor
Quinoline
+ = HOt + H20 C(1V) + HO? = U(V) + HzOz U ( \ ) + HOz C(V1) + H2Ot IIO + HO? Hz0 + 02 H202
+ 0,
U(V1) $- HOz
3.670
9.711 11.653 14.567 16.648 19.423
355
10.502 12.602 15,753 18.004 21.004
355
21.400 29,270 36.590 48.780 58.550
875
(6)
(7)
(8) (9) (10)
Since Co(I1) does not affect HOz radicals, their postulated existence in solutions with large R values is consistent with the observation that small concentrations of Co(I1) have very small inhibiting effectsin such solutions. The effect of oxygen may derive from the reaction I;( V )
of donor
(5)
The mechanism given by reactions 3-6 does not show all that happens since it fails to account for a-Methyldeviations from a second-order rate law, the slight quinoline lack of stoichiometry, and the increase in the apparent second-order rate constant a t high H20? concentrations. These complirations might tw explained by the additional reactions i,S-BenzoH’O
T.C.P.A.. moles/l. X 10’
(11)
Acknowledgment.-The authors wish to thank Helen D. Cowan for technical assistance in the rate measurements. They also acknowledge many helpful discusmms with Prof. Henry Taube, Dr. C. E. Holley, tJr., and especially with Dr. J. F. Lemons under whose general direction this work was done. (8) A recent Russian report describes some preliminary experiments made in IliSO* solutions: E. A. Kanevsky a n d L: At Federova, Radiokhimiya, 2, 559 (1960).
quinoline
3.17i
2.819
(~/CD), I./mole
Wavr length, mF
-~
J?
K,
- C1 mole-’ X 102 1.
1.01,5 1 080
26
1.152 1.211 1 .‘ai5
If.79U
15
,887 978 I .ocit? 1.13: 0.759 .865 975 1.182 1.3iG
18
The equilibrium constants of azahydrocarbon complexes are somewhat higher than those of corresponding hydrocarbon This indicates that in addition to X--R interaction there occurs some specific interaction nith nitrogen non-bonding electrons. This also explains why thr stabilities of complexes of methylquinolirie and 7,8-benzoquinoline are less than that of quinoline in spite of thc higher ?r-ionizat.ionpotential of quinoline. Substituents in quinoline possibly offer hindrance t’othe localized interaction, thereby decreasing t,he stability. I t should be noted that the equilibrium constants of T.C.P.A.-hydrocarhoii and T.C.P.X.-azahydroc a r h i complexes are doser to one another than the rase of 12-hydrocarbon3 and T2-azahydrocarl)on coniplexes.* In the case of 12-complexesof aza-compounds, where n-electrons are supposed t o be iir(1) Whitmore Chemical Laboratory, Pennsylvania State University, University Park, Pennsylvania. ( 2 ) M. Chowdhury and S.Basu, Traw. Faraday SOC.,6 6 , 335 (19tin). (3) R. Bhattacharya and S. Basu. ibid.. 64, 1286 (1938). ( 4 ) .J. Nag Chaudhuri and S. Basu, ibid., 6 6 , 898 (1D60).
1900
- 1
L 1-
-
Vol. 65
hTOTES
30
-31
1
I
-
0
-L
10
20
30
en-'. E!g. 1.
rolved in bondiiiq, the equilibrium constant is of different orrlt r from that of IL-hydrocarbon complexes. 111thc present case the aza-complexes and hydrocarbon complexes approach each other in stability more closely. It seems that both n-7r and T - A interaction play their part in the formation of aza-complexes, the contribution of A-ir interaction being more i r ~the present case than for 11complexes. For the determination of equilibrium constants, we have used tbe modified Benesi-Hildebrand relation, i.e.
have been determined at 2 5 O on several occasions, there is very little information avltilable about the enthalpy changes of these reactions. Gallagher and King' recently have reported enthalpy results determined calorimetrically for the case with chloride ions. We present here similar results for the reactions with chloride, bromide and iodide ions, all at 2 5 O . The results were obtained by measuring the heat effects produced by the dilution of small amounts of potassium halide solution with large volumes of acidified mercuric perchlorate solutions. The details of the method have been described previously.a The ionic strengths of the various final solutions were between 0.10 and 0.14. In al! cases the concentration of the species HgXf was between twenty and a hundred times greater than that of the species HgX2, and the concentrations of the higher complexes were negligibly small.
extinction coefficients of complex and ncreptor sut)strate, reap. 0ptic:tl density oi niisture (A) E = - _,; . OpKl\!:tl path leiigth (1) X added initin1 concn. of arccptrt: (C,O) CD= conreiitration of donor
Experimental Apparatus -The calorimeter has been described already .* I n each elperiment 10 ml. of potassium halide solution was diluted with 990 ml. of mercuric perchlorate solution The calorimeter was calibrated again on two separate occasions during the course of this work by measuring the heats of the one-hundredfold dilution of two different sodium chloride solutions. The average values obtained for these heats of dilution a t 25" (with the corresponding literaturea values shown in brackets) were 302 (297) and 316 (326) cat. mole-' of sodium chloride for solutions with initial concentrations 2 50 and 2.73 molal. The calorimeter was immersed in a thermostat a t 2 5 0.01'. Materials.-A.R. Chemicals were used throughout the uork. All stock solutions were rrade up by weight. -4 stork solution of mercuric perchlorate was prepared by dissolving a known weight of mercuric oxide in an excess of 60% perchloric acid. The pH of the solution wa8 measured and the concentration of mercuric ion wss checked by analysis. The concentrations of mercuric ion in the solutions used for the dilution of the potassium halide solutions ranged from 0.032 to 0 050 molal.
The plot of ] / ( I < - &) zzgaiiist 1 j C ~givos x straight line, the intercept of which on the I / ( E - E,) axis gives 1/(& - E,) and slope of which gives l / K ( E , - E=). Recently Xash5 criticized the Benesi-Hildebrand relation on the ground that an auxlliary quantit,y, the molar absorptivity of the comp!es, need h e determined before the object of primary iritercst, the equilibrium mnst.nnt, can be determincd. We have avoided this diiEculty by dcterrnining .K directly from tlie intercept on thc I/CD nsis. Xash's eyuation5 can, however, be derived very easily by simple algcbrai r transformatioil of the nbove e q w t ion. T1larlk.c arc d u c to Professor S. Basu for helpful cliscassioiis.
Results The results of the dilution experiments are \hewn in Table I. The concentrations of the ions in each solution IT ere calculated from the total concentrations of mercuric and halide ions gwen in Tablr 1 and the nworintion constants of the cornplex ions The assoriation constants used were those drtermincd by Marcus4 at 2 5 O and ionic strrngth 0 50, and nere roorrccted to the particular ionic strength\ of the yoliitions by the equation of Davies.5 Thc rrlntion betwcwi the heat of associntion of reaction 1 at zero ionic strength and at any other ionic strength for which the association constant is K' is g v c n I n 7 the formula6
whcre E,, E,
-
-
(5) C. P. \id!, J . Pitun. Cham., 64, 9 3 (1960). I -
The calculrited values of A H 0 for the thrce nrFociiI(1) P. K. Gsllagher a n d E. L. King. J. A m . Cham. S o c . . 82, 3510 (1960).
J., P h w (2) n. W. Anderson, G . N. Malcolm and 1%. P;. P a r t ~ ~ n Chrm., 64, 494 (1960). (3) "Selected Values of Chemical Thermodynnrnic I'ropcrtirs." Piatl. Bur. Standarch Circ. No. 500, 1952. :4) Y. Marcus, Acto Chem. SronJ., 11, 599 (195;) ( 5 ) C. u'. Davies, J. Chem. Soc., 2093 (1'338\. ( 6 ) J. ?VI. Austin, R. A . &%:heson a n d H. N. Parton, "l'i~e Structure of E!ectrolytio Solutions," edited by \$'. J. IIainer, J i ~ l i uWiley nnd Sonu, New York, N. Y.,1059.
NOTES
Oct., 1961 TABLE I THEHEAT(Q)OF THE REACTION diluted with KX(m) --+KX(O.01 m) Hg(ClO&(m') 711
1 .oo
0.50
0.25
Ionio
Moles of IC1 x 1 0 %
x 10.11 10.00 9.97 4.74 4.72 4 69 2.48 2 48 2.48
=
.047
.046
=
10.88 10.79 10.80 4.74 4.73 4.75 2.28 2.29 2.25
0.046
1 00
10.90 11.10 11.0
0.042
0.50
5.41
.050
0 23
5.28 5.42 2.70
.050
0.50
I),
25
atrength of final Boln.
c1
0.047
X 1 .OO
Q,cal.
m'
-57.07 -57.07 -57.09 -29.3 -29.2 -28.1 -113.4 -13.0 -13.0
0.13
-96.6 -95.9 -96.0 -48.5 -48.1 -48.4 -25.0 -24.6 -24.0
0.13
-184.7 -181.1 177.8 -89.7 -89.8 -92.5 -49.2 -52.3 -55.1 -10.7 -18.4 -1!1.2
0.12
.13
.13
Br
.032
.032
.09
.09
x-I
0.10
-
3 i3 L' 74 0 04 0 (32
.0'12
0.92
.I5
.15
.I2
TABLE I1
FENCTIONS FOR THE REACTIONS Hg+? 4-X - = HgX+ at 25'
'rlIICRMODYNANIIC
CI Br ;
.3G'.
kcal. r n o l t - ~ ~ 1
-10.0 -13.3 -18.4
AHQ
AS0 expt..
kcal. m6le-1
e.u. mole-1
-- 4 . 8 f 0 . 5 -10.6 f .5 -17.6 f . 5
17 f 3 9 i3 3 5 3
whereas their values are required at ionic strengths 0.1 and zero. In the absence of more reliable information the Davies equation was used to extrapo!ate the association constant values from ionic strength 0.5 to zero. The reliability of this equation a t ionic strengths greater than 0.1 is not known. Nevertheless it is of interest that Vanderzee and Dawson7 have used an equation very similar to the Dsvies equation to fit association constant values for CdCl+ measured a t ionic strengths 3, 2, 1 and 0.5, and obtained an extrapolated value a t zero ionic strength in good agreement with an independent determination hy Harned and Fitzgerald.8 In all the final solutions the halide ion is present, predominantly as HgX+ so that, the values of A H do not depend to any great extent on the values of the association constants. The correction of the enthalpy values to zero ionic strength involves the use of the Davies equation. The magnitude of this correction is never more than 0.4 kcal. mole-'. It is probable therefore that the uncertainty in the standard enthalpy values is within f 0.5 kcal. mole-'. The standard entropies of association will be in error by =k 2 e.u. mole-' from this cause. If the uncertainty in the free energy values arising from the use of the Davies equation i s taken into account, the error in the stand;trd entropy figures rises to as much as f 3 e.u. mole-'. Gallagher and King' obtained a value for thc enthalpy of association of HgC1+ of -5.9 kcal. mole-l at ionic strength 0.5. Correction of t h k figure to zero ionic strength by means of equation 2 gives a value of -5.4 kcal, mole-'. The agrc.tbment between the two result,s for AH*(IlgCI+) i: not close, but is within the combined experimental rincertaint ies. Entropies of Association.-C'omup:~risoii ( J f t,hc results in T&le 11 with those for the correspontliri:: reactions with cadmium(I1) and lead(I1) ions9 rcve:ile that the entropy c,hangcs for the association of eac,h cation with a given ha!itle :inion arc: f,li(. same within t.he cxperiment,nl responding cu t,halpics of assocint i c ~ i i :we markedly different. ]:or exnmp!e, for t,he rcnctions invol..iiig the formition of HgBr+, Cdl3r-+:ind PbRri, the entropies of association are all ai)o\it 9 e.u. moic--l but the enthalpies of associ:tt ion are -. 10.6, --0.3 aid f0.3 kcal. mole-l, rcspcctivcly. Thew results arc consistent with t,he that the species HgX+, CdX+ and PbX+ are cornplcx ions with direct contact bctwcen the cation and the mion, RO that their formation involvcs R i.h:lriFP in hydi~3tion structure of the ions. For ~1 qi;-i.n anion thi: strengths of the various cation-anion horids will influence the enthalpies of the association re:i,ctions but. necd have iittle effect on thc mtropica of association if these are Inrgely detwmmcd by the change of hydration of t,he inns CJR association. George'O has shown hhat t,he ent,rripicx 14 a aumher of cation-anion association rescf-iorls C R T ~tic r u m st.1itt.d by the equ at'ion (1
t,ion reactions i rivolving the inercury(I1)-halide c+omplexions are shown in Table IT. The st,and:u.d free energice of association ealculatcd from the msociation constants a t zero ionic strength, and the standa'rd entropies of association calculated from t'lw free energies and enthalpies are shcwn in t,hc same table.
S
1901
A S 0 calc.. e.u. mole-'
17 11 4
Discussion Errors.--Th;is method of determining enthalpies of association of complex ions depends for its ,-:i.iccess on a knowledge of the corresponding association constants at, the appropriat,e i o r k ,strengths. In the prwent case the msociatior! t-rmstants w e known only at ionic strength 0.5
v i m 7
( i ) C. E. Yaridersec and €I. J. D s w e a u , J . Am. Chsm. Sot 7 6 . 5659 11953). (8) 11. E. Hsrned and M. E. Fitrperdd. i b d , 68, 2624 (j'W3. (9) G. 11. Nancullaa, QUUT~. Reca. (London:. 1 4 , 402 (1WC). (IO) J. D. U. George.. J . Am. Chem. s'cc.. 6 2 . 5530 (1959'
NOTES
1902 ilS"
:=
A S (hydration of anion)
+ constant
(3)
The entropy results for the mercury(I1)-, cadmium(11)- and lead(I1)-halide systems conform to this pattern with constants -9, -10 and -9, respectively. The implication of these results according to George is that these association reactions involve the loss of water of hydration. An equation for the entropy of association has been derived by Matheson6 based on an entropy cycle involving the replacement of one water molecule in the hydration shell of the cation by the anion. The form of the equation for water a t 25' is
where Z1 m d Z 2 are the charges on the central cation and the complex ion, respectively, and T + is the crystal ___ __ radius of the cation in hgstroms. SOH~Oand SOX- are the standard partial molar entropies of the water molecule and t,he anion in aqueous solu.tion.ll When the value pf r+ for equathe mercury(I1) cation is taken as 1.10 tion 4 leads to the entropy values shown in the fifth column of Table 11. Alt,hough the equation involves several approximations there is good agreement between calculat'ed and observed values. Siniilar agreement was found when the entropies of association calculated by equa.tion 4 were compared with the observed values for the I : 1 coniplex .halides of cadmium and lead.6 (11) R. E. Powell and W. M. Latimer, J . Chem. Phys.. 19, 1139
(1951).
(12) It. -4. Robinson and R. 11. Stokes, "Electrolyte Solutions," Butterworths, London. 1959.
PROTON MAGNETIC RESONANCE OF GLYCYLGrLYCINATE, GLYCINEAMIDE, AND T H E I R METAL COMPLEXES BY
NORMAN
c. LI,
I & h Y JOHNSON AND JAMES SHOOLERY
Varian A8socaates, l'nlo Alto, Calif.,and Duqzlesne Universitz, Pitlsburgh. P a . aecewed April IS, I961
Several papers' , z have appeared recently dealing with proton resonance studies of dipeptides. Takeda and Jardetzky' found that in glyoylglycine there are two peaks from the two CH2 groups, indicating tha,t they are non-equivalent. It, would be of interest t'o cont~inuesuch studies of dipeptides and determine the effect of chelating metal ions 011 the chemical shifts of these two CH2 groups. A Varian Associates Model A-60 NMR spectrometer which operates a t 60 Mc. was used. Solutions containing 1 .Z M concentrations of each coiistitueiit in 9Y.Sr)& TIzO were placed iii 3 mm. i d . , 4 rnm. o.d. precision Pyrex tubes and the tubes were filled to a depth of about 40 mm. Pure tetramethylsilane in a 5 nun. 0.d. annular cell was used as an external reference compound. The accuracy in measuring the peak positions is estimated to be =tO.Ol p.p.m. The resonance signal of the protons being measured occurs a t lower field than that of the reference, that is, H < H,, and the ( 1 ) M. Takeda and 0. Jardetzky, J . Chem. Phya., 96, 1346 (1957). (2) F. A . Borey and G. V. D. Tiers, J . Am. Chem. Soe., 81, 2870 (1959).
Vol. 65
dowiifield chemical shift is taken to be positive, in accord with the definition 6
=
1O6(H, - H ) / H ,
where Hr and H are the resonant fields for the reference substance and the sample, respectively. PROTON
TABLE I DUE TO METALIONS
CHEMICAL SHIFTS
Volume susceptibility,
THE
PRESENCE OF
Sample
x6
CHz chemical shifts, p.p.m.. (values in parentheses corrected to xv for GG-)
Sodium glycylglycinate, GGZnClz GGhfgClz GGGlycineamide, GA ZnClz GA
0.756 .785 .768 .749 ,766
4 . 2 5 ( 4 . 2 5 ) 3.84(3.84) 4 , 4 2 ( 4 . 3 6 ) 4.20(4.14) 4.35(4.33) 4.00(3.98) 3.80(3.82) 4.20(4.18)
+ + +
Since an external reference is used, bulk-diamagnetic susceptibility corrections must be applied.3 For a cylinder of length large compared with the radius, the equation applicable here is &om.
=
Bobsd.
+ 2 ~ / 3 ( x v-
Xv.r)
where xv and x ~ are , ~the volume magnet.ic susceptibilities of the sample and reference, respectively. The volume susceptibilit'y of each of the solutions was determined immediately after the n.m.r. spectrum was taken, using the same spectrometer, following the method of Reilly, et u Z . , ~ based on the use of the non-rotating concentric The sample whose susceptibility was to be determined was placed in the inner cylindrical space and a substance with a single sharp line (in this case tetramethylsilane) was placed in the annulus. The resonance from the material in the annulus displays two maxima whose separat,ion ( n c.P.s.) is a linear function of the volume susceptibility of the liquid in the inner cell. The values of n for three pure liquids: chloroform, carbon tetrachloride, methanol (at room temperature, -xv = 0.735,6 0.684,3 0.530,' respectively) were determined and a linear plot of n vs. xv was obtained for the three liquids. From this graph, and from the values of n obtained for the solutions, the magi l o t i r susceptibilities of the latter were obtained. Tiiese are included in Table I. The use of an 1i.m.r. spectromet'er for determining magnetic susceptibilities consumes much less time than the use of the c*onventional Gouy method, and is recomnicnded for routine susc:ept,ibilitymeasurements. In Table I, the two lines in GG- are due t o tho two iioii-equivalent CH2 groups and the dowiifield chemical shifts are increased by 0.11 and 0.:30 p.p.ni., respectively, in the prcscnce of 1.2 M ZnCln. When ZnClz is added to a solut,ion of sodium glycylglycinatc iii D20, a chelate is formed, and there is either a reduction of electron deiisit'y (3) J. A. Pople, W. G . Schneider and H. J. Bernstein. "High Resolution Nuclear Magnetic Resonance," McGraw-Hill Bonk Co., New York, N. Y., 1959. (4) C. R. Reilly, H. M. McConnell and R. C. Meisenheimer, Phys. Res., 98, 264A (1955). ( 5 ) Rilmad Glass Company. (6) A. A. Rothner-By and R. E. Glick, J. Chem. f'hy8., 56, 1652 (10.57).
( 7 ) S . Broersina, ibid., 17, 873 (1949).
NOTES
Oct., 1961
1903
on the CH2 m-hich would reduce the shielding, or electric dipole moments of nitrate esters. Cowley the presence of positively charged zinc ion distorts and Partingto~i~ and de Kreuk3 measured the the bonding electron orbitals about the CH:! pro- moments of several nitrate est.ers, but, apparently, tons and generates a paramagnetic shift which is there is no published work on an extended series of in the direction of decreasing field. Li, et u Z . , ~ dinitrate esters. Measurements were made, therehave shown that the coordination sites in GG- fore, on benzene solutions of several dinitrate esters. toward zinc are the terminal amino group and the Experimental peptide oxygen in the immediately adjacent amide Materials.-Ethyl nitrate (East,man Kodak) was purified group. Since in GG- one of the CH2 groups is \)y ,distillation through a 20-cm. column packed with glass situated between t,he two coordination sites, one helices. All other esters were obtained from the research of this Station. 2,2-Bis-(nitroxymethpl)-3might, expect unequal shifts upon formation of a laboratories roxy-1-propanol was purified by chromatography on chelate and this corresponds to the observations. nit silica gel. The other nitrate esters were purified by distilFrom Table I it also is seen that when MgC1, 1:ttion a t reduced pressure. All of the nitrate esters wcrc is added to GG-- the corrected downfield chemical stored over silica gel in stoppered flasks. Solutions \\-ere shifts produced are 0.08 and 0.14 p.p.m. The .hIg prepared with thiophene-free benzene4 which was dried with hydride.6 chelate of glycylglpcinate is less stable than the Zn calcium Apparatus and Procedure.-E1ect)ric nioments were' chelate (log kl = 3.4 and 1.3,yrespectively, for Zn dctermined from benzene solutions a t 25". Dielectric coiland Mg chelates with GG-). The weaker binding stants were measured by a heterodyne beat, method with :$ ed in the smaller chemical shift crystal controlled oscillator similar in design to the apparatus by Hudson and Hobbs.0 The cell consisted of frequencies when MgC1, replaces tlescrihed t wn concentric gold-plated brass tubcs to which glass tuhw %nC12. were soldered at either end. The spare between the tiilws It is seen front Table I that in going from GG- W M provided with inlet and outlet, tubes which pcrmittcd to ZriGG and t'o MgGG, the effect on the chemical connection t o increased or diminished pressure. A calibrat c:d precision condenscr, General Radio Type 72 shift for the second CH, is greater than for the was used in parallel with the cell and the low capaci first CH,. From the above discussion regarding rangc of 25 to 100 ppf. was employed. Purified benzen binding sit,es jn M - toward the metal ions, we IINYI to ca,librate t,he cell. Temperature was controlled by come to the conclusion that of the two CH, fre- niiuns of a tiiermostal; jacket on the outside of the rtsil. were measured with a modified Ostwald pycquencies in GG-, t8hesignal at 6 = 4.25 comes from Dimities l~orneter'and refractive indices with an Abbe ~efrnctonittrr. the CH2 adjacent to t,he carboxylate group, while XIeaeurements were made on a srries of five or six soliiThe signal at 6 = 3.84 comes from the C,H2 whic:h is lions ranging from 0.01 to 0.10 $1. Dicklectrir consta.nts and specific volumes --ere plotted against nGght, fraction of adjacent to the amino group. and solute values for these quantitics wcrc ot~taineil In glycjncarnilde there is one CH2 group, hrnct. soliite bl. the method of least squares. Molar rofract,ioxi,;. Rn,wcw only one (31, frequency. The change in 6, nft>er calculated from the solutio11 refractive index data. Solute, susceptibility correction is applied, amounts t o 0.36 po1:iriaations at, infinite dilution, P a ,were cnlculat rd from and (lipoic, p"p.ni. when Z n U 2 is added. This is in line with tht. Halverstadt arid Kumler monicnts from tjhe Debye equation I he iinding that Z i i chclates stroiigly to glyc*~iiep = 0.01281 x 10-16 ((Pm- K 2 ) T ) l h (1) (log k , 3.3'). The prot'ori niagn4ic resonance shift ineasiir~- w h w e 1' is the absolute temperature. The c:tlculattd d i t t :i meats alone (lo not provide uneyui\rocal evidcnc~: arc presented in Table I. The probable error in the niolxr is 1k0.03 and the probable error in the tlipnlr: for raompltrx formation. The n.ri1.r. spectra do, polarizations rnoinents is 3~0.03D . home\w, support t,he evidence derived from ot hrr Discussion and Results :vpcs of nirasurrtments, for instance, pI1 titrations. I Iigh resol~it~ion t i . n r . spectroscopy therefore repThe electric moment obtaimd for ethyl nitratr resciit)s an additional approach to t,he study of com- ( p = 2.96 D)is in agreement ait'h the value fouud plex forrnzttim, analogous to t>hemeasurement, of by Cowley and Part,ingtonZ ( p = 2.91 D ) for flit. .AWHEN(:E A N D
-4, J. hIA'rL-SZhO
i) .i. S. Erown. P. M. Levin and I?.
,t,:Ln,=,ln, .I, C i e , n
tcren,c:, imrj Deoeiop:nsnt Department, U . S. Naval I'ropslla?it ; ' i a x l , Indiwz Head, Maryland Heceitdd . J p r t l 15, l Y 6 l
ii! c:oniievtioir 1v:f.h .studies on plasti(hizatrio:ia.:itl t ic-izing agsciitx, iiitwcs1 was devclopcd i i i t ' r:
2988
r lR*2),
( 9 , ,i,R. PartinRtim,
'.ionr:ttes. \Vl!en, inst,ead of cupric nitrate, cupric alkanoa tcs n-ere introduced, considerable vhangw in oxidntioii rate were noted. Figure 1 siiows the relative rnt.t.s of oxidation islope of r&:' mtcwt, plot.: wit11 per r.&tive to t.hosr wit,hout (.oypcr.'i. A s t l l t ) o i l i~uniherof the fat,ty acid iiic.rr*:ises,.A) t,hc oxidation rate, in thc case of ethyl ii i i o l w t,r, i. t ~ i i 1 i n n c . d . Rough measarement,s of ilic p::rt it.icsn ,diic.itwts of the cupric tilkanoat,es t ) f ~ t u - f T riiw t l - p ( J h v solvents and water shon-ed that a; st ~ a d yincrcxsr along the same sfries 13y cwiit,rast.,thr tht,;i for squdent. show a progr~suiv(>r.t.cliiviion i n osidat,ion rate. l'h t,fftict- of F'ig. 1 rcflcet, tlhc incrt&iig solii-
.___
llfl? EFlW "1 OF (>OPPEK.BLTi~~NOxrEH ON OK ID IZING OLEFIN 5 !%Y
P. W. ALLES
The .Ynlurol Ruh5t.r ,Producers' Research Associa!,on, H'a!iiyii Ourde!! Crty, Hprfs., England
K L L F I L ACp~r i l I F , 1961
, ~ l I.. ) G e t i r n r ; r t i i l :A. 1
ilc,hcrtsoit. Tratt.~. Faradiiy S o c . , 42, 217
ll,blti).
( 2 ) 11. S. 1 3 L n ~ i . ; ~ . !I . , A n t . Chcnr. Snc.. 82, 2014 ( l Y l i 0 ) . , : 3 i .A. .I.
1907
0.05 0.1 0.15 0.2 Oxygen absorbed, moles 1.-1. Fig. 3.-Rate of oxidation of squalene in chlorobenzene solution a t 55", catalyzed by cupric heptanoate a t concentrations of 1.9 (A), 4.0 (B) and 42.8 (C) X 10-5 moles/l. Other symbols are as for Fig. 3.
* I
d
10 1. 3
m
z
3
$ a .+
0
O 0
L1
fi
+ 0
E
X
2
r'
i ./
/
, 0.1 0.2 0.3 Oxygen absorbed, moles 1.-'. Fig. 2.-Rate of oxidation of ethyl linoleate in chlorobenzene solution a t 65', catalyzed by cupric heptanoate a t concentrations of 0.7 (A), 4.4 (B), 14.4 (C) and 29.1 ( D ) X 10-6 molles/l. The point of addition of copper is indicated by the arrow and dotted line. The broken line shows the course of the substrate reaction.
bility in the olefin phase of cupric alkanoates as the carbon number increases. While the result for ethyl linoleate is the expected one, the effect noted with squalene is unusual. It was considered to be necessary to establish that this effect was not associated with the presence of an aqueous phase. The second series of experiments, therefore, was conducted in homogeneous solution with chlorobenzene ;is solvent and cupric heptanoate3 as the copper compound. The experiments were arranged so that cupric heptanoate could be added to the oxidizing olefin a t some chosen extent of reaction. Figure 2 shows some data for ethyl linoleate (1.5 molar a t 65'). As the amount of copper added is increased, so the initial abrupt catalysis is enhanced, after which the rate falls off before resuming a course parallel to that of the substrate reaction. With squalene as substrate (Fig. 3) there is an initial (but not very pronounced)
0.05 0.1 Oxygen absorbed, moles I.-*. Fig. 4.-Relative rates of oxidation of gutta percha at 75" (A), squalene a t 55" (B), 2,6-dimethylocta-2,6-diene a t 65" (C) and ethyl linoleate a t 65" (D) in chlorobenzene solution in the presence of cupric heptanoate (9 X lo-' moles/l.).
rate increase followed by a rapid decline to rates below those for the substrate reaction. As the concentration of copper is increased, so the catalysis is suppressed and a retarded reaction sets in. This parallels the data of Fig. 1 and shows that the difference between ethyl linoleate and squalene is not confined to aqueous dispersions. A further series of experiments studied the relative rates of oxidation with and without, a fixed concentration of cupric heptanoate with four olefins as substrates in chlorobenzene (Fig. 4). Temperatures were chosen to give comparable substrate rates. There is a close similarity betlweeii all the curves, whether the reaction is catalyzed (relative rate >1) or retarded. It was further noted that changing the reaction temperature from 5CL70' did not increase or diminish the retardation shown by squalene. The over-all activation energy of the retarded reaction (18 kcal./mole) is dose to that (19.4 kcal./mole) for the substrate reaction.
NOTES Discussion Cupric heptanoate does not exhibit simple catalytic behavior toward the autoxidation of the four olefins examined. Even for the case where there is marked catalysis (ethyl linoleate in Figs. 2 and 4) the shape of the rete/extent curve indicates that there is an accompanying retarding effect. In the case of squalene the balance is reversed and the retardation is predominant. The simplest explanation of a fall-off in rate during a catalyzed reaction is that the concentration of one or more of the reactants (metal compound or olefin hydroperoxide, for example) becomes depleted. In fact, hydroperoxide yields were found to be unaffected by the addition o€ the copper compound. Depletion of the metal compound cannot account for reaction rates which are lower than the substrate reaction rate, unless the depletion reaction produces a retarding species. No plausible mechanism can be offered yet for the effects observed. Examination of the effect of cupric heptanoate on the decomposition rates of simple hydroperoxides has shown that the behavior is complex! It seems likely that nietal compounds which exhibit variable valency may interact with almost any species present during an autoxidation chain reaction, and the effect of adding metal compounds to oxidizing olefins cannot therefore be predicted. (81
C. L.
Vol. 65
Apparently, compound C is soluble in benzene and a photostationary e whbrium is established; however, .in hexane compound is only very sparingly soluble, it precipitates in a stable crystalline form, and the reaction goes almost to completion. On standing in room light the purple solid (C) gradually reverted to a greenish solid, similar in appearance to the starting compound (B). If stored in the dark, the colored form did not change appearance during several months. I t also was possible to recover the highly colored form (m.p. about 100") by irradiating hexane solutions of 1,3,:Jtrimethylindolino 6' nitro-8' methoxybenzopyrylospiran (1n.p. about 155'). Although previous workers',' have postulated a heterolytic photochemical cleavage of a bond to give the ionic "open form" shown &B C2, electron spin resonance studies on the solid colored state (C) reveal a weak to medium absorption which indicates an unpaired electron (g = 2.001f 0.0027 H = 20.4 f 4 gauss). Hence this colored form might be a stable biradiclrl ( C l ) fornied by a homolytic cleavage of a bond or an ionic free radical (C4) similar to that in Wurster's salt. Quantum Yields.-The number of quanta was measurrd by an RCA 935 phototube using the specifications in the RCA Tube Handbook. These checked with the output specifications for the Hanovia Sl00 lamp and with two runs with uranyl oxalate.' A Cary Model 14 spectrophotometer waa used to measure the change from uon-colored (J3) to colored (C) form. A 1 cm. path quartz crIl waa used fur most of the studies.
8
- -
-
-2'
Rl. BeU. lunpublished work.
I
c1
i'tfJ
I'BOTOCIIROhIISM IJY
A. F I N E RONAI,I)A. H E N ~ Y
~ A l i l ,,!l. I I E L I & R ,
L)WIOI1T
or ANI)
Chemistry Dzvrsion. I:. ,T. Vnvnl Ordnnnca Trsl S t a t i o n , C h i n a h k e , Cul.
CH, B
J:eceiced A p r i l ? I . 106'1
The compound investigated was 1,3,3-trimethylindolino-6'-nit,robenzopyrylospiran(B) which has been prepared and studied previously. Photochromism has bcen studied in several spicansZa and some other compoundszb which apparently exist in two forms, one of which is colored. Ultraviolet irradiation wi!l change solutions from almost colorless tto strongly colored. The colored form, so made, will facic to give the thermal equilibrium mixture. General Observations.--When a sature.t.ed, colorlcss, hexane solution of B was irradiatcd the solut,ion rapidly became dark-blue (almost black) and a dark-purple, crystalline solid sepnmted. Since the analyses3 on this latter compound were t,he same as those for the starting compound, and since a benzene soiiition of this purple solid rapidly changed from t>hekllllf? to colorless in the dark, it is Iconcluded that t,he purple solid is the colored state C. In benzene no precipitation of purple solid occurred, although the solution developed the typical blue color; after long irradiation a yellow-orange solid, m.p. 185-187' clec., crystallized. * (1) E. Berman, R. 1'. Fox and F. D. Thomson. J . Am. Chem. Soe.. 81, 5605 (1959). (2) (a) R. Heiligmari-Rim Y . Hirshberg and E. Fisher. J . Chem. Soc.. 156 (1961); tb) It. Hardwick. H. S. Mosher and R. Passailaigue. Trans. Faraday Soc., 66, 44 (1960). (3) Calculated for CIOIIUNZOI: C. 70.78; H. 5.68; N, 8.69. Found: C, 70.39; H. 5.35; K, 8.95. In contrast to the starting compound which melted 175 .180". this compound partially melted about 150D, rosolidificd. changed fr.o~ti purple to pink, and finally melted about 175'. When p11111gediuto a hut bath at 16&-l(i5° this colored form melted eoalyi&e!y chr n reuuiidi6ed.
The major problem was to lind the extinction cocficic t i t for form C at 5400 A., where it has a brotrd peak aa reported in ref. 1 . A peak at 2650 A. due to the form B decreases
as form C increases. This 2650 K. peak could, by high intensity irradiation, be decreased to 49% of its original value. The concentration of colored form must lie betwern 51 and 100% of the prepared molarity of the compound. This gives an extinction coefficient between 4.68 and 2.40 (4) The structure and nature of this product have n o t been determined; however, the analyses correspond approximately to those required for the reaction of two molecules of oxygen with each molecule of the starting compound. Cslcd. for CioHaN~O~f2Ot:C . 59.06; H, 4.70; N,7.2.5. Found: C, 60.37. 60.45; H, 4.60. 4.71: N,6.97, 7.13. (5) Y. Hirshberg, J . Am. Chem. Soc., 78, 2304 (1956). (6) The parent compound (B),88 a solid, showa an exceedingly a c a k absorption. probably heclnuse of contaminatioo by a small atnotint 0: the colored form. (7) W. A. Noyes, Jr., and P. A. h i g h t o n "The Photochemiatry u i Gama," Reinb4d P:ibL Corp.. New York, N. Y.. 1941.
NOTES
Oct., 1961 i< 104 i. mole-' cm.-l. The mean value of 3.54 X IO4 cm.*niole-l was used to get quantum yields. The Hanovia S-I00 lamp was used with Corning 7-51 and 0.52 filters to solate the 3660 A. light. A compound filter F t h single crystal salt slices* wae used to isolate the 3130 A. line. Short exposures were used to produce form C which was measured in the t,pectrophotometer. Thermal fading of form C during exposure and before measurement waa insignificant for this compound. From the light absorbed and the amount of colored form produced the quantum 1 xlrl way calculated.
rvhc,re = absorbance = log I,,/I. -4,and A , = find and initial absorbance t = 3 54 >: I O 4 1. mole-' cm.-' 3 = time in sec. hi = current chrtnge ($amp.) for spiran soln. in cell S = phototube sensitivity 10.9 x 109 /.ramp. sec. einstein-1 v = 3.14 X 10-8 1. 1.11 = corr. for light beam area on cathode
I:
1909
DETERMINATION OF THE IONZATION CONSTANTS OF SOME PHENYLMERCURY
COMPOUNDS BY S. S. PARIKS AND THOMAS R. SWEET Department of Clre?nistrg, The Ohio State Unircrsify. Colunibira IO, Ohio Receised April 80, 1061
For many years the effects of various phenylmercury compounds on plants and animals have been studied. The purpose of this work is to determine the ionization constants for phenylmercuric acetate and phenylmercuric propionate. These constants were obtained by potentiometric titrations of phenylmercuric hydroxide solutions with standard acetic acid and propionic acid solutions. In order to calculate these values, the ionization constant for phenylmercuric hydroxide was needed. This value has been report,ed by Waugh, Walton and LnuwicP and also was redetermined in the present work.
Experimental Reagents. Phenylmercuric Hydroxide .-This was obtained as a chemically pure sample from Metalsalts Corporation. Its purity waa checked by the following analyses: %C, theor., 24.45; exptl., 24.61, 24.67; %€I, theor., 2.05; exptl., 2.21, 2.23; %Hg, theor., 68.07; exptl., T.4BLE O F rrYPICAL RUNS 67.80, 68.07, 68.25. An equivalent weight of 295.66 was Phototube determined by adding an excess of KBr and titrat.ing with current, standard acid.* The theoretical equivalent weight is 204.72. Concn. --PamAbmilliA, Time, De- sorbance Phenylmercuric Hydroxide Solution .--A weighed quanmolarity 1. ~ c . Initial crease change .$ tity of phenylmercuric hydroxide was transferred to a volumetric flask. Carbon dioxide-free nitrogen gas was 0.0 3130 . . . . 10.8 1.0 ... :3130 60.09 11.2 ,03977 5 . 3 0.055 0.15 passed into the flask for one hour to flush out all carbon dioxide gas in the flask. Carbon dioxide-free water was ,3977 3130 100.07 10.65 9 . 2 .140 . I 3 added until the flask was about three quarters full and the 3660 60.08 14.7 5.1 .!I3977 ,045 .13 liquid was brought t o near boiling until all the phenylmer:36iiO 80.12 1 5 . 6 13.4 ,3977 .122 .10 curic hydroxide went into solution. During this time and 56 36 17.3 .3977 Above5400 .I1 also while the solution way cooled to room temperature and 11.3 .073 diluted to the mark with distilled water, nitrogen waa passed in a t a slow rate. The solution was standardized The average quantum yield for the change from form B by adding about 4 g. of potassium bromide to 50 ml. of the to form C of an ethanol solytion of this spiran a t 3130 A. solution. Phenylmercuric bromide precipitated and the free hydroxide was titrat,ed with a standard perchloric acid is 0.15 i 0.07 and a t 3660 A. is 0.12 f 0.06 mole/einstein. The uncertainty results from uncertainty in the extinction solution. Preliminary Studies.--It was apparent that a saturated coefficient. If :I. highly colored solution is formed and allowed to potassium chloride salt bridge should not, be placed directly stand, the color fo.dc>s. At 26.5' the first-order rate con- into the phenylmercuric hydroxide solut,ion since this stant is 7.5 i J..0 X 10-4sec.-1. At 6'' it was found t o would result in the forniation of insoluble phenylmercuric he 4.18 X sei:.-l which gives an act,ivation energy of chloride and a consequent change in pH. For this reason a saturated sodium nitrate salt bridge wm used. In order 23 kcal. When the colored solution is irradiated with visible light, to prevent contact of carbon dioxide with the solution, photofading occur8. The quantum yield for the visible nitrogen gas was continuously passed into the cell compartlight from the S I 0 0 !:imp through a 3-69 filter was 0.10 =t ment. However, when a 0.01 Af phenylmercuric hydroxide solution was titrated with 0.01 M H N O sunder these condi0.06. Sirice visible light P L I ~effect, the fading it tieems possible tions, a small amount of precipitate was observed in the solution and also on the t,ip of the 3alt bridge. This was th:it ultreviolet si:sorh(a,i I Jthe ~ color form might also be effwtivc. If it, is, the photostationary state would have attributed to the formation of phenylmercuric nitrate. perchloric acid solution was uscd with When a 0.01 l e ~ scolorcd form than otherwise. The equation for the a 0.01 M p!ien,ylmercuric hydroxide solution, no precipitate pl1otoststion:wy rate can be written was observed in the solut,ion. A thin film of precipitate C B I S = k[CI &IC could be seen on the tip of the sodium nitrate salt bridge (2) if the salt. bridge was left in contact with the phenylmercuric where hydroxide soiution for severs1 minutes. The error resulting . . from this source was minimized by taking only one pH thc tiiermrtl fading rate constant reading witli each dalt bridge-phenylmercuric hydroxide . !!gilt .Liisoi,!xd psr liter per second solut,ion combinntion. When 0.01 r?l phenylmercuric uttd siihscripts refer to the non-colored (B) and colored (C) hydroxide was titrated with 0.01 AI acetic acid or 0.01 Af iorms. Pb,ptoat,stionary states were measured for 3660 propionic acid, the solution appeared thc ssme as described .inti 31XU -4.. light which gave quantum yields +c, of 0.4 ahove for the 0.01 Af perchloric acid. .:.iicl '3.7 ;!.,(;le,'eiiistein. This increase in +C with energy Procedure.-.$ %O-ml. beaker with a. fii-: hole rub!)?r .pi-,.jiirtntn w i n s reaponahie. For 3130, 3660 and 5400 A. , 78.1 and 59.2 kca!./einstrin. (11 Taken in part from the master's thesis of S. S. Pnrikh jireseqted shows only a small shift to the i n the G r a d u a t e School of The Ohio State Univcraity. 1)eeernbrr. ! 5 0 ' s ~ )that AH must be small.
Values were obtained for 3130 and 3660 A. and for soluuons 0.3977 and 0.03977 millimolar a t 26.5' and typical data are elioivn in the table.
+
1'. D. V a i r g h , €I. Y. W a l t o n and J. A . Laswiok. J . I'hys. Ch-n.,
I): T;illiorn MrBriile of thew !aboratorieci.
NOTES
1910
stopper and a 150-ml. beaker were placed in a water-bath maintained a t 25 f 0.04". The holes in the stopper of the larger beaker were used for a nitrogen gas inlet tube, a buret tip, a gas outlet tube, a Beckman blue bulb glass electrode (KO. 40495), and one arm of a sodium nitrste salt bridge. The other arm of the salt bridge was placed in the 150-ml. beaker This beaker contained a saturated solution of sodium nitrate and a Beckman saturated calomel electrode. Carbon dioxide-free nitrogen was passed into the dry 250-1nl. beaker to displace all carbon dioxide. Nitrogen was passed into the beaker continuously until after the reading was obtained, thus keeping it carbon dioxide-free and stirring the solution. Fifty ml. of phenylmercuric hydroxide solution was transferred to the 250-ml. beaker and a known volume of standard acid solution was added from the buret. After the solution was thoroughly mixed, a saturated sodium nitrate salt bridge was introduced and the pH was determined immediately. The entire procedure was repeated for the next volume of acid. Each time a pH reading was obtained, a new salt bridge and a new phenylmercuric hydroxide solut ion wPre used.
Results Table I shows the data and results for the ionization constant of phenylmercuric hydroxide. The average of the three values closest to the half equivalence point was 1.3 X 10-lo and this is taken as the best value for Kb. I n the table, THCIO,and Z'P~,H~OH indicate, respectively, the total added perchloric acid concentration and the total added phenylmercuric hydroxide conrentration after correction for dilution.
Vol. 65
+
+
The ionization constant K for phenylmercury acetate was calculated from the equations [PhHg+] = TPhHgOH
0.00 5.02 10 00 18.38 25.09 28.24 34.00
7.63 5 04 4 66 4.29 4.04 3.94 3.71
0.998 1.755 2.824 3.512 3.793 4.254
8.610 7.920 6.949 6.328 6.073 5.656
1.42 1.02 1.29 1.33 1.31 1.30
In developing the equation used for calculating results, these equilibria were considered
+
PhHgOH = PhHg+ OHH20 = H + OH-
+
By combining equat,ions for coriservation of species, electrical neutrality and ionization constants, an equation was obt,ained Kw2
Kb =
~
[H+I [H+!Ti,hH,on
+ K W T " ~ i o-r K,[H+l
.-~_________
- KW -
[H+l T H C I O 4-~[H'F
A K , value of 1.008 X was used. Table I1 shows the data and results for the ionization constant for phenylmercuric acetate. The average of t,he three values closest to the half equivalence point is 1.5 x and is taken as the best value. In developing the equations for calculating the results, these eq,uilibria were considered (3) H. S. Harned and B. B. Owen. "The Physical Chemistry of Electrolytic; Bolutionu," Reinliold Publ. Corp.. N e w York, N. Y., 1958, p. 638,
[H+12 KW - THO&+ - - -- + K, K.
-
KW -~ [Hi]
w_-_ [H+] _K_ _
[H'IKb [OAC-] = [H+] [PhHgOAc]
=
K,
KW + [PhHg+] - 7 [H I
TrloAo- [OAc-]
+ 11
[PhHg+][OAc-] K = [ PhHgOAc]
A Ka value of 1.754 x low5for acetic acid4 and a Kb value of 1.31 X for phenylmercuric hydroxide were used in these calculations. TABLE I1 IONIZATION CONSTANT O F PHENYLMERCURIC ACETATE Normality of acetic acid = 0.01008; molarity of phenylmercuric hydroxide soln. = 0.009503; initial vol. of phenylmercuric soln. = 50.00 ml. Volume of acetic acid added
0.00 TABLE I 2.67 IONIZATION CONSTANT FOR PHENYLMERCURIC HYDROXIDE Normality of perchloric acid = 0.01055; molarity of phenylmercuric hydroxide solution = 0.009503; initial vol. of phenylmercuric hydroxide sol. = 50.00 ml.
+ +
PhHgOH = PhHg+ OHOAcPhHgOAc = PhHg+ HOAc = H + OAcH20 = H + OH-
10.74 20.00 24.71 28.02 33.00
PH
7.63 6.25 5.70 5.43 5.29 5.20 5.02
TBOAC
x
101
0.510 1.784 2.881 3.334 3.620 3.780
TPI~HPOH
x
103
K
9.030 7.821 6.788 6.360 6.090 5.399
x
105
0.82 1.57 1.45 1.47 1.48 1.48
Table 111shows the data and results obtained for the ionization constant of phenylmercuric propionate. The average of the three values closest to the half equivalence point is 3.1 X lop5 and is taken as the best value. TABLE I11 IOXIZATION COXSTANT FOR PHENYLMERCURIC P R O P I O X A T E Normalitv of propionic acid = 0.01321 ; niolxrit,y of phenylmercuric hydroxide soln. = 0.009299; initial vol. of phenylmercuric soln. = 50.00 ml. V d u m e of propionic acid added
PH
0.00 5.91 10.31 16.00 17.64 19.60 24.95 29.79 34.43
7.56 5.68 5.47 5.25 5.19 5.12 4.92 4.73 4.55
THOP?
x
108
1.396 2.258 3.202 3.445 3.720 4.397 4.932 5.387
TPhHgOA
x
108
8.316 7.700 7.045 6.874 6.680 6.203 5.827 5.507
K
x
106
3.11 3.12 3.11
3 14 3.11 3.13 3.13 3.00
The equilibria considered and the equations used to calculate the ionization constant for phenylmercuric propionate correspond to those for phenylmercuric acetate. 1.336 X was used as the ionization constant for propionic acid.4 The ionization constants were found t,o be 1.3 (4) Ref. 3. pa 755.
x IO-'" for phenylmercuric hydroxide, 1.5 X 10-5 for phenylmercuric acetate, and 3.1 X for phenylmercuric propionate. Each of these values was found to be constant for experimental points taken 01 er a wide range of (Tacld/TPhHgOH) values. The constant for the hydroxide is further substantiated since it was used in calculating both of the other constants. The value for phenylmercuric hydroxide found in the present work agrees with the value of 1.0 X that has been reported.2 Thwe constants may be used in estimations of equilibrium constants for the interaction of phenylmercury compounds with groups such as the sulfhydryl group in biological materials.
t-
I
SORPTION OF SULFUR HEXAFLUORIDE RY ARTIFICIAL ZEOLITES BY DANIELBERGA N D WILLIAMM. HICKAM V e s t i n g h o u s e Cmtrnl Laboratories, I'ittstiuryh Q 5 , Pmnsylz'ania R e c e i i d .April 86, 1961
The small and uniform size of the intracrystalline pores left, by t,he removal of water from hydrated crystalline zool.ites results in "molecular sieve" act'ion. 31Iolecules of large diameter are excluded from the pores,; smaller diameter molecules may be sorbed in the intracrystalline pores. With mixtures of molecules of diameter smaller than the sieve pore :sizes preferential sorption may take place. Habgood' 'has studied the kinet'ics of Molecular Sieve action for nitrogen-methane mixtures. In this case nitrogen diffuses more rapidly into the small pores and is preferentially taken up in the early stages of sorption but methane has a higher affinity for the sieve and is preferentially sorbed a t equilibrium. In this work we are interested in studying the sorption of a rnolecule of roughly 5 A. diameter, namely, SF,, :is a function of t'emperature and pressure and the selective sorption in mixt'ures of SF6and air. Experimental Procedure A gas circulatclry system wit,h mass spectrometer2 was used in measuring sorption of SF6 by the Molecular Sieve. wveral advantages over other equipment This system ofler~~ used for similar measurements. The calibrated volume and manometer permits accurate determination of the quantity of gas used. The system is attached to the mass spectromtter which providcs analysis of the gas when desired. When mixturw of gases are being used the circulating pump ensures ihiLt the gas abow the sieve is of uniform composition. 'This sondition, eomhined with the mass spectrometric analysis, permits calculation of the quantity of each cornponeizt adsorbed even though there are large teniper3turc gradients in the system due to heating of the sieve. .A weighetl amount of sieve was plxed in the system. T h r sieve vv)~.r ~ \ x t i l t t dand heated to ahout 200" to re. Sulfur hexafluoride wv&sintroduced
r I O it known pressim of i: tn the romvcjir was
SF, :tnd then repeated sev-
: I ) R.W. K>L,goc?, Can. J . Chem., 36, 1384 (1958). (2) L.,;.!
