Oct., 1952
KINETICSOF EUROPIUM(II)-EUROPIUM(III) EXCHANGE REACTION
853
THE KINETICS OF THE EUROPIUM(I1)-EUROPIUM(II1)EXCHANGE
HEACTION la BY DALEJ. MEIER'"AND CLIFFORD S. GARNER l'he Department of Chetnislry, University of,California, Los Angeles 2.4, California Received March 19, 1966
A kinetic study of the isotopic exchange reaction between Eu(I1) and Eu(II1) ions in aqueous hydrochloric acid was carried out, leading to the following rate law in the concentration range 0.026-0.068 f EuC12, 0.035-0.12j EuCln, 0.098-1.9 f total CI-, 1.0f H+, p adjusted to 2.0 with NaC104, 32-50', R = 6.5 X 10" X exp( -20,80O/RT) X [Eu(II)][Eu(III)][C1-] moles liter-' sec.-l. A variation in the hydrogen-ion concentration from 0.1 f to 1.6 f a t constant ionic strength 2.0 is shown to affect the exchange rate by only 13%, suggesting that hydrolyzed species are of little importance in the rate-determining step, which is proposed as the exchange between hydrated E u + + and a hydrated EuCl++ ion pair. Separation of the two euro ium species was achieved by the use of dilute ammonium hydroxide to precipitate Eu(II1) hydroxide. In chloride-free percffloric acid solutions the rate of oxidation of Eu( 11) by perchlorate ions was too rapid to allow the exchange rate to be determined. The fact that the exchange rate extrapolates to zero at zero chloride-ion concentration is taken as evidence that the rate of electron transfer between the uncomplexed europium ions is a very slow reaction.
The authors2 have previously reported that the rate of the isotoQic exchange reaction between europium(I1) and europium(II1) in aqueous hydrochloric acid solution is measurable. Preliminary data were shown which indicated that the reaction is first order each with respect to theconcentration of europium(II), europium(II1) and chloride ions. This paper describes in more detail our studies of the kinetics of this exchange reaction.
Experimental Chemicals.-The europium, lent by Mrs. Ethel Terry McCoy, was of reported 99.9% purity. The purity was confirmed by us by determining the oxalat,e-oxide and oxalate-permanganate ratios.3 Europium(II1) solutions were prepared by dissolving in tBheappropriate acid the oxide, prepared by ignit>ingat 850' europium(II1) oxalate which had been precipitated from a hot 1 f nitric acid solution of europium(II1) nitrate by the slow addition of a saturatcd oxalic acid solution. Europium( 11) solutions were preparcd by dissolving europium( 11) carbonate in the appropriRte acid. The carbonate was prepared by mctat,hcsizing curoium(I1) sulfate with hot 1 f sodium carbonate solution.' Europium( 11) sulfate was made by reduction of curopiuni(111) chloride solutions with a Joncs rcductor, followed by precipitation of the sulfate by passing the reduced solution into hot dilutc sulfuric acid to which a small amount of acetic acid had been added to reduce the solubility of the precipitate.6 The over-all yield for thc conversion of puropium(II1) oxide to europium( 11) earbonat,c was 90-957,. Although europium( 11) solutions arc quickly osidizcd by air, europium(I1) carbonate rcsists oxidation by air if kcpt dry. For example, a sample of the pure carbonate which had been stored in air in a desiccator ovcr calcium chloride for three months was found to be 96% curopium(I1) carbonate upon analysis by the method of McCoy.6 Because of the ease of oxidation of europium(I1) solutions, all solutions which were to contain this ion had to bc completely free of oxygen. Large quantities of cach reagent solution were freed from oxygen by passing a stream of osygen-free nitrogen through cach for a period of a t least 12 hours, then storing each in all-glass closed-buret systems. The water used in this investigation was redistilled from an alkaline prrmanganate solution in an all-glass still. The (1) (a) Abstracted from the Ph.D. thesis of Dale J. Meier, University of California, Los Angeles, June, 1951: (b) University Fellow, 1949-1951. Present address: Shell Development Company, Emeryville, California. (2) D . J. Meier and C. S. Garner, J . A m . Chem. Soc., 7 3 , 1894 (1951). (3) D. W. Pearce, G. L. Barthauer and R. 0 . Russell, "Inorganic Synthasw." Vol. 11, McGraw-Hill Book Co., Ino., New York, N. Y., 1948, p. 58. (4) R. A. Cooley and D . M. Yost, ibid.. p. 71. (5) J. K. Marsh, J . Cham. Soc.. 398, 925 (1942). (6) H.N. McCoy. J . Am. C h m . Sac.. Si,2466 (1989).
perchloric acid was C.P. grade, redistilled at 5 mm. pressure. C.P. hydrochloric acid was redistilled to give t,he constantboiling mixture. Baker C.P. fused sodium chloride, found to be iron free when tested with potassium thiocyanate, was used without further purification. C.P. sodium perchlorate was twice recrystallized from water to remove an iron impurity. The concentration of the stock sodium perchlorate solution was calculated from the solubility data of Cornec and Dickeley.' Oxygen-free ammonium hydroxide solutions were prepared by saturating oxygen-free water with tank ammonia. The nitrogen used in this study was scavenged of oxygen by passage through a continuously regenerated chromium( 11) chloride solution.8 Radioeuropium Tracer .-The small amount of europium available required that, it be re-used in later experiments. Thus it was desimble in prcliminary experiments to use a short-lived europium activity, which would decay out after measurement and thus avoid radioactive contamination of the whole supply. Radioeuropium for the preliminary experiments was produced by Srn(p,zn) reactions in the U.C.L.A. cyclotron, using an internal beam of ea. 16-Mev. protons. Chemical separation of europium from tho samarium oxide targct,s by the procedure of Meinkeg gave europium activitics of thc following half-lives: 53 days, 4.8 days, 20 hours and 8 hours. The latter two activities predominated greatly for the bombardment times used. At the timc of this investigation only bhe 4.8-day activity had not. been reported. Latcr Hoff, Rasmussen and Thompsonlo reported this activity among those they observed after proton bombardmcnt of samarium. Altcr completion of thc preliminary cxperiment,s it was dcsirable to use a Iongcr-livcd act,ivit,y to avoid frequent processing of cyclot,ron targcts and the large decay corrections necessary with t.he short-livcd activities. Through the cobperatidn of Professor Don M. Yost and Dr. David L. Douglas of the California Institute of Technology we obt>aineda sample of 5.2-ycar Eu1b2preparcd a year earlier by irradiation of spect.roscopically pure europium oxide in the Oak Ridgc pile. Radiochcmical expcrimcnts made by us on the aged sample showcd t,hat curopium was the only radioactive element present. The stock solut,ion of Eu162 (ca. 1 inc.) consistcd of 1.5 mg. of europium(II1) chloride in 5 nil. of dilute hydrochloric acid. For each exchange expcrinient ca. 20 pl. of this stock solution was sufficient to give a satisfactory counting ratme,thus requiring no correction for the volumc of solution or curopium added in introducing the activity. Measurement of Radioactivity.-Othcr than in a few preliminary experimcnts in which mounted precipit,ates were counted, all radioassays wcre made on solutions, using dipping counters with a scale-of-64 circuit. The solutions werc diluted to a definite volume, usually 20 ml., in a 30ml. test-tube, and the dipping Geiger tube was accurately adjusted with respect to the test-tube with a rack and pinion (7) E. Cornea and J. Dickeley, BdE. aoc. chim. France, 41, 1017 (1927). ( 8 ) H. W. Stone, J . A m . Chem. SOC.,68, 2591 (1Y36). (9) W. W. Meinke, AECD-2738(1949). (10) R. W. Hoff. J. 0. Ranrnwsen and Sa 0.Thornpaon, Phus. Rev., 88, 1068 (1951).
D. J. MEIERAND C. S. GARNER
854
asscmbly so as to give reproducible geometry. Since the solutions to be counted were of constant chemical composition, absorption and scattering corrections' were unnecessary. Background (ca. 12 c./min.) and small coincidence corrections werc applied. Earh sample was rounted for a time sufficicnt to reduce statistical counting errors to less than 1% standard deviation. Separation Methods.-Unsuccessful methods tried for the separation of europium( 11) from europium(II1) included the use of sulfate ion to precipitate europium( 11) sulfate, and the use of buffered phosphate, oxalate and formate solutions to precipitate one or the other of the reactants. Separation of the two reactants with sulfate ion gave unreproducible (and often very large) zero-time exchange traced to coprecipitation of europium( 111) by the europium(I1) sulfate. The sulfate method was used in preliminary experiments to show that the exchange was slow. Dilute ammonium hydroxide, which precipitates europium(111) hydroxide, was found to give a quantitative separation of the two ions, and gave reproducible amounts of zero-time exchange. Procedure .-Portions of europium( 111) oxide and europium(I1) carbonate were weighed into separate 50-ml. glassstoppered erlenmeyer flasks fitted with nitrogen inlet and outlet tubes. After the flasks were flushed with nitrogen, calculated quantities of reagents were added from the reagent storage systems to each flask to give a volume of 6.00 ml. Tracer europium(II1) chloride (cu.20 pl.) was added to the flask containing the europium(II1). The resulting europium( 11) and europium( 111) solutions were separately equilibrated a t the desired temperature in an oil-bath. Nitrogen was passed through the europium( 11) solution to remove the carbon dioxide resulting from the reaction of the carbonate with the acid, since the presence of carbon dioxide would have interfered with the separation method. After temperature equilibrium was reached 5.00 ml. of each solution was transferred with an automatic pipet to the reaction vessel and mixed quickly by bubbling nitrogen through the solution. The reaction vessel was similar to one described by Lewis.11 It consisted of a 20-ml. spherical reaction chamber fitted with a pipet from which aliquots could be withdrawn without opening the vessel to the air. Nitrogen could be passed through the system. Surrounding the reaction chamber was an opaque jacket through which oil from a thermostated oil-bath 'was pumped to maintain a temperature constant to fO.l The concentrat,ion of europium( 11) in the exchange solution was determined during each exchange run by the method of McCoy.6 A known volume of the exchange solution, withdrawn at ea. the half-time of the reaction, was delivcred into 10 ml. of oxygcn-free 0.1f iron(II1) ammonium sulfate solution. After the addition of a few drops of 85% phosphoric acid to reduce the color of the iron(II1) ions, the iron( 11) resulting from the reduction of iron( 111) by the divalent europium was titratcd with a standard potassium permanganate solution. The europium( 11) concentration was usually ca. 95% of that calculated from the amount of europium( 11) carbonate used to prepare thc exchange solution, the difference resulting from partial oxidation of divalent europium by air during the transfer to the reaction vessel. The concentration of europium(II1) was taken as the difference between the total europium concentration and the europium( 11) concentration. At desired times 2-ml. aliquots were withdrawn from the reaction vessel and delivered in a current of nitrogen into 10 ml. of 1.5 f ammonium hydroxide solution in a 50-ml. centrifuge tube equipped with a ground-glass stopper. After the centrifuge tube was capped the mixture was shaken gently and centrifuged. Tests showed that europium( 111) was quantitatively precipitated under these conditions, with little or no coprecipitation of europium(I1). The supernatant solution was poured off, and a few drops of 6% hydrogen peroxide was added to it to oxidize the divalent europium. The precipitate which formed was centrifuged, washed with water, then dissolved in 2 ml. of 6 f hydrochloric acid and diluted to volume in the test-tube used for the radioassay with the dipping counter. The europium content of each radioassayed fraction was determined as follows by the method of M e i t ~ k e . ~Europium was precipitated by addition of an excess of ammonium
Vol. 56
hydroxide, the prccipitat,c centrifuged, washed, and then dissolved in 10 ml. of 0.6f hydrochloric acid. This solution was heated to boiling and 10 ml. of a saturated oxalic acid solution was added. After cooling on an ice-bath for 10 minutes the prccipitste of europium( 111) oxalate dccahydrate was filtercd onto a tared fritted-glass crucible. The precipitate wm washed in turn with 10 ml. each of water, absolute ethanol, and diethyl ether, then dried in a vacuum desiccator for 10 minutes, and weighed as the decahydrate. With the above experimental conditions it was necessary to follow the exchange by measurement of the change in specific activity of the europium(I1) fraction. A change of specific activity of the europium( 111) fraction can arise from three sources: the exchange reaction, partial oxidation of europium( 11) to europium( 111) during their separation, and partial coprecipitation of europium( 11) in the europium( 111) hydroxide precipitate. The exchange reaction is the only factor causing a change in the europium(I1) specific activity.
Results Exchange in the Absence of Chloride Ion.-Originally wr planned to study thc exchange reaction in aqueous perchloric acid solution to avoid effects arising from com lex ion formation. However, this plan had tb be modii!edwhen it was found that europium(I1) was completely oxidized in perchlorate solution before the exchange had proceeded more than a few per cent. The oxidation was faster than the exchange under all conditions tried, namely, 0.1-2 f perchloric acid, 25-50', absence of light. As shown below, the exchange is catalyzed by chloride ion, and it proved possible to study the exchange as a function of the chlorideion concentration by substituting perchlorate for chloride ion. The rate of oxidation of europium(I1) by perchlorate, relative to the exchange rate, was slow a t chloride-ion concentrations above 0.1 f. By studying the rate as a function of the chloride-ion conccntration it was hoped that the data could be extrapolated to zero chloride-ion concentration, and in this manner the exchange rate for the non-complexing solutions determined. However, the extrapolated rate was essentially zero relative to the rate at 0.1 f chloride ion. Apparent Zero-Time Exchange.-The exponential exchange law12J3J4 applied to the europium( II)+urapium(111) exchange, takes the form
.
(11) W. B. Lewis,Technical Report No. 19, Laboratory for Nirnlamr Minncu and Hnght~urin~, M ~ s n ,T H M ~ , ob ~ a o h , 1, ~ 4 %
where R represents the rate a t which europium( 11)becomes europium(III), and vice versa; this rate is independent of the presence of tracer levels of radioeuropium, and is a constant when all conditions otjhcr than thc isotopic distribution are hcld constant. F is the fraction exchange at time t , and is given by =
- S(1I)O - S(1r)o
S(I1)
S,
(2)
where S(II) and S(II)O are the specific activities of europium(11). at time t and time zero, respectively, and AS, is the equilibrium specific activity. I n this investigation specific activities were arbitrarily expressed in counts per minute per milligram of europium( 111) oxalate decahydrat,e. The equilibrium specific activity was usually determined as the average specific activity of all the europium in the exchange mixture, although it was experimentally determined for europium(I1) in a few exchange runs as a check. Concentrations (bracketed quantities) are expressed in terms of gram atoms of europium per liter of solution at 25'. The expected exponential exchange behavior was exhibited in all runs. There was a reproducible zero-time exchange, which increased from ca. 7 to 20% as the hydrogenion concentration of the exchange solution was decreased from 1.6 to 0.1 f a t an ionic strength of 2.0. The work of Prestwood and Wahl" showed that when separation-in(12) H. A. C. McKay, Nature, 143, 097 (1938). (13) G. Friedlander and J. W. Kennedy, "Introduction to Rsdiochemistry," John Wiley and Eons, Inc., New York. N. Y.,1949, P. 287. (14) A. C. Wahl and N. A. Bonner (editors), "Radioactivity Applied t o Chemistry," John Wiley and Sona, Inc., New York, N. Y., 1951, Cham . 1 by . 0. E. Myers and R. J. Prestwood. (15) R. J. Prestwood and A. C. Wahl, J . A m . Chcm. S o n . , 71, 3137 lQ4Q). Alao nos Anpenrlix X of reference 14.
KINETICSOF EUROPIUM(II)-EUROPIUM(III) EXCHANGE REACTION
Oct., 1952
duced exchange and the degree of separation arc constant the slope of In (1 - F ) versus t is unchanged and the data may be corrected for such effects. Dependence of Exchange Rate on Europium(II) and Europium( 111) Concentrations.-If the exchange rate is first order each in europium(I1) and europium(II1)
R = ki[Eu(II)][Eu(III)] (3) where kl is a specific rate constant when all possible ratedetermining species other than europium are maint,ained at constant concentration. The exchange half-time, t i / * , is then inversely proportional to the total europium concentration 0.693 1 1% = (4) [ E u ( I I ) ] + [Eu(III)]
an excellent confirmation of first-order dependence for each species. Dependence of Exchange Rate on Chloride-Ion Concentration.-The greatly enhanced rate resulting when chloride was substituted for perchlorate suggests that chlorideassociated species wereimportant in determining the exchange rate. Representative data obtained when the chloride-ion concentration was varied over wide limits are presented in Table IT. If we assume that both the uncomplexed and
TABLE I1 ORDERWITH RESPECT TO CHLORIDE ION [H+] = 1.00f, 39.4", p = 2.0 C1- concn., f
Total Eu ooncn.,
Half-time,
0.0982 0.716 1.51 1.86
0.0600 .0677 .1465 .0653
870 132 29 53
TABLE I ORDERWITH RESPECT TO EUROPIUM(II) AND EUROPIUM(III) [€I+] = l.oof, 39.40, p = 2.0 Total Eu concn., f
Eu(I1) concn., f
0.0258 .0244 ,0258 .0683 .0270
0.0603 ,0653 .0894 .lo55 .1465
Eu(II1) concn., f
0.0345 .0409 .0636 .0372 .1195
c1concn., f
1.88 1.86 1.82 1.84 1.51
Half-time. minutes
59 53 40 33 29
Table I presents the data concerning this dependence. The chloride-ion concentration varied slightly, for which effect a correction may be applied by anticipating later data which show the exchange rate to be dependent on the first power of the chloride-ion concentration and to occur through only one exchange path. Equation 4 may then be generalized to 0.693 1 "" = { [Eu(II)] [Eu(III)] [Cl-] %. (5) Figure 1 shows data plotted as suggested by equation 5. The intercept is zero within the experimental error, and the first-order dependence for europium( 11) and for europium(111) assumed in deriving equations 4 and 5 is indicated.
+
+
+
7' 0.24
.A
E& 0.20
3 .+ c
*
h
n n
t?. 50 ci Y
n
40
minutes
phl?ride-complexed europium species can exchange, t,hen R given as R = [Eu(II)][Eu(III)]{ki k[C1-ln) (6) and the half-time is relat,ed to chloride-ion Concentration by the equation 0.693 x 2- = kl k[C1-]" (7) [Eu(lI)l [Eu(III)I tl/r If the left side of equation 7 is plotted against the chlorideion concentration raised to the appropriate power n, a straight line will be obtained of slope k and intercept kl. Figure 2 gives a plot with n = 1. A straight line is obtained over a 19-fold change in chloride-ion Concentration, and with zero intercept within experimental error.
70
-2
f
1s
+
60
855
: _I_ I
I
I
I
1
1
v
1
Y
$
0 0.4 0.8 1.2 1.6 2.0 2.4 Chloride-ion concentration, moles liter. - 1 Fig. 2.-Eff ect of chloride-ion concentration on the exchange rate ([H+] = l.OOf, 39.4", p =: 2.0).
-330
2
G 20 10
0
2
+
4
6
8
1 0 1 2
I/{ [Eu(II)] [Eu(III)]] [Cl-1, mole-' litera. Fig. 1.-Effect of total europium (and chloride) concentration on the half-time ([H+] = l.OOf, [Cl-I = 1.5-1.9f,
39.40, /.I= 2.0).
The usual method of determining dependence of the reaction rate on the concentration of a reactant, L e . , b evaluating b log R / b log [Eu(lI)] and of ? log I R/d log fEu(IIl)l, with the other concentrations held constant, was also carried out as a confirmatory check of equation 5. The values found were 1.02 and 1.03 for the order of the reaction with raspeot to europium( 11) and europium( IIX), tnopectively,
Dependence of Exchange Rate on Hydrogen-Ion Concentration.-Table I11 summarizes the data obtained when the hydrogen-ion concentration was varied with the total europium and chloride-ion concentrations held roughly constant. If the exchange rate is independent of hydrogen-ion concentration over the range of interest, equation 7 is applicable, with kl = 0 and n = 1 as experimentally demonstrated, giving 0.693 1 k (8) 1 [Eu(II)I [Eu(III)I 1 IC]-1 The rate constant k, calculated from equation 8 and given in the last column of Table 111, increases systematically by ca. 13% over the 16-fold increase in hydrogen-ion coficentration. This trend is op osite to the hydrolysis effect observed bg Presttvood and $ahli' and by Hrrbottle end Dod#onlo for the thallium( I)fthalliurn(III) exchange. It seema
+
-~ (16)
d. Yhtbbttla nad R. W
tiQtii1,
D o ~ Rda~ Avh,l i
Ch#ma bqdcii 711144e
856
D. J. MEIERAND C. S. GARNER
reasonable to believe that a primary result of hydrolysis might be to lower the coulombic barrier separating the reacting ions, thus allowing closer average approach with a concomitant increase in the exchange rate. These considerations, coupled with the small effect observed and the known strongly basic properties of europium(II1) and europium(11), suggest that hydro1 sis does not play an important part in the exchange a t d e concentrations used in this inveetigation. We believe that the variation in rate can best be explained by specific electrolyte activity effects resulting when sodium ion is substituted for hydrogen ion in maintaining constant ionic strength.
TABLE I11 ORDERWITH RESPECT TO HYDROQEN ION 39.40, p = 2.0 H+
concn..
f
c1-
Total
Eu concn.,
f
Halftime, minutes
concn.,
/
1000 6 moles-$ d.er2 800. --I
I
\
Vol. 56
Discussion The first-order dependence of the exchange rate on chloride-ion concentration indicstes that the activated state includes a monochloro-complexed europium species. From the kinetic studies it is not possible to conclude which of the two ions forms the complex, but it seems most reasonable to assume that it is europium(I.11). Halo-complexes of trivalent rare earths have been reported by Connick and Mayer"; for the monochloro-complex of cerium(II1) they report an association constant of ca. 2 (that for europium(II1) would be expected to be nearly the same). It seems unlikely that the association constant is this high, for the spectroscopic work of Freed1*and the polarographic work of Hollecklg have indicated no detectible complexion formation between europium(II1) and chloride. Moreover, the absorption spectra of aqueous solutions of eleven rare-earth(II1) perchlorates and chlorides, including europium and cerium, are essentially identicaLZ0 In the present study, the exchange rate is proportional t o the chloride-ion concentra.tion over the 19-fold range of concentration investigated (see Fig. 2); saturation effects which should be observable if the association constant were as great as 2 are not apparent. The apparent zero intercept of Fig. 2 shows that the rate of electron transfer between the uncomplexed ions is very slow compared to the rate of the chloride-catalyzed reaction. The exchange mechanism can be formulated in at least two ways: (1) A small fraction of one of the reactants, probably Eu(III), is complexed with chloride ion, and in the ratedetermining step one electron is transferred between the chloro-complexed species and the other reactant
+ +
Eu*+++ C1Eu*Cl++ Eu*Cl++ E u + + +Eu*CI+ EU+++ Eu*Cl+ +Eu*++ CI-
+
+
1.5 3.0
3.1
3.2
3.3
1000/T. Fig. 3.-Temperature dependence of exchange rate ([H+] = l.OOf, p = 2.0); A, 31.7" 0.0634ftotal Eu, 1.88 f C1-, 124 min. half-time; B, 39.4', average of data from Table I; C, 49.5', 0.0523f total Eu, 1.89fC1-, 22 min. halftime. Catalysis by Light.-In one of the later experiments the opaque covering of the reaction vessel was removed and light from a 300-watt incandescent lam was focused on the vessel. A slight increase in the spec& rate constant was noted, and traced to a n observed increase in temperature (cu. l o )of the solution. Hetero eneous Catalysis.-In one series of experiments glass W O was ~ placed in the solution. The glass wool, which was carefully cleaned with boiling nitric acid before use, was calculated to have increased the surface exposed to the solution by a factor of five. The specific rate constant in the presence of glass wool was found to agree with that observed in the absence of the glass wool to within 1%. An attempt was made to determine the effect of a iece of platinized platinum wire on the exchange rate. &drogen was rapidly liberated a t the wire surface in contact with the exchange solutionA A few minutes lst& europirim(l1) aould not he dotwted in the Rolutinn.
(rapid equilibrium) (rate determining) (rapid follow reaction)
(2) This reaction mechanism is identical to that of 1 except that an electron and a chloride ion are interchanged between reactants in the activated complex, either simultaneously or in two steps
+ +
Eu*+++ ClEu*Cl++ Eu*Cl++ E u + + + Eu*++ EuCI + +
+
(rapid equilibrium) (rate determining)
Comparison of the true energy and entropy of activation for this reaction with that predicted from absolute rate theory for reactions between ions of this charge type is not possible at present as the experimental rate constants and activation energy include the unknown association constant and heat of formation of the chloro-complexed ion. Acknowledgment.-The authors wish to express their sincere appreciation to Mrs. Ethel Terry McCoy for the loan of the europium used in this investigation, and to Professor Don M. Yost and Dr. activity. David L. Douglas for the (17) R. E. Connick and 8. W. Mayer. J . Am. Chcm. Soc., 7% 1176 (1951).
(18) 8.Freed, Rsv. Mod. Phya., 14, 105 (1942). (19) L. Hollsok. Z. Blcktrochsm,, 46, 59 (1940). (10) T.Modlmr and J , 0. Brnnthy, A n d . Cham., 49, 4 W (1950).
Oat,., 10Ei2
KINETICS OF EUROPIUM(II)-EUROPIUM(III) EXCHANGE REACTION DISCUSSION
N. URI (Chicago).-I would like to ask the authors how much is known on ion pair complex formation between Eu3+ and CI-? If complex formation is very weak I think it should be attributed to an accidentally large endothermicity. It appears that in practically all cases ion pair complex formation in aqueous solution is accompanied by a considerable positive entropy change which goes hand in hand with the ositive entropy changes accompanying the desolvation of t f e negative ion. On the other hand, the heat changes result, as can be demonstrated by a thermodynamic cycle, from differences between very large values. The free energy change of ion pair complex formation is thus composed of a regular positive entropy change and an accidental heat change. If the latter happens to be very endothermic, ion pair complex formation would not be observed, at least a t room temperature. c. s. GARNER.-cOmpk!X formation between Eu3+ and CI-is discussed in the aper. Apparently no direct evidence exista for chloro-compgxed species of Eu3 +. Spectroscopic and polarographic studies indicate no detectable complex-ion formation in this system, and no thermodynamic data are available for an: halo-complexes of europium(II1). S. E. VOLTZ(Houdry Process Gorp.).-You have demonstrated that increasing the amount of glass surface does not affect the reaction rate, which indicates that glass does not catalyze this reaction. This result is listed under “Heterogeneous Catalysis” and I would like to know, if you consider this articular result as indicative of the possible behavior of otier heterogeneous catalysts. Would you also comment on the effect observed in the presence of platinized platinum wire? GARNER.-The experiments performed with an increased glass surface were undertaken solely to find out if the glass container and transfer pipets were having an appreciable
857
effect on the rate of the exchange reaction; no such effect was found. Inasmuch as there are certain features common to electron-transfer exchange and electrode reactions, one may assume that there may well be certain substances that act as heterogeneous catalysts for the exchange just as there exist such catalysts for electrode reactions. The rapid liberation of hydrogen at the surface of a platinized platinum wire in contact with the exchange sohtion, and the consequent rapid disappearance of europium(I1) from the solution presumably imply that the platinum catalyzes the oxidation of europium(I1) by water (or hydrogen ion). Inasmuch as the formal reduction potential of the Eu(I1)-Eu(II1) couple is approximately 4-0.43 volt a t 25’, there is a large thermodynamic driving force for the oxidation of europium(I1) by water. I n such an oxidation there is a transfer of an electron from each europium(I1) ion, and apparently the metallic platinum provides an easy path for this electron tramfer. ORLO E. MYERS(Oak Ridge Nat. Lab.).-Would you care to amplify your discussion of the observed high activation energy for Eu(I1)-Eu(II1) exchange? It is about twice that for other one-electron transfer systems we know and of the order of that gencrally observed for two-electron transfer. GARNER.-The observed activation energy for the europium(I1)-euro ium(II1) exchange does appear to be somewhat large reLtive to the few Values which have been obtained so far for one-electron-transfer systems. However, the apparent activation energy includes an unknown contribution from the heat of formation of the chloro-complexed ion (or ion pair); the true activation energy would presumably be appreciably smaller after correction for this factor. E.g., a correction of 7 kcal./mole would bring the activation energy down to that found for the one-electron exchange between the tris-(ethylenediamine) complexes of cohalt( 11) and cobalt(II1).
