Ultrasonic absorption in aqueous salts of the ... - ACS Publications

for phosphorescence measurements, and to Dr. M. Nepras of Research Institute of Organic Synthesis,. Rybitvl, Czechoslovakia for measurements ofthe dic...
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a significant change of the LCI excitation energy, we recommend the use of the same value as for&oo. These values will undoubtedly change in the course of further studies, especially if we attempt to achieve a more general applicability.

Acknowledgment. We are indebted to Professor M. S. Newman of the Ohio State University, Columbus,

Ohio, for samples of the amino and hydroxy derivatives benzo[c]phenanthrene, to Professor Z. R. Grabowski of Academy of Science, Warsaw, Poland, and to Dr, M. J. Janssen of the University, Groningen, the Netherlands, for phosphorescence measurements, and to Dr. M. Nepra; of Research Institute of Organic Synthesis, Rybitvf, Czechoslovakia for measurements of the dichroic curves.

Ultrasonic Absorption in Aqueous Salts of the Lanthanides by Douglas P. Fay, Daniel Litchinsky, and Neil Purdiel Department of Chemtstry, Oklahoma State University. Stdllwater, Oklahoma


(Received June 8 6 , 1968)

The rates of formation of the complexes MS04t have been measured using the ultrasonic pulse teclhnique for the trivalent ions of praseodymium, neodymium, samarium, terbium, holmium, erbium, thulium, and lutetium. These results complete the series of trivalent lanthanide sulfates. To test the independence of the rate-controlling step from the identity of the entering ligand, the nitrates of trivalent lanthanum, cerium, praseodymium, samarium, and erbium have been studied. The spectra are more complex than those of the corresponding sulfates but a qualitative interpretation is presented which suggests that the third step in the overall complex formation mechanism is observed.


Experimental Section

In the continuation of previous work on the kinetics of complexation of the lanthanide ions in solution,2 the sulfate series has been concluded and a number of the nitrates have been examined. The three-step complex formation mechanism of Eigen3 was adopted in the earlier treatment of the data, the chemical relaxation having been attributed to the third step of the mechanism, that of substitution into the hydration sphere of the cation. Because of the limitations in identifying a relaxation with a specific step in the mechanism by this technique, some valid arguments were presented t o reinforce the somewhat intuitive assumption made. If the absorption could be shown to be independent of the anion, experimental justification for the assumption would be obtained. To this end a number of the lanthanide nitrates have been studied. Although the sound absorption spectra are not typical of a single relaxation, a qualitativc interpretation is presented which suggests that the original interpretation is correct. In one respect the treatment of the data in the original presentation for the sulfates was inconsistent. This has been remedied in this work and the results for La(III), Ce(III), Sm(III), Eu(III), G d ( I I I ) , Dy (111), and Yb (111) have been recalculated and are included for comparison with those for the remaining ions in the series. The results are discussed in the light of other available data.

Measurements of the absorption of ultrasonic energy were made at selected frequencies between 5 and 75 MHz using the pulse m e t h ~ d . ~The system is similar to that described previously2 for the low-frequency range. A temperature of 25 f 0.05” was maintained for all measurements reported. Solutions. Rare earth nitrates and oxides with a purity of 99.9% were obtained from the American Potash and Chemical Corp. The nitrates were used without further purification. The hydrated rare earth sulfates were prepared from the corresponding oxide. The oxides were dissolved in 6 N hydrochloric acid and 6 N sulfuric acid added t o yield a quantitative amount of the sulfate which was precipitated by the addition of a large excess of absolute ethyl alcohol. Stock solutions were prepared and analyzed for cation concentration by the titration with standard sodium hydroxide of the sulfuric acid or nitric acid produced by ion exchange on Dowex 5OW-X8, 20-50 mesh resin. Sulfate was estimated gravimetrically as barium sulfate. Dilutions

The Journal


Physical Chemistry

(1) To whom communications should be directed. (2) N. Purdie and C. A. Vincent, Trans. Faraday Soc., 6 3 , 2745 (1967). * (3) M . Eigen and K . Tamm, 2. Eleklrochem., 6 6 , 93, 107 (1962). (4) M. Eigen and L. DeMaeyer. “Technique of Organic Chemistry,” Interscience Publishers Inc., New York, N. Y., Vol. V I I I , Part 2, Chapter 18.




-500 -400





000961 F 00084 F


.9 v





- 40 - 20

- 20 5


2’5 35

55 75

5 I


“5 I





573 I




Figure 1. Excess sound absorption us. frequency as a function of concentration for neodymium and samarium sulfates. Samarium data from ref 2. (Concentration for lower curve is 0.0021 F.)

Vv MHz.

Figure 2. Excess sound absorption os. frequency a8 a function of concentration for terbium and holmium sulfates.

- 20




25 35

55 75



25 35

55 7 5

Volume 78, Number 9 March io69



of the stock solutions were made with deionized water, a t least three concentrations being exanlined for each salt.

sociation constants for steps I, 11,and 111, respectively.2 O(C) is given by the expression

Results I . The Lanthanide Sulfates. The relaxation spectra of the sulfates of Pr (111), Nd (111), Sm (111), Tb (111), 130(111), Er (111), T m (111), and Lu (111) at various concentrations are given in Figures 1-3 and 4a-8a. The solid lines are calculated from the equation for a single relaxation. The excess sound absorption produced by the chemical process is expressed as (Yohen,h. The attenuation of a plane progressive sound wave traversing a solution is given by the expression



loexp( - 2 4

(1) where I is the sound intensity at distancc x , Io is thc , a is the absorption sound intensity at distance z ~ r oand cocfficicnt for thc solution. ‘I’hc cxperiiiieiitally nicasured absorption is a, and aol,~m is obtained hy subtracting the solvent contribution ( Y I I ~ O . The units of (Yohcn, in eq 1 are nepcrs/ci-n. Only one iizaxiniuiii is o1)scrvcd in thc frequency mngc (~xaniiiicd,and Ihc nzagnitudc of the exccSs ahsorption is concenl ration depcndcnt. Reasons for believing lhal the absorption is due to substitiit ion int o thc iniicr hydration splicre of the cation have 1,ccii enumerated in tlic first paper of this scrics.2 ‘Phis is thc third stcp in the overall nicchanisin for coiiiplex forination proposed by Eigen3


where C is the concentration of solute in moles per litel. of i\[email protected])3, p is the degree of association at equilibrium, and IIf is the activity quotient fa+ f i - / f ~ . p was calculated for each value of C by standard iteration procedure using the total thermodynamic equilibrium constants, where available, derived from conductivity measurem e n t ~ . The ~ activity coefficients f3+, f2-, ft: wcre evaluated using a modified form of thc Davies equation.B

+ Xh2-(aq)

whcrc f, and z z arc the activity coefficient and chargc of the ion i, respectively, p is the ionic strcngth = 3C l 2 ( l - p ) C , B -- 0.33 X lo8, and It is the distance of closest approach of thc ions. .Is a rule It is takcn to bc cqual to 3 A which is too small for a 2 : 3 electrolyte. To he coiisistciit with thr thcorctical calculation of K12, required in the solution of cq 8, a value of 8.80 A, cqual to the suni o f thc ionic radii plus two mater moleculc dianictcrs, mas used. This rr1,rcsc~ntsa major dcparturc from 1 hc original trcatmcnt.2 The solution of thc derivative in the expression for 0 (C) was evaluated according to7


d In 111



a In p





a 111 lIf




+ zso4*-2- Z>fSO4+2)2


brsj ika2

(11) ha4



(111) Thc relaxation time for the rate-controlling step is given by the equation2 7111-1





-t- k’3r


where YmIII is the frequency of the maximum absorption. Recause of the coupling of €he successive equilibrium reactions, the correction for the activities of the ions in the first step is transmitted to k3d4. k’84 in eq 2 includes this correction. The rate expression for step 111 can be rewritten in terms of concentration and equilibrium constants 2~v~m = 1k43


+ @(C)kaa

K12, K23,

If it can hc argued that the chcniical relaxation is a of step 111, thc rate constants, and lience Ku, can be obtaiiied from a graphical solution of eq 3 , The unknowns in this equation arc K 1 2 and K23, h’la was calculated to be 2.3 X ill using the Rjerrum equations

.so4+3r _ ~ . 1 1 3 + ( ~ 1 ~ 0 ) s o ~ ~ - ~result



The Journal of Physical Chemietry


are the thermodynamic dis-

and was cssciitially independent of the small changes in 6 for the series, Previously,*K 2 3 was taken to be 0.51, the value for MgS04.9 Additional calculatioizs showed (6) A. E. Martell and L. G. Sillen, “Stability Constants,” Special Publication No. 17,The Chemical Society, London, 1964,p 237. (6) 0. W. Davies, “Ion Association,” Butterworth and Co. Ltd., London. 1962. (7)J. Steuhr and E. Yeager, “Physical Acoustics,” W.P. Mason, Ed., Academic Press, New York, N.Y., 1965 Vol. 11, Part A, p 388. (8) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed, Butterworth and 00.Ltd., London, 1959,p 392. (9) G.Atkinson and S . K. Kor, J. P h y s . Chem., 69, 128 (1966).



00545 F










-200 0 0088 F

d U f -


* -100

$ - eo

2 - 60 - 40

- 20








515 7;3


25 35 I


55 75 I




Figure 5. Excess sound absorption us. frequency as a function of conoentration for (a) cerium sulfate and (b) cerium nitrate. from ref 2. Single relaxations for nitrate calculated from T~~~ values for samarium nitrate.







." U n

Sulfate data

0.0512 F


-"OS c

-1 00 80

- 8s


- 60

t; - E O

- A0



- 20






25 35

55 75


V , MHz.

Figure 6. Excess sound absorption us. frequency as a function of concentration for (a) praseodymium sulfate and (b) praseodymium nitrate. Single relaxations for nitrate calculated from T~~~ values for samarium nitrate. Volume 7.9, Number 9 March 1960



IJ, MHz.

Figure 7. Excess sound absorption us. frequency as a function of concentration for (a) samarium sulfate (this work) and (b) samarium nitrate.



Er,(SO,), (a)


EdNOJ, (b)


Figure 8. Excess sound absorption vs. frequency as a function of concentration for (a) erbium sulfate and (b) erbium nitrate. In Figure 8(b) curve a is the theoretical single relaxation for outer substitution using T~~~ for samarium nitrate; curve b is the theoretical single I relaxation for inner substitution which added to curve a would give the experimentally observed chemical absorption curve c. Y ~ I I is outside the frequency range.

that k43 was insensitive to K23within the range 0.1 to 1.0, but that the value of ka4increased proportionately. The overall thermodynamic association constants K T , as determined by conductometric techniques, are related to the equilibrium constants for the individual steps by the expression

KT =


+ Kaa + K23K34 KizKz&a4



By a reiterative procedure, using eq 4 and 9, a value of K z was ~ calculated for samarium to give the best agreement between KT obtained kinetically with that obtained conductometrically. This value, K23 = 0.72, was subsequently used in the determination of k~ and k43 for the remaining ions. The agreement between the overall kinetic constants and the literature conductoThe Journal of Phyeical Chemistry

metric values5 mas good in all cases. Typical niagnitudes for the terms in eq 5 are shown for samarium in Table I. The characteristic relaxation frequency ~ I was I determined graphically from the linear plots of us. aohem according to the equationlo =




where the slope is - l/vn?III and A is the amplitude of the absorption. Concentrations and relaxation frequencies are given in Table 11. The values of IC34 (sec-l) and IC43 (sec-l) are given in Table I11 for all the lanthanide(II1) ions. (10) J. H.

Andreae, P. L. Joyce, nnd R. J. Oliver, Proc. P h y s . Soc.,

82, 75 (1960).





Table I: Samarium Sulfate ( d = 8.86A; K12 = 2.3X mol l.-l; K H = 0.72) 101



0.0822 0.1173 0.1636

10.60 5.29 2.64


101 X P

a In nf/a In p

0.1967 0.2503 0.3162

4.21 2.33 1.31

-0.095 -0.132 -0.157 ~~

The values are in some cases 40% lower than in the original work. This is in part due to the more exact solution of the derivative term in O(C) but more so to the use of the modified Davies equation in the calculation of activity coefficients. II. The Lanthanide Nitrates. The sound absorption spectra for the nitrates of La(III), Ce(III), P r ( I I I ) , Sm(III), and Er(II1) are shown in Figures 4b-8b. Again the solid lines are theoretical single relaxation curves. The solution concentrations had to be much higher for a relaxation to be observed. Quantitative interpretation of the data is complicated by the existence of more than one relaxation. However, the differences in the spectra from each other and from the corresponding lanthanide sulfates allow a qualitative description to be made. The curves differ in three respects. (1) Whereas the sulfate solutions of equal concentration absorb to the same order of magnitude, a comparison among the nitrates shows an order of magnitude difference, Table IV. (2) The low-frequency maximum of the sulfates has been shifted to a


x e(c) 2.53 1.73 1.21


0.507 0.457 0.403


Table 111: Rate Constants

2.1 3.3 4.4 5.2 7.4 6.6 6.7 5.2 4.2 2.8 1.9 1.4 -0.8 -0.6

Lao Cea Pr

Nd Sm Eua Gda Tb DY"

Ho Er Tm Ybo


5.6 7.0 6.4

0.870 0.902 0.918 0.926 0.962 0.970 0.981 1.000 1.010 1.030 1.043 1.053 1.065 1.077


14.0 14.6 12.8 9.6 5.2 3.5 3.0

1 .o

-3.7 ~1.8

Recalculated from data of Purdie and Vincent?

higher value in all cases except erbium. (3) In the nitrates, the relaxations are considerably broadened with respect to those of the corresponding sulfates. These observations are indicative of multiple relaxations.

Table TI: Relaxation Frequency Data

Discussion Ion



Formal concn X 101



~/KT mol , 1.-1

0.54 0.27 0.13 0.96

120 97 76 141 117 92 229 195 158 153 136 108 69 57 43 51 43 34 35 30 22 30 27 25

2.38x 10-4

0.48 Sm


Ho Er

Tma Luo


0.24 1.06 0.53 0.26 1.00 0.50 0.28 1.08 0.54 0.27 0.99 0.54 0.27 2.02 1.00 0.50 4.04 1.68 1 .oo

2.57X 10-6

The most serious limitation of the sound absorption technique in studying kinetics of chemical reactions is the uncertainty in assigning a relaxation to a specific step in a multistep mechanism. It, is essential to the present discussion that the characteristic frequency for substitution into the sulfate hydration sphere be identified. If step I11 is rate controlling, the relaxation for step 11 should occur at a higher frequency. The independence of an observed 200-MHz relaxation on the cation, in a number of divalent transition metal sulfates, led to its identification with step I1 in the overall nie~hanism.~In the earlier work on the lanthanide

2.56X lo-'

Table IV: Absorption Peak Maxima

2.29x 10-4 2.17x 10-4 2.56 x 10-4


2.56x 10-4

La Ce Pr Sm Er

Stability constants obtained by interpolation from a plot of

KT us. atomic number.

Nitrate *


2.36x 10-4



0.0120c*d 0.0120c5d 0.0190 0.0155 0.O14Oc




0.0115 0.0145 0.0130 0.0022

=Solutions are approx 5 X F. *Solutions are approx 5 X 10-8F. Peaks are not complete. d Data from ref 2. 0

Volume 75, NumbeT PI March 1060

550 sulfates, where data were available up to 230 MHz, this appeared as a broadening of the observed relaxation on the high-frequency side, an observation confirmed by Atkinson." The amplitude of this high-frequency relaxation was at least one order of magnitude smaller than those absorptions assigned to cation inner-sphere substitution. From the point of view of relaxation amplitudes, therefore, the low absorptions observed in the nitrates of lanthanum and erbium, around 35-55 MHz, might be interpreted as due to the equilibrium involving substitution into the nitrate ion, that is step 11. This relaxation manifests itself because the lanthanide nitrates are known to form predominantly outer-sphere complexes with some inner-sphere substitution occurring only a t high concentrations.'Z A similar small relaxation has been observed in preliminary studies of calcium nitrate and uranyl nitrate, but, as expected, is absent in potassium nitrate. Lanthanide sulfates on the other hand are predominantly innersphere complexes consistent with another trivalent The relaxation amplitude metal sulfate, Crz(S04) is correspondingly higher. Since in aqueous solution the effect of the pressure wave on the volume change in the equilibrium reaction is more important than the temperature effect on AH, more energy is absorbed in the process with the largest volume change. That process would be inner complex formation where the resultant reorganization of the solvent is greater. The amplitude of the absorption in the sulfates increases with increasing concentration of inner complex. To produce an appreciable concentration of inner complex in the nitrate solutions, that a similar relaxation might be observed, higher concentrations of the nitrates had to be used. For the most part a tenfold increase in cation concentration was sufficient. The magnitude of the absorption for lanthanum nitrate, for example, is much smaller than that of a more dilute lanthanum sulfate solution, consistent with a lower concentration of inner complex. If then the assignment of the relaxations in the sulfates to cation substitution is correct, the characteristic frequencies for inner substitution in the nitrates mould be rxpected to occur in approxiinntely the same frequency range, 4 4 0 MHz. From the sulfate data the relaxation frequency for innersphere substitution increases from Er (111) through L a ( I I I ) , Ce(III), P r ( I I I ) , and Sm(II1). For the nitrates the absorption amplitude increases in the same order. If the low absorption observed for erbium nitrate can be interpreted as due to substitution into the outer-sphere complex with the minimum of interference from the inner-sphere substitution process, then as the relaxation frequency for inner-substitution for the other cations increases, the extent of interaction and the absorption also increases. This model satisfactorily describes the differences in the spectra described before and provides qualitative experimental prooi that the original assignment of relaxations to step 111 i s correct.2 The Journal of Physical Chemistry

D. P. FAY,D. LITCHINSKY, AND N. PURDIE Using this interpretation, it should be possible to obtain quantitative confirmation by determining the rate constants for steps I1 and 111; the values of ka4 should be the same for both anions if the mechanism is Sn-1. To determine these rate constants the relaxation times for both processes must be separated from the complex sound absorption spectrum, and the quantity S(C) (see eq 13) must be evaluated. S(C) is related to rII by the expression

Let us consider these in turn. It is theoretically possible to analyze a complex spectrum for multiple processes if the conditions are satisfactory. Accurate experimental data is a necessity and T , ~should a t least be equal to 2 ~ ~ ~ For1 .cerium, praseodymium, and samarium nitrates, however, TIII 'v T I I and the spectra are almost typical of a single relaxation. Describing the absorption in terms of a/v2 rather than as ah, the simplest general equation for multiple relaxations mill become

for two steps, where A is the classical solvent absorption and the sum of the remaining two terms is the excess absorption due to chemical processes. From erbium nitrate the magnitude of the second term on the righthand side is very small and the measurements cannot be too accurate. To partition the total excess absorption, for samarium say, into a sum of two contributions when one is very small, cannot be done with confidence and the relaxation times therefore may not be considered unique solutions. The validity of this argument assumes the additivity of terms in the multiple relaxation equation which is the case only in the limit when the individual relaxations are independent, If, homever, as a first approximation, the value of RII for erbium nitrate is assumed to be constant for all of the nitrates, BIII for samarium is a very large quantity, almost equal in magnitude to RIII for samarium sulfate. This is unreasonable because the molar concentration of inner SmN032+complex is much less than the molar concentration of inner SmS04+ complex. The total observed excess absorption a o h e m / v 2 , therefore, is much greater than the simple sum of the individual contributions. General equations have been derived which account for this effectI4but the large number of undetermined constants involved makes it very difficult to interpret, except in extreme cases of simplification. The relaxa(11) G. Atkinson, private communication. (12) I. -4brahamer and Y. Marcus, Inorg. Chem., 6 , 2103 (1967). (13) IFogel, -, J. Yi.J. Tai, and J. Yarborough, J. Amer. Chem. Soc., 84, 1145 (1962). (14) R. T. Beyer, J. Acoust. Soc. Amer., 29, 243 (1957).



tion frequency, however, approximates to that of the slower step. Since the curve, as far as it is possible to tell, is symmetrical, an estimate of 711 equal to 7111 for samarium cannot be far from the real value. The 711 values for samarium could then be used €or the remaining salts at equal concentrations. This estimate is probably as good as would be obtained by a more protracted analysis. To evaluate e(c), the degree of association ,B must be calculated from the overall association constants. The thermodynamic

association constants for the rare earth nitrates are not available. A study of the reactions at high ionic strength, for which stability constants are available for a few of the rare earth nitrates, might be more advantageous at this time rather than attempting to measure the thermodynamic constants. It is conceivable that the rclaxations may be separated under these conditions although there is no precedent for making this observation, Such a study is at present in progress. Quantitative proof of the assignment of relaxations is therefore beyond our reach at present. Nevertheless, based upon the intuitive arguments outlined in previous work2 and the qualitative evidence described here, the interpretation of the results as indicative of inner substitution of the cation will be considered correct for the following discussion. Rate constants for Ftep 111, IC34 and ha,are given in Table I11 for the monosulfates. The estimated error is =tl5%, but for the slower rates, which approach the experimental limits of the instrument, the error is probably greater than this. In the graphical solution of eq 3, the intercept at @(C) = 0 is negative. However, when A (from eq 10) is plotted as a function of @(C) a zero value of A is obtained at a finite value of CP (C) , which means that a certain limiting concentration of inner complex must be present in solution before absorption in excess of the solvent is observed. This limiting value of @(C) was used to determine kra. The rates of complex formation are 40% lower than the original estimatej2but are still an order of magnitude greater than those observed by Geier'j for the rates of formation of murexide complexes, by Swinehart16for a number of lanthanide anthranilate complexes, and by Reuben and Fiat," in proton magnetic resonance studies of dysprosium. However, good agreement was found with results by Grecsek'* for P r ( I I I ) , N d ( I I I ) , and Dy(II1) from sound absorption and by bIarianelli19 for Gd(II1) by oxygen-17 nmr line broadening studies of the rate of water exchange. This controversy can only be resolved when results are available for the lanthanide ions with a certain ligand by more than one technique, and under similar conditions of ionic strength. If this fails then a possible explanation may transpire from



further consideration of the mechanism. I t is possible that the complexity of the entering ligand, or its ability to chelate with the cations, would preclude a simple Srvl mechanism. From Figure 9, the rate constants h4 are seen to reach a maximum around samarium as before2 and to

04 087









Figure 9. Dependence of logarithm of the rate constant for the rate-controlling step on the reciprocal cation radius. Data from: 0 , ref 2; 0, this work; A, ref 18; 0, ref 19.

(Values in ordinate X 10-8.)

fall on a smooth curve when plotted as a function of the reciprocal cation radius. The dependence is quite different from the linear behavior for log 1234 us. l/reatlo,, for the alkali metal and the alkaline earth metal series,20 but resembles somewhat the rather complicated trends in 4G, A H , and 4s for complex formation for a number of ligands with the rare earth cations.21 These trends have heen interpreted as a consequence of a change in coordination number as the radius of the cation decreases across the series. If the mechanism is SN1 so that the participation of sulfate anion in forming the transition state is small, trends in the rates of substitution may be expected to reflect changes in the intimate structure of the cations, namely a change in coordination number. Based upon this observation a feasible kinetic model was described in the first paper on the lanthanide sulfates.2 Unlike the murexide results, no minimum is observed around erbium. There is, perhaps, more error in the calculated kac values for the (15) G. Geier, Ber. Bunsenges. P h y s . Chem. 69, 617 (1965). (16) J. H. Swinehart, private communication. (17) J. Reuben and D. Fiat, Chem. Commun., 729 (1967). (18) J. J. Grecsek, M.S. Thesis, University of Maryland, 1966. (19) R. Rlarianelli, Ph.D. Thesis, University of California, Berkeley. 1966. (20) hl. Eigon, Ber. Bunsenges. Phys. Chem., 67, 783 (1963). (21) G. H . Nancollas, "Interactions in Electrolyte Solutions," Elsevier Publishing Go., New York, N.Y., 1966, p 108. Volume 78, Number 8 March 1060



slower reactions at the end of the series, but some credibility is given to the observation if, from Table 11, the characteristic frequencies are compared for solutions of equal concentration. The dependence of the rates of formation on the reciprocal cationic radius follows, to some extent, that of the heats of hydration of the trivalent lanthanide ions,22which might indeed suggest an Srvl mechanism.

Acknowledgment. We wish to acknowledge the financial assistance of the Research Corporation and the

Research Foundation, Oklahoma State University, which made it possible to construct the equipment; Dr. A. J. Barlow, Department of Electrical Engineering, University of Glasgow for its assembly; and Mr. H. Hall of this department for his help in the design and fabrication of the mechanical arrangement. We are also grateful to NASA for providing a fellowship to

D. P. F. (22) G . Choppin, unpublished results.

Radiolysis of Cyclic Fluorocarbons. 11. Perfluoroaromatics at Elevated Temperatures1 by F. W. Bloch and D. R. MacKenzie Brookhaven N a t i o n a l Laboratory, U p t o n , hrew Y o r k


(Recetved July 1 5 , 1 9 6 8 )

Aromatic fluorocarbons are known to have high thermal stability but only moderate stability to radiation at or slightly above room temperature. I n the present study of perfluorobenzene, perfluoronaphthalene, perfluorobiphenyl, and perfluoro-o-terphenyl, we find that this moderate radiation stability is maintained at elevated temperatures where most organic compounds undergo rapid decomposition by heat alone. For example, perfluorobiphenyl at 500' is several orders of magnitude more stable thermally and many times more stable to y radiation than are the hydrocarbons biphenyl and o-terphenyl. A t 100" polymerization is almost the only process occurring and G values range from 1.5 to 3. At 450' it is still the most important process, with G values from 2 to 6, but formation of Fz-addition products has become appreciable. Fragmentation is insignificant. Slthough the temperature coefficient of overall radiolytic decomposition is small, quite large changes occur in the product distributions in going from 100 to 450'. This is particularly apparent in the case of CeFawhere products have been characterized to the greatest extent. The implications regarding mechanism and changes in mechanism with temperature are discussed.

I. Introduction It has been known for some time that the commoiier cyclic fluorocarbons, both alicyclic and aromatic, arc extremely stable thernially.2a3 Our own ~ ~ hasr shown that a t temperatures from room to 100" the aromatic and alicyclic fluorocarbons have similar G values for radiation decomposition, and their resistance to radiation is intermediate between aromatic and alicyclic hydrocarbons. Because of their known thermal stability, it was of interest to find out whether these compounds retained their reasonably good radiation stability up to temperatures as high as those a t which hydrocarboizs undergo severe degradation. Thus we extended our irradiation experiments to elevated temperatures (up to 500O).

11. Experimental Section The compounds studied were perfluorobeizzenc (Cap,), perfluorobiphenyl (C12F10), perfluoronaphthaThe Journal of Physical Chemistry

lene (CloFs), and perfluoro-o-terphenyl (CI8FI4), The perfluoroterphenyl mas synthesized by ,J. W, Dale and G. J. O'Neil of Jlonsanto Research Corp., Everett, brass. k ~ It was available in only very limited amounts and was irradiated as received (98y0pure). The other compounds, obtained froin Imperial Smelting Corp., Ltd., Bristol, England, were highly purified by preparative scale gas-liquid partition cliroiiiatography k1PC) A Karl E'ischer spectrophotoinetric analysis showed a water content of <40 ppni for all the purified compounds. This small amount of water and probably e4

(1) This work mas performed under the auspices of the U. S. Atomic Energy Commission. (2) I . B. Johns, E. A . McElhill, and J. 0. Smith, I n d . Eng. Chenl.. Prqd. Res. Develop., 1 , 2 (1962). (3) L. A. Wall, R . E . Donadio, and W.J. Pummer. J. Amer. Che?ft. Soe. 82, 4846 (1960). (4) D. R. MacKenzie, F. I T . Bloch, and R . H. Wiswsll, Jr., J. P l W . Chem., 69, 2526 (1965). Paper I of this series.