THEACTIVITYCOEFFICIENT OF &BUTYL CHLORIDE
June 20, 1962
2295
tion AS,* for the aqueous solvolysis of gaseous tbutyl chloride derived from this result are 14.3 f 0.3 kcal./mole and 24.9 i 1.0 e.u., respectively, a t 8’. The solvent isotope effect, kDIO/kH,or for the rate COMPARISON OF THE CALCULATED ENERGYOF ACTIVATION of solvolysis of gaseous t-butyl chloride a t unit AH,*, WITH LITERATURE VALUES FOR SOLVOLYSIS OF pressure at 14.85’ is 0.71. This figure may be com&BUTYLCHLORIDE I N WATER pared with the solvent isotope effect for the rate Method Ref. AHc k, cal 1. mole of solvolysis of dissolved t-butyl chloride at unit 4e, 9 2 4 . 1 i:0.2 Titration concentration a t 14.03’ of 0.74.4d-11The essential 10 24.4 f 0 . 2 Conductivity agreement of these two figures (which is consistent This work 25.2 f 1.1 kp/h with the fact that the distribution constant for A plot of log kp us. l / T is linear within the ex- t-butyl chloride is the same within experimental perimental error for the temperature range meas- error in DzOl as HzO cf. Table 11) does not support ured, 1-15’. Thus the value of l o 4 k, calculated the suggestion of Heppolette and Robertson12 a t 5.00’ by the Arrhenius equation from measure- that the solvent isotope effect resides primarily ments a t 0.99’ and 14.85’ is 1.15 & 0.04 compared in the isotope effect on the initial state (dissolved t o the experimental value of 1.18 f 0.03. The t-butyl chloride). Instead, in the present instance energy of activation AHp* and the entropy of activa- a t least, our results require the effect to apply primarily to the solvolysis reaction transition (9) Calculated from the value given for 2 5 O using ACp* = -50 st ate. cal /deg., cf. ref. 4d and e
of gaseous t-butyl chloride in water to be 8.2 kcal./ mole a t 25’, in quite reasonable accord with the above value. lob TABLE IV
+
(SO),
*,
(10) (a) Private communication from Dr. R. E. Robertson; (b) R. E. Robertson, R. L. Heppolette and J M. W. Scott, Can. J . Cham., S I , 803 (1959).
[CONTRIBUTION FROM
THE
(11) Cf.also C. G. Swain, R. Cardinaud and A. D. Ketley, J . Am. Chem. Soc., 7 7 , 934 (1955). (12) R. L. Heppolette and R. E. Robertson, ibid., 88, 1834 (1061).
COLLEGE OF CHEMISTRY AND PHYSICS,THE PENNSYLVANIA STATE UNIVERSITY,UNIVERSITY PARK, PA.]
The Effects of Aqueous Electrolytes on the Activity Coefficient of f-Butyl Chloride and of its Solvolysis Transition Statel BY GEORGEA. CLARKEAND ROBERTW. TAFT RECEIVED OCTOBER 2, 1961 An extensive study has been made of the effects of aqueous electrolytes (up to 1 M )on the rate of solvolysis of gaseous t-butyl chloride at unit pressure k , and on the solubility distribution constant h, for gaseous t-butyl chloride. From the former measurement, the relationship, log f+ = K C, (where f+ is the activity coefficient of the aqueous t-butyl chloride solvolysis transition state, K is the transition state salt parameter and C, is the electrolyte concentration), is demonstrated. From the latter measurement the following relation holds, log fp = K&, (where f, is the activity coefficient for dissolved t-butyl chloride and K, is the salt parameter for the reaction ground state). Neither log fg nor log f obey the expression obtained by Ingold and co-workers based upon the ion atmosphere model. In‘stead, quite specific salt effects are found due to important contributions to the activity coefficient from salt induced effects upon the internal pressure of where p.6,is the ion atmosphere (coulombic) contribution the solvent. For inorganic salts, it is shown that f E S.Y.P., and f’.”. is the salt induced internal pressure (non-coulombic) contribution. However, Ingold’s ion atmosphere treatment of kinetic salt effects is shown to be applicable. This demonstration is made possible in the present instance by the essential cancellation between ground and transition state of the specific salt induced solvent effects. Thus, the inorganic salt parameters for the conventional rate constant k , a t unit concentration are the same (as predicted by the theory) for all salts of a given charge type. In terms of Ingold’s equation a charge separation in the aqueous t-butyl chloride transition state of about 0.8 electronic units is indicated. Specific effects have been obtained for organic electrolytes and for the acidic cations, Ha0+ and NH(+, and these results are discussed. The distinct contrast of present results in water with that of previous studies in mixed solvents is attributed to the greater complexity of salt induced medium effects in the latter solvents with the attendant non-cancellation of these effects between the ground and transition states for solvolysis reactions.
*
*
*
Electrolyte effects on reaction rates have played an important role in the investigation of mechanism.2 However, the validity of theoretical treatment of salt effects has remained uncertain (cj. following Discussion.) Several investigations have been made recently in which the effects of electrolytes on a reaction rate and on the activity coefficient of the reaction substrate have been independently determined. The results of such studies have been especially enlightening. By this procedure Long, McDevit
trolyte solutions on the rate of the acid-catalyzed hydrolysis of y-butyrolactone t o y-hydroxybutyric acid was the direction of the electrolyte effect on the activity coefficient of the l a ~ t o n e . ~ Paul4 and Long and M ~ c I n t y r ehave ~ , ~ shown that the markedly different electrolyte effects on the acid-catalyzed hydrolysis rate of methylal and on the acidity function H0 is largely a question of the difference
(1) The work was supported in part by the Office of Naval Research. (2) A . A. Frost and R . 0. Pearson, “Kinetics and Mechanism,” John W l e y and Sons, I n ? , S e w Y n r k , S . Y , 1 U S 3 , Chapter 7 ; a ! s o cf. ref. S.i.
(3) F.A. Long, W. F. McDevit and F. Dunkle, J . Pkys. C h e n . , 66, 829 (1951). (4) 31. A Paul, J . A m . Chcm. Soc., 7 6 , 3236 (1954). is) U . McIntyre and F. A . Long, i b i d . , 76, 3240 (1954). (0) F .4 1.ong and 11. Alclntyre. i b i d . , ‘76, 9243 (1954)
and Dunkle showed that the direction of the effects of several salts in moderately concentrated elec-
2296
GEORGEA. CLARKEAND ROBERTW. TAFT
Vol. 84
in the electrolyte effect on the activity coefficients in the electrolyte effect on an activity coefficient: (1) The direct electrical (ion-atmosphere) interof methylal and the Hammett reference base. In the present study, the salt effects of aqueous action and ( 2 ) the indirect salt-induced medium electrolyte solutions (up t o 1 M ) on the activity effect (effect of the salt on the internal pressure of . investigations have made coefficients of t-butyl chloride and its solvolysis the s o l ~ e n t ’ ~ )Recent transition state have been determined. This has abundantly clear the importance of salt-induced been accomplished by determining the electrolyte medium effects. 20,21 The present investigation was carried out in effect on the solubility and on the rate of solvolysis kp of gaseous t-butyl chloride a t unit pressure. This water as the solvent. This choice was dictated rate constant gives a direct measure of the activity not only by practical considerations of the method of investigation but by the availability of detailed coefficient of the solvolysis transition state. The solvolysis of t-butyl chloride was chosen for knowledge of salt effects in this r n e d i ~ m . ‘ ~ J ~ . ~ ~ study because its vapor pressure and solubility in Our results on the solubility of t-butyl chloride proaqueous solution lend themselves to this type of vide the first detailed examination of aqueous salt investigation. Further, the mechanism for this effects on a non-aromatic organic non-electrolyte. reaction has been the subject of extensive scrutiny From the combined knowledge of the salt effect on and is considered to be a prototype of a S N ~ ’the rate constant kp and on the solubility distribuor limiting8 solvolysis mechanism. tion constant h we are able t o demonstrate (under Our results permit a substantial test of the ion- our especially favorable conditions) the essential atmosphere treatment proposed by Bateman, cancellation of inorganic salt induced solvent Church, Hughes, Ingold and Tahergfor the activity effects between ground and transition states. coefficient of a highly dipolar transition state. That is, the effects of inorganic salts on the conThis rather widely adopted theory, analogous t o the ventional rate constant kc of aqueous t-butyl Debye-Hiickel theory for the activity coefficient chloride a t unit concentration do follow the preof an ion,1° was initially employed t o account for dictions of the ion atmosphere model. the observed salt effects on the rate of solvolysis of Experimental t-butyl bromide a t unit concentration in 90% Materials.-Tetramethylammonium hydroxide ( lOy0 in acetone-water s ~ l v e n t . The ~ treatment was made water), tetraethylammoniumhydroxide (10% in water), possible by the assumption that there are no tetramethylammonium bromide and chloride were Eastman electrolyte effects on the ground state activity Rodak Company “White Label” chemicals. Lithium sulcoefficient. The theory predicts that the rate fate (monohydrate), potassium sulfate, potassium fluoride will increase with the first power of the ionic (dihydrate), lithium nitrate, lithium chloride, potassium and sodium thiocyanate were Fisher “Certified strength with no specific electrolyte effects. The chloride Reagent” chemicals. Sodium perchlorate (monohydrate) theory, however, was put to no stringent experi- and sodium benzoate (U.S.P.) were Fisher Laboratory chemicals. Sodium trifluoroacetate was a Sapon Laboratories mental tests. Inc. chemical. All other electrolytes were Baker Analyzed Subsequent studies have consistently demon- Reagents. strated that electrolyte effects on unimolecular All electrolytes were used without further purification. reactions are generally quite unique for a given The solids, with the exception of sodium thiosulfate, were electro1yte.”-l6 In fact, some electrolytes were heated in an oven a t temperatures varying from 115-130” dried in a vacuum desiccator to constant weight. The discovered to have a retarding effect on the re- or per cent. weight loss of the hydrated compounds dried by action rate.” Further, the ordering of the elec- these methods agreed with the calculated water content t o trolyte effects has been shown to be dependent better than 1% with the exception of sodium dichromate. upon the medium.I3 These observations together All salt solutions were made up by weight of the dried salts and the concentrations of the solutions of sodium thiosulfate with the failure of electrostatic theories to account, and sodium dichromate were determined accurately by even qualitatively, for salt effects on the activity titration using standardized potassium iodate for sodium coefficient 01 polar non- electrolyte^'^ have cast thiosulfatez4 and standardized sodium thiosulfate for sodium strong doubt on the ability of the ion-atmosphere dichromate.26 The concentration of the acid and base soluwere determined by titration with standardized reamodel to adequately describe salt effects in re- tions gents. The solutions of sodium thiosulfate and one of the action kinetic^.^^^^^^^^ Bronsted pointed out long solutions of sodium carbonate were made up with about 0.1 ago1* that two factors must be considered t o act M base; potassium hydrogen sulfate was made up with a 1 (7) (a) J. L. Gleave, E. D. Hughes and C. K. Ingold, J . Chem. Soc., 236 (1935). (b) E. D . Hughes and C. K . Ingold, i b i d . , 244 (1935). (8) E. Grunwald, S. Winstein and H. H. Jones, J . A m . Chem. Soc.. 73, 2700 (1951). (9) L. C. Bateman, Sf.G. Church, E. D . Hughes, C. K . Ingold and N . A . Taher, J. Chem. SOL.,979 (1940). (10) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” Reinhold Publishing Corporation, New York. N . Y., 1950, Chapters 2 and 3. (11) G. R. Lucas and L. P. Hammett, J . A m . Chem. Soc., 64, 1928 (1942). (12) 0. T . Benfey, E. D . Hughes and C. K. Ingold, J . Chem. Soc., 2488, 2494 (1952). (13) F. Spieth and A. R . Olson, J . A m . Chem. Soc., 77, 1412 (1955). (14) A. H. Fainberg and S. Winstein, ibid., 78, 2763 (1056). (15) Y . Pocker, Chcm. and Ind. (London), 1599 (1957). (16) E. F . J. Duynstee, E. Grunwald and M. L. Kaplan, J . A m . Chem. Soc., 6’2, 5654 (1960). (17) F . A. Long and W. F. McDevit, Chcm. Rcvs., 81, 119 (1952). (18) J. N. Bronsted, J. A m . Chem. Soc., 44, 877 (1922).
iZd hydrochloric acid solution. Procedure.-This has been previously described in dep d / p i , was generally tail.26 The apparent olefin fraction,Z@ (19) G. Tammann, Z . arzorg. IL. allgem. C h o n . , 188, 25 (1928). (20) (a) W. F. ZfcDevit and F. A. Long, J . A m . Chcm. SOC.,T4, 1773 (1952); (b) R . I,. Bergen, Jr., and F. A. Long, J. Phys. Chem., 60, 1131 (1956). (21) E. Grunwald and A. F. Butler, J. A m . Chcm. Soc., 82, 6847 (1960). (22) R . W, Gurney, “Ionic Processes in Solution,” McGraw-Hill 1953. Book Co., Inc., New York, N. Y.. (23) Discussions of the Faraday Soc., “Interaction in Ionic Solutions,” h-0. 24, 1957, Aberdeen University Press, Ltd., Great Britain, 1958. (24) I. M. Kolthoff and E. B. Sandell, “Textbook of Quantitative Inorganic Analysis,’’ The Macmillan Company, New York, N . Y., 1952, p. 594. (25) Ref. 24, pp. 594-595. (26) G. A. Clarke, T. R. Williams and Robert W.Taft. J. A m . Chcm. Soc., 84, 2292 (1962).
THEACTIVITY COEFFICIENT OF [-BUTYL CHLORIDE
June 20, 1932
SETSCHENOW SALTPARAMETERS K, FOR
THE
22'37
TABLE I DISSOLVED t - B r r m CHLORIDE IN VARIOUS ELECTROLYTE SOLUTIONS AT 14.85' K. Electrolyle K. Electrolyte NaOH +0.41" NaClO, +0.24" NazCOs .9v .43" KOH (CHa)dNOH .24" KaC02H .24" NaCOzCeHS .26" (CzHshNOH .24" Ka COZCH3 .29" .51d KF KaCOCF, .17" KBr .19" (CH3)ZBr .02" NaI .25' SazCrpO7 .64" h-41 .00" Na&Lh .87b*a HC104 .005 KaSCN .16"
K. H C1 +0.04" LiCl .2gb NaCl .35b KC1 .26' NHaCl .18" HNOa - .lo" LiNOs .25' His04 .44* .8g5 NazSO, KzSO4 .98" KHSO, .4651**' Estimated uncertainty in the parameter is lt0.05. e Estimated a Estimated uncertainty in the parameter is =k0.03. Estimated uncertainty in uncertainty in the parameter is f0.08. Estimated uncertainty in the parameter is ltO.10. Corrected for the added acid (1 A!f HC1) by assuming an additive Setschenow relationship. the parameter is f0.15. Corrected for the added base (0.1 M NaOH/KOH) by assuming an additive Setschenow relationship. Electrolyte
+ + + + + + + + +
+ + + + +
+ + + + + + + +
+ +
TABLE I1 t - B m n CHLORIDE TRANSITION STATEIN VARIOUS ELECTROLYTE SOLUTIONS AT 14.85' K+ f(* Electrolyte Electrolyte K* Electrolyte HC1 -0.03' KHSOd +o. 18b.d HClO4 -0.14" Sac104 - ,045 LiCl .03" NaOH .20b Na?COa .41" .19" NaCl .06" KOH NaC02H .14" .22b KC1 (CH3)rNOH .06b (CzHs)aNOH .24b NaCOGH6 .ll" "4Cl .02" NaF .22b KaCOlCH3 .21° HNOa .05b LiNOa .Olb KF .20b XaC02CFa .11" NaiYOt .035 KBr - .02" ( CH3)4NBr 00" XanCrlO7 .lo" &SO4 .Olb NaI .Ogh NazSzOa + *25",C? LiL304 .29' NazS04 .27b KHII .14b XaSCN .lo" (CHa)dNI - .12' KzSOi .33* Estimated uncertainty in the parameter is f 0 . 0 2 . b Estimated uncertainty in the parameter is f 0 . 0 3 . Estimated uncertainty in the parameter is lt0.04. Corrected for the added acid (1 M HCl) assuming an additive Setschenow reCorrected for the added base (0.1 41 NaOH/KOH) by assuming an additive Setschenow relationship. ationship.
SETSCHENOW SALT PARAMETERS K FOR
THE
+ + + + + + + + + +
+ + + + + + -
+ + + + +
.
-
unaffected ( = 0.04 f 0.01) by salt concentrations up to 1 M. Exceptions were noted for salts which strongly salt-out aqueous t-butyl chloride. The following values of p0ie/pi were obtained for the exceptional electrolytes a t 1 M:KOH (0.08 3z 0.01), (CHs)aNOH(O.lO f 0.01), KF(0.07 3z 0.01), Li2S04(0.07i 0.01), Na~CO~(0.10 f 0.01) and Na~S203(0.14 f 0.01). Since the salt effects on the olefin fraction tend to parallel (probably within the accuracy of the data) the electrolyte order for either K + or Kg,it is probable that these effects are medium effects acting to a greater extent on the substitution than the elimination transition state. The intervention of a measureable E2 reaction does not appear to be indicated by the results of Table VII.
The results for the distribution constant h for t-butyl chloride follow this relationship within their precision measures. Typical results are illustrated in Fig. l. The slope of these plots is denoted as the Setschenow parameter for dissolved t-butyl chloride, K,. The rate quantity, -log (kp/kpO), which is shown in the Discussion section to equal the logarithm of the activity coefficient for the t butyl chloride solvolysis transition state, also follows the Setschenow relationship within the Results precision measures. Typical results are illustrated Electrolyte effects on k, and h were obtained in Fig. 1. The slope of these plots is denoted as in each case a t 14.85' for electrolyte concentrations the Setschenow parameter for the t-butyl chlofrom approximately 0.1 M t o 1.0 M . The values ride transition state, K+. obtained are recorded in Dr. Clarke's Ph.D. t h e s i ~ . ~ ' Tables I and I1 summarize the results obtained. The results may be summarized (as follows) in The former gives values of K , and the latter values terms of Setschenow salt effect parameters. of K+. I n general, the K+ values are considerably Electrolyte effects on activity coefficients of non- more precise because of the experimental limitations electrolytes generally follow the empirical Setsche- in determination of precise values of h.26 now equation'r.2112a Discussion logf = log ( h O / h )= KC. (1) According t o the transition state theory, the where C, is the molar electrolyte concentration and electrolyte effects on k, may be attributed to the K is the Setschenow salt effect parameter. The effects on the following hypothetical "equilibrium" superscripto refers t o the distribution constant for t-BuCI(,> + (aq.) I_ t-BuCI(.,., (transition state) pure water in terms of the rate equation: (27) G. A. Clarke, Ph.D. Thesis, The Pennsylvania State Univer-
*
.
sity, June, 1960. (28) J. Setschenow, Ann. Chim. Phrs., (61 26, 226 (1891).
Rate --
P - kp =
(RmT )( K * ) (A)
2298
GEORGE
A. CLIRKEAND KOXERTi f - .T.ZFT
Vol. 84
mately a constant, independent of electrolyte. The high concentration of the rather nucleophilic solvent employed, water, tends to justify this assumption since a high value of ks (by definition k, includes the concentration of water) will cause this quantity to approach unity. It may be noted too that the external return contribution (common ion effect) to kz has been found by recent tracer studiesz9to be small even in 80% methano1-20% water. Further, the condition of constancy of this quotient is satisfied if the individual ratios, ke/'ks and kz/k,, are both constant. The olefin fraction, k,/k,, is in fact found to be generally independent of electrolyte concentrations to 1 M (cf. Experimental for exceptions). Finally, the nature of the results in section 5 also appear t o support the above assumption. With this assumption log ( k p O / k p ) = log I
,
020
040 060 C80 l0C I20 ELECT 9 OLYTE CONCE\TR,VTIOh
I
I
I40
(MOCES/ii:ER).
Fig. 1.-Typical Setschenow plots for the transition state ( A ) and the ground state ( B ) activity coefficients.
where K* denotes the transition state theory "equilibrium constant" for the above reaction, and f* is the activity coefficient of the transition state (standard state-hypothetical one molal in pure water). No term in t-butyl chloride appears on the right hand side of this equation since a t the low pressures involved the fugacity of the chloride may be taken as equal to its partial pressure. The salt effect on the solvolysis of gaseous t-butyl chloride, k p / k P o (where 0 refers to the reference state of pure water), is accordingly given as
log ( k p o / k D )= logs* = K*C,
(3)
ke i--)
olefin
L--+ ki
alcohol
Applying the steady state approximation to the intimate ion pair, R6+- - - Cia-, one obtains the rate equation
Rate - - - k, = hkl
P
[kn ke+ k e + ks ] + ks
(4)
For the purposes of this discussion we shall assume that the quantity,
k c + ks
kz
+ k e + ks, is approxi-
= logs* =
K+Ca (3')
where f+ refers to the transition state for formation of the intermediate. (1) Effect of Electrolytes on the Dissolved Tertiary Butyl Chloride.-From the results given in Table I, it is seen that the presence of an electrolyte generally salts-out (decreases the solubility or increases the activity coefficient) the dissolved t-butyl chloride. The effect is a characteristic The only electrolyte which has a definite salting-in effect is nitric acid. For purposes of comparison of cation and anion effects on t-butyl chloride, the salt parameters may be referred to arbitrary reference ions. Table 111 gives cation salt parameters for 1-1 electrolytes relative to the sodium ion. Table IV lists salt TABLE I11 RELATIVE CATION S A L T PARAMETERS FOR &BUTYLCHLORIDE IN 1-1 ELECTROLYTE SOLUTIONS
Further, combining equations 1 and 2 we have
A final consideration is necessary before the experimental results for the salt parameters K* and K , are considered. According t o the generally accepted mechanism for the solvolysis of t-butyl chloride in aqueous solution,,the experimental rate constant k , (or k,) is potentially a composite rate constant.8
(2)
Cation
AKg
(0.00) - .28" - .0ga .06< - .21b .17" - .17b
Na + H+ Li + K+
+
"I+
-
(CHs)&+ (CzH5)aN+ a Estimated maximum uncertainty zkO.10. Estimated Estimated maximum maximum uncertainty i 0.12. uncertainty + 0.15.
*
parameters for anions relative t o the chloride 01 carbonate ions as the reference anion ion. It has been assumed throughout that the effect of the ions of the electrolyte are additive and the results appear to support this treatment. It is clear from Table I11 that salting-out by an electrolyte shows the familiar relation to the cation size-the larger cations reduce the observed saltingout with respect to the smaller reference cation, Na+. The effect observed with hydronium or ammonium ions is a reasonable one in terms of well hydrated ions. The influence of the anion (cf. Table IV) is qualitatively similar t o that of the cation, ;.e., with increasing size of the anion t-butyl chloride is increasingly salted-in. (29) C. A. Bunton and B. Sayak,
J. Chem.
SOC.,3854 (1959)
June 20, 1962
THEAICTIVITV COEFFICIENT
TABLE IV RELATIVE ANION S A L T PARAXETERS
FOR &BUTYL
Anion
A&
(0.00) .2Zid .12d
c1FOHBrINO3HCOe-
+ + -
,075 .14" .04* .11" .OSb .18" - .06"
Cl0,-
CHLORIDE
Anion
A&
cos-
(0.00)
SO,'
s?o8Cr207-
-
.08" .0gc
- .:3zh
OF
Anion
c1FOH Br-
(2) Effect of Electrolytes on the Solvolysis Transition State.-Tables V and VI, respectively, record the relative cation (standard, sodium ion) and anion (standard, chloride and carbonate ions) salt parameters for the t-butyl chloride transition state. It is apparent from Table V that, in contrast t o the ground state, the transition state is nearly insensitive t o differences in cation size. Of the cations studied, only the hydrogen ion (and probably NH4+) shows a salting-in effect which is outside of experimental error. For the inorganic anions of the same charge type, the qualitative trend of increasing salting-in by the larger ions is pronounced and clearly evident. However, for the organic acid anions, this trend is strikingly absent and is replaced by a variable salting-out effect (cf. Table VI).
2299
TABLE1.1 RELATIVEAXIOXSALTPARAMETERS FOR THE SOLVOLYSIS TRANSITIOX STATE OF &BUTYL CHLORIDE
so3-
SCN CH3C02- .18" CHsC02CeHbC02- .09" Estimated max. uncertainty & 0.08. b Estimated max. uncertainty i. 0.12. Estimated rnax. uncertainty i 0.14. Estimated max. uncertainty j= 0.18.
BUTYL CHLORIDE
c10,1SCK HSOaAnion
COS'
AX*
0.00
+ + -
. lt9 .14h .O2*
- .OF - .11" - .ISC - .16"
+
Anion
CI HCO2CH3CO2CFaCOgCsHsCOz-
AK +
0.00 .08 .15" . 05" .E"
+ + + -+
.EC
AK -c
0.00 .14' .16" .51" Estimated uncertainty a Estimated uncertainty =t 0.04. Z!C 0.05. Estimated uncertainty f 0.06.
SO,' WOa" Cr207=
-
-
the solution; k = Boltzmann's constant; T = absolute temperature. For sufficiently dilute solutions equation 5 may be written as logf =
- apZ*dp
(6)
where (Y and 0 are the Debye-Hiickel constants for water a t a given temperature. At 14.85' the product (ab) is equal to 0.164 X lo8. According to the theory, the presence of ions in solution results in an ionic atmosphere about a dipole which lowers the free energy. The reduction in free energy is proportional to the ionic strength. It will be noted that both the empirical eqn. TABLE V 3 and the theoretical eq. 5 have a first power RELATIVECATIONSALTPARAMETERS FOR THE SOLVOLYSIS dependence on the concentration of 1-1 electrolytes. TRANSITION STATE O F &BUTYL CHLORIDE However., the latter equation anticipates that elecCation A K +" trolytes of a given charge type will all give the same Na + 0.00 value of .K+. In this respect the results in Table H+ - .09 I1 in no way conform to the predictions of equaLi - .03 tion 5 , Specific electrolyte effects are observedK+ - .01 both accelerations and retardations (negative and KHa - .05 positive Setschenow parameters. respectively) of (CHs)rS+ .00 the solvolysis rate a t unit pressure are observed. .04 (C&)4N+ It is clear, therefore that eqn. 5 applied directly Estimated uncertainty f 0.06. to the t-butyl chloride transition state is totally inadequate in describing the nature of the observed (3) The Ion-atmosphere Model Applied to the salt effects. Ground and Transition States of Tertiary Butyl Aqueous t-butyl chloride may also be regarded Chloride.-Ingold and co-workers9have considered as a dipole. The K , parameters therefore provide the effect of an electrical atmosphere of pointcharge ions about a point dipole in a continuous a further test of eq. 5 . The complete failure of this dielectric medium. Their treatment was analogous equation to account for the observed behavior of to that of Debye-Hiickel'O for the activity co- t-butyl chloride in the ground state is obvious from efficient of an ion. The expression obtained for the the results given in Table I. It appears that the principal reason for the activity coefficient of the dipole is failure of eq. 5 to reproduce the observed effects 4rNe4 of electrolytes on the activity coefficient of dipolar logf = (5) 2.303( 1Oa)(DkT)* molecules arises because of the assumption that Z = the fraction of an electronic charge a t the only effect of the electrolyte is the ion ateach point-pole; d = the fixed distance of sepa- mosphere effect (cf. subsequent section). ration of the point charges of the dipole; p = A second potential source of difficulty is the the ionic strength of the electrolyte solution con- fact that the concentration range over which a taining the point-dipole; N = Avogadro's number; Debye-Hiickel type ion-atmosphere treatment can c = electronic charge; D = dielectric constant of be expected t o apply is a very limited one. Fuoss +
+
+
GEORGEA.
2300
CLilRKE AND
ROBERT\.v. TAFT
Vol. 84
am-
040
I80
c
I
-023
1
020
3 3 N,
-
0 3
TER
320
530
24c
-
0%
I
MOLES
Fig. 2.-The relationship between the ground state parameter K g and the transition state parameter K+.
0301
I
-020
0
220
;40
,s' o
a ;3
0
K g L I T E R S / VOLE.
and Onsager3" consider the region of reliability for the Debye-Hiickel theory t o extend t o no more than 0.01 M . However, Ingold and co-workersg anticipated that eq. 5 for dipoles should be valid over a more extensive range because of the smaller density variation of the ionic atmosphere around dipoles. In reaching the conclusions above, we have assumed that the ion-atmosphere theory for dipoles is a t least approximately valid over the entire concentration range of the present investigation. The results in a subsequent section appear to justify this assumption. (4) Comparison of Ground and Transition State Salt Parameters-Evidence of Salt Induced Medium Effects.-A plot of corresponding Kg ns. K+ parameters (Fig. 2) does not reveal any general relationship between the ground and transition state salt effects. However, if one considers only the inorganic salts (;.e. excludes organic cations and anions and the acidic cations H30+ and NHk+), distinctly linear relationships of near unit slope are apparent (dark circles in Fig. 2 ) which separate accordingly to the charge type of the anion. In this section anti the following one, we shall confine our attention to this orderly behavior exhibited by the inorganic salts. In a subsequent section, the behavior of the organic cations and anions and the acidic cations will be considered. The separation of inorganic salts according to charge type in Fig. 2 provides clear evidence of the operation (at least qualitatively) of the ion atmosphere effect. However, i t is equally clear that there is a t least one other important contribution to the activity coefficients. According to the ion atmosphere model, one should obtain the same salt parameter for all electrolytes of a given charge type (so that Fig. 2 would consist (30) R. M.Fuoss and L. Onsager,
J. Phys. Chcm., 61, 668 (1951).
Fig. 3.-The relationship between the salt parameters for benzene and dissolved tertiary butyl chloride.
of two points, one for all 2-1 electrolytes and another for all 1-1 electrolytes). Instead, there is ordering of 2-1 and 1-1 electrolytes along separate approximately parallel lines. The order of anions along these lines is a common one which has been frequently observed in a number of phenomena associated with aqueous electrolyte solutions. Namely, i t is that associated with increased internal pressure of the water s t r ~ c t u r e ~ ~ ~ ~ ~ resulting from decreased ion size. Nearly the same quantitative order of anion effects has been observed in all previous s t u d i e ~ ' ~ . ~ ~ of aqueous electrolyte effects on the activity coefficients of non-electrolytes. This is illustrated in Fig. 3 which shows a typical correlation (of nearly unit slope) between the Setschenow param) t-butyl chloride eters for benzene ( K c ~ H ~and ( K g ) . Deviations from this correlation are again to be noted for organic anions and cations. Similar plots are obtained for the less abundant data on the salt parameters for any pair of the non-electrolytes: nitrous oxide, benzoic acid, aniline and tbutyl chloride. Linear correlations resembling Fig. 3 are also obtained in plots of K, (or K+) vs. partial molal absolute entropies,32 9,viscosity p-c~efficients~~ and shifts of the NMR proton resonance frequencies of electrolyte solutions.34 I n the correlation of K+ parameters and partial molal (31) H.S. Frank and W. Wen, Discussions Faraday Soc., 24, 133 (1957). (32) W. M. Latimer, "Oxidation Potentials," R e n t i c e H a l l , Inc., N e w York, N. Y., 1952,p. 368. (33) R. W.Gurney, ref. 25, Chapter 10. (34) J. N. Shoolery and B. Alder, J . Chon. Phys., as, 805 (1955).
THEACTIVITYCOEFFICIENT OF L-BUTYLCHLORIDE
June 20, 1062
entropies separate lines appear for monatomic and oxyanions. It is of interest t o note in connection with these various correlations that the order of anion effects on K* clearly does not parallel the order of nucleophilicities observed in alkyl halide displacement reactions, As shown by the correlation of Fig. 3, the salt order for t-butyl chloride follows that anticipated by the theory of McDevit and Long for the salt induced internal pressure effects on non-polar non-electrolytes. Important contributions of salt induced medium effects to the activity coefficient of t-butyl chloride and its solvolysis transition state are therefore clearly established by the present results. The fact that the correlation lines of Fig. 2 ( K , vs. K+) are approximately of unit slopes indicates that the salt induced solvent effects on the activity coefficients of t-butyl chloride and its solvolysis transition state are nearly equivalent. Accordingly, the medium effect should nearly disappear in the ratio of these activity coefficients, fg/f+. (A referee has pointed out that for perfect cancellation, AV* = 0 is implied. Since this is unlikely, the cancellation is an approximation, but one which is clearly useful (cf. section 5 ) in most practical applications.) This ratio of activity coefficients gives the salt effect on the conventional rate constant k, for solvolysis of dissolved t-butyl chloride a t unit concentration. That is kc = kco(fg/f+) and log (kClkco)= logf, - logf+, or log ( k c / k c o ) = ( K , - K+)Cs, where the superscript 0 refers to the reference state pure water. The values of Kg - K* are examined in the following section. (5) The Ion-Atmosphere Model and ( K , - K+) Values.-Table VI1 summarizes the values of K , K+ obtained for the various inorganic salts. The positive values of K , - K+ denote the logarithm of the rate acceleration in one molar salt solution.
23C1
(within experimental error) in quantitative agreement with the expectations of the ion atmosphere model. Eqn. 5 may be applied to both the ground and transition states for the t-butyl chloride solvolysis, giving log ( k o / k o o )
log (fa/f+) 4nNe' (Z+* d + 2.303(DkT)a(I O a )
[
=
- Zgpd,)]
nCa (7)
where nCs = p = ionic strength; n = 1 for 1-1 salt; n, = 3 for 2-1 salts.
*,2, = fractional charges of the dipole of the transition and ground states, respectively d *, d, = distance of charge separation in the transition and ground states, respectively
Z
Equating n times the quantity in brackets on the right hand side of eq. 7 with the quantity K,- K+ gives
The ratio of the average values of ( K , - K*) is 2.6 f 0.2, in striking agreement with the theoretical ratio. One may conclude (in accord with Bronsted'8) that the effects of the aqueous inorganic salts (of Table VII) on the activity coefficient of the aqueous t-butyl chloride transition state are due primarily to two causes: the ion atmosphere effect and the salt induced medium effect. That is,f+ S f+'.". f*i,*., where f+e.s. is the coulombic ion atmosphere contribution and f*'.P. is the salt induced internal pressure contribution. Similarly, for dissolved t-butyl chloride, f g g fge.s;fpi.p. (the fact that a salt of the 2-1 as well as the 1-1 type apparently falls on the correlation line of Fig. 3 suggests that fge.".is rather inconsequential). Since we deduce i t follows f+'.P., from Figs. 2 and 3 that fgi,p. that f g / f + E fge.s./f*e.s., for most practical purposes a purely coulombic ion atmosphere quanTABLE VI1 tity appropriate to treatment by the theoretical eq. SALTEFFECTPARAMETERS (K,- K+) FOR THE RATEOF 7. Substituting the average values of K , - K+ SOLVOLYSIS OF DISSOLVED I-BUTYL CHLORIDEAT UNIT from Table VI1 eq. 7 gives the average value of CONCENTRATION, WATER15" 1.44 X for the quantity (Z+2d+ - Z,*dg). 1-1 Salts If the transition state charge separation is taken KBr +0.21" = (0.19 + 0.02) LiNOa 0.24b (0.25 - 0.01) .245 = ( .43 - . l 9 ) NaF' KOH .2QC= ( .51 - .22) to be about 0.5 A. greater3j than the equilibrium NaC104 .2ga = ( .24 + .04) N a I .33' = ( .25 $ .OS) carbon chlorine bond distancea6of 1.76 A,, and the NaSCN .26" = ( ,115 + . l o ) KC1 .20' = ( .26 - .06) fractional charge 2, for the ground state is taken LiCl .26* = ( .29 - .03) NaOH .2lC = ( .41 - .20) to be about 0.2 of an electronic charge,37 the 1NaCl .2Qb = ( . 3 6 - .06) K F .3lC ( .51 - .20) butyl chloride transition state is found to behave Av. 0.25 f 0.02d (in terms of eq. 7) as if about 0.8 of a n electronic 2-1 Salts charge is localized on the carbon and chlorine atoms. NasCrrCh $0.74" = (0.64 4- 0.10) NarSO, .61b = ( .88 - .27) This result is relatively insensitive to small values NatSzOd .62) = ( .87 - .25) chosen for 2,. The result is, of course, only of Kit304 .65' = ( .9S - . 3 3 ) qualitative significance in view of the delocalization NatCOs .55' ( .96 - .41) of positive charge into the methyl groups of the Av. +0.66 I 0.05d t-butyl "cation" and accompanying uncertainty a Estimated max. uncertainty i 0.06. Estimated of the significance of the d+ distance. (A referee Estimated rnax. uncertainty f max. uncertainty =k 0.08. has appropriately pointed out that a theoretical 0.13. Based upon (a) and (b) only. treatment more rigorous than that of the point It is clear from Table VI1 that the specificities dipole, such as the cavity model treatment of in K , and K* essentially disappear and all of the (36) A. Streitwieser, Jr., Chem. Rem., 56, 571 (1956). values K , - K* collect into two distinct groups, (36) J. H i m , "Physical Organic Chemistry," McGraw-Hill Book one for 1-1 salts (avg. value 0.25) and the other Company, Inc., New York, N. Y., 1956, p. 30. for 2-1 salts (avg. value 0.66). This result is (37) R. W. Taft, Jr., J . Chem. Phys., 26, 93 (1967).
-
-
-
++
GEORGEA. CLARKE AND ROBERT W. TAFT
2302
VOl. 81
Kirkwood ( J . Ciwm. Phys., 2 , 3.51 (1931))) would also undoubtedly give a somewhat different figure.) The reaction constant,38p * , for the solvolysis in water of t-carbinyl chlorides, XCH2CCl(CHS)z (X is a polar substituents*) has been found to be -3.8. Apparently p* values of this sign and magnitude may be associated (in water) with the kind of charge separation process indicated by the kinetic salt effect. ( 6 ) Specific Electrolyte Effects.--k series of electrolytes has been investigated which give ( K , - K+) values which do not conform to the ion atmosphere behavior exemplified by the salts of Table VII. The deviant behavior was found for the acidic cations, ammonium and hydronium, and for organic electrolytes. These results are listed in Table VIII. All of the electrolytes of Table VI11 show appreciably smaller rate accelerations ("03 apparen tly produces a small rate retardation) than the inorganic salts of the same charge type in Table VII. The ( K , - K+) values for the ammonium and alkylammonium salts are apparently independent of the anion (avg. value for the former is 0.15 and for the latter 0.01). The carboxylate ions as a class also appear to give approximately the same value (avg. 0.10). The values for mineral acids, on the other hand, show an apparent dependence on the anion.
salt effects of inorganic electrolytes. Further, Grunwald and Butlerz1 have noted that although the inorganic salt order for naphthalene is dramatically inverted between water and 50% (wt.) dioxane-water as solvent, organic electrolytes salt-in in both media. They conclude that specific short-range (direct) dispersive interactions take place between non-electrolyte and organic electrolytes. Our results are consistent with this interpretation. Since these interactions are apparently less important for the t-butyl chloride transition than ground state, i t is clear that quotient fg/f+ for organic electrolytes cannot be a purely coulombic ion atmosphere quantity. We turn now to a consideration of the acidic cations, H30+ and NH4+. The electrolytes of these cations do not show notable deviations from plots such as that shown in Fig. 3. One may infer then that there is no special effect of these cations (either direct or indirect) on dissolved t-butyl chloride. On the other hand, the only notable cation effects on the t-butyl chloride transition state (cf. Table V) are those for the salting-in exhibited by these ions, especially HeO+. The importance of front-side participation by hydroxylic solvents as a driving force in the solvolysis of t-butyl halides has been recognized since the conception of the SN1 mechanism.' Hydrogen bonding by water to the incipient chloride ion in the transition state is definitely indicated by the findTABLE VI11 ing2* that the solvolysis rate k, gaseous t-butyl ELECTROLYTE EFFECTS ( K R - K*) ON THE RATE OF chloride a t unit pressure is reduced in D?O comSOLVOLYSIS OF DISSOLVEDL-BUTYLCHLORIDEAT UNIT pared to HzO by essentially the same factor (0.72) CONCENTRATIOS, WATER,15'" as is the conventional rate constant kc. Thus the Carboxylate salts Ammonium salts Mineral acids enhancement in the rate constant k , exerted by the SaOzCH $0.10 "4Cl +0.16 HK03 -0.05 acidic cations may be associated with the transi+ .I4 tion state hydrogen bonding effect.4O Since these HCl $ .07 Ka02CCCH6 $ . I 5 ?rTH4I .02 cations do not enter the rate equation, their effect HCIOa + .14 NaOzCCH3 + .08 (CH3)aNOH H2SOa + .43* Na02CCF3 + .06 (CH3hSBr f .02 is presumably an indirect one in which an increase (C2H&)&OH ,00 in the concentration of the acidic cations increases 5 Estimated maximum uncertainty + 0.06 unless otherthe ability of the solvent to enter into the hydrogen wise indicated. b Estimated maximum uncertainty bonding interaction. Since the proposed effect 0.08. is specific to the t-butyl chloride transition state, We shall consider first the possible cause of the it must also follow for the acidic cations that f g if+ deviant behavior of the organic electrolytes. I n cannot be a purely coulombic ion atmosphere Fig. 3 and similar plots, these electrolytes deviate quantity. (7) Regarding the General Cancellation of substantially from the linear correlation shown results of by inorganic electrolytes (including, incidentally, Salt-induced Medium Eff ects.-The the acidic cations, H30+ and NH4+). Benzene is Table VI1 are in distinct contrast to those resubstantially more salted-in by these organic ported11-16 in which quite specific salt effects have electrolytes than is t-butyl chloride. Further, on been observed on the rates of solvolysis of dissolved the one hand, the t-butyl chloride transition state substrate a t unit concentration. The later results is virtually unaffected by the alkyl ammonium ions all have the common feature that a mixed aqueous relative to the Na+ ( c j . Table V) and the K+ organic solvent was employed. It is perfectly salt parameters for the carboxylate anions (cf. clear from the detailed treatment of Grunwald Table VI) also tend to be small, whereas, on the and Butler2' that salt induced medium effects are other hand, Bergen and L o n g 2 0 a have found distinctly more complex in mixed than pure solthat phthalic acid and salicylic acid, for example, vents of high dielectric constant. These authors are even more strongly salted-in solution by organic show that a t least one more term (and probably electrolytes than is benzene. It is thus abundantly others) must contribute to the salt induced medium clear that the susceptibility to salting-in by the effects in mixed solvents. I t is therefore expected organic electrolytes varies enormously more among that in mixed solvents the salt induced medium efvarious non-electrolytes than does that for the fects will not cancel between ground and solvolysis
+
+ +
+
*
(38) R. W, T a f t , Jr.. in M . S. Kewman, "Steric Effects in Orgenic Chemistry," John Wiley and Sons, New York, N. Y., 1956, Chapt. 13. (39) R. H. Boyd, Ph.D. Thesis, The Pennsylvania State University, 1959.
(40) For recent evidence of the importance of hydrogen bonding interactions on activity coefficient behavior cf. (a) R. W. T a f t , Jr. J . Am. Chem. SOL.,8 3 , 2965 (1960); (b) N . C. Deno and H. Berkheimer, J . Chcm. Eng. Dolo, 6, 1 (1960).
TRIVALENT METALBENZOYLACETONATES
June 20, 1962
2303
transition states (because the salt effect on prefer- of the dipolar transition state (i.e., through couential solution will differ for these two states). I t lombic interactions). Since the internal pressures is in fact quite interesting that the 1-1 electrolyte of the solvents investigated vary widely, the effects for the rate of solvolysis of neophyl p-tolu- partial molal free energy of the dissolved reaction enesulfonate in 50% (wt.) dioxdne-waterle parallel substrate must also vary widely.4z The success quite closely (a roughly linear plot results) corre- of the treatment implies a t least substantial cansponding salt effects, K+ (or Kgbut not Kg-K*) for cellation between ground and transition states of t-butyl chloride solvolysis in water. the non-coulombic contributions to the partial molal free energy of these two states. That is, ( 8 ) A Comment on Pure Medium Effects.Smith, Fainberg and Winstein41 have recently there is apparently involved a cancellation of inutilized the conventional rate constant for solvolysis ternal pressure effects between ground and transiof p-methoxy-neophyl p-toluenesulfonate as a tion states which leaves (for most practical purmeasure of the ionizing power of a solvent. A close poses) the rate constant as a coulombic quantity, correspondence of the logarithm of the rate con- a situation analogous to the cancellation reported stant of pure polar non-hydroxylic solvents with in this paper for kinetic salt effects in pure water. other independent measures of ionizing power is The complication of kinetic salt effects in mixed found. The results support the notion that these solvents appears to support the conclusion of the rates are governed by the ability of the solvent to above authors that i t can be less certainly accepted support a charge separation through non-specific that the effects of mixed solvents on the solvolysis induced polarization of the effective dielectric rates are purely coulombic quantities. constant of the medium by the local electric fields (42) J. H. Hildebrand and R L Scott, “Solubility of Non-Elec(41) S. G. Smith, A. H. Fainberg and S. Winstein, J . A m . Chem. .Yoc., 83, 618 (1961).
[CONTRIBUTION FROM
THE
WILLIAM
ALBERT
trolytes,” 3rd Ed., Reinhold Publishing Co., L-ew York, N. Y., l R Z O , pp. 46, 94, 129, 435.
NOYESLABORATORY, UNVERSITY O F ILLINOIS, URBANA,
ILLINOIS]
Coordination Compounds of Trivalent Metals with Unsymmetrical Bidentate Ligands. I. Benzoylacetonatesl BY ROBERTC. FAY^
AND
T. S. PIPER
RECEIVED OCTOBER 23, 1961 The preparation of benzoylacetonates of a number of trivalent metals is described ( M = Cr, Co, Rh, Al, Mn, Fe). cis and trans isomers of the inert complexes ( M = Cr, Co, Rh) were separated by utilizing solubility differences in ethanol and column chromatography on alumina. Somewhat less than the statistical amount of cis isomer was obtained. In the case of the labile complexes ( M = Al, Mn, Fe), only the more stable trans isomer could be isolated. The configuration of the diamagnetic complexes was determined by nuclear magnetic resonance; the configuration of the paramagnetic complexes was established on the basis of isomorphism of crystals of the diamagnetic and paramagnetic complexes. Small intensity differences in the electronic spectra suggest that the forbidden d-d transitions obtain intensity by mixing in odd molecular orbitals rather than odd atomic orbitals. Cis to trans isomerization of the cobalt and chromium complexes takes place a t elevated temperatures. The potential importance of n.m.r. in stereochemical studies of metal coordination complexes is emphasized.
Introduction Coordination compounds in which the central atom is bound to three unsymmetrical bidentate ligands may exist in two geometrically isomeric forms. If the distinguishable ends of an unsymmetrical bidentate ligand are represented by X and Y (see Fig. l), the X groups may be in the 1,2,3- or 1,2,6-positions. We shall refer to the isomer with X groups in the 1,2,3-positions as the cis isomer and the isomer with X groups in the 1,2,6-positionsas the trans i ~ o m e r . ~ The majority of the work reported on isomeric tris-complexes has been concerned with the trisamino acid chelates of cobalt(II1). The glycine and alanine complexes have been known for a long time,4 but only recently has the configuration of the two isomeric glycine complexes been de-
termined on the basis of visible absorption spectra.s Carassiti and c o - w o r k e r ~ using ~ , ~ the same technique have determined the configuration of the alanine complexes and also a number of other amino acid chelates of cobalt(II1). A tris-(ethanolamine)-cobalt(II1) complex has been isolated and assigned the cis configuration, also on the basis of absorption spectra.8 L’ery recently two tris(8-quinolinolo)-chromium(III)complexes (one solvated) have been assigned cis and trans structure^.^ The assignment, however, must be viewed as only tentative because of the lack of experimental evidence. One class of compounds which ought to give geometric isomers and which has not been previously investigated are tris-complexes of trivalent
(1) Presented in part a t the Symposium on Magnetic Phenomena in Inorganic Chemistry, 140th National Meeting, American Chemical Society, September, 1961, Chicago, Illinois. (2) National Science Foundation Fellow, 1960-1962. (3) W. C. Fernelitis, Adv. in Chem. Series, 8 , 9 (19.53). (4) H. Ley and H. Winkler, Ber., 44, 3894 (1909); ;bid., 45, 372 (1912); H. Ley and K. Ficken, ibid., SO, 1123 (1917).
( 5 ) F. Basolo, C . J . Ballhausen and J. Bjerrum, Acta Chcm. Scand., 9, 810 (1855). (6) V. Carassiti and M. Claudi, A n n . Chim. (Rome),SO, ,581 (1960). (7) V. Carassiti and A. M. Marinangcli. i b i d . , SO, ,593 (1960). (8) H. Yoneda and S. Kida, J . A m . C h c m S o c . , 34,?I39 (1960). (9) F. Umfand, G. H. Gudmundsson and K. Adam, Nn6urwisscnschuffcn, 48, 49 (1961).
