1244
The Journal of Physical Chemistry, Vol. 83, No. 10, 1979
A. M.
Afanassiev, K. Okazaki, and G. R. Freeman
Effect of Solvation Energy on Electron Reaction Rates in Hydroxylic Solvents' Alexei
M. Afanassiev,
Kiyoshi Okazaki, and Gordon R. Freeman"
Chemistry Department, University of Alberta, Edmonton, Alberta T6G 2G2, Canada (Received January 22, 1978) Publication costs assisted by the University of Alberta
-
The rate constant of esol; + S product (1)for inefficient reactants S is strongly dependent upon the solvation energy of electrons in the solvent. Values of k l for inefficient scavengers, such as 4-penten-2-01,were two orders of magnitude greater in 2-methyl-2-propanol than in methanol; kl increases in the order CH,OH < C2H50H < 2-C3H70H< t-C,HgOH, and the electron solvation energy changes in the opposite direction. The relative differences of kl for efficient scavengers are much smaller, and the above solvents fall in the reverse order; kl for nitrobenzene increases in the order t-C4HgOH< 2-C3H70H< C2H50H< CH,OH, in agreement with the relative fluidities of the solvents. Measurements in water, ethylene glycol, and l-propanol indicate that the relative solvation of S also affects the rate. The free energies of activation AGA* of a series of scavengers in an alcohol A were compared with those in methanol, AGM*. The difference (AGM* - AGA*)reaches a limiting value [AAGI],, for inefficient scavengers. The value of [AAG'],, corresponds to the difference of free energies of activation for reaction of electron/scavenger encounter pairs in methanol and alcohol A. The values of [ AAG*],,, vary approximately linearly with the electron solvation energy.
Introduction The reactivity of electrons depends on a number of factors, such as the nature of the c o r e a ~ t a n t , the ~ - ~nature of the s ~ l v e n t , ~ the - ~ energy state of the electron,lk12 temperature,13J4 and phase.I5 All of the factors relate to the transport properties and relative energy levels of the reactants and (intermediate) products. When reactants have a large affinity for each other the rate of their reaction is limited by the rate a t which they diffuse together. The change of rate upon changing the solvent reflects mainly the change in the diffusion coefficients. However, when reactants have a small affinity for each other the reaction rate can be greatly affected by the relative solvation energies of the reactants and products.16J7 The relative contributions of transport and solvation properties are illustrated by the fact that, for efficient scavengers such as HBO1,+ in methanol and ethanol kM/kE > 1.0, whereas for inefficient scavengers such as ~ k M and k~ are the rate constants of benzene k ~ / 1.0 is due t o the greater fluidity of methanol; kM/kE < 1.0 is due to the lower solvation energy of electrons in ethanol. The present work extends the investigation of solvent effects in electron reaction rates. Data are included for C1 to C4alcohols, ethylene glycol, and water. The reactants include a number of unsaturated alcohols. Experimental Section Materials. Solvents. The following were the suppliers and grades of the solvents: methanol, Fisher Scientific Go., spectroanalyzed; ethanol, US Industrial Chemical Co., absolute reagent, anhydrous; l-propanol, Aldrich Chemical Co., analyzed 99+ 70; 2-propanol, Matheson Coleman and Bell, spectroquality, or Aldrich, spectral grade, 99+ 70; 2-methyl-2-propano1, Aldrich, 99.5% ; and ethylene glycol, Fisher Scientific Co., 99%. The ethanol was used as received,ls but was kept under an atmosphere of ultra-high-purity argon (Matheson Co.). Ethylene glycol was distilled under reduced pressure a t -100 "C. The first 15% and last 35% fractions were 0022-3654/79/2083- 1244$01.OO/O
discarded, and the middle 50% was kept under an atmosphere of ultra-high-purity argon. Methanol, 1-and 2-propanol, and 2-methyl-2-propanol were treated for 1 day under argon with sodium borohydride (1g/L) at 40,60,50, and 50 "C, respectively. The alcohol was then fractionated through an 80 X 2.3 cm column packed with glass helices. The middle 5070 was collected and kept under argon. An argon pressure and syphon system was used to transfer the alcohol without contacting air. Water was triply distilled.l8 The solvated electron half-life after a 0.1-ys pulse of -1 X 10l6 eV/g a t 25 "C was -20 ps in water, 3.8 ps in methanol, 7.7 ps in ethanol, 5.4 ps in l-propanol, 3.5 and 17 ys in 2-propanol from Matheson Co. and Aldrich Co., respectively, 14.5 ps in 2-methyl-2-propano1, and 0.65 ps in ethylene glycol. Solutes. Allyl alcohol (99%), propargyl alcohol (97%), 3-butene-1-01 (99%), and 4-pentene-2-01 (97%) were analyzed grade from Aldrich Co. Acetone was 99.5% certified grade from Fisher Scientific Co. Propionaldehyde was from Eastman Kodak Co., nitrobenzene (analyzed reagent) from Baker Co., carbon tetrachloride (99+%, spectroquality) from Matheson Coleman and Bell, and perchloric acid (70%) from Baker and Adamson. Acetone, carbon tetrachloride, nitrobenzene, and perchloric acid were used as received. Propionaldehyde was distilled under argon and the middle, colorless fraction kept. Unsaturated alcohols were treated with sodium borohydride and distilled as the solvents above. Allyl and propargyl alcohols were treated twice. The borohydride treatment was at 65 "C for allyl alcohol and 40 "C for the others. All reagents were taken from freshly opened bottles or ampules. Techniques. The methods of sample preparation,16319 irradiation, dosimetry, and optical measurement2"2 were essentially the same as described in the indicated references. Solutions of the more reactive scavengers were prepared by injecting 1-50 pL of a stock solution into 5 mL of solvent, under air-free (argon-saturated) conditions. The stock solution had been previously deaerated by bubbling
0 1979 American Chemical Society
TABLE I: Values of h , ( l o 8M-I s-l) in Different Solvents at 298
-
no. 1 2
scavenger ___.4-pentlen-2-01 3-buten-1-01
H*O
0.00s O.OOEI
0.041" 0.20 0.12" 2.1 41 72+5 5 9 i 2b 95c 65d 130 30Of
3
allyl alcohol
4 5 6
propargyl alcohol propionaldehyde acetone
7
carbon tetrachloride
8
nitrobenzene
390 420d
9
perchloric acid
240.i: 20b 220'
1245
The Journal of Physical Chemistry, Vol. 83, No. 10, 1979
Electron Reaction Rates in Hydroxylic Solvents
2K
i
__.
CH,OH 0.009 0.008
C2H50H 0.033 0.032 0.036'" 3.8 3.4'" 16 36 52 47a3e 50 f 2OC
0.67 3.9 25 44 4 3"
110 130" 12OC
160
1-C3H,0H
2-C3H,0H t-C,H,OH 0.71 0.14 0.71 0.20
0.050 0.053
(CH,OH),
0.018 0.01
4.2
9.1
14
0.73
12 25 36
19 35 56
18 36 60
3.2
80
120
150
21 120 i: 4OC
110
130
110
30
120
59
9.3 50 i: 2OC
11og
8.9 16 6 f,In
230 230e 19O.i 40h 650: 5401
150 120 i: 4OC 15Odle 270 360' 33w
150 * 7OC
5.4 188m 190" 144m
10.8 43 173 19.2 19.4 172m 211p llop 185m 146"J' 171" 930 161"J' 208" 131m 12gm 164p 8P 132m 108p a -293 K, ref 17. Room temperature, concentrated solutions, ref 24. 294 K, ref 2. 298 K, ref 25. e 296 K, ref 293 K, ref 18. 293 K, ref 26. 293 K, ref 27. 295 K, ref 14. Viscosity, 298 16. f Floom t,emperature, ref 23. K, ref 28. Elnergy at the maximum of the optical absorption spectrum of solvated electrons, 298 K. Reference 20b. " Reference 29. Reference 32. ]Reference 30. 4 E, = (E,,, - W,) is the energy on the red side of the absorption band at half the maximum absorbance. J
it with argon that had been passed through a solution of the same strength. The solvent was ethanol, so the final solutions of perchloric acid, carbon tetrachloride, nitrobenzene, and acetone contained 0.02-0.2 vol % ethanol, and those of propionaldehyde contained 0.1-1 % ethanol. Although propionaldehyde was stable in water, it tended to form hemiacetals and acetals in the alcoholic solutions. The concentration of aldehyde in the stock solutions was therefore measured directly, by optical absorption at X 280 nm. The radiation pulse was 100 ns of 1.8-MeV electrons, delivering a dose of 1-2 X 10I6 eV/g. Solvated electrons in hydroxylic solvents can react as follows: esolv-+ S esolv-+ I
-
1
2.0
1.5
-2
c Iv)
Y.
Results and Discussion
-
-
product
(1)
product
(2)
ROsolv- + H
(3) where I is a reactive impurity, and the esol; structure is understood t o contain the ROH required for material balance in (311. The observed first-order rate constant is ,k , = kJS1 -t. kz[I] t k3 The value of kl for a given solute S is obtained from the slope of a plot of kohsd against solute concentration. Four to seven different concentrations were used in each case. The nature and quality of the results are illustrated for an inefficient scavenger in Figure 1 and an efficient scavenger in Figure 2. The values of klfor these and other scavengers aire listed in Table I. Values from the are included for comparison where possible. Agreement is satisfactory, with the exceptions of 3 - b ~ t ~ e n - l -and 0 1 ~carbon ~ t e t r a ~ h l o r i d ein~water ~ and all the earlier rate constants in ethylene In these cases the earlier values seem to be too high by factors of 2-6. For example, the viscosity of ethylene glycol is 16-fold es01v-
I
_Yo
1.0
0.5
0
Axis I
Axirn
001 [3- Buten-1 -
-011
002
003 AxisIII
(M)
Flgure 1. Pseudo-first-orderdecay constants of solvated electrons in 3-buten-1-01 solutions at 298 f 1 K in different solvents: (axis I) +, methanol; A, ethylene glycol; (axis 11) 0 , water; 0, ethanol; A, 1propanol;,. 2-propanol; (axis 111) 0 , 2-methyl-2-propanol.
greater than that of ethanolz8and the electron trap depth is greater in the g l y c 0 1 , ~so~ it~ ~is~highly improbable16p31 that the rate constants for solvated electron reaction with acetone are equal in the two solvents as reported earlier.24
1246
The Journal of Physical Chemistry, Vol. 83, No. 10, 1979
A. M. Afanassiev, K. Okazaki, and G. R. Freeman
[C6H5 NO,] ( 10-4M)
17
/ 0
6
10
log (k, /M-’S’)
I
5
IO
[C6H5 NO,] (10-5M) Figure 2. Pseudo-first-order decay constants of solvated electrons in nitrobenzene solutions at 298 f 1 K in different solvents: 0 , water; +, methanol; 0, ethanol; A, l-propanol; W, 2-propanol; 0 , 2methyl-2-propanol; A, ethylene glycol. Values of the solvent viscosityzs and solvated electron optical absorption energy21~29~30~32 are included in Table I for reference. Soluent Effect. The rate constant for solvated electron reaction with a given scavenger can vary greatly from one solvent to another. For example, the values of hl for the olefinic alcohols are two orders of magnitude greater in t-C4HgOHthan in water; kl increases in the order H20< CH30H < (CH20H)z < C2H50H < l-C3H70H < 2C3H70H< t-C4HgOH(Table I). The relative differences for more reactive scavengers are much smaller and the solvents fall in a different order. For example, the values of k, for nitrobenzene increase in the order (CH20H), < t-CdHgOH z=z 1-C3H70H 5 2-C3H,OH 5 CZHbOH < CH30H < H 2 0 (Table I). To illustrate the solvent effect more clearly methanol was chosen as the basis for comparison, because it is the simplest alcohol and has the lowest viscosity (Table I). The ratio of the rate constant in methanol, kM, to that in another alcohol, h ~was , calculated for each scavenger. Then hM/kA was plotted against kM. The plot for methanol/ethanol, hM/kE,is shown in Figure 3. Data from Table I and ref 14, 16, 22, and 33 are included; ratios are only taken of rate constants determined by the same authors. This plot is a log-log representation of Figure 4 of ref 16, with more data added. For inefficient scavengers hM/kE = 0.2. The value less than unity has been attributed16 to the greater solvation energy of electrons in methanol, which makes it more difficult in methanol than in ethanol for an electron to jump from a solvation site to a scavenger molecule of low electron affinity. The solvation energies are reflected in the optical absorption energies E,, of the solvated electrons (Table I). For progressively more efficient scavengers, that is, those with less negative or more positive electron affinities, the value of kM/kE increases. The ratio passes through unity at k M = 7 X lo9 M-ls-l (Figure 3). The most efficient scavengers have such
Figure 3. Ratio of rate constants for a given scavenger in methanol and ethanol, k,lkE, plotted against k,. The numbers 1-9 represent the scavengers in Table I; 10 is 11 is 12 is phenol;14qi613 is SFB;”14 is oxy en;” 15 is nitrous oxide;” 16 is carbon dioxide;” 17 is 1,3-butadiene;’B 18 is acetonitrile;” 19 is o-yxlene;” 20 is naphthalene.’e~aThe asterisk is the assumed diffusion-controlled limit for a neutral scavenger. The curve corresponds to that in Figure 6, by way of eq 7.
21
I
5
I
IO
I
20
I
I
I
50
I I I I I
loo
1
200
7) (mP) Figure 4. Plot of k l for nitrobenzene (nearlydiffusion controlled)against the solvent viscosity. Data from Table I.
large electron affinities that the solvated electron trap depth is unimportant; the rate constants are then limited by the diffusion coefficients of the reactants. Electrons solvated in alcohols are in sufficiently deep traps that their diffusion rates are an inverse function of the solvent viscosity q.31934 Thus for efficient scavengers hM/kE > 1.0 because qM/qE < 1.0 (Table I). The dependence of kl on 7 for an efficient scavenger, nitrobenzene, in the alcohols is shown in Figure 4. The equation of the line is kl(M-l s-l) = 62 X 109q4’59, where q is in mP. The viscosity exponent reported by CerEek for a number of alcohols and glycerol/water solutions was -0.66.31 Molecular diffusion coefficients in alcohols are roughly proportional to so the fact that kl 74.6 for an efficient scavenger indicates that the diffusion coefficient of solvated electrons in alcohols varies approximately as V-O.~. Electron mobility measurements in a number of alcohols at 293 K lead to a similar conclusion.M Pure water as solvent does not fit the correlation in Figure 4 because the diffusion coefficient of solvated electrons in water is greater than that in an alcohol of the same viscosity.34 Plots of log (kM/kA) against log k M for other alcohols (Figure 5) have shapes similar to that for ethanol (Figure
The Journal of Physical Chemistry, Vol. 83,
Electron Reaction Rates in Hydroxylic Solvents
I
-4 10
I
I
I
I
I
20
AGi
No. 10, 1979 1247
I
I
I
I
' 1
I
40
30 (kJ/mol)
Flgure 6. Difference of free energies of activation in methanol, AGM*, and ethanol, AG;, plotted against AGM*.The numbers represent the same as those in Figure 3. 10
8
6
log ( kM/M-'s-')
Figure 5. Ratio of rate constants for a given scavenger in methanol and another alcohol, k , / k , , plotted against k,. The alcohol A is as follows: 0 , ethylene glycol; 0, ?-propanol; A , 2-propanol; 0 ,2methyl-2-propanol. The numbers represent the same as in Figure 3. The curves correspond to those in Figure 7, by way of eq 7.
3). However, there is an upward or downward displacement that is more severe for the inefficient than for the efficient scavengers. The ratio kM/kA for inefficient scavengers decreases with decreasing solvation energy of the electrons in A (see Errnuvalues in Table I). The ratio for efficient scavengers increases with the viscosity of A (Figure 5 and Tablle I). The value of log (kM/kH20) hovers near zero over the range log k M = 6-1 '1, with considerable scatter. The average value and mean deviation of log (kM/hH20) is 0.04 f 0.23 for the 1.4 compounds that have been measured in the two solvents by the same authors. Energetics. Electron capture by a scavenger in the liquid phase is accompanied by solvent rearrangement about the reactant pair as follows:
e&
S-*+ solvent
+ S + S-.*
-
Ssol; or product
(4,-4)
(5)
The overall capture rate constant k l is
kf = k4k5/(k-4
+ k5)
(6)
The reaction rate can be limited by either or both of (4,-4) and (5). However, the process of solvent relaxation (eq 5 ) probably occurs at a rate similar in magnitude to the intermediate relaxation time 72 in the alcohol,16which is s at -300 IK in simple alcohols37 and ethylene The rate limiting step for most of the present scavengers is therefore probably (4). One may write kl = ( k T / h ) exp(-AG1*/RT) (7) where AGl* =s AG4* is the free energ.1 of activation. At 298 K the value of k T / h is 6.2 X 10l2,which is convenient for the preexpoinential factor of kl in units of M-I for a neutral scavenger. For example, lit gives an estimated diffusion-controlled limit for neutral S in methanol to be 6 X 1O1O M-l s-l with AG1* = 11.5 kJ/mol. The latter is the activation energy of self-diffusion in methanol.39 Values of AGl* were calculated for the neutral solutes in Figure 3 and are denoted AGM* for methanol and AGE*
for ethanol solutions. A plot of (AGM* - AGE*) against AGM* is shown in Figure 6. For discussion purposes kl and AG1*may be divided into two portions, one representing the diffusion together of the reactants, kd and AGd*, and another representing reaction after the reactants have formed an encounter pair, k, and AG,*, respectively. kl-l = hd-l kL1 (84
+
exp(AGl*/12T) = exp(AGd*/RT) + exp(AG,*/RT) (8b) In the diffusion controlled limit AG? = 0 and AG1* = AG:. Taking AGd,M*= 11.5 kJ/mol, Figure 6 indicates that AGd,e* = 14 kJ/mol. For less reactive solutes AG,* contributes to limiting the reaction rate. When AG,' < AGd* one has AG1*= AG: AG,", with the contribution of AGd* decreasing as AG,' increases. When AG: >> AGd* one has AG1* = AG;. On going from very reactive to less reactive solutes, the free energy of activation increases more rapidly in methanol than in ethanol because the electron has to escape from an approximately 17 kJ/mol deeper solvation trap in the former (see the values of E,,,, in Table I). However, it is the electrons in the shallower traps, represented by the red side of the optical absorption band, that have the greatest probability of reacting per unit time. For this reason we choose E, = (E,,,, - W,),the energy on the red side of the band at half the maximum absorbance, as a more appropriate estimate of the relative trap depths (Table I); the difference for methanol and ethanol = 15 kJ/mol. Furthermore, the electron is Er,v solvation energy is partially compensated by interactions between the solvent and S-*,so the maximum difference between AGM* and AGE* is less than 15 kJ/mol. The curve in Figure 6 is described by (9), in units of kJ/mol. (AGM* - AGE*) = 0 . 3 8 A G ~ '- 6.5 AGM* < 28 (9a)
+
(AGM* - AGE*) = 4.2
AGM* > 28
(9b)
Equation 9 is equivalent to eq 18 in ref 16. Thus the maximum difference between AG1* in methanol and ethanol is 4.2 kJ/mol. Values of AG1* were also calculated for the neutral scavengers in the other alcohols (Table 11). The differences (AGM' - AGA*) are plotted against AGM* in Figure 7; AGA* represents AG1* in the alcohol specified in the figure legend. The curves are of the same general shape as that in Figure 6. The diffusion-controlled limit of (AGM* - AGA*) is estimated by a linear extrapolation of the curve to AGM* = 11.5 kJ/mol.
1248
The Journal of Physical Chemistry, Vol. 83, No. 10, 1979
A. M. Afanassiev, K. Okazaki, and G. R. Freeman
TABLE 11: Solvent Dependence of the Free Energy of Activationa AG,
a
no.
scavenger
H,O
CH,OH
C,H,OH
1 2 3 4 5 6 7 8
4-penten-1-01 3-buten-1-01 allyl alcohol propargyl alcohol propionaldehyde acetone carbon tetrachloride nitrobenzene
41 40 32 26 18 17 15 13
39 40 29 24 19 18 15 14
36 36 24 21 19 18 16 15
*, kJ/mol
l-C,H,OH 2-C,H,OH t-C,H,OH 35 35 24 21 19 19 17 16
32 '32 22 20 19 17 15 16
28 28 21 20 19 17 15 16
(CH,OHr 37 39 28 25 22 20 19
Based on eq 7.
TABLE 111: Values of Constants in Eq 10 A
a
b, kJ/ mol
ethanol 1-propanol 2-propanol 2-methyl-2-propanol ethylene glycol
0.38 0.47 0.55 0.60 0.41
6.5 8.8 9.3 10.1 11.0
c,
kJ/ [AAG*],,, mol kJ/mol 28 28 32 35 30
4.2 4.3 8.2 11.0 1.1
d c 0 7 I
I1
I
I
I
I
I
I
I
+
1 1
7
1
n
3
4
6 5
\
r e
g 4
SCAVENGER NO. 12
-
0
a a
I
U
0 -20
0
20
40
60
AEr (kJ/mol) 0
Figure 8. Maximum difference between the free energies of activation for electron reaction with a weak scavenger in methanol and another alcohol, [AAGI],, plotted against the difference in electron trap depths, E, = (Er,M- Er,A).See eq 10 and 11.
-L
and follows eq 11 (see Figure 8), in units of kJ/mol. The fact that the line does not pass through the origin indicates that another factor affects the correlation.
4
7
Y I
+a
[AAG*],,, = 3.1 + 0.13AE,
0
a I
+c5
0
a
The compounds that give rise to [AAG*],,,, and fall along the horizontal part of the curve in Figure 6, are all unsaturated. In each case the product anion is unstable and is probably stabilized by protonation, as occurs with b e n ~ e n e . ~ OThe ~ ~ lreaction sequence may then by represented by ( 12,-12) es01v- + s F= %01;
0
-A
10
20
AG;
30
40
(kJ/mol)
Figure 7. Difference of free energies of activation in methanol, AGM*, and other alcohols, AGA*, plotted against AGM*. The numbers and symbols represent the same as those in Figure 5.
The curves in Figure 7 are described by eq 10; the values of the constants for each alcohol are listed in Table 111. (AGM*- AGA*) = uAGM* - b (AGM'
(11)
-
AGA*) = [AAG*],,,
AGM* < c (loa)
AGM*> c (lob)
The limiting value [AAGI],, corresponds to (AG,,M* AG,,A*), the maximum difference of free energies of activation for reaction of the electron/solute encounter pairs in methanol and alcohol A. It is affected by the electron solvation energies, as mentioned earlier. A plot of [AAG*],,, against the difference of electron trap depths, is approximately linear represented by AE, = (Er,M -
S801c + ROH ----* SH + RO,ol[ When klz>> k13, kl becomes kl = k12k13/k-12 = k13 exp(-AG12"/RT) = ( k T / h )exp(-[AG13*
+ AGl2']/RT)
(13)
(14)
The positive intercept in (11)could arise from AGl3' being much smaller in ethylene glycol than in methanol. The flat tops of the curves in Figures 6 and 7 imply that the ratio of protonation rates of the unsaturated anions is dependent on the solvent, but is relatively independent of the particular anion in the present series. Acknowledgment. We thank the staff of the Radiation Research Center for help with the electronics.
References and Notes (1) Assisted financially by the National Research Council of Canada and
Canada Council.
(2) S. Gordon, E. J. Hart, M. S. Matheson, J. Rabani, and J. K. Thomas, Discuss. Faraday Soc., 36, 193 (1963). (3) . . I. A. Taub, M. C. Sauer, Jr., and L. M. Dorfman, Discuss. Faraday SOC.,36,.206 (1963).
The Journal of Physical Chemistry, Vol. 83, No. 10, 1979
Enthalpies of Solution and Transfer Enthalpies (4) J. H. Baxendale, E. M. Fielden, C. Capellos, J. M. Francls, J. V. Davies, M. Ebert, C.W. Gilbert, J. P. Keene, E. J. Land, A. J. Swallow, and J. M. Nosworthy, Nature (London), 2011, 468 (1964). (5) L. M. Dorfnnan and M. S. Matheson, Rcg. React. Kinet., 3, 237 (1965). (6) K. Horacek and G. 13. Freeman, J. Chem. Phys.. 53, 4486 (1970). (7) M. G. Robinson and (3. R. Freeman, J. Chem. phys., 55, 5644 (1971). (8) M. G. Robinson and G. R. Freeman, Can. J. Chem., 51,650 (1973). (9) K. Ito and Y. Hatano, J . Phys. Chem., 78, 853 (1974). (10) P. L. T. Bwan annd W. H. Hamill, Trans Faraday SOC.,66, 2533 (1970). (11) D. Raiem and I. Dvornik, Proc. Tihany Symp. Radiat. Chem., 3rd (1972). (12) 0. Bakale, U. Sowadh, and W. F. Schmidt, J. phys. Chem., 80,2556 (1976). (13) E. J. Hart rind M. Anbar, “The HydratedElectron”, Wiley-Interscience, New York, 1970. (14) G. L. Bolton, K. N. Jha, and G. R. Freeman, Can. J. Chem., 54, 1497 (1976). (15) A. Henglein, Can. J. Chem., 55, 2112 (1977), and referencestherein. (16) G. L. Bolton and G. R. Freeman, J. Am. Chem. Soc., 98,6825 (1976). (17) D. Raiem and W. H. Hamill, J. Phys. Chem., 81, 1625 (1977). (18) S. M. S. Althtar and (2. R. Freeman, J. phys. Chem., 75, 3756 (1971). (19) K. N. Jha, G. L. Bolton, and G. R. Freeman, J. Phys. Chem., 76, 3876 (197’2). (20) (a) F.-Y. Jou and G. R. Freeman, Can. ,I. Chem., 54, 3693 (1976); (b) J. Phys. Chem , 81, 909 (1977). (21) K. Okazaki and G. H. Freeman, Can. Y. Chem., 56, 2305 (1978). (22) K. Okazakl and G. 13. Freeman, Can. Y. Chem., 56, 2313 (1978). The rate constants for benzene and toluene were found to be nearly equal in any given alcohol except methanol, for which the experimental scatter was large. We have therefore taken the average value, k, = 1.3 X IO’ M” s-’, for both benzene and toluene in methanol at 299 K. (23) E. J. Hart, S. Gordon, and J. K. Thomas, J . Phys. Chem., 68, 1271 (1964). (24) K. Y. Lam and J. W. Hunt, Int. J. Radiat. Phys. Chem., 7,317 (1975).
1249
(25) F. Barat, L. Gilles, B. Hickel, and B. Lesigne, J . Phys. Chenr., 77, 1711 (1973). (26) D. W. Johnson and 0. A. Salmon, Can. J. Chem., 55, 2030 (1977). (27) K. N. Jha, G. L. Bolton, andG. R. Freeman, Can. J. Chem., 50, 3073 (1972). (28) R. W. Gallant, “Physlcal Properties of Hydrocarbons”, Vols. 1 and 2, Gulf Publishlng Co., Houston, Tax., 1968. (29) M. S. Sauer, Jr., S. Arai, and L. M. Dorfman, J . Chem. Phys., 42, 708 (1965). (30) F.-Y. Jou and G. R. Freeman, J. Phys. Chem., 83, 261 (1979). (31) B. EerEek, Int. J . Radiat. Phys. Chem., 7, 223 (1975). (32) R. R. Hentz and G. A. Kenney-Wallace, J . Phys. Chem., 78, 514 (1974). (33) G. L. Bolton, M. G. Robinson, and 0. R. Freeman, Can. J . Chem., 54, 1177 (1976). (34) Water does not fit the alcohol serjes. Mobilities displayed in Figure 4 of J.-P. Dodelet, K. Shinsaka, and G. R. Freeman, J. Chem. Phys., 59, 1293 (1973), indicate that the diffusion coefficient of solvated electrons is higher in water than in ethanol, although the viscosities are similar. This correlates with the relative values of k, for nitrobenzene (Table I). (35) D. W. McCall, D. C. Douglass, and E. W. Anderson, J. Chem. Phys., 31, 1555 (1959). (36) A. V. Rudnev, A. V. Vannlkov, and N. A. Bakh, High Energy Chem. (Engl. Trans/.),6, 416 (1972). (37) S. K. Garg and C. P. Smyth, J. Phys. Chem., 89, 1294 (1965). (38) B. P. Jordan, R. J. Sheppard, and S. Szwarnowskl, J. Phys. D ,11, 695 (1978). (39) J. R. Partington, R. F. Hudson, and K. W. Bagnall, J . Chim. Phys., 55, 77 (1958). The trend in the values for other alcohols indicates that the reported actlvatlon energy for self-diffusion in ethanol, E, = 19 kJlmol, is too high and should be 15 kJ1mol. The diffusing species actually measuredwas not the alcohol molecule (except in methanol), but the hydroxylic deuteron in RODlROH solutions. (40) T. Shlda and W. H. Hamill, J. Am. Chem. Soc., 88, 3689 (1966). (41) C. Chachaty, J . Chim. Phys., 64, 614 (1967).
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Enthalpies of Solution and Transfer Enthalpies. An Analysis of the Pure Base Calorimetric Method for the Determination of Hydrogen Bond Enthalpies J. N. Spencer,* Jeffrey E. Gleim, Charles H. Blevlns, Robert C. Garrett, and Fred J. Mayer Department of Chemistry, Lebanon Valley College, Annville, Pennsylvania 17003 (Received August 28, 1978; Revised Manuscrlpt Received January 16, 1979) Publication costs assisted ,bythe Petroleum Research Fund
Relative magnitudes of the enthalpy terms contributing to enthalpies of solution and transfer enthalpies are examined. Enthalpy of cavity formation, van der Waals interactions, and specific interactions are considered. Transfer enthalpies calculated from simple dipole expressions and regular solution theory are in good agreement with experiment. Am analysis of the pure base calorimetric method for the determination of hydrogen bond enthalpies is carried out. Cavitation and dispersion interactions are shown to be important. Enthalpies of solution and transfer enthalpies of substituted anilines were determined for the solvents CC14, CHCl,, and cyclohexane. Specific interaction of CC14 and CHC13 with the anilines is found. Transfer enthalpies of gas phase pyridine and DMF to various solvents are calculated and compared to experiment. Anamolous behavior of pyridine and DMF in cyclohexane is suggested.
Introduction In 1967 Armett et n1.l introduced the pure base method for the calorimetric determination of the enthalpies of hydrogen bond formation. The method was viewed with skepticism by its originators but was later shown to give enthalpies in good agreement with published values for about 20 different systems.2 The authors also compared pure base enthalpies to those obtained by the high dilution calorimetric imethod and found agreement within experimental error in most cases. Duer and Bertrand3 in 1970 carried out experiments to test some of the assumptions of the pure lbase method and concluded that absolute enthalpies of formation determined by pure base may be in error relative to reported uncertainties because of the 0022-3664l79/2083-1249$0 1.OO/O
choice of model compound and reference solvents but that relative enthalpies might be more reliable than those obtained by the high dilution method. Arnett, Mitchell, and M u r t 9 reexamined the pure base method in 1974 and reported that for 29 compounds the average difference between the pure base and literature values for hydrogen bond enthalpies for phenol was 0.5 kcal mol-l. The most serious discrepancies were found when very polar bases were used in the pure base method.
Results and Discussion The pure base method assumes that when the pure base serves as solvent each acid molecule will be hydrogen bonded to a base molecule. The interactions between the @ 1979 American Chemical Society
