THEKINETICS OF
July 5 , 1957
SOLVOLYSIS O F
Furthermore, the possibilities of the other mechanisms will be discussed below. A rate-determining reaction of protonated phenylhydroxylamine with nitrosobenzene i
C~,I-IINH*OH t- CaHbNO
ks
ACYLESTERS OF
S A L I C Y L I C ,4CID
340 1
or respectively; hence neither mechanisni alone or combined with equation 5, where
--+ +
+ IT ’
C ~ H ~ ~ = - S C 4G IIIzO €~ I
(8)
I
0-
or a rate-determining condensation of both pro-
I
+
C~HSN=~C~H H20 ~
I
+ 2H+
(9)
0-
tonated reactants should lead to the rate constant expression ksKi [H+I k =_-(1 + K i [ H + l ) ( 1 +Ka[H+l)
[CONTRIBCTIOS FROM
THE
agrees with the fact, since each k value should diminish to zero as the hydrogen ion concentration approaches zero. Also, equations 8 and 9 do not seem rational in the light of the electronic theory, since they involve the attack of a cation on the positive nitrogen of nitrosobenzene. This electronic point of view can also be used to eliminate a mechanism involving the simultaneous reaction of equations 8 and 2, although the mechanism leads to a rate equation consistent with the experiment. Acknowledgments.-The authors wish to thank Profs. R. Oda, S. Tanaka and M. Okano for their aid in performing these experiments. KYOTO, JAPAN
RESEARCH LABORATORIES OF THE UPJOHX Co.]
The Kinetics of Solvolysis of Acyl Esters of Salicyclic Acid BY EDWARD R. GARRETT RECEIVED DECEMBER 28, 1956 The acid, neutral and alkaline hydrolysis of acetyl-, trimethylacetyl-, 8-cyclopentylpropionyl- and diethylacetylsalicyclic acids were studied and the rate constants determined as functions of dissociation and of hydrogen and hydroxyl ion concentration. The Arrhenius parameters of entropy and heat of activation also were determined and mechanisms were considered in light of these values and the structures of the esters. Disagreement with values given in the literature was noted. A n unusual increase in “spontaneous” rate was observed with increasing alcohol concentration whereas no such increase was noted with increasing dioxane content.
In order to determine the inhibition of hydrolysis in the hydrolysis rates of the anion in the pH-indeof acyl salicylates by varying alkyl substitution on pendent region was observed. This pH-independthe acyl group, the acid and base hydrolysis of es- ent hydrolysis of the completely ionized compound ters selectively chosen to scan substituent effects, (as represented by the plateau of rate as. p H a t pH acetyl-, trimethylacetyl-, P-cyclopentylpropionyl- 5.5-9 in Fig. 1) had been ascribed by Edwards1V2to and diethylacetyl salicylic acids were studied. ,4 the water attack on the aspirin ion and would imply very complete study of the hydrolysis of acetyl general acid-base ~ a t a l y s i s . ~I n this light ethanol salicylic acid (aspirin) had been made by Edwards. l,? would have to be considered a more active base However, since these papers had reported an unu- than water. However, Edwards’ had also shown sually high alkaline hydrolysis heat of activation of by studies in various concentrations of acetate buf25 kcal./mole, whereas similar acyl phenates had fers that the “spontaneous” hydrolysis was not subactivation energies of an order of magnitude of 12 ject to general acid-base catalysis. k ~ a l . / m o l e the , ~ hydrolysis studies of aspirin were The extraordinary magnitude of this “spontanerepeated and the Xrrhenius parameters determined. ous” or water hydrolysis of aspirin has led to the I t was desired to carry out the hydrolysis studies formulation of cyclic mechanisms involving interin solutions as nearly aqueous as possible, but the action of both carbonyls and their carbons in a rateinadequate water solubility of P-cyclopentylpro- determining ~ t e p . ~ J pionyl- and diethylacetylsalicylic acids necessitated In order to clarify the anomalous enhancement of use of water-alcohol mixtures as solvents to main- “spontaneous” hydrolysis with increasing alcohol tain sufficient material in solution for spectropho- content of the solvent and to determine if the rates tometric analysis. In order to place the hydrolysis were dependent on water content or were functions studies of all these esters on comparative grounds, of dielectric constant, the “spontaneous” hydrolysis the hydrolysis of aspirin was studied in various was studied in various dioxane-water mixtures. water-alcohol mixtures. An unexpected increase Surprisingly, there was small variation in hydrolysis rates. (1) L. J . Edwards, Trans. Faraday SOC.,46, 723 (1950). (2)
L.J. Edwards, i b i d . , 48, 696 (1952).
(3) “Tables of
Chemical Kinetics: Homogeneous Reactions,” National Bureau of Standards Circular ,510, U. S . Department of Commerce, Washington, D. C., 1951.
(4) J.L. Hockersmith and E. S. Amis,Anal. C h i m . A c t a . 9 ,101 (1953) (5) J. D. Chanley, E. M. Gindler and H. Sobotka, THISJOURNAL, 1 4 , 4347 (1952). (6) D. Davidson and L. Auerbach, i b i d , 1 6 , 5984 (1953).
3402
Yo]. 79
EDWARD R . GARRETT
T.4BULATION OF
TABLE I RATECONSTANTS (k I N SEC.-', ki I N L./MoLE/SEC.)FOR CALCULATION SALICYLIC ACIDIN ACIDAND NEUTRALMEDIA^
OF
HYDROLYSIS O F ACYL ESTERS OF
lOska
oc.
pH 1.10
1OekC pH 2.50
pH 5.05
10aki
108kkr
10'kaCHzo
Acetyl-
25.0 35.0 50.3
Trimethylacetyl-
60.3 25.0 35.0 50.3 00.3
2.24 5.00 20.5 44.2 0.113 0.273 1.13
0.562 1.31 4.83 12.1 0.0883 ,213 ,847 2.02
3.63 8.03 38.3 80.8 n ,933 2.23 10.00 25.5
25.0 55.2 237 492 n . 960 2.33 10.6 27.3
1.02 2.55 7.13 31.2 0.190 0,465 1.46 2.80
3.72 8.33 36.3 83.8 0.9% 2.35 10.5 26.8
pH 1.10
pH 2.60
p H 5.30
0.922 2.15 7.62 17.8
0.207 0.462 2.03 5.12
Ester of salicylic acid
t,
3.62 10.3 1.80 3 . SL5 8 .05 24.3 3.97 8.80 82.2 19.4 36.7 33.7 51.6 87.5 188 80.0 0.187 0.131 0.117 0.223 25.0 0.0262 0.0110 Diethylacetyl35.0 .OS50 ,0362 0.338 0.713 0.633 0.387 50.3 ,458 ,215 1.64 3.53 4.00 1.82 60.3 1.45 ,662 4.38 11.2 12.5 4.92 a The acetyl- and trimethylacetylsalicyclic acids were studied in 0.570 ethanol; the P-cyclopentylpropionyl- and diethylacetylsalicylic acids were studied in 205;lo ethanol by volume. 6 The solutions were 0.0978 M in HCI, 0.0469 J I in Kc1. c The solutions were 0.534 dl in acetic a&l. The solutions were 0,1091 M in acetic acid, 0.0785 dl in NaOH.
p-Cyclopentylpropionyl-
25.0 35.0 50.3 60.3
2.77'
10w
+
d( [CeH4(OOCCH3)COOH] [C,H,(OOCCH3) COO-]) /dt = -ki [H+][CsHkr(OOCCH3)COOH] - (kr[H+) f ~ ~ C H ~k~[OH-])[C~H,(OOCCH,)COO-] O = -k ( [CsHc(OOCCH3)COOH] [C,H,(OOCCHa)COO-] ) (1)
+
+
where k ki[H+1/(1 K'e./[H+I) (k,[H+] 4- ~ ? C H , O ks[OH-])/(l [H+l/K'd (2) where K , is the dissociation constant of aspirin and k is the pseudo first-order rate constant a t a constant pH. The pseudo first-order rate constants obtained a t PH values of 1.1,2.5 and 5.1 and the determined K,' values were used to calculate the k l , kq and k5CH20values for acetyl-, trimethylacetyl-, 0-cyclopentylpropionyl- and diethylacetylsalicylic acids a t 25.0') 35.0") 50.3' and 60.3'. These rate constants are given in Table I. The spectrophotometrically determined pKa and pKa' are given in Table TI. Their relative significance is discussed in the Experimental section. The experimentally determined bimolecular rate constants, k 6 , are in Table 111.
+
+
+
+
TABLE I1 DISSOCIATION COXSTANTS OF ESTERS OF SALICYLIC ACID Ester of salicylic acid
P H. Fig. 1.-The PH profile o f the logarithm of apparent firstorder rate constants for hydroll-sis. The points are experimental; the drawn lines a r e calculated. Curve
A €3
C
r)
Ester of salicpclic acid
Acetyl8-Cyclopentvl!,roDiori?.lTriniethylacetyiDiethylacetyl-
Calculations and Results According to Edwards1,*the hydrolysis of aspirin, C6H4(00CCHs)COOH,may be expressed by
% EtOH
1,
pK'.
0 . 5 3.62 20 3.97 Acetyl30 4.26 40 4.71 Trimethylacetyl- 0 . 5 3.74 P-Cyclopentyl20 4.26 propionyl20 4.35 Diethylacetyl-
pK.
3.69 4.15 4.47 4.94 3.87
K"
O C .
H%O
25 35 50 60
14.00 13.68 13.26 13.02
20 % EtOH
14.33 14.00 13.57 13.32
4.36 4.45
Actually, in these equations 1 and 2 , kl = ki'jix+ and k4 = k4'f11', where the ki' are true cpecificrate c o w t x t s acd f~ is the activity coefficient of hydrogen ion. The values actually used for [H +] were
THEKINETICSOF SOLVOLYSIS OF ACYLESTERSOF S.~LICYLIC ACID
July 5, 1937
3403
TABLE IJI TABULATION OF BIMOLECULAR RATECOXSTANTS (ks I N L./MoLE/SEC.)FOR ALKALINE HYDROLYSIS' OF 5 x ACID AS DETERMINED AT VARIOUS INITIAL ALKALICONCENTRATIONS, [ XaOH]: ESTERSOF SALICYLIC -T
Ester of salicylic acid
"C
lOr[NabH];'
Acetyl-d
TrimethylacetylP-Cyclopentylpropionyl-
100 40 20 20 20 10 200 100
' 0 . 5 3
25.3"-
lOzka
11.3 10.6 11.0 10.8 0.862 0.867 4.03 4.05
5 0 . 3 ' 10%
104[NaOH]o
50 20
400 200 100 50
-'3.0-
104[NaOH]o
20.2 20.7
1.67 1.66 8.05 7.78
hf AC\71,
1OZka
lO*ka
104[NaOH]a
20 10
53.8 54.3
10 5
200 100 50 20
4.75 4.50 14.7 16.0
100 50 30 20
95.0 113"
8.05 7.70 26.8 26.3
400 0.144 400 0.328 400 0.662 400 1.26 200 1.18 0.638 200 0.297 200 0.139 200 The acetyl- and trimethylacetylsalicylic acids were studied in 0.5% ethanol; the P-cyclopentylpropionyl- and diethylThe bimolecular rate constants were determined from pseudo firstacetylsalicylic acids were studied in 20y0 ethanol. At this low concentration of alkali, it is exA f and greater. order rate plots a t [NaOHIoconcentrations of 100 X pected that catalysis other than hydroxyl would be significant. Thus, the observed ks should be higher than the true k s . The apparent bimolecular rate constants (ke) for alkaline hydrolysis were 0.116 and 0.114 I./mole/sec. in 20 and 40?,;, ethanol by volume with 10 X lo-' M [NaOH]. Diethylacetyl-
OTHER A C I D
Ester of salicylic acid T, 'C.
HYDROLYSIS' STUDIES AcetylActual
10Sk
9H
OF
Calcd #H
TABLE IV ACYL ESTERSOF SALICYLIC ACIDWITH ACTUAL P H b [H+] = ( k - k4Ka)/KI; k IN S E C . - ~ TrimethylacetylActua Calcd. fiH PH
10'k
AND
8-CyclopentylpropionylActual Calcd. #K fiH
pH CALCULATED F R O M
look
1O'k
DiethylacetylActual Calcd. PH BH
11.8 0.40' 0.34 0.415 0.32" 0.40 4.51 0.43" 0.36 0.116 0.50" 0,:32 1.57 1.30d 1 . 2 7 .0885 1.27d 1 . 2 . : 0.593 1 . 3 s d 1.32 .0224 1.4Ud 1 . 2 ' ) 35.0 6.88 0.86' 0.95 .347 0.93" 0.95 2.70 0.96" 0.99 ,101 0.99e 1.00 50.3 26.9 .92" .97 1.37 .9Se .98 9.40 1.03" 1.00 ,556 .97" 0.97 60.3 199 .34" .40 9.75 .34"' .48 28.2 0.67' 0.87 2.52 .71/ 0.76 * The acetyl- and trimethylacetylsalicylic acids were studied in 0.5% ethanol; the P-cyclopentylpropionyl- and diethylPH was averaged from five samples taken a t room temperaacetylsalicylic acids were studied in 2070 ethanol by volume. 0.4708 111 H2S01. 0.04708 AI' HaS04. ture during the hydrolysis, the range in pH values was of the order of 0.1 pH. 0.2354 M HzSOI. e 0.1178 M HnSOd. 25.0
mined on the basis of the previously derived kl and k4 values, [X+] = ( k - k4Ka)/'k1; the agreement of the actual and calculated p H values lends support to the validity of the data and methods of calculation. The Arrhenius plots of the logarithm of the rate constants against the reciprocal of the absolute temperature as per log k = -(AHa/2:303R)(1/T) log P = S/T + log P
the activities as related to the experimentally determined p H values. Anion and acid concentraexperimentally detions are defined in terms of kr&' rived from the spectrophotometric titrations. The rate constants kg and ke do not need activity corrections since they are based on anion concentrations similarly derived ; the latter, k6, is determined from hydroxyl concentration given in terms of molarity. I t should also be realized that Edwards' term, k 4 [ H + ] [anion] is kinetically indistinguishable from k4K', [acid]. The logarithmic profiles of pseudo first-order rate constants against 9 H a t 25' for these esters of salicylic acid as from the derived k l j k41 k:,cH20,k6 and K,' values are drawn in Fig. 1. The plotted points are experimentally determined rate constants a t their determined p H values and confirm the derived values. Similar profiles may be constructed for other temperatures using the proper autoprotolysis constant, K,, of the solvent. Such values may be estimated from the l i t e r a t ~ r e . ~ - ~ Several acid hydrolysis studies were run a t the various temperatures and pH values and pseudo first-order rate constants were determined (Table IV). The experimental rate constant data were used to calculate the theoretical GH values deter-
Spectrophotometric Determination of Dissociation Constants.-The dissociation constants given in Table I1 of
(7) B. Gutbezahl and E. Grunwald, THISJOURNAL, 75, 5 6 5 (1953). (8) H . S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic S,,lutions," 2nd E d . , Reinhvld Publishing Corp., New Yvrk, N . I ' , 1950, p i Y i .
(9) S.Glasstone, K.J. Laidler and H . Eyring, "The Theory of Rate Processes," McGraw-Hill Book Co., Inc., New York, N. Y ,1941. (10) The esters of salicylic acid were supplied by Mr. Maxton F. Alurray of the Chemistry Department, The Upjohn Co.
+
(3)
are given in pig. 2 for k l , ~
i 3 for ~ h4, . ~
i for ~ . k&H20 determined a t p H 5.1 and Fig. 5 for k6. The thermodynamic quantities derived from this function are given in Table v where AS+, the entropy of activation, is calculated from absolute pactio11 rate theoryg as A S * = 2.303R[log P - log ( k B T / h ) ] (4) where kg is Boltzmann and h is the Planck constant. The pseudo first-order rate constants for the hydrolysis of aspirin in various percentages by volume of 95yo ethanol and dioxane in water mixtures, as well as the apparent pH values are given in Table VI. Experimental '0
3404
EDWARD R. GARRETT
Vol. 79
TABLE V T I i E R M O D Y S A Z I I C QI:AKTITIESa
k a t 9 H 1.10 katpH2.5
3640 3700 3690 3590 3840 2730
hi
ki
kd.moC I. P
16.6 6.55 1 7 . 0 0.17 1 6 . 7 7.78 l6.4:9.05 17.6 7 . 4 5 12.5 8 . 1 8
-28.8 -30.2 -22.9 -17.4 -24.7 -21.0
3610 3880 3560 4150 3820 2480
FOR THE H Y D R o I . Y S I S * OF ESTERS 0 1 1$.41.IC\ ~
16.6 17.8 16.3 19.0 17.5 11.3
-30.7 -29.3 0 . 9 3 -27.0 11.14 - 7 . 8 7.40 - 2 4 . 7 6.89 - 2 7 . 0 8.08 (i.40
3880 3860 4130 3295 3980 2730
17.7 8 . 0 5 17.7 3 90 18.9 7 . 8 0 15.1 i 3 6 18.2 i.34 12.5 7.08
r.rc
;\c,ll)
4910 4980 4710 5110 4420 2700
-31.2 -3l.(; -22 9 -217 -25.2 -2b.l
2% 5 22 8 2l.(i 23.4 20.2 12.4
8.88 8.74 9.16 13.41 7.95 0.21
-17,!+ -18.3 -l(i,5
i- 2 . 7 -22.0 -30 2
+
The quantities are derived from the logarithmic form of the -%rrheniusrelation: log k = -(AHsj2.303R)(1/T) log -S/T log P where the k is in sec.-l and the k , in l./mole/sec.. the AH, is the heat of activation in kcal./n:ole, I? is the gas constant in cal./degree and T i s the absolute temperature. The A S $ entropy of activation is calculated from absolute reaction rate theory: A S $ = 2.303R [log P - log ( k B T / h ) ] where k g is the Boltzman constant and h is Planck’s constant. The acetyl- and trimethylacetylsalicylic acids were studied in 0.5% ethanol: the 8-cyclopeutylpropionyl- and The values for ksCH,o are determined for the pseudo diethylacetylsalicylic acids were studied in 2070 ethanol by volume. first-order rate constant of the presumed product in set.-'.
p
’
+
=
the various acyl esters of salicylic acid were determined a t 26 + 1 by a procedure similar to Edwards’’ by the expression
p1i =
=
pKa
PK,’ - log
- log
{
- At)/(.4, - U . 4 - C ) j - At)/(.4t - U.4 - c) f
{(UHAC
(UHAC
fz&.fi
of acid, L e . , C = [HA] i- [A-1, and A t is the total absorbance a t any pH, i.e., A t =: ah- ( C - [HA]) a ~ . 4[HA]. The value off in water is 0.51, butf a 1/Ds/2 where D is the dielectric constant so that ( D H ~ O / D Ethanol)’/z % 0.51 = f% Ethanol where values of D for various percentages of ethanol a t 25’ are given in the literature.’* The ionic
+
( 5) where HAC is the asymptotic absorbance in acid media, a.4- C is the asymptotic absorbance in alkaline media, UHA and nA- are the molar absorptivities, respectively, of the un-
-401
,
3.0
31
-6 4
1
I
\O
30
31
32
lo3 x
33
I/T.
Pig 2.--.2rrhenius plots for acid hydrolysis of undissociated esters of salicylic acid (kl in l./mole/sec.). Curve
B
C
n
Ester
i3cetylO-CyclopentylpropionylTrimethylacetylDiethylacetyl-
dissociated and dissociated acid, p is the ionic strength, jZ,Z2 is the coefficient of the ionic strength correction in the slniple Debye-Hdckel equation,” Cis the t o t d concentration
x
I/T.
Fig 8.--krrhenius plots for acid hydrolysis of dissocia t v t l esters of salicylic acid (kd in l./rnole/sec.). Curve
a i
-6 8
33
32 lo3
x B
C D
Ester
.%cetyl,~-CyclopentylpropionvlTrimethylacetylDiethylacetyl-
strength of the acetate buffers in various ethanol concentrations also was calculated from the p K , of acetic acid in aqueous alcohol.13 Sufficient acyl ester of salicylic acid was weighed in volumetric flasks, dissolved in the appropriate amounts of eth(11) F. H. Getman and F. Daniels, “Outlines of Physical Chemist r y , ” 7th Ed., John Wile? and Sons, Inc., Xew York, S . Y., 1945, pp. lifi7-670. (12) H. S.Harned and B . € 3 Owen, “The PhyFical Chemistry o f Electrolytic Solutions,” 2nd Ed., Reinhold Publishing Corp., Pieiv York. N. Y . , 1950, p. 118. ( 1 3 ) J . I\I Vandenbelt, C . I1 Sliurlock, \ I C.iRels xn
