A KINETIC STUDST OP THE DECOMPOSITION OF SOME «C-AiHrLSULFONYLACETIC ACIDS A1 ID THEIR SALTS HI VARIOUS SOLVENTS
DISSERTATION sented in Partial Pulfillxncnt of the Requirement for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University
By Donald Joseph CJ’Connor, B.S. The Ohio State University 1952
Approved by:
i J d b l Advioer ACKNOWLEDGEMENT
The author would Ilka to aeknowladgs his indebtedness to Dr. Frank Varhoek for hid halpful eounsal during the eouraa of this work.
-1
S 0 9 4 GO TABLE OF CONTENTS
INTRODUCTION 1 EXPERIMENTAL Preparation of Solvents 3 Preparation of p-Toluenemercaptoacetic Acid 3 Preparation of p-ToluenesulfonylacetIc Acid k Preparation of Sodium p-Toluenesulfonylacetate 5 Apparatus 5 Procedure 6 RESULTS The Reaction 8 Order of the Reaction 8 Dissociation Constant of p-Toluenesulfonyl- acetic Acid 9 Reaction Rate Constants 10 Effect of Solvent 1^ Effect of Concentration 23 Effect of Added Ease 23 Energy of Activation 24- Entropy Factor 26 Decomposition of the Free Acid 23 DISCUSSION Effect of Solvent on the Activation Energy 30 Sffect of Solvent on the Sntrooy Factor 3Jl Effect of Ion Association 1u2
SUMMARY h 5 SUGGESTIONS FOR FURTHER WORK I4.6
EXPGRIHGNTAL DATA REFERENCES 61 AUTOBIOGRAPHY
-ii- A KINETIC STUDY OP THE DECOMPOSITION OP SOME og-ARYLSULPONYLACETIC ACIDS AND THEIR SALTS IN VARIOUS SOLVENTS
INTRODUCTION 1 2 Recent review articles * have highlighted the great Interest in decarboxylation reactions of many types of organ ic acids. In this paper, the major concern is for that tyne of reaction which involves first order decomposition of the anion of the acid as examplified by trinitrobenzoic acid,^*^ trihaloacetic acids^'^*?»8 and X-nitroa].Vcylcarboxylic^' ^ ! r-/ acids. Verhoek ^ * has suggested that the mechanism of this reaction involves only a unimolecular decom oosition of the
acid anion to form a carbanion and carbon dioxide with a sub sequent rapid reaction of the carbanion with a nroton from
the solvent. Eyring^ has suggested a bimolocular reaction between the acid anion and a solvent molecule to give direct ly the product and carbon dioxide.
On the basis of available evidence no clear cut flee is ion between these hypotheses could be reached. It was hoped, how
ever, that study of another system of acids of a similar re action type would help in the resolution of this problem.
Professor G.L. Wilson suggested that af-arylsulfonylfat by acids might be of interest. In the early literature phenyl-
sulfonylacetIc acid and the toluene analogue wero reoorted 1 ° to yield the methylarylsulfones and carbon dioxide when heat ed at 170° In strong alkali. Substituted of-nltroarylsulfonyl - 2 - fatty acids have been transformed * i f likewise into «i-nitro- arylalkylsulfones by heating the alkali metal salts of the acids in aqueous feebly acid solutions. In no case, however, were any quantitative kinetic investigations carried out. Therefore, it x*as decided to investigate the decomposi tion of p-toluenesulfonylacetic acid and some of its salts In water and water-glycol, water-dioxane mixtures under vari ous conditions of temperature, concentrations and added sub stances. The results of this investigation Trill be found in the pages following.
0 - 3 - SXPSilU'EJIiTAL Preparation of Solvents.
Ethylene glycol, obtained from the Carbide and Carbon Chemical Corporation, was dried over drierite for three days and distilled at 1,7 nun. of mercury through a 1|.0 cm. column packed with monel metal scroening. The middle half was col lected and stored in a glass stopnered bottle under nitrogon. Analysis with Fischer Beagent showed loss than 0.05 percent wator.
Dioxanc, obtained from the Carbide .and Carbon Chemical Corporation, was purified by the mobhod of 'less and Prohm 15 as described by Fieser. ^ The distillation was carried out through a 90 cm column packed with glass helices under nitro gen. The middle three fifths (101.1-101.2°) was collected and stored under nitrogen in a glass stoppered bottle. The mixed solvents wore prepared by transferring r. weighed quantity of the particular solvent to a glass stop pered bottle previously flushed with nitrogen. To mis uas then added a weighed quantity of double distilled wa.,er mb'- ficient to yield the solvent mixture desired. All weighin. -j were made to the nearest centigram.
Preparation of p-toluenemercaptoacetlc acid. Cd-j-
p-Toluenenercaptoacetic acid was prowared accoi’diug ; It) a modified procedure of Anwers and Thies. In 5^0 ml. c " water 100 g. (0.8 moles) of p-thlocrosol (Eastman Kodak C.., - k - white label) was neutralized with. 32 g. (0.8 moles) of sodium hydroxide. The resulting salt solution was poured rawidly
with stirring into ij.00 ml. of a water solution entaining 76 g. (0.8 moles) of monochloroacetic acid (Ilallinkrodt A.R. grade) neutralized with sodium hydroxide. In a few minutes, sodium
p-toluenemercaetoacetate precipitated forming a white pasty mass. The precipitate was allowed to stand in contact with
the supernatant liquid 3 to Ij. hours before filtration. The salt was washed on the filter nanor with one 100 ml. wortion
of cold water, then resuspended in fj'OO ml. of water. The free acid was formed fro i the salt by the add. it ion of 1^0 ml. of concentrated hydrochloric acid. After filtration, the acid was recrystallized once from -water, from, which It separates as an oily liquid before solidifying, and twice from benzene forming beautiful white flakes. (n.p. 93* 0° G ) The yield was 90 percent of the theoretical.
Preparation of p-toluenesulfonylacetic acid. CH^-^^SQo-CrigCQQH The purified mercaptoacetic acid was oxidized in glacial
acetic acid by 30 percent hydrogen peroxide by a modification of the method of Siebert and Fromm.^ In 200 ml. of rc -,1c acid was dissolved lj.3 g« (0 .21|. moles) of the acid. Over a one hour period ml. of 30 percent hydrogen peroxide . • ■: .w ’
slowly while the temperature was maintained bo two on 1':° el 20°.
After standing 2l\. hours, the solvent was distilled almost to dryness under the vacuum of a water as virator. The rosl e as - 5 - then poured into a small amount of water and recrjstallized rapidly. Two recrystallizations from benzene gave white crystals of p-toluenesulfonylacetic acid (m.p. 117• 0°-117»5>° ) • The melting point was taken as an indication of purity. The neutral equivalent was determined to bo 211|.*1 compared to a theoretical value of 2H4..2.
Preparation of sodium p-toluonosulfonylacotate. — TO^e^-'Stfo-gBp-doofa ------The sodium salt of this acid was prepared by mixing equi valent amounts of alcohol solutions of sodium hydroxide and the acid. The salt precipitated im odiatcly and w:i3 recrystal lized from alcohol. The salt was then dried and stored over Drierite.
Apparatus.
The thermostatically controlled baths wore those described 18 by Trivich. Mineral oil was used as heat transfer medium. Temperature control, using a mercury in gloss regulator to control the grid of a FG5>7 thyratron tube, was maintained at iO.Of?° in all three baths, 7^«0°» 8£.0° and 95»0°* In addi tion to the intermittent heater a permanent 2^0-watt heater controlled to about 1S>° below the bath temperature by moans of a powerstat was used.
The sample tubes used on most runs were those described l8 by Trivich consisting of a tube of about ij.0 ml. capacity fitted with a standard taper, 19/.3 J3 * ground glass cover, end - 6 - around this another large cover, standard taper, 3h A !-£, which prevented seepage of oil into the sample tube.
Procedure. The reaction samples were prepared by dissolving a weighed sample of the appropriate salt or acid in 2%0 ml. of the ap propriate solvent. Ten milliliter sai^mles were pipetted into the reaction flasks which were then closed. All of the re action flasks, suspended from bars, wore immersed in the baths at the same time. This nroceduro is justified by the similar ity in construction of all the tube3. After one half hour had elapsed to insure that the temperature of the bath had been attainod, the first sample was withdrawal and olun.god into an ice bath to halt the reaction. After cooling and opening the reaction flask, 10 ml. of standard hydrochloric acid was added to the flask to mako the solution acid. ITitrogon was then bubbled through the solution for 10 minutes to remove the carbon dioxide formed during trie reaction. Jhe excess hydrochloric acid was titrated to a phonolphthaloin end point with standard sodium hydroxide solution from a 10 ml. nicro- burette. All analytical solutions used were obtained from the Ohio
State University Control and deagent Laboratory.
When the reacting solutions contained only acid rather than salt, the procedure was altered slightly. After samples were removed from the bath and cooled, nitrogen was nassed - 7 - bhrou^h the solution Tor 10 minutes cl remove carbon dioxido, and the acid v;as then titrated directly with the standard sodium hydroxide solution. RESULTS The Reaction To verify that the reaction proceeds as reported by O t t o , ^ viz. , ° 0 -
0 , 6 ^ 7 1 g» of sodium p-toluenesulfonylacetate wus dis solved in a 0.06 N sodium hydroxide solution and refluxed for six days, sufficient for 96-99 percent of reaction to occur. From, the cooled alkaline solution 0 .iiJpG3 g. of methyl-p-tol- uenesulfone corresponding to 97 percent reaction was recovered by means of an ether extraction followed by evaporation of the solvent on a steam bath. The molting oolnt of the sulfone was 8£*5>°“6£>.20 compared to a literature"^ value of u6o-07°. The carbon dioxide liberated was not determined.
Order of the Reaction. To determine the order of uho decai’boxylatIon reaction, the concentrations of anion were plotted against rime and slope were measured at various concentrations. By plotting the loga rithm of the slope against concentration, the order of the re action was obtained from the slope and was found to be first order with resnect to the anion In water and IgO percent ethyl ene glycol. Since first order plots gave straight lines for all other solvent mixtures, this was assumed to oe the order in the other solvents too. The order did not change with con centration on going from 0.01 I] to 0.1 IT initial conooutra- t ior.3. - 9 - Dissociation Constant of p-Toluene3ulfonylacetic Acid. Sinco the acid dissociation constant, K^, for p-toluene- sulfonylacetic acid has not been reported and since it will be useful in subsequent calculations, it was determined in water at 25>° by the half titration method. A solution of O.J4J4.32 g. (0.00207 moles) of the acid in 3 ^ i'nl. of water was prepared and titrated with 0.1007 N sodium hydroxide solution.
The hydrogen ion concentrations, measured with a Beckman pH meter at various poinbs in the titration curve, were plotted, as pH values, against the anount of alkali added. from the expression
where fr-l is the concentration of the salt, assumed equal to tho amount of alkali added (corrected to the proper volume),
and is the concentration of the unneutralizod acid, and [if*] is determined by the pH value, can be calculated.
This was done for five poinbs in bhe titration curve as shown in Table I. TAELS I Approximate Dissociation Constant for p-Toluenesulfonyl- acetic Acid by Titration Hethod. 3 fraction Neutralized pH x 1CK
l A 2.314- 1.S2 1/3 2.1)4 1.02 1/2 2.63 2.34- 2/3 2.83 2.96 3A 2.98 3.15 Avg. 2.36 - 10 - Reaction Rate Constanta, With the values of the weighings and the volumes obtain ed as described in the procedure and listod under Tables IX to XXIX it is possible to calculate the velocity constants fox' the reaction at various temperatures and in various nixed solvents. These values are listed in Table II.
The solutions are made up at room temperatui’e, and the concentrations at that temperature are known from tho weight of salt and the volume of the final solution. ifJhen these solu tions are put into the heated baths, there will bo a decrease in concentration due to thermal expansion of the liquid. It Is not necessary to correct the conccntrv.tions to their high temperature values, since for each lermeraturc the concentra tion actually existing at that temoerature Is tho concentration at room temperature multiplied by a constant factor. The order of the reaction has been determined to be first order with respect to the anion concentration. Therefore a plot of the logarithm of concentration against time allows she deter mination of the velocity constant from the slope of the result ing straight line. Thus the logarithm of the constant factor cancels out of the calculations. The volume of sodium hydroxide needed to titrate the ex cess hydrochloric acid before tho samples are heated corres ponds to the initial calculated concentration. If the 'dif ference between this initial volume and the volume of sodium hydroxide solution required at any subsequent time, t, is Table II Reaction Rate Constanta Tor the Decomposition of Sodium p-Toluenesulfonylacetate in Water-tithylene Glycol and Water-DIoxane Mixtures ~ _1 Calc'd Weight Mole Dieletrlc / \ k x 10 hours” from percent percent Const. 85° 75° 85° 95° Table
STHYL3Hji GLYCOL Initial ConcentratIon 0 .1 molar 0 .0 0 . 0 0 59.2 0.186 0.790 2.97 IX 0 .201+ 0.806 3.07 X 20.5 6.97 53.6 0 .21^2 1 .0 1 1+.02 XI 39.9 16.15 Il8 . 1+ O .383 1.11-7 5.33 XII 59.2 29.7 4-2.9 0.14.95 2.15 8.19 XIII 79.6 51.9 35.8 0.937 3.61 13.28 XIV 100.0 100.0 27.8 1.732 0 .11-7 23.75 XV Initial Concentration 0.01 molar 0.0 0.00 59.2 0.223 0.775 2.86 XVI 20.5 6.97 53.6 0.223 0 .O99 3.25 XVII 39.9 16.15 k8.k 0.161 0.866 k.1+9 XVIII 59.2 29.7 1+2.9 o.h33 1.73 7 • ll'r XIX o.i+33 1.98 6.37 XX 79.6 51.9 35.8 0.720 2.1+8 11.39 XXI 100.0 100.0 27.8 1.51+ i+. 21+ 21.19 XXII Initial Concentration 0.05 molar 0.0 0.00 3.01 XXIII DIOXAlTti Initial Concentration 0.1 molar
2i. h 5.63 *1-3.5 O .383 1.603 :cxiv 1+0.0 12.79 30.9 0.967 k .05 x:cv 60.1 21+. 8 2.31 B.i-5 xx/i 6 0 .0 2k. 8 18.3 2.69 9.36 XXVII 79.7 I+0 .3 8.2 5.89 22.18 XXVII ETTIA1T0L Initial Concentrat ion 0.1 noi.'ir V2.5 82.0 19.3 8.9 XXIX (a) See Table III, p . 16 for values at 75° nnd 95°. - 12 - multiplied by the normality of the alkali, and divided by the volume of the sample, the amount of decomposed anion, x, is determined and thus also the concentration of the remaining salt, (a-x). The value for a is determined from the titration value for tho first sample withdrawn from the bath and the time of plunging this first sample into an ice-water bath be comes zero time. According to the integrated form of the differential equa tion for a first order reaction log (a-x) = -kt/2.303 + log a a plot of log (a-x) against t will yield a sti'aight line of slope -k/2.303. A typical plot of this tyne is shoT-/n on Figures 1 and 2 . All the reactions of the sodium salt in water-glycol mixtures at 95>° were carried out to about 7 to 80 percent of completion. At the lower temperatures this was not convenient because of the time involved. Therefore, the completeness of the reaction at 75° varies from about 10 percent in water to ^0 percent in ethylene glycol and dioxane mixed solvents, and at 85° from about 10 percent in water to about 7p percent in mixed solvents. Over this range however, good straight lines were obtained. The most reliable data reported here are those concerned with initial concentrations 0.00 molar or greater. The runs made with concentrations initially about 0 . 0 1 molar were easily affected by carbon dioxide absorption from the atrnos- 11(11U) if5K:S: h :i ; - 15 - phere and insufficient precautions wore taken in the early runs.
Effect of Solvent on the Reaction Hate Constants.
If the data in fable II are examined, a v^ry definite and narked increase in the rate constants is observed as the content of glycol or dioxane is increased. Tho value of the reaction rate constant in glycol is eight tines that in water at all three temperatures, and in dioxane, the rate cons traits are about fifteen tines that in water.
To see just what relation exists between the solvent and rate, it was first assumed that the dielectric constant of the solvent is the determining factor and not the specific effect of the solvent. Plots of the logarithm of the r^te constants were made against the logarithm of dielectric constant (Figures
3-a and 3-b) and against tho rociorocal of the dielectric con stant (Figures J.j.-a and k-b ) . fee uieiectric constant v alues at the apnronriate concentrations and to ineraturos were ob tained by interpolation of Akerloff 1 s1^ *1-0 data. In both cases, straight linos were obtained over the range of dielectric constant from 65 to 35 * Thus the deviation from the straight line is more pronounced in dioxane because of the much lower dielectric constant in the 60 and 80 percent dioxane mixtures.
However, in the straight line nortion it was observed that the slopes were approximately equal for both solvent systems. Therefore, the rate constants were nlotted against dielectric Table III Dielectric Constantsq on
Calculated from Akerloff's , e ~ Data
We igilt Percent 75° 85° Glycol 0 6 k . 1 59.2 56.5 20.5 56.3 53.6 52.1 39.9 50.6 1+8. k *1-5.8 59.2 U5-3 ■VO. 7 79.6 37.9 33.6 100.0 29.3 26.3 Dioxane 0 C)\.. 1 21. k >j.6 .0 i-l-0.0 32.6 60.0 19. *1- 79.7 8.50 constant at two temperatures ( j'i^ures 5-a and 5-b ). At each temperature the rate constants for both systems lie along the same line. We might conclude from these curves that tho di electric constant of tho solvent is more important than the nature of the solvent. The sodium salt was decomposed at 75° In 95 percent al cohol (used as obtained from stororoom). The first order rat re — 2 —1 constant at initial concentration 0 . 0 0 9 was 8.9 x 10” hrs” which is about five times greater than tho value in glycol at 75°• This value appears to.fit the curve of k against D only approximately (Figure 5-b). However, this is on the very steep part of the curve and more data at low dielectric constants would be desirable. :: v . u Uii
ft
1“
f t
I ...
TrfTrHlH
m trH :M z
1
I! ft
!
thi
B
1
I
*••• •«* # • # s
II uSis'iH! itiiss** Effect of Concentration. The data in the glycol-water solvents present the only picture of concentration effect. In all cases tabulated in
Table II except at 7f>° in water, the lower concentration coincides with a lower rate. The difference in rate doesn't seem to follow any pattern either in absolute difference or in percentage difference. No data are available for this determination in dioxane-water solvents. The increase in rate with increase in concentration Is difficult to explain.
Effect of Added Base. The influence of added basos was studied to evaluate the 2 suggestion by Schenkel that a base attack on the carboxyl carbon would speed the reaction. The decomposition of the
3odium salt was therefore carried out in water at 9£.0° in the presence of 0.01 molar sodium hydroxide and alao in the presence of 0.008 molar pyridine. In neither case did the added base affect the velocity constant (Table IV). Pyridine
Table IV Effect of Bases on Rate of Decomposition of Sodium p-Toluenesulfonylacetate , Calc'd Init. Cone. k x 10 hrs” from tab
.089S No added base_ 3.02 (avgj .09214. 0.01 molar OH” 3.02 .0892 0.008 molar pyridine 3.01 - 2J+ - ^uould be more likely to show any catalytic effect since it is a neutral base and would not be hindered as would the hydroxyl group by the negative field of the sulfone group*-20 .
3ffect of Solvent on the Activation linorgy. The activation energy for the decomposition of the sod ium salt in ethylene glycol-water solvent was determined from a plot of the logarithm of the rate constant against bhe re ciprocal of the absolute temperature (Figure 6) according to the equation
d log k _ -iii d'(Y/TJ' 1 = 2.303R For dioxane, since only two toiTi^raturos were available, the activation energy was calculated directly from the integrated form of the equation,
iS = -2.303 R log k^ - log k£ l/Ti - 1 / T 2 The results of these calculations are tabulated in Table V for two initial concentrations in glycol-water solvents and one in dioxane-water solvents.
Within, the experimental error, the activation energy does not change with increasing glycol content and apparently doesn't change with dioxane content although at 6 0 and 7 0 p e r cent dioxane, a slight decrease may occur. The data are not adequate to determine this with certainty. The average value in the glycol-water solvents is kcal per mole. In diox- ane-water solvents the average value i3 3 3 • The constancy W* f*
I
1
i':: r'
:.iit' - 26 - of activation energy I3 the first relation of this type ob served in studies of anionic decarboxylation. In the decom position of tr ichor acetate ion in ethanol-uater, dioxane- uater, and formamide-water solvents, and of trinitrobenmoate ion in dioxane-wator solvent, the activation onorgy increased with increasing amounts of irater. An eiqjlanation of the dif ference In behavior Is offered in the discussion.
If the activation energy for each system. Is assumed con stant and equal to the average value, It may be used to de termine the change in the "entropy factor," log s, In the
Arrehenius equation — V /n'T* k = s e The values of log s are tabulated in Table VI. Values for log k were taken from Figure 6 at 1/T equal to 2.7if0 for cal culating log s In the glycol-water solvents. In dioxano- water solvents, log k at 05° and davg. of 33*7 "an used. This leads to different values of log s in water for each system but tho trend which is more Important than the absolute value Is better observed In this way. In ooth solvent systems, log s decreases with increasing water content and it is this factor alone which accounts for tho change In rate with change in solvent* This trend Is, like the trend in acti vation energy, unusual for this type of reaction. In the four systems previously mentioned, the log s values increase with increasing water content. - 27 - Table V Variation of Activation t&aergy with. Solvent Composition Weight Percent kcal./mole
G\lycol init. cone. 0.1 molar 01 molar 0.0 3^.6 31'-.5 20.5 35.5 35-«5(a) 39.9 33.6 hi.8^a; 59.2 35.7 35.6 79.6 33.8 35.1 100. 0 33.6 33.1
Avg. (.1 + .01) = 3i;..5
Weight Percent Dioxane 0 3.’i-.6 21 . Ll 35.6 lj. o. o 35.6 60.1 32.2 60.0 31.0 79.7 3^*0
Avg. = 33.7
(a) This value was rejected on ■.sis of 00or data for determining rate constants.
Table VI Variation of Lo ■;ith the Solvent v^om )0 a j. oion ed from Data of Initial Ooncentr •at ion C .1 mol Ight Percent Sthylene Glycol log s 0 .0 18.97 20.5 1 9 .0 8 39.9 19.21 59.2 19.39 79.6 19.71 1 0 0 .0 19.37 Dioxane 0 .0 lO.tj.8 2 1 . U. 18.79 LlO.O 1 9 .1 8 60.1 19.51*. 6 0 .0 19.50 79.7 19.92 - 28 - DecompositIon of p-Tolueneaulfonylacetic Acid. The decomposition of the free acid was studied In water, ethylene glycol, and dioxane. In water, the rate of the re action was much, slower than the rate for the sodium salt. In dioxane, no reaction occurred In hours within the experi mental error. However, no comparison with tho salt in this solvent can bo made as it was insoluble in this solvent. The decomposition of the free acid in ethylene glycol was compli cated by a very rapid esterificatIon. However, qualitatively, •men the amount of ester formed, determined by rawlu saponi fication during titration, was allowed for, very little de carboxylation could be detected. This supports tho concept that the anion Is the decomposing subs banco s ince in ethylene glycol the degree of dissociation rrould be considerably less than In water even though tho rate of decomposition of what ever ions are formed Is eight tines greater.
The data for the acid in water (Table ICCCII) wore treated as if the anion was the unstable substance, i.e., the rate of decomposition of the acid is proportional to tho square root of the acid concentration.
For the equilibrium of a weakly dissociated acid HA If1" + A" a constant can be written,
K = - 29 -
[A*] = K1/2 I S A ] 1 / Z Since d ^ L a k fn"] = k K1/2 [HA]1/2 , (1) integration leads to
[HA]1/ 2 - t^lo^2 ~ " V 2 k Kl>/2t« (2) If a plot of the square root of the acid concentration against tine is made, a straight line should be obtained, the slope of which will yield a value for k. When this was done, a straight line was obtained. dinco tho acid dissociation con stant was determined earlier at 26°» it was dotorriino a at 95° from the initial slope of a concentration against time graph and equation (1) using the initial stoichiometric acid concen tration. The value of K determined in this mannor Is 1.62 x
10”3. If tils value Is then used In equation (2), graphically, at the other concentrations, k Is determined to be 2.95 x 10~2 hrs“^ compared to 3*02 x 10”2 for the salt. Tills is good evidence for the belief that It is the anion which decomnoses. 30 DISCUSSION affect of Solvent on the Activation Energy Before considering the effect of solvent on the activa tion energy of this reaction, It is necessary to examine first the essential feature of the decarboxylation reaction, namely the separation of the carboxyl group from the electronegative residue R with retention of the bonding electron pair by R.
R * ^ C R* + CO2
Any change in R which increases tho electronegativity of the
group will aid the cleavage of the carbon-carbon bond, and any change in the carboxyl group which Increases its electro negativity will hinder the transfer of tho carbon-carbon eloc
tron pair to the R group. 11 j ^ iy In previous Investigations, * * * it was found that the activation energy of the decarboxylation reaction decreased with decreasing water content in mixed solvents. Some re presentative data are reproduced in Table VII. The decroase
in activation energy has been explained^ uby assuming that the Ion is solvated In each solvent and that the ions sol vated by different substances and to a different extent show
different degrees of stability. If solvation exerts a stabi
lizing influence on the ion, in the some manner as, but to a
less degree than, the addition of a proton to form the undis sociated and completely stable trichloroacetic acid, the dif
ference in reaction velocity is immediately accounted for. - 31 - Table VII Activation Energies for the Decomposition of Various Anions In Sevoral Mixed. Solvents
TrIchloroacetate Ion 2,2}.,6,trinitro- ln in benzoate Ion in Percent'1(a) ; Organ- alcoho}.- formamidj^- dioxane-watorH- ic Solvent waterb ixater
0 36.23 36.20 35.8 20 35.51 35.10 32).. 0 35.18 3U.10 30.5 48 3^. 10 33.20 27.0 75 - 32.5 80 3 2 .ii-6 ■ - 23.7 8 5‘ - 32.00 90 3 2 .I1-7 31.75 22.0 10 0 3 0 .I4.V
(a) Weight percent for the alcohol and dioxr.no mixtures; mol percent for formamide.
The more completely unsolvatod tho ion, the 10re rapidly would it decompose." If decomposition involves dosolvating the ion, the energy of dosolvation would awooar In the energy O r' of activation.-^ In water, ethanol, and aniline, Vorhook^ found the activation energies for the decomposition of tri- chloroacetato ion to be 3 6 ,02l5 , 31>250 and 2 6 ,7 6 0 calories respectively, decreasing with decreasing ability of the sol vent to solvate the Ion.
In the case of the two solvent systems I11 h i 0V1 the de- 7 composition of sodium trIchloroacetate was studied, Cochran1 pointed out that at the same mol-norcent of water, the acti vation energy Is the same in doth the formamide- rater end aleohol-water system. Thus there would ao -.ear to be a rela tion between the activation energy and thewater concentration - 32 - and to bo no relation between tho decrease In activtion onor0y and tho dielectric const,ant of the solvent. Tho dielectric constant of formamido-water solvents has boon as3uried to vary from 80 In water to 81j. In pure formamide. Extrapolation of
Loader's values to 55>°C indicates a value of 9 7 for euro formamide, but small amounts of -rater ’markedly lower tho value to tho value used. Thus the dielectric constant for this sol vent is essentially constant while tho value In tho alcohol- uator system ranges from GO In ’.rater to about 20 in alcohol.
This relation to the quantity of water present would lead one to consider the specific solvent present and its Interaction with the decomposing anion.
As Verhoek^ pointed out, the solvent interaction -with the anion is similar to the addition of a nroton to tho carboxyl p group. Schenkel further refined this idea with tho concept of hydrogen bonding of tho solvent molecule with the oxygen of the carboxyl group. Hydrogen bonding tends to attract elec trons, I.e., to shift the negative charge to the site of the
drogon bond. In the carboxyl group, therefore, tho carbon atom bee02.10s positive and the attraction of that carbon for the carbon-carbon electron oair Is increased thus hindering the cleavage necessary in this reaction. Thus when trichlor- acetate ion i3 being decomposed in a solvent -with Increasing
H-bonding ability, e.g., water content increasing, the acti vation energy increases. - 33 - However, In the case of the p-toluonesulfonylacotate ion there are t\g-o centers of interaction with tho solvent, the carboxyl group and. the sulfone grou". If this interact ion is considered to involve hydrogen b o n d h ; through the oxygons of tho carboxyl group and also through tho oxygons of the
3ulfone group, then the results obts-inecl can be explained.
Hydrogen bonding tends to- attract electrons, i.e., to shift the negative charge to the bond iite. In the carboxyl group, therefore, the carbon ator.i becomes positive and the attraction of that carbon for the carbon-carbon electron >air is increas ed thus hindering tho cleavage necesaaiy In tho reaction.
In the case of hydrogen bonding of the solvent to the sulfone group, however, the reverse of this electron atti*act- ion occurs, i.e., the carbon-carbon electrons are dravm closer to tho «t-carbon thus aiding the decarboxylation reaction.
Therefore it can be seon that solvent interaction with tho acid anion operates in two op--osin,g directions for the
•(-arylsulf ony 1acetate ion but in only one direction for the trichloroacetate Ion. When a solvent with less ability to form hydrogen bonds is Introduced as part of the solvent, both of these c -nters are affected, apparently, equally.
Thus the gain brought about by lowering the solvation of the carboxyl group is balanced by the loss when the sulfone-sol vent interaction Is decreased and the actlv tion energy re mains constant.
To assess the effect of sulfone-solvent internction, some - 31*. - substituent could be introduced into either tho anion or the solvent with the object of selectively react in- with tho sul
fone group ^nd thus leaving the carboxyl group alone to in
teract with the solvent. In this way it might be possible to
determine the dependence of activation energy on increasing
water content. Introduction of an hydroxyl, group ortho to the sulfone would probably allow strong intranolocular hydro
gen bonding such that bonding to the solvent would be negli
gible. The activation energy of such an anion should there fore increase with increasing ’water content in a nixed solvent.
A similar but smaller effect would be noticed If the el ectronegative character of the sulfone group was modified by
introduction of nitro groups, halides, or methyl groups in tho ortho or para positions in the ring. fliese groups -would change the electronegative character of the sulfone oxygen without affecting the carboxyl oxygen. Thus tho effect of changing solvent would be different on the two centers in the decomposing anion and a change in the activation energy would appear. These experiments have not been done in the present work.
Effect of Solvent on the Sntrouy Factor.
Comparison of the log s values for the decomposition of trIchloroacetate in three mixed solvents and tuose for she decomposition of p-toluenesulfonylacotate in two mixed sol vents shows some significant differences. These are a ■ m i ’ont - 35 - from looking at Figure 7 (The data are found in Table VIII). For the trichloracetate case, log s increases with increasing water content passing through a maximum in both alcohol-water and formamide-water solvents at about 8£ mol-percent water but decreases rapidly in dioxane-water. For the p-toluonesul- fonylacetate ion, however, log s decrea3os with increasing water content with no apparent maximum. Cochran^ has calcu- lated from Hall and Verhoek's data that at 60 C and 0.1 molar solutions the rate of decomposition of the trIchloro acetate ion increases nearly 33 tines in going fron euro water to 90 mol-percent ethanol. In water formanide at the same temperature and concentration, this same change in sol vent composition increases tho rate nearly 132 times. This difference in the change in rate is due to the difference in the log s term because the activation energies are equal at equal mol-percent. In dioxano-water, the rate increases 170 time3 In going from water to gO mol-percent dioxane. however this change cannot be compared directly with that in the other solvents because the activation energy change is five times greater in dioxane-water at £0 mol-percent than in alcohol-watar and formamide-water at 50 mol-percent. It should also be noted that the change in log s Is four or five times greater. The change in rate for the p-toluonosulfonylacotato ion in the two solvent systems studied is also due to a change - 36 - Table VHI Relationship between Xntropy factor (log s) and Solvent for Various Anions OGl^OOg in p-CH^G^H^— 2^4*6—(NOg)^” . sogcn^ooo- OgHgCoo- alcohol6- forma- diox- in in •lvent nide?- ans22- glycol- dioxane- dioxane^ 0 18.23 18.29 20.97 18. 97 1 8.48 17.89 5.6 - - - - 18.79 17.50 6 .0 - - 19.50 -- 7.0 -- - 19.08 -- 8*6 18.15 -_-- - 12.5 -- 18.88 --- 1 2 .8 ---- 19.18 16.15 1 6 .2 --- 19.21 -- 1 6 .7 - 17.43 --- 2 0 .0 18.63 18.35 ---- 2 4 .8 --- 1 9.5 2 (avg) 14*68 2 5 .0 - - 16.67 --- 29.7 - - - 19.32 - - 33.3 - - 16.72 - - t m 36.2 18.35 ---- - 38.1 ---- - 13.16 4 0 .0 - 18.26 16 .06 --- 46.3 19.92 5 0 .0 16.25 51.9 19.71 6 0 .0 17.30 18.09 75.0 17.89 77.0 17.31 85.0 17.73 88.1 16.96 9 0 .0 17.63 «w 9 2 .8 16.33 100.0 16.06 19.87 a - Log a values were calculated from data found in the references given 0 3K; ** $ * . m ■ •*£ i_> * * - V V- K 4*«& * -■" «r*5 '. <, ff-. ;■• ... 9’ 9* rrssj# ®l '.. ' "■% % *r« «»•** . ».' -<■ ■ %• *■ t,. i**p « ; * £- «t i.* * *+' A.b. *•«*#ml W' 1 ' »*ir <»aai ■•&»«»► „ v •-- >■ *49 "“ *■*.» ^&- j *t*B.*if!®8fsa»i s- aiiavva : ««*> . .v»«e-^•r#' k’ ■■^. * V.,■ ■**#:*. * * * : - ,. I.' **fe tmadRssakvciiaBi.. ••**« «M«ttare t1-''.■* "ti1'«t1= , *..•- ;• ' ■-% 5 jh - -• ■-'«% % v • tr9«» *»&*»•» - .... *»■■! M » « M B ,.• ■.•>« ■• < '. •••';*«■••.•-* ■r'- V. ’•! r .■;• ,r’" . av*,--- v^, .-.. •»*-»•■*. _... ^ » gr» k «Sv a«I «8«*««M ■'.• ■** *■ •' ■ s*i*'»j s»»»fc i«aai ■r- -* >• •'- * “» * * ■ •>*; •■ * 4 m r > > *• « . v - r - . * aai?p ■■<*■; - • . * ,-. .. F.«s«ia.’ilB ; " . a--*a » :f. * --*»*» Faaoar * m 9 * • ■ ■ ^ . ■ <•' •*.%- n y *»•;» ««wcat»v ram y. ’aav^ *««-' .--■ •• ?..•- »<•«* -■ <■>■ r* ?* r/* ? v. *• ft-*« 'Awaaaaa^.' ' *«« w.. sewv r«fi*-:S'»»ii>%--ia'a.#w*a«*»am**aa»«*» *«•» • — m r -.»•'•* .* v^Haaaaaaaaaaiir ^ke 1 •* •■'.•> + * < - * a♦■«•'••■ >«aaaRfi« . .. «aiar«pr -E-». :* '*■-,;-» • »•-{(-(**•■* t?**- a - V . ", ^ ^ - aaasaaiii - iBaiaaBa * ’.« v « •» » ^ ^ »•' ■,, .Mar . *£» fV -■ ; * S: '■ ’V w — '■> *« ■ 'V *.: .*• •.• f iiiai9«r«iiB«iii»«*«
1 0 0 mol-percent glycol compared to a lp-fold increase in going from pure water to 1-1-6 mol-percent dioxane-water,
Log s for the tr initrobenzoate ion dec --nos sharnly with decreasing water content in dioxane-water solvents without passing through a maximum. Since no data are available for this ion in a sccmid solvent system, it is not possible to compare its beoavior with the results of the iresont study except to note the largo change in log s for trio trinitrobon- soate ion as comparod to p-toluenosulfonylaootate ion. hat for the trinitrobenzoate ion the activation energy cmango is large.
Comparison of tho trichloroacetate ion decomposition in three solvent systems show specific solvent effects to bo very important. Comparison of tho throe anion3 in tho same dioxane-water solvonts show specific anion effects. both these facts must bo considered in discussing change in rates with changes in solvent composition. Log s is generally related to steric effects in organic reactions. Changes in position or r . - l . z o of .groups cause changes in the rate of reaction. Such storic effects usually bring about changes in the entropy factor, log s, and in tho - 39 - activation energy. But the decomposition of p-toluonesul- fonylacetate ion is accompanied in a changing solvent with a variation in log s but no corresponding change in activa tion energy. Thus it would apoear that with this anion, a differentlation can be made between those steric effects which are accompanied by changes in both activation and en tropy factor, and those effects which are accompanied oy a change in entropy factor alone. If the steric effect is due to the orientation of the solvent which is bound to tho de composing ion, we may be able to relate tho orientation of various solvents to the value of log s in those so?, vent s.
The degree of orientation is probably proportional to the strength of the hydrogen bond formed between tho beco nosing anion and the solvent. If we examine the expression for the rate constant where the activation energy is constant but log s is not we may bo ber.ter able to see the influence of solvent. In the expression for the reaction rate conscant h de rived by the absolute reaction rate theory,
AS* -AH*
where Ain is similar to the exoer ime nt a 1 ly de t e rmine cL, ac t i- AS vat ion energy, the s terra includes o.n entropy term e .
The expression A S * is defined by tho equation
AS* = S* - SR - 14-0 - in which S* represents the molar entropy of the activ ted com plex, and represents the molar entropy of the reactants. For the case of the unimolecular decomposition of a sol vated ion, in solvent X, the concentration of activ .ted ions, I*, is determined by the equilibrium
IXh ^ I* + nX The number of solvent molecules bound to Lhe reacting ion is
Indicated by n and may vary for each solvent. Then for this reaction
A S * = S* - S~u = S I± + nSx - S when the subscripts indicate the species whose entropy the symbol represents. It can bo safely assumed that tho ac tiv tod, desolvated ion I* is in a state independent of its previous con dition of solvation and its can also be considered depend ent only on the ion. However, both SY and STV will vary A ±An with solvent composition and the difference between their values will determine the value of AS*.
If W and G represent water and ethylene glycol respect ively, tho entropy equations can bo written
AS* = Spt. + mSw - S
A S £ = SI± + nSQ -
The vc.Iu o g for Sy and Sq may bo obtained from the literature'” ^ at 2f?° and are 19.7 and 39 • 9 respectively. At the temperature of the reaction they can be calculated from heat capacity - Ill - data to be about 20 and lj.0. This difference is a reflection of the greater orientation or lac’c of freedom on the mart of water compared to ethylene glycol. This same orientation i3 probably found in the solvated ion; the hydrated ion i3 more oriented than the glycol solvated ion and, will have therefore, a smaller entropy value. Thus both Sx and Sj^ are smaller in the case of water than for glycol but, as is determined ex perimentally, is less than ASq. At oresent, unfortunately, there ’ s no ’ray to eilculate values for Sj^ from theory so that AS^ cannot bo calculated and checked against experiment• There is also the further complication of determing n for each solvent. however, if the reverse calculation is carried out, i.e., from values for
A S ± and S-w- and estimating n, calculate 3TV for v rious sol- ^ lit, vents, the difference in values for v riou3 solvents might give a quantitative measure for solvating strength relative, say, to water in so far as steric rigidity may be considered to be a measure of solvation. To verify the validity of this concept, data on the decomposition of p-toluenesulfonyl- acetate in formamide-water and alcohol-water systems would be necessary, and data on the decomposition of trichloroacetate in the ethylene glycol-water system would be desirable. We are forced to conclude on the basis of the work re ported here and. that of previous invostigators that two ef fects of solvent are important in decarboxylation mechanisms. - k 2 - The specific solvent effect, usually the larger, is reflected in the change of activation energy, or lack of change deoend- ing on particular anion Involved. The general effect of di electric constant is noticeable only if the specific solvent effect can be minimized as it is when the decomposing sub stance is the type investigated here. It ^ uld be of interest to determine the rate of decom position of p-toluenesulfonylacotace ion in formanide-water mixtures for it has boon shorn that the ach*v tion energy is constant ana independent of solvent composition nv!. from fig ure 5 it can be seen that tho rate constant den ends on the dielectric constant on! y, part icularly at blue hi. her dielectric constants. Tnerefore, in the foimai ’.ide-water system, in which the dielectric cons bant ic constant, no appreciable change in rate would take 10 lace over .ho comp la to range from oure forma mide to pure vater,
Effect of Ion Association. In any work dealing 'rith non-aquoous solvents and mix tures of lower dielectric constants, the 'Possibility as sociation of electrolytes Into ion >airs must bo considered. Since no data are available for sodium o-toluenesulfonylacet- ate, the magnitude of association for this o.Loc troly te can only be estimated by comparison with other y stems. BezraanP i 1 r has shown that even in etnunol sodium triniuro- benzoate has a dissociation constant of 0 . 0 7 3 5 and that - h 3 - hydrochloric acid has a dissociation constant of 0 . 0 1 1 3 in ethanol. The latter value changes rapidly with increasing waser constant until at a concontration of 2.39 nodes of water per liter no association id detectable. Similarly for ammonium chloride, the value of the dissociation constant in creases with water content ruite rapidly from 0.0167 in ethan ol to 0 .0 3 li in an alcoholic solution containing k.kll moles of water pel' liter, changing more ran idly at higher orator concentrations. In 90 oercent ethanol, hall estimated the dissociation constant \,o he 0.06 for sodium trichloroac state. o r' ifuoss ot al"-° iound K for fcetraisoamyl am ’onium nitrate in
79.0 oorcent d lore me at 2 5 ° to bo 0.0009. In general, sn.ll amounts of water rapidly increase the dissociation constant values. Lihowiso for large ions as sociation is less so that ion pairs lihe trini t r obemzoato or P-toluenesulfonylacet-.to would have larger dissociation con stants tho.n ammonium chlorido • The effect of a .Large sulfone group is to sore ad tie charge out, effectively iacrenslng the ion size and decreasing the amount of association. Use of values corresponding to those given above indicate that sodium p-toluenesulfonylacetate night be as little as
3 0 percent dissociated in the solvent mixtures of lower di electric constant. Assuming b >at ‘die ion -"airs do not do- / compose as has been sho’.m for other cases,0 bhis would in crease the values of the velocity constant if calculated for - [jli - the decomposition of the free ion. Little change would be made in tho activation energies since the dissociation con stant is not greatly temperature dependent, and the effect would be to increase log s in the solvents of lower ’.rater content still more. It should bo noted that if ion association is a factor, the velocity constant would be expected to decrease with in creasing concentration, contrary to the observation in the ethylene glycol-water mixture. - h$ - SUl-n-IARY The decomposition of sodium p-toluenesulfonylacetate in ethylene glycol-water find dioxane-water solvents has boon studied.
The reaction has been found to be first order with res pect to the anion forming n-toluenemethylsulfone and carbon dioxide.
The rate constants at throe t ervo e rat lire s have boon de termined and have been found to decrease rith inerons ing -.rater content in the solvent mixture. The ac tivation energy has been calculated and :;ho:m to be constant over the complete range of glycol-water mixtures,
and to be apparently constant from 0 to oO aercent dioxane. Tho activation energy is 3h»P lccal./mole. Added base has been sho-.m to have no effect; on tho rate of decomposition. The effect of solvent on the decomposition has been dis cussed with a distinction made between dielectric constant effects and specific solvent effects.
The decomposition of the free acid las been found to be one-half order in water, but does not proceed at all in
dioxane. The acid dissociation constant is x 1 0 " in -water at 2p°.
i - Ip6 - SUGGESTIONS TON FU-JTTTOd bOIiK The effect of concentration on the rate of decomposi- tion in dioxane-'.rater solvents should be investigated. Trie decomposition should be studied in other nixed sol vents, e»g.» formariide-vjator and alcohol-vater, to deton.iine whether the effects so far observed are fortuitous or signi ficant with regard to the rate and dielectric constant re lationship and the constancy of tho activ tion energy. The effect of changing substituents on the bens:one ring could lead to inberesting results since the or oport ies of TOo sulfone group, particularly with resnect to 'I-bond in c an be thus changed. The introduction of s. -ehenol "roue ortho to the sulfone onens th.e sonsioiliby of In brand ocular hydrogen bonding which s’oull proo ,bly almost cor.p.htely eliminate sol vent bonding to the sulfone. Thus in changing solvent, such as an alcohol-water series of solvents, the ortho byr:.roxy- sulfonyl anion night be expected to boaavo similar to uhe trichloroacetate Ion in t'n.t solvent. Substituents on the benzene ring such as methyl, -slilon, or nitro groups would not have as dramatic effeet out should alter tho erooorties somewhat in a solvent series. In any given solvent, it would also be of interest to determine the effects of ring sii.bstitu.ents in an attorn h to evaluate the structure and nature of the sulfone roue. W - #5 The effect of cations 3hould bo investiyatod to deter mine if ion association can becomo important. Use of sub stituted amine salts in solvents of hiyh> .glycol > o or alcohol concentration in *rhich the acid mould act as a 'realor acid niyht indic.ate a dependency of the docor:p o sition rate on th.o base strengths of the .'mines. - >1-9 -
Table IX Run 3 l \ . . Sodium p-toluenesulfonylacetate in water. Calc, initial conc. 0.09068 molar; requires 9.91 ml. 0.1C0[ir HIIaOH
95° 85° 75° Time ml. Time ml. T llTlQ ml. (Hour s) UaOH (Ho u r s ) IJaOH (H o u r s } HaOII
0 9.79 0 9.07 0 9.90 3 9.07 5 9.£6 8 9.75 8 7.93 13.5 8.98 22.^ 9.52 13.5 6.83 23.5 8 .I4.O 31l.O 9.35 22.g 5J»-6 >..0 7.79 Is-7.0 9 .1 5 33.0 Is-. 10 h - i . 0 7 . 1K 73.5 3.7 5
Table X Run 26. Sodium p-toluenesulfonylacetate in water. Calc, initial conc. O.O9136 molar; reqaires 9.55 ml. 0.1006 HlaOH 9^0 Q ^ O rj^O
Time ml. Time ml. T ime m l . Hours) ITaOH (Hours) HaOII (Hours) HaOII .'-x ! i—* 0 9.2J- 0 9 J 4-I 0 9 • HP 3.1 8.59 5.5 9.05 10 9.29 7.5 7.55- 11.0 8.67 214-.5 9.03 11.0 214-.5 7.61 37.6 0. GO 6.75r* nl. 0 / 0 20.1 . <-*+ 37.7 7.12 51.0 29.0 it-. OS 51.0 6.55 69.5 3.29 - 50 -
Table XI
Run 31. Sodium p-toluenesulfonylacetate in 20.£/£ Glycol-water Calc, initial conc. 0.09l6Lj. molar; requires 9.98 ml. 0.1000 UNaOII
95° 05° 7£° Time ml. Time ml. Time ml. (Hours ) fc&OH (Hour s) HaOII (Hours) HaOII
0 9.6k 0 9.90 0 9.93 2 8.99 3.£ 9.£9 £ 9.87 £.25 8.03 9.0 9. Hi 13 9.68 9.0 7.C4 19.0 8.32 30 9.32 19.0 £.01}. 31.0 7.li-6 i-i-£ 9.01 31.0 3.33 70 G. 90
Tablo XII
Run 21}.. Sodium p-toluenesulfonylacetate in 39.9£ lycol-uator Calc. Initial conc. 0.09162 molar; requires 9.80 ml. 0.09990 UNaOH 95° 85° 75° Time ml. Time ml. Time m l . H o u r s ) ITaOH (Hours) HaOII (Hours) HaOII
0 9.37 0 9.£6 0 9.62 2 3.38 h 9.02 l'-.7£ 9.1i-7 il- 7.60 9.2£ 0.38 11 9.27 7 6.k7 i£ 7.79 20 G.9£ 11 £.09 22 7.07 32 u.Uu'' /. n 20 3.i|-£ 3 b - 6 . 01}.
L - 51 - Table XIII Hun 22. Sodium p-toluenesulfonylacetate in 59.2^0 clycol-water Calc. Initial conc. 0.0QIl21^ molar; requires 9.76 ml. 0.0999$ NNaOI-I
95° 85° 75° Time ml. T ime ml • Time m l . (Hours) HaOII (Hours) HaOII (Hours) HaOH 0.0 9.2|-0 0 .0 9.11-5 0.0 9.72 1.7 8 . % 3.0 8.95 5.0 9.51s- 3.0 7.73 6.0 8. lj.6 12.0 9.28 5*5 6.55 10.75 7.79 21.5 8.82 9.1 5.19 19.75 6.60 31.7 8.57 12.5 5.26 I'Il.O 8.08
\ Table XIV
Run 25. Sodium p-toluenesulfonylacetato in 79.65 slycol-vrator Calc. Initial conc. 0.09051 molar; requires 9.65 ml. 0 .09998 IHTaOII
95° 85° 75°
Time m l . Time ml. T ime ml. (Hours ) HaOII (H o u r s ) HaOII (H o u r s ) HaOII
0. 8.56 0 . 9.33 0 . 9.58 1.1 7.1* 3.25 8.38 9.12 2.0 6.65 7.0 7.21-0 11.0 3.72 3.9 5.33 11.0 6.11.9 19.75 3.07 6.0 5.16 19.75 11-.90 30.5 7.21-0 9.0 3.02 30.5 3.50 [[.6.0 6.2|-5 - 52 -
Table XV
Run 32. Sodium p-toluenesulfonylacetate in 10 0,$ Glycol Calc, initial conc . 0.09222 molar; requires 9.96 ml. 0.1000 IHTaOH 0 0 tr\ CO O' 75° Time ml. Time ml. Time ml. (Hours) UaOH (Hours) HaOII (Hours ) NaOH 0. 9.00 0. 9.62 0.0 9.08 1, 7.2^ 1^.1 7-ljJU- 2.9 9.38 2. 5.81 7 . 6.28 9.5 3.1,4 3. I}..72 13.2 kJl-2 1^— 1 L • 6.90 k > 3*92 21. 3. Oli. 30. 5.32 5.25 3.1)+ 30. 1.91 53.5 J-.28
Table XVI Run 14. Sodium p.-toluenesulfonylacetate in vater Calc, initial conc. O.OO9I1.O9 molar; requires 9.80 ml. 0 .01000 NilaCII 0 95° 85 75° Time ml. Time ml. T irae . il. (Hours) HaOH (Hours) HaOH (Hours) HaOII 0. 9.80 0. 9.80 0. 9.80 3. 9.20 Ij.. 9.56 7. 9.68 7.5 8.02 15.5 8.72 19. 9.1-0 15.5 6.1.6 31.0 7.80 1-2. 9.10 27.0 J1-.72 lt-5. 7.1'-° 67. 8.30 lj.2.0 3.ii5 1|.8. 6 .$2 5i|..5 2.36 67.25 6.05 67.0 1.77 - 53 -
Table XVII
Run 15• Sodium p-toluenesulfonylacetate In 20.5.5 f.*lycol-wate] Calc. Initial conc. 0.009389 molar; requires 9.71 ml. 0 .01000 NNaOH - f r-'O 95° 85° 7 p Time ml. Time ml • Time ml. Hours ) NaOH (Hours ) HaOH (Hour s) HaOII
0. 9.71 0. 9.71 0. 9.71 2 . 9.26 3.25 9.80 1 6 . 9.60 5. 8.35 17 .So 8.51 2 0 . 9.31 9. 7.38 28.0 7.68 US • 5 9.03 16.25 5.90 US.5 6 .63 66.5 G.lt.O 20.0 U.22 US. 5 2.73 66.75 1.32
Table XVIII
R u n 2 1 . Sodium p-toluenesu.Lfo^ylacetate In 39.9.5 qlycol-v/ater Calc. Initial conc. 0.009059 molar; requires 10.22 ml. 0.009908 ITlTaOII 95° 05° 75° Time ml. Time ml. Time ml. H o u r s ) NaOH (Hours) NaOH (H o u r s ) HaOII 0. 10.10 0 . 10.15 0 . 10.13 1. 9.83 2. 10.10 3. 10.12 2. MS s. 9.85 8.5 10.03 3.S 8.9S 11. 9.51 1U. 10.02 s .s 8.20 2k. 8 .S3 21l. 9.73 8 .S 7.37 36. 7 .75 36. 9.70 - $Ur -
Table XIX Run 17* Sodium p-toluenesulfonylacetate in 59.2,2 glycol-water Calc. Initial conc. 0.008792 molar; requires 10.50 ml. 0.009988 IlNaOH
95° 85° 75° Time ml. Time ml. Time ml. (Hours) NaOH (Hours) NaOH (Hours) NaOH
0.0 10.50 0.0 10.50 0.0 10.50 1.7 9.55 3.0 1 0 . 0 9 5 . 0 10.16 5.0 7.71 9.0 9.09 12.0 9.96 10.0 6.31 20.5 7.14-9 21.0 9.60 20.5 1^.05 31.5 6 .L|.2 31.5 9.23 27.0 3 .J4.9 l\. 7.0 5.50 J1-7.0 3. 3l[. 31.5 3 . Ill- 75.0 3.30
Table XX
Run 2 3 . Sodium p-toluenesulf on3rlacetate in 59.2,2 glycol-water Calc, initial conc. 0.009l|-50 molar; requires 10.13 ml. 0.1000 NNaOH
9 5 ° 85° 75° Time m l . Time ml. Time m l . (Hours) NaOH (Hours) NaOH (Hours) NaOH
0. 9.90 0. 10.03 0.0 10.10 1.5 9.61 3. 9.77 5.1 9.98 3.0 8.32 7. 9.10 11.0 9.88 5*5 7.09 1 1 . 5 8 . lj.9 20.0 9.57 9 . 0 5.70 20.0 7.39 32.0 6.97 O ^-7 1 1 . 5 5.02 32.0 6.20 i|.b. 0 0.^7 - 25 -
Table XXI
Run 18. Sodium p-toluenesulfonylacetate in 79.6,0 glycol-water Calc, initial conc. 0.009158 molar; requires 10.30 ml. 0.009988 NHaOII
95° 85° 75° Time ml. Time ml* Time ml. (Hours) NaOH (Hours) NaOH (Hours) NaOH
0. 9.55 0.00 10.06 0.0 10.10 2.5 7.38 5.25 8.85 13*0 9. £-|.l 7*20 k § • § 5.70 15.0 23.75 8.6 8.5 ^.32 25.75 5.83 38.75 8.02 7.60 16.25 3 . ^ 1 11-0.75 k .63 k0.75 f f O 23.0 2.97 50.75 3.81 6!;.. 30 O • OO
Table XXII Run 19. Sodium p-toluenesulfonylacstate in 100,5 glycol Calc, initial conc. O.OO887O molar; requires 10.36 ml. 0 .009988 NNaOH 0 CO in 95° 75° Time m l . Time ml. Time m l . Hours) NaOH ( Hours) NaOH {Hours) HaOII
0 . 7.90 0 . 7.38 0.0 9.77 1.17 6 .I4.O 6.30 6.0 8.98 3 . 6 7 6 . 5 5.61 15.0 3.01 6.67 3.61 13.75 k.77 27.0 6.90 23.25 k.15 ii-3.1 7.80 51.7 5.6-8 - 56 -
Table XXIII Hun 1+0. Sodium p-toluenesulfonylacetate in water Calc. Initial conc. 0 .01l681 molar*; requires 9.93 ml. 0.05050 NlTaOII 95° T ime m l . (Hours) HaOH 0. 9.99 5.25 8.37 11.5 7.16 22.0 5.52 3k. 0 3.96 11-7.5 2.85
Table XXIV Hun 37* Sodium p-toluenosulfonylacetate in 21.'|,o dioxane-watnr Calc, initial conc. 0.093k5 molar; requires 9.93 ml. O.lOOlj. ililaOII 85° 75°
Time m l . Time m l . (Hours) NaCII (Hours) HaOH
0.0 9.87 0. 9.90 k.O 9.26 5. 9.70 8.0 8.82 11. 9.k7 15.25 7.86 22.5 9.10 22.5 7.10 35.5 8.80 35.5 5.77 - 57 -
Table XXV Hun 39* Sodium p-toluenesulf onylacetate in ij.0.0,3 dioxane- water. Calc, initial conc. 0 .0914-62 molar; requires 9.93 ml. 0.100i|. NNaOH Q£° 75° Time ml • Time ml. (Hour s) NaOH (Hour s) HaOII
0. 9.72+ 0. 9.86 2. 9.01+ 5.25 9.38 6 . 7.77 13.0 8.71 13. 5.96 2lj., 0 7.91 2i|.. Ip. 2k 37.0 7.06 37. 2.57 53.5 5.91].
Table XXVI
Run 2+1 • Sodium p-toluenesulfonylacotate in 60.1,o diorcano- water. Calc, initial conc. 0 . 0 9 5 0 9 molar; requires 9.9 3 ml. 0.1001*. NNaOII 8 5 ° 7 5 ° Time m l . Time m l . (Hours) HaOH (Hours) HaOH 0. 8.1k 0. 9 .02-t, 2 . 5 6.1*3 3.5 9.50 7.0 U - U 7 8.5 8.63 17.0 2 • 2 1|. 20.5 6.1^2 29.5 l.ll-O - 58 -
Table XXVII
Run lj.3. Sodium p-toluenesulfonylacetate In 6 0 .0,* dioxane- water. Calc, initial conc, 0.09591]. molar; requires 9*91 ml, O.lOOl]. NNaOH
85° Time ml. Time m l . (Hours) NaOH (Hours) HaOII
0. 9.55 0. 9.84 2.5 7.65 3. 9.22 5.0 6.08 6 • 0 . 5 1 9.0 l|—28 1 2 . 7.II-0 1 2 . 5 3.23 21.5 5.93 21.5 1.68 29.5 2;..93
Table XXVIII
Run 14.2 . Sodium p-toluenesulfonylacetate in 7 9 , 7 / e dioxane- water. Calc, initial conc. 0.09771 molar; requires 9.97 ml. O.IOOI4. NNaOII
85° 5° Time ml. Time m l . (H o u r s ) NaOH (Hour3 ) ImOII 0.0 9.09 0. 9.62 1 . 7 5 6.35 2.0 8.58 3.0 I4..82 3.5 7.91 5..0 3.93 6.0 6.85 5.5 2.88 9.0 6 . Oil. 8.0 1.72 11].. 0 - 59 -
Table XXIX Run 13* Sodium p-toluenesulfonylacetate in 92.5;£ ethnnol- water. Calc, initial conc. 0.09^4-0 molar; requires 9.93 Ml. 0.100k NNaOH
75° Time ml. (Hours) NaOH 0. 9.93 1.0 9.20 2.0 Q.ko 3.0 7.78 *4-.75 6.72
Table XXX
Run 35. Sodium p-toluenesulfonylacetate in 0.01 sodium hydroxide solution; Gale, initial conc. 0.0936k molar; requires 8 . 9 k ^ 1 . 0.100k NNaOH
95° Time ml. (Hours) NaOH
0 . 8.02 3. 8.00 7. 7.10 12. 6.07 20.25 IN 6k 31.5 3.19 — 60 «
Table XXXI Run 36 • Sodium p-toluenesulfonylacetate In water solution, 0.008l(.9 molar In pyridine. Calc. Initial conc. 0.09Qiil molar; requires 9.9lj- m l . O.lOOIj. NNaOH
95° Time ml. (Hours) NaOH 0 . 9.82 3. 9.08 7. 8.15 13. 6.98 23.5 5.36 33.5 J+.19
Table r_XXII Run 30. p-Toluenesulfonylacetic acid in water Calc. Initial conc. 0.0837 molar; requires 8.37 ml. 0.1000 NNaOH 950
Time m l. (Hours) NaOH 0. 8.37 2 .S 8.31 6.5 8 .1^ 9.5 8 . 0 6 18.75 7.77 28.75 7.1|5 - 61 - REFERENCES
1. Brown, B.R., Quarterly Rev* £, 131-11+7 (1951). 2. Schenkel, H. and Schenlcel-Rudin, M., Helv* Chlm. Acta 21, 511+-521+ (191+8).
3* Verhoek, F.H., J. Am. Chem. Soc. 61, 186 (1939). 1+. Trivich, D. and Verhoek, F.H., J. Am. Chem. Soc. 65, 1 9 1 9 (191+3). 5. Verhoek, F.H., J. Am.. Chem. Soc. 56, 571-7 (1931+). 6. Hall, G. and Verhoek, F.H., J. Am. Chem. Soc. 69, 613 (191+7 ) •
7. Cochran, C.N. and Verhoek, F.H., J. Am. Chem. Soc. 6 9 , 2987 (191+7).
8. Auerback, I., Verhoek, P.M., and henne, A.L., J. Am . Chem. Soc. 22, 299-300 (1950).
9. Pedersen, K.J., Trans. Faraday Soc., 22, 322 (1927).
10. Pedersen, K.J., J. Phys. Chem. 2§,» 5 5 9 (193!+) - 11. Glasstone, S., Laider, K.J. and Gyring, II. The Theory of Rate Processes, New York: IIcGraw Hill Book Co., Inc., 191+1, p. 1+15. 12. Otto, R. Ber. 18, 151+-162 (l88l). 13* Schimmelschmidt, K. and Thomas, h., U.S. Patent l,939,i'.l6 (1931+) C.A. 28, P17165 (1931+). 11+. Felix, F. and Riat, II., U.S. Patent 2,1+32,1+03 (191+8) C.A. 1±2, P2l+l+0h (191+8). 15. Fieser, L.F., Experiments in Organic Chemistry, Second Edition, New York: D.C. Heath oc Co., 191+1, p. 3&9. 16. Anwers, K. v. and Thies, W. Ber. 2296 (1920).
17. Siebert, E. and Fromm, E. Ber. 55, 1025 (1922). 18. Trivich, D., Ph.D. Dissertation, The Ohio State University, 191+2 .
19. ikerlWff, G., J. Am. Chem. Soc. 1+125 (1932). - 62 -
20. Ikerltfff, G., J. Am. Chem. Soc, £Q, 12l|.l (1936).
21. Bordwell, P.O. and Cooper, G.D., J. Am. Chem. Soc. H* 518U- (1951).
22. Salmi, B.J. and Korte, R., Suomen Kemistlleht1 13, 28- 30 (191*5). ~
23. Parka, G.S., Kelley, K.K. and Huffman, H.I1., J. Am,. Chem Soc. 5 1 , 1 9 6 9 (1929).
2 l \ . m Bezman, I.I., Ph.D. Dissertation, The Ohio State Univer sity, 1 9 ^ 2 .
25, Puoss, R.M. and Kraus, C.A., J. Am. Chem. Soc. 93, 1019 (1933). - 63 - AUTOBIOGRAPHY I, Donald Joseph O ’Connor, was boro In New York, N.Y., March 10, 192i|.. I received my secondary school education
in the public schools of Hex* York, IT.Y. ity undergraduate training was obtained at The College of the City of New York from which I received the degree Bachelor of Science in
19iUl* In 191^9, I received an aooointment as an Assistant in the Chemistry Department at The Ohio State University which
I held until 1952 when I received an arveolntment as a DuPont
Research Fellow. While I held these positions I was complet
ing the requirements for the degree Doctor of Philosophy.