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1948 Dielectric and Solvent Effects on the Alkaline Hydrolysis of Ethyl-Acetate. James Edward Potts Louisiana State University and Agricultural & Mechanical College
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Recommended Citation Potts, aJ mes Edward, "Dielectric and Solvent Effects on the Alkaline Hydrolysis of Ethyl-Acetate." (1948). LSU Historical Dissertations and Theses. 7910. https://digitalcommons.lsu.edu/gradschool_disstheses/7910
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LOUISIANA STATE UNIVERSITY LIBRARY
simsGfEic Aim sqelyss* m m osr m ALKALIES HTiaOLTSIS of K m kQMAm
A Dissertation
Submitted to the Graduate Faculty of the Louisiana State University and Agricultural and Mechanical College in partial fulfillment of the requirements for the degree of Doctor of Philosophy
in
She Department of Chemistry
Vy James Edward Potts, Jr. B.S., Louisiana College, 1939 M.S., Louisiana State University, 1941
1947 UMI Number: DP69288
Alt rights reserved
INFORMATION TO ALL USERS The quality of this reproduction is dependent upon the quality of the copy submitted.
In the unlikely event that the author did not send a complete manuscript and there are missing pages, these will be noted. Also, if material had to be removed, a note will indicate the deletion.
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Dissertation Publishmg
UMI DP69288 Published by ProQuest LLC (2015). Copyright in the Dissertation held by the Author. Microform Edition © ProQuest LLC. All rights reserved. This work is protected against unauthorized copying under Title 17, United States Code
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She author wishes to express his deep appreciation to
Br. S. S. Amis, under whose direction this work was carried out, for his sympathetic under standing, effective criticism end assistance. He is also indebted to Br. Georgs J&ffe* for
M s contribution to the theoretical aspects of this work.
He further wishes to thank: Br. A* B. Choppia, Br. S. L.
Compere and other members of the faculty of the Louisiana
State University Chemistry Bepariment with whom he diecussod various phases of this work, for their advice and criticism.
j. i ii car gobs$B2s
Page I* Introduction .*,..... 1
IX. Heriew of the Literature . . in. BxperiEaentul ...... 1 6
IT* D a t a ...... 2 0
T. Sugary ......
TI* Bibliography...... %2
711* V i t a ......
ill m s ? m t w m s
Page
I* Bepreseatative Calculations of Telocity Constants from Experimental Bata at fbree Temperatures •
a. 0*00° Centigrade 27
9*80° Centigrade * . * , f ...... * 28
c. 1 9 . 9 ° Centigrade...... 29
II. Experimental Values of Velocity Constants as a Function of Ionic Strength; Corresponding Energies of Activation and Prequeney f a c t o r s ...... 3C
III. Experimental Values of Telocity Constanta as a Function of Dielectric Constant ...... 31
IV. Constants tfe©d la Plotting Curve In Figure 4 at the temperatures Indicated • ...... 32
?* Constants for Plotting Curve In Figure 5 at the Temperatures Indicated ...... 33
VI* Enhancement of Moment of Ithyl Acetate as Civen by Couloablc Energy Goasl dsrat 1 one 34 Page 1. Plot of log k versus X/T for various Ionic strengths at D - 6 5 . . » - .35
2. Plot of log k versus l/$ for various dielectric constants at « 0 . 0 2 ...... 3 6
3. Plot of log & versus 1/f for pure w a t e r ...... 3?
Plot of If versus a? showing agreement of Ionic strength data with theore tical curve ...... 38
5* Plot of ¥• versus 1/D2 showing agreement of dielectric constant data with theoretical curve ...... 39
v ABSTRACT
There are two possible approaches to the prediction of reaction rates in solution. One involves the nee of the theory of absolute reaction rates, as developed "by Syring and ethers, in which interest is centered in the nature of the transition state or as it is sometimes called, the activated complex* The ether approach attempts to evaluate as many as possible of the external factors which are known to influence reaction rates, such, as temperature, ionic strength, dielectric constant of the solvent, as well as the presence or absence of electro static charges or dipole moments associated with the reactants,
She first method, while general in its applicability to reaction rates which are a function of temperature, becomes very difficult to apply to reactions in solution due to the complexity of calculating partition functions under such condi tions* In the field of ion-ion reactions, considerable pro gress has been made using the second approach, as represented by the Bronsted— Christisnsen-Scatchard reaction rate equation.
The prediction of reaction rates for ion-dipolo type reac tions in solution using the second approach has been developed by gyring on the one hand, and Amis and Jaffa" on the other,
While the equations derived by the two groups are similar In some respect®, they lead to quite different predictions, The
Barring equation for ion—dipole reaction rates predicts that an
vi increase of dielectric constant will result in a decreased rate constant, and an increase of ionic strength will cause an increased rate constant, regardless of the sign of the charge on the ion. fhe Amis-JTaffe^ equation predicts different results depending on
the sign of the charge of the ion. According to their theory,
in a negative ion-dipole reaction an increase of dielectric con
stant will cause an Increase in the rate constant, while in & positive ion-dipole reaction an Increase of dielectric constant
will cause a decrease In the rate constant. Also, In a negative
ion-dipole reaction, an increase In Ionic strength should result
In a decrease In the rate constant, while a positive ion-dipole
reaction should demonstrate an increase in the rat© constant.
fhe experimental work was undertaken in an effort to decide
the applicability of the Amis-Jaffa'" equation to the prediction
of reaction rates of negative ion-dipole reactions. Since the
findings of this experimental work do not agree with the predictions
of the Byring equation, its use was not pertinent.
•The alkaline hydrolysis of ethyl acetate was chosen as a
typical negative ion-dipole equation which could he used to ob
tain data to he compared with theoretical predictions concerning
this type of kinetic process as related to the variables temper
ature, dielectric constant, and ionic strength.
The experimental part consisted of fixing the ionic strength
of the reaction mixtures by adding pre-determined quantities of
sodium nitrate, and of controlling the dielectric constant of the
media by mixing ethyl alcohol and water in definite proportions
and fell owing the rate of hydrolysis by determining the concentration vii of sodium hydroxide present at measured intervals of time using a ti trims trie procedure. Brest thymol blue m s used as an Indicator,
$he reaction velocity constant was measured at dielectric constants of SO, 6 5 and 5 0 at a constant Ionic strength value of
0 ,0 2, and at a constant dielectric constant value of 6 5 in which the ionic strength was 0 .0 2, 0.05, 0.10, 0 .2 0 and O.3 0, at the three temperatures 0°, 9>B°* and 1 9 .1 °C.
She Ionic strength affect upon the reaction rate w found to he in conformity with the predictions of the theory proposed by dads end Jaffe^ for ion-dipolar molecule reactions.
dielectric constant dependence of the reaction rate while in the direction predicted by theory resulted in large values of the parameters involved in the ion-dipolar molecule rate theory.
the magnitudes of these parameters are discussed from the standpoint of hydrogen tends in the intermediate complex, the extrapolation of dielectric constant data to higher dielectric constants, and coulomb!c energy of activation.
vili » * 3 £ § * I n 3 © © 33 4* O rf© ffi ©1 % rH & s $ iH
I I P I S
5 | :r the u© of « i-4 4» 1 I © I 5 1 © a i I Vl O g © i M 1 3 i S *2S * I i ao 4rtn
I rate constant, regardless of the sign rate constant, of sign regardless the to reactions ia solution due to the complexity of complexity calculating the reactions to solution ia due to 2
The Amis-daffe equation predicts different results depending on the sign of the charge on the ion* According to their theory, ia a negative ion-dipole reaction an increase of dielectric con stant will cause an Increase in the rate constant, while in a positive ion-dipole reaction an Increase of dielectric constant will cause a decrease in the rate constant. Also, in a negative ion-dipole reaction, an increase In Ionic strength should result in a decrease In the rate constant, while a positive ion-dipole reaction should demonstrate an increase in the rate constant.
The experimental work herein described was undertaken in an effort to decide the applicability of the Araie-Jaffe" equation to the prediction of reaction rates of negative ion-dipole reactions.
Since the findings of this experimental work do not agree in trend with the predictions of the gyring equation it will not be dis cussed further.
Although the alkaline ester hydrolysis reaction has been studied extensively for a variety of purposes, in only one instance has a comprehensive study of the effects of mixed solvents of fairly high dielectric constant on the rate been made (26). Like wise, the study of the ionic strength effect on the reaction has received scant attention. It is to be hoped that this work will contribute in some small measure to a better understanding of the effects of dielectric constant and ionic strength on ion- dipole reactions. CM r> © p & © O © © *W s 4* e© « *3 © I s *4 aa © 01 #4 0 # a © » 0 ft M © 1 © iH CM i s i v4 I c I 1 » 0 I M £ *5 *d I I * 3 % g ^ s i § 1 § 8 n 4» © rlg S < O 4* © i g 3 , • © H CM S' *< I % < t* O 1 expreseed form logarithmic expreseed la the © 1 s 5 > £ H 1 E II i ! ©
M may he %4 ! r i J? M J © 4® 3 3 ©I * o i § i r4y « I© &« 1 «t4 $ « $ 3
« equation ¥i *!
I I The S ? O I Syring (18) and others have developed in recent years* As pointed eat in the Introduction, due to the complexity of computing parti tion functions of reactants In the liquid phase, the theory of absolute reaction rates has had hut limited application in the field of solution kinetics to date, for this reason. It will not he discussed in detail* The development of successful quantitative theories explain ing ionic reaction rates in solution m s hindered for many years by the lack of a satisfactory theory of electrostatics in solution* Those theories which were based on the assumption that solute molecules in a solution could ho treated as if they were in a gas phase, and that the kinetic theory of gases was therefore appli cable were Inadequate because they ignored the electrostatic effects between reactants themselves, the electrostatic effects between a reactant ion or molecule and its surrounding Ionic at mosphere, and the dielectric in which the reaction occurs. In 1923 Bohye and Euiskel (15) proposed an Interlonlc attrac tion theory which was found to be adequate in extremely dilute solutions. They postulated the existence of an Ionic atmosphere surrounding each. Ion in which there are on the average more ions of opposite charge than of the same charge as the ion In question. This tendency toward an orderly arrangement of Ions In solution is opposed by thermal agitation. Applying the Boltzmann distri bution law and the relation between electrostatic potential and charge density (Foisson*s equation) they computed the potential 5 on the ion itself and the electrical energy of the ion da© to its surrounding ionic atmosphere. Equating this energy to the differ ence in free energy of dilution of a real and an ideal solution they obtained the relationship between the activity coefficient and the ionic strength, w where (5) and f^ = the ionic activity coefficient A = the Bebye-Hdckel Constant 2^ “ the ionic charge. The success of the Xtebye-Hiickel theory in explaining the effect of electrostatics on Ionic activities led to its application by Breasted (12) and Christiansen (1 3) to the problem of ionic reactions in solution* They postulated that the rate of a himol- ecul&r ionic reaction was a function of the activities of the ions and of an intermediate Hcritieal complex’* which was la ther- modynaale equilibrium with the reactants. This is expressed by the equation k = k, fa f„/ f* (6) where k « the rate constant , k * the rate constant at zero /X. 0 ' V * , * ' * the activity coefficients of the ions A, 3, and the Complex* By suitable combination of equations {^) and (6 ) they derived the limiting equation In k * la + *a A 1 ~ \ljv T (?) where z& , are the loaic chargee and A* is a constant. A plot of la k Yersus-yju. should produce a straight line of slope *a *h A ® e»d intercept la kQ, which represents the velocity con stant at sere ionic strength, This equation is valid in extremely dilute solutions in accordance with the Dehye-Suckel theory. fhe next advance in the field is due to Scatchard {25). Making use of the B o m term (11) for electrostatic attraction between ions in high dielectric media, he corrected for the dielectric constant effect the value of In kQ which Breasted and Christiansen had obtained, and obtained the equation * » % € 2 * o - 1 * * * - B te,«P (ra •*. rjj) <8> where B — the solvent dielectric constant k^ = the Boltsmann constant € = the elemental charge The term la k ^ represents the velocity constant at infinite dielectric constant and aero ionic strength, and (ra *• r^} Is the radius of the Bronsted critical complex. Combining equations (7 ) and (8 ) after making the necessary substitutions for the con stant A* in equation (7 ) we obtain the Bronsted-Ghrlstianssn- Scatehard equation ? 7 which le also as the general kinetic equation for ienle reactions la solution. In this equatlon/f is the Bebye kappa and. equals (8T C f 1000)"* The validity of the general kina tie equation has been attested to by the work of Ba Mer, Amis and ethers (5, &, ?* 8 * 9 >* If one of the reactants is & neutral molecule the general kinetic equation {equation 9 ) Is no longer applicable since the charge on the molecule is sere; i.e., s^ ~ 0. Amis and Jaffe (3) have derived an equation for the rate of such an ion-dipole reac tion. Following Be bye and Buckel they compute the potential of the ionic atmosphere around a dipolar molecule, making use of Gnsager*s (2 1) theory of electric moments of molecules in polar liquids to properly account for the interaction between the dipole and its dielectric solvent, fhls potential Is applied to the calculation of the rate of formation of an Intermediate complex X formed from a dipolar molecule A of moment^ and an im M of charged z^. A description of this derivation follows. If the interaction between dipoles is disregarded thediffer ential equation forthe potential % in the neighborhood of any one dipolar molecule will be the same as for the potential around an ion. For 5 different species of ions, let z^ be the valence of species i and be their number per cap. The differential equation for % then becomes V %Vo=Jtt'Po {1 0 ) where S 2 = (4 ^ 6 2/33k^!P)^n^2i^, (the Bebye kappa) ^ (1 1 ) 8 Here 6 is the elemental charge , the Boltzmann constant , and D the solvent dlelsetrie constant, treating the dipole as a point singularity a particular solu tion of (1 0 ) can he obtained by ''polarizing® the Debye-Huckel solution for the potential of the ionic atmosphere around an ion, vtidi is V i = A* e ^ / r (12) This is done by differentiating (1 2) with respect to z and subsli- tuting for z the expression r cos from the positive z axis), The dipole Is Idealized by a sphere of radius &, The resulting equation is In order to properly account for the interaction between the dipole and the dielectric solvent It is necessary to introduce Oneager’e (2 1) concept of °external moment*, which Is different from the moment in a vacuum due to the enhancing of both the per manent and the induced dipole moments by the intervening medium* Following Onsager the molecule is characterized hy a permanent moment Cf0 and a polar!zability o( which is related to an "internal refractive index, a” by * ^ z)e^ (1^) The modified form of equation (1 3) is then shown to be whereU . * 9 the "external moment* is given by 9 U * = Ut iu z ♦ 2) £b(2 + z t f a j a 1) + n2 + ^ > J (x6) Shis equation represents the potential of the ionic atmosphere around a dipole of momently, It la now possible to calculate the rat© of formation of an intermediate complex X from a dipole molecule A of moment and an ion £ of charge € s^» Let CXr C^, and 0^ he the concentrations of X, A, and B, respectively. Planing a dipole A at the origin, with rotational symmetry about the Z axis, we Introduce the polar coordinates r, 2?*, and ^ • According to Boltzmann* s theorem, the probability of finding an ion £ in an element of relume defined by tee limits of r and {r 4 * dr), 2^*and ^ and { ^ will he proportional to 6 * 6 , axp {-If/ € •^lcb5>*2 sia*^to (1?) In order to obtain tee rate of formation of,X * and thereby the relocity of tee reaction it le assumed that there are sensi tive zone* on tee exterior of tee molecule which when touched by an ion Bt permit the formation o f X » In other words, r t Z?*, and ( p moat assume specified values between rQ, and (r0 +A i ’0) t «=4 ( i % * A l # and ^ a a d (£^Q *4* d ^ Q). Since we hare assigned to tee molecule rotational symmetry about tee z axis need not be specified and equation (1 ?) can be integrated with regard to if** From this integration is obtained the concentration of the complex §* = 2*k* ca Cfc t \ s i n # a r 0 exp (-6 zb % / kbf) (18) where k* is a constant. Collecting the constant terms and equating 1© to S, vs have K a 2 »k* r§ sia'Lj'ar( A ^ (1 9 ) faking the logarithm of "both sides of equation (18) produces ^ = la X - <*oVo)/*b *} <2 0) In the limit kappa equals zero {jt = o) , reduces to In (Cr0 /Ca0 G1j0) = l a * - ^ g * cos r#2] (21) where V * s % Combining (20) and <21} we obtain for the rate of formation of the complex; e *w eosT^ * ^ /*_ la k = In + - 5 - ^ 2 ---- - « - a*,-'"*) ( 1 + / r Q) (2 3) w O 5his equation for ion-dipole reactions corresponds to equation (9 ) for ion-loa reactions. Xu agreement with empirical data cos *1% is chosen to he 41. This equation predicts that In k will Increase with an increase of 1 / 0 for positive ions, and decrease with an increase of l/D for negative ions. Also In & should increase for positive loss as the ionic strength Increases and decrease for negative ions as the ionic strength increases. Experimental verification of equation (23) has been afforded by the work of Amis and Jaff^ (3)• and Amis, Jaffa*, and Overman, (h). The former applied data already in the literature for positive ionr-dipole reactions while the latter applied similar data for negative ion-dipole reactions. Amis (2) introduced the concept of coulombiG energy of acti vation to account for the differences in energies of activation of reactions studied in isodielectric media of various dielectric cons taint s. If too energies of activation of a reaction are measured in two ieodielectric solvents of tla® same solvent pair, the differ ences in the activation energies should he due to a large extent to a change in electrostatic energies, which can he calculated from the expression for the change of coulombic energy between two electrically charged particles* for an ion-dipole pair, the energy necessary to bring the ion from Infinity to within the distance x of the dipole is given by the equation B ( 2 * ) where Z % 6 2 la the charge on the Ion, u^ the dipole moment of the charge in the dipole and the line drawn from the ion to one of these centers of charge. If the energy change necessary to bring the two particles together at two different constant dielectric constants of the same solvent pair is the change only of coulomMe energy, this alteration of the energy of activation can he calcu lated. from equation (24) as follows (25) % » 2 * Z 2?or head-on alignment of the dipole and ion cos -f- / and the equation reduces to A £ = k f -Of c - A P (26) ®l »2 * $he alkaline hyttrelyois of ©ifcyl &e©ia%ef which is generally 1 » I I t I ss CSJ 1 r\ i i 5 § ^II i i I 5 u & $ © f$ © 8 1 * g * © t A ® I j j p-© © © j§ © p - s © $ H £ w £H » *H •H § & *rt 3 G * -<9 H ♦G T* © *H O © © I 5 © A © 5» • < * « © » * : I 4» ! « 5 « § t! K ® » u 4» M q $ t I 5 I I 0 $ 0 H » % I | I 't ® } i H • «n a I S4 ** Jfl ® H * » e | & G 3 H $ *& & O M s o 1 € "s 1 § fe 4* 0 13 hydrolysis of a number of aliphatic esters including ethyl acetate, at 25°G* 350 G# and 5Q°C, in 85 percent aqueous ethyl alcohol . A calculation of the reaction velocity constant at 0°G using their energy of activation and k ^ o gave teQo a 0 .0 ^ 5 liters per mol minute which Is la fair agreement with the value 0 .0 5 2 found by the writer experimsatally at the same temperature and alcohol concentration. Uhey suggest a mechanism similar to that proposed by Hammett (20). Envies and Evans (lb) studied the alkaline hydrolysis of ethyl acetate and other esters in ?Q percent aqueous acetone at b4.7°C, 3 5*0°G, and 2b>.8 ®£. $h«y observed velocity constants which ere lower than those for pure water, in agreement with the ion-dipole theory under consideration here. Smith and Levenson (28) obtained values of k higher than previous workers in 85 percent aqueous ethyl alcohol by elimina ting evaporation losses during the preparation of ethyl acetate solutions. The only extensive study of the effects of solvents on the kinetics of this saponification which was found in the literature was that of SelIvanova and Syrkin (26). fhey measured the rates of saponification at 1$°G, 2$°G, and 35°G, for solutions 0.023 molar in HaOH and ethyl acetate, the solvents being mixtures of water with methyl alcohol, ethylene glycol, glycerol, raannitol and erythritol. As discussed later, their k values la pure water do not agree with those found by others, including the writer. How ever, their k values in mixed solvents decrease as the dielectric constant of tlie solvent mixture decreases in conformity with, the Amis-J&ffe lon-dipole theory described above. Sel Ivanova and Syrkin found a definite correlation between the number of OH groups in the solvent alcohol end the velocity constant of the ester hydrolysis* They waste able to show that* for instance, ten percent solutions of glycerol and mannliol which contain different numbers of mols per liter have the same number of 0 H groups per liter* In both the solutions the rate constant Is the same* Other similar examples are given* This phenomenon Is discussed from the standpoint of hydrogen bond lag between the OH groups of the alcohol and the reactants, and the resulting effect on the acti vated complex* Salmi and Herts (24) studied the hydrolysis of aliphatic esters including ethyl acetate in water and 5 0 percent water- dioxane solutions at 25°C. Their value of the velocity constant in pure water at 25° is in agreement with the data of Eeieher (2 3) and the writer* The velocity constant in $0 percent dioxane solu tion was lower than that in pure water as predicted by ion-dipole theory (3). Lowry (3l)» considers the complex to be a bipolar compound, similar In nature to & Hzwitter ion”. It is evident that such a compound would have a very large dipole moment, as compared with that of ethyl acetate itself* He postulates the following mechan ism v » p 1 1 <** « m O m 1 5 *» *• £3 IS ** r*J 1 R | •H0 £ O © & H 0 * « u $ $ 3 3 *d•r* i g | » 45 # & i i 1 3 1 ? £ * * The ethyl acetate used in these experiments was a colorless sample which was purified as described by Weissberger and Proe- kaoer (3 2). The ethyl acetate was dried with anhydrous potassium carbonate and sue gran of water added for each. 2$Q grans of ester* after which it was distilled through a 3 6 inch column packed with 1 /8 inch single turn stainless steel helices and equipped with a total reflux partial take off still head. Any alcohol remaining in the mixture passed ewer first in the form of a water, alcohol, ester aseotrope. The ethyl acetate which distilled at 7?.1°G, was stored is a pyrex-glass-steppered Srlenmeyer flask. One gram samples of the material were analysed by adding twice the theore tical amount of HaOH solution* allowing the mixture to stand for one hour at 7 0°G and titrating the m u s e d alkali with standard acid. The average of four analyses gave the ratio of mols of Sa&E to m i s of ethyl acetate to he 1 .0 0 2 5* The ethyl alcohol was purified by the method of Toting, as described by Weissbergsr and Proskauer (32). One liter of 95 percent grain alcohol was re fluxed for four hours with 3 0 0 grams of GaO and distilled. The 99*5 percent alcohol was refluxed with calcium metal and distilled, yielding absolute alcohol. F r e s h l y boiled distilled water was always used in preparing aqueous solutions• Carbonate fro© sodium hydroxide stock solutions were prepared 16 1? by d&tsolviag c.p. M S in water in a one to one weight ratio, The concentrated solution was allowed to stand in a covered eoa^- taimer for 24 hours after which the clear supernatant liquor wan decanted and filtered in order to remove suspended sodium carbonate crystals. Stock solutions of 0*500 and O.OSOGlf HaCH were prepared by diluting the one to one solution and were stored in paraff in-lined 1 0 liter bottles ©quipped with soda- 1 im© tubes to exclude carbon dioxide, The 0.0200® SaOH solution used as a titrating agent was prepared in two liter hatches at frequent intervals by diluting the 0,055 solution. The 0.10005 hydrochloric acid stock solutions were prepared from c.p* concentrated acid (3? percent) and standardised against sodium carbonate, The 0.02005 EC1 used in the titration was pre pared in four liter batches at frequent intervals by dilution of the stock solution, The analytical weights and volumetric apparatus used in this investigation were calibrated for accuracy. The thermometers were calibrated against B. S. thermometers. Sine® all of the experiments were performed at temperatures lower than room temper ature It was necessary to circulate Ice water through coils in the thermostatic water bath. The temperature was controlled to + 0,01°G with a Beckmann type mercury in glass thermoregulator anA a Sargent Zero Current Belay which activated the electric heaters in the bath, The rate of the ester hydrolysis was studied at 0°, 9 .8 °* ana 19.1°C in both pure water and in mixed ethyl alcohol-water solutions. Duplicate runs were performed In every case, using the data of Akerlof (1) to compute the proper weight percentages of the aleo- hol-w&ter solutions to give dielectric constants of 80, 6 5, and 50 at d l three temperatures. In addition to the above, runs were made at all three temper atures at a dielectric constant B of 65* in which the ionic strength pU} of the solution was increased when necessary hy the addition of Ha§0^, so as to give/^ values of 0.02, 0.05, G.l, 0.2, and 0.3* The following procedure was used in following the course of the reaction between ethyl acetate and sodium hydroxide. Three hundred ml. of the solvent mixture at reaction tem perature were pipetted into a 6 0 0 ml. kjeld&hl flash. To this was added 130 ml. of Q.0 5H ethyl acetate at the same temperature. She flash was allowed to reach temperature equilibrium in a ther mostat after which the reaction was initiated by pipetting 1 0 0 ml. of O.Uf HaOH at the same temperature into the kjel&ahl flask and shaking vigorously. The solution was initially approximately 0.02H in SaGB and approximately 0 .0 1N in ethyl acetate. After a few minutes a 50 ml. sample was withdrawn and drained into an excess of O.OZOOH HG1 at 0°C. Brom thymol blue Indicator was added and the acid solution titrated with 0 .0 2 00$ BaOE to an emerald green endpoint. The NaGE volume was measured with a 1 0 ml. sacblette auto burette calibrated to the nearest 0.05 ®1* Samples were removed and titrated at suitable intervals during the first 8 0 percent of the reaction after which the flask© were 19 kept at room temperature (approximately 28°0 ) for 2h hours or until the reaction had. gone to completion. The concentration of the KaOH remaining in the flank after hydrolysis m s complete m s deter mined hr replacing the flask in the thermostat and withdrawing and titrating two or three samples., She average of these values was taken as C • The reaction velocity constant* k* was calculated using the equation t * elapsed time (ia- t^) in minutes — HaOH concentration at t^ Ca = SaOH concentration at 0 ® U&OH concentration at end of reaction The derivation of this equation is given by Eeicher (23) * The time t9 of mining was not observed since the actus! concentra tion of reactants at an initial time t^ was determined after mix ing by the same analytical means used in all subsequent analyses of samples. The equation given above does not require an extra polation to zero time. a m The ethyl acetate-hydroxide ion reaction m e selected as a typical negative ion-dlpole reaction* the kinetics of which, would afford a test of the ion-dipole theory of reaction rates proposed b y Amis and Jaffe^ (3) • In Table I are presented represen tat i ve data taken on the reaction Telocity constant at different temper atures In both pure water and water-eihyl alcohol media, The tables are lllustratiTe of both the precision obtained in the experimental work, and of the well established fact that the reaction is free of complications, with respect to order, being strictly second order. The Telocity constants at a dielectric constant of 65 and ionic strengths j(U} ranging from 0 .0 2 to 0 .3 0, and at temperatures ranging from 0° to 19*1° 0 are recorded in Table II. Here are also recorded the energies of activation 4 B and the Arrhenius fre quency factor 3 which were calculated from equation (3). The data indicate the slight dependence of both A 3 and 3 on ionic strength. The decreasing values of these factors with decreasing ionic strength are reminiscent of like trends observed and explained theoretically in the case of ion-ion reactions(7). Figure 1 is a graphical representation of the dependence of the logarithms of the rate constants on l/T. from the differences in slopes of the lines At different ionic strength® arise the variations of with ionic strength ohserred in Table II. 20 21 velocity constants obtained in various iso&ielectric media at an ionic strength of 0 .0 2 and at three different temperatures are given in fhble XII together with the values of A and cal culated from equation (3). The energies of activation and the Arrhenius frequency factors decrease with increasing dielectric constant, The decrease of the energy of activation is predicted by the theory of coulomb!c energy of activation (2). the mag nitudes of the variations are larger than would normally he ex pected from theory as will he discussed later. An Arrhenius plot of the data is given in figure 2. The differences in the slopes of the lines account for the variations in energies of activation noted in fable III. the rate constants obtained in pure water and recorded in fable III are plotted in Figure 3 together with data by other isves tlgaters In the same solvent, The observed values of k in pure water in the dielectric range 8 8 *3 to 8 0 .7 are lower at all temperatures than the k values obtained In ethyl alcohol- water mixtures at a dielectric constant of 80.0. In spite of the fact that the pure water constants are opposite in trend with respect to dielectric constant to those obtained in the mixed solvents, the values are in agreement with those obtained by ether investigators in pure water over the temperature range studied, as is evident from Figure 3» where the author's data together with those of Belcher (23) and Salmi and Korfce (2h) are plotted. Two points obtained by SelIvanova and Syrfein (2 6) are shown in the figure to fall on a separate line of apparently nearly the same slope. She preponderance of data would indicate that the values of k obtained by Selivanova and Syr kin are too low. The reversal of trend of the water points obtained by the author calls to mind the difference in the reaction Telocity constants of the acid inversion of sucrose in pure water as compared with dioxane-water mixtures which occurred in Holmes1 research (6) wad were noted by M i a and Jaffa" (3)* In the present investigation the varia tions in the data in the case of pure water as compared to water- ethyl alcohol are not so great as to necessitate the elimination of the water data from consideration; therefore, in comparison of the experimental data with theory the water runs are included* T o test the theory of ion-dipole reactions represented by equation(23) 6 % e&m2% s* * * » x In k = In + - — .j r-A - ue the equation was transformed to a dimension!ess form by substi tuting in it s = / f a = / r 0 (31) w = (1* k - la ^ Xsn^ftp^/e *b u* cos'll; (3 2) giving W = S2/(l + i + *2/2) + (a2/2»)(l + z) , (3 3) and. a theoretical curre (Jfignre h) of W Tarsus z2 was drawn and the data for the dependence of the velocity constant on ionic strength (recorded in Sable II) were fitted to the curve using the constants of fable IV* It is observed from Figure h that the data fit the theoretical curve well. The radius rQ of the critical complex, the enhanced moments and the squares of th© 2 3 latersal refractive im&ieesi are to tSi® values of these eaaskats f o M in former applications of M u t&tser^ (3} W . ft- 1b significant that the «*£t*» of the intermediate com plex vat the ssee afc dll temperatures and resulted la reasonable and comparable values of «* and a2. fits i m & m & of Severe obtained exir&polatia# k as a fraction of/ < at each te®^ rater© to would require an extrapolation of k t«wa# “\^JT. She enhanced momenta wars calculated using tit© aquation * M e t Is derivable from equation (3 3)- $&» n^ values were 4b** t&tnsd from the expression u£ * ^ ( a 2 ♦ 2}X$2£ ♦ a2) (3 5) using the value 1 * 0 Behfe units for tils ®&®e»t of atfcyl aeatate la v&eao {2 7}* Xu tost lag the dependence of tit© velocity constants for the »Tfai»i» hydrolysis of ethyl sestets oa the dielectric constant of the medla» equation (2 3} was again transformed to a dI^asion- less expression tgr making the eabsfcit-aiion !«**•*'*•/ (3 6) and K' : (In i£ - lfi 1^ X 2**?)/6 ■£ e»8 ?S0 (3?) which resolto in the foil owing function S' = { i f & m / l +"&/# + -i/2S) (3B) 03lag this equation a plat of W* versus l/D2 (figure 5) was and the dielectric dependent data which is recorded in fable III was fitted to the curve using the constants given in fable V value of rQ used was 7 Angstroms comparable to the value of 7.5 Angstroms resulting from the ionic strength dependence of the reaction rates* At all three temperatures the resulting values e of U q calculated by the equation * 6 »bW 8 ^ll A W ' 2. and of n calculated by equation (3 5) show enhancement but to a ma t h larger extent then those found in the application of this theory to other ioa-dipole reactions, and to the corresponding factors obtained in this research from the ionic strength dependence. s o A^jr attempt to adjust uQ and nr to acceptable values resulted in unreasonable values of rQ. Seven Angstroms is a reasonable value for for the reaction complex if the fisher-Hirschfelder-T&ylor molecular model of ethyl acetate is representative of the dimensions of this molecule, This molecular model results In a maximum diameter of 9 Angstroms for the ester and if the reaction complex is composed of the OH ion, the ester molecule and perhaps even a molecule or molecules of solvent, It would not be unreasonable to suppose its radius to be seven Angstroms. That the Intermediate complex does involve molecules of solvent agrees with both the theory proposed by Lowry (3 1) and that of Sellvanova and Syrkin (2 6 ). The mechanism proposed by the latter two workers leads to a reason-* able explanation of the large enhanced moments obtained in the dielectric dependence of the rate constants. The complex pro posal by these authors la of the nature of a H*witt«r ion» and involves hydrogen beads* as illustrated by the earntion m o * x V bsooh* ♦ tar *- p 0 4* fiOH {bO) i V j ' H * \ O&sager, in his theory of electric moments of molecules la liquids {2 1) emphasises the statement *the formation of a hydrogen bead Increases the electric moment of the group which, carries the hydrogen, * Amis and Jaff<^(3) also state that though the limit of a** as B approaches infinity exists, the point D = 0° is cer tainly not an adequate point of reference since it accentuates rather then eliminates the effect of the dielectric constant. These factors of hydrogen bonds and extrapolation toward higher dielectrics (to obtain ^extrap. giTfea in Table 7) both accentuate the effect of the dielectric constant upon the external moment of the reaction complex. In cases where reasonable enhanced moments have been found from dielectric effects, extrapolations t^#C= 0 for ionic strength effects and to high dielectrics for dielectric constant effects have resulted in entirely similar k constants being obtained in both instances. Xn this instance the dielec tric extrapolation resulted in entirely different k constants from those obtained from the ionic strength extrapolation. This would indicate that the dieleetric effect is abnormally large as compared ft 5 4 9 ® V* © ft © vt I s I T£ « w*f « $ 4* 1«P <"*s. I 1 1 © I o ft 13 3? .*«1 © s ) It Is evident that a large enhancement of the «rl « © •H 2 3 I * Jf 4 1 I © § «rt15 Jj 4 £ I t n u lion ( £ 7 &BL& la Representative Calculations of Velocity Constants from Experiments 1 Data B=80.3, pur© water, HaGH= .02M, lt0Ae = G.GlM Tiae s u o g 3EU02M Ga Telocity const* (Min) M HOI Ha OH & (11 ter s/aiol. rain 0 60*0 11.50 0.01940 15 50.0 7.35 0.01706 .1.10 SO 50.0 11.20 0.01552 1.09 45 50.0 13.95 0.01442 ,1.09 60 50.0 15.95 0.01362 .1.06 75 50.0 17.55 0.01290 1.09 50.0 24.00 0.01040 Average k.«. • 1 . 0 9 ± 0*01 2? TABLI l b Representative Calculations of Telocity Constants from Bxperimental Bata Temp. = 9.©Q°G 0 = 65, 51.5^3 tOH, JB*0H=0.02M. ttOAesO.OlM Time Ml. 02 M 1 . 0 2 M C Telocity coast. (Mia.) M HCl NaGH n k{liters/mol*mim.} 0 60.0 11.45 0.01942 10 50.0 6.35 0.01746 1.41 20 50.0 9.90 0.01604 1 • 42 SO 50.0 12.40 0.01504 1.40 40 50.0 14* 40 0.01424 1.41 SO 50.0 16. GO 0.01080 Average k.... 1.41 ±0.01 28 TABLB XC Representative Calc til at ions of Telocity Constants from Ixperlmental Data Temp.S 19.1G°G Ds50, 51.1$Rt0H. Ha0H«0.Q2M# It0Ae«0.OlM Time Ml.02 M1.02M Cat Telocity const. (Min.} M. HOI NaOH k(liters/mol min. ) 0 60.0 10.70 0.01972 10 50.0 5.03 0.01799 1.09 SO 80*0 6.20 0.01672 1.06 so 50.0 10.70 0.01572 1.08 45 50.0 13.60 Q.014S6 1.08 55 50.0 16.25 0.01350 1.07 50.0 23.93 0.01043 Average k .... l.QStO.Ol 2$ TABLT£ IX Experimental Values of Veloc ity Constants as ao Function of Ionic Strength; Corresponding Energies of Activation and Frequency Factors D = 65, Ha OS =0.02M, EtOAesO.GlM, Initially. Ionic strength varied by adding B&N0$ when necessary. M k k © 0 ° C 9# 8 0 1 8 * 1 0 4 E B 3 6 . 7 % 31. 5$ s e . s ^ S t O H E t O H I t O H 0 . 0 2 0. 519 1 . 4 1 3 . 0 2 1 4 , 6 2 5 1 1 . 4 2 0 . 0 5 0 . 5 0 5 1 . 3 5 2 . 9 6 1 4 , 6 7 0 1 1 . 4 4 0 . 1 0 0 . 4 9 5 1.29 2 . 9 2 1 4 , 7 5 0 1 1 . 4 9 0 .2 0 0 . 4 6 8 1 . 2 5 2 . 8 0 1 4 , 8 6 0 1 1 . 5 6 0 . 0 5 —r m* 1 . 1 8 2. 55 30 TABLS XII Experimental Values of Velocity Constants as a function of Dielectric Constant} Corresponding Energies of Aeti- uatlou and Ifrequency Factors'’ >as0*02f SaOHffOiOS M, St0&osOi,Ol M, Initially 9.8°C 19.1°C B o°c AS B * fc * k * k I t O H St OH StOH 88.3 O O i O 1*09 ——------— -* 11*740 9*43 6 4 . 3 00*0 2*54 w + m n» mm -<— 80.7 ------00.0 4*34 80i0 13.3 1*16 7*50 2*41 lil 4. 51 11,300 9*09 85.0 36.7 .519 51*5 1*41 26. 5 3.02 14,623 11*42 5 0 . 0 60*2 .161 55*6 .426 51.1 1.09 15j880 11*92 ---- 3 S . 8 85*0 .052 —— ---- —• 31 m i L i w Sottstaats In Fitting JW^Ngj » * « t > U * 1 Visa 7. 66 0*631 4 * 5 1 7a 72 1*46 i x * o 7a 46 5« OS fe"i«?• - 52 a Wt4'#I 0O*3T AS6 998 0T*6I 16*9 060T 686 08*6 *98*1 09*® 809 69® 00*0 ST_0t X v; {*a*x*x»)* ga f^OT T 0 0 ‘Aa°JE o § ©,mST£ n% OAJino Snx34I.£ *ioj 8 9 .8 *9 . 8 1 1 0 0 A n s ? £ 2ABUI 71 Enhancement of Moment of Etiiyl Acetate as Given by Coulombla Energy Considerations. Dielectric Ag{elect.) elect*} Moment to make range (Calculated) (Observed) A3£(Calo« ) agree with AE(Obs*) $5-60 51*5 3,325 116.8 $0*05 82*0 1,255 27*5 3.4 35 Logarithm k versus l/T at various ionic strengths. D = 65 LEGEND 0 M = .0 2 -O- M = .05 - O M = .1 0 O- yM = .20 1 A« = .50 c-t I—t O o -.1 3.4 3.7 l/T x 103 Figure 1 LOGARITHM 3.45 Logarithm k versusl/T at variousdielectric constants. M = .02 3.60 3.65 LOGARITHM 3.35 3.40 Logarithm k versus l/T for runs in pure water. pure in runs for l/T versus k Logarithm 3.45 l/T x 10; x l/T Figure 3 Figure 3.50 -O- A SELIVANOVA AND SYBKIN AND SELIVANOVA A 3.55 j REICHES (j) O LEGEND POTTS SALMI AMD KORTE AMD SALMI 3.60 3.65 Agreement ofdata with theoretical curve, 1? = (f)Z 58 .5 LEGEND “Q “ 9*8 C 19.1 C .1 9.8 C LEGEND "O* Agreement of dielectric constant data with theoretical curve, W f =W (f)l/D2 f curve, with theoretical data constant dielectric of Agreement 01 X m m m the reaction of ethyl acetate and the 0H~ ion was studied fciaetieally at 0°C, 9.8°C, and 19*1°C, both in pure water and In constant dielectric constant water-ethyl alcohol mixtures, in order to obtain information with which to test the validity of the Arais-Jaff/ theory pertaining to the prediction of reaction rates of ion-dipolar molecule type reactions, The reaction velocity constant was measured at dielectric constants of 80, 65* and 50* at a constant ionic strength value of 0*02, and at a dielectric constant of 65 In which the ionic strength was 0.02* 0.05* 0.10, 0 *2 0 , O .30 at all three temperatures* She velocity constants obtained in pure water are compared with similar data by other authors* fhe dependence of the velocity constants upon ionic strength was accounted for satisfactorily on the theory of the rates of ion-dipolar molecule reactions proposed by Amis and Jaffe. The comparison of the dielectric constant dependence of the rate constants with the theory of Amis and JaSfi indicates obe dience with respect to trend but results in very large values of the enhanced moment, w£, and internal refractive Index, n. The magnitudes of u£ and n2 are discussed in the light of hydrogen bonding in the reaction complex, the possibility that the reaction complex may be & bipolar compound, similar to a "awibier ho iefii* extrapolation of log k versus 1/B2 to higher dielectric eon* stanta, and the ceuloobi© energy of activation theory* It to suggested that a combination of the above factors would resnlt In abnormally large enhanced moments for the react ion complex. BX32»I3$3£PKr 1* Akerlof, 0. Dielectric constants of some organic solvent-water mixtures at various temperatures . J. Am. Chssu So©., 5^; **135 (1932) 2. M s , S* S. Goolomhlc energy of activation. J. Aa. Choa. Soc., 6 3 * 1606 (1943.) 3« M s , 2. S., and Jaffdf 0. fhe derivation of a general kinetic equation for reaction 'between lone and dipolar molecules. J. Chen. Fhy*. 10: Bo. 9, 598-604 (19^2) 4. Amis, 3* S., Jaffe, G., end Overman, H» T. The diaeetone aleohol-hy&roxide ion reaction from the standpoint of len-dlpele theory. J. Am. Chea. See., 66, 1823* <19***0 5* Amis, 2. S., and Cook, S. 3. Dielectric end solvent effects upon the kinetics of the fading of hrom phenol blue in mixed solvents. J. Am. Chsm. Soc. 6 3 $ 2621 (1941) 6. Amis, 2. S., and Holmes, 2. G. Dielectric and solvent effects upon the rate of sucrose Inversion by hydrochloric add. J. Am. Ghem. So©. 63* 2231 (1941) 43 ?. Amis, S. S., aad halier, d. S. fhe entropies end energies of activation of ionic reactions: the kinetics of the alkaline fading of hroa phenol "blue in iso&ieleetrlc media, d. Am. Chem. Sec. 61: 905-12 (1939) S. Amis, S. S., and Potts, J. E. Dielectric and solvent effects upon the iodide-per- salfate reaction. J. as. Cham. See. 6 3, 2B83 (1941) 9* Amis, S. S., and Price, J. B. Effect of dielectrics and solvent open the regeneration in acid eolation of alkali-faded hrom phenol blue. J. Phys. Cham. 4?: 3 38 (19^3) 10. Arrhenius, S. Uber die Dissociation der in Passer Gelosten St of fa. Z. Phyaik Chem. 1, 631 (188?) 11. B o m , M. Volumes and heats of hydration of ions, s. Phyaik, 1: 45-8 (1 9 2 0) 12. Breasted, J. B. Zur theorie der chemischen reaktionggeschwindigkeit. Z. Pfcyeik. Chem. 102: 169-207 (1922) 115037-64 (1925) 13* Christiansen, J. A. Ufcer die himolekularern reaktioaon In loeimgen. %* Physik. Chem. 113: 35-52 (1925) 44 14. Davies, ft.., sad Evans. B. P. Influence of alkyl groups upon reaction velocity in solution. IP. the alkaline and acid hydrolysis of the ethyl eaters of lever saturated aliphatic acids In aqueous acetone, y. Ched u Soo. 339-345* (1940) 1 5 * Bekye, P.# and Suekel, P. tke theory of electrolytes. Phyaik. 2. 24: 105-206 (1923) 16. Seans, S. P. , Gordon, J., and Watson, C. Alkaline hydrolysis of saturated aliphatic esters, y. Ghea. See. 1434-9 (1930) 17* 3iana, S. P., Gordon, J., and Watson, 0. Alkaline Hydrolysis of henzsle esters, y. She*. Soo. 1430, (193?) 10. filasstone, S., L&idler, K. J., and gyring, H. •the theory of 'Bate Processes,» McGraw-Hill, 1941, pp. 401, 442. 19. Goldberg, aad Waage Shades sur les affinites chimiques. J. Prakt. Chem. (2) 19, 69 <1878) 2G. Hammett, Louis P. “Physical Organic Chemistry” McGrow Hill, S. f*. 1940, pp. 355* 21* Oasager, L. Blscirls Moment 3 of Molecules in Li quids, y. Am. Chem. Soo. 58, I486, (1936) m 22. Balatasr** & and Snake* A. L. tke aeehaaism of hydrolysis. Th© alkaline sapoitfieation of aayl acetate. Trans. Faraday See. 30* 508 (193*0 23. Belcher, 1*. Th. die geschwindigkeit dor rersoifung. Annalen, Bead 226: 257 (1885) 24. Salgi, B. J. t and Berts, Baymond The alkaline hydrolysis ot eaters In water and water dlexaae seltttlens. l u . Acad. Set. Fennieae A 54* Bo. 1 2* 13 pp. (1940) 25* Seatofca*ft, t. Statistical mechanics and reaction rates In liquid eolations. Cham. Bar. 10s 229-40 (1932) 2 6. SellTaaoYa, A* S.* and Syrfcin, J. X. Effect of solvent on kinetics. Contes Bendas (DQXLADT) d© L*Academic des Sciences do L»CRSS. Vol. 2 3. 8 6. 1, p. *5' (1939) 27. Sidgvlck, 31. Y. Appendix. A Table of Dipole Moments Trans. Faraday Soo. 30* Appendix (193*0 28. Salth, H. A., and Levensom, H. S. Kinetics of the saponification of ethyl esters of a-aliphatie acids. J. Jto. O M h Soe., 6 1, 1172-5 (1939) 46 2 9 % Kamila, Sere fhe action of solvents la eater hydrolysis. Suosea Kamlstilehti 1 5 B, 9 - 1 0 (1 9 4 2 ) 3 0 . 7o»»ila, Sero the factors responsible for chemical reaction velo cities. Suenen Keaistilehti 17 A, 1 - 1 5 (1 9 4 4 ) 3 1 . tfetters, W. A. •Physical Aspects of Organic Chemistry. ** G. Bontledge and Sons. London, 1 9 3 5 * P* 268 3 2 . Weisaberger. A., and Proskauer, B. •Organic Solvents11 Oxford University Press (1935) James Edward Potts, Jr. was born in Alexandria, Louis iana, on October 28, 1918* He received his elementary and high school education in the public schools of that city. In September, 1935 He entered Louisiana College, Pineville, Louisiana, from which he m s graduated in June, 1939* with the Bachelor of Science Degree* Be entered graduate training at Louisiana State Univer sity in September, 1939* where he held a teaching fellowship until his graduation with the Has ter of Science Degree in chemistry in August, 19A1. Be re-entered Louisiana State University the following September, resigning at the end of the fall semester to accept a position as a research chemist with the Tennessee Talley Authority Department of Chemical Engineering* After four years employment with the Tennessee Talley Authority, he returned to Louisiana State University at the beginning of the spring sem ester, 1 9 4 6 . He is now a candidate for the Degree of Doctor of Philosophy* EXAMINATION AND THESIS REPORT Candidate: James Edward Potts, Jr. Major Field: Chemistry Title of Thesis: Dielectric and Solvent Effects on the Alkaline Hydrolysis of Ethyl Acetate Approved: Major Professor Chairni Dean of the Graduate School EXAMINING COMMITTEE: " 7 •t . ■ • c <- Date of Examination: October 22, 1 9 4 7