, n £~ 21 1965

WHAL-E 59-0368

PRESSURE EFFECT AND MECHANISM IN ACID CATALYSIS

PART 1

BY

E. WHALLEY

Offprinted from the Transactions 0/ the Faraday Society No. 437, VoL 55, Part 5, May, 1959

PRESSURE EFFECf AND MECHANISM IN ACID CATALYSIS PART ·l

PRESSURE EFFECf AND MECHANISM IN ACID CATALYSIS* PART 1

By E. WHALLEY

Division of Applied Chemistry, National Research Council of Canada, Ottawa

Received 10th October, 1958

The effect of pressure on the rates of acid-catalyzed reactions is discussed. Uni­molecular (A-I) decompositions will have a volume of activation close to zero and probably positive, whereas bimolecular (A-2) substitutions and additions and A-I intramolecular rearrangements will have a volume of activation that is negative by at least several cm3 mole-I. Thus, it is possible to distinguish between the A-I and A-2 mechanisms, at least for some reactions. Relevant data in the literature are discussed. The measurement of the effect of pressure on the acid-catalyzed hydrolysis of esters and their derivatives can give some information about the location of the proton in the re­active conjugate acid of the ester, etc. The entropies of activation of acid-catalyzed reactions are also discussed; measured entropies of activation are not reliable for de­tecting A-I mechanisms, but they may be much more reliable for detecting A-2 mechanisms.

1. lNrRODUcnON

Solvolysis reactions that are catalyzed specifically by oxonium ions, i.e. re­actions in which a pre-equilibrium proton-transfer to the substrate, so forming its conjugate acid, occurs, are usually classified into two groups according to whether the conjugate acid decomposes unimolecularly (A-I mechanism in the nomen­clature of Ingold 1) or bimolecularly, a solvent molecule being covalently bound

. in the transition state (A-2 mechanism). Several methods have been used to distinguish between the two mechanisms, and these are summarized in table 1. The use of optically-active substrates is of limited applicability and, while it may reliably detect an A-I mechanism, it may not reliably detect an A-2 mechanism. The steric repulsion of substituents is often valuable for determining mechanism but as we shall see in part 3 of this series, which follows, it is sometimes over­whelmed by the effect of electron release or withdrawal and it then becomes useless for the purpose. In addition, study of substituent effects leaves uncertain the mechanism for the unsubstituted compound; and frequently the effect of sub­stituents cannot be studied, for example in the hydrolysis of sucrose. The use of Hammett's acidity function ho is essentially empirical and it relies for its validity upon a few reactions whose mechanisms have been determined by other methods. Until recently no reaction was known for which it failed, but this is mainly because there are few methods of determining the molecularity of reference reactions. We will show in later papers of this series that a number of reactions usually thought to be unimolecular because, among other reasons, log (rate) ex: Ho are actually bimolecular. The method based on the entropy of activation is semi-empirical in the sense that entropies of activation for both A-I and A-2 mechanisms can frequently be neither predicted nor, when known, understood even semi-quanti­tatively. In addition there are insufficient experimental values for reliable emr·irical correlations. A discussion of the information derivable from entropies of activ­ation is given in § 3, and it will be shown that entropies of activatioll have

* N.R.C. DO. 5115. 798

E. WHALLEY 799

frequently been misinterpreted. The deuterium isotope effect as a method of dis­tinguishing between A-I and A-2 mechanisms is, at present, purely empirical. When enough information about the vibration frequencies of the conjugate acids has been obtained it may be more useful and more reliable, but it cannot be given any weight at present.

It is evident from this brief discussion that soundly-based methods of dis­tinguishing between the A-I and A-2 mechanisms are badly needed. The pur­pose of this series of papers is to apply a new method, in which the volume of activation is used to distinguish between them, which appears to be generally applicable and to be reliable. It has been used recently by Whalley 8 to show that the acid-catalyzed cyclization of citraldehyde to 3: 8-carvomenthenediol, whose rate under pressure was measured by Harris and Weale,9 is bimolecular. In this, the first paper of the series, the theory of the effect of pressure on acid-catalzyed reactions and the use of measurements of the pressure effect in determining mechanism are discussed. In establishing the validity of the method use is made only of measurements that are independent of a reaction mechanism. In § 4 measurements on acid-catalyzed reactions that are already in the literature are discussed in the light of the theory.

TABLE I.-PROPOSED METHODS OF DISTINGUISHING BETWEEN A-I AND A-2 MECHANISMS

result expected for method

A-I A-2

substrate optically-active inversion plus some inversion of at centre of substitution racemization of configuration

configuration * - bulky substituents near rate little affected rate decreased

reaction site

electron-releasing sub- rate increased rate increased stituents near reaction site

electron-withdrawing rate decreased rate decreased substituents near

reaction site

concentrated acid rate a: ho rate a: CH+t

entropy of activation t small and positive large and negative

deuterium (D20) rate increased by rate increased by isotope effect factor of 2-3 factor of 1'4-1 '8

* In special circumstances retention of configuration may occur. t Cf{+ is the concentration of hydrogen ion.

references

I, p. 372

I, p. 335

I, p. 335

1, p. 335

2,3

4,5,6

7

t The entropy of activation is defined as the sum of the standard entropy change of the proton transfer and the entropy of activation of the slow step.

"2. THE EFFECT OF PRESSURE ON THE RATE OF ACID-CATALYZED REACTIONS

We consider a reaction that is known, from evidence other than the pressure effect, to be cataly.::ed specifically by the hyd.(ogen ion. Let A be the acid used as catalyst, S the proton-accepting solvent, and R the substrate that solvolyzes to give the products. The mechanism in a single solvent can be written

A + S .= B- + SH+, R + SH+ .= RH+ + S,

RH+ + nS -+ products.

(2.1) (2.2) (2.3)

800 MECHANISM IN ACID CATALYSIS

Reactions (2.1) and (2.2) are at equilibrium and have equilibrium constants Kl and K2 respectively, and reaction (2.3) is slow and has rate constant k3.

Let a be the initial concentration (i.e. before ionization) of A, b be that of B- (b = 0 if no B- is added initially), and CR, CRH+, cs, CS!{+, be the concentrations of R, RH+, Sand SH+ respectively at the temperature and pressure of the reaction. We note that

din k3 t..3V~ <lP = - RT' (2.4)

where Ll1 VO and 6,2Vo are the standard partial-volume changes of reactions 1 and 2, t..3 V~ is the standard volume of activation of reaction 3, R is the gas constant, and T the temperature.

The experimental first-order rate constant k is defined by the equation

rate = - d(CR + CRl{+)/dt = k(CR + CRli+)' (2.5)

In terms of concentrations at room temperature and pressure, denoted by super-script 0, k is given by .

- d(cR + cRH+)/dt = k cR(l + CRW/cR.}. (2.6)

By considering reaction (2.3) and t1sing Bronsted's rule, that the rate is pro­portional to the concentration c~ of the transition state, we find that

rate = k~c~,

where k~ is the rate constant. The equilibrium constant K~ for the formation of the transition state is

(2.8)

where a denotes activities and I~ is the activity coefficient of the transition state. The equilibrium constant K2 of reaction (2.2) is

By inserting eqn. (2.9) and (2.8) into (2.7), and writing

k3 = K~k+

we find

rate = K2k3aRaSH+ asn- 1/1+.

By comparing eqn. (2.10) with (2.5) we find

k = K2k3csWasn-l I RI SH+

1 + CRoH+/ CR 1* According to the Debye-Huckellirniting law for 1 : 1 electrolytes,

IR=l, ISH+ = 1+,

and so eqn. (2.11) becomes

k = K2k3CSH+ asn-1/(l + CRW/CR).

If a. is the degree of dissociation of A then

and so

CSH+ = aa./(l + CRH+/CSH+)'

(2.9)

(2.10)

(2.11)

(2.12)

(2.13)

(2.14)

(2.15)

E. WHALLEY SOl

Inserting eqn. (2.15) into (2.13) and expanding the bottom line, and introducing into it eqn. (2.9) and (2.12) we find

k = K2k3arxasn-l 1 + K2(CR + aa.)/as

(2.16)

Eqn. (2.13) is useful for discussing strong acids as catalysts and eqn. (2.16) for discussing weak acids as catalysts. .

In order to calculate d 10 k/dp from eqn. (2.16) we need to know dIn aa./dp. For strong acids ex = 1 and d 10 ex/dp = O. For weak acids

Kl = ex(aex + b) fa- ISH+ (1 - ex)(1 + CRH+/CSH+) fA

(2.17)

where IB-, etc., are the activity coefficients of B-, etc. According to the Debye­HUckellimiting law for 1 : 1 electrolytes

fA = 1, (2.1S)

where A oc(ET)- 3/2, E being the dielectric constant, but is otherwise a constant, and fL is the ionic strength. Inserting eqn. (2.1S) into (2.17), taking logarithms, differ­entiating, rearranging, introducing the pressure coefficient ~ of the dielectric constant and the compressibility X,

~ OE = ~ E op •

1 op pop = x.

where p is the density, and noting that

d 10 a/dp = d 10 b/dp = d In p/dp = X. (2.19)

we find that

d In a. = {d In Kl (1 + K2 CR/CS)/k'!. A !(3; _ ) l {_1_ + 1 }-l . (2.20) dp dp fL X J 1 - ex 1 + b/arx

Two extreme conditions of eqn. (2.20) that are important are as follows.

(i) The acid is unbuffered, i.e. b = o. Then

d In IX = 1 - IX {d In K1(1 + K2cR)/p _ AIl)(3Y _ X)}. dp 2 - (J; dp r .. (2.21)

(ii) The acid is buffered, i.e. b ~ a(/.. Then

d In IX = (1 _ ex){d In K j (1 + K2cR)/p _ AfL1(3~ _ X)} . (2.22) dp dp

In order to determine the pressure coefficient of In k we take logarithms of eqn. (2.16), differentiate. and substitute the appropriate values of d In ex/dp. For K2CR ~ 1 and K2aa ~ 1 we obtain the following.

(i) The acid is strong. Then

din k/dp = - /). V+/RT + X. (2.23)

where we have written

(2.24)

t1 V* will be called the volume of activation of the reaction. It is frequently

802 MECHANISM IN ACID CATALYSIS

convenient, when the acid is strong, to use the second-order rate constant ka defined thus,

ka = k/CSH+, (2.25)

rather than k. We have immediately

d In ka/dp = - Ll V +/RT. (2.26)

(ii) The acid is weak and unbuffered. Then

dInk =_ LlV* + X _ 1 - (f.{LlVIO + X + Ap.~(3~ - X)}, dp RT 2- rx. RT

(2.27)

(iii) The acid is weak and buffered. Then

d In Ie = _ Ll V* + X __ 1 _ {Ll VI ° + X + Ap.!(3~ - X)} . dp RT 1- rx. RT

(2.28)

Two quantities of interest that can be determined from eqn. (2.23)-(2.28) are Ll V* and Lli Va. Ll V* is obtained from measurements with strong acids, or with weak acids if Lli V O is known; Lli V O is obtained from measurements both with strong acids and with weak acids, extrapolated to zero ionic strength.

3. VOLUMES AND ENTROPIES OF ACTIVATION IN ACID CATALYSIS

In this section we discuss the theory underlying the use of volumes of activation for distinguishing between the A-I and A-2 mechanisms. The entropy of activ­ation has been used 4. 5. 6 in attempts to distinguish between the two mechanisms, and we discuss this method also.

l[ the slow step (2.3) is a unimolecular decomposition then the transition state is more loosely bound than RH+. and, because the charge will usually be more dispersed in the transition state, the solvent will also be more loosely bound. Consequently, the volume of the transition state will be greater than that of the reactants, and t.3 V + will be positive. If step (2.3) is bimolecular or is a unirnolec­ular rearrangement then in the transition state a partial valence bond has replaced a van der Waals bond. The decrease in volume accompanying this is likely to overwhelm otber contributions and will result in a negative Ll3 V*.

From a simple model we can learn something of tbe magnitude of the volume changes expected. It is frequently assumed that during a simple unirnolecular decomposition the length of the breaking bond in the transition state is about 10 % greater than that in the reactant. If a bond of length 1· 5 A and cross-sectional area of 10 A2 lengthens by 10 % then the increase in volume is about I cm3 mole-I. The dispersal of tbe charge will probably cause a similarly small increase in volume because the charge is usually distributed between two adjacent atoms and the solvent will be aware of this mainly as a slight increase in the volume of tbe ion.

If the slow step is an A-2 addition or substitution or an A-I isomerization the contriuutions to the volume of activation can be roughly divided into (i) a decrease in volume caused by tbe conversion of a van der Waals bond to a partial valence bond, (ii) an increase in volume due to the lengtbening of the bond (if any) that breaks, and (iii) an increase in volume due to the dispersal of the charge. Extending our rough calculations of the preceding paragraph, if a van der Waals bond of length 3·6 A and cross-sectional area 10 A2 decreases in length to 10 % greater than 1'5 A then the accompanying decrease in volume is about 12 cm3 mole-I. This decrease, though only a very rough approximation to the decreases expected in real reactions, is overwhelmingly greater than the increase due to contributions ~ and'~ Consequently, Ll3 V* will be negative by several cm3 mole-t .

1I i) (ii iI.

E. WHALLEY 803

Arguments exactly similar to those in the second paragraph, mutatis nIutandis, show that the entropy of activation ~3S+ of step (2.3) will be positive for an A-I decomposition and negative for an A-2 reaction or an A-I isomerization.

Now we have seen that

~V+ = ~2Vo + ~3V+, and similarly, if the entropy of activation ~S+ for the complete reaction is defined by

ka = (kT/Iz) exp (6.S+/R) exp (- 6.H+/RD,

it is given by

6.S+ = 6.zso + 6.3S"",

where 6.2so is the st(1.uciard entropy change of reaction (2.2). To determine 6.3 V+ and 6.3S*, and hence to learn about the mechanism, requires knowledge of 6. V+ and 6.S*, which are experimentally measured, and of 6.2vo and 6.2so. We now examine the available information about 6.zvo and 6.2so, i.e. about the volume and entropy changes in isoelectronic proton transfers.

The only substrates R for which Ll1 VO and 6.1so are known are amines. The data are collected in table 2. The volume changes were calculated assuming that the volume of ionization17 of water is - 23·4 cm3 mole-i.

VOLUME CHANGES

The volume changes listed in table 2 are all negative, fairly small but certainly not zero, and of ~bout the same magnitude. The volume changes for the reaction

R3N (1 M aq.) + NH;t(l M aq.) -+ R3NH+ (1 M aq.) + NH3 (aq.),

are given in table 3; they are almost zero and slightly positive. Evidently, the presence of groups attached to the nitrogen atom makes little difference to the electrostatic contribution to the volume of the ion. The decreases in volume in table 2 are, perhaps, due more to local van der Waals interactions than to purely electrostatic forces. For, consider the relative volumes of NHt and H30+. Both ions will orient the neighbouring water molecules so that the lone pairs of the water molecules are near the ions. If a lone pair is close to a hydrogen

TABLE 2.-VOLUME AND ENTROPY CHANGES rN THE PROTON TRANSFER REAcnONS

RJN (1 M aq.) + H10+ (1 M aq.) --+ RJNH+ (l M aq.) + H 20 (I)

temp. 6,V· temp. 6,S· R3N ref. ref.

·c em'moie- 1 ·C cal deg. -I mole-1

NH3 25 - 6·6 10 25 + 0·6 13 -5·5 11

CHJNH2 25 -306 10 25 + 4·7 14

(CH3hNH 25 -3-9 10,12 25 + 9·5 14

(CH3hN 25 -4·7 10,12 25 + 15·2 14

C6HSNH2 25 - 5·2 10,12

CsHsN 25 -4·7 10,12

CSHllN 25 -1·0 10,12

CzHsNH2 30 + 3·1 15

n-C3H7NH2 30 + 1·7 15

n-C4H9NH2 30 + 1·3 15

NH2(CH2hNH2 25 + 5·8 16

NH2(CHV6NH2 2S + 3·3 16

804 MECHANISM IN ACID CATALYSIS

TABLE 3.-VOLUME AND ENTROPY CHANGES IN THE PROTON TRANSFER REACTIONS

R3N (1 M aq.) + Nm (I M aq.) -7 R3NH+ (I M aq.) + NH3 (I M aq.) From data in table I, tl 2V o for

NH3 + H30+ -7 ~ + H20

is taken as - 6·0 cm3 mole- I

CH3NH2

(CH3)2NH (CH3hN C6H SNH2

CsHsN CSHIIN C2H SNH2 n-C3H7NH2 n-C4H9NH2 NH2(CH2hNH2 Nl-h(CH2)6NH2

em) mole-1

+2-4 +2'1 + 1'3 + 0·8 + 1'3 + 5'0

AS·

cal deg. -I moJe- 1

+ 4·1

+ 8·9 + 14·6

+ 2·5 + 1'1 + 0'7 + 5'2 + 2·7

atom in the ion an interaction due to hydrogen-bonding, which will add to the electrostatic interaction, will occur, a'ld the water molecule will be pulled even closer to the ion. When a lone pair of the water molecule is near a lone pair on the ion, as can happen if the ion is H30+, hydrogen-bonding cannot occur and there will probably be a local repulsion due to interaction between lone pairs. We conclude, therefore, tbat local hydrogen-bonding will tend to cause NHl", which has no lone pair, to have a smaller partial molar volume than H30+, which has one lone pair; this agrees with observation. This effect will not occur in the reaction

R3N + NH;t ~ R3NH+ + NH3•

and, as we see in table 3, the volume changes are small. They are actually positive; perhaps there is a small effect due to shielding of the solvent by groups attached to nitrogen.

It seems very likely therefore on the evidence available that if R of reaction (2.2) is an oxygen-containing compound then ~2 V o will be small and, if S is iess substituted than R, perhaps slightly positive. Consequently, we deduce the following. If Rand S contain the same proton-accepting atom and R is not less substituted than S then

(i) for an A-I reaction, .:1 V* will be positive, or, in the apparently unlikely event of .:12 V O being slightly negative it may be zero or slightly negative;

(ii) for an A-2 reaction, .:1 V* will be negative by several cm3 mole-I,

If Rand S do not contain the same proton-accepting atom, .:1 V* may be somewhat more positive. Measurement of the volume of activation should, therefore, be a good method for distinguishing between the A-I and the A-2 mechanisms.

ENTROPY CHANGES

The entropy changes listed in table 2 are all positive, and certainly not close to zero. They are not uniform, but extend over the range + 0·6-15·2 cal deg.-1

mole-I. Furthermore, the entropy changes for the reaction

R3N (l M aq.) + NHt (l M aq.) -'>- R3NH+ (1 M aq.) + NH3(1 M aq.),

E. WHALLEY 805

which are given in table 3, cover a similar range. It is clear that entropy changes in isoelectronic proton transfers are not even approximately zero. The large increase in I:1so with increase in the number of substituents is probably due in part to the substituents displacing the firmly-held water from the first hydration shell.15 The same effects are expected if Rand S are oxygen-containing.

Therefore we draw the following conclusions about the entropy of activation I:1S'* for acid-catalyzed reactions. It is assumed that R is not less-substituted tj-.?~ S; if it is then modifications are called for.

(i) For an A-I reaction, 1:1 2So and 1:13S'* will probably always be positive and consequently I:1S* will be positive.

(ii) For an A-2 reaction, 1:12SO will be positive and 1:13S* will be negative. There seems to be no reason to assume that necessarily 11:12SO 1 < 11:13SF I. Consequently, I:1S* can have either sign depending upon whether 1:12So or 1:13S+ is dominant.

We deduce from this that

(i) if I:1S" is less than zero then the reaction is probably A-2, the probability increasing as I:1S* decreases, and

(ii) if I:1S+ is greater than zero the mechanism can be either A-lor A-2 and the probability of its being A-I increases as I:1S* increases. However, because of the high values of 1:12So that can occur, even a I:1S* of + 10 cal deg.-I mole-1 cannot be considered to prove, or even to suggest strongly, an A-I mechanism. We must infer, therefore, that conclusions that hav« been drawn from the entropy of activation, to the effect that several reactions have an A-I mechanism,4. 5.6 are not valid. They may, nevertheless, be correct; other evidence is needed.

4. DISCUSSION OF DATA 1N THE LITERATURE

The effect of pressure on acid-catalyzed reactions has been measured a number of times and from some of these measurements information about mechanism can be obtained. .

1. CYCLIZATION OF CITRALDEHYDE.-The cyclization of citraldehyde has been discussed previously 8 and the A-I mechanism eliminated. The volume of activ­ation,9 - 25·2 cm3 mole-1 at 25'2°C and I atm, is consistent with an A-2 mech;\ . ism; a mechanism in which water is rapidly added to the conjugate acid of citralcl. hyde, and the molecule thus formed slowly cyclizes, is also consistent with th·' volume of activation. The entropy of activation calculated from the results of Harris and Weale,9 - 37'1 cal deg.-1 mole-I, agrees with either mechanism.

2. INVERSION OF SUCROSE.-The acid-catalyzed inversion of sucrose is the first reaction whose pressure coefficient was studied.l8-21 There is undoubtedly a pre-equilibrium proton transfer 22 though a recent paper 23 argues, not very con­vincingly, that the proton transfer to sucrose is the slow step. The volume of activation, obtained graphically from the results of Cohen and de Boer,21 is + ,6'0 ±--- 0·3 cm3 mole-l at 25°C; hence the mechanism is A-I. That the hydrolysis is A-I was deduced by Hammett 2 from the proportionality of log (rate) to Ho, and Long and Paul 3 have pointed out that this was the only evidence of an A-I mechanism. Because we show in later papers that some reactions that follow ho rather than CH+ are actually A-2, the volume of activation is important evidence for the A-I inversion of sucrose.

3. HYDROLYSIS OF ESTERS.-The acid-catalyzed hydrolysis of methyl and ethyl acetates by dilute hydrochloric acid was also an early reaction to be studied under pressure.19. 24 The rate is increased by pressure, and the results of Bogojawlensky

806 MECHANISM IN ACID CATALYSIS

and Tammann 24 on methyl acetate catalyzed by hydrochloric acid at I and 500 atm yield

Do v'" = - 8'3 cm3 mole-l at 30·S°C.

Measurements 24 with acetic acid as catalyst confirm this. From the first-order rate constants at I and SOO atm we find, using eqn. (2.27),

Do V+ + tD.I Vo - h = - 15·3 em3 mole-I,

and so, using the above value of Do v'" we find for the volume of ionization of acetic acid

Dol Vo =-12·8 cm3 mole-t at 30'SoC,

in good agreement with recent values at 2S°C.Z5a It is well known that the initial proton transfer is at equilibrium. It is not possible to tell from the volume of activation alone whether the mechanism is A-lor A-2 because there are two oxygen atoms to which the proton can be attached, the carbonyl oxygen and the ether oxygen. If the mechanism is A-I and the proton is on the carbonyl oxygen in the reactive conjugate acid then the transition state will be more polar than the conjugate acid, as shown in 1, table 4, and a negative volume of activation is expected.2Sb However, there is abundant evidence I that the mechanism is A-2. If we accept this, then it is possible to learn something about the location of the proton of the reactive form of the conjugate acid. We first note that there are two possible mechanisms, according as an intermediate molecule is produced or not. Experiments on 0 18 exchange between hydrolyzing ethyl benzoate and the solvent 26 appear to show that a molecule is formed from water and the conjugate acid, but we will leave the question open whether this is an intermediate in the hydrolysis. However, we note that the rate of exchange is several times smaller than the rate of hydrolysis. Consequently, if this is true for methyl acetate and if an intermediate molecule is formed in the hydrolysis then the formation of the intermediate is problbly the slowest step of the hydrolysis and the measured volume of activation is determined mainly by this step. In table 4 we summarize the possible transition states for the hydrolysis.of methyl acetate by the A-I and A-2 mechanisms. If an intermediate molecule is formed the 0 18 exchange implies, as is expected, ready mobility of the protons in the intermediate.

TABLE 4.-HYDROLYSlS OF METHYL ACETATE

transition state reactant bimolecular with bimolecular with no unimolccular intermediate molecule intermediate molecule

+ °OH 8+ 8+

OH OH2 OHz / .~ / : (1- 8)+ +

C-H3C CH3-C. CH3--C:=OH CH3-C=OH

'" ".,8- I 8-eCH3 OCH3 OCH3 OCH3

2 3

0 8+ (I - 8)+ 0 / OH2 OHz

/ (1-8)+ 8-CH3-C CH3-C CH3-C-····O CH3- C=O

'" ·· ... 8+ I 8+: 0+CH3 OCH3 0+CH3 OCH3 H H H H

4 5 6

I I

E. WHALLEY 807

The volume of activation is - 8·3 em3 mole-I. This is approximately the value expected for a simple bimolecular reaction of H20 with a substrate in which no major changes of electrostriction occur. Consequently, it is unlikely that there is a major change in electrostriction in the slow step and so transition state 3 and intermediate molecule 5 can probably be excluded. We conclude therefore that if a transition state only is produced then the proton is probably on the ether oxygen in the active form and if an intermediate molecule is produced then the proton is probably on the carbonyl oxygen in the active form. It is quite possible of course that both reactions occur together. Because of the very limited experi­mental data this conclusion must be considered to be only tentative at present.

4. REARRANGEMENT OF CX:-PHENYLALLYL ALCOHoL.-cx:-Phenylallyl alcohol readily rearranges in the presence of acid to cinnamyl alcohol.27 The mechanism generally accepted is the prior formation of an oxonium ion, Ph-CH-CH=CH2.

1+ OH2

followed by re-arrangement, either by bimolecular substitution

CH / , +

Ph-CH CH2 + H20 ~. Ph-CH=CH-CH20H2,

"'-+ OH2 + H20

by unimolecular rearrangement

Ph-C/ C~, CH21Ph--QI/ CH" CH'): PhCH=CH-CH20H2,

1+ .. .' OH2 'OH~

or by unimolecular decomposition

CH / , slow

Ph-CH CH2 ~ [Ph-CH-CH-CH2]+ + H20 ~

OH2 fast + ~ PhCH =CH-CH20H2,

and followed finally by the rapid loss of a proton from the conjugate acid of cinnamyl alcohol. It is not known which of the slow steps occurs. Harris and Weale 9 have found from measurement of the pressure effect that

t:. V+ = - 5·0 em3 mole-l at 30°C in 45·7 % w/w aqueous acetone.

Because there is a volume decrease the unimolecular decomposition is probably not the most important mechanism. It is not possible to choose on the basis of the volume of activation between the bimolecular substitution and the uni­molecular rearrangement; both would have a negative volume of activation.

1 Ingold, Structure and Mechanism in Organic Chemistry (Cornell University Press Ithaca, 1953).

2 Hammett, Physical Organic Chemistry (McGraw-Hill Book Co., New York, 1940). 3 Long and Paul, Chern. Rev., 1957, 57, 935. 4 Taft, Purlee, Riesz and de Fazio, J. Amer. Chern. Soc., 1955, 77, 1584. 5 Long, Pritchard and Stafford, J. Amer. Chern. Soc., 1957,79,2362. 6 Farley, et al., J. Amer. Chern. Soc., 1958, 80, 3458.

MECHANISM IN ACID CATALYSIS

7 Pritchard and Long, J. Amer. Chem. Soc., 1956, 78, 6008. 8 Whalley, Can. J. Chem., 1958, 36, 228. 9 Harris and Weale, J. Chem. Soc., 1956, 953.

10 Hamann and Lim, Austral. J. Chem., 1954, 7, 329. 11 Buchanan and Hamann, Trans. Faraday Soc., 1953, 49, 1425. 12 Hamann and Strauss, Trans. Faraday Soc. , 1956, 51, 1684. 13 Everett and Wynne-Jones, Proc. Roy. Soc. A, 1941, 177, 499. 14 Everett and Wynne-Jones, Proc. Roy. Soc. A, 1938,169, 190. 15 Evans and Hamann, Trans. Faraday Soc., 1951, 47, 34. 16 Everett and Pinsett, Proc. Roy. Soc. A, 1952, 215, 416. 17 Owen and Brinkley, Chem. Rev., 1941, 29, 461. 18 Rontgen, Ann., 1892, 45, 98. 19 Rothmund, Olver. Konig. Vet. Akad. Forhand. Stockholm, 1896, 53, 25. 20 Stem, Wied. Ann., 1896, 59, 652. 21 Cohen and de Boer, Z . physik. Chem., 1913, 84, 41. 22 see Bell, Acid-Base Catalysis (Oxford University Press, 1941), p. 79. 23 Zakharyevsky and Vaslenko, Zhur. Obshchi. Khim., 1956, 26, 2304, Consultant's

Bureau translation, p. 2577. 24 Bogojawlensky and Tammann, Z . physik. Chem., 1897, 23, 13. 25 Hamann, Physico-Chemical Effects 0/ Pressure (Butterwortbs Scientific Publications,

. London, 1957), (a) chap. 8; (b) chap. 9. 26 Bender, J. Amer. Chem. Soc., 1951,73, 1626. Bender, Ginger and Unik, J. Amer.

Chem. Soc., 1958, 80, 1044. 27 see Braude, Quart. Rev., 1950, 4, 404, for a review of such reactions.

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