Acidity in aqueous mixed solvent systems—III

Acidity in aqueous mixed solvent systems—III

ACIDITY IN AQUEOUS MIXED SOLVENT SYSTEMS-III INFLUENCE OF THE ORGANIC ON THE ACIDITY WATER SOLVENT OF AQUO-ORGANIC STRUCTURE, MOLECULAR MIXTURE...

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ACIDITY IN AQUEOUS MIXED SOLVENT SYSTEMS-III INFLUENCE

OF THE ORGANIC

ON THE ACIDITY WATER

SOLVENT

OF AQUO-ORGANIC

STRUCTURE,

MOLECULAR

MIXTURES

HYDROPHILIC

STRUCTURE

(ACIDITY

FUNCTION,

AND HYDROPHOBIC

EFFECTS)

J. P. H. BOYER, R. J. P. CORRIU,R. J. M. PERZ et C. G. REYE Laboratoire

des

Organom&alliques, Universitt des Sciences et Techniques du Languedoc, Place Eugene Bataillon, 34060, Montpellier-Cedex,

France

(Receivedin UK lOFebruary W75;Accepfedforpublicalion

14March 1975)

Abstract--The acidity functions Ho and HR have been determined for different aqueous organic mixtures (diols, carbohydrates) at constant acid concentration. We have checked, with a reaction involving a slow proton transfer, that Ho and HI vary in direct proportion to the acidity of the systems. The results are interpreted in terms of the variations of the proton activity which is dependent on the modifications of the water structure b; the organic cosolvent.

In a preceeding paper,’ we studied the acidity of different aqueous mixed solvents at constant acid concentration. We attributed theobserved decrease in acidity to the increase in proton solvation. The organic solvent, when added to the water, broke the H-bonds and freed the water molecules which in turn became available to solvate the proton. It was apparent that the basicity of these mixtures was a functionof the hydrocarbongroupof theorganic solvent. We intend to clarify the influence of the hydrophobic and hydrophilic groups on the water structure. As evidence of the role of a long carbon chain on the structure of water, we chose a series of dials, since the diols aremoremiscibleinwaterthanthecorrespondingalcohols. Then we studied the acidic properties of aqueous solutions of polyols and carbohydrates. If the basic character of the binary aqueous mixtures depends on the hydrophobic part of the organic product, then one should expect the reverse to be true (i.e. the aqueous system will have a more acidic character when the organic product is hydrophilic). We determined the functions Ho and HRat constant acid concentration with these different aqueous mixed solvent systems. Under the same conditions, we checked that the reaction involving the proton activity in the slow step is modified by the relative proportions of the reactants. Finally, we determined by the isopiestic method, the variations of the activity coefficient of water in the aqueous mixed solvents.

CH,OH H-C-OH HO $4.

-ALREsuLTs

hfeaturement of Ho. We have determined the function Ho’ with constant acid concentration (HCI= 0.1 MI-‘) in the following aqueous mixtures (Fig. I): ethane diol-I,3 propane diol-I.4 butane

% molar of awnic

Ho Is

Fig. 1. Valuesof theacidityfunctionHoin thefollowingmixtures: I, HDglycerol; 2, HIO-ethane diol; 3, HIO-I,3 propane diol; 4, H&1,2,6 hexane triol; 5, H&l,4 butane diol; 6, H20-I,5 pentane diol;7,H20-1,6hexanediol;&H&-1,7heptanediol.

1

I

2

HO-C-H %$H

H$j

H$

$1

H CHzOH Xylitol

: 1 I

CHzOH Adonitol

sdvenf

H-C-OH CHlOH Sorbitol 2075

H-C-OH JH~OH Fructose

H--C-OH CHzOH WIlllit

J. P. H. BOYER ef al.

2076 dial-I.5 pentane diol-I,6 hexane dial-I,7 heptane 1,2,6 hexane triol. Ho was also measured in aqueous solutions carbohydrates

dial-glycerol

and

of the indicated

Me0

(Fig. 2).

01

pK>,$ is the negative logarithm constant of BH’ in water.

-

of the thermodynamic

+

H’

Me

slow

(BH’) - tog (B)

Ho = pK:$

C=CHz

Me0

dissociation

I

SLw”ot..

M*O~yCH1 Y.

0075

0 050

L

E

d

---

_c__-em---

0 025

xylitd

0

I210

5

IO

% moles

of carbohydrates

IO

IS

Fig. 2. Values of acidity functions Ho (-)and HR (---) in aqueous solutions of carbohydrates. (HCI = 0. I MI-‘). Measurement

30

Fig. 3. Hydration

of p methoxy o-methyl styrene mixtures. (HCI = 0.2 Ml-‘).

in aquo organic

of the H,, function HR=

-loge=pKR*+logE

We measured the function Ha (3,4) in several aqueous mixed solvents (HCI = 0.2 MI-‘) and found the same sequences for HR as with Ho (Pig. 2 and Table I).

Table

20

% motor of organic poduct

0

I. H,, values at 25°C in some aqua-organic mixtures(HCI = 0.2 MI-‘)

% moles glycerol HR % moles ethyleneglycol HR % moles 1.2 propanediol H” % moles I ,5 pentanediol HR % moles I ,6 hexanediol HR

0 0.72

9.57 068

15.6 0.56

24.9 0.29

0 0.72

5.3 0.82

12 090

19.6 096

0 0.72

7.5 0.95

I4 I.12

26.6 I .32

0 0.72

4.8 1.18

IO 1.5s

20 I.80

0 0.72

I O.%

I.5 1.13

0

zE r’

0

0

%

6 1.70

I3 I.87

Studies ofreaction rates. We studied p-methoxy a-methyl styrene hydration rates as a function of the proportions of the organic solvent with a constant acid concentration (Figs. 3 and 4). The mechanism of this reactiowthe slow transfer of a proton to an olefin-has already been demonstrated’”

moles

of

corbohydrote

Fig. 4. Hydration of p methoxy (I methyl styrene in aqueous solutionsof carbohydrates. (HCI = 0.1 MI-.‘). We checked that sucrose inversion was negligable throughout our studies of the reaction rates of aqueous solutions of sucrose. Measurement ofwoferactiuity. In Fig. 5, we reported the variations of the activity coefficient of water as a function of the percentage of organic compound.

2078

J. P. H. BOYER cl aL

Wen,” by an essentially cooperative process and further one would expect agreater number of agregates with a larger values of n. Schematically:

K.

'0"

H

/J-I’ ‘H._ H-9

‘Q-H

A

A

The solvation of protons is reduced. Consequently, the numberofthewatermoleculesandtheactivityoftheproton increases. Thus, the formation of various important agregates can explain the progressive evolution of the acid properties of different aqua-organic solutions. (2) Systems with hydrophobic character. The relative experimental data of these systems were more difficult to interpret. Fist, the acidity decreases regularly with the increase in size of the hydrocarbon group of the organic compound (Fig. 1). The water activity in those aqua-organic mixtures increases with the increase of the hydrophobic part in the organic compound (Fig. 5). Furthermore, it seems reasonable to consider that the entropy of the water diol mixtures is negative by analogy to the results obtained with aqueous alcohol mixtures”. This seems a pn’ori in contradiction to variations of Ho and fH* which denote agreater freedom between water molecules. One can consider that entropy is a mean measure of the interaction between the substituents of a mixture. At the same time, there is not a qualitative distinction between the different types of bonds in the system. (bonds between solute-solute, solute-solvent or solvent-solvent are not differentiated). This fact leads one to consider, for the aqua-organic systems, the possibility of structure of a different nature than those which exist in the aqueous solutions of carbohydrates. Weexplain theseresultsas theprogressivedestructionof three dimensional clusters in water by the hydrocarbon groups of the organic molecules. The thermodynamic equilibrium existing between the monomeric water molecules and the agregates is then shifted towards the monomers. The rupture of the agregates is a function of the increase of the hydrocarbon group of the organic compound. Thus, there is a greater solvation of the proton due to the liberation of non bonded water molecules (Increase in f&. According to Frank and Wen,” the unique properties of water are due to the existence of these 3- dimensional clusters. Consequently, their destruction gives the systems properties different than those of water. This is what we observed experimentally. The system’sincreasedorder (AS < 0) would nolonger be duetoH-bondingasinthecaseoftheaqueouscarbohydrate solutions but instead to hydrophobic forces. Hydrophobic bonding may be defined” as an interaction of molecules with each other which is stronger than the interaction of the separate molecules with water and which cannot be accounted for by covalent, electrostatic, H-bond or charge transfer forces. Thus the hydrocarbon groups of the tThis apparatus has been constructed technician (C.N.R.S.).

by M. G. Guiraud

molecules come together in aqueous solutions; the water molecules group around these agregates. In doing so, a new water structure is formed. These interactions do not imply H-bonding and the interactions themselves contribute to lessen H-bonds. This fact agrees with the variations in activitycoefficientsof waterinthedifferent systems(Fig.5). Weaccountedfortheacidityoftheaquo-organicsystems by structural modifications of water because of addition of the organic solvent. We have shown that these structural modifications depended on the organic solvent and more particularly on the relative hydrophilic and hydrophobic character of the solvent.

p- Methoxy -(I - methylstyrencp -Methoxyphenyl-2hydroxy-2 - propane was prepared by addition of MeMgBr on pmethoxyacetophenone followed by hydrolysis. The alcohol was distilled in the piesence of hydroquinone, giving p - methoxy - a methylstyrene (EL = 90°C). This was recrystallized from EtOH (mp = 32°C) and refrigerated. Tri - p - anisylmethonol. This was prepared by oxidation of tri p -anisylmethaneP with an equimolar amount of lead oxide in acetic acid (mp 82-82,5”). hfeasumnen~ of Ho. (BH’) Ho = pKBH* - log (B) Ho=pK.,--log&$ 4, DBH+ and D represent respectively the optical densities of the indicatorsolutionsunderbasic,acidicandmeasuredconditions.The indicators usedwerep-nitraniline(pK,, = 099) and m-nit.raniline ([email protected]= 2.50). Measurements of Ho were made with a thermostated D. B. Beckman UV Spectrophotometer al 25”. Measure of H. H.=pK,++log$+) we used tri p-anisylmethanol

(pK = 0.82) as the indicator.

(ROH) -=+ (R 1

D., - D D-Dm,

by+. 410~ et D represent respectively the optical densities of solutions of the indicator under acidic, basic and measured conditions. Measurements of &+ and D were made with a maximum absorption of the indicator in the tested solutions. Kineticmeasumnmls.Thehydrationreactionofp -methoxy-cr methylstyrene in the aqueous acid homogeneous mixture was followed directly on the W spectrum by the variation in the optical density of the oletin during transformation in the alcohol. We considered the reaction as a lirst order reaction at least during two half-lives. v=k,Jolefinel,=

-dy

Log&=Log$=klt The rate constant of hydration can be easily calculated by using the slope of the linear part of the mph log D, = f(t). We used the Beckman UV Spectrophotometer at u” for the kinetic measurements. Deferminotfon of actiaity co&ients of water in dilg’cmt water alcohol solukms. We used the isopiestic method. Processes and apparatus used were as descrii by Scatchard.? As reference__ solns, we employedeithersaccharosesolnPorglycero1solutions.”

2079

Acidity in aqueous mixed solvent systems-111 Table 2. Isotonic concentrations at 25°C (moles/kg water)

Exp. 1 2 3 4 5 6 7 8 9 10 11 12 13 14

Saccharose Glucose

Sorbitol

Xylitol

2.82

366 3.91 392 3.% 4.26

Glycerol

I,3 proPanediol

3.45 3.57 4.43 4.76

4.65

4.91

4.91

5.02

5.14

I,4 butanediol

I,5 pentanediol

I,6 Hexatanediol

I,7 Heptanediol

387 4.22 5.29

4.13 4.70 6.11

4.42 5.65 966

6.49 II.19 23.78

688

IO.91

5.20

5.47

6.14

6.04

744

5.73 5.82 5.94 8.32 IO.65 1590

5.79

640 6.59 6.83

6.59

7.88 8.42 9.61 17.m

A correspondance between these last solns and the solns of saccharose was established by Scatchard.= In Table 2 are indicated the concentrations of ditferent isopiestic solutions and in Fig. 5 are found the rational activity coefficients of water fHP, as a function of the molality of the alcohol. Products. Ethylene glycol and I.4butane diol weredistilled using a column packing with glass rings. 1,2-propane diol, 1,3-propane diol, glycerol, 1j-pentane diol, 1.2.6hexane trio1 and I,bhexane dial (Merck products). IJ-heptane diol (Schuchardt product). D( + ) glucose, xylitol, adonitol, D( - ) fructose D( - ) sorbitol, D( - ) mannitol (Merck products). REFIXKNCES ‘J. Boyer, R. Corriu. R. Perz and C. Reye. Tetrahedron Letters 4115 (1974). ‘L.P.HammettandA.J.Deyrup,J.Am. Chem.Soc54,2721(1932). ‘F. H. Westheimer and M. S. Kharasch, Ibid 68, 1871 (1946). ‘A. M. Lowen, M. A. Murray and G. Williams, 1. Cgcm. Sot. 3318 (1950). ‘W. M. Schubert, B. Lamm and J. B. Keeffe, 1. Am. Chcm. Sot. 86, 4727 (1964). “J. C. Simandoux, B. Torck. M. Hellin and F. Coussemant, Bull. Sot. Chim. 4402 (lW2). ‘F. Coussemant. M. Hellin and B. Torck, Lssfoncfions d’aciditttct lcurs utilisations en catalyst acido basiquc. p. 105. Gordon & Breach (1%9).

TJITVOL3lNO.17-H

Ethylene glycol

20.50

7.20 11.67 15.30 25.47

15.12 1860

‘R. Gaboriaud. 1. Chim. Phys. 349 (1970). ‘E. A. Btaude and E. S. Stem, 1. Chem. Sot. 1976 (1948). ‘“P. Salomaa, Acta Chcm. Scand. 11, I25 (1957). “R. W. Gurney, Ionic Processesin Solutions. McGraw-Hill, New York (1953). “H. Strehlow, 2. Phys. Chcm., Frankjurt, 24, 240 (1960). “C.E.NewallandA.M.Eastham,Can.J. Chem.39,1752(1%1). ‘J. Koskikallio and J. Suomen, Knistilehti, B36, 111 (1957). “W. Gerrard and E. D. MackJen, Chcm. RNS 59. 1105 (1959). ‘“J. L. Kavanau, Water and Solute- Water Solutions, p. 10. Holden Day, New York (1964). “F. Franks and D. J. G. Ives, Quart. Reus 20, 11 (1966). “H. S. Frankand W. Y. Wen.Disc. FaradaySoc.24,133 (1957). ‘7. B. Taylor and J. S. Rowlinson, Trwts. Faraday Sot. 51, 1183 (1955). ‘OR.A.RobinsonandRH.Stokes,J.Phys.Chem.65,662(1%1). “W. P. Jencks. Cafalysis in chemistry and enzymology. McGrawHilt Series in Advanced Chemistry, p. 394. “A.BayerandV.Villiger,Bcr.Dtsch. Chmt.Gcs.35,1198(1902). UG. Scatchard and W. J. Hamer, S. E. Wood, J. Am. Chcm. Sot. 60, 3061 (1938). %R. A. Robinson and R. H. Stokes, Electrolyte Solutions. Butterworths. London, p. 478 (1959). =A. K. Covington and P. Jones, Hydrogen Bondcdsolucnt sysfcms, Taylor & Francis (1968).