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Vapour pressures at six temperatures between 278.15 K and 323.15 K and excess molar functions atT= 298.15 of (butanone + methanol or ethanol)
J. Chem. Thermodynamics 1996, 28, 567–576
Vapour pressures at six temperatures between 278.15 K and 323.15 K and excess molar functions at T = 298.15 K of (butanone + methanol or ethanol) R. Garriga, F. Sa´nchez, P. Pe´rez, and M. Gracia a Departamento de Quı´ mica Orga´nica-Quı´ mica Fı´ sica (Area de Quı´ mica-Fı´ sica) , Facultad de Ciencias, Universidad de Zaragoza, Spain Vapour pressures of (butanone + methanol or ethanol) at six temperatures between 278.15 K and 323.15 K were measured by a static method. Excess enthalpies and volumes of mixtures containing methanol were also measured at T=298.15 K. Reduction of the vapour pressures to obtain activity coefficients and excess molar Gibbs free energies was carried out by Barker’s method. Azeotropic mixtures with a minimum boiling temperature were observed and the azeotropic coordinates were calculated. The experimental results are analyzed in terms of the specific interaction between the hydroxyl hydrogen of the alcohol and the ketoxy group. 7 1996 Academic Press Limited
1. Introduction Following our thermodynamic study of binary mixtures(1–5) of (alcohol+a polar component), we report here vapour pressures at six temperatures between 278.15 K and 323.15 K of (butanone + methanol or ethanol). We have also measured excess molar enthalpies H mE and volumes VmE , at T = 298.15 K, for (butanone + methanol). Excess enthalpies and volumes, at T = 298.15 K, of (butanone+ethanol) have been taken from the literature.(1,6,7) This type of thermodynamic information is necessary for a better understanding of hydrogen bonding in the liquid state.
2. Experimental Butanone, methanol, and ethanol were Fluka products (better than mass fraction 0.995 per cent). All the liquids were kept over a molecular sieve (0.3 nm) and used without further purification. The vapour-pressure measurements were performed by a static method. The apparatus was similar to Marsh’s(8) except for experimental details which are described elsewhere.(9,10) Manometric readings were made with a cathetometer to a
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7 1996 Academic Press Limited
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20.01 mm, and the pressure reproducibility was estimated to be better than 15 Pa. The temperature of the liquid sample was constant within 10 mK and the uncertainty in the mole fraction was Q3·10−4 . The HmE measurements were carried out with a calorimeter described elsewhere,(11) and VmE was calculated from density measurements made with a densimeter (Anton Paar DMA 60/DMA 602).
3. Results and discussion The second virial coefficients, at T = 325.0 K, of CH3COC2 H5 (BBB = −1840 cm3·mol−1 ), CH3OH (BAA = −1180 cm3·mol−1 ), and C2 H5OH (BAA = −2620 cm3·mol−1 ), were taken from reference 12. We assume that the vapour phase is an ideal mixture of imperfect gases: BAB=(BAA+BBB )/2. Table 1 contains the experimental vapour pressures of the pure compounds which are compared with values calculated from equations found in the literature. Table 2 lists our total vapour-pressure measurements along with the activity coefficients fA and fB , and the values of GmE fitted by Barker’s method(15) with a Redlich-Kister polynomial: m
GmE /RT=x(1−x)· s Af (1−2x) f .
(1)
f=0
The activity coefficients fi are obtained through differentiation of equation (1):
$
m
%
(2)
%
(3)
ln fA=(1−x)2 A0+ s {Af (1−2x) f−2fAf x(1−2x) f−1} ,
$
f=1
m
ln fB=x 2 A0+ s {Af (1−2x) f+2fAf (1−x)(1−2x) f−1} , f=1
TABLE 1. Vapour pressures (p*) of the pure compounds T K
278.15 288.15 298.15 303.15 308.15 313.15 323.15
CH3COC2 H5
CH3OH
C2 H5OH
This work
Lit.(13)
This work
Lit.(14)
This work
Lit.(14)
4.277 7.334 12.071 15.281 19.110 23.683 35.540
4.278 7.342 12.060 15.231 19.060 23.649 35.548
p*/kPa 5.488 9.903 16.960 21.904 27.991 35.450 55.566
5.512 9.876 16.945 21.868 27.962 35.444 55.567
2.246 4.314 7.867 10.531 13.823 17.980 29.419
2.248 4.312 7.875 10.462 13.756 17.905 29.494
569
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
TABLE 2. Experimental vapour pressures p, deviations dp=p−p(calc), activity coefficients fi , excess molar Gibbs free energy GmE , smoothing equations, and standard deviations s(p) x
p Pa
dp Pa
fA
fB
GmE J·mol−1
x
p Pa
dp Pa
fA
fB
GmE J·mol−1
0.0963 0.2051 0.2548 0.3303 0.4108
{(1−x)CH3COC2 H5+xCH3OH} at T=288.15 K 8727 11 2.1160 1.0097 194 0.5711 10938 13 1.1839 1.3395 9661 −13 1.7654 1.0425 359 0.6939 10982 6 1.0932 1.5367 9994 14 1.6436 1.0650 416 0.7437 10966 21 1.0657 1.6402 10320 −17 1.4919 1.1085 482 0.8652 10679 −12 1.0187 1.9762 10602 −12 1.3642 1.1684 525 0.9639 10163 −28 1.0014 2.3779 GmE /RT=x(1−x){0.9047+0.0001(1−2x)+0.0382(1−2x)2}; s(p)=16 Pa
531 463 417 258 78
0.0963 0.2050 0.2548 0.3303 0.4108
{(1−x)CH3COC2 H5+xCH 3OH} at T=298.15 K 14411 17 2.0596 1.0098 194 0.5711 18337 21 1.1815 1.3198 15999 −21 1.7242 1.0419 358 0.6939 18476 −6 1.0945 1.5061 16559 6 1.6095 1.0635 414 0.7437 18484 19 1.0674 1.6059 17177 −12 1.4682 1.1046 480 0.8652 18123 −21 1.0198 1.9402 17697 0 1.3499 1.1606 523 0.9639 17400 −8 1.0015 2.3599 GmE /RT=x(1−x){0.8719−0.0115(1−2x)+0.0582(1−2x)2}; s(p)=16 Pa
531 466 421 264 81
0.0963 0.2049 0.2547 0.3301 0.4107
{(1−x)CH3COC2 H5+xCH3OH} at T=303.15 K 18219 17 2.0120 1.0091 190 0.5711 23435 28 1.1807 1.3084 20290 −19 1.7028 1.0392 352 0.6939 23662 −5 1.0942 1.4921 21012 −2 1.5949 1.0597 409 0.7437 23671 5 1.0672 1.5907 21857 −6 1.4605 1.0989 474 0.8652 23291 −25 1.0197 1.9208 22551 1 1.3460 1.1529 519 0.9639 22453 9 1.0015 2.3338 GmE /RT=x(1−x){0.8534−0.0245(1−2x)+0.0519(1−2x)2}; s(p)=16 Pa
530 466 422 264 81
0.0962 0.2048 0.2546 0.3298 0.4105
{(1−x)CH3COC2 H5+xCH3OH} at T=308.15 K 22801 20 1.9770 1.0088 188 0.5710 29587 31 1.1767 1.2995 25446 −28 1.6817 1.0379 349 0.6939 29944 −2 1.0922 1.4780 26387 2 1.5780 1.0578 405 0.7437 29970 −5 1.0657 1.5736 27486 −4 1.4487 1.0958 470 0.8652 29590 −27 1.0193 1.8929 28399 3 1.3377 1.1483 515 0.9639 28638 18 1.0015 2.2907 GmE /RT=x(1−x){0.8335−0.0261(1−2x)+0.0499(1−2x)2}; s(p)=19 Pa
526 463 419 263 80
0.0961 0.2046 0.2545 0.3296 0.4104
{(1−x)CH3COC2 H5+xCH3OH} at T=313.15 K 28266 2 1.9440 1.0086 187 0.5709 37014 12 1.1739 1.2893 31656 −6 1.6594 1.0371 345 0.6938 37582 4 1.0915 1.4619 32836 8 1.5596 1.0564 401 0.7437 37646 −5 1.0654 1.5550 34241 −11 1.4357 1.0931 466 0.8652 37288 −25 1.0193 1.8681 35444 3 1.3290 1.1438 510 0.9639 36218 33 1.0015 2.2625 GmE /RT=x(1−x){0.8133−0.0299(1−2x)+0.0547(1−2x)2}; s(p)=16 Pa
522 461 417 263 81
0.0958 0.2042 0.2542 0.3287 0.4095
{(1−x)CH3COC2 H5+xCH3OH} at T=323.15 K 42483 54 1.8713 1.0073 179 0.5706 56693 150 1.1586 1.2780 47793 −15 1.6225 1.0330 335 0.6936 57538 −73 1.0804 1.4412 49607 −83 1.5305 1.0511 390 0.7436 57779 −38 1.0565 1.5260 51993 8 1.4140 1.0858 454 0.8653 57587 −24 1.0160 1.7939 53934 10 1.3105 1.1352 499 0.9639 56402 25 1.0012 2.1006 GmE /RT=x(1−x){0.7700−0.0168(1−2x)+0.0199(1−2x)2}; s(p)=67 Pa
509 445 401 248 75
570
R. Garriga et al. TABLE 2—continued
x
p Pa
dp Pa
0.0697 0.1515 0.1956 0.3029 0.4110
4400 4450 4453 4421 4349
−1 −6 −9 −5 8
fA
fB
GmE J·mol−1
p Pa
x
{(1−x)CH3COC2 H5+xC2 H5OH} at 2.5298 1.0060 163 0.5101 2.1280 1.0278 318 0.6648 1.9578 1.0460 388 0.7597 1.6401 1.1091 514 0.8168 1.4177 1.2024 583 0.9126
dp Pa
T=278.15 K 4244 19 3942 1 3650 −21 3454 2 2956 10
fA
fB
GmE J·mol−1
1.2720 1.1202 1.0609 1.0354 1.0081
1.3190 1.5803 1.8083 1.9801 2.3487
597 529 433 355 190
GmE /RT=x(1−x){1.0344+0.0321(1−2x)+0.0297(1−2x)2}; s(p)=11 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=288.15 K 0.0697 0.1515 0.1956 0.3029 0.4111
7575 7694 7743 7717 7581
9 −6 16 11 −12
2.3474 2.0223 1.8772 1.5941 1.3869
1.0050 1.0238 1.0399 1.0979 1.1861
154 303 371 494 563
0.5102 0.6649 0.7598 0.8169 0.9127
7394 6970 6517 6225 5325
−24 0 −12 52 −37
1.2487 1.1056 1.0516 1.0292 1.0064
1.2973 1.5423 1.7457 1.8914 2.1819
577 508 412 336 177
GmE /RT=x(1−x){0.9643+0.0367(1−2x)−0.0096(1−2x)2}; s(p)=25 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=298.15 K 0.0697 0.1515 0.1955 0.3029 0.4111
12532 12807 12888 12920 12767
−7 −12 −2 12 5
2.2565 1.9308 1.7939 1.5379 1.3580
1.0056 1.0251 1.0411 1.0954 1.1736
153 299 363 481 546
0.5101 0.6649 0.7600 0.8171 0.9128
12502 11867 11140 10687 9519
−3 33 −51 6 17
1.2385 1.1099 1.0572 1.0338 1.0080
1.2696 1.4844 1.6749 1.8210 2.1423
560 500 412 339 183
GmE /RT=x(1−x){0.9048+0.0150(1−2x)+0.0492(1−2x)2}; s(p)=22 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=308.15 K 0.0697 0.1514 0.1954 0.3029 0.4111
19941 20438 20628 20789 20676
23 −24 −4 2 7
2.1076 1.8382 1.7216 1.4976 1.3350
1.0048 1.0220 1.0361 1.0850 1.1564
145 283 345 459 524
0.5102 0.6650 0.7599 0.8170 0.9127
20361 19466 18510 17840 16140
4 17 −37 18 1
1.2244 1.1035 1.0538 1.0317 1.0075
1.2451 1.4441 1.6192 1.7524 2.0415
540 483 398 328 177
GmE /RT=x(1−x){0.8429+0.0006(1−2x)+0.0363(1−2x)2}; s(p)=19 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=313.15 K 0.0696 0.1514 0.1953 0.3028 0.4111
24765 25461 25725 25993 26038
22 −32 −15 −24 94
2.0561 1.8042 1.6943 1.4806 1.3237
1.0045 1.0209 1.0345 1.0817 1.1510
142 278 339 452 516
0.5102 0.6650 0.7599 0.8171 0.9128
25586 24585 23473 22682 20693
−28 5 −49 19 13
1.2166 1.0994 1.0514 1.0301 1.0070
1.2370 1.4293 1.5967 1.7231 1.9933
532 476 392 322 173
GmE /RT=x(1−x){0.8171+0.0012(1−2x)+0.0293(1−2x)2}; s(p)=41 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=323.15 K 0.0696 0.1513 0.1951 0.3027 0.4110
37274 38582 38999 39714 39807
−23 −1 −34 61 66
1.9571 1.7220 1.6229 1.4347 1.2996
1.0046 1.0206 1.0334 1.0762 1.1369
137 268 325 431 492
0.5102 0.6650 0.7601 0.8172 0.9129
39333 38204 36854 35793 33184
−93 8 −16 15 7
1.2066 1.1009 1.0543 1.0327 1.0080
1.2112 1.3812 1.5378 1.6615 1.9435
GmE /RT=x(1−x){0.7586−0.0208(1−2x)+0.0640(1−2x)2}; s(p)=46 Pa
510 462 385 320 175
571
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
x being the mole fraction of alcohol in the liquid phase. The vapour pressure is then given by: p(calc)=xA fA p* A RA+xB fB p* B RB ,
(4)
using for nonideality of the vapour-phase the corrections: 2 RA=exp[{(V* A −BAA )(p−p* A )−pdAB yB }/RT ];
(5)
2 RB=exp[{(V* B −BBB )(p−p* B )−pdAB yA }/RT],
(6)
where yA and yB are the vapour-phase mole fractions of alcohol and butanone, respectively, and dAB is given by: dAB=2BAB−BAA−BBB ,
(7)
which in our case is taken to be equal to zero. Molar volumes of the pure compounds used in the Barker analysis were taken from the literature.(16,17) For a given composition, the sample temperature is changed and a slight variation of the true liquid mole fraction may be detected in table 2, according to the variable composition of the vapour phase. Analytic equations for GmE at the lower and higher temperatures are plotted in figure 1. The HmE and VmE experimental results for {(1−x)CH3COC2 H5+xCH3OH}, at T=298.15 K, are given in table 3, and are plotted in figure 2 together with TSmE at the same temperature, calculated from TSmE =(HmE −GmE ). The curves for HmE and VmE were found by fitting the experimental results with a polynomial: m
QmE =x(1−x)· s Af (1−2x) f ,
(8)
f=0
where QmE is either HmE /(J·mol−1 ) or VmE /(cm3·mol−1 ). We have tested the consistency of the molar enthalpies and excess molar Gibbs free energies by means of the Gibbs-Helmholtz equation. By assuming that Af varies linearly with temperature, the calculated values of HmE are shown as curves in figure 2, together with HmE experimental data, at T=298.15 K. The parameters were observed to be acceptably consistent and were more so for mixtures containing methanol. Azeotropic mixtures with a minimum boiling temperature were observed over the whole range of temperature. Azeotropic mole fractions z were calculated graphically, assuming ideal behaviour of the vapour, from the well-known equation fB /fA = p* A /p* B . Azeotropic compositions show a linear relation to temperature, according to the equation: (9)
z=a+b(T/K).
For {(1−x)CH3COC2 H5+xCH3OH}, a=−0.3032 and b=3.386·10 ; and for {(1−x)CH3COC2 H5+xC2 H5OH}, a=−0.9698 and b=4.136·10−3. Along the azeotropic line, assuming both ideal behaviour of the vapour phase and negligible volume of the liquid phase, the equation: −3
d ln pz /dT=Dvap Hm /RT 2,
(10)
572
R. Garriga et al.
FIGURE 1. Excess molar Gibbs free energies GmE for {(1−x)CH3COC2 H5+xROH}. R, CH3 : curve 2, T=288.15 K; curve 3, T=323.15 K. R=C2 H5 : curve 1, T=278.15 K; curve 4, T=323.15 K. TABLE 3. Excess molar enthalpies and volumes (XmE ), smoothing equations, and standard deviations s(XmE ) for {(1−x)CH3COC2 H5+xCH3OH} at T=298.15 K x
0.1061 0.1788 0.2349
0.0193 0.0414 0.1267 0.1346
XmE
x
XmE
x
XmE
x
QmE =HmE /(J·mol−1 ) 0.2594 592 0.4842 686 0.7278 0.3670 662 0.5811 650 0.8316 0.4332 685 0.6816 532 0.9123 E E Hm=x(1−x){2716+577(1−2x)+106(1−2x)2}; s(Hm )=9 J·mol−1 QmE =VmE /(cm3·mol−1 ) −0.020 0.1535 −0.134 0.4832 −0.264 0.7771 −0.038 0.2427 −0.189 0.5935 −0.263 0.8705 −0.117 0.3482 −0.235 0.6439 −0.252 0.9570 −0.118 0.3968 −0.249 0.7027 −0.238 E E 2 Vm=x(1−x){−1.061+0.127(1−2x)−0.121(1−2x) }; s(Vm )=0.002 cm3·mol−1 299 460 545
XmE
485 341 190
−0.201 −0.140 −0.051
573
P, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
FIGURE 2. Excess molar functions, at T=298.15 K, for {(1−x)CH3COC2 H5+xROH}. (a) ——, T·SmE ; – – – –, HmE . R=CH3 : w, this work; W, reference 6. R=C2 H5 : q, reference 1; Q, reference 6. —·—·—·, calculated from Gibbs-Helmholtz equation. Arrows indicate T·SmE and HmE for the same systems. (b) VmE . R=CH3 : w, this work; W, reference 7. R=C2 H5 : q, reference 1; Q, reference 7.
is satisfied. If we accept that the enthalpy of azeotropic vaporization is constant, the azeotropic pressure is related to the temperature in a similar way to that for a pure substance: ln (pz /kPa)=A+B(T/K)−1 . (11) For {(1−x)CH3COC2 H5 + xCH3OH}, A = 17.74 and B = −4419; and for {(1−x)CH3COC2 H5+xC2 H5OH}, A = 17.23 and B = −4377. Azeotropic compositions and pressures, both experimental and calculated from equations (9) and (11), are compared in table 4 and plotted in figure 3. From equations (10) and (11), the molar enthalpy of azeotropic vaporization for the system containing TABLE 4. Azeotropic pressures and mole fractions {(1−x)CH3COC2 H5+xCH3OH}
a
T K
za
pz /Pa a
zb
pz /Pa c
278.15 288.15 298.15 303.15 308.15 313.15 323.15
0.670 0.709 0.725 0.740 0.757 0.790
10973 18478 23668 29968 37607 57727
0.673 0.706 0.723 0.740 0.757 0.791
11033 18453 23564 29852 37533 58086
{(1−x)CH3COC2 H5+xC2 H5OH} za
pz /Pa a
zb
pz /Pa c
0.185 0.221 0.259
4452 7737 12907
0.181 0.222 0.263
4474 7724 12856
0.301 0.326 0.371
20786 26002 39773
0.305 0.325 0.367
20700 25969 40022
Experimental value; b calculated from equation (9); c calculated from equation (11).
574
R. Garriga et al.
FIGURE 3. (a) Azeotropic mole fractions z, and (b), azeotropic vapour pressures p2 , for {(1−x)CH3COC2 H5+xROH}. w, R=CH3 ; W, R=C2 H5 .
methanol is Dvap Hz=36.7 kJ·mol−1 , and with ethanol Dvap Hz=36.4 kJ·mol−1 . By using the azeotropic compositions together with the vaporization enthalpies for the pure compounds(18) butanone (Dvap H B* = 34.5 kJ·mol−1 ), methanol (Dvap H A*=37.4 kJ·mol−1 ), and ethanol (Dvap H A*=42.3 kJ·mol−1 ), the equation: Dvap Hz=(1−z)·Dvap H B*+z·Dvap H A* (12) is satisfied. Nagata et al.(19) measured vapour pressures for {(1−x)CH3COC2 H5 + CH3OH} at T=323.15 K, and from these data we have calculated GmE(x=0.5)= 558 J·mol−1 , z=0.783, and pz=58.82 kPa; at the same temperature, our experimental results are: GmE(x=0.5)=517 J·mol−1 , z=0.790, and pz=58.27 kPa. The HmE of this system has also been measured by Nagata et al.(6) at T=298.15 K; their values, at x=0.5, are higher than ours by about 4 per cent. The VmE measured by Letcher(7) compare satisfactorily with ours, figure 2(b). TABLE 5. Thermodynamic excess functions for {0.5 R OH+0.5CH3COC2 H5 or +0.5n-C6 H14 }, at T=298.15 K n-C6 H14 (23)
CH3COC2 H5
Alkanol
HmE J·mol−1
TSmE J·mol−1
VmE cm3·mol−1
HmE J·mol−1
TSmE J·mol−1
VmE cm3·mol−1
CH3OH C2 H5OH
679 1029 b
139 468
−0.265 −0.071 b
500 555
(−1160) a −850
0.41
a
Extrapolated value. b Reference 1.
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
575
For {(1−x)CH3COC2 H5 +C2 H5OH} at T = 298.15 K, Ohta et al.(20) obtained G (x=0.5) = 554 J·mol−1 , z = 0.261, and pz = 13.04 kPa; these values compare favourably with ours: GmE(x = 0.5) = 561 J·mol−1 , z = 0.259, and pz=12.91 kPa. For this system, Nagata et al.(6) found, at T = 298.15 K, HmE(x=0.5)=1071 J·mol−1 , while In˜arrea et al.(1) give, at the same temperature, HmE(x=0.5)=1029 J·mol−1 . The VmE data taken from the literature(1,7) show discrepancies as shown in figure 2(b). In mixtures of CH3OH or C2 H5OH with CH3COC2 H5 as a strong polar solvent (dipole moment p=1.1·10−29 C·m)(21) strong solvent–solvent and hydroxyl group–solvent interactions come into play and the HmE and SmE , at T=298.15 K, are much more positive than those for inert-solvent solutions owing to the interaction between the hydroxyl hydrogen and the oxygen of the carbonyl group. According to Barker,(22) in systems with specific interactions, more hydrogen bonds are broken at a given mole fraction. In table 5, experimental values of these thermodynamic properties are briefly summarized and compared with those for n-hexane solutions.(23) Presumably, the strong entropy increase is associated with a large gain in configurational freedom on breaking hydrogen bonds by the specific interaction. In figure 2(b), VmE are plotted. The excess volumes are the result of opposing effects: a positive one due to the break-up of the alcohol structure, and a negative one due to the interaction between the hydroxyl hydrogen and the oxygen of the carbonyl group. A third contribution, due to the destruction of the strong ketoxy-to-ketoxy interactions by the alkyl chain of the alcohol, is less important for short alkanol molecules. E m
R.G. wishes to thank the Diputacio´n General de Arago´n for the award of a predoctoral grant (BCB49/93). REFERENCES 1. In˜arrea, J.; Valero, J.; Pe´rez, P.; Gracia, M.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1988, 20, 193–199. 2. Garriga, R.; Putze, I.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 481–491. 3. Garriga, R.; Sa´nchez, F.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 887–895. 4. Garriga, R.; Sa´nchez, F.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 1057–1066. 5. Garriga, R.; Ilarraza, J.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1996, 28, 230–243. 6. Nagata, I. Int. Data Ser., Selec. Data Mixtures, Ser. A 1984, 2, 81–82. 7. Letcher, T. M.; Nevines, J. J. Chem. Eng. Data 1995, 40, 293–295. 8. Marsh, K. N. Trans. Faraday Soc. 1968, 64, 883–893. 9. Gracia, M.; Sa´nchez, F.; Pe´rez, P.; Valero, J.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1992, 24, 463–471. 10. Pardo, J.; Pe´rez, P.; Royo, F.; Gracia, M.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1987, 19, 521–526. 11. Gutie´rrez Losa, C.; Gracia, M. Rev. Acad. Cienc. Exactas Fis.-Quim. Natur. Zaragoza XXVI 1971, 26, 1: 101–135. 12. Dymond, J. H.; Smith, E. B. The Virial Coefficients of Pure Gases and Mixtures. Clarendon Press: Oxford. 1980, pp. 42, 80 and 110. 13. Ambrose, D.; Ghiassee, N. B. J. Chem. Thermodynamics 1987, 19, 505–519. 14. Ambrose, D.; Ellender, J. H.; Sprake, C. H. S. J. Chem. Thermodynamics 1974, 6, 909–914. 15. Barker, J. A. Aust. J. Chem. 1953, 6, 207–210. 16. TRC Thermodynamic Tables. Non-Hydrocarbons. Thermodynamics Research Center, The Texas A&M University System: College Station, TX. 1991, p. d-5870. 17. TRC Thermodynamic Tables. Non-Hydrocarbons. Thermodynamics Research Center, The Texas A&M University System: College Station, TX. 1966, p. d-5030.
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18. Riddick, J. A.; Bunger, W. B.; Sakano, T. K. Organic Solvents. Physical Properties and Methods of Purification. Wiley: Chichester. 1986, pp. 190, 192 and 338. 19. Gmehling, J.; Onken, U. Vapor-Liquid Equilibrium Data Collection. DECHEMA Chemistry Series, Vol. I, Parts 2a. DECHEMA: Frankfurt. 1986, p. 138. 20. Ohta, T.; Koyabu, J.; Nagata, I. Fluid Phase Equilib. 1981, 7, 65–73. 21. McClellan, A. L. Tables of Experimental Dipole Moments. Freeman: San Francisco. 1963. 22. Barker, J. A. J. Chem. Phys. 1952, 20, 1526–1532. 23. Brown, I.; Fock, W.; Smith, F. Aust. J. Chem. 1969, 1, 273–291.
(Received 29 November 1995; in final form 5 January 1996)
WA-088
Vapour pressures at six temperatures between 278.15 K and 323.15 K and excess molar functions at T = 298.15 K of (butanone + methanol or ethanol) R. Garriga, F. Sa´nchez, P. Pe´rez, and M. Gracia a Departamento de Quı´ mica Orga´nica-Quı´ mica Fı´ sica (Area de Quı´ mica-Fı´ sica) , Facultad de Ciencias, Universidad de Zaragoza, Spain Vapour pressures of (butanone + methanol or ethanol) at six temperatures between 278.15 K and 323.15 K were measured by a static method. Excess enthalpies and volumes of mixtures containing methanol were also measured at T=298.15 K. Reduction of the vapour pressures to obtain activity coefficients and excess molar Gibbs free energies was carried out by Barker’s method. Azeotropic mixtures with a minimum boiling temperature were observed and the azeotropic coordinates were calculated. The experimental results are analyzed in terms of the specific interaction between the hydroxyl hydrogen of the alcohol and the ketoxy group. 7 1996 Academic Press Limited
1. Introduction Following our thermodynamic study of binary mixtures(1–5) of (alcohol+a polar component), we report here vapour pressures at six temperatures between 278.15 K and 323.15 K of (butanone + methanol or ethanol). We have also measured excess molar enthalpies H mE and volumes VmE , at T = 298.15 K, for (butanone + methanol). Excess enthalpies and volumes, at T = 298.15 K, of (butanone+ethanol) have been taken from the literature.(1,6,7) This type of thermodynamic information is necessary for a better understanding of hydrogen bonding in the liquid state.
2. Experimental Butanone, methanol, and ethanol were Fluka products (better than mass fraction 0.995 per cent). All the liquids were kept over a molecular sieve (0.3 nm) and used without further purification. The vapour-pressure measurements were performed by a static method. The apparatus was similar to Marsh’s(8) except for experimental details which are described elsewhere.(9,10) Manometric readings were made with a cathetometer to a
To whom correspondence should be addressed.
0021–9614/96/050567+10 $18.00/0
7 1996 Academic Press Limited
568
R. Garriga et al.
20.01 mm, and the pressure reproducibility was estimated to be better than 15 Pa. The temperature of the liquid sample was constant within 10 mK and the uncertainty in the mole fraction was Q3·10−4 . The HmE measurements were carried out with a calorimeter described elsewhere,(11) and VmE was calculated from density measurements made with a densimeter (Anton Paar DMA 60/DMA 602).
3. Results and discussion The second virial coefficients, at T = 325.0 K, of CH3COC2 H5 (BBB = −1840 cm3·mol−1 ), CH3OH (BAA = −1180 cm3·mol−1 ), and C2 H5OH (BAA = −2620 cm3·mol−1 ), were taken from reference 12. We assume that the vapour phase is an ideal mixture of imperfect gases: BAB=(BAA+BBB )/2. Table 1 contains the experimental vapour pressures of the pure compounds which are compared with values calculated from equations found in the literature. Table 2 lists our total vapour-pressure measurements along with the activity coefficients fA and fB , and the values of GmE fitted by Barker’s method(15) with a Redlich-Kister polynomial: m
GmE /RT=x(1−x)· s Af (1−2x) f .
(1)
f=0
The activity coefficients fi are obtained through differentiation of equation (1):
$
m
%
(2)
%
(3)
ln fA=(1−x)2 A0+ s {Af (1−2x) f−2fAf x(1−2x) f−1} ,
$
f=1
m
ln fB=x 2 A0+ s {Af (1−2x) f+2fAf (1−x)(1−2x) f−1} , f=1
TABLE 1. Vapour pressures (p*) of the pure compounds T K
278.15 288.15 298.15 303.15 308.15 313.15 323.15
CH3COC2 H5
CH3OH
C2 H5OH
This work
Lit.(13)
This work
Lit.(14)
This work
Lit.(14)
4.277 7.334 12.071 15.281 19.110 23.683 35.540
4.278 7.342 12.060 15.231 19.060 23.649 35.548
p*/kPa 5.488 9.903 16.960 21.904 27.991 35.450 55.566
5.512 9.876 16.945 21.868 27.962 35.444 55.567
2.246 4.314 7.867 10.531 13.823 17.980 29.419
2.248 4.312 7.875 10.462 13.756 17.905 29.494
569
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
TABLE 2. Experimental vapour pressures p, deviations dp=p−p(calc), activity coefficients fi , excess molar Gibbs free energy GmE , smoothing equations, and standard deviations s(p) x
p Pa
dp Pa
fA
fB
GmE J·mol−1
x
p Pa
dp Pa
fA
fB
GmE J·mol−1
0.0963 0.2051 0.2548 0.3303 0.4108
{(1−x)CH3COC2 H5+xCH3OH} at T=288.15 K 8727 11 2.1160 1.0097 194 0.5711 10938 13 1.1839 1.3395 9661 −13 1.7654 1.0425 359 0.6939 10982 6 1.0932 1.5367 9994 14 1.6436 1.0650 416 0.7437 10966 21 1.0657 1.6402 10320 −17 1.4919 1.1085 482 0.8652 10679 −12 1.0187 1.9762 10602 −12 1.3642 1.1684 525 0.9639 10163 −28 1.0014 2.3779 GmE /RT=x(1−x){0.9047+0.0001(1−2x)+0.0382(1−2x)2}; s(p)=16 Pa
531 463 417 258 78
0.0963 0.2050 0.2548 0.3303 0.4108
{(1−x)CH3COC2 H5+xCH 3OH} at T=298.15 K 14411 17 2.0596 1.0098 194 0.5711 18337 21 1.1815 1.3198 15999 −21 1.7242 1.0419 358 0.6939 18476 −6 1.0945 1.5061 16559 6 1.6095 1.0635 414 0.7437 18484 19 1.0674 1.6059 17177 −12 1.4682 1.1046 480 0.8652 18123 −21 1.0198 1.9402 17697 0 1.3499 1.1606 523 0.9639 17400 −8 1.0015 2.3599 GmE /RT=x(1−x){0.8719−0.0115(1−2x)+0.0582(1−2x)2}; s(p)=16 Pa
531 466 421 264 81
0.0963 0.2049 0.2547 0.3301 0.4107
{(1−x)CH3COC2 H5+xCH3OH} at T=303.15 K 18219 17 2.0120 1.0091 190 0.5711 23435 28 1.1807 1.3084 20290 −19 1.7028 1.0392 352 0.6939 23662 −5 1.0942 1.4921 21012 −2 1.5949 1.0597 409 0.7437 23671 5 1.0672 1.5907 21857 −6 1.4605 1.0989 474 0.8652 23291 −25 1.0197 1.9208 22551 1 1.3460 1.1529 519 0.9639 22453 9 1.0015 2.3338 GmE /RT=x(1−x){0.8534−0.0245(1−2x)+0.0519(1−2x)2}; s(p)=16 Pa
530 466 422 264 81
0.0962 0.2048 0.2546 0.3298 0.4105
{(1−x)CH3COC2 H5+xCH3OH} at T=308.15 K 22801 20 1.9770 1.0088 188 0.5710 29587 31 1.1767 1.2995 25446 −28 1.6817 1.0379 349 0.6939 29944 −2 1.0922 1.4780 26387 2 1.5780 1.0578 405 0.7437 29970 −5 1.0657 1.5736 27486 −4 1.4487 1.0958 470 0.8652 29590 −27 1.0193 1.8929 28399 3 1.3377 1.1483 515 0.9639 28638 18 1.0015 2.2907 GmE /RT=x(1−x){0.8335−0.0261(1−2x)+0.0499(1−2x)2}; s(p)=19 Pa
526 463 419 263 80
0.0961 0.2046 0.2545 0.3296 0.4104
{(1−x)CH3COC2 H5+xCH3OH} at T=313.15 K 28266 2 1.9440 1.0086 187 0.5709 37014 12 1.1739 1.2893 31656 −6 1.6594 1.0371 345 0.6938 37582 4 1.0915 1.4619 32836 8 1.5596 1.0564 401 0.7437 37646 −5 1.0654 1.5550 34241 −11 1.4357 1.0931 466 0.8652 37288 −25 1.0193 1.8681 35444 3 1.3290 1.1438 510 0.9639 36218 33 1.0015 2.2625 GmE /RT=x(1−x){0.8133−0.0299(1−2x)+0.0547(1−2x)2}; s(p)=16 Pa
522 461 417 263 81
0.0958 0.2042 0.2542 0.3287 0.4095
{(1−x)CH3COC2 H5+xCH3OH} at T=323.15 K 42483 54 1.8713 1.0073 179 0.5706 56693 150 1.1586 1.2780 47793 −15 1.6225 1.0330 335 0.6936 57538 −73 1.0804 1.4412 49607 −83 1.5305 1.0511 390 0.7436 57779 −38 1.0565 1.5260 51993 8 1.4140 1.0858 454 0.8653 57587 −24 1.0160 1.7939 53934 10 1.3105 1.1352 499 0.9639 56402 25 1.0012 2.1006 GmE /RT=x(1−x){0.7700−0.0168(1−2x)+0.0199(1−2x)2}; s(p)=67 Pa
509 445 401 248 75
570
R. Garriga et al. TABLE 2—continued
x
p Pa
dp Pa
0.0697 0.1515 0.1956 0.3029 0.4110
4400 4450 4453 4421 4349
−1 −6 −9 −5 8
fA
fB
GmE J·mol−1
p Pa
x
{(1−x)CH3COC2 H5+xC2 H5OH} at 2.5298 1.0060 163 0.5101 2.1280 1.0278 318 0.6648 1.9578 1.0460 388 0.7597 1.6401 1.1091 514 0.8168 1.4177 1.2024 583 0.9126
dp Pa
T=278.15 K 4244 19 3942 1 3650 −21 3454 2 2956 10
fA
fB
GmE J·mol−1
1.2720 1.1202 1.0609 1.0354 1.0081
1.3190 1.5803 1.8083 1.9801 2.3487
597 529 433 355 190
GmE /RT=x(1−x){1.0344+0.0321(1−2x)+0.0297(1−2x)2}; s(p)=11 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=288.15 K 0.0697 0.1515 0.1956 0.3029 0.4111
7575 7694 7743 7717 7581
9 −6 16 11 −12
2.3474 2.0223 1.8772 1.5941 1.3869
1.0050 1.0238 1.0399 1.0979 1.1861
154 303 371 494 563
0.5102 0.6649 0.7598 0.8169 0.9127
7394 6970 6517 6225 5325
−24 0 −12 52 −37
1.2487 1.1056 1.0516 1.0292 1.0064
1.2973 1.5423 1.7457 1.8914 2.1819
577 508 412 336 177
GmE /RT=x(1−x){0.9643+0.0367(1−2x)−0.0096(1−2x)2}; s(p)=25 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=298.15 K 0.0697 0.1515 0.1955 0.3029 0.4111
12532 12807 12888 12920 12767
−7 −12 −2 12 5
2.2565 1.9308 1.7939 1.5379 1.3580
1.0056 1.0251 1.0411 1.0954 1.1736
153 299 363 481 546
0.5101 0.6649 0.7600 0.8171 0.9128
12502 11867 11140 10687 9519
−3 33 −51 6 17
1.2385 1.1099 1.0572 1.0338 1.0080
1.2696 1.4844 1.6749 1.8210 2.1423
560 500 412 339 183
GmE /RT=x(1−x){0.9048+0.0150(1−2x)+0.0492(1−2x)2}; s(p)=22 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=308.15 K 0.0697 0.1514 0.1954 0.3029 0.4111
19941 20438 20628 20789 20676
23 −24 −4 2 7
2.1076 1.8382 1.7216 1.4976 1.3350
1.0048 1.0220 1.0361 1.0850 1.1564
145 283 345 459 524
0.5102 0.6650 0.7599 0.8170 0.9127
20361 19466 18510 17840 16140
4 17 −37 18 1
1.2244 1.1035 1.0538 1.0317 1.0075
1.2451 1.4441 1.6192 1.7524 2.0415
540 483 398 328 177
GmE /RT=x(1−x){0.8429+0.0006(1−2x)+0.0363(1−2x)2}; s(p)=19 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=313.15 K 0.0696 0.1514 0.1953 0.3028 0.4111
24765 25461 25725 25993 26038
22 −32 −15 −24 94
2.0561 1.8042 1.6943 1.4806 1.3237
1.0045 1.0209 1.0345 1.0817 1.1510
142 278 339 452 516
0.5102 0.6650 0.7599 0.8171 0.9128
25586 24585 23473 22682 20693
−28 5 −49 19 13
1.2166 1.0994 1.0514 1.0301 1.0070
1.2370 1.4293 1.5967 1.7231 1.9933
532 476 392 322 173
GmE /RT=x(1−x){0.8171+0.0012(1−2x)+0.0293(1−2x)2}; s(p)=41 Pa {(1−x)CH3COC2 H5+xC2 H5OH} at T=323.15 K 0.0696 0.1513 0.1951 0.3027 0.4110
37274 38582 38999 39714 39807
−23 −1 −34 61 66
1.9571 1.7220 1.6229 1.4347 1.2996
1.0046 1.0206 1.0334 1.0762 1.1369
137 268 325 431 492
0.5102 0.6650 0.7601 0.8172 0.9129
39333 38204 36854 35793 33184
−93 8 −16 15 7
1.2066 1.1009 1.0543 1.0327 1.0080
1.2112 1.3812 1.5378 1.6615 1.9435
GmE /RT=x(1−x){0.7586−0.0208(1−2x)+0.0640(1−2x)2}; s(p)=46 Pa
510 462 385 320 175
571
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
x being the mole fraction of alcohol in the liquid phase. The vapour pressure is then given by: p(calc)=xA fA p* A RA+xB fB p* B RB ,
(4)
using for nonideality of the vapour-phase the corrections: 2 RA=exp[{(V* A −BAA )(p−p* A )−pdAB yB }/RT ];
(5)
2 RB=exp[{(V* B −BBB )(p−p* B )−pdAB yA }/RT],
(6)
where yA and yB are the vapour-phase mole fractions of alcohol and butanone, respectively, and dAB is given by: dAB=2BAB−BAA−BBB ,
(7)
which in our case is taken to be equal to zero. Molar volumes of the pure compounds used in the Barker analysis were taken from the literature.(16,17) For a given composition, the sample temperature is changed and a slight variation of the true liquid mole fraction may be detected in table 2, according to the variable composition of the vapour phase. Analytic equations for GmE at the lower and higher temperatures are plotted in figure 1. The HmE and VmE experimental results for {(1−x)CH3COC2 H5+xCH3OH}, at T=298.15 K, are given in table 3, and are plotted in figure 2 together with TSmE at the same temperature, calculated from TSmE =(HmE −GmE ). The curves for HmE and VmE were found by fitting the experimental results with a polynomial: m
QmE =x(1−x)· s Af (1−2x) f ,
(8)
f=0
where QmE is either HmE /(J·mol−1 ) or VmE /(cm3·mol−1 ). We have tested the consistency of the molar enthalpies and excess molar Gibbs free energies by means of the Gibbs-Helmholtz equation. By assuming that Af varies linearly with temperature, the calculated values of HmE are shown as curves in figure 2, together with HmE experimental data, at T=298.15 K. The parameters were observed to be acceptably consistent and were more so for mixtures containing methanol. Azeotropic mixtures with a minimum boiling temperature were observed over the whole range of temperature. Azeotropic mole fractions z were calculated graphically, assuming ideal behaviour of the vapour, from the well-known equation fB /fA = p* A /p* B . Azeotropic compositions show a linear relation to temperature, according to the equation: (9)
z=a+b(T/K).
For {(1−x)CH3COC2 H5+xCH3OH}, a=−0.3032 and b=3.386·10 ; and for {(1−x)CH3COC2 H5+xC2 H5OH}, a=−0.9698 and b=4.136·10−3. Along the azeotropic line, assuming both ideal behaviour of the vapour phase and negligible volume of the liquid phase, the equation: −3
d ln pz /dT=Dvap Hm /RT 2,
(10)
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R. Garriga et al.
FIGURE 1. Excess molar Gibbs free energies GmE for {(1−x)CH3COC2 H5+xROH}. R, CH3 : curve 2, T=288.15 K; curve 3, T=323.15 K. R=C2 H5 : curve 1, T=278.15 K; curve 4, T=323.15 K. TABLE 3. Excess molar enthalpies and volumes (XmE ), smoothing equations, and standard deviations s(XmE ) for {(1−x)CH3COC2 H5+xCH3OH} at T=298.15 K x
0.1061 0.1788 0.2349
0.0193 0.0414 0.1267 0.1346
XmE
x
XmE
x
XmE
x
QmE =HmE /(J·mol−1 ) 0.2594 592 0.4842 686 0.7278 0.3670 662 0.5811 650 0.8316 0.4332 685 0.6816 532 0.9123 E E Hm=x(1−x){2716+577(1−2x)+106(1−2x)2}; s(Hm )=9 J·mol−1 QmE =VmE /(cm3·mol−1 ) −0.020 0.1535 −0.134 0.4832 −0.264 0.7771 −0.038 0.2427 −0.189 0.5935 −0.263 0.8705 −0.117 0.3482 −0.235 0.6439 −0.252 0.9570 −0.118 0.3968 −0.249 0.7027 −0.238 E E 2 Vm=x(1−x){−1.061+0.127(1−2x)−0.121(1−2x) }; s(Vm )=0.002 cm3·mol−1 299 460 545
XmE
485 341 190
−0.201 −0.140 −0.051
573
P, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
FIGURE 2. Excess molar functions, at T=298.15 K, for {(1−x)CH3COC2 H5+xROH}. (a) ——, T·SmE ; – – – –, HmE . R=CH3 : w, this work; W, reference 6. R=C2 H5 : q, reference 1; Q, reference 6. —·—·—·, calculated from Gibbs-Helmholtz equation. Arrows indicate T·SmE and HmE for the same systems. (b) VmE . R=CH3 : w, this work; W, reference 7. R=C2 H5 : q, reference 1; Q, reference 7.
is satisfied. If we accept that the enthalpy of azeotropic vaporization is constant, the azeotropic pressure is related to the temperature in a similar way to that for a pure substance: ln (pz /kPa)=A+B(T/K)−1 . (11) For {(1−x)CH3COC2 H5 + xCH3OH}, A = 17.74 and B = −4419; and for {(1−x)CH3COC2 H5+xC2 H5OH}, A = 17.23 and B = −4377. Azeotropic compositions and pressures, both experimental and calculated from equations (9) and (11), are compared in table 4 and plotted in figure 3. From equations (10) and (11), the molar enthalpy of azeotropic vaporization for the system containing TABLE 4. Azeotropic pressures and mole fractions {(1−x)CH3COC2 H5+xCH3OH}
a
T K
za
pz /Pa a
zb
pz /Pa c
278.15 288.15 298.15 303.15 308.15 313.15 323.15
0.670 0.709 0.725 0.740 0.757 0.790
10973 18478 23668 29968 37607 57727
0.673 0.706 0.723 0.740 0.757 0.791
11033 18453 23564 29852 37533 58086
{(1−x)CH3COC2 H5+xC2 H5OH} za
pz /Pa a
zb
pz /Pa c
0.185 0.221 0.259
4452 7737 12907
0.181 0.222 0.263
4474 7724 12856
0.301 0.326 0.371
20786 26002 39773
0.305 0.325 0.367
20700 25969 40022
Experimental value; b calculated from equation (9); c calculated from equation (11).
574
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FIGURE 3. (a) Azeotropic mole fractions z, and (b), azeotropic vapour pressures p2 , for {(1−x)CH3COC2 H5+xROH}. w, R=CH3 ; W, R=C2 H5 .
methanol is Dvap Hz=36.7 kJ·mol−1 , and with ethanol Dvap Hz=36.4 kJ·mol−1 . By using the azeotropic compositions together with the vaporization enthalpies for the pure compounds(18) butanone (Dvap H B* = 34.5 kJ·mol−1 ), methanol (Dvap H A*=37.4 kJ·mol−1 ), and ethanol (Dvap H A*=42.3 kJ·mol−1 ), the equation: Dvap Hz=(1−z)·Dvap H B*+z·Dvap H A* (12) is satisfied. Nagata et al.(19) measured vapour pressures for {(1−x)CH3COC2 H5 + CH3OH} at T=323.15 K, and from these data we have calculated GmE(x=0.5)= 558 J·mol−1 , z=0.783, and pz=58.82 kPa; at the same temperature, our experimental results are: GmE(x=0.5)=517 J·mol−1 , z=0.790, and pz=58.27 kPa. The HmE of this system has also been measured by Nagata et al.(6) at T=298.15 K; their values, at x=0.5, are higher than ours by about 4 per cent. The VmE measured by Letcher(7) compare satisfactorily with ours, figure 2(b). TABLE 5. Thermodynamic excess functions for {0.5 R OH+0.5CH3COC2 H5 or +0.5n-C6 H14 }, at T=298.15 K n-C6 H14 (23)
CH3COC2 H5
Alkanol
HmE J·mol−1
TSmE J·mol−1
VmE cm3·mol−1
HmE J·mol−1
TSmE J·mol−1
VmE cm3·mol−1
CH3OH C2 H5OH
679 1029 b
139 468
−0.265 −0.071 b
500 555
(−1160) a −850
0.41
a
Extrapolated value. b Reference 1.
p, VmE , and HmE of (CH3 COC2 H5+CH3 OH or C2 H5 OH)
575
For {(1−x)CH3COC2 H5 +C2 H5OH} at T = 298.15 K, Ohta et al.(20) obtained G (x=0.5) = 554 J·mol−1 , z = 0.261, and pz = 13.04 kPa; these values compare favourably with ours: GmE(x = 0.5) = 561 J·mol−1 , z = 0.259, and pz=12.91 kPa. For this system, Nagata et al.(6) found, at T = 298.15 K, HmE(x=0.5)=1071 J·mol−1 , while In˜arrea et al.(1) give, at the same temperature, HmE(x=0.5)=1029 J·mol−1 . The VmE data taken from the literature(1,7) show discrepancies as shown in figure 2(b). In mixtures of CH3OH or C2 H5OH with CH3COC2 H5 as a strong polar solvent (dipole moment p=1.1·10−29 C·m)(21) strong solvent–solvent and hydroxyl group–solvent interactions come into play and the HmE and SmE , at T=298.15 K, are much more positive than those for inert-solvent solutions owing to the interaction between the hydroxyl hydrogen and the oxygen of the carbonyl group. According to Barker,(22) in systems with specific interactions, more hydrogen bonds are broken at a given mole fraction. In table 5, experimental values of these thermodynamic properties are briefly summarized and compared with those for n-hexane solutions.(23) Presumably, the strong entropy increase is associated with a large gain in configurational freedom on breaking hydrogen bonds by the specific interaction. In figure 2(b), VmE are plotted. The excess volumes are the result of opposing effects: a positive one due to the break-up of the alcohol structure, and a negative one due to the interaction between the hydroxyl hydrogen and the oxygen of the carbonyl group. A third contribution, due to the destruction of the strong ketoxy-to-ketoxy interactions by the alkyl chain of the alcohol, is less important for short alkanol molecules. E m
R.G. wishes to thank the Diputacio´n General de Arago´n for the award of a predoctoral grant (BCB49/93). REFERENCES 1. In˜arrea, J.; Valero, J.; Pe´rez, P.; Gracia, M.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1988, 20, 193–199. 2. Garriga, R.; Putze, I.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 481–491. 3. Garriga, R.; Sa´nchez, F.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 887–895. 4. Garriga, R.; Sa´nchez, F.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1995, 27, 1057–1066. 5. Garriga, R.; Ilarraza, J.; Pe´rez, P.; Gracia, M. J. Chem. Thermodynamics 1996, 28, 230–243. 6. Nagata, I. Int. Data Ser., Selec. Data Mixtures, Ser. A 1984, 2, 81–82. 7. Letcher, T. M.; Nevines, J. J. Chem. Eng. Data 1995, 40, 293–295. 8. Marsh, K. N. Trans. Faraday Soc. 1968, 64, 883–893. 9. Gracia, M.; Sa´nchez, F.; Pe´rez, P.; Valero, J.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1992, 24, 463–471. 10. Pardo, J.; Pe´rez, P.; Royo, F.; Gracia, M.; Gutie´rrez Losa, C. J. Chem. Thermodynamics 1987, 19, 521–526. 11. Gutie´rrez Losa, C.; Gracia, M. Rev. Acad. Cienc. Exactas Fis.-Quim. Natur. Zaragoza XXVI 1971, 26, 1: 101–135. 12. Dymond, J. H.; Smith, E. B. The Virial Coefficients of Pure Gases and Mixtures. Clarendon Press: Oxford. 1980, pp. 42, 80 and 110. 13. Ambrose, D.; Ghiassee, N. B. J. Chem. Thermodynamics 1987, 19, 505–519. 14. Ambrose, D.; Ellender, J. H.; Sprake, C. H. S. J. Chem. Thermodynamics 1974, 6, 909–914. 15. Barker, J. A. Aust. J. Chem. 1953, 6, 207–210. 16. TRC Thermodynamic Tables. Non-Hydrocarbons. Thermodynamics Research Center, The Texas A&M University System: College Station, TX. 1991, p. d-5870. 17. TRC Thermodynamic Tables. Non-Hydrocarbons. Thermodynamics Research Center, The Texas A&M University System: College Station, TX. 1966, p. d-5030.
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18. Riddick, J. A.; Bunger, W. B.; Sakano, T. K. Organic Solvents. Physical Properties and Methods of Purification. Wiley: Chichester. 1986, pp. 190, 192 and 338. 19. Gmehling, J.; Onken, U. Vapor-Liquid Equilibrium Data Collection. DECHEMA Chemistry Series, Vol. I, Parts 2a. DECHEMA: Frankfurt. 1986, p. 138. 20. Ohta, T.; Koyabu, J.; Nagata, I. Fluid Phase Equilib. 1981, 7, 65–73. 21. McClellan, A. L. Tables of Experimental Dipole Moments. Freeman: San Francisco. 1963. 22. Barker, J. A. J. Chem. Phys. 1952, 20, 1526–1532. 23. Brown, I.; Fock, W.; Smith, F. Aust. J. Chem. 1969, 1, 273–291.
(Received 29 November 1995; in final form 5 January 1996)
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