--0
i3 3
J. F.MATHEWASD J. J. ~VICKETTA
Vol. 65
THE THERhlODYX;A&lIC PROPERTIES OF TZ-PROPYLALCOHOL B Y J. F.;\IATHEWS AND J. J. MCKETTA Department of Chemical Engineering, University of Texas, Austin, Texas Recezzed September I , 1,960
Vapor heat capacities of n-propyl alcohol have been measured from 371.2 to 451.2"K. at selected pressures from 1/3 to 5/3 atmospheres. Latent heats of vaporization have also been measured from 343.9 to 384.5"K. The model of an equilibrium formation of dimers and tetramers has been utilized to represent the behavior of the vapor heat capacity as a function of pressure and ideal heat capacities have been determined. These values of CO , have been used, along with molecular structure and spectroscopic data and barrier assignments for the methyl and hydroxyl rotating tops, to determine the height of the transbarrier for the ethyl group and the energy difference between rotational tautomers. Thermodynamic properties in the ideal gaseous state have been calculated a t selected temperatures from 0 to 1000°K.
In order to make thermodynamic calculations for Rossini, et a1.8 ,4 molecular weight of 60.094 was used for n-propyl alcohol and 0°C. = 273.16"K. All reported industrial reactions it is helpful to know the ther- temperatures are on the defined International Temperature modynamic properties of the compounds involved. Scale. Yormal propyl alcohol is both industrially useful Vapor Pressures.-Vapor pressures of n-propyl alcohol and theoretically interesting. This compound be- were measured a t eight temperatures from 70 t o 112". vapor pressures were fitted by the method of least longs to the c l a s of hydrogen-bonding substances These squares to an Antoine equation and, as such, would be expected to exhibit highly 1277.2 non-ideal behavior. It is also the first member of log p = 7.4572 (1) 181.89 +t the alcohol series in which rotational isomerism might be expected to affect significantly the ther- where p is in mm. and t is in "C. The maximum deviation the experimental vapor pressures from equation 1is 0.26%, modynamic properties. The investigation of the of and the average deviation is 0.19%. The accuracy uncerlatter two pheiiomeiia is helpful in the formulation tainty of these data is estimated as A 2 mm. The normal of a hetter understanding of the thermodynamic boiling point predicted by equation 1 is 97.20°, which compares quite well with the value of 97.209' reported by beha\-ior.of fluids. ciechowski.9 Tabulations of the thermodynamic functions of \Toj An Antoine equation for n-propyl alcohol has been pren-propyl alcohol have been presented in the litera- sented by the -4.P.I. Project 4410 as t ~ r e l -but, ~ in all cases, the estimations have been 1375.14 l o g p = 7.61924 - ~(la) b w d on insufficient experimental data or on no 193.00 + t data n-hatsoever. where the units are the same as in (1). Vapor pressures 'the purpo$e of this n-ork was to investigate the calculated from equation l a agree with the data taken in behavior of the vapor heat capacities of n-propyl this work, within the accuracy of the latter. In the followalcohol as a function of temperature and pressure: ing sections, where it is desired to calculate a vapor prespoint outside of the range of equation 1, equation l a is then to determine from this beha.i-ior the ideal heat sure employed. capacities, and to calculate statistically the thermoT.4BLE 1 dynamic. fiinctions in the ideal gavous state.
Experimental Apparatus.-The flow calorimeter used in this work was orivinally constructed by Pennington4 and was of the same de& as the apparatus described by Waddington, et a1.,6 and McCullough, et a1.6 Material.-The sample of n-propyl alcohol used in the measurements was obtained from the Celanese Corporation of America. The analysis (infrared) was 98.49% n-propyl alcohol, 1.51y0 sec-butyl alcohol and negligible amounts of water and other impurities. This original material was distilled twice through a 1.2 meter column packed with glass rings. The final distillate had a boiling range of less than 0.02'. This distillate was then vacuum distilled iblilb-to. bulb) into receivers for introduction into the apparatus. KO significant changes were noted in the physical properties of this compound during the course of the exprrimental work. Physical Constants.-Calculations in this work were based on the 1951 International Atomic Weights' and on the values of the fundamental and derived constants given by (1) H. A. G. Chermin, Petroleum Refiner (to be published). 12) 11.E. Dyatkina, Zhur. F i z . K h i m . , 28, 377 (1954). ( 3 ) K. A . Kobe, R. H . Harrison and R. E. Pennington, Petroleum
Refiner, 3 0 , N o . 8 , 119 (1951). (4) R. E. Pennington and K . 4.Kobe, J . A m . Chem. Soc., 79, 301 (1957). ( 3 ) (a) G . Waddington. 8 . S. Todd and H. M. Huffman, ibid., 69, 22 (1947); ib) G. Waddington and D. R. Douslin, i b i d . , 69, 2275 (1947). ( 6 ) J. P. McCullough, D. W. Scott, R. E. Pennington, I. A . Hossenlopp and G. Waddington, i b i d . , 76, 4791 (1954). (7) E. Wichers, ibid.. 74, 24.47 (1952).
AVERAGEVALUESOF LATENT HEATOF VAPORIZATIOX Temp., OK.
a
Pressure,
A H v , cal./mole
343 9 252.2 10487 i 3 360.0 505 2 10121 i 10 370 3 758 0 9852 i 5 1010 4 9637 i 3 378 1 384 5 1262 2 9478 =k 0 I-apor pressures were calculated from equation 1.
Latent Heat of Vaporization.-Latent heats of vaporization of n-propyl alcohol have been measured a t pressures atmospheres, a t saturation temof l j 3 , 2 / 3 , 3 1 3 , 4 i 3 and peratures. The average values of a t leaEt two measurements at each set of conditions are presented in Table I, along with the maximum deviations of these measuremente from the averages. I t is estimated that the accuracy uncertaintJ- of these values is not greater than 0.1%. The measured latent heats may be represented within this accuracy uncertainty by the equation A H v = 9707 26.75T 0.07117T2 (343.9-384.5"K.)
+
-
where AH,, is in calories per mole and T is in OK. Some values of the latent heat of vaporization have been reported in the literature. For example, the value a t the normal boiling point of 9852 cal./mole found in this work (8) F. D. Rossini, F. T. Gucker, H. L. Johnston, L. Pauling and G. st'. Vinal, ibid., 74, 2699 (1952). (9) 51. Wojciechowski, J. Research Natl. Bur. Standards, 17, 721 (1936). (10) A.P.I. Research Project 44, "Selected Values of Properties of Hydrocarbons a n d Related Compounds," Table 203k, p . 1, 1956.
may be compared with 10012 cal./mole reported by Plewes, et U Z . , ~ ' and 9879 cal./mole given in the International Critical Tables.'* Vapor Heat Capacities.-Vapor heat capacities of npropyl alcohol have been measured a t selected pressures horn 1/3 to 5 / 3 atmospheres and a t temperatures from 3 71.2 to 451.2'K. These experimental values are presented in Table I1 and plotted as a function of pressure in F i g . 1. The estimated accuracy uncertainty of thesedata is 1 0 . 2 yo.
VAPOR
29.0
I -
28.0
TABLE I1 HEATCAPACITY I N CALORIES PER hIOLE DEGREE
Temp., OK. C, (5/3 atm. C , (4/3 atm. C, (3/3 atm. C, ( 2 / 3 atm. C , (1/3 atm.
371.2 abs.) abs.) abs.) abs.) Rbs.)
26.44 25.16 1.833 2.664 24.44
391.2 28.19 27.15 26.40 26.91 1.304 0.417 25.44 25.42
411.2 28.35 27.83 27.43 27.02
1
26.0 -
I 25.0
a Paranietw in equation 7 , [(cal./mole "K.)/atm.]. Parameter in eqmtion 7 , [(cal./mole "K.)/atrn.l].
24.4
Cb Cpo (exptl.) cpo(eq. 15)
2!4.45
0.71.5 0.017 27.33 27.36
451.2 29.27 29.09 28.90
0.547 0 004 28.35 28.34
245
0.953 0.078 26.39 26.38
431.2 28.61 28.34 28.04
Direct vapor heat capacity measurements have been made by Sinlre and DeVrie@ from 373 to 437°K. a t a pressure of 750 mm. with an estimated accuracy uncertainty of &0.2 cal./(inole)( "K.). Bennevit'z and Rossnerl* have report,ed a da.tum at 410%. and a pressure of 748 mm., and Jatkar and Lakshimarayan'5 have deduced the vapor heat capacity at 407OK. from measurements of sonic velocity. For ease of compariijon, these reported values are presented along with those of the present work in Fig. 2 , after being corrected to ideal heat capacities by the gas imperfection equations which will be developed in the next section. Jatksr's value was not corrected, as it is presumed that a value obtained from the velocity of sound would be an "ideal" heat capacit'y.
Correction for Gas Imperfection In order to determine the ideal heat capacities of a compound the observed heat capacities must be corrected t'o z8ro pressure by a consideration of the effects of gas imperfection. This may be done by using the Tvell-known thermodynamic relationship
if a proper representation of [b2V,bT2]pis known. 1:or most compounds which have been studied the vapor heat capacity is nearly a linear function of temperature and can be represented by C,, = CPo
+ Ap f
ACp
(3a)
where A is a function of the second virial coefficient and ACp is a small correction term which may be det'ermined by a t'rial and error procedure. However, as Fig. l shows, n-propyl alcohol does not belong to this class of compounds; and another method must be used to evaluate t'he effects of gap imperfection. Gas Imperfection as an Association ProblemIrit,eractione between two or more molecules in the gas phase may be looked upon as a chemical reac( 11 ) A. C. Pleaes. D. A. Jardine and R . 54. Butler. Can. J . Teclinnl., 32, 133 (1954). (12) "Internarional Critical Tables," T'ol. T'. McGrarv-Hill Book To.. lnc., New York. A-. I-,, 1929. 113) C . C . Sinke and T . DeVries. J . Am. Ciiern. Soc., 76, 1815 (19531. (14) K. Bennawitz and W. Rossner, 2. p h y s i k . Chem., B23, 126 (19x8). (15) S. Jatkar and D. Lakshimarayan, C. A., 41, 19010 (1947); J . I n d i a n Inst. S c i . . 28A. 1 (1946).
P'
eo E X P E R I M E N T A L -EQUATION
/'
L
,d
T= 37'
I
(7)
OK
I
L
1
I
I
0 Fig. 1.-Vapor
29.0
1/3 2/3 3/3 413 S / S Pressure, atmospheres. heat capacity as a function of pressure.
F----
* I 28.0
&
r
,,.0/ *
26.0
B
2 25.0
c
-
24.0 350
P
v
J
C
--
JATKAR AND LAKSWARAYAN BENNEWITZ AND ROSSNER
0 THIS WORK
- EOUATION 1151
375 400 425 450 Temperature, "K. Fig. 2.-Ideal vapor heat capacity.
tion between molecules and treated from an equilibrium standpoint. The molecules are considered to form dimers, trimers, tetramerst etc., by association. This approach has been discussed by several worker^^^-^^ and has been used primarily t o describe the behavior of the experimental second virial coefficients of hydrocarbons and other more or less ideal substances. Weltner and I'it.zerZ0measured the vapor heat capacit,ies of methyl alcohol and used this concept of associat'ion to explain the large deviations from ideality which were encountered with tha,t' substanee. These workers found that an assumption of dimer and t'etramer formation best described the variat,ion of heat capacity wit'h pressure. Based on this assumption of an equilibrium between monomers and tetramers and dimers, t'he equation of state of methyl alcohol gas becomes (16) 0. R. Foz and ,I. M. Tidal, A n d e s 3 s . y quim. ( M a d r i d ) , 48, 842 (1947). (17) ,J. 0 . Hirahfelder, F. T. McClure and I. I'. Weeks, J . Chem. Phye., IO, 201 (1942). (18) J. D. Lambert, G. A . H. Roberts, J. S. Rowlinson and V Wilkinson, Proc. Roy. SOC.( L o n d o n ) , A196, 113 (1949). (19) H. N', Woolley, J . Chem. Phys., 21, 236 (1952). (20) W. Weltner and K. S. Pitzer, J. Am. Chem. Soc., 73, 2606 ( 1951).
J. F. MATHEWS AND J. J. MCKETTA
760 RT
V=-+B++p' P
(4)
' = [ 2*506T
- RT e
(5)
' = [0'54 Txz10-9]
in which B =b
-
~ pA&[ ~exp[$] ]
Vol. 65
loa] exp
[';-'I
[(cal./mole 0 ~ ) l a t m . l
[
exp l2?] - [(cal./mole OK.)/atm.a] (14)
and
and AH2 = 3400 cal./mole A& = 15.4 e.u.
AHc = 25176 cal./mole AS4 = 75.4 e.u.
where b is the covolume and ASZ,AHZ and AS4,AHq The representation of the experimental heat are the entropy and enthalpy changes of dimeriaa- capacities by the above equations is quite reasontion and tetramerization, respect'ively. Substitu- able, with the maximum deviation being 0.12% tion of equations 4,5 and 6 into equation 3 gives and the average deviation 0.0670. Curves calcuC, = C O, + Ap Cp3 lated from equation 7, using equations 13 and 14, (7) are shown in Fig. 1, along with the experimental where points. The extrapolated values of the ideal heat (Cpo)are presented in Table I1 and plotted capacity (8) A = e x p [ T ] exp[i;?] versus temperature in Fig. 2. The observed CPo and may be represented by the equation
+
[$I
c
=
[iir:;] AH exp[?]
exp[g]
2
(9)
The constants in the above equations were determined by simultaneously fitting equation 4 to some available P-V-T data and equation 7 to the experimental vapor heat capacities. Barrowz1 has used the same approach on ethyl alcohol. Kretschmer and WiebeZ2have since made an extensive study of the P-V-T properties of methyl alcohol, ethyl alcohol and isopropyl alcohol and have determined that equation 4 does indeed represent the observed gaseous behavior very well and is much better than an equation of the form RT P
V=-+B++p
(10)
for correlating vapor heat capacity. n-Propyl Alcohol.-Essentially the same method as described above has been used in this work to determine the effects of gas imperfection on vapor heat capacity. The method of evaluation of the constants in equations 4 through 9 depends on the type of P-V-T data available. Foz, et have made some P-V-T measurements on n-propyl alcohol and present some values of B which they obtained by fitting their data to an equation of state of the same form as equation 4 (following the suggestion of Weltner and PitzerZ0). For n-propyl alcohol the constants in equation 5 were adjusted to give the best fit to the B values of FOB, et a1.,23and then using these constants in equation 8, the constants of equation 9 were determined to give the best representation of all the heat capacity data. The equations thus derived are B = 130
D
=
-4.168
- 0.0353T exp X
(3 --
(cc./mole atm.) (11)
T exp (Gg3) (cc./mole atm.8) (12)
(21) G. M. Barrow, J . Chsm. Phys., 20, 1739 (1952). (22) C. B. Kretschmer and R. Wiebe, J . Am. Cham. Soc., 76. 2579 (1954).
(23) (a) 0 . R. Foz,J. Moroillo and A. Mendes, Anales Teal 8oc. espan. fis. guim. (Madrid), SOB, 17 (1954); (b) 0.R. Foz, J. Moroillo. A. P. Ma& a s d A. Mender, {bid., SOB, 23 (1954).
CO ,
=
7.365
+ 4.400 X
+
5.507 X 10-6T2 (cal./ mole OK.) (371.2-451.2'K.) (15)
within =kO.l%. Total Gas Imperfection.-From equations 4, 11 and 12 the total gas imperfection (V - R T / p ) may be caIcuIated for any temperature and pressure by the equation V
-
P
= 130
- 0.0353Texp(":') 7
- 4.168
X
lo-" T e x p ( y ) p 2 (cc./mole) (16)
The saturation curve calculated from equation 16 is presented in Fig 3. Values of the saturated vapor volume ( V ) may be calculated from the Clapeyron equation as
with the different quantities in consistent units. The latent heats of vaporization (AH,) have been taken from the experimental values of this work, values of ( d T / b p ) v have been determined from equation 1, and the liquid molar volumes (V,) have been taken from Young's work on orthobaric densitiesaZ4 Values of the total gas imperfection thus calculated are presented in Fig. 3 also. Finally, direct measurements of total gas imperfection have been taken from YouiigZ4and plotted in Fig. 3. It must be emphasized that these calculated values of total gas imperiection are all for saturation conditions since ( T i - 12?'/p) is a function of pressure as well as temperature. The agreement among the values, calculated by the three different methods described above, is very satisfactory (Fig. 3). This is evidence that the procedure used in this work for the estimation of gas imperfection is a reasonable one. Calculation of the Thermodynamic Functions I n this work: principal moments of inertia and reduced moments of inertia of the rotating tops were calculated from molecular structure; frequencies from published infrared and Raman spectra were assigned to the fundamental vibrations of n-propyl alcohol ; potential barrier func(24) 5. Young, Sci. Proc. Rog. D u b h SOC.,12, 374 (1910).
May, 1961
THERMODYNAMIC PROPERTIES OF PROPYL ALCOHOL
tions were assumed for two of the rotating tops, and the parameters of the potential barrier of the third top were rsdjusted to give the best representation of the ex:perimental ideal vapor heat capacities. Rotational Tiautomerism.-Rotation about the central carbon-to-carbon bond in n-propyl alcohol results in one trans and two skew rotational tautomers. These tautomers would be expected to have different inomeints of inertia and different normal vibrational frequencies. P i t ~ e has r ~ ~shown that the t’hermodynamic properties are approximately the same for each tautomer, and thus it is sufficient to make t,he calculations for only one form. According to BertlnelotZ6and Golik, et aZ.,27the most stable configuration of the n-propyl alcohol molecule is the planar, or trans form. Calculations in this work are based on the trans form. Moments of Inertia of the n-Propyl Alcohol Molecule:-The product of the principal moments of inertia, and the reduced moments of the rotating tops were calculated using the formalized procedure of Kilpatrick an’dPitzer.28 Since a direct determination of bond distances and bond angleis in the n-propyl alcohol molecule has not been published, these quantities were estimated from those reported for methyl alcohol, ethyl alcohol and isopropyl a 1 ~ 0 h o 1 . ~The ~ values used were Bond 1eng;ths GH 1.09 A. C-C 1.54 A. G O 1.43 k. 0-H 0.96 A.
Bond angles
C-0-H
110’
all others
Tetrahedral
76 1
-“--I
I
0
EQUATION (17) EOUATION I161
500c I
340
I
360
Fig. 3.-Total
I
I
380 TEMPEd??JRE,%
460
440
uzo
gas imperfection for the saturation curve.
hydroxyl rotating tops, and the barrier heights were selected by an inspection of the assignments for these barriers in similar compounds -Barrier CHr
CzHaSHS I-C3HSH* CzH60HZ1 1-C3H70H(this work)
3310 3100 3300 3100
heights, cal./mole-
-SH
-OH
1640 1650
800 800
TABLE I11 THEVIBRATIONAL ASSIGNMENT Fundamental, om. - 1
Deaignation
Fundamental, om. - 1
Designation
463 730 760 860 890 971 1052 1066 1103 1220
C-C-0 bend CHI rock C-C-C bend C-C stretch C-C stretch CHs rock CHI rock C-0 stretch CH1rock CH2 twist
1272 1299 1341 1381 1393 1450 [a] 1463 1478 2940 [ 7 ] 3680
CHI wag CH2 wag CHs twist CHs sym. bend C-0-H bend 2 CH3 bend and 2 CH2 bend C-H stretches 0-H stretch
The calculated value of the product of the principal moments of inertia of the n-propyl alcohol g.3 cm.6; and the remolecule is 1695 X duced moment’s of inertia are 1.26, 4.59 and 16.34 >< lo4”)for the -OH, CH3- and C2Hs(g. tops, respective1,y. Fundamental ‘VibrationalFrequencies.-Inf rared potential barrier hindering the ethyl group and Raman spectra reported in the l i t e r a t ~ r e ~ O - was ~ ~The assumed to have the same form as the barrier have been used to make the vibrational assign- used for 1-propanethiol by Pennington, et aLa6 ment as shown in Table 111. Different designa- This barrier has unequal minima and is described tions of the fundamental vibrational modes could by t1f-o parameters: V O ,the height of the trans of course be made, but the calculation of the barrier and AEt, the energy difference between the thermodynamic properties depends only on the trans and skew rotational tautomers. The contrimagnitude of th’e vibrations and is not affected by bution of the rotational tautomerism to the incorrect designa,t,ion. thermodynamic properties was calculated by the Internal Rotation.-Simple threefold cosine hin- use of a model of an equilibrium between rotational dering potentials were assumed for the methyl and tautomers.37 V O and A E t were determined by (25) K. S. I’itaer, J . Chem. Phys., 14, 239 (1946). fitting the calculated ideal vapor heat capacities to (26) C. Berthelot, Compt. rend., 231, 1481 (1950). the experimentally determined values. The values (27) A. Z. Golik, A. F. Skryshevskii a n d S. D. Ravikovich, C. A., 60, selected were 2310 cal./mole for Vo and 850 cal./ 7534d (1956) [Dopouidi Aknd. N a u k , Ukr. E. S. R . , 336 (1954)l. mole for A&. (28) J. E. ICilprrtrick and K. S. Pitaer. J . Chem. Phys., 17, 1064 (1949). From measurements on the variation of the in(29) “Table:i of Interatomic Distances a n d Configuration in Moletensities of Raman lines as a function of temperacules and Ions.” Special Publication No. 11, T h e Chemical Society, ture, BerthelotZ6has deduced the energy difference London, 1958. (30) A.P.I. Xesearcti Project 44, C a t d o g of Infrared Speotral D a t a , between the tautomers of liquid n-propyl alcohol Serial No. 427. (31) W. Braun, D. Spooner and M. Feneke, Anal. Chem., 22, 1074 (1950). (32) K. W. F. Kolilrausch, “Ramanspektren,” Akademische Verlagesellschaft Becker a n d Erler, Kom.-Ges., Leipsig, 1943. (33) E. K. F’lyler, J . Resmrch Natl. Bur. Standards, 48, 281 (1952). (34) J. R. Quinan and S. E. Wiberley. Anal. Chcm., 26, 1762 (1054).
(35) J. P. McCullough, D. W. Scott, H . L. Finke, I f . E. Gross, K. D. Williamson, R. E. Pennington, G. Waddington and H. M. Huffman, J . A m . Chem. Soc., 74, 2801 (1952). (36) R. E. Pennington, D . W. Scott, H. L. Finke, J. P. McCullough, J. F. Messerly, I. A. Hossenlopp and G. Waddington, ibid., 78, 3266 (1956). (37) K. S. Pitaer, J . Cham. Phue., 6, 473 (1937).
J. F. MATHEWS ASD J. J. MCKETTA
762 IJIQUID
TABLE IV ENTROPY I T 298.16"B. IN DEGREE
CALORIES PER
by Parks, et u Z . , ~ O was not used to select the internal parameters. The accuracy of this datum as given by these authors is ". . . probably accurate to within 1 or 2% . . . . " A value of entropy as calculated in this work is compared with that of Parks, et al., in Table IV. The agreement is within the accuracy uncertainty of the latter. Thermodynamic Functions.-The Cpo,So,(HTO Hoo)/T,and (FTO - H,O)/T functions of normal propyl alcohol in the ideal gaseous state were calculated at selected temperatures from O-IOOOOK.
MOLE rotational
So(calcd.) R In (PY (;as iinpt>rfwtioil"
77.63 7.18 - 0.04 AS"' -38.09 S (liquid) 46 68 S (liquid, Third Law)d 46.1 0 . 7 a Vapor pressure a t 298.16'K. calculated from equation l a . * Calulated from equations 11 and 12. Latent heat of vaporization calculated from eqiiation 2. Parks, et d.40
TABLE V THERMODYNAMIC PROPERTIES IN THE Temp., OK.
CPQ.
cal./mole OK.
SO, cal./mole OK.
Vol. 651
IDEAL G . 4 s E o u s STATE'
- Hoe)/
( H T ~- HoQ)/ -(FTQ
T,
oal./mole OK.
T.
cal./mole OK.
- HfQ,
kcal./mole
AFfD, kcal./mole
log KfO
-56.35 Infinite 56.35 0 0 0 0 0 32.99 61.51 -41.24 273.16 19.67 75.86 13.95 61.91 -39.32 28.82 61.92 298.16 20.82 i7.63 63.11 14.48 -39.18 28.54 61.95 i7.75 63.19 300 20.91 14.52 -31.29 17.10 63.49 400 25.86 84.31 16.74 67.57 -23.07 10.08 500 30.51 64.78 90.54 19.04 71.50 5.32 65.82 -14.61 600 34.56 96.44 21.29 75.15 - 6.00 1.87 38.03 102.02 23.43 66.67 700 78.58 800 41.04 107.29 25.46 81.83 67.31 2.72 -0.74 900 43,65 112.27 27.32 84.94 67.80 +11.51 -2.80 1000 45.93 116.98 29.08 87,91 68.11 +20.36 -4.45 To retain internal consistencv some of the values in this table are given t'o more decimal places than are justified by their absolute accuracy.
+
5
to be 820 f 120 cal.,'mole (mean value from -80 t o +30°). The restricted rotator contributions to the thermodynamic properties vere taken from the tables of Pitzer and G ~ i n and n ~ from ~ the extension of these tables by Li and P i t ~ e r . ~ ~ Comparisons with Experimental Data.-With the information about rotation, vibration, internal rotation, etc., as discussed in the preceding sections, ideal gaseous heat capacities were calculated at temperatures corresponding to the experimental runs. These calculations (and all other calculations of thermodynamic properties in this work) are based on the rigid rotator, harmonic oscillator approximation, with interactions between rotation and vibration and among internal rotations neglected. It is desirable to include a term for the effects of anharmonicity on the thermodynamic properties, but in the present case there are insuffirient data for this purpose. A comparison betveen the calculated and experimental heat capacities in the ideal gaseous qtate is shown in Table 11. The maximum deviation is o.14yO and the average deviation is 0.0670. The calculated CPofit the smoothed values of equation 15 somewhat better, with the maximum deriation being 0.08% and the average 0.06Vc7,. The value of entropy of n-propyl alcohol given (38) K. S.Pitser and W. D G u i n n , J . C h e m . Phys., 10, 428 (1942). (39) J. C. hI LI and K. S. Pitzer, Jr , J . Phzjs Chem.. 60, 466 'lc)56)
and are presented in Table V. The standard heat, standard free energy, and common logarithm of the equilibrium constant for the formation of n-propyl alcohol by the reaction 3C (graphite)
+ 4H2 (gas) + 21 0, (gas)
=
1 - C3H?OH(gas) (18)
have been computed at selected temperatures from 0 to 1000OK. and are presented in Table V. These functions were calculated by use of the heat of combustion of liquid n-propyl alcohol a t 298.16'K. given by R ~ s s i n i(482.15 ~~ f 0.24 kral./mole), the thermodynamic properties of carbon (graphite), hydrogen (gas) and oxygen (gas) given by IJ'agman, et ~ 1and. the ~ energy ~ values AHP[H?O(gas)]2g816%. = -57797.9 cal./mole42 AHfO[CO, (gas)]pss1 6 0 ~ .= -95041.8 ~ a l . / r n o l e ~ ~ AHv[H20]2~8 16% = 10519.5 cal./mole42 Afl,[CaH,0H]29s 16% = 11356 cal /mole (eq. 2)
Acknowledgments.-This project was supported by a grant from the Esso Research and Engineering Company. The authors are grateful for this assistance. Mr. Mathews was the Humble Oil and Refining Fellow for one year of this work. (40) G S. Parks, I< I< Kelle, arid H. Soc , 61, 1969 (1929).
M. Huffman, J Am. Chem
(41) F. D Rossini, J . Research Natl. Bur. Standards, lS, 189 (1934) (42) D D. Wagmsn, J. E Kilpatnok, W. J. Taylor, K. S.Pitzer a n d F D. Rossini, tbzd , 34, 143 (1945).
