f-reeLnergy or Lthylene Hydration E. R. GILLILAND, R. C. GUNNESS, AND v. 0.BOWLES' Massachusetts Institute of Technology, Cambridge, Mass.
ably accurate only to ~ 2 0 0 0calories. Francis and Kleinschmidt (3) give an equation for hydration with liquid water, which, combined with an equation for the free energy of vaporization of water based on Keenan's tables (6),gives AF'
Within an accuracy of *500 calories, the free energy change for the hydration of ethylene,
+ HiO(g)
+
+ 0.00235T + 7.05
log,, K p = 'OgO - 4.71 loglOT T
CzHaOH(g)
is given by the expression, AF' = 26.93 - 8300
Since their derivation used existing thermal data and the specific heats of Parks and Huffman, their formula offers no check on existing data except a t the temperature they employed. Bowles (I), as a result of a series of batch liquidphase hydration experiments a t temperatures from 225" to 32.5' C. and a pressure of 400 atmospheres, using dilute mineral acids as catalysts, recommends the expression
This equation is based on all available experimental data and is applicable over a range of temperature from 150" to 380" C. The newer heat of combustion of ethanol offered by Rossini is probably more accurate than the older value of 328,700 calories. The specific heat of ethylene used by Francis and KleinSchmidt is not as accurate as that used by Parks and Huffman.
AF'
Other published data, including those of Swann, Snow, and Keyes (15) and Kleaver and Glaser (6), have not yielded results suitable for free-energy determinations.
Apparatus The copper-lined steel reactor, L (Figure l), was 14.25 inches
(36.2 cm.) long, with a capacity of 346 cc. The contents were agitated by three disks operated by a solenoid, and the reactor was maintained at constant temperature by a vapor bath. There was a bottom opening for withdrawal of liquid-phase samples, one 5.5 inches (14 cm.) from the top for vapor-phase samples, and one at the top for maintenance of constancy of pressure during sampling. Vapor-phase sampling was accomplished through an electrically heated line and valve, I , preventing condensation prior to pressure reduction. Four thermocoupleswere provided to insure maintenance of the necessary temperature. A glass condenser, J, cooled by ice water in a Dewar flask, K , was used to remove ethanol and water vapor; gas-measuring bottles, Q , collected the
-4350 f 26.1T
for the gas-phase hydration of ethylene. Because of a more recent determination of the heat of combustion of alcohol ( I g ) , they suggest that the equation might better be AF'
= -6350
f 26.1T
Because of inaccuracies in the data on heat of combustion of ethylene, their recommended free-energy equation is prob1
= -10,120 f 30.2T
His calculations made use of hydrocarbon vapor fugacities. Stanley, Youell, and Dymock (1.4) have published the results of a series of experiments in which they studied the vaporphase hydration reaction, approaching equilibrium from both sides a t atmospheric pressure and temperatures from 145' C. to 250' C., recommending
HERE is a considerable discrepancy between the free-energy change of hydration of ethylene as predicted from existing thermodynamic data and that calculated from the two existing experimental determinations. Using available data on free energy of formation, Parks and Huffman (IO) compute the expression, AF'
f 39.2T
This equation indicates a greater free-energy temperature coefficient than that of Parks and Huffman, chiefly because of the difference in the specific heat of ethylene gas. On the basis of an experimental determination of the vaporphase equilibrium of the hydration of ethylene using an alumina catalyst at 69.6 atmospheres and 380" C., Saunders and Dodge ( I S ) recommend
FIGURE 1. DIAQRAM OF APPARATUS
C2Hdg)
= -12,140
Present address, Humble Oil and Refining Company, Houston, Texas.
370
INDUSTRIAL AND ENGINEERING CHEMISTRY
.
.
-
~
~
~
371
~
removed and &&;red in bottles, C. Pr&ure waa maintained in the reactor by connecting it to an 11-liter high-pressure storage vessel, B, in which the ethylene gas was stored. A Bourdon gage, R. oalibrsted aeainst t~ "dead-weirht" ea=. WBS used for nressw miasurements. The ethylene used was the medical grade and analyzed better than 99.5 per cent ethylene.
-
~
Experimental Procedure Seven runs at a constant t.emperature and four different pressures were made, and 8. second series at a constant pressnre but four different temperatures varying from 176" to 307" C. The catalyst wm sulfuric acid averaging about 3 mole per cent. In each run liquid-phase samples were withdrawn at interval8 (a.pproximatelyhalf-hour) and analyzed for ethanol. When the concentration became eonstant, a vapor sample was removed. The final liquid-phase sample was analyzed for ethylene, ethanol, and sulfuric acid. The ethylene was determined by measuring t.he volume of gas which &sled on pressure reduction and the ethanol bv the method of Pondorf ( I I ) , which was found to be rapid and" satisfactory. In analysis of 'the gas phase, the total condensate was weighed and nnalyaed For ethanol, and the et.hvlene hv meesurinz the volume and comnosition " - ~ -determined -~ of~nancondensedgas. " ~
r/
Discussion of Results The results of these determina,tions are presented in Table A~SEMBLEDAPPARATns
I. E'roni the vapor-phase equilibrium data the free energy of hydration was calculated for each run,
Data are not available for the evaluation of the fugacities of water and cthanol at pressures above the vapor pressure and a t temperatures helow the critical. For this reason it was necessary to resort to the use of fugacity d a h for bydrocarbon vapors (2,7, 8, 9) based on equivalent conditions of re-
duced temperature and pressure. Calculations were made to check this method in regions where actual data were available. The following is a. comparison of fugacity of saturated vapors based on hydrocarbon vapor data with that calculated from actual data (4,6): -(~)CIHIOA--
OC. 170
True Afm. 14.0
2W 243.1
24.0 46.5
Temp.
Caicd. Atm.
13.21 22.6
41.1
Error % .. - 5.0 - 5.8
-11.2
--u*)n&-True
Calod.
Alm. 6.80 13.90
Aim. 7.42 ?4.22
30.0
dO.7
Error % .~ 19.1 +2.8 +Z.B
It is worthy of note that the use of pressures in place of fugacities produces relatively slight changes in the calculated AF". This fact is not surprising when it is noted that under the conditions of the experiments, ethylene existed at such high reduced temperatures that its behavior approached that of a perfect gas. Whereas tho ethanol and water existed at far lower reduced temperatures, their deviatioirs from perfect gas behavior tend to cancel, since they occur in tho free-energy equation as a ratio. At a pressure of 200 atmospheres the value of AF" calciilated with pressures deviates about 10 per cent from the value calculated with fugacities. Figure 2A is a plot of tho experimentally determined values of the free energy of hydration for the reaction, CzHdd
+ HzO(g) +GHsOH(g)
The recommended equation based on this plot is AF-
FIGUAE 2
-
-7780
+ 26.2T
Two side reactions were found to occur appreciably urider favorable conditions. Ahove 250' C. and 200 atmospheres, polymerization of ethylene was found to be considerable, when using approximately 3 mole per cent sulfuric acid c a h lyst. At temperatures below 220" C . ether formation was very noticeable. In Figure 2B the equations for free energy of hydration recommended by previous investigators have been plotted with the present equation. The result8 of the present invrsti-
INDUSTRIAL AND ENGINEERING CHEMISTRY
372
TABLEI.
EQUILIBRIUM COMPOSITION
-Mole Fraction in Li uid PhaseRun No. Temp. Pressure EtOH CzHa &so4 Hz0
OK. Atm. 4 5 6 7
527 264.2 0.0755 527 196.9 0.063 527 129.7 0.046 527 82.6 0.014 8 527 196.9 0.060 9 527 264.2 0.084 129.7 0.0445 10 527 196.9 0.055 11 491 196.9 0.039 12 449 13 580 197.8 0.028 a Remaining fraction, ether.
.... 0.0115 0.0076 0.0023 0.0079 0.0197 0.0075 0.0110 0,0093 0,0170
Mole Fraction in Gas Phase EtOH CzHa Hz0
0.885 0.903 0.929 0.956 0.918 0.881 0.937 0.901 0.920 0.953
0:023 0.017 0.028 0.014 0.015 0.011 0.033 0.032 0.002
...
...
... ...
0:OiS 0.104 0.075 0.095 0.226 0.083 0.033
0:ifS 0.380 0.250 0.226 0.225 0.475 0.394
4
.... .... .... 0.449 0.516 0.675 0.679 0.385a 0.442 0.573
VOL. 28, NO. 3
well have industrial possibilities. If dilute sulfuric acid catalysts are to be used, it is apparent that the temperature must lie between 220” and 250 O C. in order to avoid side reactions to ether or polymerized ethylene. On the other hand, the temperature should be as near 250” C. as possible to secure increased reaction rate.
Nomenclature AFO =
standard free-energy change
= fugacity of satd. vapor at temp. in question fr = fugacity of pure gas at temp. in question fp
at I)ressure eaual to total Dressure
i.
gation correlate well with the experimental work of Saunders and Dodge, of Bowles, and of Stanley, Youell, and Dymock. In contrast, the computed values of Parks and Huffman and of Francis and Kleinschmidt indicate far higher values. The fact that the values of AF” calculated by Parks and Huffman, using the newer heat of combustion of ethanol, are 2000 calories closer to the experimental values indicates that the newer heat of combustion of ethanol (326,610 calories) is probably more accurate. Although the curve of Francis and Kleinschmidt is higher than the experimental curves a t the higher temperatures, because of its greater slope it intersects the extrapolated experimental curves a t about 25” C. This would indicate that although the AF0288 used by Francis and Kleinschmidt is accurate, their specific heat data used in obtaining AF” as a function of temperature were in error. Accordingly, it is probable that the specific heat equation used by Parks and Huffman is the more reliable. An expression for the vapw-phase hydration that satisfactorily correlates all available data (including the present work) is given by hhe equation, A F o = 26.9T
- 8300
The close agreement of the experimental data indicates that this expression is probably accurate within * 500 calories. The fact that liquid-phase concentrations of ethanol as high as 18 weight per cent of alcohol were attained in the experiments indicates that the direct hydration of ethylene may
(g) = gas phase
K p = equilibrium constant expressed with pressures in atmos-
Dheres R = gas constant T = temp., O K. 2 = mole fraction of substance in liquid phase y = mole fraction of substance in vapor phase
Literature Cited (1) Bowles, S.M. thesis, Mass. Inst. Tech., 1933. 23,887(1931).’ (2) Cope, Lewis, and Weber, IND.EN^. CHEM., (3) Francis and Kleinschmidt, Am. Petroleum Inst. Bull. 11, 93 (1930). (4) International Critical Tables, Vol. 111, p. 437, New York, McGraw-Hill Book Co., 1928. (5) Keenan, “Steam Tables,” A. S. M. E., 1930. (6) Kleaver and Glaser, M.H. chem. tech. Inst. Tech. Hochschule Karlsruhe, 1, l(1923). (7) Lewis and Kay, Oil Gas J.,33 (45),40 (1934). (8) Lewis and Luke, Trans. Am. SOC.Mech. Engrs., 54,55 (1932). (9) Lewis and Randall, “Thermodynamics,” New York, McGrawHill Book Co., 1923. (10) Parks and Huffman, “Free Energies of Some Organic Compounds,” A. C. S. Monograph No. 60, pp. 76-81, 109, 126, New York, Chemical Catalog Co., 1932. (11) Pondorf, 2.anaE. Chim., 80, 401 (1930). (12) Rossini, Bur. Standards J . Research, 8, 119 (1932). 26,208(1934). (13) Saunders and Dodge, IND.ENG.CHEM., (14) Stanley, Youell, and Dymock, J. SOC.Chem. Ind., 53, 205T (1934). (15) Swann, Snow, and Keyes, IND.ENG.CHEM.,22, 1048 (1930). RECEIVED August 27, 1935.
CRANBERRY PECTIN PROPERTIES H E cranberry is interesting to jelly workers on account of the high viscosity of solutions made from its extracted pectin. It is also interesting because of the formation of jellies of low sugar content. The reason for the formation of low-sugar jellies lies in the low p H of the fruit extract and in the quality and quantity of the pectin present. Attempts to make 65 per cent sugar jellies from the extracted juice have failed. If the pectin present in the extract is precipitated with alcohol to get rid of excess acid and is then put back into water solution, jellies containing 65 per cent sugar can be made with ease. It seemed desirable fora better understanding of the fundamentals of jelly making to investigate properties of precipitated cranberry pectin and its relationship to alcohol-precipitated apple and lemon pectins.
Jelly from Heat-Extracted Juice
In the course of the preparation of cranberry pectin a correlation was found between the viscosity of heat-extracted
GEORGE L. BAKER AND RALPH F. KNEELAND Delaware Agricultural Experiment Station, Newark, Del.
cranberry juice and the strength of jelly produced from these extracts. The cranberries used were purchased in the local market without regard to variety a t intervals from October through January. The juice for the jelly strength-viscosity correlation was extracted by boiling the cranberries with an equal weight of water in a covered dish until the berries were soft. The boiling period approximated 10 minutes. The juice was pressed out of the pulp through muslin cloth and then filtered through flannel. Viscosities of the extracted juice were measured at 26” C. with a large capillary Ostwald viscometer. The relative viscosities of the juice vaned from 132 to 15. Contrary to ex ectation the early season extractions had relative viscosities or from 40 to 60, while the extract with a relative viscosity of 132 was obtained from berries purchased in December. Relative viscosities below 40 were obtained by boiling the extract under a reflux condenser, a process which easily depolymerizes the pectin.
