WILFRED C. LING' and CHRISTIE J. GEANKOPLIS

The Ohio Stdte University, Columbus 10, Ohio I Liquid-Phase Esterification of Oleic Acid and Isobutyl Alcohol

Studies on the rate of esterification indicate that conversion can be predicted within 2 6% using starting conditions. A comparison is made with reaction rates obtained using n-butyl alcohol

IN THE last few years the manufacture and importance of aliphatic esters have increased appreciably. Liquid phase esterification using an acid and alcohol is an important method for producing these esters. Much experimental work remains to be done to make it possible to predict the kinetics of such reactions and the effects of the structure of the alcohol on the reaction rates. Groggins (2) states that the reaction rates of is0 alcohols are generally slower than those of normal alcohols, because of the steric hindrance. The conversions a t equilib- rium, however, are the same. Ingold (3) summarized the kinetics of the acid- catalyzed reaction between an organic acid and an alcohol as a reaction between an acid-hydrogen ion complex and an alcohol.

n-Butyl oleate, an important ester, was studied by Othmer and Rao (5), who reported the effects of the ratio of alcohol to oleic acid, temperature, and sulfuric acid catalyst concentration on the reaction of n-butyl alcohol with oleic acid. They used oleic acid of 97.37% purity obtained from crude oleic acid by esterification, hydrolysis, and vacuum distillat ion,

Because extensive studies using the n-butyl alcohol were available, the present work was concerned with the reaction of isobutyl alcohol and oleic acid. The effects of the process vari- ables-temperature, concentration of the catalyst sulfuric acid, and molar ratio of alcohol to oleic acid-were studied. These reaction rates were compared with those using n-butyl alcohol. An empirical equation was derived relating these process variables and the reaction velocity constant.

Experimental Procedures The starting material for purification

was C.P. oleic acid from the Fisher Scien- tific Co., which on analysis gave the following results: iodine value, 96.91 ; neutralization value, 198.18; melting point, 8-9' C.; n?, 1.4612. The method of purification was similar to that used by Brown and Shinowara

Present address, Abbott Laboratories, North Chicago, 111.

(7), who obtained ve'ry pure acid. The. oleic acid was dissolved in acetone, cooled, and crystallized three times. The product was then fractionally dis- tilled at 1-mm. pressure using a Snyder column. The fraction at 195' C. was collected and analyzed.

The purity expressed on the basis of molecular weight was calculated to be 99.4%. The isobutyl alcohol, reagent grade (Eastman Kodak Co.), was re- distilled. The acid used contained 95.6% sulfuric acid by weight.

The esterification method used here was essentially that of Othmer and Rao

(5) . The oleic acid, alcohol, and sul- furic acid were weighed separately, chilled, and poured together into a flask. The contents of the flask were quickly mixed and aliquots were taken and added to a series of 13-mm. test tubes drawn out to a point. The tips were rapidly sealed in a flame and the vials placed in a constant temperature bath. At a given time interval a vial was removed.and quickly chilled, and a sample was removed for weighing. This sample was then dissolved in acetone and oleic acid was determined by titra- tion with potassium hydroxide (Table I).

Analysis of Purified Oleic Acid Theoretical

Exptl. Value Neutralization No. 199.89 198.58 Saponification No. 199.46 198.58 Iodine No. 88.38 89.87 Melting point, C. 12.4 13-34" Refractive Index, ng 1.4596 1.45970 Mean molecular weight (from neutralization

No.) 280.65 282.45 a Literature value (8).

Table I. Kinetic Data for Esterification of Oleic Acid and Isobutyl Alcohol Time, Min.

Run A. 0 30 60 120 180 300 420

Run B. 0 30 60 100 150 200 300

Run C. 0 25 50 75 100 140 200

A - X HzS04 = 0 Wt.%.

0.1544 0.1518 0.1497 0.1458 0.1418 0.1361 0.1300

His04 = 0.1505 n%, 0.1534 0.0522 0.0267 0.0159 0.0124 0.0116 0.0111

X Temp., 100' C.

0 0.0026 0.0047 0.0086 0.0126 0.0183 0.0244

X

B / A = 4.955. 0 1.04 2.41 4.95 7.56 11.27 15.22

Ao = 1.193 0

0.01435 0.0266 0.0494 0.0745 0.1125 0.1572

.%.

HzSOc = 0.2375 n&%. 0.1524 0.0424 0.0202 0.0141

0.0114 0.0113

O.OJ22

Temp., 100' C. B / A = 4.997. 0 0

0.1012 66.0 0.1267 82.5 0.1375 89.6 0.1410 91.9 0.1418 92.4 0.1423 92.8

Ao = 1.187 0 1.63 4.00 7.29 9.58 10.30 10.81

Temp., 100' C. B / A = 5.043. Ao = 1.180 0 0 0

72.3 2.20 86.9 5.55 90.8 8.33 92.1 9.75 92.6 10.50 92.6 10.60

0.'1101 0.1323 0.1384 0.1403 0.1411 0.1412

VOL. 50, NO. 6 JUNE 1958 939

Table 1. Kinetic Data for Esterification of Oleic Acid and (Continued)

Time, Min. A - X X (X/A)UOO)

Run D. &SO4 = 0.5626 wt.%. Temp., looo C. B / A = 4.998. 0 0.1520 0 0 10 0.0525 0.1002 66.7 22 0.0204 0.1323 86.7 34 0.0143 0.1384 90.8 46 0.0126 0.1401 91.9 58 0,0121 0.1406 92.1 90 0.0119 0.1408 92.3

Run E. HzSOI = 0.931 wt.%. Temp., 100" C. B / A = 5.017. 0 0.1518 0 0 6 0.0450 0.1068 70.4 21 0.0132 0.1386 91.3 31 0.0125 0.1393 91.9

Run E'. HzSOa = 1.076 wt.%. Temp., 100' C. B / A = 5.066. 0 0.1507 0 0 10 0.0245 0.1262 83.9 29 0.0126 0.1381 91.8 44 0.0122 0.1385 92.0 59 0.0119 0.1388 92.1

Run F. HzSOa = 0.2613 wt.%. Temp., 100" C. B / A = 1.017. 0 0,2801 0 0 10 0.1814 0,0987 35.2 30 0.1591 0,1210 43.1 63 0.1505 0.1296 46.3 120 0.1403 0.1398 49.9 164 0.1369 0.1432 51.1

Run G. HzS04 = 0.2418 at.%. Temp., looo C. B / A = 2.016. 0 0.2318 0 0 10 0.1217 0.1101 47.8 28 0.0808 0.1510 65.5 43 0.0707 0.1611 70.0 60 0.0653 0.1665 72.3 90 0.0582 0.1736 75.3 136 0.0534 0.1784 77.5 180 0.0506 0.1812 78.7

Run H. HzS04 = 0.2605 Bt .%. Temp., 100' C. B / A = 9 993. 0 0.0976 0 0 10 0.0517 0.0459 47.0 20 0.02996 0.0679 69.5 31 0.01788 0.0798 81.6 50 0.00906 0.0886 90.7 91 0.00475 0.0929 95.1 152 0.00410 0.0935 95.8

Run I. HzSO4 = 0.2522 wt.%. Temp., 100' C. B / A = 20.13. 0 0.0563 0 0 10 0.03052 0.02575 45.7 20 0.01670 0.0396 70.3 31 0.00964 0.0466 82.9 44 0.00532 0.0510 90.5 75 0.00215 0.0541 96.2 150 0.00148 0.0548 97.3

Run J. HzS04 = 0.2715 at.%. Temp., 8OOC. B / A = 5.042. 0 0.1524 0 0 10 0.1081 0.0443 34.4 20 0.0866 0.0658 43.1 38 0.0649 0.0875 57.4 63 0.0446 0.1078 70.7 90 0.0322 0.1202 78.9 150 0.0196 0.1328 87.1 2 10 0.0150 0.1374 90.2

Run K. HzSOc = 0.2506 wt.%. Temp., 150' C. B / A = 5.019. 0 0.1529 0 0 10 0.01 108 0.1419 92.8 20 0.01034 0.1426 93.3 29 0.01023 0.1427 93.4

Run K'. HzSOc = 0.2489 at.%. Temp., 150" C. B / A = 5.037. 0 0.1525 0 0 5 0.02486 0.1276 83.5 15 0.01019 0.1423 93.4 45 0.01008 0.1424 93.5

Isobutyl Alcohol

X Ao(A - X)

Ao = 1.184 0

1.61 5.48 8.18 9.39 9.81 9.99

A0 = 1.179 0

2.01 8.90 9.45

Ao = 1.172 0

4.40 9.35 9.69 9.95

A0 = 2.275 0

0.239 0.334 0.378 0.437 0.460

do = 1.848 0

0.390 1.01 1.23 1.38 1.61 1.81 1.94

Ao = 0.740 0 1.20 3.08 6.02 13.20 26.4 30.8

Ao = 0.420 0 2.01 5.64 11.51 22.8 59.9 88.1

Ao 1.207 0

0.339 0.629 1.19 2.00 3.09 5.60 7.59

A0 = 1.112 0

11.50 12.39 12.52

A0 = 1.110 0

4.62 12.60 12.73

Density data at high temperatures for water, n-butyl alcohol, sulfuric acid, and oleic acid were available (5) and density of isobutyl alcohol was de- termined using a pycnometer.

These data and data at 17.5' C. from Perry (6) were plotted and extrapolated to 150' C.

Experimental Densities for Isobutyl Alcohol

Temp., O C. G./CC. 80.5 0.7512 84.7 0.7467 90.1 0.7418 94.9 0.7369 100.2 0.7312 105.1 0.7256

Treatment and Discussion of Data

Calculation of Conversion. By using rhe data of Table I, the conversion was calculated as follows: An equivalent weight of 94.3 for sulfuric acid in isobutyl alcohol was used, as found by experiment. This compares with the value of 94.0 when n-butyl alcohol was used, as found by Leyes and Othmer (4). 'This amount of acid was subtracted from the total acid found by titration with potas- sium hydroxide to give the amount of oleic acid remaining or ( A - X). The moles of oleic acid per 100 grams of solution at 'the start of the reaction, A , were calculated from the weights of the starting constituents. The amount of isobutyl oleate formed, X, was the dif- ference between A and ( A - X). The oleic acid at the start in moles per liter, Ao, was calculated assuming additive volumes of the three components. The per cent conversion to ester was cal- culated as ( X / A ) (1 00).

Determination of Order of Reaction. The order of reaction can often be deter- mined by plotting the appropriate con- centration function against time. Using data from Table I, functions of concen- tration for first-, second-, and third-order equations were plotted against time. The resulting plots indicated that a second-order equation best represented the experimental data over the range of variables studied. Othmer and Rao (5), in their studies with oleic acid and n-butyl alcohol, showed that their data satisfied the following equation, which is second order with regard to the oleic acid concentration:

X Ao(A - X) k t =

for the cat- alyst series runs is plotted in Figures 1 and 2. In Figure 1 the n-butyl alcohol data of Othmer are plotted for comparison with the isobutyl alcohol

&(A - X) The function

940 INDUSTRIAL AND ENGINEERING CHEMISTRY

40f

b 0 0.26

- 6 X

a x .4 a

I 0 P 0.48

/"\ \ ] FOR RUN L, MULTIPLY / \

201- / - I - : -

0 100 200 300

Figure 1 . The curves show how the function X/Ao (A-X) , which is related to conversion, increases almost linearly as time increases. However, the curves of Othmer and Rao (5) for n-butyl alcohol show an abrupt peak and do not level off

TIME, MINUTES

data in Figure 2 . When there was no catalyst present, run L for n-butyl alcohol (Figure 1) gave an initial reac- tion rate about 80% greater than run A (Figure 2 ) for isobutyl alcohol. This can be explained, because steric hin- drance of the isobutyl alcohol would be expected to slow down the reaction. However, runs C, D, and EE' (Figure 2) with isobutyl alcohol and sulfuric acid catalyst showed initial reaction rates at least 15% higher than did runs 0, P, and Q (Figure 1) with comparable amounts of catalyst. This could be due to different impurities in the oleic acid or other unrecognized factors.

The curves in Figure 2 and plots of the remaining molar ratio and tempera- ture series data show that the reaction reaches a maximum above 90% con- version. These curves show the ex- pected approach to esterification equi- librium where the reverse hydrolysis reaction enters in. This is in contrast to the curves in Othmer's Figure 1 for n-butyl alcohol, which show more abrupt peaks and drop readily.

To determine the order of reaction, the slopes should be measured at low conversions to avoid the reverse reaction complications. Sufficient data (Figure 2) at low conversions are available only for run A (no catalyst), which give a straight line and indicate a second-order

OLEIC ACID ESTERIFICATION

12

IC

8

2

C

Figure 2.

0 / A = 5 RUN %CAT.

B 0.1505 -

C 0.2375 D 0.5626 EE' I .003(AV.)

100 200 300 400 5 TIME, MINUTES

The data plotted for isobutyl alcohol show the expected approach to esterification equilibrium and level off above 90% conversion. A small amount of sulfuric acid catalyst has a large effect on the conversion

3

reaction. Data for other runs with cat- alyst present are insufficient to draw any conclusions regarding the order of reaction because of the reverse hydrolysis a t high conversions. Runs B, C , D, EE', H, I, J, and KK' appear to give relatively straight lines even up to high conversions and not to pass through zero time. This effect was also found by Leyes and Othmer (4) in their study of the esterification of butanol with acetic acid, particularly at high alcohol- acid ratios. This may indicate an in- duction period, or it could be caused by errors in analytical methods or in measuring zero time, or by slight im- purities in the oleic acid.

To determine an average reaction rate, k, which best represents the data

+ over the range of variables studied, the

slopes were measured of the best straight lines drawn through zero time and the points up to 90% conversion. The data are given in Table 11. Because of curvature, the best straight lines through the first three points (below 90% con- version) of runs H and I were used. Runs F and G were not used because they had very small BIA ratios and the products yielded a heterogeneous mix- ture. Hence, diffusion may be a factor in these two runs.

Effect of Process Variables on Rate Constant. To determine an empirical equation which would predict the rate constant and conversion from the process variables catalyst concentration, molar ratio, and temperature, various plots were made. The rate constant was plotted against the per cent sulfuric acid for runs A, B, C, D, and EE', in which the other variables were kept constant. The straight-line equa-

Run A B C D EE' H I J KIC'

Table II. Rate Constants for Esterification Rate Constant, k, Liters/G. Mole-Min.

Exptl. Value calod. '% Deviation of Calcd. and

value from Es. 6 Exptl. Value 0.000368 0.0667 0.107 0.229 0.400 0.175 0.315 0.0337 1.00

0.000375 0.0622 0.0986 0.231 0.415 0.173 0.297 0.0375 0.975

1.8 -6 .5 -7 .9

1.0 4.0

-1 .2 -5 .7 11.2

-2 .5

VOL. 50, NO. 6 JUNE 1958 941

Table 111. Comparison of Experimental and Calculated Con- versions, Using Equations 5 and 1

Run Min. Exptl. Calcd. Deviation 5 Time, Conversion, yo

PREDICT ED CONVERSION, % Figure 3. The figure shows that the predicted conver- sions using the empirical equation compare favorably with the actual conversions. The average deviation for all runs is 9 2 . 5 conversion units

tion representing the data is

k = 0.401C + 0.000368 (2)

Othmer and Rao (5) found similar trends with n-butyl alcohol.

Runs C , H, and I were plotted to show the effect of molar ratio, B/A, a t a cor- rected 0.25 weight yo sulfuric acid cat- alyst concentration. The equation re- lating k to B/A is

k = 0.0133B/A + 0.042 (3)

Othmer (5) found a somewhat similar linear relation for B/A values between 1 and 15.

Following the classical method, a plot of log k us. 1/T was made; it shows a straight line for the 80”, loo”, and 150” C. runs as follows:

log k = 7.28 - 3085/T (4) This corresponds to an energy of ac- tivation of 14,100 cal. per mole. Smith and Reichardt (7) determined the energy of activation of esterification of methanol with acids from caproic to pelargonic with hydrochloric acid as a catalyst. They obtained an average of 10,000 cal. per mole for all reactions.

The final equation to predict the average reaction velocity, k , was de- termined by combining the separate equations relating C, B/A, and T. The combined equation is

k = m(0.401C f 0.000368)(0.0133 B/A 4- 0.042)(e16.76--7100/T) ( 5 )

m was evaluated as 93.0 by substituting experimental values of k , C, B/A, and T in Equation 5 . Equation 5 was then used to calculate k for the experimental runs. The calculated and experimental values of k in Table I1 show an average deviation of &4.7y0 and a maximum of ll.2yo for run J for a temperature of 80” C.

A 30 60 120 180 300 420 30 60 100

B

C

D

EE’

H

1.04 2.41 4.95 7.56 11.27 15.22 66.0 82.5 89.6

1.33 2.62 5.10 6.96 11.77 15.75 68.9 81.6 88.1

21.7 8.0 3.0

-8.7 4.0 3.5 4.0

-1.0 -1.7

25 72.3 74.5 2.9 50 86.9 85.3 -1.9 10 66.7 73.3 9.0 22 86.7 85.8 -1.0 6 70.4 74.8 6.3 10 83.9 83.0 -1.1 10 47.0 56.1 20 69.5 71.8 31 81.6 80.0 10 45.7 55.5 20 70.3 71.3 31 82.9 79.5

16.2 3-0

-1.9 17.7 1.3

-4.2 10 34.4 31.2 - 10.2 20 43.1 47.6 9.5 38 57.4 63.3 9.5 63 70.7 74.1 5.5 90 78.9 80.4 2.8 , 150 87.1 87.2 0.1

KK’ 5 83.5 84.4 1.0

Using Equation 5 to calculate k and Equation 1 to calculate X, the per cent conversion, (X/A)(100) a t time t was calculated and compared with the actual conversion in Table 111. This was done for the runs up to 90% con- version, the point where the reaction velocity constant decreased appreciably because of reverse or side reactions. The calculated conversions appear to agree fairly well with the experimental values. The average per cent deviation for all runs is f 5 . 8 % , with the maximum 21.7% for run A having a very low con- version of 1.04%. This corresponds to an average deviation of 1 2 . 5 con- version units. The actual us. predicted conversions, plotted in Figure 3, show that the calculated data do not deviate appreciably from the predicted. This equation holds only for the range of given conditions covered in the experi- ments, and extrapolations should be done with caution. For example, the effect of catalyst concentration was de- termined at 100’ C. but not a t 80” or 150” C. The effect of B / A was de- termined only a t 100” C. and not at the other temperatures. The final equation does not justify any conclusions regarding the order of reaction when catalyst is present, but can be used to relate the process variables to an average reaction rate.

Nomenclature A = initial oleic acid concentra-

tion, gram moles per 100 grams of solution

AJ = initial oleic acid concentra- tion, gram moles per liter of solution

( A - X) = oleic acid remaining after reaction, gram moles per 100 grams of solution

= initial iso-butyl alcohol con- centration, gram moles per 100 grams of solution

C = H2S04 catalyst concentra- tion, weight per cent

k = reaction velocity constant, liters per gram mole-min.

rn = constant t = time, min. T = temperature, O K. X

B

= amount of iso-butyl oleate ester formed, gram moles per 100 grams of solution

(X/A)(100) = per cent conversion to ester

Literature Cited

Brown, J. B., Shinowara, G. Y . , J. Am. Chem. SOC. 59, 6 (1937).

Groggins, P. H., “Unit Processes in Organic Synthesis,” 4th ed., p. 605, McGraw-Hill, New York, 1952.

Inaold. C. K.. “Structure and Mecha- ;ism in Organic Chemistry,” p. 769, Cornell University Press, Ithaca, N. Y., 1953.

Leyes, C. E., Othmer, D. F., IND. ENC. CHEM. 37, 968 (1945).

Othmer, D. F., Rao, S. A., Ibid., 42, 1912 (1950).

Perry, J.’H., ed., “Chemical Engineers’ Handbook,” 3rd ed., p. 132, Mc- Graw-Hill, New York, 1950.

Smith, H. A., Reichardt, C . H., J. Am. Chem. SOC. 63, 605 (1941).

Smith, J. C., J . Chem. SOC. 1939, p. 974. RECEIVED for review March 12, 1956

ACCEPTED October 19,1957

942 INDUSTRIAL AND ENGINEERING CHEMISTRY