KINETICS IN ACID MEDIA Condensation of o-Benzoylbenzoic Acid Charles W . Deanel with John R . Huflrnan NEW YORK UNIVERSITY, UNIVERSITY HEIGHTS, N. Y.
Catalysts f o r the condensation of obenzoylbenzoic acid to anthraquinone in the presence of concentrated and .fuming sulfuric acid have been sought by kinetic measurements; flasks were held at constant temperature in an oil thermostai and other conditions were varied systematically. The only added substance which promoted the reaction velocity was sulfur trioxide. A study of this reaction in fuming acids of various concentration established that the reaction is unimolecular, practically quantitative, inhibited by the water and t o a greater extent by the anthraquinone formed during the reaction. These lastphe-
A
LTHOUGH catalysis was recognized more than EI hundred years ago, our knowledge of catalytic phenomena is for the most part still uncoordinated. Reactions of organic materials in sulfuric acid media disclose widely varying effects and useful products obtainable in achieving esterification (63), dehydration (69),hydration (1, $3, 68, 64), alkylation (4, 6, 28, 35, 36, 37, 40, 57), condensation (20, 34, %), polymerization (3, 32, 39, 41, @, 62) and molecular rearrangement (9, H),but no comprehensive explanation of the kinetics has been advanced. Changing solvents for reactions have been noted to change rates sixfold (7), two thousand fold (@), and one and a half million fold (84, 25). To explain these phenomena, investigators have used different properties of the solvent to account for the results: cohesion (66), refraction (11), viscosity (29), solvent power (IQ), and dielectric constant (31, 43, 67). No general interpretation, therefore, exists to explain the retarding or accelerating action of solvents. The tendency for catalysis may be less a function of the general factors in experimental conditions than of temperature, time, pressure (when phase is gaseous and order is higher than unimolecular) , and solvent factors. Kinetic studies of the decomposition of organic acids by sulfuric acid were initiated in 1906 by Bredig and Lichty (5). The unusual effects noted in the sulfuric acid decomposition of oxalic acid led Taylor (60) to develop his theory of negative catalysis. Wiig (72) summarized the work in this field to 1930, discussing in detail the double retarder theory. Schierz (51) studied the effects of various nitrogen compounds on the decomposition of formic acid, and in 1938 two Russians (47) noted the inhibitory effects of sulfur dioxide, hydrogen sulfide, phosphorus pentoxide, and arsenic pentoxide. Work on the decomposition of malic acid (l5,16,68) showed the effects of fourteen inhibitors. 1
Present address, Colgste-Palmolive-Peet Company, Jersey City, N. J.
684
nomena are contrary to results previously reported in the literature. Various reaction mechanisms are discussed. The energy of activation as measured in the fuming sulfuric acid media is 26,100 calories. The temperature coeflcient for a 10’ C. rise is approximately 3. The commercial implications and engineering values of the above data are: to indicate methods of determining desirable operating conditions, to suggest means f o r continuous processing t o yield anthraquinone, and to establish desirable short cuts and new economies in producing several anthraquinone dyes and intermediates. Gleason and Dougherty (21)summarized the history of o-benzoylbenzoic acid condensation to anthraquinone with the accompanying removal of water, from the first work in 1874 (2,42). All experiments reported were carried out with 100 per cent sulfuric acid or less. None of these data indicate the possibility of a m a x i m a decomposition rate at a certain sulfuric acid concentration, although for malic acid (68) and for citric acid (70) maximum rates were observed a t 99.8 per cent sulfuric acid concentration. Wiig’s work (71) indicated a maximum rate for oxalic acid at about 14 per cent fuming sulfuric acid. Gleason and Dougherty (bl), however pronounced the o-benzoylbenzoic acid reaction unusual in the organic series and believe that it differs fundamentally from the “oxalic acid type’,. Wiig (72) showed that a plot of the logarithm of velocity constant us. molality of water gave a straight line within the accuracy of the data, which is contrary to Gleason and Dougherty’s reported direct straight-line correlation of molality of water os. reaction velocity constant. Four fifths of the reaction rate measurements by Gleason and Dougherty occurred above the 0.5 life point, and some were above the 0.9 life fraction. In the light of criteria (48, t o ) for the interpretation of kinetic figures, care must be used in making conclusions from these data. The practical need for verification and extension of data on the production of anthraquinone is emphasized b y the stimulus of World War I to domestic dye manufacture and the heavy demands of the present war for the valued anthraquinone vat dye derivatives. Both the oleum and o-benzoylbenzoic acid (90) raw materials are now economically available for manufacturing, as a result of processing advances. ADVANTAGES FOR KIXETIC STUDY
This work was undertaken to investigate materials known to be promoters for similar reactions, to establish data on theoretical catalytic implications, to develop practical economic processing, to secure check data, and to extend data into the important region of fuming sulfuric acid concentrations.
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
This transformation affords almost ideal characteristics because the reaction produces only negligible side products. It is one of the few condensation reactions having both this feature and freedom from reaction by induced reaction. The product for analysis (anthraquinone) precipitates from the sulfuric acid media b y dilution with water. After filtering, washing, and drying, the yields are quantitative. Of the kinetic studies on organic acids, this reaction is exceptional in that no gas is released and all the original atoms remain in the liquid phase, although they are differently configured. EXPERIMENTAL METHODS. The o-benzo lbenzoic acid was purified according to the Eastman Kodak Laborator method (28) by crystallization from glacial acetic acid, wasging, and drying., The dry pure product melted at 127.2" to 128.5" C. (corrected). International Critical Tables and Richter's Lexikon give a melting point of 127" for water-free material. Drying the urified material for 12 hours at 110" to 115" C. caused a loss of fess than 0.2 per cent. The water-absorbing power of this material on exposure to ordinary conditions of humidity was negligible (0.002 per cent). Unpurified material was 3-5 per cent volatile. Although dark colored solutions were formed in SUIfuric acid, no charring occurred, and check tests roved the analytical procedure within 0.1 to 0.4 per cent. T t e sulfuric acid and o eum were analyzed by precipitation as barium sulfate; due precautions were observed in weighing and mani ulating samples to prevent sulfur trioxide loss and moisture atsorption. Checks were made by specific gravity determinations. Fuming acid concentrations were ralculated from the weight and strength of two lots of acid, and check determinations were made by hydrometer and temperature readings and the 1939 data (8) on fuming sulfuric acid strength. In addition, calculated proportions of the two stock acids to give 100 per cent acid were mixed, and the melting point was determined. This served to check the mixing and the strength calculating procedures for the acid. The average of a triplicate melting point was 10.38" C. Since 100 per cent sulfuric acid has a maximum melting point of 10.46' C., the acid proved to be within 0.15 per cent on the fuming side or within 0.05 per cent on the water side by comparison with a plot of the Hantzsch data (27). APPARATUSAND PROCEDURE. I n general, solutions of o-benzoylbenzoic acid in sulfuric acid of the desired strength were prepared in 50-cc. Erlenmeyer flasks or in 200-cc. round-bottomed flasks fitted with calcium chloride tubes and mercury-sealed stirrers. Usually 8.1021 grams of ortho acid were added t o 35 cc. of sulfuric acid (or a proportionate fraction of each was used) at the reaction temperature attained in an automatic oil thermostat. Independent tests showed that within 6-9 minutes in the most severe cases (i. e., high temperature and high fuming acid strength), the tem erature was within 0.1" C. of that prevailing in the thermostat ($onstant within 0.07') ; at that point an initial sample was removed. Subsequent Sam les were taken by a 5-cc. pipet at 0.25,0.40, 0.60, and 0.75 of the fife period of the reaction, drlivered into 100 cc. of water to stop the reaction, and analyzed in accord with the procedure outlined below. The initial concentration, a, in the unimolecular rate equation was determined either by allowing the reaction to go to completion in triplicate samples; or by calculation from the original amounts of ortho acid added, where correction was made in each case for the expansion due to temperature change and to the addition of ortho acid. In order to start with known weights of sulfuric acid and obenzoylbenzoic acid, to withdraw a 5-cc. volhetric sample, and to know the actual extent of its organic content, the correction factor had to be applied. These corrections were at 75" C., 19.6 and 19.9 per cent, and at 85' C., 20.1 and 20.4 per cent, for 98.0 and 105.9 per cent sulfuric acid, respectively. At intermediate acid percentages, correction values were obtained by interpolation between proper points of the foregoing expansion data. In the search for expected promoters, one individual run, usually a check run, and a comparison run were made. The comparison consisted of placing 5 or 10 cc. of sulfuric acid in 50-cc. Erlenmeyer flasks. Owing to comparison under similar conditions, it was not necessary to determine a in these experiments. T o compute the reaction rates, the addition time until the reaction was stopped by quenching in water was used. METHODOF ANALYSIS.The procedure of Gleason and Dougherty (21) was followed except that filter papers were reglaced by tared Gooch crucibles prepared with dried asbestos ber mat. Check runs proved the analytical procedure as better than *0.4 per cent. The experimental procedures outlined afforded several advantages: June, 1943
The use of reaction flasks instead of test tubes minimized the tendency of the solid o-benzoylbenzoic acid to stick undissolved on the walls of the vessel. By the older methods a quick and complete solution of small lumps of o-benzoylbenzoic acid was difficult, and this was practically overcome in the flasks. The experiments weke made on one batch (except in the preliminary work on promoters) of the same original material; hence the same conditions of life history for the reaction were assured. This is not true of small test tube runs. The carefully pre ared Gooch crucibles eliminated the uncertainties in using fiker paper for the quantitative analysis of anthraquinone. The number of weighings and measurements of starting samples was decreased; by using larger samples, a considerable increase in accuracy and reliability was possible. The use of a flask permitted glass paddle agitation, which assured homogeneity of the'mixture if there was a tendency to stratification; tests, however, showed no difference in the reaction rate with or without stirring. In cases where better heat transfer was an advantage, such as rapid reactions at high concentrations of fuming acid and cases where the heat of solution & t h e start of the reaction raised the temperature of the flask for a short time, stirring was advisable. The larger volume also permitted checks on the interior temperature; for in a test tube such a check would not be practicable owing to the lowering of the temperature of such a small volume by the cool thermometer. Successive runs gave checks within one per cent. In reaction, rate work this is unusual. I One of the objectives of the present study was to establish whether the reaction follows the unimolecular law. Another was the systematic study of velocity phenomena. Maximum rates of decomposition for other organic acids have been found in concentrations slightly under 100 per cent sulfuric acid. To ascertain whether the maximum velocity point predicted b y Wiig (72)actually eyisted for o-benzylbanzoic acid, the confirmation in this investigation of Gleason and Dougherty's results showing no maximum at concentrations near or above 100 per cent sulfuric acid was especially relevant. This led to the experiments in the fuming acid concentration range which yielded unexpected reaction rate data. PROOF O F UNIMOLECULARITY
Yields of anthraquinone b y the action of concentrated or slightly fuming sulfuric acid have been reported (21) which are quantitative or approximately 100 per cent. Up to 8 per cent fuming acid at 85" C. the present work showed yields above 98 per cent. At concentrations above 20 per cent fuming acid, however, the yields fell to 95-96 per cent (8590" (2.). A general test for unimolecularity is that the velocity constant obtained in reducing experimental data b y the unimolecular formula shall hold rather constant throughout the course of the reaction. I n the decomposition of o-benzoylbenzoic acid, more than one molecular species undoubtedly influences whatever arrangement or key mechanism is controlling the reaction rate; however, assuming that sulfuric acid is the principal kinetically effective reactant, in a large or
EXPERIMENTS IN 1.8 PERCENT TABLE I. DATAFROM TYPICAL SULFURTRIOXIDE (100.4 PERCENTSULFURIC ACID) Temp., C . a , Gram* 75 0.7599
85
0.7974
t
a--2
k
x
104
20.0 40.2
50.0 65.0
0.5894 0 4572 0.4081 0.3541
127 0 126 4 124 4 121.8 Av. 124.9
5.4 10.0 15.0
0.6547 0.5430 0.4555
382.5 379.1 373.0 365.9 Av. 375.1
20.2
* Gram of anthraquinone per 5 0 0 .
INDUSTRIAL AND ENGINEERING CHEMISTRY
0,3804 of solution.
685
686
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
TABLE11. DETERMINATION OF REACTION ORDERAT 85' C a (100.4 per cent sulfuric acid= 1.8 er cent sulfur trioxide from runs with molality of o-benzoyfbenzoic acid = 0.274) a, Gram* t a--2 k X 104 0.3839 11.0 0.2517 406.6 14.0 0.2263 395.3 0.3810 16.0 0.2168 342.0 19.0 0.1934 356.3 0.3806 16.0 0.2096 372.4 19.0 0.1883 370.1 Av. 373.86
*
Gram of anthraquinone per 5 cc. of solution. a k was calculated by the unimolecular formula from data on o-benzoylbenzoin acid at haif the usual initial concentration. b In Table IV at o-benzoylbenzoic acid molality of 0.55, average IC X 104 373.7.
-
TABLE111. WATER AS INHIBITOR OF THE REACTION HeS04,
%
96.0 98.0
100.0
98.0
100.0
(Molality of o-benroylbenzoio acid, 0.55) k X 104 Molality Av. for each Final of H20 expt. av. TemDerature. 75' C. 48.1 2.314 48.1 1.6822 1.134 77.3 77.7 77.5 1.8893 118.2 2.0726 0.000 118.7 117.5 Temperature, 85' C. 1.134 227.1 229.3 228 2.3579 343 0.000 344.3 342.8 2.5353
. -
25
--.
Av.. at l/z lifeb
Oxides Chlorides
Sulfates
47 76.8 117 223 335
Log values of final average k (column 5). b One-half life values of lines drawn through the points used to plot k life fraction for various acid strengths.
AGENTSTESTED,Types were studied which influence the related reactions, alkylation and condensation. The most pertinent consideration is the solubility of these materials in strong and weak sulfuric acid, water, and alkali. Since small amounts of promoters were added, they would have to be completely soluble to facilitate molecular and ionic action, unless the reaction unexpectedly proved to be heterogeneous and not unimolecular. The purpose of the work, however, was to investigate only homogeneous liquid phase reactions. I n general, intermediates used in dye production must be very pure; addition agents should not only be soluble in oleum or concentrated sulfuric acid, but also reasonably soluble in any one of the following, hot or cold: weak sulfuric acid, water, or alkali up to approximately 2 N strength. The agents tested were separated into the following seven classifications:
a
88.
Dehydrating agent (oxide, acid anhydride) Dissolved gas (chloride, acid) Complex inorganic oxidizing agents Organic agents (chloride, acid anhydride, ketone)
PPOS HCI IGCrzOT, KMnO4, Mg(ClOd2, HaBOs CClr, acetic anhydride, anthraquinone
A graphical comparison of the results of experiments in fuming and in 96 per cent sulfuric acid shows that they may be broadly correlated over the range studied by the equation: sufficient excess of this reactant the measured reaction should be unimolecular. I n Table I the velocity constant k is computed from the unimolecular formula. The results show a consistent slowing of the reaction as it proceeded. Moelwyn-Hughes, however, pointed out (M) that the constancy of velocity constant k throughout the reaction is not an infallible index to reaction order; expressed dserently, conformity of the results for a given reaction with one of the well known velocity constant equations is not conclusive evidence that the order of the reaction in question is the same as that for which the equation was deduced. I n practically all cases, instead of this test (or applying the common unimolecular criterion of a straight-line plot of log concentration against time) the most reliable method of determining reaction order is to study the variation of half-life time (or half completion) with different initial concentrations of the reactant. Unimolecularity is proved by tests made with the o-benaoyl benzoic acid concentration equal to half that used in the usual kinetic experiments. These values were averaged and shown a t one point of initial concentration of 0.28 molality (Table 11) ; the second point is that resulting from runs a t 0.55 molality. Despite some variations in the results of six separate determinations, the arithmetic mean corresponds within 0.2 per cent with the velocity constant determined a t the usual o-benzoylbenzoic acid concentration. Furthermore, the velocity constants obtained at 100 per cent suIfuric acid agree with those of Gleason and Dougherty a t another slightly different initial concentration of o-benzoylbenzoic acid. CATALYTIC ACTION OF YARIOUS SUBSTANCES
The effect of twenty-one materials as promoters or inhibitors of the reaction rate was studied. They were added both to fuming sulfuric acid (25 per cent sulfur trioxide) and to 95 per cent sulfuric acid in 3 and 5 per cent concentration. Typical results are shown in Tables I11 and IV. T o determine the temperature coefficient, runs were made a t 25', 65', 70', 78", and 85" C. (Table V). June, 1943
where T = ratio of rate observed with inhibitor to rate with no initial inhibitor lt = molality of inhibitor REACTION MECHANISM
The reaction rate (as compared with rates for other classes
of addition agents) is not significantly influenced by dehydrating agents such as phosphorus pentoxide in fuming sulfuric acid. This hiehlv effective dehvdrator behaves as do other inhibitors; a i oiher materials generally retard the reaction except anthraquinone and water, the products. Both of these retard the liquid phase reaction to a considerably greater extent. Sulfur trioxide, however, is a positive catalyst; since it acts as a promoter up to 28.8 per cent concentration, Wiig's prediction (72) based on Taylor's theory
TABLEIV. SULFUR TRIOXIDE AS PROMOTER OF THE REACTION HzSOc.
%
100.0
SOU,
%
0.0
100.4 101.8 103.16 104.5 105.9
1.8 8.0 14.0 20.0 26.2
100.0
101 8
0.0 l.S 8.0
103.2 104.5 105.4 105.9 106.5
14.0 20.0 24.0 26.2 28.8
100.4
(Molality of o-benzoylbenzoic acid, 0.55) k X 104 Molality Av. for eaoh Final Lo5 of SO8 expt. av. &v, Temperature, 76' C. 0.000 118.7 117.5 118 2.0726 0.233 124.9 124.9 125 2.0965 1.086 129.1 129 2.1108 2.033 135.8 136.3 136 2.1335 3.122 145.5 144.7 145 2.1624 4.434 170.4 171.0 171 2.6606 Temoerature. 85' 0-. . 0.000 34h.l 342.8 343 2.5353 373 3 0.233 375.1 374 2.5729 383.1 381 2.5809 1.086
{!;:}
{g:: ti
2.033 3.122 3.944 4.434 5.065
{E:;' :if::) 416.7 441.2 458.1 477.8
419.1 436.8 456.4 486.2
384 418 439 457 481
2.5843 2.6212 2.6425 2.6699 2.6821
Av., at lifeb
1/z
117 121.6 128 134 145 168 334
366
375 378 414 433 448 474
Log values of fins1 average k (column 6 ) b One-half life values of linea drawn thro&h the points used to plot k m. life fraction for various mid strengths. a
INDUSTRIAL AND ENGINEERING CHEMISTRY
687
% WAT ER
%SULFUR TRIOXIDE
that E should decrease only slightly with temperature increase and thus be a reference point for observations. Table VI gives values obtained a t 75” and 85”; the values of k2/kl were taken from a large scale graph of Figure 1. Since E in Table VI decreases as the sulfur triovide concentration increases, the pure sulfur trioxide reaction proceeds more easily than the water-inhibited one; the former would therefore be expected to have a steeper slope. Figure 1, however, shows that the 1%-aterreaction has the steeper slope and that LTater is a better inhibitor than sulfur trioxide is a catalyst; H. A. Taylor suggests that this anomaly is probably due to the masked inhibition referred to later. In his survey of the decomposition of organic acids in sulfuric acid, Wiig (72) predicted a maximum decomposition rate for o-benzoylbenzoic acid, in accord n-ith his deductions from previous studies and v i t h Taylor’s theory of negative catalysis. On Figure 1, however, the rate of decomposition increases with higher sulfuric acid concentration, and additional sulfur trioxide accelerates the decomposition rate of o-benzoylbenzoic acid. Since water acts as a retarder and sulfur trioxide does not, the basic mechanism for the reaction of both molecules acting as retarders appears to have been altered in this case. Possibly this is due to the large and stable phenyl rings in the organic acid, both of which are adjacent to the portions of the molecule undergoing change; thus the special steric effect of inhibition may be caused only by the more mobile water molecules (or ions). AUALYSIS OF GLEASOIV AND DOUGHERTY’S D&T4
4.0 3.0
2.0
1.0
OTO1.0
M O L A L I T Y H,O
ZP
3.0
4.0
5.0
MOLALITY SO3
Figure 1 . Velocity Constant k f o r Condensatio? of o-Benzoylbenzoic Acid by Sulfuric Acid a t 75 and 85’ C . 0 Deane. 0 Gleaaon and Dougherty.
does not seem to be confirmed, “Complex” materials containing oxygen appear t o have a higher inhibiting or hastening effect. Examples are negative catalysts such as water, potassium permanganate, potassium dichromate, and sulfur trioxide (the promoter). On the other hand, “ordinary” oxides such as vanadium pentoxide, phosphorus pentoxide, and arsenic trioxide appear to be relatively inert. The more strongly inhibiting materials may dissociate more readily or to a higher degree in fuming acid than the others and thus attain a higher effective concentration. Another basis for theoretical interpretation is the effect on the polarity of the solvent media. The influence of water is confirmed in the tests with salts containing water of crystallization; the water seems to become dissociated from the salt, since it exerts an independent influence in retarding the reactions. The accelerating effect of sulfur trioxide is also evident from a comparison of the rates in fuming sulfuric acid n-ith those in 96 per cent acid. Gleason and Dougherty (21) showed the effects of water as an inhibitor belox 100 per cent sulfuric acid concentration. The energy of activation, E , may be evaluated in calories from the evpression :
Gleason and Dougherty recorded the decomposition rate of o-benzoylbenzoic acid in sulfuric acid a t concentrations of 100 per cent and lower. Their data were reduced to a plot of the life fraction os. the reaction velocity constant, k. The two curves for 65” and 78” C. resulted from a least squares correlation of these data; the respective equations are: k = 0 0738 k = 0.2595
TABLE V.
PROMOTER TESTRUNS
TEMPERATURE, 25’ C. ?fold-4ddition Rate Addition Rate Ratioa ity of Agent k X 103 kr/kio Agent Agent k X 103 None l o d 0 . 6 7 4 1.00 s - o n e 3 2 d 0.762 CuSOi 0.369 0.548 0:33 .LszOs 0.620 cc14 0.362 0 . 5 3 7 0.33 C U S O ~ 0.651 K2C1.20. 0.237 0 . 3 5 1 0.36 PzOh 0.651 Acetic anhydride 0 516 TEMPERATURE, 65’ C. M ol alAddition Rate Addition Rate Ratioa i t y of Agcnt ii X 10 Agent k X 10 k r / k i n 4gent S o n e 31 d 0 , 6 0 3 None 1 6 d 0.740 1.00 ... V2Os
0.507
AS103
0.632
P?O6 TICls NiCh.6H20 AICh
0.698
0.720
0.29 0.27
0.519 0.700 0 537 0 . 5 7 3 0.573 0.728
0.17 1.54 0.40
0.480 0 . 6 5 0
v205
AsnOs TlCh KaiSOb
0.37
TEXPERATURE,
hlolal-
Addition Agent None 446 cU?ch
688
10-j
with a n average slope of 5.10 X
KMnOl HaBOa
E , the “critical increment’’ or energy of activation, is defined by Tolman (63) as the “difference between the average energy of the molecules and the modes of electromagnetic vibration which actually take part in the reaction, and the total average energy of these same elements, whether in reactive condition or not.” Upon this basis it would be expected
+ (t) X + 27 26 94 ( t ) X
Cu2C12,
-
~H?O HC1
Rate IC
1.00 0.187 0.47 0,281 0.70 0.279 0.70
O:i67 0.43
0.94 0.48
0.465 0.15b
0.400
0.376 0.192
a ICi is reaction b Estimated. C
d
Ration ity of kr/kaa Agent
0.13
velocity oonstant
78’
c.
0.484
0..573 0.573 0 532
LIolalRation ity of JcI/I:3~ Agent 1.00 ... 0.81 0 85 0.8J
0.13 0.17 0.18
0 08
0 34
Ratioa
1Iolality of Agent
IcI/I;31
1.00 0.80 0.95 0.95 0.88
0:ik 0.133 0.085
0.185
AIoIal-
Addition Rate Ratio“ ity of Agent I; kdlcae Agent None 56d 0 . 3 5 9 1.00 ... Anthraquinone 0.222 0.62 0.25 None 56‘d 0 . 1 9 ~ 1~. 0 0 ... Anthraquinone 0 . 1 2 4 C 0.61 0.278 0.09lc 0.46 0.556 Same 0 . 0 6 3 c 0.32 1.112 Rame f o r the run w i t h o u t additive.
At 70’ C. Run number.
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
TABLEVI. H&Oa,
%
EFFECTOF SULFURIC ACID CONCENTRATION ON TEMPERATURE COEFFICIENT (75-85" C.) Molality
1.13 (H20)
98
100 100.4
101.8 103.2 104.5 105.4 105.9
0.00 (HzO)
0.233 1.086 2.033 3.122 3.944 4.434
(Sod
(so8) (SOs) SOS)
/so8) (SOS)
SOs
%
.. ..
1.8 8.0
14.0 20.0 24.0 26.2
k~/k~ 2.94 2.92 2.97 2.96 2.98 2.84 2.74 2.70
Log loz/kl
E, calories
0.4691 0.4649 0.4732 0.4707 0.4743 0.4528 0.4381 0.4309
26,700 26,400 26,900 26,800 26,900 26,800 24,900 24,500 26,100
Av.
A surprising point brought out is that in both sets of time data the reaction rate appears to increase as the reaction proceeds. This trend was not found in any of the present experiments; the runs consistently demonstrated a decrease in velocity constant as the reaction proceeded. Of the ten points of Gleason and Dougherty a t 65' C. and the ten at 75', only two of each group were taken below the half-life decomposition point. At 75' C. the reaction rates varied from the mean by *1.68 per cent, and a t 85" by an average of *2.96per cent. The retardation in the present work i s attributed to the inhibiting influence of the anthraquinone formed during the reaction. To express the retardation mathematically, the data were reduced to the following equation, in which the logarithm of the over-all reaction velocity constant k is a logarithmic function of the initial sulfur trioxide concentration less rnx times the log of kl, the initial reaction velocity. I n these expressions 2 represents the life fraction, and rn is the slope or a proportionality constant for the variation of k with increase in x and t: log k = 6 log CsoS- mx (log ICO)
Solving for m,
log k = log ko - mx (log bo) logk = l o g k o (I - mx)
(3)
All the data may be correlated in this manner for engineering use (Table VII). TEMPERATURE COEFFICIENT.According to Gleason and Dougherty's interpretation of their work, the temperature coefficient was more than 3. An average of their data gives 3.22 for the 85-75' C. interval and 3.41 for the 75-65' C. interval. Although the present study indicates a temperature coefficient about 8 per cent lower, the agreement is reasonably satisfactory. The difference is probably due to experimental technique and t o a different degree of accuracy in temperature control of the reactor flask. (A difference of 0.1' C. affects the reaction rate by approximately 3 per cent.) The present velocity constants obtained under conditions comparable to those of Gleason and Dougherty agree well with their values. The former are consistently higher. VELOCITY CONSTANT AT VARIOUS LIFE FRACTIONS
Since the reaction was retarded as it proceeded, the reaction velocity constantk was determined for the start of thereaction (zero time), the half-life point, and the 95 per cent decomposition point; the last is useful for commercial application. These values are important in analyzing the inhibiting effects of anthraquinone and water produced by the reaction, and in evaluating the inhibition induced by other materials. The usual k values were multiplied by the corresponding experimental values of t. The corresponding life fraction was obtained from a plot of kt against 2, derived from the unimolecular kinetic decomposition formula, IC = 2.303 log [ a / ( a - x ) ] / t
(1)
Since ko = q5 Cso,and log ko = q5 log CBOs, the latter expression can be substituted in Equation 1 to obtain:
221
1
m = X- [ ~ -
From kt, the life fraction z is obtained and plotted against k on a semilog plot; a line is drawn through these points t o show the best values for the runs a t an initial oleum or sulfuric acid concentration and the corresponding reaction temperature (Figure 2). PROPOSED REACTION MECHANISM
TABLE VII.
Acid strength, % 808 Log IC (750 C . ) Log ko (75' C . ) Log k/log ko 1 (logk/logko) m (750 C . )
-
= ma!
Log k (85O C . ) Log ko (85' C.) Log k/log ko 1 (log k/log ko) = m (85' C.)
-
Log k (750 C.) Log ko (75' C . ) Log k/log ko 1
- (lo
m (750
8,IC/logko)
Log k (85' C . ) Log ko (85' C . ) Log k/log ko 1 - (lo k/logko) m (85'
6.)
0
3
-
Per cent Hd304.
wm
mz
REDUCTION OF DATATO m IN THE BASICEQUATION: m = (1,'s) (1-log k/log ICo) Molality of Water Molality of Sulfur Trioxide 1.13 0.0 0.233 1.086 2.033 3.122 3.944 98Q 100a 1.8 8 14 20 24 Life Fraction, z 3 0.50 1.8865 2.06595 2.10755 2.1196 2.1235 2.1741 2.2164 1.9186 2.1035 2.1265 2.1517 2.1556 2.2084 2.2528 0.984 0.981 0.99 0.985 0.986 0.985 0.982 0.016 0,019 0.01 0.015 0.014 0.015 0.018 0.032 0.038 0.020 0.030 0.028 0.030 0.036 Averagem = 2.3483 2.5238 2.5647 2.5752 2.5781 2.6120 2.6365 2.3766 2.5557 2.5933 2.6031 2.6064 2.6401 2.6646 0.989 0,995 0.989 0.989 0.989 0.989 0.988 0.011 0.005 0.011 0.011 0.011 0.011 0,012 0.022 0,010 0.022 0.022 0.024 0.022 0.022 Averagem = 1.8573 1.9286 0.967 0.033 0.0348
Life Fraction. x = 0.95 2.0386 2.0763 2.0888 2.1035 2.1265 2.1517 0.967 0.975 0.971 0.033 0.025 0.029 0.0348 0.0264 0.0306
2.3222 2.3766 0.980 0.020 0.021
2.4991 2.5557 0.978 0.022 0.023
2.5366 2.5933 0.977 0.023 0.024
2.0917 2.1467 2.1937 2.1556 2.1861 2.2528 0.971 0.971 0.973 0.029 0.029 0.027 0.0306 0,0306 0,0284 Averanem = 2.5478 2.5502 2.5922 2.6212 2.6031 2,6064 2.6401 2.6646 0.979 0.978 0.981 0.984 0.021 0.022 0.019 0.016 0.022 0 023 0.020 0.017 Average m =
4434 26.2 2.2413 2.2792 0.983 0.01'l' 0.034 0.031 2.6513 2.6794 0.99 0.01 0,020 0.019
Gleason and Dougherty proposed a mechanism for the decomposition of o-benzoylbenzoic acid by sulfuric acid in two stens: ....
Fast:
2.2098 2.2792 0,970
OR
f &SO4
*
0
0.030
0.0315 0.031 2 6263 2: 6794 0.982 0.018 0.019 0.021
OHSOs
0 June, 1943
INDUSTRIAL AND ENGINEERING CHEMISTRY
689
mechanism is based on the fact that it predicts an increased rate of anthraquinone formation by increasing the sulfur trioxide concentration, because extra sulfur trioxide will push the equilibrium represented in step I to the right. This follows when it is assumed that the speed of step I is the controlling factor in the over-all rate of ring closure. Hammett and co-workers made a similar assumption ($6). Since this mechanism is complex and no ready means exist t o measure the intermediate ionic concentrations, it is not now possible to reduce the reaction phenomena observed to a quantitative basis. O li
COOH
0 4 4
1
+ /
I
(large excess)
.os
0.0
0.2
0.4
0.6
0.8
1.0
0.8
1.0
0-C=O
>
k
3
.04
w
> .03
.02
I
I
I
I
0.0
0.2
0.4
0.6
I
LIFE FRACTION Figure 2 . Decomposition Rate VS. Life Fraction of Reaction a t 7 5 " C . (ahove) and a t 85 O C . (below)
Slow:
COMPARISON WITH OTHER ORGANIC ACIDS
Figure 3 compares the effects of corresponding molalities of water and of sulfur trioxide upon the rate of three organic acids with sulfuric acid. The unusual feature for o-benzoylbenzoic acid is the behavior of sulfur trioxide as a positive catalyst, particularly in the fuming acid concentrations above 14 per cent:
OHSOs
Their velocity measurements did not disclose the slowing down of the reaction because the speed of reaction is relatively insensitive to the slight amount of water formed. Inasmuch as the decrease of k as shown in Table I occurred in all these experiments, the fast reaction (postulated because the rate did not decrease) receives no support here. The proposed intermediate reaction may take place but is probably not instantaneous (nor the governing one). The reaction rate of o-benzoylbenzoic acid, compared with other organic acids studied, is relatively insensitive to change in water concentration, and is even less affected by variation in sulfur trioxide concentration. Kewman has presented a plausible ionic mechanism for the formation of anthraquinone from o-benzoylbenzoic acid (46)in steps I, 11, and I11 which follow. Step I is the only one for which experimental evidence has been obtained, but the steric and ionic behavior advanced give credence t o steps I1 and 111. Further support for this
690
Organic Acid
Water
Oxalic Malic (and citric) o-Beneoylbeneoic
Very strong Very strong Quite weak
Inhibiting Power Sulfur trioxide
7
Weak, if at all Very strong Acceleration (not inhibition)
Because organic oxygen compounds are highly ionized basically in strong sulfuric acid, it is possible that acid-base catalysis is effective. Assuming that hydrogen-ion concentration is related to reaction velocity, the kinetic data on o-benzoylbenzoic acid indicate an increasing acidity in fuming sulfuric acid. An interesting question is the validity of this correlation above 100 per cent sulfuric acid. To decide whether hydrogen ion concentration is important, it should be measured in fuming sulfuric acid. Because, however, of the kinetic features of the oxalic and malic acid decompositions, other phenomena seem to predominate, and these reactions are not governed by simple ionic factors. For comparison of the kinetics of o-benzoylbenzoic acid with other decompositions up to 28.8 per cent fuming sulfuric acid, the previous data are limited. Of the five organic acids, only the oxalic acid study covers a sufficient range to permit general comparisons. (However, for citric and malic acids, the reaction in fuming sulfuric acid a t 14 per cent sulfur trioxide approaches zero rate a t room temperatures.) In Figure 3 the ortho acid rate remains nearly constant from
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
approximately 2 to 14 per cent fuming acid; 14 per cent seems to represent the point of maximum decomposition velocity for oxalic acid. Above 14 per cent sulfur trioxide, however, a marked difference is noted: the oxalic acid decomposition is probably slightly retarded, while that of o-benzoylbenzoic acid is definitely accelerated by additional sulfur trioxide. To explain these phenomena molecularly, the oxalic acid behavior up to the 14 per cent point is due to the decreased water (a very strong inhibitor here) in the fuming acid by the increase of sulfur trioxide. Above 14 per cent sulfur trioxide the reaction is that of oxalic acid with sulfuric acid, slightly inhibited by sulfur trioxide in accord with Taylor’s theory of negative catalysis. To rationalize the exceptional o-benzoylbenzoic acid, a new case may be postulated in that the predominant reaction above the 14 per cent point is caused by sulfur trioxide or by pyrosulfuric acid. No inhibiting complex with sulfur trioxide is effective, so that above 14 per cent sulfur trioxide Taylor’s theory of negative catalysis does not here apply. Below 14 per cent sulfur trioxide the same mechanism may apply for o-benzoylbenzoic acid as for oxalic acid. According t o Figure 3 and Table I11 water is a relatively weak inhibitor for o-benzoylbenzoic acid. An explanation as t o why its curve remains flat to about 14 per cent sulfur trioxide follows: The inhibiting water is furnished by the equilibrium HzSO~ SO3 H20. Owing t o displacement of this equilibrium to the right by temperatures of 75’ and 85’ C., o-benzoylbenzoic acid requires about 14 per cent sulfur trioxide t o overcome the water dissociating a t these higher temperatures. The preceding explanation of the slope of the curves shown in Figures 1 and 3 has been developed about the simplest possible case, assuming that only one mechanism yields the
+
~
%SULFUR TRIOXIDE 84
$lO
2$0
?$O
28J8
product. It is possible, as H. A. Taylor suggests, that several mechanisms or paths may form the product. Because the specificity of the reaction may be only circumstantial evidence for a single mechanism, it is well to consider a multiple path. For example, sulfur trioxide may retard the water-inhibited reaction; this inhibition may, in turn, be masked as a result of acceleration caused by an alternative mechanism. I n such a case the resultant rate curves would show the net effect of combined acceleration and inhibition. There is little evidence to preclude such a possibility. These considerations indicate the present difficulty of complete analysis of nearly all proposed reaction mechanisms and, except for the simplest cases, the need for the development of experimental means t o establish the reaction paths postulated. ACID OR OLEUM STRENGTH VARIATIONS
The reaction rates a t various life fractions (Figure 2) disclose a consistent reduction as the reaction proceeds. Thus the question arises as t o whether this retardation can be attributed solely t o the water formed. The effect of water only was noted by a plot of the velocity constant a t zero time us. initial water concentration (or its equivalent sulfuric acid and oleum strength). This analysis showed that the water formed accounts for only a small part of the inhibition; the presence of the anthraquinone formed in molal concentra-
TABLE VIII. ANTHRAQUINONE AS INHIBITOR Relation of Reaotion Rate k t o Rate without Inhibitor, ko Initial molality Relative Temp., Conoen- of anthrarate C. tration quinone k/ki 0.254 0.62 70 0.278 0.621 70 0.458 0.656 70 1.112 0.254 70 0,250 0.62 78 0.612 0.278 70 0.462 0.656 70 1 112 0.318 70
Comparison with Effect of Water
&j2$y of
inhibitor 0.25 0.4 0.6 0.8 1.0 1.1
Relative rate k / k o Anthraquinone Water 0.62 (0.97) 0.92 0.51 0.83 0.43 0.77 0.36 0.72 0.32 0.70 0.30
tion equal to the water produced shows, in comparison, a definitely greater inhibiting power. T o establish this more definitely, various amounts of initial anthraquinone were added t o the usual reaction mixture of o-benzoylbenzoic acid. The experimental conditions and results are correlated in Table VIII. CORRELATION O F DATA FOR COMMERCIAL RATE PREDICTION
,O 2.0
1.0
e 0
MOLALITY HO ,
1.0
2.0
3.0
4.0
5.0
MOLALITY SO3
Figure 3. Comparison of Effects of Water and Sulfur Trioxide on Decomposition Rate of Organic Acids by Sulfuric Acid June, 1943
From the values for inhibition by water and by anthraquinone, it is possible to predict the over-all reaction rate when using concentrations of o-benzoylbenzoic acid other than those employed in the experiments here reported. The over-all average reaction rate is computed by applying t h e proper reaction rate correction factor t o the initial rate, or the rate of zero time. From the initial concentration, the decomposition point z, to which it is proposed t o run the reaction, and the value of m (for 95 per cent decomposition) shown a t 76” C. to be 0.031 and a t 85” t o be 0.021 in Table VII, the product, mz, is evaluated. The desired value of k is computed from Equation 2. To prove this method, ko values were used t o predict the over-all rates on reactions run above 0.6 life. A comparison of actual values and those computed showed that the proposed method gives predictions with an accuracy of *2.2 per cent. It is possible that some of the materials may act as promoters in sulfuric acid of less than 100 per cent concentration. Other points to be studied are the promoting effects
INDUSTRIAL AND ENGINEERING CHEMISTRY
691
100
50
2t-
10
3
z 3
5
I
W
-II10 .
0,: 80
90
100
110
I20
TEMPERATURE Figure 4 .
- “C.
130
I4 0
T i m e of 0.95 Completion VS. Temperutrcre and Sulfuric Acid Concentration
of sulfur trioxide in higher fuming concentrations (for example, up to 40 per cent) and an evaluation of the economies in producing particular anthraquinone intermediates. Special promise is afforded for processes which use highly concentrated fuming acids to produce sulfonated anthraquinones, since approximately 40 to 65 per cent of the commercially valuable anthraquinone products are derived through sulfonation.
theory and a growth of the matheniatical basis for the extension of the deductions t o phenomena in solutions will aid materially in developing the theory of positively and negatively catalyzed acid reactions. IN ENGII~EERING. The industrial importance of anthraquinone and fifty of its commercially valuable derivatives has grown steadily both in the United States and abroad (66). Table IX shows that from 1927 to 1935 domestic sales increased 102 per cent in gross volume; the total sales value increased 62 per cent, since the average price of anthraquinone compounds decreased from $1.42 per pound in 1927 to $1.14 in 1935. Certain of the chemical data of this study are correlated with the engineering and process economics in the production of anthraquinone. The materials are sulfonated intermediates derived from anthraquinone by sulfuric acid treatment. I n fact, sulfonation is the first step in the preparation of most anthraquinone intermediates, including the important dyestuffs of the alizarin and indanthrene classes. Since more than three fourths of the useful derivatives of anthraquinone are obtained through the sulfonic acids, sulfonation is the most important chemical operation in the production of anthraquinone intermediates. The nitro compounds produced from mixed acid are also important. I n nitration reactions, the economic proportions of nitric and sulfuric acid vary with reaction and with market conditions (82). The nitro compounds may be reduced t o the corresponding amino materials. The hydroxy derivatives of anthraquinone are obtained from the sulfonated anthraquinones. An important example is alizarin, which results from the fusion of “silver” salt (anthraquinone-2-sulfonic acid) with sodium hydroxide.
TABLEIx. PRODUCTION AND SALE O F AKTHRAQUINOSE AXD DERIVATIVES I N THE UKITED STATES (66)
--
Year
ProductionPounds 2,905,785 3,587,766 4,917,296 4,031,946 3,532,834 4,288,381
1927 1928 1929 1930 1933 1934 1935
Sales
, -
Pounds
Value
2,423,960 3,119,852 4,188,501 5,216,402 3,705,976 4,158,870 4,894,117
$3,440,432 4,266,603 6,684,812 5,216,402 4,635,688 4,834,488 5,572,187
7
Value/L b. $1.42 1.37 1.35 1.29 1.n9 1.16 1.14
APPLICATIONS
INCHEhfIsTRY. Further use of these new data may arise in the estimation of the chemical strengths of concentrated inorganic acids; that is, the extent of ionic dissociation and activity may be deduced from kinetic determinations. For checking and extending the theory, use might be made of these data, particularly if o-benzoylbenzoic acid proves reasonably ideal in this respect. Future studies like the above on substituted o-benzoylbenzoic acids, as well as effectively devised dipole determinations of the relative electronic displacements may be in order. At present, however, the results of such work on water and simple solutes seem baffling, and the conclusions are, in general, based upon very broad assumptions (88). The effect of definite added amounts of concentrated hydrofluoric acid will more fully indicate the status of this reaction as a case of acid-base catalysis. Boric acid serves as an industrial catalyst, but its action has not been explained. Related to this, Hantzsch’s determinations have shown that in sulfuric acid, organic oxygen compounds (0-benzoylbenzoic acid and anthraquinone) are ionized strongly as bases. Extension of the work in this field requires careful analysis of the existing results, as well as the development of kinetic data on other organic acids in the range of sulfur trioxide concentrations. I n addition, the physicists’ development of atomic 692
INTERMEDIATE SYNTHESIS.This work has considerable importance for optimum manufacturing practice since it demonstrates the economic possibilities of utilizing special oleum concentrations for the condensation of o-benzoylbenzoic acid to anthraquinone as the first step in synthesizing various desired dye intermediates. By choosing the proper initial oleum concentration, it is possible to eliminate expensive and time-consuming process steps, such as diluting, cooling, recrystallizing, filtering, washing, purifying, drying, and then redissolving in oleum the intermediate product, anthraquinone. I n eliminating these steps, it is possible to proceed almost directly with the synthesis of the desired sulfonated derivative. The concentration of fuming sulfuric acid given in Table X can be computed back t o the initial stoichiometric proportions of o-benzoylbensoic acid and fuming sulfuric acid. These results can be combined with a suitable choice of reaction time and temperature (Tables X I and XII) t o obtain economic yields (95+ per cent) of anthraquinone prior to adjusting the process to conditions for the derivative sought. As the basis for Tables XI and XII, the derivations used to correlate the effects of temperature and of acid concentration are developed as follows: The velocity constants a t the
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
INTERMEDIATES DERIVED FROM ANTHRAQUINONEBY SULFURIC ACIDTREATMENT TABLE X. SULFONATED Product Anthraquinone-1sulfonic acid
Anthraquinonedsulfonic acid
Use To prepare 1-aminoanthraquinone
Fuming SOs Strength 20% , 3 0 8 ( 1 : l b y wt,) 0.38 of 40%
+
so8
Temp., Time, OC. Hr. Catalyst 50-135 1 HgO or HgSOa (1% Hg) 135
To prepare alizarin, purpurin, 2-aminoanthraquinone
25% SOs (1: 1 by wt.)
+ 0.4 of 40%
2
-1 4
135 135 140
SOa
l'/a
3
None
1
5=/2
1.5-Anthraquinonedisulfonic acid 1 &Anthraquinonedisulfonic acid 2,6-Anthraquinonedisulfonic acid 2,7-Anthraquinonedisulfqnic acid uinizarin purpurin) Anthrarufin
'(0;
To prepare flavopurpurin, alizarin YCA, diaminoanthraquinones
40% SOs ( 1 , l : l by wt.)
130
6
40% SO8 (1.1 : 1 by wt.)
130
5
None
40% SO8 ( 1 . 3 : l b y wt.)
180
4
None
120
4
None
..
Boric acid, 1 part
To prepare anthrapurpurin, alizarin SC, 20% SO8 (1.2: 1 by wt.) diamino compounds To prepare aliaarin saphiron-B and Se Conc. or fuming HzS04 (20: 1 by wt.) (Bolvay) blue 30% SO8 (100:5 by wt.), 5-7 atm. pressure
260-80
100 36
HgO or HgSOa (1% H g )
Boric acid, 2 parts
sulfonation, when anthraquinone is sought, particularly in sulfuric acid concentrations below or slightly above 100 per cent. Figure 4 is derived from Tables XI and XI1 and sum-E marizes the effect of temperature and sulfuric acid concenlogk = 2.303 X 1.987 + tration on the time required for 95 per cent yield. I n industrial operations for producing 100 per cent anthraFrom the data in Tables 111,IV, and VI this becomes: quinone, some saving of sulfuric acid and a considerable increase in yield capacity of the reaction vessel can be made -5480 l o g k = - 14.834 through adjustments involving the following considerations: Dependent on the stability and power of the mixer drive, as much o-benzoylbenzoic acid should be added as the equipThe latter formula is the basis for the calculated specific reacment will safely carry. I n too concentrated a mixture the tion rates in 100 per cent sulfuric acid given in column 3 of mass tends t o become pasty, and the removal of sulfuric and Table XI. Column 4 shows the times of 95 per cent yield unreacted o-benzoylbenzoic acid from the product becomes to.^^), computed through the unimolecular equation from difficult. If i t is desired to obtain another intermediate, the which, to.sJ= 3.00/k. maximum required concentration of o-benzoylbenzoic acid is 4.4 molality, and the minimum is 0.24, for quinizarin and anthrarufin end products, as shown in Table X. I n setting TABLE XI. EFFECTOF TEMPERATURE ON TIMEFOR PRODUC- up processes, a stoichiometric balance should be made of the TION OF 95 PERCENTYIELDSOF ANTHRAQUINONE FROM O-BENsulfur trioxide concentration t o be maintained in consequence ZOYLBENZOIC ACID IN 100 PERCENTSULFURIC ACID of the water released in the condensation of o-benzoylbenzoic Temp., C . Log k X 104 k X lo4 f 0 4 6 , Min. acid. 85 2.532 340 88.3 higher temperatures to be used in commercial operation are computed from the following integrated equation:
('T ')
+
O
90 100
120 140
2.719 3.186 3.904 4,548
524 1,533
8,010
35,250
57.3 19.6 3.74 0.85
TABLE XII. EFFECTOF SULFURIC ACID CONCENTRATION ON TIMEFOR EQUIVALENT (0.95) COMPLETION RELATIVE TO TIME FOR 100 PERCENTSULFURIC ACID Relative time
--% 85 34 4
HB04-
90 95 11 1 3 0
100 10
--%
5
0 92
Free SOa15 25 28 0 89 0 71 0 63
Table X I shows the marked influence of temperature on the rate of anthraquinone formation. From the viewpoint of industrial practice, anthraquinone is readily formed, prior to sulfonation at relatively high temperatures during the warming-up period, in the concentrations of fuming sulfuric acid employed (Table X). Table XI1 shows the reduced time required by fuming acid t o reach 95 per cent completion, and indicates the desirability of using sulfuric acid a t 100 per cent concentration (or slightly higher) t o complete quickly and economically the production of 95 per cent or more anthraquinone. Table XI shows that operating temperature is a primary rate factor. Temperatures as high as 120"C. may be employed without appreciable loss of anthraquinone through June, 1943
CONCLUSIONS
1. The reaction is unimolecular and the yields of "infinity" experiments are practically quantitative. 2. Of twenty-two substances tested, the specific and only positive promoter for the reaction is sulfur trioxide. 3. No maximum point of decomposition velocity occurs in fuming sulfuric acid, and the rate is continuously catalyzed by additional sulfur trioxide up t o the maximum concentration employed, 28.8 per cent fuming acid. 4. The products of the reaction are water and anthraquinone; the inhibiting effect of the latter on reaction rate is more pronounced than the effect of water. 5. The ionic mechanism recently proposed for the formation of anthraquinone from o-benzoylbenzoic acid appears t o be plausible in view of the rate findings of this study. The complexity of the reaction, however, prevents a reduction of the kinetic phenomena t o final terms owing to the lack of methods for the detection and measurement of large organic ions in solution. It is possible t h a t spectrographic or similar studies might yield useful ionic data on these points. 6. Data on the amount of inhibition caused by water and by anthraquinone 'are submitted as the means of predicting the velocity of anthraquinone formation under operating conditions differing from the experimental work.
INDUSTRIAL AND ENGINEERING CHEMISTRY
693
BIBLIOGRAPHY (1) Balandin and Nesvizhskiii, Sci. Communic. (Moscow State UnC.), 2 , 233 (1934) ; Chem. Zentr., 1935, 11, 1528. (2) Behr and Van Dorg, Ber., 7, 578 (1874). (3) Boeseken and Max, Rec. trav. chim., 48, 486 (1929). (4) Brauss, Bayer, and Fieser, Ann., 459, 287 (1927). (5) Bredig and Lichty, 2. Elektrochem., 12, 459 (1906). (6) Brochets, Compt. rend., 117, 115 (1893). (7) Burk and Daus, J. Phys. Chem., 35, 1461 (1931). (8) Chemical Construction Co., Handbook, 1939. (9) Chalmers, Can. J. Research, 7, 422 (1932). (10) Corson and Ipatieff, Org. Syntheses, 17, 36-7 (1939). (11) Cox, J. Chem. Soc., 119, 142 (1921). (12) Deutsche Hydierwerke A.-G., Frenoh Patent 809,128 (1937). (13) DeRight, J . Am. C h m . Soc., 55, 4761 (1933). (14) Dimroth, Ann., 377, 127 (1919). (15) Dittmar, J . Am. Chem. Soc., 52, 2746 (1930). (16) Dittmar, J.Phys. Chem., 3 3 , 5 3 3 (1929). (17) Dolique, Bull. SOC. chim., [5] 2, 1489-91 (1935). (18) Dunstan, Birch, Fidler, Pim, a n d T a i t , Oil Gas J.,37, 49 (1938). (19) Fischer, E., Ber., 26, 2400 (1893); 27, 615 (1894). (20) Gibbs and Conover, U. S. Patent 2,143,493 (1939). (21) Gleason and Dougherty, J . Am. Chem. SOC.,51, 310 (1929). (22) Groggins, “Unit Processes in Organic Synthesis”, pp. 10-58 (1935). (23) Gutyrja and Brimitzkaja, Azerbaidzhanskoe Neftyanoe Khoz., 17, NO.-l, 50 (1837). (24) Halban and Heoht, 2.Elektrochem., 24, 65 (1918). (25) Halban and Kersch, 2. physik. Chem., 82, 325 (1913). (26) H a m m e t t , J. Am. Chem. Soc., 59, 1708 (1937). (27) Hantzsch, 2. physik. Chem., 61, 260 (1908). (28) H a r t m a n n , W. W., private communication. (29) Hawkins, J. Chem. SOC.,121, 1170 (1922). (30) Hinshelwood, “Kinetics of Gaseous Reactions”, 1933. (31) Huckel, Liepziger Vortrage, 1927, 83. (32) Hurd, IND. ENC.C H E W ,26, 50 (1934). (33) Ingold, Chem. Rea., 15, 225 (1934).
(34) Ipatieff, “Catalytic Reactions a t High Pressures and Temperatures”, New York, Macmillan Co., 1936. (35) Ipatieff, Corson, and Pines, J. Am. Chem. Soc., 58, 919 (1936). (36) Ipatieff, Orlov, and Petrow, Ber., 60, 1006 (1927). (37) Ipatieff, Pines, and Schmeding, J . Am. Chem. SOC., 60, 353 (1938). (38) Kendall and Cargentes, Ibid., 36,2498 (1914). (39) Klever and Fritz, Mitt. chem-tech. Inst. tech. Hochschule Karlsruhe, 1, 1 (1923). (40) Konigs, Ber., 23, 3145 (1890). (41) Kraemer and Spilker, Ibid., 23, 3276 (1890). (42) Liebermann, Ibid.,7, 805 (1874). (43) Michael and Hilbert, Ibid., 41, 1080 (1908). (44)Moelwyn-Hughes, “Kinetics of Reaotions in Solution”, 1933. (45) Moelwyn-Hughes and Hinshelwood, Proc. Roy. SOC.(London) A131, 186 (1931). (46) Newman, J . Am. C h m . Soc., 64, 2324 (1942). (47) Platonow and Tomilov, J . Gen. Chem. (U. S . S.R.), 8 (70), 354 (1938). (48) Reid a n d Theriault, J. Phys. Chem., 35, 673, 950 (1931). (49) Rice, F. O., and Rice, K . K., “Aliphatic Free Radlcals”, 1935. (50) Roseveare, J . Am. Chem. Soc., 53, 1651 (1931). (51) Schierz, Ibid.,45, 447 (1923). (52) Senderens, Compt. rend., 184, 856 (1927). (53) Ibid., 194, 809 (1932). (54) Ibid., 196, 979 (1933). (55) Smith and Cass, J . Am. Chem. SOC.,54, 1603, 1609, 1614 (1932). (56) Soper and Williams, J . Chem. SOC.(London), 1931, 2297. (57) Spilker, Ber., 23, 3169 (1890). (58) Standard Oil Co. of Calif., French Patent 799,076 (1936). (59) Stanley, Minkoff, and Youell, U.S. Patent 2,143,493 (1939). (60) Taylor, J. Phys. Chem., 27, 322 (1923), (61) Taylor, “Treatise on Physical Chemistry”, 1928. (62) Taylor and Jones, J . Am. Chem. Soc., 52, 1111 (1930). (63) Tolman, “Statistical Mechanics”, 1927. (64) Union Carbide Co., Brit. Patent 156,162 (1920). (65) Union Carbide Co., French Patent 479,656 (1919). (66) U. S. Tariff Comm., Census of Dyes, etc , 1927-41. (67) Walden, “Elektrochemie Nichtwlissriger LBsungen”, p. 39, Liepzig, J. A. Barth, 1924. (68) Whitford, J. Am. Chem. SOC.,47, 953 (1925). (69) Whitmore, “Organic Chemistry”, pp. 104, 112. 1937. (70) Wiig, J . Am. Chem. Soc., 52, 4729 (1930). (71) Ibid.,52, 4737 (1930). (72) Ibid., 52, 4742 (1930). (73) Wunderly, Sowa, and Nieuwland, Ibid.,58, 1007 (1936). PRESENTED as part of the Symposium on Industrial Reaction Rates held under the auspices of the Division of Industrial and Engineering Chemistry, AMERICAB CHEMICAL SOCIETY, a t Chicago, Ill. Condensed from the dissertation of Charles W. Deane for the degree of doctor of engineering science, New York University.
Removing Crude Phthalic Anhydride f r o m a Condensing Bin a t a Plant of E. I . d u Pont de Nemours & Company, Inc. Phthalic anhydride is not only the base for anthraquinone dyes used for army uniforms and other military purposes, but is a l s o an important intermediate in the manufaoture[of plasticizer for smokeless powder, and the alkyd synthetic resins now widely used in paints for naval vessels, military aircraft, tanks, and other motorized equipment.
694
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 35, No. 6
