M E C H A N I S M OF T H E A M I N E - C A T A L Y Z E D REACTION OF ISOCYANATES WITH H Y D R O X Y L COMPOUNDS ADALBERT
FARKAS A N D PAUL F. S T R O H M
Houdry Process and Chemical Go., M a r c u s Hook, Pa. The kinetics of the triethylamine- and triethylenediamine-catalyzed reaction of phenyl isocyanate with 2ethylhexanol, phenol, hexanethiol, and thiophenol were determined in benzene at 25" C. Both the basicity of the amines in ethyl acetate and their association constants with hydroxyl compounds and thiols were measured in benzene or carbon tetrachloride. Increased basicity of triethylenediamine in ethyl acetate, high association constants for the triethylenediamine-phenol complex, and independence of the rate of the phenol-phenyl isocyanate reaction on the phenol concentration were observed. The results indicate that the amine-catalyzed reaction involves the initial formation of the amine-phenol complex, followed b y the ratedetermining reaction with isocyanate. hydroxyl compounds or thiols.
This mechanism should be applicable to other reactions involving
HE currently accepted mechanism for the tertiary amineTcatal\.zed reaction of isocyanate with alcohols (and related active hydrogen compounds) ( 7 . 2, 3. 27) involves the formation of an amine-isocyanate complex a n d a subsequent reaction of the complex \vith alcohol. in which the urethane is formed and the catalyst is regenerated. According to this mechanism the I ate of reaction is given by the expression
\$here a. b , and c are the concentrations of the alcohol. isocyanate, and amine, respectively, kl and k , are the rate constants for the formation and decomposition of the complex, and k3 is the rate constant for its reactian Lvith the alcohol. There are a number of shortcomings in the kinetic treatment and difficulties arise in applying the mechanism (5.9, 7 7 ) . O n e of the most serious problems is the observation that rate constants k , vary with the nature of the alcohol, although the theory would require that the rate of the formation of the isocyanate-amine complex be independent of the nature of the alcohol with which the complex reacts in the subsequent step. An alternative mechanism in which the first step is the formation of a complex between the amine and the hydroxy compounds was discussed but rejected by Baker and coworkers ( 7 , 2 . 3 ) .even though they observed amine-hydroxyl interaction in the infrared spectra. According to Burkus (5, 6), this mechanism does not provide the activation in the aminecatalyzed reaction of isocyanates \cith alcohols but is operative in the reaction of isocyanates with more acidic reactants such as phenol, thiophenol, and mercaptan with which strong amines form complexes possessing a degree of ionic character. Smith and Friedrich ( 2 3 ) referred to the ionizing effect of amines on thiols to explain the preferential formation of Surethanes in the amine-catalyzed reaction between Z-mercaptoethanol and phenyl isocyanate. Similarly, Iwakura and Okada ( 7 6 ) postulate a n amine-thiol complex as being in32
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FUNDAMENTALS
volved in the mechanism of the amine-catalyzed reaction of isocyanates and thiols. However, Dyer and coworkers ( 8 ) reject a mechanism involving a proton transfer from a thiol to the amines because of steric considerations. Schieler ( 2 2 ) suggested that the base-catalyzed reaction between alcohols and isocyanates is similar to a base-catalyzed aldol condensation and involves an amine-alcohol complex in which a proton from the alcohol is transferred to the amine. This suggestion was made to overcome shortcomings of Baker's mechanism. I n a recent paper: Flynn and Xenortas (73) discuss the mechanism of the amine-catalyzed reaction of phenyl isocyanate with alcohols in connection \vith a study of the catalytic activity of heptamethylisobiguanide and related amines. These authors favor Baker's mechanism on the basis of a correlation beticeen the basicity of the catalysts on the one hand and their catalytic activity and the kinetic deuterium isotope effect on the other. For some time we have endeavored to find more direct approaches to the elucidation of the mechanism of the aminecatalyzed reaction or the nature of the intermediate complex formed by the catalyst. As one approach Jve have attempted to check into the nature of the complex formed between phenyl isocyanate and the amine catalyst reported by Pestemer and Lauerer ( 2 0 ) . According to these authors the bands that appear in the region of 1635 to 1652 cm.-I in solutions of phenyl isocyanate and tertiary amines in hydrophobic solvents are caused by the formation of the complex:
S o such bands were observed in the system phenyl isocya-
nate-triethylenediamine in benzene solution, if care was taken to eliminate traces of moisture ( 7 2 ) . .As another approach the amine-catalyzed reaction of phenyl isocyanate hvith hydroxyl compounds and thiols was studied. in the expectation that the widely different acidity of the reactants \vould sholv a marked change in the rates or kinetics. 'The present paper is concerned v,ith the results of this approach.
Experimental
Materials. Phenyl isocyanate (Eastman Kodak) and 2eth!-l-l -hexanol (Cnion Carbide Corp.) were fractionally distilled under reduced pressure prior to use. Phenol (Baker and .Adamson). thiophenol, and n-hexanethiol (Eastman Kodak) \\.ere triply distilled under reduced pressure. T r i ethylenediamine (DABCO. Houdry Process & Chemical Co.) !vas recrystallized three times from hexane and stored in a vacuum desiccator. Triethylamine (Pennsalt Chemical Corp.) \vas triply distilled over sodium hydroxide. Carbon tetrachloride and benzene (Spectroquality grade. Matheson. Coleman & Bell) Lvere dried over calcium hydride before using. Ethyl acetate a n d p-dioxane (Spectroquality grade, Matheson, Coleman eL Belli and 707~perchloric acid (Baker a n d Adamson reagent) \%.ereused without further purification. Kinetic Experiments. Stock solutions \vere prepared by weighing out the correct amount of material and diluting to the proper level, using volumetric flasks. Reaction mixtures \\.ere prepared by adding the desired amount of catalyst and active hydrogen compound to a volumetric flask. diluting with benzene short of the calibration mark. and adding the proper amount of phenyl i s o c y n a t e solution and then benzene to the mark. l ' h e reaction flasks Lvere stoppered with rubber caps, so that samples could be 1vithdrav.m a t the proper time interval with a hypodermic syringe. T h e flasks were immediately placed in a water bath controlled a t 25' C. Samples uithdraivn from the reaction flasks were placed in a 0.1-mm. cell (Connecticut Instrument Corp., Type FT) using Irtran-2 crystals. ?-he infrared spectrum was scanned in the 4.5-micron region using a Perkin-Elmer Infracord, Model 137 spectrometer a n d when the isocyanate peak \vas reached the time was noted. Determination of Catalytic Constants. T h e activity of various amine catalysts was determined by measuring the second-order rate constants a t equimolar concentration (0.07,W) of 2-ethylhexanol and phenyl isocyanate for the uncatalyzed (k,) a n d the catalyzed ( k ? ) reactions. T h e catalytic constant, h,. is then defined by k , = ( k 2 - k,) IC,where c = catalyst concentration. Determination of E, '2 Values. T o obtain information on the basicity of various amine catalysts in nonaqueous systems, the procedure of Hall ( 7 . f ) was used. Potentiometric titrations were performed using 50 ml. of ethyl acetate in a 250-ml. beaker provided with magnetic stirring. T h e quantity of amine was such as to require ca. 1.0 ml. of 0 . 5 5 perchloric acid (made from 727, perchloric acid in p-dioxane) for neutralization. ?'he standard buffer solution was prepared by weighing out 1.5000 grams of diisobutylamine, and making it u p to 250 ml. with ethyl acetate in a volumetric flask, then adding 5.00 ml. of 0.5.1' perchloric acid in dioxane. A Beckman Model G p H meter was used for the titrations. After standardizing (buffer = 170 m v . ) , a potentiometric titration was performed: a plot of E (millivolts) us. milliliters
Table I .
Catalyst Activity, pK,, and Basicity in Ethyl Acetate (fl,*) for Various Amines
Relatiue Amine
Trieth) lenediamine 2-Methvl-1,4-diaza-(2 2 2)bic\clooctane Trieth) lamine
.\'-EthyIrnorpholine
1 00
pK, 8 60
El 2 124
0 92 0 19 0 05
8 86 10 64 7 15
168 186 291
k, 5600
Activity
5130 500 315
of acid was made, and E l l 2was read directly from the graph as the voltage a t half-neutralization. Association Constants. Solutions were prepared by adding the desired amount of tertiary amine and reactive hydrogen compound to a volumetric flask and diluting with solvent to the calibration mark. \Yhen carbon tetrachloride was used as solvent, the samples withdrawn from the volumetric flask were placed in a 2.5-cm. cylindrical cell (Connecticut Instrument Corp., Type H) using NaCl crystals. Equilibrium concentrations were obtained by measuring the decrease in absorbance of the OH or SH band. IVhen benzene was used as solvent, the samples were placed in a 0.05-cm. KaCl cavity cell (Connecticut Instrument Corp., Type C) a n d equilibrium concentrations were obtained as mentioned above. I n calculating the association constants allowance was made for the fact that triethylenediamine has two amine groups, but for the sake of simplicity. differences between the association constants for the first a n d second amine groups were disregarded. Results
Basicity and Catalytic Activity. I n the past-repeated efforts have been made to correlate the activity of amines in the catalysis of the reaction of isocyanates with alcohols and related compounds with the basicity of these amines as indicated by pK, values. While in a general way higher activity was shown by more basic amines, the correlation was far from satisfactory, especially when amines having various structures were compared Since the pK,'s are determined in aqueous solutions (5, 8) while the isocyanate reactions are carried o u t in nonaqueous systems, we have determined the basicities of certain amines by Hall's technique, using as measure of the basicity the El/ value, the potential a t the half-neutralization point in a potentiometric titration in ethyl acetate. High potential indicates low basicity. Figure 1 illustrates how the catalytic activity varies for different amines in the reactions of 2-ethylhexanol with phenyl isocyanate. Table I lists the pK, a n d E l l S for several amines! together with their catalytic constants for the phenyl isocyanate-ethylhexanol reaction. T h e correlation between catalytic constant (k,) a n d basicity as defined by E l l 2 rather than by pK, is much better than the correlation between the k , and pK, values. Association Constants of Amines with Hydroxyl Compounds and Thiols. T h e association constants were measured in carbon tetrachloride and benzene solution.
Table II. Equilibrium Constant for Association of "Active Hydrogen" Compounds with Triethylamine at 25' C.
Initial Concentration .Component Triethyl~~
Component A Ethyl alcohol
CC1,
Phenol
CCI,
Thiophenol n-Hexanethiol
Solvent
CCI, CCI,
2-Ethylhexanol
Benzene
Phenol
Benzene
VOL. 4
A 0 0043 0 0039 0 0020
0 0020 0 0020 0 0234 0 0234 0 0234 0 0400 0 0400 0 0'00 o 0700 0 0750 0 0750
NO. 1
amtnp
0 0 0 0 0
0 0 0
0 0 0
o
0 0
0226 0121 0590
0284 0513 0100 0200 0400 0400 0600 0700 1400 0750 0500
FEBRUARY
v LL,
Lzterl
Mole
3 0 3 7 85 0 97 0 101 0 3 3 3 4 2 7
3 3 3 2 35 35
1965
5 4 0 0 4 5
33
160
45
14 0
40 120
35 100
n
>
n 30 >
80
25
20
15 40
50
150
100
200
250
120
80
160
200
280
240
300 TIME, MINUTES
TIME,
MINUTES
Figure 2. Triethylenediarnine-catalyzed reactions of phenyl isocyanate at equimolar reactant concentrations in benzene at 2 5 ” C.
Figure 1. Reaction of 0.07M phenyl isocynate with 0.07M 2 ethylhexanol in benzene at 25OC. Catalyzed b y vurious amines (0.001 4 M )
Initial reactant concentration With phenol
1.
Triethylenediamine 2. 2-Methyl- 1,4-diaza(2.2.2)bicyclooctane 3. Triethylamine 4. N-Ethylmorpholine b. Phenyl isocyanate concentration, moles/liter
1.
4. 0.04M 5. O.07M b.
Equilibrium Constant for Association of “Active Hydrogen” Compounds with Triethylenediamine at 25’ C.
Table Ill.
Initial Concentration, .Mole, Lite7 -~ TriethjleneComponent A diamine ~~~
Cornfionent 4
Ethyl alcohol
Sol L ent CCI,
Phenol
CC1,
Thiophenol
CClr
n-Hrxanethlol 2-Cth\lheuanol
CCI,
Benzene
~ _ _ _ _ _ _ _ _ _
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Phenyl isocyanate concentration, moles/liter
enediamine-phenol-benzene system are listed in Table I\-, Once again the strikingly high association constant for phenol is noted, rvhich is appreciably higher than those found for the system triethylamine-phenol. The association constants were calculated under the assumption that the association constant for one amine group of the triethylenediamine molecule remains unchanged when the other amine group becomes associated w.ith a phenol. Since this assumption is probably not valid and the association constant for the first phenol is higher than that for the second. it is very likely that the fraction of free triethylenediamine molecules is, lower than indicated by the association constants given in Tables 111 and IV. T h e association constants for the thiophenol and hexanethiol
The data obtained in carbon tetrachloride for the association constants of triethylamine ivith ethanol. phenol, n-hexylmercaptan, and thiophenol are shotvn in Table 11. The values for the systems ethanol-triethylamine and phenol-triethylamine are in satisfactory agreement with those obtained by Barrov and Yerger (-I), Davis (7), Joester and Drago (77), and Julg and Bonnet ( 7 8 ) . Table I1 also shotvs the associatiop constants of triethylamine \vith 2-ethylhexannl and with phenol measured in benzene solution. The constants for phenol are appreciably higher than for the alcohols or thiols. T h e association constants for triethylenediamine \vith active hydrogen compounds in carbon tetrachloride and benzene are summarized i n Table 111. \vhile the data for the triethyl-
34
0.035M
2. 0.07M With 2-ethylhexanol 3. 0.02M
FUNDAMENTALS
0 0 0 0 0 0 0 0 0 0 0 0 0
0060 0048 0043 0020 0020 0235 0156 0235 0400 0400 0400 0’00 0700
0 0 0 0 0 0 0 0 0 0
0 0 0
0054 0042 0069 0123 0044 0120 0080 0120 0200 0400 0600 1400 0700
Equilibrium Concentration. Equir alerif Lzler TriethjleneComponent ‘4 Comp1e.u diamine 0 0056 0 0004 0 0105 0 0045 0 0002 0 0081
0 0 0 0 0 0 0 0 0
0 0
0040 0004 0010 0214 0145 0214 0391 0368 03’3 0352 0482
0 0003 0 0016 0 0010 0 0021 0 0011 0 0021 0 0009 0 0032 0 0027 0 0348 0 0218
0 0 0 0 0 0 0
0135 0230 0078 0219 0149 0219 0391 0 0768 0 1173 0 2452 0 1182
K. LitersiAiSole 6 9
8 5 174 128 4 5
4 0 1
0 4 3
2 6 0 0 5 1 5 59 13 62 04 84
Table IV.
Equilibrium Constant for Association of Phenol with Triethylenediamine in Benzene at 25' C.
Initial Concrritrotlon .I l o l r ~ Liter ~ E,H,OH rriufh~lenr~znrriirze -
0 0'51 0 0-50 0 1000 0 1000 0 0-00 0 0200
0 0751
0 0600 0 0'51 0 1000 0 0350 0 009'
Equilibrium Concentration, EquivalentlLiter TriethyleneCeH,OH Complrx diamine 0 0090 0.0661 0.0841 0 0160 0 0590 0.0610 0.0815 0 0185 0.0687 0 0065 0.0935 0.1065 0.0455 0 0245 0.0245 ~-
-
-
~~
~
~~
0 0114
are not highcr than those for the alcohols, in spite of the fact that in a q w o u s .;olutionj thiols are more acidic than the alcohols. Kinetic Studies. In the reaction of phenyl isocyanate {vith equimolar concentrations of alcohols or other active hydrogen compounds. second-ordrr kinetics require proportionality brt\\.cen reciprocal isocyanate concentration and time. 'l'his relationship is satisfactorily obeyed in the reaction of 2-ethylhexanol in the presence of either triethylamine (Figure 1 ) or triethylenediamine (see Figure 2), in agreement Lvith observations of Farkas a n d Flynn (70). Figure 2 also s h o w the progress of the triethylcnediaminecatalyzed phenol-phenyl isoc)-anate reaction for equimolar rractant concentrations. This run shows significant deviation from second-order kinetics. l ' h e upward curvature of the plot of the reciprocal concentration zss. time indicates a n order lo\ver t h a n 2 . Actually the reaction follows first-order kinetics, as indicated by Figure 3 > in Lvhich the logarithms of the concentrations are plotted against time. T h e identical slope of the lines representing three runs carried out a t three different equimolar concentrations is a further proof of first-order kinetics. '1.~~ additional 0 runs carried out \iith unequal reactant concentrations are also plotted on Figure 3. I n one r u n the phenol concentration \vas half of the isocyanate concentration ; in the other it was twice as high. T h e points corresponding to these runs folloiv the lines representing the runs carried out with equimolar reactant concentrations, indicating that the rates depend only on the isocyanate concentration and are independent of the phenol concentration. Deviation from the straight line appears in the r u n with the deficient phenol concentration only bvhen most of the phenol has been used up. Second-order kinetics \vas also observed for the reaction with n-hrxanethiol (Figure 4) a n d thiophenol. T h e latter reaction was unusually fast \vith either triethylenediamine or triethylamine as the catalyst, which is contrary to the finding of Iwakura and Okada ( 7 6 ) . Similar kinetic relations \ 1964 ACCEPTED July 29. 1964 Division of Industrial and Engineering Chemistry, 147th Meeting, ACS, Philadelphia, Pa., April 1964.
DIFFERENTIAL THERMAL ANALYSIS AND REACTION KINETICS R O N A L D L. R E E D , ’ LEON WEBER, AND B Y R O N S. G O T T F R I E D
Gulj Research 3 De~elopmentCo., Pzttshurgh, Pa.
Methods a r e presented for the use of differential thermal analysis in the quantitative determination of chemi-
cal reaction kinetic parameters. The theory of Borchardt and Daniels, developed to describe irreversible reactions in stirred systems, is extended, allowing the use of different portions of one or more thermograms in determining activation energies and frequency factors. A widely accepted method devised b y Kissinger is shown to b e incorrect for stirred systems and of questionable value under any circumstances. The Borchardt arid Daniels and the Kissinger methods a r e applied to experimental data, resulting in favorable values for kinetic constants using the Borchardt and Daniels technique; the results of the Kissinger method a r e shown to b e in serious error. The Borchardt and Daniels equations a r e integrated numerically, producing theoretical thermograms which a g r e e well with the corresponding experimental curves. The effects of various pararneters on the thermograms a r e also established b y numerical integration.
differential thermal analysis (DTA) and thermogravimetric analysis (‘PGA) are used for the qualitative characterization of complex chemical reactions, such as the thermal dehydration of clays (79, 20: 25) and the oxidation of crude oil (24). I n the past few years several methods have been suggested for using D‘PA to determine quantitatively the kinetic parameters for certain types of reactions (2. 4. 5, 7 7 , 72). T h e most fundamental approach is that of Borchardt and Daniels (1: 5). where the DT.4 method is modified in such a way that the experiments are accurately described by a simple theory. This paper develops additional theoretical aspects of the Borchardt and Daniels method. Several results obtained with this method are compared with results obtained by follo\ving the method of Kissinger ( 7 7 , 72) and by conventional kinetic studies. OTH
Description of Method
Detailed descriptions of equipment and techniques are not given here, since there is a n abundance of literature on the subject (73: 76-78. 22). Although specific designs of DT.4 equipment vary \videly. depending on what is to be studied, most share the general features which follow. Briefly. the reactive substance to be studied is placed in a sample cell and an inert substance is placed in a reference cell. The inert substance is chosen so that its heat capacity and thermal conductivity approximate those of the reactive material. Both cells are then immersed in a heat bath, and the reaction is initiated by supplying heat to the bath in a prescribed fashion. For simplicity. this is usually done in such Present address. Department of Mechanical Engineering, Drvxel Institute of Technology, Philadelphia, Pa. 38
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FUNDAMENTALS
a \\.a)- that the temperature of the reference cell, T,? increases linearly \vith time. I n the absence of reaction, the sample temperature, T ! \vi11 also increase linearly and equal T,. \Vhen reaction commences in the sample cell, the heat liberated (or absorbed) causes 7‘ to differ from T,. This difference, A T = 7’ - T,? is recorded by means of a differential thermocouple. T h e actual reference temperature, T,. is simultaneously recorded. l h u s : the outcome of the experiment is a graph of AT LIS. time, t , as shown in Figure 1. T h e rate of reaction is not measured directly, but must be inferred from the D T A curve by suitable theoretical analysis. Theory of Borchardt and Daniels
When the heat bath consists of a metal block and the contents of the cells are a t rest. there is a temperature gradient betlveen the walls of the cells and their centers where the temperatures are measured. \Yhen the appropriate differential equations and auxiliary conditions are specified, this situation is taken into account and the D T X curve A T LIS. t can be computed if the kinetic parameters are known. This procedure was carried through by Sewell (27) for a first-order reaction using a n approximation to the reaction-rate function. T h e result !vas an integro-differential equation ivhich was treated using series approximations. Borchardt and Daniels avoided the mathematical complexity introduced by the finite thermal conductivities of a metal block. stagnant sample. and reference materials by changing the conditions of the experiment (4. 5). They employed a n all-liquid system with a stirred heat bath in which kvere immersed a sample cell containing a stirred reactive liquid and a reference cell containing a stirred inert liquid. LVith such a system the temperatures of reference. sample, and bath are
