Edward C. M. Chen and Stephen L. Sjobergl University of Houston at Clear Lake City Houston, TX 77058
The Kinetics and Thermodynamics of the Phenol from Cumene Process
I
A physical chemistry experiment
The differences between thermodynamics and kinetics are emphasized in every physical chemistry textbook. However, these differencesare seldom illustrated in laboratory experiments carried out on the same reaction ( 1 , 2 ) .~ort-unately, the formation of phenol and acetone from cumene hydroperoxide is a reaction which is ideally suited for lahoratory experiments displaying these differences. In addition, the reaction is a good example of an industrially significant process. The ~ h e n ofrom l cumene wrocess. which Droduces over 2.5 billion pounds of phenol pe; year, is typical of modern petrochemical processes which have been developed over the past thirty years. There are three hasic reactions in the process, each of which is a verv interesting example - of a different tvwe .. of catalysis. T h e a1k;lation reaction: CaHs + CsHs = CsHs(C3H7) (1) is an acid catalyzed condensation reaction using a solid or liquid catalyst. The formation of the hydroperoxide:
+
C9H12 0% = CgHs(C3H600H) (2) is a dual phase, base catalyzed, autocatalytic, free radical oxidation. T h e formation of phenol and acetone from the cumene hydroperoxide: (3) CsHs(C3HeOOH) = CsHsOH + CaHeO is a non-aaueous. acid catalvzed. 1.2-phenvl shift reaction which exhibits bdth negativiandp&iiive catalysis. The activation enerev and the heat of reaction for this last step are both easily measured. The experimental rate expression can be derived from a proposed mechanism so that this difference between kinetics and thermodynamici can he emphasized. 'I'he measurement of the heat reaction for the cleavage reaction (eqn. (:%)leads to theevaluation of the heat oiformstion of cumene hydroperoxide and hence to thecalculation of the heat 01' rearrion ror reaction 121. The heat of reaction for reactiun ( 1 1 ran be ohtainrd from standard tablessu that a11 of the enthalwies for tht. maior reactions are obtained. Additional details of the process c&be found in article by Armstrong (31, and Puiado. Salazar. and Beraer ( 4 ) . The thermodynamic expeAments are essentially calorimetric studies and can be completed in one period. However, because of the number of variables investigated in the kinetics experiment, two periods must he allotted or the work can be divided up among two or more groups. There are five separate kinetic reaction sequences, three a t the reflux temperature of acetone to measure the effect of hydroperoxide concentration, acid concentration, and water concentration a t a fixed temperature, and two a t lower temperatures to determine the activation energy of the reaction. If the project is to be a group effort, each group should do one high temperature experiment and the two lower temperature runs. At the end of the rotation period, the results of the group effort can he pooled and a final report written by each student. This type of project is illustrative of group~effortswhich are often encountered in industrial laboratories.
Work done in partial fulfilment of the requirements for the degree of Masters of Science, University of Houston, CLC. Current address: Exxon Research and Development, Baytown, TX. 458 / Journal of Chemical Education
H+ + ~~0& H ~ O +
Equilibrium (4)
H+ + ROOH &ROOH:
Equilibrium (5)
k-IB k-,A
ROOH%+&RO+
+ HzO
Equilibrium (6)
k-1
RO+ 3 R , ~ R , , R ~ R+"H ~ O %R'OR"OH
slow (7)
+ H+
Fast (8)
n
In this mechanism, reactions (4), (51,and (6)are equilibrium reactions, reaction (7) is the rate determining step, and reactions (8) and (9) are fast. Under these conditions, the overall rate expression given below can be obtained. (This is left as a student exercise.) d[Phenol] - K I A K Z[ROOH][H30+] -(10) dt ksKIB [HzOIZ When the svstem is comoletelv anhvdrous. the rate does not become i n f h t e but is governed by the dksociation of the sulfuric acid in acetone. The experimental rate data agrees well with the above rate expression, but this does not unambiguously establish the reaction mechanism. Analytical Procedures Because of decompmition of the hydroperoxide solutions, the exact concentration of the stock solution must be determinedjust prior to the experiments.This determination is done hv a non-aqueous iodine. ~uplicatedeterminations should be done for accuracy. ~pproaimately 0.1 eof the 50-75s cumene hvdroneroxide is weiehed accuratelvinto a 2 j 0 ml Erlenmever flask ebuiobed with a standard taoer ground
heated to the reflux temperature. Ten milliliters of an 80%isopropanol-water solution which has been saturated with potassium or sodium iodide is then added and the solution refluxed for 5 mi". The heat is then removed and all traces of iodine washed down the condenser with isopropanol, and the sample is cooled. Then thesample is titrated with a 0.025 F sodium thiosulfate solution usine a 50 ml turette. The end p g h t ipshnrp hut thr rltrarion should"& done rapidly to minimize air oridatim. The f m t complrw drsappmmrlc~ ui rhe rohr should hc used ns the end pcunr. There m m be some cloudiness due to water which can be remedied by adding more isapropanol. The densities of the solutions, 5%HzSOl in acetone, and the cumene hydroperoxide solution, can be obtained with a pycnometer. Because of the large number of analyses required for the kinetics experiment, a spectrophotometric determination of the phenol product is used. A Varian Technitron UV-VIS spectrometer, set at 284 nm, the peak for the phenolate ion was actually used; but a Spectronic 20 can also be used on the side of the band if the amount of sample is adjusted. Far our determinations, 0.05 ml samples were withdrawn from the reaction vessel at various times and dropped into a 20 ml test tube previously filled with 10.0 ml of a 0.1%NaOH solution. Not only does this caustic produce the phenolate ion but also it
immediately quenches the reaction. A 0.20 ml graduated pipette was used to obtain the samples. Obviously, larger samples could be taken but would require additional dilutions. Because of the small sample, special care must he taken to avoid contamination by color bodies. For this reason, cork stoppers should be avoided. The samples are poured into the spectrometer cells after being mixed thoroughly. The reference is a sample taken from the reaction flask before the acid catalyst is added. The cells and all other glassware should be rinsed thoroughly with distilled water, acetone, and then dried. As a precaution against a had analysis, duplicate samples should be taken.
Experimental Apparatus 'rhechrmicnls that are nrcded are cumene hvdropen,xide,availahle from Matheson. Coleman and Bell. rragent grade acetone, nmninally tr4T water. !A? s u l f u r i ~wld, wdium h\drox~de.p~tassiumiudide. sodium thiosulfate, and isopropyl alcohol. For the calorimetric studies, a thermistor bridge and recorder or an accurate thermometer is required. A stirred insulated beaker was used successfullyas a calorimeter since the reaction is so rapid. Simple volumetric glassware for titration and density measurements are
to reach infinite nhwrbanw. Alternatively, the infinitr trine value fur the lower temperaturr runs rau he calculnted from the hich ternperature infin~telime nhkcdrsnce ur ma)' he obtained the fdluwiny week. In the later case, the samples must be sealed tightly to prevent evaporation. While one student is doing a high temperature run, once the low temperature run is started, the other student can begin the analyses. Once the high temperature run is completed, the first student can help with the analvses and olot the data versus time to check for consistency. If a data point is "out of line," then the contingency sample can be analyzed.
Calculations The weight percent of the cumene hydroperoxide is calculated as ml Thio X F thia X 7.61 (11) wt sample (g) The heat absorbed hy the calorimeter and solutions is defined as q and is given by the relationship wt% ROOH =
nppded ....-.-
qbul = c,,
~
For the kinetic studies, besides the spectrometer, a 100 ml3-neck distillation flask with ports for thermometer, condenser, and sampling is required for a reactor. A hot water bath is used to heat the mixture t o reflux temperature to prevent "hot spots." The lower temperature runs are carried out in 50 ml Erlenmeyer flasks placed in a room temperature or ice bath and mixed with a TeflonQoated stirring bar. The temperatures in the samples were measured with thermometers. For the analysis, a large numher of 20 ml test tubes are needed for sample collection. Pipettes needed are as follows: 50 ml, 25 ml, 10m1, 2 ml, graduated 0.2 ml, and a graduated 5.0 ml. Syringes may be used for the sampling andlor catalyst addition. The volumetric flasks that are needed are a 100 ml, a 500 ml, and a 1000 ml.
AT
qmta1 = c ,,,,,,A T + q.~,
. .
~~~
~
~~~~~
~
(13)
The heat of reaction, AH, is then given by qtota1
AH,=
The thermodynamic experiments consist of collecting temperature versus time data for variousamountsof solvents. Earlier, the hydroperoxide content, the densities of the solutions and a calibration curve for the thermistor must he established. Also the standard calorimetric reaction solution, a 5% solution of H&04 in acetone must he prepared. This preparation is done by takings 200-ml volumetricflask, adding about 100 ml of acetone, then adding 5 ml of HzSO4 (98%) while cooling, and making up to mark with acetone. This solution is best oreoared hv the instructor , iust before the class since i t is not stahle for mure than a day nt n u m temperature. Extreme care must he tnken to moid ptting the acid ruluricm m u m a c t w ~ t hthe hydropennide stack wlulim 'I'hus the instructor s h d d pour uut ahout 10 mlofrhe solution for each group doing the experiment and keep the supply in another area. The temperature rise data should he obtained far a t least four solutions using 1.0 ml of the hydroperoxide solution and varying amounts of the acid solution. Initially pipette 15ml of theacidsolution into the calorimeter, in our ease a 100-ml insulated beaker. The probe or thermometer is then placed in the calorimeter and the stirring started. When a constant temperature is reached, (almost immediately), record it and add the 1.0 ml of cumene hydroperoxide from a pipette. Record the temperature every 3 see until the maximum is reached and then every 10 sec until the temperature drops by several degrees. If a thermistor is used, the temperature r-rd is obtained automatically. Repeat this procedure with 20 ml, 25 ml, 30 ml, and 40 ml of the acid solution. The calorimeter should be cleaned and rinsed with acetone between each run. There are three solutions to be prepmed for the kineticstudies, a 0.1% NaOH solution for analysis purposes, a dilute ROOH solution, 10 ml of the standardized concentrated ROOH (--75%setive) made up t o 500 ml, and a 4 % sulfuric acid in acetone solution prepared as . . above. Table 1summarizes the conditions for the various kinetic erneriments. The three hieh temoerature runs are carried nut in the 160ml llnak. First 50 ml of the stock hydrs,peroxidr rolutim is hrought tu the reflux temperature A sample is xithdmwn and quenched to be used for a reference. Then the acid is added and the stopwatch started. At about 25 min, either the second acid addition or the water addition can be made. While onestudent is doine the hiah temperature run, the other can be setting up the law temperature r u n s . ~ ~ hice e bath will provide a temperature of about 5 T in the flask. Time can he measured with a wrist watch for these runs. The samnles are agitated ,. with n Ttflon" covered stirring har. Once enough data points are rollrcr~dtu define thr c u r w , the low temperature snmples may br healed or ndditwnal acid may be added to decrease thr time required
(12)
where cpI,, is the heat capacity of the physical system, c,, p, and u (sol) refer to the heat capacity, density, and volume of the sulfuric acid-acetone solution and e,, p, and "(ROOH) refer to these parameters for the hydroperoxide solution. The last term can he qalculated from the physical properties of the pure materials and the measured densities and the experimental heat rise and can be defined as q,,, t o give upon substitution into eqn. (12)
Experimental Procedures
.~
+ (c,. p ."(sol) + c, .p .u(R0OH))AT
(14)
- moles ROOH
It is obvious that to find the heat ofreaction, the heat capacity of the system must he ohtained. The heat capacity e m be done by plotting qappversus AT and extrapolating to zero. The intercept is qt-I while the slope is the heat capacity of the system. The heat of formation for the hydroperoxide is obtained from
AH:,
(ROOH) = AH!,
(Phenol)
+ W , (Acetone) - AH, (15)
The heat of combustion of ROOH. the heat of reaction (2) and the hydroperoxide oxygen-oxygen bond dissociation energy can be calculated by standard techniques. Initially, the studentsshould plot the kinetic dateasLn(A- -A,) versus time. If the water concentration and the acid concentration nre conitant, and the reaction ir first order in hgdruperuxide euncentration, then this should yield a straight line (see eqn. ~1101. The do1)e of the line wdl he the preudo t m t order rate constant ihr and by carryingout the experiment under various water and acid &xs, the exponents, a and b, of the expression:
k = K [Ht]' [HsO]
(16)
can he determined. The empirical rate constant is defined by
Table 1. Klnetlc Reactlon Conditions
Reaction
T 'C
At mi"
ROOH
H~SOI
sol mi
sol ml
H20
ml
~~~~
~~
Reactions 1A and 18 are separate runs. Remion 28 is a Mtinvatim of2A in which 0.2 ml ol me %SO, soiutim Is added after about 20 min.
Resctiwl 38 is a cominuatlon of 3A
in which 0.1 ml
of
water
ia added after about 20
min.
Volume 57, Number 6, June 1980 1 459
Figure 1.TemperaturevwsusTime for 20 mi. 25 mi. 30 ml. and 40 mi of sulfuric acid solution.
Figure 3. Ln(A.
- A3 versus time as a function of hydroperoxide(lA.18)and
ass function of temperature (1.4.5)(See Table 1 for exact cmditions.)
Figure 2.Apparent heat release versus temperature change.
or in the integrated form: Ln(A,
- A3 = k t
Thus Exoeriments 1A and 1B (see Table 1) test the first order relarionahip-, experiments 21\ and 2H esrnhlish thc vnlue d o , and ex. prrimmrqR.4 and RBrital,liih the \,alurot h. tSeeTahle 1 f o r Reac. tion conditions.) Once the relationship is established, then the intrinsic rate constant, K, (not to he confused with an equilibrium constant), can he ealeulated and the temperature dependency examined. A plot of LnK versus l i T eives a d o ~ of e E*IRT and an interce~tof LnA aeeordina to the kinetic ~ r r h e n Expression. k
K = A erp(-E*lRT)
(19)
Typical Results The eumene hydroperoxide content of our sample was determined to be 51.0%. The experimental temperature nwes are shown in Figure 1, and the data is summarized in Table 2. The heat capacities are corrected for temperature. The extrapolation data is shown in Figure 2 where the intercept gives a qbhl value of 181.0 cal and a slope of 1.105 cal OC-'. From this the value of -56.11 keal mole-' is obtained using eqn. (4). This value is compared t o literature values of -54.0 and -60.4 kcalmole-' (3.4). With additional thermodynamic data taken from standard sources such as Perry and Chilton's "Chemical Engineering Handbook" (5) and the "Handbook of Chemistry and Physics" ( 6 ) .This data and the other calculated thermodynamic properties are summarized in Tables 3 and 4. The kinetic data are plotted in Figures 3 and 4. Table 5 gives the 460 1 Journal of Chemical Education
Table 3. Thermodynamic Data
AH:. Compound Pmpylene Benzene Cumene Cumene-OOH Phenol Acetone Sulfuric Acid Imiutionl
298
(kcdl mole)
4.879 11.63 -9.848
-
-AHcomb (kcail mole)
490.2 782.3 1247.3
-
-37.80 -59.32
732.2 426.8
-
-
C, callg - O K
0.35 0.45 0.50 0.535-0.540 0.350
mlml Experimental solution
0.9615
0.8163
Table 4. Heats of Reactlono (kcal-mole-') This Experiment
Ref (3)
Ref (4)
31.26 56.11
31.0 54.0
23.8 27.7 60.4
-Awl) -AM(?) -AM31
Figure 5. Ln K versus reciprocal Temperature for Me formation of phenol h a m cumene hydroperoxide.
Table 5. Typical Klnetlcs Results OMalned slope
t lmin ) Figure 4. Ln(A., - A,) versus time as a function of water (3A. 38). and function of acid (1, 2A, 20). (See Table 1 for the exact conditions.)
as a
values of the slope, the concentrations of the various species, and the intrinsic rate constant K. These were calculated as follows.
Thuso
= +1
&= k3*
(=)* 00205
( [ H ~ O I ~=B ) ~ = [H2013* 0.222 0.0438
Thus b = -2
,.- , A numher id'bnd poinLs can twsrrn in the kinetic plow 'I'hept.~)uld have hwn checked if rontmgenn. cnmplca had hen ohtamed.'l'hc exprrimrntal d a t a agrwswith the rateleu &tamed hum the mcch-
anism. Figure 5 shows the Arrhenius plot fur the K data. Theslope yields a value of E* = 19.23 keal mole-' and can be compared with avalue of 21.3 kcal mole-' reported by Seubuld and Vaughan (7). Conclusions This experiment has been carried out in the University of Houston Physical Chemistry Laboratories in one form or another for the past four years. The final form developed last year a t the University of Houston a t Clear Lake City is pre-
C H+
CR-
Reaction
miK1
molesll
1A 10 2A 20 3A 30 4 5
0.0450 0.0470 0.0195 0.0405 0.0438 0.0205 0.0155 0.0033
0.025 0.010 0.025 0.025 0.025 0.025 0.023 0.021
lo3 X
moles11
5.88 5.88 2.74 5.88 5.88 5.88 67.28 123.34
C H20 moles/l
0.222 0.222 0.222 0.222 0.222 0.332 0.222 0.222
Kmin-' mole11
0.377 0.394 0.327 0.339 0.367 0.384 1.14 X 1.32 x
lo-' lo-'
Ef = 19.63 Lullmole In A = 24.82
sented here. In general, the students obtain a much more realistic " erasr, - of the similarity and differences betu,ecn kinetin and thermodynamics by actually measuring these quantities on the same reaction, Acknolwedgment T h e authors are grateful to the University of Houston a t Clear Lake City for the support of this research. The aid of Dr. J. A. Ayala and Dr. B. H. White in running these experiments in the Physical Chemistry Laboratories a t the University of Houston Central Campus is also recognized. Literature Cited (1) Beftelheim.F. A.,"Experimental PhysicslChemistry..l W. B. Saunders, Philadelphia, Ps., 1971. (2) Danielr, F.,el al.. '"ExperimentalPhysical Chemistry," MeCraw-Hill, Near York, New Yurk. 1962. 131 ~ ~ ~ p t ,in ~..Modern ~ eChemi3tryin . ~ .Industry." (EdifocJ.G.Grewry),I.U.P.A.C., I., m td 196~.I I inr, t o P . g ~ ~I,.. K ~ ,. s d l ~ a Ic .H ,..d uerper.l' \' . H ! d m c . r n , m R o e . x . w Y l 5.5. I976 ,i, "I~err'.L'tlcrncrl Fn*,npmng Han.1~ k; Mrlir.u.ll,ll. h u Y r i 1963. $C 'IlcndboA ,.l('l.rm 4 r \ .na p h ) ~ ~ ~('n.tn.~al .,' H Y ~ ~ IPuhl P C rhlng I'(. ~lrveland 1970. (7) Seubuld, F. H.,and Vaughan, W. E., J Amer Chem. Soc., 75,3794 (1953)
n.
.
Volume 57, Number 6,June 1980 1 461
