Experiments Illustrating Heterogeneous Catalysis Maria A. Martin Luengo and Paul A. Sermon Brunel University, Uxbridge, Middx., UB8 3PH, UK
A catalvst accelerates a reaction hv orovidine a reaction pathway bf lower free energy than thk noncacalyzed one, although the catalyst does not itself appear in the reaction. Catalysis may be homogeneous (I) or heterogeneous (2) deen dine uDon whether the catalvtically active component is i f the same phase as the reactants or not. The importance of catalysis is seen in the fact that catalyst sales in p;ocess chemistriand environmental chemistry are likely to reach $US 5 billion per annum worldwide, with use in pollution abatement, oilrefining, and chemical processing (3). Consider one example of this chemistry: methanol synthesis CO 3Hz = CH30H H20. Catalysts began to he used for this conversion (4,5) a t 873 K and 900 atm, but by the mid-1920's Patart's Oxylite plant near Paris operated a t 573-673 K with CO/H? = 0.5 a t 200 atm, and in 1923 the Nerseherg plant of ~ a d i s c h ewas producing 10-20 tons methanol Der dav (5) with 33% conversion of CO to CHsOH. As a resuli in 1931 an Eneter schoolmaster (6)was to write suggesting a remarkably far-sighted use for this chemistry:
+
+
Another great service thet the chemist of tomorrow will render to his fellows is in connection with oower. In Eneland - our su~nlies .. of oower are derived chieflv,~ from coal.. .our cars of course deoend on p e t r d . .thr world's uil supplies are rapidly giving out. How will rhey hrrplared"Therhcmist isalready busy on theproblrm and there are indications that it will be solved by the preparation of a sort of artificial petrol.. .water gas can easily be manufactured (from coal1
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~
~
ZnO and C u I Z n O l A l ~ 0(11) ~ as the reaction benefitted from advances in catalysis (11) with the result that operating costs starting with methane for the lower pressure process were 41% lower (12) than with the high-pressure process. The ICI work eventually led to the award of the 1974 MacRohert Prize (12). Under standard conditions the CO hydrogenation reaction to yield methanol C0 + 2H,
= CH,OH
has a negative enthalpy of -91 kJ/mol, hut the reaction is difficult to demonstrate to undergraduates since the equilihrium constant is only sufficiently large for measurable conversions say a t 523 K and 20 atm. However, the reverse reaction occurs readily over P t or Cu surfaces a t atmospheric pressure and 523 K (5). Naturally, one has to be careful about comparing catalysts in methanol decomposition a t atmospheric pressure (13) and then deducing relative activities in methanol synthesis at higher pressures (14). The vapor pressure of methanol (see Fig. 1) is about 4 kPa a t 273 K. Thus a simple flow reactor may he used a t atmornheric oressure in which a N? stream of constant known -=--flow rate is saturated with met&nol a t this partial pressure. The extent of methanol decomoosition (and the formation of methane) can then he analyzed as a function of reactor temperature using samples of the gas stream injected onto a gas chromatograph fitted with a flame ionization detector.
.
By the use of asuitable catalyst (Cu mixed with ZnO) chemists are now able to turn water-gas into. . .methyl alcohol C0 + 2H, = CH,OH There is no doubt that motor car engines could be modified if necessary so as to work with this substance instead of petrol. The conversion of coal to COIH2 to methanol fuel is still the subject of discussion (7) with its use as a substitute for vehicle fuel possible (8) as aresult of the drivingforce for less dependence on unstable oil supplies. Related to this is the methanol conversion over ZSM5 and other zeolites in the Mohil "methanol-to-gasoline" process. The thermodvnamics of the methanol svnthesis reaction from C O / H ~(with a decrease in volume on-going from left to right for the exothermic reaction) favor methanol synthesis a t high pressure and low temperature, although in reality the latter is a comoromise hetween the requirements of kinetics and thermodynamics. Therefore, the search was on for improved catalvsts operatine a t milder conditions. First, Cul i n 0 (90:lO) had greater artivity in this reaction than either Cuor ZnOalone (5J.Then.in 1965 Imperial Chemical Industries (9) initiated a process in which methane was steamreformed to CO/C0&2, which was then converted to methanol; the latter step occurred over a Cu/ZnO/Alz03 (60/30/ 10) catalyst (10) a t 50-100 atm and 523 K. These conditions are clearly much more modest than those prevailing earlier and as a result are much more energy-efficient. ZnO alone was thus followed by ZnO-CmOn, Zn0-Mn0-Cr203, Cu-
Figure 1. Variation of vapor pressure of nnrmanol wnh absolute temperature.
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Number 3 March 1991
25 1
The number of pmol methanol converted per gram catalyst per minute is then given by:
catalyst has achieved essentially complete conversion of methanol. Thus knowing Wand F then a t 563 K P = 88.6% and so in fact 521.8 pmol methanol are converted per minute per gram catalyst. Since the same weight of catalyst adsorbs 168 pmol each adsorption site converts 3.1 molecules/min.
where W is the weight of catalyst used, P is the percentage methanol conversion, and F is the gaseous flow rate in cm3/ min. If the number of adsorption sites per gram of catalyst is known then the turnover frequency (TOF) per adsorption site can be estimated. Consider the results in the table for the hlank reactor and the reactor containing 0.15 g (W) of prereduced Pt/Si02 (sample K (15),which adsorbs on average 168 pmol CO or Hz per gram catalyst) if F = 50 cm3/min. Clearly when the activity of the empty reactor is just showing conversion (via the gaseous noncatalyzed reaction), the Convenlons ol Methanol Observed [MeOHI* Imml
T iK1
[CHI]' lmml
%MeOH con~ersion
Pt/S1O2sample 413 433 453 463 503 523 543 563
~ l g u r e3. Activation energy plot of noneatalyzed ( E d and catalysed (Ed methanol decomposltion. Change in AG IkJlrnol)
I "aC peak hslghk.
In rate CHS OH decompo~ltlon
Figwe 2. EHect of reciprocal temperatureon the l~garlthrnlcrate of methanol d a w m p i t i o n in the presence(open symbols) and absence (closed symbols) of a sllicaaupponed Ftcatalyst.
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Journal of Chemical Education
Figure 4. Free energies of fwmation of melhanol and methane horn COIHz with water produced in the latter case.
This is not a large number when it is remembered that hundreds of methanol molecules will be striking each square centimeter of the catalyat surface per minute. This conversion can now be related to the limiting conversion predicted from thermodynamics. Kelley (16) found that the equilibrium constant for decomposition of methanol to COIH2 (relative to a total pressure of 1 atm) varied with absolute temperature in the following manner:
Figure 2 shows the Arrhenius plots for the catalyzed and the noncatalyzed decomposition of methanol; from this it is clear that the noncatalyzed process has a higher activation energy (77.1 kJ1mol; ENC) than the catalyzed value (24.8 kJ/ mol;Ec). This is in keeping with normal definitions of catalysis (see Fig. 3, where ENC > Ec). Returning now to the table, the selectivity of the reaction can be considered at least in the sense that at the higher temperatures (i.e., 500450 K) methane is also produced in small concentrations. In Fiaure - 4 it can be seen that in a thermodynamic sense methanol is unstable with respect to COand hvdrorenat hirher temperatures hut isalso unstable with respect -to metKane at intermediate temperatures. Therefore, it is not surprising that some methane is produced. However, formation of metbanemeans thatsome Ois released (possibly transiently) to the catalyst surface:
+
2CH,OH = C 0 + 2H2 CH,
+0
either to be removed by CO or surface hydrogen (left in the
orereduction). In other words, this illustrates the transient modification ofcatalyst surfaces that canoccur. If theextent of methane and C 0 j release is monitored as a function of time, then the net 0 content of the catalyst during the reaction can be deduced. Therefore, this simple inversion of the industrially relevant reaction reveals some of the intricacies of catalysis. In addition the CuIZnO catalysts that model the commercial catalysts can readily and inexpensively be prepared by (1) coprecipitation from a mixed aqueous solution of the zinc acetate and copper nitrate using NaOH-NaHC03 at pH 910, (2) drying, (3) calcining at 573 K i n air, and reducing in 6% H2/N2at 473 K. Such CuIZnO catalysts produce methyl formate (HCOOCH3) in addition to the products seen here for Ptlsilica in decomposition of CH30H. Whether the predictions of the Exeter schoolmaster become true or not now depends upon economics and politics, for chemistry has already made its contribution in facilitating the "art of the possible". Literature Clted 1. Hanson, B.C.:Davis, M.E.J.ChmxEduc. 1987,64,92&930. 2. Copperthwaite.R.G.:Hutehings,G.J.:VanDerRiet,M.J.Chem.Edue. 1986.63.632634. 3. Chem. E n s News 1989. (29 May). 29-56. 4. BadischeBrifirhpatent237030; Patsrt.G. Compl. Rand. 1924.179.1330-1332; Chimie et lnduslr 1925.13.179-186:Bull.Soc.Encour.1nd.Nnt.1925,137,141-173. 5. L0rmand.C. lnd.Eng. Chem. 1925.17.430-432. 6. Littler, W. A Junior Chemistry: Bell. 1931:p318. 7. ~ n e i l i n gG. , Haus. ~ e r h Vorrrogsuerorii . 1978,405, 11-15. 8. Moihonol: Promise ond Probi~mr: SF-726SAE; Alcohols as Motor Fuels; SAE 1980. 9. Gallagher, J.T.:Kidd. J. M.British patent 119651. 1159035. 10. Bawkor, M. Vacuum 1983,33,669-685. 11. Sermon. P. A. Chzm. Ind. (London) 1979.681490:Twigg, M. V. Cofalyrl Hondbaok Wolfe. 1989. 12. Nolmn,D. H.:Hanson, D. Chem. Ens 1969, (22Sopt.lNewSci1974,929. 13. Pospekhov,D.A.J.Appl. Chrm. (USSR11967.20.769. 14. Bardet, R.: Thivolle-Carat,J.; Trarnhouze, Y. J. Chim. Phya. Chim. Rml. 1981, 78, 135-138. W.:Sermbn,P.A,:Wuria,A.T.Adsn.Sei. Tech. 1984, 15. Berzin8.A. R.;LauVang,M.S. 1 , 51-76. 16. Kelley. K. K. 1nd.Eng.Chrm 1926,18.78.
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