Chemistry in Catalysis

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3 4.6 Carbon 108 4.7 Preparation of supported catalysts 109 5. Kinetics 5.1 Langmuir-Hinshelwood model 115 5.2 Monomolecular reaction 117 5.2.1 Surface reaction rate determining 118 5.2.2 Adsorption of A rate determining 121 5.2.3 Desorption of B rate determining 121 5.3 Bimolecular reaction 121 5.4
4.6 Carb 108 4.7 Preparation of supported catalysts 109 5. kineti 1 Langmuir-Hinshelwood model 115 5.2 Monomolecular reaction 117 5.2.1 Surface reaction rate determining 118 5.2.2 Adsorption of a rate determining 121 5.2.3 Desorption of B rate determining 121 5.3 Bimolecular reaction 121 5.4 fluence of diffusion 122 6. Metal Surface 6.1 Surface struct 130 6.2 Surface analysis 134 6.2.1 XPS 6.2.2 Surface sensitivity 139 6.2.3 Auger electron spectroscopy 140 6.3 Surface enrichment 142 6.4 Metal binding 147 7. Metal Catalysis 7. 1 Dissociation of h 154 7.2 Hydrogenation of ethene 158 7.3 Ammonia synthesis 162 7.4 Co hydrogenation 7.4.1 Fischer-Tropsch and methanation reaction 166 7.4.1.1c0 dissociation 167 7.4.1.2 Methanation 172 7.4.1.3 Fischer-Tropsch reaction 175 3 7.4.2 Methanol synthesis 7.5 Volcano curves 183 7.6 Bronsted-Evans-Polanyi relations 186 8. Catalysis by Acids 8.1 Reactions catalyzed by solid acids 197 8.2 Reactions of carbenium lons 199 8.3 Applications of carbenium lons 207 4 Bifunctional catalysis 8.4.1 Principles 209 8.4.2 Industrially relevant bifunctional catalysis 218 4 1. Introduction 1.1 Kinetics and thermodynamics The word catalysis was introduced in 1826 by Berzelius to indicate the action of a substance (the catalyst) that increases the rate of a reaction but is not changed during the reaction today we know that a catalyst does change when it catalyzes a reaction but that it is, in principle, changed ick to its starting form. That means that a catalyst can be used over and over again. the following example of the oxidation of So2 to SO3 illustrates this SO2+ v2O5 SO3 +V2O4 V2O4+02 V2O5 SO2+12 O2 △H=-100k/mol The catalyst transforms so 2 in two steps into SO3 In the first step, the v2Os catalyst oxidizes so2 to SO3 and changes to V2O4, which is transformed back into V2O5 in the second step. The first step is an oxidation of So, and a reduction of V2O5. in the second step the reverse reactions occur now a molecule is reduced and the catalyst is oxidized. Thus, after one full catalytic cycle, it looks as if the catalyst has not changed. Schematically this can be indicated by a circle with the catalyst stages on the circle, reactants entering the circle and products leaving the circle. It shows that the catalyst can be used an infinite"number of times. VoO 12 外 Of course, in reality life is not so simple. Often the catalyst becomes progressively covered with coke and less and less catalyst surface is accessible for adsorption and catalysis When the deactivation of the catalyst has become too severe one can regenerate the catalyst for instance by burning off the coke in air and re - reducing the catalyst if needed another reason for loss of catalyst activity can be sintering, growth of the catalyst particles. Small particles have a large surface area and energy which is a thermodynamic instable situation. At elevated reaction temperature particle growth may take place by diffusion and particle- particle growth or by scission of atoms from particles and diffusion to other particles. In both cases the catalyst activity decreases but not because of a change in the catalyst structure but because of a loss in accessible surface area. the catalytic reactions thus are still the same only the kinetics has slowed down. a better definition of catalysis than given by berzelius thus could be that a catalyst hanges the kinetics but not the thermodynamics of a reaction Catalysis has to do with kinetics, with increased rates of reaction. Rates can only be increased, however, if thermodynamics allows. the difference between kinetics and thermodynamics is illustrated in the following examples The equilibrium of the reaction 2 NCI3 n2+3 Cl2 is at the right hand side at room temperature. thus, NCl3 is instable (refers to thermodynamics), but only decomposes (explosion when shock is applied or when a catalyst is added The equilibrium of the reaction SiCl, Si+ 2 Cl2 is at the left hand side at room temperatur Thus, SiCla is stable (refers to thermodynamics) and will not decompose, even if a catalyst is The equilibrium of the reaction CCl4 +2 H2O+ CO2+4 Hcl is at the right hand side at room temperature. Thus CCl4 is instable but, as all chemists know tetrachlorocarbon does not react with water at room temperature. Ccla is instable (refers to thermodynamics), but inert (refers to kinetics) In the absence of a catalyst, the oxidation of so2 to SO3 has a low reaction rate and T >600C would be necessary to obtain a reasonable rate(Arrhenius equation, r=A. exp(-AE/RT) Unfortunately, the reaction is exothermic and the onversion is limited at high temperature by thermodynamics △G0=AHP-T△S0=- RTIn k and d In K/aT=△HP/RT2 An increase in t has a negative influence on the position of the equilibrium of an exothermic reaction(because AH <O). For such reaction, thermodynamics tells us to keep t low to obtain a high conversion but low T means low rate k= A. exp(-AE/RT)). The solution is to use a catalyst and accelerate the reaction so that one can run the reaction at low t. this suggests that endothermic reactions do not need a catalyst, because thermodynamics requires them to be run at elevated T and then kinetics may be sufficient even without catalyst. In reality catalysts are 6 also used in many endothermic reactions because material problems do not always allow reactions to be run at high t The exo-or endothermicity of a reaction can often (but not always )be predicted from the reaction stoichiometry, because in many cases the sign of the change in enthalpy AHis the same as the sign of the change in the number of molecules An during reaction When the number of product molecules is larger than the number of reactant molecules(An>0), then in principle bonds are broken and enthalpy will have to be provided (endothermic reaction, AH>0). The reverse often holds for exothermic reactions(An <0, AH<0). Thus, dehydrogenation reactions such as CnH2n 2 -CnH2n H2 and csh12 -ch6 +3 H2 are endothermic and the reverse hydrogenation reactions are exothermic. Oxidation reactions such as 2 H2+02>2 H2O(AH 400 kJ/mol)and 2 SO2+ 02+2 SO3(AH =-180 kJ/mol)are exothermic and the reverse reduction reactions are endothermic. When An a 0, AH may be small, as in the water-gas-shift reaction co+H2O→CO2+HO(△HP=-10k/mo) and in so2+Vo5→S03+VOn(△HP=10 kJ/mol) But an exception is the reaction n2+02= NO with An =0 and Ah= 160 k]/mol! The following table summarizes the effects of temperature and pressure for exo- and endothermic reactions. The effect of pressure is predicted with the formulae Kx=Ko p-n and d Ink / d InP=-An Exothermic reactions Endothermic reactions Low t, high x owT,lowⅹ ncreasing T, decreasing x increasing t, increasing x AS/R<o high T, low x △S/R>0 high t, high x An<0 increasing P, increasing X An>0 increasing P, decreasing x atalyst important catalyst relatively less important As the example of the oxidation of So2 to SO3 shows, the catalyst is not present in the total reaction SO2+ 102- SO3. The equilibrium of the reaction is therefore independent of the catalyst A B AG =-RTIn K=-RT In(kl/k-1) A+KH B+K AG=-RT In K=-RT In(k'/k1 In these equations k, and k-1 are the rate constants of the forward and backward reactions Because AG=AG k,/k,=k'/k,. This means that the ratio of the forward and backward reactions is independent of the presence of a catalyst. In other words the catalyst increases the backward reaction rate as much as it increases the forward reaction rate. this is a very important conclusion, because it allows us to explore catalysts for a reaction by studying their effect on the reverse reaction. For instance, 100 years ago this principle was used in the search for the optimum catalyst for the synthesis of ammonia N2+3H2→2NH3 △H=-92kJ/mo The forward reaction must be run at high pressure to attain reasonable conversion(an <0), but the reverse reaction is favoured by low pressure Experiments are easier to perform at low than at high pressure and thus, in a few years, more than 1000 catalysts could be explored and Fe wa found to be the best ammonia-synthesis catalyst. Analogously a catalyst that is good for the methanation of carbon monoxide is also good for the reverse reaction, the steam reforming of methane to synthesis gas: CO+3H2→CH4+H2O△HP=-167k/mo Again, the conditions under which one operates the forward or backward reaction are different because of the different thermodynamics. Thus, the methanation reaction (An <0, AH<0 should be performed at low T and high P, while the steam reforming conversion will be best at high T and low P, but in both cases the catalyst of industrial choice is Ni. To prevent sintering of the Ni particles, the Ni is supported on an alumina support(Ni/ Al O3). In steam reforming the temperature is high(700C). Therefore, the catalyst support must be temperature stable and a Al2O3 is used, as we will discuss in Chapter 4. At the lower temperature at which the methanation process is operated (300C), the lower stability of y-Al2O3 is not a problem, but its much higher surface area is a big advantage in stabilizing small and thus highly active, Ni particles 1. 2 Activity, selectivity and stability The rate of a reaction is equal to r=-dX/dt=kXY with k= A. exp(-AE/RT) which means that an acceleration of the reaction rate can be obtained by an increase in the concentrations of the reactants (as may occur in enzymatic catalysis and with immobilized catalysts) as well as by an increase in the rate constant k Statistical mechanics tells us that the pre-exponential factor a of a catalyzed reaction cannot be larger than that of the non-catalyzed reaction. That leaves a decrease in the activation energy as a possibility to increase the reaction rate and thus the activity of the catalyst. The next figure shows that two reactants a and b bind to the catalyst to a* and b and from there react to the products C* and d'* adsorbed on the catalyst surface. Desorption then gives the products C and D. the activation energy of the non-catalyzed reaction is e and the activation energy of the catalyzed reaction is e, E or e depending on the fact which is the largest. When the binding between catalyst and reactants is large then e" is large and when the binding between catalyst and products is large then e" is large. As we will see in chapter 5, metal catalysis, early transition metals normally bind reactants and products strongly. This leads to large E and not very active catalysts. Late transition metals on the other hand may bind reactants and products weakly and this may lead to a large e and again not very active catalysts The best catalyst is then somewhere in the middle, as the fe catalyst for ammonia synthesis is better than a v catalyst (too high e and a ni catalyst too high e. A ‘B餐 C+ D 黄十 a catalyst should not only be active(accelerate the reaction but also selective and stable to be a success in industry. a high selectivity to the desired product(s) means that less reactant nolecules are wasted on byproducts and that less energy is needed for separation of desired products from byproducts. an example of the strong influence that a catalyst can have on the selectivity is the transformation of synthesis gas to hydrocarbons and oXo-products Fe Co. Ru co+ CH3OH Cu/Zno, Pd-Cao/Sio, CH CHO Rh Metals such as Fe, Co, Ni and Ru lead to hydrocarbons. Ni produces mainly to methane and ru mainly higher hydrocarbons. Fe produces lower alkanes, alkenes and alcohols. Pd, Pt, Ir and cu are good catalysts for methanol formation in the presence of basic supports or promoters such as Zno and CaO rh is flexible and able to make hydrocarbons methanol as well as higher oxo products such as ethanol and acetaldehyde Interest in increasing the selectivity of rh to c2 and higher oxo-products is especially high in countries rich in coal or natural gas, which can be transformed into synthesis gas In addition to high activity and selectivity in several processes industry needs stable catalysts. a high stability means that the catalyst can be used for a long time and that fewer regenerations or even replacements, of the catalyst are needed. Regeneration and replacement not only add to the catalyst cost, but also to operation cost because during regeneration or replacement operation is interrupted. Stability is especially important in the refinery and bulk chemicals industry, where the catalyst must be used for extended periods because of the lower margins than in the fine chemicals industry, where one can afford to use the catalyst only once. Also, in the refinery and bulk chemicals industry continuous processes are used to produce the required hundreds or thousands of tons per day and taking these processes out of operation for catalyst regeneration or replacement diminishes production. Therefore, it is not only important to know the initial activity of a catalyst, but also how long it can run this is expressed in the turn-over number (ton), the number of molecules that react per atom on the catalyst surface before the catalyst has deactivated. Whereas batch processes without reuse of the catalyst may afford TON's as low as 10, continuous processes that run for one year may need ton 10. A ton of one, of course, means a stoichiometric, non-catalytic, reaction 1.3 Heterogeneous, homogeneous and enzymatic catalysis In heterogeneous catalysis the catalyst is in a different phase(solid phase)than the reactants and products (gaseous or liquid phase). the catalysts that were mentioned in the foregoing are all solid catalysts(V,Os, Fe, Ni/Al203)used to convert gaseous reactants (SO,, N2+H2, Co+H2, 10

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