Catalysis [PDF]

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Catalysis

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Table of content Sr.No.

Topic

1

Catalysis positive catalysis Negative catalysis promoters Inhibitors History

1.1 1.2 1.3 1.4 2 3 3.1 3.2

Page no. 3 3 3 4 4 6

Types of catalysis

8

What is a phase?

8

Homogeneous catalysis

9

3.2.1

The destruction of atmospheric ozone

10

3.2.2

Homogenous catalysis in liquid phase

12

3.2.3 3.3 4

Enzymes catalysis

12

Heterogeneous catalysis

13

Autocatalysis

16

5

Characteristic of catalysis

18

6

Theories of catalysis

19

6.1

Intermediate compound formation theory 6.1.1

6.2 7 7.1

Limitations of intermediate compound formation theory

7.1.2

21

Adsorption theory

21

Applications

23

Catalysis in Fluid Catalytic Cracking 7.1.1

20

Catalyst Specifications Chemistry involve in catalytic cracking

23 24 26

7.2

Hydrogenation

27

7.3

Oxidation

29

7.4

Water-gas shift (WGS) reaction

30

7.4.1

High temperature shift (HTS) catalyst

31

7.4.2

Low temperature shift (LTS) catalyst

32

2 7.5

Production of Carbon Monoxide

7.6

8 8.1

Methanol production

34

7.6.1

Steam reforming

36

7.6.2

Methanol decomposition

37

7.6.3

Partial oxidation

37

7.6.4

Combined reforming

37

7.6.5

Ammonia synthesis

37

Catalyst characterization techniques

38

Scanning Electron Microscopy

38

8.2

Physisorption analyses

39

8.3

Chemisorption analyses

40

8.4 8.5 9

33

Activity measurements

42

Chemical analyses

45

References

46

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1. Catalysis An acceleration of the rate of a process or reaction, brought about by a catalyst, usually present in small managed quantities and unaffected at the end of the reaction. A catalyst permits reactions or processes to take place more effectively or under milder conditions than would otherwise be possible. [1] Unlike other reagents that participate in the chemical reaction, a catalyst is not consumed by the reaction itself. A catalyst may participate in multiple chemical transformations. 1.1 .

Positive catalysis

A catalyst which increase the rate of a reaction is called a positive catalyst and the is called positive catalysis. There are million of examples of positive catalysis For example (1) In the preparation of oxygen MnO 2 at as a positive catalyst in the decomposition of KCLO3 . MnO

2 2KCLO3   2KCL  3O 2

(2) Manufacture of ammonia by Heber’s process using finally divided iron as catalyst Fe N 2  3H 2  2NH 3

1.2.

Negative catalysis

A catalyst which retards the rate of a reaction is called negative catalysis and the phenomenon is called negative catalysis. For example (1) Decomposition of H2O2 is retarded in the presence of traces of acetanilide. 2H 2 O 2 acetanilid  e 2H 2 O  O 2

(2) Oxidation of chloroform is retarded in the presence of a small quantity of ethyl alcohol.

4 C H OH

2 5 2CHCl3  O 2   2COCl 2  2HCL

1.3. Promoters (Activators) Some times the presence of some other substances which are not necessarily catalysts increases the efficiency of a catalyst. Such substances thus act as catalyst for the catalyst and are called promoters or activators. Molybdenum act as promoter for the iron catalyst in the manufacture of ammonia by Haber’s process. 1.4. Inhibitors (Catalytic poisons) In some cases the presences of small quantities of impurities in the reacting substances make the catalyst in active. Such substances are called catalytic poisons or inhibitors. Catalytic poisons are usually also poisonous to catalytic poisons. Arsenious oxide, As 2O3, and hydrogen oxide poisons platinized asbestos used as catalyst in the manufacture of H 2 SO4 by contact process.[2]

Catalytic reactions have a lower rate-limiting free energy of activation than the corresponding uncatalyzed reaction, resulting in higher reaction rate at the same temperature. A catalyst can do this by participating in the activated complex for the rate-limiting step, even though the catalyst itself is neither a reactant nor a product in the overall stoichiometric equation. A good example of catalysis is provided by the effect of I2 on the rate of isomerization of cis-2-butene. The rate law for the catalyzed reaction involves the concentration of I2 raised to the one-half power. This implies that half a molecule of I2 (that is, an I atom) is involved in the activated complex, which probably has the structure

Since there is no double bond between the two central C atoms, one end of the activated complex can readily twist around the other.

5 The currently accepted mechanism for this catalyzed reaction involves three steps: ½ I2

I

I + cis-C4H8 → C4H8I‡ → trans-C4H8 + I I

½ I2

The first and last steps have coefficients of one-half associated with I 2 because each I2 molecule that dissociates produces two I atoms, only one of which is needed to help a given cis-2-butene molecule to read. Note also that for every I2 molecule which dissociates in the first step of the mechanism, an I2 molecule is eventually regenerated by the last step. As a result, the concentration of I2 after the reaction remains exactly the same as before. If we consider the energetics of each step in the proposed mechanism, we find that much less than the 262 kJ mol–1 activation energy of the uncatalyzed reaction is required. A complete energy profile for the catalyzed process is compared with that of the uncatalyzed one in Fig. The bond enthalpy is 151 kJ mol–1 and so 75.5 kJ mol –1 will be required for formation of the I atom in the first step. An additional increase in energy occurs as the I atom collides with cis-2-butene and bonds with it. Then about 12 kJ mol –1 is required for twisting around the C—C single bond in the activated complex. All told 115 kJ mol –1 is required to go from the initial molecules to the activated complex. When rotation to a trans structure is complete, the I atom dissociates fromtrans-2-buten and eventually reacts with another I atom to form I2. These last two processes involve an overall decrease in energy which is nearly the same as the increase required to achieve the activated complex.

The reduction in activation energy illustrated in Fig

6

Fig. Energy profile for the cis-trans conversion of 2-butene catalyzed by iodine. Iodine atoms must first be produced from I 2 molecules after which the addition of an I atom to a butene molecule produces the activated complex. Since the activated complex contains no double bond, one part of the molecule can now swivel easily around the other, producing the trans isomer after the loss of the I atom. Since this pathway requires much lower activation energy, it allows the reaction to occur much more rapidly. [3] Catalysts may affect the reaction environment favorably, or bind to the reagents to polarize bonds, e.g. acid catalysts for reactions of carbonyl compounds, or form specific intermediates that are not produced naturally, such as osmate esters in osmium tetroxidecatalyzed dihydroxylation of alkenes, or cause lysis of reagents to reactive forms, such as atomic hydrogen in catalytic hydrogenation. 2. History This term derived from Greek καταλύειν, meaning "to annul," or "to untie," or "to pick up." The word "catalysis" (Katalyse) was coined by Berzelius in 1835: "Catalysts are substance that increase the rate of reaction but remain unchanged after the reaction."

7 The Chinese Tsoo Mei is more picturesque; it also means "the marriage broker," and so implies a theory of catalytic action. The idea of catalysis extends far back in to chemical history. The quest of the alchemist for the philosopher's stone seems like the search of the modern petroleum chemist for the magical catalyst that will convert crude petroleum into high octane fuel. In a fourteenth-century Arabian manuscript, Al Alfani described the "Xerion, aliksir, noble stone, magisterium, that heals the sick, and turns base metals into gold, without in itself undergoing the least change" The earliest consciously used catalysts Were the ferments or enzymes. "In the seeding and growth of plants, and in the diverse changes of the fluids of the animal body in sickness and in health, a fermentative action takes place ". Its action has been likened to that of a coin inserted in a slot machine that yields valuable products and also returns the coin. In a chemical reaction the catalyst enters at one stage and leaves at another. The essence of catalysis is not the entering but the falling out. Other early chemists involved in catalysis were Alexander Mitscherlich who referred to contact processes and Johann Wolfgang Döbereiner who spoke of contact action and whose lighter based on hydrogen and a platinum sponge became a huge commercial success in the 1820s. Humphry Davy discovered the use of platinum in catalysis. In the 1880s, Wilhelm Ostwald at Leipzig University started a systematic investigation into reactions that were catalyzed by the presence of acids and bases, and found that chemical reactions occur at finite rates and that these rates can be used to determine the strengths of acids and bases. For this work, Ostwald was awarded the 1909 Nobel Prize in Chemistry. Wilhelm Ostwald was the first to emphasize that the catalyst influences the rate of a chemical reaction but has no effect on the position of equilibrium. His famous definition was: "A catalyst is a substance that changes the velocity of a chemical reaction without itself appearing in the end products." Ostwald showed that a catalyst cannot change the equilibrium position, by a simple argument based on the First Law of Thermodynamics. Consider a gas reaction that proceeds with a change in volume. The gas is confined in a cylinder fitted with a piston;

8 the catalyst is in a small receptacle with in the cylinder, and can be alternately exposed and covered. If the equilibrium position were altered by exposing the catalyst, the volume Would change, the piston would move up and down, and a perpetual-motion Machine would be available. Since a catalyst can change the rate but not the equilibrium, it follows that a catalyst Must accelerate the forward and reverse reactions in the same proportion, since k  k f kb . Thus catalysts that accelerate the hydrolysis of esters must also accelerate the

esterification of alcohols; dehydrogenation catalysts like nickel and platinum are also good hydrogenation catalysts; enzymes like pepsin and papain that catalyze the splitting of peptides must also catalyze their synthesis from the amino acids. A distinction is generally made between homogeneous catalysis, the entire reaction occurring in a single phase, and heterogeneous catalysis at phase interfaces. The latter is also called contactor surface catalysis.

[8]

3. Types of catalysis: There are two major classes of catalysis Homogeneous Heterogeneous Before discussing homogeneous and heterogeneous catalysis first we discuss phase. 3.1. What is a phase? If we look at a mixture and can see a boundary between two of the components, those substances are in different phases. A mixture containing a solid and a liquid consists of two phases. A mixture of various chemicals in a single solution consists of only one phase, because we can't see any boundary between them.

9

we might wonder why phase differs from the term physical state (solid, liquid or gas). It includes solids, liquids and gases, but is actually a bit more general. It can also apply to two liquids (oil and water, for example) which don't dissolve in each other. we could see the boundary between the two liquids.

If we want to be fussy about things, the diagrams actually show more phases than are labelled. Each, for example, also has the glass beaker as a solid phase. All probably have a gas above the liquid - that's another phase. We don't count these extra phases because they aren't a part of the reaction. [4] 3.2. Homogeneous catalysis homo implies the same (as in homosexual). Homogeneous catalysis has the catalyst in the same phase as the reactants. Typically everything will be present as a gas or contained in a single liquid phase.[9]

10 Examples The reaction between ethanol and ethanoic acid is very slow even when heat is applied. C2H5OH + CH3COOH

CH3COOC2H5 + H2O

The addition of a few mL of sulfuric acid greatly increases the rate of reaction. Since ethanol, ethanoic acid and sulfuric acid are all liquids, sulfuric acid, as used in this example, is a homogeneous catalyst. [10] An example of homogeneous catalysis in the gas phase is the effect of iodine vapor On the decomposition of aldehydes and ethers. The addition of a few percent of iodine often increases the rate of pyrolysis several hundredfold. The reaction velocity follows the equation,

 d (ether )

dt

 k2 ( I 2 )(ether )

Dependence of the rate on catalyst concentration is characteristic of homogeneous catalysis. The catalyst acts by providing a path for the decomposition that has considerably lower activation energy than the uncatalyzed path. In this instance the Uncatalyzed pyrolysis has an Ea  53 kcal, where as with added iodine the E a drops to34 kcal. 3.2.1. The destruction of atmospheric ozone This is a good example of homogeneous catalysis where everything is present as a gas. Ozone, O3, is constantly being formed and broken up again in the high atmosphere by the action of ultraviolet light. Ordinary oxygen molecules absorb ultraviolet light and break into individual oxygen atoms. These have unpaired electrons, and are known as free radicals. They are very reactive.

The oxygen radicals can then combine with ordinary oxygen molecules to make ozone.

11

Ozone can also be split up again into ordinary oxygen and an oxygen radical by absorbing ultraviolet light.

This formation and breaking up of ozone is going on all the time. Taken together, these reactions stop a lot of harmful ultraviolet radiation penetrating the atmosphere to reach the surface of the Earth. The catalytic reaction we are interested in destroys the ozone and so stops it absorbing UV in this way. Chlorofluorocarbons (CFCs) like CF2Cl2, for example, were used extensively in aerosols and as refrigerants. Their slow breakdown in the atmosphere produces chlorine atoms chlorine free radicals. These catalyse the destruction of the ozone. This happens in two stages. In the first, the ozone is broken up and a new free radical is produced. The chlorine radical catalyst is regenerated by a second reaction. This can happen in two ways depending on whether the ClO radical hits an ozone molecule or an oxygen radical. If it hits an oxygen radical (produced from one of the reactions we've looked at previously):

Or if it hits an ozone molecule:

Because the chlorine radical keeps on being regenerated, each one can destroy thousands of ozone molecules.[5]

12 3.2.2. Homogenous catalysis in liquid phase Important examples of homogenous catalysis in liquid phase are acid base catalysysis. The common acid catalyst in water is the hydronium ion and the most common base catalyst is the hydroxyl ion. If a acid catalyses a reaction the reaction is said to be the subject of acid catalysis. Inversion of cane sugar and hydrolysis of esters are some common examples of acid catalysed reaction. However it was shown that different acids have different catalytic activity; hydrochloric acid has greater activity than acetic acid.so it is evident that the actual catalysts are H+ ions. The rate of reaction is found to be proportional to the concentration of H+ ions and the concentration of the reacting molecules or ions. 

H C12 H 22O11  H 2O  C6 H 12O6  C6 H 12O6

H CH 3COOC 2 H 5  H 2O  CH 3COOH  C2 H 5OH

Such reaction which are catalysed by certain acids are said to be specific acid catalysis. Similarly there are reactions which are catalysed by OH - ions only and are said to be specific hydroxyl ion catalysis. Conversion of acetone in to diaacetonyl alcohol is example of hydroxyl ion catalysis. 

CH 3COOCH3  CH 3COOCH 3 OH   Chs CCH 2C (CH 3) 2 OH

3.2.3. Enzymes catalysis Enzymes are homogeneous biological catalysts. They are complex protein molecules with three dimensional structures. These are responsible for catalyzing the chemical reaction in living organisms. The diameter of the enzyme molecules fall in the range of 10-100 nm. Enzymes are often present in colloidal states and are extremely specific in their catalytic function. Various enzyme catalyzed reactions are known. Some important examples are (1) Urease, an enzymes that catalyses the hydrolysis of urea but has no effect on the hydrolysis of substituted urea, e.g. methyl urea. NH 2CONH 2  H 2O urease  2 NH 3  CO2

(2) peptide, glycyl-L-glutamy1-L-tyrosine, is hydrolysed by an enzyme known as pepsin.

13

(3) Hydrolysis of starch in to maltose by diastase. 2(C6 H10O5 ) n  nH 2O diastase  nC12 H 22O11 Starch

Maltose

(4) Conversion of glucose in to ethanol by zymase present in the yeast C6 H12O6  H 2O zymase  2C2 H 5OH  2CO2

(5) Conversion of maltose in to glucose by maltase C12 H 22O11  H 2O maltase  2C6 H12O6

(6) Oxidation of alcohol to acetic acid by micoderma aceti

Micoderma aceti

C2 H5OH  O2 CH3COOH  H 2O Almost all enzymes fall in to two classes, the hydrolytic enzymes and the oxidation reduction enzymes. The hydrolytic enzymes appear to be complex acid base catalysis which accelerates the ionic reaction mainly due to the transfer of hydrogen ions. The oxidation reduction enzymes catalyze electron transfer perhaps through the formation of an intermediate radical. [6] 3.3. Heterogeneous catalysis: The catalyst is in a different phase from the reactants. . Usually the catalyst is a solid and the reactants are gases, and so the rate-limiting step occurs at the solid surface. Thus heterogeneous catalysis is also referred to as surface catalysis. A solid catalyst adsorbs reacting species onto its surface where the reaction takes place 2SO2(g) + O2(g) Solid V2O5 is the catalyst for this reaction.

2SO3(g)

14 Other example of heterogeneous catalysis is hydrogenation of an unsaturated organic compound such as ethane (C2H4) by metal catalysts such as Pt or Ni:

The currently accepted mechanism for this reaction involves weak bonding of both H2 and C2H4 to atoms on the metal surface. This is called adsorption. Ethene molecules are adsorbed on the surface of the nickel. The double bond between the carbon atoms breaks and the electrons are used to bond it to the nickel surface.

Hydrogen molecules are also adsorbed on to the surface of the nickel. When this happens, the hydrogen molecules are broken into atoms. These can move around on the surface of the nickel.

If a hydrogen atom diffuses close to one of the bonded carbons, the bond between the carbon and the nickel is replaced by one between the carbon and hydrogen.

15

That end of the original ethene now breaks free of the surface, and eventually the same thing will happen at the other end.

As before, one of the hydrogen atoms forms a bond with the carbon, and that end also breaks free. There is now space on the surface of the nickel for new reactant molecules to go through the whole process again.

[11]

These adsorbed H atoms can move across the metal surface, and eventually they combine with a C2H4 molecule, completing the reaction. Because adsorption and dissociation of H2 on a Pt surface is exothermic (ΔHm° >= –160 kJ mol–1), it can provide H atoms for further reaction without a large activation energy. By contrast, dissociation of gaseous H 2

16 molecules without a metal surface would require the full bond enthalpy (ΔHm° = +436 kJ mol–1). Clearly the metal surface makes a major contribution in lowering the activation energy. [5] 4. Autocatalysis This is the case of the generation of catalyst during the course of reaction. When one of the products of a reaction itself act as a catalyst for that reaction, the phenomenon is called autocatalysis. For example, oxidation of oxalic acid by KMnO 4 in the presence of sulphuric acid is initially slow, but the rate increase with progress of reaction, since managanous sulphate produce during the reaction acts as a catalyst for the reaction. 2 KMnO4  5 H 2C2O4  3H 2 SO4  2MnSO4  K 2 SO4  8H 2O  10CO2

Another example is the hydrolysis of ester by water where the product of hydrolysis the acid is a catalyst in the reaction. CH 3COOC 2 H 5  H 2O  CH 3COOH  C2 H 5OH

Heat, light, electricity or energy may alter the rate of reaction, but they are not catalyst. A catalyst must be a material particle. A catalyst takes part in the formation of the activated complex and is regenerated at the end. Thus a catalyst is both the reactant and a product of a reaction, a catalyst decreases the activation energy, by taking part in the activated complex formation and thus increases the rate. We can measure this effect by plotting the concentration of one of the reactants as time goes on. We can get a graph quite unlike the normal rate curve for a reaction.[6] Most reactions give a rate curve which looks like this

17

Concentrations are high at the beginning and so the reaction is fast - shown by a rapid fall in the reactant concentration. As things get used up, the reaction slows down and eventually stops as one or more of the reactants are completely used up. An example of autocatalysis gives a curve like this:

Note

18 Don't assume that a rate curve which looks like this necessarily shows an example of autocatalysis. There are other effects which might produce a similar graph. For example, if the reaction involved a solid reacting with a liquid, there might be some sort of surface coating on the solid which the liquid has to penetrate before the expected reaction can happen. A more common possibility is that we have a strongly exothermic reaction and aren't controlling the temperature properly. The heat evolved during the reaction speeds the reaction up. [7]

5. Characteristic of catalysis 

A catalyst remains unchanged in mass and chemical composition at the end of the reaction. However, its physical forms may be altered. Thus granular MnO 2 used as a catalyst in the thermal decomposition of KClO3 is left as a fine power at the end of the reaction.



A small amount of catalyst is sufficient to bring about a considerable extent of reaction. Hence a small amount may cause large quantities of materials to react. However this is not an essential criteria and also cannot be approved. The rate of catalytic reaction always dependent on the concentration of the catalyst when the amount of catalyst is small.



The rate of catalytic reaction depends upon the concentration of the catalyst. This is true when the concentration of catalyst is small, or concentration of reactants and catalyst are comparable.



A catalyst is more effective when finely divided. Thus a lump of platinum will have much less catalytic activity than colloidal platinum. Finally divided Ni is better catalyst than lumps of solid nickel.



A catalyst cannot affect the position of equilibrium, although it shortens the time required to establish equilibrium. However, a homogeneous catalyst can change

19 the equilibrium composition, but the effect is usually very small since a catalyst is generally present in small amount. 

A catalyst is in its action and thus it eliminate undesirable side product. When a several product are possible from same reactant, a catalyst can lower the activation energy of specific path. Thus it is able to select and direct the course of the reaction in a specific path. This property is specificity or selectivity. This is a property of changing the direction.



A catalyst cannot start a reaction, which is not thermodynamically feasible.



A catalyst is universal in its use. If it catalyses a reaction, once, it will catalyse that reaction always. What is more, catalysts are general. For example, silver oxide is generally used as a catalyst for hydration, and alumina and zircon is used for the dehydration.



Change of temperature alters the rate of catalytic reaction as it would do for the same reaction without a catalyst. Some catalysts are, however, physically altered by rise in temperature and hence their catalytic activity may be decreased. Thus as obvious from all the criteria for all purposes a catalyst can

(i)

increase the velocity of a reaction

(ii)

direct the reaction toward a specific product and

(iii)

Eliminate undivisable side product.

In short catalysis is the process of giving higher yield of purer product in shorter reaction time. 6. Theories of catalysis A stated earlier, the essential requirement for a reaction to occur is that the reacting molecules must acquire sufficient energy. In case of catalyst is added to the reaction, the energy required to activate the molecules is less than in the absence of a catalyst. Dye to the lower activation energy more molecules will take part in the reaction and hence the of catalyzed reaction would increase.the action of a catalyst can be explained by two different mechanism (1)

Intermediate compound formation theory

(2)

Adsorption theory

6.1. Intermediate compound formation theory

20 In this theory essentially two steps are involved. (i)

Combination of the catalyst with one or more of the reactants forming intermediate compound; and

(ii)

Decomposition of the intermediate compound or its combination with other reactant yielding the product and the catalyst back. Consider a reaction between reactant A and B giving the product AB. A + B

AB

This reaction is very slow and is catalysed by the presence of a catalyst X. the reaction will therefore proceed as A + X AX + B

AX

( intermediate compound)

AB + X

The formation of an intermediate compound AX is an easy reaction and needs low energy of activation thereby accelerating the rate of chemical reaction. Examples 

In lead chamber process for the manufacturing of sulphuric acid, the catalyst NO first forms the intermediate compound with oxygen. 2 NO  1 / 2O2  2 NO2 NO2  SO2  SO3  NO



In the preparation of diethyl ether from ethanol using concentrated H 2 SO4

C2 H 5 HSO4 is the first formed as an intermediate. C2 H 5OH  H 2 SO4  C2 H 5 HSO4  H 2O C2 H 5 HSO4  C2 H 5OH  C2 H 5OC2 H 5  H 2 SO4



The formation of water by combination of hydrogen and oxygen in the presence of copper as a catalyst is as follows.

21 2Cu  1 / 2O2  Cu 2O Cu 2O  H 2  H 2O  2Cu

6.1.1. Limitations of intermediate compound formation theory

This theory does not explain the cases of heterogeneous catalysis in general and more specifically the deactivation by a catalytic poison and activation by a promoter. 6.2. Adsorption theory A large number of gaseous reactions take place in the presence of solid catalysts. The surface of the catalyst has certain active centers due to the unsaturation of valences. Appreciable quantities of a reactant molecules are adsorbed or retained by solid surface at these active centers and the reaction occur at the surface of a solid. For this reason this type of catalysis is sometime referred to as the contact catalysis. The adsorb molecules from some sort of an active complex on the surface, which then decompose forming the products. The products are ultimately desorbed from the surface. A catalytic reaction involves the following steps (i) Diffusion of the reactant from the bulk on to the surface. (ii) Adsorption of the reactant on the surface of the catalyst. (iii)

Activation of the adsorbed reactants leading to a reaction in the adsorbed phase.

(iv)

Description of the product from the surface of the catalyst.

(v)

Diffusion of the product away from the surface of the catalyst.

Any one of these steps may be slowest and consequently the rate determining but generally step (iii) is the rate controlling step, Due to the adsorption the concentration of the reactant tend to increase on the surface of the catalyst and according to the law of mass action the rate of the reaction will increase. Further more adsorption being an exothermic process the heat evolved during adsorption is utilized in the activation of the surface reaction. Adsorption may also lead to proper orientation of the reacting molecules, partial loosening of the bonds in the finally divided state and thus requiring only small energy to form the activated complex.

22 Adsorption theory can explain the enhanced catalytic action of a catalyst in the finely divided state. It is due to the larger surface area available for adsorption and also the formation of more active centers.

It has observed that chemisorption plays an important role in the surface catalysis. The stronger forces are operative in the chemisorptions which tend to bring about a partial loosening of the bonds in the adsorbed reacting molecules and very small energy is required to form the activated complex for the reaction. It has been found that there are some active centers on the surface of the catalyst where the adsorption take place, X-ray studies of catalysts show that there are unsatisfied forces present on the surface of the catalyst which can be used to attach molecules. The surface of the nickel catalyst may be represented as follow.

Isolated nickel atoms and discontinuities probably act as active centers. This explains the enhanced catalytic activity of finally divided catalysts.[2,6] 7. Applications

23 The rate of reaction for some chemical process may be sped up in one of two ways. Either temperature can be increased or the mechanism can be altered so as to lower the activation energy of the reaction. The latter method may be can be done with the implementation of a catalyst. A catalyst is an intermediate in a chemical reaction that participates in the activated complex for the rate limiting step. There are three main features of catalysts: 1. The catalyst allows a reaction to proceed via an alternative mechanism. 2. For every step in the mechanism in which the catalyst appears as a product, there is another step in which the catalyst is a reactant. 3. The catalyst increases the rate of reaction by lowering the activation energy of the reaction. 7.1. Catalysis in Fluid Catalytic Cracking One very important example of the use of catalysis is fluid catalytic cracking. This process has been fundamental in the last few decades in meeting the high demands for high octane liquid fuels, such as gasoline. In fact, without FCC, methods of transportation such as cars and buses in their present forms would not be sustainable due to such high costs in fuel. From this fact alone, it is obvious that FCC is an essential process that has an important impact on many areas, including industry, transportation, and multiple aspects of everyday life. When petroleum is collected from the stores in the earth, its raw form contains many different hydrocarbon compounds. These hydrocarbons can contain anywhere from five to forty carbon atoms. Historically, this small fact has caused a bit of a dilemma for the energy industry. Hydrocarbons with less carbons in the chain are much more valuable than the long hydrocarbons with carbon chains of twenty or more.

24

These more compact hydrocarbons are especially valuable in a culture highly dependent on cars buses, jet planes, and other such aspects of life that require a compact and potent source of fuel.Due to this imbalance in demand and the excess of heavier hydrocarbons, the process of fluid catalytic cracking was implemented. Generally speaking, a longer carbon chain will have a significantly higher boiling point. This basis forms the means by which the various hydrocarbons in petroleum can be distinguished and separated. As the goal of FCC is to break down the larger hydrocarbons into smaller, more valuable compounds, the portion of the crude oil that is used in FCC (the feedstock) is all those hydrocarbons with an initial boiling point of roughly 340 degrees Celsius or higher. In an equivalent sense, the feedstock contains all hydrocarbons with a molecular weight ranging from roughly 200 to 600 grams per mole. In fluid catalytic cracking, these hydrocarbons are vaporized and broken at high temperatures in the presence of a special catalyst. The purpose of the catalyst is to make the process more efficient and thus economically viable. 7.1.1. Catalyst Specifications The most efficient form for the FCC process is highly desired and has been avidly searched for. One factor that effects the overall efficiency very significantly is the substance that is used as the catalyst. After many years of trial and error, a set of criteria

25 for an efficient FCC catalyst has been established. The best form is a fine powder with a density ranging from 0.80 to 0.96 g/ml and an average particle size of 60-100 microns. The catalyst should be highly reactive, stable at high temperatures, and retain large pore sizes. All FCC catalysts contain a crystalline zeolite. In essence, the zeolite component acts as a molecular sieve that only allows a certain size range of hydrocarbons to enter its lattice.

Fig.

A schematic flow diagram of a Fluid Catalytic Cracking unit as used in

petroleum refineries The catalytic activity sites are provided by the matrix component of the catalyst, which contains amorphous alumina that are capable of cracking the larger feedstock molecules. Other components are present in order to ensure that the catalyst maintains strength and stability. Due to the nature of the catalyst components, it is highly important that the catalyst is not introduced to feedstock with metal contaminants. Even

26 concentrations in the range of a few ppm of these contaminants can have detrimental effects on the performance of the FCC catalyst. Once an efficient catalyst is found, it is used extensively as a main ingredient. For every one kilogram of feedstock in the processing unit, there are five kilograms of catalyst compound added. Thus, a comparable processing unit might circulate 55,900 metric tons of catalyst each day. Fluid catalytic cracking is a form of applied catalysis that has become essential to the energy industry. Its ability to recycle large hydrocarbons present in petroleum into more valuable and smaller hydrocarbons has allowed for demand imbalances to be reconciled. During the year 2007, in the United States alone, over 5,300,000 barrels of crude oil were processed using FCC. This astounding data proves the importance of fluid catalytic cracking to industry, and the entire process could not be implemented without a well devised incorporation of catalysis. [12] 7.1.2. Chemistry involve in catalytic cracking Olefins or alkenes, which are unsaturated straight-chain or branched hydrocarbons, do not occur naturally in crude oil.

Fig: Diagrammatic example of the catalytic cracking of petroleum hydrocarbons

27 In plain language, the fluid catalytic cracking process breaks large hydrocarbon molecules into smaller molecules by contacting them with powdered catalyst at a high temperature and moderate pressure which first vaporizes the hydrocarbons and then breaks them. The cracking reactions occur in the vapor phase and start immediately when the feedstock is vaporized in the catalyst riser. Figure is a very simplified schematic diagram that exemplifies how the process breaks high boiling, straight-chain alkane (paraffin) hydrocarbons into smaller straightchain alkanes as well as branched-chain alkanes, branched alkenes (olefins) and cycloalkanes (naphthenes). The breaking of the large hydrocarbon molecules into smaller molecules is more technically referred to by organic chemists as scission of the carbonto-carbon bonds. As above in Figure, some of the smaller alkanes are then broken and converted into even smaller alkenes and branched alkenes such as the gases ethylene, propylene, butylenes, and isobutylenes. Those olefinic gases are valuable for use as petrochemical feedstocks. The propylene, butylene and isobutylene are also valuable feedstocks for certain petroleum refining processes that convert them into high-octane gasoline blending components. As also depicted in Figure, the cycloalkanes (naphthenes) formed by the initial breakup of the large molecules are further converted to aromatics such as benzene, toluene, and xylenes, which boil in the gasoline boiling range and have much higher octane ratings than alkanes. By no means does Figure include all the chemistry of the primary and secondary reactions taking place in the fluid catalytic process. There are a great many other reactions involved. 7.2. Hydrogenation An example of catalytic action is found in the hydrogenation of alkenes. The alkene (2) adsorbs by forming two bonds with the surface ( 3), and on the same surface there may be adsorbed H atoms. When an encounter occurs, one of the alkene–surface bonds is broken (forming 4 or 5) and later an encounter with a second H atom releases the fully hydrogenated hydrocarbon, which is the thermodynamically more stable species. The evidence for a two-stage reaction is the appearance of different isomeric

28 alkenes in the mixture. The formation of isomers comes about because, while the hydrocarbon chain is waving about over the surface of the metal, an atom in the chain might chemisorb again to form (6) and then desorb to ( 7), an isomer of the original molecule. The new alkene would not be formed if the two hydrogen atoms attached simultaneously.

A

major

catalytic

industrial

application

hydrogenation

is

to

of the

formation of edible fats from vegetable and animal oils. Raw oils obtained from sources such as the Soya bean have the structure CH 2 (OOCR)CH(OOCR′) CH 2 (OOCR″ ), where R, R ′, and R ″ are long-chain hydrocarbons with several double bonds. One disadvantage of the presence of many double bonds is that the oils are susceptible to atmospheric oxidation, and therefore are liable to become rancid. The geometrical configuration of

29 the chains is responsible for the liquid nature of the oil, and in many applications a solid fat is at least much better and often necessary. Controlled partial hydrogenation of an oil with a catalyst carefully selected so that hydrogenation is incomplete and so that the chains do not isomerize (finely divided nickel, in fact), is used on a wide scale to produce edible fats. The process, and the industry, is not made any easier by the seasonal variation of the number of double bonds in the oils. [13] 7.3. Oxidation Catalytic oxidation is also widely used in industry and increasingly in pollution control. Although in some cases it is desirable to achieve complete oxidation, as in the elimination of nitrogen oxide from engine emission nitric acid from ammonia, in others partial oxidation is the aim. For example, the complete oxidation of propane to carbon dioxide and water is wasteful, but its partial oxidation to propenal (acrolein, CH2=CHCHO) is the start of important industrial processes. Likewise, the controlled oxidations of ethane to ethanol, ethanal (acetaldehyde), and (in the presence of acetic acid or chlorine) to chloroethene (vinyl chloride, for the manufacture of PVC), are the initial stages of very important chemical industries.Some of these oxidation reactions are catalysed by d-metal oxides of various kinds. [6] 7.4. Water-gas shift (WGS) reaction The water gas shift reaction was first used industrially at the beginning of the 20th century. Today the WGS reaction is used primarily to produce hydrogen that can be used for further production of methanol and ammonia. WGS reaction: CO + H2O ↔ H2 + CO2 The reaction refers to carbon monoxide (CO) that reacts with water (H2O) to from carbon dioxide (CO2) and hydrogen (H2). The reaction is exothermic with ΔH= -41.1 kJ/mol and have an adiabatic temperature rise of 8–10 °C per percent CO converted to CO2 and H2. The most common catalysts used in the water-gas shift reaction are the high temperature shift (HTS) catalyst and the low temperature shift (LTS) catalyst. The HTS catalyst consists of iron oxide stabilized by chromium oxide, while the LTS catalyst is based on copper. The main purpose of the LTS catalyst is to reduce CO content in the reformate which is especially important in the ammonia production for high yield of H 2. Both

30 catalysts are necessary for thermal stability, since using the LTS reactor alone increases exit-stream temperatures to unacceptable levels. The equilibrium constant for the reaction is given as: Kp= (pH2 x pCO2)/ (pCO x pH2O) Kp= e((4577.8K/T-4.22)) Low temperatures will therefore shift the reaction to the right, and more products will be produced. The equilibrium constant is extremely dependent on the reaction temperature, for example is the Kp equal to 228 at 200 °C, but only 11.8 at 400 °C. The WGS reaction can be performed both homogenously and heterogeneously, but only the heterogeneously way is used commercially. 7.4.1. High temperature shift (HTS) catalyst The first step in the WGS reaction is the high temperature shift which is carried out at temperatures between 320 °C and 450 °C. As mentioned before, the catalyst is a composition of iron-oxide, Fe2O3(90-95%), and chromium oxides Cr2O3 (5-10%) which have an ideally activity and selectivity at these temperatures. When preparing this catalyst, one of the most important step is washing to remove sulfate that can turn into hydrogen sulfide and poison the LTS catalyst later in the process. Chromium is added to the catalyst to stabilize the catalyst activity over time and to delay sintering of iron oxide. Sintering will decrease the active catalyst area, so by decreasing the sintering rate the life time of the catalyst will be extended. The catalyst is usually used in pellets form, and the size play an important role. Large pellets will be strong, but the reaction rate will be limited. In the end, the dominate phase in the catalyst consist of Cr 3+ in α-Fe2O3 but the catalyst is still not active. To be active α-Fe2O3 must be reduced to Fe3O4 and CrO3 must be reduced to Cr2O3. This usually happens in the reactor start-up phase and because the reduction reactions are exothermic the reduction should happen under controlled circumstances. The lifetime of the iron-chrome catalyst is approximately 3–5 years, depending on how the catalyst is handled. Even though the mechanism for the HTS catalyst has been done a lot of research on, there is no final agreement on the kinetics/mechanism. Research has narrowed it down to

31 two

possible

mechanisms:

a

regenerative

redox

mechanism

and

an

adsorptive(associative) mechanism. The redox mechanism is given below: First a CO molecule reduces an O molecule, yielding CO2 and a vacant surface center: CO+(O) →CO2 + (*) The vacant side is then reoxidized by water, and the oxide center is regenerated: H2O+(*)→H2+ (O) The adsorptive mechanism assumes that format species is produced when an adsorbed CO molecule reacts with a surface hydroxyl group: H2O →OH(ads)+ H(ads) CO(ads)+ OH(ads)→COOH (ads) The format decomposes then in the presence of steam: COOH(ads)→CO2+H(ads) 2H(ads)→H2 7.4.2. Low temperature shift (LTS) catalyst The low temperature process is the second stage in the process, and is designed to take advantage of higher hydrogen equilibrium at low temperatures. The reaction is carried out between 200 °C and 250 °C, and the most commonly used catalyst is based on copper. While the HTS reactor used an iron-chrome based catalyst, the copper-catalyst is more active at lower temperatures thereby yielding a lower equilibrium concentration of CO and a higher equilibrium concentration of H2. The disadvantage with a copper catalysts is that it is very sensitive when it comes to sulfide poisoning, a future use of for example a cobalt- molybdenum catalyst could solve this problem. The catalyst mainly used in the industry today is a copper-zinc-alumina (Cu/ZnO/Al2O3) based catalyst. Also the LTS catalyst has to be activated by reduction before it can be used. The reduction reaction CuO + H2 →Cu + H2O is highly exothermic and should be conducted in dry gas for an optimal result.

32 As for the HTS catalyst mechanism, two similar reaction mechanisms are suggested. The first mechanism that was proposed for the LTS reaction was a redox mechanism, but later evidence showed that the reaction can proceed via associated intermediates. The different intermediates that is suggested are: HOCO, HCO and HCOO. In 2009 there are in total three mechanisms that are proposed for the water-gas shift reaction over Cu(111), given below. Intermediate mechanism (usually called associative mechanism): An intermediate is first formed and then decomposes into the final products: CO + (species derived from H2O) →Intermediate→CO2 Associative mechanism: CO2 produced from the reaction of CO with OH without the formation of an intermediate: CO + OH →H + CO2 Redox mechanism: Water dissociation that yields surface oxygen atoms which react with CO to produce CO2: H2O→O (surface) O (surface) + CO → CO2 It is not said that just one of these mechanisms is controlling the reaction, it is possible that several of them are active. Q.-L. Tang et al. has suggested that the most favorable mechanism is the intermediate mechanism (with HOCO as intermediate) followed by the redox mechanism with the rate determining step being the water dissociation. For both HTS catalyst and LTS catalyst the redox mechanism is the oldest theory and most published articles support this theory, but as technology has developed the adsorptive mechanism has become more of interest. One of the reasons to the fact that the literature is not agreeing on one mechanism can be because of experiments are carried out under different assumptions. 7.5. Production of Carbon Monoxide

33 CO must be produced for the WGS reaction to take place. This can be done in different ways from a variety of carbon sources such as: -passing steam over coal: C + H2O = CO +H2 -steam reforming methane, over a nickel catalyst: CH4 + H2O = CO +3H2 -or by using biomass. Both the reactions shown above are highly endothermic and can be coupled to an exothermic partial oxidation. The products of CO and H2 are known as syngas. When dealing with a catalyst and CO, it is common to assume that the intermediate COMetal is formated before the intermediate reacts further into the products. When designing a catalyst this is important to remember. The strength of interaction between the CO molecule and the metal should be strong enough to provide a sufficient concentration of the intermediate, but not so strong that the reaction will not continue. CO is a common molecule to use in a catalytic reaction, and when it interacts with a metal surface it is actually the molecular orbitals of CO that interacts with the d-band of the metal surface. When considering a molecular orbital(MO)-diagram CO can act as an σ-donor via the lone pair of the electrons on C, and a π-acceptor ligand in transition metal complexes. When a CO molecule is adsorbed on a metal surface, the d-band of the metal will interact with the molecular orbitals of CO. It is possible to look at a simplified picture, and only consider the LUMO (2π*) and HOMO (5σ) to CO. The overall effect of the σ-donation and the π- back donation is that a strong bond between C and the metal is being formed and in addition the bond between C and O will be weakened. The latter effect is due to charge depletion of the CO 5σ bonding and charge increase of the CO 2π* antibonding orbital. When looking at chemical surfaces, many researchers seems to agree on that the surface of the Cu/Al2O3/ZnO is most similar to the Cu(111) surface. Since copper is the main catalyst and the active phase in the LTS catalyst, many experiments has been done with copper. In the figure given here experiments has been done on Cu(110) and Cu(111). The figure shows Arrhenius plot derived from reaction rates. It can be seen from the figure

34 that Cu(110) shows a faster reaction rate and a lower activation energy. This can be due to the fact that Cu(111) is more closely packed than Cu(110). 7.6. Methanol production Production of methanol is an important industry today and methanol is one of the largest volume carbonylation products. The process uses syngas as feedstock and for that reason the water gas shift reaction is important for this synthesis. The most important reaction based on methanol is the decomposition of methanol to yield carbon monoxide and hydrogen. Methanol is therefore an important raw material for production of CO and H 2 that can be used in generation of fuel.

BASF was the first company (in 1923) to produce methanol on large-scale, then using a sulfur-resistant ZnO/Cr2O3 catalyst. The feed gas was produced by gasification over coal. Today the synthesis gas is usually manufactured via steam reforming of natural gas. The most effective catalysts for methanol synthesis are Cu, Ni, Pd and Pt, while the most common metals used for support are Al and Si. In 1966 ICI (Imperial Chemical Industries) developed a process that is still in use today. The process is a low-pressure process that uses a Cu/ZnO/Al2O3 catalyst where copper is the active material. This catalyst is actually the same that the low-temperature shift catalyst in the WGS reaction is using. The reaction described below is carried out at 250 °C and 5-10 MPa: CO+2H2→CH3OH (l) CO2+3H2→CH3OH (l) +H2O (l) Both of these reactions are exothermic and proceeds with volume contraction. Maximum yield of methanol is therefore obtained at low temperatures and high pressure and with use of a catalyst that has a high activity at these conditions. A catalyst with sufficiently high activity at the low temperature does still not exist, and this is one of the main reasons that companies keep doing research and catalyst development. A reaction mechanism for methanol synthesis has been suggested by Chinchen et al.:

35 CO2→CO2* H2→2H* CO2*+ H*→HCOO* HCOO*+3H*→CH3OH+ O* CO+ O*→CO2 H2 + O*→H2O Today there are four different ways to catalytically obtain hydrogen production from methanol, and all reactions can be carried out by using a transition metal catalyst (Cu, Pd): 7.6.1. Steam reforming The reaction is given as: CH3OH(l)+ H2O (l) →CO2+ 3H2 ΔH= +131 KJ/mol Steam reforming is a good source for production of hydrogen, but the reaction is endothermic. The reaction can be carried out over a copper-based catalyst, but the reaction mechanism is dependent on the catalyst. For a copper-based catalyst two different reaction mechanisms have been proposed, a decomposition-water-gas shift sequence and a mechanism that proceeds via methanol dehydrogenation to methyl formate. The first mechanism aims at methanol decomposition followed by the WGS reaction and has been roposed for the Cu/ZnO/Al2O3: CH3OH+ H2O →CO2+ 3H2 CH3OH→CO+ 2H2 CO+ H2O →CO2+H2 The mechanism for the methyl format reaction can be dependent of the composition of the catalyst. The following mechanism has been proposed over Cu/ZnO/Al2O3: 2CH3OH→CH3OCHO+ 2H2

36

CH3OCHO+H2O→HCOOH+CH3OH HCOOC→CO2+H2 When methanol is almost completely converted CO is being produced as a secondary product via the reverse water-gas shift reaction.

7.6.2. Methanol decomposition The second way to produce hydrogen from methanol is by methanol decomposition: CH3OH(l)→ CO + 2H2 ΔH= +128 KJ/mol As the enthalpy shows, the reaction is endothermic and this can be taken further advantage of in the industry. This reaction is the opposite of the methanol synthesis from syngas, and the most effective catalysts seems to be Cu, Ni, Pd and Pt as mentioned before. Often, a Cu/ZnO-based catalyst is used at temperatures between 200 and 300 °C but a production of by-product as dimethyl ether, methyl format, methane and water is common. The reaction mechanism is not fully understood and there are two possible mechanism proposed (2002) : one producing CO2 and H2 by decomposition of formate intermediates and the other producing CO and H2 via a methyl formate intermediate. 7.6.3. Partial oxidation Partial oxidation is a third way for producing hydrogen from methanol. The reaction is given below, and is often carried out with air or oxygen as oxidant : CH3OH(l) + 1/2 O2 → CO2 + 2H2 ΔH=-155 KJ/mol The reaction is exothermic and has, under favorable conditions, a higher reaction rate than steam reforming. The catalyst used is often Cu (Cu/ZnO) or Pd and they differ in qualities such as by-product formation, product distribution and the effect of oxygen partial pressure. 7.6.4. Combined reforming Combined reforming is a combination of partial oxidation and steam reforming and is the last reaction that is used for hydrogen production. The general equation is given below:

37 (s+p)CH3OH(l) +sH2O(l) + 1/2pO2→ (s+p)CO2 +(3s+2p)H2 s and p are the stoichiometric coefficients for steam reforming and partial oxidation, respectively. The reaction can be both endothermic and exothermic determined by the conditions, and combine both the advantages of steam reforming and partial oxidation. 7.6.5. Ammonia synthesis Ammonia synthesis was discovered by Fritz Haber, by using an iron catalysts. The ammonia synthesis advanced between 1909 and 1913, and two important concepts were developed; the benefits of a promoter and the poisoning effect (see catalysis for more details). Ammonia production was one of the first commercial processes that required the production of hydrogen, and the cheapest and best way to obtain hydrogen was via the water-gas shift reaction. The, Haber-Bosch process is the most common process used in the ammonia industry. A lot of research has been done on the catalyst used in the ammonia process, but the main catalyst that is used today is not that dissimilar to the one that was first developed. The catalyst the industry use is an promoted iron catalyst, where the promoters can be K 2O (Potassium oxide), Al2O3 ( Aluminium oxide) and CaO (Calcium oxide ) and the basic catalytic material is Fe. The most common is to use fixed bed reactors for the synthesis catalyst. The main ammonia reaction is given below: N2+ 3H2↔ 2NH3 The produced ammonia can be used further in production of nitric acid via the Ostwald process.[14] 8. Catalyst characterization techniques Experimental surface characterization methods were used to study ageing-induced changes in the active metals and washcoat oxides. The characterization techniques included both microscopic and spectroscopic methods, such as Scanning Electron Microscopy (SEM) and X-ray Diffraction (XRD). Several characterization techniques are available to study solid surfaces and the properties of catalysts, and no single

38 characterization method can be used to explain the basis for the catalyst deactivation phenomena of three-way catalysts . 8.1. Scanning Electron Microscopy Scanning Electron Microscopy (SEM) was used for high magnification imaging and elemental analysis. A Jeol JSM-6400 scanning electron microscope equipped with an energy dispersive spectrometer (EDS) was used for the analysis. In the pretreatment stage, flat pieces of fresh and aged catalysts were cut, and either potted in epoxy or fastened with a carbon tape in order to obtain side or top views of the catalyst respectively. Catalysts were polished down to 1 µm using diamond paste and coaled prior the analysis to avoid the accumulation of charge. The accelerating voltage and current in the measurements were 15 kV and 12 nA, respectively, and the resolution of the instrument was 3.5 nm (35 kV). 8.2. Physisorption analyses Measurements of gas adsorption isotherms are widely used for determining the surface area and pore size distribution of solids. The use of nitrogen as the adsorptive gas is recommended if the surface areas are higher than 5 m2/g (Serwicka 2000). The first step in the interpretation of a physisorption isotherm is to identify the isotherm type. This in turn allows for the possibility to choose an appropriate procedure for evaluation of the textural properties. Non-specific Brunauer-Emmett-Teller (BET) method is the most commonly used standard procedure to measure surface areas, in spite of the oversimplification of the model on which the theory is based. The BET equation is applicable at low p/p0 range and it is written in the linear form (Wachs 1992): Equation

where

39 p is the sample pressure, p0 is the saturation vapour pressure, na is the amount of gas adsorbed at the relative pressure p/p0, nam is the monolayer capacity, and C the so-called BET constant. Equation can be applied for determining the surface areas and pore volumes from adsorption isotherms. The pore size distributions can be calculated from desorption isotherms. The pores are usually classified according to their widths as micropores (diameter less than 2 nm), mesopores (diameter between 2 and 50 nm) and macropores (diameter exceeding 50 nm) (Hayes & Kolaczkowski 1997). Several approaches have been developed to assess the micro- and mesoporosity, and to compute pore size distribution from the adsorption-desorption data. All of these involve a number of assumptions, e.g. relating to pore shape and mechanism of pore filling. (Serwicka 2000) In this physisorption measurements were carried out to characterize catalysts before and after the ageings. Specific surface areas (m2/g) and pore volumes were measured according to the standard BET method, as described above, by using a Coulter Omnisorp 360CX. The specific surface areas and pore volumes were obtained from N2 adsorption isotherms at -196°C by assuming the cylindrical shape of pores. Catalysts were outgassed in a vacuum at 140°C overnight before the measurements. All the BET values in this study were measured within a precision of ± 5%. Pore size distributions for micropores as well as meso- and macropores were calculated from N2 -desorption isotherms by differential HK (Horvath-Kawazoe) and BJH (Barrett-Joyner-Hallender) methods respectively (see Anon 1992). Since the monoliths showed systematically lower BET values than the crushed samples after similar ageing procedures, all the BET values presented in this thesis have been determined for the metallic monoliths with a standard shape and mass.

40 8.3. Chemisorption analyses Chemisorption measurements were carried out in order to determine the dispersions of Pd and Rh metal particles, monolayer capacities and the amount of active metal in the catalysts. Hydrogen and carbon monoxide were used as the adsorbate gases. H 2chemisorption and CO-chemisorption experiments were carried out close to room temperature (30°C) by volumetric adsorption method by using a Coulter Omnisorp 360CX and a Sorptomatic 1900), respectively. The accuracy of the measurements was estimated to be better than ± 5%. The experimental procedure for the H 2-chemisorption measurements is presented in Table . In the chemisorption procedure, the temperature ramping rate of the furnace was 10°C/min. As shown in Table, the adsorption of H 2 was measured twice. The difference between these two measurements was assumed to be the amount of irreversibly adsorbed H2, which is further used to calculate the dispersion values. The experimental procedure for the H2-chemisorption measurements 1. Flow of He at 150°C for 5 minutes followed by 10 minutes at 375°C 2. Evacuation at 375°C for 10 minutes 3. Reduction in flowing H2 at 375°C for 10 minutes followed by 5 minutes at 400°C 4. Evacuation at 400°C for 20 minutes followed by 10 minutes at 30 °C 5. Leak test at 30 °C 6. First analysis with H2 at 30 °C 7. Evacuation at 30 °C for 30 minutes 8. Second analysis with H2 at 30 °C Chemisorption has long been employed as a valuable technique for rapid evaluation of the active metal dispersions and hence the particle sizes of supported metals (Gasser 1985). This method has, however, undergone severe criticism, since the underlying assumptions of the stoichiometry between adsorbate gas and precious metal and the particle geometry may not be true, in particular in the case of small particles (Di Monte et al. 2000). Furthermore, in the case of metal oxides (such as CeO 2 and Ce-Zr-mixed oxides) in contact with active metals, adsorbed H 2-molecules can also diffuse from the

41 active metal particles to the washcoat. This spillover effect can be reduced by lowering the adsorption temperature, as has been reported by Bernal et al. (1993) and Fornasiero et al. (1995). Active metal dispersions and particle sizes are calculated by assuming the stoichiometry factor between chemisorbed gas molecules and surface metal atoms. In this thesis, chemisorption measurements are based on the assumption of the stoichiometry of 2:1 for H2 and the stoichiometry 1:1 for CO adsorption, respectively, and regardless of the particle size. The stoichiometric ratio may depend on the precious metal particle size, a reason why caution should be exercised when comparing the dispersion values of different catalysts. However, it is assumed that all the aged catalysts as well as the fresh catalyst exhibit rather large metal particle sizes due the low dispersion values (below 30%). Therefore, the changes in dispersion values presented in this thesis reflect the structural changes induced by ageings, such as the sintering of the precious metals. As well, the chemical correctness behind the stoichiometry assumptions is not relevant because, in this case, the relative dispersion values are more interesting than the absolute ones. 8.4. Activity measurements Catalytic activities were determined by laboratory scale light-off experiments to compare the catalysts after the ageings. Catalyst light-off is determined as the temperature of 50% conversion, which is used to indicate the efficiency of an automotive exhaust gas catalyst (the lower the light-off temperature, the more active the catalyst is). In addition to lightoff temperatures, the conversions of CO and NO at 400°C were also determined. The experimental set-up for the activity measurements is presented in Figures. Catalytic activities were determined by using a simple model reaction: the reduction of NO by CO in lean and rich conditions. The composition of test gas mixture is presented in Table. Before the measurements, the catalysts were reduced in a hydrogen flow at 500°C for 10 minutes, followed by 15 minutes at 550°C. Activity measurements were carried out at atmospheric pressure by using a cylindrical catalyst with a volume of 1.4 cm 3 (length 28

42 mm and diameter 8 mm). In the measurements, the gas-solid reactor system equipped with mass flow controllers, magnetic valves for flow selection, tubular furnace with a quartz reactor and analysis instruments were used. The total gas flow during the experiments was 1 dm3/min corresponding to the feed gas hourly space velocity (GHSV) of 43 000 h-1. The temperature of a catalyst was increased from room temperature up to 400°C, with a linear heating rate of 20°C/min. The concentrations of CO, NO, CO 2, N2O and NO2 as a function of temperature were measured every 5 seconds by an FTIR gas analyser and the gas flow was controlled by mass flow controllers. Furthermore, the effect of poisoning on the catalytic activity was evaluated by changing the flow direction in the catalyst. Blank tests were carried out with the uncoated metal foil to ensure the inactivity of metal foil in the thermal treatments. In the following discussion, differences larger than ± 5°C in the light-off temperatures and ± 1% in the conversions of CO and NO can be regarded as statistically significant.

Fig. Experimental set-up for the activity measurements

43

Composition of the test gas mixture for the activity measurements Component Lean

Rich

CO

800 ppm

1200 ppm

NO

1000 ppm 1000 ppm

N2

Balance

Balance

Fig. Activity measurement system equipped with the GASMETTM gas analyzer

The activity of some aged pre-catalysts and main catalysts was also tested at Kemira Metalkat Oy, Finland. In these measurements, the conversions of CO, HC and NO X were

44 measured as a function of catalyst’s temperature using a test gas mixture simulating the real exhaust gas composition. Catalyst light-off temperatures (T50 values) as well as conversions at 400°C were determined. Furthermore, OSC was measured for fresh and aged monoliths by CO-O2 exchange experiments at constant adsorption temperatures of 450°C, 600°C and 750°C. The consumption and the adsorption of O 2 and CO were determined by mass spectrometer. (Härkönen et al. 2001) 8.5. Chemical analyses Chemical analysis provides the information on elemental composition of the catalyst. The ‘wet’ and ‘dry’ analyses were performed due to low quantities of poisons present in the catalysts. The dry analysis was carried out beyond the SEM-EDS sensitivity, as described in section 8.1. In a wet analysis, the fresh and aged catalysts were dissolved in an acidic solution in order to determine the quantities of the most important catalytic poisons (Ca, P, S, Pb, Mg and Zn) and the amounts of active metals (Pd and Rh). These elements were typically present in small quantities, and unevenly distributed in the catalyst. Therefore, separated samples were prepared from the inlet and outlet parts of the engine-aged and vehicle-aged catalysts. For the chemical analysis, 0.010 to 0.050 g of the washcoat (scraped from the monolith) was dissolved in an acidic solution (HNO 3, HCl, HF and H3BO3) and subjected to the digestion of the sample in the microwave oven (Milestone MLS 1200). This resulted partly in an incomplete dissolution of the analysed solids. The decomposed sample was analysed quantitatively by plasma atomic emission spectrometry (Pye Unicam 7000 ICP-AES). [15]

45

References 1. http://www.chemguide.co.uk/physical/catalysis/introduction.html 2. Chaudhry. G. Rasool, A Text Book of Physical Chemistry, (2009). Lahore: Azeem Publishers. 3. http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/HeterogeneousCatalysis/chemprime/CoreChem3ACatalysis-1005.html 4. http://www.chemguide.co.uk/physical/catalysis/introduction.html 5. http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/HeterogeneousCatalysis-1005.html 6. Bhatti. H. Nawaz, Principles of Physical Chemistry, (2008). Lahore: The Caravan Book House. 7. http://www.chemguide.co.uk/physical/catalysis/introduction.html 8. Moora. J. Walter, Physical Chemistry, (1991). 9. http://www.chemguide.co.uk/physical/catalysis/introduction.html 10. http://www.chem.canterbury.ac.nz/letstalkchemistry/electronicversion/ electronicversionnew/chapter16/catalyst.shtml 11. http://www.chemguide.co.uk/physical/catalysis/introduction.html 12. http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/HeterogeneousCatalysis/chemprime/Catalysis_in_Everyday_Life-1005.html 13. Atkin. P. Physical Chemistry, (2006). New York: W. H. Freeman & Company. 14. http://en.wikipedia.org/wiki/Industrial_catalysts 15. http://herkules.oulu.fi/isbn9514269543/html/x900.html