Catalysis: Gold as an Oxidation Catalyst

1. Catalysis Gold oxidation catalyst Photcatalyst in the form of nanoflower Berzelius is credited with identifying the chemical elements silicon, selenium, thorium, and cerium. Students working in Berzelius's laboratory also discovered lithium, and vanadium Berzelius is credited with originating the chemical terms "catalysis", "polymer", "isomer" and "allotrope" Jöns Jacob Berzelius Born 20 August 1779 Väversunda, Östergötland, Sweden Died 7 August 1848 (aged 68) Stockholm, Sweden

2. Estimates are that 90% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture.[1] In 2005, catalytic processes generated about $900 billion in products worldwide.[2] Anything that increases the rate of a process is a "catalyst", a term derived from Greek καταλύειν, meaning "to unite" The phrase catalyzed processes was coined by Jöns Jakob Berzelius in 1836 to describe reactions that are accelerated by substances that remain unchanged after the reaction. Humphry Davy discovered the use of platinum in catalysis. Probably the most important metal in catalysis. Wilhelm Ostwald at Leipzig University started a systematic investigation into reactions that were catalyzed by the presence of acids and bases; Ostwald was awarded the 1909 Nobel Prize in Chemistry. Other recent Noble prices in Chemistry for Catalysis: 2011 for palladium-catalyzed cross couplings in organic synthesis, 2007 for chemical processes on solid surfaces, 2005 for the development of the metathesis method in organic synthesis, 2001 for chirally catalysed oxidation and reduction reactions • Recognizing the Best in Innovation: Breakthrough Catalyst". R&D Magazine, September 2005, pg 20. • http://www.climatetechnology.gov/library/2005/tech-options/tor2005-143.pdf

3. 1. Catalysts work by providing an (alternative) mechanism involving a different transition state and lower activation energy. Consequently, more molecular collisions have the energy needed to reach the transition state. Hence, catalysts can enable reactions that would otherwise be blocked or slowed by a kinetic barrier. The catalyst may increase reaction rate or selectivity, or enable the reaction at lower temperatures. 2. In the catalyzed elementary reaction, catalysts do not change the extent of a reaction: they have no effect on the chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are both affected. 3. The productivity of a catalyst can be described by the turn over number (or TON) and the catalytic activity by the turn over frequency (TOF), which is the TON per time unit. 4. The catalyst stabilizes the transition state more than it stabilizes the starting material. It decreases the kinetic barrier by decreasing the difference in energy between starting material and transition state.

4. Principles of Catalysis • A catalyst opens a new pathway with a lower activation barrier for reaction to follow. • The Gibbs Energy of the reaction is unchanged. • There are no stable intermediates in the catalytic pathway.

8. Homogeneous Catalysis “A catalyst accelerates a chemical reaction without appearing in any of the products. An equilibrium is equilibrated faster, but the position of the equilibrium will not be changed” The world market for catalysts is estimated to be more than $ 2x109 and the total value of chemicals produced by means of catalysis exceeds $ 1500x109

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17. Heterogeneous Catalysis

18. Typical catalytic materials • The chemical nature of catalysts is as diverse as catalysis itself, although some generalizations can be made. • Proton acids are probably the most widely used catalysts, especially for the many reactions involving water, including hydrolysis and its reverse. • Multifunctional solids often are catalytically active, e.g. zeolites, alumina, higher-order oxides, graphitic carbon, nanoparticles, and facets of bulk materials. • c. Transition metals are often used to catalyze redox reactions. Many catalytic processes, especially those used in organic synthesis, require so called "late transition metals", which include palladium, platinum, gold, ruthenium, rhodium, and iridium. • Chemical species that improve catalytic activity, without themselves being active, are called co-catalysts or promoters.

19. Temperature and the Rate Constant • The rates of chemical reactions are sensitive to temperature: most reactions slow down at lower temperatures and speed up at higher temperatures. – This temperature dependence is contained in the rate constant, k. Rate = k [A] – Increasing the value of k increases the rate of the reaction. For many reactions, every increase in temperature by 10°C doubles the reaction rate. • The temperature dependence of k is given by the Arrhenius equation:

20. Collision Theory • Collision theory views the reaction rate as the result of particles colliding with a certain frequency and minimum energy. – Particles must collide in order to react, but most collisions do not result in a reaction, either because the particles do not hit each other hard enough, or they are turned the wrong way, etc. – As the number of colliding reactants increases, the chances of two reactants colliding also increases. Thus, increasing concentration increases the rate of the reaction. – Anything that increases the number of effective collisions increases the rate.

21. For a general reaction A + BC → AB + C as A and BC collide, their electron clouds repel each other. The energy needed to overcome this repulsion comes from the kinetic energy of the particles, and is converted to the potential energy of the A---B---C complex.

22. Factors that Influence Effective Collisions • Not every collision between reactant molecules leads to the formation of a product molecule. The number of effective collisions, which actually lead to the formation of a product molecule, depends on three factors: – the exponential factor, f — the fraction with enough energy to react (related to the activation energy). – the collision frequency, Z —the number of collisions per unit of time. – the orientation factor, p —the fraction of collisions with the correct orientation.

23. Activation Energy, Ea • The height of the barrier is the activation energy, Ea, and the configuration of the atoms at the maximum potential energy is the transition state or activated complex (++). – If the reactant particles collide with an energy less than Ea, they bounce apart. – If the collision energy is greater than Ea (and orientation is right), there is enough energy to overcome the repulsions, and they react. – In the transition state, the reactant bonds are in the process of breaking, and the product bonds are in the process of forming. – The higher Ea is, the slower the reaction will be Winger and Polanyi’s representation of Arrhenius model of activation barriers to reactions

24. The Frequency Factor • The exponential factor, f, is the fraction of collisions with enough energy to react: where R is the gas constant 8.314 J K-1 mol-1. • At higher temperatures, the distribution of collision energies broadens and shifts to higher energies, enlarging the fraction of collisions with energy greater than Ea. This makes f a larger number.

25. Reaction Rate and Temperature • The collision frequency, Z, is the number ofcollisions which occur in a given unit of time. • For a gas at room temperature and a pressure of 1 atmosphere, each molecule undergoes about 109 collisions per second, or 1 collision every 10-9 s. – If every collision resulted in a reaction, every gas phase reaction would be over in 10-9 s. Most reactions are obviously much slower than this. – For a reaction where Ea is 75 kJ/mol, at 298 K f= 7 x 10-14 only 7 collisions in 100 trillion are energetic enough to cause a reaction to occur! • The collision frequency is directly proportional to the concentration of the reactants.

26. Molecular Orientation • Not all collisions with energy greater than Ea lead to a reaction: the molecules have to be facing each other the right way when they hit each other. • The fraction of collisions having the right orientation is called the orientation factor, p.

27. The Arrhenius Equation • All of these factors can be combined into a single equation: • p and Z are often combined into a frequency factor, A (A = pZ); in this form, this equation is known as the Arrhenius equation (Svante Arrhenius, 1889): • Rearranging the Arrhenius equation, we can obtain the form of an equation of a line:

28. Rate-Determining Steps • Usually one step in a mechanism is much slower than the other steps, and acts as a “bottleneck” for the reaction; the rate of this step limits how fast the overall reaction can occur, and is known as the rate determining step. • The rate law for the rate-determining step represents the rate law for the overall reaction.

30. Overcoming unfavorable thermodynamic (water splitting)

31. H2 production from water (O2 release is the bottleneck) H2O  H2 + ½ O2DG +286 kJ/mol, 2.3 eV, T = 3000 °C HP and HT electrolysis Solar thermal (Almeria, Spain) Hyrosol-2 (100kW)

32. Solar thermal

33. Supporting catalysts Dispersion - nano In an eight-atom cluster, all of the atoms are on the surface. However, the dispersion, D, defined as the number of surface atoms divided by the total number of atoms in the cluster, declines rapidly with increasing cluster size.

34. The making of Ammonia N2 + 3 H2 2 NH3 Nothing happens in the system without a catalyst as T raised until 1000oC or higher. Above this temperature some H2 molecules are dissociated to H atoms. H2  2 H (atoms) For example at 1430oC with p(H2) = 150 atm., the partial pressure of H atoms is ca. 0.1 % only above 3000 oC where N2 molecules dissociates to N atoms and ammonia can be synthesized in reasonable quantities. N2  2 N (atoms) The role of the catalyst in ammonia synthesis is that of making the reaction go sufficiently fast (by facilitating the dissociation of molecular nitrogen) so that significant rates are obtained. Fritz Haber's successful synthesis of ammonia in 1909, capturing nitrogen from the air, brought him fame and wealth.  In 1911, he moved to Berlin to head the Kaiser Wilhelm Institute for Physical Chemistry and Electrochemistry.  In Berlin, he became friend with Albert Einstein.

35. Gas phase reaction + N N N N N + Catalytic reaction N N N N N N N N Energy profiles for the series of reaction steps to make ammonia from N2 and H2 by both homogeneous gas-phase and iron-catalyzed reactions. The role of the catalyst in decreasing the energy barrier to reaction can be seen (vales are in kJ mol-1)

36. Kinetic definition of catalysis Paul J. Crutzen Born: 3 December 1933, Amsterdam, the Netherlands. The Nobel Prize in Chemistry 1995

37. Hydrogenating organic compounds in the presence of finely disintegrated metals As you can also see in the figure The catalyst has not changed the thermodynamics, DH and therefore DG and Kp are unchanged, it only affected the transition state. Born: 5 November 1854, Carcassonne, France Died: 14 August 1941, Toulouse, France The Nobel Prize in Chemistry 1912

38. 50 nm x 50 nm Rh/Al2O3 Examples of Catalysts Au/TiO2 PtRu/CeO2

39. The active sites: acidity in zeolites There is one acid hydrogen for every tetrahedrally bonded aluminium. These active sites are distributed uniformly throughout the bulk and are bridging hydroxyl groups. These are the classic Bronsted acid sites, the intrinsic strength of which is a function both of the particular local environment and also the Si/Al ratio.

40. The active sites: bi-functional catalysts Example: the Pt/Al2O3 catalysts used in the hydroprocessing of petrochemicals, the metal serves to dissociate H2, while the acid support serves to catalyze the build-up of vital carbonium ion intermediates. H:H methyl cyclo-propane Pt 2-butene butane Al3+ O2-

41. The concept of “active site” is therefore very wide. Some examples of adsorbed surface complex are showed, you can observe how reactants interacts with the catalysts surface depending on the nature and distribution of the active sites. A. NH3 (Lewis base) coordinately linked to Al+3 ions (Lewis acid) on Al2O3 surface. B and C. Linear and bridge adsorption of CO on Pt. D and E. Dissociative adsorption on Pt of H2 or ethane. F. Dissociative adsorption of N2 on Fe. G. Heterolytic dissociative adsorption of H2 on the ZnO surface. H. Adsorbed complex with charge transference. I. Adsorption of isobutene on silica alumina where the acid surface proton (σ-OH) was transferred to the isobutene. J and K. Possibilities of ethylene adsorption on Pt. L. Adsorption of O2 on metal oxides with charge transference. M. Dissociative adsorption of O2. N. Heterolitic dissociative adsorption of propylene on ZnO.

42. CH4 Hydrogen Production (by “Steam Reforming” and “Water-gas Shift”) CO2 H2O (160 million tons/year) Heat H2 with CO and CO2 impurities 1 % of the World’s energy production CO clean-up (by “Methanation”) CH4 and H2O Pure H2 Ammonia synthesis (by “Haber-Bosch”) N2 NH3  Fertilizer  Food for 2-3 billion people