CHEM1612 - Pharmacy Week 12: Kinetics – Catalysis Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au

Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille,      Chemistry, John Wiley & Sons Australia, Ltd. 2008      ISBN: 9 78047081 0866

Energy Landscape in Chemical Reactions A + B C + D Endothermic reaction A + B C + D Exothermic reaction Ea (forw) (rev) Activated state Ea (forw) Ea (rev) A + B Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. C + D Forward reaction is faster than reverse Reverse reaction is faster than forward Larger Ea smaller k lower rate

By the way: why the product? elementary reaction A + B→ C rate = k [A][B] Why does rate depend on the product of reactant concentrations? Rate proportional to the number of collisions of A and B No. collisions = product of the number or particles present 2×2 =4 3×2 =6 3×3=9

Transition State If reactants come together with enough energy and the right orientation, they combine to form a transition state (or activated complex). This species is half-way between the reactants and the products but is not neither. Transition states are very unstable (cannot be isolated). Blackman Figure 14.10

Transition State Nature of the transition state in the reaction between CH3Br and OH-. CH3Br + OH- CH3OH + Br - Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. transition state or activated complex

Reaction Energy Diagram Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Transition state in elementary steps k = A e – Ea / R T Ea1 > Ea2, therefore Ea1 is the slow step and Ea2 is the fast step. Two transition states. Blackman Figure 14.11

Transition state in elementary steps Step 1 NO2 + F2 →NO2F + F Step 2 NO2 + F → NO2F Overall 2 NO2 + F2 → 2NO2F Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

Catalysis A catalyst increases the rate of a chemical reaction. A catalyst is not consumed or changed in the overall process. A catalyst provides an alternative reaction pathway of lower activation energy: more molecules have the minimum energy required for successful reaction and the reaction proceeds at a faster rate. The catalyst may react to form an intermediate, but it must be regenerated in a subsequent step of the reaction Catalyst speeds up a reaction by providing a set of elementary steps with more favourable kinetics (faster) than those existing in its absence Diagrams show the energetics of the reaction with and without a catalyst. Note that the catalyst lowers the energy barrier to the reaction without changing the energies of the reactants and products. Reducing Ea increases reaction rate k A + B C + D

Catalysis A catalyst speeds both the forward and reverse reaction, so does NOT affect the position of the equilibrium. Does not change the equilibrium constant Keq = k1/k-1; even though k1 and k-1 may be much larger for the catalyzed reaction. Does not change the G0 for the reaction. A catalyst can be homogeneous (one phase with reactants and products) or heterogeneous (more than one). Many catalysed reactions are zero-order. A small quantity of catalyst affects the reaction rate for a large amount of reactant. Once the surface is entirely covered with reactants, increasing further the reactant concentration does not affect the rate.

Demo: Catalytic Decomposition of Hydrogen Peroxide Manganese dioxide MnO2 is used to catalyse the decomposition of H2O2. Prepare a slurry of MnO2 and NaOH, and add H2O2. The rapid decomposition of H2O2 occurs with production of a grey foam: 2 H2O2 → 2 H2O + O2

Catalysis Let’s look at the decomposition of H2O2 again. 2 H2O2(aq) 2 H2O(l) + O2(g) Catalyst Ea (kJ mol-1) Rel. rate of reaction None 75.3 1 I– 56.5 2.0 x 103 Pt 49.0 4.1 x 104 Catalase 8.0 6.3 x 1011 So far the rates of reaction have been considered as well as the effect of temperature on the reaction. It is clear that the magnitude of Ea determines the efficacy of a reaction. Ea corresponds to the formation of an intermediate between the reactants and products. If there was a way of decreasing Ea then reactions could be made to occur more quickly Demonstration: Shows that I- increases the rate of reaction at constant temperature. Iodide ions facilitates the decomposition of H2O2. The experimentally determined rate law shows a linear dependence on the iodide concentration I- facilitates reaction through decrease in Ea Finer examination shows that two separate reactions occur. Note that when both equations added together end up with initial equation and that I- not consumed Iodide acts as a catalyst and reduces Ea for the conversion. It can be envisaged that the conversion of H2O2 to H2O + O2 via the formation of IO- provides an intermediate which is easier to form (energetically) than for the direct decomposition of H2O2 Bombardier beetle (Brachinus fumans)

Homogeneous Catalysis Let’s closely examine the reaction of H2O2 with I–: Rate law: rate = k[H2O2][I–]. Reaction occurs in two steps: Step 1: H2O2 + I– H2O + IO– Step 2: H2O2 + IO– H2O + O2(g) + I– Note that I– regenerated during reaction and it does not appear in overall reaction. I– acts as a homogeneous catalyst for H2O2 decomposition. So far the rates of reaction have been considered as well as the effect of temperature on the reaction. It is clear that the magnitude of Ea determines the efficacy of a reaction. Ea corresponds to the formation of an intermediate between the reactants and products. If there was a way of decreasing Ea then reactions could be made to occur more quickly Demonstration: Shows that I- increases the rate of reaction at constant temperature. Iodide ions facilitates the decomposition of H2O2. The experimentally determined rate law shows a linear dependence on the iodide concentration I- facilitates reaction through decrease in Ea Finer examination shows that two separate reactions occur. Note that when both equations added together end up with initial equation and that I- not consumed Iodide acts as a catalyst and reduces Ea for the conversion. It can be envisaged that the conversion of H2O2 to H2O + O2 via the formation of IO- provides an intermediate which is easier to form (energetically) than for the direct decomposition of H2O2

Heterogeneous Catalysis Heterogeneous catalysis: most important industrially. Catalytic converters: First converter (A): Rh Catalyses: 2 NO(g) N2(g) + O2(g) Second converter (B): Pt/Pd Catalyses: 2 CO(g) + O2(g) 2 CO2(g) A B Catalytic converter There are three types of catalysis: homogeneous, heterogeneous and enzyme catalysis In homogeneous catalysis: reactants, products and catalyst all dispersed in single phase (eg H2O2 decomposition using I-) Heterogeneous catalysis: the reactants and catalysts are in different phases (usually the catalyst is solid and reactants liquid or gas) The key to the success is that intermediates species formation promoted by the surface, usually through some type of intermolecular bonding with the surface Catalytic converters are a good example of heterogeneous catalysis in action. Fuel which is not efficiently burnt contains residual CO - which is toxic. First converter contains a mixture of Pt, Pd and Rh and operates at high temperature (exhaust gas temp) to catalyse the conversion of CO (and other unburnt hydrocarbons) to CO2 which is less harmful The reaction would not occur without the transition metal to reduce the activation energy to an acceptable value (cf. kT). This occurs due to a stabilization of the transition state due to adsorption of CO etc. at the catalyst surface The second converter (chamber) contains a different transition metal catalyst that specifically adsorbs NO and catalyses its transformation to nitrogen and oxygen. The principle of specifically binding the NO (and not CO) will be operative in this chamber

The metal-catalyzed formation of ammonia Fe N2(g) + 3H2(g) → 2NH3(g) Both substrates must bind to a free active site on the Fe surface before the reaction can proceed. Increasing the concentration of either gas cannot increase the rate of reaction (i.e. rate independent of concentration). The industrial hydrogenation of C2H4 uses a Pt metal catalyst.

Enzymes Catalysts of biological reactions Complex 3D structure Huge molar mass Active site attracts substrates through intermolecular forces Haber process (500 atm and 450 °C; Nitrogenase (1 atm and 25°C) Enzyme-substrate complex of elastase and small peptide

Enzymatic Catalysis All enzymes are proteins, but not all proteins are enzymes. Enzymes must possess catalytic activity. The part of the enzyme tertiary structure that is responsible for the catalytic activity is known as the “active site”. Active site Structure of the enzyme Hexokinase from X-ray data. Each enzyme catalyses a single chemical reaction on a specific substrate molecule with high selectivity. The active site: makes up only 10 - 20% of the total volume of the enzyme. is where all the catalytic chemistry occurs. is usually a cleft or cavity containing an array of amino acid side-chains. binds the substrate and carries out the enzymatic reaction.

Enzymatic Catalysis Efficient: rate enhancements of 108 to 1020 possible Specific (one enzyme per reaction) Low tolerance to temperature and pH changes Lock-and-key model (E. Fischer, 1894) Induced fit model (D. Koshland, 1958) Michaelis-Menten mechanism: enzyme-substrate complex ES. Product + enzyme OOO… Bring in nifty examples…

Enzymatic Catalysis Hexokinase alters its conformation to fit around the substrate molecule (D-glucose). Enzyme and substrate adapt to accommodate one another. “Enzymes are molecules that are complementary in structure to the transition states of the reactions they catalyze”. Linus Pauling (1948)

Enzymatic Catalysis Enzymes can distinguish between enantiomers: Only one of the enantiomers can be used as a substrate for this enzyme.

Enzymatic Catalysis Free energy E a(uncat) – a(cat) = D a ….. uncatalyzed reaction catalyzed reaction Reaction co-ordinate If DEa = 10 kJ mol-1 55-fold rate acceleration (at 25°C). If DEa = 20 kJ mol-1 3000-fold rate acceleration (at 25°C). If DEa = 40 kJ mol-1 107-fold rate acceleration (at 25°C). By reducing the activation energy for a chemical reaction, the catalyst (enzyme) can enhance the rate of reaction dramatically. An alternative, lower energy pathway for the reaction becomes available as a result of the catalyst. In terms of enzymatic catalysis, the transition state stabilization energy (DEa) is given by the difference in free energy between the (altered) substrate in the transition state and that of the enzyme-substrate complex in the transition state.

Enzymatic Catalysis k = A e – Ea / R T The Arrhenius equation indicates that in order to increase the rate of a reaction: The temperature must be increased, Ea must be decreased, and/or The reactants must be positioned so as to maximise the reaction efficiency. Increasing the temperature is not an option for most biological reactions, so the remaining options are exploited by Nature.

Summary CONCEPTS Elementary reactions, reaction mechanisms Dependence of reaction rate on temperature and orientation Arrhenius equation and its implications Activation energy and transition states Catalysis CALCULATIONS Express reaction rate in terms of reactant/product concentrations Derive rate law of a reaction from experimental data on reactant consumption/product formation Derive rate law of a reaction, knowing its elementary steps Calculate k, A, T or Ea for a reaction using Arrhenius equation

The Ozone System In the atmosphere: O2 + hn (<200 nm)  O + O k1 O + O2  O3 assume very fast O3 + hn (210-300 nm)  O2 + O k2 O3 + O  2 O2 k3 Kinetics very complicated since UV intensity will vary so much in time and place. At equilibrium, rate of ozone creation and destruction will be the same: steady state approximation.

The Ozone Hole Chlorofluorocarbons provide an additional pathway for ozone decomposition. CF2Cl2 + hn  CF2Cl• + Cl• Cl• + O3  ClO• + O2 removing ozone ClO• + O•  Cl• + O2 Cl is regenerated during the reaction Cl will stick around until eventually reacts to HCl and is precipitated out. Other reactions are possible – happen on stratospheric dust and ice particles