Lecture Slides on Chemical Kinetics - General Chemistry | CHEM 162, Study notes of Chemistry

Download Lecture Slides on Chemical Kinetics - General Chemistry | CHEM 162 and more Study notes Chemistry in PDF only on Docsity! Chapter #15 – Chemical Kinetics 15.1) Reaction Rates 15.2) Rate Laws: Introduction 15.3) Determining the Form of the Rate Law 15.4) Integrated Rate Law 15.5) Rate Laws: Summary 15.6) Reaction Mechanisms 15.7) The Steady-State Approximation 15.8) A Model for Chemical Kinetics 15.9) Catalysis A Model for Chemical Kinetics (Zumdahl Section 15.8) • Concentrations of reactants affect rates • So does temperature • So do catalysts Obviously, a rate constant is not really a universal constant, but is constant only for a fixed set of experimental conditions. How do we explain the dependence of reaction rates on temperatures and catalysts? i.e. Where do Rate Constants come from? Collision Theory of Reactions Consider a gas phase reaction: 2A(g) products • This reaction requires the collision of two reactants. • We can calculate the collision frequency from the ideal gas law. • We can also measure the reaction rate experimentally. • We typically find that the reaction rate is several orders of magnitude slower than that predicted from the collisional frequency alone. Thus, we can conclude that not every collision of reactants results in a successful reaction. What factors might keep the two reactants from reacting once they have collided? One obvious candidate is molecular orientation: only the correct orientation of reactants will lead to successful product formation This could account for perhaps 1 order of magnitude, but not for the several orders of magnitude reduction in experimentally observed rates. The answer must be more complex. Figure15.11 Finding the activation energy from 2 data points Instead of a plot of of ln k vs. 1/T and getting Ea from the slope, you can calculate Ea from: ln k1 = ln A - Ea/RT1 ln k2 = ln A - Ea/RT2 Taking the difference of these two equations: ln (k2/k1) = (Ea/R)[(1/T1 - 1/T2)] T1 = 500K k1 = 9.51x10-9M-1s-1 T2 = 600K k2 = 1.10x10-5M-1s-1 ln (k2/k1) = (Ea/R)[(1/T1 - 1/T2)] ln (1.1x10-5/9.51x10-9) = (Ea/8.3145J/molK)*[(1/500K - 1/600K)] rearrange and solve: Ea = 176kJ/mol Example: 2HI(g) H2(g) + I2(g) Rate = k[HI]2 Chapter #15 – Chemical Kinetics 15.1) Reaction Rates 15.2) Rate Laws: Introduction 15.3) Determining the Form of the Rate Law 15.4) Integrated Rate Law 15.5) Rate Laws: Summary 15.6) Reaction Mechanisms 15.7) The Steady-State Approximation 15.8) A Model for Chemical Kinetics 15.9) Catalysis Catalysis It is not always practical or convenient to increase reaction rates by increasing the temperature. A Catalyst is a substance that speeds up a reaction without being consumed during by the reaction Catalysis - The use of catalysts to speed up reactions without changing the temperature Homogeneous catalysts- catalysts that are in the same phase (e.g. solution or gas) as the reacting molecules Heterogeneous catalysts- catalysts that are in a different phase from the reacting molecules Catalysts • Catalysts are used in a huge variety of ways because they can enhance reaction rates by many orders of magnitudes! • In general, they work by lowering the activation barrier to a reaction. Note that the energies of the reactants and products do not change, only the energy barrier changes. The transition state is stabilized by the catalyst! Example of a Homogeneous Catalyst Ethylene Polyethylene (polymerization) Catalyst: TiCl4/Al(C2H5)3 This is one member of a large class of polymerization catalysts developed by Ziegler and Natta (Nobel Prize in Chemistry, 1965) A lot of research is currently directed toward developing polymerization catalysts that make new polymers with attractive properties (e.g. teflon, kevlar, conductive polymers) (we'll talk more about polymers in chapter 22)