Download Chemical Kinetics & Intermolecular Forces: Reaction Rates, Catalysis & Forces in Chemistry 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 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 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 Chapter #16 – Liquids and Solids 16.1) Intermolecular Forces 16.2) The Liquid State 16.3) An Introduction to Structures and Types of Solids 16.4) Structure and Bonding of Metals 16.5) Carbon and Silicon: Network Atomic Solids 16.6) Molecular Solids 16.7) Ionic Solids 16.8) Structures of Actual Ionic Solids 16.9) Lattice Defects 16.10) Vapor Pressure and Changes of State 16.11) Phase Diagrams Macroscopic Properties of Gasses, Liquids, and Solids Densities of Water in each state Solid (0°C, 1atm) 0.9168g/cm3 Liquid (25°C, 1atm) 0.9971g/cm3 Gas (400°C, 1atm) 0.000326g/cm3 Gas Liquid Solid Condensation H0vap = - 40.7 kJ/mol Vaporization H0vap= 40.7 kJ/mol Freezing H0fus = -6.02 kJ/mol Melting H0fus = 6.02 kJ/mol - H0subH 0 sub Phase Changes for Water, and their Enthalpies - - Sublim ation D eposition Where do these enthalpies come from? Intermolecular Forces Intramolecular forces: chemical bonds holding atoms together in a molecule ionic bonds, e.g. Na+Cl- ~400 - 4000 kJ/mol covalent bonds, e.g. H2, CH4, etc. ~150 - 1200 kJ/mol (the O-H bond energy in H2O is 934 kJ/mol) metallic bonds, e.g. Au, Fe, etc. ~75 - 1000 kJ/mol Intermolecular forces: forces between molecules Much weaker: ~0.05 - 40 kJ/mol Types of Intermolecular Forces Dipole-Dipole forces • based on charge-charge interactions • occur in molecules with polar bonds and dipole moments • Typically 5-25 kJ/mol Figure 16.2 A special type of dipole-dipole force. Occurs when a Hydrogen atom is bound to the highly electronegative atoms O, N, or F, which have lone pairs (let’s call these atoms A or B). These H-A bonds are very polar. The partially positive H atom from one unit is attracted to the partially negative A atom from another unit (or to another partially negative atom, B). This is the hydrogen bond. To occur, the atom sequence must be -B:…H-A- where both B and A are O, N, or F. Even though the strength of one H bond may be small, the combined strength of many H bonds may be large. Hydrogen Bonding
