kinetics

1. kinetics

2. Collision Theory • Reactants must collide with proper orientation and sufficient energy

3. Transition State Theory • Explains what happens once colliding particles react • Transition state is the in-between state when reactants are being converted to products • Kinetic energy is converted to potential energy (think about bouncing a basketball) Kinetic Energy Potential Energy

4. Potential Energy Diagrams Negative

5. Exo- Vs Endothermic • Exothermic reactions- give off heat (products more stable, less potential energy) • Endothermic reactions require heat (products less stable than reactants, higher PE) Endothermic Reaction Exothermic Reaction

6. Draw a PE diagram. Include: axes labels, the transition state, activated complex, Ea (forward and reverse), and ΔE. Try It • CO reacts with NO2 to form CO2 and NO. The activation energy of the forward reaction is 134 kJ and the ΔE is -226 kJ.

7. Reaction Mechanisms • Chemical reactions typically occur as a series of steps. This series of steps that make up the overall reaction is called the reaction mechanism. • Elementary reactions are a single step in the overall reaction mechanism. • Singular molecular event, such as a simple collision of atoms, molecules or ions. • Cannot be broken down into further simpler steps.

8. Elementary Reactions • For example, the reaction 2NO(g) + O2(g)  2NO2(g) involves a two-step reaction mechanism: • Step 1: NO(g) + O2(g)  NO3(g) • Step 2: NO3(g) + NO(g)  2NO2(g) • Each step is an elementary reaction, both steps together give the overall reaction mechanism. • Notice NO3(g) • Not a product or reactant of overall reaction • Produced then consumed = reaction intermediate

9. Molecularity • Describes the number of reactant particles in an elementary step • Unimolecular = one reactant • (CH3)3CBr(aq)­  (CH3)3C+ + Br- • Bimolecular = 2 reactants come together. Ex: • Step 1: NO(g) + O2(g)  NO3(g) • Step 2: NO3(g) + NO(g)  2NO2(g) • Termolecular = 3 reactants • Extremely rare!!! Why?

10. Rate Determining Step • One step in a reaction mechanism is always much slower than the others • Since it is so much slower, it determines the rate of the overall reaction. • Hence, this slow step is called the rate determining step • Consider the process of making toast: slow fast

11. RDS and PE Diagrams • Activation energy for rds is always higher • 2 STEP REACTION MEANS 2 TRANSITION STATES AND 1 INTERMEDIATE

12. Catalysts • Increases the rate of a reaction • Homogeneous catalyst: Same phase as reactants • Heterogeneous catalyst: Different phase as reactants Pd

13. Catalysts • Not consumed! • There in beginning and end of reaction • Works by lowering the activation energy • Therefore, greater number of collisions have sufficient energy to react • Provide alternative reaction mechanism

14. Kinetics Taboo • Rate Constant • Elementary Reaction • Mechanism • Catalyst • Rate Law • Reaction Order • Molecularity • Activation Energy • Reverse Reaction • Potential Energy • Transition State • Activated Complex • Rate Determining Step • Biomolecular • Termolecular • Reaction Rate • Concentration • Surface Area • Catalyst • Reactivity • Temperature • Collision Theory • Orientation • Rate Expression • Exothermic • Endothermic • Homogeneous • Heterogeneous • Kinetic Energy • Unimolecular