Rates of Reaction

1. Rates of Reaction Unit 8

2. Reaction Times • Time of reaction varies • Examples: • HCl + NaOH  NaCl + H2O Instantaneous • 2 CO + O2  2CO2Thousands of Years • It is therefore important to not only know if a reaction will occur, but also what the reaction rate is

3. Collision Theory • Collision Theory: • Atoms, ions, and molecules can react to form products when they collide, provided that particles have enough kinetic energy. • Concentration is very important • If the individual particles that collide do not have enough KE, they will bounce apart when they collide.

4. Collision Theory • Activation energy – the minimum amount of energy particles need in order to react • Activated complex – the temporary arrangement of atoms at peak of activation-energy barrier • Called the transition state – can re-form reactants or form products

5. Collision Theory

6. Kinetics vs. Thermodynamics • Kinetics – How FAST a reaction will occur • Thermodynamics – WHETHER a reaction will occur in the first place

7. Temperature Concentration Particle Size Catalysts Pressure Factors Affecting Reaction Rates

8. Temperature • Temperature = Kinetic Energy • By increasing the temperature of a reaction, the kinetic energy is increased • More particles have enough energy to collide at higher temperature (above activation-energy barrier)

9. Concentration • Number of particles in a given area • Related to pressure (gas) • Increase collision frequency

10. Particle Size • Smaller particles have larger surface area • Increased surface area means more reactant exposed for reaction • Ways to increase surface area: • Dissolve solid reactants to increase surface area • Grind into powder

11. Catalyst • Catalyst – substance that increases rate of reaction without being used up itself during the reaction • Lowers activation energy

12. Catalyst