Kinetics

1. Kinetics Topic 6

2. 6.1 Rates of reactions • Different reactions react at different rates. • A rate is a measure of speed of any change that occurs within an interval of time. • Rate is expressed per unit time or in SI units per sec ( s-1 ) • For example, the formation of rust is slow compared to a fire burning.

3. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions.(p.116-119) • Change in volume of a gas produced • Change in mass • Change in transmission of light • Change in concentration using titration • Change in concentration using conductivity • Non-continuous methods of detecting change during a reaction “clock reactions”

4. 6.1.3 Analyze data from rate experiments • Design an experiment showing the rate of reaction of vinegar and baking soda. • Using the equipment provided, decide which type of procedure you should use from slide #3.

5. 6.2 Collision Theory • Particles can react to form products when they collide with one another, provided that the colliding particles have enough kinetic energy. • Particles lacking enough kinetic energy to react will simply bounce apart unchanged (they also need to have the correct “geometry”) • The minimum energy colliding particles must have to react is called activation energy.

6. Activation Energy Molecules need a minimum amount of energy to react. Visualized as an energy barrier - activation energy, Ea.

7. Activation Energy and Temperature Reactions are faster at higher T because a larger fraction of reactant molecules have enough energy to convert to product molecules. In general, differences in activation energy cause reactions to vary from fast to slow.

8. Factors affecting reaction rates • Particle size = The smaller the particle size, the larger the surface area. An increase in the surface area increases the amount of the reactant exposed for reaction. For example, coal dust in a mine is much more explosive than chunks of coal.

9. Temperature = Usually, raising the temperature speeds up reactions, and lowering the temperature slows down reactions. For example, charcoal does not burn in air, but will continue burning once it has been lit by a flame.

10. Concentration = Putting more particles into a fixed volume increases the frequency of collisions which leads to a higher reaction rate. *Pressure is another factor with gases only. For example, fire burns better in pure oxygen as compared to burning in air.

11. Catalysts = A substance that increases the rate of reaction without being used up in a reaction. It lowers the activation energy for a reaction. For example, enzymes in our bodies break down proteins during digestion. Without these enzymes, this process would take years!

12. Uncatalyzed reaction Catalyzed reaction CATALYSTS MnO2 catalyzes decomposition of H2O2 2 H2O2 2 H2O + O2

13. Maxwell-Boltzmann curve • Shows the # of particles that have a particular value of kinetic energy plotted against values for kinetic energy. • Imagine a busy interstate highway with the cars in the different lanes going different speeds. (example on board)

14. Half-Life HALF-LIFE is the time it takes for 1/2 a sample to disappear

15. Half-Life example: • Decomposition of H2O2.

16. Half-Life • Reaction after 1 half-life. • 1/2 of the reactant has been consumed and 1/2 remains.

17. Half-Life • After 2 half-lives 1/4 of the reactant remains.

18. Half-Life • A 3 half-lives 1/8 of the reactant remains.

19. Half-Life • After 4 half-lives 1/16 of the reactant remains.