1. Topic 6 Kinetics Rates of reaction Collision theory

2. 6.1 Reaction rate Definition Experimental procedures Analyse data Rate = (conc. of A at time t 2 - conc. of A at time t 1 ) / t 2 -t 1 =  [A]/  t Unit: M/s = mol*dm -3 *s -1

3. The rate of which products are formed = the rate of which reactants are consumed. But the value can vary according the number of moles involved: MnO 4 - + 8 H + + 5 Fe 2+  Mn 2+ + 4 H 2 O + 5 Fe 3+ The formation of Fe 3+ is five time faster than the consumption of MnO 4 -.

4. Measurement of reaction rate Any property that differ between a product or a reactant can be used. e.g. absorption of light, electrical conductivity, gas production, removal of samples for chemical studies and so on What property can be measured in this reaction: Mg (s) + 2 HCl (aq) MgCl 2 (aq) + H 2 (g)

5. Examples of measurements Mass or volume change for gaseous reactions Change in pH for acid-base reactions Change in conductivity for reactions with electrolytes Use of a spectrophotometer for reactions involving colour change

6. A graph is normally drawn : property against time. The rate of the reaction is proportional to the gradient of the curve Change in reaction conditions => change in gradient. Often no need to convert the “property” to M/s

7. Changes in mass against time Marble CaCO 3 + acid = carbondioxide + CaCl 2

8. Changes in mass and volume against time Marble CaCO 3 + acid = carbondioxide + CaCl 2

9. Changes in concentration against time

10. Activation energy (E a ) Kinetic theory Collision theory Factors affecting kinetics Maxwell-Boltmann Catalysts 6.2 Collision theory

11. Activation energy (E a ) When two particles going to react they must have energy enough to overcome the repulsive force in their negatively charged electron clouds. They must also have enough energy to break the required bonds

12. The amount of activation energy needed for a reaction differ widely from reaction to reaction. e.g. H + + OH -  H 2 O reacts immediately at room temperature Wood + O 2  CO 2 + H 2 O doesn’t react at room temperature

13. Collision theory In order to react, the two particles involved must: 1.Collide with each other 2.The collision must be energetic enough to overcome the activation energy. 3.The collision must bring the reactive parts of the molecule into contact in the correct way. (Hard if large organic molecules )

14. Maxwell-Boltzmann energy distribution curve The average kinetic energy of a collection of molecules is proportional to their absolute temperature. Lots of collisions between them give a wide range of different molecular speeds.

15. The maximum on the curve: represent the kinetic energy possessed by the largest fraction of the molecules. Most probably kinetic energy. The average kinetic energy occurs at a higher value. For a reaction the kinetic energy must overcome the E a. At a higher temperature a bigger fraction of molecules have kin. energy > E a => faster reaction.

16. Factors that affect the reaction rate Factors mainly affecting collision rate Factors mainly affecting E a Concentration/pressureTemperature Surface areaCatalyst

17. If a catalyst is used the reaction take an other path with a lower E a => the E a shift to left in the graph => bigger fraction of molecules have kin. energy > E a => faster reaction

18. Concentration: Zinc reacts faster with concentrated than diluted acid Surface area: Zinc reacts faster with zinc powder than a big lump. Temperature: Number of particles with E > E a is higher at a higher temperature (recall Maxwell- Boltzmann energy distribution curve) Also a small effect on the collision rate. Catalyst: A compound that increases the reaction rate without being consumed. Catalyst reduce the E a by taking an other reaction path.

19. Exercises Page 116