Kinetics Topic 6

1. KineticsTopic 6 http://www.wwnorton.com/college/chemistry/gilbert2/chemtours.asp#16

2. Rate of reactions • can be fast- explosions • can be slow- rusting • as a reaction takes place • concentration of reactants decreases (negative) over time • concentration of products increase (positive) over time • normally measure in Molarity per sec (M/s) • will continue until reaches equilibrium or one of the reactants is used up

3. Measuring rate of reactions • can measure: • mass or volume change for gaseous reactions • mass would go down as gas escapes • volume would increase at constant temp. and pressure • pressure would increase at constant temp. and volume • change in pH if acids and bases are involved • change in electrical conductivity • if produces ions in solution, conductivity will increase • using a spectrometer to detect color changes

5. Determining Rate of Reactionfrom reactions (Concentration, Volume, and Mass) • usually involves a graph of properties over time • usually a curve, and the reaction rate is determined from the slope of the line at a time (also known as a tangent) • reaction rates tend to slow with time as reactants are converted to products

6. reaction slows down with time because the Concentration of the reactants decreases “rise over run” .040M/200s = .0002M/s .025M/400s = .000063M/s Example C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)

8. the change in concentration of a reactant or product per unit of time • [ ] refer to the concentration of the reactants

9. 2NO2(g)  2NO(g) + O2(g) product [NO2] t product reactant

10. gas is being released in the reaction so Massdecreases over time reaction is creating gas so Volume increases over time

11. Kinetic theory (6.2.1) • energy of particles is proportional to the temperature (Kelvin or Celsius) • all particles have same energy if the same temperature • lighter particles would have greater speed than larger particles given the same energy

12. Activation Energy Ea. • a minimum amount of energy required for reaction to occur • bonds need to be broken first • the molecules must posses sufficient energy to get over the activation energy barrier.

13. Collision theory (Topic 6.2) • in order for particles to react • particles must collide • must collide in the correct orientation/angle • must collide with enough kinetic energy to overcome the activation energy (Ea) • if the previous conditions are “enough”, particles can overcome the activation energy and reaction will occur • meaning the bonds holding the reactants together will be broken

14. http://phet.colorado.edu/en/simulation/reactions-and-rates

15. Factors That Affect Reaction Rates • any factor that increases the frequency of collisions or increases the energy with which particles collide will make the reaction go faster: • temperature • pressure • surface area • concentration • catalysts

16. 1. Temperature • increase temp • increases number of collision per unit time • reaction rate approximately doubles for each 10oC or K rise in temperature • increases energy of the collisions

17. 2. Pressure • only for gasses • reducing volume while keeping temp constant • forcing them together will increase number of collisions

18. 3. Surface area • smaller particles have more surface area • only the particles on the surface can come in contact with a reactant • more collisions per unit time

22. 4. Concentration • increasing concentration will increase more collisions per unit time

24. 5. Catalysts • lowers the activation energy (Ea) for the reaction • provide an alternate reaction (rxn) pathway • increase the rate of a reaction • are not used up or chemically changed in the reaction

26. Maxwell–Boltzmanenergy distribution curve • another way to look at what particles can react • area under the curve shows the number of gas particles • not all gas particles have the same energy • only some gas particles (blue area) have enough energy to react

27. The affect of temperature

28. the area under the curve remains the same because the number of particles doesn't change • higher temps. shifts the curve to the right (therefore, the peek must be lower) resulting in an increase in collision frequency and thus more successful collisions

29. The affect of a catalyst

30. Never move the new activation energy to the left of the peak. Catalysts don’t help out that much!