Collision Theory and Reaction Rates

1. Collision Theory and Reaction Rates

2. The Collision Theory is a theory which states how reacting particles (like atoms and molecules) must interact to start a chemical reaction.

3. The Collision Theory States 3 Things: 1st: Reactant particles must collide in order for the reaction to occur.

4. The Collision Theory Continued . . . 2nd: Reactant particles must collide the “right way” with proper orientation.

5. 3rd: Reactant particles must collide with enough energy to “stick together.” • Ea is called activation energy. • Activation energy is the minimum amount of energy required for a reaction to occur.

6. High Activation Energy (Ea) • LOTS of energy needed for the reaction to happen. • It takes a long time to get this energy, so there is a slow reaction rate.

7. Low Activation Energy (Ea) • LESS energy needed for the reaction to occur. • It takes less time to get needed energy so there is a fast reaction rate.

8. If the reactant particles meet ALL THREE requirements then an activated complex forms when the reactant particles “stick together”.

9. An Activated Complex Represents: • a temporary, unstable structure • the transition between the breaking of old bonds and the forming of new bonds

10. Endo- or Exothermic?

11. Label reactants, products, Ea, activated complex, E Exothermic

12. Endo- or Exothermic?

13. Label reactants, products, Ea, activated complex, E Endothermic

14. What is Ea? • Activation energy – the minimum amount of energy required to start a reaction. • Why does a high Ea mean there will be a slow reaction? • It takes a while for particles to get enough energy to overcome the Ea barrier. • Why does a low Ea mean there will be a fast reaction? • It won’t take as long for particles to get enough energy to overcome the Ea barrier.

15. Graph 2 has smaller Ea barrier. • Which of the following reactions would occur fastest: 1 or 2 and why?

16. Catalyst • Catalysts are substances added to a chemical reaction that speed up the rate of reaction to make it occur faster. • How does it do this? • It lowers activation energy! Draw the new Ea on the diagram below WITH a catalyst.

17. What is “proper orientation?” • The direction/orientation in which reactants must collide. • Example: NO3 + CO → NO2 + CO2 In this chemical reaction, when NO3 and CO react together, they must collide with sufficient energy and with the nitrogen atom in NO3 colliding with the carbon atom in CO. In other words, the nitrogen and carbon atoms must “bump” together!

18. What would happen if . . . • particles collide with sufficient energy but do not have proper orientation?

19. What would happen if . . . • particles collide with proper orientation but not enough energy?

20. What would happen if . . . • a reaction were to occur?

21. Which of the parts of the collision theory can be manipulated by us to cause a reaction to occur faster (or slower)? • Collision frequency • Collision energy

22. How can collision frequency and collision energy be manipulated?

23. Reaction Rate A reaction rate is how fast or how slow a reaction occurs over time.

24. Certain factors can be changed to make the reaction occur faster or slower. *Concentration *Surface Area *Temperature *Catalyst or inhibitor

25. Concentration is the amount of reactant(s) used measured in terms of molarity. • Increasing the concentration of a substance increases the number of particles involved in the reaction. • This makes it easier for the reacting substances to “find” each other which leads to an increase in the frequency of collisions. • More frequent collisions means a higher probability of collisions with proper orientation!

26. Example: Will a strip of Mg have a faster rxn rate with 10. M HCl or 0.50 M HCl? Explain why – watch terminology! Answer: 10. M HCl will consume Mg faster because it is more concentrated. Because there are more moles of acid per liter, there will be more frequent collisions b/w the acid and Mg and more frequent collisions w/proper orientation as a result.

27. Surface area is the amount of the surface the reactant has exposed for collisions. • By pulverizing or crushing a substance, you increase the amount of the substance’s surface that is exposed. • This makes more surface area of the particles available for the reaction. • More particles available for the reaction means more frequent collisions and more frequent collisions with proper orientation.

28. big chunk of Na crushed

29. Example Which would work faster: Aspirin in powder form or aspirin in tablet form? Explain why. Answer: The aspirin would work faster if taken in powder form because the fact that it has been crushed increases surface area. This will make more frequent collisions and more frequent collisions with proper orientation – which will make the aspirin work faster.

30. Temperature is the measure of kinetic energy of particles in a substance. • An increase in temperature increases the kinetic energy (or movement) of particles. • This increases the frequency of collisions because particles are bouncing around more. • This will increase the frequency of the collisions with proper orientation. • Because the particles are moving faster, they are more likely to achieve enough energy to overcome the Ea barrier.

31. Example: Will a strip of Mg react faster in 1 M HCl at 75C or in 1 M HCl at 23C? Explain. Answer: The Mg reacts faster in the acid at 75C because the Mg will absorb more kinetic energy in the acid at the higher temperature causing more frequent collisions and more frequent collisions with proper orientation and more collisions with sufficient Ea.

32. Catalyst Energy Diagram

33. A catalyst is a substance that speeds up a reaction. • It does this by lowering the activation energy. • This makes the Ea barrier easier to overcome so that that reaction can take place more quickly. • Catalysts form an activated complex of lower energy than a reaction without a catalyst.

34. More Catalyst Info . . . • Catalysts are NOT used up or consumed in a chemical reaction so they can be used over and over again like recycling– very cost efficient for pharmaceutical companies! • Catalysts do NOT make more of a product, they only are used to speed up a chemical reaction.

35. inhibitors • are catalysts that are used to SLOW DOWN reactions • Draw an energy diagram showing how an inhibitor works to reduce a reaction rate.

36. Inhibitors RAISE the activation energy, which makes the reaction SLOW

37. Why would a reaction want to be slowed down? • Food preservatives added to prevent/slow down spoilage. • Weed killers slow down the growth of weeds. • Medications used to treat HIV and cancer slow the spread of each.

38. Enzymes • An enzyme is a biological catalyst found in nature – which means enzymes speed up chemical reactions by lowering activation energy. • Example: Amylase is an enzyme in saliva that breaks down starches into sugars. • Lipase is an enzyme that breaks down fats.

39. You Survived!