18. a.-393.51 kJ/mol; exothermic, c. +176 kJ/mol; endothermic

19. b. H2(g) + 1/2 O2 (g) --> H2O (l) + 285.83 kJ; exothermic

24. a. 115 kJ/mol; not spontaneous, b. -135 kj/mol; spontaneous

26. a. E forward = +80 kJ/mol, E reverse = -80 kJ/mol, Ea = 100 kJ/mol, Ea' = 20 kJ/molRates of chemical reactions.

5.06 Analyze the factors that affect the rates of chemical reactions.
∑ The nature of the reactants.
∑ Temperature.
∑ Concentration.
∑ Surface area.
∑ Catalyst.
Students should be able to:
∑ Explain collision theory – molecules must collide in order to react, and they must collide in the correct or appropriate orientation and with sufficient energy to equal or exceed the activation energy. (Goal 4.02)
∑ Understand qualitatively that reaction rate is proportional to number of effective collisions.
∑ Explain that nature of reactants can refer to their complexity and the number of bonds that must be broken and reformed in the course of reaction.
∑ Interpret potential energy diagrams.
∑ Explain how temperature (kinetic energy), concentration, and/or pressure affects number of collisions.
∑ Explain how increased surface area increases number of collisions.
Optional: Rate law, rate constant, overall order
Determine mechanisms from rate law
∑ Rate determining step
Explain how a catalyst lowers the activation energy, so that at a given temperature, more molecules will have energy equal to or greater than the activation energy.
Practice:
Which of the following have a higher reaction rate? Why?
1. Using smaller or larger pieces of wood. ______________________________
2. Which will digest faster, powdered sugar or lumps of sugar? __________________
3. Increasing or decreasing temperature. ______________________________
4. Adding or removing a catalyst. ______________________________
5. Adding or removing an inhibitor. ______________________________
6. Burning of methane or ethanol. ______________________________