1. Exothermic release energy into the surroundings temp of surroundings increases products have less energy - ΔE R  P Chapter 13 Endothermic absorb energy from the surroundings temp of surroundings decreases products have more energy +Δ E P  R

2. A + B ----> C + D + 30 joules A + B ----> C + D ΔE = -30 joules Forward reaction is exothermic Reverse reaction is endothermic 30 joules + C + D ----> A + B C + D ----> A + B ΔE = + 30 joules

3. Reactions involve Bond breaking –requires energy +ΔE Bond formation –releases energy -ΔE If more energy is needed to break bonds than is released +ΔE ( endothermic) If more energy is released in bond formation than is absorbed -ΔE (exothermic)

4. Reaction mechanism- steps by which a reaction occurs Each step of a reaction mechanism involves the collision of two molecules.

5. Colliding molecules need 1. enough energy to break the bonds Activation Energy ( A.E) or threshold energy 2. the proper geometry or orientation (correct angle)

6. Energy (P.E.) reaction coordinate c

7. Factors affecting rate of reaction Nature of reactants (number of bonds) Surface area (solids and liquids) Temperature Concentration of reactants (solutions and gases) Catalyst

8. Rate =(# of coll/time) (fract with A.E.) (fract with orientation) If orientation factor is 1 (orientation does not matter) TEMPERATURE affects fraction with A.E.

9. Temp - measure of the average K.E. At a given temp all molecules do not have the same K.E. # of molecules Kinetic energy threshold energy (A.E)

10. Rate =(# of coll/time) (fract with A.E.) (fract with orientation) Concentration affects the # of collisions At a given temp Rate = k (# of collisions /time)

11. k large - fast rxn k small - slow rxn Concentration of reactants raised to some power (order) Rate = (k ) [A] x [B] y X and Y (orders) Orders found only by experiment

12. Rate = (k ) [A] x [B] y If concentration of a reactant is doubled and the rate doubles the order is 1 If concentration of a reactant is doubled and the rate quadruples the order is 2 If changing concentration of a reactant has no effect on the rate of reaction it is not included in the rate law Sum of all orders is the order of the reaction

13. Reaction coordinate Potential energy Catalysts and rate

14. A catalyst does not change the K.E of the molecules # of molecules Kinetic energy  catalyst no catalyst

15. Catalyst -changes the pathway (steps) Steps require less energy Changes the orientation requirement- more molecules have required orientation Not consumed in the reaction

16. Reaction mechanism and Rate Law Each step involves the collision of two molecules –adding steps gives net reaction Each step has its own rate law. In the steps the coefficients are the orders of the rate law Slowest step determines the overall rate -rate determining step The rate law for the slowest step is the rate law for the reaction.

17. Reaction Mechanism A + B-  C + 2D D+ B  DB DB + D  F A + 2B -  C + F

18. Determining Reaction Mechanism 1. Do an experiment to determine the rate law. Experimental rate law 2. Postulate possible steps by which the reaction could take place. Reaction Mechanism

19. A 2 + B 2  2AB rate = k [A 2 ] 2 (from experiment) Possible mechanism 1 A 2 + B 2  2AB rate = k [ A 2 ] [ B 2 ] Predicted does not match experimental Not the reaction mechanism

20. A 2 + B 2 -> 2AB rate = k [A 2 ] 2 (from experiment) Possible mechanism 2 Step 1 A 2 + A 2  2A + A 2 Step 2 A + B 2  AB 2 Step 3 AB 2 + A  2AB Step 1 rate = k [A 2 ][A 2 ] or rate =k [A 2 ] 2 Step 2 rate = k [A][B 2 ] Step 3 rate = k [AB 2 ] [A] Do any rate laws match the experiment rate law? rate = k [ A 2 ] 2

21. 2NO + O 2  2NO 2 ExperimentInitial concentration [NO] Initial concentration [O 2 ] Rate of formation of NO 2 (M/s) 1 0.015 M 0.048 2 0.030 M 0.015 M 0.192 3 0.015 M 0.030 M 0.096 4 0.030 M 0.384