1. What is this?

2. Kinetics Reaction Rates: How fast reactions occur

3. How do we measure rxn rates?  Rates must be measured by experiment  Indicators that a reaction is happening Color change Gas evolution Precipitate formation Heat and light  Many ways to measure the rate Volume / time Concentration / time Mass / time Pressure / time

4. How do we measure rxn rate? A  B  How fast product appears  How fast reactant disappears

5. Forward vs Reverse Rxn  Some rxns are reversible  After a sufficient amount of product is made, the products begin to collide and form the reactants  We will deal only w/ rxns for which reverse rxn is insignificant 2 N 2 O 5 (aq)  4 NO 2 (aq) + O 2 (g)  Why is reverse rxn not important here?

6. Rate Law  Math equation that tells how reaction rate depends on concentration of reactants and products  Rates = k[A] n K = rate constant / proportionality constant n = order of reaction  Tells how reaction depends on concentration Does rate double when concentration doubles? Does rate quadruple when concentration doubles?

7. 2 kinds of rate laws  Both determined by experiment  Differential Rate Law How rate depends on [ ]  Integrated Rate Law How rate depends on time

8. Differential Rate Law  2 methods Graphical analysis Method of initial rates

9. Graphical Analysis 1. Graph [ ] vs. time 2. Take slope at various pts 3. Evaluate rate for various concentrations

10. [N 2 O 5 ] (M) Rate (M/s) 1.02 0.51.0 0.250.5 Graphical Analysis  When concentration is halved… Rate is halved Order = 1 Rate = k[N 2 O 5 ] 1

11. [NO 2 ] (M) Rate (M/s) 1.02 2.08 4.032 Graphical Analysis  When concentration is doubled… Rate is quadrupled Order = 2 Rate = k[N 2 O 5 ] 2

12. Method of Initial Rates  Initial rate calculated right after rxn begins for various initial concentrations  NH 4 + (aq) + NO 2 - (aq)  N 2 (g) + 2H 2 O(l)  Rate = k [NH 4 + ] n [NO 2 - ] m [NH 4 + ][NO 2 - ]Rate (M/s) 0.1 2 0.24 6

13. [NH 4 ][NO 2 - ]Rate 0.1 2 0.24 8 [NH 4 ][NO 2 - ]Rate 0.1 2 0.24 6 When [NO2] doubles, rate doubles, First order with respect to (wrt) NO2 m = 1 When [NO2] doubles, rate doubles, First order with respect to (wrt) NO2 n = 1 Rate = k[NH 4 +] [NO 2 -]

14. Try this one: Rate = k [NO 2 - ] 2 [NH 4 + ][NO 2 - ]Rate (M/s) 0.1 2 0.28 8 Calculate k, using any of the trials, you should get the same value

15. Integrated Rate Law  Tells how rate changes with time  Laws are different depending on order  Overall reaction order is sum of exponents Rate = k  zero order Rate = k[A]  first order Rate = k[A] 2  second order Rate= k[A][B]  second order

16. First order integrated rate law  Rearrange and use some calculus to get:  This is y = mx + b form A plot of ln[A] vs time will give a straight line  If k and [A] 0 (initial concentration) known, then you know the concentration at any time

17. Second order integrated rate law  Rearrange and use some calculus to get:  This is y = mx + b form A plot of 1/[A] vs time will give a straight line  If k and [A] 0 (initial concentration) known, then you can now the concentration at any time

18. Zero order integrated rate law  Rearrange and use some calculus to get:  This is y = mx + b form A plot of [A] vs time will give a straight line  If k and [A] 0 (initial concentration) known, then you can now the concentration at any time

19. Graphs give order of rxn  Use graphs to determine order If [A] vs time = zero order If ln [A] vs time = first order If 1/ [A] vs time = second order

20. Half-life  Def’n: time it takes for concentration to halve  Depends on order of rxn  At t 1/2 [A]=[A] 0 /2

21. Half-life: First order

22. Half-Life  First order  Second order  Zero Order