1. Reaction Rates Measures concentration (molarity!) change over time Measures concentration (molarity!) change over time Example: Example: 2H 2 O 2  2H 2 O + O 2 2H 2 O 2  2H 2 O + O 2 [http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/TB15_01.JPG]

2. [http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/FG15_02.JPG] -1.7 x 10 -3 M/s Slopes of the tangent lines give instantaneous reaction rates

3. [http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/FG15_02.JPG] -6.3 x 10 -4 M/s (y 1 – y 2 ) (x 1 – x 2 ) = (1.75 – 0) (0 – 2800) =

4. Rate Laws An equation that calculates the reaction rate of a reaction. An equation that calculates the reaction rate of a reaction. A + B  C + D A + B  C + D Rate = k[A] x [B] y Rate = k[A] x [B] y

5. Rate Laws [A] & [B] – Molarity of reactants [A] & [B] – Molarity of reactants k – Rate constant (experimentally derived) k – Rate constant (experimentally derived) x & y – Exponents of reactants (experimentally derived, not coefficients!) x & y – Exponents of reactants (experimentally derived, not coefficients!)

6. Rate Laws

7. Example: Example: 2 NO + O 2  2 NO 2 2 NO + O 2  2 NO 2 Rate = k[NO] 2 [O 2 ] Rate = k[NO] 2 [O 2 ] If each reactant starts out at 1.00 M, how will the rate increase if [NO] is doubled? If each reactant starts out at 1.00 M, how will the rate increase if [NO] is doubled?

8. Rate Laws Initially: Initially: Rate = k[1.00] 2 [1.00] Rate = k[1.00] 2 [1.00] Rate = k Rate = k [NO] Doubled: [NO] Doubled: Rate = k[2.00] 2 [1.00] Rate = k[2.00] 2 [1.00] Rate = 4.00k (increases by factor of 4) Rate = 4.00k (increases by factor of 4)

9. Rate Laws Example: Example: What if [O 2 ] is doubled instead? What if [O 2 ] is doubled instead? Rate = k[1.00] 2 [2.00] Rate = k[1.00] 2 [2.00] Rate = 2.00k (only increased by factor of 2) Rate = 2.00k (only increased by factor of 2)