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Unit 13

Unit 13. Reaction Rates. Unit 13 Goals. Understand collision theory Understand how concentration, surface area & temperature, effect reaction rate. Understand potential energy diagrams Predict reaction rate change when concentration, temperature or surface area ∆’s

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Unit 13

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  1. Unit 13 Reaction Rates

  2. Unit 13 Goals • Understand collision theory • Understand how concentration, surface area & temperature, effect reaction rate. • Understand potential energy diagrams • Predict reaction rate change when concentration, temperature or surface area ∆’s • Construct and interpret PE diagrams

  3. Thinker: • Make a list of RATES that you have heard about in your daily lives. • Rate of growth • Rate of pay • Rate of decay • Rate of aging • Rate of flow • Rate of power usage • Rate of speed

  4. Rates of Chemical Change • Rate is how fast something changes with time • Sprinters • Human 28.9 mph • Cheetah 62 mph • Quarter horse 55 mph • Reaction Rate – rate at which reactants are changed to products • Chemical Kinetics – study of Rx rates http://en.wikipedia.org/wiki/Footspeed

  5. Karl Wenzel German Scientist 1777 First to study Rx rates Claude-Louis Berthollet French – 1803 Published first book on Rx rates & concentration Chemical Kinetics http://en.wikipedia.org/wiki/Claude-Louis_Berthollet

  6. Rate = Change over time • CH3OCH3(g) CH4(g) + CO(g) + H2(g) • At 500 C, dimethyl ether decomposes • As Rx continues: • Conc. of dimethyl ether will decrease • Time will continue to change • Therefore: • Note: all reactants are -

  7. CH3OCH3(g) CH4(g) + CO(g) + H2(g) • We can talk about the rates of reactants OR products: • Note: all products are +

  8. What about coefficients? • Consider this Rx: • 2N2O5(s) 4NO2(g) + O2(g) • How do coefficients affect reactant/product? • Divide by each coefficient: • Think about how many per time unit

  9. Measuring Reactions • Need to keep track of at least 1 concentration: • Color change • Pressure change • Temperature change • Rates is measured as is indicated by the rate you write.

  10. Calculate the rate of reaction between 0 and 20.0 s: 0.0000585 M/s Calculate the rate of reaction between 20.0 and 40.0 s 0.0000528 M/s Calculating Rx Rates:

  11. Will Reaction Rate always be the same? • 2N2O5(s) 4NO2(g) + O2(g) • No! • Reactant is constantly decreasing • Product is constantly increasing slope of graph = rate of Rx

  12. Factors Affecting Rate • Concentration • Pressure • Temperature • Surface Area

  13. Concentration • Almost always: • Increase in concentration of reactants yields in increase of reaction rate NO2(g) + CO(g) NO(g) + CO2(g) • If [NO2] is doubled • More collisions with CO will occur • More collisions means more successful collisions • More successful collisions means more product • The inverse is also true

  14. Non-collision reactions • Decomposition does not require collisions (CH2)3(g) CH2=CH-CH3(g) • The more cyclopropane you have: • The more is available to decompose • Thus a faster reaction rate

  15. Pressure • In liquid or solid, has almost no affect • In gasses: • Doubling pressure essentially doubles the concentration. •  by increasing pressure, we increase reaction rates in gasses. • Inverse is true

  16. Temperature • In General: • Increase in temperature will increase reaction rate. • Often increase of 10% temp yields in 2 times the reaction rate! • Our bodies work best around 37 C • Small changes cause great distress! • Snakes, lizards & other reptiles

  17. Surface Area • Increased surface area yields an increase in reaction rate. • Wood & fires • Grain Elevator

  18. Questions? • HW: Read 16.1 • P. 585 Section Review: • 1-3, 6, 8, 11 – 13

  19. Rate Law • Describes how reactant concentration affect the reaction rate. • Provides the ability to form a reaction mechanism. • Model of how the reaction occurs • Discusses # of steps necessary

  20. Determining a General Rate Law Equation • When there is a single reactant: • Rate of reaction is proportional to the concentration raised to some power (n) • The power is the order of the Rx. Rate = k[reactant]n • k is the rate constant • Proportionality that varies with temperature • Order is usually a whole number (1, 2) • Can be fractions • Occasionally is 0 • Reaction rate is independent of concentration

  21. Determining Order (example) • Experiments were preformed to measure the initial rate of the reaction 2HI(g)  H2(g) + I2(g). Conditions were identical except that the HI concentrations were varied. • rate = k[HI]n • n=? • Find ratio of the reactant concentrations b/t experiments. • Find ratio of the reaction rates. • Plug into rate = k[HI]n • Verify the Results (compare 3 and 1)

  22. Other questions • Given the previous reaction: • What will happen if [HI] is increased from 4 M to 8 M? • What will happen if [HI] is decreased from 6 M to 2 M? • What will happen if [HI] is quadrupled?

  23. Questions? • Homework: • Worksheet

  24. Rate Laws for Several Reactants • When there are several reactants the rate law applies to each. •  an order applies to each reactant NO(g) + O3(g) NO2(g) + O2(g) • rate = k[NO3]n1[O3]n2 • It turns out: • n1 = n2 = 1 • Since each order is equal, it suggests that this reaction has a 1 step mechanism. • If they weren’t 1, then the mechanism could be longer than 1 step.

  25. Rate Determining Step • Despite being written as 1 step, many Rx, aren’t quite that simple. • The mechanism is more complex. • Example: 2Br-(aq) + H2O2(aq) + 2H3O+(aq) Br2(aq) + 4H2O(l) • Has 4 separate steps (each reactants order 1) • Br-(aq) + H3O+(aq) HBr(aq) + H2O(l) • HBr(aq) + H2O2(aq) HOBr(aq) + H2O(l) • Br-(aq) + HOBr(aq)  Br2(aq) + OH-(aq) • OH-(aq) + H3O+(aq)  2H2O(l) • Each step produces intermediates. • Species that are produced, but consumed • If one of the steps is slower, it is known as the rate determining step • In this case it is step 2

  26. Collision Theory: • Reactions occur because: • Substances (atoms, ions, molecules) must collide • As particles get closer, they repel, so they must have enough KE to actually collide • Collisions must occur in correct orientation • Collisions must have sufficient energy to form the activated complex • This energy is known as activation energy (Ea)

  27. A + B  C + D Imagine trying to bowl uphill This is what molecules need to do Activation energyis required to formactivated complexat top of curve. # peaks = # of steps Potential Energy (PE) Diagram http://commons.wikimedia.org/wiki/Image:Coordenada_reaccion.GIF

  28. Exothermic vs. Endothermic

  29. Collision Theory • Reacting substances (atoms, ions, or molecules) must collide • Collisions must occur with the correct orientation • Collisions must have sufficient energy to form the activated complex. • Activated complex is a temporary unstable arrangement of atoms that may form products or break apart to re-form the reactants

  30. Unsuccessful Collision (not enough energy)

  31. Unsuccessful Collision (wrong orientation)

  32. Successful Collision

  33. Catalyst – chemical in a reaction that lowers the activation energy, but never is used. Save energy Usually costly Biological catalysts are called enzymes Catalysts & Potential Energy http://en.wikipedia.org/wiki/Image:Activation_energy.svg

  34. Homework: • PE Diagram Worksheet

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