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Reaction Rates

Reaction Rates. AP chapter 14.3. Reaction Rates. Describe how quickly concentration of reactants or products are changing Units typically D M/ D t for aqueous reactants and products Could be units of D P/ D t for gaseous products

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Reaction Rates

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  1. Reaction Rates AP chapter 14.3

  2. Reaction Rates • Describe how quickly concentration of reactants or products are changing • Units typically DM/Dt for aqueous reactants and products • Could be units of DP/Dt for gaseous products • Effective concentration solids and liquids does not change over the course of a reaction, so it will be more difficult to model these changes

  3. Reaction equations • Rate=k[A]m[B]n • [A] is a symbol roughly meaning “concentration of” (in units of molarity or partial pressure) • More precisely, it means “Activity of” or “effective concentration”. Remember, that even in an aqueous solution, not all of the ions are available to react. • This is given only as one example • The equation for each type of reaction must be tested experimentally. The values of m and n are determined experimentally

  4. Reaction orders • Rate=k[A]m[B]n • The reaction above is “m” order for reactant A, “n” order for reactant B, and “m+n” order for the reaction overall. “Reaction order” describes the influence which increasing concentration has on the reaction rate

  5. You Try • Describe the reaction orders for the following rate equation: • Rate=k[A]2[B]1 • Sketch the shape of the graph showing the relationship between • [A] and rate • [B] and rate • If [A] and [B] are measured in molar units, what is the unit for the rate constant?

  6. What’s the relationship between concentration and time? • Note that as time passes, the concentrations change, so the rate changes. • Based on this analysis will rate increase or decrease as time passes? • The general solutions for these relationships require some calculus, but even without calculus, you can learn the equations for some simple cases.

  7. Integrated rate laws • These show the relationship between [A] and time over the course of a single experiment • First order: Rate=k[A] • First order: ln [A] = -kt + ln[A]0 • Other integrated rate laws • Sketch the graph of ln[A] against t

  8. Collision model • Reactions occur when molecules collide • Even for unimolecular reactions, collisions are necessary to change the kinetic energy of colliding particles • Example: 3 O2 + hn 2 O3 • Example: Cl2 + F2  2 ClF

  9. Distribution of kinetic energies • Temperature is proportional to average kinetic energy for a collection of molecules • However, the kinetic energy of any individual molecule could be a little less or a little more. • Distribution of molecular energies is a predictable function • http://intro.chem.okstate.edu/1314F00/Laboratory/GLP.htm

  10. Reactions at a molecular level • Molecules only react if they possess enough combined energy to overcome the activation energy. • The orientations of the molecules also matters. • http://www.mp-docker.demon.co.uk/chains_and_rings/mechanisms/index.html

  11. Arrhenius equation • Relates Temperature, reaction rate and activation energy. • Various forms of the Arrhenius equation

  12. Reaction pathway diagrams • Show the relative potential energy of reactants, products, intermediates and transition states. • “intermediates” are the various molecular forms which appear as reactant becomes product. • Depending on context, “intermediate” could mean only the stable intermediates, or could also include short lived (transient) “transition states” • Energy diagram

  13. Catalysts • Catalysts are substances which speed up a reaction, without being altered or consumed in the process • Catalysts may be temporarily altered as part of one of the reaction intermediates. • How Catalysts work

  14. How catalysts work at a chemical level • Catalysts lower activation energy, either by bringing reactants closer together, or otherwise stabilizing the reaction transition state. • A catalyst speeds up both forward and reverse reactions, so the mixture comes to equilibrium more rapidly. A catalyst can not change the equilibrium concentrations • What do you think would be the result of adding a catalyst to a mixture already at equilibrium?

  15. Enzymes • “Enzyme” is a term for a catalyst found in, or obtained from a biological system. • Enzymes are primarily made of protein, but may also include metal ions, nucleic acids, or other structural materials. • How enzymes work

  16. Non-enzyme catalysts • Both enzyme and non-enzyme catalysts are important in manufacturing processes, such as the synthesis of ammonia • Catalysts can be homogeneous (in the same phase as the reaction) or heterogeneous (a finely divided metal in solution, for example.

  17. Heterogeneous catalysts • Why do you think heterogeneous catalysts must be finally divided? • What other methods do chemical engineers to increase the surface area of a catalyst?

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