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Chemical Kinetics

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  1. Chemical Kinetics Reaction Rates

  2. Kinetics: Reaction Rates • Rate of a chemical reaction = change in concentration/time • Rate = [ A ] t • = change (take the slope!) • [ ] = molar concentration

  3. Reaction Rate: Take the slope of the curve At which time is the reaction rate greater? [A] t1 t2 time

  4. Reaction Kinetics: AB [A] rate=slope [B] concentration time

  5. Measuring reaction rates

  6. Measured Reaction Rates • The rate of reaction changes with time, as the concentrations of reactants and products change. • The initial rate is usually faster than later in the reaction. Why?

  7. Collision Theory • In order to react, particles must collide with enough energy to break and reform bonds • They must collide with the proper orientation (Hit in the right way)

  8. Factors affecting the rate of a reaction • Surface area if a solid is present • Concentration of solutions • Pressure of a gas • Temperature • Presence of a catalyst

  9. Size of particles • Which will react faster? • Hydrochloric acid and large pieces of magnesium metal or finely ground pieces of magnesium metal? • Relate this to collision theory

  10. Size of particles/surface area

  11. Size of particles/ surface area

  12. Concentration of solution • Which will react faster? 6 M hydrochloric acid or 1 M hydrochloric acid and magnesium metal? Relate this to collision theory

  13. Which contains a higher concentration?

  14. Pressure of a gas • Which will react faster? • Hydrogen and oxygen gases at high pressure or low pressure? • Relate this to collision theory

  15. Pressure of a gas

  16. Temperature • Which will react faster? • A reaction at high temperature or at low temperatures? • Note: When the temperature is higher , the average kinetic energy is higher. • Relate to collision theory

  17. If the temperature increases by 10oC, The reaction rate is twice as fast

  18. Catalyst • A substance that increases the rate of a reaction without being consumed in the reaction. • Catalyst provides an alternative pathway from reactant to product. The mechanism of the reaction changes. • The activation energy is lowered sot the reaction will proceed faster at a given temperature. • The catalyst is not consumed during the reaction, so only small amounts are needed.

  19. A Catalyst lowers the activation energy

  20. Homogeneous and Heterogeneous Catalysts • Homogeneous catalyst: catalyst existing in the same phase as the reactants. • Heterogeneous catalysis: catalyst existing in a different phase than the reactants.

  21. Homogeneous Catalyst: CFC’s in Catalyzing depletion of the Ozone layer

  22. Heterogeneous Catalyst • The catalytic hydrogenation of ethylene is an example of a heterogeneous catalysis reaction CH2=CH2(g) + H2(g) --Ni--> CH3CH3(g) Nickel is solid while the reactants are gases.

  23. 2 CO + 2 NO  2 CO2 + N2

  24. H2O2 H2O + O2 • Manganese dioxide is a catalyst that lowers the activation energy of the decomposition of hydrogen peroxide.

  25. Rates and Stoichiometry For the general reactionA + 2B + 3C + …  … X + 2Y + 3Z

  26. Example • For the reaction: N2 + 3H2 2 NH3 If the rate of reaction is the change in the concentration of nitrogen, what is the rate of the reaction in terms of hydrogen and ammonia?

  27. Example • The reaction of hydrogen with nitrogen to produce ammonia is 3H2 + N2 2NH3 • If the rate of appearance of NH3 is 4.0 x 10-6 M/sec, what is the rate of disappearance of N2? Answer: -2.0 x 106 M/sec

  28. Edited from: http://www.cdli.ca/sampleResources/chem3202/unit01_org01_ilo03/b_activity.html