Reaction Rate

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					Chapter 16
 The time taken for the disappearance of the
  reactant or the appearance of the product .
  Rate is a ratio as the amount of reactant
  disappeared divided by the time.
 Average rate: The change in the
  concentration divided by the total time
 Rate = amount reacted or produced/ time
 units: g/s, mol/s, or %/s
 Instantaneous   rate: rate measured
 between very short interval
 Initial rate: instantaneous rate at the
 beginning of an experiment
 Page592
 Concept check
 Rate depends on the concentration:
 In some reactions doubling the
  concentration doubles the rate of
  reaction. In some doubling the reaction
  increases the reaction four folds.This
  happens in the decomposition of HI to
  form H 2 and I2.
 An  expression for the rate of a reaction as
  a function of the concentration of one or
  more of the reactants.
 Rate=k [A]n
 This equation is the general rate law. The
  exponent n , is called the order with
  respect to substance A and must be
  determined from experimental data.
 Order of a chemical reaction can be said
 as the exponent on the concentration for
 a specified reactant in a rate law
 Determine  the rate law equation for the
  following reaction , given the
  experimental data
 3AC
 Concentration of A       Reaction rate
 0.2M                         1.0M/s
 0.4M                         4.0M/s
 Page  596
 Practice problems
 The  rate law equations you have looked
  so far have been for reactions involving
  only one reactant.
 If more than one reactant is found to
  contribute to the rate of the reaction, then
  all contributing reactants must appear in
  the rate law.
 The rate law equation for this will be
 rate=k [A] n [B] m
 The  value of n is the order with respect to
  reactant A. The value of m is the order with
  respect to reactant B. The overall reaction
  order will be the sum of n and m.
 From the above equation if you double the
  concentration of A and the rate doubles then
  the reaction is first order with respect to A. If
  you double the conc. of B (keeping the conc
  of A constant, and the rate quadruples the
  the rate of the reaction is second order with
  respect to B.
 For the reaction A and B for this example
  the rate law would be rate=k[A][B]2
 Rate laws cannot be derived from a
  chemical equation.
 2N2O₅↔4NO2+O2
 Keq=[NO2] 4 [O2]/[N2O₅]2

 rate=k[N2O₅]1
 The  slowest step in a mechanism, the step
  that determines the overall rate of
  reaction is the rate determining step.
 Mechanism is a proposed sequence of
  steps that describes how reactants are
  changed into products.
 Each step in the mechanism is called as
  elementary step.
 Page  598
 Critical thinking 2,3,5
 Practice problems 7and 8 all
 Temperature: An increase in temperature is
  accompanied by an increase in the reaction rate.
  Temperature is a measure of the kinetic energy
  of a system, so higher temperature means higher
  average kinetic energy of molecules and more
  collisions per unit time.
 For most chemical reactions the rate at which the
  reaction proceeds will approximately double for
  each 10°C increase in temperature. Once the
  temperature reaches a certain point, some of the
  chemical species may be altered (e.g.,
  denaturing of proteins) and the chemical
  reaction will slow or stop.
 Concentration: A higher concentration of
 reactants leads to more effective
 collisions per unit time, which leads to an
 increasing reaction rate (except for zero
 order reactions). Similarly, a higher
 concentration of products tends to be
 associated with a lower reaction rate.
 Medium: The    rate of a chemical reaction
 depends on the medium in which the
 reaction occurs. It sometimes could make
 a difference whether a medium is
 aqueous or organic; polar or nonpolar; or
 liquid, solid, or gaseous.
 Surface   area: It is easier to dissolve sugar
 if it is crushed. Crushing the sugar
 increases its surface tension.The larger
 surface area allows more sugar
 molecules to contact the solution.
 Catalyst: A catalyst is a substance that alters
  the rate of a chemical reaction without
  being used up or permanently changed
 A catalyst works by changing the energy
  pathway for a chemical reaction. It provides
  an alternative route (mechanism) that
  lowers the Activation Energy meaning more
  particles now have the required energy
  needed to undergo a successful collision.
 What  is activation energy?
 The least amount of energy needed to
  permit a particular chemical reaction.
 There are 2 types of catalysts:
 Homogeneous catalyst: Homogeneous
  catalysts are in the same phase as the
 Heterogeneous catalyst: Heterogeneous
  catalysts are present in different phases
  from the reactants (for example, a solid
  catalyst in a liquid reaction mixture),
  whereas homogeneous catalysts are in the
  same phase (for example, a dissolved
  catalyst in a liquid reaction mixture).
 Example    of Homogeneous catalyst
 2H2O2(aq)+ KI(aq)2H2O(l)+O2(g)
 Example of Heterogeneous catalyst
 Decomposition of H2O2 in presence of MnO2
 Hydrogen peroxide is a solution while
  manganese dioxide is a solid and can be easily
 Term  review all
 Page 614
 13,22, 23 and 25
 Test prep all

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