Chemistry 100 Chapter 14 - Chemical Kinetics Chemistry 10

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Chemistry 100

• Chapter 14 - Chemical Kinetics

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The Connection Between Chemical Reactions and Time

• Not all chemical reaction proceed
  instantaneously!!!
• 2 A B combination reaction
• Can we quantify the length of time it takes for
  this (or any) chemical reaction to occur

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Practical examples of how long!

• H2(g) ½ O2 (g) H2O (l) Þ Very Slow
• N2O(g) N2 (g) ½ O2 (g) Þ SLOW
• combustion reactions fast process
• TNT exploding very fast reaction
• Food spoilage
• Drug decomposition

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Chemical Kinetics

• Chemical kinetics is concerned with determining
  the speed or rate at which a reaction occurs.
• How is the reaction rate affected by
• temperature
• states of reactants
• amount of reactants
• catalyst
• surface area of the reacting species

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Example

• Br2 (aq) HCOOH (aq) 2 Br - (aq) 2 H (aq)
  CO2 (g)
• Define the average rate

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A Sample

• Reaction
• Br2 (aq) HCOOH (aq) 2 Br (aq) CO2 (g) 2
  H (aq)

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Average Rate Data
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Instantaneous Rates

• Its best to define an instantaneous speed of
  reaction

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Instantaneous Rate Data
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The Rate Constant

• The rate constant relates the speed of the
  chemical reaction to the instantaneous reactant
  concentration.
• k constant for constant temperature
• The rate of the reaction is dependent on reactant
  concentration
• RATE CONSTANT IS INDEPENDENT OF THE REACTANT
  CONCENTRATION.

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Reaction Rates and Reaction Stoichiometry

• Look at the reaction
• O3(g) NO(g) NO2(g) O2(g)

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Another Example

• 2 NOCl (g) 2 NO 1 Cl2 (g)

WHY 2 moles of NOCl disappear for every 1 mole
Cl2 formed.
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The General Case

• a A b B c C d D
• rate -1 A -1 B 1 C 1
  D
• a t b t c t
  d t
• Why do we define our rate in this way
• removes ambiguity
• obtain a single rate for the entire equation not
  just for the change in a single reactant or
  product.

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The Rate Law

• Relates rate of the reaction to the reactant
  concentrations and rate constant
• For a general reaction
• a A b B c C d D e E
• rate kAxByCz

The exponents (xy and z) are called the
reactant orders.
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The Reaction Orders

• Determine the superscripts (x y and z) for a
  non-elementary chemical reaction by
  experimentation.
• S x y z reaction order
• e.g. x 1 y 1 z 0
• 2nd order reaction (x y z 2)
• x 0 y 0 z 1 (1st order reaction)
• x 2 y 0 z 0 (2nd order)

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The Isolation Method of Obtaining Rate Laws

• a A b B c C d D
• rate kAxBy
• Fix the concentration of one reactant (say
  reactant A).
• We then perform a series of experiments to
  examine how changing the B affects the initial
  reaction rate
• rate (constant) By

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Isolation Method (contd)

• Now we fix the concentration of reactant B.
• We then perform another series of experiments to
  examine how changing the A affects the initial
  reaction rate
• rate (constant) Ax

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Types of Reactions

• The rate law gives us information about how the
  concentration of the reactant varies with time
• How much reactant remains after specified period
  of time

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First Order Reaction

• A product
• rate -DA/D t kA
• How does the concentration of the reactant
  depend on time
• k has units of s-1

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The Half-Life of a First Order Reaction

• For a first order reaction the half-life t1/2 is
  calculated as follows.

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Radioactive Decay

• Radioactive Samples decay according to first
  order kinetics.
• This is the half-life of samples containing e.g.
  14C 239Pu 99Tc.
• Example

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Second Order Reaction

• A B products Rate kAB
• A products Rate kA2
• Reaction 1 is 1st order in A and B and 2nd order
  overall
• Reaction 2 is 2nd order in A

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The Dependence of Concentration on Time

• For a second order process where rate kA2

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Half-life for a Second Order Reaction.

• A at t t½ ½ A0

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A Pseudo-First Order Reaction

• Example hydration of methyl iodide
• CH3I(aq) H2O(l) CH3OH(aq) H(aq) I-(aq)
• Rate k CH3I H2O
• Since we carry out the reaction in aqueous
  solution
• H2O gtgtgtgt CH3I / H2O doesnt change by a lot

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• Since the concentration of H2O is essentially
  constant
• rate k CH3I constant
• k CH3I where k k H2O
• Pseudo first order since it appears to be first
  order but it is actually a second order process.

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Collision Theory of Kinetics

• With few exceptions the reaction rate increases
  with increasing temperature.
• Chemical reactions take place due to collisions
  between reactant molecules
• i.e. rate µ number of collisions / unit time
• A2 B2 product rate kA2B2

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The Reaction Profile

• How does the energy of the reactants vary during
  the reaction sequence

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The Activation Energy

• The minimum amount of energy need for initiation
  of a chemical reaction is the activation energy
  (Ea).
• Colliding reactant molecules possess kinetic
  energy gt the activation energy or Ea.

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The Activated Complex

• The species temporarily formed by the reactant
  molecules the activated complex.
• A small fraction of molecules usually have the
  required kinetic energy to get to the transition
  state
• The concentration of the activated complex is
  extremely small.

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The Arrhenius Equation

• Arrhenius showed how the rate constant depended
  on temperature.

Ea is the activation energy
A is called the frequency factor an estimate of
the number of reactive collisions in the system
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Activation Energies and the Arrhenius Equation

• Reaction possesses a large activation energy
  small rate constant
• Slow reaction!!
• Measure k at several different temperatures

R 8.314 J/(K mole) T in Kelvin units!!!
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Arrhenius Equation (contd)

• The Arrhenius equation is best suited for
  studying reactions between simple species (atoms
  diatomic molecules).
• The orientation of the reactants (how they
  collide) becomes very important when the species
  get bigger.

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Catalysts

• So far we have considered one way of speeding up
  a reaction
• increasing T usually increases k.
• Another way is by the use of a catalyst.
• A catalyst - a substance that speeds up the rate
  of the reaction without being consumed in the
  overall reaction.

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• look at the following two reactions
• AB C rate constant k
• AB C rate constant with catalyst is kc
• NOTE RATE WITH CATALYST gt RATE WITHOUT CATALYST

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Types of Catalyst

• We will briefly discuss three types of catalysts.
  The type of catalyst depends on the phase of the
  catalyst and the reacting species.
• Homogeneous
• Heterogeneous
• Enzyme

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Homogeneous Catalysis

• The catalyst and the reactants are in the same
  phase
• e.g. Oxidation of SO2 (g) to SO3 (g)
• 2 SO2(g) O2(g) 2 SO3 (g) SLOW
• Presence of NO (g) the following occurs.
• NO (g) O2 (g) NO2 (g)
• NO2 (g) 2 SO2 (g) 2 SO3 (g) NO (g) FAST

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• SO3 (g) is a potent acid rain gas
• H2O (l) SO3 (g) H2SO4 (aq)
• Note the rate of NO2(g) oxidizing SO2(g) to
  SO3(g) is faster than the direct oxidation.
• NOx(g) are produced from burning fossil fuels
  such as gasoline coal oil!!

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Heterogeneous Catalysis

• The catalyst and the reactants are in different
  phases
• adsorption the binding of molecules to the
  surface to a surface.
• Adsorption on the surface occurs on active sites.
• An active site is a place where reacting
  molecules are adsorbed and physically bond to the
  metal surface.

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• The hydrogenation of ethene (C2H4 (g)) to ethane
• C2H4 (g) H2(g) C2H6 (g)
• Reaction is energetically favourable
• rH -136.98 kJ/mole of ethane.
• With a finely divided metal such as Ni (s) Pt
  (s) or Pd(s) the reaction goes very quickly .

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Common Heterogeneous Catalysts

• Two other important heterogeneous catalysis
  processes
• petroleum cracking (refining crude oil)
• catalytic converters very efficient in
  reducing exhaust emission when hot cold is
  another story!

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Enzyme Catalysis

• Enzymes - proteins (M gt 10000 g/mol)
• High degree of specificity (i.e. they will react
  with one substance and one substance primarily
• Living cell gt 3000 different enzymes

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The Lock and Key Hypothesis

• Enzymes are folded into fixed configurations.
• According to Fischer active site is rigid.
• The substrates molecular structure exactly fits
  the lock (hence the key).

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Simplified Model for Enzyme Catalysis

• E º enzyme S º substrate P º product
• E S ES
• ES P E
• rate k ES
• The reaction rate depends directly on the
  concentration of the substrate.

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Rate Laws for Multistep Processes

• Chemical reactions generally proceed via a large
  number of elementary steps - the reaction
  mechanism
• The experimentally established rate law must
  reflect the reaction rate of the slowest
  elementary step Þ the rate determining step (rds)

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What do we mean by an rds

• A commuter goes through a two step process to get
  to work in Halifax.
• (1) highway MacKay Bridge Toll booth
• (2) toll booth downtown

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MacKay Bridge Toll Booth

• Overall reaction
• highway downtown
• Situation 1. highway clogged toll booth is
  fast.
• Situation 2. fast highway clogged toll booth.

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The Rate Determining Step (rds)

• Situation 1 - clogged highway is the slowest step
  in the commuting process (rds).
• Situation 2 - the clogged toll-booth is the
  slowest step in the commuting process (the rds).
  
• Speed of overall process (highway downtown)
  depends which step is slowest!
• Which is the rate-determining step.

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Elementary steps and the Molecularity

• Any chemical reaction occurs via a sequence of
  elementary steps.
• Kinetics of the elementary step only depends on
  the number of reactant molecules in that step!
• Molecularity the number of reactant molecules
  that participate in elementary steps

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The Kinetics of Elementary Steps

• Classes of elementary steps

A unimolecular step
A bimolecular step
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• For the step

A termolecular (three molecule) step.