Chemical Kinetics

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Chemical Kinetics

• Collision Theory
• How reactions takes place
• Reaction Rates
• How fast reactions occur
• Reaction Mechanisms

Resource www.mwiseman.com
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Why are kinetics important?

• In order to control processes.
• speed up useful reactions that occur too slowly
• slow down reactions that are harmful
• Example
• Catalysts are used in our cars to
  rapidly convert toxic substances into safer
  substances
• Refrigerators are used to slow the process of
  spoiling in food

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Collision Theory

• How do reactions occur at the molecular level?
• Molecules collide with each other
• Form activated complex
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/animations/NOO3singlerxn.html
• collisions
• http//www.mhhe.com/physsci/chemistry/essentialche
  mistry/flash/collis11.swf
• correct and incorrect collisions

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The area under the curve is a measure of the
total number of particles present.
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• Svante Arrhenius
• Did some fancy math to figure out that number of
  collisions alone dont account for reaction rates
• He found that reactants also require
• Activation energy (Ea - energy to break bonds)
• Right orientation
• http//www.mhhe.com/physsci/chemistry/essentialche
  mistry/flash/activa2.swf
• transition state

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Not all collisions leads to a reaction For
effective collisions proper orientation ofthe
molecules must be possible
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What affects reaction rate?

• Temperature
• http//www.sciencepages.co.uk/keystage4/GCSEChemis
  try/rate5
• concentration and temperature
• Increased number of collisions
• More molecules have enough activation energy
• Remember Maxwell-Boltzmann distribution
• Increased temperature, distribution flattens out
• More molecules
  have Ea

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What affects reaction rate?

• Higher concentration
• Number of collisions increased
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/animations/O2NO220kinetics8.htm
  l
• concentration
• Increased surface area
• Number of collisions increased

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What affects reaction rate?

• Catalysts
• Defn substance that speeds up a rxn w/o being
  used up itself
• Number of collisions with Ea increase
• Ea lowers
• Catalysts hold molecules in right orientation
• Homogeneous catalyst (same phase of matter)
• Demo Catalysis by Co2
• Heterogeneous catalyst (different phase)
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/animations/Catalyst2NOO2N28.html
• catalyst

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What is this?
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How do we measure rxn rates?

• Rates must be measured by experiment
• Indicators that a reaction is happening
• Color change
• Gas formation
• Precipitate formation
• Heat and light
• Many ways to measure the rate
• ?Volume / time
• ?Concentration / time
• ?Mass / time
• ?Pressure / time

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How do we measure rxn rate?

• A ? B
• How fast product appears
• How fast reactant disappears

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Forward vs Reverse Rxn

• Some rxns are reversible
• After a sufficient amount of product is made, the
  products begin to collide and form the reactants
• We will deal only w/ rxns for which reverse rxn
  is insignificant
• 2 N2O5(aq) ? 4 NO2(aq) O2 (g)
• Why is reverse rxn not important here?

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Rate Law

• Math equation that tells how reaction rate
  depends on concentration of reactants and
  products
• Rates kAn
• K rate constant / proportionality constant
• n order of reaction
• Tells how reaction depends on concentration
• Does rate double when concentration doubles?
• Does rate quadruple when concentration doubles?

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2 kinds of rate laws

• Both determined by experiment
• Differential Rate Law
• How rate depends on
• Integrated Rate Law
• How rate depends on time

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Differential Rate Law

• 2 methods
• Graphical analysis
• Method of initial rates

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Graphical Analysis

• Graph vs. time
• Take slope at various pts
• Evaluate rate for various concentrations

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Graphical Analysis
N2O5 (M) Rate (M/s)
1.0 2
0.5 1.0
0.25 0.5

• When concentration is halved
• Rate is halved
• Order 1
• Rate kN2O51

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Graphical Analysis
NO2 (M) Rate (M/s)
1.0 2
2.0 8
4.0 32

• When concentration is doubled
• Rate is quadrupled
• Order 2
• Rate kN2O52

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Method of Initial Rates

• Initial rate calculated right after rxn begins
  for various initial concentrations
• NH4(aq) NO2-(aq) ? N2(g) 2H2O(l)
• Rate k NH4nNO2-m

NH4 NO2- Rate (M/s)
0.1 0.1 2
0.1 0.2 4
0.2 0.2 6
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NH4 NO2- Rate
0.1 0.1 2
0.1 0.2 4
0.2 0.2 8
NH4 NO2- Rate
0.1 0.1 2
0.1 0.2 4
0.2 0.2 6
When NO2 doubles, rate doubles, First order
with respect to (wrt) NO2 n 1
When NO2 doubles, rate doubles, First order
with respect to (wrt) NO2 m 1
Rate kNH4 NO2-
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Try this one
NH4 NO2- Rate (M/s)
0.1 0.1 2
0.1 0.2 8
0.2 0.2 8

• Rate k NO2-2

• Calculate k, using any of the trials, you should
  get the same value

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Integrated Rate Law

• Tells how rate changes with time
• Laws are different depending on order
• Overall reaction order is sum of exponents
• Rate k ? zero order
• Rate kA ? first order
• Rate kA2 ? second order
• Rate kAB ? second order

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First order integrated rate law

• Rearrange and use some calculus to get

• This is y mx b form
• A plot of lnA vs time will give a straight line
• If k and A0 (initial concentration) known, then
  you know the concentration at any time

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Second order integrated rate law

• Rearrange and use some calculus to get

• This is y mx b form
• A plot of 1/A vs time will give a straight line
• If k and A0 (initial concentration) known, then
  you can now the concentration at any time

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Zero order integrated rate law

• Rearrange and use some calculus to get

• This is y mx b form
• A plot of A vs time will give a straight line
• If k and A0 (initial concentration) known, then
  you can now the concentration at any time

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Graphs give order of rxn

• Use graphs to determine order
• If A vs time zero order
• If ln A vs time first order
• If 1/ A vs time second order

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Half-life

• Defn time it takes for concentration to halve
• Depends on order of rxn
• At t1/2 AA0/2

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Half-Life

• First order
• Second order
• Zero Order

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Reaction Mechanism

• Reactions occur by a series of steps
• Reaction mechanism
• Example
• Overall reaction NO2 CO ? NO CO2
• occurs by following steps
• Step 1
• Step 2

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Intermediates

• Two molecules of NO2 collide
• Oxygen is transferred, making NO3, the
  intermediate
• Intermediates are temporarily formed during a
  reaction
• They are neither a reactant nor a product
• Get used up in reaction

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Rules for Reaction Mechanisms

• Sum of elementary steps overall balanced rxn
• Mechanism must agree with experimental rate law

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Elementary Step

• Steps in reaction from which a rate law for step
  can be directly written
• 2 molecules of NO2 need to collide, therefore
• Rate k NO22

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Molecularity

• Rate law written based on molecularity
• Number of things that have to collide
• Unimolecular rxn depends on 1 molecule
• Bimolecular rxn depends on 2 molecules
• Termolecular rxn depends on 3 molecules
• Very rare!

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Give molecularity and rate law

• Unimolecular (first order)
• ratekA
• Bimolecular (second order)
• ratekAB

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Rate Determining Step

• The slowest step in mechanism determines overall
  rate
• Rate cannot be faster than slowest step
• Demo Filling bottle with funnel
• Overall rate law can be written from molecularity
  of slowest step

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How are mechanisms determined?

• Rate law is determined using experiment (method
  of initial rates, etc.)
• Chemist uses intuition to come up w/ various
  mechanisms
• Narrows down choices using rules for mechanisms
• No mechanism is ever absolutely proven