CH 11 Rate of Reaction /Kinetics

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CH 11Rate of Reaction /Kinetics

• http//www.youtube.com/watch?v_HA1se_gyvs

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(No Transcript)
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Reaction Rate

• Expresses how concentration (of a reactant or
  product) changes over time.
• In typical reactions, over time the reactants
  decrease, and products increase
• The reactants are used up/ the products are
  created.

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N2 3 H2 ? 2 NH3

• Using coefficients we can compare the relative
  rates formation and consumption.
• Nitrogen Ammonia
• Ammonia is formed at twice the rate Nitrogen gas
  is used up.
• Hydrogen Nitrogen
• Hydrogen is used up 3x the rate of Nitrogen

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Generic reactionaA bB ? cC dD
This allows us to relate the degradation of
reactants to the formation of products.
Express the rates of the reaction for nitrogen
hydrogen forming ammonia in terms of nitrogen.

• Rate -? A / a ? t - ? B / b ? t
• ? C / c ? t ? D / d ? t

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Collision Theory

• Reactions occur as a result of collisions between
  molecules.
• The molecules must have sufficient energy, and
  correct orientation to react.
• increased results in more frequent collisions
• Increased T results in higher energy collisions

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Reactions and Collision Theory

• To react molecules must
• Collide
• Have sufficient energy to overcome activation
  energy barrier. (Ea)
• Have correct orientation (stearic factor)

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Energy barrier

• The input energy required to start a reaction is
  called the activation energy
• Ea

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For reactions to occur

• The existing chemical bonds must break.
• New bonds must be formed.
• The original breaking of bonds requires energy

• The breaking is endothermic, the formation of new
  bonds can make the overall reaction exothermic.

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Fluorine Ion reacting w/ CH3Cl
As F- approaches, the C-Cl bond lenthens/stretches
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The energy required to reach the transition state
is called the Activation Energy Ea
Configuration of maximum energy reached
transition state highly unstable.
The energy decreases, as F-C is formed and C-Cl
is broken
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Boltzman Distribution of Kinetic Energies

• As T increases, more and more moelcules have a
  high Ek, . The proportion that with energy gt Ea
  increases significantly

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Rate expression N2 3 H2 ? 2 NH3

• Rate k N2xH2y
• The reaction rate is dependent upon
• N2
• H2
• k (a rate constant)
• the order of the reaction x and y

• The order of each reactant is determined
  experimentally.
• You cannot determine the order from looking at
  the coefficients.

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Determining reaction order

• Measure the initial rate (t 0) at 2 different
  concentrations.
• (either double or ½ the concentration)

• If doubling doubles the rate then the reaction
  is 1st order for that reactant
• If doubling the quadruples the rate, then the
  reaction is
• 2nd order for that reactant
• Add them together for overall order.

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Determining Reaction order

• Rate2 A 2
• Rate1 A 1
• The order can be determined from a ratio of the
  rates and concentrations

m
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Determine the reaction order for CH3CHO ? CH4
CO
Rate 2.0 M/s 0.50 M/s 0.08 M/s
CH3CHO 1.0 M 0.50 M 0.20 M
Use the equation from the last slide. You can use any two starting concentrations to determine the order. Use the equation from the last slide. You can use any two starting concentrations to determine the order. Use the equation from the last slide. You can use any two starting concentrations to determine the order. Use the equation from the last slide. You can use any two starting concentrations to determine the order.
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Reaction concentration and Time Using calculus

• Integrating the rate equations give us an
  equation that will allow us to determine the
  at any time, as long as we know the rate
  constant(k)

• Rate equations

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To relate concentration to time we must know the
order of the reaction

• Then choose the correct form of the integrated
  rate equation.

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• Rate equation
• Can tell us the initial rate of the reaction

• Integrated rate equation
• Can tell us the concentration at any time.

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Using the integrated rate equations 1st order
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T1/2 half life
of half lives passed Fraction remaining
1 .5000
2 .2500
3 .1250
4 .0625
5 .0312

• The amount of time required for 50 (one-half) of
  a sample to be used.

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• FIRST order integrated rate law.
• A 1
• A0 2
• -.693 /-k t1/2

• T ½ .693/k
• The decomposition of N2O5 is first order.
• _at_ 67C, k 0.35/min
• What is the half life?

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Half life and radioactivity
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Using the integrated rate equation 2nd Order
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• Zero order t1/2
• t1/2 Ao /2k

• 2nd Order t1/2
• t1/2 1/ kAo

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Energy diagram w/ and w/o catalyst.

• A catalyst allows for a new reaction
  mechanism with a lower Ea barrier

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Catalysts
Chem puzzler 1, 13 pg 2

• Increases the rate of reaction without being
  consumed.
• It changes the reaction path to a lower energy
  process.
• Heterogeneous-
• Reactants and catalyst are in a different phase.
  
• Gas reactant, solid catalyst
• Homogenous-
• Reactants and catalyst are in the same phase

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Reaction Mechanisms

• The path or sequence of steps by which a reaction
  occurs at the molecular level.
• Reaction order depends upon the mechanism

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Reaction Mechanism

• Broken down into a series of elementary steps
• of molecules involved
• Unimolecular (1)
• Bimolecular (2)
• Termolecular (3)

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Rate determining step
Rate limiter (like school Buses on warrior way)

• The slowest step is the one that determines the
  rate.
• The rate expression of the rate determining step
  is the rate expression for the entire reaction.

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Rate expression (multi-step mechanism)

• Focus on slow step
• Rate for slow step overall rate
• In rate determining step, coefficients give the
  order.
• Eliminate unstable intermediates from the rate
  expression.

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Reaction between X2 and A2

• X2 ? ? 2 X fast
• X A2 ? ? AX A slow
• A X2 ? ? AX X fast