06523 Kinetics' Lecture 1 Introduction and review of basic concepts

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06523 Kinetics. Lecture 1Introduction and
review of basic concepts

• Dr John J. Birtill

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The importance of reaction kinetics

• Knowledge of reaction kinetics is essential for
  control of the reaction in the laboratory and in
  the chemical and biochemical industries
• to ensure that reaction goes to completion in a
  reasonable time
• to ensure that the reaction does not go out of
  control at any time
• to achieve optimal yield without waste or
  undesirable by-products.
• Also provides valuable insight when studying
  reaction mechanism.

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Chemical reaction and equilibrium

• All chemical reactions can move in a forward or
  reverse direction.
• Some reactions are very favourable and go forward
  virtually to completion
• with a negligible degree of reverse reaction.
• Some reactions proceed at a significant rate in
  both forward and reverse directions and reach a
  dynamic equilibrium.
• At equilibrium the rates of forward and reverse
  reactions are equal
• there is no net overall change taking place
• the overall Gibbs reaction energy ?rG 0
• Thermodynamics controls the position of
  equilibrium.

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Chemical reaction and equilibrium continued
Consider the simple reversible reaction in
solution
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Reaction kinetics

• Consider the same simple reaction in solution
• The concentration of A (CA) decreases with time
  and the concentration of B (CB) increases with
  time until equilibrium is reached.
• The reaction kinetics describes how fast the
  reaction proceeds.
• The gradients give the net rate of loss of A and
  the net rate of creation of B at that time
• different ways of expressing the overall reaction
  rate (mol s-1) .
• The reactions might go faster with a catalyst but
  the equilibrium point does not change.

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Kinetics of reversible reactions continued.

• Same simple reaction
• At equilibrium, reaction is still going on but
  the forward rate the reverse rate
• The net (overall) forward rate 0

equilibrium
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Reaction rate

• Rate of change of amount of a reactant or product
  ( mol s-1)
• For work at constant volume, can use intensive
  units such as rate of change of concentration (
  mol dm-3 s-1) or partial pressure ( bar s-1)
• Can define rate in terms of any reactant or
  product but need to adjust for the stoichiometry
• Example 1

Example 2
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Example Nitrobenzene hydrogenation
25 C
MeOH, Pd cat

• Can measure rate of consumption of H2 gas with a
  flowmeter ?VH2 /?t dm3 s-1
• PV nRT Hence
  mol s-1
• Reaction rate rate of reaction of
  nitrobenzene (main reactant) rNB rH2 /3 mol
  s-1
• If reaction vessel contains volume of liquid VL
  dm-3 , then rate of change of concentration of
  nitrobenzene nNB / VL mol dm-3 s-1

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Conversion and degree of reaction

• Fractional conversion ? is often used as a
  measure of the degree of reaction that has taken
  place.
• For the simple reaction insolution at constant
  volume
• if the starting concentration of A was CA0 mol
  dm-3
• and CA is the concentration at time t
• then at time t the conversion of A (CA0 - CA
  )/ CA0 .
• The percentage conversion 100 (CA0 - CA ) / CA0
  
• Conversion is usually defined in terms of the
  limiting reactant.

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Reaction rate laws and reaction order

• Reaction rates usually depend on the
  concentrations or pressures of the reactants.
• Consider the reaction
• The empirical reaction law may be found by
  experiment to be
• k is the reaction rate constant. Units depend on
  the overall reaction order.
• The reaction is of order a with respect to
  reactant A and order b with respect to reactant
  B. If a 1 or 2 then we say 1st order or 2nd
  order with respect to a
• The overall reaction order is a b.
• Reaction order is a convenient classification.
  Values can be negative. They are often not whole
  numbers (non-integral). Products can be involved
  as well.

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Rate laws continued

• The rate law and reaction order are empirical
  (determined by experiment). They cannot be
  predicted from the reaction stoichiometry.
• The rate law derives from the fundamental
  reaction mechanism.
• Consider two similar looking bromination
  reactions, both in aqueous solution

C6H5NH2 Br2 ? BrC6H4NH2 H Br- CH2(CN)2
Br2 ? BrCH(CN)2 H Br-

• Similar reaction stoichiometries but very
  different rate laws. Reaction 2 has a complex
  rate expression.
• What are the reaction orders in each case?

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Temperature dependence of rate constant
Arrhenius Law
Gradient - Ea/R

• A pre-exponential factor
• E activation energy /
• T absolute temperature / K

Data ex Cox, Modern liquid phase kinetics
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Reaction mechanism
Transition state

• Reaction mechanism describes the nature of the
  reaction route.
• Activated complex theory the reactants in a
  reaction step come together in a loose structure
  of higher energy. The maximum is called the
  transition state and difference from the ground
  state is the activation energy Ea of the reaction
  step.
• Note that the reverse step also has an activation
  energy, in this case higher than the forward
  step.
• The molecularity of a reaction step is the number
  of molecules coming together to react in that
  step
• not the same as reaction order.

Ea
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Problems to think about

• (1) Calculate conversion at end of the reactions
  below
• C6H5NH2 Br2 ? BrC6H4NH2 HBr
• Start 10 12 moles
• End 0.01 2.01 moles
• (2) If the reaction rate doubles for a rise in
  temperature from 230 ºC to 240 ºC what is the
  activation energy?