Temperature Dependence Of Rate Coefficient

1
Temperature Dependence Of Rate Coefficient

• Rate coefficients have strong temperature
  dependence
• Arrhenius equation

2
Example

• For a 2nd-order reaction we can find (look up) A
  108 M1 s1 , Eact 100 kJ mol1.
• Find rate coefficients at 25 and 35º C.
• k A exp( Eact / R T )
• Exponent must be dimensionless, butunits not
  consistent

Eact 100 kJ mol1 100 ? 103 J mol1.
Factor 3.7!!
Similarly, k(35º C) 1.10 ? 109 M1 s1
3
Transition State or Activated Complex

• BrCH3 OH- ? Br - CH3OH

• TRANSITION STATE or ACTIVATED COMPLEX
• geometry through which system must pass between
  reactant and product
• ? higher energy than reactants or products
• not a stable molecule!

4
Fig 16.19 p 690
BrCH3 OH-
Br- CH3OH

• Forward reaction is EXOTHERMIC
• Reverse reaction is ENDOTHERMIC

5
Rationalization of Arrhenius Equation

• Consider a bimolecular collision A B ? products
• must go via transition state to react, therefore
  must have energy of at least Eact
• fraction of collisions with energy greater than
  Eact

• For gases, collision frequency ?1030 collisions
  s-1
• If each collision produced reaction, reaction
  rate ?106 M s-1
• Experimental gas phase reaction rates ?10-4 M s-1
• ? Only a very small fraction of collisions lead
  to reaction

6
The importance of molecular orientation
Fluorine
See also textbook Figure 16.17 page 689 F2 NO2 ?
FNO2F ?
Nitrogen
Oxygen
7
Example

• Assuming the activation energies are equal,
  predict which of the following reactions will
  occur at a higher rate at 50º C. Explain.
• NH3(g) HCl(g) ? NH4Cl(s)
• N(CH3)3(g) HCl(g) ? NH(CH3)3Cl(s)
• At the same temperature NH3 and N(CH3)3 have the
  same average kinetic energy
• Due to their larger mass, N(CH3)3 molecules will
  be travelling more slowly
• This will decrease the number of collisions per
  unit time for N(CH3)3 relative to NH3
• also the higher molecular complexity of N(CH3)3
  would probably reduce its orientation probability
  factor

8
Catalysis

• Catalysts increase rate coefficient by providing
  alternative reaction path (or mechanism) with
  lower activation energy.
• Catalyst does not change equilibrium, only rate
  of attaining equilibrium.
• Enormous practical importance all modern
  chemical products.
• Examples
• nitrogeneous fertilizers (N2 3H2 ? 2NH3 ) Pt
  catalyst
• removal of NO in vehicle exhaust Pd oxide
  catalyst
• cracking petroleum ? petrol, nylon,
• hydrogenating natural oils for margarine

9
Catalyst Lowers Barrier
10
Catalyst Can Significantly Change Mechanism
Fig 16.22 p 698
11
Heterogeneous Catalysis

• Homogeneous reactants and catalyst in same phase
• Heterogeneous reactants and catalyst in
  different phases
• eg H2 ? H H is rate limiting step in the
  synthesis of ammonia N2 3H2 ? 2NH3 -
  catalyzed by Pt
• HH bond too strong to react with N2 Pt has
  dangling bond forms weak bond with Hsweakens
  HH, now H is more reactive

12
Summary
Series of plots of concentration vs time
Fig 16.12 p 686
Initial Rates
Reaction Orders
Rate Constant k and actual rate law
Integrated rate law (reaction order half-life)
Activation Energy Ea
13
Reaction Mechanisms
Determine Rate Law by Experiment