CHEM 2 - Kinetics

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  • Created by: Franklin
  • Created on: 07-04-14 15:37

Collision Theory

Kinetics is the study of the factors which affect the rate of a chemical reaction

  • For a reaction to take place between two particles they must collide with sufficient energy to break bonds
  •  These reactions happen at certain parts of the molecule which means the orientation of the molecule also has to be correct.

Factors affecting rate of reaction

  • Change in temperature - particles have more kinetic energy
  • Increase in concentration of solution - more particles in a given volume at any given time leads to more collisions
  • Increase in pressure (gas) - more particles in a given volume
  • Increase in surface area of solid reactants - Greater total surface area means more particles are available to collide. More sites for reaction 
  • Use a catalyst - provides an alternative pathway for which there is a lower activation energy. Doesn't get used up during the course of the reaction
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Collision Theory 2

Activation Energy, Ea - The minimum energy that a particle needs in order to react.    

  • At the transition state bonds are being both formed and broken. 
  • In an exothermic reaction energy of products is < reactants


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Maxwell-Boltzmann Distribution

Maxwell-Boltzmann Distribution shows the distribution of energy amongst particles

  • Peak indicates most probable energy (Emp)
  • Slight to the right of this is the mean energy

The graph shows that:

  • No particles have zero energy
  • Most particles have intermediate energies (around peak)

The line Ea inducates the activation energy and the area below the curve beyond this line indicates the no. of particles with E>Ea

  • If the temperature is increased the peak (Emp) moves to the right and down but area underneath curve should still be the same. Area beyond Ea line will be greater.
  • Catalysts provide an alternative pathway for which there is a lower Ea (smaller amount of energy needed to start reaction). With cataylst Ea line moves back. 
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With a catalyst the transition state for a chemical reaction is lower.

  • Catalysts do not affect the enthalpy change of reactions or position of equilibrium

Heterogenous catalyst - The catalyst is in a different phase (state) to the reactants. Usually solid catalyst and liquid/ gas reactants

Homogeneous catalyst - The catalyst is the same phase as the reactants 

Catalysts for well-known reaction

  • Haber process - Iron
  • Ostwald process for making HNO3 - Platinum
  • Cracking hydrocarbons - Aluminium oxide/ Silicon oxide/ zeolite

Catalytic Converters

  • Honeycomb shape - provides a large surface area for reaction to take place
  • Coated in Rhodium/Platinum (catalysts)
  • Reactions: CO + NOx ---> N2 + CO2
  • Metal forms weak bonds with reactants to hold in place - ADSORPTION
  • Products then break away from metal - DESORPTION
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