Not all chemical reactions proceed at the same rate. Some reactions are extremely slow whereas other reactions are extremely fast. Why is this?

Likely you are familiar with the "Collison-Reaction Theory". This theory is used to explain chemical reactions and can be summarized as follows:

• The atoms, ions, or molecules of a chemical sample are in continuous, random motion at various speeds. This results in collisions between the entities.
  
  
• If the reactant entities (atoms, ions, or molecules) collide with sufficient energy and at the correct orientation, a chemical reaction occurs.
  
  
   

  
  

  Reactant entities (atoms, ions, or molecules) must possess a minimum amount of energy before they can collide with sufficient energy to form products. This minimum amount of energy, known as the "activation energy," functions as an energy barrier and must be added to the system before a reaction can be initiated. Not surprisingly, slow-moving reactions have larger activation energies than fast-moving reactions.
  
  As you may have suspected, energy to initiate a reaction can come from a spark, the flame from a match, or, in the case of a spontaneous combustion, from the heat of a hot day. Watch the video below to increase your understanding of this important concept.
  

  
  

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