Thursday, December 29, 2011

Collision Theory and Transition State Theory

1. Collision theory
    Collision theory is the first theory proposed to explain the rate of chemical reactions involving gaseous reactants.
    According to this theory, not all collisions between the molecules result in the formation of products. Effective collisions between molecules, which result in the formation of products, only occur when the following two conditions are fulfilled:
    (a) the colliding molecules possess a minimum energy, termed the activation energy, to initiate the chemical reaction.
    (b) the reactant molecules collide at the correct orientation.
   

    It is important to note that effective collision only results when two reactants molecules having activation energy collide at the correct orientation. Chemical reaction does not occur if only one of these conditions is fulfilled.

2. Activation energy
    Activation energy (Ea) is the minimum energy required to initiate a chemical reaction. It is a constant for a particular reaction.
   
    It appears as a potential energy 'hill' or barrier between the reactants and products. Only colliding molecules that are properly oriented can deliver kinetic energy into potential energy as least as large as Ea to be able to climb over the 'hill' and produce products.

3. Transition state theory
    (a) Transition state theory is used to explain in detail what happens when reactant molecules come together in a collision.
    (b) A collision between reactant molecules may or may not result in a successful reaction.
    (c) The outcome depends on the factors such as the relative kinetic energy, relative orientation and internal energy of the molecules. Even if the collisions partners form an activated complex, they are not bound to go on and form products, but instead the complex may fall back to the reactants.
   
    During the collision, there is a moment when the reactant bond is partially broken and the new product bond is partially formed. This brief moment during a successful collision is called the reaction's transition state, and the unstable chemical species that exists at this instant with its partially broken bonds in called the activated complex.

    (e) The potential energy diagram (or energy profile diagram) is used to visualise the relationship between activation energy and the development of total potential energy.
   
    The size of the activation energy tells us the relative importance of bond breaking and bond making during the formation of the activated complex. For example, a very high activation energy may suggest that the activated complex involves a significant amount of bond breaking as bond breaking is an energy-absorbing process.

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