Chemical Kinetics II

Collision Theory

Collision Theory attempts to explain reaction rates.  It is based on the basic ideas:

• Molecules must collide to react.
• Concentration affects the rate.  A higher concentration allows for more particles to have a chance of colliding, therefore increasing the rate.  Fewer particles would lessen the chance of collision, therefore slowing the rate.
• Molecules must collide hard enough (but not too hard) to cause a reaction.
• Temperature and rate are related.
• Only a small number of collisions actually result in a reaction.

Activation Energy is the energy needed to form an activated complex or transition state so a reaction can occur.  In other words, how much energy is needed to get the reaction started.  Think of it as the amount of energy it takes to cut up all the vegetables you need to prepare a meal.

Take the synthesis reaction that forms hydrogen iodide gas.

Before the hydrogen and iodine can combine, the diatomic molecules must be broken apart in a redox reaction.  After the positive hydrogen ions and the iodine ions have been formed, the electrostatic force (attraction of a + for a -) will pull the ions together.

This reaction needs energy to break up the diatomic molecules and form ions.  This is the activation energy.  The H+ and I- are the activated complex (the transition state or intermediate step) in the mechanism of the reaction.