Collisions between reactants are necessary for chemical reactions to occur. The reactants must collide with the proper orientation and with sufficient kinetic energy to overcome the activation energy barrier. Adding a catalyst can also increase the reaction rate by lowering the activation energy needed for collisions between reactants to be productive.

1. How do reactions occur?Collision TheoryCollision TheoryCollision TheoryIn order for a chemical reaction to take place, the reactants must collide.The collision transfers kinetic energy needed to break the necessary bonds so that new bonds can be formed.

2. Collision requirementsRequirement 1Must have the proper orientation.2HCl + Mg  MgCl2 + H2H---ClMgH-ClMgH Cl-MgWrong OrientationCorrect Orientation

3. Collision requirementsRequirement 2Must have enough kinetic energy to reach a threshold of energy called activation energyMgH---ClH---ClMgH Cl--Mg

4. Energy of Activation

5. Energy of Activation

6. Increasing the Rate of ReactionsWhat needs to happen in order for the rate of the chemical reaction to increase (go faster)?More collisions= Faster reaction rate

7. 5th way in increase Rxn RateAdd a Catalyst= Speeds up a reaction but is not used in the reactionLowers the activation energy