1. Chemical Equilibrium

2.  So far we have been assuming that chemical reactions go to completion (all the reactants are used up).  Some reactions do go to completion, but most reach a state of equilibrium  These are called reversible reactions a chemical reaction in which the products reform the original reactants

3. Chemical Equilibrium  Chemical equilibrium – a state of balance in which the rate of a forward reaction equals the rate of the reverse reaction and the concentrations of products and reactants remain unchanged!  Example: Ca +2 (aq) + SO 4 -2 (aq) ↔ CaSO 4(s)

4. Chemical Equilibrium

5.  The theory that explains why chemical reactions occur is called collision theory.  Just like the moose hitting this car, molecules and atoms have to collide with each other to form new products.

6. Collision Theory – page 592 Not all collisions result in new products; sometimes molecules just bounce off of each other. To make new products, molecules must collide with enough energy to make a product. This energy is called activation energy.

7. Activation Energy  The activation energy is the minimum amount of energy needed to start a chemical reaction.

8. Collision Theory  Activation Energy the minimum energy required to start a chemical reaction.

9. Collision Theory  Reaction Pathways As two molecules approach each other, the outer electrons of each molecule repel For two molecules to react, they most collide violently enough to overcome this repulsion The molecules that are most likely to react must have especially high kinetic energy; other molecules must wait until collisions bring their kinetic energy up to the required amount

10. Activation Energy Diagrams  Also called potential energy diagrams  When molecules collide, they form an activation complex Definition: a molecule in an unstable state intermediate to the reactants and the products in the chemical reaction Example:  2HI → H 2 I 2 → H 2 + I 2 initial state activated final state complex (See attached graph)