1. 6.3

3.  Why does the rate of a reaction increase with Increased concentration of reactants Increased temperature Increased surface area

4.  In order for a reaction to occur, reacting particles (atoms, molecule, or ions) must collide with one another.  If collision is necessary for a reaction to occur, then it makes sense that the rate of the reaction will increase if there are more collisions per unit time.

5.  Increasing surface area of a solid-phase reactant speeds up a reaction.  With greater surface area, more collisions can occur. Starting a fire: small twigs rather than logs.

6.  Not every collision between reactants results in a reaction.  Collision must be EFFECTIVE: One that results in the formation of products.

7.  Collision geometry: the correct orientation of reactants relative to one another.

8.  Reactants must collide with energy that is sufficient to begin to break the bonds in the reactants and to begin to form the bonds in the products.  In most reactions, only a small fraction of collisions have sufficient energy for a reaction to occur.  ACTIVATION ENERGY, E a : the minimum collision energy that is required for a successful reaction.

9. Graph shows the distribution of kinetic energy in a sample of reacting gases at two different temperature s, T1 and T2, where T2>T1. Important Observations 1)At both temperatures, a relatively small fraction of collisions have sufficient kinetic energy – the activation energy – to result in a reaction. 2)As the temperature of a sample increases, the fraction of collisions with sufficient energy to cause a reaction increases significantly.

10.  Examines the transition, or change, from reactants to products.  Kinetic energy of reactants is transferred to potential energy as the reactants collide: law of conservation of energy. Basketball analogy: kinetic energy of ball converted to potential energy, which is stored in the deformed ball as it hits the floor. The potential energy is converted to kinetic energy as the ball bounces away.

11.  Diagram that charts the potential energy of a reaction against the progress of the reaction.

12.  The ‘hill’: Ea barrier.  Notice the difference between exothermic and endothermic reactions.  There is no way to predict the activation energy of a reaction from its enthalpy change: a highly exothermic reaction may be very slow because of a high activation energy.  Activation energy is determined by analyzing the reaction rate.

13.  For an exothermic reaction, the activation energy of the reverse reaction, E a(rev) equals E a(fwd) + ∆H.  For an endothermic reaction, E a(rev) equals E a(fwd) - ∆H.

14.  `Change-over`point  Activated Complex: chemical species that exists at the transition state neither product nor reactant Partial bonds  highly unstable.  Activated complex: can either break down to form products or decompose to re-form the reactants. Like a rock teetering on top of a mountain  can fall either way.

17.  PPs, page 294, #13-16  Read page 295 and make brief notes  Section Review, page 296, #1- 8.