1. Enthalpy Changes in Chemical Reactions

2. Enthalpy Enthalpy (H) The “heat content” of a substance It is the total KE and PE of a substance at constant pressure KE + PE = constant (enthalpy)

3. Kinetic & Potential Energy KE: it is the energy of MOTION… ex; atoms and electrons moving within the molecules, or in the entire system (reaction) PE: energy existing due to position and the sum of all attractions/repulsions between particles. PE is directly related to BOND energies.

4. Bond energy Bond energy is the amount of energy required to break/form a bond between two atoms. Ex:

5. Enthalpy changes Chemists interested in enthalpy changes (  H )  H = H products- H reactants

6. Enthalpy Vs Rxn Proceeding

7. Exothermic Reactions H 2 + S ---> H 2 S  H = - 20 KJ negative  H means exothermic H 2 + S ---> H 2 S +

8. Endothermic Reactions CH 3 OH  C(s) + 2H 2 (g) + ½ O 2 (g)  H = + 201 KJ positive  H means endothermic CH 3 OH +  C(s) + 2H 2 (g) + ½ O 2 (g)

9. Provincial Exam Questions

10. Kinetic Energy Distributions In general, molecules at R.T and pressure undergo about 10^10 collisions/second! Yet, there are only so many successful collisions…

11. KE Vs Number of Molecules

12. KE Distributions Increased Temperature = increased number of molecules that have enough energy to react/cross the activation energy barrier successfully. Yes, collisions would also increase, but only 1% more collisions for every 10C increase in Temp.

13. Activation Energy The minimum potential energy needed in a collision before a reaction can take place. It can also be defined as the minimum energy colliding particles must have in order to have a “successful” collision.

14. Activation Energy Graph

17. Homework for tomorrow! Try solving the questions on those pages p.12 p.16 p.19 and p.20