1. Potential Energy Main Concept:
Potential energy is associated with a particular geometric arrangement of atoms or ions and the electrostatic interactions between them.

2. Potential Energy Energy: Attraction vs Repulsion
Potential Energy Graph
Bond Strength

3. - attraction between electrons of one atom and protons of another explains tendency for atoms to approach one another.
- repulsion between the nuclei (or core electrons) explains why atoms repel one another at close distance
- distance at which the energy of interaction is minimized is called the bond length, and atoms vibrate about this minimum energy position

4. Potential Energy Simulation(

5. - A graph of energy versus the distance between atoms can identify bond length and bond energy
- bond making and bond breaking are opposing processes that have the same magnitude of energy associated with them.
- bond energy - the energy required to break a bond.

7. - Because chemical bonding arises from electrostatic interaction between electrons and nuclei, larger charges tend to lead to larger strengths of interaction. Ex: triple bonds are stronger than double or
single bonds because they share more pairs
of electrons.
- Stronger bonds tend to be shorter bonds. 

8. Question:
What is the ΔH for the reaction below given the average bond energies below?
Bond: ∆H (kJ/mol):
CΞC
C–C
H–I
C–I
C–H
Answer:
( ) +
–( )
-217 kJ

9. Question:
What is the ΔH for the reaction below given the average bond energies below?
Bond: ∆H (kJ/mol): H–CΞC–H + H–I  H2C=CHI
CΞC
C=C
H–I
C–I
C–H
Answer:
( ) +
–( )
-129 kJ

10. Net Energy Change Main Concept:
The net energy change during a reaction is the sum of the energy required to break the bonds in the reactant molecules and the energy released in forming the bonds of the product molecules. The net change in energy may be positive for endothermic reactions where energy is required, or negative for exothermic reactions where energy is released.

11. Net Energy Change Net Energy Change: Bonds Forming vs Bonds Breaking
Exothermic Reactions vs Endothermic Reactions

12. - During a chemical reaction, bonds are broken and/or formed  changes potential energy of reaction system

13. - average energy required to break all bonds in reactant molecules = total average bond energies/bond enthalpies for all bonds in reactant molecules
- average energy released in forming bonds in products = total average bond energies/bond enthalpies for all bonds in product molecules
- If energy released > energy required, reaction  exothermic
- If energy required > energy released, reaction  endothermic.

15. - In exothermic reactions, products have lower potential energy compared with reactants
- In endothermic reactions, products have higher potential energy than reactants

16. - In isolated systems, energy is conserved
- if potential energy of products is lower than reactants, then kinetic energy of products must be higher
- In exothermic reactions, products have higher kinetic energy  higher temperature.
- In endothermic reactions, products have lower kinetic energy  lower temperature.

17. - Because products have higher or lower temperatures than surroundings, products move toward thermal equilibrium with surroundings
- Thermal energy is transferred to surroundings from hot products in exothermic reactions
- Thermal energy is transferred from the surroundings to cold products in endothermic reactions

18. - Hess’s law ideas:
- When reaction is reversed, sign of enthalpy of
reaction is changed
- when two (or more) reactions are summed to
obtain an overall reaction, enthalpies of reaction
are summed to obtain net enthalpy of reaction
- Tables of standard enthalpies of formation can
be used to calculate standard enthalpy of
reactions