1. Enthalpy

2. Internal Energy Equation  ΔE = Q + W = Q + PΔV  If the reaction is carried out at a constant volume (ΔV = 0), then ΔE = Q  If volume is constant, any heat added or removed changes the internal energy

3. Constant Pressure  If pressure is constant, Q = ΔE + PΔV  Heat needed to bring about any change is the sum of internal energy plus P-V work  Constant pressure reactions are common in chemistry

4. Enthalpy  Sum of internal energy (E) and work (PV)  It is a state function  ΔH = Q = ΔE + PΔV  Flow of heat is equal to change in enthalpy  Enthalpy is called the heat of reaction  Since in many reactions the change in volume is small, ΔH is very often the same as ΔE

5. Enthalpies in Reaction  ΔH = H(products) – H(reactants)  Exothermic Reactions ΔH is negative 2H 2 (g) + O 2 (g) → 2H 2 O(g) + 484 kJ 2H 2 (g) + O 2 (g) → 2H 2 O(g) ΔH = -484 kJ

6. Endothermic Reaction  ΔH is positive  68 kJ + N 2 (g) + 2O 2 (g) → 2NO 2 (g)  N 2 (g) + 2O 2 (g) → 2NO 2 (g) ΔH = + 68 kJ

8. Enthalpy is Extensive  It depends on the amount  For the reaction below CH 4 (g) + O 2 (g) → CO 2 (g) + H 2 O(g) ΔH = - 802kJ how much heat is produced when 4.50 g of methane gas is burned?

9. Reverse Reactions  If a reaction is reversed its ΔH is numerically the same, but opposite in sign  CO 2 (g) + H 2 O(g) → CH 4 (g) + O 2 (g) ΔH = +802kJ

10. Different State Have Different ΔH  Enthalpy change is different for different states of matter of reactants and products  2H 2 (g) + O 2 (g) → 2H 2 O(g) ΔH = -484 kJ  2H 2 (g) + O 2 (g) → 2H 2 O(l) ΔH = -572kJ

11. Potential Energy Diagrams  Used to represent the general change in energy over the course of a reaction

14. Activation Energy  Minimum amount of energy required to initiate a chemical reaction