Enthalpy

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 • 2H2(g) + O2(g) → 2H2O(g) + 484 kJ • 2H2(g) + O2(g) → 2H2O(g) ΔH = -484 kJ

6. Endothermic Reaction • ΔH is positive • 68 kJ + N2(g) + 2O2(g) → 2NO2(g) • N2(g) + 2O2(g) → 2NO2(g) ΔH = + 68 kJ

8. Enthalpy is Extensive • It depends on the amount • For the reaction below CH4(g) + O2(g) → CO2(g) + H2O(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 • CO2(g) + H2O(g) → CH4(g) + O2(g) ΔH = +802kJ

10. Different State Have Different ΔH • Enthalpy change is different for different states of matter of reactants and products • 2H2(g) + O2(g) → 2H2O(g) ΔH = -484 kJ • 2H2(g) + O2(g) → 2H2O(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