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  1. 1. Thermochemistry: Enthalpy Lecture 16
  2. 2. Most chemical and physical changes occur at virtually constant atmospheric pressure.
  3. 3. Important types of chemical work: <ul><li>Electrical work – the work done by moving charged particles. </li></ul><ul><li>PV work – the work done by an expanding gas. </li></ul>
  4. 4. PV work can be found by multiplying the external pressure (P) by the change in the volume of the gas (ΔV = V final - V initial ).
  5. 5. Enthalpy H is a thermodynamic variable for reactions occurring at constant pressure: H= E + PV
  6. 6. The change in enthalpy ΔH is a the change in internal energy ΔE plus the product of the constant pressure P and the change in volume ΔV: ΔH= ΔE + PΔV
  7. 7. Let us derive: <ul><li>Recall: ΔE = q + w </li></ul><ul><li>Recall: w = -PΔV </li></ul><ul><li>ΔE = q + (-PΔV) = q - PΔV </li></ul><ul><li>At constant pressure, we denote q = q P </li></ul><ul><li>ΔE = q – PΔV, q P = ΔE + PΔV </li></ul><ul><li>Recall: ΔH = ΔE + PΔV </li></ul><ul><li>ΔH = q P </li></ul>
  8. 8. The change in enthalpy equals the heat gained or lost at constant pressure.
  9. 9. Cases with ΔE and ΔH: <ul><li>Reactions that do not involve gases. 2KOH (aq) + H 2 SO 4(aq)  K 2 SO 4(aq) + 2H 2 O (l) </li></ul><ul><li>ΔV≈ 0 and ΔH ≈ ΔE </li></ul><ul><li>Reactions in which the amount of gas does not change. </li></ul><ul><li>N 2(g) + O 2(g)  2NO (g) </li></ul><ul><li>ΔV= 0, so PΔV = 0 and ΔH = ΔE </li></ul>
  10. 10. Cases with ΔE and ΔH: <ul><li>Reactions in which the amount of gas does change. </li></ul><ul><li>2H 2(g) + O 2(g)  2H 2 O (g) </li></ul><ul><li>PΔV ≠ 0 </li></ul><ul><li>q P is so much larger than PΔV that ΔH is very close to ΔE </li></ul><ul><li>The key point: for many reactions, ΔH equals, or is very close, to ΔE. </li></ul>
  11. 11. E, P, and V are state functions. H is also a state function. ΔH depends only on the difference between H final and H initial .
  12. 12. The enthalpy change of a reaction Is also called the heat of reaction ΔH rxn and always refers to H final minus H initial: ΔH = H final – H initial = H products - H reactants
  13. 13. We determine the sign of ΔH by imagining the heat as a “reactant” or “product”.
  14. 14. An exothermic (“heat out”) process releases heat and results in a decrease in the enthalpy of the system. Exothermic: H final < H initial , ΔH < 0
  15. 15. An endothermic (“heat in”) process absorbs heat and results in an increase in the enthalpy of the system. Endothermic: H final > H initial , ΔH > 0
  16. 16. Endo- and exothermic processes
  17. 17. A sample problem on drawing enthalpy diagrams and determining the sign of ΔH.
  18. 18. In general, the value of an enthalpy change refers to reactants and products at the same temperature.
  19. 19. Important types of enthalpy change: <ul><li>Heat of combustion (ΔH comb ) </li></ul><ul><li>Heat of formation (ΔH f ) </li></ul><ul><li>Heat of fusion (ΔH fus ) </li></ul><ul><li>Heat of vaporization (ΔH vaporization ) </li></ul>
  20. 20. Enthalpy change
  21. 21. Enthalpy change
  22. 22. Enthalpy change
  23. 23. Heat of fusion for gallium is low:
  24. 24. THE END