2. Advanced Chemistry Ms. Grobsky Enthalpies of Reactions and Hess’ Law

3. Recall that: Enthalpy is a measure of the heat that is released or absorbed in a chemical reaction at constant pressure You cannot measure the actual enthalpy of a substance You can measure an enthalpy CHANGE Written as the symbol , “delta H ” THERMODYNAMICS - ENTHALPY CHANGES

4. Why is there a standard? Enthalpy values vary according to the conditions A substance under these conditions is said to be in its standard state: Pressure1 atmosphere Temperature298K (25°C) Concentrationmol/L As a guide, just think of how a substance would be under normal lab conditions Assign the correct subscript [(g), (l) or (s) ] to indicate which state it is in To tell if standard conditions are used we modify the symbol for  : Enthalpy Change Standard Enthalpy Change (at 298K) STANDARD ENTHALPY CHANGES

5. STANDARD HEATS (ENTHALPIES) OF FORMATION Definition The enthalpy change when ONE MOLE of a compound is formed in its standard state from its elements in their standard states Symbol  ° f Values Usually, but not exclusively, exothermic Example(s) C(graphite) + O 2 (g) ———> CO 2 (g) H 2 (g) + ½O 2 (g) ———> H 2 O(l) 2C(graphite) + ½O 2 (g) + 3H 2 (g) ———> C 2 H 5 OH(l) Note: Only ONE MOLE of product on the right-hand side of the equation Elements In their standard states have zero enthalpy of formation

6. STANDARD HEATS (ENTHALPIES) OF FORMATION

7. Definition The enthalpy change when ONE MOLE of a substance undergoes complete combustion under standard conditions. All reactants and products are in their standard states Symbol  ° comb Values Always exothermic Example(s) C(graphite) + O 2 (g) ———> CO 2 (g) H 2 (g) + ½O 2 (g) ———> H 2 O(l) C 2 H 5 OH(l) + 3O 2 (g) ———> 2CO 2 (g) + 3H 2 O(l) Notes Always only ONE MOLE of what you are burning on the LHS of the equation To aid balancing the equation, remember: You get one carbon dioxide molecule for every carbon atom in the original and one water molecule for every two hydrogen atoms When you have done this, go back and balance the oxygen STANDARD HEATS (ENTHALPIES) OF COMBUSTION

8. Heat of Fusion The heat absorbed by ONE MOLE of a substance in melting from a solid to a liquid  ° fus If ONE MOLE of a substance releases heat in freezing from a liquid to a solid,  ° fus is negative! Heat of Neutralization We will soon find out in lab! Heat of Vaporization The amount of heat necessary to vaporize ONE MOLE of a given liquid Heat of Condensation The amount of heat released when ONE MOLE of a vapor condenses Heat of Solution Heat change caused by dissolution of ONE MOLE of substance OTHER HEATS (ENTHALPIES) OF REACTION

9. Theory Imagine that, during a reaction, all the bonds of reacting species are broken and the individual atoms join up again but in the form of products Breaking a bond REQUIRES energy ENDOTHERMIC Forming a bond RELEASES energy Exothermic The overall energy change will depend on the difference between the energy required to break the bonds and that released as bonds are made If energy released making bonds > energy used to break bonds EXOTHERMIC Products are lower in energy than the reactants If energy used to break bonds > energy released making bonds ENDOTHERMIC Products are greater in energy than reactants Relation of Enthalpy to Bonds

10. Turn to FRONT of page 165 Enthalpy change (  ) = Enthalpy of products - Enthalpy of reactants Enthalpy of reactants > products Enthalpy of reactants < products  = -  = + EXOTHERMIC Heat given out ENDOTHERMIC Heat absorbed REACTION CO-ORDINATE ENTHALPY REACTION CO-ORDINATE ENTHALPY THERMODYNAMICS - ENTHALPY CHANGES

11. Enthalpy and Hess’ Law

12. Enthalpy is a state function As such,  H for going from some initial state to some final state is pathway independent Hess’ Law  H for a process involving the transformation of reactants into products is not dependent on pathway Therefore, we can pick any pathway to calculate  H for a reaction

13. Hess’s Law 12 Start Finish Both lines accomplished the same result, they went from start to finish Net result = same

14. Campsite to Illustrate Altitude as a State Function

15. Hess’s Law Hess’s law of heat summation states that for a chemical equation that can be written as the sum of two or more steps, the enthalpy change for the overall equation is the sum of the enthalpy changes for the individual steps.

16. A → B + C  H = x B + C → D  H = y A → D  H = ?  H = x + y

17. “The enthalpy change is independent of the path taken” How The enthalpy change going from A to B can be found by adding the values of the enthalpy changes for the reactions A to X, X to Y and Y to B.  r =   +  2 +  3 HESS’S LAW

18. “The enthalpy change is independent of the path taken” How The enthalpy change going from A to B can be found by adding the values of the enthalpy changes for the reactions A to X, X to Y and Y to B.  r =   +  2 +  3 If you go in the opposite direction of an arrow, you subtract the value of the enthalpy change  2 = -   +  r -  3 The values of   and   3 have been subtracted because the route involves going in the opposite direction to their definition HESS’S LAW

19. “The enthalpy change is independent of the path taken” Use Applying Hess’s Law enables one to calculate enthalpy changes from other data, such as: Changes which cannot be measured directly Lattice Enthalpy Enthalpy change of reaction from bond enthalpy Enthalpy change of reaction from  ° c Enthalpy change of formation from  ° f HESS’S LAW

20. Enthalpy diagram illustrating Hess’s law

21. Hess’s Law For example, suppose you are given the following data: Could you use these data to obtain the enthalpy change for the following reaction?

22. Hess’s Law If we multiply the first equation by 2 and reverse the second equation, they will sum together to become the third.