Understanding Energy and Disorder in Chemical Systems

1. Chapter 20 Energy and Disorder

2. Objectives • Understand the concepts of enthalpy, entropy, and free energy and the relationship among them • Use these concepts to solve problems dealing with any of them • Understand energy as it is related to chemical systems

3. Why Reactions Occur • Exothermic reactions generally take place spontaneously • Endothermic reactions generally do not take place spontaneously • Natural processes tend to go from an orderly state to a disorderly one • High energy  low energy • Order  disorder

4. Isothermal Processes – reactions taking place at constant temperature • Isobaric Processes – reactions taking place at a constant pressure • Thermodynamics – the study of the flow of energy

6. State Functions • A state function is one whose value depends only on the current state of the system • T = T2 – T1 • V = V2 – V1 • P = P2 – P1

7. Internal Energy • Every system has some internal energy • Internal energy, U, is a state function • Ways to transfer energy • By heating the system/surroundings • By doing work

8. U = q + w • q = heat gained by the system • w = the amount of work done on the system • Neither q nor w are state functions • q has a positive sign if heat is flowing in, negative if heat is flowing out • w is negative if work is done on the surroundings, and positive if done on the system

9. Assignment • 1-5 page 394 • Read through the rest of chapter 20 • Due: Tomorrow

10. Enthalpy • Enthalpy (H) = U + PV • Enthalpy is a state function • H = H2 – H1 • Exothermic Rxn: H < 0 • Endothermic Rxn: H > 0

11. Enthalpy Change

12. Which is Which?

13. What about this one?

14. Standard States • Standard state refers to the enthalpy substances have at 298.15 K and 100.000 kPa • Not the same as with the gas laws • In measuring enthalpy, set the enthalpy of any free elements to be equal to zero • A free element is one that is not in a compound

15. Enthalpy of Formation • Enthalpy of formation is the change in enthalpy when one mole of a compound is produced from free elements in their standard states • Units: KJ/mole • Symbol: Hfº • º means at standard state • Thermodynamic stability depends on the amount of energy that would be required to decompose the compound • See table A-6 in appendix for values • Thermodynamically stable compounds have large negative enthalpies of formation

16. Calculation of Enthalpy of Reaction • Hfº (products) = Hfº (reactants) + Hrº •  means summation • Hrº means change in enthalpy • If the enthalpy of formation of each product is known, you can calculate the amount of energy produced or absorbed, which then tells you if the reaction will be endothermic or exothermic • Assignment: Due at end of class • 6-7 page 397

17. Hess’s Law • Hess’s Law – The enthalpy change for a reaction is the sum of the enthalpy changes for a series of reactions that add up to the overall reaction

18. Consider reaction A  C • Break into two parts • (1)AB and (2)B  C • Hr(1) = Hf°B - Hf°A • Hr(2) = Hf°C - Hf°B • So, the enthalpy change for the overall change of A to C is • H° = Hr(1)° + Hr(2)°

19. Entropy • Entropy, S, is derived from the second law of thermodynamics. This law places limits on the conversion of heat into work and prohibits perpetual motion • Entropy (S) is a measure of disorder in a system • Entropy is a state function

20. Entropy cont. • Examples of • - • - • - • S > 0 = Increase in disorder • S < 0 = Decrease in disorder • Assignment: Due at end of class • 8-10 pg 400 • 10 points

21. Gibbs Free Energy • Gibbs free energy determines whether a reaction will occur or not • G = H - T S • If G < 0 the reaction is exergonic (spontaneous) • If G > 0 the reaction is endergonic • The reaction can only occur is T S is very large • If G = 0 the system is at equilibrium

22. Gibbs Free Energy Calculations • Appendix A-6 • Gr° =  Gf°(products) -  Gf°(reactants) • Assignment • Problems 11-14 pg 403-404

23. Chapter Review • Complete questions