CHEM 163 Chapter 20

1. CHEM 163Chapter 20 Spring 2009

2. 3-minute exercise Is each of the following a spontaneous change? • Water evaporates from a puddle • A small amount of sugar dissolves in hot tea • Methane burns in air • A hamburger becomes uncooked

3. Thermodynamics First Law: Law of Conservation of Energy Limitation: Internal E of a system Explains change, but not direction heat work Second Law: Systems change towards more disorder

4. Spontaneous Change • Change that occurs without continuous E input • Change can only be spontaneous in one direction, under a given set of conditions • Enthalpy (∆H): • heat gained or lost at constant P • Sign of ∆H • Exothermic or endothermic • No information about spontaneity • Entropy (∆S): • Freedom of particle motion (dispersed E of motion)

5. Energy Levels • Each atom or molecule has quantized E levels • Electronic • Kinetic • Vibrational, rotational, translational • Microstate: combined E at any given point • each microstate is equally possible (equal E) for a given set of conditions Number of microstates entropy Boltzmann constant (J/K) = R/NA = 1.38 x 10-23 J/K

6. Entropy Change: Microstates for 1 mol

7. Entropy Change: Heat Changes • Remove 1 grain of sand • Gas does work on piston • absorbs heat to maintain E • Works for tiny changes (totally reversible)

8. Always Increasing Entropy All real processes occur spontaneously in direction that increases the entropy of the universe. • Perfect crystal at T = 0 K has S = 0 • 1 microstate • Standard Molar Entropy (S°) • S increase from 0 to standard state • 1 atm (gases) • 1 M (solutions) • Pure substance, most stable form (liquids/solids)

9. What affects S°? As # of microstates (or kinetic E) increases, S increases • Temperature change ↑ T ↑ S • Phase change absorb heat ↑ S

10. Dissolution Ions: increased S, except small, highly charged ions Molecules (solid or liquid): ∆S ≈ 0 Gases: decreased S • Atomic Size • heavier atoms • allotropes  closer E levels  more microstates Molecular Complexity more complex  more types of movement  more microstates Only applies to molecules in same physical state

11. 3-minute practice What is the sign of ∆Ssys? • A pond freezes in winter • Atmospheric CO2 dissolves in the ocean • 2 K (s) + F2 (g)  2KF (s)

12. Standard Entropy of Reaction Increasing disorder: Decreasing disorder: Predict ∆Srxn: Change in # moles of gas Calculate ∆Srxn: N2 (g) + 3H2 (g)  2 NH3 (g)

13. ∆Suniverse Decrease in ∆Ssys only if greater increase ∆Ssurr System acts as heat sink or drain • Exothermic • Endothermic at constant P Measure ∆Hsys to determine ∆Ssurr

14. Spontaneous at 298 K? 3 2 Balance equation! Calculate ∆Ssys Calculate ∆Hsys Calculate ∆Ssurr Calculate ∆Suniv N2 (g) + H2 (g)  NH3 (g)

15. Entropy at Equilibrium Approaching equilibrium: At equilibrium: No net change

16. 5-minute Practice Calculate ∆Ssys for the combustion reaction of ammonia (producing nitrogen dioxide and water vapor).

17. Gibbs Free Energy • Measure of spontaneity • Combines enthalpy and entropy Spontaneous if… ∆Suniv > 0 ∆Gsys < 0 T∆Suniv > 0 -T∆Suniv < 0

18. ∆G Spontaneous process Nonspontaneous process Process at equilibrium Standard Free Energy Change Standard Free Energy of Formation : E change when 1 mol of compound is made from its elements in standard states Element in standard state:

19. ∆G and work (constant T & P) • Nonspontaneous process: • Process may occur if work is done to the system • How much work is needed? w = ∆G • Spontaneous process: • ∆G = maximum useful work done by the system • wmax only if process is totally reversible • Actually does less w <wmax • Extra E lost as heat

20. Useful Work • Excludes work done by or on atmosphere • Some free energy is always lost to heat ∆Hsys? < 0 ∆Ssys? > 0 ∆Gsys = wmax that can be done by system ∆Gsys? < 0 ∆Gsys > w actually done by system In some multistep reactions, ∆G from one reaction can cause an otherwise nonspontaneous reaction to occur. “coupling of reactions”

21. What about T? • Typically ∆H > T∆S • For ∆G to be negative, need ∆H to be… • What about at high T? negative • T∆S term can dominate • Sign of ∆S becomes important 4 situations: ∆H < 0 & ∆S > 0 ∆H > 0 & ∆S < 0 ∆H > 0 & ∆S > 0 ∆H < 0 & ∆S < 0 ∆G < 0 ∆G > 0 ∆G > 0 at low T; ∆G < 0 at high T ∆G < 0 at low T; ∆G > 0 at high T

22. High T v. Low T? Spontaneous “limit” at what temperature?

23. Equilibria and ∆G • If Q < K, reaction… • If Q > K, reaction… • If Q = K, at equilibrium → Q/K < 1 ∆G < 0 ← Q/K > 1 ∆G > 0 Q/K = 1 ∆G = 0 proportional 0 Make Q standard state (all values = 1)

24. Homework due MONDAY, May 11th Chap 20: #19, 26, 30, 41, 50, 52, 58, 70, 78, 106