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Chapter 17

Chapter 17. Spontaneity of Reaction Thermodynamic functions of enthalpy, entropy, and Gibbs free energy. Chapter 17. Spontaneous reactions: examples - CH 4 (g) + 2O 2 (g) --> CO 2 (g) + 2 H 2 O(l) H 2 O(s) --> H 2 O(l) at 25 o C. Spontaneity Factors.

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Chapter 17

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  1. Chapter 17 • Spontaneity of Reaction • Thermodynamic functions of enthalpy, entropy, and Gibbs free energy.

  2. Chapter 17 • Spontaneous reactions: examples - • CH4(g) + 2O2(g) --> CO2(g) + 2 H2O(l) • H2O(s) --> H2O(l) at 25oC.

  3. Spontaneity Factors • 1.) Energy factor: at 25oC, 1 atm, exothermic reactions are ordinarily spontaneous (nature tends toward stability)

  4. Spontaneity Factors • 2.) Randomness factor: other things being equal, systems tend move from a more ordered to a more random state.

  5. Spontaneity Factors • Nature tends toward disorder (the entropy of the universe is increasing).

  6. Entropy Changes • S = S products - S reactants; measure of change in order (randomness) • Ex. Solid --> liquid; S = positive

  7. Entropy Changes • S is usually positive for a reaction in which the number of moles of gas increases.

  8. Entropy Changes Determine if the entropy of the following reactions will be positive or negative. 2 SO3(g)-->2 SO2(g) + O2(g) N2(g) + 3 H2(g) --> 2 NH3(g) first reaction = positive, second = negative.

  9. Entropy Changes • So calculations (S at 1atm, 25oC - see chart). • So = So products - So reactants. • S is positive for elements and compounds.

  10. Entropy Changes • Reactions/processes for which So is positive tend to be spontaneous, at least at high temperatures

  11. Entropy Changes • H2O(s)-->H2O(l),So > 0 • H2O(l)-->H2O(g),So > 0 • Fe2O3(s)+3H2(g)--> 2Fe(s) + 3 H2O(g), So > 0. • All are endothermic, but become spontaneous at high temps.

  12. Free Energy Changes • Go = represents the useful energy available for work. Negative = energy available from the reaction to do work. • Positive = energy needed to get the reaction to occur.

  13. Free Energy Changes • Go =is a state function and is dependent only on the initial states of the reactants and not the pathway of the reaction.

  14. Free Energy Changes • Go = o - TSo • Note that G, like S, is dependent on pressure and concentration. Unlike Ho and So, Go is strongly temperature dependent.

  15. Free Energy Changes • If Go < 0, reaction is spontaneous at standard conditions (1 M conc., 1 atm P). • If Go > 0, reaction is nonspontaneous at standard conditions.

  16. Free Energy Changes • If Go = 0, reaction is at equilibrium at standard conditions.

  17. Free Energy Changes • Effect of Ho and So on spontaneity: • If Ho > 0, So < 0, and Go > 0 at all temps? • If Ho < 0, So > 0, and Go < 0 at all temps?

  18. Free Energy Changes • If Ho > 0, So > 0, and Go > 0 at low temps? High temps? • If Ho < 0, So < 0, and Go < 0 at low temps? High temps?

  19. Free Energy Changes Calculation of Go from So and Ho. • Fe2O3(s)+3H2(g)--> 2Fe(s) + 3 H2O(g). What is Go at 25oC? 500oC? 56.6kJ, -10.6 kJ.

  20. Free Energy Changes At what temp does the reduction of iron (III) oxide by hydrogen become spontaneous at 1 atm? T = 96.8/0.1387 = 698 K.

  21. Free Energy Changes • Calculation Go at 25oC from, Gof: • Go = ( Gof products) - (  Gof reactants).

  22. Free Energy Changes • Calculation G from Go: • G= Go+ RT ln Q. • R = 0.00831 kJ/K

  23. Free Energy Changes Given that Golead (II) chloride ionizing in solution is 27.3 kJ and that the concentration of Pb2+ and Cl- are each 0.0010 M,calculate= G of this reaction (at 25oC). -24.0 kJ

  24. Free Energy Changes • Relationship of G and K (Keq). • Go= -RT ln K. • If K > 1, Go< 0, reaction is spontaneous at standard conditions.

  25. Free Energy Changes • Relationship of G and K (Keq). • Go= -RT ln K. • If K < 1, Go> 0, reaction is nonspontaneous at standard conditions.

  26. Free Energy Changes • Relationship of G and K (Keq). • Go= -RT ln K. • If K = 1, Go= 0, reaction is at equilibrium at standard conditions.

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