 G 0 system =  H 0 system - T  S 0 system  G 0 rxn =  m  G 0 products -  n  G 0 reactants

Free Energy and Pressure  G =  G  + RT ln(Q) Q = reaction quotient from the law of mass action.

Free Energy and Equilibrium  G  =  RT ln(K) K = equilibrium constant This is so because  G = 0 and Q = K at equilibrium.

Free Energy, Equilibrium and Reaction Direction If Q/K < 1, then ln Q/K < 0; the reaction proceeds to the right (  G < 0) If Q/K > 1, then ln Q/K > 0; the reaction proceeds to the left (  G > 0) If Q/K = 1, then ln Q/K = 0; the reaction is at equilibrium (  G = 0)  G =  G o +RT ln Q  G 0 = - RT lnK

Figure 16.7: Schematic representations of balls rolling down two types of hills.

Figure 16.9: (a) The change in free energy to reach equilibrium, beginning with 1.0 mol A(g) at P A = 2.0 atm. (b) The change in free energy to reach equilibrium, beginning with 1.0 mol B(g) at P B = 2.0 atm. (c) The free energy profile for A(g) B(g) in a system containing 1.0 mol (A plus B) at P TOTAL = 2.0 atm. Each point on the curve corresponds to the total free energy of the system for a given combination of A and B.

Figure The relation between free energy and the extent of reaction  G 0 1  G 0 > 0 K <1 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

FORWARD REACTION REVERSE REACTION Table 20.2 The Relationship Between  G 0 and K at 25 0 C  G 0 (kJ) KSignificance x x x x x x10 1 6x10 8 3x x10 35 Essentially no forward reaction; reverse reaction goes to completion Forward and reverse reactions proceed to same extent Forward reaction goes to completion; essentially no reverse reaction

 G and the Work a System Can Do For a spontaneous process,  G is the maximum work obtainable from the system as the process takes place:  G = work max For a nonspontaneous process,  G is the maximum work that must be done to the system as the process takes place:  G = work max An example

The Law of Mass Action For jA + kB  lC + mD The law of mass action is represented by the equilibrium expression:

Figure 13.2: The changes in concentrations with time for the reaction H 2 O(g) + CO(g) H 2 (g) + CO 2 (g) when equimolar quantities of H 2 O(g) and CO(g) are mixed.

Figure 13.4: The changes with time in the rates of forward and reverse reactions for H2O(g) + CO(g)H2(g) + CO2(g) when equimolar quantities of H2O(g) and CO(g) are mixed. The rates do not change in the same way with time because the forward reaction has a much larger rate constant than the reverse reaction.

Figure 13.5: A concentration profile for the reaction N 2 (g) + 3H 2 (g)-->2NH 3 (g) when only N 2 (g) and H 2 (g) are mixed initially.

Figure 13.6: The position of the equilibrium CaCO 3 (s) CaO(s) + CO 2 (g) does not depend on the amounts of CaCO 3 (s) and CaO(s) present.

Figure 13.7: (a) A physical analogy illustrating the difference between thermodynamic and kinetic stabilities. (b) The reactants H2 and O2 have a strong tendency to form H2O.

Figure 13.8: (a) The initial equilibrium mixture of N2, H2, and NH3. (b) Addition of N2.(c) The new equilibrium position for the system containing more N2 (due to addition of N2), less H2, and more NH3 than in (a).