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Question

If you calculate a value for ∆G° for a reaction using the values of in Appendix 4 and get a negative number, is it correct to say that the reaction is always spontaneous? Why or why not? Free energy changes also depend on concentration. For gases, how is G related to the pressure of the gas? What are standard pressures for gases and standard concentrations for solutes? How do you calculate ∆G for a reaction at nonstandard conditions? The equation to determine ∆G at nonstandard conditions has Q in it: What is Q? A reaction is spontaneous as long as ∆G is negative; that is, reactions always proceed as long as the products have a lower free energy than the reactants. What is so special about equilibrium? Why don’t reactions move away from equilibrium?

Solution Preview

For a spontaneous reaction, Δ G 0 is always negative.

The negative sign of Δ G 0 indicates that the reaction is favorable thermodynamically. But for determining whether a reaction is fast enough to be useful, temperature should be considered. Therefore, if a process must be carried out at high temperatures to be fast enough to be feasible, Δ G 0 must be recalculated at that temperature from the Δ H 0 and Δ S 0 values for the reaction.

Free energy changes depends upon pressure of the gas. The relationship between free energy changes and pressure of the gas is, Δ G = Δ G 0 + R T ln ( Q )

Where,

  • G 0 is the standard free energy change.
  • G is the free energy change of the gas.
  • P is the pressure of the gas.
  • T is the temperature.

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