Question

I have a doubt about using Gibbs free energy to predict the spontaneity of a reaction. It is shown that ∆G = ∆G° + RT ln (Q). That said, in order to predict which direction the reaction spontaneously goes, should ∆G or ∆G° be used? As I understand it, ∆G is what actually allows one to predict the spontaneity of the reaction, that is, whether the reaction proceeds spontaneously from the reactants to the products (in the direction written, ∆G < 0, exoergonic) or whether it proceeds spontaneously from the products to the reactants (in the opposite direction, ∆G > 0, endoergonic), and all of this depends on the value of Q, since ∆G° is constant. Instead, ∆G° should be used to understand how far the reaction is shifted to the right or left when it is at equilibrium (∆G = 0). However, very often I see that ∆G° is used to predict the spontaneity of a reaction, so I see a contradiction (especially in biochemistry, when it comes to reactions in which ATP is involved). P.S. again referring to Gibbs free energy: can it be said that a process of mixing two substances is (almost) always characterized by a ∆S > 0 and therefore that the spontaneity of the process depends essentially on ∆H? A process is called spontaneous when reacting in the forward direction would approach equilibrium (no matter whether a reaction in that direction is observed or not). For a chemical reaction, this depends on the activities (concentrations and/or partial pressures) of reactants and products. To figure out whether a reaction under a given set of conditions (activities, temperature, pressure) is spontaneous in the forward or reverse direction, you have to knowΔrG. Sometimes, people ask whether a certain reaction is spontaneous or not. The answer depends on the conditions, i.e. the activities of the reactants and products. If no set of conditions is given, you can answer that question by stating whether at standard conditions (i.e. Q = 1) and a temperature of choice the reaction is spontaneous in the forward or in the reverse direction. In that case, you would look atΔrG∘to determine which direction is spontaneous at standard conditions. When a reaction proceeds in the absence of work, it approaches equilibrium;ΔrGwill continuously increase until it reaches zero at equilibrium. You are correct in supposing thatΔrGis what is important. A criterion for a spontaneous reaction isΔrG<0and for equilibrium is whenΔrG=0so thatΔGo=−RTln(Ke). Only here isQ≡Ke. At equilibrium thechangeof free energy is zero, it must be so, and the entropy is at a maximum. Suppose thatΔGo=0then the mixture of reactants and products is such that all have unit activity and are at equilibrium, and thenKe=1as required byΔGo=−RTln(Ke). IfΔGo<0the mixture is not at equilibrium and this is now produced by reactants converting to products and the reaction proceeding to the right. AsΔGo<0thenKe>1which is just another way of saying that the reaction moves to the right. The largerKebecomes the further to the right the reaction reaches equilibrium. IfΔGo>0thenKe<1and the reaction does not proceed spontaneously and the equilibrium lies on the left. (Your p.s. Mixing always does increase the entropy which via−TΔScan have a big effect on the free energy but additionally reactions may increase or reduce the number of different gaseous species which also has a big effect on entropy changes.)

Answer 1

Predicting reaction spontaneity involves ΔG for actual conditions (with ΔG < 0 indicating spontaneity) and ΔG° for standard conditions (with ΔG° < 0 indicating product-favoring equilibrium). Mixing substances typically raises entropy, influencing ΔG and spontaneity, but ΔS must be considered alongside ΔH and temperature.

Step-by-step explanation:

To predict the spontaneity of a chemical reaction, both ΔG (Gibbs free energy change) and ΔG° (standard Gibbs free energy change) are used, but they serve different roles. The spontaneity under any set of conditions is determined by ΔG, where ΔG < 0 indicates a spontaneous reaction in the forward direction, and ΔG > 0 indicates spontaneity in the reverse direction. When ΔG = 0, the system is at equilibrium.

On the other hand, ΔG° provides insight into the spontaneity of a reaction at standard conditions, and its relationship with the equilibrium constant K. If ΔG° < 0, then K > 1, suggesting products are favored at equilibrium.

For mixing two substances, the process is generally characterized by an increase in entropy (ΔS > 0), which often contributes to spontaneity; however, it's the combined effect of ΔS and the enthalpy (ΔH) along with the temperature of the system that dictates whether the total Gibbs free energy (ΔG) will be negative for a spontaneous process.