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According to Wikipedia: "A chemical reaction is endergonic when non spontaneous. Thus in this type of reaction the Gibbs free energy increases."

My problem with that statement is the wording of "non-spontaneous" reactions. According to my knowledge, for every chemical reaction there is a ΔG, be it positive or negative, and is conncected to its equilibrium constant:

enter image description here

So in essence, if we only start with reactants, every reaction should yield some products, even if its ΔG is >0. Basically, every reaction is spontaneous to some extent. So is it only about semantics?

For example, in the biological context, even "non-spontaneous" reactions (with ΔG>0) can be driven to completion if its products are coupled to an exergonic reaction that consumes them and removes them from the endergonic equilibrium of the first reaction. That would mean that under those cirumstances the "non-spontaneous" reaction still produced some products (when starting with only reactants), is that right?

(What I also read though is that when stoechiometric molar quantities of reactants and products are mixed, more reactants are formed when ΔG is positive, so the reaction goes "in the wrong direction".)

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    $\begingroup$ It's only a "rule of thumb" for neophytes. But it's taken on a life of its own from repetition by irresponsible authors. You are entirely right in your assessment; all reactions will proceed to some extent. $\endgroup$ Oct 14, 2023 at 11:30
  • $\begingroup$ You might be confusing $\Delta G$ and $\Delta G^\circ$ $\endgroup$
    – Buck Thorn
    Oct 16, 2023 at 6:45
  • $\begingroup$ See chemistry.stackexchange.com/questions/133585/… and many many more.... $\endgroup$
    – Buck Thorn
    Oct 16, 2023 at 6:46
  • $\begingroup$ Taking a thermodynamic approach to whether a reaction is spontaneous needs to be compared to an approach that considers kinetics. In normal english usage "spontaneous" means the reaction will happen if the components are mixed. But many chemicals do not. Boranes catch fire when mixed with air; hydrogen and oxygen mixtures do not, though both reactions are exothermic. Kinetic barriers prevent many reactions from happening without a kick to push them over the activation energy barrier. $\endgroup$
    – matt_black
    Oct 17, 2023 at 18:23
  • $\begingroup$ See also chemistry.stackexchange.com/questions/43258/… and chemistry.stackexchange.com/questions/47262/… and chemistry.stackexchange.com/questions/31047/…, and, as Buck Thorn says, many others. $\endgroup$
    – Curt F.
    Oct 18, 2023 at 14:34

3 Answers 3

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"Spontaneous" is a confusing term, especially because different sources use it differently (either it makes a statement about the standard Gibbs energy of reaction, or the Gibbs energy of reaction with respect to the current state).

The term "spontaneous reaction" has been around since about 1910:

enter image description here

In John Parker's 1891 Elementary Thermodynamics, the term "spontaneous" is used in the following way:

enter image description here

(found via this page and accessed here)

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For evaluation of spontaneity, two possible approaches can be chosen:

  • A reaction is spontaneous, if $\Delta_\mathrm{r} G^\circ = -RT \ln{K} \lt 0$.

    • This is to be interpreted as a system with all components in their standard state with unit thermodynamic activities undergoes spontaneously the net reaction in forward direction.
    • This is the evaluation of spontaneity of reaction regardless of the initial system state, as the implicit presumption.
  • A reaction is spontaneous, if $\Delta_\mathrm{r} G = \Delta_\mathrm{r} G^\circ + RT \ln{Q} = RT \ln{\left(\frac{Q}{K}\right)}\lt 0$.

    • This is to be interpreted as a system with components in general states and component activities undergoes spontaneously the net reaction in forward direction.
    • This is the evaluation of spontaneity of reaction considering the initial system state.
    • If $Q \lt K$ then the equilibrium is in the forward direction of the reaction and the reaction is spontaneous in the latter sense regardless of if it is spontaneous in the former sense.
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At equilibrium the forward and reverse reactions have equal rates and the activities [approximated by concentrations] satisfy the equilibrium constant and the cell potential when the reaction is done as an electrical cell equals zero volts. If the activities are changed from the equilibrium condition in EITHER direction by a change in chemical activities, temperature, or applied voltage, the reaction will be spontaneous to reestablish equilibrium. This is approaching the situation from an active equilibrium, the equilibrium concentrations, appropriate temperature, mixing and provided there is a possible mechanism and catalysis.

A reaction: Ag+ + Cl- = AgCl or AgCl = Ag+ + Cl-. When approached from the first equation adding Ag+ to a solution of Cl- will produce a precipitate of AgCl as soon as the equilibrium constant is satisfied. This will be equilibrium and does not imply completion of the reaction. The reaction is driven by the enthalpy change in forming the crystal; adding more chloride or silver ion will increase the amount of AgCl and give a new equilibrium position. Removing either chloride or silver ion will reverse the reaction; the reaction will become spontaneous in the reverse direction. This is why ammonia dissolves AgCl, it removes silver ion.

When AgCl is added to pure water there are originally no silver or chloride ions. The second equation is now spontaneous and is entropy driven until the equilibrium concentrations are established. The interplay between enthalpy and entropy can be complicated. At equilibrium for a simple process one direction is enthalpy driven the other entropy driven; at equilibrium they balance: Delta H = TDelta S.

Stoichiometry is WHAT HAPPENS not necessarily what is present. Changing activities may also change the chemical reactions taking place. This does happen in the AgCl reaction because higher concentrations involve complex formation.

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