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In High School I learned that an exothermic reactions releases energy, while an endothermic reaction needs energy to occur. Now I learned that there is a separate, somewhat similar classification scheme of exergonic and endergonic reactions.

What is the difference between these two classification schemes? Are exothermic reactions always exergonic, and if not, can you give me an example?

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The classifications endothermic and exothermic refer to transfer of heat $q$ or changes in enthalpy $\Delta_\mathrm{R} H$. The classifications endergonic and exergonic refer to changes in free energy (usually the Gibbs Free Energy) $\Delta_\mathrm{R} G$.

If reactions are characterized and balanced by solely by heat transfer (or change in enthalpy), then you're going to use reaction enthalpy $\Delta{}_{\mathrm{R}}H$.

Then there are three cases to distinguish:

  1. $\Delta{}_{\mathrm{R}}H < 0$, an exothermic reaction that releases heat to the surroundings (temperature increases)
  2. $\Delta{}_{\mathrm{R}}H = 0$, no net exchange of heat
  3. $\Delta{}_{\mathrm{R}}H > 0$, an endothermic reaction that absorbs heat from the surroundings (temperature decreases)

In 1876, Thomson and Berthelot described this driving force in a principle regarding affinities of reactions. According to them, only exothermic reactions were possible.

Yet how would you explain, for example, wet cloths being suspended on a cloth-line -- dry, even during cold winter? Thanks to works by von Helmholtz, van't Hoff, Boltzmann (and others) we may do. Entropy $S$, depending on the number of accessible realisations of the reactants ("describing the degree of order") necessarily is to be taken into account, too.

These two contribute to the maximum work a reaction may produce, described by the Gibbs free energy $G$. This is of particular importance considering reactions with gases, because the number of accessible realisations of the reactants ("degree or order") may change ($\Delta_\mathrm{R} S$ may be large). For a given reaction, the change in reaction Gibbs free energy is $\Delta{}_{\mathrm{R}}G = \Delta{}_{\mathrm{R}}H - T\Delta{}_{\mathrm{}R}S$.

Then there are three cases to distinguish:

  1. $\Delta{}_{\mathrm{R}}G < 0$, an exergonic reaction, "running voluntarily" from the left to the right side of the reaction equation (react is spontaneous as written)
  2. $\Delta{}_{\mathrm{R}}G = 0$, the state of thermodynamic equilibrium, i.e. on a macroscopic level, there is no net reaction or
  3. $\Delta{}_{\mathrm{R}}G > 0$, an endergonic reaction, which either needs energy input from outside to run from the left to the right side of the reaction equation or otherwise runs backwards, from the right to the left side (reaction is spontaneous in the reverse direction)

Reactions may be classified according to reaction enthalpy, reaction entropy, free reaction enthalpy -- even simultaneously -- always favouring an exergonic reaction:

  1. Example, combustion of propane with oxygen, $\ce{5 O2 + C3H8 -> 4H2O + 3CO2}$. Since both heat dissipation ($\Delta_{\mathrm{R}}H < 0$, exothermic) and increase of the number of particles ($\Delta_{\mathrm{R}}S > 0$) favour the reaction, it is an exergonic reaction ($\Delta_{\mathrm{R}}G < 0$).
  2. Example, reaction of dioxygen to ozone, $\ce{3 O2 -> 2 O3}$. This is an endergonic reaction ($\Delta_{\mathrm{R}}G > 0$), because the number of molecules decreases ($\Delta_{\mathrm{R}}S < 0$) and simultaneously it is endothermic ($\Delta_{\mathrm{R}}H > 0$), too.
  3. Water gas reaction, where water vapour is guided over solid carbon $\ce{H2O + C <=> CO + H2}$. Only at temperatures $T$ yielding an entropic contribution $T \cdot \Delta_{\mathrm{R}}S > \Delta_{\mathrm{R}}H$, an endothermic reaction may become exergonic.
  4. Reaction of hydrogen and oxygen to yield water vapour, $\ce{2 H2 + O2 -> 2 H2O}$. This is an exothermic reaction ($\Delta_{\mathrm{R}}H < 0$) with decreasing number of particles ($\Delta_{\mathrm{R}}S < 0$). Only at temperatures at or below $T$ with $|T \cdot \Delta_{\mathrm{R}}S| < |\Delta_{\mathrm{R}}H|$ there is a macroscopic reaction. In other words, while the reaction works fine at room temperature, at high temperatures (e.g. 6000 K), this reaction does not run.

After all, please keep in mind this is about thermodynamics, and not kinetics. There are also indications of spontaneity of a reaction.

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  • $\begingroup$ So they're just synonyms for spontaneous and nonspontaneous? $\endgroup$ – user3932000 May 5 '17 at 21:22
  • $\begingroup$ @user3932000 No, they are not synonyms for spontaneous, or nonspontaneous. They assess the energy difference, comparing the energy state of the starting material(s) with the one of the product(s). $\endgroup$ – Buttonwood May 6 '17 at 3:02
  • $\begingroup$ Then are they two ways of expressing the same states? Exergonic/endergonic when describing energy differences, and spontaneous/nonspontaneous when describing reaction thermodynamics. $\endgroup$ – user3932000 May 7 '17 at 0:21
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Both exergonic and exothermic reactions release energy, however, the energies released have different meanings as follows:

  • Exothermic reaction

    • Energy released is just called energy
    • Energy of reactants is greater than that of products
    • Energy of the reaction system decreases relative to that of the surounding, i.e. the surrounding becomes hotter.
  • Exergonic reaction

    • Energy released, has a special name called Gibbs energy or Gibbs free energy
    • Energy reactants is greater than that of the products
    • It has nothing to do with how hot or cold reactants become. Has a more chemical meaning - it relates to the spontaneity of the reaction; thus it always means that a reaction is feasible, i.e. reaction will always happen.

In summary, whereas, an exergonic reaction means that a reaction is spontaneous, an exothermic reaction has nothing to do with spontaneity, but that an energy is released to the surrounding.

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For a exothermic reaction, $\Delta H\lt0$. For a exergonic reaction constraint is (from Gibbs-Helmholtz eqn): $\Delta G\lt0 \Rightarrow \Delta H-T\Delta S\lt0 \Rightarrow \Delta H\lt T\Delta S$ Hence, even if $\Delta H>0$ (endothermic reaction), a reaction can be exergonic provided it follows the constraint for it ($\Delta H\lt T\Delta S$; high temperature or greater no. of degree of freedom). So there is no such imposition that a reaction has to be exothermic if it is exergonic or vice versa.

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    $\begingroup$ Please edit your answer - as written, it's incomplete. See this style guide for how to typeset your posts. $\endgroup$ – Todd Minehardt Oct 27 '15 at 16:14
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Yes, all exergonic reactions are exothermic. Consider a reaction occurring spontaneously we know that energy would be released i.e. '$\ce{\Delta H}$ is negative' (since a reaction or process absorbing energy makes it non-spontaneous) and according to the second law of thermodynamics, entropy(or disorder) of the system must increase.

Negative $\ce{\Delta H}$ and increasing, positive entropy together make $\ce{\Delta G}$ negative according to the equation: $\ce{\Delta G = \Delta H~ - ~T\Delta S}$ (where $\ce{\Delta }$ = change; G= Gibb's free energy; H= enthalpy; T= Thermodynamic temperature and S= entropy). Therefore, if enthalpy change is negative and change in free energy is negative, they are both (respectively) exothermic and exergonic. The same applies for endothermic and endergonic.

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    $\begingroup$ Your first sentence is incorrect. See here for a spontaneous (i.e. exergonic) yet endothermic reaction. Examples are not so common because at low temperatures the entropic factor often turns out to be small, so free energy changes are mostly influenced by enthalpy changes. $\endgroup$ – Nicolau Saker Neto Aug 9 '14 at 15:03

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