Exergonic and endergonic response

Chemical reactions are referred to  as exergonic or endergonic reactions with regard to whether the free enthalpy G of the components involved in the reaction increases or decreases :

• exergon: ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G<0}$
• endergon: ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G>0}$

These terms should not be confused with exothermic and endothermic (see below and delimitation ).

Exergonic and endergonic reactions

Reactions that are thermodynamically favorable are referred to as exergonic or exergonic , whereas endergonic reactions are thermodynamically unfavorable .

Both exergonic and endergonic reactions take place in principle "voluntarily", provided that the reaction kinetics allow this. In exergonic reactions, the equilibrium is only further on the side of the products than in endergonic reactions . It should be noted that on closer inspection, every reaction is an equilibrium reaction .

An example of an endergonic process is the formation of a protein in an aqueous solution of amino acids . This reaction can only be realized if it is coupled to other, exergonic processes, so that in total the sign of is negative; in biological systems this is mostly achieved through the hydrolysis of ATP . Without this coupling, the reaction would take place, but only to a marginal extent, which would be completely inadequate to cope with the biochemical task. ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G}$

Since the reverse reaction of an endergonic reaction is always exergonic (and vice versa), proteins should actually spontaneously break down into their amino acids again. However, the rate of the disintegration reaction under physiological conditions is so slow that it can be neglected; H. In this case, peptide bonds are kinetically stable . An argument from the reaction kinetics is decisive here.

Systems always strive towards the state of equilibrium , because here the free enthalpy assumes the minimum value. Once a system has reached its equilibrium, the concentrations of the reactants no longer change because G can not be reduced any further, and it applies . ${\displaystyle \mathrm {\Delta } G=0}$

Important distinctions

• ${\displaystyle \mathrm {\Delta } G}$ is generally the change in free enthalpy during a process
• ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G}$ describes the change in the free enthalpy when the reaction is complete.

It is hypothetical that a reaction takes place completely, since every reaction only proceeds to chemical equilibrium. Nevertheless, it is an important quantity, since with its help the equilibrium constant K can be calculated: ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G}$

${\displaystyle K=\exp \left({\frac {-\Delta _{\mathrm {R} }G^{0}}{R\cdot T}}\right)}$

With

• the gas constant  R
• the absolute temperature  T.

Determination of the free enthalpy of reaction

${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G}$is given by the following relationship (often also referred to as the Gibbs-Helmholtz equation ):

${\displaystyle \Delta _{\mathrm {R} }G=\Delta _{\mathrm {R} }H-T\cdot \Delta _{\mathrm {R} }S}$

With

• ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }G}$: free enthalpy of reaction
• ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }H}$: Enthalpy of reaction (i.e. change in the enthalpy of substances during the course of the reaction)
• ${\displaystyle \mathrm {\Delta } _{\mathrm {R} }S}$: Entropy of reaction (i.e. change in entropy of substances as the reaction proceeds).

${\displaystyle \mathrm {\Delta _{\mathrm {R} }G} }$can be calculated for standard conditions with the help of tabulated values ​​(standard reaction enthalpies and standard reaction entropies ) . One then speaks of the free standard enthalpy of reaction . A conversion to other temperatures can be done using the Van-'t-Hoff equation . ${\displaystyle \Delta H^{\circ }}$ ${\displaystyle \Delta S^{\circ }}$ ${\displaystyle \Delta G^{\circ }}$

Interpretation of the equation Δ G = Δ H - T · Δ S

Thermodynamics of the chemical reaction

The driving force for a chemical reaction to take place is the increase in entropy S in the universe ( cf. 2nd law of thermodynamics ).

If one considers a system that cannot exchange heat with the environment ( adiabatic system ), the only condition that has to be placed on a spontaneously occurring process is . The decisive factor here is not whether the individual reaction is exergonic or endergonic, but that the system is not yet in equilibrium. ${\displaystyle \Delta S_{\rm {System}}>0}$

Exothermic and endothermic reaction

If the system is allowed to exchange heat with the environment ( diabatic system ), the change in entropy in the environment must also be taken into account. This can be recorded via the change in enthalpy of the system: ${\displaystyle \Delta H}$

• as a negative contribution if the reaction is exothermic , thermal energy is distributed to the environment and in this way the entropy increases in the environment,
• as a positive contribution if the reaction is endothermic and the entropy in the environment decreases because thermal energy is concentrated in the reaction vessel.

The system strives for states with minimum free enthalpy, since this is the state of maximum entropy.

Individual evidence

1. ↑ ^{a b} Florian Horn: Human biochemistry, the textbook for medical studies . Georg Thieme Verlag, 2009, ISBN 978-3-13-130884-9 , p. 63 ( limited preview in Google Book search).
2. ↑ P. Stephan, K. Schaber, K. Stephan, F. Mayinger: Thermodynamics - Fundamentals and technical applications - Volume 2: Multi-substance systems and chemical reactions 15th edition, Springer, Heidelberg 2010, ISBN 978-3-540-36709-3 .
3. ↑ Ulf Dettmer and Malte Folkerts: Biochemistry . Elsevier, Urban & FischerVerlag, 2005, ISBN 978-3-437-44450-0 , pp. 6 ( limited preview in Google Book search).
4. ^ HA Harper, G. Löffler, PE Petrides, L. Weiss: Physiological chemistry An introduction to medical biochemistry for students of medicine and doctors . Springer-Verlag, 2013, ISBN 978-3-662-09766-3 , pp. 257 ( limited preview in Google Book search).
5. ↑ Walter J. Moore; "Fundamentals of physical chemistry", Walter de Gruyter, 1990.