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Endergonic and Exergonic Reactions#

Endergonic and exergonic reactions

Any system of molecules that undergos a physical transformation/reorganizaton (aka. reaction) will have an associated change in internal energy and entropy.  If we examine a single isolated reaction, in which unique reactants are converted into unique products the Gibbs energy of the system will be dependent several factors, key among which are (a) the internal energy and entropy differences associated with the molecular rearrangements and (b) the degree to which the reaction is out-of-equilibrium. 

If, for the sake of simplicity we begin by considering only the contribution of the molecular transformations in the system on ∆G, we conclude that reactions with ∆G < 0, the products of the reaction have less Gibbs energy than the reactants. Since ∆G is the difference between the enthalpy and temperature-scaled entropy changes in a reaction, a net negative ∆G can arise in through changes largely of enthalpy, entropy or most often both. The left panel of Figure 1 below shows a common graphical representation of an exergonic reaction. This type of graph is called a reaction coordinate diagram. Gibbs energy is plotted on the y-axis, and the x-axis in arbitrary units shows the progress of a reaction. In the case of an exergonic reaction, the figure on the left indicates two key things: (1) the difference between the free energy of the reactants and products is negative and (2) the progress of the reaction requires some input of free energy (shown as an energy "hill" or barrier). This graph does not tell us how the energy in the system was redistributed, only that the difference between enthalpy and temperature-scaled entropy is negative. Exergonic reactions are said to occur spontaneously. Understanding which chemical reactions are spontaneous is extremely useful for biologists who are trying to understand whether a reaction is likely to "go" or not.

It is important to note that the term "spontaneous"—in the context of thermodynamics—does NOT imply anything about how fast the reaction proceeds. The change in free energy only describes the difference between beginning and end states, NOT how fast that transition takes place. This is somewhat contrary to the everyday use of the term, which usually carries the implicit understanding that something happens quickly. As an example, the oxidation/rusting of iron is a spontaneous reaction. However, an iron nail exposed to air does not rust instantly—it may take years.

A chemical reaction with a positive ∆G means that the products of the reaction have a higher free energy than the reactants (see the right panel of Figure 1). These chemical reactions are called endergonic reactions, and they are NOT spontaneous. An endergonic reaction will not take place on its own without the transfer of energy into the reaction or increase of entropy somewhere else.

Figure 1. Reaction coordinate diagrams of exergonic and endergonic reactions.  Exergonic and endergonic reactions are characterized by changes in Gibbs energy. In the equilibrium state of an exergonic reaction, the Gibbs energy of the products is lower than that of the reactants. Meanwhile, the equilibrium state of an endergonic reaction in, the Gibbs energy of the products is higher than that of the reactants. Attribution: Marc T. Facciotti (own work)

The building of complex molecules, such as sugars, from simpler ones is an anabolic process and is endergonic. On the other hand, the catabolic process, such as the breaking down of sugar into simpler molecules, is generally exergonic. Like the example of rust above, while the breakdown of biomolecules is generally spontaneous, these reactions don’t necessarily occur instantaneously (quickly). Remember, the terms endergonic and exergonic only refer to the difference in Gibbs energy between the products and reactants; they don't tell you about the rate of the reaction (how fast it happens). The issue of rate will be discussed in later sections.

An important concept in the study of metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can proceed in both directions, often transferring energy into their environment in one direction and transferring energy in from the environment in the other direction. The same is true for the chemical reactions involved in cell metabolism, such as the breaking down and building up of proteins into and from individual amino acids, respectively. Reactants within a closed system will undergo chemical reactions in both directions until a state of equilibrium is reached. Equilibrium in a chemical reaction is the state in which both reactants and products are present in concentrations that have no further tendency to change with time. Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction. NOTE THIS LAST STATEMENT! Equilibrium means that the relative concentrations of reactants and products are not changing in time, BUT it does NOT mean that there is no interconversion between substrates and products—it just means that when the reactant(s) are converted to product(s) that product(s) are converted to reactant(s) at an equal rate (see Figure 2). The state of equilibrium is also one of the lowest possible free energy states for the reaction and is a state of maximal entropy.

If a reaction is kept or started far out of equilibrium this state of the system also contributes to the overall Gibbs energy of a reaction. Either a rebalancing of substrate or product concentrations (by adding or removing substrate or product) or a positive change in free energy, typically by the transfer of energy from outside the reaction, can to move a reaction to an out-of-equilibrium state. Note that in a living cell, most chemical reactions do not reach a state of equilibrium—this would require that they reach their lowest free energy state, a state that is almost by definition incompatible with life. Energy is therefore required to keep biological reactions out of their equilibrium state. In this way, living organisms are in a constant, energy-requiring, uphill battle against equilibrium and entropy. This also means that the Gibbs energy of most biological reactions as they occur in the cell must also include a contribution from this out-of-equilibrium state. The Gibbs energy of these reactions, therefore, is often different from that reported under standard conditions.

Figure 2. At equilibrium, do not think of a static, unchanging system. Instead, picture molecules moving in equal amounts from one area to another. Here, at equilibrium, molecules are still moving from left to right and right to left. The net movement however, is equal. There will still be about 15 molecules in each side of this flask once equilibrium is reached. Source: