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The second law of thermodynamics states that entropy is always increasing in the universe. Entropy was defined as a random, disordered condition. When a reaction has reached the end of the road and entropy is at a maximum, the reaction has reached equilibrium. At chemical equilibrium there is no longer any net change from reactant to product. At equilibrium, then, there is no further potential to do work. In other words, the further from equilibrium the reactants of a process are, the more potential use the reaction has.
The importance of the equilibrium constant in determining the free energy change for a reaction is illustrated in Figure 1, where A is the reactant, B is the product, and each of the reactions proceeds to equilibrium.

Figure 1 - Equilibrium constant and free energy

In reaction (a), where very little or no product is formed, it is obvious that little has happened, so there is little opportunity for energy to be exchanged. At the other end, in reaction (c), almost all the reactant has become product. In this reaction, much has happened chemically, and there has been opportunity to capture some of the energy given off.
The equilibrium constant Keq is used to denote the ratio of concentrations of products to reactants at equilibrium:

Here the brackets indicate “concentration of.” The of a reaction is an immutable constant at specified conditions of temperature and pressure.
In examining K and looking at the reactions in Figure 1, we can see that if the quotient is large, the reaction has potential use in driving a biological system.
Empirically, it has been determined that the change in free energy in a reaction is simply related to its equilibrium constant:

where R is the gas constant (1.99 cal.mol.deg^-1), T is absolute temperature, and ln is the natural logarithm of the equilibrium constant determined under standard laboratory conditions.
The symbol .G°’ thus refers to the standard free energy change determined in the laboratory when a reaction takes place at 25 C, at 1 atmosphere of pressure, and where the concentrations are maintained at 1 molal and in an aqueous medium at pH 7. In the notation the superscript’ refers to pH 7; the refers to the other standard conditions. Although these “standard conditions” seem to be far removed from those in the body, categorizing reactions by this system allows us to compare free energy potentials of various reactions.
The relationship between K and .G°’ is further illustrated in Table 1.

table 1 - relationship between Keq and free energy

This entry was posted on Tuesday, December 29th, 2009 at 7:40 pm and is filed under Bioenergetics, Physiology, Sports Medicine. You can follow any responses to this entry through the RSS 2.0 feed. You can leave a response, or trackback from your own site.