Gibbs Free Energy (\(G\)) is a fundamental concept in chemistry and biology, measuring the usable energy within a system. It represents the maximum amount of non-mechanical work that can be extracted from a system operating at constant temperature and pressure. The change in this energy, \(\Delta G\), is the primary tool for determining the inherent feasibility of any chemical reaction or physical process. A process with thermodynamic potential can occur without a continuous external energy input.
Predicting if a Reaction Will Happen
The sign of the change in Gibbs Free Energy (\(\Delta G\)) directly predicts whether a chemical process is possible under given conditions. A negative \(\Delta G\) indicates an exergonic reaction, meaning it is energetically favorable and can proceed spontaneously. In this case, the products contain less free energy than the reactants, and the excess energy is released.
Conversely, a positive \(\Delta G\) signifies an endergonic process that is not energetically favorable. Such a reaction will not occur without a continuous supply of external energy, such as the input required for synthesizing large molecules. If \(\Delta G\) is exactly zero, the system is in a state of balance called equilibrium.
This prediction only concerns the thermodynamic possibility of a reaction, not the speed at which it occurs. A reaction could have a highly negative \(\Delta G\) but still happen extremely slowly, such as the conversion of diamond to graphite. The \(\Delta G\) value only indicates whether a process can happen, not how fast it will happen.
The Role of Energy and Disorder
Gibbs Free Energy integrates two separate, often competing, thermodynamic driving forces into a single predictive value. The relationship is summarized by the equation \(\Delta G = \Delta H – T\Delta S\), combining the change in enthalpy (\(\Delta H\)) and the change in entropy (\(\Delta S\)). \(\Delta H\) represents the change in the total heat content, measuring the drive toward lower energy. Reactions that release heat have a negative \(\Delta H\), which favors the process.
The second term, \(T\Delta S\), accounts for the tendency of systems toward greater disorder or randomness. \(\Delta S\), the change in entropy, measures this molecular chaos; a positive \(\Delta S\) (increase in disorder) favors a reaction’s feasibility. Since the entropy term is multiplied by the absolute temperature (\(T\)), the influence of disorder becomes greater at higher temperatures.
The final \(\Delta G\) value represents the balance between the system’s drive to minimize energy (\(\Delta H\)) and maximize disorder (\(T\Delta S\)). If the energy-releasing component or the increase in disorder is significant enough, the overall \(\Delta G\) becomes negative. This balance makes Gibbs Free Energy a comprehensive measure of a system’s propensity for change.
How Cells Use Free Energy
Living cells constantly perform thousands of chemical reactions, many of which are endergonic and would not occur spontaneously. To overcome this energetic barrier, cells employ reaction coupling, utilizing energy released from highly exergonic processes. Adenosine triphosphate (ATP) serves as the universal energy currency for this purpose.
The breakdown of ATP into adenosine diphosphate (ADP) and an inorganic phosphate group (P\(_{\text{i}}\)) is a highly exergonic reaction, releasing a large amount of free energy. Under standard conditions, this hydrolysis yields a \(\Delta G\) between -28 and -34 kJ/mol, though the magnitude is often larger inside a cell. This substantial energy release is directly linked to an endergonic process, such as building a protein or contracting a muscle fiber.
Coupling often occurs through a shared intermediate, where the phosphate group released from ATP is temporarily transferred to a reactant in the endergonic pathway. This transfer changes the free energy landscape of the second reaction, resulting in a net negative \(\Delta G\) for the overall combined process. By summing the \(\Delta G\) values of the two reactions, the coupled reaction becomes energetically favorable, powering necessary life functions.
When Reactions Reach Balance
The state of equilibrium is defined thermodynamically by a \(\Delta G\) value of zero. At this point, the concentrations of reactants and products are no longer changing, creating a balance. This indicates that the rate of the forward reaction is exactly equal to the rate of the reverse reaction, meaning chemical activity has not stopped.
The magnitude of the \(\Delta G\) value at any non-equilibrium state relates directly to how far the system is from the final balance point. A reaction with a large negative \(\Delta G\) is far from equilibrium and has a high capacity to do work. As the reaction proceeds, \(\Delta G\) continually decreases until it reaches the minimum value of zero at equilibrium.
The \(\Delta G\) value is also directly linked to the equilibrium constant (\(K\)), which is the ratio of product concentration to reactant concentration at balance. If the standard change in free energy (\(\Delta G^{\circ}\)) is negative, \(K\) will be greater than one, meaning products are favored at equilibrium. Conversely, a positive \(\Delta G^{\circ}\) means \(K\) is less than one, and the reactants are favored.