Free energy Knowledge-Based Study

Study notes

1. What is Free Energy?

   In biological systems, Gibbs free energy (represented by the symbol G) is the portion of a system's total energy that can perform work when temperature and pressure remain uniform throughout the system, as is typical within a living cell.

   When a metabolic process takes place, the overall change in free energy is symbolized as $\Delta G$. This value is determined by changes in enthalpy (total internal heat energy, $\Delta H$), entropy (the measure of disorder, $\Delta S$), and absolute temperature in Kelvin (T):

   $$\Delta G = \Delta H - T\Delta S$$

   The value of $\Delta G$ allows biologists to predict whether a biological reaction will occur spontaneously (without requiring external energy input) or non-spontaneously.

   Classifying Metabolic Reactions

   Based on the sign and magnitude of $\Delta G$, biochemical reactions are categorized into two major classes:

   • Exergonic Reactions: These reactions have a negative change in free energy ($\Delta G < 0$). The products contain less free energy than the reactants, meaning free energy is released into the system. Exergonic reactions are thermodynamically spontaneous. Catabolic processes, such as the overall breakdown of glucose during cellular respiration, are classic examples.
   • Endergonic Reactions: These reactions have a positive change in free energy ($\Delta G > 0$). The products store more free energy than the initial reactants, which requires a net input of energy from the surroundings. These reactions are non-spontaneous. Anabolic pathways, such as synthesizing polypeptides from amino acids, are endergonic.

   When a system reaches chemical equilibrium, $\Delta G = 0$, and the system can do no work. Living cells continuously import substrates and export wastes to prevent their metabolic pathways from reaching thermodynamic equilibrium, which would result in cellular death.

   Driving Endergonic Processes

   Because cells constantly perform mechanical, transport, and chemical work that is endergonic, they rely on energy coupling—the use of an exergonic process to drive an endergonic one.

   The primary energy currency enabling this coupling is adenosine triphosphate (ATP). The hydrolysis of ATP into adenosine diphosphate (ADP) and an inorganic phosphate group ($\text{P}_i$) is highly exergonic:

   $$\text{ATP} + \text{H}_2\text{O} \rightarrow \text{ADP} + \text{P}_i \quad (\Delta G \approx -30.5\text{ kJ/mol under standard conditions})$$

   Cells frequently couple this release of free energy directly to an endergonic reaction through the transfer of a phosphate group to a recipient molecule, forming a phosphorylated intermediate that is less stable and more reactive.

   Thermodynamics vs. Reaction Kinetics

   It is critical to distinguish between the thermodynamic feasibility of a reaction (determined by $\Delta G$) and the speed at which the reaction occurs (kinetics).

   • A negative $\Delta G$ indicates that a reaction is favorable, but it does not specify how rapidly it will occur. Even strongly exergonic reactions may proceed imperceptibly slowly if the activation energy ($E_a$) barrier is high.
   • Enzymes function as biological catalysts by lowering the activation energy barrier, thereby accelerating the rate of the reaction.
   • Crucially, enzymes do not alter the overall $\Delta G$ of a reaction; they cannot turn an endergonic reaction into an exergonic one, nor do they change the final equilibrium ratio between products and reactants.