Living organisms rely on countless chemical reactions for essential processes like growth, movement, and reproduction. Understanding how these reactions proceed and whether they require or release energy is fundamental to comprehending life. This energy accounting is captured by the concept of Gibbs free energy, and in biological contexts, a specific adaptation known as Delta G prime ($\Delta G’$) is particularly relevant. It provides a framework for analyzing the energetics of biochemical reactions within the cell’s unique environment.
The Concept of Free Energy
Gibbs Free Energy ($\Delta G$) represents the energy available in a system that can be harnessed to do useful work at constant temperature and pressure. The sign of $\Delta G$ indicates the spontaneity of a reaction. Reactions that release free energy are termed exergonic reactions, and they proceed spontaneously without external energy input. Conversely, reactions that require an input of energy to occur are called endergonic reactions, and they are non-spontaneous.
Chemists use the standard Gibbs free energy change, $\Delta G^\circ$, to establish a baseline for comparison. This value is measured under specific, idealized standard conditions: 25 degrees Celsius, 1 atmosphere of pressure, and all reactants and products at an initial concentration of 1 M. While $\Delta G^\circ$ is a fixed, intrinsic property for a given chemical reaction, these conditions rarely reflect the dynamic environment inside living cells.
Adjusting for Biological Reality
The standard conditions used for $\Delta G^\circ$ are unsuitable for biological systems, which operate under very different circumstances. For instance, a 1 M concentration of hydrogen ions would correspond to a highly acidic pH of 0, a condition incompatible with cellular life. Biochemists therefore employ “transformed standard conditions” to make the free energy calculations more biologically relevant. This adjusted value is known as $\Delta G’$.
Under these transformed standard conditions, several parameters are redefined. The pH is set to a physiological value of 7.0, which is typical for the cellular cytoplasm. The temperature is commonly standardized to 25 degrees Celsius. Additionally, the concentration of water is considered constant at 55.5 M, and for reactions involving magnesium ions, their concentration is often set to 1 mM, reflecting their common physiological levels. These adjustments ensure that $\Delta G’$ provides a more accurate representation of the energy changes that occur within a living cell.
What the Values Mean
Interpreting the numerical value of $\Delta G’$ offers direct insights into how a biochemical reaction behaves under transformed standard conditions. A negative $\Delta G’$ signifies an exergonic reaction, meaning it releases free energy and can proceed spontaneously. This indicates that the products of the reaction have less free energy than the reactants, making the conversion energetically favorable. Conversely, a positive $\Delta G’$ identifies an endergonic reaction, which requires an input of free energy to proceed. In such cases, the products possess more free energy than the reactants, making the reaction non-spontaneous under these conditions.
When $\Delta G’$ is zero, the reaction is at equilibrium, meaning there is no net change in the concentrations of reactants and products. While $\Delta G’$ indicates spontaneity under standard transformed conditions, the actual spontaneity of a reaction within a cell ($\Delta G$) also depends on the real-time, non-standard concentrations of reactants and products. However, $\Delta G’$ still serves as a valuable reference point for understanding the intrinsic energy favorability of a reaction in a physiological context.
Driving Life’s Reactions
Understanding $\Delta G’$ is important for comprehending cellular metabolism, where many essential reactions are endergonic and would not proceed spontaneously on their own. Cells overcome this challenge through a process called reaction coupling. This involves linking an energy-requiring (endergonic) reaction with a highly energy-releasing (exergonic) reaction. The overall $\Delta G’$ for the combined, coupled reaction must be negative for the process to proceed spontaneously.
A prime example of this coupling mechanism involves adenosine triphosphate (ATP), often referred to as the cell’s energy currency. The hydrolysis of ATP, where ATP is broken down into adenosine diphosphate (ADP) and inorganic phosphate (Pi), is a highly exergonic reaction, releasing a significant amount of free energy. This released energy can then be used to drive endergonic processes, such as muscle contraction, active transport across membranes, or the synthesis of complex molecules. By coupling these reactions, cells ensure that vital, otherwise non-spontaneous, biochemical pathways can occur efficiently, sustaining all aspects of life.