Gibbs Free Energy (\(G\)) measures the energy within a chemical system available to perform useful work at a constant temperature and pressure. It is often referred to as “available energy,” since it represents the theoretical maximum amount of non-expansion work a reaction can yield. This thermodynamic concept is fundamental because it predicts the direction and feasibility of a chemical transformation. Understanding the change in Gibbs Free Energy (\(Delta G\)) across a reaction is the basis for comprehending physical and chemical processes, including the complex metabolic pathways that sustain life.
The Difference Between Free Energy and Standard Free Energy
Free energy is most often discussed in terms of the change that occurs during a reaction, represented by \(Delta G\). This value reflects the energy difference between the final products and the initial reactants under a specific set of actual, real-world conditions. The \(Delta G\) is determined by the balance between the change in enthalpy (\(Delta H\), or heat content) and the change in entropy (\(Delta S\), or disorder), alongside the absolute temperature (\(T\)). Calculating \(Delta G\) using the equation \(Delta G = Delta H – TDelta S\) measures a reaction’s tendency to proceed under the current conditions.
The \(Delta G^circ\), or Standard Free Energy Change, is a specialized reference point calculated under a fixed set of agreed-upon “standard conditions.” For most chemical reactions, these conditions are defined as 25 degrees Celsius (298 Kelvin), a pressure of 1 atmosphere (or 1 bar), and a concentration of 1 Molar (M) for all dissolved reactants and products. The degree symbol (\(circ\)) indicates that the calculation uses these specific, theoretical starting concentrations. This standardization allows scientists worldwide to compare the inherent thermodynamic potential of different reactions on an equal footing.
The value of \(Delta G^circ\) is a fixed constant for a given reaction at a specified temperature, providing a baseline for its thermodynamic favorability. It is mathematically related to the reaction’s equilibrium constant, describing the reaction’s tendency to reach equilibrium when starting from the idealized standard state. The actual \(Delta G\) for a reaction under non-standard conditions is calculated by adjusting the \(Delta G^circ\) based on the actual concentrations of the reactants and products.
Interpreting Spontaneity and Equilibrium
The sign of the change in Gibbs Free Energy (\(Delta G\)) predicts the behavior of a chemical reaction under current conditions. A negative \(Delta G\) indicates the reaction is exergonic, meaning it releases free energy and is considered thermodynamically favorable, or spontaneous. Conversely, a positive \(Delta G\) signifies an endergonic, or non-spontaneous, reaction that requires an input of free energy to proceed. For example, the oxidation of glucose has a large negative \(Delta G\), showing it will proceed without a continuous energy input.
The term “spontaneous” in thermodynamics does not imply the reaction will be fast. It simply means the transformation is favorable and will eventually occur on its own, though it may still require an initial energy boost, known as activation energy, to start. A reaction is at equilibrium when the change in free energy is zero (\(Delta G = 0\)). In this state, the forward and reverse reaction rates are equal, resulting in no net change in the concentration of reactants or products.
While the \(Delta G^circ\) provides a fixed reference, the actual behavior within a system, especially a living cell, is dictated by the actual \(Delta G\). The actual \(Delta G\) accounts for the current concentrations of all molecules, which are rarely at the 1 M standard condition. A reaction that has a positive \(Delta G^circ\) (non-spontaneous at standard conditions) can become spontaneous (negative \(Delta G\)) inside a cell if the ratio of reactants to products is significantly skewed. This means a cell can drive an otherwise unfavorable reaction forward simply by maintaining a high concentration of reactants or a low concentration of products.
How Cells Use Free Energy to Power Life
Living organisms maintain complex internal environments far from equilibrium, utilizing free energy to perform necessary biological work. Since cells rarely operate under the theoretical chemical standard conditions (1 M concentrations), they rely on the actual \(Delta G\) value for their processes. A specialized biochemical standard free energy, \(Delta G^{circ’}\), is sometimes used in biology, which adjusts the standard conditions to a physiological pH of 7.0, a more realistic reference for the cell’s environment.
To overcome endergonic reactions (those with a positive \(Delta G\)), cells employ a mechanism called reaction coupling. This involves pairing an unfavorable reaction with a highly exergonic reaction. The combined reactions proceed because the overall change in free energy for the two reactions summed together is negative. This strategy allows the energy released from one process to directly power another.
Adenosine triphosphate (ATP) is the central energy currency that makes coupling possible. ATP stores potential energy in its phosphate bonds, and its hydrolysis (breakdown into ADP and inorganic phosphate, \(text{P}_i\)) is a highly exergonic reaction. Under standard conditions, the \(Delta G^circ\) for ATP hydrolysis is approximately \(-30.5\) kilojoules per mole (\(text{kJ/mol}\)). Inside a living cell, the actual \(Delta G\) is much more negative (around \(-57\) to \(-69\) \(text{kJ/mol}\)). This large negative value drives essential endergonic processes, such as synthesizing biomolecules or powering muscle contraction, by transferring a phosphate group to a reactant.