Entropy is a fundamental concept in physics and chemistry, often described as a measure of disorder or the dispersal of energy within a system. The idea that a system might spontaneously move toward a state of higher organization, or “negative entropy,” appears to challenge basic physical laws. Scientists confirm that a temporary or local increase in order can occur. However, this local ordering must always be accompanied by a greater increase in disorder elsewhere, ensuring the total entropy of the entire process always increases.
Understanding Entropy and Spontaneity
The scientific definition of entropy relates to the number of ways energy can be distributed among the particles in a system. Systems naturally move toward arrangements with the highest probability, such as a messy room being more probable than a perfectly ordered one. When energy spreads out from a concentrated source to a dispersed state, like heat moving from a hot object to a cold one, the entropy of the universe increases.
A spontaneous process is one that occurs on its own without requiring a continuous external input of energy. The predictor of spontaneity is the change in Gibbs Free Energy (Delta G). A process is defined as spontaneous only if Delta G is negative, indicating that the system’s potential energy decreases during the transformation.
The change in Gibbs Free Energy links the system’s enthalpy (heat content), temperature, and the system’s change in entropy (Delta S system). This relationship shows that the entropy change of the system alone is insufficient to predict spontaneity. Processes that move toward greater order (decreasing system entropy) can still proceed spontaneously if other thermodynamic factors compensate appropriately.
Local Decreases in Entropy (The System)
Processes where the system’s entropy decreases, often called local ordering, are common occurrences in the natural world. This happens whenever matter or energy becomes more concentrated or organized. A well-known example is the freezing of liquid water into solid ice, where mobile water molecules arrange themselves into a rigid, crystalline structure. This transition involves a significant decrease in the system’s entropy.
Crystallization is another example where entropy decreases. When salt precipitates out of a saturated solution, the ions move from a dispersed state to an ordered, fixed lattice structure. In both freezing and crystallization, the system moves toward a state of higher organization, resulting in a negative change in system entropy.
These ordering processes only occur under specific environmental conditions. Water freezing, for instance, only happens spontaneously below zero degrees Celsius. The internal ordering is allowed because the system sheds energy, often as heat, to the surroundings.
The Thermodynamic Rule of the Universe
The Second Law of Thermodynamics states that the total entropy of the universe must always increase for any spontaneous change. The total entropy change is the sum of the entropy change of the system (Delta S system) and the entropy change of the surroundings (Delta S surroundings).
For a process to be spontaneous, the entropy change of the universe (Delta S universe) must be positive. If a system undergoes a local decrease in entropy (negative Delta S system), the surrounding environment must experience a larger, positive increase in its entropy. This ensures the sum of the two changes remains positive overall.
The increase in the surroundings’ entropy is achieved through the dispersal of energy released by the system. When water freezes, the formation of the ordered ice structure releases heat into the surrounding environment. This released heat causes the molecules in the surroundings to move faster and more randomly, increasing the environment’s disorder.
Any process that locally decreases entropy is necessarily coupled to a larger, entropy-increasing process in the environment. The dispersal of energy into the surroundings is the thermodynamic cost required to create order within a localized region of space.
Applying the Concepts to Biological Order
Living organisms represent the most compelling example of local decreases in entropy, as they maintain highly complex and organized structures over time. A cell, with its precisely arranged organelles, represents a state of low entropy compared to the simple raw materials it consumes.
Life is an open system that continuously exchanges both energy and matter with its surroundings. To maintain internal organization, organisms must constantly import energy, such as sunlight or chemical energy consumed as food. This energy is utilized in metabolic pathways to build and repair cellular components.
The spontaneous processes of life are coupled to energy-dissipating reactions that increase the entropy of the surroundings. Organisms break down complex fuel molecules into simple, high-entropy waste products like carbon dioxide and water. A substantial amount of consumed energy is also released into the environment as heat, which increases the dispersal of energy in the surroundings. The local decrease in entropy associated with maintaining biological order is always paid for by a net increase in the entropy of the universe.