Thermodynamics in Biology: Energy, Entropy, and Free Energy

Core Concept

Thermodynamics is the study of energy transformations. Cells are open systems that continuously take in energy (light or food), transform it, and release heat and disorder into the environment. The first law (energy is conserved) and the second law (every energy transfer increases the total entropy of the universe) govern every metabolic reaction. A reaction's free-energy change (ΔG) tells us whether it can proceed spontaneously — that is, without a net input of energy — but says nothing about how quickly it will occur.

Key Concepts

Forms of Energy

• Kinetic energy — energy of motion (a moving molecule, heat, light).
• Potential energy — stored energy due to position or structure (a molecule's chemical bonds, a proton gradient across a membrane).
• Chemical energy — potential energy stored in the arrangement of atoms in molecules; it is released when bonds are rearranged in reactions.

The First Law of Thermodynamics

Energy cannot be created or destroyed, only transferred or transformed. In a cell, chemical energy in glucose is transformed into ATP, then into motion or biosynthesis — the total amount of energy is conserved even as its form changes. (This is the conservation of energy.)

The Second Law and Entropy

Every energy transfer or transformation increases the entropy (disorder) of the universe. Some energy is always lost as heat, which disperses and cannot be fully recovered for useful work. Entropy (S) is a measure of disorder or randomness. The second law explains why no energy conversion is 100% efficient and why heat is the "least useful" form of energy.

Life and the Second Law

Living things are highly ordered (low entropy), which seems to contradict the second law — but it does not. Cells are open systems that import energy (sunlight or food) and export disorder (heat and waste, like CO₂ and H₂O). The decrease in a cell's entropy is more than offset by the increase in entropy of its surroundings, so the entropy of the universe still rises. Life does not break the second law; it runs on it.

Free Energy (Gibbs Free Energy, ΔG)

Gibbs free energy (G) is the portion of a system's energy that can do work at constant temperature and pressure. The change in free energy during a reaction is:

ΔG = ΔH − TΔS

where ΔH is the change in enthalpy (total energy), T is temperature, and ΔS is the change in entropy. ΔG determines spontaneity:

• Exergonic reaction: ΔG < 0 — releases free energy, proceeds spontaneously.
• Endergonic reaction: ΔG > 0 — requires an input of free energy, non-spontaneous.
• ΔG = 0: the system is at equilibrium; no net change.

Spontaneous ≠ Fast

"Spontaneous" is a thermodynamic term meaning the reaction can proceed with a net release of free energy; it says nothing about speed. The hydrolysis of ATP is spontaneous but slow without an enzyme, and the oxidation of sugar is hugely exergonic yet sugar sits inert on a shelf. Speed is governed by the activation-energy barrier and by enzymes, not by ΔG.

How It Works

A reaction's ΔG is the net of two contributions: the change in bond energies (ΔH) and the change in disorder (TΔS). (1) When strong bonds in reactants are replaced by even stronger bonds in products, or when a system becomes more disordered (e.g., a large molecule splits into many small ones), ΔG is negative and the reaction is exergonic. (2) Exergonic reactions can, in principle, do work. (3) Endergonic reactions, such as building proteins or pumping ions uphill, cannot proceed unless they are coupled to an exergonic reaction (usually ATP hydrolysis) so that the net ΔG is negative. (4) Enzymes speed reactions by lowering the activation-energy barrier but never change ΔG — a reaction that is thermodynamically unfavorable stays unfavorable no matter what enzyme is present.

How it works

A reaction's ΔG is the net of two contributions: the change in bond energies (ΔH) and the change in disorder (TΔS). (1) When strong bonds in reactants are replaced by even stronger bonds in products, or when a system becomes more disordered (e.g., a large molecule splits into many small ones), ΔG is negative and the reaction is exergonic. (2) Exergonic reactions can, in principle, do work. (3) Endergonic reactions, such as building proteins or pumping ions uphill, cannot proceed unless they are coupled to an exergonic reaction (usually ATP hydrolysis) so that the net ΔG is negative. (4) Enzymes speed reactions by lowering the activation-energy barrier but never change ΔG — a reaction that is thermodynamically unfavorable stays unfavorable no matter what enzyme is present.

Common confusions

• "Living things violate the second law." Wrong — they are open systems that export enough entropy (heat + waste) to keep the universe's total entropy rising.
• "Spontaneous means fast." Wrong — spontaneous is about ΔG (thermodynamics), not speed (kinetics); many spontaneous reactions are slow without enzymes.
• "Endergonic reactions never happen." Wrong — they happen constantly when coupled to exergonic reactions (e.g., ATP hydrolysis) so the overall ΔG is negative.
• "Energy is 'used up' in reactions." Wrong — the first law says energy is conserved; it is transformed and partly dispersed as heat, not destroyed.
• "ΔG positive means the reaction releases energy." Wrong — positive ΔG means the reaction requires an input of free energy (endergonic).

Quick review

• Energy: kinetic (motion) vs. potential (stored, chemical).
• First law = conservation of energy; second law = entropy increases.
• Cells are open systems importing energy and exporting disorder.
• ΔG < 0 exergonic (spontaneous); ΔG > 0 endergonic (needs energy).
• Spontaneous ≠ fast; enzymes lower activation energy but not ΔG.
• Life runs downhill thermodynamically by coupling unfavorable reactions to favorable ones.