20.2 Thermodynamics of biological systems

Thermodynamic Properties and Energy in Biological Systems

Biological systems are thermodynamic outliers. They maintain extraordinary order and complexity while surrounded by an environment that trends toward disorder. They pull this off by operating far from equilibrium, constantly importing energy and exporting waste. Understanding how they do this requires applying core thermodynamic principles to the messy, compartmentalized reality of living cells.

Thermodynamics in Biological Systems

Living organisms are open systems, meaning they continuously exchange both energy and matter with their surroundings. Nutrients flow in, waste products flow out, and the organism maintains a stable internal environment (homeostasis) despite fluctuations in temperature, pH, and other external conditions.

This high degree of organization seems to violate the Second Law of Thermodynamics, which states that the total entropy of an isolated system always increases. But organisms aren't isolated systems. They maintain internal order by increasing the entropy of their surroundings, typically by releasing heat and disordered waste products. The total entropy of the universe still increases.

Two strategies make this possible:

• Coupled reactions: Exergonic reactions (energy-releasing, negative $\Delta G$) are paired with endergonic reactions (energy-requiring, positive $\Delta G$) so that the overall process is thermodynamically favorable. For example, ATP hydrolysis releases energy that drives muscle contraction.
• Compartmentalization: Biological membranes with selective permeability create and maintain concentration gradients of ions and molecules. Mitochondria and chloroplasts exploit these gradients to store and convert energy. Without membrane-bound compartments, the gradients that power oxidative phosphorylation and photosynthesis would dissipate.

ATP as Energy Currency

ATP (Adenosine Triphosphate) is the primary energy carrier in cells. It consists of an adenosine base, a ribose sugar, and three phosphate groups linked by phosphoanhydride bonds. Hydrolysis of the terminal phosphate bond releases energy:

$\Delta G = -30.5 \text{ kJ/mol (under standard biochemical conditions)}$

That value is for standard conditions. In the actual cellular environment, the free energy of ATP hydrolysis is often closer to $-50$ to $-54$ kJ/mol because cells maintain ATP concentrations far above equilibrium levels.

ATP coupling drives thermodynamically unfavorable processes by linking ATP hydrolysis to endergonic reactions. Three major categories of work depend on this coupling:

1. Mechanical work — Muscle contraction through actin-myosin interactions
2. Transport work — Active transport against concentration gradients (e.g., the $\text{Na}^+/\text{K}^+$-ATPase pumps 3 $\text{Na}^+$ out and 2 $\text{K}^+$ in per ATP hydrolyzed)
3. Biosynthetic work — Assembly of complex molecules like proteins and nucleic acids from simpler precursors

ATP is regenerated through two main mechanisms:

• Substrate-level phosphorylation: Direct transfer of a phosphate group to ADP, occurring in glycolysis and the citric acid cycle
• Oxidative phosphorylation: ATP synthesis driven by the proton motive force across the inner mitochondrial membrane, generated by the electron transport chain and harnessed by ATP synthase

A typical human cell turns over its entire pool of ATP roughly every 1–2 minutes. This rapid cycling maintains a high $\text{ATP}/\text{ADP}$ ratio, which keeps metabolic reactions displaced far from equilibrium. That displacement is what allows cells to regulate metabolic flux through mechanisms like allosteric enzyme regulation and feedback inhibition.

Free Energy and Biomolecular Stability

Free Energy in Biochemical Reactions

Gibbs free energy ($G$) is the thermodynamic potential that determines whether a reaction will proceed spontaneously at constant temperature and pressure:

$G = H - TS$

where $H$ is enthalpy, $T$ is absolute temperature, and $S$ is entropy. The change in free energy for a reaction, $\Delta G$, tells you its direction:

• $\Delta G < 0$: Spontaneous (exergonic). Products are at lower free energy than reactants. Examples include glucose oxidation and ATP hydrolysis.
• $\Delta G = 0$: System is at equilibrium. Forward and reverse reactions proceed at equal rates with no net change in concentrations.
• $\Delta G > 0$: Non-spontaneous (endergonic). Requires energy input to proceed.

A critical distinction for biochemistry: the standard free energy change ($\Delta G^{\circ'}$, using the biochemical standard state at pH 7) tells you the free energy change when all reactants and products are at 1 M concentration. But actual cellular concentrations are rarely 1 M. The real free energy change depends on those concentrations:

$\Delta G = \Delta G^{\circ'} + RT \ln Q$

where $R$ is the gas constant, $T$ is temperature in Kelvin, and $Q$ is the reaction quotient (ratio of product to reactant concentrations). This equation is why a reaction with a positive $\Delta G^{\circ'}$ can still proceed forward in a cell if the concentration ratio is favorable.

Thermodynamics of Biomolecule Stability

Protein folding is governed by free energy minimization. The native (functional) state of a protein sits at a free energy minimum, representing the most thermodynamically stable conformation accessible under physiological conditions. Reaching this state involves a trade-off between enthalpic and entropic contributions.

Enthalpic contributions stabilize the folded state through favorable noncovalent interactions:

• Hydrogen bonds between backbone atoms and between side chains
• Van der Waals interactions from tight packing of residues in the protein interior
• Electrostatic interactions (salt bridges) between charged residues
• Favorable contacts between hydrophilic residues and the aqueous solvent at the protein surface

Entropic contributions are dominated by the hydrophobic effect, which is actually the single largest driving force for protein folding. When nonpolar residues are exposed to water, surrounding water molecules form ordered "cages" (clathrate-like structures), which is entropically costly. Burying these residues in the protein core releases those ordered water molecules back into the bulk solvent, producing a net entropy increase that favors folding.

Protein denaturation occurs when external factors destabilize the native state enough that the unfolded ensemble becomes thermodynamically favored. Common denaturants include:

• Heat: Increases the $TS$ term, favoring the higher-entropy unfolded state
• Extreme pH: Disrupts electrostatic interactions and hydrogen bonds by altering residue charge states
• Chemical denaturants (urea, guanidinium chloride): Stabilize the unfolded state by forming favorable interactions with the peptide backbone

Denaturation results in loss of tertiary and secondary structure, exposing buried hydrophobic residues and typically abolishing biological function.

These same thermodynamic principles extend to other biomolecules:

• Nucleic acids: DNA double helix stability arises from Watson-Crick base pairing (hydrogen bonds) and base stacking interactions (van der Waals and hydrophobic contributions). RNA secondary structures like hairpins and stem-loops follow similar logic.
• Lipid bilayers: Membrane formation is driven by the hydrophobic effect. Nonpolar fatty acid tails sequester away from water, while polar head groups face the aqueous phase. This self-assembly is spontaneous because it maximizes the entropy of surrounding water molecules.