Summary of "Introduction to Biological Thermodynamics"

Summary of “Introduction to Biological Thermodynamics“

This lecture by Professor Jeff Yarger from Arizona State University introduces the fundamental concepts of thermodynamics with a focus on biological systems. The discussion covers the nature of energy, how it is partitioned in thermodynamics, and its application in understanding biological processes at macroscopic and molecular levels.

Main Ideas and Concepts

Thermodynamics Overview

Thermodynamics is the study of energy and its transformations.
In biological systems, the universe is divided into the system (the part of interest) and the surroundings.
Energy bookkeeping tracks how energy moves as heat or work.

Energy Types in Thermodynamics

Unlike physics where energy is split into kinetic and potential, thermodynamics splits energy into:
Heat (q): Energy transferred due to temperature difference.
Work (w): Various forms such as mechanical, chemical, electrical, magnetic, or light-induced work.
Internal energy (U) represents the total energy within the system; focus is usually on changes in internal energy (ΔU).

Fundamental Thermodynamic Equation

Change in internal energy equals heat component plus sum of work components:

[ \Delta U = q + \sum w ]

Common work terms in biology include:
Mechanical work (pressure-volume changes)
Chemical work (changes in mole numbers and chemical potential)
Additional work terms (e.g., magnetic) can be added as needed.

State Functions and Variables

Energies like internal energy, enthalpy, and free energy are state functions; only changes matter, not absolute values.
Internal energy depends on entropy (S), volume (V), and number of moles (n).
Enthalpy (H) is defined as:

[ H = U + PV ]

and is useful under constant pressure conditions common in biology. - Under constant pressure, change in enthalpy (ΔH) corresponds directly to heat absorbed or released: - ΔH < 0: exothermic (heat released) - ΔH > 0: endothermic (heat absorbed)

Entropy (S) and Heat

Entropy relates to the number of microscopic configurations (microstates) available to a system.
Boltzmann linked entropy to the logarithm of the number of microstates, connecting macroscopic thermodynamics to molecular statistics.
Entropy is often associated with disorder or randomness but more precisely relates to the probability distribution of microstates.
Entropy tends to increase spontaneously, providing directionality to natural processes (“arrow of time”).
Total entropy of the universe (system + surroundings) always increases in spontaneous processes, even if system entropy decreases.

Free Energy and Spontaneity

Gibbs Free Energy (G) combines enthalpy and entropy effects:

[ G = H - T S ]

G depends on temperature (T), pressure (P), and mole numbers (n).
Biological systems are typically at constant temperature and pressure, making Gibbs free energy the most practical energy function.
Spontaneous processes have ΔG < 0.
Reactions can be:
Enthalpy-driven (bonding and molecular interactions dominate)
Entropy-driven (disorder and configurational freedom dominate)
Or a balance of both, depending on temperature and conditions.

Biological Examples

Protein-ligand binding is analyzed thermodynamically by considering:
Enthalpic changes (bonding interactions)
Entropic changes (loss of configurational freedom and solvent ordering)
The equilibrium of such binding depends on Gibbs free energy.

Utility and Power of Thermodynamics

Thermodynamics simplifies complex biological processes by reducing them to a few key variables.
Experimental and computational methods (e.g., calorimetry) can measure these variables to characterize biological systems.
Understanding thermodynamics helps predict and explain biological behavior at molecular and macroscopic levels.

Future Topics in the Series

Mathematics of thermodynamics
Calorimetry techniques
Computational thermodynamics
Specific biological applications and examples

Detailed Methodology / Key Equations

Energy Partitioning

Change in internal energy:

[ \Delta U = q + w ]

where heat (q) and work (w) contribute to energy changes. - Work can be mechanical (pressure-volume work), chemical (chemical potential times change in mole number), electrical, magnetic, etc.

Enthalpy

Definition:

[ H = U + PV ]

Change in enthalpy:

[ \Delta H = \Delta U + P \Delta V + V \Delta P ]

At constant pressure (ΔP = 0), enthalpy change approximates heat transferred:

[ \Delta H \approx q_p ]

Interpretation:
ΔH < 0: exothermic process (heat released)
ΔH > 0: endothermic process (heat absorbed)

Entropy and Heat

For a reversible process:

[ dQ_{rev} = T dS ]

where ( dQ_{rev} ) is the reversible heat transfer, ( T ) is temperature, and ( dS ) is the change in entropy.

This summary provides a foundational understanding of biological thermodynamics, setting the stage for deeper exploration of thermodynamic principles and their applications in biology.