Laws of thermodynamics in living systems Knowledge-Based Study

Study notes

1. Thermodynamic Systems in Biology

   Thermodynamics is the study of energy transformations that occur in a collection of matter. In biological contexts, organisms must constantly acquire, convert, and expend energy to sustain cellular processes and avoid decay.

   System vs. Surroundings

   To analyze biological energy transfers, scientists classify physical systems into three distinct categories:

   • Isolated systems: Cannot exchange either matter or energy with their surroundings.
   • Closed systems: Can exchange energy, but not matter, with their surroundings.
   • Open systems: Can freely exchange both matter and energy with their surroundings.

   Living organisms are fundamentally open systems. Cells absorb energy-rich organic molecules or radiant energy from sunlight and release metabolic byproducts such as carbon dioxide, water, and heat back into the environment. Because cells constantly interact with their surroundings, they do not reach static thermodynamic equilibrium while alive.

   The Principle of Energy Conservation

   The First Law of Thermodynamics, also known as the law of conservation of energy, states that energy can be transferred and transformed, but it cannot be created or destroyed. The total amount of energy in the universe remains constant.

   Energy Transductions in Living Systems

   Biological entities operate by continually converting one form of energy into another to perform biological work:

   • Photosynthesis: Photoautotrophs capture radiant light energy and transform it into chemical potential energy stored in the covalent bonds of carbohydrates.
   • Cellular Respiration: Heterotrophs and autotrophs extract that chemical potential energy from glucose and other fuels, storing usable portions in adenosine triphosphate (ATP).
   • Mechanical and Transport Work: Cells convert the chemical energy of ATP into mechanical movement (e.g., muscle contraction, vesicle trafficking) or electrical gradients across membranes.

   During every transformation, the total quantitative energy before and after the reaction is conserved; none vanishes, though some becomes unavailable to do work.

   Entropy and Spontaneous Processes

   The Second Law of Thermodynamics states that every energy transfer or transformation increases the entropy ($S$) of the universe. Entropy is a measure of molecular disorder or randomness. In physical terms, energy spontaneously disperses from localized, ordered states to more widely distributed, disordered states.

   Heat as Inevitable Waste

   No biological energy conversion is 100% efficient. During any biochemical reaction, a portion of the system's usable energy is inevitably converted to thermal energy (heat) and dissipated into the surroundings. Because biological cells are generally isothermal (operating at relatively uniform temperatures), this released heat cannot be harnessed to drive cellular work, thereby contributing to the entropy increase of the environment.

   How Living Things Maintain Complex Order

   Organisms are highly ordered structures, which might seem to contradict the Second Law. However, the Second Law applies to the universe (the system plus its surroundings), not just an open system alone:

   • A cell can locally decrease its internal entropy by synthesizing complex polymers (proteins, nucleic acids) and maintaining organelles.
   • To achieve this, the cell releases heat and metabolic waste, producing an entropy increase in the surroundings that exceeds the local decrease in entropy within the organism.
   • Thus, biological order is preserved strictly at the expense of increasing the overall entropy of the universe.

   Gibbs Free Energy and Reaction Spontaneity

   To determine whether a chemical reaction occurs spontaneously without an external input of energy, biochemists use Gibbs free energy ($G$). The change in free energy ($\Delta G$) reflects the portion of a system's energy capable of performing work at constant temperature and pressure:

   $$\Delta G = \Delta H - T\Delta S$$

   Where $\Delta H$ represents change in enthalpy (total energy), $T$ is absolute temperature in Kelvin, and $\Delta S$ is change in entropy.

   Exergonic vs. Endergonic Reactions

   • Exergonic reactions ($\Delta G < 0$): Release free energy spontaneously. Catabolic pathways, such as the breakdown of glucose, are predominantly exergonic.
   • Endergonic reactions ($\Delta G > 0$): Absorb free energy and require an energy input to proceed. Anabolic pathways, such as protein synthesis from amino acids, are non-spontaneous and endergonic.

   Energy Coupling via ATP

   Living cells drive essential endergonic processes through energy coupling, using an exergonic reaction to power an endergonic one. The primary mediator of cellular energy coupling is the hydrolysis of ATP:

   $$\text{ATP} + \text{H}_2\text{O} ightarrow \text{ADP} + \text{P}_i \quad (\Delta G^\circ = -30.5\text{ kJ/mol})$$

   Because the overall free-energy change of coupled reactions remains negative ($\Delta G_{net} < 0$), cells successfully drive vital biosynthetic, mechanical, and transport processes while strictly adhering to thermodynamic laws.