Thermodynamics

Thermodynamics Overview

  • Thermodynamics is the study of energy changes associated with physical processes (involving matter).

  • Key Question: "Will a process proceed spontaneously without additional energy?"

  • Important Concept: It does not determine how fast a process will occur.

  • Bioenergetics: Application of thermodynamics to biological systems.

Energy (U)

  • Energy is the ability to cause predictable changes.

  • Sources of energy in biological systems:

    • Chemotrophs: Derive energy from fermentation or oxidation of organic molecules.

    • Phototrophs: Obtain energy from sunlight.

  • Energy (U) is measured in units of joules (J) or kilojoules (kJ).

Systems and Surroundings

  • A system is the part of the universe of interest; surroundings are everything else.

  • Types of systems:

    • Isolated Systems: Cannot exchange matter or energy.

    • Closed Systems: Exchange only energy.

    • Open Systems: Exchange both matter and energy.

  • Thermodynamic parameters are defined in relation to the system, not surroundings.

Thermodynamic States

  • A thermodynamic state defines a condition of a system with constant properties:

    • Pressure (P)

    • Volume (V)

    • Temperature (T)

  • Biological systems often maintain stable P, V, and T, requiring more parameters in bioenergetics.

Thermodynamic State Functions

  • State Function: Defined mathematically, unique to any state, independent of the pathway taken.

  • Important for measuring changes in thermodynamic parameters.

  • Examples include internal energy (U) and enthalpy (H), which show path independence.

Criterion of Spontaneity

  • Goal: to find a reliable criterion (a number) to determine if a process is spontaneous.

  • This explores thermodynamic laws and state functions.

First Law of Thermodynamics: Conservation of Energy

  • Energy in the universe remains constant.

  • Energy can change forms but cannot be created or destroyed.

  • Internal energy change (DU) is equal to heat (q) absorbed or liberated plus work (w) done by or on the system.

  • Note: DU is a state function; q and w are not.

Distinguishing Work vs. Useful Work

  • Heat (q) is thermal energy gained or lost, often measured as absolute temperature.

  • Types of useful work in biological systems include:

    • Synthetic Work: Generating new molecules via chemical bonds.

    • Mechanical Work: Changing position/orientation of cells.

    • Concentration Work: Moving molecules against concentration gradients.

    • Electrical Work: Moving charge species against gradients to establish potential.

    • Work of Heat: Generating heat for maintaining temperature.

Mathematical Values of Heat and Work

  • Heat and work determine DU's sign (+/-):

    • If q > 0: System gains heat (endothermic).

    • If q < 0: System loses heat (exothermic).

    • If w > 0: System does work on surroundings.

    • If w < 0: Surroundings do work on system.

Enthalpy

  • Enthalpy (H) reflects total heat of reaction in a system.

  • Changes in enthalpy are indicated by heat absorbed or released as bonds break/make.

  • Often not a good spontaneity criterion due to positive/negative values.

Second Law of Thermodynamics: Disorder

  • Total entropy of the universe must increase in spontaneous processes.

  • Change in Entropy (DS) relates to heat divided by temperature (K).

  • Entropy reflects molecular arrangements: greater arrangements mean higher entropy.

Gibbs Free Energy

  • Gibbs Free Energy (G) measures work potential and equilibrium distance:

    • DG < 0 (negative): Spontaneous (exergonic).

    • DG > 0 (positive): Non-spontaneous (endergonic).

    • DG = 0: Equilibrium.

  • Units of DG: joules (J) or kilojoules (kJ) per mole.

Gibbs Free Energy Influences

  • The signs of DH and DS influence DG:

    • DH (-) and DS (+): Spontaneous at all temperatures.

    • DH (+) and DS (-): Non-spontaneous at all temperatures.

    • DH (+) and DS (+): Spontaneous at high temperatures.

    • DH (-) and DS (-): Spontaneous at low temperatures.

Standard States and Gibbs Energy

  • Chemical and biochemical standard states define reference conditions:

    • Chemical Standard State: 1 atm, 25 °C, 1 M concentrations.

    • Biochemical Standard State: Similar but with pH = 7.

Thermodynamic Equilibrium vs. Steady State

  • Equilibrium: No net change; forward and reverse rates equal, no work possible.

  • Steady State: Possible net change; concentrations not at equilibrium; allows for useful work.

Cellular Strategies Against Equilibrium

  • Strategies include thermodynamic coupling, enzyme catalysis, and maintaining gradients.

  • Thermodynamic Coupling: Uses high-energy conformations to drive endergonic reactions with exergonic ones.

Key Concepts Summary

  • Understand definitions of thermodynamic terms, useful work, spontaneity, Gibbs Free Energy, and the differences between equilibrium and steady states.

  • Importance of calculating and applying principles of thermodynamics and Gibbs Free Energy in biochemical contexts.