Concept 8.2 (The free-energy change of a reaction tells us whether or not the reaction occurs spontaneously)

Free Energy and Spontaneity of Reactions

  • Free Energy Change (ΔG): Indicates whether a reaction occurs spontaneously.

    • Spontaneous reactions have a negative ΔG, meaning they are energetically favorable without external energy.

Thermodynamics in Biological Systems

  • Laws of Thermodynamics: Apply universally, helping biologists understand chemical reactions in life.

    • Key to assess energy and entropy changes for reactions.

Gibbs Free Energy

  • J. Willard Gibbs (1878): Introduced Gibbs free energy (G) to describe energy without considering surroundings.

    • Free energy (G): Portion of a system's energy available to do work under uniform temperature and pressure.

Calculating Free Energy Change

  • Equation for ΔG: ΔG = ΔH - TΔS

    • ΔH: Change in enthalpy (total energy).

    • ΔS: Change in entropy.

    • T: Absolute temperature in Kelvin.

  • Factors affecting ΔG: pH, temperature, concentrations of reactants/products.

Predicting Spontaneity with ΔG

  • Negative ΔG: Indicates spontaneity; occurs if:

    • ΔH is negative (loss of enthalpy) and/or ΔS is positive (increase in entropy).

    • Consequently, all spontaneous processes decrease free energy.

  • Positive or Zero ΔG: These processes are non-spontaneous.

Stability and Free Energy

  • Free energy represents a system's instability; systems tend to move towards lower free energy (greater stability).

    • Examples:

      • Diver on a platform vs. floating in water.

      • Concentrated dye spreading in liquid.

      • Glucose molecule breakdown.

Equilibrium in Chemical Reactions

  • Chemical Equilibrium: When forward and reverse reactions occur at the same rate.

    • At equilibrium, free energy is at its lowest, denoting maximum stability.

    • ΔG is zero, indicating no spontaneous change.

  • Disruptions in equilibrium lead to increases in free energy, making systems less stable.

Free Energy in Metabolic Reactions

  • Exergonic Reactions: Release free energy (ΔG < 0); spontaneous and can perform work.

    • Example: Cellular respiration—glucose decomposition.

    • For each mole of glucose, ΔG = -2,870 kJ/mol.

  • Endergonic Reactions: Absorb free energy (ΔG > 0); non-spontaneous reactions that require external energy input.

    • Example of reversal: Converting CO2 and H2O back into glucose is endergonic.

Metabolism and Work Performance

  • Living Systems: Never reach metabolic equilibrium due to constant influx and efflux of materials.

  • Processes keep moving towards stability, preventing equilibrium which would halt work in cells.

    • Cellular respiration illustrates a series of reactions maintaining energy flow and avoiding equilibrium.

Energy Sources in Ecosystems

  • Ecosystems rely on sunlight for free energy, captured by photosynthetic organisms.

  • Non-photosynthetic organisms depend on consuming organic products of photosynthesis for energy.

Concept Check Questions

  1. Cellular Respiration: Is it spontaneous and exergonic?

    • Yes, it uses glucose and O2 (high free energy) to produce CO2 and H2O (low free energy).

  2. Relation to Catabolism and Anabolism: Consider their roles in energy transfer and entropy changes.

  3. Chemiluminescent Necklace: Is the reaction exergonic or endergonic?

    • Requires analysis of energy release upon activation.