Free energy

  1. Predict the direction of a reaction given the free energy, or the entropy and enthalpy parameters of the reaction.

  2. Predict the direction of a reaction given an equilibrium constant and the initial concentrations of the substance and product of the reaction.

  3. Calculate free energy of a reaction at the standard condition given its equilibrium constant and temperature.

Introduction to Biochemical Reactions

  • Understanding the transformation of biomolecules

  • Key roles of free energy and other factors in driving biochemical reactions

Free Energy in Biochemical Reactions

Definition of Free Energy

  • Free energy: Portion of system energy that can perform work

  • High free energy => unstable state that tends to evolve towards lower free energy

Spontaneous vs. Non-Spontaneous Reactions

Spontaneous Reactions

  • Characterized by:

    • Decrease of free energy

    • Thermodynamically favorable (exergonic)

  • Example: Compressed gas in a container

    • Initially high free energy, unstable

    • When opened, gas disperses, increasing entropy

    • Energy is required to maintain compression

  • In biochemical context:

    • Substrate with high potential energy converts to products with lower potential energy

    • Characterized by negative change in free energy (ΔG)

Non-Spontaneous Reactions

  • Signified by positive ΔG

  • Example: Compressing air into a bicycle tire

    • Requires energy input

    • Result: Products have higher potential energy than substrates

Gibbs Free Energy Equation

  • ΔG = ΔH - TΔS

    • ΔG: Change of free energy

    • ΔH: Change in enthalpy (energy in bonds)

    • ΔS: Change in entropy

  • Conditions:

    • Exothermic: ΔH is negative (heat released)

    • Endothermic: ΔH is positive (heat consumed)

Characterizing Biochemical Reactions

  • Besides spontaneity, three features are important:

    1. Equilibrium constant

    2. Directionality of reaction

    3. Velocity of reaction (to be covered later)

Equilibrium and Equilibrium Constant

  • Equilibrium: No net change in concentrations of substrate (S) and product (P)

  • Continuous interconversion occurs at equal rates

  • Equilibrium constant (K): Ratio of concentrations at equilibrium

    • K = [P]{eq} / [S]{eq}

Directionality of Reactions

  • Many biochemical reactions are reversible

  • Directionality shaped by the Le Chatelier principle:

    • When equilibrium is disturbed, the reaction shifts to counteract the disturbance

  • Example: Conversion between carbon dioxide (CO2) and carbonic acid (H2CO3)

Illustration of Directionality

  • Free energy vs. reaction progress graph:

    • Rightward movement shows CO2 converting to H2CO3

    • At equilibrium, lowest free energy point reached

    • Ratio of carbonic acid to carbon dioxide matches equilibrium constant

    • Spontaneous reaction towards equilibrium: CO2 converts to H2CO3 when excess CO2 is present

    • Conversely, when H2CO3 concentration is high, the reaction may shift left toward CO2.