Principles of Bioenergetics and Thermodynamics

Principles of Bioenergetics

• Bioenergetics and thermodynamics quantitatively study energy transductions in living cells.
• It explores the nature and function of chemical processes behind these transductions.
• Energy is the ability to do work.
• Chemical reactions are studied from thermodynamics and kinetics perspectives.
• Thermodynamics (bioenergetics) focuses on the energy transformation by analyzing the starting and ending points of reactions.
• Kinetics examines the processes occurring between these points.
• Thermodynamics is the focus.

Laws of Thermodynamics

• The first law (conservation of energy) states the total amount of energy in the universe remains constant for any change noting that energy cannot be created nor destroyed.
• Energy can be changed from one form to another.
• The second law of thermodynamics dictates the universe tends towards increasing disorder.
• Entropy (ΔS) measures how random or disordered a system is, describing how molecules are arranged within a sample.
• Natural systems tend to become more disordered over time, as molecules spread apart.
• Disorder happens spontaneously without needing energy.
• Organizing a system requires energy.

Concept of Free Energy

• Gibbs free energy (ΔG) is the usable energy stored in a molecule's bonds.
• It indicates whether a biochemical reaction will occur spontaneously through energy absorption/release.
• ΔG= ΔH – TΔS, where:
  • ΔG is the free energy in the system (negative in spontaneous reactions).
  • ΔH is enthalpy (heat change).
  • T is temperature.
  • ΔS is entropy (randomness/disorder in the system).
• If ΔG is negative, the reaction releases energy and occurs spontaneously (exergonic).
• If ΔG is positive, the reaction requires energy input to proceed (endergonic).
• When water (H2O) exists as gas, liquid, or solid, energy and disorder differ.
• ΔG stays constant because the molecule is still H2O with the same bonds, yet can change regarding temperature and pressure.
• Enthalpy (ΔH) varies andis lowest in solid state (ice), higher in liquid, and highest in gas because of the energy in molecular bonds.
• Entropy (ΔS) measures disorder and is lowest in the solid state (more ordered), higher in liquid, and highest in gas.
• Each state has different ΔH and ΔS, unlike ΔG
• By rearranging the equation to ΔH = ΔG + TΔS, it is inferred that ΔH is the total energy of the system.
• Thermodynamics indicates if a reaction is possible.

Chemical Reactions

• Chemical reactions occur to reach stability with the lowest energy amount.
• Logically, biochemical reactions go from high to lower energy scales.
• Exergonic reactions:
  • ΔG is negative.
  • They are spontaneous.
  • They are favorable.
• Endergonic reactions:
  • ΔG is positive.
  • They are non-spontaneous.
  • They are non-favorable.
• Reactions need added energy to occur, like adding ATP.

Standard ΔG

• ΔGo represents ΔG at standard conditions: 1 atm, 1 M concentrations of reactants/products, 25°C, and pH=7. It compares different reactions in the laboratory
• Altering concentrations will change the ΔG

State Function

• ΔG is a state function, meaning it relies only on the starting and ending points (concentrations) not the reaction mechanism or intermediates.

Glucose

• Combustion of glucose in a calorimeter:
  • Glucose + O2 → CO2 + H2O.
  • ΔG = -680 kcal/mol.
• Glycolysis and Krebs cycle are complex, multi-step processes that do not affect the overall ΔG.
  • Glucose → CO2 + H2O. ΔG = -680 kcal/mol.
• The pathway order does not matter, so ΔG is still the same.

Concentration Impact

• ΔG calculation in the reaction Glucose → Glucose 6-phosphate uses the formula ΔG = Gf - Gi.
• Gf is free energy of final product (glucose 6-phosphate).
• Gi is free energy of initial reactant (glucose).
• Altering glucose concentration changes ΔG because concentration affects molecular number/bonds.

Reaction Spontaneity

• ΔG determines reaction spontaneity.
• Adjusting reactant/product concentrations (e.g., glucose/glucose 6-phosphate) can make the reaction spontaneous (negative ΔG) or non-spontaneous (positive ΔG).
• Changing concentrations influences if reactions occurs naturally or needs energy input
• Equilibrium happens when the energy of reactants equals the energy of products (not concentrations).

Relationship of ΔGo and ΔG

• ΔG = ΔGo + RT ln ([Products]/[Reactants]).
• ΔG = ΔGo + RT 2.3 log ([Products]/[Reactants]).
• Temperature and atmosphere are constant, so concentrations are the only variables.
• When [products] < [reactants], then ΔG < ΔGo.
• When [products] > [reactants], then ΔG > ΔGo.
• When [products] = [reactants], then ΔG = ΔGo.
• Equilibrium occurs at the end.
• At equilibrium concentrations are fixed; rates are equal.
• Concentrations define the equilibrium constant.
• Altering concentrations modifies ΔG.
• It is possible to make a nonspontaneous reaction spontaneous by concentration adjustments.
• Increase reactants or decrease can result a more negative value of ΔG.

Equilibrium Consideration

• At equilibrium, ΔG is zero (ln 1 = 0).
• You can have an unfavourable situation and still make it favorable by manipulating the conditions.
  • Non-standard conditions drive the reaction in the required direction.

The End