Thermodynamics and Biochemical Pathways

First Law of Thermodynamics

  • Basic Concept:

  • Energy conservation: Energy cannot be created or destroyed.

  • Changes in internal energy (9;E) can be described by the equation:

    9;E = q + w

  • Where:

    • q = heat transferred
    • w = work done on or by the system

Understanding Heat and Work

  • Heat (q):
  • Energy flow from high to low temperature.
  • Positive heat indicates energy flow into the system.
  • Work (w):
  • Can take forms such as translational, rotational, and vibrational; contributes to changes in internal energy.
  • Positive work means energy input from surroundings.

Enthalpy and Chemical Reactions

  • Enthalpy:
  • 9;H < 0: Exothermic, reaction favored (e.g., H2SO4 + H2O).
  • 9;H > 0: Endothermic, less favored (e.g., KCl + H2O).
  • Reaction changes involve bond energy and changes in molecular motion (translation, vibration).

Second Law of Thermodynamics and Entropy

  • Entropy (S):
  • A measure of disorder; directional predictor of reactions.
  • For any process, total entropy change is greater than zero; indicates spontaneous reactions.
  • Positive change in entropy favors product formation.

Gibbs Free Energy (G)

  • Gibbs Free Energy Equation:
  • 9;G = 9;H - T9;S
  • 9;G < 0: Exergonic, reaction favorable, mainly enthalpy-driven.
  • 9;G > 0: Endergonic, reaction unfavorable, but reverse is exergonic.
  • 9;G = 0: Equilibrium condition.

Standard Free Energy Changes

  • Standard conditions: Temperature of 0°C (273 K) and pressure at 1 atm, with all substances at 1 M concentration.
  • Equation:
  • 9;G = 9;G° + RT ln([C][D]/[A][B]).
  • 9;G° calculated for equilibrium states.

Metabolic Reactions: Near Equilibrium vs. Irreversible

  • Near Equilibrium Reactions:
  • 9;G ~ 0, sensitive to concentration changes, not controlled externally.
  • Metabolically Irreversible Reactions:
  • 9;G << 0, cannot reverse; crucial for metabolic control through allosteric and covalent modifications.

ATP in Metabolism

  • ATP: An energy intermediate critical for metabolic reactions.
  • ATP hydrolysis drives endothermic reactions.
  • Interaction with Mg2+ is essential for stability and function.
  • Energy Coupling:
  • Combined exergonic and endergonic processes create overall exergonic reactions.
  • Example: Aspartate + ATP + NH3 → Asparagine + ADP + Pi3

NADH and Redox Reactions

  • NADH: Main electron carrier in metabolism, crucial for redox reactions.
  • Redox Basics:
    • Oxidation: loss of electrons;
    • Reduction: gain of electrons.
  • Half-Reactions Example:
  • CH3CH2OH → CH3CHO + 2e- + 2H+ (oxidation).
  • NAD+ + 2e- + H+ → NADH (reduction).

Mobile Cofactors and Pathways

  • Energy Intermediates: Include ATP, NADH, and others; link reactions in metabolic pathways.
  • Mobile cofactors are not permanently bound and can facilitate different reactions.
  • Metabolic Pathways:
  • Each product of one reaction serves as a substrate for the next; fundamental in energy transfer and chemical processes.

Conclusion

  • The interconnectedness of energy transformations, thermodynamic laws, and biochemical processes is key to understanding metabolism and energy utilization in living organisms. Each element plays a critical role in sustaining life and facilitating cellular functions.