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.