Thermodynamic Overview

Introduction to Thermodynamics in Biochemistry

Thermodynamics provides the foundational principles for understanding energy transformations in biological systems. Key concepts include enthalpy, entropy, Gibbs free energy, and the role of coupled reactions in cellular processes.

• Enthalpy (H): The heat content of a system, representing the energy exchanged as heat at constant pressure.

  • Exothermic reactions: Release heat to the surroundings (negative ΔH).

  • Endothermic reactions: Absorb heat from the surroundings (positive ΔH).

• Entropy (S): A measure of disorder or randomness in a system. Systems tend to move toward higher entropy (greater disorder).

• Gibbs Free Energy (G): A thermodynamic quantity that combines enthalpy and entropy to predict the spontaneity of a reaction.

• Predicting Reaction Spontaneity: Gibbs free energy change (ΔG) determines if a reaction will occur spontaneously.

• Coupled Reactions: Unfavorable reactions can be driven by coupling them to favorable ones, a principle essential in metabolism.

Thermodynamic Systems

Definition and Types of Systems

A system is any part of the universe chosen for study, while everything else constitutes the surroundings. Systems are characterized by their composition, temperature, pressure, and volume.

• Closed Systems: Exchange energy (heat or work) but not matter with surroundings. Example: Earth (approximate)

• Open Systems: Exchange both energy and matter with surroundings. Example: Biological systems

• Isolated Systems: Exchange neither energy nor matter with surroundings. Example: The universe as a whole

Living systems typically operate at constant temperature and pressure.

The First Law of Thermodynamics and Enthalpy

Energy Conservation and Heat Exchange

The first law of thermodynamics states that energy is conserved. In closed systems, energy can be exchanged with the surroundings as work or heat.

• Enthalpy (H): The heat exchanged between the system and surroundings at constant pressure, typical of biological conditions.

• Exothermic: Reactions that release heat (ΔH < 0).

• Endothermic: Reactions that absorb heat (ΔH > 0).

Enthalpy as a State Function

Enthalpy depends only on the initial and final states, not the path taken. For example, the breakdown of glucose to CO2 and H2O is exothermic and can be measured by combustion or biochemically; the energy change is the same in both cases.

Glucose Metabolism

Pathways and Energy Yield

Glucose metabolism involves multiple pathways (e.g., glycolysis, citric acid cycle, fermentation), but the total energy generated is the same due to the state function property of enthalpy.

• Example: Glucose breakdown via glycolysis and aerobic respiration both ultimately yield the same energy per mole of glucose.

Reversible and Irreversible Processes

Equilibrium and Process Directionality

• Reversible Processes: Occur near equilibrium, where the system is at its lowest energy state and the forward and reverse rates are equal. Example: Ice melting at 0°C (both ice and water coexist)

• Irreversible Processes: Start far from equilibrium and proceed toward it. Example: Burning paper (irreversible conversion to ash)

The Second Law of Thermodynamics and Entropy

Entropy and Disorder

The second law of thermodynamics states that the entropy of an isolated system tends to increase to a maximum value. Entropy (S) quantifies the degree of disorder or randomness.

• Systems become more disordered as temperature increases.

• The universe naturally progresses toward greater disorder.

• Hydrophobic effect: Protein folding decreases the entropy of the protein but increases the entropy of surrounding water molecules.

Reaction Spontaneity and Gibbs Free Energy

Determining Whether Reactions Occur

Neither enthalpy nor entropy alone can predict reaction spontaneity. Gibbs free energy (G) combines both:

• Gibbs Free Energy Equation:

• Spontaneity Criteria:

  • Exergonic: (spontaneous)

  • Endergonic: (not spontaneous)

  • Equilibrium: (no net reaction)

Effect of Temperature on ΔG

The sign and magnitude of ΔG depend on both ΔH and ΔS, and temperature can shift reaction favorability.

Free Energy, Standard State, and Equilibrium

Standard State and Equilibrium Calculations

• Standard State: Defined as 1 M concentration and 1 atm pressure.

• Relationship to Equilibrium: Where Q is the reaction quotient.

• At equilibrium, and (equilibrium constant):

Biochemical Standard State

Adjustments for Biochemical Reactions

Biochemical reactions often involve water and protons, which are not at 1 M concentration. The biochemical standard state (ΔG°′) accounts for physiological conditions (e.g., pH 7).

• Example: ATP hydrolysis at pH 7.4 and 25°C, with specific concentrations for ATP, ADP, and phosphate.

• Calculation:

Coupling Unfavorable and Favorable Reactions

Driving Cellular Processes

Cells couple unfavorable (endergonic) reactions to favorable (exergonic) ones to drive essential processes such as metabolism, transport, and muscle contraction.

• Example: Unfavorable: , kJ/mol Favorable: , kJ/mol Coupled: , kJ/mol

• As long as the overall ΔG is negative, the coupled reaction is spontaneous.

Summary Table: Types of Thermodynamic Systems

Summary Table: Spontaneity and Gibbs Free Energy

Chapter Summary

Key Principles

• Thermodynamics explains how living systems extract and use energy.

• Spontaneity is determined by the sign of ΔG.

• Living systems operate far from equilibrium, and ΔG calculations must consider actual cellular conditions.

• Unfavorable reactions are driven by coupling to favorable ones.

Additional info: Standard biochemical conditions typically refer to pH 7, 25°C, 1 atm, and 1 M concentrations for solutes except H+ and H2O.