Revision sheet: Bioenergetics and Thermodynamics in Biology

Bioenergetics Revision Sheet

1. 📌 Essentials

• Life relies on extraction, transformation, and dissipation within cells- Systems: open (exchange matter/energy), closed (energy only), isolated (no exchange).
• First law: energy conserved; ΔU = Q + W.
• Enthalpy (H): heat exchange at constant pressure; ΔH > 0 endothermic, ΔH < 0 exothermic.
• Spontaneous reactions: ΔG < 0; depend on ΔH and ΔS.
• Entropy (S): measure of disorder; increases in spontaneous processes.
• Gibbs free energy (G): ΔG = ΔH - T ΔS; determines reaction spontaneity.
• Equilibrium: ΔG = 0; ΔG° relates to equilibrium constant K.
• Redox reactions: electron transfer; characterized by reduction potentials (E°).
• Energy forms: caloric, electrical, chemical, nuclear, mechanical.
• Biological systems are open thermodynamic systems constantly exchanging energy and matter.

2. 🧩 Key Structures & Components

• Cells — basic units performing energy exchange and metabolic reactions.
• ATP — primary energy currency for cellular work.
• Redox couples (e.g., NAD/NADH) — facilitate electron transfer.
• Metabolic pathways — series of enzyme-catalyzed reactions transforming energy.
• Thermodynamic variables — temperature, pressure, volume, entropy.
• Reaction intermediates — molecules formed during biochemical reactions.
• Enzymes — catalyze reactions, influence energy barriers.
• Energy transfer molecules — ATP, NADH, FADH2.
• Heat — energy dispersal increasing entropy.
• Standard conditions — 1 atm, 25°C, 1 M solutions.

3. 🔬 Functions, Mechanisms & Relationships

• Cells extract energy via oxidation of organic molecules or light absorption.
• Energy transformations follow thermodynamic laws, with some energy lost as heat.
• ΔU (internal energy) changes with heat (Q) added and work (W) done.
• Enthalpy (H) reflects heat exchange at constant pressure; ΔH indicates reaction heat flow.
• Reactions tend toward lower free energy (ΔG < 0) and higher entropy (ΔS).
• Spontaneity depends on the balance of ΔH and TΔS.
• Redox reactions transfer electrons; potentials (E°) determine energy yield.
• Coupling reactions (e.g., ATP hydrolysis) enable energetically unfavorable processes.
• Reaction directionality is driven by ΔG; at equilibrium, ΔG = 0.
• Energy flow in pathways is hierarchical: substrate → intermediates → products.
• The universe's entropy increases; biological systems maintain order locally by increasing surroundings' entropy.

4. Comparative Table

Item	Key Features	Notes / Differences
System Type	Open: exchanges matter & energy; Closed: energy only; Isolated: none	Biological systems are open
Variables	Extensive: volume, energy; Intensive: temperature, concentration	State-dependent
Transformation Types	Isothermal, isobaric, isochore, adiabatic	Thermodynamic processes
Enthalpy (ΔH)	Heat exchanged at constant pressure	ΔH > 0 endothermic, ΔH < 0 exothermic
Standard ΔH (ΔH°)	At 1 atm, 25°C, 1 M solutions	Reference conditions
Hess's Law	Total ΔH = sum of ΔH of steps	Allows indirect calculation
Glucose combustion	ΔH° = -2808 kJ/mol	Major metabolic energy release
ΔG (Gibbs free energy)	ΔG = ΔH - T ΔS	Determines spontaneity
Equilibrium constant (K)	ΔG° = -RT ln K	Relationship with ΔG°
Redox potentials (E°)	Electron transfer drives energy release	NAD/NADH as common pair
ΔG for redox reactions	ΔG = -nFΔE	n: electrons transferred, F: Faraday

5. 🗂️ Hierarchical Diagram

Bioenergetics
 ├─ Energy Sources
 │   ├─ Phototrophs: light → chemical energy
 │   └─ Chemotrophs: oxidation of organics
 ├─ Thermodynamics
 │   ├─ Systems: open, closed, isolated
 │   ├─ Variables: extensive, intensive
 │   └─ Transformations: isothermal, adiabatic, etc.
 ├─ Equilibrium & Reversibility
 │   ├─ Equilibrium: static/dynamic
 │   └─ Reversible: infinitesimal steps
 ├─ First Law
 │   └─ ΔU = Q + W
 ├─ Enthalpy & Heat
 │   ├─ ΔH: heat at constant P
 │   └─ ΔH reaction: ΔH(products) - ΔH(reactants)
 ├─ Standard & Hess
 │   ├─ ΔH°, ΔG°
 │   └─ Hess's Law
 ├─ Reaction energetics
 │   ├─ Spontaneous if ΔG < 0
 │   └─ ΔG = ΔH - T ΔS
 ├─ Entropy
 │   ├─ Measure of disorder
 │   └─ ΔS = ΔQrev / T
 ├─ Second Law
 │   └─ Universe entropy increases
 ├─ Gibbs Free Energy
 │   ├─ ΔG = ΔH - T ΔS
 │   └─ Equilibrium when ΔG = 0
 └─ Redox & Coupling
     ├─ Electron transfer: Aox/Ared
     └─ ΔG = -nFΔE

6. ⚠️ High-Yield Pitfalls & Confusions

• Confusing ΔH (enthalpy) with ΔU (internal energy); ΔH includes PV work.
• Assuming all exergonic reactions are spontaneous without considering entropy.
• Overlooking the role of entropy (ΔS) in reaction spontaneity.
• Misinterpreting ΔG as only dependent on ΔH; temperature and entropy are critical.
• Forgetting that ΔG° relates to equilibrium constant K, but actual ΔG depends on reaction quotient Q.
• Confusing standard conditions (ΔG°, ΔH°) with actual reaction conditions.
• Ignoring the importance of coupling reactions for energetically unfavorable processes.
• Misjudging redox potentials: higher E° means more positive reduction potential.
• Overestimating the energy yield of redox reactions without considering electron transfer number (n).

7. ✅ Final Exam Checklist

• Understand the different system types and their relevance to biology.
• Know the first law: ΔU = Q + W, and how it applies to biochemical reactions.
• Distinguish between ΔH, ΔS, and ΔG; their signs and implications.
• Be able to calculate ΔH and ΔG for reactions, including standard conditions.
• Comprehend Hess's Law for indirect enthalpy calculations.
• Recognize the significance of ΔG in reaction spontaneity and equilibrium.
• Understand how energy coupling via ATP or redox reactions drives cellular processes.
• Know the basics of redox reactions, electron transfer, and reduction potentials.
• Be familiar with the thermodynamic principles governing metabolic pathways.
• Recognize the importance of entropy and the second law in biological systems.
• Be able to interpret energy diagrams, reaction coordinate graphs, and hierarchical structures.

End of Revision Sheet