Gibbs free energy is one of the most important thermodynamic concepts in IB Chemistry. It connects enthalpy, entropy, and temperature into a single value that tells you whether a reaction is spontaneous. Understanding Gibbs free energy allows you to predict reaction behavior, evaluate chemical feasibility, and analyze energy changes with confidence. This article explains the meaning of Gibbs free energy and shows how to use it effectively.
What Is Gibbs Free Energy?
Gibbs free energy (G) is a thermodynamic quantity that measures the amount of usable energy available to do work in a chemical system at constant temperature and pressure.
The key value used in IB Chemistry is ΔG, the change in Gibbs free energy for a reaction.
The core idea:
- ΔG < 0 → reaction is spontaneous
- ΔG > 0 → reaction is non-spontaneous
- ΔG = 0 → system is at equilibrium
Gibbs free energy determines whether a reaction can proceed without continuous energy input.
The Gibbs Free Energy Equation
The IB equation you must know is:
ΔG = ΔH – TΔS
Where:
- ΔG = Gibbs free energy change
- ΔH = enthalpy change
- T = temperature in Kelvin
- ΔS = entropy change
This equation expresses spontaneity as the competition between enthalpy (heat energy) and entropy (disorder).
How ΔH, ΔS, and T Affect ΔG
1. Enthalpy (ΔH)
- Negative ΔH → exothermic → usually favorable
- Positive ΔH → endothermic → may still be favorable at high temperatures
2. Entropy (ΔS)
- Positive ΔS → increased disorder → promotes spontaneity
- Negative ΔS → decreased disorder → opposes spontaneity
3. Temperature (T)
Temperature determines whether the entropy term (TΔS) outweighs the enthalpy term.
The Four Key Spontaneity Cases
IB Chemistry often tests these four scenarios:
1. ΔH negative, ΔS positive
- Always spontaneous
- ΔG is always negative
Examples: combustion reactions
2. ΔH positive, ΔS negative
- Never spontaneous
- ΔG always positive
Examples: decompression of gases in reverse, some precipitation reactions
3. ΔH negative, ΔS negative
- Spontaneous at low temperatures
- Enthalpy helps; entropy opposes
Example: freezing water
4. ΔH positive, ΔS positive
- Spontaneous at high temperatures
- Entropy helps overcome endothermicity
Example: melting ice
Temperature often determines which term dominates.
What Does ΔG Tell Us About a Reaction?
Gibbs free energy answers three essential questions:
1. Will the reaction occur on its own?
If ΔG < 0, the reaction is thermodynamically favorable.
2. How does temperature affect feasibility?
Raising or lowering temperature can change ΔG by affecting TΔS.
3. Is the system at equilibrium?
At equilibrium, ΔG = 0 and no net reaction occurs.
Calculating ΔG Using Standard Values
IB data booklet provides standard Gibbs free energies of formation (ΔGf°).
You can calculate reaction ΔG° using:
ΔG° = ΣΔGf°(products) – ΣΔGf°(reactants)
This method is common in Paper 2 questions.
Why Gibbs Free Energy Is Better Than ΔH Alone
A reaction may be:
- Endothermic (ΔH positive) but still spontaneous
- Exothermic (ΔH negative) but non-spontaneous if entropy decreases enough
ΔG considers both energy and disorder, making it more complete than enthalpy alone.
Real-World Examples
1. Melting Ice at 25°C
- ΔH positive
- ΔS positive
- TΔS > ΔH → ΔG negative → spontaneous
2. Rusting of Iron
- ΔG negative
- Very slow (high activation energy)
Shows spontaneity ≠ speed
3. Photosynthesis
- ΔG positive
- Requires sunlight
Example of a non-spontaneous reaction driven by external energy
FAQs
Does spontaneous mean fast?
No. Thermodynamics determines whether a reaction is possible; kinetics determines how fast it happens.
Can ΔG change with temperature?
Yes. Because of the TΔS term, spontaneity can switch depending on temperature.
What does ΔG = 0 mean?
The system is at equilibrium—no net change occurs.
Conclusion
Gibbs free energy brings together enthalpy, entropy, and temperature to determine whether a chemical reaction is spontaneous. By applying the equation ΔG = ΔH – TΔS, you can analyze reaction feasibility, predict temperature dependence, and better understand the thermodynamic forces that drive chemical change. Mastering this concept is essential for achieving high marks in IB Chemistry.
