H = U + PV
While enthalpy itself doesn't directly determine the spontaneity of a reaction, its change (ΔH) plays a crucial role. Here's how:
1. Exothermic Reactions (ΔH < 0):
* Release heat to the surroundings.
* Favorable in terms of enthalpy, as the system loses energy, making it more stable.
* However, not always spontaneous, as other factors like entropy can influence the process.
2. Endothermic Reactions (ΔH > 0):
* Absorb heat from the surroundings.
* Unfavorable in terms of enthalpy, as the system gains energy, making it less stable.
* Usually not spontaneous, requiring external energy input to proceed.
Gibbs Free Energy (G):
To accurately predict spontaneity, we need to consider both enthalpy change (ΔH) and entropy change (ΔS) using Gibbs free energy (G):
ΔG = ΔH - TΔS
* ΔS > 0: Increased disorder or randomness in the system, generally favorable.
* T: Temperature in Kelvin.
Spontaneity and Gibbs Free Energy:
* ΔG < 0: Reaction is spontaneous (favorable) under given conditions.
* ΔG > 0: Reaction is non-spontaneous (unfavorable) under given conditions.
* ΔG = 0: Reaction is at equilibrium, where forward and reverse rates are equal.
In Summary:
Enthalpy change alone does not guarantee spontaneity. Gibbs free energy, incorporating both enthalpy and entropy, is the ultimate indicator of whether a reaction will proceed spontaneously under specific conditions.
