* Equilibrium Favors the Side with Lower Free Energy: Equilibrium is reached when the forward and reverse reaction rates are equal. This doesn't mean the concentrations are equal. Equilibrium favors the side of the reaction with the lower Gibbs free energy (ΔG). This is influenced by factors like:
* Enthalpy Change (ΔH): Exothermic reactions (releasing heat) are favored at equilibrium.
* Entropy Change (ΔS): Reactions that increase disorder (higher entropy) are favored.
* Equilibrium Constant (K): The equilibrium constant (K) quantifies the relative amounts of reactants and products at equilibrium.
* K > 1: The products are favored at equilibrium.
* K < 1: The reactants are favored at equilibrium.
* K = 1: The reactants and products are present in roughly equal amounts.
* Reaction Stoichiometry: The balanced chemical equation dictates the mole ratios of reactants and products. For example, in a reaction like:
```
A + B <=> 2C
```
Even if the equilibrium constant is 1, you won't have 50% A + B and 50% C. You'll have a different distribution to satisfy the 1:1:2 mole ratio.
Example:
Imagine a reaction with an equilibrium constant (K) of 10. This means the products are strongly favored at equilibrium. At equilibrium, you'll likely find a much higher percentage of products than reactants.
In Summary:
The position of equilibrium is determined by the relative free energy of the reactants and products, which is reflected in the equilibrium constant. This often leads to unequal concentrations of reactants and products at equilibrium. The 50/50 split is just a special case that only occurs in specific scenarios with a K value close to 1.
