Introduction to Equilibrium

Introduction

Reversible Reactions: Reactions that can go both forward and reverse directions. In a closed system, reactants of a reversible reaction can form products, and those products can naturally react back into reactants.

Example: A + B ⇌ C + D

Forward Reaction: A + B → C + D
Reverse Reaction: C + D → A + B

Requirements for Reversible Reactions:

Kinetically-Feasible Activation Energies: Reactants and products must have similar and achievable Eₐ. A huge difference/requirement in Eₐ would make the reaction less probable to naturally revert.
Closed System: No products or reactants may escape the system. Precipitates or escaped gases will remove crucial components of the reaction.
Gaseous/Aqueous Phase: At least one reactant must be gaseous or aqueous to facilitate the reaction, as particles in these states have enough kinetic energy to freely move and collide.

What is Chemical Equilibrium?

In a closed system, reversible reactions eventually stabilize in rate.

Equilibrium: The state where the reaction rate of reactants (forward reaction) and products (reverse reaction) equalize.

Microscopically, the reaction still goes on in both directions during equilibrium.
Macroscopically, the concentrations of products and reactants remain constant during equilibrium.

Sneak Peek:

Equilibrium Constant (K)

To find the ratio of products to reactants in a system at equilibrium, we use the equilibrium constant (K).

Basically, K is determined by the division of the amount of products over reactants (in molarity or atm to the power of their coefficients) at equilibrium.

K < 1 implies that the reaction favors the formation of reactants at equilibrium (larger denominator).
K > 1 implies that the reaction favors the formation of products at equilibrium (larger numerator).

This also means that the tendency of a reaction to favor product formation, such as in solubility processes, varies across a spectrum depending on conditions.