Equilibrium

Concept of Equilibrium

Everyday Meaning of Equilibrium:
- Describes a balance or state of no change.
- Example: Two teams tied in a tug of war.
Chemical Perspective:
- Confusion arises as reactions may seem to reach a balanced state contrary to completion principles learned in general chemistry.
- Traditional understanding: Chemical reactions go to completion.

Reversible Reactions:
- Bio chemical reactions can progress both forward (reactants to products) and backward (products reforming reactants).
- Most chemical reactions are reversible, except those treated as irreversible in stoichiometry problems.
Example:
- Haber Process:
  - Reaction: N2(g) + 3H2(g) ⇌ 2NH3(g)
  - Indicates the reaction is reversible.
  - Begins with reactants only, as NH3 is absent initially.
  - As NH3 forms, its concentration increases, causing the reverse reaction rate (NH3 to reactants) to increase.
  - Eventually, forward and reverse rates equalize – this is termed equilibrium.

Characterized by:
- Equal rates of forward and reverse reactions but not stopped entirely.
- Reactant and product concentrations may not be equal, but they remain constant over time.

Thermodynamics & Reaction Spontaneity:
- Spontaneous reactions favor products (high product concentrations at equilibrium).
- Non-spontaneous reactions favor reactants (high reactant concentrations staying).
- At equilibrium, entropy is maximized, and ΔG = 0.
Kinetics vs. Thermodynamics:
- Important to differentiate between reaction speed (kinetics) and spontaneity (thermodynamics).

K_eq Expression:
- For a generic reaction:
- K_eq = [products]^coefficients / [reactants]^coefficients.
Haber Process Example:
- Reiterates how K_eq is determined by reactant and product concentrations.
Deriving K_eq:
- Based on the rates of forward and reverse reactions, K_eq can be defined.
- Example Rate Relationships:
  - Rate_forward = k_f[N2][H2]^3; Rate_reverse = k_r[NH3]^2.
  - Equal rates yields relation: k_f / k_r = [NH3]^2 / ([N2][H2]^3) → Leading to K_eq.

If K_eq > 1, products favored.
If K_eq < 1, reactants favored.
K_eq cannot be negative; must be positive values between 0 and 1,
K_eq changes with temperature but not with concentration changes or catalysts.
Example: K_eq for H2O and CO2 forming H2CO3 excludes water due to being a pure liquid.

Q vs. K_eq:
- Q same expression as K_eq but uses concentrations from any point in the reaction.
- Q < K_eq → Reaction shifts toward products to reach equilibrium.
- Q > K_eq → Reaction shifts toward reactants.
- Q = K_eq → Already at equilibrium.

Stress on Equilibrium:
- When equilibrium disturbed (change in concentration, temperature, pressure, or volume), reaction shifts to relieve stress.
- Example: Adding SO2 to SO2 + O2 ⇌ SO3 shifts right to produce more SO3.
- Removing products can increase yield as it drives reaction toward product side.

Adding a reactant shifts toward products; adding a product shifts toward reactants.
Removing reactants shifts toward reactants; removing products shifts toward products.

Increasing volume = decreasing pressure (shifts toward side with more moles of gas).
Decreasing volume = increasing pressure (shifts toward side with fewer moles of gas).
Only count moles of gas: solids and liquids do not affect pressure/volume shifts.

Temperature viewed as a reactant (endothermic) or product (exothermic).
- Increasing temperature shifts toward products in endothermic reactions; decreases yield in exothermic.

Catalysts:
- Speed up reactions but do not shift equilibrium.
Inert Gas Addition:
- Does not change partial pressures of equilibrating substances; no shift in equilibrium.