EQUILIBRIUM
Key Definitions and Concepts
• Dynamic equilibrium: Rates of forward and reverse processes are equal; macroscopic
properties remain constant though microscopic change continues.
• Equilibrium mixture: Coexisting reactants and products at equilibrium.
• Law of mass action: For a balanced reaction aA+bB ⇌ cC+dD, Kc = [C]^c [D]^d /
[A]^a [B]^b (concentrations at equilibrium).
• Reaction quotient Q: Same form as K but using instantaneous concentrations;
compare Q and K to predict direction:
o Q < K → reaction proceeds forward (to products).
o Q > K → reaction proceeds reverse (to reactants).
o Q = K → system at equilibrium.
• Homogeneous vs heterogeneous equilibrium:
o Homogeneous: all species same phase (e.g., gases).
o Heterogeneous: multiple phases; pure solids/liquids omitted from K
expression.
• Dynamic demonstrations: isotope scrambling (H/D) shows forward and reverse
reactions continue at equilibrium.
Equilibrium Constants and Relations
• Kc: equilibrium constant in concentration units (mol L–1).
• Kp: equilibrium constant in terms of partial pressures. Relation:
o Kp = Kc (RT)^(Δn) where Δn = moles gaseous products − moles gaseous
reactants; R = 0.0831 bar·L·mol–1·K–1.
• Changing equation stoichiometry:
o Reverse reaction: K(rev) = 1 / K(forward).
o Multiplying equation by n: K(new) = K(old)^n.
• Temperature dependence: K is temperature dependent; not affected by catalysts or
initial concentrations.
Table: Equilibrium Constant Relations
Equation Change Effect on K
Reverse reaction K' = 1 / K
Multiply balanced equation by n K' = K^n
Use partial pressures (gases) Kp = Kc (RT)^(Δn)
Factors Affecting Equilibrium (Le Chatelier’s Principle)
, • Concentration: addition shifts equilibrium to consume added species; removal shifts
to produce it.
• Pressure/Volume (gaseous systems): increasing pressure (decreasing volume) shifts
toward side with fewer moles of gas; opposite for decreasing pressure.
• Temperature:
o Exothermic reaction (ΔH < 0): increasing T shifts equilibrium left (K
decreases).
o Endothermic reaction (ΔH > 0): increasing T shifts equilibrium right (K
increases).
• Inert gas at constant volume: no effect on equilibrium (partial pressures of
reactants/products unchanged).
• Catalyst: speeds both forward and reverse rates equally; does not change equilibrium
position or K.
Homogeneous Gaseous Equilibria — Kp and Kc Example
• For H2(g) + I2(g) ⇌ 2HI(g):
o Kc = [HI]^2 / ([H2][I2])
o Kp = (pHI)^2 / (pH2 pI2)
o If Δn = 0, Kp = Kc.
• Use ideal gas relation p = [gas] RT to convert between Kp and Kc.
Ionic Equilibrium and Electrolytes
• Electrolytes: substances whose aqueous solutions conduct electricity (ions present).
• Strong electrolytes: nearly complete ionization (e.g., HCl, NaOH).
• Weak electrolytes: partial ionization establishing equilibrium between ions and
molecules (e.g., CH3COOH, NH3).
Acid–Base Theories
• Arrhenius:
o Acid → produces H+ (or H3O+) in water.
o Base → produces OH– in water.
• Brønsted–Lowry:
o Acid → proton donor.
o Base → proton acceptor.
o Conjugate acid-base pair differ by one proton.
• Lewis:
o Acid → electron-pair acceptor.
o Base → electron-pair donor.
• Conjugate relation: Ka × Kb = Kw (for conjugate acid–base pair).
Ionization Constants and p-Scale
• Acid dissociation (weak acid HA):
o HA + H2O ⇌ H3O+ + A–
o Ka = [H+][A–] / [HA]
Key Definitions and Concepts
• Dynamic equilibrium: Rates of forward and reverse processes are equal; macroscopic
properties remain constant though microscopic change continues.
• Equilibrium mixture: Coexisting reactants and products at equilibrium.
• Law of mass action: For a balanced reaction aA+bB ⇌ cC+dD, Kc = [C]^c [D]^d /
[A]^a [B]^b (concentrations at equilibrium).
• Reaction quotient Q: Same form as K but using instantaneous concentrations;
compare Q and K to predict direction:
o Q < K → reaction proceeds forward (to products).
o Q > K → reaction proceeds reverse (to reactants).
o Q = K → system at equilibrium.
• Homogeneous vs heterogeneous equilibrium:
o Homogeneous: all species same phase (e.g., gases).
o Heterogeneous: multiple phases; pure solids/liquids omitted from K
expression.
• Dynamic demonstrations: isotope scrambling (H/D) shows forward and reverse
reactions continue at equilibrium.
Equilibrium Constants and Relations
• Kc: equilibrium constant in concentration units (mol L–1).
• Kp: equilibrium constant in terms of partial pressures. Relation:
o Kp = Kc (RT)^(Δn) where Δn = moles gaseous products − moles gaseous
reactants; R = 0.0831 bar·L·mol–1·K–1.
• Changing equation stoichiometry:
o Reverse reaction: K(rev) = 1 / K(forward).
o Multiplying equation by n: K(new) = K(old)^n.
• Temperature dependence: K is temperature dependent; not affected by catalysts or
initial concentrations.
Table: Equilibrium Constant Relations
Equation Change Effect on K
Reverse reaction K' = 1 / K
Multiply balanced equation by n K' = K^n
Use partial pressures (gases) Kp = Kc (RT)^(Δn)
Factors Affecting Equilibrium (Le Chatelier’s Principle)
, • Concentration: addition shifts equilibrium to consume added species; removal shifts
to produce it.
• Pressure/Volume (gaseous systems): increasing pressure (decreasing volume) shifts
toward side with fewer moles of gas; opposite for decreasing pressure.
• Temperature:
o Exothermic reaction (ΔH < 0): increasing T shifts equilibrium left (K
decreases).
o Endothermic reaction (ΔH > 0): increasing T shifts equilibrium right (K
increases).
• Inert gas at constant volume: no effect on equilibrium (partial pressures of
reactants/products unchanged).
• Catalyst: speeds both forward and reverse rates equally; does not change equilibrium
position or K.
Homogeneous Gaseous Equilibria — Kp and Kc Example
• For H2(g) + I2(g) ⇌ 2HI(g):
o Kc = [HI]^2 / ([H2][I2])
o Kp = (pHI)^2 / (pH2 pI2)
o If Δn = 0, Kp = Kc.
• Use ideal gas relation p = [gas] RT to convert between Kp and Kc.
Ionic Equilibrium and Electrolytes
• Electrolytes: substances whose aqueous solutions conduct electricity (ions present).
• Strong electrolytes: nearly complete ionization (e.g., HCl, NaOH).
• Weak electrolytes: partial ionization establishing equilibrium between ions and
molecules (e.g., CH3COOH, NH3).
Acid–Base Theories
• Arrhenius:
o Acid → produces H+ (or H3O+) in water.
o Base → produces OH– in water.
• Brønsted–Lowry:
o Acid → proton donor.
o Base → proton acceptor.
o Conjugate acid-base pair differ by one proton.
• Lewis:
o Acid → electron-pair acceptor.
o Base → electron-pair donor.
• Conjugate relation: Ka × Kb = Kw (for conjugate acid–base pair).
Ionization Constants and p-Scale
• Acid dissociation (weak acid HA):
o HA + H2O ⇌ H3O+ + A–
o Ka = [H+][A–] / [HA]
