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Study Notes: chapter 6 Chemical Equilibrium

Chemical Equilibrium

The Equilibrium Constant

  • Equilibrium constant (K) is defined for a general reaction: ( aA + bB \rightleftharpoons cC + dD ) as ( K = \frac{[C]^c [D]^d}{[A]^a [B]^b} ).

  • K is unitless; concentrations are relative to standard states (1 M for solutes, 1 bar for gases).

  • If a reaction's direction is reversed, K is the reciprocal of the original K. When reactions are added, K is the product of their individual K values.

Equilibrium Thermodynamics and Kinetics

  • Enthalpy (ΔH) and entropy (ΔS) affect K.

    • ΔH: heat absorbed (endothermic) or released (exothermic).

    • Standard enthalpy change (ΔH°) occurs under standard conditions.

  • Entropy (ΔS) change: ( \Delta S = \frac{q_{rev}}{T} ). Higher entropy favors spontaneous reactions.

  • Gibbs free energy (ΔG) relates to K: ( K = e^{-\frac{ΔG°}{RT}} ).

Le Chatelier’s Principle

  • If a system at equilibrium is disturbed, it shifts to restore equilibrium. Reaction quotient (Q) can differ from K unless at equilibrium.

Solubility Product

  • Ksp indicates the solubility product for saturated solutions of salts.

  • Common ion effect: solubility decreases when a common ion is added to the system.

Protic Acids and Bases

  • Protic requires H+ transfer. Acids increase [H3O+] and bases decrease [H3O+]. Brønsted-Lowry: acids are proton donors, bases are proton acceptors.

  • Acid-base reactions produce salts, and resulting products can also act as acids or bases (conjugate pairs).

pH and Strengths of Acids/Bases

  • pH defined as ( pH = -\log[H^+] ). In pure water at 25 °C, pH is approximately 7. Acidic solutions have [H+] > [OH−] (pH < 7); basic solutions have the opposite. Acid strength is determined by the extent of dissociation (strong vs. weak).

  • Relationship between Ka (acid dissociation constant) and Kb (base dissociation constant) for conjugate acid-base pairs: ( Ka Kb = K_w ).

Summary of Acid-Base Strengths

  • Strong acids and bases completely dissociate in water, while weak acids and bases partially dissociate.

  • Examples of strong acids include HCl, HBr, and H2SO4; strong bases include NaOH and KOH. Weak acids and bases have low Ka and Kb values, respectively.