Lesson 6: Acid-Base Equilibrium

• Kw

  Constant when Ka and Kb are multiplied
• Do not include spectator ions, if needed, do a dissociation equation before writing the reaction equation

  
• Base Ionization Constant (Kb)

  B + H2O ↔ BH+ + OH-

  Kb = [BH+] (OH-]

  [B]

  Larger Kb = stronger base

  smaller Kb = weaker base

  BH+ is the conjugate acid
• Ion-Product Constant for Water (Kw)

  Autoionization = the transfer of one H+ ion from one water molecule to another

  H2O (l) + H2O (l) ↔ OH- (aq) + H3O+

  Kw = [OH-] [H3O+]

  @25℃ Kw = 1 x 10-14 mol/L→ [OH-] = 1.0 x 10-7 mol/L; [H+] = 1.0 x 10-7 mol/L

  Acidic - [H3O+] > [OH-]; neutral - [H3O+] = [OH-]; basic - [H3O+] < [OH-]

  
• water reacts with itself, forming OH- and H3O+
• if hydronium concentration is greater than hydroxide, it is acidic
• if hydronium concentration is equal to hydroxide, it is neutral

  
• if hydronium concentration is less than hydroxide, it is basic

  
• pH

  the negative logarithm of the concentration of hydrogen ions in solution

  
• pOH

  the negative logarithm of the concentration of hydroxide ions in solution
• pH = -log[H+]

  pOH = -log[OH-]
• the pH is usually around the range of the positive version of the exponent
• units for Ka and Kb: mol/L
• [OH-] = 10^-pOH
• the equilibrium position for a strong acid favors the products (shifts right)

  the equilibrium position for a weak acid favors the reactants (shifts left)