Chapter 6 - Acid/Base Equilibria

• Acid/Base theories
• Arrhenius
• Bronsted - Lowry
• Lewis
• Solvern/System
• pH and pOH
• -log = p function
• based on activities of species (pH = -log aH+)
• For 'working' definition assume activities = concentrations
• -log[H+] = pH
• -log[OH-] = pOH
• -logKa = pKa
• neutral - pH = 7.00 ([H+] = 1.0 x 10-14)
• pH + pOH = 14.00
• pH < 7, acidic
• pH > 7, basic
• [H+] = 10^-pH
• If acidic pH greater than 6.5 OR basic pH less than 7.5 we must consider H+ or OH- from water!!
• Calculation of pH
• Strong acid
• assume complete dissociation, [H+] = [strong acid]
• Strong base - calculate hydroxide ion concentration from formula of salt
• soluble hydroxide - 1 OH- for every hydroxide
• soluble oxide - 2 OH- for every O2-
• If temp different from 25C, Kw may change - change 'neutral' pH.

• Dissociation of weak acids and bases
• HA H+ + A-
• Ka =[H+][A-]/[HA]
• Ka - acid dissociation constant
• B + H2O HB+ + OH-
• Kb = [HB+][OH-] / [B]
• Kb - basic equilibrium constant (hydrolysis constant)
• NOTE - B can be 1 of 2 types of base
• 'Molecular' base - look up Kb in table
• Conjugate base of weak acid
• Kb = Kw / Ka
• note OH- comes from reaction with water.
• Calculation of pH
• Acid
• Strong acid
• assume complete dissociation, [H+] = [strong acid]
• Weak acid
• Calculate CHA / Ka (CHA = original concentration of acid)
• If CHA / Ka 100, then [H+] = the square root of (Ka*CHA)
• If CHA / Ka < 100, then use the equation based on the quadratic formula
• Base
• Strong base - calculate hydroxide ion concentration from formula of salt
• soluble hydroxide - 1 OH- for every hydroxide
• soluble oxide - 2 OH- for every O2-
• Weak base - Calculate CB / Kb (CB = original concentration of base)
• If CB / Kb 100, then [OH-] = SQRT(Kb*Cb)
• If CB / Kb < 100, then use formula based on the quadratic equation
• Conjugate acids/bases
• Calculate "missing" K value from Kw = Ka * Kb
• Treat like any other weak a/b

• Common Ion Effect
• Common ion - common to the equilibrium system
• causes a shift away from the ionic side of the equilibrium
• For HA H+ + A-
• Addition of H+ or A- causes shift back to HA
• 'Suppresses ionization'
• Buffers
• Typically mixture of weak conjugate a/b pair
• Weak acid reacts with strong base
• Weak base reacts with strong acid
• prepared by:
• addition of weak acid and soluble salt of the weak acid
• HA + NaA (NaA -> Na+ + A-)
• partial neutralization of weak acid with strong base
• base is limiting reactant
• reverse is true for weak base / conj. acid systems
• Resistant to a change in pH
• Derive Henderson-Hasselbalch eqn
• pH dependent on ratio of base to acid

• Buffer capacity - amount of material to get ratio of base/acid outside limits
• 10 base/acid 0.1
• ideal buffer - b/a = 1
• Selection of buffer
• pH < 7, weak acid/conj. base
• pH > 7, weak base/conj. acid
• Choose one that pH pKa (or [H+] Ka)

• Essentially all first proton comes off before any second
• H2A H+ + HA- (K1)
• HA- H+ + A2- (K2)
• Solution of weak acid - treat like monoprotic acid
• Use K1 - describes equilibrium for proton donation
• Solution of 'stripped' anion (A2-) treat like conjugate base of monoprotic acid
• must have all donatable protons removed
• use 'last' K value
• equilibrium w/formation of stripped base
• Acid salt
• partially neutralized polyprotic acid (HA-)
• Amphiprotic species
• conjugate base of polyprotic acid
• weak acid w/protons left to donate
• Use K value for dissociation of proton (K2 for diprotic) and K value for formation (K1 for diprotic)
• Same types of calculations true for triprotic acids
• Caution choosing K values for acid salt
• Buffers from polyprotic acids
• Choose Kx that has both buffering species in equation
• Treat like other buffer

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