Definitions Arrhenius Acids and Bases • Acids – produce hydronium ion in solution • Bases – produce hydroxide ion in solution Note: These definitions are specifically written for aqueous solutions. HA B - + H2O H3O+ + H2O HB + + A- OH Acid and Base Reactions HCl + H2O H2O NaOH(s) H3O+ Na+(aq) + + Cl- OH-(aq) Table 18.1 Fig. 18.2 Hydronium Ion and Intermolecular Interactions Acid – Base Neutralization + H3O Acid 2H2O OH + + + NaOH NaCl Base Salt H HCl - + - H2O OH + H2O Water Fig. 18.3 Strong vs. Weak Acids Fig. 18.4 Acid Strength Effects [H3O+] and Reaction Rate which Depends upon [H3O+] 1M HCl 1M CH3COOH HA(aq) Zn(s) + + H2O(l) H3O+(aq) + A-(aq) 2H3O+(aq) Zn+2(aq) + H2(g) Strong Acids and Strong Bases • Strong Acids • Hydrohalic Acids: HCl, HBr, and HI • Oxoacids: • # of O’s exceed # acid protons (H’s) by two or more • Strong Bases • Soluble compounds with O2- or OH• M2O or MOH Group 1A metals • MO or M(OH)2 Group 2A metals K = [Products]/[Reactants] Law of Mass Action (Law of Chemical Equilibrium) aA + bB cC + dD c d [C] [D] Qc = = K c a b [A] [B] For Acid Dissociation Reaction HA(aq) Qc = + [H3O H3O+(aq) H2O(l) + (aq) ][A (aq) A-(aq) + ] [HA (aq) ] [H2O(l) ] K c [H2O(l) ] = Ka = [H3O + (aq) ][A [HA (aq) ] (aq) ] + [H3O(aq) ][A (aq) ] [HA (aq) ] = Ka Acid Dissociation Constant Table 18.2 Stronger Acid Higher [H3O+] Larger Ka Smaller Ka Lower % HA Dissociated Weaker Acid [Acid Strength: % Dissociation Ka Magnitude] Prob. 18.1 Page 764 Autoionization of Water Ka = K W = [H3O + (aq) [H3O + (aq) ][OH (aq) ] [H2O(l) ]2 ][OH (aq) ] Ion Product of Water Fig. 18.5 Aqueous Solutions + 1 x 10-14 = [H3O(aq) ][OH (aq) ] Neutral Solution: + [H3O(aq) ] = [OH (aq) ] = 1 x 10-7 Prob. 18.2 Fig. 18.6 pH Scale + K W = [H3O(aq) ][OH (aq) ] 1 x 10 -14 = [H3O + (aq) ][OH pH = - log[H3O (aq) + (aq) ] ] Fig. 18.7 Aqueous Solution Characteristics pOH = - log[OH (aq) ] Prob. 18.3 Table 18.3 pK a = - log K a Fig. 18.9 Brønsted-Lowry Definitions Fig. 18.10 Conjugate Acid-Base Pairs Table 18.4 Table 18.5 Fig. 18.11 Base Dissociation Reactions Conjugate Base Conjugate Acid Base Dissociation Constant Conjugate Base(aq) + H2O(l) Conjugate Acid(aq) + OH-(aq) Product of [Products] Kb = Product of [Reactants] [Conj. Acid][OH - ] [HA][OH - ] Kb = = [Conj. Base] [A - ] A-(aq) + H2O(l) HA(aq) + OH-(aq) Table 18.6 Kb pKb [OH-] pOH A-(aq) + H2O(l) HA(aq) HA(aq) + H2O(l) A-(aq) 2H2O(l) H3O+(aq) + + + OH-(aq) OH-(aq) Kb H3O+(aq) Ka Ka x Kb = Kw Fig. 18.12 Acid – Base Strength and the Periodic Table Fig. 18.13 Salts of Weak Acids and Weak Bases Consider Cations and Anions Table 18.8 Lewis Acids and Bases Lewis Acid – electron pair acceptor Lewis Base – electron pair donor Page 793 Metals act as Lewis acids in water Equilibria of Acid-Base Buffer Systems Buffer: resists a change in pH upon addition of Acid or Base conjugate acid-base pair of a weak acid HA(aq) + Ka = A-(aq) H2O(l) [H3O + (aq) ][A [HA (aq) ] (aq) ] + H3O+(aq) Acid Dissociation Constant Table 19.1 Common Ion Effect Fig. 19.2 HA(aq) H2O(l) + HC2H3O2(aq) + H2O(l) A-(aq) + C2H3O2-(aq) H3O+(aq) + H3O+(aq) Henderson-Hasselbalch Equation Ka = [H3O + (aq) ][A (aq) ] [HA (aq) ] [A ] + -log K a = -log [H 3O ] - log [HA] [A - ] pK a = pH - log [HA] - [A ] pH = pK a + log [HA] Definition of pKa for a buffer solution. pKa = pH when [conjugate acid] = [conjugate base] Prob. 19.2 How to prepare a buffer at a specific pH: - [A ] pH = pK a + log [HA] Calculate the proportions of NaH2PO4 and Na2HPO4 needed to Prepare a buffer solution of pH=7.00. The pKa of H2PO4- is 7.21 HPO4−2 = 7.21 + log 7.00 H 2 PO4−1 HPO4−2 = log 7.00 − 7.21 −1 H 2 PO4 take antilog of both sides: HPO4−2 0.62 = 1 H 2 PO4−1 Use 0.62 mol of Na2HPO4 for every mol of NaH2PO4 in any total amount. Fig. 19.4 Acid-Base Indicators Color changes at pKa values. Fig. 19.6 Titration of a Strong Acid with a Strong Base Fig. 19.7 Titration of a Weak Acid with a Strong Base Fig. 19.9 Titration of a Weak Polyprotic Acid with a Strong Base
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