3 AgCl(s) Ag+(aq) + Cl(aq) Common Ion EffectThe shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction.AgCl(s) Ag+(aq) + Cl(aq)See problems – 17.18
4 Problem Calculate the % Ionization and pH of a solution that is M in formic acid (HCHO2) Ka=1.8x10-4a solution that is M in sodium formate (NaCHO2) and M in formic acid (HCHO2)a solution that is made by combining 125 mL of M hydrofluoric acid (Ka=6.8x10-4)with 50.0 mL of 0.10 M sodium fluoride.
5 ProblemCalculate the percent ionization of M lactic acid (HC3H5O3) (Ka = 1.4 × 10−4).Calculate the percent ionization of M lactic acid in a solution containing M sodium lactate.
7 A Buffered Solution. . . resists change in its pH when either H+ or OH are added.1.0 L of 0.50 M H3CCOOHM H3CCOONapH = 4.74Adding mol solid NaOH raises the pH of the solution to 4.76, a very minor change.
9 Key Points on Buffered Solutions 1. They are weak acids or bases containing a common ion.2. After addition of strong acid or base, deal with stoichiometry first, then equilibrium.
10 Henderson-Hasselbalch Equation Useful for calculating pH when the [A]/[HA] ratios are known.
11 ProblemCarbonic Acid (H2CO3 )Ka ValuesKa14.5 x 10-7 Ka24.7 x 10-11Calculate the pH of a buffer that is M in NaHCO3 and M in Na2CO3Calculate problem both ways – equilibrium method and using Henderson Hasselbalch equation
12 Buffered Solution Characteristics Buffers contain relatively large amounts of weak acid and corresponding base.Added H+ reacts to completion with the weak base.Added OH reacts to completion with the weak acid.The pH is determined by the ratio of the concentrations of the weak acid and weak base.
14 Making a Buffer Solution Problem How many moles of NaBrO should be added to 1.00 L of 5.0 x 10-2 M hypobromous acid (HBrO) to form a buffer solution of pH = 9.14? Assume that no volume change occurs when the NaBrO is added.See problems – 17.26
15 Buffering Capacity. . . represents the amount of H+ or OH the buffer can absorb without a significant change in pH.
18 Solubility Product For solids dissolving to form aqueous solutions. Bi2S3(s) 2Bi3+(aq) + 3S2(aq)Ksp = solubility product constantandKsp = [Bi3+]2[S2]3See problem 17.49
19 Solubility Product“Solubility” = x = concentration of Bi2S3 that dissolves, which equals 1/2[Bi3+] and 1/3[S2].Note: Ksp is constant (at a given temperature)x is variable (especially with a commonion present)See problem 17.50
25 Titration (pH) CurveA plot of pH of the solution being analyzed as a function of the amount of titrant added.Equivalence (stoichiometric) point: Enough titrant has been added to react exactly with the solution being analyzed.
31 Weak Acid - Strong Base Titration Step 1 - A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species.Step 2 - An equilibrium problem - determine position of weak acid equilibrium and calculate pH.
41 ProblemIf the molar solubility of CaF2 at 35°C is × 10−3 mol/L, what is Ksp at this temperature?The Ksp of Ba(IO3)2 at 25°C is 6.0 × 10−10. What is the molar solubility of Ba(IO3)2?
42 Problem – 17.52 pure water 0.010 M KF solution 0.050 M LaCl3 solution. Calculate the solubility of LaF3 (Ksp = 2 x 10-19) in grams per liter inpure water0.010 M KF solution0.050 M LaCl3 solution.
43 Problem – 17.56Calculate the molar solubility of Fe(OH) (Ksp = 7.9 x 10-16) when buffered at pH8.010.012.0.
44 Equilibria Involving Complex Ions Complex Ion: A charged species consisting of a metal ion surrounded by ligands (Lewis bases).Coordination Number: Number of ligands attached to a metal ion. (Most common are 6 and 4.)Formation (Stability) Constants: The equilibrium constants characterizing the stepwise addition of ligands to metal ions.