ch16blank Page 1 Chapter 16: Aqueous ionic equilibrium Topics in this chapter: 1. Buffers 2. Titrations and ph curves 3. Solubility equilibria Buffersresist changes to the ph of a solution. Consider a normal weak acid equilibrium: HA + H 2 O H 3 O + + A Which direction will the reaction shift if more A is added? What happens to the acidity? What happens to the % ionization of HA? The equilibrium concentrations of HA and A will be very close to their initial concentrations: K a = [H 3 O + ] =
ch16blank Page 2 Henderson-Hasselbalch Equation ph = pk a + log [base] [acid] But be careful This only works for a buffer solution (high concentrations of both parts of a conjugate acid-base pair) Don't flip the fraction! The conj. base is on top! Use the pk a of the conjugate acid. What is the ph of a solution of 0.50 M acetic acid and 0.50 M sodium acetate?
ch16blank Page 3 Buffer ph calculation What is the ph of a solution of 0.10 M ammonium chloride and 0.15 M ammonia?
ch16blank Page 4 Adding a strong acid or base to a buffer Buffers resist changes in ph because they contain both an acid and a base. In a buffer of HA and A : What happens when you add H 3 O + to the buffer? What happens when you add OH to the buffer?
ch16blank Page 5 Adding a strong acid or base to a buffer Calculate the ph when 0.0050 mol strong acid is added to a buffer which contains 0.010 mol NH 4+ and 0.015 mol NH 3? 1. Stoichiometry calculation 2. Equilibrium calculation (H-H equation)
ch16blank Page 6 Adding a strong acid or base to a buffer A buffer contains 0.47 M HNO 2 and 0.61 M NaNO 2. What is the ph of this buffer? What is the ph when 8.7 ml of 0.128 M NaOH is added to 243 ml of the buffer solution?
ch16blank Page 7 Buffer capacity and buffer range Relative amounts of acid and base: A buffer is most effective when the ratio of A to HA is between 0.1 and 10. What ph range does this correspond to? So, when creating a buffer, make sure the pk a of the acid is as close to the desired ph as possible. What is the ph range of a buffer created by formic acid (pk a = 3.74) and its conjugate? What ratio of the base to acid would be needed to make a ph 4.00 buffer? Absolute concentrations of acid and base: A buffer is effective when the concentrations of A and HA are at least 10x greater than the concentration of an added strong acid or base.
ch16blank Page 8 Titrations and titration curves Known concentration Unknown concentration Titration curve: graph of ph vs volume of solution added Strong acid titrated with strong base ph Vol NaOH added (ml)
ch16blank Page 9 Strong acid titrated with strong base 50.0 ml HCl of unknown concentration is titrated with 0.150 M NaOH. The titration curve's only inflection point occurs at 40.0 ml of added base. What was the initial HCl concentration? What was the initial ph of the HCl solution? What is the ph after 60.0 ml NaOH is added?
ch16blank Page 10 Weak acid titrated with strong base ph Vol NaOH added (ml) HA + OH Buffer region: At the midpoint (half-equivalence pt):
ch16blank Page 11 Weak acid titrated with strong base 40.0 ml of 0.0788 M HA are titrated with 0.100 M NaOH. What is the volume of NaOH added at the equivalence point? If the ph = 4.88 at the midpoint, what is K a and what is the ph at the endpoint?
ch16blank Page 12 Weak acid titrated with strong base What is the ph after 10.0 ml 0.100 M NaOH is added to 40.0 ml of 0.0788 M HA (the titration on the last page)?
ch16blank Page 13 Weak acid titrated with strong base 75.0 ml of a weak acid are titrated by 0.125 M NaOH. The endpoint occurs at 62.1 ml of added base, and the ph at the midpoint is 3.12. What is the initial acid concentration? What is the K a? What is the ph initially and at the equivalence point?
ch16blank Page 14 Weak base and polyprotic acid titrations Weak base titr by strong acid ph Vol HCl added (ml) Polyprotic acid titr by strong base ph Vol NaOH added (ml)
ch16blank Page 15 ph Indicators Make sure the color range of your indicator is within the most vertical part of the titration curve!
ch16blank Page 16 Solubility equilibria Compounds we called "insoluble" before, are actually very slightly soluble, governed by an equilibrium: CaF 2 (s) Ca 2+ (aq) + 2 F (aq) Solubility product constant, K sp for CaF 2 = Molar solubility, S= mol solute (when saturated) L solution If the molar solubility, S,of PbF 2 is 2.61 x 10 3 M, what is the K sp of PbF 2? PbF 2 (s) What are the ion concentrations?
ch16blank Page 17 Calculating molar solubility If K sp for Mg(OH) 2 is 2.06 x 10 13, calculate the molar solubility, S.
ch16blank Page 18 K sp and relative solubility Direct comparison of K sp values only works between compounds that dissociate into the same number of ions. Which is the most soluble? BaF 2 K sp = 2.45 x 10 5 CaF 2 K sp = 1.46 x 10 10 Fe(OH) 2 K sp = 4.87 x 10 17 PbCl 2 K sp = 1.17 x 10 15 PbBr 2 K sp = 4.67 x 10 6
ch16blank Page 19 Common Ion Effect How is solubility affected when more ions are added to the solution? Ca(OH) 2 (s) Ca 2+ (aq) + 2 OH (aq); K sp = 4.68 x 10 6 To this solution, we add Ca(NO 3 ) 2 Na +, K +, NH 4+, NO 3, C 2 H 3 O 2 make compounds totally soluble The reaction shifts: A common ion in solution will a compound's solubility.
ch16blank Page 20 Common ion effect calculation What is Sfor Ca(OH) 2? K sp = 4.68 x 10 6 What is Sfor Ca(OH) 2 in 0.10 M Ca(NO 3 ) 2 solution?
ch16blank Page 21 Effect of ph on solubility If a compound's dissociation produces a basic anion (usually OH, CO 3 2, S 2 ), solubility is dependent on ph. CaCO 3 (s) Ca 2+ (aq) + CO 3 2 (aq) Addition of strong acid to this solution A similar effect occurs when base is added to a solution which contains an acidic cation.
ch16blank Page 22 Precipitation To calculate whether or not a precipitate will form, take the given ion concentrations and calculate Q(the reaction quotient). If Q= K sp : If Q> K sp : If Q< K sp : A newly mixed solution has these ion concentrations: [Pb 2+ ] = 0.25 M, [Cl ] = 0.017 M. Will a precipitate form? K sp (PbCl 2 ) = 1.17 x 10 5.