Chapter 19. Ionic Equilibria in Aqueous Systems
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19.1Equilibria of Acid-Base Buffers
o Common-Ion Effect
o Henderson-Hasselbalch Equation
o Buffer Capacity and Range
o Preparing a Buffer
19.2Acid-Base Titration Curves
o Acid-Base Indicators
o Strong Acid–Strong Base Titrations
o Weak Acid–Strong Base Titrations
o Weak Base–Strong Acid Titrations
o Polyprotic Acid Titrations
o Amino Acids as Polyprotic Acids
19.3Equilibria of Slightly Soluble Ionic Compounds
o Solubility-Product Constant (Ksp)
o Calculations Involving Ksp
o Effect of a Common Ion
o Effect of pH
o Formation of a Limestone Cave
o Predicting Precipitate Formation: Qsp vs. Ksp
o Selective Precipitation and Simultaneous Equilibria
19.4Equilibria Involving Complex Ions
o Formation of Complex Ions
o Complex Ions and Solubility
o Amphoteric Hydroxides
Concepts and Skills to Review Before You Study This Chapter
1.
solubility rules for ionic compounds (Section 4.3)
2.
equilibrium nature of a saturated solution (Section 13.4)
3.
effect of concentration on equilibrium position (Section 17.6)
4.
conjugate acid-base pairs (Section 18.3)
5.
calculations for weak-acid and weak-base equilibria (Sections 18.4 and 18.6)
6.
acid-base properties of salt solutions (Section 18.7)
7.
Lewis acids and bases (Section 18.9)
1. What is the pH of a solution that consists of 1.5 M NH3 and 2.5M NH4Cl? Kb = 1.8 × 10-5
Titration
250.ml of 0.50 M CH3COOH is titrated with 0.65M NaOH.
a) What is the pH of 250.ml 0.50 M CH3COOH?
b) What is the pH of the solution after 20.mL of NaOH is added?
c) What is the pH of the solution at half point?
d) What is the pH of the solution at equivalent point?
e) What is the pH after 230.mL of NaOH is added?
f)
Sketch the titration curve and locate (a) –(e) on the curve.
H
H
H
H 3N
H 2N
C
H
COOH
R
C
COOH
H 2N
C
R
Acidic
COO
R
Basic
H 3N
C
COO
R
1. Lead(II) iodide, PbI2, is an ionic compound with a solubility product constant Ksp of 7.9 × 10-9.
a. Write a chemical equation that goes with Ksp of PbI2.
b. Calculate the solubility of this compound in pure water.
c. Calculate the solubility of this compound in 0.50 mol L-1 KI solution.
2. The solubility of magnesium phosphate is 2.27 × 10-3 g/1.0 L of solution. What is the Ksp for Mg3(PO4)2
(262g/mol)?
3. At what pH will an aqueous solution of 0.022 M Cu2+ begin to precipitate as Cu(OH)2 at 25 °C? Ksp for
Cu(OH)2 = 1.6x10-19
4. Will a precipitate form when 20.0 mL of 1.8 × 10-3 M Pb(NO3)2 is added to 30.0 mL of 5.0 × 10-4 M Na2SO4?
The Ksp of (PbSO4) is 6.3 × 10-7.
5. Calculate the Molar solubility of Ag2S in 2.0M NH3 given Ksp=6.0 x 10-51 for Ag2S and Kf=1.59x107 for
Ag(NH3)2+ .
Ag2S(s) + 4NH3(aq)  2 Ag(NH3)2+ (aq) + S-2 (aq)