Chem. 31 – 10/16 Lecture

Announcements Quiz 3 on Wednesday Water Hardness Resubmission Due today HW 2.1 Solutions Posted (But will not get to Ch. 7 today) Today’s Lecture Chapter 6 Sparingly Soluble Salts Selective Precipitation Complex Ions

Precipitations Used for Separations Example: If we wanted to know the concentrations of Ca2+ and Mg2+ in a water sample. EDTA titration gives [Ca2+] + [Mg2+]. However, if we could selectively remove Ca2+ or Mg2+ (e.g. through titration) and re-titrate, we could determine the concentrations of each ion. Determine if it is possible to remove 99% of Mg2+ through precipitation as Mg(OH)2 without precipitating out any Ca(OH)2 if a tap water solution initially has 1.0 x 10-3 M Mg2+ and 1.0 x 10-3 M Ca2+. 3

Complex Ions Example Reaction: Ag+ + 2NH3(aq) ↔ Ag(NH3)2+ Metal Ligand Complex Ion Why does reaction occur? Metal is a Lewis acid (electron pair acceptor) NH3 is a Lewis base (electron pair donator) Metal-ligand bonds are intermediate strength

Complex Ions – Why Study? Crown ether (12-crown-4) Useful in separations Complexed metals become more organic soluble Effects on metal solubility (e.g. addition of NH3 on AgCl solubility) Complexometric titrations (e.g. water hardness titration) Some Complexes are Colored (use as indicators or for spectroscopic measurements) Na+ Crown ether added Diethyl ether Sodium conc. given by gray shading water

Complex Ions Step-wise vs. full reactions: Example: addition of NH3 to Ag+ Reaction occurs in steps: 1) Ag+ + NH3(aq) ↔ AgNH3+ K1 (= β1) 2) AgNH3+ + NH3(aq) ↔ Ag(NH3)2+ K2 Net) Ag+ + 2NH3(aq) ↔ Ag(NH3)2+ β2 = K1·K2

Complex Ions Due to large exponents on ligand concentration, a small change in ligand concentration has a big effect on how metal exists Example: Al3+ + 3C2O42- ↔ Al(C2O4)33- β3 = 4.0 x 1015 [C2O42-] [Al(C2O4)33-]/[Al3+] 10-4 M 4000 10-5 M 4 10-6 M 0.004

Complex Ions – “U” Shaped Solubility Curves Many sparingly soluble salts release cations and anions that form complexes with each other Example: calcium oxalate (CaC2O4) CaC2O4(s) ↔ Ca2+ + C2O42- (Ksp = 1.3 x 10-8M) increased [C2O42-] decreases Ca2+ solubility for above reaction only, but ... Ca2+ + C2O42- ↔ CaC2O4(aq) K1 = 46 CaC2O4(aq) + C2O42- ↔ Ca(C2O4)22- K2 = 490 β2 = K1·K2 = 2.3 x 104 = [Ca(C2O4)22-]/([Ca2+][C2O42-]2)

Complex Ions – “U” Shaped Solubility Curves Solubility in water Complex ion effect Common ion effect Note: looks “U” shaped if not on log scale (otherwise “V” shaped)

Some Questions In the reaction: Ca2+ + Y4- ↔ CaY2- (where Y4- = EDTA), which species is the Lewis acid? List two applications in which the formation of a complex ion would be useful for analytical chemists. List two applications in the lab in which you used or are using complex ions. AgCN is a sparingly soluble salt. However, a student observed that adding a little of a NaCN solution to a saturated solution of AgCN did not result in more precipitation of solid. Addition of more NaCN solution resulted in total dissolution of the AgCN. Explain what is happening.

One More Question Cu2+ reacts with thiosulfate (S2O32-) to form a complex which is most stable when two moles of thiosulfate to one mole of Cu2+ are present. The b2 value is found to be 2.00 x 106. If a solution containing both Cu2+ and S2O32- is prepared and found to contain 1.7 x 10-3 M free (uncomplexed) S2O32- at equilibrium, what is the ratio of complexed to free Cu? Assume that little CuS2O3 forms.

Acids, Bases and Salts Definitions of Acids and Bases - Lewis Acids/Bases (defined before, most general category) - Brønsted-Lowry Acids/Bases: acid = proton donor base = proton acceptor (must have free electron pair so also is a Lewis base) - definitions are relative

Brønsted-Lowry Acids - examples HCO2H(aq) + H2O(l) ↔ HCO2- + H3O+ acid base conjugate conjugate base acid CH3NH2(aq) + H2O(l) ↔ CH3NH3+ + OH- base acid conjugate conjugate acid base H2SO4 + CH3CO2H(l) ↔ HSO4- + CH3CO2H2+ acid base conjugate conjugate

Brønsted-Lowry Acids Note: for most acids, the reaction with water is simplified: Example: HNO2 (nitrous acid) HNO2 ↔ H+ + NO2-

Autoprotolysis and the pH Scale Autoprotolysis refers to proton transfer in protic solvents like water: H2O(l) ↔ H+ + OH- K = Kw = [H+][OH-] = 1.0 x 10-14 (T = 25°C) In pure water [H+] = [OH-] = Kw0.5 = 1.0 x 10-7 M pH = -log[H+] = 7.0 Acidic is pH < 7; basic is pH > 7

Strong Acids Strong acids completely dissociate in water (except at very high concentrations) HX(aq) → H+ + X- (no HX(aq) exists) Ka > 1 Major strong acids: HCl, HNO3, H2SO4 Note: For H2SO4, 1st dissociation is that of a strong acid, but 2nd dissociation is that of a weak acid (Ka ~ 0.01)

Weak Acids Partially dissociate in water Most have H that can dissociate HX(aq) ↔ H+ + X- (HX(aq) exists) Example: HNO2 ↔ H+ + NO2- Degree of dissociation given by Ka value Ka = [H+][NO2-]/[HNO2] Metal cations can be acids through the reaction: Mn+ + H2O(l) ↔ MOH(n-1)+ + H+ (although for +1 and some +2 metals the above reactions favor reactants so strongly the metals can be considered “neutral”)

Ionic Compounds in Water First step should be dissociation to respective ions: example: NaCl(s) → Na+ + Cl- In subsequent steps, determine how anion/cation react: - anions usually only react as bases - cations may react as acids - see if ions are recognizable conjugate acids or bases - polyprotic acids are somewhat different

Ionic Compounds in Water Conjugate bases of weak acids are basic. NO2- + H2O(l) ↔ HNO2 (aq) + OH- Conjugate bases of weaker weak acids are stronger bases. Kb = Kw/Ka CN- is a stronger base than NO2- because Ka(HCN) = 6.2 x 10-10 and Ka(HNO2) = 7.1 x 10-3