Common Ion Effect CH 3 COOH H + (aq) + CH 3 COO  (aq) pH of 0.1 M soln = Add 0.1 M CH 3 COONa: CH 3 COONa  Na + + CH 3 COO  (aq) pH = What happened to [CH 3 COO  ]? [CH 3 COOH]?

Buffer Solutions A buffer is a solution that “resists” a change in pH E.g. blood contains substances that keep its pH fixed at important for life functions Buffer solutions consist of either: A weak acid + salt of its conjugate base or A weak base + salt of its conjugate acid

Buffers HX(aq) H + (aq) + X  (aq) Note: If [HX] = [X  ], pH = pK a Buffers work best near pH = pKa Henderson- Hasselbalch equation

Buffer Capacity CH 3 COOH H + (aq) + CH 3 COO  (aq) [CH 3 COOH] = [CH 3 COO  ] = 1.0 M 1 L solution, pH = pK a = Add a dropper (~20ml) of 1M HCl 2.Add a dropper (~20ml) of NaOH Repeat calculation starting from a 1.8 x M HCl solution (pH = 4.74)

Buffer Capacity A CH 3 COOH + CH 3 COONa (both 1 M) pH = B dilute solution A 10x pH = Repeat with 1 M HCl + 1 M NaCl solution

SOLUBILITY Solubility: quantity of a substance that dissolves to form a saturated solution Solubility: g/L Molar solubility:mole/L Some salts are very soluble (> 0.1 M). Recall solubility rules. Some salts are sparingly soluble (< 0.1 M) sometimes referred to as ‘insoluble’. Precipitation and Solubility of ionic salts and their equilibrium in water MX(s) M + (aq) + X  (aq)

SOLUBILITY EQUILIBRIA Precipitation Ag + (aq) + Cl  (aq)  AgCl(s) Dissolution AgCl(s)  Ag + (aq) + Cl  (aq) At equilibrium when forward rate = backward rate AgCl (s) Ag + (aq) + Cl  (aq) K eq = [Ag + ][Cl  ]/ [AgCl(s)] K eq [AgCl(s)] = [Ag + ][Cl  ]  K sp = [Ag + ][Cl  ] The concentration of solid does not change at equilibrium

Solubility Product: K sp AgCl1.8 x AgBr5.0 x AgI8.3 x CdS8.0 x ZnS1.1 x Mg(OH) x Ca(OH) x CaF x BaCO x BaSO x SOLUBILITY K sp is constant for a given solid at a given temp.

SOLUBILITY CALCULATION Calculate [Ca 2+ ] and [F - ] for a saturated CaF 2 solution. CaF 2 (s)  Ca +2 (aq) +2F  (aq) K sp = [Ca 2+ ][F  ] 2 =3.9 x 10  11 at 25 o C What is the solubility? solubility = amount of CaF 2 dissociated

COMMON ION EFFECT If Q = [Ag + ][Cl - ] > K sp, AgCl precipitates (Ion product > solubility product) If [Ag + ][Cl - ] < K sp ; AgCl dissolves When [Ag + ][Cl - ] = K sp, the solution is saturated Adding either [Ag + ] or [Cl - ] will precipitate AgCl(s) precipitation dissolution saturation

What about CaF 2 in 0.01 M NaF solution? [F  ] = 0.01 M K sp = 3.9 x 10  11 CaF 2 (s)  Ca 2+ (aq) + 2F  (aq) SOLUBILITY CALCULATION Common Ion Effect

Water Chemistry (Ch ) Water in State College/UP Campus Predominantly well water 23 wells + 1 open reservoir [Ca 2+ ] ppm ppm Mixed to reduce water hardness Cl 2 injected to kill bacteria F - added Alum (K 2 SO 4. Al 2 (SO 4 ) 3. 24H 2 O) added to improve clarity NaOH added to neutralize pH

Remediation of Water Ca 2+ (hard water), Pb 2+ (toxic) are precipitated by CO 3 2- K sp CaCO x PbCO x Which compound(s) could we use to supply CO 3 2- ? How much do we need to add ?

If we add a stoichiometric amount of Na 2 CO 3, [Ca 2+ ] = [CO 3 2- ] = K sp 1/2 [Ca 2+ ] = 3 x M 2.9 x mol/L x 40 g/mol x 1000 mg/g = 116 mg/L (116 ppm) If [CO 3 2- ] = 3 mM, what is [Pb 2+ ]? [Pb 2+ ] = K sp /[CO 3 2- ] = 3.3 x /3 x = 1 x M = 2 parts per trillion

Effect of pH What is the solubility of Mg(OH) 2 in pure water? K sp = 1.8 x What is the solubility of Mg(OH) 2 in a solution with a pH of 9?

Effect of pH on common ions If either the anion or the cation is involved in an acid base equilibrium, then it is a common ion problem. Basic metal hydroxides Low pH increases solubility e.g., Mg(OH) 2 Salts of weakly basic anions Low pH increases solubility examples (write out equilibria for practice) Mg(HCO 3 ) 2 ZnCO 3 Ca 3 (PO 4 ) 2 NaF Contrast with NaCl, Ca(NO 3 ) 2

There are amphoteric hydroxides of Al 3+ Cr 3+ Zn 2+ Sn 2+ and many transition metal ions Al(OH) 3, Cr(OH) 3, Zn(OH) 2, Sn(OH) 2, … Dissolution involves formation of complex ions: Al(OH) 3 (OH 2 ) 3 (s) + H +  Al(OH) 2 (OH 2 ) OH   Al(OH) 4 (OH 2 ) 2  Amphoteric hydroxides: Both low and high pH increases solubility AMPHOTERIC METAL HYDROXIDES

Hydration of metal ions Cu 2+ (aq) + 4 H 2 O(l)  [Cu(OH 2 ) 4 ] 2+ (aq) Lewis + Lewis  Lewis Acid/Base Acid BaseAdduct = Metal Complex Other Lewis bases react with metal ions to form complexes Cu 2+ (aq) + 4 NH 3 (aq)  [Cu(NH 3 ) 4 ] 2+ (aq) Cu 2+ (aq) + 4 CN  (aq)  [Cu(CN) 4 ] 2  (aq) Cu 2+ (aq) + 4 Cl  (aq)  [Cu(Cl) 4 ] 2  (aq) FORMATION OF COMPLEX IONS

METAL COMPLEX STABILITY Cu(NH 3 ) H 2 O  Cu(OH 2 ) NH 3 Cu 2+ (aq) Cu(OH 2 ) NH 3  Cu(NH 3 ) H 2 O [H 2 O] = constant Dissociation constant Formation constant

K f VALUES OF SOME COMPLEXES Ag(NH 3 ) x 10 7 Cu(NH 3 ) x Cu(CN) x Ag(CN) x Ag(S 2 O 3 ) x 10 13

What is the conc of free Cu 2+ ions in a 1 L solution that contains 1 x10 -3 moles total Cu 2+ and is 0.1 M in NH 3 ? Cu 2+ (aq) + 4 NH 3 (aq)  Cu(NH 3 ) 4 2+ (aq) K f = 5 x Complex Ion Formation

What is the equilibrium constant for the following reaction? CuCO 3 (s) + 4CN  (aq)  CO 3 2  (aq) + Cu(CN) 4 2  (aq) K sp CuCO 3 = 2.3 x 10  10 K f Cu(CN) 4 2  = 1 x CuCO 3 is a sparingly soluble salt? K sp CuCO 3 = 2.3 x 10  10 How can I get it to dissolve?