Atkins & de Paula: Elements of Physical Chemistry: 5e Chapter 8: Chemical Equilibrium: Equilibria in Solution

End of chapter 8 assignments Discussion questions: 3, 5, 6 Exercises: 1, 5, 7, 8, 9, 10, 11, 12, 13, 24, 28, 29, 34, 35, 36 Use Excel if data needs to be graphed

Homework assignments Did you: Read the chapter? Work through the example problems? Connect to the publisher’s website & access the “Living Graphs”? Examine the “Checklist of Key Ideas”? Work any end-of-chapter exercises? Review terms and concepts that you should recall from previous courses Study the introductory pages closely!

Brønsted-Lowry theory Note the use of activity, a, for [J] Activity is technically correct; ions interact over distances, making [J] “hazardous” Conjugate acid=accepted the H+ Conjugate base=anion remains from acid So… pH = -log[H3O+] is pH = -log a H3O+

Brønsted-Lowry theory conjugate acid conjugate base base acid

Protonation and deprotonation Strong & weak electrolytes, acids, bases Acidity constants & basicity constants HA + H2O  H3O+ + A- B +H2O  BH+ + OH- aHA a H3O+ A- Ka = [H3O+][A-] [HA] = pKa = -log Ka aB a BH+ OH- Kb = [BH+][OH-] [B] = pKb = -log Kb

Protonation and deprotonation Autoionization (autoprotolysis) of water H2O + H2O  H3O+ + OH- Kw = Ka x Kb & pKw = pKa + pKb pKw = pH + pOH pH = -log a and pOH = -log a H3O+ OH-

Protonation and deprotonation The fraction (f) of molecules of a weak acid (HA) that has donated a proton (deprotonated): f = equilibrium molar concentration of conjugate base molar concentration of acid as prepared [A-]equilibrium [HA]as prepared

Protonation and deprotonation The fraction (f) of molecules of a weak base (B) that has accepted a proton (been protonated): f = equilibrium molar concentration of conjugate acid molar concentration of base as prepared [BH+]equilibrium [B]as prepared

Table 8.1 Acidity and basicity constants* at 298.15 K (1) Weakest weak acids p.182 Weakest weak acids

Table 8.1 Acidity and basicity constants* at 298.15 K (2) p.182 Table 8.1 Acidity and basicity constants* at 298.15 K (2)

Work through these Example 8.1, p.176 Example 8.2, p.176f Do any of you need help working I C E problems?

Polyprotic acids H2A(aq) + H2O(l)  H3O+(aq) + HA-(aq) HA-(aq) + H2O(l)  H3O+(aq) + A2-(aq) Work Example 8.3 & Example 8.4 (p.178f) Ka1 = aH2A a H3O+ HA- Ka2 = aHA- a H3O+ A2-

Polyprotic acids Work Example 8.3 (p.178) Work Example 8.4 (p.178f) Understand the graphs in Fig 8.1 & 8.2

Polyprotic acids For Example 8.4 For pH < pKa, the acid form dominates For pH = pKa, the conjugate pair have equal concentrations For pH > pKa, the base form dominates

Table 8.2 Successive acidity constants of polyprotic acids at 298.15 K

Amphiprotic systems Amphoteric is the old term Can act as an acid or a base Examples: HCO3- and amino acids Is NaHCO3(aq) acidic or basic? pH = ½ (pKa1 + pKa2) [See Derivation 8.1 p.180] (Back up one slide and calculate the pH) Does this “make sense”?

Acid-base titrations Terms you must understand: Stoichiometric point (equivalence point) Analyte Titrant pH curve Work through Example 8.5 (p.182) and Illustration 8.3 (p.183)

pH curve Fig 8.3 (181) Titration of a strong acid with a strong base

pH curve Fig 8.4 (181) Titration of a weak acid with a strong base

pH curve Fig 8.5 (183) Titration of a weak base with a strong acid

Buffer action Buffer solutions change pH very little when a small amount of acid or base is added to the solution Acid buffer stabilizes pH < 7 – a weak acid and a salt that contains its conjugate base Base buffer stabilizes pH > 7 – a weak base and a salt that contains its conjugate acid

pH curve of a buffer solution Fig 8.6 (191) pH of a solution changes slowly in the region halfway to the stoichiometric point Here, the solution is buffered pH ~ pKa acid base

Buffer action Consider an aqueous solution of CH3COOH CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+ conjugate base acid Consider an equimolar mixture of CH3COOH and CH3COO-Na+ Add strong acid H+ (aq) + CH3COO- (aq) CH3COOH (aq) Add strong base OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)

Acid-base indicators A water-soluble organic molecule with acid (HIn) and conjugate base (In-) forms with different colors. Examples: Phenolphthalein Cresol red Methyl red Thymol blue Litmus

Acid-base indicators HIn (aq) H+ (aq) + In- (aq) [HIn]  10 Color of acid (HIn) predominates  1 [HIn] [In-] If [HIn]  [In-], combination color  10 [HIn] [In-] Color of conjugate base (In-) predominates

Table 8.3 Indicator color changes

pH curve Fig 8.7 (194) Strong acid, strong base titration Which indicators work well here? Why will several indicators be adequate here? acid base

pH curve Fig 8.8 (194) Weak acid, strong base titration Which indicators work well here? Why will bromo-thymol blue not work well here? acid base

Solubility Solubility: the maximum quantity of solute that will dissolve in a solvent (saturation) Does T or p have to be specified? Explain The solubility constant, A/K/A “solubility product constant.” Why? MX  M+ + X- Ks = a a or Ks = [M+][X-] M+ X-

The solubility constant Ks varies according to the formula of the salt AgCl (s) Ag+ (aq) + Cl- (aq) Ks = [Ag+][Cl-] MgF2 (s) Mg2+ (aq) + 2F- (aq) Ks = [Mg2+][F-]2 Ag2CO3 (s) 2Ag+ (aq) + CO32- (aq) Ks = [Ag+]2[CO32-] Ca3(PO4)2 (s) 3Ca2+ (aq) + 2PO43- (aq) Ks = [Ca2+]3[PO33-]2

Relationship between Ks and Molar Solubility (s) Ks expression Compound Cation Anion Relation between Ks & s

Table 8.4 Solubility constants at 298.15 K (1)

Table 8.4 Solubility constants at 298.15 K (2)

Table 8.4 Solubility constants at 298.15 K (3)

The common ion effect The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider a mixture of CH3COO-Na+ (strong electrolyte) and CH3COOH (weak acid). CH3COO-Na+ (s) Na+ (aq) + CH3COO- (aq) CH3COOH (aq) H+ (aq) + CH3COO- (aq)

The common ion effect The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The addition of a common ion decreases the solubility of a salt. “It is very dangerous to neglect deviations from ideal behavior in ionic solutions…. [Q]uantitative calculations are unreliable.” (p.196)

Solubility Rules for Common Ionic Compounds in water at 25°C Soluble Compounds Exceptions Compounds containing alkali metal ions & NH4+ NO3-, HCO3-, ClO3- Cl-, Br-, I- Halides of Ag+, Hg22+, Pb2+ SO42- Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, Pb2+ Insoluble Compounds CO32-, PO43-, CrO42-, S2- Compounds containing alkali metal ions and NH4+ OH- Compounds containing alkali metal ions and Ba2+ You should know these!

Solubility Equilibria Dissolution of an ionic solid in aqueous solution: Q < Ksp Unsaturated solution No precipitate Q = Ksp Saturated solution No precipitate Q > Ksp Supersaturated solution Precipitate will form

Key Ideas

Key Ideas

The End …of this chapter…”