AQUEOUS EQUILIBRIA AP Chapter 17

Common Ion Effect Whenever a weak electrolyte and a strong electrolyte contain a common ion, the weak electrolyte ionizes less than it would if it were alone in solution.

Common Ion Effect In a solution that contains both sodium acetate (a strong electrolyte) and acetic acid, (a weak electrolyte) they share a common ion, the acetate ion. CH3COONa(aq) → Na+(aq) + CH3COO-(aq) CH3COOH(aq) ↔ H+(aq) + CH3COO-(aq)

Calculating pH – Common Ion Identify the strong electrolytes and the weak electrolytes. Identify the major species in the solution. Identify the equilibrium that is the source of the H+, which determines the pH. Tabulate the ions concentrations involved in the equilibrium. Calculate Ka and then pH.

Buffer Solutions A buffer solution is a solution that resists a change in pH upon addition of small amounts of strong acid or base. A mixture of a weak acid and its conjugate base, both existing in significant quantities in the same solution, makes an effective buffer.

How buffers work Strong acid + strong base → salt + water (neutral) HCl(aq) + NaOH(aq) →NaCl(aq) + HOH(l) Net ionic: H+(aq) + OH-(aq) → HOH(l) The solution is neutral because both the Na ions and the Cl ions are neutral.

Buffers, continued Strong acid + weak base → weak acid + salt HCl(aq) + NaC2H3O2(aq) → NaCl(aq)+ HC2H3O2(aq) Net ionic: H+(aq) + C2H3O2-(aq) → HC2H3O2(aq) Weak acid

Buffers, continued Strong base + weak acid → weak base NaOH(aq) + HC2H3O2(aq)→ NaC2H3O2(aq) + HOH(l) OH-(aq) + HC2H3O2(aq) → C2H3O2-(aq) + HOH(l) Net ionic:

Buffers, continued A weak acid does not react appreciably with a conjugate weak base.

Calculating the pH of a Buffer It is calculated like the pH of a weak acid solution. Reference the example on page 725.

Calculating the pH of a Buffer Example: What is the pH of an aqueous mixture containing 0.20 M acetic acid and 0.10 M sodium acetate? HC2H3O2(aq) + HOH(l) ↔ C2H3O2-(aq) + H+(aq) I 0.20M 0.10M + 0 C -x +x +x E 0.20 – x 0.10 + x x

pH = -log[H3O+] = -log(3.6 x 10-5) = 4.44 Ka = = If x is small compared to 0.10, then 0.10 + x ≈ 0.10 and 0.20 – x ≈ 0.20. So, Ka = 1.8 x 10-5 = x = [H3O+] = 3.6 x 10-5 pH = -log[H3O+] = -log(3.6 x 10-5) = 4.44 [C2H3O2-][H+] [HC2H3O2] (0.10 + x)(x) (0.20 – x) 0.10x 0.20 0.10x 0.20 From Appendix D!

Calculation of the pH of a buffer after adding an acid or base.

Calculating the pH of Buffers Because the ICE tables for all buffer solutions are the same, the equation used in all buffer calculations can be generalized: Ka = pH = pKa + log([base]/[acid]) Ka is the ionization constant of the weak acid, x is the molar concentration of [H+], [base] is the initial concentration of the weak base, and [acid] is the initial concentration of the weak acid. x[base] [acid]

Buffer Capacity and pH Range Buffer capacity is the amount of acid or base that the buffer can neutralize before the pH begins to change an appreciable amount. pH range of any buffer is the pH range over which the buffer acts effectively. Buffers most effectively resist change in pH in either direction when the concentration of weak acid and conjugate base are about the same.

Acid-Base Titrations An acid-base titration is a method used to determine an unknown concentration of an acid or a base. It determines the volume of a standard solution of base of unknown concentration that’s required to completely react with an acid sample. The reverse can also be done to react with a base solution.

Acid-Base Indicators An acid-base indicator changes color at the endpoint of the titration. The indicator signals the equivalence point, the point at which there are equal molar amounts of acid and base. A titration curve is a graph of pH vs. mL of titrant. Various acid-base titrations produce distinctive titration curves.

Adding a strong base to a strong acid. Both indicators change color at the equivalence point.

Adding a strong base to a weak acid. Only phenolphthalein changes color here.

Adding a strong acid to a weak base.

Acid Strength on Titration Curves The weaker the acid, the higher the initial pH and the smaller the pH change at the equivalence point.

Solubility Equilibria When a precipitate forms from the mixing of two solutions, an equilibrium is established between the solid precipitate and the dissolved ions. (Ksp is the solubility product constant.) AgCl(s) ↔ Ag+(aq) + Cl-(aq) The equilibrium constant is: Ksp = [Ag+][Cl-]

Solubility Product Constant For silver chloride, Ksp = 1.8 x 10-10 (from Appendix D Table D-3) AgCl(s) = Ag+(aq) + Cl-(aq) I I 0 0 C -x +x +x E I – x x x

Molar Solubility The molar solubility is the number of moles of silver chloride that dissolve in a liter of water. The molar solubility is x. For every mole of silver chloride that dissolves, one mole of Ag+ and one mole of Cl- are formed.

SO . . . x = [Ag+] = [Cl-] Ksp = [Ag+][Cl-] Substituting: Ksp = x2 1.8 x 10-10 = x2 x = 1.3 x 10-5 M, the molar solubility of silver chloride

The solubility product of a compound equals the product of the concentration of the ions involved in the equilibrium, each raised to the power of its coefficient in the equilibrium equation. The values of many ionic solids are tabulated in Appendix D. (example – The Ksp for BaSO4 is 1.1x10-10, which is very small. Only a very small amount of the solid will dissolve in water.)

Examples BaSO4(s) ↔ Ba+2(aq) + SO4-2 (aq) Ksp = [Ba+2][SO4-2] CaF2(s) ↔ Ca+2(aq) + 2 F-(aq) Ksp = [Ca+2][F-]2

Solubility and pH The solubility of a substance can change considerably as the concentrations of other solutes change, such as pH. It can be affected by the presence of other ions in the solution.

Factors that affect solubility Common ion effect – the solubility of a slightly soluble salt is decreased by the presence of second solute that furnishes a common ion.

Solubility and pH If a compound contains a basic anion, (the anion of a weak acid) the solubility will increase as the solution becomes more acidic. The solubility of slightly soluble salts containing basic anions increases as [H+] increases (or as pH is lowered.)