Common Ion Effect, Buffers, Titrations, and Aqueous Solubility

I. Common Ion Effect:

A. Definition-

B. Example: HC2H3O2 + H2O D C2H3O2- + H3O+

Adding NaC2H3O2 to a solution of HC2H3O2 where C2H3O2- is the common ion.

C. Problems

1. Suppose we add 8.20 grams, or 0.100 mol, of sodium acetate to 1.00 L of a 0.100 M solution of acetic acid, HC2H3O2. What is the pH of the resultant solution?

2. Calculate the fluoride ion concentration and the pH of a solution containing 0.10

mole of HCl and 0.20 mole of HF in a liter of solution.

3. What is the pH of a solution made with 0.050 M NH3 and 0.10 M NH4Cl?

II. Buffers

A.  Buffer Solutions

1.  Definition-Solutions that contain an acid-base conjugate pair. They are characterized by the fact they can resist pH change upon the addition of small amounts of acid or base.

2.  Two important characteristics:

a. Buffering Capacity

  1. pH

B.  Buffers when a strong acid or base is added:

1.  With a strong acid:

2.  With a strong base:

3.  Example 1:

  1. What is the pH of 1.00 L of a buffer when 1.00 mole of lactic acid is added to 1.00

mole of sodium lactate? Ka = 1.4 x 10-4.

  1. Calculate the pH and concentrations of all species when 0.10 mole of HCl is

added to the above buffer.

  1. Calculate the pH and concentrations of all species when 0.10 mole of NaOH

is added to the buffer in a.

4.  Example 2:

a. Calculate the pH of 0.500 L of a buffer solution composed of 0.50 M formic

acid and 0.70 M sodium formate before and after adding 10.0 mL of 1.00M HCl.

5.  Example 3:

a. A buffer solution is prepared by adding100. mL of 0.050 M NH3 with 200. mL

0.0090 M NH4Cl. Calculate the pH of this buffer.

b. Calculate the pH of this buffer the pH and concentration of all species when

10.0 mL or 0.010 M HCl is added to 150. mL of this buffer

c. Calculate the pH and the concentration of all species when 10.0 mL of 0.0010 M

NaOH is added to the other 150. mL of the original buffer.

B.  Preparing a Buffer – This is easy if you know the ratio of acid to base.

  1. Using acetic acid/sodium acetate buffer solution, what ratio of acid to conjugate base will you need to maintain a pH at 5.00? Explain how you would make up such a solution. The Ka of acetic acid is 1.8 x 10-5.
  1. A chemist needs a solution buffered at pH 4.30 and can choose from the following acids (and their sodium salts):

1.  chloroacetic acid (Ka = 1.35 x 10-3)

2.  propanoic acid (Ka = 1.3 x 10-5)

3.  benzoic acid (Ka = 6.4 x 10-5)

4.  hypochlorous acid (Ka = 3.5 x 10-8)

III. Titrations

B.  Definition-A procedure for determining the amount of acid (or base) in a solution by determining the volume of base (or acid) of known concentration.

C.  Acid-Base Titration curve- a plot of

C.  Equivalence Point-

D.  Titration of a Strong Acid by a Strong Base

1. Titration Curve

2.  Calculating the pH

  1. Example 1: Calculate the pH of a solution in which 10.0 mL of 0.100 M NaOH is added to 25.0 mL of 0.100 M HCl.
  1. Example 2: Calculate the pH of a solution in which 49.5 mL of 0.100 M

NaOH is added to 50.0 mL of 0.100 M HCl.

E.  Titration of a Weak Acid by a Strong Base

1.  Titration Curve

2.  Calculating pH

  1. Example 1: Calculate the pH of a solution at the equivalence point when 50.0 mL of 0.100 M acetic acid is titrated by 0.l00 M sodium hydroxide. Ka= 1.7 x 10-7.
  1. Example 2: Calculate the pH of a solution at the equivalence point when 25.0 mL of 0.10 M HF is titrated by 0.15 M NaOH.

F. Calculation of Ka

C.  Solubility Equilibria

A.  The Solubility Product Constant – Ksp

1.  Definition- the equilibrium constant

2.  Equilibrium Expressions

  1. PbI2
  1. AgCl
  1. Pb3(AsO4)2

B.  Calculation Ksp from the solubility

1.  A liter of solution saturated at 25oC with calcium oxalate is evaporated to dryness, giving a 0.0061 g residue. Calculate the solubility product constant of this salt.

2. By experiment, it is found that 1.2 x 10-3 mole of PbI2 dissolves in 1.00 L of

aqueous solution. What is the solubility product constant at this temperature?

3. Lead (II) Arsenate has been used as an insecticide. It is only slightly soluble in

water. If the solubility is 3.0 x 10-5 g/L, what is the solubility product constant?

C.  Calculating the solubility from Ksp

1.  Calculate the solubility in g/L of barium sulfate. Ksp = 1.1 x 10-10.

2.  Calculate the solubility in g/L of calcium fluoride in water from the known solubility constant, 3.4 x 10-11.

3.  What is the solubility of calcium phosphate in grams/L? Ksp is 1.0 x 10-26.

D.  Solubility and the Common-Ion effect

1.  When attempting to dissolve an insoluble solid in a solution containing a common ion the solubility of the solid ______. Example:

2.  Calculating molar solubility in a solution containing a common ion:

  1. What is the molar solubility of calcium oxalate in 0.15 M CaCl2?
  1. Calculate the molar solubility of barium fluoride in 0.15 M NaF.

Ksp = 1.0 x 10-6.

E.  Precipitation Calculations (Q)

1.  Ion product (not reaction quotient)- the product of ion concentrations in a solution, each concentration raised to a power equal to the number of ions in the formula of the ionic compound.

2.  When:

Q < Ksp :

Q = Ksp :

Q > Ksp :

3.  Example 1: The concentration of calcium ion in blood plasma is 0.0025 M. If the concentration of oxalate ion is 1.0 x 10-7 M, do you expect calcium oxalate to precipitate? Ksp = 2.3 x 10-9.

4.  Example 2: A solution of 0.00016 M lead (II) nitrate was poured into 456 mL of 0.00023 M sodium sulfate. Would a precipitate of lead (II) sulfate be expected to form if 255 mL of the lead nitrate solution were added?