Bluffer’s Guide for 17.1
Section 1 dealt most significantly with the concept of the Common-Ion Effect.

Common-Ion Effect: The extent of ionization of a weak electrolyte is decreased by adding to the
solution a strong electrolyte that has an ion in common with the weak electrolyte.
For Example – think of the dissociation of acetic acid (weak electrolyte):
HC2H3O2 (aq)
H+(aq) + C2H3O2-(aq)
If one were to add NaC2H3O2 (strong electrolyte) to the solution, the equilibrium would shift to the
left (in accordance to Le Chatelier’s Principle). Therefore, the [H+] is decreased and the pH is
increased.
Remember for problem-solving that the initial concentration for the common-ion ISN’T zero (like we’re
used to putting); it’s the Molarity of the strong electrolyte, and the final concentration is going to be the
ionization of the acid (or base) PLUS the initial concentration.
Bluffers' Guide 17-2: Buffer Solutions
A buffer solution (buffer) is a solution formed from a weak acid/base pair that can resist drastic changes in pH on
addition of strong acids or bases.
- Examples - human blood (buffered at 7.4), seawater (8.1-8.3)
- Has both acid species (neutralize OH-) and base species (neutralize H+)
- Prepared usually by the addition of weak acid conjugate base to a solution of the weak acid.
+ e.g. - adding NaF to a solution of HF
- Resist change in pH in either direction
Buffer capacity - the amount of acid or base a buffer can neutralize before the pH changes appreciably.
- Depends on the amount of acid and base present
- The more concentrated the acid and base, the greater the buffer capacity
Calculating the pH of a buffer:
- May be found using ICE charts or
-Use the Henderson-Hasselbalch equation:
pH = pKa + log([base]/[Acid])
Calculating pH change as a result of addition of acid/base
1. Determine moles of acid and base in the buffer as well as moles of strong acid/base added
2. Set up a diagram of the resulting reaction. Acid will react with the conjugate base, base with the conjugate
acid. Use stoichiometry to find final moles of conjugate acid and base, and recalculate concentrations
3. Use Henderson-Hasselbalch to calculate final pH.
Chapter 17-3: Acid-Base Titrations
Titration is defined as the process of reacting a solution of unknown concentration with one of known concentration.
End Point is when the amount of acid and base are schoichiometrically equivalent in the titration.
Indicators: added to the substance before titration and used to indicate the end point of the titration. (Phenolphthalein
changes color from pH 8.3 to 10.0 and Methyl Red changes color from pH 4.2 to 6.0).
* A pH meter can be used to determine the change in pH (and therefore the change in concentration of the acid and
base) and can create a pH titration Curve.
Strong Acid- Strong Base Titration
Initial pH: determined by the concentration of the strong acid before any base is added.
pH between the initial pH and the Equivalence Point: determined by the amount of acid and base added. The
concentration of the acid will be less than that of the concentration of the original pH, but it will still be larger than the
concentration of the base.
pH after the Equivalence Point: determined by the concentration of the excess amount of base in solution. All of the
acid has been neutralized and therefore the pH will rise dramatically with each excess drop of the strong base.
* Equivalence point is at pH of 7 because it completely neutralizes.
Weak Acid-Strong Base Titrations
Initial pH: determined by the concentration of the weak acid.
pH between the Initial and Equivalence Point: determined by the amount of weak acid and strong base in solution. Part
of the weak acid has been neutralized and therefore contains the acid and its conjugate base.
pH after the Equivalence Point: determined by the concentration of the strong base.
*The pH at the equivalence point is higher than 7 because of the strength of the weak acid.
Strong Base-Weak Acid Titrations
Initial pH: determined by the concentration of the strong base
pH between the Initial and Equivalence Point: determined by the amount of the concentration of the strong base and
weak acid in solution. The weak acid is present with its conjugate base in solution.
pH after the Equivalence Point: determined by the concentration of the weak acid and the pH will be lower than 7
because the acid is being added to the solution, which increases the [H+] concentration, and therefore decreases the pH.
* The graph of this particular titration graph will be a reflection of the graphs in the previous two titrations.
* The weaker the acid, the higher the pH of the equivalence point.
Titrations of Polyprotic Acids
Acids or bases that contain more than one ionizable protons or electrons neutralize in more than one step. The titration
curves of Polyprotic acids and bases have multiple equivalence points (the number of equivalence points depend on the
amount of ionizable protons or electrons).
Bluffer’s guide 17.4 Solubility Equilibria
1.Solubility equilibriums are only present in saturated solutions
-Saturated solutions are those in which the solution is in contact with undissolved
solute.
2. The solubility-product constant is denoted as Ksp.
-Remember that solids, liquids, and solvents do not appear in equillibrium
expressions.
3. The value of Ksp is the product of the ions in the solution, raised to their stoichiometric coefficients.
4. Solubility is affected by factors such as temperature, pH, and as the concentrations of other solutes
changes.
5. Generally Ksp is the equilibrium constant for an ionic solid and its constitute ions in solutions.
BaSO 4  Ba 2  SO4
2
Ksp = [Ba2+] [SO2-]
Bluffers guide 17.5
Factors that affect solubility
1. The common ion effect
a. Solubility is decreased when a common ion is added
b. As pH decreases the [H] increases and so does solubility
c. The more basic the anion, the more solubility is influenced by pH
2. Complex ions
a. Metal ion and Lewis base bonded together
i. Ag(NH3)
b. NH3 is called a ligand
c. Formation constant kf
d. Solubility of a metal salt increases in the presence of a suitable Lewis base
i. NH3, CN-,OH-, if the metal forms a complex with the base
3. Amphoterism
a. Soluble in strong acids and bases
b. Can behave as an acid or base
i. Hydroxides and oxides of Al3+, Cr3+ ,Zn2+
ii. The hydroxides generally form complex ions
c. Hydrated metal ion act as weak acids
i. Thus amphoterism is interrupted
17.6
Much like with the equilibrium constant Q can be used to determine which way solubility will go
If Q>Ksp precipitation occurs until Q=ksp
If Q=Ksp equilibrium exists
If Q<Ksp solid dissolves until Q=Ksp
17.7
Separation of ions follows solubility rules
1. Insoluble chloride like Ag+ and Hg+
-Add HCL to remove
2. Acid soluble sulfides such as Cu+2
-add Sulfurous Acid
3. Insoluble Hydroxides and Sulfides
-add Ammonium Sulfide
4. Insoluble Phosphates
-add Ammonium Phosphate
5. Use flame test on remaining ions