# 15.1 - Solutions of Acids or Bases Containing Containing a Common Ion

• The common ion effect is an application of Le Châtelier’s principle

• This effect makes a solution of NaF and HF less acidic than a solution of HF alone

• The common ion effect is also important in solutions of polyprotic acids.

• It is more important than ever to be systematic and to focus on the chemistry occurring in the solution before thinking about mathematical procedures

• The way to do this is always to write the major species first and consider the chemical properties of each one

# 15.2 - Buffered Solutions

• A buffered solution is one that resists a change in its pH when either hydroxide ions or protons are added.

• The most important practical example of a buffered solution is our blood, which can absorb the acids and bases produced in biological reactions without changing its pH.

• Contains a weak acid (HA) and its salt (NaA) or a weak base (B) and its salt (BHCl)

• The pH in the buffered solution is determined by the ratio of the concentrations of the weak acid and weak base.

• As long as this ratio remains virtually constant, the pH will remain virtually constant

# 15.3 - Buffering Capacity

• A buffer with a large capacity contains large concentrations of the buffering components

• The capacity of a buffered solution is determined by the magnitudes

• The pKa of the weak acid to be used in the buffer should be as close as possible to the desired pH

# 15.4 - Titrations and pH Curves

• The progress of a titration is represented by plotting the pH of the solution versus the volume of added titrant; the resulting graph is called a pH curve or titration curve

• Strong acid: A strong base titrations show a sharp change in pH near the equivalence point

• The shape of the pH curve for a strong base–strong acid titration is quite different before the equivalence point from the shape of the pH curve for a strong base–weak acid titration

• The strong base–weak acid pH curve shows the effects of buffering before the equivalence point

• For a strong base–weak acid titration, the pH is greater than 7 at the equivalence point because of the basic properties

• Indicators are sometimes used to mark the equivalence point of an acid-base titration

• The endpoint is where the indicator changes color

• The goal is to have the endpoint and the equivalence point be as close as possible

• Equivalence (stoichiometric) point: The point in the titration where an amount of base has been added to exactly react with all the acid originally present

• The equivalence point is defined by the stoichiometry, not by the pH

• The amount of acid present, not its strength, determines the equivalence point

# 15.5 - Acid-Base Indicators

• The most common acid-base indicators are complex molecules that are themselves weak acids

• The endpoint is defined by the change in color of the indicator. The equivalence point is defined by the reaction stoichiometry

• Choosing an indicator is easier if there is a large change in pH near the equivalence point of the titration

• The dramatic change in pH near the equivalence point in a strong acid–strong base titration produces a sharp endpoint; that is, the complete color change usually occurs over one drop of added titrant

• There is a wide choice of suitable indicators. The results will agree within one drop of titrant, using indicators with endpoints as far apart as pH 5 and pH 9

# 15.6 - Solubility Equilibria and the Solubility Product

• An important consequence of solubility involves the use of a suspension of barium sulfate to improve the clarity of X rays of the gastrointestinal tract

• Pure liquids and pure solids are never included in an equilibrium expression

• The solubility product is an equilibrium constant and has only one value for a given solid at a given temperature

• On the other hand, if a common ion is present in the solution, the solubility varies according to the concentration of the common ion

• Ksp is an equilibrium constant; solubility is an equilibrium position

• Carbon dioxide dissolved in groundwater makes it acidic, increasing the solubility of calcium carbonate and eventually producing huge caverns.

• As the carbon dioxide escapes to the air, the pH of the dripping water goes up and the calcium carbonate precipitates, forming stalactites and stalagmites

# 15.7 - Precipitation and Qualitative Analysis

• We will use the ion product, which is defined just like the expression for Ksp for a given solid except that initial concentrations are used instead of equilibrium concentrations.

• Mixtures of metal ions in an aqueous solution are often separated by selective precipitation by using a reagent whose anion forms a precipitate with only one or a few of the metal ions in the mixture

• The classic scheme for qualitative analysis of a mixture containing all the common cations involves first separating them into five major groups based on solubilities

• Each group is then treated further to separate and identify the individual ions

• Group I—Insoluble chlorides

• Group II—Sulfides insoluble in acid solution

• Group III—Sulfides insoluble in basic solution

• Group IV—Insoluble carbonates

• Group V—Alkali metal and ammonium ions

# 15.8 - Equilibria Involving Complex Ions

• A complex ion is a charged species consisting of a metal ion surrounded by ligands.

• Metal ions and ligands one at a time in steps characterized by equilibrium constants called formation constants for stability constants.

• When reactions are added, the equilibrium constant for the overall process is the product of the constants for the individual reactions

• If the anion of the solid is a good base, the solubility is greatly increased by acidifying the solution.

• In cases where the anion is not sufficiently basic, the ionic solid often can be dissolved in a solution containing a ligand that forms stable complex ions with its cation