Chapter 15 - Acid-Base Equilibria

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.
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