Chapter 16 - Acid-Base Equilibria and Solubility Equilibria

16.1 - Homogeneous versus Heterogeneous Solution Equilibria

• A weak acid solution, for example, contains nonionized acid, H+ ions, and the conjugate base at equilibrium.
  • Nonetheless, all of these species are dissolved, making the system a homogeneous equilibrium example.
• The dissolution and precipitation of mildly soluble substances is another type of equilibrium.
  • These processes are examples of heterogeneous equilibrium, that is to say they relate to reactions in more than one phase of the components.

16.2 - The Common Ion Effect

• The ionization of a weak acid or base is suppressed by the presence of a common ion.
  • The shift in equilibrium generated by the addition of a chemical with a common ion with the dissolved component is known as the common ion effect.
• In determining the pH of a solution and solubility of a salt with a small amount of solubility, the joint ion effect is important.
• It must be remembered that the equation between Henderson and Hasselbeal comes from the constant expression of equilibrium.
  • Whether it is from alone acid or is supplied both with acid and its salt, it is valid regardless of the source of the conjugate base.

16.3 - Buffer Solutions

• A buffer solution is made up of two components: (1) a weak acid or base and (2) it’s salt; both must be present.
  • When a tiny amount of acid or base is added to the solution, it canthe the the resist changes in pH.
• By adding a comparable amount of molar acetic acid (Ch3COOH) and salt sodium acetate (CH3COONa) to water, a simple buffer solution can also be prepared.
• The acid and conjugate base (from CH3COONa) balance concentrations are supposed to be the same as the initial concentrations.
  • A solution which contains these two substances is capable of either adding acid or adding base neutralization.
  • A salt acid or a conjugate base acid buffer system can be represented.
• So the recently discussed sodium acetate acid buffer can be written as CH3COONa/CH3COOH or simply as CH3COO−/CH3COOH.

16.4 - Acid-Base Titrations

• Titrations with a strong acid and a strong base, titrations with a weak acid and a strong base, and titrations with a strong acid and a weak base are the three types of reactions.
• Acid pH is indicated by −log (0.100) or 1.00 before the addition of NaOH. The solution's pH slowly increases at first when added to NaOH.
• The pH starts to increase steeply at the point of equivalence and the curve rises nearly vertically at the point of equivalence.
• The hydrogen ion and hydroxide ion concentrations are very low in a strong-acid titration at the equivalence point (about 1 to 10-7 M); therefore, the added value is one drop.

16.5 - Acid-Base Indicators

• When the indicator changes color, the titration is complete.

  • However, not all indicators change color at the same pH, therefore the indicator to employ for a given titration is determined by the nature of the acid and base.

• The endpoint of an indicator is not at a specific pH; instead, the endpoint is within the pH range.

  • In practice, an indicator with the end point in the steep section of the titration curve is chosen.

• As the equivalence point is located on the steep part of the curve, the pH on the equivalence point is inserted in the range that changes the color of the indicator.

• In acidic and neutral solutions phenolphthalein is colorless but in basic solutions reddish pink.

  • According to measurements, the pH is less than 8.3 but the indicator is reddish pink, as long as the pH is more than 8.3.

• The steepness of the pH curve near the equivalence point means that a large increase in the pH of the solution is caused by the addition of a very small amount of NaOH

  

16.6 - Solubility Equilibria

• The product of the molar concentrations of the constituent ions, each raised to the power of their stoichiometric coefficient in the equilibrium equation, is the solubility product of a compound.

• Molar solubility and solubility are two more ways to express a substance's solubility.

• Any of the following conditions may exist for the dissolution of an ionic solid into an aqueous solution:

  • (1) The solution is insaturated
  • (2) the saturated solution is saturated.
  • (3) the saturated solution.

• We use the reaction quotient for concentrations of ions that do not match balance conditions.

  

16.7 - Separation of Ions by Fractional Precipitation

• It is occasionally advantageous in chemical analysis to precipitate one type of ion from solution while leaving other ions in solution.
• It is also possible to separate these ions in insoluble silver halides.
• The solubility of halides decreases from AgCl to AgI, as the Ksp values in the margin show.
• When this solution is gradually added to a soluble compound such as silver nitrate, AgI begins precipitation, followed by AgBr and AgCl.

16.8 - The Common Ion Effect and Solubility

• The solubility product is an equilibrium constant; whenever the ion product exceeds Ksp for that substance, precipitation of an ionic compound from solution occurs.
• To restore balance, some AgCl will break away from the solution, as the principle of Le Châtelier would predict, until the ion product again matches Ksp.
  • Thus a decrease in salt solubility (AgCl) in solution is the effect of adding a common ion.

16.9 - pH and Solubility

• Because of the common ion (OH) impact, if the pH of the medium was higher than 10.45, [OH] would be higher and Mg(OH)2 solubility would decrease.
• pH has little effect on the solubilities of salts containing non-hydrolyzing anions.
• A higher [H+] and therefore lower [OH−] indicate a lower pH as would have been expected by Kw = [H+][OH−].
  • As a result, [Mg2+] increases and more Mg(OH)2 dissolves to maintain balance.

16.10 - Complex Ion Equilibria and Solubility

• Complex ions are formed when a metal cation mixes with a Lewis base in a Lewis acid-base reaction.
  • A complex ion has a core metal cation that is linked to one or more molecules or ions.
  • The formation constant Kf, which is the equilibrium constant for complex ion formation, is a measure of a metal ion's tendency to create a specific complex ion.
• Al(OH)3 is increased solubility in a medium by the formation of a complex ion Al(OH)4 − Al(OH)3 acting as Lewis acid and OH− acting as the Lewis base.
  • Similarly, other amphoteric hydroxides act.

16. 11 - Application of the Solubility Product Principle to Qualitative Analysis

• The determination of the types of ions present in a solution is known as qualitative analysis.

  • Cations in Group 1 Only the Ag+, Hg2 2+, and Pb2+ ions precipitate as insoluble chlorides when dilute HCl is added to the unknown solution. The other ions remain in the solution because their chlorides are soluble.
  • Cations of Group 2. Hydrogen sulfide is reacted with the unknown acidic solution after the chloride precipitates have been removed by filtration. The concentration of the S2 ion in the solution is minimal under these conditions.
  • Cations in Group 3 To make the solution basic, sodium hydroxide is added at this point. The aforementioned equilibrium swings to the right in a basic solution.
  • Cations in Group 4 Following the removal of all group 1, 2, and 3 cations from the solution, sodium carbonate is added to the basic solution to precipitate Ba2+, Ca2+, and Sr2+ ions as BaCO3, CaCO3, and SrCO3. Filtration is used to remove these precipitates from the solution as well.
  • Cations in Group 5 Na+, K+, and NH4 + are the only cations that may still be present in the solution at this point. Adding sodium hydroxide to a sample can reveal the presence of NH4 +.