Chapter 17 Notes- Additional Aspects of Aqueous Equilibria

17.1 The Common-ion Effect [p.720]

  1. Whenever a weak electrolyte and a strong electrolyte contain a common ion the weak electrolyte ionizes less than if it would if it were alone in solution

17.2 Buffered Solutions [p.723]

  1. Composition and Action of Buffered Solutions
  2. Buffered Solutions- contain weak conjugate acid-base pair that drastically reduces a pH change on addition of a strong acid or base
  3. Since the [H+] = Ka ([HX]/[X-]), the pH of a buffer does not change much as long as [HX]/[X-] is small.
  4. Calculating the pH of a Buffer
  5. Method 1:Use same procedure for calculating common-ion effect
  6. Method 2: Henderson-Hasselbalch equation
  7. Typically use the starting concentrations of the acid and base components because the acid and base of the buffer are ignored
  8. Buffer Capacity and pH Range
  9. Buffer capacity –amount of acid or base the buffer can neutralize before the pH begins to change a lot
  10. Dependent on the amount of acid and base used to make buffer
  11. Greater the amounts the more resistant it is to change
  12. pHrange- range over which the buffer acts effectively
  13. Buffers are most effective when the concentration of the weak acid and conjugate base is the same.
  14. Buffers usually have a useable range of +/- 1
  15. Addition of Strong Acids or Bases to Buffers
  16. Step 1:Stoichiometry calculation: consider the acid-base neutralization and determine its effect on the [HX] and [X-]
  17. Step2: Equilibrium calculation: use Ka and the new concentrations of [HX] and [X-] to calculate [H+]

17.3Acid – Base Titrations [p. 730]

  1. Strong Acid-Base Titrations
  2. Titration curve graph of the pH as a function of the volume of the added titrant
  3. Curve is broken down into four regions
  4. The initial pH (initial acid)
  5. Between the initial pH and the equivalence point (remaining acid)
  6. The equivalence point- moles of base = moles of acid. Should be 7.00 for a strong acid and strong base
  7. After the equivalence point (excess base)
  8. Endpoint- point where indication changes color in a titration
  9. Indicator most commonly used for strong acid-base titrations is phenolphthalein
  10. Titration of a strong base is the reverse curve
  11. Weak Acid- Base Titrations
  12. Curve also has four regions
  13. Initial pH (initial acid)
  14. Between the initial pH and the equivalence point (buffer mixture)
  15. Equivalence point
  16. After the equivalence point (excess base)
  17. Differences between curves for strong acid-base titrations and weak acid-base titrations
  18. Solution of a weak acid has a higher pH than the solution of a strong acid
  19. The pH change at the rapid-rise portion curve near the equivalence pt is smaller for a weak acid
  20. The pH at the equivalence pt is above 7.00
  21. Choice of indicator for a a weak acid-base titration is more critical
  22. Titrations of Polyprotic Acids
  23. Weak acids that are polyprotic or polybasic have multiple equivalence point
  24. Example would be H3PO4

17.4 Solubility Equilibria [p.737]

  1. General information
  2. Heterogeneous solutions contain a precipitate
  3. Use solubility rules to make quantitative predictions about the amount of a given compound dissolved
  4. Equilibria is used to analyze factors that affect solubility
  5. The Solubility- Product Ksp
  6. Ksp – describes the dissolution of a solid which indicates how soluble the solid is in water
  7. Calculating Ksp = product of ion concentrations involved in equilibrium
  8. Smaller the Ksp, the lower the solubility
  9. Solubility and Ksp
  10. Solubility- quantity that dissolves to form a saturated solution
  11. Expressed in g/L
  12. Molar solubility = mol/L
  13. Can change considerably as the concentrations of other solutes change
  14. Ksp
  15. Equilibrium constant is unitless
  16. Magnitude indicates how much dissolves to form a saturated solution
  17. Only one value for a given solute at a specific temperature
  18. Agreement between solubility and Ksp is best for salt with low ionic charges

17.5Factors that Affect Solubility[p.741]

  1. Common-Ion Effect

Solubility of a slightly soluble salt is decreased by the presence of a second solute that furnishes a common ion

  1. Solubility and pH
  2. Compounds containing a basic anion ( weak conjugate-acid) increases solubility as it becomes more acidic
  3. Solubility of slightly soluble salts containing basic anions increases as [H+] increases

Examples: CO32-, PO43-, CN-, S2-

  1. Formation of Complex Ions
  2. Solubility of metal salts increases in the presence of suitable Lewis Bases, like NH3, CN-, or OH-, if metal forms a complex ion with the base
  3. Example
  4. Amphoterism
  5. Amphoteric oxides and hydroxides- some metal oxides and hydroxides that are insoluble in water but soluble in a strong acid and strong base solutions
  6. Examples Al3+, Cr3+, Zn2+, Sn2+

17.6Precipitation and Separation of Ions [p. 750]

  1. Formation of a precipitate during a double –replacement reaction
  2. If Q > Ksp precipitation occurs Q = Ksp
  3. If Q = Ksp equilibrium (saturated solution)
  4. If Q < Ksp solid dissolves until Q = Ksp
  5. Selective Precipitation of Ions

Separation of ions in an aqueous solution by using a reagent that forms a precipitate with one or a few ions

17.7Qualitative Analysis for Metallic Elements [p. 753]

  1. Qualitative Analysis- determines presence of absence of ions
  2. Quantitative Analysis – determines amount present
  3. Three Stages of Analysis
  1. Ions separated into broad groups
  2. Individual ions within each group are then separated by selectively dissolving members in the group
  3. Ions are Identified
  1. Analysis Scheme
  2. Insoluble chlorides (HCl)
  3. Acid insoluble sulfides (H2S)
  4. Base insoluble sulfides and hydroxides ((NH4)2S)
  5. Insoluble Phosphates
  6. Alkali Metal Ions and NH4+

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