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Chapter 17

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  1. Chapter 17 buffers- resist changes in pH by neutralizing added acid or base -contain significant amounts of a weak acid and its conj base or a weak base and its conj acid -acid will neutralize added OH-(base) and base will neutralize added H+(acid)

  2. Ex: CH3COOH and CH3COONa -both have a common acetate ion #1 CH3COONa(aq)  Na+(aq) + CH3COO-(aq) #2 CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)

  3. -when mixed, acetate from #1 causes a shift to left in #2, dec [H+] -causes acetic acid to ionize less than it normally would produces higher pH (less acidic) common-ion effect- weak electro. and strong electro. with common ion in a solution causes weak electro. to ionize less than it would if it were alone

  4. Calculating pH of a Buffer 1. identify equilibrium that is source of [H+]- determines pH 2. ICE table- be sure to include initial [ ] for acid and its conj base • use Ka to find [H+] and pH Ka= ([H+][A-])/[HA] *if initial [ ] of acid and its conj base are 102 or 103 times >Ka you can neglect the x

  5. Can also use: Henderson-Hasselbalch equation: pH = pKa + log([base]/[acid]) -base and acid are [ ] of conj acid-base pair -when [base] = [acid] pH = pKa *can only be used for buffers and when Ka is small compared to [ ]

  6. buffer capacity- the amount of acid or base the buffer can neutralize before the pH begins to change -inc with inc [ ] of acid and base used to prepare buffer -the pH range of any buffer is the pH range which the buffer acts effectively

  7. -buffers are most effective when [ ] of weak acid and base are about the same *remember when [base] = [acid], pH = pKa -this gives optimal pH of any buffer -if [ ] of one component is more than 10X(1 pH unit) the other, the buffering action is poor -effective range for a buffering system is: pH = pKa± 1

  8. Ex: -a buffering system with pKa= 5.0 can be used to prepare a buffer in the range of 4.0-6.0 -amounts of acid and conj base can be adjusted to achieve any pH within this range *most effective at pH=5.0

  9. Addition of Strong Acids or Bases to Buffers -if strong base is added: OH-(aq) + HX(aq) H2O(l) + X-(aq) *OH- reacts with HX (weak acid) to produce X- *[HX] will dec and [X-] will inc *inc pH slightly

  10. -if strong acid is added: H+(aq) + X-(aq) HX(aq) *H+reacts with X-(conj base) to produce HX *[X-] will dec and [HX] will inc *pH dec slightly

  11. Acid-Base Titrations -a base in a buret is added to an acid in small increments (or acid added to base) pH titration curve- graphs pH vs volume of titrant added *page 714 equivalence point- moles base = moles acid

  12. Strong Acid-Strong Base Titrations

  13. Titration Curve (finding pH values) • initial pH: pH before any base is added *initial conc of SA ([H+]) = initial pH • between initial pH and equivalence pt: as base is added pH inc slowly and then rapidly near equiv pt. *pH = conc of excess acid (not yet neutralized)

  14. equiv. pt.: mol base = mol acid, leaving only solution of the salt *pH = 7.00 4. after equivpt: pH determined by conc of excess base

  15. Weak Acid-Strong Base Titrations How does this differ from titration curve of strong-strong? 1. weak acid has higher initial pH than strong • pH change in rapid-rise portion is smaller for weak acid than strong 3. pH at equiv. is > 7.00 -page 717 fig 17.9

  16. Titration Curve (finding pH values) • initial pH: use Ka to calculate 2. between initial pH and equiv. pt: acid is being neutralized and conj base is being formed *done using ICE table and [ ] 3. equiv. pt.: mol base = molacid *pH > 7 b/c anion of salt is a weak base 4. after equivpt: pH determined by conc of excess base

  17. Titrations of Polyprotic Acids -occurs in a series of steps -has multiple equiv. pt. (one for each H+) -page 720 fig 17.12 Acid-Base Indicators endpoint- point where indicator changes color (closely approximates equivpt) -must make sure the equiv. pt. falls within the color-change interval -page 721 and 722

  18. Solubility Equilibria -looking at dissolution of ionic compounds -heterogeneous reactions ex: BaSO4(s) BaSO4(s) ⇌ Ba2+(aq) + SO42-(aq) -to determine solubility (saturated solution): solubility product constantKsp = [ions]coeff Ex: Ksp = [Ba2+][SO42-] *the smaller the Ksp, the lower the solubility

  19. Ex: Write the Ksp expression for: • calcium fluoride CaF2(s) ⇌ Ca2+(aq) + 2F-(aq) Ksp= [Ca2+][F-]2 • silver sulfate Ag2SO4(s) ⇌ 2Ag+(aq) + SO42-(aq) Ksp= [Ag+]2[SO42-]

  20. Factors Affecting Solubility • Common-Ion Effect *solubility of an ionic compound is lower in a solution containing a common ion than in pure water

  21. Solubility and pH *solubility of a compound with a basic anion (anion of weak acid) inc. as the solution becomes more acidic

  22. Formation of Complex Ions *involves transition metal ion in solution and a Lewis base complex ion- contains a central metal ion bound to one or more ligands ligand- a neutral molecule or ion that acts as a Lewis base with the central metal ion -forms a coordinate covalent bond (one atom contributes both electrons for a bond) page 731 and 732

  23. Amphoterism -behave as an acid or a base amphoteric oxides/hydroxides- soluble in strong acids and bases b/c they can behave as an acid or a base ex: Aℓ3+, Cr3+, Zn2+, Sn2+

  24. Precipitation and Separation of Ions -if two ionic compounds are mixed, a precipitate will form if product of initial ion [ ] > Ksp -use Q (reaction quotient) *if Q > Ksp, precipoccurs, dec ion [ ] until Q = Ksp *if Q = Ksp, equilibrium exists (saturated solution) *if Q < Ksp, solid dissolves, inc ion [ ] until Q = Ksp