1. Chapter 17
buffers- resist changes in pH by neutralizing added acid or base -acid will neutralize added OH- (base) and base will neutralize added H+ (acid) -a WA/WB by itself does not contain enough CB/CA to be a buffer -contain significant amounts of WA and its CB or a WB and its CA

2. Ex: a simple buffer can be made by dissolving: CH3COOH and NaCH3COO -both have a common acetate ion #1 NaCH3COO(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 b/c adding to product side, dec [H+] #1 NaCH3COO(aq)  Na+(aq) + CH3COO-(aq) #2 CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq) -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. Ka = [H+][A−] [HA] Calculating pH of a Buffer
1. identify equilibrium that is source of [H+] determines pH (the acid)
2. ICE table- be sure to include initial [ ] for acid and its CB (common ion)
use Ka to find [H+] and pH
Ka = [H+][A−] [HA]
*if initial [ ] of acid and its CB are 102 or 103 times > Ka you can neglect the x value (change)

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 WA and CB are about the same
-buffers are most effective when [ ] of WA and CB 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
Ex: -a buffering system with pKa= 5.0 can be used to prepare a buffer in the range of amounts of acid and CB can be adjusted to achieve any pH within this range *most effective at pH=5.0 *b/c pH = pKa

9. Addition of Strong Acids or Bases to Buffers -if SB is added: OH-(aq) + HX(aq)  H2O(l) + X-(aq) *OH- reacts with HX (WA) 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)
-if strong acid is added: H+(aq) + X-(aq)  HX(aq) *H+ reacts with X-(CB) 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 equiv pt: 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 CB is being formed
*done using ICE table and [ ]
3. equiv. pt.: mol base = mol acid
*pH > 7 b/c anion of salt is a weak base
4. after equiv pt: 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 Acid-Base Indicators endpoint- point where indicator changes color (equiv pt) -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 = [products]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
-pH of a solution affects the solubility of any substance whose anion is basic
*solubility of a compound with a basic anion (anion of WA) 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, precip occurs, dec ion [ ] until Q = Ksp
*if Q = Ksp, equilibrium exists (saturated solution)
*if Q < Ksp, solid dissolves, inc ion [ ] until Q = Ksp