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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)
Ex: CH3COOH and CH3COONa
-both have a common acetate ion
#1 CH3COONa(aq)  Na+(aq) + CH3COO-(aq)
#2 CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)
-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
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 conj base
3. 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
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 [ ]
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
-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
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
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
-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
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
Strong Acid-Strong Base Titrations
Titration Curve (finding pH values)
1. initial pH: pH before any base is added
*initial conc of SA ([H+]) = initial pH
2. 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)
3. 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
Weak Acid-Strong Base Titrations
How does this differ from titration curve of
strong-strong?
1. weak acid has higher initial pH than strong
2. 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
Titration Curve (finding pH values)
1. 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 = mol acid
*pH > 7 b/c anion of salt is a weak base
4. after equiv pt: pH determined by conc of
excess base
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 equiv pt)
-must make sure the equiv. pt. falls within the
color-change interval
-page 721 and 722
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
Ex: Write the Ksp expression for:
1) calcium fluoride
CaF2(s) ⇌ Ca2+(aq) + 2F-(aq)
Ksp = [Ca2+][F-]2
2) silver sulfate
Ag2SO4(s) ⇌ 2Ag+(aq) + SO42-(aq)
Ksp = [Ag+]2[SO42-]
Factors Affecting Solubility
1. Common-Ion Effect
*solubility of an ionic compound is lower in a
solution containing a common ion than in pure
water
2. 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
3. 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
4. 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+
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