Transcript Chapter 17 PPT

Chapter 17
Additional Aspects of Aqueous
Equilibria
Prentice Hall © 2003
Chapter 17
17.1: The Common Ion
Effect
• The solubility of a partially soluble acid is decreased
when a common ion is added
• HC2H3O2(aq) + H2O(l)
H3O+(aq) + C2H3O2-
(aq)
• Consider the addition of C2H3O2• This is a common ion
• From a salt such as NaC2H3O2
• Therefore, [C2H3O2-] increases and the system is no
longer at equilibrium
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Chapter 17
+
• So, [H ] must decrease (shift left…LeChâtelier!)
17.2: Buffered Solutions
Composition and Action of Buffered
Solutions
• A buffer consists of a mixture of a weak acid (HX) and
its conjugate base (X-):
HX(aq)
H+(aq) + X-(aq)
• The Ka expression is
[H  ][ X - ]
Ka 
[HX]
[HX]

[H ]  K a
[X ]
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Chapter 17
• A buffer resists a change in pH when a small amount of
OH- or H+ is added
• When OH- is added to the buffer, the OH- reacts with HX
to produce X- and water
• The [HX]/[X-] ratio remains more or less constant, so
the pH is not significantly changed
• When H+ is added to the buffer, X- is consumed to produce
HX
• the pH does not change significantly
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Chapter 17
Text, P. 665
Buffer Capacity and pH
• Buffer capacity is the amount of acid or base neutralized
by the buffer before there is a significant change in pH
• It depends on the composition of the buffer
• The greater the amounts of conjugate acid-base pair
(molar concentration), the greater the buffer capacity
• The pH of the buffer depends on Ka
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Chapter 17
• Recall:
[H  ][ X - ]
Ka 
[HX]
[HX]

[H ]  K a
[X ]
• If Ka is small (the equilibrium concentration of the
undissociated acid is close to the initial concentration),
then
[HX]

 log[ H ]   log K a  log
[X- ]
[X- ]
 pH  pK a  log
[HX
]
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Chapter
17
the HendersonHasselbalch
Equation!
Addition of Strong Acids or Bases to Buffers
• The amount of strong acid or base added results in a
neutralization reaction:
X- + H3O+  HX + H2O
conjugate
HX + OH- 
X- + H2base
O
pH  pK a  log
acid
Text, P. 668
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Chapter 17
• Problems 3, 5, 9, 15, 17, 19
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Chapter 17
17.3: Acid-Base Titrations
Strong Acid-Strong Base Titrations
• A plot of pH versus volume of acid (or base) added is
called a titration curve
• Consider adding a strong base (NaOH) to a solution of a
strong acid (HCl):
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Chapter 17
Text, P. 672
Appropriate
indicator: dramatic
color change in the
desired range
pH is determined by ?
pH is determined by ?
pH is determined by ?
pH is determined by ?
• The equivalence point in a titration is the point at which
the acid and base are present in stoichiometric quantities
• The end point in a titration is the observed point
• The difference between equivalence point and end point
is called the titration error
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Chapter 17
Strong Base-Strong Acid
Titrations
• Add HCl to NaOH
Text, P. 674
Weak Acid-Strong Base Titrations
• Consider the titration of acetic acid, HC2H3O2 and NaOH
• Before any base is added, the solution contains only weak
acid
• As strong base is added, the strong base consumes a
stoichiometric quantity of weak acid:
HC2H3O2(aq) + NaOH(aq)  C2H3O2-(aq) + H2O(l)
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Chapter 17
Text, P. 674
pH is determined by
pH is determined by ?
• There is an excess of acid before the equivalence point so there is a
mixture of weak acid and its conjugate base
– The pH is given by the buffer calculation
• First the amount of C2H3O2- generated is calculated, as well as
the amount of HC2H3O2 consumed (Stoichiometry)
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Chapter
17
• Then
the pH is calculated
using
equilibrium conditions (H-H)
Text, P. 675
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Chapter 17
pH is determined by ?
pH is determined by ?
Note that pH is above 7
…the acetate ion is a weak
base
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Chapter 17
Weak Acid/Strong Base Curve
Strong Acid/Strong Base Curve
Compare pH
values at eq.
points
Compare pH
change near
eq. points
Compare
initial pH
values
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Chapter 17
Text, P. 676
The influence of
acid strength on
the shape of the
curve for the
titration
with
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NaOH
Chapter 17
Text, P. 677
The titration
of a weak
base with a
strong acid
Titrations of Polyprotic Acids
• In polyprotic acids, each ionizable proton dissociates in
steps
• Therefore, in a titration there are n equivalence points
corresponding to each ionizable proton
• In the titration of H3PO3 with NaOH,
– The first proton dissociates to form H2PO3– Then the second proton dissociates to form HPO32Prentice Hall © 2003
Chapter 17
Text, P. 677
• Problems 25, 27, 29, 31, 33
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Chapter 17
17.4: Solubility Equilibria
The Solubility-Product Constant, Ksp
• Consider equilibria that are heterogeneous
• Some common applications:
• Tooth enamel and soda, salts and kidney stones,
stalactites and stalagmites
• Example:
BaSO4(s)
Ba2+(aq) + SO42-(aq)
• for which
K sp  [Ba 2 ][SO24- ]
• Ksp is the solubility product constant
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Chapter 17
• In general: the solubility product is the molar concentration
of ions raised to their stoichiometric powers
• Solubility is the amount (grams) of substance that
dissolves to form a saturated solution
• Affected by
• pH
• concentrations of other ions in solution
• Molar solubility is the number of moles of solute
dissolving to form a liter of saturated solution
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Chapter 17
Solubility and Ksp
• To convert solubility to Ksp:
• Solubility needs to be converted into molar solubility (via
molar mass)
• Molar solubility is converted into the molar concentration
of ions at equilibrium (equilibrium calculation)
• Ksp is the product of equilibrium concentration of ions
Text, P. 697
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Chapter 17
• Sample Problems # 37 & 39
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Chapter 17
17.5: Factors that Affect
Solubility
The Common Ion Effect
• Solubility is decreased when a common ion is added
• Le Châtelier’s principle:
CaF2(s)
Ca2+(aq) + 2F-(aq)
• as F- is added (from NaF), the equilibrium shifts away
from the increase
• CaF2(s) is formed and precipitation occurs
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Chapter 17
Solubility and pH
• If the F- is removed, then the equilibrium shifts right and
CaF2 dissolves
• F- can be removed by adding a strong acid:
CaF2(s)
Ca2+(aq) + 2F-(aq)
– As pH decreases, [H+] increases and solubility increases
F-(aq) + H+(aq)
HF(aq)
• The effect of pH on solubility is dramatic
• The more basic the anion, the more solubility is
influenced by pH
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Chapter 17
Formation of Complex Ions
• The formation of Ag(NH3)2+:
Ag+(aq) + 2NH3(aq)
Ag(NH3)2(aq)
• The Ag(NH3)2+ is called a complex ion
• NH3 (the attached Lewis base) is called a ligand
• Lewis bases share their nonbonded electron pairs with
vacant orbitals on the metal atom
• The equilibrium constant for the reaction is called the
formation constant, Kf:
[Ag(NH 3 )2 ]
Kf 
 17 NH ]2
[Ag
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Chapter][
3
Text, P. 687
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Chapter 17
Amphoterism
• Amphoteric oxides will dissolve in either a strong acid or
a strong base
• Examples: hydroxides and oxides of Al3+, Cr3+, Zn2+,
and Sn2+
• The hydroxides generally form complex ions with four
hydroxide ligands attached to the metal:
Al(OH3)(s) + OH-(aq)
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Chapter 17
Al(OH)4-(aq)
• Hydrated metal ions act as weak acids
Thus, the amphoterism is interrupted:
Al(H2O)63+(aq) + OH-(aq)
Al(H2O)5(OH)2+(aq) + H2O(l)
Al(H2O)5(OH)2+(aq) + OH-(aq)
Al(H2O)4(OH)2+(aq) + H2O(l)
Al(H2O)4(OH)2+(aq) + OH-(aq)
Al(H2O)3(OH)3(s) + H2O(l)
Al(H2O)3(OH)3(s) + OH-(aq)
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Al(H2O)2(OH)4-(aq) + H2O(l)
Chapter 17
• Problems 41, 43, 49
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Chapter 17
17.6: Precipitation and
Separation of Ions
BaSO4(s)
Ba2+(aq) + SO42-(aq)
• At any instant in time, Q = [Ba2+][SO42-]
– If Q > Ksp, precipitation occurs until Q = Ksp
– If Q = Ksp, equilibrium exists
– If Q < Ksp, solid dissolves until Q = Ksp
• Based on solubilities, ions can be selectively removed
from solutions
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Chapter 17
Selective Precipitation of Ions
• Ions can be separated from each other based on their salt
solubilities
• Example: if HCl is added to a solution containing Ag+
and Cu2+
• the silver precipitates (Ksp for AgCl is 1.8  10-10) while
the Cu2+ remains in solution
• Removal of one metal ion from a solution is called
selective precipitation
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Chapter 17
• Problems 51, 53, 55
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Chapter 17
17.7: Qualitative Analysis
for Metallic Elements
• Qualitative analysis is
designed to detect the
presence of metal ions
• Quantitative analysis is
designed to determine
how much metal ion is
present
• See Text, P. 692-695