Chapter 18 Principles of Chemical Reactivity: Other Aspects of Aqueous Equilibria
Chapter 17 Additional Aspects of Aqueous Equilibria
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Transcript of Chapter 17 Additional Aspects of Aqueous Equilibria
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Prentice Hall © 2003 Chapter 17
Chapter 17Chapter 17Additional Aspects of Aqueous Additional Aspects of Aqueous
EquilibriaEquilibria
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• 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
• So, [H+] must decrease (shift left…LeChâtelier!)
17.1: The Common Ion 17.1: The Common Ion EffectEffect
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Composition and Action of Buffered Solutions
• A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X-):
• The Ka expression is
17.2: Buffered Solutions17.2: Buffered Solutions
HX(aq) H+(aq) + X-(aq)
]X[
]HX[]H[
]HX[]X][H[
-
-
a
a
K
K
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• 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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Text, P. 665
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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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• Recall:
• If Ka is small (the equilibrium concentration of the undissociated acid is close to the initial concentration), then
.
]HX[]X[logpKpH
]X[
]HX[loglog]Hlog[
-a
-
aK
]X[
]HX[]H[
]HX[]X][H[
-
-
a
a
K
K
the Henderson-Hasselbalch Equation!
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acidbase conjugatelogpKpH a
Addition of Strong Acids or Bases to Buffers
• The amount of strong acid or base added results in a neutralization reaction:
X- + H3O+ HX + H2OHX + OH- X- + H2O
Text, P. 668
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• Problems 3, 5, 9, 15, 17, 19
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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):
17.3: Acid-Base Titrations17.3: Acid-Base Titrations
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Text, P. 672
pH is determined by ?
pH is determined by ?
pH is determined by ?
pH is determined by ?Appropriate
indicator: dramatic color change in the
desired range
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• 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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Strong Base-Strong Acid Titrations• Add HCl to NaOH
Text, P. 674
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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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Text, P. 674
• 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)
• Then the pH is calculated using equilibrium conditions (H-H)
pH is determined by ?
pH is determined by
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Text, P. 675
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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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Compare initial pH
values
Compare pH values at eq.
points
Compare pH change near
eq. points
Weak Acid/Strong Base Curve Strong Acid/Strong Base Curve
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The influence of acid strength on the shape of the
curve for the titration with
NaOH
Text, P. 676
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Text, P. 677
The titration of a weak
base with a strong acid
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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 HPO32-
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Text, P. 677
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• Problems 25, 27, 29, 31, 33
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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: • for which
• Ksp is the solubility product constant
17.4: Solubility Equilibria17.4: Solubility Equilibria
BaSO4(s) Ba2+(aq) + SO42-(aq)
]SO][Ba[ -24
2spK
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• 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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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 ionsText, P. 697
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• Sample Problems # 37 & 39
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The Common Ion Effect• Solubility is decreased when a common ion is added
• Le Châtelier’s principle:
• as F- is added (from NaF), the equilibrium shifts away from the increase
• CaF2(s) is formed and precipitation occurs
17.5: Factors that Affect 17.5: Factors that Affect SolubilitySolubility
CaF2(s) Ca2+(aq) + 2F-(aq)
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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:
– As pH decreases, [H+] increases and solubility increases
• The effect of pH on solubility is dramatic• The more basic the anion, the more solubility is
influenced by pH
CaF2(s) Ca2+(aq) + 2F-(aq)
F-(aq) + H+(aq) HF(aq)
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Formation of Complex Ions• The formation of Ag(NH3)2
+:
• 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+(aq) + 2NH3(aq) Ag(NH3)2(aq)
23
23]NH][Ag[
])Ag(NH[
fK
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Text, P. 687
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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) Al(OH)4-(aq)
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• Hydrated metal ions act as weak acidsThus, 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) Al(H2O)2(OH)4-(aq) + H2O(l)
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• Problems 41, 43, 49
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• 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
17.6: Precipitation and 17.6: Precipitation and Separation of IonsSeparation of Ions
BaSO4(s) Ba2+(aq) + SO42-(aq)
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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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• Problems 51, 53, 55
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• 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
17.7: Qualitative Analysis 17.7: Qualitative Analysis for Metallic Elementsfor Metallic Elements