Chapter 15: Applications of Aqueous Equilibria

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Chapter 15 Applications of Aqueous Equilibria

• The common ion effect
• Buffered solutions
• Buffer capacity
• Titrations pH curves strong acid-strong base
• Solubility equilibria solubility product
• Precipitation complex ions

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The Common Ion Effect

• The shift in equilibrium that occurs because of
  the addition of an ion already involved in the
  equilibrium reaction.
• AgCl(s) ? Ag(aq) Cl?(aq)

• Concept test What happens if you add NaF(s) into
  a solution of HF?

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The Common-Ion Effect

• Consider a solution of acetic acid
• If acetate ion is added to the solution, Le
  Châtelier says the equilibrium will shift to the
  ..?

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The Common-Ion Effect

• The extent of ionization of a weak electrolyte
  is decreased by adding to the solution a strong
  electrolyte that has an ion in common with the
  weak electrolyte.

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The Common-Ion Effect

• Calculate the fluoride ion concentration and pH
  of a solution that is 0.20 M in HF and 0.10 M in
  HCl.
• Ka for HF is 6.8 ? 10-4.

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The Common-Ion Effect
Because HCl, a strong acid, is also present, the
initial H3O is not 0, but rather 0.10 M.
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The Common-Ion Effect

• x
• 1.4 ? 10-3 x

• Therefore, F- x 1.4 ? 10-3
• H3O 0.10 x 1.01 1.4 ? 10-3 0.10 M
• So, pH -log (0.10)
• pH 1.00

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A Buffered Solution

• ? Resists change in its pH when either H or OH?
  are added.
• Example 1.0 L of 0.50 M H3CCOOH
• 0.50 M H3CCOONa
• pH 4.74
• Adding 0.010 mol solid NaOH raises the pH of the
  solution to 4.76, a very minor change.
• Conceptual question Why solid NaOH?

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A Buffered Solution
Most effective when balanced and relatively high
concentration of acid base conjugate pair.
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Henderson-Hasselbalch Equation
Useful for calculating pH when the
A?/HA ratios are known.
Remember where this came from? Re-derive if
necessary
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HendersonHasselbalch Equation
What is the pH of a buffer that is 0.12 M in
lactic acid, HC3H5O3, and 0.10 M in sodium
lactate? Ka for lactic acid is 1.4 ? 10-4.
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HendersonHasselbalch Equation
pH 3.85 (-0.08) pH 3.77
Best to choose an acid with a pKa close to
desired pH
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Characteristics of Buffer solutions

• Buffers contain relatively large amounts of weak
  acid and corresponding base.
• Added H reacts to completion with the weak base.
• Added OH? reacts to completion with the weak
  acid.
• The pH is determined by the ratio of the
  concentrations of the weak acid and weak base.
• Buffer capacity is represents the amount of H or
  OH? the buffer can absorb without a significant
  change in pH.

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When Strong Acids or Bases Are Added to a Buffer

• safe to assume that all of the strong acid or
  base is consumed in the reaction.

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Calculating pH Changes in Buffers

• A buffer is made by adding 0.300 mol HC2H3O2 and
  0.300 mol NaC2H3O2 to enough water to make 1.00 L
  of solution. The pH of the buffer is 4.74.
  Calculate the pH of this solution after 0.020 mol
  of NaOH is added.

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Calculating pH Changes in Buffers

• Before the reaction, since
• mol HC2H3O2 mol C2H3O2-
• pH pKa -log (1.8 ? 10-5) 4.74

The 0.020 mol NaOH will react with 0.020 mol of
the acetic acid HC2H3O2(aq) OH-(aq) ???
C2H3O2-(aq) H2O(l)
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Calculating pH Changes in Buffers
Now use the HendersonHasselbalch equation to
calculate the new pH
pH 4.74 0.06 pH 4.80
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Buffered Solutions

• More practice,group problem solving
• See Handout

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Titrations strong acid,strong base.

• A plot of pH of the solution being analyzed as a
  function of the amount of titrant added.
• Equivalence (stoichiometric) point Enough
  titrant has been added to react exactly with the
  solution being analyzed.

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The pH curve for the titration of 100.0 mL of
0.50 M NaOH with 1.0 M HCl.
Strong base with a strong acid
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Titration of a Weak Acid with a Strong Base
With weaker acids, the initial pH is higher and
pH changes near the equivalence point are more
subtle.
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Titrations of Polyprotic Acids
In these cases there is an equivalence point for
each dissociation.
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The Solubility Product
Ex The precipitation of bismuth sulfide.
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The Solubility Product

• Solubility s concentration of Bi2S3
  that dissolves, which equals 1/2Bi3 and
  1/3S2?.
• Ksp is constant (at a given temp)
• s is variable (especially with a common ion
  present)
• Lets Practice

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Factors that Affect Solubility

• Solubility and pH
• Again we apply Le Châteliers principle
• If the F- is removed, then the equilibrium shifts
  towards the decrease 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 can be dramatic.

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Factors Affecting Solubility

• The Common-Ion Effect
• If one of the ions in a solution equilibrium is
  already dissolved in the solution, the
  equilibrium will shift to the left and the
  solubility of the salt will decrease.

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Factors Affecting Solubility

• pH
• If a substance has a basic anion, it will be more
  soluble in an acidic solution.
• Substances with acidic cations are more soluble
  in basic solutions.

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Factors Affecting Solubility

• Complex Ions
• Metal ions can act as Lewis acids and form
  complex ions with Lewis bases in the solvent.

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Factors Affecting Solubility

• Complex Ions
• The formation of these complex ions increases the
  solubility of these salts.

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Factors that Affect Solubility

• Formation of Complex Ions Continued
• Consider the formation of Ag(NH3)2
• The equilibrium constant for the reaction is
  called the formation constant, Kf
• The stability of the complex ion can be judged by
  the size of the Kf

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Factors that Affect Solubility
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Equilibria Involving Complex Ions

• Complex Ion A charged species consisting of
  a metal ion surrounded by ligands
• (Lewis bases).
• Coordination Number Number of ligands
  attached to a metal ion.
• (Most common are 6 and 4.)
• Formation (Stability) Constants The
  equilibrium constants characterizing the
  stepwise addition of ligands to metal ions.

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Factors Affecting Solubility

• Amphoterism
• Amphoteric metal oxides and hydroxides are
  soluble in strong acid or base, because they can
  act either as acids or bases.
• Examples of such cations are Al3, Zn2, and Sn2.

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Will a Precipitate Form?

• In a solution,
• If Q Ksp, the system is at equilibrium and the
  solution is saturated.
• If Q lt Ksp, more solid will dissolve until Q
  Ksp.
• If Q gt Ksp, the salt will precipitate until Q
  Ksp.

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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.