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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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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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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Buffered Solutions

• Lets practice with 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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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

• 2. 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 that Affect Solubility

• 3. Formation of Complex Ions
• Consider the addition of ammonia to AgCl (white
  precipitate)
• The overall reaction is
• The Ag(NH3)2 is called a complex ion
• NH3 (the attached Lewis base) is called a ligand.
• Effectively, the Ag(aq) has been removed from
  solution.

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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 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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Precipitation and Separation of Ions

• At any instant in time, Q Ba2SO42-.
• If Q gt Ksp, precipitation occurs until Q Ksp.
• If Q Ksp, equilibrium exists.
• If Q lt Ksp, solid dissolves until Q Ksp.
• Based on solubilities, ions can be selectively
  removed from solutions.

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Precipitation and Separation of Ions

• 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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• Qualitative analysis is designed to detect the
  presence of metal ions.
• Quantitative analysis is designed to determine
  how much metal ion is present.