Applications of Aqueous Equilibria

1
Applications of Aqueous Equilibria

• Chapter 15

2
Common Ion Effect

• When a solution of a weak electrolyte is altered
  by adding one of its ions from another source,
  the ionization of the weak electrolyte is
  suppressed. This is referred to as the common ion
  effect.
• A shift in equilibrium occurs because of the
  addition of an ion already involved in the
  equilibrium reaction.

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Effect on pH of an Acid

• Solutions that contain a weak acid plus a salt of
  a weak acid are always less acidic than solutions
  that contain the same concentration of the weak
  acid alone. (pH increases).
• The presence of the ion from the salt shifts
  equilibrium to the left which reduces the
  concentration of H3O and increases the pH.

4
Effect on pH of a Base

• Solutions that contain a weak base plus the salt
  of a weak base are always less basic than
  solutions that contain the same concentration of
  the weak base alone. (pH decreases).
• The presence of the ion from the salt shifts
  equilibrium to the left which reduces the
  concentration of OH- and increases the pH.

5
Buffered Solutions

• A buffered solution is one that resists a change
  in the pH when either hydroxide or hydronium ions
  are added.
• A buffered solution may contain a weak acid and
  its salt (ex. HF and NaF) or a weak base and its
  salt (ex. NH3 and NH4Cl)

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pH of a Buffered Solution

• Calculate the pH of a buffered solution
  containing 0.75 M lactic acid, HC3H5O3 (Ka1.4 x
  10-4) and 0.25 M sodium lactate, NaC3H5O3.

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Henderson-Hasselbalch Equation

• pH pKa log conjugate base/acid
• Using the information from the previous problem,
  calculate the pH using this equation.

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Selecting Buffers

• When choosing the buffering components for a
  specific application, the ratio of A-/HA
  should be as close to 1 as possible.
• The pKa of the acid to be used in the buffer
  should be as close as possible to the desired pH.
• As long as the ratio of A-/HA remains the
  same, the pH will not be affected.
• Increasing the concentrations of A- and HA
  determines the buffering capacity (the amount of
  H or OH- that can be absorbed without a
  significant change in pH.

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A chemist needs a solution buffered at pH 4.30
and can choose from the following acids (and
their sodium salts)

• Chloroacetic acid (Ka 1.35 x 10-3)
• Propanoic acid (Ka 1.3 x 10-5)
• Benzoic acid (Ka 6.4 x 10-5)
• Hypochlorous acid (Ka 3.5 x 10-8)
• Calculate the ratio HA/A- required for each
  system to yield a pH of 4.30. Which system will
  work best?

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Solubility Equilibria

• Calcium fluoride dissolves in water as follows
• CaF2 ? Ca2 2F-
• When the salt first dissolves, no ions are
  present. However, as the dissolution proceeds,
  the concentrations of the ions increase making it
  more likely that these ions will collide and
  reform the solid phase.
• Ultimately, dynamic equilibrium is achieved.
• CaF2 lt -- gt Ca2 2F- At this point no more
  solid will dissolve and the solution is said to
  be saturated.

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Solubility Product Constant

• An equilibrium expression can be written for the
  process
• Ksp Ca2F-2
• The Ksp is called the solubility product constant
  or simply the solubility product.
• CaF2 is a pure solid and is not included in the
  expression.

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Relative Solubilities

• The greater the Ksp for a salt, the greater the
  solubility (if the salts being compared have the
  same number of ions).
• The solubility of a salt is lowered if the
  solution already contains ions common to the
  solid (common ion effect).

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pH and Solubility

• The pH can affect a salts solubility.
• Mg(OH)2 lt -- gt Mg2 2OH-
• The addition of OH- (an increase in pH) will
  decrease the solubility (equilibrium shifts to
  the left).
• The addition of H (a decrease in pH will
  increase the solubility (equilibrium shifts to
  the right)

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To precipitate or not precipitate

• Consider the following expression
• Ksp Ca2F-2
• If we add a solution containing Ca2 to a
  solution containing F-, a precipitate may or may
  not form depending on the concentrations of the
  solutions.
• To predict whether a precipitate will form,
  consider the relationship between Q and Ksp

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Reaction Quotient

• When given a set of reaction components, it is
  helpful to know if the mixture is at equilibrium
  or, if not, in what direction the system must
  shift to reach equilibrium.
• To determine the direction of the move toward
  equilibrium, we use the reaction quotient (Q).
• The reaction quotient is obtained by applying the
  law of mass action to the initial concentrations
  instead of equilibrium concentrations.

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Reaction Quotient

• If QgtKsp , precipitation occurs and will continue
  until the concentrations are reduced to the point
  that they satisfy Ksp (equilibrium shifts to
  the left).
• If QltKsp , precipitation will not occur.
  (equilibrium shifts to the right)

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Titrations and pH Curves

• A pH curve is the plot of the pH of the solution
  being analyzed as a function of the amount of
  titrant added.
• The titrant is the solution of known
  concentration.
• The equivalence point occurs when the moles of
  acid are equal to the moles of base and is often
  signaled by a color change of an indicator.

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Titration Curves

• A titration curve is a plot of pH vs. amount
  (volume) of acid or base added.
• Used to find the molarity of an unknown solution.
• Endpoint-the point at which the indicator changes
  color
• Equivalence Point-point at which chemically
  equivalent amounts of acid and base have reacted.
• Ideally, the endpoint and equivalence point
  should coincide.

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Acid-Base Indicators

• Indicators are organic dyes whose color depends
  on the H3O or pH.
• Common examples include litmus and
  phenolphthalein. (see p. 715 for others)
• Indicators typically change color over a range of
  1.5-2.0 pH units.

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Interpreting Color Changes in Indicators

• Bromthymol Blue
• HIn H2O ? H3O In- Ka 1.0 x
  10-7
• yellow blue
• Color 1 Color 2
• pHlt6 pHgt8
• General rule In-/HIn gt 10 color 2
• In-/HIn lt .10 color 1

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Universal Indicators

• Universal indicators are a mixture of several
  indicators.
• These indicators display a continuous range of
  colors over a wide range of pH values.

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Strong Acid-Strong Base Titrations

• The net ionic equation for a strong acid-strong
  base titration is
• H OH- ? H2O
• To determine the concentration of H at any given
  point, the amount of H that remains must be
  divided by the total volume of the solution.
• Since titrations usually involve small
  quantities, we will use the millimole as our
  unit.
• Number of millimoles volume in mL x molarity

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Case Study-Strong Acid-Strong Base Titration

• Situation 50.0 mL of 0.200 M HNO3 is being
  titrated with 0.100 M NaOH.
• We will calculate the pH of the solution at
  various points during the course of the
  titration.
• We will then draw a pH curve of the data.

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A. No NaOH has been added.

• HNO3 ? H NO3-
• 0.200 M? 0.200 M 0.200M
• HNO3 is a strong acid so it completely ionizes.
• pH -log(0.200) 0.699

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Additional Steps

• B 10.0 mL of 0.100 M NaOH has been added.
• C 20.0 mL of 0.100 M NaOH has been added.
• D 50.0 mL of 0.100 M NaOH has been added.
• E 100.0 mL of 0.100 M NaOH has been
• added.
• F 150.0 mL of 0.100 M NaOH has been added.
• G 200.0 mL of 0.100 M NaOH has been added.

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Characteristics of the pH curve for the titration
of a strong acid with a strong base.

• Before the equivalence point, the H can be
  calculated by dividing the number of millimoles
  of H remaining by the total volume of the
  solution.
• At the equivalence point, the pH 7.00
• After the equivalence point, the OH- can be
  calculated by dividing the number of millimoles
  of excess OH- by the total volume of the
  solution. The H (and pH) can be determined
  using Kw.
• The titration of a strong base with a strong acid
  requires similar reasoning except it will be
  reversed.

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Titration Curve for the reaction of a strong acid
and strong base.
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Affect of procedural errors during titrations

• How would the following lab errors affect the
  calculated values of the molarity of the NaOH?
• If the buret was rinsed with distilled water
  immediately prior to the titration with the NaOH.
• If bromophenol blue was used as the indicator
  instead of phenolphthalein.

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Titrations of Weak Acids and Strong Bases

• Calculation of the pH curve for a titration of a
  weak acid and strong base really amounts to a
  series of buffer problems.
• Reminder Even though the acid is weak, it reacts
  essentially to completion with the strong base.

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Case Study Weak Acid and Strong Base Titration

• Consider the titration of 50.0 mL of 0.10 M
  HC2H3O2 (Ka - 1.8 x 10-5) with 0.10 M NaOH.
  Calculate the pH at the following points in the
  titration
• A No NaOH has been added.
• B 10.0 mL of 0.10 M NaOH has been added.
• C 25.0 mL of 0.10 M NaOH has been added.
• D 40.0 mL of 0.10 M NaOH has been added.
• E 50.0 mL of 0.10 M NaOH has been added.
• F 60.0 mL of 0.10 M NaOH has been added.
• G 75.0 mL of 0.10 M NaOH has been added.

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Conclusions

• It is the amount of acid, not its strength, that
  determines the equivalence point.
• The pH value at the equivalence point is affected
  by the strength. (The stronger the acid, the
  lower the pH at the equivalence point). (see p.
  707)
• For the titration of a weak acid and strong base,
  the equivalence point will be greater than 7.
• At the halfway point in the titration, the pH
  pKa.

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pH curve for strong base and weak acid.
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Titrations of Weak Bases and Strong Acids

• See the example on page709-710.
• For the titration of a weak base and strong acid,
  the equivalence point will be lower than 7.
• At the halfway point in the titration, the pOH
  pKb

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pH curve for titration of strong acid and weak
base.