Applications of Aqueous Equilibria

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Applications of Aqueous Equilibria

• Acid-Base Equilibria
• More Acid-Base Equilibria
• Solubility of Salts
• Formation of Complex Ions

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Applications of Aqueous Equilibria

• Natural processes that involve these equilibria
• weathering of minerals
• uptake of nutrients by plants
• tooth decay
• formation of limestone caverns, stalactites, and
  stalagmites

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Applications of Aqueous Equilibria

• Common Ion Problems
• Involve solutions that contain a weak acid, HA,
  and its salt, NaA
• Ex Given a solution containing HF
  (Ka 7.2 x 10-4) and its salt, NaF. What
  happens to the dissociation of HF?
• NaF is a strong electrolyte, i.e. it ionizes 100
• NaF ---gt Na F-

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Applications of Aqueous Equilibria

• HF is a weak acid
• HF ltgt H F-
• Therefore, the major species in the solution is
  Na, F-, HF, and H2O
• F- is a common ion, since it is produced by both
  HF and NaF.

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Applications of Aqueous Equilibria

• Apply Le Chateliers Principle. The F- present
  from the ionization of NaF forces the HF to
  ionize less as the reverse reaction is favored.
• The acidity of a solution with a common ion
  present is less than a solution of HF alone, i.e.
  the H is less, or the pH is greater.

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Applications of Aqueous Equilibria

• Buffered Solutions
• consists of a solution that contains a weak acid
  and its salt or a weak base and its salt.
• resists changes in pH upon the addition of an
  acid or base
• Ex Our blood is a buffered solution which can
  absorb acids or bases without changing its pH,
  which is important because our cells can only
  survive in a narrow pH range.

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Applications of Aqueous Equilibria

• A solution can be buffered to any pH by choosing
  the appropriate components and concentrations.
• Buffered solutions are still only solutions of
  weak acids or weak bases containing a common ion.
  The pH calculations are the same as presented
  before.

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Applications of Aqueous Equilibria

• How does buffering work?
• Buffered solution contains HA and A-.
• Add OH-
• OH- reacts with the HA
• OH- HA --gt H2O A-
• The OH- gets eaten up as it forms water, so
  with no extra OH- in the system, the pH remains
  the same.

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Applications of Aqueous Equilibria

• Add H
• The H reacts with the A- to reform HA
• H A- --gt HA
• The added H gets eaten up as it reforms HA,
  with no extra H in the system, the pH remains
  the same.

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Applications of Aqueous Equilibria

• Buffered solutions can be formed with a weak
  base, B, and its salt, BH.
• Add OH-
• the OH- will react with the BH
• OH- BH --gt B H2O
• the extra OH- gets eaten up by the H to form
  water.
• With no extra OH- in the system, the pH stays the
  same.

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Applications of Aqueous Equilibria

• Add H
• The H will react with the B to form the salt
• H B --gt BH
• The extra H is eaten up by the B, with no
  extra H in the system, the pH remains the same.

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Applications of Aqueous Equilibria

• Henderson-Hasselbach equation
• Start with Ka expression
• Ka HA-/HA
• Rearrange to get H KaHA/A-
• Take the negative log of both sides
• pH pKa - log(HA/A-) or
• pH pKa log (A-/HA) or
• pH pKa log (base/acid)
• Useful for calculating the pH when the HA/A-
  ratio is known.

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Applications of Aqueous Equilibria

• Important characteristics of buffered solutions
• they contain large concentrations of a weak acid
  and the corresponding weak base
• HA and A- or
• B and BH

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Applications of Aqueous Equilibria

• Important characteristics of buffered solutions
• Add H, it will react with the weak base to form
  a weak acid
• H A- --gt HA or
• H B --gt BH
• Add OH-, it will react with the weak acid to form
  a weak base and water
• OH- HA --gt H2O A- or
• OH- BH --gt H2O B

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Applications of Aqueous Equilibria

• Important characteristics of buffered solutions
• the pH of a buffered solution depends on the
  ratio of the concentrations of the weak acid and
  the weak base.
• While this ratio remains constant, the pH will
  remain constant.

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Applications of Aqueous Equilibria

• This ratio will remain constant as long as the
  concentrations of the buffering materials are
  large compared to the amounts of H or OH- added

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Applications of Aqueous Equilibria

• Buffer Capacity
• the amount of H or OH- a buffer can absorb
  without a significant change in the pH
• a buffer with a large capacity can absorb a large
  amount of H or OH- with only a little change in
  the pH
• a buffer with a large capacity contains a large
  concentration of the weak acid and its salt

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Applications of Aqueous Equilibria

• pH of a buffered solution
• determined by the ratio of A-/HA
• optimal buffering (least change in pH) will occur
  when A- HA
• so pH pKa log (1)
• so pH pKa 0
• so pH pKa
• Pick an acid with a Ka closest to the desired pH
  of the buffered solution.

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Applications of Aqueous Equilibria

• Titration and pH curves
• Titration
• used to determine the concentration or amount of
  an unknown acid or base
• uses a solution of known concentration (the
  titrant)
• uses a buret (for precision)
• uses an indicator to show the endpoint
• can be monitored by plotting pH vs. amount of
  titrant added

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Applications of Aqueous Equilibria

• Strong Acid-Strong Base Titrations
• millimoles may be used because of the small
  quantities used in a titration
• 1000 millimoles 1000mmoles 1 mole
• Molarity moles/liter mmoles/mL
• mmoles Volume (in mL) x Molarity

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Applications of Aqueous Equilibria

• Strong Acid-Strong Base Titration
• Ex Titration of 50.0 mL of 0.200 M HNO3 with
  0.100 M NaOH
• What is the pH at various stages of the
  titration?
• Initially...No NaOH has been added
• pH is determined by H from the 0.200 M HNO3
• pH - log (0.200) 0.699
• 50.0 mL x 0.200 M H 10.0 mmol H present

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Applications of Aqueous Equilibria

• Still the beginning of the reaction10.0 mL of
  0.100 M NaOH has been added
• The added OH- will neutralize an equivalent
  amount of H, I.e., 10.0 ml x 0.100 M or 1.00
  mmole.
• How much H is left? 10.0 mmole - 1.00 mmole
  9.0 mmole H
• What is the concentration of H now?
• 9.0 mmole/ (50.0 mL 10.0 mL) 0.15 M
• pH -log (0.15) 0.82pH is increasing

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Applications of Aqueous Equilibria

• 20.0 mL of NaOH has been added
• pH 0.942
• 50.0 mL of NaOH as been added
• pH 1.301

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Applications of Aqueous Equilibria

• Strong Acid-Strong Base Titration
• At the equivalence point100.0 mL of NaOH has
  been added
• 100.0 mL x 0.100 M 10.0 mmole OH-
• This is enough OH- to completely react with the
  H in solution.
• At the equivalence point, the pH is 7, the
  solution is neutral

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Applications of Aqueous Equilibria

• 150.0 mL of NaOH has been added
• now OH- is in excess, the pH is determined by the
  excess OH-
• 150.0 mL x 0.100 M 15.00 mmoles OH-
• 15.00 mmoles OH- - 10.0 mmoles H 5.0 mmoles
  OH- in solution
• 5.0 mmoles/ (50.0 150.0 mL) 0.025 M OH-.so
  H 4.0 x 10-13pH 12.40

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Applications of Aqueous Equilibria

• 200.0 mL of NaOH has been added
• still excess OH-
• pH 12.60

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Applications of Aqueous Equilibria

• Strong Acid-Strong Base Titration
• pH changes very little initially until close to
  the equivalent point (lots of H, added OH-
  doesnt change the pH much)
• Near the equivalence point, there is less H, so
  added OH- changes the pH a lot
• At the equivalence point, the pH 7.00
• Just after the equivalence point, added OH- also
  changes the pH a lot

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Applications of Aqueous Equilibria

• Weak Acid-Strong Base Titration
• Calculating the titration curve is like solving a
  series of buffer problems.
• Involves a stoichiometry problem where the
  reaction goes to completion and the
  concentrations of the weak acid and the conjugate
  base are calculated
• Involves an equilibrium problem, calculate pH
  from this

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Applications of Aqueous Equilibria

• pH curve for this titration is different before
  the equivalence point from the strong acid-strong
  base titration
• after the equivalence point, the titration curves
  are the same

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Applications of Aqueous Equilibria

• For a weak acid-strong base titration, the pH
  rises more rapidly in the beginning of the
  titration, then levels off at the halfway point
  due to buffering effects.
• pH at the equivalence point is higher than for a
  strong acid-strong base titration

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Applications of Aqueous Equilibria

• Weak acid-Strong Base titration
• The amount of acid determines the equivalence
  point
• The pH value at the equivalence point depends on
  the acid strength
• the weaker the acid, the higher the pH at the
  equivalence point

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Applications of Aqueous Equilibria

• Acid-Base Indicators
• Two common methods for determining the
  equivalence point of a titration
• Use a pH meter and then plot the titration curve.
  The center of the vertical region of the pH
  curve indicates the equivalence point.
• Use an acid-base indicator to mark the end point
  with a change in color.

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Applications of Aqueous Equilibria

• Acid-Base Indicators
• The endpoint (when the indicator changes color)
  may not be the same as the equivalence point.
• An indicator must be chosen based on the acid and
  base used in the titration so that the endpoint
  is as close to the equivalence point as possible.

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Applications of Aqueous Equilibria

• Acid-Base Indicators
• Indicators are usually complex molecules that are
  weak acids (HIn)
• HIn is one color while In is another color.
• For example, phenolphthalein is colorless in the
  HIn form, but pink in the In form.

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Applications of Aqueous Equilibria

• HIn (Red) --gt In (blue) H
• Ka HIn-/HIn
• Rearrange to get
• Ka In-
• H HIn
• Lets say the indicators Ka 1 x 10-8
• Add a few drops of the indicator to an acidic
  solution of pH 1.0

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Applications of Aqueous Equilibria

• 1 x 10-8 In-
• 1 x 10-1 HIn
• 1 x 10-7 1 In-
• 10,000,000 HIn
• So in an acidic solution, most of the Indicator
  will be in the red HIn form.
• As OH- is added, the H dcreases, shifting the
  equilibrium so more In- and less HIn is present.
• When will the color change, or rather, when can
  the human eye detect the color change?
• At a pH when In- /HIn 1/10

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Applications of Aqueous Equilibria