Aqueous Equilibria

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Aqueous Equilibria

• By Chris Via

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Common-Ion Effect

• C.I.E.- the dissociation of a weak electrolyte by
  adding to the solution a strong electrolyte that
  has an ion in common with the weak electrolyte.
• The addition of this strong electrolyte causes
  the equilibrium to shift as well as the pH to
  change.

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Example

• The dissociation of HC2H3O2 is
• HC2H3O2 H C2H3O2-
• but if NaC2H3O2 is added to the solution, the
  addition of C2H3O2- from NaC2H3O2 causes the
  equilibrium to shift to the left and also
  reducing the concentration of H, therefore
  raising the pH of the solution.

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Practice Problem

• Calculate the fluoride -ion concentration and pH
  of a solution containing .1 mol of HCl and .2 mol
  of HF in 1.0L solution. (Ka of HF 6.810-4)
• Now you want to set up an ice box of the
  dissociation of HF.

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Continued
HF
H
F-
Ka6.810-4 HF-/HF(.1x)x/(.2-x)
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Solution

• In this case we can assume x is too small to make
  a difference so (.1x) becomes just .1
• So we are left with .1x/.2 6.810-4
• x1.410-3 M F-
• H (.1x).1M
• -log(H)-log(.1)1
• The pH1.00 and the F- 1.410-3 M

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Buffered Solutions

• Buffers are solutions that resist changes in pH
  upon the addition of small amounts of acids or
  bases.
• A key example of a buffer is human blood which
  will stay around a pH of 7.4.
• Buffers resist changes in pH because they contain
  both an acidic species to neutralize OH- ions and
  a basic species to neutralize H ions.

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Buffers

• However, it is key that the acid and base species
  of the buffer do not neutralize each other.
• To do this, buffers are often prepared by mixing
  a weak acid or base with a salt of that acid or
  base.
• Ex, the HC2H3O2-C2H3O2- buffer can be prepared by
  adding NaC2H3O2 to a solution of HC2H3O2.

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Buffer Capacity

• Buffer Capacity is the amount of an acid or base
  the buffer can neutralize before the pH begins to
  change significantly.
• The greater the amounts of conjugate acid-base
  pair, the more resistant the solution is to
  change in pH.
• Henderson-Hasselbalch equation is used to
  calculate the pH of buffers
• pH pKa logbase/acid

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Sample Exercise

• What is the pH of a buffer that is .12 M lactic
  acid and .1 M sodium lactate? (Ka for lactic acid
  1.410-4)
• You want to start out with an ice box of the
  dissociation of lactic acid
• lactic acid hydrogen lactate

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Problem Continued

• Ka 1.410-4x(.1x)/(.12-x)
• Once again in this case x will be too small so it
  can be ignored.
• X1.710-4H
• pH -log(1.710-4) 3.77
• With the Henderson-Hasselbalch equation
• pH pKa log(base/acid)
• pH 3.85 (-.08) 3.77

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Acid-Base Titrations

• An acid-base titration is when an acid is added
  to a base or vice versa.
• Acid-Base indicators are usually used to
  determine the equivalence point which is the
  point at which stoichiometrically equivalent
  quantities of acid and base have been brought
  together.
• When strong acids and bases are mixed the pH
  changes very dramatically near the equivalence
  point
• a single drop a this point could change the pH by
  a number of units.

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Titrations
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Polyprotic Acids

• When titrating with Polyprotic acids or bases the
  substance has multiple equivalence points.
• So for example, in a titration of Na2CO3 with
  HCl. There are two distinct equivalence points on
  the titration curve.

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Solubility-Product Constant

• Ksp Solubility-Product Constant
• It expresses the degree to which the solid is
  soluble in water.
• The equation for Ksp is Ksp ionaionb
• An example is the expression of the Ksp of BaSO4
• Ksp Ba2SO42-

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Factors that Affect Solubility

• They are
• the presence of common ions
• the pH of the solution
• presence of complexing agents
• The presence of common ions in a solution will
  reduce the solubility and make the equilibrium
  shift left.

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pH and Solubility

• The general rule of solubility and pH is
• The solubility of slightly-soluble salts
  containing basic anions increases as the pH of
  the solution is lowered.
• This is because the OH- ion is insoluble in
  water while the H ion is very soluble, therefore
  when a basic solution has a low concentration of
  OH- ions the salt will be easier to dissolve.

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Continued

• The more basic the anion, the more the solubility
  is influenced by the pH of the solution.
• However, salts with anions of strong acids are
  unaffected by changes in pH.

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Complex Ions

• A characteristic of most metal ions is their
  ability to act as a Lewis acid when interacting
  with water (Lewis base).
• However, when these metal ions interact with
  Lewis bases other than water, the solubility of
  that metal salt changes dramatically.

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Complex Ion

• Complex Ion-when a metal ion (Lewis acid) is
  bonded together with a Lewis base other than
  water.
• Ex. Ag(NH3)2, Fe(CN)6-4
• The stability of a complex ion can be judged by
  the size of the equilibrium constant for its
  formation.
• This is called the formation constant, Kf
• the larger Kf the more stable the ion is.

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Amphoterism

• Amphoteric-a metal hydroxide that is capable of
  being dissolved in strong acids or strong bases,
  but not in water because it can act like an acid
  or a base.
• Ex. Al3, Cr3, Zn2, and Sn2
• However, these metal ions are more accurately
  expressed as Al(H2O)63.
• This is a weak acid and as it is added to a
  strong base it loses protons and eventually forms
  the neutral and water-soluble Al(H2O)3(OH)3

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Precipitation of Ions

• The reaction quotient, Q, can with the
  solubility-product constant to determine if a
  precipitation will occur.
• If Q gt Ksp-precipitation occurs until Q Ksp
• If Q Ksp-equilibrium exists (saturated
  solution)
• If Q lt Ksp-solid dissolves until Q Ksp
• Q is sometimes referred to as the ion product
  because there is no denominator.

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Selective Precipitation

• Selective Precipitation-the separation of ions in
  an aqueous solution by using a reagent that forms
  a precipitate with one or more of the ions.
• The sulfide ion is widely used to separate metal
  ions because the solubilities of the sulfide
  salts span a wide range and are dependent on the
  pH of the solution.

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Continued

• An example is a solution containing both Cu2 and
  Zn2 mixed with H2S gas.
• CuS ends up forming a precipitate, but ZnS does
  not.
• This is because CuS has a Ksp value of 610-37
  making it less soluble than ZnS which has a Ksp
  value of 210-25

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Qualitative Analysis

• Qualitative Analysis determines the presence or
  absence of a particular metal ion, whereas
  quantitative analysis determines how much of a
  certain substance is present or produced.
• There are 5 main groups of metal ions
• Insoluble Chlorides, Acid-insoluble sulfides,
  base-insoluble sulfides and hydroxides, insoluble
  phosphates, and the alkali metal ions and NH4
• These can be found on page 673 in your textbook