Additional Aspects of Aqueous Equilibria

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Additional Aspects of Aqueous Equilibria

• BLB 11th Chapter 17

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• Buffered Solutions (sections 1-2)
• Acid/Base Reactions Titration Curves (3)
• Solubility Equilibria (sections 4-5)
• Two important points
• Reactions with strong acids or strong bases go to
  completion.
• Reactions with only weak acids and bases reach an
  equilibrium.

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17.1 The Common Ion Effect
Weak acid HA H2O ? H3O A- Salt of conj. Base NaA ? Na(aq) A-(aq)
two sources of A- Common Ion! two sources of A- Common Ion! two sources of A- Common Ion!

• What affect does the addition of its conjugate
  base have on the weak acid equilibrium? On the
  pH?
• Used in making buffered solutions

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Calculate the pH of a 0.60 M HF solution. The Ka
of HF is 7.210-4.
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Calculate the pH of a solution containing 0.60 M
HF and 1.00 M KF.
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17.2 Buffered Solutions

• Resist a change in pH upon the addition of small
  amounts of strong acid or strong base
• Consist of a weak conjugate acid-base pair
• Control pH at a desired level (pKa)
• Examples blood (p. 729), physiological fluids,
  seawater, foods

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How do buffers work?
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Calculating pH of a Buffer
Henderson-Hasselbalch equation
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Calculate the pH of a solution containing 0.60 M
HF and 1.00 M KF. (again, but the easy way)
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Adding strong acid or base to a buffer

• Adding acid H3O HA or A- ?
• Adding base OH- HA or A- ?
• Calculating pH
• Stoichiometry of added acid or base
• Equilibrium problem (H-H equation)

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Calculate the pH after adding 0.20 mol of HCl to
1.0 L of the 0.60 M HF and 1.00 M KF buffer.
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Calculate the pH after adding 0.10 mol of NaOH to
1.0 L of the 0.60 M HF and 1.00 M KF buffer.
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Calculate the pH for a 1.0-L solution that
contains 0.25 M NH3 and 0.15 M NH4Br. Kb1.8x10-5
for NH3
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Calculate the pH for a 1.0-L solution that
contains 0.25 M NH3 and 0.15 M NH4Br after the
addition of 0.05 mol of RbOH.
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Calculate the pH for a 1.0-L solution that
contains 0.25 M NH3 and 0.15 M NH4Br after the
addition of 0.35 mol of HCl.
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Buffers (wrap up)

• H-H equation
• No 5 check
• When strong acid or base is added, start reaction
  with that acid or base.
• Making buffers of a specific pH? H-H equation
• Buffer capacity exceeded when added acid or
  base totally consumes a buffer component (p. 726)

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How would you prepare a phenol buffer to control
pH at 9.50? Ka 1.3x10-10 for phenol
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17.3 Acid-Base Titrations

• Titration a reaction used to determine
  concentration (acid-base, redox, precipitation)
• Titrant solution in buret usually a strong
  base or acid
• Analyte solution being titrated often the
  unknown
• _at_ equivalence point (or stoichiometric point)
  mol acid mol base
• Found by titration with an indicator
• Solution not necessarily neutral
• pH dependent upon salt formed
• pH titration curve plot of pH vs. titrant volume

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Acid-base Titration Reactions and Curves
Type Acid Base
1 strong strong
2 weak strong
3 strong weak

• Recognize curve types
• Calculate pH at various points on curve.

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Type 1 Strong acid strong base

• Goes to completion
• Forms a neutral salt
• Equivalence point - neutral solution, H3O
  1.0 x 10-7 M, pH 7.00
• pH calculations involve only stoichiometry and
  excess H3O and OH-

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Strong acid Strong base
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Type 1 Strong acid strong base20.0 mL 0.200
M HClO4 titrated with 0.200 M KOH
Initial mmol acid
mL base mmol base added mmol acid remain total mL H3O pH
0.00
10.00
20.00
30.00
40.00
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Another SA/SB titration10.0 mL 0.20 M KOH
titrated with 0.10 M HCl
Initial mmol base
mL acid mmol acid added mol base remain total mL OH- pH
0.00
15.00
20.00
35.00
50.00
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Type 2 Weak acid strong base

• Titration reaction goes to completion
• Forms a basic salt (from conj. base of the weak
  acid)
• Equivalence point - basic solution, pH gt 7.00
• pH calculations involve stoichiometry and
  equilibrium

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Weak acid Strong base
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Type 2 Weak acid strong base25.0 mL 0.100M
HC3H5O2 titrated with 0.100 M KOHKa 1.3x10-5

• Calculate the pH at the following points
• Initial (0.00 mL KOH)
• 10.00 mL KOH
• Midpoint (12.50 mL KOH)
• Equivalence pt. (25.00 mL KOH)
• 10.00 mL after eq. pt. (35.00 mL KOH)

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Polyprotic Weak acid Strong base
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Type 3 Weak base strong acid

• Titration reaction goes to completion
• Forms an acidic salt (from conj. acid of the weak
  base)
• Equivalence point - acidic solution, pH lt 7.00
• pH calculations involve stoichiometry and
  equilibrium

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Strong base Strong acid Weak base Strong acid
Strong base
Weak base
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Type 3 Weak base strong acid 25.0 mL 0.150 M
NH3 titrated with 0.100 M HCl Kb 1.8x10-5

• Calculate the pH at the following points
• Initial (0.00 mL HCl)
• Midpoint (______ mL HCl)
• 25.00 mL HCl
• Equivalence pt. (______ mL HCl)
• 10.00 mL after eq. pt. (______ mL HCl)

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Types 2 3 pH Calculations

• Initial pH same as weak acid or base problem
  (chapter 16)
• Before equivalence point Buffer
• _at_ midpoint half of the weak analyte has been
  neutralized
• weak acid conj. base or weak base
  conj. acid
• H3O Ka and pH pKa
• _at_ equivalence point mol acid mol base
• Beyond equivalence point pH based on excess
  titrant

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Test 2 Summary for Acid/Base problems

1. Weak acid or weak base only (ch. 16)
2. Buffer
3. SA SB Titration
4. WA SB or WB SA Titration

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17.4 Solubility Equilibria

• Solubility maximum amount of material that can
  dissolve in a given amount of solvent at a given
  temperature units of g/100 g or M (ch. 13)
• Insoluble compound compound with a solubility
  less than 0.01 M also sparingly soluble
• Solubility rules are given on p. 125 (ch. 4)
• Dissolution reaches equilibrium in water between
  undissolved solid and hydrated ions

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Solubility Product Constant, Ksp

• Equilibrium constant for insoluble compounds
• Solid salt nor water included in expression
• Appendix D, p. 1116 for values
• BaSO4(s) ? Ba2(aq) SO42-(aq)
• PbCl2(s) ? Pb2(aq) 2 Cl-(aq)

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Solubility Product Calculations

• In concentration tables, x solubility
• Problem types
• solubility ? Ksp
• Ksp ? solubility

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Comparing Salt Solubilities

• Generally solubility ? Ksp ?
• Can only compare Ksp values if the salts produce
  the same number of ions
• If different numbers of ions are produced,
  solubility must be compared.

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17.5 Factors that Affect Solubility

• Common-Ion Effect
• LeChateliers Principle revisited
• Addition of a product ion causes the solubility
  of the solid to decrease, but the Ksp remains
  constant.
• pH
• LeChateliers Principle again!
• Basic salts are more soluble in acidic solution.
• Acidic salts are more soluble in basic solution.
• Environmental example CaCO3 limestone
• Stalactites and stalagmites form due to changing
  pH in the water and thus solubility of the
  limestone. (p. 964)