Applications of Aqueous Equilibria

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Applications of Aqueous Equilibria

• Chemistry 4th Edition
• McMurry/Fay

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Buffer Solutions 01

• A Buffer Solution is a solution of a weak acid
  and a weak base (usually conjugate pair) both
  components must be present.
• A buffer solution has the ability to resist
  changes in pH upon the addition of small amounts
  of either acid or base.
• Buffers are very important to biological systems!

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Buffer Solutions 02
Acid is neutralized by the weak base.
Base is neutralized by the weak acid.
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Buffer Solutions 03

• Buffer solutions must contain relatively high
  concentrations of weak acid and weak base
  components to provide a high buffering
  capacity.
• The acid and base components must not neutralize
  each other.
• The simplest buffer is prepared from equal
  concentrations of an acid and its conjugate base.

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Buffer Solutions 04

• a) Calculate the pH of a buffer system containing
  1.0 M CH3COOH and 1.0 M CH3COONa.
• b) What is the pH of the system after the
  addition of 0.10 mole of HCl to 1.0 L of buffer
  solution?

a) pH of a buffer system containing 1.0 M CH3COOH
and 1.0 M CH3COO Na is 4.74 (see
earlier slide)
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Buffer Solutions 04

• b) What is the pH of the system after the
  addition of 0.10 mole of HCl to 1.0 L of buffer
  solution?

CH3CO2H(aq) H2O(aq) CH3CO2(aq)
H3O(aq) I 1.0 1.0 0
add 0.10 mol HCl 0.10
0.10 C -x -x
x E 1.1 x 0.9 x x
(0.9x)(x) (0.9)x (1.1 x) (1.1)
Ka 1.8 x 10-5
pH 4.66
x (1.1/0.9) 1.8 x 10-5 x 2.2 x 10-5
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Water No Buffer

• What is the pH after the addition of 0.10 mole of
  HCl to 1.0 L of pure water?

HCl strong acid, completely ionized. H
concentration will be 0.10 molar. pH will be
log(0.10) 1.0
The power of buffers!
Adding acid to water ?pH 7.0 - 1.0
6.0 pH units
Adding acid to buffer ?pH 4.74 - 4.66
0.09 pH units!!
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AcidBase Titrations 01

• Titration a procedure for determining the
  concentration of a solution using another
  solution of known concentration.
• Titrations involving strong acids or strong bases
  are straightforward, and give clear endpoints.
• Titration of a weak acid and a weak base may be
  difficult and give endpoints that are less well
  defined.

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AcidBase Titrations 02

• The equivalence point of a titration is the point
  at which equimolar amounts of acid and base have
  reacted. (The acid and base have neutralized
  each other.)
• For a strong acid/strong base titration, the
  equivalence point should be at pH 7.

H (aq) OH (aq) ? H2O (l)
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AcidBase Titrations 02

• The equivalence point of a titration is the point
  at which equimolar amounts of acid and base have
  reacted. (The acid and base have neutralized
  each other.)
• Titration of a weak acid with a strong base
  gives an equivalence point with pH gt 7.

HA (aq) OH (aq) ? H2O (l) A (aq)
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AcidBase Titration Examples 03

• Titration curve for strong acidstrong base

Note the very sharp endpoint (vertical line) seen
with strong acid strong base titrations. The
pH is changingvery rapidly in this region.
add one drop ofbase get a BIG change in pH
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Titration of 0.10 M HCl 03
pH

• Volume of Added NaOH
• 1. Starting pH (no NaOH added) 1.00
• 2. 20.0 mL (total) of 0.10 M NaOH. 1.48
• 3. 30.0 mL (total) of 0.10 M NaOH. 1.85
• 4. 39.0 mL (total) of 0.10 M NaOH. 2.90
• 5. 39.9 mL (total) of 0.10 M NaOH. 3.90
• 6. 40.0 mL (total) of 0.10 M NaOH. 7.00
• 7. 40.1 mL (total) of 0.10 M NaOH. 10.10
• 8. 41.0 mL (total) of 0.10 M NaOH. 11.08
• 9. 50.0 mL (total) of 0.10 M NaOH 12.05

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AcidBase Titrations 04

• Titration curve for weak acidstrong base

The endpoint (vertical line) is less sharp with
weak acid strong base titrations.
weak acid equivalence point
strong acid equivalence point
pH at the equivalence point will always be gt7 w/
weak acid/strong base
weak acid
strong acid
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AcidBase Titration Curves
very weak acid
With a very weak acid, the endpoint may be
difficult to detect.
weak acid
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AcidBase Titrations 09

• Strong AcidWeak Base

The (conjugate) acid hydrolyzes to form weak
base and H3O. At equivalence point only the
(conjugate) acid is present. pH at equivalence
point will always be lt7.
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Solubility Equilibria 01
Aqueous Solubility Rules for Ionic Compounds A
compound is probably soluble if it contains the
cations a. Li, Na, K, Rb (Group 1A
on periodic table) b. NH4 A compound is
probably soluble if it contains the anions a.
NO3 (nitrate), CH3CO2 (acetate, also written
C2H3O2) b. Cl, Br, I (halides)
except Ag, Hg22, Pb2 halides c. SO42
(sulfate) except Ca2, Sr2, Ba2, and Pb2
sulfates Other ionic compounds are probably
insoluble.
Old way to analyze solubility. Answer is
either yes or no.
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Solubility Equilibria 02

• New method to measure solubilityConsider
  solution formation an equilibrium process
• MCl2(s) ? M2(aq) 2 Cl(aq)

Give equilibrium expression Kc for this equation
KC M2Cl2
This type of equilibrium constant Kc that
measures solubility is called Ksp
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Solubility Equilibria 03

• Solubility Product is the product of the molar
  concentrations of the ions and provides a measure
  of a compounds solubility.
• MX2(s) ? M2(aq) 2 X(aq)
• Ksp M2X2

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Solubility Equilibria - Ksp Values 04
CuI 5.1 x 1012 Cu(OH)2 2.2 x 1020 CuS 6.0 x
1037 Fe(OH)2 1.6 x 1014 Fe(OH)3 1.1 x
1036 FeS 6.0 x 1019 PbCO3 3.3 x 1014 PbCl2 2.4
x 104 PbCrO4 2.0 x 1014 PbF2 4.1 x
108 PbI2 1.4 x 108 PbS 3.4 x 1028 MgCO3 4.0 x
105 Mg(OH)2 1.2 x 1011

• Al(OH)3 1.8 x 1033
• BaCO3 8.1 x 109
• BaF2 1.7 x 106
• BaSO4 1.1 x 1010
• Bi2S3 1.6 x 1072
• CdS 8.0 x 1028
• CaCO3 8.7 x 109
• CaF2 4.0 x 1011
• Ca(OH)2 8.0 x 106
• Ca3(PO4)2 1.2 x 1026
• Cr(OH)3 3.0 x 1029
• CoS 4.0 x 1021
• CuBr 4.2 x 108

MnS 3.0 x 1014 Hg2Cl2 3.5 x 1018 HgS 4.0 x
1054 NiS 1.4 x 1024 AgBr 7.7 x 1013 Ag2CO3
8.1 x 1012 AgCl 1.6 x 1010 Ag2SO4 1.4 x
105 Ag2S 6.0 x 1051 SrCO3 1.6 x 109 SrSO4
3.8 x 107 SnS 1.0 x 1026 Zn(OH)2 1.8 x
1014 ZnS 3.0 x 1023
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Solubility Equilibria 05

• The solubility of calcium sulfate (CaSO4) is
  found experimentally to be 0.67 g/L. Calculate
  the value of Ksp for calcium sulfate.
• The solubility of lead chromate (PbCrO4) is
  4.5 x 105 g/L. Calculate the solubility
  product of this compound.
• Calculate the solubility of copper(II) hydroxide,
  Cu(OH)2, in g/L.

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Equilibrium Constants - Review

• The reaction quotient (Qc) is obtained by
  substituting initial concentrations into the
  equilibrium constant. Predicts reaction
  direction.Qc lt Kc System forms more products
  (right)Qc Kc System is at equilibriumQc gt
  Kc System forms more reactants (left)

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Equilibrium Constants - Qc

• Predicting the direction of a reaction.

Qc gt Kc
Qc lt Kc
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Solubility Equilibria 06

• We use the reaction quotient (Qc) to determine if
  a chemical reaction is at equilibrium
  compare Qc and Kc

Ksp values are also a type of equilibrium
constant, but are valid for saturated solutions
only.
We can use ion product (IP) to determine
whether a precipitate will form compare IP
and Ksp
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Solubility Equilibria 06

• Ion Product (IP) solubility equivalent of
  reaction quotient (Qc). It is used to determine
  whether a precipitate will form.

IP lt Ksp IP Ksp IP gt Ksp
Unsaturated (more solute can dissolve) Saturated
solution Supersaturated precipitate forms.
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Solubility Equilibria 07

• A BaCl2 solution (200 mL of 0.0040 M) is added to
  600 mL of 0.0080 M K2SO4. Will precipitate
  form? (Ksp for BaSO4 is 1.1 x 10-10)

Ba2
0.200 L x 0.0040 mol/L .00080 moles Ba2
Ba2 0.00080 mol ? 0.800 L 0.0010 M
SO42
0.600 L x 0.0080 mol/L .00480 moles SO42
SO42 0.00480 mol ? 0.800 L 0.0060 M
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Solubility Equilibria 07

• A BaCl2 solution (200 mL of 0.0040 M) is added to
  600 mL of 0.0080 M K2SO4. Will precipitate
  form? (Ksp for BaSO4 is 1.1 x 10-10)

IP Ba2 1 x SO42 1 (0.0010)
x (0.0060) 6.00 x 10-6
IP gt Ksp, so ppt forms
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Solubility Equilibria 07

• Exactly 200 mL of 0.0040 M BaCl2 are added to
  exactly 600 mL of 0.0080 M K2SO4. Will a
  precipitate form?
• If 2.00 mL of 0.200 M NaOH are added to 1.00 L of
  0.100 M CaCl2, will precipitation occur?

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The Common-Ion Effect and Solubility

• The solubility product (Ksp) is an equilibrium
  constant precipitation will occur when the ion
  product (IP) exceeds the Ksp for a compound.
• If AgNO3 is added to saturated AgCl, the increase
  in Ag will cause AgCl to precipitate.
• IP Ag0 Cl0 gt Ksp

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The Common-Ion Effect and Solubility
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The Common-Ion Effect and Solubility
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The Common-Ion Effect and Solubility

• Calculate the solubility of silver chloride (in
  g/L) in a 6.5 x 103 M silver chloride solution.
• Calculate the solubility of AgBr (in g/L) in
  (a) pure water (b) 0.0010 M NaBr