CHEMISTRY Chapter 15 Applications of Aqueous Equilibria

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CHEMISTRYChapter 15Applications of Aqueous
Equilibria
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• 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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The pH scale

• The pH scale ranges from 0 to 14.
• Acids have a pH less than 7.
• A base has a pH greater than 7.
• Pure water has a pH equal to 7.

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Acids and bases in your body

• Many reactions, such as the ones that occur in
  your body, work best at specific pH values.

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pH and blood

• The pH of your blood is normally within the range
  of 7.37.5.
• Holding your breath causes blood pH to drop.
• High blood pH can be caused by hyperventilating.

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Acids Bases

• Strong acids
• Know the names and formulas of the 7 common
  strong acids
• HCl (aq) hydrochloric acid
• HBr (aq) hydrobromic acid
• HI (aq) hydroiodic acid
• HClO3 chloric acid
• HClO4 perchloric acid
• HNO3 nitric acid
• H2SO4 sulfuric acid

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Acids Bases

• Examples of Weak Acids
• HF (aq) hydrofluoric acid
• H3PO4 phosphoric acid
• CH3COOH acetic acid

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Acids Bases

• Strong Bases
• Know the names and formulas of the strong bases
• Alkali metal (1A) hydroxides
• LiOH lithium hydroxide
• NaOH sodium hydroxide
• KOH potassium hydroxide
• RbOH rubidium hydroxide
• CsOH cesium hydroxide

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Acids Bases

• Strong bases to know (cont)
• Heavy alkaline earth metal (2A) hydroxides
• Ca(OH)2 calcium hydroxide
• Sr(OH)2 strontium hydroxide
• Ba(OH)2 barium hydroxide

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Acids Bases

• Examples of Weak Bases
• ammonia (NH3)
• sodium bicarbonate (NaHCO3)
• baking soda
• a component of Alka-Seltzer

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Acids Bases

• Generally, when solutions of an acid and a base
  are combined, the products are a salt and water.

• CH3COOH (aq) NaOH (aq) ??CH3COONa (aq) H2O
  (l)
• HCl (aq) NaOH (aq) ?? NaCl (aq) H2O (l)
• All neutralization reactions are double
  displacement reactions

Neutralization Reaction
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Neutralization Reactions

• When a
• strong acid reacts with a strong base,
• the net ionic equation is
• HCl (aq) NaOH (aq) ?? NaCl (aq) H2O (l)
• H (aq) Cl- (aq) Na (aq) OH-(aq) ??
• Na (aq) Cl- (aq) H2O (l)
• H (aq) OH-(aq) ?? H2O (l)

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Strong Acids and Bases
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Weak Acids and Strong Bases
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Strong Acid Weak Base
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Weak Acid Weak Base
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COMMON-ION EFFECT

• A shift in equilibrium due to addition of same
  ion salt to an aqueous weak acid or weak base is
  the common-ion effect.
• This is an example of
• Le Chatliers Principle.

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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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The Common-Ion Effect

• Common Ion Two dissolved solutes that contain
  the same ion (cation or anion).
• The presence of a common ion suppresses the
  ionization of a weak acid or a weak base.
• Common-Ion Effect is the shift in equilibrium
  caused by the addition of a compound having an
  ion in common with the dissolved substance.

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The Common-Ion Effect

• Q Calculate the pH of a 0.20 M CH3COOH solution
  with no salt added.
• Q Calculate the pH of a solution containing 0.20
  M CH3COOH and 0.30 M CH3COONa.
• Q What is the pH of a solution containing 0.30 M
  HCOOH, before and after adding 0.52 M HCOOK?

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Buffer Solutions

• A Buffer Solution is a solution of
• (1) a weak acid or a weak base and
• (2) its salt
• 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
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How do buffers work?
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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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HendersonHasselbalch equation

• To determine the pH, we apply I.C.E. and then the
  HendersonHasselbalch equation.
• When the concentration of HA and salt are high
  (0.1 M) we can neglect the ionization of acid
  and hydrolysis of salt.

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HendersonHasselbalch equation
Henderson-Hasselbalch equation
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Buffer Solutions

• Buffer solutions must contain relatively high
  acid and base component concentrations, the
  buffer capacity.
• Acid and base component concentrations must not
  react together.
• The simplest buffer is prepared from equal
  concentrations of acid and conjugate base.

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Buffer Solutions

• Buffer Preparation Use the HendersonHasselbalch
  equation in reverse.
• Choose weak acid with pKa close to required pH.
• Substitute into HendersonHasselbalch equation.
• Solve for the ratio of conjugate base/acid.
• This will give the mole ratio of conjugate base
  to acid. The acid should always be 1.0.

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Buffer Solutions

• Q Describe how you would prepare a phosphate
  buffer with a pH of about 7.40.
• Q How would you prepare a liter of carbonate
  buffer at a pH of 10.10? You are provided with
  carbonic acid (H2CO3), sodium hydrogen carbonate
  (NaHCO3), and sodium carbonate (Na2CO3).

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Buffer Solutions

• Q Calculate the pH of a buffer system containing
  1.0 M CH3COOH and 1.0 M CH3COONa. What is the pH
  of the system after the addition of 0.10 mole of
  gaseous HCl to 1.0 L of solution?
• Q Calculate the pH of 0.30 M NH3/0.36 NH4Cl
  buffer system. What is the pH after the addition
  of 20.0 mL of 0.050 M NaOH to 80.0 mL of the
  buffer solution?

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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
• 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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Solubility Equilibria

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

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Solubility Equilibria

• 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

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
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

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

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Solubility Equilibria

• Ion Product (Q) solubility equivalent of the
  reaction quotient. It is used to determine
  whether a precipitate will form.
• Q lt Ksp Unsaturated Q Ksp Saturated Q gt
  Ksp Supersaturated precipitate forms.

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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.

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Solubility Equilibria

• Q Exactly 200 mL of 0.0040 M BaCl2 are added to
  exactly 600 mL of 0.0080 M K2SO4. Will a
  precipitate form?
• Q 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 exceeds the Ksp for a compound.
• If AgNO3 is added to saturated AgCl, the increase
  in Ag will cause AgCl to precipitate.
• Q 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

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

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Complex Ion Equilibria and Solubility

• A complex ion is an ion containing a central
  metal cation bonded to one or more molecules or
  ions.
• Most metal cations are transition metals because
  they have more than one oxidation state.
• The formation constant (Kf) is the equilibrium
  constant for the complex ion formation.

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Complex Ion Equilibria and Solubility
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Complex Ion Equilibria and Solubility
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Complex Ion Equilibria and Solubility

• ION Kf
• Ag(NH3)2 1.5 x 107
• Ag(CN)2 1.0 x 1021
• Cu(CN)42 1.0 x 1025
• Cu(NH3)42 5.0 x 1013
• Cd(CN)42 7.1 x 1016
• CdI42 2.0 x 106

ION Kf HgCl42 1.7 x 1016 HgI42 3.0 x
1030 Hg(CN)42 2.5 x 1041 Co(NH3)63 5.0 x
1031 Zn(NH3)42 2.9 x 109
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Complex Ion Equilibria and Solubility

• Q A 0.20 mole quantity of CuSO4 is added to a
  liter of 1.20 M NH3 solution. What is the
  concentration of Cu2 ions at equilibrium?
• Q If 2.50 g of CuSO4 are dissolved in 9.0 x 102
  mL of 0.30 M NH3, what are the concentrations of
  Cu2, Cu(NH3)42, and NH3 at equilibrium?

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Complex Ion Equilibria and Solubility

• Q Calculate the molar solubility of AgCl in a
  1.0 M NH3 solution.
• Q Calculate the molar solubility of AgBr in a
  1.0 M NH3 solution.

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Complex Ion Equilibria and Solubility
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Qualitative Analysis Scheme
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Qualitative Analysis