Principles of Reactivity: Other Aspects of Aqueous Equilibria

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Principles of Reactivity Other Aspects of
Aqueous Equilibria

• Chapter 17

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Learning Objectives

• Students understand
• The common ion effect
• The control of pH in aqueous solutions with
  buffers
• Students will be able to
• Calculate the pH of buffer solutions
• Evaluate the pH in the course of acid-base
  titrations
• Apply equilibrium concepts to the solubility of
  ionic compounds

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17.1 Common Ion Effect

• The common ion effect is the limiting of the
  ionization of an acid (or base) by the presence
  of a significant concentration of its conjugate
  base (or acid).
• The extent to which the acid can ionize is
  affected, therefore affecting the pH of the
  solution.

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Practice Problem

• Assume you have a 0.30M solution of formic acid
  (HCO2H) and have added enough sodium formate
  (NaHCO2) to make the solution 0.10M in the salt.
  Calculate the pH of the formic acid solution
  before and after adding sodium formate.

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Practice Problem

• What is the pH of the solution that results from
  adding 30.0mL of 0.100M NaOH to 45.0mL of 0.100M
  acetic acid?

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17.2 Controlling pH Buffers

• A buffer causes solutions to be resistant to a
  change in pH on addition of a strong acid or
  base.
• Two substances are needed an acid capable of
  reacting with added OH- ions and a base that can
  consume added H3O ions.
• The acid and base must not react with each other

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Buffers

• A buffer is usually prepared from a conjugate
  acid-base pair.
• The action of a buffer is a special case of the
  common ion effect.

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Buffers

• Weak acid and its conjugate base
• H3O acid/conjugate base Ka
• Another form of the same equation
• pH pKa log conjugate base/acid
• known as the Henderson-Hasselbalch equation

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Henderson-Hasselbalch equation

• To use this equation, you assume that the
  equilibrium concentrations of the acid and its
  conjugate base are approximately equal to their
  initial concentrations
• The pH of the buffer falls within 3 to 11
• The initial concentrations of the acid and the
  conjugate base are large

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Practice Problem

• What is the pH of a buffer solution composed of
  0.50M formic acid (HCO2H) and 0.70M sodium
  formate (NaHCO2)?

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Practice Problem

• Suppose you dissolve 15.0 g of NaHCO3 and 18.0g
  of Na2CO3 in enough water to make 1.00L of
  solution. Use the H-H equation to calculate the
  pH of the solution. (Consider this buffer as a
  solution of the weak acid HCO3- with CO32- as its
  conjugate base.)

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Preparing Buffer Solutions

• pH control The solution should control the pH at
  the desired value. Choose an acid with Ka near to
  the intended value of H3O. The exact value of
  H3O can be achieved by adjusting the
  acid/conjugate base ratio.
• Buffer capacity The buffer should be able to
  control the pH after the addition of reasonable
  amounts of acid and base.

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Practice Problem Not Tested!

• Using an acetic acid/sodium acetate buffer
  solution, what ratio of acid to conjugate base
  will you need to maintain the pH at 5.00?

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Preparing Buffer Solutions

• It is the relative number of moles of acid and
  conjugate base that is important in determining
  the pH of a buffer solution. (the solution volume
  is the same for both components)
• Diluting a buffer solution will not change its pH!

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Practice Problem Not Tested!

• Calculate the pH of 0.500L of a buffer solution
  composed of 0.50M formic acid (HCO2H) and 0.70M
  sodium formate (NaHCO2) before and after adding
  10.0mL of 1.0M HCl.

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Homework

• After reading sections 17.1-17.2, you should be
  able to do the following
• P. 677 (7-25 odd)

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17.3 Acid-Base Titrations

• A titration is one of the most important ways of
  determining accurately the quantity of an acid, a
  base, or some other substance in a mixture.
• The pH at the equivalence point of a strong
  acid/strong base titration is 7.
• If the substance is a weak acid or base, then the
  pH at equivalence is not 7 depends on conjugate
  base or acid.

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Titration Strong Acid/Strong Base

• The equivalence point in any acid-base titration
  is identified as the midpoint in the vertical
  portion of the pH versus volume of titrant curve.
  (titrant refers to the substance being added,
  analyte is the substance being tested)
• The pH of the solution at the equivalence point
  in a strong acid/strong base reaction is always
  7. (at 25oC)

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Practice Problem

• What is the pH after 25.0 mL of 0.100M NaOH has
  been added to 50.0mL of 0.100M HCl? What is the
  pH after 50.50 mL of NaOH has been added?

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Titration Weak Acid/Strong Base

• Before titration begins, the pH is found from the
  weak acid Ka value and the acid concentration.
• At the equivalence point, the pH is controlled by
  the conjugate base.
• The pH at the halfway point of the titration is
  equal to the pKa of the weak acid.

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At the halfway point in the titration of a weak
acid with a strong base H3O Ka and pH pKa
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Practice Problem

• The titration of 0.100M acetic acid with 0.100M
  NaOH is described in the text. What is the pH of
  the solution when 35.0mL of the base has been
  added to 100.0mL of 0.100M acetic acid?

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Titration of Weak Polyprotic Acids

• Multiple equivalence points in the graph as each
  successive H ion is titrated.

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Titration Weak Base/Strong Acid

• Similar to Weak Acid/Strong Base, but the
  titration curve goes from high pH to low
  (opposite).

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Practice Problem

• Calculate the pH after 75.0mL of 0.100M HCl has
  been added to 100.0mL of 0.100M NH3. See Figure
  17.7 on page 651.

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pH Indicators

• An acid-base indicator is a weak acid or a weak
  base whose color is sensitive to pH. The acid
  form (HInd) has one color and the conjugate base
  (Ind-) has another.
• Although the indicator reacts with substances in
  solution, so little indicator is present that the
  analysis is not significantly affected.

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Homework

• After reading section 17.3, you should be able to
  do the following
• P. 677a-b (27-35 odd)

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17.4 Solubility of Salts

• The equilibrium constant that reflects the
  solubility of a compound is referred to as its
  solubility product constant, Ksp.
• The solubility of a salt is the amount present in
  some volume of saturated solution. The Ksp is an
  equilibrium constant.

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Practice Problem

• Write Ksp expressions for the following insoluble
  salts and look up numerical values for the
  constant in Appendix J.
• AgI
• BaF2
• Ag2CO3

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Practice Problem

• The barium ion concentration Ba2 in a
  saturated solution of barium fluoride is 3.6 x
  10-3M. Calculate the value of the Ksp for BaF2.
• BaF2(s) ?? Ba2(aq) 2F-(aq)

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Practice Problem

• Using the value of Ksp 5.5 x 10-5, calculate
  the solubility of Ca(OH)2 in moles per liter and
  grams per liter.

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Solubility and Ksp

• You can compare Ksp for two salts, but only if
  they have the same ion ratio!
• Larger Ksp means that the salt is more soluble.

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Practice Problem

• Using Ksp values, predict which salt in each pair
  is more soluble in water.
• AgCl or AgCN
• Mg(OH)2 or Ca(OH)2
• Ca(OH)2 or CaSO4

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Solubility and the Common Ion Effect

• The ionization of weak acids and bases is
  affected by the presence of an ion common to the
  equilibrium process and the effect of adding an
  ion to a saturated solution (with the same ion)
  will shift the equilibrium back to the formation
  of the compound (decrease the solubility).

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Practice Problem

• Calculate the solubility of BaSO4 (a) in pure
  water and (b) in the presence of 0.010M Ba(NO3)2,
  Ksp for BaSO4 is 1.1x10-10.

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Practice Problem

• Calculate the solubility of Zn(CN)2 at 25oC (a)
  in pure water and (b) in the presence of 0.10M
  Zn(NO3)2. Ksp for Zn(CN)2 is 8.0x10-12.

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Effect of Basic Anions on Salt Solubility

• Any salt containing an anion that is the
  conjugate base of a weak acid will dissolve in
  water to a greater extent than given by Ksp.
• due to reaction of the anion with water, which
  decreases its concentration and shifts the
  equilibrium
• Insoluble salts in which the anion is the
  conjugate base of a weak acid will dissolve in
  strong acids.

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Homework

• After reading section 17.4, you should be able to
  do the following
• P. 677b-c (41-61 odd)

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17.5 Precipitation Reactions

• Recall that the difference between Q (the
  reaction quotient) and K (the equilibrium
  constant) is that the concentrations used in the
  reaction quotient may or may not be those at
  equilibrium.
• If KspQ the solution is saturated.
• If KspgtQ the solution is not saturated.
• If KspltQ the solution is supersatured.

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Practice Problem

• Solid PbI2 (Ksp 9.8x10-9) is placed in a beaker
  of water. After a period of time, the lead (II)
  concentration is measured and found to be
  1.1x10-3M. Has the system yet reached
  equilibrium? That is, is the solution saturated?
  If not, will more PbI2 dissolve?

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Practice Problem

• If the concentration of strontium ion, Sr2, is
  2.5 x 10-4M, will precipitation of SrSO4 occur
  when enough of the soluble salt Na2SO4 is added
  to make the solution 2.5 x 10-4M in SO42-? Ksp
  for SrSO4 is 3.4x 10-7.

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Practice Problem

• What is the minimum concentration of I- that can
  cause precipitation of PbI2 from a 0.050M
  solution of Pb(NO3)2? Ksp for PbI2 is 9.8 x 10-9.
  What concentration of Pb2 ions remains in
  solution when the concentration of I- is 0.0015M?

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Practice Problem

• You have 100.0mL of 0.0010M silver nitrate. Will
  AgCl precipitate if you add 5.0mL of 0.025M HCl?

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17.6 Solubility and Complex Ions

• Metal ions exist in aqueous solution as complex
  ions.
• The equilibrium constant for the formation of a
  complex ion such as Ag(NH3)2 is called a
  formation constant, Kform.
• Knet KspKform

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Practice Problem

• Silver nitrate (0.0050 mol) is added to 1.00 L of
  1.00 M NH3. What is the concentration of Ag
  ions at equilibrium?
• Ag(aq) 2NH3(aq) ?? Ag(NH3)2(aq)
• Kf 1.1 x 107

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17.7 Solubility, Ion Separations, and Qualitative
Analysis

• It is often necessary to use Ksp information in
  order to identify ions in solution.
• You will need to find a reagent that will form a
  precipitate with one or more of the cations while
  leaving the others in solution.
• Also consider which salts are soluble in water
  and which arent.
• You can then add chemicals in order to identify
  the salts that have known colors.

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Homework

• After reading sections 17.5-17.7, you should be
  able to do the following
• P. 677c (65-73 odd)