Announcements

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Announcements

• Homework Assignments
• Volumetric Chloride Lab due Monday, Oct 20

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Chapter 11 Chemical Equilibrium

• Although we have already used concepts of
  equilibrium in some of the previous chapters, we
  want to take a more detailed look at this
  important topic.

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Chapter 11 Chemical Equilibrium

• Chemical equilibrium describes the distribution
  of the products and reactants
• in a reaction mixture. Frequently these
  relationships are more complicated than first
  meets the eye. For example, the equilibria
  (plural) for the solution of slightly soluble
• lead(II) iodide in water is shown in the next
  slide.

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The percent values give the approximate
contribution of each species to the whole.
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Chapter 11 Chemical Equilibrium

• Although the model of chemical equilibrium
  predicts that inert, that is, non-reactive
  salts should not affect the solubility, the
  actual effect is shown in the next slide, where
  the addition of KNO3 has caused an increase in
  the solubility of PbI2. In general the addition
  of such an inert salt causes an increase the
  solubility.

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Increased solubility of PbI2. in KNO3
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Chapter 11 Chemical Equilibrium

• Our understanding of the phenomena shown on the
  proceeding slide is that the inert salt increases
  the ionic atmosphere (environment), allowing each
  cation or anion to be surrounded by species of
  the opposite charge, but farther separated from
  the counter ion which caused its original
  chemical reaction.

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Ionic Environment around Cations or Anions.
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Chapter 11 Chemical Equilibrium

• The effect of the ionic environment within the
  solution is known as the ionic strength and may
  be represented as ? or I.
• ? I ½ (c1z12 c2z22 ) ½ ? cizi2
• The sum of terms includes all of the ions in
  solution. An example of this calculation is
  shown in Problem 1

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What is the ionic strength of a solution that is
0.0100 M in KNO3 and 0.0100 M Na2SO4?

• All of the ions present contribute to the ionic
• environment.
• I ? ½ ? cizi2
• ? ½ 0.01(1)2 0.02(1)2 0.01(-1)2
  0.01(-2)2
• K Na
  NO3? SO4 ?2
• ? ½ 0.08 0.04M

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Activity Coefficients

• The actual or effective concentration of an
  ionic species in solution is known as the
  activity your author uses the symbol A (more
  commonly used is simply a lower case a)
• I will use the later symbol, so that his
• Equation 11-2 is written as
• aC C ?C

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Activity Coefficients

• The exact equilibrium constant K is then
  expressed in terms of the activities of the
  species involved instead of the more commonly
  concentrations.
• For the reaction aA bB lt gt cC dD
• K (ac)c(aD)d / (aA)a(aB)b
• or K (C ?C)c(D ?D)d / (A ?C)a(B ?C)b

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Activity Coefficients

• The individual values for the activity
  coefficients ? of each of the species is a
  function of the ionic strength as shown by the
  extended Debye-Huckel equation
• __
  __
• log 10 ? -0.51z2? ? / 1 (??? (305))
• where ? is the size of the ion in pm
  (picometers). Examples of ? for the F- and I-
  ions are shown in the next slide.

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? Values for the F- and I- ions where ?
represents the hydrated radius of the ions.
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Figure 11-4 ? as a function of ? for different
ionic charges, values, z.
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The activity coefficient ? for the H ion as a
function of the ionic strength.
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The Refined definition of pH

• Because of the effect ionic strength may have on
  the activity of an ion in an environment of other
  ions, the definition of pH is refined as
• pH - log10 aH - log10 H ?H

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Charge balance in a solution that contains 0.0250
M KH2PO4 and 0.0300 M KOH. The sum of positive
charges sum of negative charges.
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