ADDITIONAL ASPECTS of EQUILIBRIA
Common-Ion Effect
 The dissociation of a weak electrolyte is decreased by the
addition of a strong electrolyte that has an ion in common with
the weak electrolyte
 Consider : HC2H3O2  H+ + C2H3O2- If we add sodium acetate
, which increases the acetate ion and shifts the equilibrium to
the left.
 This causes a reduction in [H+] and a decrease in the %
ionization of the acetic acid
 This is the common ion effect
Buffered Solutions
 A buffered solution resists a change in pH when small
amounts of acid or base are added
 A buffer consists of a weak acid(HX) and its conjugate
base(X-)
 A buffer contains an acidic species and a basic species
Buffer Capacity and pH
 Buffer capacity is the amount of acid or base that can be added
before there is a significant change in pH
 Buffer capacity depends on the concentration of the buffer
components
 The pH of the buffer is related to the Ka and the relative
concentrations of acid and base
 The Henderson-Hasselbach equation shows the relationship
between conjugate acid-base concentrations, pH and the Ka
 PH = pKa + log(base/acid)= pKa + log([A-]/[HA] )
Addition of Strong Acids or Bases to buffers
 Break the calculation into 2 parts
o A stoichiometric calculation
o An equilibrium calculation
 Addition of a strong acid or base results in a neutralization
reaction: X- + H3O+ HX + H2O
HX + OH-  H2O + X By knowing how much acid or base was added, we know how
much HX or X- is formed. This is a stoichiometric calculation
 From the concentrations of HX and X- (note the volume change
in the solution) we can use the Henderson-Hasselbach equation
to calculate pH
 See example 17.5 page 648-649
Acid Base Titration
 The plot of pH versus added volume during a titration is called
a titration curve
 Consider adding a strong base to a strong acid
 We can divide the titration into 4 regions
1. Initial pH(before any base is added)
 pH 7
2. Between initial pH and the equivalence point
 When base is added before the equivalence point
the pH is determined by the excess strong acid
 pH 7
3. At the equivalence point
 The amount of base is stoichiometrically equal to
the amount of acid
 PH is determined by hydrolysis of the salt
 Therefore, pH 7
4. After the equivalence point
 PH is determined by the excess base
 pH 7
 How do we analyze a titration?
o To detect the equivalence point we choose an indicator which
changes color at a pH near the equivalence point
o The equivalence point is when base and acid are
stoichiometrically equivalent. End point of the titration is when
the indicator changes color. The difference is the titration error.
Weak Acid – Strong Base Titration
 Consider a titration with HC2H3O2 and NaOH
 Divide into 4 regions
1. Before base is added
 Solution is only weak acid
 PH is given by the equilibrium calculation
2. Between the initial pH and the equivalence point
 As base is added it consumes a stoichiometric quantity of
the weak acid
 Since there is an excess of weak acid, we have a mixture
of a weak acid and its conjugate base
 This is the composition of a buffer
 PH is given by the buffer calculation
o First calculate the amount of acetate generated
o Calculate amount of acid consumed
o Calculate pH using the Henderson-Hasselbach
equation
3. At the equivalence point all of the acid has been consumed
and the base has been consumed
 However, C2H3O2- has been generated
 The pH of the solution is given by the acetate solution
 Therefore the pH 7
 The pH of the equivalence point 7 for a weak acidstrong base titration
4. After the equivalence point the pH is given by the
concentration of the excess strong base
 The pH curve of a weak acid-strong base titration differs significantly
from a strong base-strong acid titration
o For a strong acid-strong base titration
 The pH begins at less than 7 and gradually increases as
base is added
 Near the equivalence point, the pH rises dramatically
 PH at equivalence equals 7
o For a weak acid-strong base titration
 The initial pH rise is steeper than for strong acid
 Then there is a leveling off due to buffer effects
 The inflection point is not as steep as the strong/strong
 ph at equivalence is greater than 7
o The curve for a weak base – strong acid also varies significantly
 The equivalence point is at pH= 5.28, so phenolphthalein
should not be used as an indicator. Methyl red is better
since its color changes in a range of pH4.2-6.0
Titrations of Polyprotic Acids
 Titrations of polyprotic acids have multiple equivalence points
Solubility Equilibria
 Consider a saturated solution of BaSO4 in contact with solid BaSO4
 BaSO4(s) Ba+2 + SO4- because BaSO4 is a pure solid the equilibrium
expression depends only on the concentration of the ions
 Ksp is called the solubility product constant or the solubility product
 Ksp BaSO4 = [Ba+2] [SO4-]
 In general, the solubility product is equal to the molar concentrations of
ions raised to powers corresponding to their stoichiometric coefficients
Solubility and Ksp
 Solubility is the amount of substance in grams that dissolves to
form a saturated solution. This is often expressed as grams per
liter of solution
 Molar solubility is the number of moles which dissolve to form
a liter of solution
 To convert solubility to Ksp
o Convert solubility into molar solubility
o Convert molar solubility into molar concentration of ions
at equilibrium
o Use the equilibrium concentrations in the Ksp expression
o Example 17.10 page 661
 To convert Ksp to solubility
o Write the Ksp expression
o Let x=the molar solubility of the salt
o Use the stoichiometry of the reaction to express the
concentration of each species in terms of x
o Substitute and solve for x
o Example 17.11 page 662
Factors That Effect Solubility
 Common ion effect
o Solubility is decreased when a common ion is added
 pH
o consider Mg(OH)2  Mg+2 + 2OH- if OH is removed by adding
a strong acid the equilibrium shifts to the right. Increasing the
solubility of the magnesium hydroxide
o pH effects are most significant if one or both of the ions are
somewhat basic or acidic
 Formation of Complex Ions
o The solubility of AgCl is very low. However if ammonia is
added solubility increases dramatically. Why?
o Ag++2 NH3  Ag(NH3)2+ this is a complex ion
o The equilibrium constant ,the constant of formation, is:
Kf = [Ag(NH3)2+] / [Ag+] [NH3+] = 1.7 x 107
o Obviously the formation of the complex is favored
o When ammonia is added to AgCl the Ag+ is effectively
removed from solution. This causes the equilibrium to shift to
the right, favoring the dissolution of the silver chloride
 Amphoterism
o Amphoteric oxides and hydroxides will dissolve in either acid
or base : Examples are oxides and hydroxides of Al +3,Cr+3 ,
Zn+2, Sn+2
Precipitation and Separation of Ions
 Consider BaSO4(s)  Ba+2 + SO4 At any instant in time, Q= [Ba+2] [SO4-]
o If Q Ksp precipitation occurs until Q=Ksp
o If Q = Ksp equilibrium exists
o If Q  Ksp solid dissolves until Q = Ksp
 Ions can be separated based on solubility
 Qualitative analysis uses solubility to separate and then identify ions
o Separation is based on 5 major groups
 Insoluble chlorides
 Acid insoluble sulfides
 Base insoluble sulfides and hydroxides
 Insoluble phosphates
 Alkali metals and ammonium ion