Transcript Chapter 15 Common Ion Effect

Chapter 15
Application of Aqueous Equilibria
Common Ion Effect
The shift in equilibrium that occurs
because of the addition of an ion already
involved in the equilibrium reaction. This
effect makes a solution of NaF and HF less
acidic than a solution of HF alone.
AgCl(s)  Ag+(aq) + Cl(aq)
adding


NaCl(aq)shifts equilibrium position
A Buffered Solution
. . . resists a change in its pH when either
H+ or OH are added.
1.0 L of 0.50 M H3CCOOH
+ 0.50 M H3CCOONa
pH = 4.74
Adding 0.010 mol solid NaOH raises the pH
of the solution to 4.76, a very minor change.
Blood can absorb the acids and bases
produced in biological reactions without
changing its pH. The buffering system in the
blood involves HCO3- and H2CO3.
Key Points on Buffered
Solutions
1. They are weak acids or bases
containing a common ion.
2. After addition of strong acid or
base, deal with stoichiometry first,
then equilibrium.
Henderson-Hasselbalch Equation
Useful for calculating pH when
the [A]/[HA] ratios are known.
HA  H+ + A


[ H ][ A ]
[ HA]
Ka 
 [ H ]  Ka 
[ HA]
[A ]
[ HA]

  log[ H ]   log Ka  log 
[A ]
pH  pKa  log( A  / HA ) 
pKa  log( base / acid )
Buffered Solution Characteristics
• Buffers contain relatively large amounts of
weak acid and corresponding weak base.
• Added H+ reacts to completion with the weak
base. H+ + A-  HA
or H+ + B  BH+
• Added OH reacts to completion with the
weak acid. OH- + HA  A- + H2O
or OH- + BH+ B + H2O
• The pH is determined by the ratio of the
concentrations of the weak acid and weak
base.
Buffering Capacity
. . . represents the amount of H+ or OH
the buffer can absorb without a significant
change in pH.
A buffer with a large capacity contains
large
concentrations
of
buffering
components and can absorb a relatively
large amount of protons or hydroxide ions
without much pH change.
Titration (pH) Curve
• The pH changes during an acid-base
titration. The progress of an acid-base
titration is monitored by a plot of pH of the
solution being analyzed as a function of
the amount of titrant added. Such a plot is
called a pH curve or titration curve.
• Equivalence
(stoichiometric)
point:
Enough titrant has been added to react
exactly with the solution being analyzed.
Strong Acid-Strong Base Titration
• Strong acids and strong bases are completely
dissociated.
• The pH can be calculated at selected points
during the course of the titration from
concentration and volume used.
Net ionic equation for strong acid-strong base is
H+(aq) + OH(aq)  H2O(l)
mol solute mmol solute
Molarity 

L solution mL solution
Number of mmol = Volume (mL) x molarity
Figure 15.1 The pH Curve for the Titration of 50.0 mL of 0.200 M HNO3 with 0.100 M NaOH
Figure 15.2 The pH Curve for the Titration of 100.0 mL of 0.50 M NaOH with 1.0 M HCI
Weak Acid - Strong Base Titration
Step 1 - A stoichiometry problem reaction is assumed to run to
completion
then
determine
remaining species.
Step 2 - An equilibrium problem determine position of weak acid
equilibrium and calculate pH.
Figure 15.3 The pH Curve for the Titration of 50.0 mL of 0.100 M HC2H3O2 with 0.100 M NaOH
Figure 15.4 The pH Curves for the Titrations of 50.0-mL Samples of 0.10 M
Acids with Various Ka Values with 0.10 M NaOH
Acid-Base Indicator
. . . marks the end point of a titration by
changing color. The equivalence point
(defined by the stoichiometry) is not
necessarily the same as the end point
(where the indicator changes color).
Indicators exhibit one color when the
proton is attached to the molecule and a
different color when the proton is absent.
Phenolphthalein is colorless in its HIn
(acidic) form and pink in its In- (basic) form.
pH = pKa  1
Figure 15.5 The pH Curve for the Titration of 100.0 mL of 0.050 M NH3 with 0.10 M HCI
Figure 15.6 The Acid and Base Forms of the Indicator Phenolphthalein
Figure 15.8 The Useful pH Ranges for Several Common Indicators
Figure 15.9 The pH Curve for the Titration of 100.0 mL of 0.10 M HCI with 0.10 M NaOH
Figure 15.10 The pH Curve for the Titration of 50 mL of 0.1 M HC2H3O2 with 0.1 M NaOH
Solubility Product
For solids dissolving to form aqueous
solutions.
Bi2S3(s)  2Bi3+(aq) + 3S2(aq)
Ksp = solubility product constant
and
Ksp = [Bi3+]2[S2]3
Solubility Product
continued..
“Solubility” = s = concentration
of Bi2S3 that dissolves, which
equals1/2[Bi3+] and 1/3[S2].
Note: Ksp is constant (at a given
temperature)
s is variable (especially with a
common ion present)
Precipitation and Qualitative Analysis
• Precipitation is the reverse process of
dissolving solids in solutions. From the ion
product we can predict whether a
precipitate will form when two solutions are
mixed.
• For CaF2, the expression for ion product is
Q = [Ca2+]o[F-]o2
If Q is greater than Ksp  precipitation
occurs
If Q is less than Ksp  no precipitation
occurs
Qualitative Analysis
• Qualitative analysis of a mixture containing
all the common cations involves first
separating them into five major groups
based on solubilities. Each group is then
treated further to separate and identify the
individual ions.
Equilibria Involving Complex Ions
• Complex Ion: A charged species
consisting of a metal ion surrounded by
ligands [Lewis bases H2O, NH3, Cl-, CN].
• Coordination Number: Number of
ligands attached to a metal ion. [Most
common are 6 and 4 ie. Co(H2O)62+,
Ni(NH3)62+, CoCl42-, Cu(NH3)4 2+]
• Formation (Stability) Constants: The
equilibrium constants characterizing the
stepwise addition of ligands to metal
ions.