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COMMON IONEFFECT
COMMON IONCOMMON IONan ion common with one an ion common with one
in a system at in a system at equilibrium which equilibrium which
places a stress on the places a stress on the equilibriumequilibrium
Common IonCommon Ion
Uses of Common Uses of Common Ion EffectIon Effect
1. control pH of a 1. control pH of a weak acid or baseweak acid or base
2. control formation 2. control formation of a precipitateof a precipitate
BUFFERBUFFERExampleExample
Non-exampleNon-example
A solution which resists a A solution which resists a change in pH when an change in pH when an acid or base is addedacid or base is added
consists of a consists of a weakweak acid acid or base and a salt or base and a salt
containing a common containing a common ion of its conjugateion of its conjugate
How does LeChatelier’s Principle
explain the operation of a buffer?
Example of a buffer systemExample of a buffer system
CHCH33COOH + HOH COOH + HOH CHCH33COOCOO-- + H+ H33OO++
NaCHNaCH33COOCOO(aq) (aq) Na Na++ + + CHCH33COOCOO--
Characteristics of a Characteristics of a Good BufferGood Buffer
1. operates over a narrow 1. operates over a narrow pH range (< 1 pH unit)pH range (< 1 pH unit)
2. no reactions between 2. no reactions between buffers in a multiple buffer buffers in a multiple buffer systemsystem
3. range can be extended 3. range can be extended using more than one bufferusing more than one buffer
]HA[
]A[logpKpH a
Henderson-HasselbalchEquation
]HA[
]A[logpKpH a
Maximum buffering will occur when ratio is close to 1, or when
pH = pKa
1. What is the pH of a 0.20 M acetic acid solution?
Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the
preceding solution. What is the pH?
Add 5.0g sodium acetate (MM 82.05) to 500. mL of the 0.20 M acetic acid solution.
What is the pH?
Add 10.0 mL of 0.20 M NaOH to 50.0 mL of the
preceding solution. What is the pH?
2. Calculate the mass of ammonium chloride (MM 43.6) needed to buffer 250. mL of 2.0 M ammonia to a pH of 10.
TITRATIONCURVES
Titration CurveTitration CurveA graphical history of a A graphical history of a
titrationtitration
typically a plot of the pH typically a plot of the pH (dependent variable) (dependent variable) and volume titrant and volume titrant
(independent variable)(independent variable)
Uses of Titration CurvesUses of Titration Curves1. determine equivalence 1. determine equivalence
pointpoint
2. determine number of 2. determine number of ionization reactionsionization reactions
3. determine optimum 3. determine optimum buffer regionbuffer region
4. determine possible 4. determine possible indicatorsindicators
Shape of Titration CurveShape of Titration Curve
Strong acid - strong baseStrong acid - strong base
Weak acid - strong baseWeak acid - strong base
0
2
4
6
8
10
12
14
0 5 10 15 20 25 30 35 40 45 50Volume NaOH (mL)
pH
HAc-NaOH
Shape of Titration CurveShape of Titration Curve
0
2
4
6
8
10
12
14
0 5 10 15 20 25 30 35 40 45 50Volume NaOH (mL)
pH
HCl-NaOH
Shape of Titration CurveShape of Titration Curve
Equivalence PointEquivalence Point
1. Midpoint between 1. Midpoint between points of inflectionpoints of inflection
2. Plot of the slope of 2. Plot of the slope of each point of the curve each point of the curve against volume titrant against volume titrant ((pH/pH/V vs VV vs Vavgavg))
Number ofNumber of Ionization ReactionsIonization Reactions
CHCH33COOH - NaOHCOOH - NaOH
HH22CC22OO44 - NaOH - NaOH
Titration Curve of Acetic Acid
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00
0 5 10 15 20 25 30 35 40
Volume NaOH (mL)
pH
Titration Curve of Phosphoric Acid
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00
0 5 10 15 20 25 30 35 40 45 50
Volume NaOH (mL)
pH
Optimum Buffer RegionOptimum Buffer RegionArea where the Area where the concentration of concentration of
molecules and their molecules and their conjugate ions are conjugate ions are
relatively highrelatively high
IndicatorsIndicators
Need to choose for each Need to choose for each titration systemtitration system
Dependent on pH at Dependent on pH at equivalence pointequivalence point
ACID-BASEINDICATORS
Acid-base indicators are weak Bronsted-Lowry compounds that are different
colors in acid and base form.
Acid-base indicators are all large organic
molecules.
HIn <===> H+ + In- Color 1 Color 2
OHHO
C
OH
C
O
O-
PhenolphthaleinColorless acid form, HIn
O-O
C
C
O
O-
PhenolphthaleinPink base form, In-
The color change The color change occurs at a different pH occurs at a different pH for different indicators.for different indicators.
The pH at which the The pH at which the indicator changes color indicator changes color is dependent on the Kis dependent on the Kaa
of the indicator as a of the indicator as a weak acid.weak acid.
][
][log
][
]][[
HIn
InpKpH
or
HIn
InHK
a
a
HIn <===> H+ + In-
Experiments have shown that the minimum
amount of change of HIn <==> In-
that can be detected visually is
1
10 or
10
1
]HIn[
]In[
1pHpK or 1pKpH
then ,1
10 or
10
1
]HIn[
]In[ ifbut
]HIn[
]In[lgpKpH
aa
a
Thus, from the Henderson-Hasselbalch equation, one can select an appropriate indicator
for a titration based upon the Ka of the
indicator and the pH at the equivalence point.
What is the pH at the equivalence point of a titration of 25.0 mL each of 0.10 M HCl and 0.10 M NaOH?
What is the pH at the equivalence point of a
titration of 25.0 mL each of 0.10 M
CH3COOH and 0.10 M NaOH?
PhenolphthaleinKa = 1 x 10-9
pH of perceptiblecolor change?
SOLUBILITYEQUILIBRIA
Saturated SolutionSaturated Solution
Maximum amount of solute Maximum amount of solute dissolved in a specific dissolved in a specific volume of solvent at a volume of solvent at a specific temperaturespecific temperature
Saturated SolutionSaturated Solution
Equilibrium Equilibrium is is established between a established between a solid solute and ions solid solute and ions
from the solutefrom the solute
Super-Saturated SolutionSuper-Saturated Solution
More than the normal More than the normal maximum amount of maximum amount of
solute is dissolved in a solute is dissolved in a solution.solution.
QuestionQuestionat a constant temperature, at a constant temperature,
what is the difference in what is the difference in concentration of a concentration of a saturated solution:saturated solution:
(1 mL vs 1 ML solution)(1 mL vs 1 ML solution)
(1 mg vs 1 kg solid)(1 mg vs 1 kg solid)
The concentration of a saturated solution
remains the same, no matter how much solid is present, as long as
the temperature remains constant.
The “concentration” of a solid remains the
same at a constant temperature.
By convention, equations for the
formation of saturated solutions are written in
the format solid <===> solution
AgCl(s) <===> Ag+ + Cl-
AgCl(s) <===> Ag+ + Cl-
]AgCl[
]Cl][Ag[Keq
AgCl(s) <===> Ag+ + Cl-
]Cl][Ag[K
]Cl][Ag[]AgCl[K
constant, a is ]AgCl[ ,but
]AgCl[
]Cl][Ag[K
sp
eq
eq
3. What is the solubility of silver chloride in water at 25oC? (Ksp = 1.6 x 10-10)
4. What is the solubility of lead(II) bromide at 25oC? (Ksp = 4.6 x 10-6)
6. What mass of nickel is 6. What mass of nickel is dissolved in 100. mL of dissolved in 100. mL of
saturated nickel(II) saturated nickel(II) hydroxide? hydroxide?
(K(Kspsp = 1.6 x 10 = 1.6 x 10-16-16))
What is the pH of this What is the pH of this solution?solution?
Which is more soluble?
Ag2CO3 [Ksp = 8.5 x 10-13]or
CaCO3 [Ksp = 3.4 x 10-9]
SOLUBILITY----
ACIDITY----
PRECIPITATION
8. If 0.581 gram of 8. If 0.581 gram of magnesium hydroxide magnesium hydroxide (MM 58.3) is added to (MM 58.3) is added to
1.00L of water, will it all 1.00L of water, will it all dissolve? dissolve?
(K(Kspsp = 8.9 x 10 = 8.9 x 10-12-12))
Below what pH would the Below what pH would the solution be buffered so solution be buffered so that it does all dissolve?that it does all dissolve?
9. Calculate the concentration of NH4
+ from ammonium chloride required to prevent the precipitation of Ca(OH)2 in a liter of solution that contains 0.10 mole of ammonia and 0.10 mole of calcium ion.
10.10. If 50. mL of 0.012M If 50. mL of 0.012M barium chloride are mixed barium chloride are mixed with 25 mL of 1.0 x 10with 25 mL of 1.0 x 10-6-6M M
sulfuric acid, will a sulfuric acid, will a precipitate form?precipitate form?
HINT: use the concentration HINT: use the concentration quotient “Q” as we used it quotient “Q” as we used it beforebefore
11.You have a aqueous solution of Zn2+ and Pb2+ both 0.0010 M. Both form
insoluble sulfides. Approximately what pH will
allow maximum precipitation of one ion and leave the
other in solution?[Ksp ZnS = 2.5 x 10-22][Ksp PbS = 7 x 10-29]
SOLUBILITY----
COMMON IONS----
COMPLEX IONS
12. Calculate the molar solubility of silver thiocyanate, AgSCN, in pure water and in 0.010M NaSCN.
Complex IonComplex Ion
A charged species A charged species consisting of a metal ion consisting of a metal ion surrounded by ligandssurrounded by ligands
LIGANDLIGAND
An ion or molecule, An ion or molecule, acting as a Lewis acting as a Lewis
base, attached to the base, attached to the central metal ion central metal ion
using the d-orbitals of using the d-orbitals of the metalthe metal
Coordination NumberCoordination Number
The number of ligands The number of ligands attached to the central attached to the central
metal ion.metal ion.
2, 4, or 6 are most 2, 4, or 6 are most common CNcommon CN
Metal ions add ligands one Metal ions add ligands one step at a time.step at a time.
AgAg++ + NH + NH33 <==> Ag(NH <==> Ag(NH33))++
KKf1f1 = 2.1 x 10 = 2.1 x 1033
Ag(NHAg(NH33))++ + NH + NH33 <==> Ag(NH <==> Ag(NH33))22++
KKf2f2 = 8.2 x 10 = 8.2 x 1033
where Kwhere Kff = = formation constantformation constant
2f1f23
23f
233
KK]NH][Ag[
])NH(Ag[K
)NH(Ag2NH Ag
You need to familiarize yourself with “typical”
complex ions, Appendix K
Note that a formation constant reflects the
stability of the complex.
13. Calculate the equilibrium constant for
AgI(s) + 2NH3(aq) <===> [Ag(NH3)2]+(aq) + I-(aq)
14. Will 5.0 mL of 2.5 M NH3 dissolve 0.0001 mole AgCl?
15.A solution is prepared by adding 0.10 mole Ni(NH3)6Cl2 to 0.50 L of 3.0 M NH3. Calculate the [Ni(NH3)6
2+] and [Ni2+] in the solution.
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