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.