Chemistry 102(001) Fall 2012
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Transcript of Chemistry 102(001) Fall 2012
17-1CHEM 102, Fall 2010, LA TECH
CTH 328 10:00-11:15 am
Instructor: Dr. Upali Siriwardane
e-mail: [email protected]
Office: CTH 311 Phone 257-4941
Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu, Th, F 8:00 - 10:00am..
Exams: 10:00-11:15 am, CTH 328.
September 27, 2012 (Test 1): Chapter 13
October 18, 2012 (Test 2): Chapter 14 &15
November 13, 2012 (Test 3): Chapter 16 &18
Optional Comprehensive Final Exam: November 15, 2012 :
Chapters 13, 14, 15, 16, 17, and 18
Chemistry 102(001) Fall 2012
17-2CHEM 102, Fall 2010, LA TECH
Chapter 16.Aditional Aqueous Equilibria17.1 Buffer Solutions17.2 Acid-Base Titrations17.3 Acid Rain17.4 Solubility Equilibria and the Solubility Product Constant,
Ksp
17.5 Factors Affecting Solubility / Precipitation: Will It Occur?
17-3CHEM 102, Fall 2010, LA TECH
Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-Lowry acid-
base reaction.
CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)
NH4+
(aq) + H2O(l) NH3 (aq) + H3O+
(aq)
This type of reaction is given a special name.
Hydrolysis
The reaction of an anion with water to produce the conjugate acid and OH-.
The reaction of a cation with water to produce the conjugate base and H3O+
.
Hydrolysis
17-4CHEM 102, Fall 2010, LA TECH
Acid-Base Properties of Typical Ions
17-5CHEM 102, Fall 2010, LA TECH
What salt solutions would be acidic, basic and neutral?1) strong acid + strong base =
neutral 2) weak acid + strong base =
basic 3) strong acid + weak base =
acidic 4) weak acid + weak base =
neutral, basic or an acidic solution
depending on the relative strengths of the acid and the base.
17-6CHEM 102, Fall 2010, LA TECH
What pH? Neutral, basic or acidic?
•a)NaCl • neutral•b) NaC2H3O2
• basic•c) NaHSO4 • acidic•d) NH4Cl• acidic
17-7CHEM 102, Fall 2010, LA TECH
1) If the following substance is dissolved in pure water, will the solution be acidic, neutral, or basic?
a) Solid sodium carbonate-(Na2CO3):
b) Sodium chloride- (NaCl):
c) Sodium acetate- (NaC2H3O2):
d) Ammonium sulfate-((NH4)2SO4):
17-8CHEM 102, Fall 2010, LA TECH
How do you calculate pH of a salt solution?Find out the pH, acidic or basic?If acidic it should be a salt of weak baseIf basic it should be a salt of weak acidif acidic calculate Ka from Ka= Kw/Kb
if basic calculate Kb from Kb= Kw/Ka
Do a calculation similar to pH of a weak acid or base
17-9CHEM 102, Fall 2010, LA TECH
What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5)Find out the pH, acidic
if acidic calculate Ka from Ka= Kw/Kb
Ka= Kw/Kb = 1 x 10-14 /1.8 x 10-5)
Ka= 5.56. X 10-10
Do a calculation similar to pH of a weak acid
17-10CHEM 102, Fall 2010, LA TECH
Continued NH4
+ + H2O H 3+O + NH3
[NH4+] [H3
+O ] [NH3 ]Ini. Con. 0.5 M 0.0 M 0.00 MChange -x x xEq. Con. 0.5 - x x x [H 3+O ] [NH3 ] Ka(NH4
+) = -------------------- = [NH 4+] x2
---------------- ; appro.:0.5 - x . 0.5 (0.5 - x)
17-11CHEM 102, Fall 2010, LA TECH
x2 Ka(NH4
+) = ----------- = 5.56 x 10 -10
0. 5 x2
= 5.56 x 10 -10 x 0.5 = 2.78 x 10 -10
x= 2.78 x 10 -10 = 1.66 x 10-5
[H+ ] = x = 1.66 x 10-5 MpH = -log [H+ ] = - log 1.66 x 10-5
pH = 4.77pH of 0.5 M NH4Cl solution is 4.77
(acidic)
Continued
17-12CHEM 102, Fall 2010, LA TECH
2) What is the pH of a 0.05 M aqueous NH4Cl solution? (Kb (NH3) = 1.8 x 10-5)
a) equilibrium reaction for the hydrolysis of salt:b) Ka for the conjugate acid NH4
+:c) ICE set-up:I:___________________________________________C:__________________________________________E___________________________________________d) Calculation of x, [H3O]+:e) pH of the solution:
17-13CHEM 102, Fall 2010, LA TECH
3) What is the pH of a 0.05 M aqueous NaC2H3O2 solution? (Ka (HC2H3O2) = 1.8 x 10-5)
a) equilibrium reaction for the hydrolysis of salt:b) Kb for the conjugate base C2H3O2
-:c) ICE set-up:I:___________________________________________C:__________________________________________E___________________________________________d) Calculation of x, [OH]-:e) pOH and pH of the solution:
17-14CHEM 102, Fall 2010, LA TECH
Acid-Base Chemistryof Some Antacids
17-15CHEM 102, Fall 2010, LA TECH
4) A 50.00-mL sample of 0.100 M KOH is being titrated with 0.100 M HNO3. Calculate the pH of the solution after 52.00 mL of HNO3 is added.
a) acid base reaction:
b) moles of KOH: c) moles of HNO3:
d) [H3O]+:
e) pH of the solution:
17-16CHEM 102, Fall 2010, LA TECH
Reaction of a basic anion with water is an ordinary Brønsted-Lowry acid-base reaction.
CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)
This type of reaction is given a special name.
Hydrolysis
The reaction of an anion with water to produce the conjugate acid and OH-.
The reaction of a cation with water to produce the conjugate base and H3O+.
Hydrolysis
17-17CHEM 102, Fall 2010, LA TECH
Common Ion Effect
Weak acid and salt solutionsE.g. HC2H3O2 and NaC2H3O2
Weak base and salt solutionsE.g. NH3 and NH4Cl. H2O + C2H3O2
- OH- + HC2H3O2 (common ion)
H2O + NH4 + H3+O + NH3
(common ion)
17-18CHEM 102, Fall 2010, LA TECH
Solutions that resist pH change when small amounts of acid or base are added.
Two types
Mixture of weak acid and its salt
Mixture of weak base and its salt
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
Add OH- Add H3O+
shift to right shift to left
Based on the common ion effect.
Buffers
17-19CHEM 102, Fall 2010, LA TECH
The pH of a buffer does not depend on the absolute amount of the conjugate acid-base
pair. It is based on the ratio of the two.
Henderson-Hasselbalch equation.
Easily derived from the Ka or Kb expression.
Starting with an acid
pH = pKa + log
Starting with a base
pH = 14 - ( pKb + log )[HA]
[A-]
[A-]
[HA]
Buffers
17-20CHEM 102, Fall 2010, LA TECH
Henderson-Hasselbalch Equation
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
[H3O+] [A-] Ka = ---------------- [HA]
[H3O+] = Ka ([HA]/[A-])
pH = pKa + log([A-]/[HA])
when the [A-] = [HA]pH = pKa
17-21CHEM 102, Fall 2010, LA TECH
Calcualtion of pH of BuffersHenderson Hesselbach Equation
[ACID]pH = pKa - log --------- [BASE] [BASE]pH = pKa + log --------- [ACID]
17-22CHEM 102, Fall 2010, LA TECH
Control of blood pH
Oxygen is transported primarily by hemoglobin in the red blood cells.
CO2 is transported both in plasma and the red blood cells.
CO2 (aq) + H2O H2CO3 (aq)
H+
(aq) + HCO3-(aq)
The bicarbonate
buffer is essential
for controlling
blood pH
Buffers and blood
17-23CHEM 102, Fall 2010, LA TECH
Buffer Capacity
Refers to the ability of the buffer to retard changes in pH when small amounts of acid or base are added
The ratio of [A-]/[HA] determines the pH of the buffer whereas the magnitude of [A-] and [HA] determine the buffer capacity
17-24CHEM 102, Fall 2010, LA TECH
Adding an Acid or Baseto a Buffer
17-25CHEM 102, Fall 2010, LA TECH
Buffer Systems
17-26CHEM 102, Fall 2010, LA TECH
5) For the (buffer effect) of HC2H3O2/NaC2H3O2
a) Acid dissociation reaction:
b) Salt hydrolysis reaction:
c) Common ions in both equilibria:
d) Which way salt hydrolysis equilibrium move adding
H3O+:
e) Which way salt hydrolysis equilibrium move adding
OH-:
17-27CHEM 102, Fall 2010, LA TECH
6) Describe the (buffer effect) of NH3/NH4Cl
a) Buffer type: (weak acid or base)/soluble salt):
b) Base dissociation reaction:
c) Salt hydrolysis reaction:
d) common ions in both equilibria:
e) Which way salt hydrolysis equilibrium move adding
H3O+:
f) Which way salt hydrolysis equilibrium move adding
OH-:
17-28CHEM 102, Fall 2010, LA TECH
7) What is the pH of a solution that is 0.2 M in acetic acid (Ka = 1.8 x 10-5) and 0.2 M in sodium acetate?
a) Is it a acid, base, salt or buffer solution?
b) Henderson-Hesselbalch equation:
c) pKa: d)
e) pH of the solution:
17-29CHEM 102, Fall 2010, LA TECH
Titrations ofAcids and Bases
TitrationAnalyteTitrant
analyte + titrant => products
17-30CHEM 102, Fall 2010, LA TECH
Acid-base indicators are highly colored weak acids or bases.
HIn In- + H+
color 1 color 2
They may have more than one color transition.
Example. Thymol blue
Red - Yellow - Blue
One of the forms may be colorless - phenolphthalein (colorless to pink)
Indicators
17-31CHEM 102, Fall 2010, LA TECH
Acid-Base Indicator
HIn + H2O H3O+ + In-
acid base color color
[H3O+][In-]Ka =
[HIn]They may have more than one color transition.
Example. Thymol blue Red - Yellow - Blue
Weak acid that changes color with changes in pH
17-32CHEM 102, Fall 2010, LA TECH
What is an Indicator?Indicator is an weak acid with different Ka, colors to
the acid and its conjugate base. E.g. phenolphthalein
Hin H+ + In-
colorless pinkAcidic colorlessBasic pink
17-33CHEM 102, Fall 2010, LA TECH
Selection of an indicator for a titration
a) strong acid/strong base b) weak acid/strong base c) strong acid/weak base d) weak acid/weak base
Calculate the pH of the solution at he equivalence point or end point
17-34CHEM 102, Fall 2010, LA TECH
pH and Color of Indicators
17-35CHEM 102, Fall 2010, LA TECH
Red Cabbage as IndicatorC O H
O
N N N
CH3
CH3
(aq)
C O-
O
N N N
CH3
CH3
(aq) + H3O+ (aq)
yellow
red
+ H2O (l)
17-36CHEM 102, Fall 2010, LA TECH
Acid-base indicators are weak acids that undergo a color change at a
known pH.
phenolphthalein
pH
Indicator examples
17-37CHEM 102, Fall 2010, LA TECH
8) If 50 ml of a 0.01 M HCl solution is titrated with a 0.01 M NaOH solution, what will be the concentration of salt (NaCl) the pH at the endpoint?
a) NaCl solution acidic, basic or neutral?
b) Concentration of [NaCl]:
c) pH of the solution?
d) Suitable indicator for the titration:
17-38CHEM 102, Fall 2010, LA TECH
Titration Apparatus
Burette delivering base to a flask containing an acid.
The pink color in the flask is due to the
phenolphthalein indicator.
17-39CHEM 102, Fall 2010, LA TECH
Endpoint vs. Equivalence Point
Endpointpoint where there is a physical change, such
as color change, with the indicator
Equivalence Point# moles titrant = # moles analyte
#molestitrant=(V M)titrant
#molesanalyte=(V M)analyte
17-40CHEM 102, Fall 2010, LA TECH
9) If 50 mL of a 0.01 M HCl solution is titrated with a 0.01 M NH3 (Kb = 1.8 x 10-5) solution, what will be
a) The initial pH (0.01 M NH3):
b) Concentration of NH4Cl at the endpoint:
c) pH at the endpoint:
d) Suitable indicator for the titration:
17-41CHEM 102, Fall 2010, LA TECH
10) If 50 ml of a 0.01 M HC2H3O2 solution is titrated with a 0.01 M NaOH solution, what will be the
a) Molarity of NaC2H3O2 at the endpoint:
b) The pH at the endpoint:
c) What indicator would be most suitable for this titration:
17-42CHEM 102, Fall 2010, LA TECH
Polyprotic Acids
17-43CHEM 102, Fall 2010, LA TECH
Organic or Carboxylic Acids
H C
H
H
C
H
H
C
H
H
C
O
O H
nonacidic hydrogens
butanoic acid
acidic hydrogen
CH 3 C
O
acetic acid
OH
electron-attractingoxygen atom
acidic hydrogenCH 3 C
O
OH
-
CH 3 C
O
O-
acetate ion
17-44CHEM 102, Fall 2010, LA TECH
FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid).
Acid Ka pKa
HCOOH (formic acid) 1.78 X 10-43 0.75CH3COOH (acetic acid) 1.74 X 10-54 0.76CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86
Organic or Carboxylic Acids
17-45CHEM 102, Fall 2010, LA TECH
Acid-Base in the Kitchenvinegar - acetic acidlemon juice (citrus juice) - citric acidbaking soda - NaHCO3
milk - lactic acidbaking powder - H2PO4
- & HCO3-
17-46CHEM 102, Fall 2010, LA TECH
Household Cleaners
CH 3CH2CH 2CH2CH 2CH2CH 2CH 2CH 2CH 2CH2CH 2CH2CH 2 SO3-Na+
Oil-soluble part(hydrophobic)
Water-soluble part(hydrophilic)
A Typical Synthetic Detergent Molecule
CH 3(CH 2)4COO(CH 2)2O( CH2CH 2O) 2CH 2CH 2OH
esterlink
(hydro-philic)
etherlink
etherlink
(hydrophilic)
hydrocarbonchain
(hydro-phobic)
alcohol group(hydrophilic)
A nonionic detergent
17-47CHEM 102, Fall 2010, LA TECH
Dishwashing Detergent
17-48CHEM 102, Fall 2010, LA TECH
Acid-Base Indicator Behavioracid color shows when
[In-] 1
£ [HIn] 10
[H3O+
][In-] 1
= [H3O+
] = Ka [HIn] 10
base color shows when
[In-] 1 ³ [HIn] 10
[H3O+][In-] = 10 [H3O+] = Ka [HIn]
17-49CHEM 102, Fall 2010, LA TECH
Indicator pH Range
acid color shows when
pH + 1 = pKa
and base color shows when
pH - 1 = pKa
Color change range is
pKa = pH 1 or pH = pKa 1
17-50CHEM 102, Fall 2010, LA TECH
Titration curves
Acid-base titration curveA plot of the pH against the amount of acid or base
added during a titration.Plots of this type are useful for visualizing a titration.
It also can be used to show where an indicator undergoes its color change.
17-51CHEM 102, Fall 2010, LA TECH
Indicator and Titration Curve 0.1000 M HCl vs 0.1000 M NaOH
17-52CHEM 102, Fall 2010, LA TECH
EXAMPLE: Derive the titration curve for the titration of 35.00 mL of 0.1000 M HCl with 0.00, 15.00, 35.00, and 50.00 mL of 0.1000 M NaOH.at 0.00 mL of NaOH added, initial point
[H3O+
] = [HCl] = 0.1000 M
pH = 1.0000
17-53CHEM 102, Fall 2010, LA TECH
EXAMPLE: Derive the titration curve for the titration of 35.00 mL of 0.1000 M HCl with 0.00, 15.00, 35.00, and 50.00 mL of 0.1000 M NaOH.
at 15.00 mL of NaOH added
Va Ma > Vb Mb thus
(Va Ma) - (Vb Mb)
[H3O+] = (Va + Vb)
((35.00mL) (0.1000M)) - ((15.00mL) (0.1000M))
= (35.00 + 15.00)mL
= 4.000 10-2
M pH = 1.3979
17-54CHEM 102, Fall 2010, LA TECH
EXAMPLE: Derive the titration curve for the titration of 35.00 mL of 0.1000 M HCl with 0.00, 15.00, 35.00, and 50.00 mL of 0.1000 M NaOH.
at 35.00 mL of NaOH added
Va Ma = Vb Mb , equivalence point
at equivalence point of a
strong acid - strong base titration
pH º 7.0000
17-55CHEM 102, Fall 2010, LA TECH
EXAMPLE: Derive the titration curve for the titration of 35.00 mL of 0.1000 M HCl with 0.00, 15.00, 35.00, and 50.00 mL of 0.1000 M NaOH.
at 50.00 mL of NaOH added
Vb Mb > Va Ma , post equvalence point
(Vb Mb) - (Va Ma)
[OH-] = (Va + Vb)
((50.00mL) (0.1000M)) - ((35.00mL) (0.1000M))
=
(35.00 + 50.00)mL
= 1.765 10-2
M pOH = 1.7533
pH = 14.00 - 1.7533 = 12.25
17-56CHEM 102, Fall 2010, LA TECH
0
5
10
15
0 10 20 30 40 50
Volume of Base Added
pH equivalence point x
Titration of Strong Acid with Strong Base
17-57CHEM 102, Fall 2010, LA TECH
Titration of Weak Acid with Strong Base
17-58CHEM 102, Fall 2010, LA TECH
Effect of Acid Strength on Titration Curve
17-59CHEM 102, Fall 2010, LA TECH
Titration of Weak Base with Strong Acid
17-60CHEM 102, Fall 2010, LA TECH
Titration of Diprotic Weak Acid with Strong Base
17-61CHEM 102, Fall 2010, LA TECH
pH range of Indicatorslitmus (5.0-8.0)bromothymole blue (6.0-7.6) methyl red (4.8-6.0)thymol blue (8.0-9.6) phenolphthalein (8.2-10.0) thymolphthalein (9.4-10.6)
17-62CHEM 102, Fall 2010, LA TECH
Acid Rainacid rain is defined as rain with a pH < 5.6pH = 5.6 for rain in equilibrium with
atmospheric carbon dioxide
17-63CHEM 102, Fall 2010, LA TECH
Sulfuric Acid from Sulfur burning
SO2
S + O2 => SO2
SO3
2 SO2 + O2 => 2 SO3
Sulfuric AcidSO3 + H2O => H2SO4
17-64CHEM 102, Fall 2010, LA TECH
Nitric Acid2 NO2(g) + H2O(l) => HNO3(aq) + HNO2(aq)
17-65CHEM 102, Fall 2010, LA TECH
How Acid Precipitation Forms
17-66CHEM 102, Fall 2010, LA TECH
Acid Precipitation in U.S.
17-67CHEM 102, Fall 2010, LA TECH
Solubility Productsolubility-product - the product of the solubilitiessolubility-product constant => Ksp
constant that is equal to the solubilities of the ions produced when a substance dissolves
17-68CHEM 102, Fall 2010, LA TECH
Solubility Product Constant
In General:AxBy xA+y + yB-x
[A+y
]x [B-x
]yK =
[AxBy]
[AxBy] K = Ksp = [A+y
]x
[B-x
]y
Ksp = [A+y
]x [B
-x]y
For silver sulfate
Ag2SO4 2 Ag+
+ SO4-2
Ksp = [Ag+
]2[SO4
-2]
17-69CHEM 102, Fall 2010, LA TECH
Solubility Product Constant Values
17-70CHEM 102, Fall 2010, LA TECH
Dissolving Slightly Soluble Salts Using AcidsInsoluble salts containing anions of
Bronsted-Lowry bases can be dissolved in solutions of low pH
17-71CHEM 102, Fall 2010, LA TECH
Calcium CarbonateDissolved in Acid
Limestone Dissolving in Ground WaterCaCO3(S) + H2O + CO2 => Ca+2
(aq) + 2 HCO3-(aq)
Stalactite and Stalagmite FormationCa+2
(aq) + 2 HCO3-(aq) => CaCO3(S) + H2O + CO2
17-72CHEM 102, Fall 2010, LA TECH
The Common Ion Effectcommon ionsecond source which is completely dissociatedIn the presence of a second source of the ion,
there will be less dissolved than in its absencecommon ion effecta salt will be less soluble if one of its constitutent
ions is already present in the solution
17-73CHEM 102, Fall 2010, LA TECH
EXAMPLE: What is the molar solubility of AgCl in pure water and in 1.0 M NaCl?
AgCl Ag+ + Cl-
Ksp = [Ag+][Cl-] = 1.82 10-10M2
let x = molar solubility = [Ag+] = [Cl-](x)(x) = Ksp = [Ag+][Cl-] = 1.82 10-10M2
x = 1.35 10-5M
17-74CHEM 102, Fall 2010, LA TECH
EXAMPLE: What is the molar solubility of AgCl in pure water and in 1.0 M NaCl?
AgCl Ag+ + Cl-
Ksp = [Ag+][Cl-] = 1.82 10-10M2
let x = molar solubility = [Ag+][Cl-] = 1.0 M
Ksp = [Ag+][Cl-] = (x)(1.0M) = 1.82 10-10M2
x = 1.82 10-10M
17-75CHEM 102, Fall 2010, LA TECH
Formation of Complexesligand - Lewis basecomplexes - product of Lewis acid-base reaction
Ag+(aq) + 2 NH3(aq) [Ag(NH3)2(aq)]+
Ag+(aq) + Cl-
(aq) AgCl(s)
AgCl(s) + 2 NH3(aq) [Ag(NH3)2(aq)]+ + Cl-(aq)
17-76CHEM 102, Fall 2010, LA TECH
Sodium Thiosulfate Dissolves Silver Bromide
17-77CHEM 102, Fall 2010, LA TECH
11) For a saturated solution of a) AgCl in water:
i) Solubility equilibrium reaction:
ii) Ksp expression:
17-78CHEM 102, Fall 2010, LA TECH
11) For a saturated solution of b) CaF2 in water:i) Solubility equilibrium reaction:
ii) Ksp expression:
c) Fe2S3 in wateri) Solubility equilibrium reaction:
ii) Ksp expression:
17-79CHEM 102, Fall 2010, LA TECH
12) Which of following has the highest molar solubility (mole/L)?
Salt Ksp
a) CaCO3 5 × 10-9
b) PbCO 3 1.4 × 10-13
c) Li2CO3 2 × 10-3
d) NiCO3 1.2 × 10-7
17-80CHEM 102, Fall 2010, LA TECH
Formation Constants
17-81CHEM 102, Fall 2010, LA TECH
Amphoterism
17-82CHEM 102, Fall 2010, LA TECH
Reactant Quotient, Q
Reactant Quotient, Qion product of the initial concentrationsame form as solubility product constant
Q < Ksp - no precipitate forms an unsaturate solutionQ > Ksp - precipitate may form to restore
condition of saturated solutionQ = Ksp - no precipitate forms, saturated solution
17-83CHEM 102, Fall 2010, LA TECH
Will Precipitation Occur?
17-84CHEM 102, Fall 2010, LA TECH
13) For Li2CO3, Ksp is 2 × 10-3 M3. What is the concentration of Li+ in a saturated solution of Li2CO3?
17-85CHEM 102, Fall 2010, LA TECH
14) Chemical analysis gave [Pb2+] = 0.012 M, and [Br-] = 0.024 M in a solution. From a table, you find Ksp for PbBr2 has a value of 4 x 10-5 M3.
a) Solubility equilibrium reaction:
b) Qsp:
c) Ksp:d) Qsp < Ksp, Qsp = Ksp or Qsp > Ksp?e) Is the solution saturated, oversaturated or
unsaturated?
17-86CHEM 102, Fall 2010, LA TECH
Kidney StonesKidney stones are normally composed of:
calcium oxalatecalcium phosphate
magnesium ammonium phsphate
17-87CHEM 102, Fall 2010, LA TECH
Calcium OxalatePrecipitate formed from calcium ions from food rich
in calcium, dairy products, and oxalate ions from fruits and vegetables
Ca+2 + C2O4-2 CaC2O4
17-88CHEM 102, Fall 2010, LA TECH
PRECIPITATION REACTIONS
17-89CHEM 102, Fall 2010, LA TECH
Analysis of Silver Group
These salts formed are insoluble, they do dissolve to some SLIGHT extent.
AgCl(s) Ag+(aq) + Cl-(aq)When equilibrium has been
established, no more AgCl dissolves and the solution is SATURATED.
Ag+ Pb2+ Hg22+
AgCl PbCl2 Hg2Cl2
17-90CHEM 102, Fall 2010, LA TECH
Analysis of Silver Group
AgCl(s) Ag+(aq) + Cl-(aq)When solution is SATURATED, expt.
shows that [Ag+] = 1.67 x 10-5 M.(SOLUBILITY) of AgCl.What is [Cl-]?
Ag+ Pb2+ Hg22+
AgCl PbCl2 Hg2Cl2
17-91CHEM 102, Fall 2010, LA TECH
Analysis of Silver Group
AgCl(s) Ag+(aq) + Cl-(aq)Saturated solution has
[Ag+] = [Cl-] = 1.67 x 10-5 MUse this to calculate Ksp
Ksp = [Ag+] [Cl-] = (1.67 x 10-5)(1.67 x 10-5) = 2.79 x 10-10
Ag+ Pb2+ Hg22+
AgCl PbCl2 Hg2Cl2
17-92CHEM 102, Fall 2010, LA TECH
Because this is the product of “solubilities”, we call it
Ksp = solubility product constant See Table in the Text
Ksp = solubility product
17-93CHEM 102, Fall 2010, LA TECH
Lead(II) ChloridePbCl2(s) Pb2+(aq) + 2 Cl-(aq) Ksp = 1.9 x 10-5
17-94CHEM 102, Fall 2010, LA TECH
Consider PbI2 dissolving in waterPbI2(s) Pb2+(aq) + 2 I-(aq)Calculate Ksp if solubility = 0.00130 MSolution1. Solubility = [Pb2+]
= 1.30 x 10-3 M [I-] = 2 x [Pb2+] = 2.60 x 10-3 M2. Ksp = [Pb2+] [I-]2
= [Pb2+] {2 • [Pb2+]}2
= 4 [Pb2+]3 = 4 (solubility)3
Ksp = 4 (1.30 x 10-3)3 = 8.8 x 10-9
Solubility of Lead(II) Iodide
17-95CHEM 102, Fall 2010, LA TECH
Precipitating an Insoluble SaltHg2Cl2(s) Hg2
2+(aq) + 2 Cl-(aq)
Ksp = 1.1 x 10-18 = [Hg22+] [Cl-]2
If [Hg22+] = 0.010 M, what [Cl-] is req’d to just begin
the precipitation of Hg2Cl2?
(maximum [Cl-] in 0.010 M Hg22+ without forming
Hg2Cl2?)
17-96CHEM 102, Fall 2010, LA TECH
Precipitating an Insoluble SaltHg2Cl2(s) Hg2
2+(aq) + 2 Cl-(aq)
Ksp = 1.1 x 10-18 = [Hg22+] [Cl-]2
Recognize that
Ksp = product of maximum ion concs.
Precip. begins when product of ion Concs. EXCEEDS the Ksp.
17-97CHEM 102, Fall 2010, LA TECH
Precipitating an Insoluble SaltHg2Cl2(s) Hg2
2+(aq) + 2 Cl-(aq)
Ksp = 1.1 x 10-18 = [Hg22+] [Cl-]2
Solution
[Cl-] that can exist when [Hg22+] = 0.010 M,
If this conc. of Cl- is just exceeded, Hg2Cl2 begins to precipitate.
[Cl ] = Ksp0.010
= 1.1 x 10-8 M
17-98CHEM 102, Fall 2010, LA TECH
Precipitating an Insoluble SaltHg2Cl2(s) Hg2
2+(aq) + 2 Cl-(aq)Ksp = 1.1 x 10-18
Now raise [Cl-] to 1.0 M. What is the value of [Hg2
2+] at this point?Solution[Hg2
2+] = Ksp / [Cl-]2
= Ksp / (1.0)2 = 1.1 x 10-18 MThe concentration of Hg2
2+ has been reduced by 1016 !
17-99CHEM 102, Fall 2010, LA TECH
Separating Metal Ions Cu2+, Ag+, Pb2+
Ksp ValuesAgCl 1.8 x 10-10
PbCl2 1.7 x 10-5
PbCrO4 1.8 x 10-14
Cu2+
Ag+
Pb2+
Cl-
Insoluble
PbCl2 AgClSoluble
CuCl2
Heat
Insoluble
AgCl
Soluble
PbCl2
CrO4-2
Insoluble
PbCrO4
17-100CHEM 102, Fall 2010, LA TECH
Separating Salts by Differences in Ksp
A solution contains 0.020 M Ag+ and Pb2+. Add CrO42-
to precipitate Ag2CrO4 (red) and PbCrO4 (yellow). Which precipitates first?
Ksp for Ag2CrO4 = 9.0 x 10-12
Ksp for PbCrO4 = 1.8 x 10-14
SolutionThe substance whose Ksp is first exceeded
precipitates first. The ion requiring the smaller amount of CrO4
2- ppts. first.
17-101CHEM 102, Fall 2010, LA TECH
Separating Salts by Differences in Ksp
[CrO42-] to ppt. PbCrO4 = Ksp / [Pb2+]
= 1.8 x 10-14 / 0.020 = 9.0 x 10-13 M[CrO4
2-] to ppt. Ag2CrO4 = Ksp / [Ag+]2 = 9.0 x 10-12 / (0.020)2 = 2.3 x 10-8 M PbCrO4 precipitates first.
Solution
Calculate [CrO42-
] required by each ion.
17-102CHEM 102, Fall 2010, LA TECH
How much Pb2+ remains in solution when Ag+ begins to precipitate?
SolutionWe know that [CrO4
2-] = 2.3 x 10-8 M to begin to ppt. Ag2CrO4. What is the Pb2+ conc. at this point?[Pb2+] = Ksp / [CrO4
2-] = 1.8 x 10-14 / 2.3 x 10-8 M = 7.8 x 10-7 MLead ion has dropped from 0.020 M to < 10-6 M
Separating Salts by Differences in Ksp
17-103CHEM 102, Fall 2010, LA TECH
Common Ion EffectAdding an Ion “Common” to an Equilibrium
PbCl2(s) Pb2+
(aq) + 2Cl-(aq)
NaCl Na+
(aq) + Cl- (aq)
17-104CHEM 102, Fall 2010, LA TECH
Calculate the solubility of BaSO4 in (a) pure water and (b) in 0.010 M Ba(NO3)2.
Ksp for BaSO4 = 1.1 x 10-10
BaSO4(s) Ba2+(aq) + SO42-(aq)
Solutiona) Solubility in pure water = [Ba2+] = [SO4
2-] = xKsp = [Ba2+] [SO4
2-] = x2
x = (Ksp)1/2 = 1.1 x 10-5 M
The Common Ion Effect
17-105CHEM 102, Fall 2010, LA TECH
BaSO4(s) Ba2+(aq) + SO42-(aq)
Ksp = 1.1 x 10-10
Solutionb) Now dissolve BaSO4 in water already containing
0.010 M Ba2+. Which way will the “common ion” shift the
equilibrium? ___ Will solubility of BaSO4 be less than or greater than
in pure water?___
The Common Ion Effect
17-106CHEM 102, Fall 2010, LA TECH
BaSO4(s) Ba2+(aq) + SO42-(aq)
Solution [Ba2+] [SO4
2-]initialchangeequilib.
The Common Ion Effect
17-107CHEM 102, Fall 2010, LA TECH
Calculate the solubility of BaSO4 in (a) pure water and (b) in 0.010 M Ba(NO3)2.
Ksp for BaSO4 = 1.1 x 10-10
BaSO4(s) Ba2+(aq) + SO42-(aq)
Solution [Ba2+] [SO4
2-]initial 0.010 0change + y + yequilib. 0.010 + y y
The Common Ion Effect
17-108CHEM 102, Fall 2010, LA TECH
Ksp = [Ba2+] [SO42-] = (0.010 + y) (y)
Because y << x (1.1 x 10-5 M) 0.010 + y 0.010. Therefore,
Ksp = 1.1 x 10-10 = (0.010)(y)y = 1.1 x 10-8 M =
solubility in presence of added Ba2+ ion.
Le Chatelier’s Principle is followed!
The Common Ion Effect