Chemistry 102(01) spring 2009
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Transcript of Chemistry 102(01) spring 2009
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Instructor: Dr. Upali Siriwardanee-mail: [email protected]: CTH 311 Phone 257-4941
Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.m.; Tu,Th,F 9:00 - 10:00 a.m.
Test Dates: March 25, April 26, and May 18; Comprehensive Final Exam: May 20,2009 9:30-10:45 am, CTH 328.
March 30, 2009 (Test 1): Chapter 13 April 27, 2009 (Test 2): Chapters 14 & 15 May 18, 2009 (Test 3): Chapters 16, 17 & 18
Comprehensive Final Exam: May 20,2009 :Chapters 13, 14, 15, 16, 17 and 18
Chemistry 102(01) spring 2009
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Chapter 16. Acids and Bases
16.1 The Brønsted-Lowry Concept of Acids and Bases
16.2 Types of acids/bases:Organic Acids and Amines
16.3 The Autoionization of Water
16.4 The pH Scale
16.5 Ionization Constants of Acids and Bases
16.6 Problem Solving Using Ka and Kb
16.7 Molecular Structure and Acid Strength
16.8 Acid-Base Reactions of Salts
16.9 Practical Acid-Base Chemistry
16.10 Lewis Acid and Bases
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Types of Reactionsa) Precipitation Reactions. Reactions of ionic compounds or saltsb) Acid/base Reactions. Reactions of acids and basesc) Redox Reactions. reactions of oxidizing & reducing
agents
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What are Acids &Bases?
Definition?
a) Arrhenius
b) Bronsted-Lowry
c) Lewis
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Arrhenius, Svante August (1859-1927), Swedish chemist, 1903 Nobel Prize in chemistry
• AcidAcidAnything that produces hydrogen ions in a water solution.
HCl (aq) H+ ( aq) + Cl- ( aq)
• BaseBaseAnything that producs hydroxide ions in a water solution.
NaOH (aq) Na+ ( aq) + OH- ( aq) • Arrhenius definitions are limited proton acids
and hydroxide bases to aqueous solutions.
Arrhenius Definitions
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Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions
AcidAcid Proton donor
BaseBase Proton acceptor
This definition explains how substances like ammonia can act as bases.
Eg. HCl(g) + NH3(g) ------> NH4Cl(s)
HCl (acid), NH3 (base).
NH3(g) + H2O(l) NH4+ + OH-
Brønsted-Lowry definitions
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Lewis DefinitionG.N. Lewis was successful in including acid and bases
without proton or hydroxyl ions.
Lewis Acid: A substance that accepts an electron pair.
Lewis base: A substance that donates an electron pair.
E.g. BF3(g) + :NH3(g) F3B:NH3(s)
the base donates a pair of electrons to the acid forming a
coordinate covalent bond common to coordination
compounds. Lewis acids/bases will be discussed later in
detail
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Dissociation
Strong Acids:
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
Dissociation Equilibrium Weak Acid/base:
H2O(l) + H2O(l) H3+O(aq) + OH-(aq)
This dissociation is called autoionization of water.
HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2
-(aq)
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
Equilibrium constants: Ka, Kb and Kw
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Conjugate acid-base pairs.Conjugate acid-base pairs.
Acids and bases that are related by loss or gain of H+ as H3O+ and H2O.
Examples.Examples. Acid Base
H3O+ H2O
HC2H3O2 C2H3O2-
NH4+ NH3
H2SO4 HSO4-
HSO4- SO42-
Brønsted-Lowry Definitions
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Bronsted acid/conjugate base and base/conjugate acid pairs in
acid/base equilibria
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
HCl(aq): acid
H2O(l): base
H3+O(aq): conjugate acid
Cl-(aq): conjugate base
H2O/ H3+O: base/conjugate acid pair
HCl/Cl-: acid/conjugate base pair
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Select acid, base, acid/conjugate base pair,base/conjugate acid pair
H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4
-(aq)
acid
base
conjugate acid
conjugate base
base/conjugate acid pair
acid/conjugate base pair
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Types of Acids and BasesBinary acids: HCl, HBr, HI, H2S
More than two elements: HCN
Oxyacid: HNO3, H2SO4, H3PO4
Polyprotic acids: H2SO4, H3PO4
Organic acids: R-COOH, R= CH3-, CH3CH2-
Acidic oxides: SO3, NO2, CO2,
Basic oxides: Na2O, CaO
Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary
R2-NH : secondary, R3-N: tertiary
Lewis acids & bases: BF3 and NH3
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Strong Acid vs. Weak AcidsStrong acidcompletely ionized
Hydrioidic HI Ka ~ 1011 pKa = -11Hydrobromic HBr Ka ~ 109 pKa = -9Perchloric HClO4 Ka ~ 107 pKa = -7Hyrdrochloric HCl Ka ~ 107 pKa = -7Chloric HClO3 Ka ~ 103 pKa = -3Sulfuric H2SO4 Ka ~ 102 pKa = -2Nitric HNO3 Ka ~ 20 pKa = -1.3
Weak acidpartially ionized
Hydrofluoric acid HF Ka = 6.6x10-4 pKa = 3.18Formic acid HCOOH Ka = 1.77x10-4 pKa = 3.75Acetic acid CH3COOH Ka = 1.76x10-5 pKa = 4.75Nitrous acid HNO2 Ka = 4.6x10-4 pKa = 3.34Acetyl Salicylic acid C9H8O4 Ka = 3x10-4 pKa = 3.52Hydrocyanic acid HCN Ka = 6.17x10-10 pKa = 9.21
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Strong Base vs. Weak BaseStrong Basecompletely ionizedLithium hydroxide LiOHSodium hydroxide NaOH
Potasium hydroxide KOH Kb~ 102-103
Rubidium hydroxide RbOHCesium hydroxide CsOHBoarder-line Bases
Magnesium hydroxide Mg(OH)2
Calcium hydroxide Ca(OH)2
Strotium hydroxide Sr(OH)2 Kb~ 0.01 to0.1
Barium hydroxide Ba(OH)2
Weak Base
partially ionized
Ammonia NH3 Kb=1.79x10-5 pKb = 4.74
Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25
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• Strong acidsStrong acids Ionize completely in water. HCl, HBr, HI, HClO3,
HNO3, HClO4, H2SO4.
• Weak acidsWeak acids Partially ionize in water. Most acids are weak.
• Strong basesStrong bases Ionize completely in water. Strong bases are metal
hydroxides - NaOH, KOH
• Weak basesWeak bases Partially ionize in water.
Acid and Base StrengthAcid and Base Strength
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Common Acids and BasesAcidsAcids Formula Molarity*
nitric HNO3 16
hydrochloric HCl 12
sulfuric H2SO4 18
acetic HC2H3O2 18
BasesBases
ammonia NH3(aq) 15
sodium hydroxide NaOH solid
*undiluted.
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AutoionizationAutoionization When water molecules react with one another to form ions.
Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle
Kw = [ H3O+ ] [ OH- ]
= 1.0 x 10-14 at 25oC
Note:Note: [H2O] is constant and is included in Kw.
ion productof water
ion productof water
H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M)
Autoionization of Water
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We need to measure and use acids and bases over a very large concentration range.
pH and pOH are systems to keep track of these very large ranges.pH = -log[H3O
+]
pOH = -log[OH-]
pH + pOH = 14
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
pH and other “p” scales
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A logarithmic scale used to keep track of the large changes in [H+].
0 7 1410-14 M 10-7 M 10-14 M
Very Neutral Veryacidic BasicWhen you add an acid to, the pH gets smaller.
When you add a base to, the pH gets larger.
pH scale
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Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
pH of some common materials
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pH of Aqueous Solutions
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What is pH?
Kw = [H3+O][OH-] = 1 x 10-14
[H3+O][OH-] = 10-7 x 10-7
Extreme cases:
Basic medium
[H3+O][OH-] = 10-14 x 100
Acidic medium
[H3+O][OH-] = 100 x 10-14
pH value is -log[H+]
spans only 0-14 in water.
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pH, pKw and pOHThe relation of pH, Kw and pOH Kw = [H+][OH-] log Kw = log [H+] + log [OH-] -log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14 since Kw =1 x 10-14
14 = pH + pOH pH = 14 - pOH pOH = 14 - pH
14 = pH + pOH pH = 14 - pOH pOH = 14 - pH
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pH and pOH calculations of acid and base solutions
a) Strong acids/bases
dissociation is complete for strong acid such as HNO3 or base NaOH
[H+] is calculated from molarity (M) of the
solution
b) weak acids/bases
needs Ka , Kb or percent(%)dissociation
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pH of Strong Acid/bases
HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)
Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.
[HNO3] = [H+] = 0.2 mole/L
pH = -log [H+]
= -log(0.2)
pH = 0.699
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
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pH of 0.5 M H2SO4 Solution
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
[H3+O][HSO4
-]
H2SO4 ; Ka1 = -------------------
[H2SO4]
[H3+O][SO4
2-]
H2SO4 ; Ka2 = ------------------- ; Ka2 ignored
[HSO4-]
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H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning.
[H2SO4] = [H+] = 0.5 mole/L
pH = -log [H+]
pH = -log(0.5) pH = 0.30
pH of 0.5 M H2SO4 Solution
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1.5 x 10-2 M NaOH.1.5 x 10-2 M NaOH.
NaOH is also a strong base dissociates completely in water.
[NaOH] = [HO- ] = 1.5 x 10-2 mole/L
pOH = -log[HO-]= -log(1.5 x 10-2)
pOH = 1.82
As defined and derived previously: pKw= pH + pOH; pKw= 14
pH = pKw + pOH
pH = 14 - pOH
pH = 14 - 1.82 ; pH = 12.18
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Mixtures of Strong and Weak Acids
• the presence of the strong acid retards the dissociation of the weak acid
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Measuring pH
Arnold Beckman
• inventor of the pH meter
• father of electronic instrumentation
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Equilibrium, Constant, Ka & Kb
Ka: Acid dissociation constant for a equilibrium reaction.
Kb: Base dissociation constant for a equilibrium reaction.
Acid: HA + H2O H3+O + A-
Base: BOH + H2O B+ + OH-
[H3+O][ A-] [B+ ][OH-]
Ka = --------------- ; Kb = -----------------
[HA] [BOH]
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HCl(aq) + H2O(l) H3
+O(aq) + Cl-(aq)
[H3+O][Cl-]
Ka= ----------------- [HCl]
[H+][Cl-] Ka= ----------------- [HCl]
Acid Dissociation Constant
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Base Dissociation Constant
NH3 + H2O NH4+ + OH-
[NH4+][OH-]
K = [NH3]
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Hydrated Metal Ions as Acids
[Fe(H2O)6]3+ (aq) + H2O () [Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq)
Ka [Fe(H2O)5 (OH)2 ][H3O ]
Fe(H2O)63 6.310 3
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Ionization Constants for Acids
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Comparing Kw and Ka & Kb
• Any compound with a Ka value greater than Kw of water will be a an acid in water.
• Any compound with a Kb value greater than Kw of water will be a base in water.
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WEAKER/STRONGER Acids and Bases & Ka and Kb values
• A larger value of Ka or Kb indicates an equilibrium favoring product side.
• Acidity and basicity increase with increasing Ka or Kb.
• pKa = - log Ka and pKb = - log Kb
• Acidity and basicity decrease with increasing pKa or pKb.
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Which is weaker?
• a. HNO2 ; Ka= 4.0 x 10-4.
• b. HOCl2 ; Ka= 1.2 x 10-2.
• c. HOCl ; Ka= 3.5 x 10-8.
• d. HCN ; Ka= 4.9 x 10-10.
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What is Ka1 and Ka2?
• H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
• HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
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H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
[H3+O][HSO4
-]
H2SO4 ; Ka1 = -------------------
[H2SO4]
[H3+O][SO4
2-]
H2SO4 ; Ka2 = -------------------
[HSO4-]
Ka Examples
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HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2
-(aq)
[H+][C2H3O2-]
H C2H3O2; Ka= ------------------
[H C2H3O2]
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
[NH4+][OH-]
NH3; Kb= --------------
[ NH3]
Ka Examples
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How do you calculate pH of How do you calculate pH of weak acids/basesweak acids/bases
From % dissociation
From Ka or Kb
What is % dissociation
Amount dissociated
% Dissoc. = ------------------------- x 100
Initial amount
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How do you calculate % How do you calculate % dissociation from Kdissociation from Kaa or K or Kbb
1.00 M solution of HCN; Ka = 4.9 x 10-10
What is the % dissociation for the acid?
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1.00 M solution of HCN; Ka = 4.9 x 10-10
First write the dissociation equilibrium equation:HCN(aq) + H 2O(l) <===> H 3
+O(aq) + CN-(aq)
[HCN] [H+ ] [CN- ]
Ini. Con. 1.00 M 0.0 M 0.00 M
Cha. Con -x x x
Eq. Con. 1.0 - x x x
[H 3+O ][CN-] x2
Ka = ----------- = ----------------
[HCN] 1.0 - x
1.00 M solution of HCN; Ka = 4.9 x 10-10
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1.0 - x ~ 1.00 since x is small
x2
Ka = -----------; Ka = 4.9 x 10-10 = x2
1.0
x = 4.9 x 10-10 = 2.21 x 10 -5
Amount disso. 2.21 x 10 -5
----------------- x 100 =- ------------- x 100 Ini. amount 1.00
% Diss. =2.21 x 10 -5 x 100 = 0.00221 %
1.00 M solution of HCN; Ka = 4.9 x 10-10
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% Dissociation gives x (amount dissociated) need for pH calculation
Amount dissociated
% Dissoc. = ------------------------- x 100
Initial amount/con.
x
% Dissoc. = --------------------------- x 100
concentration
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1 M HF, 2.7% dissociated
Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027
Calculate the pH of a weak acid from % dissociation
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• HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq)
• [H+][F-] • Ka = -----------• [HF]• [HF] [H+ ] [F- ]• Ini. Con. 1.00 M 0.0 M 0.00 M• Chg. Con -x x x• Eq.Con. 1.0-0.027 0.0270.027• pH = -log [H+] • pH = -log(0.027) • pH = 1.57
Calculate the pH of a weak acid from % dissociation
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Weak acid EquilibriaExampleExample
Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = 6.28 x 10-5
HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq)
The first step is to write the equilibrium expression
Ka = [H3O+][Bz-]
[HBz]
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Weak acid Equilibria HBz H3O+ Bz-
Initial conc., M 0.10 0.00 0.00
Change, M -x x x
Eq. Conc., M 0.10 - x x x
[H3O+] = [Bz-] = x
We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.
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Weak Acid Equilibria
Solve the equilibrium equation in terms of Solve the equilibrium equation in terms of xx
Ka = 6.28 x 10-5 =
x = (6.28 x 10-5 )(0.10)
H3O+ = 0.0025 M
pH = 2.60
x2
0.10
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pH from Ka or Kb
1.00 M solution of HCN; Ka = 4.9 x 10-10
First write the dissociation equilibrium equation:
HCN(aq) + H 2O(l) H 3+O(aq) + CN-(aq)
[HCN] [H+ ] [CN- ]
Ini. Con. 1.00 M 0.0 M 0.00 M
Chg. Con -x x x
Eq. Con. 1.0 - x x x
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[H 3+O ][CN-] x2
Ka = --------------- = ----------------
[HCN] 1.0 - x
1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2
1.0
x = 4.9 x 10-10 = 2.21 x 10 -5
pH = -log [H+]
pH = -log(2.21 x 10-5)
pH = 4.65
Weak Acid Equilibria
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The Conjugate Partners of Strong Acids and Bases
The conjugate acid/base of a strong base/acid has no net effect on the pH of a solution
The conjugate base of a weak acid hydrolyze in water and basic or
pH of a solution > 7.00 E.g. Na+C2H3O2- sodium acetate
The conjugate acid of a weak base hydrolyze in water and acidic or
pH of a solution < 7.00 E.g NH4Cl
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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.
HydrolysisHydrolysis
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
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Acid-Base Properties of Typical Ions
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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.
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What pH? Neutral, basic or acidic?
•a)NaCl • neutral•b) NaC2H3O2
• basic•c) NaHSO4 • acidic•d) NH4Cl• acidic
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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 base
• If basic it should be a salt of weak acid
• if 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
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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
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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)
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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
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Types of Acids and Bases
• Binary acids
• Oxyacid
• Organic acids
• Acidic oxides
• Basic oxides
• Amine
• Polyprotic acids
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Influence of Molecular Structure on Acid Strength
Binary Hydrides– hydrogen & one other element
• Bond Strengths– weaker the bond, the stronger the acid
• Stability of Anion– higher the electronegativity, stronger the acid
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Binary Acids
Compounds containing acidic protons bonded to a more electronegative atom.
e.g. HF, HCl, HBr, HI, H2S
The acidity of the haloacid (HX; X = Cl, Br, I, F)Series increase in the following order: HF < HCl < HBr < HI
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Oxyacids
Compounds containing acidic - OH groups in the molecule.
Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens)
The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend.
HClO4 > HClO3 > HClO2 > HClO
perchloric chloric chlorus hyphochlorus
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Influence of Molecular Structure on Acid Strength
Oxyacids– hydrogen, oxygen, & one other element
H-O-E– higher the electronegativity on E, stronger the
acid as this weakens the bond between the O and H
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< <<
<
Oxo Acid
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Acidic Oxides
These are usually oxides of non-metallic elements such as P, S and N.
E.g. NO2, SO2, SO3, CO2
They produce oxyacids when dissolved in water
SO3 + H2O ---> H2SO4
CO2 + H2O ---> H2CO3
NO2 + H2O ---> HNO3
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Basic Oxides
Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.
e.g. CaO + H2O ---> Ca(OH)2
Na2O + H2O ---> 2 NaOH
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Protic Acids
Monoprotic Acids: The form protic refers to acidity due to protons. Monoprotic acids have only one acidic proton. e.g. HCl.
Polyprotic Acids: They have more than one acidic proton.
e.g. H2SO4 - diprotic acid
H3PO4 - triprotic acid.
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Polyprotic Acids
• acids where more than one hydrogen per molecule is released
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Polyprotic Acids
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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 hydrogen
CH 3 C
O
OH
-
CH 3 C
O
O-
acetate ion
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FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid).
Acid Ka pKa
HCOOH (formic acid) 1.78 X 10-43 0.75
CH3COOH (acetic acid) 1.74 X 10-54 0.76
CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86
Organic or Carboxylic Acids
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Amines
Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted-Lowery or Lewis acid/base definitions.
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Amines
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Acid-Base Chemistryof Some Antacids
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Acid-Base in the Kitchen
vinegar - acetic acid
lemon juice (citrus juice) - citric acid
baking soda - NaHCO3
milk - lactic acid
baking powder - H2PO4- & HCO3
-
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Household Cleaners
CH 3CH2CH 2CH2CH 2CH2CH 2CH 2CH 2CH 2CH2CH 2CH2CH 2SO3
-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
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Dishwashing Detergent
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Lewis DefinitionG.N. Lewis was successful in including acid and bases
without proton or hydroxyl ions.
Lewis Acid: A substance that accepts an electron pair.
Lewis base: A substance that donates an electron pair.
E.g. BF3(g) + :NH3(g) F3B:NH3(s)
the base donates a pair of electrons to the acid forming a
coordinate covalent bond common to coordination
compounds. Lewis acids/bases will be discussed later in
detail
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Lewis Acids and Bases Reactions
H+ + NH3
acid base
Cu+2 + 4 NH3 [Cu(NH3)4+2]
acid base
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What acid base concepts (Arrhenius/Bronsted/Lewis) would best
describe the following reactions:
•a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)
•b)HCl(g) + NH3(g) ---> NH4Cl(s)
•c)BF3(g) + NH3(g) ---> F3B:NH3(s)
•d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)