Hoa huu co_chuong 17

Philip Dutton
University of Windsor, Canada
N9B 3P4

Prentice-Hall © 2002

General Chemistry
Principles and Modern Applications
Petrucci • Harwood • Herring
8th Edition
Chapter 17: Acids and Bases
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General Chemistry: Chapter 17
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Contents
17-1 The Arrhenius Theory: A Brief Review
17-2 Brønsted-Lowry Theory of Acids and Bases
17-3 The Self-Ionization of Water and the pH Scale
17-4 Strong Acids and Strong Bases
17-5 Weak Acids and Weak Bases
17-6 Polyprotic Acids
17-7 Ions as Acids and Bases
17-8 Molecular Structure and Acid-Base Behavior
17-8 Lewis Acids and Bases
Focus On Acid Rain.


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General Chemistry: Chapter 17
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17-1 The Arrhenius Theory:
A Brief Review
HCl(g) → H+(aq) + Cl-(aq)
H2O
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → H2O(l) + Na+(aq) + Cl-(aq)
H+(aq) + OH-(aq) → H2O(l)
Arrhenius theory did not handle non OH- bases such as ammonia very well.
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17-2 Brønsted-Lowry Theory of
Acids and Bases
An acid is a proton donor.
A base is a proton acceptor.
NH3 + H2O  NH4+ + OH-
NH4+ + OH-  NH3 + H2O
base
acid
base
acid
conjugate acid
conjugate base
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Base Ionization Constant
NH3 + H2O  NH4+ + OH-
base
acid
conjugate
acid
conjugate
base
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Acid Ionization Constant
CH3CO2H + H2O  CH3CO2- + H3O+
base
acid
conjugate
acid
conjugate
base
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Table 17.1 Relative Strengths of Some Brønsted-Lowry Acids and Bases
HClO4 + H2O  ClO4- + H3O+
NH4+ + CO32-  NH3 + HCO3-
HCl + OH-  Cl- + H2O
H2O + I-  OH- + HI
H2O + I-  OH- + HI
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17-3 The Self-Ionization of Water and
the pH Scale
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Ion Product of Water
H2O + H2O  H3O+ + OH-
base
acid
conjugate
acid
conjugate
base
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pH and pOH
The potential of the hydrogen ion was defined in 1909 as the negative of the logarithm of [H+].
pH = -log[H3O+]
pOH = -log[OH-]
-logKW = -log[H3O+]-log[OH-]= -log(1.010-14)
KW = [H3O+][OH-]= 1.010-14
pKW = pH + pOH= -(-14)
pKW = pH + pOH = 14
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pH and pOH Scales
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17-4 Strong Acids and Bases
HCl CH3CO2H
Thymol Blue Indicator
pH < 1.2 < pH < 2.8 < pH
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17-5 Weak Acids and Bases
Acetic Acid
HC2H3O2 or CH3CO2H
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Weak Acids
Ka=
= 1.810-5
[CH3CO2H]
[CH3CO2-][H3O+]
pKa= -log(1.810-5) = 4.74
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Table 17.3 Ionization Constants of Weak Acids and Bases
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Example 17-5
Determining a Value of KA from the pH of a Solution of a Weak Acid.
Butyric acid, HC4H7O2 (or CH3CH2CH2CO2H) is used to make compounds employed in artificial flavorings and syrups. A 0.250 M aqueous solution of HC4H7O2 is found to have a pH of 2.72. Determine KA for butyric acid.
HC4H7O2 + H2O  C4H77O2 + H3O+ Ka = ?
Solution:
For HC4H7O2 KA is likely to be much larger than KW. Therefore assume self-ionization of water is unimportant.
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Example 17-5
HC4H7O2 + H2O  C4H7O2 + H3O+
Initial conc. 0.250 M 0 0
Changes -x M +x M +x M
Eqlbrm conc. (0.250-x) M x M x M
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General Chemistry: Chapter 17
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Example 17-5
Log[H3O+] = -pH = -2.72
HC4H7O2 + H2O  C4H77O2 + H3O+
[H3O+] = 10-2.72 = 1.910-3 = x
Ka= 1.510-5
Check assumption: Ka >> KW.
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Percent Ionization
HA + H2O  H3O+ + A-
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Percent Ionization
Ka =
[H3O+][A-]
[HA]
Ka =
n
H3O+
A-
n
HA
n
1
V
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17-6 Polyprotic Acids
H3PO4 + H2O  H3O+ + H2PO4-
H2PO4- + H2O  H3O+ + HPO42-
HPO42- + H2O  H3O+ + PO43-
Phosphoric acid:
A triprotic acid.
Ka = 7.110-3
Ka = 6.310-8
Ka = 4.210-13
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Phosphoric Acid
Ka1 >> Ka2
All H3O+ is formed in the first ionization step.
H2PO4- essentially does not ionize further.
Assume [H2PO4-] = [H3O+].
[HPO42-] Ka2 regardless of solution molarity.
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Table 17.4 Ionization Constants of Some Polyprotic Acids
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Example 17-9
Calculating Ion Concentrations in a Polyprotic Acid Solution.
For a 3.0 M H3PO4 solution, calculate:
(a) [H3O+]; (b) [H2PO4-]; (c) [HPO42-] (d) [PO43-]
H3PO4 + H2O  H2PO4- + H3O+
Initial conc. 3.0 M 0 0
Changes -x M +x M +x M
Eqlbrm conc. (3.0-x) M x M x M
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Example 17-9
H3PO4 + H2O  H2PO4- + H3O+
[H3O+] [H2PO4-]
[H3PO4]
Ka=
x · x
(3.0 – x)
=
Assume that x << 3.0
= 7.110-3
x2 = (3.0)(7.110-3) x = 0.14 M
[H2PO4-] = [H3O+] = 0.14 M
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Example 17-9
H2PO4- + H2O  HPO42- + H3O+
Initial conc. 0.14 M 0 0.14 M
Changes -y M +y M +y M
Eqlbrm conc. (0.14 - y) M y M (0.14 +y) M
y << 0.14 M y = [HPO42-] = 6.310-8
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Example 17-9
HPO4- + H2O  PO43- + H3O+
[H3O+] [HPO42-]
[H2PO4-]
Ka=
(0.14)[PO43-]
6.310-8
=
= 4.210-13 M
[PO43-] = 1.910-19 M
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Sulfuric Acid
Sulfuric acid:
A diprotic acid.
H2SO4 + H2O  H3O+ + HSO4-
HSO4- + H2O  H3O+ + SO42-
Ka = very large
Ka = 1.96
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General Approach to Solution Equilibrium Calculations
Identify species present in any significant amounts in solution (excluding H2O).
Write equations that include these species.
Number of equations = number of unknowns.
Equilibrium constant expressions.
Material balance equations.
Electroneutrality condition.
Solve the system of equations for the unknowns.
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17-7 Ions as Acids and Bases
NH4+ + H2O  NH3 + H3O+
base
acid
CH3CO2- + H2O  CH3CO2H + OH-
base
acid
Ka Kb = Kw
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Hydrolysis
Water (hydro) causing cleavage (lysis) of a bond.
Na+ + H2O → Na+ + H2O
NH4+ + H2O → NH3 + H3O+
Cl- + H2O → Cl- + H2O
No reaction
No reaction
Hydrolysis
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17-8 Molecular Structure and
Acid-Base Behavior
Why is HCl a strong acid, but HF is a weak one?
Why is CH3CO2H a stronger acid than CH3CH2OH?

There is a relationship between molecular structure and acid strength.
Bond dissociation energies are measured in the gas phase and not in solution.
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Strengths of Binary Acids
HI HBr HCl HF
160.9 > 141.4 > 127.4 > 91.7 pm
297 < 368 < 431 < 569 kJ/mol
Bond length
Bond energy
109 > 108 > 1.3106 >> 6.610-4
Acid strength
HF + H2O → [F-·····H3O+]  F- + H3O+
ion pair
H-bonding
free ions
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Strengths of Oxoacids
Factors promoting electron withdrawal from the OH bond to the oxygen atom:
High electronegativity (EN) of the central atom.
A large number of terminal O atoms in the molecule.
H-O-Cl H-O-Br
ENCl = 3.0 ENBr= 2.8
Ka = 2.910-8 Ka = 2.110-9
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Ka 103 Ka =1.310-2
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Strengths of Organic Acids
C
O
C
O
H
H
··
··
H
H
O
C
H
H
··
··
H
H
C
H
H
Ka = 1.810-5 Ka =1.310-16
acetic acid ethanol
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Focus on the Anions Formed
O
C
H
··
··
H
H
C
H
H
··
-
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Structural Effects
C
O
C
O
H
-
··
H
H
C
H
H
H
Ka = 1.810-5

Ka = 1.310-5
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Structural Effects
C
O
C
O
H
-
··
H
H
Ka = 1.810-5

Ka = 1.410-3
C
O
C
O
H
-
··
H
Cl
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Strengths of Amines as Bases
N
H
H
H
··
N
Br
H
H
··
pKb = 4.74 pKa = 7.61
ammonia bromamine
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Strengths of Amines as Bases
C
H
H
H
C
H
H
C
H
H
H
C
H
H
C
H
H
H
C
H
H
pKb = 4.74 pKa = 3.38 pKb = 3.37
methylamine ethylamine propylamine
NH2
NH2
NH2
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Resonance Effects
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Inductive Effects
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17-9 Lewis Acids and Bases
Lewis Acid
A species (atom, ion or molecule) that is an electron pair acceptor.
Lewis Base
A species that is an electron pair donor.
base
acid
adduct
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Showing Electron Movement
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CO2 + H2O  H2CO3
H2CO3 + H2O  HCO3- + H3O+
3 NO2 + H2O  2 HNO3 + NO
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Chapter 17 Questions

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Choose a variety of problems from the text as examples.

Practice good techniques and get coaching from people who have been here before.