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Chapter 16
Overview:
•Definitions
Arrhenius
Bronsted -- Conjugate Pairs
Hydronium Ion
•Relative Strengths
Strong/Weak acids and reactions
Strong/Weak bases and reactions
Ka’s and Kb’s
•pH and H2O Ionization
•Calculating Equilibrium Concentrations
•Hydrolysis
Acidic and Basic Salts
•Polyprotic Acids
•Molecular Structure/Bonding
•Lewis Acids and Bases
Acids
•sour
•corrosive
•reddens blue vegetable colors
•react with bases
Bases
•bitter
•soapy
•restores vegetable colors reddened by acids
•react with acids
Arrhenius Acid:
substance that releases or produces H+
HCl(aq)  H+(aq)
+
Cl -(aq)
Arrhenius Base:
substance that releases or produces OH NaOH(aq) Na+(aq)
+ OH -(aq)
Bronsted Acid:
substance that donates a H+ to another
HNO3(aq) + H2O(l)

H3O+(aq)
+ NO3-(aq)
Bronsted Base:
substance that accepts a H+ from another
CO32-(aq) + H2O(l) 
HCO3-(aq)
+ OH -(aq)
Water Dissociation:
2H2O
H3O+ +
OH -
Hydronium Ion -- because
bare protons are unlikely
d+
O H
H
d-
-
O H
O +H O H
H
H
H+ transfer
H
+
pH and Water Ionization:
2H2O(l)
H3O+(aq) +
OH -(aq)
K = [H3O+][OH -] = K [H2O]2 = [H3O+][OH -]
[H2O] 2
Kw = [H3O+][OH -] = 1.0 x 10 -14
(at 25C)
Reactant Favored
ion-product constant for water
pH = - log [H+] = - log [H3O+]
neutral solution:
2H2O
[H3O+] = [OH -]
H3O+
+
OH -
1.0 x 10 -14 = [H3O+][OH -] = [H3O+]2
[H3O+] = 1.0 x 10 -7 M
pH = - log (1.0 x 10 -7) = 7.0
non-neutral solutions:
acidic solutions:
[H+] > [OH-]
pH < 7.0
basic solutions:
[H+] < [OH-]
pH >
7.0
pH of Strong Acid and Base Solutions:
Calculate the pH of a solution containing 0.00100 M of
a strong acid such as HCl.
HCl(aq)
0.00100
H+(aq) + OH-(aq)
0.00100
0.00100
pH = - log (0.00100) = 3.00
but. . . .
What about H3O+ from water?
2 H2O
initial
change
equil.
H3O+
+
0.00100
OH 0
+x
+x
0.00100 + x
x
1 x 10 -14 = (0.00100 + x)(x)
x = H3O+ and OH- from disso. of water
Kw = [H3O+] [OH -]
0
x is very small
compared to
0.00100 & can be
neglected
1 x 10 -14 = (0.00100 + x)(x)
x = 1.00 x 10 -11 M = [OH -]
pOH = 11.0
[H+]  [H3O+]
[H3O+] (1.0 x 10-11) = 1.0 x 10 -14
[H3O+] = 1.0 x 10 -3 M
the conc. of H+ in a
solution of a strong
acid is the conc. of
the strong acid
pH = 3.0
Note: pH + pOH = 14.0
Bottom Line:
 [H+] = [H3O+] = [H5O2+] = [H9O4+] hydrated
hydrogen ions
 neutral solution [H3O+] = [OH-] pH = 7
 acidic solution [H3O+] > [OH-] pH < 7
 basic solution [H3O+] < [OH-] pH > 7
 the concentration of [H3O+] in a strong acid
is the concentration of the acid
Calculate the pH of a solution containing
0.010 M KOH.
KOH

0.010 M
K+(aq)
+
0.010 M
OH-(aq)
0.010 M
strong base, complete rxn, stoichiometric
1.0 x 10 -14 = [H+] (0.010 M)
[H+] = 1.0 x 10 -12 M
pH = 12.0
contribution of [H+] from
dissociation of H2O is
negligible
Calculate the [H+] and [OH -] in a
solution that has a pH = 8.60
- log [H+] = pH
- log [H+] = 8.60
log [H+] = -8.60
[H+] = anti log (-8.60)
[H+] = 1.8 x 10 -4 M
[OH -] = 5.6 x 10 -11 M
pH
14
[H3O+]
10 -14
[OH-]
1
pOH
0
Basic
ammonia
human blood
milk
7.0
10 -7
10 -7
7.0
vinegar, cola
0
1
10 -14
14
Neutral
Acidic
Measuring pH:
 pH meter -- electrodes measure pH
– most precise method
 acid-base indicators
– less precise but good when a pH meter is not
available
– substances which are differently colored at
different pH values
– litmus, phenolphthalein, thymol blue
Arrhenius Acid:
substance that releases or produces H+
HCl(aq)  H+(aq)
+
Cl -(aq)
Arrhenius Base:
substance that releases or produces OH NaOH(aq) Na+(aq)
+ OH -(aq)
Bronsted Acid:
substance that donates a H+ to another
HNO3(aq) + H2O(l)

H3O+(aq)
+ NO3-(aq)
Bronsted Base:
substance that accepts a H+ from another
CO32-(aq) + H2O(l) 
HCO3-(aq)
+ OH -(aq)
Examples:
H+
NH4+(aq) + H2O(l)  H3O+(aq) + NH3(aq)
acid
base
H+
conjugate
acid
conjugate
base
conjugate pair
PO43-(aq) + H2O(l)  HPO42-(aq) + OH -(aq)
base
acid
conjugate
acid
conjugate
base
conjugate pair
Some species can act as an acid or base:
H+
HCO3-(aq) + H2O(l)
acid

H3O+(aq) + CO32-(aq)
conjugate
acid
base
conjugate
base
H+
HCO3-(aq) + H2O(l)
base
acid

H2CO3(aq) + OH -(aq)
conjugate
acid
conjugate
base
HCO3- is an amphiprotic substance
You Must Know:
what an acid and a base is and how to
identify both
know definitions and properties
the reaction of an acid and a base with
water
how to identify acid, base, conjugate
acid and conjugate base
what the hydronium ion is and the
ionization reaction of water
Problems:
HX + H2O  H3O+ + X A
B
CA
•What is the conjugate base of:
H2S
NH4+
NH3
H2O
OH -
HS NH3
NH2OH O2-
•What is the conjugate acid of:
NO3HNO3
HPO42H2PO4H3SO4+
H2SO4
CB
Relative Strengths:
HCl +
stronger A
H2O
stronger B
H3O+ + Cl -
weaker A
weaker B
equilibrium is a competition between the bases
H2O and Cl - -- the equilibrium will lie toward
the direction of the weaker acid and base
in this case, H2O is a stronger base than Cl - as
it competes much more effectively for the H+
HCN +
weaker A
H2O
weaker B
H3O+ + CN stronger A
H2O < CN -
stronger B
Which is the weaker acid:
H2S + CN -
HCN + HS -
HCN
HCO3- + SO42-
HSO4- + CO32-
HCO3-
HClO4 + H2O
H3O+ + ClO4-
H3O+
NH4+ +
H3O+ + NH3
NH4+
H 2O
The stronger the acid the weaker its conjugate base:
strongest acid
weakest acid
weakest
conjugate base
strongest
conjugate base
Given the following, which is the weaker
conjugate:
acids
bases
HCl > CH3CO2H
Cl -
HCN < H3PO4
CN -
H2SO4 > H2SO3
HSO4
NH3 > H2O
NH4
H
-
>
> H2PO4 -
2-
NH3
< HSO3<
+
HSO4 < CO3
-
< CH3CO2-
H3O+
H2SO4 > HCO3 H2
<
NH4+
Predicting Direction of Acid/Base Rxns.
CH3CO2H
stronger A
stronger B
+ CN -
weaker B
HCN + CH3CO2 -
weaker A
stronger A
HSO4-
+ NH3
stronger B
weaker A
NH4+
+ SO42-
weaker B
Strong Acids and Bases
•Strong Acids
HCl
HNO3
HClO4
HClO3
H2SO4
HBr
HI
HX + H2O
H3 O+ + X -
•Strong Bases
Grp I hydroxides
Grp II hydroxides
(except Be)
MOH
essentially complete rxns.
M+ + OH -
Weak Acids and Bases:
HX + H2O
H3O+ + X -
Ka = [H3O+][X -]
[HX]
B + H2O
< 1
BH+ + OH -
Kb = [BH+][OH -]
[B]
< 1
Weak Acids Can Be:
•cations
NH4+ or [Cu(H2O)6]2+
•anions
H2PO4- or HCO3-
•neutral
CH3CO2H or HCO2H
Weak Bases Can Be:
•anions
CO32- or
•neutral
NH3
or
CN (CH3)3N
You Must:
Be able to determine direction of rxn
based on acid or base relative strengths
Know the strong acids and strong bases
Be able to recognize weak acids and bases
Know water ionization rxn, constant and
expression
Know what pH is and how to calculate it
Know how to calculate equil. conc., pH,
pOH for weak acids and bases
Problem 1: A 0.015 M solution of an
unknown base has a pH of 10.09. Is it a strong
or weak base? What is the Kb, if it is weak?
Problem 2: What are the equil. conc. of
H3O+, acetate ion, acetic acid in a 0.20 M
aqueous solution of acetic acid, CH3CO2H?
Use the approximation whenever possible: When the initial
weak acid/base conc. > 100 *Ka/b neglect x when it is added
to or subtracted from the initial conc.
Answer 1: If this were a strong base then
[OH-] = 0.015 M and pOH = 1.8 and pH = 12.2.
B + H2O
HB+ + OH-
pH = 10.09 and pOH = 3. 91  [OH-] = 1.2 x 10-4
Kb = [HB+][OH-] = (1.2 x 10 -4)2
[B]
(0.015)
= 9.6 x 10 -7
Answer 2:
CH3CO2H + H2O
initial 0.20
change -x
equil. 0.20 -x
H3O+ + CH3CO20
+x
x
0
+x
x
Ka = 1.8 x 10 -5 = [H3O+][CH3CO2-]
[CH3CO2H]
Ka = 1.8 x 10 -5 = (x)(x)
(0.20 -x)
x2 = 3.6 x 10 -6
x = 1.9 x 10 -3
[CH3CO2-]=[H3O+] = 1.9 x 10 -3 M
[CH3CO2H] = 0.20 M
pH = 2.7
What is the percent ionization?
% ion = [H3O+] x 100
[CH3CO2H]
= 0.95 %
Polyprotic Acids and Bases:
H3PO4 + H2O
H3O+ + H2PO4- Ka1
H2PO4- + H2O
H3O+ + HPO42- Ka2
HPO42- + H2O
H3O+ + PO43-
H3PO4 + 3H2O
Ka3
3H3O+ + PO43--
Ka tot = Ka1 Ka2 Ka3
(7.5 x 10 -3)(6.2 x 10 -8)(3.6 x 10 -13) = 1.7 x 10 -22
Weak Bases:
 Can be anions such as:
– CN-, HSO3-, SO32-, HCO3-, CO32-, etc.
 Can be N-containing compounds, such as:
– NH3, (CH3)NH2, (CH3)2NH, (CH3)3N, etc.
 React with water
– B- + H2O
HB
+
OH-
– B + H2O
HB+ +
OH-
or
 Have base dissociation constants, Kb
Example:
NH3 + H2O
initial 0.20 M
change -x
equil. 0.20 -x
NH4+ + OH 0
+x
x
0
+x
x
Kb = 1.8 x 10 -5 = [NH4+][OH -]
[NH3]
Kb = 1.8 x 10 -5 = (x)(x)
(0.20 -x)
x2 = 3.6 x 10 -6
x = 1.9 x 10 -3
[NH4+]=[OH-] = 1.9 x 10 -3 M
[NH3] = 0.20 M
pOH = 2.7
pH = 12.3
CO3 -2 + H2O
OH - + HCO3--
Kb1
HCO3 -- + H2O
OH - + H2CO3
Kb2
CO3 + 2H2O
2OH -- + H2CO32--
Kb tot = Kb1 Kb2
(2.1 x 10 -4)(2.4 x 10 -8) = 5.0 x 10 -12
Relationship of an acid Ka and the Kb
of its conjugate base:
conjugate pair rxns with water
NH3 + H2O
NH4+ +
NH4+ + H2O
2H2O
NH3
H 3O+
+
+
OH H3 O +
OH -
Kb = 1.8 x 10 -5
Ka = 5.6 x 10 -10
Kw = 1 x 10 -14
For Conjugate Acid/Base Pairs:
Kb Ka = K w
or
pKb + pKa = pKw
Find Ka for the conjugate acids of
the following bases: Ka = Kw/Kb
Base
Kb
Conj. Acid
NH3
1.8 x 10 -5
NH4+
C5H5N 1.7 x 10 -9
HS CO3
2-
Ka
5.6 x 10 -10
C5H5NH + 5.9 x 10 -6
1.8 x 10 -7
H2S
5.6 x 10 -8
1.8 x 10
HCO3 -
5.6 x 10 -11
-4
Acid/Base Hydrolysis
Many salts produce solutions that are acidic or basic
Why?
because either (or both) the cation or anion of the salt
acted as a weak acid or a weak base
HX + H2O
H3O+ + X -
weak acid
or
B + H2O
weak base
BH+ + OH -
Which cations will not hydrolyze?
Cations of strong bases -- Grp I and II (except
Be)
Li+ Na+ K+ Rb+ Cs+ Mg2+ Ca2+ Sr2+ Ba2+
Anions of strong acids -ClO4- ClO3- SO42-
NO3-
Cl -
Br -
I-
All other cations and anions will hydrolyze
Which cations will hydrolyze?
Conjugate acids of weak bases:
NH4+ Al3+ Cu2+ etc.
Conjugate bases of weak acids:
CO32 - CH3CO2- NH3 F - etc.
What is the pH of the following salt solutions?
•NaCl
Na+ -- neutral
Cl - -- neutral
N
•NaF
Na+ -- neutral
F - -- basic
B
•NH4Cl
NH4+ -- acidic
Cl - -- neutral
A
•Na2CO3
Na+ -- neutral
CO32- -- basic
B
•Cr(NO3)3
Cr3+ -- acidic
NO3- -- neutral
A
•KNO3
K+ -- neutral
NO3+ -- neutral
N
•NaC3H2O4
Na+ -- neutral
C3H2O4- -- basic
B
•NH4CH3CO2
NH4+ -- acidic
CH3CO2- -- basic ?
Is NH4CH3CO2 acidic or basic?
NH4+ +
H 2O
H3O+
+
NH3
Ka = 5.6 x 10 -10
CH3CO2- +
H2O
CH3CO2H + OH -
Kb = 5.6 x 10 -10
weak acid & weak base strength equal -- neutral soln.
You Must:
Know what hydrolysis is
Know which cations hydrolyze or produce acidic
solutions
Know which anions hydrolyze or produce basic
solutions
Be able to estimate and calculate the pH of a
salt solution
Know how to determine Ka’s & Kb’s of conjugate
pairs
Strength of Acids:
Can be increased by anything that facilitates
the loss of a Proton (H+)
•Bond Strength
•Polarity or Delectronegativity
•Central Atom Charge or Oxidation #
IVA
VA
VIA
VIIA
row 2
CH4
NH3
H2O
HF
row 3
SiH4
PH3
H2S
HCl
Increasing acid strength
Increasing acid strength
Examples:
HClO4
HClO3
HClO2
HClO
Cl ox. #
Increasing acid strength
Acid
Ka
+7
strong
+5
strong
+3
1.1 x 10 -2
+1
3.0 x 10 -8
HClO
HBrO
HIO
Ka
Increasing acid strength
Acid
3 x 10 -8
2.5 x 10 -9
2.3 x 10 -11
Organic Acids:
H
H
O
C
C
H
O-
H
Acetate ion
resonance
H
O-
C
C
H
O
H
H
O
C
C
OH
H
Acetic Acid
Ka = 1.8 x 10 -5
F
F
O
C
C
OH
F
Trifluoroacetic Acid
Ka = 5.0 x 10 -1
Lewis Acids and Bases:
Lewis Acid -- accepts a pair of electrons
Lewis Base -- donates a pair of electrons
formation of a coordinate covalent bond
H
H
B
N
H
F
F
F
NH3
H
H
Lewis Base -donates a pair of
electrons
+
H+
H+
N
H
NH4
+
Lewis Acid -accepts a pair of
electrons