Acids and Bases
Chapter 14
Properties of Acids
Acids:
•
taste sour (citrus fruits & vinegar)
•
affect indicators (e.g. turn blue litmus red)
•
produce H+ ions in aqueous solution
•
corrosive to metals
•
pH < 7
Classifying Acids
Organic acids contain a carboxyl group or
-COOH -- HC2H3O2 & citric acid.
Inorganic acids -- HCl, H2SO4, HNO3.
Oxyacids -- acid proton attached to oxygen
-- H3PO4.
Monoprotic -- HCl & HC2H3O2
Diprotic -- H2SO4
Triprotic -- H3PO4
Properties of Bases
Bases:
•
taste bitter
•
feel slippery
•
affect indicators (e.g. turn red litmus blue)
•
produce OH- ions in aqueous solution
•
pH > 7
•
caustic
Models of Acids and Bases
Arrhenius Concept: Acids produce H+ in
solution, bases produce OH ion.
Brønsted-Lowry: Acids are H+ donors, bases
are proton acceptors.
HCl + H2O  Cl + H3O+
acid base
Hydronium Ion
Hydronium ion is a hydrated proton -H+.H2O.
The H+ ion is simply a proton. It has a very
high charge density, so it strongly is
attracted to the very electronegative
oxygen of the polar water molecule.
Conjugate Acid/Base Pairs
HA(aq) + H2O(l)  H3O+(aq) + A(aq)
conj
acid 1
conj
base 2
conj
acid 2
conj
base 1
conjugate base: everything that remains of
the acid molecule after a proton is lost.
conjugate acid: formed when the proton is
transferred to the base.
Which is the stronger base--H2O or A-?
Acid Dissociation Constant (Ka)
HA(aq) + H2O(l)  H3O+(aq) + A(aq)
Ka 
H3O

HA
A


H

A

HA
Ka values for common acids are found in Table 14.2
on page 663.
14_02T
Formula
HSO4
HClO2
HC2H2ClO2
HF
HNO2
HC2H3O2
[Al(H2O)6]3+
HOCl
HCN
NH4
HOC6H5
Values of Ka for Some Common Monoprotic Acids
Name
Hydrogen sulfate ion
Chlorous acid
Monochloracetic acid
Hydrofluoric acid
Nitrous acid
Acetic acid
Hydrated aluminum(III) ion
Hypochlorous acid
Hydrocyanic acid
Ammonium ion
Phenol
*The units of Ka are mol/L but are customarily omitted.
Value of K a*
1.2 x 102
1.2 x 102
1.35 x 103
7.2 x 104
4.0 x 104
1.8 x 105
1.4 x 105
3.5 x 108
6.2 x 1010
5.6 x 1010
1.6 x 1010
Increasing acid strength
Table 14.2
Bronsted-Lowry Model
The Bronsted-Lowry Model is not limited to
aqueous solutions like the Arrhenius
Model.
NH3(g) + HCl(g) ----> NH4Cl(s)
This is an acid-base reaction according to
Bronsted-Lowry, but not according to
Arrhenius!
Acid Strength
Strong Acid:
-
Its equilibrium position lies far to the right.
(HNO3)
-
Yields a weak conjugate base. (NO3)
Acid Strength
(continued)
Weak Acid:
-
Its equilibrium lies far to the left.
(CH3COOH)
-
Yields a much stronger (water is relatively
strong) conjugate base than water.
(CH3COO)
14_1577
H+
+
H
A- H + A
AH+
A(a)
A- H +
H+
H+
A-
H+
HB
A-
HB
H+
HB
HB
H+
H+
H+
HB
HB
A-
AA-
A-
H+
HB
A-
HB
HB
B-
HB
HB
(b)
A strong acid is nearly 100 % ionized, while a weak acid
is only slightly ionized.
14_322
Before dissociation
HA
After dissociation,
at equilibrium
H+ A–
(a)
HA
HA
H+ A–
(b)
Diagram a represents a strong acid, while b represents a weak
acid which remains mostly in the molecular form.
14_323
Relative
acid strength
Very
strong
Relative
conjugate
base strength
Very
weak
Strong
Weak
The relationship of
acid strength and
conjugate base
strength for acidbase reactions.
Weak
Strong
Very
weak
Very
strong
Arranging Species According
to Increasing Basic Strength
H2O, F-, Cl-, NO2-, & CNUse Table 14.2 on page 663.
Cl- is weakest since it is conjugate base of
strong acid and weaker than water. Use Ka
values to arrange the remaining bases.
Cl- < H2O < F- < NO2- < CN-
Water as an Acid and a Base
Water is amphoteric (it can behave either as
an acid or a base).
H2O + H2O  H3O+ + OH
acid 1 base 2
conj
acid 2
conj
base 1
Kw = 1  1014 M2 at 25°C
Ion product Constant, Kw
Kw is called the ion-product constant or
dissociation constant.
neutral solution [H+] = [OH-] = 1.0 x 10 -7 M
acidic solution [H+] > [OH-] [H+] > 1.0 x 10-7 M
basic solution [H+] < [OH-] [OH-] > 1.0 x 10-7 M
No matter what the concentration of H+ or OH- in an
aqueous solution, the product, Kw, will remain
the same.
[H+] & [OH-] Calculations
Calculate the [H+] for a 1.0 x 10-5 M OH-.
Kw = [H+][OH-]
[H+] = Kw/[OH-]
[H+] = 1.0 x 10-14 M2/1.0 x 10-5 M
[H+] = 1.0 x 10-9 M
[H+] & [OH-] Calculations
Continued
Calculate the [OH-] for a 10.0 M H+.
Kw = [H+][OH-]
[OH-] = Kw/[H+]
[OH-] = 1.0 x 10-14 M2/10.0 M
[OH-] = 1.0 x 10-15 M
Kw & H
At 60oC, the value of Kw is 1 x 10-13 for the
dissociation of water:
2 H2O(l) <---> H3O+(aq) + OH-(aq)
Is this reaction exothermic or endothermic?
Endothermic -- Kw increased with
temperature.
The pH Scale
pH = log[H+]
pH in water usually ranges from 0 to 14.
Kw = 1.00  1014 = [H+] [OH]
pKw = 14.00 = pH + pOH
As pH rises, pOH falls (sum = 14.00).
pH &
+
[H ]
pH = 0
pH = 7
pH = 14
1x 10-14
1 x 10-7
1 x 100
OH -
+
O
H3
OHH3O+
OH
H3O+
1 x 100
1 x 10-7
1 x 10-14
Logarithms
-log 1.00 x 10-7 = 7.000
7.000
characteristic
mantissa
The number of significant digits in 1.00 x 10-7
is three, therefore, the log has three
decimal places. The mantissa represents
the log of 1.00 and the characteristic
represents the exponent 7.
14_324
[H+] pH
10–14 14
–13
10
Basic
1 M NaOH
13
10–12 12
10–11 11
Ammonia
(Household
cleaner)
10–10 10
pH scale and pH values for
common substances. A pH of
1 is 100 times more acidic
than a pH of 3.
10–9
9
10–8
8
–7
7
10–6
6
10–5
5
10–4
4
10–3
3
Acidic 10–2
2
10–1
1
1
0
Neutral 10
Blood
Pure water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
pH Calculations
What is the pOH, [H+], & [OH-] for human
blood with a pH of 7.41?
pH + pOH = 14.00
pOH = 14.00 - pH
pOH = 14.00 - 7.41
pOH = 6.59
pH Calculations
Continued
What is the pOH, [H+], & [OH-]
for human blood with a pH of
7.41?
pH = - log [H+]
[H+] = antilog (-pH)
[H+] = antilog (-7.41)
[H+] = 3.9 x 10-8 M
Note: The number of
significant figures in the
antilog is equal to the number
of decimal places in the pH.
pH Calculations
Continued
What is the pOH, [H+], & [OH-]
for human blood with a pH of
7.41?
pOH = - log [OH-]
[OH-] = antilog (-pOH)
[OH-] = antilog (-6.59)
[OH-] = 2.6 x 10-7 M
Note: The number of
significant figures in the
antilog is equal to the number
of decimal places in the pOH.
pH of Strong Acid Solutions
Calculate the pH of a 0.10 M HNO3 solution.
Major species are: H+, NO3-, and H2O
Sources of H+ are from HNO3 and H2O -amount from water is insignificant.
[H+] = 0.10 M
Note: The number of
significant figures in
the [H+] is the same as
the decimal places in
the pH.
pH = - log [H+]
pH = - log [0.10]
pH = 1.00
pH & Significant Figures
log
# Significant Figures -------> # decimal places
<------inv log
pH = - log [H+]
[H+] = inv log (-pH)
[H+] = 1.0 x 10-5 M
pH = 5.00
Solving Weak Acid Equilibrium
Problems
-
List major species in solution.
-
Choose species that can produce H+ and write
reactions.
-
Based on K values, decide on dominant
equilibrium.
-
Write equilibrium expression for dominant
equilibrium.
-
List initial concentrations in dominant
equilibrium.
Solving Weak Acid Equilibrium
Problems (continued)
-
Define change at equilibrium (as “x”).
-
Write equilibrium concentrations in terms of x.
-
Substitute equilibrium concentrations into
equilibrium expression.
-
Solve for x the “easy way.” x can be neglected
when concentration is 2 powers of 10 (100x)
greater than Ka or Kb.
-
Verify assumptions using 5% rule.
-
Calculate [H+] and pH.
pH of Weak Acid Solutions
Calculate the pH of a 0.100 M HOCl solution.
Major species: HOCl and HOH
Ka HOCl = 3.5 x 10-8 & Ka HOH = 1.0 x 10-14
 HOCl will be only significant source of [H+].
Ka = 3.5 x 10-8 = [H+][OCl-]/[HOCl]
pH of Weak Acid Solutions
Continued
ICE
[HOCl]
[OCl-]
[H+]
Initial (mol/L)
0.100
0
0
Change (mol/L)
-x
+x
+x
0+x
0+x
Equil. (mol/L)
0.100 - x
pH of Weak Acid Solutions
Continued
Ka = 3.5 x 10-8 = [H+][OCl-]/[HOCl]
3.5 x 10-8 = [x][x]/[0.100 - x]
Ka is more than 100 x smaller than concentration,
x can be neglected in the denominator.
Ka = 3.5 x 10-8 = [x][x]/[0.100]
x2 = 3.5 x 10-9
x = 5.9 x 10-5 M
pH of Weak Acid Solutions
Continued
Approximation check:
% dissociation = (x/[HA]o) (100%)
% dissociation = (x/[HOCl]o) (100%)
% dissociation = (5.9 x 10-5/0.100)(100%)
% dissociation = 0.059 %
This is much less than 5 % and therefore the
approximation was valid.
Percent Dissociation
(Ionization)
amount dissociated( M )
% dissociation 
(100%)
initial concentration( M )
The percent dissociation calculation is exactly the same as the
one to check the 5 % approximation. See Sample Exercise 14.10
on pages 678 and 679.
% Dissociation Calculations
In a 0.100 M lactic acid solution (HC3H5O3),
lactic acid is 3.7 % dissociated. Calculate
the Ka for this acid.
Major species: HC3H5O3 & HOH
HC3H5O3(aq) <---> H+(aq)+ C3H5O3-(aq)
Ka = [H+][C3H5O3-]/ [HC3H5O3]
% Dissociation Calculations
Continued
ICE
[HC3H5O3]
Initial (M)
0.10
[C3H5O3-]
0
[H+]
0
Change (M) - 3.7 x 10-3 + 3.7 x 10-3 + 3.7 x 10-3
Equil. (M)
0.10
+ 3.7 x 10-3 + 3.7 x 10-3
% Dissociation Calculations
Continued
Ka = [H+][C3H5O3-]/ [HC3H5O3]
Ka = [3.7 x 10-3]2/ [0.10]
Ka = 1.4 x 10-4
14_325
More concentrated
More dilute
Acid concentration
Percent dissociation
H+ concentration
The effect of dilution on the % dissociation and [H+] of a
weak acid solution.
Bases
Bases are often called alkalis because they
often contain alkali or alkaline earth metals.
“Strong” and “weak” are used in the same
sense for bases as for acids.
strong = complete dissociation (hydroxide ion
supplied to solution)
NaOH(s)  Na+(aq) + OH(aq)
Bases
(continued)
weak = very little dissociation (or reaction
with water)
H3CNH2(aq) + H2O(l)  H3CNH3+(aq) + OH(aq)
See Table 14.3 on page 685 for Kb values of
common bases.
Kb calculations are identical to Ka calculations.
Polyprotic Acids
. . . can furnish more than one proton (H+) to
the solution.
H 2CO3  H   HCO3
( Ka1 )
HCO3  H   CO32 
( Ka 2 )
See Table 14.4 on page 689 for Ka values for common
polyprotic acids.
Know Sample Exercises 14.15 & 14.16 on pages 689-692.
Acid-Base Properties of Salts
Cation
neutral
neutral
Acidic
or Basic
neutral
basic
Anion
neutral
conj base of
weak acid
conj acid of
neutral
acidic
weak base
conj acid of conj base of depends on
weak base weak acid
Ka & Kb
values
Example
NaCl
NaF
NH4Cl
Al2(SO4)3
Structure and Acid-Base
Properties
Two factors for acidity in binary compounds:
-
Bond Polarity (high is good)
-
Bond Strength (low is good)
14_326
Cl
O
H
Electron density
O
Cl
O
H
Electron density
O
Cl
The effect of the
number of attached
oxygen on the H-O
bond
in a series of
chlorine oxyacids.
O
H
O
Electron density
O
O
Cl
O
O
H
Electron density
Oxides
Acidic Oxides (Acid Anhydrides):
-
OX bond is strong and covalent.
SO2, NO2, CrO3
Basic Oxides (Basic Anhydrides):
-
OX bond is ionic.
K2O, CaO
Lewis Acids and Bases
Lewis Acid: electron pair acceptor
Lewis Base: electron pair donor
Al3+ + 6 O
H
H
Al
H
3+
O
H
6