CHAPTER 15: ACIDS AND BASES
Part One: Acid-Base Concepts
A. Properties of Aqueous Solutions of Acids.
1.
Sour taste. (Examples: vinegar = acetic acid; lemons - citric acid)
2.
Change the colors of many indicators: (pH-sensitive dyes)
3.
acid

→
a.
blue litmus
b.
bromothymol blue
red
acid

→
yellow
React with active metals liberating H2(g). (HNO3 is exception, an oxidizing acid)
Ca(s) + 2 HI(aq) → CaI2(aq) + H2(g)
2 H+(aq) + Zn(s) + 2 HNO3(aq) → Zn2+(aq) + 2 NO2(g) + 2 H2O(l)
4.
Neutralize metal oxides and hydroxides to form salts and water:
2 HCl + MgO → MgCl2 + H2O
2 HCl + Ba(OH)2 → BaCl2 + H2O
5.
React with salts of weaker acids to form the weaker acid and a new salt:
HCl + NaF → NaCl
sodium
fluoride
6.
B.
+ HF
hydrofluoric
acid
Wholly or partially ionize in water and conduct electric current.
Properties of Aqueous Solutions of Bases.
1.
Bitter taste. (e.g. the bitter taste of soaps and many pharmaceuticals is because of their
alkalinity)
2.
Slippery feeling. (soaps are mildly alkaline)
3.
Change the color of many indicators:
Chapter 15
Page 1
4.
base

→
a.
red litmus
blue
b.
bromothymol blue yellow
base

→
blue
Neutralize protonic acids forming salts and water.
• protonic acids - containing H+ also called “protic”
5.
C.
Ionize in water and conduct current.
Arrhenius Theory. (1884) (Section 15.1)
1.
An acid = substance that contains hydrogen and produces H+ in aqueous solution.
• protonic acids - containing H+ also called “protic” acids
2.
A base = substance that contains OH group and produces OH- ions in aqueous
solution.
3.
Neutralization is the combination of H+ and OH- to form H2O:
H+(aq) + OH-(aq) → H2O(l)
4.
Chapter 15
Theory works well for protic acids and hydroxy bases, but not all acids/bases. For
example, does not explain why NH3 is a base, nor why BF3 is an acid.
Page 2
D. Hydronium Ions.
+
1.
H = a proton, and as such cannot exist as separate entity in aqueous solution.
2.
H becomes “hydrated,” attaching to one or more H2O molecules.
+
Figure 15.2
E.
The Bronsted-Lowry Theory. (1923) (Section 15.2)
1.
More general and complete than Arrhenius Theory.
2.
Acid = a proton donor. (H+ = proton)
3.
Base = a proton acceptor.
4.
Acid/Base reaction = proton transfer reaction, transfer of H+ from an acid to a base.
Figure 15.3
5.
Chapter 15
Bronsted Theory thus explains why NH3 is a base, even though it has no OH- group.
It accepts protons.
Page 3
6.
Ionization of a strong acid HA in water is a proton transfer to H2O, in which H2O acts
as the base:
H3O+(aq) + A-(aq)
new
new
acid
base
HA(aq) + H2O(l)
acid
base
Example - 1 M HCl in water exists entirely as:
+
[H3O ] = 1 M
-
[Cl ] = 1 M
no un-ionized HCl molecules actually present
7.
Weak acid HAc:
HAc(aq) + H2O(l)
acid
base
H3O+(aq) + Ac-(aq)
Example - 1 M acetic acid HAc (abbreviation for CH3COOH) exists as:
[HAc] ≈ 1 M
[H3O+] ≈ 10-3 M
[Ac ] ≈ 10-3 M
• most HAc molecules remain un-ionized.
8.
Concept of conjugate acid/base pairs.
a.
When HA acts as an acid (ionizes) a base A- is produced.
A- is the conjugate base of the acid HA.
b.
A strong acid HA produces very weak conjugate base A-.
Example:
Cl- is conj. base of strong acid HCl
Cl- has very little tendency to accept a proton (H+) back and
re-form HCl.
Chapter 15
Page 4
c.
A weak acid like acetic HAc produces a relatively stronger conj. base Ac(acetate ion).
-
+
Ac has stronger tendency to accept proton H back to re-form HAc.
d.
When base B acts as a base (accepts a proton), its conjugate acid is formed, BH+.
e.
Identify conj. acid/base pairs in the rxn:
HCN(aq) +
acid
H 2O
base
+
-
H3O (aq) + CN (aq)
conj. acid
conj. base
of base
of acid
H2O
HCN
-
Is CN a relatively strong or weak base?
Relatively strong, because HCN is a very weak acid.
The stronger the acid, the weaker its conjugate base, and vice versa.
9.
Work the following examples:
conjugate acid of NH3
conjugate base of H2O
conjugate base of H2SO4
conjugate acid of OHconjugate acid of H2O
conjugate base of NH4+
conjugate acid of F-
10. Amphiprotism of Water.
Chapter 15
a.
Amphi = “both kinds.”
b.
Amphiprotism is when a substance can act either as an acid or a base.
Page 5
F.
can donate H+ to become OHcan accept H+ to become H3O+.
c.
H2O is amphiprotic:
d.
HSO4- is also amphiprotic:
can be acid to become SO42-,
can be base to become H2SO4.
Amphoterism.
1.
Amphoterism = more general term describing the ability of a substance to react either
as an acid or base.
2.
Most common examples – insoluble metal hydroxides.
a.
Here reacting as a base:
Al(OH)3(s) + H+(aq) → Al(OH)2+(aq) + H2O
b.
Here as an acid:
Al(OH)3(s) + OH-(aq) → Al(OH)4-(aq)
G. The Lewis Theory. (1923) (Section 15.3)
1.
An acid is any species that can accept a share in an electron pair. That is, an electron
pair acceptor.
2.
A base is an electron pair donor.
3.
Let’s see how it works:
4.
Lewis theory encompasses all of Bronsted-Lowry theory plus other substances. So it
is the most inclusive acid/base theory.
Chapter 15
Page 6
Example: BCl3 is an acid, but how?
According to Bronsted-Lowry, it would need to donate a proton. Yet it can neutralize
NH3 by accepting e- pair.
Part Two: Acid and Base Strengths
A. Strength of Acids. (Section 15.4)
1.
Strength = ease of ionization; HX(aq)
2.
Depends on:
3.
H+(aq) + X-(aq)
a.
the ease of breaking H-X bond.
b.
stability of resulting ions in solution (H+ and X-).
Group VIIA Binary Acids-Hydrogen Halides relative acid strengths.
HI > HBr > HCl >> HF
strong acids,
weak acid
completely ionize
in aqueous solution
These are monoprotic acids = give one H+.
4.
Trend is due to strengths of H-X bond:
bond strength HF >> HCl > HBr > HI
5.
Why HF so much weaker acid?
a.
Large H-F bond strength (due to smaller size of F, relative to Cl, Br and I).
b.
F- ion is very small, causes very strong orientation of H2O molecules in solution,
thus decreasing entropy.
Therefore, F-(aq) not as stable as Cl-(aq)...
Chapter 15
Page 7
6.
7.
Leveling effect of water:
a.
all the strong acids HCl, HBr, HI in water fully ionize to produce H+(aq). (i.e.
+
H3O )
b.
thus they all appear to have the same strength in water.
c.
leveling effect ⇒ H3O+ ion is the strongest acid that can exist in aqueous
solution.
d.
the strongest acid HClO4 thus appears to be no stronger than the other strong
+
acids, all producing equal numbers of H3O ions.
e.
similar leveling effect among bases ⇒ OH is the strongest base that can exist in
water.
f.
any stronger base simply produces OH ions.
-
-
Group VIA Binary Acids vary in acid strength the same way.
H2Te > H2Se > H2S >> H2O
All are weak though diprotic.
8.
Chapter 15
Table of Acid Strengths:
Page 8
B.
Ternary Acids. (oxoacids, 3 elements)
1.
Common examples:
HClO4
HNO3
H2SO4
H3PO4
2.
perchloric
nitric
sulfuric
phosphoric
Examples of their structure:
nitric acid
3.
strong acid
strong acid
strong 1st ionization
weak acid
sulfuric acid
phosphoric acid
Trends in acid strength of ternary acids:
a.
Among those containing same central atom, strength increase with increasing
number of oxygens (or increasing oxidation # of central atom).
H2SO3
weak
<
H2SO4
strong 1st ioniz.
HNO2
weak
<
HNO3
strong
HClO
weak
<
HClO2
weak
<
HClO3
strong
<
HClO4
strong
Figure 15.7
Chapter 15
Page 9
b.
Trends among same group members having same ox# of central atom.
(+5)
HNO3
(+6)
H2SO4
(+7)
HClO4
C.
>
(+5)
H3PO4
>
(+6)
H2SeO4
>
(+7)
HBrO4
Common Strong Bases.
1.
Strong soluble bases (metal hydroxides):
LiOH
NaOH
KOH
2.
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
Mg(OH)2 is not very soluble, but the small amount that does dissolve completely
dissociates giving OH- ions. Think of Mg(OH)2 as a strong base that is insoluble.
D. Acid/Base Reactions. (neutralization)
1.
Strong acid with strong base:
formal equation:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
total ionic equation:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
net ionic equation:
H+(aq) + OH-(aq) → H2O(l)
Therefore, this is the net ionic equation for ANY strong acid/strong base reaction.
2.
Weak acid with strong base:
formal equation:
HAc(aq) + NaOH(aq) → NaAc(aq) + H2O(l)
Chapter 15
Page 10
total ionic equation:
HAc(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Ac-(aq) + H2O(l)
net ionic equation:
HAc(aq) + OH-(aq) → Ac-(aq) + H2O(l)
3.
Try writing the formal, total ionic, and net ionic equations for neutralization of
perchloric acid with ammonia in aqueous solution.
4.
Do the same thing for neutralizing phosphoric acid with an equal number of moles of
NaOH. Hint: an acidic salt is formed.
Chapter 15
Page 11
5.
Do the same for neutralizing aluminum hydroxide with an equal number of moles of
HCl. Hint: a basic salt is formed.
Part Three: Self-Ionization of Water and pH
A. Auto-ionization (or self-ionization) of water.
1.
Rxn:
H3O+(aq) + OH-(aq)
H2O(l) + H2O(l)
2.
Also written shorthand as:
H+(aq) + OH-(aq)
H2O(l)
3.
Equilib. constant for this reaction:
[H O ][OH ]
=
+
Kc
−
3
[H 2O]
2
[H2O] in all aqueous solutions = 55.5 M (a constant)
€
So group it with Kc:
Kc[H2O]2 = [H3O+][OH-]
4.
Kw = [H3O+][OH-] = 1.0 × 10-14 at 25° C
5.
Kw = ion product constant for water.
6.
In pure water then:
[H3O+] = [OH-] = 1.0 × 10-7 M
Chapter 15
Page 12
7.
In aqueous solutions containing acids/bases/some salts:
[H3O+] ≠ [OH-]
but still [H3O+][OH-] = Kw = 1.0 × 10-14 M
B.
Strong Electrolyte Solutions.
1.
Complete ionization makes calculation of ionic concentrations simple.
2.
Example: Calculate [H+] and [Cl-] conc. in 0.100 M aqueous HCl solution.
+
-
HCl(aq) → H (aq) + Cl (aq)
[H+(aq)] = 0.100 M
[Cl-(aq)] = 0.100 M
[HCl(aq)] = 0.0 M
3.
Example: Calculate concentrations of predominant ionic species in 0.003 M aqueous
Ca(OH)2 solution.
Ca(OH)2 → Ca2+(aq) + 2 OH-(aq)
[Ca2+(aq)] = 0.003 M
[OH-(aq)] = 0.006 M
Example: Calculate [OH-] in 0.100 M HCl(aq).
4.
Kw
; Since [H3O+] = 0.100 M
+
H 3O
[OH-] =
[
]
1.0 ×10 −14
[OH ] =
= 1.0 × 10-13 M (very small)
0.100 M
€
-
What is [H+] in 3.0 M Ca(OH)2?
5.
€
[H+] =
Kw
; Since [OH-] = 6.0 M
−
OH
[
]
1.0 ×10 −14
[H ] =
= 0.166 × 10-14 M (very small)
6.0 M
€
+
Chapter 15
€
Page 13
6.
In acid solutions:
[H+] > 1.0 × 10-7 M
[OH-] < 1.0 × 10-7 M
7.
Reverse for basic solutions.
8.
Neutral solutions or pure water:
[H+] = [OH-] = 1.0 × 10-7 M
C.
pH and pOH Scales. (Section 15.8)
1.
Acidity expressed often as pH.
pH = -log[H3O+]
-
pOH = -log[OH ]
2.
Calculate pH of 0.100 M HNO3.
[H+] = 0.100 M
pH = -log (0.100) = - (-1) = 1
3.
Chapter 15
The Scale:
Page 14
Figure 15.8
4.
Note that always:
pH + pOH = 14
D. pH Indicators.
1.
Indicator is a weak acid itself. Symbolize it HIn:
H3O+ + In-
HIn + H2O
acid form
of indicator
(color 1)
base form
of indicator
(color 2)
pH where it changes ≈ pKa of HIn.
Figure 15.10
Chapter 15
Page 15
NOTES:
Chapter 15
Page 16