CHAPTER 15: ACIDS AND BASES
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)
acid a. blue litmus → red
acid b. bromothymol blue → yellow
3. React with active metals liberating H2(g). (HNO3 is exception, an oxidizing acid)
Ca(s) + 2 HI(aq) → CaI2(aq) + H2(g)
+ 2+ 2 H (aq) + Zn(s) + 2 HNO3(aq) → Zn (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 + HF sodium hydrofluoric fluoride acid
6. Wholly or partially ionize in water and conduct electric current.
B. 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 base a. red litmus → blue
base b. bromothymol blue yellow → blue
4. Neutralize protonic acids forming salts and water.
• protonic acids - containing H+ also called “protic”
5. Ionize in water and conduct current.
C. 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. 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.
Chapter 15 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. Bronsted Theory thus explains why NH3 is a base, even though it has no OH group. It accepts protons.
Chapter 15 Page 3 6. Ionization of a strong acid HA in water is a proton transfer to H2O, in which H2O acts as the base:
+ - HA(aq) + H2O(l) H3O (aq) + A (aq) acid base new new 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) H3O (aq) + Ac (aq) acid base
Example - 1 M acetic acid HAc (abbreviation for CH3COOH) exists as:
[HAc] ≈ 1 M + -3 [H3O ] ≈ 10 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) + H2O H3O (aq) + CN (aq) acid base 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 OH-
conjugate acid of H2O
+ conjugate base of NH4
conjugate acid of F-
10. Amphiprotism of Water.
a. Amphi = “both kinds.”
b. Amphiprotism is when a substance can act either as an acid or a base.
Chapter 15 Page 5 + - c. H2O is amphiprotic: can donate H to become OH + + can accept H to become H3O .
- 2- d. HSO4 is also amphiprotic: can be acid to become SO4 ,
can be base to become H2SO4.
F. 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) H+(aq) + X-(aq)
2. Depends on:
a. the ease of breaking H-X bond.
b. stability of resulting ions in solution (H+ and X-).
3. 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. 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.
7. Group VIA Binary Acids vary in acid strength the same way.
H2Te > H2Se > H2S >> H2O
All are weak though diprotic.
8. Table of Acid Strengths:
Chapter 15 Page 8 B. Ternary Acids. (oxoacids, 3 elements)
1. Common examples:
HClO4 perchloric strong acid HNO3 nitric strong acid H2SO4 sulfuric strong 1st ionization H3PO4 phosphoric weak acid
2. Examples of their structure:
nitric acid sulfuric acid phosphoric acid
3. 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 < H2SO4 weak strong 1st ioniz.
HNO2 < HNO3 weak strong
HClO < HClO2 < HClO3 < HClO4 weak weak strong strong
Figure 15.7
Chapter 15 Page 9 b. Trends among same group members having same ox# of central atom.
(+5) (+5)
HNO3 > H3PO4
(+6) (+6)
H2SO4 > H2SeO4
(+7) (+7)
HClO4 > HBrO4
C. Common Strong Bases.
1. Strong soluble bases (metal hydroxides):
LiOH RbOH Sr(OH)2 NaOH CsOH Ba(OH)2 KOH Ca(OH)2
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:
+ - H2O(l) + H2O(l) H3O (aq) + OH (aq)
2. Also written shorthand as:
+ - H2O(l) H (aq) + OH (aq)
3. Equilib. constant for this reaction:
+ − [H3O ][OH ] Kc = 2 [H2O]
[H2O] in all aqueous solutions = 55.5 M (a constant)
€ 2 + - So group it with Kc: Kc[H2O] = [H3O ][OH ]
4. + - -14 Kw = [H3O ][OH ] = 1.0 × 10 at 25° C
5. Kw = ion product constant for water.
6. In pure water then:
+ - -7 [H3O ] = [OH ] = 1.0 × 10 M
Chapter 15 Page 12 7. In aqueous solutions containing acids/bases/some salts:
+ - [H3O ] ≠ [OH ]
+ - -14 but still [H3O ][OH ] = Kw = 1.0 × 10 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.
2+ - Ca(OH)2 → Ca (aq) + 2 OH (aq)
[Ca2+(aq)] = 0.003 M [OH-(aq)] = 0.006 M
4. Example: Calculate [OH-] in 0.100 M HCl(aq).
K [OH-] = w ; Since [H O+] = 0.100 M + 3 [H3O ]
1.0 ×10−14 [OH-] = = 1.0 × 10-13 M (very small) 0.100 M € + 5. What is [H ] in 3.0 M Ca(OH)2?
€ K [H+] = w ; 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. The Scale:
Chapter 15 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:
+ - HIn + H2O H3O + In
acid form base form of indicator of indicator (color 1) (color 2)
pH where it changes ≈ pKa of HIn.
Figure 15.10
Chapter 15 Page 15
NOTES:
Chapter 15 Page 16