CHAPTER 9 Acids, Bases, and
Salts
9.1 ARRHENIUS THEORY
• Understand the Arrhenius and Brønsted theories
of acid-base behavior
• Use equations for neutralization reactions
ACID
Anything that produces hydrogen
ions when dissolved in water
HCl → H+ + Cl-
BASE
Anything that produces hydroxide
ions when dissolved in water
NaOH → Na+ + OH-
• Describe relative strengths of acids and bases
• Explain the hydrolysis of salts
• Calculate pH
• Understand the action of buffers and calculate
buffer pH using the Henderson-Hasselbalch
equation
Acid-Base Properties of Water
9.2 BRØNSTED-LOWRY THEORY
Water is amphiprotic, it can act either as an acid
or a base:
THIS ONE COUNTS:
ACID
H+
Proton donor (adds
to solvent)
HCl + H2O → H3O+ + ClH+ + H2O: → HO2H+ (H3O+)
HCN(aq) + H2O(l)
base1
acid1
H2O(l) + NH3(aq)
acid1
base1
H+ addition
BASE
Proton acceptor (removes H+ from
solvent)
NH3 + H2O → NH4+ + OH:NH3 + HOH → HNH3+ + OH-
↔
↔
H3O+ + CN- (aq)
acid2
base2
NH4+(aq) + OH- (aq)
acid2
base2
CONJUGATE PAIR
CONJUGATE PAIR
H+ transfer
9.4 SELF-IONIZATION OF WATER
• In pure water at 25°C,
[H+] = 1.0×10-7 M
[OH-] = 1.0×10-7 M
Water undergoes autoionization (self-ionization):
its molecules react with one another to form ions:
• A solution is ACIDIC if
[H+] > [OH-]
This means:
[H+] > 1.0×10-7 M
[OH-] < 1.0×10-7 M
H2O + H2O ↔ H3O+ + OHacid base
acid base
[H3O+][OH-]
Keq = ——————
[H2O]2
• A solution is BASIC or
[H+] <
This means:
[H+] <
[OH-]
or, since H3O+ ⇔ H+(aq):
H2O ↔ H+(aq) + OH-(aq)
ALKALINE if
[OH-]
1.0×10-7 M
> 1.0×10-7 M
• A solution is NEUTRAL if
[H+] = [OH-] = 1.0×10-7 M
+
KKw == [H
[H+][OH
][OH-]]
w
-14
== 1.0×10
at 25°C
25°C
1.0×10-14 at
Page 9-1
9.9 ACID AND BASE STRENGTH
Dissociation
Complete dissociation in water
STRONG
HX (X = Cl, Br, I), HNO3, HClO4,
ACIDS
H2SO4
WEAK
ACIDS
STRONG ACIDS AND BASES
HCl + H2O
NaOH
Partial dissociation in water
HF and most other acids
WEAK ACIDS AND BASES
OH-
STRONG Dissociation in water to yield
BASES
Metal hydroxides: NaOH, Ca(OH)2
WEAK
BASES
H3O+ + ClNa+ + OH-
CH3COOH + H2O
NH3 + H2O
H3O+ + CH3COONH4+ + OH-
Weak H+ acceptors
NH3, C6H5NH2
Acid Dissociation Constant Ka
ACID
The weak acid dissociation
HCl
HA + H2O ↔ H3O+ + A-
+
[A
][H33O
O+]]
[A-][H
KKa == —————
————— ≡≡
a
[HA]
[HA]
Ka
HCl → H+ + Cl-
∞
H2SO4 → H+ + HSO4-
H2SO4
has the equilibrium constant Ka (< 1):
DISSOCIATION
HSO4- → H+ + SO42-
HC2H3O2 HC2H3O2 → H+ + C2H3O2-
+
[A
[A-][H
][H+]]
————
————
[HA]
[HA]
KKa is
a measure of the acid strength
a is a measure of the acid strength
∞
1.2×10-2
1.8×10-5
H2CO3*
H2CO3 → H+ + HCO3HCO3- → H+ + CO32-
4.3×10-7
5.6×10-11
C6H5OH
C6H5OH → H+ + C6H5O-
1.3×10-10
*H2CO3 ↔ H2O + CO2↑
9.6-9.8 PROPERTIES of ACIDS and
BASES
• Metal hydroxides and oxides react with
(NEUTRALIZE) acids to produce the salt of
the acid and H2O:
• Many metals react with acids to produce H2
and the SALT of the acid:
NaOH + HBr → NaBr + H2O
Cu2O + 2HBr → CuBr2 + H2O
Zn + 2HCl → ZnCl2 + H2↑
• Weak bases react with (NEUTRALIZE) acids
to produce the salt of the acid:
• Metal oxides react with acids to produce the
salt of the acid (and H2O if acid other than
water), since O2- is a very strong base:
NH3 + HCl → NH4Cl
CaO + H2O → Ca(OH)2
CaO + 2HCl → CaCl2 + H2O
ACID
ACID ++ BASE
BASE →
→ SALT
SALT
Page 9-2
9.10 Acid-Base Analysis
SALTS
Examples:
Recovered by evaporation
from aqueous solution
(where their ions were
solvated), many salts are
isolated as HYDRATES:
the solid salt contains
water molecules as part of
the crystalline structure
Requires known [OH-] (or
[H+]) as standard titrant
CuSO4•5H2O
actually consists of
Cu(H2O)42+ and SO4(H2O)2ions
Vf
H +)
(or
is added
Enough
to react exactly — no extra STANDARD SOLN
Vstd = V0 - Vf
OH- (or H+) — with the
unknown acid or base
OH-
Application:
CaCO3·6H2O + heat →
limestone
CaCO3 + 6H2O
Hydrates may lose all or
part of this WATER OF
HYDRATION when they
are heated
Vx
9.5 THE pH CONCEPT
If
• We measure and use acids and bases over a wide
(!) range of concentrations.
• A logarithmic scale (pH) is a convenient way to
express these very large and very small numbers.
then
so
pH
pH == -log[H
-log[H++]]
KKww ==
V0
Titration of [H+] or [OH- ]
FeCl3•6H2O
actually consists of
Fe(H2O)63+ and Cl- ions
[H
[H++][OH
][OH--]]
==
X = 10n
0.010 = 10−2
log10X = n
log10(0.010) = −2.00
X =
0.010 = 10log(0.010)
10logX
“Logarithmic” ruler:
-14
10
10-14
1
pK
pKww == pH
pH ++ pOH
pOH == 14.00
14.00
2
3
4
5
6
10−4
10−5
10−6
Linear translation of “pH” ruler:
• [H+] = 10-pH
• If pH1 = pH2 - 1, then [H+]1 = 10[H+]2
1
2−1
6−1
3−1
10−1
10−2
10−3
pH Scale
Examples
A log based scale used to keep track of the large
change important to acids and bases
14
10-14 M
VERY
BASIC
7
0
100 M
VERY
ACIDIC
• When you add an ACID, the pH gets LOWER
• When you add a BASE, the pH gets HIGHER
Page 9-3
SUBSTANCE
pH
1M HCl
0.0
Lemon juice
2.3
Coffee
5.0
Pure H2O
7.0
Blood
7.35-7.45
Milk of magnesia
10.5
Household ammonia
11.7
1M NaOH
14.0
9.12 HYDROLYSIS REACTIONS of
SALTS
• When a salt dissolves in water, the cation(s) and
anion(s) are hydrated
CATION ANION SOLUTION
of
of
pH
• The ANION OF A WEAK ACID will hydrolyze
(react with water) to RELEASE HYDROXIDE:
A- + H2O ↔ HA + OHC2H3O2- + H2O ↔ HC2H3O2 + OH• The CATION OF A WEAK BASE will hydrolyze to
PRODUCE HYDROGEN ION:
BH+ + H2O ↔ B: + H3O+
NH4+ + H2O ↔ NH3 + H3O+
RESULT:
STRONG
BASE
ACID
> 7.0
NaC2H3O2
BASE
STRONG
ACID
< 7.0
NH4Cl
STRONG
BASE
STRONG
ACID
IF EQUAL
STRENGTH
BASE
ACID
IF EQUAL
STRENGTH
Many salt solutions are either acidic or alkaline
(basic)
pH Values:
0.100 M solutions
NaOH
13.00
NaOH
13.00
NaCl
7.00
NH3
11.12
HCl
1.00
Example of Type 1 (Acetic acid + Na acetate):
H+ + A- ↔ HA + H2O
← removes H+
OH- + HA ↔ A- + H2O
← removes OHExample of Type 2 (Ammonia + NH4+ chloride):
H+ + NH3 ↔ NH4+ + H2O
← removes H+
OH- + NH4+ ↔ NH3 + H2O
← removes OH-
Henderson-Hasselbalch Equation
• O2 transported by hemoglobin in erythrocytes
• CO2 transported in both plasma and
erythrocytes
• CO2 converted to H2CO3 by enzymatic action
and to HCO3- by hydrolysis:
[A-][H+]
Ka = ————
[HA]
Ka[A-]
[H+] = ————
[A-]
H2CO3 ↔ H+ + HCO3-
H+
High pH:
H2CO3 +
OH-
NH4C2H3O2
Two types:
1. Weak acid (HA) and its salt (source of A-)
2. Weak base (B:) and its salt (BH+A-)
Buffers and Blood pH
H2CO3
~7.0
KBr
Solutions
Solutions that
that resist
resist changing
changing pH
pH when
when small
small
amounts
amounts of
of acid
acid or
or base
base are
are added
added
NH4Cl
5.12
Low pH:
~7.0
9.13 BUFFERS
HC2H3O2
2.88
NaC2H3O2
8.88
HCl
1.00
EXAMPLE
[HA]
log[H+] = logKa + log ———
[A-]
+ HCO3–
HCO3– + H2O
[A
[A-]] ← Salt (ex: Na acetate)
pH
log ———
———
pH == pK
pKaa ++ log
[HA]
[HA] ← Weak acid (ex: acetic acid)
Page 9-4