Ch 16: Acid-Base Equilibria
Brown, LeMay Ch 16AP Chemistry
2
16.1: Acids and Bases
* Defined by Svante Arrhenius in 1880’s Arrhenius acids: produce protons;
increase [H+]HCl (aq) → H+ (aq) + Cl- (aq)
Arrhenius bases: produce hydroxides; increase [OH-]
NaOH (aq) → Na+ (aq) + OH- (aq)or
NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
3
16.2: Dissociation of Water
Autoionization of water:H2O (l) ↔ H+ (aq) + OH- (aq)
][
]][[
2OH
OHHKc
M 55.6 g 18.0
mol 1
wKOHH -1410 x 0.1]][[
KW = ion-product constant for water
H3O+ (aq) or H+ (aq) = hydronium
L 1
g 1000 (l)] O[H2
4
16.3: The pH Scale
pH = -log [H+] = -log [H3O+]
or [H+] = 10-pH
[H+][OH-] = KW = 1.0x10-14
-log ([H+][OH-]) = -log KW
-log [H+] + -log[OH-] = -log (1.0x10-14)
pH + pOH = 14.00
pOH = -log [OH-] or [OH-] = 10-pOH
pX = -log [X]
5
16.3: The pH Scale
If [H+]<[OH-], then [H+]<1.0x10-7
Ex: pH = -log[1.0x10-10] = 10.00 (basic)
If [H+] = [OH-]Since [H+][OH-] = 1.0x10-14
[H+] = [OH-] = 1.0x10-7
pH = -log[1.0x10-7] = 7.00 (neutral)
If [H+]>[OH-], then [H+]>1.0x10-7
Ex: pH = -log [1.0x10-3] = 3.00 (acidic)
pH
14
7
0
pOH
0
7
14
6
Johannes Brønsted (Denmark)Thomas Lowry (England),
1923
Brønsted-Lowry acids: H+ donor Brønsted-Lowry bases: H+ acceptor
NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
Base Acid
16.4: Brønsted-Lowry Acids & Bases
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Amphoterism Amphoteric: capable of acting as either
an acid or baseH2O (l) ↔
Acting as an acid
Acting as a base
Al(OH)3 (aq) ↔
Al(OH)4- (aq) + H+(aq)
Al(OH)2+ (aq) + OH-(aq)
Al(OH)3 (aq) + H2O(l) ↔
OH- (aq) + H+(aq)
* Amphiprotic: can accept or donate a p+
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Conjugated Acid-Base Pairs
For acid “HA”:HA (aq) + H2O (l) ↔ A- (aq) + H3O+ (aq)
acid base conjugate base
For base “B”:
B (aq) + H2O (l) ↔ HB+ (aq) + OH- (aq)
base acid
conjugate acid
conjugate base
conjugate acid
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Relative Acid-Base Strengths The stronger an acid (the greater its
ability to donate p+), the weaker its conjugate base (the lesser its ability to accept p+).
The stronger a base, the weaker its conjugate acid.
In an acid-base equilibrium, the p+ is transferred from the strongest acid to the strongest base.HSO4
- + CO32- ↔ SO4
2- + HCO3-
Stronger acidStronger base
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16.5: Strong Acids and Bases
Strong acids and bases fully ionize in water (equilibrium is shifted “entirely” toward ions).
Strong acids:HI, HBr, HCl, HClO4, HClO3, H2SO4, HNO3
Ex: In 6M HCl solution, 0.004% exist as molecules
Strong bases:LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2
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16.6: Weak Acids
Weak acids partially ionize in water (equilibrium is somewhere between ions and molecules).
HA (aq) ↔ A- (aq) + H+ (aq)
eqa HA
AHK
][
]][[
Ka = acid-dissociation constant in water
Weak acids generally have Ka < 10-3
See Appendix D for full listing of Ka values
12
Ex: Calculate the pH of 2.0 M HCl solution (Ka≈106) Strong acid, completely dissociated HCl (aq) → H+ (aq) + Cl- (aq)
HCl (aq) H+ (aq) Cl- (aq)
Initial
Change
Final
2.0 M
- 2.0 M
0 M
0 M 0 M
+ 2.0 M + 2.0 M
2.0 M 2.0 M
So:
[HCl]initial = [H+]final = [Cl-]final = 2.0 M
pH = - log [H+] = - log [2.0] = -0.30
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Ex: Calculate pH of 2.0 M HF solution (Ka=7.2x10-4) Weak acid, partially dissociated HF (aq) ↔ H+ (aq) + F- (aq)
HF (aq) H+ (aq) F- (aq)Initial
Change
Equilibrium
2.0 M- x M
(2.0 – x) M
0 M 0 M+ x M + x Mx M x M
eq
aHF
FHK
][
]][[102.7 4-
x
x
0.2
2
xHF
xx
initial
][
))((
Using quadratic eq’n, 0 = x2 + 7.2 x 10-4x – 1.44 x 10-3
x = 3.7229 x 10-2 or – 3.8669 x 10-2 = [H+]
pH = - log [H+] = - log [3.7 x 10-2] = 1.43
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Or, since weak acids partially dissociate, assume that [HF]init >> [H+]eq
Then, [HF]init – [H+] ≈ [HF]init
0.20.2102.7
224 x
x
x
pH = - log [H+] = - log [3.8 x 10-2] = 1.42 General rule: if [H+] 5% of [HA], it is better to
use quadratic formula.
23104.1 x
][108.3104.1 23 Hx
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Percent Ionization of an Acid
100%[HA]
][HIonization %
init
eq
Ex: Calculate the % ionization of:
•2.0 M solution of HCl
•2.0 M solution of HF
0%01100%[2.0]
.0]2[Ionization %
1.9%100%[2.0]
].8x103[Ionization %
-2
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Polyprotic acids: have more than one H+ to “donate”Ex: H2SO3 (aq) ↔ HSO3
- (aq) + H+ (aq)
Ka1 = 1st acid-dissociation constant = 1.7 x 10-2
HSO3- (aq) ↔ SO3
2- (aq) + H+ (aq)
Ka2 = 2nd acid-dissociation constant = 6.4 x 10-8
Ka1>Ka2; 1st H+ dissociates more easily than
the 2nd.
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* Polyprotic Acids
Ascorbic acid (Vitamin C):
Citric acid:
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16.7: Weak Bases
Partially ionize in water.B (aq) + H2O (l) ↔ BH+ (aq) + OH- (aq)
eqb B
OHBHK
][
]][[
Kb = base-dissociation constant in water
In practice,
eqb B
OHBHK
][
]][[
initialB
x
][
2
xB
xx
initial
][
))((
where x = [OH-]
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16.8: Relationship between Ka and Kb
Weak base: NH3(aq) + H2O(l) ↔ NH4+(aq)+OH-
(aq)
Conjugate acid: NH4+(aq) ↔ NH3(aq) + H+
(aq)
][
]][[
4
3
NH
NHHKa
][
]][[
3
4
NH
OHNHKb
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NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
+ NH4+ (aq) ↔ NH3 (aq) + H+ (aq)
][
]][[
][
]][[
4
3
3
4
NH
HNH
NH
OHNHKK ab
H2O (l) ↔ H+ (aq) + OH-
(aq)And:
]][[ HOHTherefore: 14100.1 wab KKK
For a conjugate acid-base pair
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In general, when two reactions are added to give a 3rd, the equilibrium constant for the 3rd reaction equals the product of the equilibrium constants of the two added reactions.
Furthermore:
)100.1log(log)log( 14 wab KKK
00.14logloglog wab KKK
00.14 wab pKpKpK
14100.1 wab KKK
For a conjugate acid-base pair
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16.9: Salt Solutions as Acids & Bases
Hydrolysis: acid/base reaction of ion with water to produce H+ or OH-
Anion (A-) = a conjugate baseA- (aq) + H2O (l) ↔ HA (aq) + OH- (aq)
Cation (B+) = a conjugate acidB+ (aq) + H2O (l) ↔ BOH (aq) + H+ (aq)
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Predicting pH of Salt Solutions
Salt type Cation AnionHydrolyze
s to produce
pH
Consider the relative strengths of the acid and base from which the salt is derived:
Strong electrolyte
Ex: Ca(NO3)2
Ca2+
conjugate acid of strong base
Ca(OH)2
NO3-
conjugate base of strong
acid HNO3
Neither H+ nor OH- 7
Salt type Cation AnionHydrolyze
s to produce
pH
ClO- (aq) + H2O (l) ↔ HClO (aq) + OH- (aq)
eqb ClO
OHHClOK
][
]][[
xClO
xx
initial ][
))((
initiala
wb ClO
x
K
KK
][
2
Weak electrolyte
Ex: NaClO
Na+
conjugate acid of strong base, NaOH
ClO-
conjugate base of weak acid, HClO
OH-
> 7
where x = [OH-]
Salt type Cation AnionHydrolyze
s to produce
pH
NH4+ (aq) + H2O (l) ↔ NH3 (aq) + H3O+ (aq)
eq
aNH
HNHK
][
]][[
4
3
xNH
xx
initial ][
))((
4
initialb
wa
NH
x
K
KK
][ 4
2
Weak electrolyte
Ex: NH4Cl
NH4+
conjugate acid of weak
base, NH3
Cl-
conjugate base of strong
acid, HCl
H+ < 7
where x = [H+]
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16.10: Acid-Base Behavior & Chemical Structure Stronger acids, HA, have:
1. H with a higher +
2. Weaker H-A covalent bond (smaller bond enthalphy)
3. More stable conjugate bases A-
Stronger oxyacids, HxOz-Y, have:
1. Central nonmetal “Y” with higher electronegativity
2. More O atoms
Ex: Rank these in order from strongest to weakest: HClO, HClO2, HCl, HBr
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16.11: Lewis Acids & Bases Lewis acid: “e- pair acceptor”
Brønsted-Lowry acid = H+ donor Arrhenius acid = produces H+
Lewis base: “e- pair donor” B-L base = H+ acceptor Arrhenius base = produces OH-
Ex:NH3 + BF3 → NH3BF3
Lewis base Lewis acid Lewis salt
6 CN- + Fe3+ → Fe(CN)63-
Lewis base Lewis acid Coordination compound
Gilbert N. Lewis
(1875 – 1946)