Acid-Base Equilibria
BLB 10th Chapter 16
Examples of acids & bases
Acids
Bases
Sour (like vinegar)
Bitter and slippery (like soap)
React with bases to neutralize React with acids to neutralize
them and form salts
them and form salts
Change indicator colors in
opposite direction from base
(e.g. litmus blue to red)
Change indicator colors in
opposite direction from acid
(e.g. litmus red to blue)
Aqueous solutions conduct
electricity
Aqueous solutions conduct
electricity
Liberate hydrogen in reactions React in aqueous solution
with active metals
with salts of heavy metals to
form insoluble hydroxides or
oxides
16.1 Acids & Bases: A Brief Review
 Arrhenius Definitions


Acid – a substance that produces
hydrogen ions (H+) in water
HA → H+ + ABase – a substance that produces
hydroxide ions (OH-) in water
BOH → B+ + OH-
16.2 Brønsted-Lowry Acids & Bases
 H+ (proton) in water:
H+ + H2O → H3O+
hydronium ion
 Hydronium ion can hydrogen bond with
more water molecules to form large
clusters of hydrated hydronium ions.
 H+ and H3O+ are used interchangeably.
16.2 Brønsted-Lowry Acids & Bases
 Brønsted-Lowry definitions
acid – proton donor


Neutral (HNO3), anionic (HCO3-), cationic (NH4+)
Must have a removable (acidic) proton
base – proton acceptor


Neutral (NH3), anionic (CO32-)
Must have a lone pair of electrons
Acid-Base Reactions
HCl(aq) + H2O(l) → H3O+ (aq) + Cl-(aq)
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
HCl(aq) + NH3(aq) → NH4+(aq) + Cl-(aq)
Acid-base reaction in non-aqueous media:
HCl(g) + NH3(g) → NH4Cl(s)
 amphiprotic – capable of behaving as a
Brønsted acid and Brønsted base
 amphoteric – capable of behaving as a
Lewis acid and Brønsted base (17.5)
 neutralization: acid + base → salt + water
 Conjugate acid/base pairs – differ by a
single proton
HA(aq) + H2O(l) → H3O+(aq) + A-(aq)
acid + base
conj. acid + conj. base
Relative Acid/Base Strength
 Strength is a measure of the ability of an
acid (or base) to donate (or accepts) a H+.
 Stronger acids donate H+ more readily.
 Completely
dissociate in water
 Conjugate bases have negligible tendency to
accept protons.
 Weaker acids donate H+ less readily.
 Partially
dissociate and establish equilibrium
 Conjugate bases have some tendency to accept
protons.
 The stronger an acid, the weaker its
conjugate base and vice versa.
p. 672
 Acid/base reactions proceed from the
stronger acid-base pair to the weaker acidbase pair.
 Common strong acids (p. 679):
HClO4, HClO3, H2SO4, HI, HBr, HCl, HNO3


Monoprotic acid – capable of donating only
one H+
Polyprotic acid – capable of donating more
than one H+
 Common strong bases (p. 680):
M(OH)n, where M = group I (n=1) & II (n=2)
metals, except Be
Acid/Base Reactions
16.3 The Autoionization of Water
H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq)
 Kw = [H3O+][OH-] = [H+][OH-] = 1.0 x 10-14 (@ 25°C)
 Kw – ion-product constant (or dissociation constant)
 Pure water is neutral. Thus,
[H3O+] = [OH-] = 1.0 x 10-7 M @ 25°C
 For an aqueous solution:
[H3O+] > [OH-]
[H3O+] = [OH-]
[H3O+] < [OH-]
acidic
neutral
basic
Working with Kw
16.4 The pH Scale
 pH represents a solution’s acidity (@
0 ← 7 → 14
acid neutral base
 See Table 16.1, p. 678 for summary.
 See Figure 16.5, p. 679 for examples.
 pH = −log[H3O+] = −log[H+]
[H3O+] = 10-pH
25°C.
pOH = −log[OH-]
[OH-] = 10-pOH
pH + pOH = 14
p. 676
More common chemicals
Basic
Neutral
Acidic
Chemical
pH
Windex
10.57
Bleach
9.58
Tap water*
7.46
Alka Seltzer (in tap water)
6.43
Distilled water**
6.37
Flat Coke
2.60
Toilet bowl cleaner
1.04
6.0 M HCl
−0.29
*CaCO3 CO3- + H2O ⇌ HCO3- + OH**CO2 + H2O → H2CO3
pH calculations
More about pH
 pH does not necessarily indicate strength.
 Measuring pH


pH meters
Acid-base indicators
p. 679
16.5 Strong Acids and Bases
 Strong acids & bases completely ionize.
[HA]0 = [H3O+] → pH
[MOH]0 = [OH-] → pOH → pH
2[M(OH)2]0 = [OH-] → pOH → pH
 H3O+ is the strongest acid that can exist in
water. (produced by all acids in water)
 OH- is the strongest base that can exist in
water. (produced by all bases in water)
pH problems
End Test #1 material
16.6 Weak Acids & 16.7 Weak Bases
 Weak acids & bases do not completely
ionize.
 Weak acids establish an equilibrium in
aqueous solution.
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
HA(aq) ⇌ H+(aq) + A-(aq)
 They do not readily donate or accept H+’s.
 [HA]0 ≠ [H3O+]
[MOH]0 ≠ [OH-]
Weak Acids & Acid-dissociation Constant
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
HA(aq) ⇌ H+(aq) + A-(aq)




[ H 3O ][ A ] [ H ][ A ]
Ka 

[ HA]
[ HA]
Ka ↑ acid strength ↑
For polyprotic acids: Ka1 >> Ka2 >> Ka3
pKa = −log[Ka]
pKa ↑ acid strength↓
From p. 682 + more in Appendix D, p. 1115-1116
Weak Bases & Base-dissociation Constant
 Weak bases establish an equilibrium in
aqueous solution.
B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)


[ BH ][OH ]
Kb 
[ B]
From p. 691 + more in Appendix D, p. 1115-1116
% Dissociation (or ionization)
amount dissociate d
% dissociati on 
x 100%
initial concentration
x

x 100%
[ HA]0
 % dissociation decreases as concentration
increases (p. 686)
Weak acid/base Problems
1) Ka (or Kb) from equilibrium pH
2) pH from Ka (or Kb)
1. Identify as weak acid or base.
2. Write the chemical equilibrium.
3. Write the equilibrium constant
4.
5.
6.
7.
expression.
Set up concentration table. (Ch. 15.5)
Solve for x.
Check with 5% rule. If greater than 5%,
use quadratic equation. (type 2 only)
Complete problem.
The pH of a 0.10 M solution of propanoic acid
(CH3CH2CO2H) is 2.94. Calculate the Ka for
propanoic acid.
Calculate the pH of a 1.0 M HF solution.
Calculate the pH of a 0.0010 M HF solution.
Calculate the pH of a 0.20 M solution of
triethylamine N(CH2CH3)3.
16.8 Relationship between Ka and Kb
 For a conjugate acid/base pair:
Ka x Kb = Kw (derivation p. 693)
pKa + pKb = pKw = 14.00
16.9 Acid-Base Properties of Salt Solutions
 Salt – ionic compound
 Salts dissolve in water to produce ions.
 Ions can also affect the pH.
 Hydrolysis – reaction between an ion
and water to produce H3O+ or OHF-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq)
NH4+(aq) + H2O(l) ⇌ H3O+(aq) + NH3(aq)
Which ions will undergo hydrolysis, i.e. react
with water and affect the pH of the solution?
 Anion:


Conjugate base of a weak acid ► basic
Conjugate base of a monoprotic strong
acid ► neutral
 Cation:



Conjugate acid of a weak base ► acidic
Group I & II metal ions ► neutral
(exceptions Be2+ and Mg2+ ► acidic)
Other metal ions ► acidic
Cation + Anion ►Acidic, basic, or neutral?
16.10 Acid-Base Behavior and
Chemical Structure
 Binary Acids (HX)

As bond strength increases, acid strength
decreases.

Group: size of X ↑ acid strength ↓
Period: electronegativity of X ↑ acid strength↑

 Oxyacids – acidic H attached to an oxygen
atom

Same # of OH groups and O atoms: central
atom electronegativity ↑ acid strength ↑
HClO > HBrO > HIO

Same central atom, Y: # O atoms ↑ acid
strength ↑
HClO4 > HClO3 > HClO2 > HClO
 Carboxylic acids – contain −COOH or CO2H

# electronegative atoms ↑ acid strength ↑
Oxyacids
16.11 Lewis Acids and Bases
 Lewis acid – electron-pair acceptor


e--poor compounds
Metal ions
 Lewis base – electron-pair donor


Amines, NR3
Ligands (see chapter 24.1)
Every Brønsted base is a Lewis base, but
not vice versa.
Lewis acid & base examples