Acid-Base Equilibria
BLB 12th Chapter 16
Expectations
 Distinguish between acids and bases


Definitions & properties
Know common strong and weak examples
 Calculate pH for strong and weak systems
 Write chemical reactions of acids and
bases
 Predict relative acid-base strength
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 – hydrogen ion (or proton) donor


Neutral (HNO3), anionic (HCO3-), cationic (NH4+)
Must have a removable (acidic) proton
base – hydrogen ion (or proton) acceptor


Neutral (NH3), anionic (HCO3- , CO32-)
Must have a lone pair of electrons
Acid-Base Reactions (H+-transfer reactions)
HCl(aq) + H2O(l) → Cl-(aq) + H3O+ (aq)
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
HCl(aq) + NH3(aq) → Cl-(aq) + NH4+(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


reaction in which mol acid = mol base
acid(aq) + base(aq) → salt(aq) + water(l)
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
HCl(aq) + NH3(aq) → NH4+(aq) + OH-(aq)
 Conjugate acid/base pairs – reactant and
product that differ by a single proton
HA(aq) + H2O(l) → H3O+(aq) + A-(aq)
acid + base
conj. acid + conj. base
Relative Strengths of Acids and Bases
 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; neutral.
 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.
HA(aq) + H2O(l) →
H3O+(aq) + A-(aq)
HA(aq) + H2O(l) ⇌
H3O+(aq) + A-(aq)
p. 657
 Acid/base reactions proceed from the
stronger acid-base pair to the weaker acidbase pair.
 Common strong acids (p. 664):
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. 665):
M(OH)n, where M = Group I (n=1) &
heavier Group II (n=2) metals
Acid/Base Reactions
16.3 The Autoionization of Water
H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq)
H2O(l) ⇌ H+(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
16.3 The Autoionization of Water
 For an aqueous solution:
Working with Kw
16.4 The pH Scale
 pH represents a solution’s acidity (@ 25°C).
0 ← 7 → 14
acidic neutral basic
 See Table 16.1, p. 661 for summary.
 See Figure 16.5, p. 663 for examples.
 pH = −log[H3O+] = −log[H+]
[H3O+] = 10-pH
pOH = −log[OH-]
[OH-] = 10-pOH
pH + pOH = 14
p. 663
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 – measures exact pH based on
electrochemistry
Acid-base indicators – estimates pH based on
the appearance of color
p. 664
Indicator
Colors.
16.5 Strong Acids and Bases
 Strong acids & bases completely dissociate.
[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
dissociate.
 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. 667 + more in Appendix D, p. 1062
p. 674
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]
Kb ↑ base strength ↑
pKb = −log[Kb]
pKb ↑ base strength↓
From p. 674 + more in Appendix D, p. 1063
% Dissociation (or ionization)
amount dissociated
% dissociation 
x 100%
initial concentration


[ H ]eq
[ HA]0
x 100%
 % dissociation increases as acid/base strength
increases. (p. 669)
 % dissociation decreases as concentration
increases.
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 propionic acid
(CH3CH2CO2H) is 2.94. Calculate the Ka for
propionic 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. 679)
Thus, at 25°C, Ka x Kb = 1.0 × 10-14
And, 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
 See p. 683 for summary of combined effect.
Effect on cations on solution pH
Cation + Anion ►Acidic, basic, or neutral?
Calculate the pH of a 0.15 M NaC2H3O2,
sodium acetate, solution.
16.10 Acid-Base Behavior and Chemical
Structure
 Binary Acids (HX)




As bond H−X bond strength increases, acid strength
decreases.
The greater the stability of the conjugate base, X-,
the stronger the acid.
Group: size of X ↑, bond strength ↓, 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 23.2)
Every Brønsted acid/base is a Lewis
acid/base, but not vice versa.
16.11 Lewis Acids and Bases
Lewis acid & base examples
Amphoterism – capable of acting as a
Bronsted base and a Lewis acid (See
p. 733 for more information.)