1. Base Theory Concepts Arrhenius Concept: base produces OH- in water
Bronsted-Lowery Model: base is a H+ acceptor
Strong Hydroxide Bases
Alkali Metal Hydroxides: NaOH, KOH, etc…
Completely Dissociated in Water
NaOH Na OH- K = very large
Alkaline Earth Hydroxides: Ca(OH)2, Mg(OH)2, etc…
Not very soluble in water
What does dissolve is completely dissociated
Ca(OH) Ca OH K = very large
Example: pH of 0.05 M NaOH

2. Non-Hydroxide Bases (Weak Bases)
Any atom with a lone pair of electrons can accept a proton = base
Ammonia in water:
Other amine molecules are also bases
The general base equation
B H2O BH OH-
Strong Base: equilibrium lies far to the right ([OH-] ≈ [B]0)
Weak Base: equilibrium lies far to the left ([OH-] << [B]0)
a) Calculations are similar to weak acids
Kb = 1.8 x10-5
conjugate base
acid
conjugate acid
base

3. Polyprotic Acids Example: pH of 15.0 M NH3 Kb = 1.8 x 10-5
We can find [H+] from KW = [H+][OH-] = 1 x 10-14
or pH pOH = 14
Percent dissociation still means the same thing
The 5% rule for approximations: x/[B] x 100% < 5%
Example: pH of 1.0 M methylamine Kb = 4.38 x 10-4
Polyprotic Acids
Carbonic Acid is a diprotic acid
Polyprotic means there are more than one ionizable proton
Carbonic acid is the acid that helps maintain body pH H2CO3
H2CO H HCO3-
HCO H CO32-
3) Ka1 and Ka2 stand for the loss of the first and second protons, respectively

4. Usually Ka1 >> Ka2
As (-) charge builds up, it is harder to remove the next proton
H+ from the first ionization forces the second ionization to the left
We can usually ignore all but the first ionization in calculations
Phosphoric Acid is a triprotic acid
Ionizations
H3PO H H2PO4- Ka = 7.5 x 10-3
H2PO H HPO42- Ka = 6.2 x 10-8
HPO H PO Ka = 4.8 x 10-13
Ka1 >> Ka2 >> Ka3
Example: pH of 5.0 M H3PO4, and the concentrations of all of the phosphoric acid derived species
Use Ka1 only to find [H+] and the pH
Then, [H+] = [H2PO4-]
[H3PO4] = [H3PO4]0 - [H+]
Find [HPO42-] and [PO43-] from what is already calculated and Ka2, Ka3

5. Sulfuric Acid is a unique diprotic acid
The first ionization of sulfuric acid is a strong acid (Ka1 = large)
H2SO H HSO4-
The second ionization of sulfuric acid is a weak acid
HSO H SO Ka2 = 1.2 x 10-2
Example: pH of 1.0 M H2SO4
Assume complete dissociation of the first proton
Use the equilibrium calculation on Ka2
[H+]0 does not = 0 because of Ka1
[H+]0 = 1.0 M
Approximate 1 + x ≈ 1 in this case (5% rule says this is ok)
When [H2SO4] > 1.0 M, you can ignore Ka2
Example: pH of 0.01 M H2SO4
The 5% rule tells us we can’t ignore Ka2 in this case
We must use the quadratic equation to solve for x
When H2SO4 < 1.0 M, we can’t ignore Ka2

6. Acid-Base Properties of Salts
A) Simple Salts
1) Salt = ionic compound = one that completely ionizes in water
2) Some salts have no effect on pH
Cations of strong bases have no effect on pH
Na+, K+, etc…
These cations have no affinity for OH- in water
Anions of strong acids have no effect on on pH
Cl-, NO3-, etc…
These anions have no affinity for H+ in water
Solutions of these combined ion salts have pH = 7.00
B) Basic Salts
Salts containing the conjugate base of a weak acid produce basic solutions
The conjugate base must be strong if the acid is weak, so it must have a strong affinity for H+, which will affect the pH of a solution
Sodium Acetate Example
NaC2H3O Na C2H3O2-
C2H3O H2O HC2H3O OH-

7. How do we find Kb for the conjugate base of a weak acid?
For any weak acid and its conjugate base, Ka x Kb = KW
Kb = KW / Ka = 1 x / 1.8 x 10-5 = 5.6 x for acetate
Example: pH of 0.30 M NaF Ka for HF = 7.2 x 10-4
Base Strength in Water
Any base in water must compete with the hydroxide anion for protons
Hydrocyanic acid example:
HCN H2O H3O CN- Ka = 6.2 x 10-10
weak acid strong base (compared to H2O)
CN H2O HCN OH- Kb = KW / Ka = 1.6 x 10-5
weak base strong acid
(compared to OH-)
OH- > CN- > H2O

8. Acidic Salts
Salts having the conjugate acid of a weak base produce acidic solutions
Ammonium chloride = NH4Cl NH Cl-
NH NH H+
Since ammonia is a weak base, ammonium is “strong” acid and will effect pH
Example: pH of 0.1 M NH4Cl Kb = 1.8 x 10-5
Find Ka from KW
Make sure the conjugate acid is stronger than water or it will have no effect
Highly charged metal ions can also be acidic
AlCl H2O Al(H2O) Cl-
Al(H2O) H2O Al(H2O)5(OH) H3O+
The higher the metal’s charge the more acidic
Example: pH of 0.01 M AlCl3 Ka = 1.4 x 10-5

9. Salts containing both acidic and basic components
NH4C2H3O NH C2H3O2-
The calculations for these compounds are complex
We can, however, at least decide if the solution is acidic or basic
If Ka > Kb, the solution will be acidic
If Kb > Ka, the solution will be basic
If Ka = Kb, the solution will be neutral
Example: Will the following solutions be acidic or basic?
NH4C2H3O NH4CN Al2(SO4)3

10. Acid-Base Properties and Molecular Structure
Polarity
Not all H containing molecules are acidic
CHCl H CCl3-
Strong, nonpolar bonds don’t dissociate easily
Polar X—H bonds are easily dissociated (acidic)
H—Cl is a polar bond
H—OH is a polar bond
Bond strength also plays a part in acidity
Hydrohalide Polarity: H—F > H—Cl > H—Br > H—I
Bond Strength (kJ.mol)
Acidity weak strong strong strong
Oxyacids: the more O on the central atom, the stronger the acid in a series
Electronegative Oxygens remove electrons from the center atom
This polarizes and weakens the O—H bond even more
HClO4 > HClO3 > HClO2 > HClO
H2SO4 > H2SO3

11. H—O—X Molecules
The more electronegative X is, the more acidic the molecule is
Electronegative X removes electrons from the H—O bond
Acidity: H—O—Cl > H—O—Br > H—O—I > H—O—CH3
Electronegativity:
Acid-Base Properties of Oxides
A generic oxide can be represented as X—O
If X—O is a strong and covalent bond, the oxide will be acidic in water
H—O—X H O—X
If X is electronegative, like O, it should form a strong covalent O—X bond
If X—O is weak and ionic, the oxide will be basic in water
H—O—X X OH-
NaOH Na OH-
The X atom in these oxides is usually not electronegative (Na+)
Examples of Acidic Oxides
SO H2O H2SO H HSO3-
CO H2O H2CO H HCO3-

12. Lewis Acid-Base Definition
Examples of Basic Oxides
CaO H2O Ca(OH) Ca OH-
K2O H2O KOH K OH-
Lewis Acid-Base Definition
Definitions
Lewis Acid = an electron pair acceptor
Lewis Base = an electron pair donor
Example: H :NH NH4+
This model includes the other acid-base concepts
This model accounts for many other chemical reactions that the others don’t
BF :NH H3N:BF3
Lewis acid Lewis Base Lewis Acid-Base complex
Lewis acid Lewis Base Lewis Acid-Base complex

13. AlCl H2O Al(H2O) Cl-
Example: Identify Lewis Acid and Lewis Base
Ni NH Ni(NH3)62+
H H2O H3O+
Summary of Acid-Base Problem Solving
1. List major species in solution
2. Look for reactions that go to completion
a. concentration of product
b. major species left
3. Identify acids and bases
4. Solve the equilibrium problem, check the approximation, find pH