Acids and Bases
Chapter 15
Some Properties of Acids
 Produce H+ (as H3O+) ions in water (the hydronium ion is a
hydrogen ion attached to a water molecule)
 Taste sour
 Corrode metals
 Electrolytes
 React with bases to form a salt and water
 pH is less than 7
 Turns blue litmus paper to red “Blue to Red A-CID”
Some Properties of Bases
 Produce OH- ions in water
 Taste bitter, chalky
 Are electrolytes
 Feel soapy, slippery
 React with acids to form salts and water
 pH greater than 7
 Turns red litmus paper to blue
“Basic Blue”
Acid Nomenclature Review
Anion
Ending
No Oxygen
Acid Name
-ide
hydro-(stem)-ic acid
-ate
(stem)-ic acid
-ite
(stem)-ous acid
w/Oxygen
An easy way to remember which goes with which…
“In the cafeteria, you ATE something ICky”
Acid/Base definitions
Definition 1: Arrhenius
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
4.3
Acid/Base Definitions
• Definition #2: Brønsted – Lowry
Acids – proton donor
Bases – proton acceptor
A “proton” is really just a hydrogen
atom that has lost it’s electron!
A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
base
acid
conjugate
acid
conjugate
base
ACID-BASE THEORIES
The Brønsted definition means NH3 is a BASE in
water — and water is itself an ACID
NH3 + H 2O
Base
Acid
NH4+ + OH Acid
Base
Conjugate Pairs
Learning Check!
Label the acid, base, conjugate acid, and
conjugate base in each reaction:
HCl + OH-  Cl- + H2O
Acid
Base
Conj.
Base
Conj.
Acid
H2O + H2SO4  HSO4- + H3O+
Base
Acid
Conj.
Base
Conj.
Acid
Acids & Base Definitions
Definition #3 – Lewis
Lewis acid - a substance that
accepts an electron pair
Lewis base - a substance
that donates an electron
pair
Lewis Acids & Bases
Formation of hydronium ion is also
an excellent example.
H
+
ACID
••
•• O—H
H
BASE
••
H O—H
H
•Electron pair of the new O-H bond
originates on the Lewis base.
Lewis Acid/Base Reaction
The pH scale is a way of
expressing the strength of
acids and bases. Instead of
using very small numbers,
we just use the NEGATIVE
power of 10 on the Molarity
of the H+ (or OH-) ion.
Under 7 = acid
7 = neutral
Over 7 = base
Calculating the pH
pH = - log [H+]
(Remember that the [ ] mean Molarity)
Example: If [H+] = 1 X 10-10
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example: If [H+] = 1.8 X 10-5
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74
Try These!
Find the pH of these:
1) A 0.15 M solution of
Hydrochloric acid
pH = - log [H+]
pH = - log 0.15
pH = - (- 0.82)
pH = 0.82
2) A 3.00 X 10-7 M
solution of Nitric
acid
pH = - log 3 X 10-7
pH = - (- 6.52)
pH = 6.52
pH calculations – Solving for H+
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both
sides and get
10-pH = [H+]
[H+] = 10-3.12 = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift”
or “2nd function” and then the log button
More About Water
H2O can function as both an ACID and a BASE.
In pure water there can be AUTOIONIZATION
Equilibrium constant for water = Kw
Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC
More About Water
Autoionization
OH-
H3O+
Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC
In a neutral solution [H3O+] = [OH-]
and so [H3O+] = [OH-] = 1.00 x 10-7 M
pOH
• Since acids and bases are
opposites, pH and pOH are
opposites!
• pOH does not really exist, but it is
useful for changing bases to pH.
• pOH looks at the perspective of a
base
pOH = - log [OH-]
Since pH and pOH are on opposite
ends,
pH + pOH = 14
pH
[H+]
[OH-]
pOH
+
[H3O ],
[OH ]
and pH
What is the pH of the
0.0010 M NaOH solution?
[OH-] = 0.0010 (or 1.0 X 10-3 M)
pOH = - log 0.0010
pOH = 3
pH = 14 – 3 = 11
OR Kw = [H3O+] [OH-]
[H3O+] = 1.0 x 10-11 M
pH = - log (1.0 x 10-11) = 11.00
What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
Start 0.002 M
HNO3 (aq) + H2O (l)
End 0.0 M
0.0 M
0.0 M
H3O+ (aq) + NO3- (aq)
0.002 M 0.002 M
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Start 0.018 M
Ba(OH)2 (s)
End 0.0 M
0.0 M
0.0 M
Ba2+ (aq) + 2OH- (aq)
0.018 M 0.036 M
pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56
15.4
Strong and Weak Acids/Bases
The strength of an acid (or base) is
determined by the amount of
IONIZATION.
HNO3, HCl, HBr, HI, H2SO4 and HClO4 are
the strong acids.
Strong and Weak Acids/Bases
• Generally divide acids and bases into STRONG or
WEAK ones.
STRONG ACID: HNO3 (aq) + H2O (l) 
H3O+ (aq) + NO3- (aq)
HNO3 is about 100% dissociated in water.
Strong and Weak Acids/Bases
• Weak acids are much less than 100% ionized in
water.
*One of the best known is acetic acid = CH3CO2H
Strong and Weak Acids/Bases
• Strong Base: 100% dissociated in water.
NaOH (aq)  Na+ (aq) + OH- (aq)
Other common strong
bases include KOH and
Ca(OH)2.
CaO (lime) + H2O -->
Ca(OH)2 (slaked lime)
Strong bases are the group I hydroxides
CaO
Calcium, strontium, and barium hydroxides are
strong, but only soluble in water to 0.01 M
Strong and Weak Acids/Bases
• Weak base: less than 100% ionized in water
One of the best known weak bases is ammonia
NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
Weak Bases
Equilibria Involving
Weak Acids and Bases
Consider acetic acid, HC2H3O2 (HOAc)
HC2H3O2 + H2O ↔ H3O+
Acid
+
C2H3O2 Conj. base
[H3O+ ][OAc - ]
-5
Ka 
 1.8 x 10
[HOAc]
(K is designated Ka for ACID)
K gives the ratio of ions (split up) to molecules (don’t split up)
Ionization Constants for Acids/Bases
Acids
Conjugate
Bases
Increase
strength
Increase
strength
Equilibrium Constants
for Weak Acids
Weak acid has Ka < 1
Leads to small [H3O+] and a pH of 2 - 7
Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the
equilibrium concs. of HOAc, H3O+, OAc-,
and the pH.
Step 1. Define equilibrium concs. in ICE
table.
[HOAc]
[H3O+]
[OAc-]
initial
1.00
0
0
change
-x
+x
+x
equilib
1.00-x
x
x
Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs.
of HOAc, H3O+, OAc-, and the pH.
Step 2. Write Ka expression
+
2
[H
O
][OAc
]
x
3
Ka  1.8 x 10-5 =

[HOAc]
1.00 - x
This is a quadratic. Solve using quadratic
formula.
or you can make an approximation if x is very
small! (Rule of thumb: 10-5 or smaller is ok)
Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs.
of HOAc, H3O+, OAc-, and the pH.
Step 3. Solve Ka expression
+
2
[H
O
][OAc
]
x
3
Ka  1.8 x 10-5 =

[HOAc]
1.00 - x
First assume x is very small because
Ka is so small.
Ka  1.8 x 10-5 =
x2
1.00
Now we can more easily solve this
approximate expression.
Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs.
of HOAc, H3O+, OAc-, and the pH.
Step 3. Solve Ka approximate expression
Ka  1.8 x 10-5 =
x2
1.00
x = [H3O+] = [OAc-] = 4.2 x 10-3 M
pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37
Equilibria Involving A Weak Acid
Calculate the pH of a 0.0010 M solution of formic
acid, HCO2H.
HCO2H + H2O ↔ HCO2- + H3O+
Ka = 1.8 x 10-4
Approximate solution
[H3O+] = 4.2 x 10-4 M, pH = 3.37
Exact Solution
[H3O+] = [HCO2-] = 3.4 x 10-4 M
[HCO2H] = 0.0010 - 3.4 x 10-4 = 0.0007 M
pH = 3.47
Equilibrium Constants
for Weak Bases
Weak base has Kb < 1
Leads to small [OH-] and a pH of 12 - 7
Relation
of Ka,
Kb,
[H3O+]
and pH
Equilibria Involving A Weak Base
You have 0.010 M NH3. Calc. the pH.
NH3 + H2O ↔ NH4+ + OHKb = 1.8 x 10-5
Step 1. Define equilibrium concs. in ICE table
[NH3]
[NH4+]
[OH-]
initial
0.010
0
0
change
-x
+x
+x
equilib
0.010 - x
x
x
Equilibria Involving A Weak Base
You have 0.010 M NH3. Calc. the pH.
NH3 + H2O  NH4+ + OHKb = 1.8 x 10-5
Step 2. Solve the equilibrium expression
[NH 4+ ][OH- ]
x2
-5
Kb  1.8 x 10 =
=
[NH3 ]
0.010 - x
Assume x is small, so
x = [OH-] = [NH4+] = 4.2 x 10-4 M
and [NH3] = 0.010 - 4.2 x 10-4 ≈ 0.010 M
The approximation is valid !
Equilibria Involving A Weak Base
You have 0.010 M NH3. Calc. the pH.
NH3 + H2O  NH4+ + OHKb = 1.8 x 10-5
Step 3. Calculate pH
[OH-] = 4.2 x 10-4 M
so pOH = - log [OH-] = 3.37
Because pH + pOH = 14,
pH = 10.63
Types of Acid/Base Reactions:
Summary
Weak Bases are weak electrolytes
F- (aq) + H2O (l)
NO2- (aq) + H2O (l)
OH- (aq) + HF (aq)
OH- (aq) + HNO2 (aq)
Conjugate acid-base pairs:
•
The conjugate base of a strong acid has no measurable
strength.
•
H3O+ is the strongest acid that can exist in aqueous
solution.
•
The OH- ion is the strongest base that can exist in aqueous
solution.
15.4
15.4
Strong Acid
Weak Acid
15.4
Ionized acid concentration at equilibrium
percent ionization =
x 100%
Initial concentration of acid
For a monoprotic acid HA
Percent ionization =
[H+]
[HA]0
x 100%
[HA]0 = initial concentration
15.5
Ionization Constants of Conjugate Acid-Base Pairs
HA (aq)
A- (aq) + H2O (l)
H2O (l)
H+ (aq) + A- (aq)
OH- (aq) + HA (aq)
H+ (aq) + OH- (aq)
Ka
Kb
Kw
KaKb = Kw
Weak Acid and Its Conjugate Base
Kw
Ka =
Kb
Kw
Kb =
Ka
15.7
Molecular Structure and Acid Strength
H X
H+ + X-
The stronger
the bond
The weaker
the acid
• Bond strength
• Polarity
HF << HCl < HBr < HI
15.9
Molecular Structure and Acid Strength
Z
dO
d+
H
Z
O- + H+
The O-H bond will be more polar and easier to break if:
•
Z is very electronegative or
•
Z is in a high oxidation state
15.9
Molecular Structure and Acid Strength
1. Oxoacids having different central atoms (Z) that are from
the same group and that have the same oxidation number.
••
••
••
••
••
••
Acid strength increases with increasing electronegativity of Z
••
••
O
O
••
••
••
••
H O Cl O
H O Br O
•• •• • •
•• •• • •
Cl is more electronegative than Br
HClO3 > HBrO3
15.9
Molecular Structure and Acid Strength
2. Oxoacids having the same central atom (Z) but different
numbers of attached groups.
Acid strength increases as the oxidation number of Z increases.
HClO4 > HClO3 > HClO2 > HClO
15.9
Acid-Base Properties of Salts
Neutral Solutions:
Salts containing an alkali metal or alkaline earth metal
ion (except Be2+) and the conjugate base of a strong
acid (e.g. Cl-, Br-, and NO3-).
NaCl (s)
H2O
Na+ (aq) + Cl- (aq)
Basic Solutions:
Salts derived from a strong base and a weak acid.
NaCH3COO (s)
H 2O
CH3COO- (aq) + H2O (l)
Na+ (aq) + CH3COO- (aq)
CH3COOH (aq) + OH- (aq)
15.10
Acid-Base Properties of Salts
Acid Solutions:
Salts derived from a strong acid and a weak base.
NH4Cl (s)
NH4+ (aq)
H2O
NH4+ (aq) + Cl- (aq)
NH3 (aq) + H+ (aq)
Salts with small, highly charged metal cations (e.g. Al3+,
Cr3+, and Be2+) and the conjugate base of a strong acid.
Al(H2O)3+
6 (aq)
Al(OH)(H2O)52+(aq) + H+ (aq)
15.10
Acid-Base Properties of Salts
Solutions in which both the cation and the anion hydrolyze:
•
Kb for the anion > Ka for the cation, solution will be basic
•
Kb for the anion < Ka for the cation, solution will be acidic
•
Kb for the anion  Ka for the cation, solution will be neutral
15.10