Unit 13
Acids and Bases
Properties
• electrolyte
• sour taste
• sticky feel
• turn litmus red
•
•
•
•
electrolyte
bitter taste
slippery feel
turn litmus blue
• react with bases to • react with acids to
form water and a
form water and a
salt (ionic
salt (ionic
compound)
compound)
Examples
ACIDS: Most citrus fruits, tea, battery acid,
vinegar, milk, soda, apples.
BASES: Common household bases include
baking soda, lye, ammonia, soap, and
antacids.
.
Indicators
• Indicators are substances that change color in
the presence of an acid or a base
– Indicators are made up of weak acids or weak
bases
– Examples of indicators include pH paper, red and
blue litmus paper, and phenolphthalein
Acids Affect Indicators:
Blue litmus paper turns red in contact with
an acid. It remains blue when in contact
with a base or neutral solution.
Bases affect indicators:
Red litmus paper
turns blue in contact
with a base. It
remains red when in
contact with an acid
or neutral solution.
Phenolphthalein
turns pink in a
base. It is colorless
in an acid or
neutral solution.
Definitions
• There are 3 definitions used to
describe acids and bases:
• Arrhenius
• BrØnsted-Lowry
• Lewis
• The most traditional is Arrhenius
acids and bases.
Definitions
• Arrhenius - In aqueous solution…
– Acids form hydrogen ions (H+)
HCl + H2O 
+
H
+
–
Cl
Also called hydronium ions (H3O+)
H
H
Cl
acid
O
H
O
H
H
–
+
Cl
H
Definitions
• Arrhenius - In aqueous solution…
– Bases form hydroxide ions (OH-)
NH3 + H2O 
+
NH4
H
H
H
N
H
base
O
H
H
–
+
O
N
H
+
OH
H
H
H
Definitions
Another common way to refer
to hydrogen ions is to call them
“protons”
• Brønsted-Lowry
– Acids are proton (H+) donors.
– Bases are proton (H+) acceptors.
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base conjugate acid
Conjugate acid – particle formed when a base gains a H+
Conjugate base – particle that remains when an acid has
donated a H+
• .
Definitions
• Lewis
– Acids are electron pair acceptors.
– Bases are electron pair donors.
Lewis
base
Lewis
acid
White Board Questions
1. When you wafted a substance your nose
burned. Would this substance be an acid or a
base? ACID
2. A hydrogen ion (H+) can also be called a
+
Proton
H
O
3
_________ or ____________.
3. Arrhenius acids are compounds that break up
H+
in water to give off _____________.
4. What color litmus paper would you use to test
an acid? What color will it turn? Blue turns red
5. If your food tastes bitter, which do you think it
could possibly be an acid or a base? BASE
White Board Questions
accepts
6. A BrØnsted-Lowry base _________
hydrogen
ions.
7. Phenolphthalein turns pink when it comes in
base
contact with a(n) _________.
8. Which of the scientists defined the typical
acid? Arrhenius
9. If you are eating and it has a sour taste, would
that be an acid or a base? acid
10. If a piece of red litmus paper turns blue than it
base
is a(n) ___________.
Naming Acids
• Binary acids
– Contains 2 different elements: H and
another
– Always has “hydro-” prefix
– Root of other element’s name
– Ending “-ic”
Examples of Binary Acids
•
•
•
•
HI is hydroiodic acid
H2S is hydrosulfuric acid
HBr is hydrobromic acid
HCl is hydrochloric acid
Naming Acids
• Ternary Acids - Oxyacids
– Contains 3 different elements: H, O, and
another
– No prefix
– Name of polyatomic ion (p. 147)
– Ending “–ic” for polyatomic ion ending in
“-ate” and “–ous” for ion ending in “-ite”
Examples of Ternary Acids
• ClO3 is chlorate so HClO3 is chloric acid
• PO4 is phosphate
so H3PO4 is phosphoric acid
• PO3 is phosphite
so H3PO3 is phosphorous acid
• NO2 is nitrite HNO2 is nitrous acid
• NO3 is nitrate HNO3 is nitric acid
Naming Acids cont.
• HC2H3O2 or CH3COOH
Name is acetic acid
Common name = vinegar
Practice Naming Acids
• H2SO3
– Sulfurous acid
• HF
– Hydrofluoric acid
• H2Se
– Hydroselenic
acid
• Perchloric acid
– HClO4
• Carbonic acid
– H2CO3
• Hydrobromic acid
– HBr
Ion Product of Water
Self- ionization of water – the simple dissociation of water
H2O
H+ +
OH-
Concentration of ea. ion in pure water:
[H+] = 1.0 x 10-7M
+
[OH-] = 1.0 x 10-7M
Ion-product constant for water (Kw), Where Kw = 1.0 x 10-14
Kw = [H+] [OH-]
Acid [H+] > [OH-]
Base [H+] < [OH-]
Neutral [H+] = [OH-]
Calculating [H+] and [OH-]
• reverse the pH equation
[H+] = 1 x 10-pH and [OH-] = 1 x 10-pOH
• The pH of a solution is 8. Find the [H+] and [OH-]
and determine whether it is acidic, basic, or
neutral.
[H+] = 1 x 10-8 M
[OH-] = 1 x 10-(14-8) M = 1 x 10-6 M
– basic
Examples
1. If the [H+] in a solution is 1.0 x 10-5M, is the
solution acidic, basic or neutral?
1.0 x 10-5 M
pH 5 = acidic
What is the concentration of the [OH-]?
Use the ion-product constant for water (Kw):
Kw = [H+] [OH-]
1.0 x 10-14 = [1.0 x 10-5] [OH-]
1.0 x 10-14 = [OH-]
1.0 x 10-5
1.0 x 10-(14-5)
1.0 x 10-9 M
Examples
2. If the pH is 9, what is the concentration of
the hydroxide ion?
14 = pH + pOH
Kw = [H+] [OH-]
1.0 x 10-14 M = [1.0 x 10-9M] [OH-] 14 = 9 + pOH
1.0 x 10-5 M = [OH-]
5 = pOH
3. If the pOH is 4, what is the concentration
of the hydrogen ion?
14 = pH + pOH
Kw = [H+] [OH-]
1.0 x 10-14 M = [H+] [1.0 x 10-4 M] 14 = pH + 4
1.0 x 10-10 M = [H+]
10 = pH
Examples
4. A solution has a pH of 4. Calculate the pOH,
[H+] and [OH-]. Is it acidic, basic, or neutral?

4
[H ]  110 M
14= pH + pOH
14= 4 + pOH
10= pOH

[OH ]  110
–Acidic since pH is 4
10
M
Practice Problems:
Classify each solution as acidic, basic or
neutral.
1. [H+] = 1.0 x 10-10 M Basic pH 10
2. [H+] = 0.001M 1.0 x 10-3 acid pH 3
3. [OH-] = 1.0 x 10-7 M Neutral
4. [OH-] = 1.0 x 10-4 M 14=pH+4 base pH 10
[OH-]
pOH
pH
[H+]
1 x 10-14
14
0
1 x 100
1 x 10-13
13
1
1 x 10-1
1 x 10-12
12
2
1 x 10-2
1 x 10-11
11
3
1 x 10-3
1 x 10-10
10
4
1 x 10-4
1 x 10-9
9
5
1 x 10-5
1 x 10-8
8
6
1 x 10-6
1 x 10-7
7
7
1 x 10-7
1 x 10-6
6
8
1 x 10-8
1 x 10-5
5
9
1 x 10-9
1 x 10-4
4
10
1 x 10-10
1 x 10-3
3
11
1 x 10-11
1 x 10-2
2
12
1 x 10-12
1 x 10-1
1
13
1 x 10-13
1 x 100
0
14
1 x 10-14
Increasing acidity
Neutral
Increasing basicity
White Board Practice
Fill in the chart.
[OH-]
pOH
pH
[H+]
1.0 X 10 -8
8
6
1.0 X 10 -6
1.0 X 10 -2
2
12
1x 10-12
1.0 X 10 -4
4
10
1.0 X 10 -10
3
11
1.0 X 10 -11
1 x 10-3
Fill in the chart.
[OH-]
pOH
pH
[H+]
1.0 X 10 -8
8
6
1.0 X 10 -6
1.0 X 10 -2
2
12
1x 10-12
1.0 X 10 -4
4
10
1.0 X 10 -10
3
11
1.0 X 10 -11
9
1.0 X 10 -9
1
1 × 10-1
1 x 10-3
1.0 X 10 -5
1.0 X 10 -13
5
13
Strength or Concentration
-
+
• Strong Acid/Base
– Ionize completely in water
– strong electrolyte
Acids
HCl
HNO3
H2SO4
HBr
HI
HClO4
Bases
NaOH
KOH
Ca(OH)2
Ba(OH)2
Strength or Concentration
• Weak Acid/Base
-
+
– ionize partially in water
– weak electrolyte
Acids
HF
CH3COOH
H3PO4
H2CO3
HCN
Base
NH3
Strength or Concentration
• How strong or weak an acid or base is,
depends on its degree of ionization.
-
+
-
+
pH Scale
14
0
7
INCREASING
ACIDITY
NEUTRAL
INCREASING
BASICITY
pH is the negative logarithm of the hydrogen ion concentration
pH =
+
-log[H ]
pouvoir hydrogène (Fr.)
“hydrogen power”
The pH Scale
pH Scale
pH of Common Substances
pH formulas
pH =
pOH =
+
-log[H ]
-log[OH ]
pH + pOH = 14
Neutralization
• Chemical reaction between an acid and a
base.
• Products are a salt (ionic compound) and
water.
Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
neutral
HC2H3O2 + NaOH  NaC2H3O2 + H2O
weak
strong
basic
– Salts can be neutral, acidic, or basic.
– Neutralization does not mean pH = 7.
Titration
standard solution
• Titration
– Analytical method in
which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
Titration cont.
• Equivalence point (endpoint)
– Point at which equal amounts of H+
and OH- have been added.
– Determined by…
• indicator color change
• dramatic change in pH
Titration formula
+
H
moles
= moles
MVn = MVn
OH
M: Molarity
V: volume
n: # of H+ ions in the acid
or OH- ions in the base
Titration example
• 42.5 mL of 1.3M KOH are required to
neutralize 50.0 mL of H2SO4. Find the
molarity of H2SO4.
H2SO4
M=?
KOH
M = 1.3M
V = 50.0 mL V = 42.5 mL
n=2
n=1
MVn = MVn
M(50.0mL)(2)=(1.3M)(42.5mL)(1)
M= 55.25
100
M = 0.55M H2SO4
Review of Acid and Base Definitions
• Arrhenius
Most specific/exclusive definition
Created by Svante Arrhenius, Swedish
Acid: compound that creates H+ in an aqueous
solution
HNO3  H+ + NO3Base: compound that creates OH- in an aqueous
solution
NaOH  Na+ + OH-
Review of Acid and Base Definitions
• Bronsted-Lowry
More general definition than Arrhenius definition
Most commonly used definition
Created by 2 scientists around the same time
(1923)
Acid: Molecule or ion that is a proton (H+) donor
HCl + H2O  H3O+ + ClBase: Molecule or ion that is a proton (H+)
acceptor
NH3 + H2O  NH4+ + OH-
Review of Acid and Base Definitions
• Lewis
Most general definition
Defined by electrons and bonding rather than H+
Created by the same scientist who electron-dot diagrams
are named after
Acid: atom, ion, or molecule that accepts an electron pair to
form a covalent bond
NH3 + Ag+  [Ag(NH3)2]+
Base: atom, ion, or molecule that donates an electron pair
to form a covalent bond
BF3 + F-  BF4-