Chemistry 25
Chapter 14
Acids and Bases
Types of Electrolytes
Types of Electrolytes

Strong electrolytes

Week electrolytes
Properties of Acids


Sour taste.
React with ―active‖ metals.
•
•



I.e., Al, Zn, Fe, but not Cu, Ag or Au.
2 Al + 6 HCl
AlCl3 + 3 H2
Corrosive.
React with carbonates, producing CO2.
•
Marble, baking soda, chalk, limestone.
CaCO3 + 2 HCl
CaCl2 + CO2 + H2O
Change color of vegetable dyes.
•
Blue litmus turns red.
React with bases to form ionic salts.
Properties of Bases





Also known as alkalis
Bitter Taste
Feel slippery
Change color of vegetable dyes
• Different color than acid
• Litmus = blue
React with acids to form ionic salts, and
often water
• Neutralization
Arrhenius Theory

Acids

Bases
Strong Acids

Definition:
•

equation :
Examples:
•
•
•
•
•
•
Hydroiodic acid
Hydrobromic acid
Hydrochloric acid
Sulfuric acid
Nitric acid
Perchloric acid
Week Acids


Strong acids in the water
Week acids in the water
• Oversaturation of sugars
• Dose the dissolving process stopped?
• Equation:
Strong Bases

Definition:

Examples:
•
•
•
•
•
•
LiOH
NaOH
KOH
RbOH
Ca(OH)2
Ba(OH)2
Neutralization reaction

Example:
• NaOH(s)
• NaOH(s)
• NaOH(s)
+ HCl(aq) →
+ H2SO4(aq) →
+ H3PO4(aq) →
Brønsted-Lowery Theory

Acids

Bases

In the reaction, a

Products are called the conjugate acid and conjugate
base
Comparing Arrhenius Theory
and Brønsted–Lowry Theory

Arrhenius theory
• HCl(aq)
H+(aq) + Cl−(aq)
• HF(aq)
H+(aq) + F−(aq)
• NaOH(aq)
Na+(aq) + OH−(aq)
• NH4OH(aq)
NH4+(aq) + OH−(aq)

Brønsted–Lowry theory
• HCl(aq) + H2O(l)
Cl−(aq) + H3O+(aq)
• HF(aq) + H2O(l)
F−(aq) + H3O+(aq)
• NaOH(aq) + H2O(l)
Na+(aq) + OH−(aq) +
H2O(l)
• NH3(aq) + H2O(l)
NH4+(aq) + OH−(aq)
Brønsted-Lowery Theory

Chemical equation
H-A + :B

A- + H-B+
Conjugate acid-base pair
Brønsted–Lowry
Acid–Base Reactions
HCHO2
H2SO4
+
+ H2O
CHO2–
+
H3O+
H2O
HSO4–
+
H3O+
Brønsted–Lowry
Acid–Base Reactions
H2O
H2SO4
+
+
NH3
HO–
+
NH4+
H2O
HSO4–
+
H3O+
Write the Formula for the
Conjugate Acid of the Following:
• H2O
• NH3
• CO32−
• H2PO41−
Write the Formula for the
Conjugate Base of the Following:
• H2O
• NH3
• CO32−
• H2PO41−
Molecular Structure and
Acid Strength
Molecular Structure and Acid
Strength

Two factors are important in determining
the relative acid strengths.
• One is the polarity of the bond to which the
hydrogen atom is attached.
• The H atom should have a partial positive charge:
• The more polarized the bond, the more easily the
proton is removed and the greater the acid strength.
Molecular Structure and Acid
Strength

Two factors are important in determining
the relative acid strengths.
• The second factor is the strength of the bond.
Or how tightly the proton is held.
• This depends on the size of atom X:
• The larger atom X, the weaker the bond and the
greater the acid strength.
Molecular Structure and Acid
Strength

Consider a series of binary acids from a
given column of elements.

Example
HF
HCl
NH3
HBr
H2O
HI
HF
Molecular Structure and Acid
Strength

Consider the oxoacids. An oxoacid has the
structure:
•
•
•
The acidic H atom is always attached to an O atom,
which in turn is attached to another atom Y.
Bond polarity is the dominant factor in the relative
strength of oxoacids.
This, in turn, depends on the electronegativity of the
atom Y.
Molecular Structure and Acid
Strength

Consider the oxoacids. An oxoacid has the
structure:
•
If the electronegativity of Y is large, then the O-H bond
is relatively polar and the acid strength is greater.
HOCl
HOBr
HOI
Molecular Structure and Acid
Strength

Consider the oxoacids. An oxoacid has the
structure:
•
With each additional O atom in Y, Y becomes
effectively more electronegative.
• As a result, the H atom becomes more acidic.
HClO
HClO2
HClO3
HClO4
Molecular Structure and Acid
Strength

Consider polyprotic acids and their
corresponding anions.
•
•
Each successive H atom becomes more difficult to
remove.
Therefore the acid strength of a polyprotic acid and its
anions decreases with increasing negative charge.
H3PO4
H2PO4-1
HPO4-2
Acid Rxns
Gas Formation Rxns

Acid + (Hydrogen) Carbonate → Salt +
H2O + CO2
Gas Formation Rxns

Acids react with metals
• Acids react with many metals.
• But not all!!
• When acids react with metals, they produce a
salt and hydrogen gas.
Acids React with Metal Oxides

When acids react with metal oxides, they
produce a salt and water.
Base Rxns
Base Reactions


The reaction all bases have in common
is neutralization of acids.
Strong bases will react with Al metal to
form sodium aluminate and hydrogen
gas.
Titration
Titration

Definition:
•
Technique uses reaction
stoichiometry to determine the
concentration of an unknown
solution.

Titrant

Indicators

The endpoint
Acid–Base Titration

The base solution is the titrant in
the buret.

As the base is added to the acid,
the H+ reacts with the OH– to form
water. But there is still excess acid
present, so the color does not
change.

At the titration’s endpoint, just
enough base has been added to
neutralize all the acid. At this point,
the indicator changes color.
Example

What Is the Molarity of an HCl Solution if
10.00 mL Is required to Titrate 12.54 mL of
0.100 M NaOH?
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
Practice

What Is the Molarity of a Ba(OH)2
Solution if 37.6 mL Is Required to Titrate
43.8 mL of 0.107 M HCl?
Ba(OH)2(aq) + 2 HCl(aq)
BaCl2(aq) +
2 H2O(l)
Acidity vs Basicity
Autoionization of Water

Amphiprotic
• Example:

The acidic/basic properties of aqueous
solutions
Autoionization of Water



Water is actually an extremely weak electrolyte.
•
Therefore, there must be a few ions present.
About 1 out of every 10 million water molecules form
ions through a process called autoionization.
H2O
H+ + OH–
H2O + H2O
H3O+ + OH–
All aqueous solutions contain both H3O+ and OH–.
•
The concentration of H3O+ and OH– are equal in water.
Ion Product of Water



The product of the H3O+ and OH–
concentrations is always the same number.
The number is called the ion product of water
and has the symbol Kw.
[H3O+] x [OH–] = 1 x 10-14 = Kw.
Acidic and Basic Solutions

Neutral solutions

Acidic solutions have a larger

Basic solutions have a
Determine the [H3O+] Concentration
and Whether the Solution Is Acidic,
Basic, or Neutral for the Following:

[OH–] = 0.000250 M

[OH–] = 3.50 x 10-8 M

Ca(OH)2 = 0.20 M
Determine the [H3O+]/ [OH-] for a 0.00020 M
Ba(OH)2 and Determine Whether the
Solution Is Acidic, Basic, or Neutral.
pH

The acidity/basicity of a solution is often
expressed as pH.
pH, Continued

The lower the pH, the more acidic the solution;
the higher the pH, the more basic the solution.
• 1 pH unit corresponds to a factor of 10 difference
in acidity.

Normal range is 0 to 14.
• pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M.
• pH can be negative (very acidic) or larger than 14
(very alkaline).
pH of Common Substances
Substance
pH
1.0 M HCl
0.0
0.1 M HCl
1.0
Stomach acid
1.0 to 3.0
Lemons
2.2 to 2.4
Soft drinks
2.0 to 4.0
Plums
2.8 to 3.0
Apples
2.9 to 3.3
Cherries
3.2 to 4.0
Unpolluted rainwater
5.6
Human blood
7.3 to 7.4
Egg whites
7.6 to 8.0
Milk of magnesia (saturated Mg(OH)2)
10.5
Household ammonia
10.5 to 11.5
1.0 M NaOH
14
pOH

The acidity/basicity of a solution may
also be expressed as pOH.
pOH, Continued

The lower the pOH, the more basic the
solution; the higher the pOH, the more
acidic the solution.
• 1 pOH unit corresponds to a factor of 10
difference in basicity.

Normal range is 0 to 14.
• pOH 0 is [OH−] = 1 M; pOH 14 is [H3O+] = 1 M.
• pOH can be negative (very basic) or larger
than 14 (very acidic).

pH + pOH = 14.00.
Calculate the pH of a 0.0010 M
Ba(OH)2 Solution and Determine if
It Is Acidic, Basic, or Neutral.
Calculate the pH of the Following
Strong Acid or Base Solutions.

0.0020 M HCl

0.0050 M Ca(OH)2

0.25 M HNO3
Example

Calculate the Concentration of [H3O+] for
a Solution with pH 3.7.
Practice—Determine the [H3O+]
for Each of the Following:

pH = 2.7

pH = 12

pH = 0.60
Practice—Calculate the pOH and pH
of the Following Strong Acid or Base
Solutions.

0.0020 M KOH

0.0050 M Ca(OH)2

0.25 M HNO3
Buffers



Definition:
•
Solutions resist changing pH when small amounts of
acid or base are added.
They resist changing pH by neutralizing added
acid or base.
Preparation:
•
•
Strong acid/base is neutralized by half is the week
base/acid
Buffers are made by mixing together a weak acid and
its conjugate base.
• Or weak base and its conjugate acid.
Nonmetal Oxides Are Acidic


Nonmetal oxides react with water to form
acids.
Causes acid rain.
What Is Acid Rain?




Natural rain water has a pH of 5.6.
•
Naturally slightly acidic due mainly to CO2.
Rain water with a pH lower than 5.6 is called
acid rain.
Acid rain is linked to damage in ecosystems
and structures.
Many natural and pollutant gases dissolved
in the air are nonmetal oxides.
•
CO2, SO2, NO2.
What Causes Acid Rain?


Nonmetal oxides are acidic.
CO2 + H2O
H2CO3
2 SO2 + O2 + 2 H2O
2 H2SO4
Processes that produce nonmetal oxide
gases as waste increase the acidity of the
rain.
•
•

Natural—volcanoes and some bacterial action.
Man-made—combustion of fuel.
Weather patterns may cause rain to be acidic
in regions other than where the nonmetal
oxide is produced.
Vocabularies








Strong/Weak electrolytes
Ionic/ Molecular compound
Acid/base
•
•
Arrhenius Theory
Brønsted-Lowery Theory
• Conjugate acid-base pair
Neutralization reaction
Titration/Titrant/Indicators/endpoint
Acid Strength
Amphiprotic
pH/pOH/Buffer