Chapter 16 Text

16.1 Acids and Bases: A Brief Review
16.2 Br0nsted-Lowry Acids and Bases
16.3 The Autoionization of Water
16.4 The pH Scale
16.5 Strong Acids and Bases
16.6 Weak Acids
16.7 Weak Bases
16.8 Relationship Between Ka and Kb
16.9 Acid-Base Properties of Salt Solutions
16.10 Acid-Base Behavior and Chemical Structure
16.11 Lewis Acids and Bases
ACIDS AND 6AMES are important in numerous chemical
processes that occur around us, from industrial
processes to biological ones, from reactions in the
laboratory to those in our environment. The time
required for a metal object immersed in water to corrode,
the ability of an aquatic environment to support fish and plant life, the fate
of pollutants washed out of the air by rain, and even the rates of reactions that
maintain our lives all critically depend upon the acidity or basicity of solutions.
Indeed, an enormous amount of chemistry can be understood in terms of acidbase
reactions.
We have encountered acids and bases many times in earlier discussions. For
example, a portion of Chapter 4 focused on their reactions. But what makes a
substance behave as an acid or as a base? In this chapter we reexamine acids and
bases, taking a closer look at how they are identified and characterized. In doing
so, we will consider their behavior not only in terms of their structure and bonding
but also in terms of the chemical equilibria in which they participate.
» What's Ahead «
We start by reviewing the definitions
of acid and base presented in Chapter
4 and learn that these are the
Arrhenius definitions.
We then learn the more general
Br0nsted-Lowry definitions for acid
and base. A Br0nsted-Lowry acid is a
proton donor and a Bronsted-Lowry
base is a proton acceptor,
The conjugate base of a Brensted-
Lowry acid is what remains after the
acid has donated a proton. Similarly,
the conjugate acid of a Br0nsted-
Lowry base is the species that results
when the base accepts a proton. Two
such species that differ from each
other only by the presence or absence
of a proton together are known as a
conjugate acid-base pair.
Autoionization of water produces
small concentrations of hydronium
and hydroxide ions in pure water.
The equilibrium constant for autoionization,
Kw, defines the relationship
between H3O+ and OH~ concentrations
in aqueous solutions.
The pH scale is used to describe the
acidity or basicity of a solution.
Strong acids and bases are those that
ionize or dissociate completely in
aqueous solution, whereas weak acids
and bases ionize only partially.
We learn that the ionization of a weak
acid in water is an equilibrium
process with an equilibrium constant
Ka, which can be used to calculate the
pH of a weak acid solution.
Likewise, ionization of a weak base in
water is an equilibrium process with
equilibrium constant K^, which we
can use to calculate the pH of a weak
base solution.
There is a constant relationship,
Ka X Kb - Kw, between the Ka and Kb
of any conjugate acid-base pair. This
relationship can be used to determine
the pH of a salt solution.
We continue by exploring the relationship
between chemical structure
and acid-base behavior.
Finally, we learn the Lewis definitions
of acid and base. A Lewis acid is an
electron acceptor, and a Lewis base is an
electron donor. The Lewis definitions
are more general and inclusive than
either Arrhenius or Br0nsted-Lowry.
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