192 DIY Science: Illustrated Guide to Home Chemistry Experiments
AT vTIC vI y ERSUS CoNCENTRATIoN
Technically, the pH of a solution depends on the activity
of the hydronium ions rather than their concentration, but
it’s easy to measure concentration accurately and very
difficult to measure activity accurately. (Activity can be
thought of as the effective concentration of hydronium
ions, which is lower than the actual concentration,
because in solution some of the hydronium ions are
“screened” by other ions present in the solution and
therefore unable to participate in reactions.) In very
concentrated solutions, activity may be considerably
lower than the actual concentration. In dilute solutions,
concentration and activity correlate very closely, so the
concentration is typically used for calculations.dAUL. Rp joNES CommENTS
Most beginning students are aware of acids and their dangers but are more or less ignorant of the
dangers of alkali (base). For instance, aqueous sodium hydroxide can blind you in a matter of minutes if
not cleansed thoroughly and I’ve seen lots of kids who are quick to put on goggles to work with 0.01 M
HCl but throw 6 M NaOH around like it’s candy. Aqueous bases are every bit as dangerous as aqueous
acids, if not more so, but many students treat aqueous bases as though they were innocuous.EvRdE y Ay ACId-BASE CHEmISTRy
Acid-base chemistry is inextricably tied with our everyday
lives. Here are just a few examples:- Without the hydrochloric acid in our stomachs, we
could not digest our food. Most people are surprised
to learn that the pH of gastric juice (see Lab 11.1)
can be as low as 1, a very strong acid capable of
“dissolving” some metals. - Mineral (inorganic) acids, including hydrochloric,
nitric, phosphoric, and sulfuric acids, are among the
most important industrial chemicals. Millions of tons
of these acids are produced worldwide, and they are
critical to the manufacture of thousands of important
compounds, including foods and food supplements,
drugs, fertilizers, soaps and detergents, dyes and
pigments, explosives, metals, and fuels. - Bases react with fats by a process called
saponification to produce soaps, a reaction that is
used to produce commercial soap products. But we
also use saponification directly every time we use
a window cleaner or other cleaning product that
contains ammonia. The basic ammonia reacts with
greasy dirt to produce soluble compounds (soaps)
that can be rinsed away.
But not all compounds that behave chemically as acids can
dissociate to produce hydrogen ions, and not all compounds that
behave chemically as bases can dissociate to produce hydroxide
ions. In 1923, the Danish chemist Johannes Nicolaus Brønsted
and the English chemist Thomas Martin Lowry redefined an acid
as a proton (hydrogen nucleus) donor, and a base as a proton
acceptor. Under the Brønsted-Lowry definition of acids and
bases, an acid and its corresponding base are referred to as a
conjugate acid-base pair. The conjugate acid is the member
of the pair that donates a proton, and the conjugate base
the member of the pair that accepts a proton. For example,
hydrochloric acid dissociates in water to form chloride ions and
hydronium ions:
HCl + H 2 o ⇔ H 3 o+ + Cl-
In the forward reaction, the acid reactant (HCl) and the base
reactant (H 2 O) form the acid product (H 3 O+) and the base product
(Cl–). In the reverse reaction, the acid reactant (H 3 O+) and the base
reactant (Cl–) form the acid product (HCl) and the base product
(H 2 O). If the conjugate acid (HCl on the left side of the equation
and H 3 O+ on the right side) is strong, the conjugate base (H 2 O on
the left side of the equation and Cl– on the right side) is weak, and
vice versa. At equilibrium, the weaker acid is favored.
For aqueous solutions, the Arrhenius definition and the Brønsted-
Lowry definition are essentially the same. The value of the
Brønsted-Lowry definition is that it extends the concept of acids
and bases to compounds that are not soluble in water.
The same year that Brønsted and Lowry defined acids and bases
as proton donors and proton acceptors, respectively, the American
chemist Gilbert N. Lewis extended the definition of acids and bases
to include compounds that behaved chemically as acids and bases
without donating or accepting a proton. Under the Lewis definition
of acids and bases, a Lewis acid (also called an electrophile) is
a substance that accepts an electron pair, and a Lewis base (also
called a nucleophile) is a substance that donates an electron pair.
(For example, the compound ferric chloride, FeCl 3 , behaves as an
acid, although it has no proton to donate, and so is classified as a
Lewis acid.) Lewis acids and bases are particularly important to
organic chemists, who make use of them in many syntheses.
In this chapter, we’ll examine the properties of acids and bases.