Unformatted text preview: Chemistry for Majors
Module 14: Acid-Based Equilibria Brønsted-Lowry Acids and Bases LEARNING OUTCOMES Identify acids, bases, and conjugate acid-base pairs
according to the Brønsted-Lowry de nition
Write equations for acid and base ionization reactions
Use the ion-product constant for water to calculate
hydronium and hydroxide ion concentrations
Describe the acid-base behavior of amphiprotic
substances The acid-base reaction class has been studied for quite some time. In
1680, Robert Boyle reported traits of acid solutions that included their
ability to dissolve many substances, to change the colors of certain
natural dyes, and to lose these traits after coming in contact with alkali
(base) solutions. In the eighteenth century, it was recognized that acids
have a sour taste, react with limestone to liberate a gaseous substance
(now known to be CO2), and interact with alkalis to form neutral
substances. In 1815, Humphry Davy contributed greatly to the
development of the modern acid-base concept by demonstrating that
hydrogen is the essential constituent of acids. Around that same time,
Joseph Louis Gay-Lussac concluded that acids are substances that can
neutralize bases and that these two classes of substances can be de ned
only in terms of each other. The signi cance of hydrogen was
reemphasized in 1884 when Svante Arrhenius de ned an acid as a
compound that dissolves in water to yield hydrogen cations (now
recognized to be hydronium ions) and a base as a compound that
dissolves in water to yield hydroxide anions.
Johannes Brønsted and Thomas Lowry proposed a more general
description in 1923 in which acids and bases were de ned in terms of the
transfer of hydrogen ions, H+. (Note that these hydrogen ions are often
referred to simply as protons, since that subatomic particle is the only
component of cations derived from the most abundant hydrogen
isotope, 1H.) A compound that donates a proton to another compound is called a Brønsted-Lowry acid, and a compound that accepts a proton is
called a Brønsted-Lowry base. An acid-base reaction is, thus, the
transfer of a proton from a donor (acid) to an acceptor (base).
The concept of conjugate pairs is useful in describing Brønsted-Lowry
acid-base reactions (and other reversible reactions, as well). When an
acid donates H+, the species that remains is called the conjugate base of
the acid because it reacts as a proton acceptor in the reverse reaction.
Likewise, when a base accepts H+, it is converted to its conjugate acid.
The reaction between water and ammonia illustrates this idea. In the
forward direction, water acts as an acid by donating a proton to ammonia
and subsequently becoming a hydroxide ion, OH−, the conjugate base of
water. The ammonia acts as a base in accepting this proton, becoming an
ammonium ion, the conjugate acid of ammonia. In the reverse direction,
a hydroxide ion acts as a base in accepting a proton from ammonium ion,
which acts as an acid. The reaction between a Brønsted-Lowry acid and water is called acid
ionization. For example, when hydrogen uoride dissolves in water and ionizes, protons are transferred from hydrogen uoride molecules to
water molecules, yielding hydronium ions and uoride ions: Base ionization of a species occurs when it accepts protons from water molecules. In the example below, pyridine molecules, C5NH5, undergo
base ionization when dissolved in water, yielding hydroxide and
pyridinium ions: The preceding ionization reactions suggest that water may function as
both a base (as in its reaction with hydrogen uoride) and an acid (as in
its reaction with ammonia). Species capable of either donating or
accepting protons are called amphiprotric, or more
generally, amphoteric, a term that may be used for acids and bases per
de nitions other than the Brønsted-Lowry one. The equations below
show the two possible acid-base reactions for two amphiprotic species,
bicarbonate ion and water: HCO 3 − HCO 3 (aq) + H2 O(l) − (aq) + H2 O(l) CO 3 2− (aq) + H3 O + − H2 CO 3 (aq) + OH (aq) (aq) The rst equation represents the reaction of bicarbonate as an acid with
water as a base, whereas the second represents reaction of bicarbonate
as a base with water as an acid. When bicarbonate is added to water,
both these equilibria are established simultaneously and the composition
of the resulting solution may be determined through appropriate
equilibrium calculations, as described later in this chapter.
In the liquid state, molecules of an amphiprotic substance can react with
one another as illustrated for water in the equations below: The process in which like molecules react to yield ions is
called autoionization. Liquid water undergoes autoionization to a very
slight extent; at 25 °C, approximately two out of every billion water
molecules are ionized. The extent of the water autoionization process is
re ected in the value of its equilibrium constant, the ion-product
constant for water, Kw: H2 O (l) + H2 O (l) ⇌ H3 O + − (aq) + OH (aq) Kw = [H3 O + The slight ionization of pure water is re ected in the small value of the
equilibrium constant; at 25 °C, Kw has a value of 1.0 × 10−14. The process
is endothermic, and so the extent of ionization and the resulting
concentrations of hydronium ion and hydroxide ion increase with
temperature. For example, at 100 °C, the value for Kw is about 5.1 × 10−13,
roughly 100-times larger than the value at 25 °C. EXAMPLE 1: ION CONCENTRATIONS IN PURE WATER What are the hydronium ion concentration and the hydroxide
ion concentration in pure water at 25 °C?
Show Solution
Check Your Learning The ion product of water at 80 °C is 2.4 × 10−13. What are the
concentrations of hydronium and hydroxide ions in pure water
at 80 °C?
Show Solution − [OH EXAMPLE 2: THE INVERSE PROPORTIONALITY OF
[H3O+] AND [OH−]
A solution of carbon dioxide in water has a hydronium ion
concentration of 2.0 × 10−6M. What is the concentration of
hydroxide ion at 25 °C?
Show Solution
Check Your Learning What is the hydronium ion concentration in an aqueous
solution with a hydroxide ion concentration of 0.001 M at 25
°C?
Show Solution EXAMPLE 3: REPRESENTING THE ACID-BASE
BEHAVIOR OF AN AMPHOTERIC SUBSTANCE Write separate equations representing the reaction of HSO
. as an acid with −
3 OH− . as a base with HI
Show Solution
Check Your Learning Write separate equations representing the reaction of H − 2 PO 4 . as a base with HBr
. as an acid with OH−
Show Solution KEY CONCEPTS AND SUMMARY A compound that can donate a proton (a hydrogen ion) to
another compound is called a Brønsted-Lowry acid. The
compound that accepts the proton is called a Brønsted-Lowry
base. The species remaining after a Brønsted-Lowry acid has
lost a proton is the conjugate base of the acid. The species
formed when a Brønsted-Lowry base gains a proton is the
conjugate acid of the base. Thus, an acid-base reaction occurs
when a proton is transferred from an acid to a base, with
formation of the conjugate base of the reactant acid and
formation of the conjugate acid of the reactant base.
Amphiprotic species can act as both proton donors and proton
acceptors. Water is the most important amphiprotic species. It
can form both the hydronium ion, H 3O + , and the hydroxide ion, OH− when it undergoes autoionization:
2H2 O (l) ⇌ H3 O + − (aq) + OH (aq) The ion product of water, Kw is the equilibrium constant for the
autoionization reaction:
Kw = [H2 O + − [OH = 1.0 × 10 −14 ∘ at 25 C Key Equations
Kw = [H2 O + − [OH = 1.0 × 10 −14 ∘ at 25 C TRY IT . Write equations that show NH3 as both a conjugate acid
and a conjugate base.
. Write equations that show H 2 PO 4 − acting both as an acid and as a base.
. Show by suitable net ionic equations that each of the
following species can act as a Brønsted-Lowry acid: a. H 3O + b. HCl
c. NH3
d. CH3CO2H
e. NH + 4 f. HSO −
4 . Show by suitable net ionic equations that each of the
following species can act as a Brønsted-Lowry acid:
a. HNO3
b. PH + 4 c. H2S
d. CH3CH2COOH
e. H 2 PO 4 − f. HS−
. Show by suitable net ionic equations that each of the
following species can act as a Brønsted-Lowry base:
a. H2O
b. OH−
c. NH3
d. CN−
e. S2−
f. H 2 PO 4 − . Show by suitable net ionic equations that each of the
following species can act as a Brønsted-Lowry base:
a. HS−
b. PO 34 c. NH − 2 d. C2H5OH
e. O2− f. H 2 PO 4 − . What is the conjugate acid of each of the following?
What is the conjugate base of each?
a. OH−
b. H2O
c. HCO −
3 d. NH3
e. HSO −
4 f. H2O2
g. HS−
h. H 5 N2 + . What is the conjugate acid of each of the following?
What is the conjugate base of each?
a. H2S
b. H 2 PO 4 − c. PH3
d. HS−
e. HSO −
3 f. H 3 O2 + g. H4N2
h. CH3OH
. Identify and label the Brønsted-Lowry acid, its conjugate
base, the Brønsted-Lowry base, and its conjugate acid in
each of the following equations:
a. HNO 3 b. CN + H2 O ⟶ HCN + OH − + H2 O ⟶ H3 O 2 SO 4 e. O 2− + NO 3 − − c. H d. HSO + −
4 + Cl − ⟶ HCl + HSO 4
− + OH ⟶ SO 4
− + H2 O ⟶ 2OH 2- − + H2 O f.
[Cu(H2 O) g. H 2S + (OH)] 3 + NH2 − 3+ + [Al(H2 O) ]
6 ⟶ HS − 2+ ⟶ [Cu(H2 O) ]
4 + [A + NH3 . Identify and label the Brønsted-Lowry acid, its conjugate
base, the Brønsted-Lowry base, and its conjugate acid in
each of the following equations:
a. NO − − + H2 O ⟶ HNO 2 + OH 2 b. HBr + H 2O c. HS − ⟶ H3 O + + Br − − + H2 O ⟶ H2 S + OH d. H − e. H − 2 PO 4 2 PO 4 − + OH ⟶ HPO 4 2- + H2 O + HCl ⟶ H3 PO 4 + Cl − f.
2+ [Fe(H2 O) g. CH 3 OH 5 (OH)]
− + H 3+ + [Al(H2 O) ] ⟶ CH3 O 6 − 3+ ⟶ [Fe(H2 O) ] + H2 . What are amphiprotic species? Illustrate with suitable
equations.
. State which of the following species are amphiprotic
and write chemical equations illustrating the
amphiprotic character of these species:
a. H2O
b. H 2 PO 4 − c. S2−
d. CO 23 e. HSO −
4 . State which of the following species are amphiprotic
and write chemical equations illustrating the
amphiprotic character of these species.
a. NH3
b. HPO −
4 c. Br−
d. NH 4 + 6 + [A e. ASO 34 . Is the self ionization of water endothermic or
exothermic? The ionization constant for water (Kw) is 2.9
× 10−14 at 40 °C and 9.6 × 10−14 at 60 °C. Show Selected Solutions Glossary
acid ionization: reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid
amphiprotic: species that may either gain or lose a proton in a reaction
amphoteric: species that can act as either an acid or a base
autoionization: reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield
hydronium and hydroxide ions
base ionization: reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base
Brønsted-Lowry acid: proton donor
Brønsted-Lowry base: proton acceptor
conjugate acid: substance formed when a base gains a proton
conjugate base: substance formed when an acid loses a proton
ion-product constant for water (Kw): equilibrium constant for the autoionization of water Previous Next ...
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