The Chemical Composition of Aqueous Solutions
Water is the
most plentiful solvent on Earth, is easily purified, and is not toxic. It is, therefore,
widely used as a medium for chemical analyses.
Classifying Solutions of Electrolytes
** Most of the
solutes we will discuss are electrolytes, which form ions when dissolved in water
(or certain other solvents) and thus produce solutions that conduct
electricity.
** Strong electrolytes
ionize essentially completely in a solvent, but weak electrolytes ionize only partially.
These characteristics mean that a solution of a weak electrolyte will not
conduct electricity as well as a solution containing an equal concentration of
a strong electrolyte.
** Table (1)
shows various solutes that act as strong and weak electrolytes in water. Among
the strong electrolytes listed are acids, bases, and salts.
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Table (1) |
*H2SO4
is completely dissociated into HSO4- and H3O+
ions and for this reason is classified as a strong electrolyte. Note, however,
that the HSO4+ ion is a weak electrolyte and is only partially
dissociated into SO42- and H3O+
Acids and Bases
** In 1923, J.
N. Brønsted in Denmark and J. M. Lowry in England proposed independently a theory
of acid/base behavior that is especially useful in analytical chemistry.
** According to
the Brønsted-Lowry theory, an acid is a proton donor, and a base is a proton
acceptor. For a molecule to behave as an acid, it must encounter a proton
acceptor (or base). Likewise, a molecule that can accept a proton behaves as a
base if it encounters an acid.
Conjugate Acids and Bases
** An important
feature of the Brønsted-Lowry concept is the idea that the product formed when an
acid gives up a proton is a potential proton acceptor and is called the
conjugate base of the parent acid. For example, when the species acid1
gives up a proton, the species base1 is formed, as shown by the
reaction:
** We refer to
acid1 and base1 as a conjugate acid/base pair, or just a conjugate pair. Similarly, every base accepts a proton to
produce a conjugate acid. That is,
** When these
two processes are combined, the result is an acid/base, or neutralization, reaction:
** This
reaction proceeds to an extent that depends on the relative tendencies of the
two bases to accept a proton (or the two acids to donate a proton). Examples of
conjugate acid/base relationships are shown in Equations (1) through (4).
** Many
solvents are proton donors or proton acceptors and can thus induce basic or
acidic behavior in solutes dissolved in them. For example, in an aqueous
solution of ammonia, water can donate a proton and acts as an acid with respect
to the solute NH3:
** In this
reaction, ammonia (base1) reacts with water, which is labeled acid2,
to give the conjugate acid ammonium ion (acid1) and hydroxide ion,
which is the conjugate base (base2) of the acid water.
** On the other
hand, water acts as a proton acceptor, or base, in an aqueous solution of
nitrous acid:
** The
conjugate base of the acid HNO2 is nitrite ion. The conjugate acid
of water is the hydrated proton written as H3O+. This
species is called the hydronium ion, and it consists of a proton covalently
bonded to a single water molecule.
** Higher
hydrates such as H5O2+, H9O4+,
and the dodecahedral cage structure shown in Figure (1) may also appear in
aqueous solutions of protons. For convenience, however, we generally use the
notation H3O+, or more simply H+, when we
write chemical equations containing the hydrated proton.
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Figure (1) |
** Figure (1)
Possible structures for the hydronium ion. (a) The species H9O4+
has been observed in the solid state and may be an important contributor in
aqueous solution. (b) The species (H2O)20H+ exhibits
a dodecahedral caged structure. The extra proton in the structure, which may be
any one of the three marked with an asterisk, is free to move around the
surface of the dodecahedron by being transferred to an adjacent water molecule
** An acid that
has donated a proton becomes a conjugate base capable of accepting a proton to reform
the original acid. Similarly, a base that has accepted a proton becomes a
conjugate acid that can donate a proton to form the original base. Thus, nitrite
ion, the species produced by the loss of a proton from nitrous acid, is a potential
acceptor of a proton from a suitable donor. It is this reaction that causes an aqueous
solution of sodium nitrite to be slightly basic:
Notes
(1) A
salt: is produced in the reaction of an acid with a base. Examples include
NaCl, Na2SO4, and NaOOCCH3 (sodium acetate).
(2) An
acid: donates protons. A base: accepts protons.
(3) An
acid donates protons only in the presence of a proton acceptor (a base).
Likewise, a base accepts protons only in the presence of a proton donor
(an acid).
(4) A conjugate base: is formed when an acid loses a proton.
For example, acetate ion is the conjugate base of acetic acid. Similarly,
ammonium ion is the conjugate acid of the base ammonia.
(5) A
conjugate acid: is formed when a base accepts a proton.
(6) A
substance acts as an acid only in the presence of a base and vice versa.
Amphiprotic Species
** Species that
have both acidic and basic properties are amphiprotic. An example is dihydrogen
phosphate ion, H2PO4- , which behaves as a
base in the presence of a proton donor such as H3O+
** Here, H3PO4
is the conjugate acid of the original base. In the presence of a proton
acceptor, such as hydroxide ion, however, H2PO4 - behaves as an acid and donates a proton to
form the conjugate base HPO42-
** The simple amino acids are an important
class of amphiprotic compounds that contain both a weak acid and a weak base
functional group. When dissolved in water, an amino acid, such as glycine,
undergoes a kind of internal acid/base reaction to produce a zwitterion—a
species that has both a positive and a negative charge. Thus,
** This
reaction is analogous to the acid/base reaction between a carboxylic acid and
an amine:
** Water is the
classic example of an amphiprotic solvent, that is, a solvent that can act
either as an acid (Equation 1) or as a base (Equation 2), depending on the
solute.
** Other common
amphiprotic solvents are methanol, ethanol, and anhydrous acetic acid. In
methanol, for example, the equilibria analogous to the water equilibria shown
in Equations (1) and (2) are:
Notes
(1) A
zwitterion is an ion that has both a positive and a negative charge.
(2) Water can
act as either an acid or a base.
(3) Amphiprotic
solvents behave as acids in the presence of basic solutes and bases in the
presence of acidic solutes.
Autoprotolysis
** Amphiprotic
solvents undergo self-ionization, or autoprotolysis, to form a pair of ionic
species.
** Autoprotolysis
is yet another example of acid/base behavior, as illustrated by the following
equations:
** The extent
to which water undergoes autoprotolysis at room temperature is slight. Thus,
the hydronium and hydroxide ion concentrations in pure water are only about 10-7
M. Despite the small values of these concentrations, this dissociation reaction
is of utmost importance in understanding the behavior of aqueous solutions.
Note: Autoprotolysis
(also called autoionization) is the spontaneous reaction of molecules of a substance
to give a pair of ions
Strengths of Acids and Bases
** Figure (2)
shows the dissociation reactions of a few common acids in water.
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Figure (2) |
(b) The rest
are weak acids, which react incompletely with water to give solutions
containing significant quantities of both the parent acid and its conjugate
base.
** Note that
acids can be cationic, anionic, or electrically neutral. The same holds for
bases.
** The acids in
Figure (2) become progressively weaker from top to bottom. Perchloric acid and
hydrochloric acid are completely dissociated, but only about 1% of acetic acid
(HC2H3O2) is dissociated. Ammonium ion is an
even weaker acid with only about 0.01% of this ion being dissociated into
hydronium ions and ammonia molecules. Another generality illustrated in Figure (2)
is that the weakest acid forms the strongest conjugate base, that is, ammonia
has a much stronger affinity for protons than any base above it. Perchlorate
and chloride ions have no affinity for protons.
** The tendency
of a solvent to accept or donate protons determines the strength of a solute
acid or base dissolved in it. For example, perchloric and hydrochloric acids
are strong acids in water. If anhydrous acetic acid, a weaker proton acceptor
than water, is substituted as the solvent, neither of these acids undergoes
complete dissociation. Instead, equilibria such as the following are
established:
** Perchloric
acid is, however, about 5000 times stronger than hydrochloric acid in this
solvent. Acetic acid thus acts as a differentiating solvent toward the two
acids by revealing the inherent differences in their acidities. Water, on
the other hand, is a leveling solvent for perchloric, hydrochloric, and nitric
acids because all three are completely ionized in this solvent and show no
differences in strength. There are differentiating and leveling solvents for
bases as well.
** The common
strong bases include NaOH, KOH, Ba(OH)2, and the quaternary ammonium
hydroxide R4NOH, where R is an alkyl group such as CH3 or
C2H5.
** The common
strong acids include HCl, HBr, HI, HClO4, HNO3, the first
proton in H2SO4, and the organic sulfonic acid RSO3H.
Reference: Fundamentals of analytical chemistry / Douglas A. Skoog, Donald M. West, F. James Holler, Stanley R. Crouch. (ninth edition) , 2014 . USA
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