REACTIONS INAQUEOUS SOLUTIONS
Several types of reactions occur in water. When water is the solvent for a reaction, the
reaction is said to occur in aqueous solution, which is denoted by the abbreviation (aq)
following the name of a chemical species in a reaction. Three important types of
reactions in water are precipitation, acid-base, and oxidation-reduction reactions.
Precipitation Reactions
A precipitation reaction is a reaction in which soluble ions in separate solutions are
mixed together to form an insoluble compound that settles out of solution as a
solid. That insoluble compound is called a precipitate.
Solubility Rules for Ionic Compounds
Solubility rules are useful summaries of information about which ionic compounds
(or combinations of ions) are soluble in water and which are not. They are also
important tools for making predictions about whether certain ions will react with
one another to form a precipitate. In addition, they are useful for figuring out
what ions might be involved when a precipitation reaction has been observed.
You will be responsible for being able to use the solubility rules to achieve the
tasks listed above, although you will not be responsible for memorizing the
solubility rules.
In this section you will use solubility rules to predict precipitation reactions and
then write equations to represent them
Acid–base reaction
An acid–base reaction is a chemical reaction that occurs between an acid and a base.
Several concepts that provide alternative definitions for the reaction mechanisms
involved and their application in solving related problems exist. Despite several
differences in definitions, their importance becomes apparent as different methods of
analysis when applied to acid–base reactions for gaseous or liquid species, or when acid
or base character may be somewhat less apparent. The first of these scientific concepts
of acids and bases was provided by the French chemist Antoine Lavoisier
Arrhenius definition
Svante Arrhenius
The Arrhenius definition of acid–base reactions is a development of the hydrogen
theory of acids, devised by Svante Arrhenius, which was used to provide a modern
definition of acids and bases that followed from his work with Friedrich Wilhelm
Ostwald in establishing the presence of ions in aqueous solution in 1884, and led to
Arrhenius receiving the Nobel Prize in Chemistry in 1903 for "recognition of the
extraordinary services... rendered to the advancement of chemistry by his electrolytic
theory of dissociation"]
As defined by Arrhenius, acid–base reactions are characterized by Arrhenius acids,
which dissociate in aqueous solution to form hydrogen ions (H+), and Arrhenius
bases, which form hydroxide (OH−) ions. More recent IUPAC recommendations now
suggest the newer term "hydronium" be used in favor of the older accepted term
"oxonium" to illustrate reaction mechanisms such as those defined in the Brønsted–
Lowry and solvent system definitions more clearly, with the Arrhenius definition
serving as a simple general outline of acid–base character The Arrhenius definition
can be summarised as "Arrhenius acids form hydrogen ions in aqueous solution with
Arrhenius bases forming hydroxide ions."
The universal aqueous acid–base definition of the Arrhenius concept is described as
the formation of water from hydrogen and hydroxide ions, or hydrogen ions and
hydroxide ions from the dissociation of an acid and base in aqueous solution:
H+ (aq) + OH− (aq)
H2 O
(In modern times, the use of H+ is regarded as a shorthand for H3O+, since it is now
known that the bare proton H+ does not exist as a free species in solution.)[citation needed]
This leads to the definition that in Arrhenius acid–base reactions, a salt and water are
formed from the reaction between an acid and a base.[6] In other words, this is a
neutralization reaction.
acid + base →salt + water
The positive ion from a base forms a salt with the negative ion from an acid. For
example, two moles of sodium ion (Na+) from the base sodium hydroxide (NaOH)
combine with one mole of sulfate ion (SO2−
4) from sulfuric acid (H2SO4) to form one mole of sodium sulfate (Na2SO4) . Two
moles of water are also formed.
2 NaOH + H2SO4 →Na2SO4 + 2 H2O
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions,
and refer to the concentration of the solvent ions. Under this definition, pure H2SO4 or
HCl dissolved in toluene are not acidic, and molten KOH and solutions of sodium
amide in liquid ammonia are not alkaline.
Solvent system definition
One of the limitations of Arrhenius definition was its reliance on water solutions.
Edward C. Franklin studied the acid–base reactions in liquid ammonia in 1905 and
pointed out the similarities to water-based Arrhenius theory, and Albert F. O.
Germann, working with liquid COCl2, generalized Arrhenius definition to cover
aprotic solvents and formulated the solvent system theory in 1925.
Germann pointed out that in many solvents there is a certain concentration of a
positive species, solvonium (earlier lyonium) cations and negative species, solvate
(earlier lyate) anions, in equilibrium with the neutral solvent molecules. For example,
water and ammonia undergo such dissociation into hydronium and hydroxide, and
ammonium and amide, respectively:
2 H2 O
H3O+ + OH−
2 NH3
NH+4 + NH−2
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into
nitrosonium and nitrate, antimony trichloride into dichloroantimonium and
tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride.
N2O4
2 SbCl3
COCl2
NO+ + NO−3
SbCl+2 + SbCl−4
COCl+ + Cl−
A solute causing an increase in the concentration of the solvonium ions and a decrease
in the solvate ions is defined as an acid and one causing the reverse is defined as a
base. Thus, in liquid ammonia, KNH2 (supplying NH−
2) is a strong base, and NH4NO3 (supplying NH+
4) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds (supplying SO2+)
behave as acids, and sulfites (supplying SO2−3) behave as bases.
The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in
water:
2 NaNH2 (base) + Zn(NH2)2 (amphiphilic amide) →Na2[Zn(NH2)4]
2 NH4I (acid) + Zn(NH2)2 (amphiphilic amide) →[Zn(NH3)4)]I2
Nitric acid can be a base in liquid sulfuric acid:
HNO3 (base) + 2 H2SO4 →NO+2 + H3O+ + 2 HSO−4
The unique strength of this definition shows in describing the reactions in aprotic
solvents, for example in liquid N2O4:
AgNO3 (base) + NOCl (acid) →N2O4 (solvent) + AgCl (salt)
Since solvent-system definition depends on the solvent as well as on the compound
itself, the same compound can change its role depending on the choice of the solvent.
Thus, HClO4 is a strong acid in water, a weak acid in acetic acid, and a weak base in
fluorosulfonic acid. This was seen as both a strength and a weakness, since some
substances, such as SO3 and NH3, were felt to be acidic or basic on their own right.
On the other hand, solvent system theory was criticized as too general to be useful; it
was felt that there is something intrinsically acidic about hydrogen compounds, not
shared by non-hydrogenic solvonium salts
Brønsted–Lowry definition
Main article: Brønsted–Lowry acid–base theory
The Brønsted–Lowry definition, formulated in 1923, independently by Johannes
Nicolaus Brønsted in Denmark and Martin Lowry in England, is based upon the idea
of protonation of bases through the de-protonation of acids – that is, the ability of
acids to "donate" hydrogen ions (H+) or protons to bases, which "accept" them.
Unlike the previous definitions, the Brønsted–Lowry definition does not refer to the
formation of salt and solvent, but instead to the formation of conjugate acids and
conjugate bases, produced by the transfer of a proton from the acid to the base.[6][10]
In this approach, acids and bases are fundamentally different in behavior from salts,
which are seen as electrolytes, subject to the theories of Debye, Onsager, and others.
An acid and a base react not to produce a salt and a solvent, but to form a new acid
and a new base. The concept of neutralization is thus absent.[
According to Brønsted–Lowry definition, an acid is a compound that can donate a
proton, and a base is a compound that can receive a proton. An acid–base reaction is,
thus, the removal of a hydrogen ion from the acid and its addition to the base.[11] This
does not refer to the removal of a proton from the nucleus of an atom, which would
require levels of energy not attainable through the simple dissociation of acids, but to
removal of a hydrogen ion (H+).
The removal of a proton (hydrogen ion) from an acid produces its conjugate base,
which is the acid with a hydrogen ion removed, and the reception of a proton by a
base produces its conjugate acid, which is the base with a hydrogen ion added.
For example, the removal of H+ from hydrochloric acid (HCl) produces the chloride
ion (Cl−), the conjugate base of the acid:
HCl →H+ + Cl−
The addition of H+ to the hydroxide ion (OH−), a base, produces water (H2O), its
conjugate acid:
H+ + OH− →H2O
Although Brønsted–Lowry acid–base behavior is formally independent of any
solvent, it encompasses Arrhenius and solvent system definitions in an unenforced
way. For example, protonation of ammonia, a base, gives ammonium ion, its
conjugate acid:
H+ + NH3 →NH+4
The reaction of ammonia, a base, with acetic acid in absence of water can be
described to give ammonium cation, an acid, and acetate anion, a base:
CH3COOH + NH3 →NH+4 + CH3COO−
This definition also explains the dissociation of water into low concentrations of
hydronium and hydroxide ions:
H2 O + H2 O
H3O+ + OH−
Water, being amphoteric, can act as both an acid and a base; here, one molecule of
water acts as an acid, donating a H+ ion and forming the conjugate base, OH−, and a
second molecule of water acts as a base, accepting the H+ ion and forming the
conjugate acid, H3O+.
Acid dissociation and acid hydrolysis are seen to be entirely similar phenomena:
HCl (acid) + H2O (base)
NH+4 (acid) + H2O (base)
H3O+ (acid) + Cl− (base)
H3O+ (acid) + NH3 (base)
as are basic dissociation and basic hydrolysis:
NH3 (base) + H2O (acid)
NH+4 (acid) + OH− (base)
CH3COO− (base) + H2O (acid)
CH3COOH (acid) + OH− (base)
Thus, the general formula for acid–base reactions according to the Brønsted–Lowry
definition is:
AH + B →BH+ + A−
where AH represents the acid, B represents the base, BH+ represents the conjugate
acid of B, and A− represents the conjugate base of AH.
Although Brønsted–Lowry calls hydrogen-containing substances like HCl acids,
KOH and KNH2 are not bases but salts containing the bases OH− and NH−
2. Also, some substances, which many chemists considered to be acids, such as SO3 or
BCl3, are excluded from this classification due to lack of hydrogen. Gilbert Lewis
wrote in 1938, "To restrict the group of acids to those substances that contain
hydrogen interferes as seriously with the systematic understanding of chemistry as
would the restriction of the term oxidizing agent to substances containing oxygen."[
redox reaction
Redox (portmanteau for reduction-oxidation) reactions refer to all chemical
reactions in which atoms have their oxidation state changed. This can be either a
simple redox process, such as the oxidation of carbon to yield carbon dioxide (CO2)
or the reduction of carbon by hydrogen to yield methane (CH4), or a complex process
such as the oxidation of glucose (C6H12O6) in the human body through a series of
complex electron transfer processes.
Redox reactions, or oxidation-reduction reactions, have a number of similarities to
acid-base reactions. Fundamentally, redox reactions are a family of reactions that are
concerned with the transfer of electrons between species.
The term comes from the two concepts of reduction and oxidation. It can be explained
in simple terms:


Oxidation is the loss of electrons or an increase in oxidation state by a molecule,
atom, or ion.
Reduction is the gain of electrons or a decrease in oxidation state by a molecule,
atom, or ion.
Like acid-base reactions, redox reactions are a matched set, that is, there cannot be an
oxidation reaction without a reduction reaction happening simultaneously. The
oxidation alone and the reduction alone are each called a half-reaction, because two
half-reactions always occur together to form a whole reaction. When writing halfreactions, the gained or lost electrons are typically included explicitly in order that the
half-reaction be balanced with respect to electric charge.
Though sufficient for many purposes, these descriptions are not precisely correct.
Oxidation and reduction properly refer to a change in oxidation state — the actual
transfer of electrons may never occur. Thus, oxidation is better defined as an increase
in oxidation state, and reduction as a decrease in oxidation state. In practice, the
transfer of electrons will always cause a change in oxidation state, but there are many
reactions that are classed as "redox" even though no electron transfer occurs (such as
those involving covalent bonds).
Non-redox reactions, which do not involve changes in formal charge, are known as
metathesis reactions