Process of Discovery: Oxidation and Reduction
The first step toward a theory of chemical reactions was
taken by Georg Ernst Stahl in 1697 when he proposed the phlogiston
theory, which was based on the following observations.
- Metals have many properties in common.
- Metals often produce a "calx" when heated.
(The term calx is defined as the crumbly
residue left after a mineral or metal is roasted.)
- These calxes are not as dense as the metals from
which they are produced.
- Some of these calxes form metals when heated with
- With only a few exceptions, the calx is found in
nature, not the metal.
These observations led Stahl to the following conclusions.
- Phlogiston (from the Greek phlogistos,
"to burn") is given off whenever something
- Wood and charcoal are particularly rich in phlogiston
because they leave very little ash when they burn.
(Candles must be almost pure phlogiston because they
leave no ash.)
- Because they are found in nature, calxes must be
simpler than metals.
- Metals form a calx by giving off phlogiston.
Metal calx + phlogiston
- Metals can be made by adding phlogiston to the calx.
Calx + phlogiston metal
- Because charcoal is rich in phlogiston, heating
calxes in the presence of charcoal sometimes produces
This model was remarkably successful. It explained why
metals have similar properties they all contained phlogiston.
It explained the relationship between metals and their calxesthey
were related by the gain or loss of phlogiston. It even
explained why a candle goes out when placed in a bell jar the air
eventually becomes saturated with phlogiston.
There was only one problem with the phlogiston theory. As
early as 1630, Jean Rey noted that tin gains weight when it
forms a calx. (The calx is about 25% heavier than the metal.)
From our point of view, this seems to be a fatal flaw: If
phlogiston is given off when a metal forms a calx, why does
the calx weigh more than the metal? This observation didn't
bother proponents of the phlogiston theory. Stahl explained
it by suggesting that the weight increased because air
entered the metal to fill the vacuum left after the
The phlogiston theory was the basis for research in
chemistry for most of the 18th century. It was not until 1772
that Antoine Lavoisier noted that nonmetals gain large
amounts of weight when burned in air. (The weight of
phosphorus, for example, increases by a factor of about 2.3.)
The magnitude of this change led Lavoisier to conclude that
phosphorus must combine with something in air when it burns.
This conclusion was reinforced by the observation that the
volume of air decreases by a factor of 1/5th when phosphorus
burns in a limited amount of air.
Lavoisier proposed the name oxygene (literally,
"acid-former") for the substance absorbed from air
when a compound burns. He chose this name because the
products of the combustion of nonmetals such as phosphorus
are acids when they dissolve in water.
|P4(s) + 5 O2(g)
|P4O10(s) + 6 H2O(l)
Lavoisier's oxygen theory of combustion was
eventually accepted and chemists began to describe any
reaction between an element or compound and oxygen as oxidation.
The reaction between magnesium metal and oxygen, for example,
involves the oxidation of magnesium.
2 Mg(s) + O2(g)
By the turn of the 20th century, it seemed that all
oxidation reactions had one thing in common
oxidation always seemed to involve the loss of electrons.
Chemists therefore developed a model for these reactions that
focused on the transfer of electrons. Magnesium metal, for
example, was thought to lose electrons to form Mg2+
ions when it reacted with oxygen. By convention, the element
or compound that gained these electrons was said to undergo reduction.
In this case, O2 molecules were said to be reduced
to form O2- ions.
A classic demonstration of oxidation-reduction reactions
involves placing a piece of copper wire into an aqueous
solution of the Ag+ ion. The reaction involves the
net transfer of electrons from copper metal to Ag+
ions to produce whiskers of silver metal that grow out from
the copper wire and Cu2+ ions.
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
The Cu2+ ions formed in this reaction are
responsible for the light-blue color of the solution. Their
presence can be confirmed by adding ammonia to this solution
to form the deep-blue Cu(NH3)42+
Chemists eventually recognized that oxidation-reduction
reactions don't always involve the transfer of electrons.
There is no change in the number of valence electrons on any
of the atoms when CO2 reacts with H2,
CO2(g) + H2(g)
CO(g) + H2O(g)
as shown by the following Lewis structures:
Chemists therefore developed the concept of oxidation
number to extend the idea of oxidation and reduction
to reactions in which electrons are not really gained or
lost. The most powerful model of oxidation-reduction
reactions is based on the following definitions.
Oxidation involves an increase in the oxidation
number of an atom.
Reduction occurs when the oxidation number of an
According to this model, CO2 is reduced when it
reacts with hydrogen because the oxidation number of the
carbon decreases from +4 to +2. Hydrogen is oxidized in this
reaction because its oxidation number increases from 0 to +1.
We find examples of oxidation-reduction or redox
reactions almost every time we analyze the reactions
used as sources of either heat or work. When natural gas
burns, for example, an oxidation-reduction reaction occurs
that releases more than 800 kJ/mol of energy.
CH4(g) + 2 O2(g)
CO2(g) + 2 H2O(g)
Within our bodies, a sequence of oxidation-reduction
reactions are used to burn sugars, such as glucose (C6H12O6)
and the fatty acids in the fats we eat.
+ 6 O2(g) 6 CO2(g) + 6 H2O(l)
+ 26 O2(g) 18 CO2(g) + 18 H2O(l)
We don't have to restrict ourselves to
reactions that can be used as a source of energy, however, to
find examples of oxidation-reduction reactions. Silver metal,
for example, is oxidized when it comes in contact with trace
quantities of H2S or SO2 in the
atmosphere or foods, such as eggs, that are rich in sulfur
4 Ag(s) + 2 H2S(g)
+ O2(g) 2 Ag2S(s)
+ 2 H2O(g)
Fortunately, the film of Ag2S that collects on
the metal surface forms a protective coating that slows down
further oxidation of the silver metal.
The tarnishing of silver is just one example of a broad
class of oxidation-reduction reactions that fall under the
general heading of corrosion. Another
example is the series of reactions that occur when iron or
steel rusts. When heated, iron reacts with oxygen to form a
mixture of iron(II) and iron(III) oxides.
|2 Fe(s) + O2(g) 2 FeO(s)
|2 Fe(s) + 3 O2(g) 2 Fe2O3(s)
Molten iron even reacts with water to form
an aqueous solution of Fe2+ ions and H2
Fe(l) + 2 H2O(l)
Fe2+(aq) + 2 OH-(aq)
At room temperature, however, all three of these reactions
are so slow they can be ignored.
Iron only corrodes at room temperature in the presence of
both oxygen and water. In the course of this reaction, the
iron is oxidized to give a hydrated form of iron(II) oxide.
2 Fe(s) + O2(aq)
+ 2 H2O(l) 2 FeO · H2O(s)
Because this compound has the same empirical formula as
Fe(OH)2, it is often mistakenly called iron(II),
or ferrous, hydroxide. The FeO · H2O formed in
this reaction is further oxidized by O2 dissolved
in water to give a hydrated form of iron(III), or ferric,
4 FeO · H2O(s) + O2(aq)
+ 2 H2O(l) 2 Fe2O3 · 3 H2O(s)
To further complicate matters, FeO · H2O
formed at the metal surface combines with Fe2O3
· 3 H2O to give a hydrated form of magnetic
iron oxide (Fe3O4).
FeO · H2O(s) + Fe2O3
3 H2O(s) Fe3O4
· n H2O(s)
Because these reactions only occur in the presence of both
water and oxygen, cars tend to rust where water collects.
Furthermore, because the simplest way of preventing iron from
rusting is to coat the metal so that it doesn't come in
contact with water, cars were originally painted for only one
reason to slow down the formation of rust.
The key to identifying oxidation-reduction reactions is
recognizing when a chemical reaction leads to a change in the
oxidation number of one or more atoms. It is therefore a good
idea to take another look at the rules for assigning
oxidation numbers. By definition, the oxidation number of an
atom is equal to the charge that would be present on the atom
if the compound was composed of ions. If we assume that CH4
contains C4- and H+ ions, for example,
the oxidation numbers of the carbon and hydrogen atoms would
be -4 and +1.
Note that it doesn't matter whether the compound actually
contains ions. The oxidation number is the charge an atom
would have if the compound was ionic. The concept of
oxidation number is nothing more than a bookkeeping system
used to keep track of electrons in chemical reactions. This
system is based on a series of rules, summarized in the table
|Rules for Assigning Oxidation
- The oxidation number of an atom is zero in a
neutral substance that contains atoms of only
one element. Thus, the atoms in O2,
O3, P4, S8,
and aluminum metal all have an oxidation
number of 0.
- The oxidation number of monatomic ions is
equal to the charge on the ion. The oxidation
number of sodium in the Na+ ion is
+1, for example, and the oxida-tion number of
chlorine in the Cl- ion is -1.
- The oxidation number of hydrogen is +1 when
it is combined with a nonmetal.
Hydrogen is therefore in the +1 oxidation
state in CH4, NH3, H2O,
- The oxidation number of hydrogen is -1 when
it is combined with a metal.
Hydrogen is therefore in the -1 oxidation
state in LiH, NaH, CaH2, and LiAlH4.
- The metals in Group IA form compounds (such
as Li3N and Na2S) in
which the metal atom is in the +1 oxidation
- The elements in Group IIA form compounds
(such as Mg3N2 and CaCO3)
in which the metal atom is in the +2
- Oxygen usually has an oxidation number of -2.
Exceptions include molecules and polyatomic
ions that contain O-O bonds, such as O2,
and the O22- ion.
- The nonmetals in Group VIIA often form
compounds (such as AlF3, HCl, and
ZnBr2) in which the nonmetal is in
the -1 oxidation state.
- The sum of the oxidation numbers of the atoms
in a molecule is equal to the charge on the
- The most electronegative element in a
compound has a negative oxidation number.
Any set of rules, no matter how good, will
only get you so far. You then have to rely on a combination
of common sense and prior knowledge. Questions to keep in
mind while assigning oxidation numbers include the following:
Are there any recognizable ions hidden in the molecule? Does
the oxidation number make sense in terms of the known
electron configuration of the atom?
Chemical reactions are often divided into two categories:
oxidation-reduction or metathesis reactions. Metathesis
reactions include acid-base reactions that involve
the transfer of an H+ ion from a Brønsted acid to
a Brønsted base.
They can also involve the sharing of a pair
of electrons by an electron-pair donor (Lewis base) and an
electron-pair acceptor (Lewis acid).
or redox reactions can involve the transfer
of one or more electrons.
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
They can also occur by the transfer of oxygen, hydrogen,
or halogen atoms.
|CO2(g) + H2(g)
CO(g) + H2O(g)
|SF4(g) + F2(g)
Fortunately, there is an almost foolproof
method of distinguishing between metathesis and redox
reactions. Reactions in which none of the atoms undergoes a
change in oxidation number are metathesis
reactions. There is no change in the oxidation number of any
atom in either of the metathesis reactions, for example.
The word metathesis literally means
"interchange" or "transposition," and it
is used to describe changes that occur in the order of
letters or sounds in a word as a language develops.
Metathesis occurred, for example, when the Old English word brid
became bird. In chemistry, metathesis is used to
describe reactions that interchange atoms or groups of atoms
When at least one atom undergoes a change in oxidation
state, the reaction is an oxidation-reduction reaction. Each
of the reactions in the figure below is therefore an example
of an oxidation-reduction reaction.