Oxidation and Reduction
The term oxidation was originally used to describe reactions in which an element combines
Example: The reaction between magnesium metal and oxygen to form magnesium
oxide involves the oxidation of magnesium.
The term reduction comes from the Latin stem meaning "to lead back." Anything that that leads
back to magnesium metal therefore involves reduction.
Example: The reaction between magnesium oxide and carbon at 2000ºC to form
magnesium metal and carbon monoxide is an example of the reduction of magnesium
oxide to magnesium metal.
After electrons were discovered, chemists became convinced that oxidation-reduction
reactions involved the transfer of electrons from one atom to another. From
this perspective, the reaction between magnesium and oxygen is written as
2 Mg + O2 2 [Mg2+][O2-]
In the course of this reaction, each magnesium atom loses two electrons
to form an Mg2+ ion.
Mg Mg2+ + 2 e-
And, each O2 molecule gains four electrons to form a pair of O2- ions.
O2 + 4 e- 2 O2-
Because electrons are neither created nor destroyed in a chemical reaction,
oxidation and reduction are linked. It is impossible to have one without
the other, as shown in the figure below.
Example: Let's determine which element is oxidized and which is reduced
when lithium reacts with nitrogen to form lithium nitride.
6 Li(s) + N2(g) 2 Li3N(s)
In the course of this reaction, the lithium atoms lose one electron to form
Li+ ions, and the nitrogen atoms gain three electrons to form N3- ions. Lithium is therefore oxidized and nitrogen is reduced.
The Role of Oxidation Numbers in Oxidation-Reduction Reactions
Chemists eventually extended the idea of oxidation and reduction to reactions
that do not formally involve the transfer of electrons.
Example: Consider the following reaction.
CO(g) + H2O(g) CO2(g) + H2(g)
As can be seen in the figure below, the total number of electrons in the
valence shell of each atom remains constant in this reaction.
What changes in this reaction is the oxidation state of these atoms. The
oxidation state of carbon increases from +2 to +4, while the oxidation state
of the hydrogen decreases from +1 to 0.
Oxidation and reduction are therefore best defined as follows. Oxidation occurs when the oxidation number of an atom becomes larger. Reduction occurs when the oxidation number of an atom becomes smaller.
Example: Let's determine which atom is oxidized and which is reduced in the following reaction
Sr(s) + 2 H2O(l) Sr2+(aq) + 2 OH-(aq) + H2(g)
We begin by assigning an oxidation number to each component of the reaction.
Nothing happens to the oxidation number of the oxygen in this reaction,
which means that oxygen is neither oxidized nor reduced. The oxidation number
of strontium becomes larger (from 0 to +2), which means Sr is oxidized.
The oxidation number of some (but not all) of the hydrogen atoms becomes
smaller (from +1 to 0). Hydrogen is therefore reduced in this reaction.
Oxidation Numbers Versus the True Charge on Ions
The terms ionic and covalent describe the extremes of a continuum of bonding. There is some covalent
character in even the most ionic compounds and vice versa.
It is useful to think about the compounds of the main group metals as if
they contained positive and negative ions. The chemistry of magnesium oxide,
for example, is easy to understand if we assume that MgO contains Mg2+ and O2- ions. But no compounds are 100% ionic. There is experimental evidence,
for example, that the true charge on the magnesium and oxygen atoms in MgO
is +1.5 and -1.5.
Oxidation states provide a compromise between a powerful model of oxidation-reduction
reactions based on the assumption that these compounds contain ions and
our knowledge that the true charge on the ions in these compounds is not
as large as this model predicts. By definition, the oxidation state of an
atom is the charge that atom would carry if the compound were purely ionic.
For the active metals in Groups IA and IIA, the difference between the oxidation
state of the metal atom and the charge on this atom is small enough to be
ignored. The main group metals in Groups IIIA and IVA, however, form compounds
that have a significant amount of covalent character. It is misleading,
for example, to assume that aluminum bromide contains Al3+ and Br- ions. It actually exists as Al2Br6 molecules.
This problem becomes even more severe when we turn to the chemistry of the
transition metals. MnO, for example, is ionic enough to be considered a
salt that contains Mn2+ and O2- ions. Mn2O7, on the other hand, is a covalent compound that boils at room temperature.
It is therefore more useful to think about this compound as if it contained
manganese in a +7 oxidation state, not Mn7+ ions.
Oxidizing Agents and Reducing Agents
Let's consider the role that each element plays in the reaction in which
a particular element gains or loses electrons..
When magnesium reacts with oxygen, the magnesium atoms donate electrons
to O2 molecules and thereby reduce the oxygen. Magnesium therefore acts as a reducing agent in this reaction.
|2 Mg ||+ O2 || ||2 MgO |
| || || |
The O2 molecules, on the other hand, gain electrons from magnesium atoms and thereby
oxidize the magnesium. Oxygen is therefore an oxidizing agent.
|2 Mg + ||O2 || ||2 MgO |
| ||oxidizing |
| || |
Oxidizing and reducing agents therefore can be defined as follows. Oxidizing agents gain electrons. Reducing agents lose electrons.
Example: Let's identify the oxidizing agent and the reducing agent in the
Ca(s) + H2(g) CaH2(g)
We start by assigning oxidation numbers to this reaction.
We then decide which element is oxidized and which is reduced. In the course
of this reaction, calcium atoms are transformed into Ca2+ ions by the loss of a pair of electrons.
Calcium metal is therefore oxidized in this reaction. The hydrogen atoms,
on the other hand, formally gain an electron to form H- ions. Hydrogen is therefore reduced.
If calcium reduces H2 molecules to H- ions, calcium metal must be the reducing agent. The H2, on the other hand, oxidizes calcium atoms to form Ca2+ ions. H2 is therefore the oxidizing agent.
|Reducing agent:||Ca |
|Oxidizing agent: ||H2|
The table below identifies the reducing agent and the oxidizing agent for
some of the reactions discussed in this web page. One trend is immediately
obvious: The main group metals act as reducing agents in all of their chemical reactions.
Typical Reactions of Main Group Metals
Conjugate Oxidizing Agent/Reducing Agent Pairs
Metals act as reducing agents in their chemical reactions. When copper is
heated over a flame, for example, the surface slowly turns black as the
copper metal reduces oxygen in the atmosphere to form copper(II) oxide.
If we turn off the flame, and blow H2 gas over the hot metal surface, the black CuO that formed on the surface
of the metal is slowly converted back to copper metal. In the course of
this reaction, CuO is reduced to copper metal. Thus, H2 is the reducing agent in this reaction, and CuO acts as an oxidizing agent.
An important feature of oxidation-reduction reactions can be recognized
by examining what happens to the copper in this pair of reactions. The first
reaction converts copper metal into CuO, thereby transforming a reducing
agent (Cu) into an oxidizing agent (CuO). The second reaction converts an
oxidizing agent (CuO) into a reducing agent (Cu). Every reducing agent is
therefore linked, or coupled, to a conjugate oxidizing agent, and vice versa.
Every time a reducing agent loses electrons, it forms an oxidizing agent
that could gain electrons if the reaction were reversed.
Conversely, every time an oxidizing agent gains electrons, it forms a reducing
agent that could lose electrons if the reaction went in the opposite direction.
The idea that oxidizing agents and reducing agents are linked, or coupled,
is why they are called conjugate oxidizing agents and reducing agents. Conjugate comes from the Latin stem meaning "to join together." It is therefore used
to describe things that are linked or coupled, such as oxidizing agents
and reducing agents.
The main group metals are all reducing agents. They tend to be "strong"
reducing agents. The active metals in Group IA, for example, give up electrons
better than any other elements in the periodic table.
The fact that an active metal such as sodium is a strong reducing agent
should tell us something about the relative strength of the Na+ ion as an oxidizing agent. If sodium metal is relatively good at giving
up electrons, Na+ ions must be unusually bad at picking up electrons. If Na is a strong reducing
agent, the Na+ ion must be a weak oxidizing agent.
Conversely, if O2 has such a high affinity for electrons that it is unusually good at accepting
them from other elements, it should be able to hang onto these electrons
once it picks them up. In other words, if O2 is a strong oxidizing agent, then the O2- ion must be a weak reducing agent.
In general, the relationship between conjugate oxidizing and reducing agents
can be described as follows. Every strong reducing agent (such as Na) has a weak conjugate oxidizing
agent (such as the Na+ ion). Every strong oxidizing agent (such as O2) has a weak conjugate reducing agent (such as the O2- ion).
The Relative Strength of Metals as Reducing Agents
We can determine the relative strengths of a pair of metals as reducing
agents by determining whether a reaction occurs when one of these metals
is mixed with a salt of the other. Consider the relative strength of iron
and aluminum, for example. Nothing happens when we mix powdered aluminum
metal with iron(III) oxide. If we place this mixture in a crucible, however,
and get the reaction started by applying a little heat, a vigorous reaction
takes place to give aluminum oxide and molten iron metal.
2 Al(s) + Fe2O3(s) Al2O3(s) + 2 Fe(l)
By assigning oxidation numbers, we can pick out the oxidation and reduction
halves of the reaction.
Aluminum is oxidized to Al2O3 in this reaction, which means that Fe2O3 must be the oxidizing agent. Conversely, Fe2O3 is reduced to iron metal, which means that aluminum must be the reducing
agent. Because a reducing agent is always transformed into its conjugate
oxidizing agent in an oxidation-reduction reaction, the products of this
reaction include a new oxidizing agent (Al2O3) and a new reducing agent (Fe).
Since the reaction proceeds in this direction, it seems reasonable to assume
that the starting materials contain the stronger reducing agent and the
stronger oxidizing agent.
In other words, if aluminum reduces Fe2O3 to form Al2O3 and iron metal, aluminum must be a stronger reducing agent than iron.
We can conclude from the fact that aluminum cannot reduce sodium chloride
to form sodium metal that the starting materials in this reaction are the
weaker oxidizing agent and the weaker reducing agent.
We can test this hypothesis by asking: What happens when we try to run the
reaction in the opposite direction? (Is sodium metal strong enough to reduce
a salt of aluminum to aluminum metal?) When this reaction is run, we find
that sodium metal can, in fact, reduce aluminum chloride to aluminum metal
and sodium chloride when the reaction is run at temperatures hot enough
to melt the reactants.
3 Na(l) + AlCl3(l) 3 NaCl(l) + Al(l)
If sodium is strong enough to reduce Al3+ salts to aluminum metal and aluminum is strong enough to reduce Fe3+ salts to iron metal, the relative strengths of these reducing agents can
be summarized as follows.
Na > Al > Fe
Example: Let's use the following equations to determine the relative strengths
of sodium, magnesium, aluminum, and calcium metal as reducing agents.
2 Na + MgCl2 2 NaCl + Mg
|Al + MgBr2|| |
Ca + MgI2 CaI2 + Mg
|Ca + 2 NaCl|| |
The first reaction suggests that sodium is a stronger reducing agent than
magnesium. The second reaction suggests that aluminum is not as strong a
reducing agent as magnesium or, conversely, that magnesium is a stronger
reducing agent than aluminum. The first two reactions therefore give the
Na > Mg > Al
The third and fourth reactions suggest that calcium is a stronger reducing
agent than magnesium but a weaker reducing agent than sodium. Thus, calcium
has to be inserted into this sequence between sodium and magnesium.
Na > Ca > Mg > Al