The term "redox" comes from two concepts involved
with electron transfer: 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.
Although oxidation reactions are commonly associated with
the formation of oxides from oxygen molecules, these are only specific examples
of a more general concept of reactions involving electron transfer.
Redox reactions, or oxidation-reduction reactions, have a
number of similarities to
acid–base reactions. 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 half-reactions, 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).
Etymology
"Redox" is a
portmanteau
of "reduction" and "oxidation".
The word oxidation originally implied reaction with
oxygen to form an oxide, since dioxygen (O2 (g)) was historically
the first recognized oxidizing agent. Later, the term was expanded to encompass
oxygen-like substances that accomplished parallel chemical reactions.
Ultimately, the meaning was generalized to include all processes involving loss
of electrons.
The word
reduction originally referred to the loss in
weight upon heating a metallic
ore such as a
metal oxide to extract the metal. In other words, ore
was "reduced" to metal.
Antoine
Lavoisier (1743-1794) showed that this loss of weight was due to the loss
of oxygen as a gas. Later, scientists realized that the metal atom gains
electrons in this process. The meaning of
reduction then became
generalized to include all processes involving gain of electrons. Even though
"reduction" seems counter-intuitive when speaking of the
gain
of electrons, it might help to think of reduction as the loss of oxygen, which
was its historical meaning.
The electrochemist
John
Bockris has used the words
electronation and
deelectronation
to describe reduction and oxidation processes respectively when they occur at
electrodes.
These words are analogous to
protonation and
deprotonation,
but they have not been widely adopted by chemists.
The term "hydrogenation" could be used instead of
reduction, since hydrogen is the reducing agent in a large number of reactions,
especially in organic chemistry and biochemistry. But, unlike oxidation, which
has been generalized beyond its root element, hydrogenation has maintained its
specific connection to reactions that add hydrogen to another substance
(e.g., the hydrogenation of unsaturated fats in saturated fats, R-CH=CH-R + H2
→ R-CH2-CH2-R).
Oxidizing and reducing agents
In redox processes, the reductant transfers electrons to the
oxidant. Thus, in the reaction, the reductant or reducing agent loses
electrons and is oxidized, and the oxidant or oxidizing agent gains
electrons and is reduced. The pair of an oxidizing and reducing agent that are
involved in a particular reaction is called a redox pair. A redox
couple is a reducing species and its corresponding oxidized form, e.g., Fe2+/Fe3+.
Oxidizers
Substances that have the ability to
oxidize other
substances (cause them to lose electrons) are said to be
oxidative or
oxidizing
and are known as
oxidizing agents, oxidants, or oxidizers. That is,
the oxidant (oxidizing agent) removes electrons from another substance, and is
thus itself reduced. And, because it "accepts" electrons, the
oxidizing agent is also called an
electron
acceptor, hence the name.
Oxygen is the quintessential oxidizer.
Reducers
Substances that have the ability to
reduce other
substances (cause them to gain electrons) are said to be
reductive or
reducing
and are known as
reducing agents, reductants, or reducers. The
reductant (reducing agent) transfers electrons to another substance, and is
thus itself oxidized. And, because it "donates" electrons, the
reducing agent is also called an
electron
donor. Electron donors can also form
charge transfer complexes with electron
acceptors.
Reductants in chemistry are very diverse.
Electropositive
elemental
metals,
such as
lithium,
sodium,
magnesium,
iron,
zinc, and
aluminium,
are good reducing agents. These metals donate or
give away electrons
readily.
Hydride transfer reagents, such as
NaBH4 and
LiAlH4, are widely used in
organic
chemistry,
[3][4] primarily
in the reduction of
carbonyl compounds to
alcohols.
Another method of reduction involves the use of hydrogen gas (H
2)
with a
palladium,
platinum, or
nickel catalyst. These
catalytic reductions are used primarily in the reduction of
carbon-carbon double or triple bonds.
Standard electrode potentials (reduction potentials)
The electrode potential of each half-reaction is also known
as its reduction potential E0red, or potential
when the half-reaction takes place at a cathode. The reduction potential is a
measure of the tendency of the oxidizing agent to be reduced. Its value is zero
for H+ + e− → ½ H2 by definition, positive for
oxidizing agents stronger than H+ (e.g., +2.866 V for F2)
and negative for oxidizing agents that are weaker than H+ (e.g.,
–0.763 V for Zn2+).
For a redox reaction that takes place in a cell, the
potential difference E0cell = E0cathode
– E0anode
However, the potential of the reaction at the anode was
sometimes expressed as an oxidation potential, E0ox
= – E0. The oxidation potential is a measure of the tendency of the
reducing agent to be oxidized, but does not represent the physical potential at
an electrode. With this notation, the cell voltage equation is written with a
plus sign E0cell = E0cathode + E0ox
(anode)
Examples of redox reactions
Illustration of a redox reaction
A good example is the reaction between
hydrogen and
fluorine in
which hydrogen is being oxidized and fluorine is being reduced:
H
2 + F
2 → 2 HF
the oxidation reaction:
and the reduction reaction:
Analyzing each half-reaction in isolation can often make the
overall chemical process clearer. Because there is no net change in charge
during a redox reaction, the number of electrons in excess in the oxidation
reaction must equal the number consumed by the reduction reaction (as shown above).
Elements, even in molecular form, always have an oxidation
state of zero. In the first half-reaction, hydrogen is oxidized from an
oxidation state of zero to an oxidation state of +1. In the second
half-reaction, fluorine is reduced from an oxidation state of zero to an
oxidation state of −1.
When adding the reactions together the electrons are
canceled:
H
2
|
→
|
2 H+ + 2 e−
|
F
2 + 2 e−
|
→
|
2 F−
|
|
H
2 + F
2
|
→
|
2 H+ + 2 F−
|
2 H+ + 2 F− → 2 HF
The overall reaction is:
H
2 + F
2 → 2 HF
Metal displacement
A redox reaction is the force behind an
electrochemical cell like the
Galvanic
cell pictured. The battery is made out of a zinc electrode in a ZnSO
4
solution connected with a wire and a porous disk to a copper electrode in a
CuSO
4 solution.
In this type of reaction, a metal atom in a compound (or in
a solution) is replaced by an atom of another metal. For example,
copper is deposited
when
zinc metal is
placed in a
copper(II) sulfate solution:
Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)
In the above reaction, zinc metal displaces the copper(II)
ion from copper sulfate solution and thus liberates free copper metal.
The ionic equation for this reaction is:
Zn + Cu2+ → Zn2+ + Cu
Zn → Zn2+ + 2 e−
And the copper is reduced:
Cu2+ + 2 e− → Cu
Other examples
2 NO3− + 10 e− + 12 H+
→ N2 + 6 H2O
Corrosion and rusting
- The
term corrosion
refers to the electrochemical oxidation of metals in reaction with an
oxidant such as oxygen. Rusting, the formation of iron
oxides, is a well-known example of electrochemical corrosion; it forms
as a result of the oxidation of iron metal. Common rust often refers to iron(III) oxide, formed in the following
chemical reaction:
4Fe + 3O2 → 2Fe2O3
- The
oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an
acid:
Fe2+ → Fe3+ + e−
H2O2 + 2 e− → 2 OH−
Overall equation:
2 Fe2+ + H2O2 + 2 H+
→ 2 Fe3+ + 2 H2O
Redox reactions in industry
Cathodic protection is a technique used to
control the corrosion of a metal surface by making it the cathode of an
electrochemical cell. A simple method of protection connects protected metal to
a more easily corroded "
sacrificial
anode" to act as the anode. The sacrificial metal instead of the
protected metal, then, corrodes. A common application of cathodic protection is
in
galvanized
steel, in which a sacrificial coating of zinc on steel parts protects them from
rust.
The primary process of reducing ore at high temperature to
produce
metals
is known as
smelting.
The production of
compact
discs depends on a redox reaction, which coats the disc with a thin layer
of metal film.
Redox reactions in biology
Many important
biological processes involve redox reactions.
C6H12O6 + 6 O2 →
6 CO2 + 6 H2O
The process of cell respiration also depends heavily on the
reduction of
NAD+
to NADH and the reverse reaction (the oxidation of NADH to NAD
+).
Photosynthesis
and cellular respiration are complementary, but photosynthesis is not the
reverse of the redox reaction in cell respiration:
Biological energy is frequently stored and released by means
of redox reactions. Photosynthesis involves the reduction of
carbon
dioxide into
sugars
and the oxidation of
water into molecular oxygen. The reverse reaction,
respiration, oxidizes sugars to produce carbon dioxide and water. As
intermediate steps, the reduced carbon compounds are used to reduce
nicotinamide adenine dinucleotide
(NAD
+), which then contributes to the creation of a
proton
gradient, which drives the synthesis of
adenosine triphosphate (ATP) and is
maintained by the reduction of oxygen. In animal cells,
mitochondria
perform similar functions. See the
Membrane potential article.
Free radical reactions are redox reactions that occur
as a part of
homeostasis and killing microorganisms, where an electron
detaches from a molecule and then reattaches almost instantaneously. Free
radicals are a part of redox molecules and can become harmful to the human body
if they do not reattach to the redox molecule or an
antioxidant.
Unsatisfied free radicals can spur the mutation of cells they encounter and
are, thus, causes of cancer.
The term
redox state is often used to describe the
balance of
GSH/GSSG, NAD
+/NADH and
NADP+/NADPH
in a biological system such as a cell or organ. The redox state is reflected in
the balance of several sets of metabolites (e.g.,
lactate
and
pyruvate,
beta-hydroxybutyrate, and
acetoacetate),
whose interconversion is dependent on these ratios. An abnormal redox state can
develop in a variety of deleterious situations, such as
hypoxia,
shock, and
sepsis. Redox
mechanism also control some cellular processes. Redox proteins and their genes
must be co-located for redox regulation according to the
CoRR
hypothesis for the function of DNA in mitochondria and chloroplasts.
Redox cycling
A wide variety of
aromatic
compounds are
enzymatically
reduced to form
free radicals that contain one more electron
than their parent compounds. In general, the electron donor is any of a wide
variety of flavoenzymes and their
coenzymes. Once
formed, these anion free radicals reduce molecular oxygen to
superoxide,
and regenerate the unchanged parent compound. The net reaction is the oxidation
of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form
superoxide. This catalytic behavior has been described as futile cycle or redox
cycling.
Redox reactions in geology
In
geology, redox is important to both the formation of minerals
and the mobilization of minerals, and is also important in some
depositional environments. In general, the
redox state of most rocks can be seen in the color of the rock. The rock forms
in oxidizing conditions, giving it a red color. It is then "bleached"
to a green—or sometimes white—form when a reducing fluid passes through the
rock. The reduced fluid can also carry uranium-bearing
minerals. Famous examples of redox conditions
affecting geological processes include
uranium
deposits and
Moqui marbles.
Balancing redox reactions
Describing the overall electrochemical reaction for a redox
process requires a
balancing of the component
half-reactions
for oxidation and reduction. In general, for reactions in aqueous solution,
this involves adding
H+,
OH−,
H2O, and electrons to compensate for
the oxidation changes.
Acidic media
In acidic media, H+
ions and water are added to half-reactions to balance the overall reaction.
Unbalanced reaction:
|
Mn2+
(aq) + NaBiO
3(s) → Bi3+
(aq) + MnO
4− (aq)
|
Oxidation:
|
4 H
2O(l) + Mn2+
(aq) → MnO−
4(aq) + 8 H+
(aq) + 5 e−
|
Reduction:
|
2 e− + 6 H+
+ BiO−
3(s) → Bi3+
(aq) + 3 H
2O(l)
|
The reaction is balanced by scaling the two half-cell
reactions to involve the same number of electrons (multiplying the oxidation
reaction by the number of electrons in the reduction step and vice versa):
8 H
2O(l) + 2 Mn2+
(aq) → 2 MnO−
4(aq) + 16 H+
(aq) + 10 e−
10 e− + 30 H+
+ 5 BiO−
3(s) → 5 Bi3+
(aq) + 15 H
2O(l)
Adding these two reactions eliminates the electrons terms
and yields the balanced reaction:
14 H+
(aq) + 2 Mn2+
(aq) + 5 NaBiO
3(s) → 7 H
2O(l) + 2 MnO−
4(aq) + 5 Bi3+
(aq) + 5 Na+
(aq)
Basic media
In basic media,
OH−
ions and water are added to half reactions to balance the overall reaction.
Unbalanced reaction:
|
KMnO
4 + Na
2SO
3 + H
2O → MnO
2 + Na
2SO
4 + KOH
|
Reduction:
|
3 e− + 2 H
2O + MnO
4− → MnO
2 + 4 OH−
|
Oxidation:
|
2 OH− + SO
32− → SO
42− + H
2O + 2 e−
|
Balancing the number of electrons in the two half-cell
reactions gives:
6 e− + 4 H
2O + 2 MnO
4− → 2 MnO
2 + 8 OH−
6 OH− + 3 SO
32− → 3 SO
42− + 3 H
2O + 6 e−
Adding these two half-cell reactions together gives the
balanced equation:
2 KMnO
4 + 3 Na
2SO
3 + H
2O → 2 MnO
2 + 3 Na
2SO
4 + 2 KOH
Memory aids
The key terms involved in redox are often confusing to
students. For example, an element that is oxidized loses electrons; however,
that element is referred to as the reducing agent. Likewise, an element that is
reduced gains electrons and is referred to as the oxidizing agent. Acronyms or
mnemonics are commonly used to help remember what is happening:
- "OIL
RIG"—Oxidation Is Loss of electrons, Reduction
Is Gain of electrons.
- "LEO
the lion says GER" — Loss of Electrons is Oxidation,
Gain of Electrons is Reduction.
- "LEORA
says GEROA" — Loss of Electrons is Oxidation (Reducing
Agent), Gain of Electrons is Reduction (Oxidizing
Agent).
- "RED
CAT" and "AN OX", or "AnOx RedCat" ("an
ox-red cat"), — Reduction occurs at the Cathode and the
Anode is for Oxidation.
- "RED
CAT gains what AN OX loses" - Reduction occurs at the Cathode
gains (electrons) what Anode Oxidation loses (electrons).
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