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Redox
From Wikipedia, the free encyclopedia

Redox (short for reduction–oxidation

reaction) is a chemical reaction in which
the oxidation states of atoms are changed.
Any such reaction involves both a reduction
process and a complementary oxidation
process, two key concepts involved with
The two parts of a redox reaction
electron transfer processes.[1] Redox
reactions include all chemical reactions in
which atoms have their oxidation state
changed; in general, redox reactions involve
the transfer of electrons between chemical
species. The chemical species from which
the electron is stripped is said to have been
oxidized, while the chemical species to
which the electron is added is said to have
been reduced. 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 Rusting iron
a decrease in oxidation state by a
molecule, atom, or ion.

As an example, during the combustion of

wood, oxygen from the air is reduced,
transferring electrons from the carbon.[2]
Although oxidation reactions are commonly
associated with the formation of oxides
from oxygen molecules, oxygen is not
necessarily included in such reactions, as
other chemical species can serve the same
function.[2]
A bonfire
The reaction can occur relatively slowly, as
in the case of rust, or more quickly, as in the
case of fire. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide
(CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such
as the oxidation of glucose (C6H12O6) in the human body.

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Contents
◾ 1 Etymology
◾ 2 Definitions
◾ 3 Oxidizing and reducing agents
◾ 3.1 Oxidizers
◾ 3.2 Reducers
◾ 4 Standard electrode potentials (reduction potentials)
◾ 5 Examples of redox reactions Sodium and fluorine bonding ionically to form sodium
◾ 5.1 Metal displacement fluoride. Sodium loses its outer electron to give it a stable
◾ 5.2 Other examples electron configuration, and this electron enters the fluorine
◾ 5.2.1 Corrosion and rusting atom exothermically. The oppositely charged ions are then
◾ 6 Redox reactions in industry attracted to each other. The sodium is oxidized, and the
◾ 7 Redox reactions in biology fluorine is reduced.
◾ 7.1 Redox cycling
◾ 8 Redox reactions in geology
◾ 9 Balancing redox reactions
◾ 9.1 Acidic media
◾ 9.2 Basic media
◾ 10 Memory aids
◾ 11 See also
◾ 12 References
◾ 13 External links

Demonstration of the reaction between a strong

oxidising and a reducing agent. When few drops
Etymology of glycerol (reducing agent) are added to
powdered potassium permanganate (strong
"Redox" is a combination of "reduction" and oxidising agent), a vigorous reaction
"oxidation". accompanied by self-ignition starts.

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. Since electrons are negatively charged, it is also helpful to think of this as reduction
in electrical charge.

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The electrochemist John Bockris has used the words electronation and deelectronation to describe
reduction and oxidation processes respectively when they occur at electrodes.[3] 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 into saturated fats, R−CH=CH−R + H2 → R−CH2−CH2−R). The word "redox" was first used in
1928.[4]

Definitions
The processes of oxidation and reduction occur simultaneously and cannot happen independently of one
another, similar to the acid–base reaction.[2] 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 general descriptions are not precisely correct. Although
oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons
may never occur. The oxidation state of an atom is the fictitious charge that an atom would have if all
bonds between atoms of different elements were 100% ionic. Thus, oxidation is best 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).

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 oxidizing 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. Oxygen is the quintessential
oxidizer.

Oxidants are usually chemical substances with elements in high oxidation states (e.g., H2O2, MnO−4 ,
CrO3, Cr2O2−7 , OsO4), or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra
electrons by oxidizing another substance.

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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,
The international pictogram for
zinc, and aluminium, are good reducing agents. These metals
oxidising chemicals.
donate or give away electrons readily. Hydride transfer reagents,
such as NaBH4 and LiAlH4, are widely used in organic
chemistry,[5][6] primarily in the reduction of carbonyl compounds to alcohols. Another method of
reduction involves the use of hydrogen gas (H2) 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)

0
Each half-reaction has a standard electrode potential (Ecell), which is equal to the potential difference or
voltage at equilibrium under standard conditions of an electrochemical cell in which the cathode reaction
is the half-reaction considered, and the anode is a standard hydrogen electrode where hydrogen is
oxidized:
1
⁄2 H2 → H+ + e−.

0
The electrode potential of each half-reaction is also known as its reduction potential Ered, 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− → 1⁄2 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+).[7]

For a redox reaction that takes place in a cell, the potential difference is:

0 0 0
Ecell = Ecathode – Eanode

However, the potential of the reaction at the anode was sometimes expressed as an oxidation potential:

0 0
Eox = –Ered.

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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

0 0 0
Ecell = Ered(cathode) + Eox(anode)

Examples of redox reactions

A good example is the reaction between hydrogen
and fluorine in which hydrogen is being oxidized
and fluorine is being reduced:

H2 + F2 → 2 HF

We can write this overall reaction as two half-

reactions:
Illustration of a redox reaction
the oxidation reaction:

H2 → 2 H+ + 2 e−

and the reduction reaction:

F 2 + 2 e− → 2 F −

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:

H2 → 2 H+ + 2 e−
F2 + 2 e− → 2 F−

H2 + F2 → 2 H+ + 2 F−

And the ions combine to form hydrogen fluoride:

2 H+ + 2 F− → 2 HF

The overall reaction is:

H + F → 2 HF

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2 2

Metal displacement

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: A redox reaction is the force behind an electrochemical cell
like the Galvanic cell pictured. The battery is made out of a
Zn + Cu2+ → Zn2+ + Cu zinc electrode in a ZnSO4 solution connected with a wire
and a porous disk to a copper electrode in a CuSO4 solution.
As two half-reactions, it is seen that the zinc
is oxidized:

Zn → Zn2+ + 2 e−

And the copper is reduced:

Cu2+ + 2 e− → Cu

Other examples

◾ The reduction of nitrate to nitrogen in the presence of an acid (denitrification):

2 NO−3 + 10 e− + 12 H+ → N2 + 6 H2O

◾ The combustion of hydrocarbons, such as in an internal combustion engine, which produces

water, carbon dioxide, some partially oxidized forms such as carbon monoxide, and heat energy.
Complete oxidation of materials containing carbon produces carbon dioxide.
◾ In organic chemistry, the stepwise oxidation of a hydrocarbon by oxygen produces water and,
successively, an alcohol, an aldehyde or a ketone, a carboxylic acid, and then a peroxide.

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:

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4 Fe + 3 O2 → 2 Fe2O3

◾ 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 Oxides, such as iron(III) oxide or

rust, which consists of hydrated iron
Redox reactions in industry (III) oxides Fe2O3·nH2O and iron(III)
oxide-hydroxide (FeO(OH), Fe(OH)
Cathodic protection is a technique used to control the corrosion 3), form when oxygen combines with

of a metal surface by making it the cathode of an electrochemical other elements

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.

Oxidation is used in a wide variety of industries such as in the

production of cleaning products and oxidizing ammonia to Iron rusting in pyrite cubes
produce nitric acid, which is used in most fertilizers.

Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or
support electrosynthesis.

The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in
chrome-plated automotive parts, silver plating cutlery, and gold-plated jewelry.

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.

Cellular respiration, for instance, is the oxidation of glucose (C6H12O6) to CO2 and the reduction of
oxygen to water. The summary equation for cell respiration is:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O

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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:

6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2

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.
Top: ascorbic acid (reduced
Free radical reactions are redox reactions that occur as a part of homeostasis form of Vitamin C)
and killing microorganisms, where an electron detaches from a molecule Bottom: dehydroascorbic
and then reattaches almost instantaneously. Free radicals are a part of redox acid (oxidized form of
molecules and can become harmful to the human body if they do not Vitamin C)
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.

Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and
quinones such as menadione.[8]

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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.
Mi Vida uranium mine, near Moab,
Utah. The alternating red and
Balancing redox reactions white/green bands of sandstone
correspond to oxidized and reduced
Describing the overall electrochemical reaction for a redox conditions in groundwater redox
process requires a balancing of the component half-reactions for chemistry.
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.

For instance, when manganese(II) reacts with sodium bismuthate:

Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO−4 (aq)

Oxidation: 4 H2O(l) + Mn2+(aq) → MnO−4 (aq) + 8 H+(aq) + 5 e−
Reduction: 2 e− + 6 H+ + BiO−3 (s) → Bi3+(aq) + 3 H2O(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 H2O(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 H2O(l)

Adding these two reactions eliminates the electrons terms and yields the balanced reaction:
+
14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 7 H2O(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.

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For example, in the reaction between potassium permanganate and sodium sulfite:

Unbalanced reaction: KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH

Reduction: 3 e− + 2 H2O + MnO−4 → MnO2 + 4 OH−
Oxidation: 2 OH− + SO2− 2−
3 → SO4 + H2O + 2 e
−

Balancing the number of electrons in the two half-cell reactions gives:

6 e− + 4 H2O + 2 MnO−4 → 2 MnO2 + 8 OH−

6 OH− + 3 SO2− 2−
3 → 3 SO4 + 3 H2O + 6 e
−

Adding these two half-cell reactions together gives the balanced equation:

2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH

Memory aids
The key terms involved in redox are often confusing to students.[9][10] 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.[11] Acronyms or
mnemonics are commonly used[12] to help remember the terminology:

◾ "OIL RIG" — oxidation is loss of electrons, reduction is gain of electrons.[9][10][11][12]

◾ "LEO the lion says GER" — loss of electrons is oxidation, gain of electrons is reduction.
[9][10][11][12]

◾ "LEORA says GEROA" — loss of electrons is oxidation (reducing agent), gain of electrons is
reduction (oxidizing agent).[11]
◾ "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 at the cathode gains (electrons) what anode
oxidation loses (electrons).

See also
◾ Bessemer process ◾ Electron transport chain
◾ Bioremediation ◾ Electrosynthesis
◾ Calvin cycle ◾ Galvanic cell
◾ Chemical equation ◾ Hydrogenation
◾ Chemical looping combustion ◾ Membrane potential
◾ Citric acid cycle ◾ Nucleophilic abstraction
◾ Electrochemical series ◾ Organic redox reaction
◾ Electrochemistry ◾ Oxidative addition and reductive elimination
◾ Electrolysis ◾ Oxidative phosphorylation
◾ Electron equivalent ◾ Partial oxidation

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◾ Pro-oxidant ◾ Reduction potential

◾ Reduced gas ◾ Thermic reaction
◾ Reducing agent ◾ Transmetalation
◾ Reducing atmosphere

References
Notes

1. "Redox Reactions". wiley.com.

2. Haustein, Catherine Hinga (2014). K. Lee Lerner and Brenda Wilmoth Lerner, eds. Oxidation-reduction
reaction. The Gale Encyclopedia of Science. 5th edition. Farmington Hills, MI: Gale Group.
3. Bockris, John O'M.; Reddy, Amulya K. N. (1970). Modern Electrochemistry. Plenum Press. pp. 352–3.
4. Harper, Douglas. "redox". Online Etymology Dictionary.
5. Hudlický, Miloš (1996). Reductions in Organic Chemistry. Washington, D.C.: American Chemical Society.
p. 429. ISBN 0-8412-3344-6.
6. Hudlický, Miloš (1990). Oxidations in Organic Chemistry. Washington, D.C.: American Chemical Society.
p. 456. ISBN 0-8412-1780-7.
7. Electrode potential values from Petrucci, R. H.; Harwood, W. S.; Herring, F. G. (2002). General Chemistry
(8th ed.). Prentice-Hall. p. 832.
8. "gutier.doc" (PDF). Retrieved 2008-06-30. (2.76 MiB)
9. Robertson, William (2010). More Chemistry Basics. National Science Teachers Association. p. 82.
ISBN 978-1-936137-74-9.
10. Phillips, John; Strozak, Victor; Wistrom, Cheryl (2000). Chemistry: Concepts and Applications. Glencoe
McGraw-Hill. p. 558. ISBN 978-0-02-828210-7. "Students often are confused when associating reduction
with the gain of electrons."
11. Rodgers, Glen (2012). Descriptive Inorganic, Coordination, and Solid-State Chemistry. Brooks/Cole,
Cengage Learning. p. 330. ISBN 978-0-8400-6846-0.
12. Zumdahl, Steven; Zumdahl, Susan (2009). Chemistry. Houghton Mifflin. p. 160. ISBN 978-0-547-05405-6.

Bibliography

◾ Schüring, J., Schulz, H. D., Fischer, W. R., Böttcher, J., Duijnisveld, W. H. (editors)(1999).
Redox: Fundamentals, Processes and Applications, Springer-Verlag, Heidelberg, 246 pp. ISBN
978-3-540-66528-1 (pdf 3,6 MB) (http://hdl.handle.net/10013/epic.31694.d001)
◾ Tratnyek, Paul G.; Grundl, Timothy J.; Haderlein, Stefan B., eds. (2011). Aquatic Redox
Chemistry. ACS Symposium Series. 1071. doi:10.1021/bk-2011-1071. ISBN 9780841226524.

External links
◾ Chemical Equation Balancer
Wikiquote has quotations
(http://www.berkeleychurchill.com/software/chembal.php) related to: Redox
– An open source chemical equation balancer that handles
redox reactions.
Wikimedia Commons has
◾ Video – Synthesis of Copper(II) Acetate
media related to Redox
(https://www.youtube.com/watch?v=rF1ls-v7puQ) 20 Feb. reactions.
2009
◾ Redox reactions calculator (http://www.shodor.org/UNChem/advanced/redox/redoxcalc.html)

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◾ Redox reactions at Chemguide

(http://www.chemguide.co.uk/inorganic/redox/definitions.html#top)
◾ Online redox reaction equation balancer, balances equations of any half-cell and full reactions
(http://www.webqc.org/balance.php)

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