Download - Redox Reactions
Redox Reactions
A “redox” reaction involves the reduction and oxidation of the reactants, thereby changing the oxidation numbers of atoms taking part in the chemical reaction, through an exchange of electrons.
Examples of well-known redox reactions include the rusting of metal, the chemical reaction inside a battery, and combustion of hydrocarbons.
The roaring fire shown to the left is an example of the rapid oxidization of the hydrocarbons making up the wood and the reduction of the Oxygen gas from the air. The, very rusty, Iron hammer to the bottom right is also being oxidized by the Oxygen in the air, but at a much slower rate than the burning wood.
Oxidation Number
Rules for determining oxidation numbers:
1. The oxidation number of a neutral element is zero.
2. Fluorine always has an oxidation number of -1 in compounds.
3. The elements of groups IA (e.g., Na, K), IIA (e.g., Mg, Ca), and IIIA (e.g., Al, Ga) always have positive oxidation numbers of +1, +2, and +3, respectively, in compounds.
4. Hydrogen has an oxidation number of +1 in all its compounds except in binary and ternary compounds where the only other atoms are metals or boron.
5. Oxygen has an oxidation number of -2 in compounds, except for some compounds when combined with F, to which rules 2-4 apply. Oxygen, as the peroxide ion (O2
2-), has an oxidation number of -1.
6. The elements of groups VA, VIA, and VIIA have oxidation numbers of -3, -2, and -1, respectively, when found in binary compounds with metals or hydrogen.
Two processes exist which can change the oxidation number of an atom, namely Oxidation and Reduction. When a substance is oxidized its oxidation number increases, and when a substance is reduced its oxidation number decreases; oxidation and reduction are reverse processes of each other.
The change in the oxidation number of an atom is the result of an exchange of electrons with another substance, either a loss or a gain of electrons. Oxidation involves the loss of electrons by a substance while reduction involves a substance gaining electrons. A common mnemonic device which is useful in remembering this is the phrase, “LEO the lion says GER”,
Lose Electrons – OxidationGain Electrons – Reduction
Reduction / Oxidation reactions always occur in pairs such that when one substance is oxidized another substance is reduced.
Oxidizing / Reducing agents
Certain substances are more likely than others to be either oxidized or reduced due to how likely they are to give away or gain electrons. The chemical property which relates how likely a substance is to gain an electron is called its “electronegativity”. Since a completed octet of electrons in the outermost shell of an atom is most stable, highly electronegative atoms tend to gain electrons and become reduced while less-electronegative atoms tend to lose electrons and become oxidized.
Periodic Table showing the electronegativity of each element using the Pauling scale
Elements located on the left-most side of the periodic table have a high tendency to give up electrons to other atoms and be oxidized in order to
achieve a completed octet in their outermost shell. These, and other, substances which have a tendency to give up electrons are often referred to as reducing agents since, as they are oxidized, they reduce other substances in the process. Reducing agents are oxidized during a redox reaction. Substances which are strong reducing agents include the Alkali and Alkaline-Earth elements (for example, Lithium, Sodium, Calcium, …) which are located in the first two columns of the periodic table.
On the other side of the periodic table, located just to the left of the Nobel gasses, are a group of elements which have a high tendency to gain electrons from other substances and be reduced. The Halogens, for example Fluorine and Chlorine are strong oxidizing agents, as are the Chalcogens with Oxygen being a prime example. During redox reactions, these substances become reduced as they oxidize other substances and are known as oxidizing agents. Oxidizing agents are reduced during a redox reaction.
Other, polyatomic, oxidizing and reducing agents exist in addition to pure elements. Common oxidizing substances include salts containing the Nitrate (NO3
-), Chlorate (ClO3-), and Permanganate (MnO4
-) ion; for example, KNO3, KClO3, and KMnO4. Ascorbic acid (also known as Vitamin C) as well as Hydrogen gas, Carbon Monoxide, and Hydrocarbons can act as a reducing agent in some reactions.
Half Reactions
Like any chemical reaction, a redox reaction must be balanced by mass, but additionally must also be balanced by charge so that the reaction obeys the laws of conservation of mass and charge. Because of this, Reduction and Oxidation reactions always occur in pairs; if one substance is oxidized, another substance must be reduced, and therefore charge is always conserved.
A redox reaction can be broken up into two parts and analyzed separately, each called a half-reaction, one involving reduction and the other involving oxidation. Although individual half-reactions will contain free charge on either the reactant or product side, when a pair of balanced half-reactions are combined into a complete redox reaction it should contain no free charge since any electrons given up by the reducing agent will be gained by the oxidizing agent.
For example, Consider the redox reaction which takes place between Zinc metal and Hydrochloric acid.
Compare the oxidation numbers / states of the reactants to the products. Initially, the Zinc is in its elemental form and is defined to have an oxidization number of zero. On the products side of the reaction the Zinc is part of an ionic compound [Zinc(II) Chloride] and now has an oxidization number of +2. During the reaction, the Zinc atom lost two electrons and was oxidized to become the Zn+2 ion. Now look at the Hydrogen on the reactants side; initially the Hydrogen’s oxidization number was +1. On the products side, however, the Hydrogen is in its elemental form and has an oxidization number of zero. During the reaction, the Hydrogen was reduced. Clearly a redox reaction is taking place. Chlorine’s oxidization number did not change during the reaction; it is merely a spectator ion and is not involving in the redox process. With this new knowledge, we can write the ionic equation,
And we can infer the form of the two half-reactions to be,
In this case, the coefficients on the reactants / products were obtained from the coefficients from the full reaction, but this is not always the case since one might not know the full reaction to begin with, sometimes the coefficients must be altered to make sure both half-reactions are balanced by mass and charge. When the two half-reactions are combined, they should yield the full reaction and no longer contain any references to free electrons (they will cancel out since the same number of electrons will appear on both sides). In this reaction, the Zinc acts as the reducing agent, and is oxidized, and the Hydrogen ions act as the oxidizing agent, and are reduced.
Redox Reactions in Aqueous Solution
Other, arguably more complex, redox reactions may also occur when the reactants are in aqueous solution where the water itself takes part in the reaction. When construction / balancing the half reactions, it may be necessary to assume (due to the fact that the reactants are dissolved in water) that excess Hydrogen (H+) or Hydroxide (OH-) ions are present; the reaction will precede either under acidic or alkaline conditions. In this case, the water itself, or the ions it breaks into, becomes one of the reactants in the redox reaction even though it may not be entirely obvious initially.
For example,Consider the reaction between the Permanganate ion (MnO4
-) and the Ferrous, Iron +2, ion (Fe+2) in, acidic, aqueous solution.
In this case, the cation of the Permanganate compound and the anion of the Fe+2 compound are, unimportant, spectator ions. The Permanganate ion is a strong oxidizer and will act as the oxidizing agent in this case while the Fe+2 ion will act as the reducing agent.
Since the Fe+2 ion is the reducing agent, it is oxidized in the process to become the Fe+3 ion.
As the Iron ion is oxidized, the Permanganate ion must be reduced. When the Permanganate ion is reduced it forms the Mn+2 ion and water. Now the fact that this reaction takes place in acidic conditions becomes important. The Oxygen from the Permanganate ion combine with the excess of H+ ions from acid solution to form water, leaving Mn+2 behind in slightly less-acidic solution.
These are the two half reactions which are balanced by mass. When these two reactions are combined in the right proportions such that the reaction is also balanced by charge we will have our complete redox reaction.
For every 1 oxidization reaction of an Fe+2 ion, 1 electron is released. It takes 5 electrons (in combination with 8 Hydrogen ions) to reduce the Permanganate ion to Mn+2 and 4 water molecules. Therefore, the Iron-oxidizing reaction must proceed at 5 times the rate as the Permanganate-
reducing reaction. Taking this into account when combining the two half-reactions we find that the complete redox reaction is,
Unless we knew the reaction took place in acidic aqueous solution, it might be surprising to see that water is produced from this reaction even though no Hydrogen atoms are contained within the two most obvious reactants, namely the Fe+2 and Permanganate ions.
Redox Demonstrations
Aluminum’s capacity to act as a strong reducing agent is excellently demonstrated during a thermite reaction where it is used to reduce another, less reactive, metal oxide. The Aluminum is oxidized in the process, leaving the metal reduced to its elemental state and liberating a great deal of energy in the process.
Potassium Permanganate’s strong oxidizing nature is demonstrated in a reaction with glycerin. Shortly after the two substances are mixed, the Potassium Permanganate will automatically begin to rapidly oxidize / burn the glycerin in a very hot fire without even the need for external ignition.
Reducing Flames
Video
Good (0.5 MB)Better (1.1 MB)Best (6.5 MB)
The above video illustrates a flame’s ability to reduce an oxidized piece of metal; these flames are known as 'reducing flames'. In the above video, a
propane torch is used to heat a heavily-oxidized piece of Copper metal. Some of the propane fuel in the torch’s flame is not fully combusted due to insufficient Oxygen flow into the torch’s nozzle. The uncombusted fuel acts as the reducing agent for the Copper Oxide and one can see that, as the flame passes over a portion of the metal, bare (unoxidized) Copper becomes visible. When then flame is removed the hot Copper is again exposed to the Oxygen in the air and quickly oxidizes again, developing a black oxide layer.Flames which have excessive amounts of oxidizing gasses present will act to oxidize metal; such flames are known as 'oxidizing flames'.
RedoxFrom Wikipedia, the free encyclopedia
Illustration of a redox reaction
Redox (reduction-oxidation) reactions include 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.
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.[1] 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).
The two parts of a redox reaction
Rusting iron
A bonfire. Combustion consists of redox reactions involving free radicals.
Contents
[hide]
1 Etymology
2 Oxidizing and reducing agents
o 2.1 Oxidizers
o 2.2 Reducers
3 Standard electrode potentials (reduction potentials)
4 Examples of redox reactions
o 4.1 Displacement reactions
o 4.2 Other examples
5 Redox reactions in industry
6 Redox reactions in biology
o 6.1 Redox cycling
7 Redox reactions in geology
8 Balancing redox reactions
o 8.1 Acidic media
o 8.2 Basic media
9 Memory aids
10 See also
11 References
12 External links
[edit]Etymology
"Redox" is a portmanteau of "reduction" and "oxidation."
The word oxidation originally implied reaction with oxygen to form an oxide, since (di)oxygen 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.Lavoisier 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 was then 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 is its historical development.
The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction
and oxidation processes respectively when they occur at electrodes [2] . 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. Hydrogen is a primary or defining reducing
agent. But unlike oxidation, which has been generalized beyond its root element, hydrogenation has
maintained is specific connection to reactions which "add" hydrogen to another substance (i.e., the
hydrogenation of unsaturated fats into saturated fats, R-CH=CH-R + H2 = R-CH2-CH2-R).
[edit]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+.
[edit]Oxidizers
Substances that have the ability to oxidize other substances 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; i.e., it oxidizes other substances, and is thus itself reduced. And, because it "accepts"
electrons, it is also called an electron acceptor.
Oxidants are usually chemical substances with elements in high oxidation states (e.g., H2O2, MnO −
4, CrO3, Cr2O 2−
7, OsO4), or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra electrons by oxidizing
another substance.
[edit]Reducers
Substances that have the ability to reduce other substances 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; i.e., it reduces others, and is thus itself oxidized. And, because it "donates" electrons, it 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 awayelectrons 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 (H2) with a palladium, platinum, or nickel catalyst. These catalytic
reductions are used primarily in the reduction of carbon-carbon double or triple bonds.
[edit]Standard electrode potentials (reduction potentials)
Each half-reaction has a standard electrode potential (E0cell), which is equal to the potential difference
(or voltage) (E0cell) at equilibrium under standard conditions of an electrochemical cell in which
thecathode reaction is the half-reaction considered, and the anode is a standard hydrogen electrode where
hydrogen is oxidized: ½ H2 → H+ + e-.
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 which are weaker than H+ (e.g. –0.763 V for
Zn2+).[5]
For a redox reaction which takes place in a cell, the potential difference E0cell = E0
cathode – E0anode
Historically, 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 = E0
cathode + E0ox (anode)
[edit]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:
the oxidation reaction:
H2 → 2 H + + 2 e −
and the reduction reaction:
F2 + 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:
H2 + F2 → 2 HF
[edit]Displacement reactions
Redox occurs in single displacement reactions or substitution reactions. The redox
component of these types of reactions is the change of oxidation state (charge) on
certain atoms, not the actual exchange of atoms in the compounds.
For example, in the reaction between iron and copper(II) sulfate solution:
Fe + CuSO4 → FeSO4 + Cu
The ionic equation for this reaction is:
Fe + Cu2+ → Fe2+ + Cu
As two half-equations, it is seen that the iron is oxidized:
Fe → Fe2+ + 2 e−
And the copper is reduced:
Cu2+ + 2 e− → Cu
[edit]Other examples
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
The reduction of nitrate to nitrogen in the
presence of an acid (denitrification):
2 NO3− + 10 e− + 12 H+ → N2 + 6 H2O
Iron rusting in pyrite cubes
Oxidation of elemental iron to iron(III)
oxide by oxygen (commonly known
as rusting, which is similar to tarnishing):
4 Fe + 3 O2 → 2 Fe2O3
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.
[edit]Redox reactions in industry
The primary process of reducing ore to
produce metals is discussed in the article
on Smelting.
Oxidation is used in a wide variety of
industries such as in the production
of cleaning products and
oxidizing ammonia to produce nitric acid,
which is used in mostfertilizers.
Redox reactions are the foundation
of electrochemical cells.
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.[clarification needed]
[edit]Redox reactions in biology
Top: ascorbic acid (reduced form of Vitamin C)
Bottom: dehydroascorbic acid(oxidized
form of Vitamin C)
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
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 involve
s the reduction of carbon
dioxide into sugars and the
oxidation of waterinto
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 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 NAD + /NADH and NAD
P + /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 acetoacetat
e), 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
signaling involves the control of
cellular processes by redox
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.
[edit]Redox cycling
A wide variety of aromatic
compounds are enzymatically red
uced 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 suc
h as menadione.[6]
[edit]Redox reactions in geology
A uranium mine, near Moab,
Utah. Note alternating red and
white/green sandstone. This
corresponds to oxidized and
reduced conditions in
groundwater redox chemistry.
The rock forms in oxidizing
conditions, and is then
"bleached" to the white/green
state when a reducing fluid
passes through the rock. The
reduced fluid can also carry
uranium-bearing minerals.
In geology, redox is important to
both the formation of minerals,
mobilization of minerals, and in
some depositional environments.
In general, the redox state of
most rocks can be seen in the
color of the rock. Red is
associated with oxidizing
conditions of formation, and
green is typically associated with
reducing conditions. White can
also be associated with reducing
conditions. Famous examples of
redox conditions affecting
geological processes
include uranium
deposits and Moqui marbles.
[edit]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.
[edit]Acidic media
In acidic media, H+ ions and
water are added to half reactions
to balance the overall reaction.
For example,
when manganese(II) reacts
with sodium bismuthate:
Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4− (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)
[edit]Basic media
In basic
media, OH − ion
s and water
are added to
half reactions
to balance the
overall
reaction.
For example,
in the reaction
between potas
sium
permanganate
and sodium
sulfite:
Unbalanced reaction: KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOHReduction: 3 e− + 2 H2O + MnO4
− → MnO2 + 4 OH−
Oxidation: 2 OH− + SO32− → SO4
2− + H2O + 2 e−
Balancing
the
number of
electrons
in the two
half-cell
reactions
gives:
6 e− + 4 H2O + 2 MnO4− → 2 MnO2 + 8 OH−
6 OH− + 3 SO32− → 3 SO4
2− + 3 H2O + 6 e−
A
d
d
i
n
g
t
h
e
s
e
t
w
o
h
a
lf
-
c
e
ll
r
e
a
c
ti
o
n
s
t
o
g
e
t
h
e
r
g
iv
e
s
t
h
e
b
a
l
a
n
c
e
d
e
q
u
a
ti
o
n
:
2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH
[
edi
t]
Memory aids
Th
e
ke
y
ter
ms
inv
olv
ed
in
red
ox
are
oft
en
co
nfu
sin
g
to
stu
de
nts
.[7]
[8]
For
ex
am
ple
,
an
ele
me
nt
tha
t is
oxi
diz
ed
los
es
ele
ctr
on
s;
ho
we
ver
,
tha
t
ele
me
nt
is
ref
err
ed
to
as
the
red
uci
ng
ag
ent
.
Lik
ewi
se,
an
ele
me
nt
tha
t is
red
uc
ed
gai
ns
ele
ctr
on
s
an
d
is
ref
err
ed
to
as
the
oxi
dizi
ng
ag
ent
.[9]
Acr
on
ym
s
or
mn
em
oni
cs
are
co
m
mo
nly
us
ed[
10] t
o
hel
p
re
me
mb
er
wh
at
is
ha
pp
eni
ng:
"OI
L
RI
G"
—
O
xid
atio
n Is
Lo
ss
of
ele
ctr
ons
, R
ed
ucti
on
Is
Gai
n
of
ele
ctr
ons
.[7][8]
[10][9]
"LE
O
the
lion
say
s
GE
R"
—
Lo
ss
of
Ele
ctr
ons
is
Oxi
dati
on,
Gai
n
of
Ele
ctr
ons
is
Re
duc
tion
.[7][8]
[10][9]
"LE
OR
A
say
s
GE
RO
A"
—
Lo
ss
of
Ele
ctr
ons
is
Oxi
dati
on
(
Re
duc
ing
Ag
ent
)
an
d G
ain
of
Ele
ctr
ons
is
Re
duc
ed
(
Oxi
dizi
ng
Ag
ent
).[9]