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−

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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]