In simple words we can say that, "Inorganic chemistry is the the branch of chemistry that deals with inorganic compounds".
Definition of Inorganic
Chemistry- Inorganic
chemistry is the branch of chemistry, which deals with the study of inorganic
compounds and organometallic compounds. Inorganic chemistry includes study of all
chemical compounds excluding organic compounds (where, organic compounds are carbon
based compounds, generally having carbon-hydrogen bonds i.e. C-H bonds.
Example of Inorganic
Compound
Salt Magnesium Chloride (MgCl2)
is an ionicInorganic Compound which contains magnesium-cations Mg2+
and chloride-anions Cl−
Branches of Inorganic
Chemistry
Descriptive Inorganic Chemistry
Descriptive inorganic chemistry based on the arrangement and
ordering of compounds based on their properties.
Coordination Chemistry
Coordination chemistry is the branch of Inorganic chemistry which deals with
the study of coordination compounds or coordination complexes. A coordination
compound or coordination complex contains of a central atom or ion, which is generally
metallic and is known as coordination center, and get surrounded by
group of bounding molecules or ions, these surrounding molecules and ions are
known as ligands.
Materials Science
Material science deal with the study of materials. Study of metals,
semiconductors, polymers, ceramics, nanomaterials, biomaterials and many other
form of materials comes under material science. A material scientist study
about any material to find out its structure, properties and use this in
certain applications.
Bioinorganic Chemistry
Bioinorganic Chemistry is the branch of inorganic chemistry which deal with
study of different metals in biological system. Bioinorganic chemistry deal in
both natural field (like study and application of metalloproteins etc.) and
artificial field (like study, application and development of medicines). Bioinorganic
chemistry is the mixture of biochemistry
and inorganic chemistry.
Organometallic Chemistry
Organometallic chemistry is the branch of inorganic chemistry
which deal with the study of chemical compounds having at least one bond
between a Metal and carbon atom of organic compound. Organometallic compounds
are broadly used in research, industrial chemical reactions, catalysis to
increase the rates of reactions.
Cluster Chemistry
Cluster chemistry is the branch of inorganic chemistry which deals with the
study of clusters. Cluster forms when atoms or molecules bound together to form
large size substances whose size is intermediate in between the size of a molecule
and a bulk solid. Example of cluster- fullerene cluster from carbon atoms and
borane cluster from boron atoms.
Hydrogen is the first element in the
periodic table and is the lightest element. It exists as a diatomic molecule (H2)
and is called dihydrogen. It has one proton and one electron and it has electronic
configuration 1s1. This electronic configuration is responsible for
dual nature of hydrogen i.e. hydrogen shows resemblance with alkali metals as
well as halogens.
Resemblance of hydrogen with alkali
metals
Like alkali metals –
·Hydrogen exhibits electropositive character i.e. it has
tendency to lose electron.
·The normal valency of hydrogen is 1.
·Hydrogen forms halides, oxides and sulphides.
·Hydrogen acts as strong reducing agent.
Resemblance of hydrogen with halogens
Like
halogens-
Hydrogen exhibits electronegative character by accepting one
electron to complete its valance shell.
Hydrogen shows same order of ionisation energy (IE1)
as that of halogens.
Hydrogen (H2) is a gas in molecular state similar
to F2 and Cl2.
Hydrogen shows -1 oxidation state in its compounds with more
electropositive element.
POSITION OF HYDROGEN IN THE PERIODIC TABLE
As
the hydrogen resembles both the alkali metals and the halogens and also differs
from them. Also hydrogen exhibits some unique behavior therefore the position
of hydrogen is uncertain and it is placed separately in periodic to table.
ISOTOPES OF HYDROGEN
Hydrogen
has three isotopes, protium (11H), deuterium (12H)and tritium (13H).
i.Protium or ordinary hydrogen (11H).
It is the most abundant isotope of
hydrogen (99.985%). Its nucleus has one proton and no neutron (mass no. = 1).
ii.Deuterium or heavy
hydrogen (12H or D).
It is less abundant and is present in heavy water (D2O). Its
nucleus has one proton and one neutron (mass no. = 2).
iii.Tritium (13H or T)
It is the rarest isotope of hydrogen and is radioactive in nature. Its
nucleus has one proton and two neutrons (mass no. = 3).
PREPARATION OF DIHYDROGEN (H2)
Laboratory methods
1.Action of dilute HCl or H2SO4
on granulated zinc.
Zn+2HCl→ZnCl2+H2↑
2.Action of dilute H2SO4
on magnesium ribbon.
Mg+H2SO4→MgSO4+H2↑
3.Action of water on sodium
hydride.
NaH+H2O→NaOH+H2↑
4.By action of KOH of scrap aluminum
or silicon. (Uyeno’s method)
2Al+2KOH+2H2O→2KAlO2+3H2↑
Commercial method
1) From water gas (Bosch process)
Water gas (CO + H2) is
mixed with the steam and the gaseous mixture is passed over heated catalytic
mixture of ferric oxide (Fe2O3) and chromium oxide (Cr2O3)
at 773K.
H2+CO+H2OFe2O3 + Cr2O3 (773K)→CO2+2H2
Water gasSteam
2) From Steam (Lanes process)
When
superheated steam is passed over iron filling heated to 1023 – 1073K hydrogen
is formed.
3Fe+4H2O(steam)(1023 – 1073)K (Fe)→Fe2O4+4H2(g)
3) Hydrocarbon steam process
Hydrogen
is prepared by the action of steam of hydrocarbon at 1270K.
CH4+H2O1270K→CO+3H2
4) By electrolysis of water
Electrolysis
of acidified water using platinum electrodes gives dihydrogen. Here dihydrogen
is librated at the cathode while dioxygen is librated at anode.
CHEMICAL PROPERTIES OF DIHYDROGEN
1)
Dihydrogen reacts with halogens (X2) to give hydrogen halide (HX).
H2+X2→2HX( X = F, Cl, Br, I )
2)
Dihydrogen reacts with dioxygen to forms water. This reactions is highly
exothermic.
2H2+O2catalyst or heating→2H2O . ฮH = -285.9 KJ.mol-1
3)
With dinitrogen dihydrogen forms ammonia. This process is coiled Hober process
of
manufacture of ammonia.
3H2+N2673K, Fe catalyst, 200atm→2NH3
4)
Dihydrogen combines with many metals at high temperature to yield the
correspondinghydrides
2Na+H2→2NaH(sodium hydride)
5)
It acts as reducing agent and thus reduces certain oxides of metals.
ZnO+H2→Zn+H2O
Fe2O4+4H2→3Fe+4H2O
6)
Hydrogenation of vegetable oils using nickel as catalyst gives edible facts
(margarine and Vanaspati
ghee)
·Dihydrogen is used in the preparation of ammonia by Haber’s
process.
·It is used in the hydrogenation of vegetable oils.
·It is used for the manufacture of metal hydrides.
·It is used as rocket fuel in space research.
·Dihydrogen is used in fuel cell for generating electrical
energy.
·It is used in the atomic hydrogen torch and oxy hydrogen
torches for cutting and welding.
HYDRIDES
The
binary compounds of hydrogen with other elements are called hydrides. These
hydrides have the formula EHx or EmHn (E =
Element). These are classified into three types.
a)
Ionic or saline hydrides
b)
Covalent or molecular hydrides
c)
Metallic or nonstoichiometric or interstitial hydrides
a) Ionic or saline hydrides
These
are formed by the combination of hydrogen with metals which have low electro
negativity values and are electropositive with respect to hydrogen. These
includes elements of S block element. Ionic hydrides are prepared by the direct
combination of metals.
e.g.
LiH, KH, CaH2, MgH2 etc.
b) Covalent or molecular hydrides
These
hydrides are formed by the combination of elements of comparatively higher
electro –negativity as p-block element. The bonds formed are mostly covalent in
character. Covalent hydrides are generally volatile.
e.g.
H2O, CH4, NH3, HF etc.
c) Metallic or nonstoichiometric or interstitial hydrides
Metallic
hydrides are formed by many d - block and f - blockelements. Elements of group 7, 8, 9 of d –
block do not form hydrides and this is referred as hydride gap. Metallic
hydrides are non – stoichiometric and show electric conductance. In these
hydrides, hydrogen occupies interstices in the metal lattice producing
distortion without any change in its type. Therefore, they are termed as
interstitial hydrides.
e.g.
SCH2, TiH2, VH, ZrH2
WATER (H2O)
Water
is essential to all forms of life. It is most common, abundant and easily
obtainable of all chemical compounds. It is regarded as universal solvent.
Water (H2O) is hydride of oxygen. In nature, water exist in three
physical states, water (liquid), Ice (Solid) and water vapour (gas).
Structure of water
Water
molecule is a bent molecule with bond angle of 104.5แต and O-H bond length is
95.7pm.
Structure of Ice
Ice
is the solid form of water. It has a highly ordered three dimensional hydrogen
bonded structure. In ice each oxygen atom is surrounded tetrahedrally by four
other hydrogen atoms at a distance of 276pm. The strength of hydrogen bonding
in ice is more than liquid water. There is empty space in crystal of ice due to
hydrogen bonding. This makes density of ice lower than liquid water, hence ice
floats on water.
Chemical properties of water
1) Dissociation of water
Water
dissociates into its ion.
H2O+H2O→H3O++OH-
Hydronium ion
2) Amphoteric nature
Water
has the ability to act as on acid as well as base. Such kind of behavior is
known as amphoteric nature.
H2O+NH3→OH-+NH4+
H2O+HCl→H3O++Cl-
3) Oxidizing and reducing nature
Water
can act both as an oxidizing and reducing agent in its chemical reactions. With
active metals water behaves as an oxidizing agent while with highly
electronegative element, it acts as a reducing agent.
2Na+2H2O→2NaOH+H2
Oxidizing agent
2F2+2H2O→4HF+O2
Reducing agent
4) Formation of Hydrates
Water
has strong ability to form compounds with some metal salts known as hydrates. This
hydrates can be classified into three types:
i) Coordinated water
[Cr(H2O)6]3+3Cl-
ii) Interstitial water
BaCl2.2H2O
iii) Hydrogen – bonded water
[Cu(H2O)4]2+SO42-.H2Oin
CuSO4.5H2O
HEAVY WATER
Heavy
water is deuterium oxide (D2O). It was discovered by Urey in 1932.
The reaction of D2O are slightly slower than H2O. Ordinary
water contains one part of heavy water in 6000 part of it. It is used as a
moderator in nuclear reactors. It is used as tracer compound in determining the
mechanism of many organic reactions.
HARD AND SOFT WATER
Natural
water contains dissolved salts. Depending upon its behavior towards soap
solution, water may be classified as hard water and soft water.
a) Soft water :
Water which produces lather with soap solution readily is
called soft water. For example : distilled water, rain water, and
demineralizedwater.
b) Hard water :
Water which does not produces lather with soap solution
readily is called hard water. For example : sea water, river water, well water
and tap water.
The
hardness of water is due to the presence of the bicarbonates, chlorides and
sulphates of calcium and magnesium.
Isotopic varieties of water
Ordinary
water contains 18 different kinds of water. Such as variety is possible due to
the different isotopic forms of hydrogen and oxygen. Hydrogen has three
isotopes H(protium), D(Deutrium), T(tritium) and oxygen also has three isotopes
16O, 17O, 18O. These isotopes of hydrogen and
oxygen combine to give 18 different kinds of water from which H2O is
most abundant.
HYDROGEN PROXIDE (H2O2)
Hydrogen
peroxide was discovered by Thenard in 1818. It is an important chemical used in
pollution control treatment of domestic and industrial effluents.
Preparation of H2O2
1)
In laboratory H2O2 is prepared by acidifying barium
peroxide and removing excess water by evaporation under reduced pressure.
BaO2.8H2O+H2SO4→BaSO4+H2O2+8H2O
2)
By the action of dilute acid on sodium peroxide (Merck’s method)
Na2O2+H2SO4→NaSO4+H2O2
3)
By bubbling CO2 through a paste of BaO2
BaO2+H2O+CO2→BaCO3 ↓+H2O2
4)
By the action of phosphoric acid on BaO2
3BaO2+2H3PO4→Ba3(PO4)2
↓+3H2O2
5)
Industrially H2O2 is prepared by the auto – oxidation of
2 – Ethyl anthraquinol.
It is non–planar open book (Skew) structure. The bond length
and bond angle are slightly different in gas and solid phase due to hydrogen bonding. Structure of H2O2 HYDROGEN PROXIDE.
Chemical Properties of H2O2
1)
It decomposes rapidly on heating in presence of finely divided metals, such as
Ca, Fe, Cu, Au, Ag, Pt, MnO2, Carbon, dust, light, etc.
H2O2+H2O2→2H2O+O2(ฮH = -196KJ)
2)
It acts as an oxidizing as well as reducing agent in both acidic and alkaline or
basic medium. Some example are given below
i)
Oxidizing action in acidic medium
PbS+4H2O2→PbSO4+4H2O
ii)
Reducing action in acidic medium
HOCl+H2O2→H3O++Cl-+O2
iii)
Oxidizing action in alkaline medium
Mn2++H2O2→2Fe3++2OH-
iv)
Reducing action in alkaline medium
Cl2+H2O2+2OH-→2Cl-+2H2O+O2
3)
Hydrogen peroxide acts as a bleaching agents due to the release of nascent
oxygen.
H2O2→H2O+O
4)
Hydrogen peroxide undergoes addition reaction with alkenes to form glycols.
Uses of H2O2
·It is widely used in environmental chemistry.
·It is used as a antiseptic and is sold in the market as
perhydrol.
·It is used in daily life as a hair bleach and as mild
disinfectant.
·It is used as a bleaching agent and to manufacture chemicals
like sodium perborate and percarbonate.
Chemical reactions are generally
classified into three types:
1) Precipitation reactions
2) Acid – base neutralization
3) Redox reactions
Redox reaction is a very important
group of reactions. These reactions are involved in large number of processes
in nature, biological and industrial.
Definition of Redox Reaction
Redox or reduction -oxidation reactions are the chemical reactions in which the oxidation state of atoms is get changed. → Redox reactions are conducted by the transfer of electrons from one chemical species (reducing agent) undergoing oxidation (losing electrons) to another chemical species (oxidizing agent) undergoes reduction (gains electrons).
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CLASSICAL CONCEPT OF OXIDATION AND
REDUCTION
Oxidation is a reaction which involves the
addition of oxygen or any other electronegative element or the removal of
hydrogen or any electropositive element from the substance.
For example,
C(s) + O2(g) → CO2(g) (oxidation of carbon)
Reduction reaction involves the addition of
hydrogen or any electropositive element or removal of oxygen or any electronegative
element. It is just the reverse of oxidation.
For example,
Cl2 + H2 →
2HCl (reduction of Cl)
ELECTRONIC CONCEPT OF OXIDATION AND
REDUCTION
OXIDATION
An oxidation is defined as the loss
of one or more electron from a substance.
For example,
Na → Na+ + e-
REDUCTION
A reduction is defined as the gain of
one or more electrons by a substance.
For example,
Cl2 + 2e- → 2Cl-
OXIDIZING AGENT OR OXIDANT
Oxidizing agent is a substance which
accepts or gain electrons and causes the oxidation of other substance.
REDUCING AGENT OR REDUCTANT
Reducing agent is a substance which
donates electrons and causes reduction of other substance.
REDOX REACTION (oxidation – reduction
reaction)
The reactions in which oxidation and
reduction reactions occur simultaneously are called redox reaction.
In redox reactions, electrons are
transferred from one substance to the other substance, therefore these
reactions are electron transfer reactions.
For example,
Loss of 2e- (oxidation)
| ↓
Zn
+ Cu2+ →
Zn2+ + Cu
| ↑
Gain of 2e- (Reduction)
OXIDATION NUMBER OR OXIDATION STATE
Oxidation number of an element in a
compound is defined as the number of electrical charges it carries when all
other atoms are removed from it as ions.
Rules to assign oxidation numbers
1)The oxidation number of an
atom in free uncombined elemental state is zero.
For example,
H2, Ca, Cl2,
O3, S8, P4, and so on has oxidation number
zero.
2)The oxidation number of an
atom in a monoatomic ion is equal to its charge.
For example,
K+ Ba2+ Cr3+ Br- S2-
↑ ↑
↑ ↑ ↑
+1 +2
+3 -1 -2
3)When hydrogen atom is
bounded to non – metals, its oxidation number is +1. And when it is bounded to
metals, its oxidation number is -1.
For example,
[O – H]- H – O – H
Li – H
H – Ca – H
-2 +1
+1 -2 +1
+1 -1
-1 +2 -1
4)The oxidation number of
oxygen is usually -2 in all of its compounds except in peroxide or peroxide ion
where it has oxidation number of -1.
For example,
Ca – O H – O – O – H [O – O]2-
+2 -2
+1 -1 -1 +1
-1 -1
5)The oxidation number of F
is -1 in all of its compounds. The other halogens Cl, Br and I usually exhibit
oxidation number of -1 in their halide compounds. However when Cl, Br and I are
bonded to oxygen then they exhibit oxidation number of +1.
For example,
H – F K – Br Cl – O – Cl H – O – Cl
+1 -1 +1
-1 +1 -2
+1 +1 -2
+1
6)The algebraic sum of the
oxidation number of all atoms in a neutral molecule is zero and for the
polyatomic ion it is equal to net charges of the ion.
For example,
In H – Cl
Oxidation number of H +
oxidation number of Cl = 0
= +1
-1 = 0
Using
above rules we can calculate oxidation number of an atom in any molecule or ion
Problems
Determine the oxidation number of
a)N in HNO3
HNO3 is neutral
molecule
: . Sum of oxidation number of all atom of HNO3
is equal to zero.
[ O.N. of H ] + [
O.N. of N ] + [ 2 X O.N. of O ] = 0
= +1
+ [O.N. of N ] + [ 2
X (-2) ]
= 0
= +1
+ [O.N. of N ] -
4 = 0
= O. N
of N - 3
= 0
O. N
of N = +
3
: . Oxidation number of N in HNO3 is
+3.
b)Pt in PtCl62-
PtCl62- is ionic species.
: . Sum of oxidation number of all atom of PtCl62- is equal to -2.
[ O.N. of Pt ] + [
O.N. of Cl ] X 6 = -2
= [O.N. of Pt ]
+ [ 6 X (-1) ]
= -2
= O.N. of Pt
= - 2 + 6 = + 4
: . Oxidation number of Pt in PtCl62- is + 4.
KEY POINTS
Oxidation
·Loss of electron.
·Increase in oxidation number of oxidized species.
Reduction
·Gain of electron.
·Decrease in oxidation number of reduced species.
Redox reaction
·Oxidation and reduction occur simultaneously together.
·Simultaneous loss and gain of electron.
·Simultaneous increase and decrease in the oxidation number.
Oxidizing agent (oxidant)
·Causes oxidation of other.
·Accept electron.
·It self undergoes reduction.
Reducing agent (Reductant)
·Causes reduction of other.
·Donates electron.
·It self undergoes oxidation.
BALANCING OF REDOX REACTIONS
Two
important methods are used for systematic balancing of redox reactions.
A] Oxidation number method
Step I
Write
the unbalanced net equation for the redox reaction. Assign the oxidation number
to all the atoms in both, the reactant and the product.
Step II
Identify
the atoms that have changed oxidation number. Draw bracket to connect atoms of
the elements that are oxidized and reduced.
Step III
Calculate
the increase or decrease in the oxidation number per atom of molecule. If
increase and decrease in oxidation number is not equal, then equalize them by
multiplying with a suitable number.
Step IV
Balance
hydrogen and oxygen. For balancing oxygen atoms add water molecules to the side
containing less oxygen atoms. Balancing of hydrogen atoms depends upon the
medium acidic or basic.
i)
For acidic solutions, add H+ ions to the side deficient in hydrogen
atoms.
ii)
For basic solutions, add H2O molecule to the side deficient in
hydrogen atoms and simultaneously add equal number of OH- ions on
the other side of the equation.
Step V
Check
the balance equation to make sure that the reaction is balanced with respect to
both the number of atoms of each element and charge.
Illustrative example
Balance the following redox equation by oxidation number method.
H2SO4(aq) + C(s) →
CO2(g) + SO2(g) + H2O(l)
Solution :
Step I
The
oxidation number of all the atoms are as follows :
H2SO4 +
C → CO2 +
SO2 + H2O
↑ ↑
+1
+6 -2 0 +4 -2 +4 -2 +1 -2
Step II
Less of 4e-(oxidation)
| ↓
H2SO4 +
C → CO2 +
SO2 + H2O
+6 0 +4
+4
| ↑
Gain of 2e- (Reduction)
Step III
Increase
in oxidation number :
C
(O) →
C (+4), net increase = +4
Decrease
in oxidation number :
S
(+6) →
S (+4), net decrease = -2
To
make total increase and decrease equal the net decrease must be multiplied by
2. Hence the coefficient 2 is required on both the side for S.
2 H2SO4 +
C → CO2 +
2SO2 + H2O
Step IV
Balance
the equation for O atoms by adding H2O molecule to the side with
less O atoms.
2
H2SO4 + C
→ CO2 +
2SO2 + H2O
H
atoms have already been balanced.
Step V
In
reaction the number of atoms and charge is balanced, thus the balanced reaction
is
2
H2SO4(aq) + C(g) →
CO2(g) + 2SO2(g) + H2O(l)
B] lon-electron method(Half reaction method)
Step I
Write
the unbalanced equation for the redox reaction and assign the oxidation numbers
to all the atoms in the reactants and the products.
Step II
Divide
the equation into two half equation. One half is oxidation in which oxidation
number of oxidized species increases and other half is reduction in which
oxidation number of reduced species decreases.
Step III
Balance
the atoms except O and H in each half equation balance O atom by adding H2O
to the side with less O atoms.
Step IV
Balance
H atoms by adding H+ ions to
the side with less H atoms.
Step V
Balance
the charges by adding appropriate number of electrons to the right side of
oxidation half equation and to the left of the reduction half equation.
Step VI
Multiplying
half equations by suitable factors to equalize the number of electrons in the
two half equations. Then add two half equations and cancel the number of
electrons on both the sides of the equation.
Step VII
If
the reaction occurs in basic medium then add OH- ions, equal to the
number of H+ ions, on both the side of the equation, H+ and
OH- ion on same side combine to give H2O molecule.
Step VIII
Check
the equation is balanced for both, the atoms and charges.
Illustrative example
Balance
the following redox equation by ion electron method. The reaction occurs in
acidic medium.
Mn2+(aq) + ClO-3(aq) →
MnO2(s) + ClO2(aq)
Solution :
Step I
Write
the unbalanced equation and assign oxidation numbers to all atoms.
Gain of 2e- (Reduction)
| ↓
Mn2+(aq) + ClO-3(aq) →
MnO2(s) + ClO2(aq)
+2 +5 -2 +4 -2
+4 -2
| ↑
Less
of 4e-(oxidation)
Step II
Divide
the equation into two half equation such that one half is oxidation and other
half is reduction.
Oxidation
half equation :
Mn2+ →
MnO2
Reduction
half equation :
ClO-3 →
ClO2
Step III
Balance
the each half equations for O atoms by adding H2O to the side with
less O atoms.
Oxidation
: Mn2+ +
2H2O → MnO2
Reduction
: ClO-3 →
ClO2 + H2O
Step IV
Balance
H atoms by adding H+ ions to the side with less H atoms
Oxidation
: Mn2+ +
2H2O → MnO2 + 4H+
Reduction
: ClO-3 + 2H+ →
ClO2 + H2O
Step V
Balance
the charge on both side of equation by adding appropriate number of electron.
Oxidation
: Mn2+ +
2H2O → MnO2 + 4H+ +
2e-
Reduction
: ClO-3 +
2H+ + e- →
ClO2 + H2O
Step VI
Multiply
reduction half equation by 2 to equalize the number of electrons in two half
equations. Then add two half equations.
Mn2+ +
2H2O → MnO2 + 4H+ + 2e-
2ClO-3 +
4H+ + 2e- →
2ClO2 + 2H2O
Mn2+ +
2ClO-3
→ MnO2 +
2ClO2
Step VII
The
reaction has equal number of atoms and charge, thus the reaction is balanced.
Mn2+ +
2ClO-3
→ MnO2 +
2ClO2
PROBLEMS FOR PRACTICE
Balance the following redox equations by oxidation number method.
a)
Ag +
NO-3 → Ag+ + NO2 (Acidic medium)
b)
Fe2+ + BrO-3 → Fe3+ + Br- (Acidic Medium)
c)
MnCl2 + HO-2 →
Mn(OH)3 + Cl- (Basic Medium)
d)
Mn2+ + H2O2 → MnO2 + H2O (Basic Medium)
Balance the following redox equations by ion - electron method.
a)
SO2 + Fe3+ → Fe2+ + SO42- (Acidic medium)
b)
2HgO →
2Hg + O2 (Acidic Medium)
c)
S2O32-
+ I2 → S4O62- + I- (Basic Medium)
d)
SeO32- + Cl2 → SeO42- + Cl- (Basic Medium)
APPLICATIONS OF REDOX REACTIONS
1) Combustion
Burning
of a substance by oxidation with oxygen in air is called combustion. Combustion
involves redox reactions.
2) Bleaching
Decolourization
or lightening of coloured materials uses redox reaction and is called bleaching.
3) Batteries
The
electricity produced in batteries or galvanic cells is due to redox reactions occurring
in them.
4) Metallurgy
The
extraction and purification of metals uses redox reactions in different steps.
5) Corrosion
The
corrosion is the destruction of metals by oxidation.
For example,
Rusting
of iron is its oxidation by oxygen of air in presence of moisture.
4Fe +
3O2 H2O →
2Fe2O3 . H2O
Corroded
iron
6) Respiration
The
process of breathing and using oxygen is biological redox reactions and is
called respiration. you read these notes on CHEMISTRYNOTESINFO & this is 11th class chemistry notes.
