Chemical reactions and equations 10 chm(1)

BY: VEENU GUPTA
(PGT CHEMISTRY)
APS RAKHMUTHI
6/5/2020BY :VEENU GUPTA(CHEM)

 The transformation of chemical substance into a new
chemical substance by making and breaking of bonds
between different atoms is known as Chemical
Reaction.
 For example:
Rusting of iron, the setting of milk into curd, digestion
of food, respiration, etc.
 A chemical reaction involves a chemical change in which
substances react to form new substances with entirely new
properties.
 During a chemical reaction, there is a breaking of bonds
between atoms of the reacting molecules to give
products.

 The burning of magnesium in the air to form
magnesium oxide(white ash) is an example of a
chemical reaction.
2Mg(s) + O2(g) →2MgO(s)[Basic in nature]
 Before burning in air, the magnesium ribbon is
cleaned by rubbing with sandpaper. This is
done to remove the protective layer of basic
Magnesium carbonate from the surface of the
magnesium ribbon.
 2Mg(s) + O2(g)+ CO2(g) →2MgCO3(s)

Reaction: 2Mg(s) + O2(g) →2MgO(s)
 Substances which take part in a chemical
reaction are called reactants.
Example: Mg and O2.
 New substance formed after a chemical
reaction is called a product.
Example: MgO.

 EVOLUTION OF GAS
 CHANGE IN COLOUR OF SOLUTION
 FORMATION OF PRECIPITATES
 CHANGE IN STATE OF SUBSTANCE
 CHANGE IN TEMPERATURE

The chemical reaction between zinc and
dilute sulphuric acid is characterised by
the evolution of hydrogen gas.
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) ↑

The gas evolved after reaction of acid
with metal can be tested by bringing a
lighted candle near it.
If the gas burns with a pop sound, then
it confirms the evolution of hydrogen
gas.
Burning with pop sound is the
characteristic test for hydrogen gas.

 When Iron reacts with copper sulphate , it forms Iron
sulphate and copper metal and the colour of solution
changes from blue to green and there is no formation of
ppts.
Fe (s) + CuSO4 (aq) → FeSO4 (aq) + Cu (s)
(Blue colour ) (green colouration)
 The chemical reaction between citric acid and purple
coloured potassium permanganate solution is
characterised by a change in colour from purple to
colourless.
 The chemical reaction between sulphur dioxide gas and
acidified potassium dichromate solution is
characterized by a change in colour from orange to
green.

 When the solution of barium chloride reacts with the
solution of sodium sulphate, white precipitates of
barium sulphate are formed along with sodium chloride.
BaCl2 (aq) + Na2SO4(aq) → BaSO4 (s) + 2 NaCl (aq)
(White ppt)
 When the solution of Lead Nitrate reacts with the
solution of Potassium Iodide, Yellow precipitates of Lead
Iodide are formed along with Potassium Nitrate.
Pb(NO3)2 (aq) + 2 KI (aq) → PbI2 ( ↓ ) + 2 KNO3 (aq)
(yellow ppt)

 When the solution of Silver Nitrate reacts
with the solution of sodium chloride then
white ppts of silver chloride are formed along
with sodium Nitrate
 AgNO3 (aq) + NaCl (aq) → AgCl ( ↓ ) + NaNO3 (aq)
(white ppt)
 When the solution of Silver Nitrate reacts
with the solution of sodium bromide then
Yellow precipitates of silver bromides are
formed along with sodium Nitrate
 AgNO3 (aq) + NaBr (aq) → AgBr ( ↓ ) + NaNO3 (aq)
(yellow ppt)

 The chemical reaction between sulphuric acid
and barium chloride solution is characterised
by the formation of a white precipitate of
barium sulphate.
BaCl2(aq) + H2SO4(aq) → BaSO4(s) + 2HCl(aq)
(white ppt)

 The combustion reaction of candle wax is
characterised by a change in state from solid to liquid
and gas (because the wax is a solid, water formed by
the combustion of wax is a liquid at room
temperature whereas, carbon dioxide produced by
the combustion of wax is a gas).
 There are some chemical reactions which can show
more than one characteristics.

(IV) CHANGE IN
TEMPERATURE

The chemical reaction between quick lime
& water to form slaked lime is
characterized by rise in temperature .
Ca0 + H2O → Ca(OH)2
 Respiration is also a chemical reaction
which involves increase in temperature.
 Most of the combinations reactions
occurs by an increase in temperature of
the solution.

 The chemical reaction between zinc granules
and dilute sulphuric acid is also characterised
by rise in temperature.
Zn + 2H2SO4 → ZnSO4 + H2
 The chemical reaction between zinc granules
and dilute Hydrochloric acid is also
characterised by a rise in temperature.
Zn + 2HCl → ZnCl2 + H2

i) Thermal Decomposition of Lead Nitrate
2Pb(NO3)2 (s) + Heat → 2PbO(s) + 4NO2(g) + O2(g)
white brown gas
ii) Thermal Decomposition of Ferrous Sulphate
2 FeSO4 (s) + Heat → Fe2O3 (s) + SO2 (g) + SO3(g)
Ferrous Sulphate Ferric oxide Sulphur Sulphur
(green) (brown) dioxide trioxide
iii) Thermal Decomposition of calcium carbonate
CaCO3(s) + Heat → CaO(aq) + CO2 (g)
(Calcium Carbonate) ( calcium oxide)

 The symbolic representation of a chemical reaction is called a
chemical equation. It is also called as representation of chemical
reaction by means of symbols of substances in the form of
formulae.
 Example: A + B → C + D
In this equation, A and B are called reactants and C and D are
called the products.
 The arrow shows the direction of the chemical reaction.
Condition, if any, is written generally above the arrow.

 The reactants are written on the left hand side with a
plus sign between them.
 The products are written on the right hand side with
a plus sign between them.
 An arrow separates the reactants from the products.
The arrow head points towards the products and
indicates the direction of the reaction.
 Condition, if any, is written generally above the
arrow

 When hydrogen reacts with oxygen, it gives water. This
reaction can be represented by the following chemical
equation:
i) Hydrogen + Oxygen → Water
ii) H2 + O2 → H2O
In the first equation, words are used and in second,
symbols of substances are used to write the chemical
equation.
 For convenience, the symbol of substance is used to
represent chemical equations.
 A chemical equation is a way to represent the chemical
reaction in a concise and informative way.

 A chemical equation which simply represents the
symbols and formulas of reactants and products
taking part in the reaction is known as skeletal
chemical equation for a reaction.
 For example:
For the burning of Magnesium in the air,
Mg + O2 → MgO is the skeletal equation.

 A chemical equation can be divided into two types:
 Balanced Chemical Equation and
 Unbalanced Chemical Equation.

 A balanced equation is a chemical equation in which
number of atoms of each element is equal on both
sides of the equation i.e
number of atoms of an element on reactant side =
number of atoms of that element on the product
side.
 Example: Zn + H2SO4 → ZnSO4 + H2
 In this equation, numbers of zinc, hydrogen and
sulphate are equal on both sides, so it is a Balanced
Chemical Equation.

 If the number of atoms of each element in reactants is
not equal to the number of atoms of each element
present in the product, then the chemical equation is
called Unbalanced Chemical Equation.
Example: Fe + H2O → Fe3O4 + H2
 In this example, a number of atoms of elements are not
equal on two sides of the reaction.
 For example; on the left-hand side only one iron atom is
present, while three iron atoms are present on the
right-hand side. Therefore, it is an unbalanced chemical
equation.

The process of equating the
number of atoms on both the sides
of a chemical equation is known
as balancing of a chemical
equation.

 write the number of atoms of each element present on the left
hand side and right hand side.
 Always start balancing with the compound that contains
maximum number of atoms. It can be reactant or a product.
Then in that compound select the element which has the
maximum number of atoms.
 While balancing a chemical equation, the molecular formulas of
the reactants and products should not change. The molecular
formulas are simply multiplied by suitable coefficients.
 To make a chemical equation more informative, the reaction
conditions such as temperature, pressure or catalyst are
written on the arrow separating the reactants and products.
 The evolution of gas is indicated by an upward arrow.
 The formation of precipitate is indicated by a downward arrow.
 Heat evolved during the reaction is written as + Heat on the
product side.
 Heat absorbed during the reaction is written as + Heat on the
reactant side.

Fe + H2O → Fe3O4 + H2
 Write the number of atoms of elements present in reactants and in
products in a table as shown here.
 Balance the atom which is maximum in number on either side of a
chemical equation. In this equation, the number of oxygen atom is the
maximum on the RHS.
 To balance the oxygen, one needs to multiply the oxygen on the LHS by
4, so that, the number of oxygen atoms becomes equal on both sides.
Fe + 4 H2O → Fe3O4 + H2
Name of atom No. of atoms in the
reactant
No. of atoms in the
product
Iron 1 3
Hydrogen 2 2
Oxygen 1 4

 Now, the number of hydrogen atoms becomes 8 on the LHS, which is
more than that on the RHS. To balance it, one needs to multiply the
hydrogen on the RHS by 4.
Fe + 4 H2O → Fe3O4 + 4 H2
 After that, the number of oxygen and hydrogen atoms becomes equal
on both sides. The number of iron is one on the LHS, while it is three on
the RHS. To balance it, multiply the iron on the LHS by 3.
3 Fe + 4 H2O → Fe3O4 + 4 H2
 Now the number of atoms of each element becomes equal on both
sides. Thus, this equation becomes a balanced equation.
 After balancing, the above equation can be written as follows:
3Fe + 4H2O → Fe3O4 + 4H2.
Name of atom No. of atoms in the
reactant
No. of atoms in the
product
Iron 3 3
Hydrogen 8 8
Oxygen 4 4

 Writing the symbols of physical states of substances in a
chemical equation
 By writing the physical states of substances, a chemical
equation becomes more informative.
 Gaseous state is represented by symbol (g).
 Liquid state is represented by symbol (l).
 Solid state is written by symbol (s).
 Aqueous solution is written by symbol (aq).
 Writing the condition in which reaction takes place: The
condition is generally written above and/or below the arrow of
a chemical equation.
 Thus, by writing the symbols of the physical state of
substances and condition under which reaction takes place, a
chemical equation can be made more informative.

 Combination Reaction
 Decomposition Reaction
 Displacement Reaction
 Double Displacement Reaction
 Neutralization Reactions
 Exothermic – Endothermic Reactions
 Oxidation-Reduction Reactions
 Redox Reaction

 A reaction in which two or more reactant combine to
form a single product is called combination and these
reactions are mostly characterised by an increase in
temperature.
Examples: A + B → AB
 The chemical reaction between quick lime & water to
form slaked lime is a combination reaction which is
characterized rise in temperature .

 When magnesium is burnt in the air (oxygen),
magnesium oxide is formed. In this reaction,
magnesium is combined with oxygen.
Magnesium + Oxygen → Magnesium Oxide
2 Mg (s) + O2 (g) → 2 MgO
Magnesium oxide (White ash)
(basic in nature)
turns Red litmus blue
 When carbon is burnt in oxygen (air), carbon
dioxide is formed. In this reaction, carbon is
combined with oxygen.
C (s) + O2(g) → CO2(g) ;
Carbon + Oxygen → Carbon dioxide

 A reaction in which a single compound decomposes or breaks
down to give two or more simpler substances is called
decomposition reaction.
 Decomposition reactions require energy in the form of heat,
light or electricity .
 A decomposition reaction is just the opposite of combination reaction.
A general decomposition reaction can be represented as follows :
AB → A + B
Examples:
 When calcium carbonate is heated, it decomposes into calcium
oxide and carbon dioxide.
CaCO3(s) + heat → CaO(s) + CO2(g) ;
 When ferric hydroxide is heated, it decomposes into ferric oxide
and water.
2Fe (OH)3(s) + heat → Fe2O3(s) + 3H2O(l)

 Thermal Decomposition
 Electrolytic Decomposition
 Photolysis or Photo Decomposition
Reaction

 The decomposition of a substance on heating is known as
Thermal or thermolytic decomposition.
Example:
white brown gas
2 FeSO4 (s) + Heat .→ Fe2O3 (s) + SO2 (g) + SO3(g)
iii) Thermal Decomposition of calcium carbonate
CaCO3(s) + Heat → CaO (aq) + CO2 (g)

 Reactions in which compounds decompose into simpler
compounds on passing electricity, are known as Electrolytic
Decomposition. This is also known as Electrolysis.
Example: When electricity is passed in water, it decomposes
into hydrogen and oxygen.
2H2O(l) + electricity →2H2(g) + O2(g)
 NOTE: As pure water is a bad conductor of electricity so before
carrying the electrolysis of water , few drops of acid, base or salt
are added ,to make water a good conductor of electricity.
 When electricity is passed in the solution of sodium chloride, it
decomposes into sodium and chlorine.
2NaCl(aq) + electricity →2Na(s) + Cl2(g)

 Reactions in which a compound decomposes because of
sunlight are known as Photolysis or Photo Decomposition
Reaction.
Example: When silver chloride salt is put in sunlight, it
decomposes into silver metal and chlorine gas.
2AgCl(s) (white) + sunlight →2Ag(s) (grey) + Cl2(g)
 Photographic paper has a coat of silver chloride, which
turns into grey when exposed to sunlight. It happens
because silver chloride is colourless while silver is a grey
metal.
 When Silver Bromide salt is put in sunlight, it decomposes
into silver metal and Bromine gas.
2AgBr(s) + sunlight →2Ag(s) (grey) + Br2(g)

 A reaction in which a more reactive element [metal]
displaces a less reactive element [metal] from its
aqueous salt solution is called displacement reaction.
Displacement reactions are also known as Substitution
Reaction or Single Displacement/ replacement reactions.
 A general displacement reaction can be represented by
using a chemical equation as follows :A + BC → AC + B
 Displacement reaction takes place only when ‘A’ is more
reactive than B. If ‘B’ is more reactive than ‘A’, then ‘A’
will not displace ‘C’ from ‘BC’ and reaction will not be
taking place.
 These reactions involves only change in colouration of
solution but there is no formation of precipitates.
 These reactions mostly involves rise in temperature so
they are considered as Exothermic reactions.

 The chemical reaction between zinc granules and dilute sulphuric acid
is characterised by the displacement of a less reactive element
[hydrogen] from its aqueous salt solution by more reactive zinc and
also by rise in temperature .
Zn(s) + 2H2SO4 (aq) → ZnSO4 + H2(g)
 The chemical reaction between zinc granules and dilute Hydrochloric
acid is also characterised by the displacement of a less reactive
element [hydrogen] from its aqueous salt solution by more reactive
zinc and also by rise in temperature .
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
 Reaction of Iron with copper sulphate is also characterised by the
displacement of a less reactive element [copper] from its aqueous salt
solution and also by change in colouration of solution.
 Fe (s) + CuSO4 (aq) → FeSO4 (aq) + Cu (s)
(Blue) (green)

 A reactions in which ions are exchanged between the aqueous solution of
two metal salts to form a new compound or precipitates is called double
displacement reactions. This reaction is also known as precipitation
reaction.
 A general displacement reaction can be represented by using a chemical
equation as follows : AB + CD → AC + BD
Examples:
When the solution of barium chloride reacts with the solution of sodium
sulphate, white precipitate of barium sulphate is formed along with
sodium chloride.
BaCl2 (aq) + Na2SO4(aq) → BaSO4 (s) + 2 NaCl (aq)
(White ppt)
 The reaction in which precipitates are formed by the mixing of the
aqueous solution of two salts is called Precipitation Reaction.
 Precipitates- The insoluble compounds formed during double
displacement reactions are called precipitate

EXAMPLES OF DOUBLE
DISPLACEMENT
REACTIONS
OR
PRECIPITATION
REACTION

 When the solution of barium chloride reacts with the
solution of copper sulphate, white precipitates of
barium sulphate are formed along with cupric
chloride.
BaCl2 (aq) + CuSO4(aq) → BaSO4 (s) + CuCl2 (aq)
(White ppt)
 When the solution of Lead Nitrate reacts with the
solution of Potassium Iodide, Yellow precipitates of
Lead Iodide are formed along with Potassium Nitrate.
Pb(NO3)2 (aq) + 2 KI (aq) → PbI2 ( ↓ ) + 2 KNO3 (aq)
(yellow ppt)

 When the solution of Silver Nitrate reacts with the
solution of sodium chloride then white ppts of silver
chloride are formed along with sodium Nitrate
 AgNO3 (aq) + NaCl (aq) → AgCl ( ↓ ) + NaNO3 (aq)
(white ppt)
 When the solution of Silver Nitrate reacts with the
solution of sodium bromide then Yellow precipitates
of silver bromides are formed along with sodium
Nitrate
 AgNO3 (aq) + NaBr (aq) → AgBr ( ↓ ) + NaNO3 (aq)
(yellow ppt)
 The chemical reaction between sulphuric acid and
barium chloride solution is characterised by the
formation of a white precipitate of barium sulphate.
BaCl2(aq) + H2SO4(aq) → BaSO4(s) + 2HCl(aq)
(white ppt).

 The reaction in which an acid reacts with a base to form salt
and water by an exchange of ions is called Neutralization
Reaction.
 ACID + BASE → SALT + WATER
 HCl (aq) + KOH (aq) → H2O (l) + KCl (aq)
(acid) (base) (Water) (Salt)
 HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
 H2SO4 (aq) +2KOH (aq) → 2H2O (l) + K2SO4 (aq)
 H2SO4 (aq) + 2NaOH (aq) → 2H2O (l) + Na2SO4 (aq)

 The Reactions in which heat or energy is evolved
are called exothermic reaction.
 The chemical reaction between quick lime & water to
form slaked lime is characterized by rise in
temperature .
 The chemical reaction between zinc granules and
dilute sulphuric acid is also characterised by rise in
temperature.
Zn(s) + 2H2SO4(aq) → ZnSO4 (aq) + H2(g)

 The chemical reaction between zinc granules and
dilute Hydrochloric acid is also characterised by
a rise in temperature.
Zn + 2HCl → ZnCl2 + H2
 Respiration is also a exothermic reaction in which
heat energy is released and thus results in increase in
temperature .
 Most of the combinations reactions occurs by an
increase in temperature of the solution so they are
also considered as exothermic reactions but all
combination reactions are not exothermic in nature.

 A chemical reaction in which heat energy is absorbed
is called Endothermic Reaction.
Example: Most of the decomposition reactions occurs by an decrease
in temperature of the solution.
white brown gas
2 FeSO4 (s) + Heat → Fe2O3 (s) + SO2 (g) + SO3(g)
iii) Thermal Decomposition of Calcium Carbonate
CaCO3(s) + Heat → CaO(aq) + CO2 (g)
iv) Photosynthesis is also a endothermic reaction in which heat energy is
absorbed and results in decrease in temperature of the system.

 The reaction in which addition of oxygen or removal of hydrogen or loss
of electron takes place is known as oxidation reaction.eg
2 Mg(s) + O2 (g) → 2MgO (s) [turns red litmus blue]
Magnesium oxide
 When a brown shiny metal copper is heated in the presence of air ,it
gets oxidised to form a black coloured substance called cupric oxide
which is found to be basic in nature and turns red litmus to blue.
2 Cu + O2 → 2 CuO (Black)
 In terms of electronic concept, “Oxidation is defined as a loss of
electrons”
 Na → Na+ + e-
 Zn → Zn2+ + 2e-
 Elements or compounds in which oxygen or non-metallic element is
added or hydrogen or metallic element is removed are called to be
Oxidized.

 The reaction in which addition of hydrogen or
removal of oxygen or gain of electron takes
place is called reduction reaction.
 In terms of electronic concept, “reduction is
defined as a gain of electrons”.
 Cu2+ + 2 e-→Cu
 Zn2+ + 2e-→Zn

 A chemical reactions where oxidation and reduction both take
place simultaneously are also known as redox reaction.
 Example: When copper oxide is heated with hydrogen, then
copper metal and water are formed.
CuO (s)+ H2 → Cu(s) + H2O (l).
 In this rxn ,CuO is reduced and H2 is oxidised simultaneously ,so
it is a redox rxn.
 Photosynthesis is also a Redox reaction.
6 CO2 + 6 H2O +sunlight +chlorophyll → C6H12O6 +6 O2
 In this rxn , CO2 is reduced and H2O is oxidised simultaneously
,so it is a redox rxn.

 When zinc oxide is heated with coke, then zinc metal
and carbon monoxide are formed.
ZnO + C (coke) → Zn (s) + CO
 In this rxn ,ZnO is reduced and C is oxidised
simultaneously ,so it is also a redox rxn.
 PbO + CO → Pb + CO2
In this rxn , Carbon monoxide is oxidised as it gains
oxygen & PbO is reduced
 (ii) H2S + Cl2→2HCl + S.
In this rxn , Chlorine is reduced as it gains hydrogen &
H2S is oxidised as it loss Hydrogen

 Oxidising Agent- A substance that loses oxygen
or gains hydrogen is known as an oxidising
agent.
 An oxidising agent gets reduced during a redox
reaction.
 Reducing Agent- A substance that loses
hydrogen or gains oxygen is known as a
reducing agent.
 A reducing agent gets oxidized during a redox
reaction.

 When copper oxide is heated with hydrogen, then
copper metal and hydrogen are formed.
CuO (s)+ H2 → Cu(s) + H2 O (l).
 In this reaction, CuO is changing into Cu. Oxygen is
being removed from copper oxide. Removal of oxygen
from a substance is called Reduction, so copper oxide
is being reduced to copper.
 In this reaction, H2 is changing to H2O. Oxygen is being
added to hydrogen. Addition of oxygen to a substance
is called Oxidation, so hydrogen is being oxidised to
water.
 The substance(H2) which gets oxidised is the reducing
agent.
 The substance (CuO)which gets reduced is the
oxidizing agent.

The reactivity series is a list of
metals arranged in the order of
decreasing reactivity.
The most reactive metal is placed
at the top and the least reactive
metal is placed at the bottom.

 The metals at the top of the reactivity series are powerful reducing
agents since they are easily oxidized. These metals tarnish/corrode
very easily.
 The reducing ability of the metals grows weaker while traversing
down the series.
 The electro-positivity of the elements also reduces while moving
down the reactivity series of metals.
 All metals that are found above hydrogen in the activity series
liberate H2 gas upon reacting with dilute HCl or dilute H2SO4.
 Metals that are placed higher on the reactivity series have the
ability to displace metals that are placed lower from their salt
solutions.
 Higher ranking metals require greater amounts of energy for their
isolation from ores and other compounds.

 The slow eating up of metals by the
action of air and moisture on their
surfaces is called Corrosion. Corrosion in
case of Iron is known as Rusting.
4 Fe(s) + 3O2 → 2Fe2O3
Iron Oxygen Ferric oxide
Fe2O3 + x H2O → Fe2O3.xH2O (Rust)
 Chemically, rust is hydrated ferric oxide
(Fe2O3.xH2O).

 Corrosion of Copper: Copper objects lose their lustre and shine
after some time because the surface of these objects acquires
a green coating of basic copper carbonate,
CuCO3.Cu(OH)2 when exposed to air.
2Cu + O2+ CO2 +H2O → CuCO3.Cu(OH)2
(GREEN )
 Blackening of Silver Metal: The surface of silver metal gets
tarnished (becomes dull) on exposure to air, due to the
formation of a coating of black silver sulphide(Ag2S) on its
surface by the action of H2S gas present in the air.
2Ag + H2S → Ag2S +H2
( BLACK)

 Presence of impurities in metal
 Air and moisture
 Presence of electrolytes
 Presence of bends
 Large surface area of metal

 Protection : Surface corrosion forms an oxide layer which
further protects the inner metal form corrosion.
In case of aluminium which on exposure to air, gets
coated with a protective layer of aluminium oxide. This
protects the metal underneath from further corrosion
and damage.
 Corrosion prevents corrosion : Some metals like zinc is
used as sacrificial anodes to prevent corrosion to other
metal.
 Galvanic corrosion : The principle on which primary
batteries work. In which one metal corrodes
preferentially to another when both metals are in
electrical contact, in the presence of an electrolyte.

 Corrosion (rusting) weakens the iron and
steel objects and structures such as railings,
car bodies, bridges and ships etc. and cuts
short their life.
 Decreases the beauty of the metal objects
 Causes Economic loss by effecting the quality
of metal objects

HOW TO
PREVENT
CORROSION?

 1. Avoid exposure to Corrosive Agents: Prevent the deterioration of
metals by limiting its contact with corrosive agents. For instance,
safeguard the metal materials from rainwater or excessive moisture by
properly storing it indoors. Moreover, exposure to high chloride
containing substances (such as salt water) must be limited. For
instance, treat the feed water inside the water boilers with softener to
prevent corrosion.
 2. Proper Monitoring of Metal Surface: Carefully monitor the metal
surface. Look for cracks and crevices. These manufacturing flaws can
also lead to corrosion. Moreover, use corrosion resistant products. For
instance, if you are buying TMT bars for construction, choose corrosion
resistant bars (such as SRMB TMT bars). Corrosion resistant TMT bars
ensure the longevity of the structure.
 3. Prevent Galvanic Corrosion: When two different metals come into
contact with each other in an electrolyte, one metal acts as anode and
the other as cathode. The presence of electrolyte acts as a conduit for
ion migration moving metallic ions from the anode to the cathode. This
leads to the anode metal corroding more quickly while the cathode
corrodes slowly. This results in Galvanic corrosion. Such type of
corrosion is common in gas or oil pipelines, the hull of the ships and
boats etc. To prevent Galvanic Corrosion, apply a protective coating on
the metal surface. This prevents the electrolytes facilitate ion flow
between the metals. Providing Cathodic protection by using a sacrificial
metal anode in the electrolytic environment also protects the metals.

 4. Protect the Metal Surface by using barriers: Paints can be
used to protect the metal surface from corrosion. The paint
forms a protective barrier between the metal surface and the
corrosive agent. For instance, coating the outdoor metal units
with a coat of paint protects them from exposure to rainwater
or snowfall etc. A number of solutions such as galvanised zinc
coating, paint or oil sealant can be used to prevent corrosion.
 5. sacrificial protection/Metallic Plating: Metallic plating can be
applied to prevent corrosion. Common processes of metallic
plating include Electroplating (covering the surface with a layer
of tin or nickel), Mechanical plating (applying zinc or cadmium
to metal surface), Electroless (coating with cobalt or nickel), or
Hot Dipping (immersing the metal in molten zinc).
 6. by applying antirust compounds on the metal surface.

 A reaction in which iron reacts with oxygen and
moisture to forms a red or brownish substance is
called rust.
 The rusting of iron is a redox reaction.
Methods to Prevent Rusting
 By painting.
 By greasing and oiling.
 By galvanisation.

 When oils and fats or foods containing oils and fats are
exposed to air for a long time, they get oxidised due to
which the food becomes stale and gives a bad taste or
smell. This is called Rancidity.
Methods to prevent rancidity:
 Adding antioxidants i.e. the substances which prevent
oxidation.
 Storing the food in air-tight containers or Vacuum Packing
 Replacing air by Nitrogen
 Refrigeration of food stuff

THANKS

Chemical reactions and equations 10 chm(1)

Recommended

Recommended

More Related Content

What's hot

What's hot (20)

Similar to Chemical reactions and equations 10 chm(1)

Similar to Chemical reactions and equations 10 chm(1) (20)

More from VeenuGupta8

More from VeenuGupta8 (17)

Recently uploaded

Recently uploaded (20)

Chemical reactions and equations 10 chm(1)