Chemical reactions and equation (chem-1)

Activity 1.1:
 Magnesium ribbon is cleaned by rubbing it with sandpaper
before burning.
 Magnesium reacts with atmospheric oxygen to develop a layer
of magnesium oxide. This layer does not allow the underlying
magnesium to undergo combustion.
 Magnesium ribbon burns with a dazzling white flame and
changes into a white powder. This powder is magnesium
oxide.
 It is formed due to the reaction between magnesium and
oxygen present in the air.
 2Mg + O2 → 2MgO
Activity 1.2:
 Lead nitrate reacts with Potassium Iodide to form Lead Iodide
(yellow precipitate) and Potassium nitrate.
 Pb(NO3)2 + 2KI → PbI2 (yellow ppt)+ 2KNO3
Activity 1.3:
 Formation of hydrogen gas by the action of dilute sulphuric
acid on zinc.
 Zn + 2HCl → ZnCl2 + H2

Chemical Equations
Balanced Chemical Equation:
Definition: A balanced chemical equation is an equation in which the
number of atoms of various elements is equal on both sides of the
equation.
Reason for balancing: An equation should be balanced due to the law of
conservation of mass which states that mass can neither be created nor
destroyed in a chemical reaction. That is, the total mass of the elements
present in the products of a chemical reaction has to be equal to the total
mass of the elements present in the reactants.
Writing Symbols of Physical States:
States Symbol

gaseous g

liquid l

aqueous aq

solid s
Sometimes the reaction conditions, such as temperature, pressure,
catalyst, etc., for the reaction are indicated above and/or below the arrow
in the equation. For example:
Types of Chemical Reactions
1. Combination Reaction
2. Decomposition Reaction
3. Displacement Reaction
4. Double Displacement Reaction
5. Oxidation and Reduction
Some other types:

 Exothermic reaction
 Endothermic reaction
 Precipitation reaction
 Redox reaction
Combination Reaction
Definition: A reaction in which a single product is formed from two or more
reactants is known as a combination reaction.
Examples:
i. Calcium oxide reacts vigorously with water to produce slaked lime
(calcium hydroxide) releasing a large amount of heat.

A solution of slaked lime (calcium hydroxide) is used for whitewashing

walls. It reacts slowly with the carbon dioxide in the air to form a thin
layer of calcium carbonate on the walls.

ii. Burning of coal: C + O2 → CO2

iii. Formation of water from H2(g) and O2(g): 2H2 (g) + O2 (g) → 2H2O (l)
Exothermic Reaction:
Definition: Reactions in which heat is released along with the formation of
products are called exothermic chemical reactions.
Examples:
i. Burning of natural gas: CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)
ii. Respiration: Carbohydrates are broken down to form glucose. This
glucose combines with oxygen in the cells of our body and provides
energy. The special name of this reaction is respiration.
iii. The decomposition of vegetable matter into compost.

Decomposition Reaction
Definition: The reaction in which a single reactant breaks down to give
simpler products is called a decomposition reaction.
Examples:
i. Ferrous sulphate crystals (FeSO4.7H2O) lose water when heated and the
color of the crystals changes. It then decomposes to ferric oxide (Fe 2O3),
sulphur dioxide (SO2), and sulphur trioxide (SO 3). Ferric oxide is a solid,
while SO2 and SO3 are gases.

ii. Decomposition of calcium carbonate to calcium oxide (quick lime) and

carbon dioxide on heating is an important decomposition reaction.

When a decomposition reaction is carried out by heating, it is

called thermal decomposition.
iii. Heating of lead nitrate. You will observe the emission of brown fumes.
These fumes are of nitrogen dioxide (NO2).

iv. Electrolysis of water is a decomposition reaction. The mole ratio of

hydrogen and oxygen gases liberated during the electrolysis of water is
2:1.

Cathode: Hydrogen; anode: Oxygen.

Hydrogen is collected in double the amount of oxygen because water is

formed by the chemical combination of hydrogen and oxygen in the ratio
2:1 by volume, so it decomposes in the same ratio.

2H2O(l) → 2H2(g) + O2(g).

v. White silver chloride turns grey in sunlight.

2AgCl (s) → 2Ag (s) + Cl2 (g)

The application of this reaction is in black-and-white photography.

Endothermic reaction:
Reactions in which energy (either in the form of heat, light, or electricity is
absorbed are known as endothermic reactions.

Displacement Reaction
Definition: The reaction in which a more reactive element displaces a less
reactive element from its salt solution is called displacement reaction.
Examples:
i. Iron displaces copper, from copper sulphate solution. Iron nail become
brownish in color and the blue colour of the copper sulphate solution
fades.

Fe + CuSO4 → FeSO4 + Cu
ii. Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
iii. Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)

Double Displacement Reaction

Definition: Reactions in which there is an exchange of ions between the
reactants are called double displacement reactions.
Examples:
i. Na2SO4 + BaCl2 → BaSO4 (ppt) + 2NaCl
You will observe that a white substance (BaSO 4), which is insoluble in
water, is formed. This insoluble substance formed is known as a
precipitate. Any reaction that produces a precipitate can be called
a precipitation reaction.
ii. When potassium iodide solution is added to a solution of lead nitrate in
a test tube, a yellow color precipitate is formed. The compound
precipitated is Lead Iodide.

Pb(NO3)2 + 2KI → PbI2 + 2KNO3

Oxidation and Reduction
Definition: If a substance gains oxygen or loses hydrogen during a
reaction, it is said to be oxidized. If a substance loses oxygen or gains
hydrogen during a reaction, it is said to be reduced.
Examples:
i. Oxidation of copper to copper oxide: When we heat copper powder, the
surface of copper powder becomes coated with black copper(II) oxide.
This is because oxygen is added to copper and copper oxide is formed.

2Cu + O2 → 2CuO
ii. Reduction of copper oxide to copper: If hydrogen gas is passed over this
heated material (CuO), the black coating on the surface turns brown as
the reverse reaction takes place and copper is obtained.

CuO + H2 → Cu + H2O
Redox Reactions
Definition: The reactions in which one reactant gets oxidized while the
other gets reduced are called oxidation-reduction reactions or redox
reactions.
Note: The substance which gets oxidized is the reducing agent and the
substance which get reduced is the reducing agent.

Examples:
i.

ii. ZnO + C → + Zn + CO

Carbon is oxidized to CO and ZnO is reduced to Zn. [Here, Carbon is the

reducing agent and ZnO is the Oxidizing agent.]

iii. MnO2 + 4HCl → MnCl2 + 2H2O + Cl2

HCl is getting oxidized to Cl2 while MnO2 is getting reduced to MnCl2. [Here,
HCl is the reducing agent and MnO2 is the oxidizing agent.]
Effects of Oxidation and Reduction
in everyday life
i. Corrosion:
When a metal is attacked by substances around it such as moisture, acids,
etc., it is said to corrode and this process is called corrosion.

Examples:
 Rusting of iron
 The black coating on silver
 The green coating on copper
Disadvantages of Corrosion:
 Corrosion causes damage to car bodies, bridges, iron railings,
ships, and to all objects made of metals, especially those of
iron.
 Every year an enormous amount of money is spent to replace
damaged iron.
ii. Rancidity
When fats and oils are oxidized, they become rancid, and their smell and
taste change. Usually, substances that prevent oxidation (antioxidants)
are added to foods containing fats and oil. Keeping food in air-tight
containers helps to slow down oxidation.

Chips manufacturers usually flush bags of chips with a gas such as

nitrogen to prevent the chips from getting oxidized.