The document provides information about various chemical reactions: 1. Decomposition reactions involve a single reactant breaking down into simpler products, such as ferrous sulfate decomposing into ferric oxide, sulfur dioxide, and sulfur trioxide when heated. 2. Displacement reactions occur when a more reactive element displaces a less reactive element from its compound, like iron displacing copper from copper sulfate solution. 3. Double displacement reactions involve the switching of ions between reactants to form new ionic compounds, exemplified by the reaction of barium chloride and sodium sulfate forming barium sulfate and sodium chloride. 4. Combination reactions form a single product from two or more reactants, such as calcium oxide react

2. Introduction

3. Activity 1.1

4. Activity 1.1

5. • 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.magnesium, etc.
Activity 1. 1

6. Activity 1.2

7. Activity 1.2

8. Activity 1. 3

9. Activity 1. 3

10. Activity 1. 3

11. From the above three activities, we can say that any of
the –
• Change in state
• Change in colour
• Evolution of a gas
• Change in temperature.
Conclusion

12.  A chemical reaction is identified by any of these 4
factors
Change in state
 change in colour
 evolution of a gas
 Change in temperature.
How do we identify a chemical reaction?

13.  A chemical reaction is a process where the reactant
gets converted into a product which may be under an
influence of a catalyst .
 A word-equation shows change of reactants to products
through an arrow placed between them.
What are chemical reactions ?

14. 1. What is a chemical Reaction?
1. DddTt
2. Yjj
3. What are the observations helps us to determine whether a
chemical reaction has taken place or not.

15.  The reactants are written on the left-hand side
(LHS) with a plus sign (+) between them. Similarly,
products are written on the right- hand side (RHS)
with a plus sign (+) between them.
 The arrowhead points towards the products,
and shows the direction of the reaction.
What are chemical reactions ?

16. What are chemical reactions ?

17. Writing a chemical reaction
 A chemical equation represents a chemical reaction. If you
recall formulae of magnesium, oxygen and magnesium
oxide, the above word-equation can be written as –
Mg+O2 →MgO

18.  According to law of conservation of mass matter
can neither be created nor be destroyed in a
chemical reaction.
ANTOINE LAVOISER
INTRODUCED
LAW OF CONSERVATION OF
MASS
Balancing a chemical reaction

19. 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.

In other words, the number of atoms of each element remains
then same, before and after a chemical reaction. Hence, we
need to balance a skeletal chemical equation
Balancing a chemical reaction

20. Balancing a chemical reaction

21. Activity 1. 3

22. Balancing a chemical reaction

23.  To balance a chemical equation, first draw boxes around each
formula. Do not change anything inside the boxes while
balancing the equation.
 List the number of atoms of different elements present in the
unbalanced equation
 Start balancing with the compound that contains the maximum
number of atoms. It may be a reactant or a product. In that
compound, select the element which has the maximum number
of atoms.
Balancing a chemical reaction

24.  Balance the unbalanced atom.
 Finally, to check the correctness of the balanced equation, we
count atoms of each element on both sides of the equation.
 To make a chemical equation more informative, the physical states of
the reactants and products are mentioned along with their chemical
formulae. The gaseous, liquid, aqueous and solid states of reactants
and products are represented by the notations (g), (l), (aq) and (s),
respectively.

25.  The word aqueous (aq) is written if the reactant or product is present as a
solution in water.

26. Balancing a chemical reaction
Mg+O2 →MgO

27. Balancing a chemical reaction
Mg+O2 →MgO

28. Balancing a chemical reaction

29. Balancing a chemical reaction

30. Balancing a chemical reaction

38. Revision

39. Revision
Name and symbols of some ions

56.  THERE ARE 5 TYPES OF REACTIONS WE ARE GOING TO STUDY :-
1. COMBINATION REACTIONS
2. DECOMPOSITION REACTION
3. DISPLACEMENT REACTIONS
4. DUBLE DISPLACEMENT REACTIONS
5. REDOX REACTIONS
TYPES OF CHEMICAL REACTIONS

57. When two or more substances (elements or compounds)
combine to form a single product, the reactions are called
combination reactions
A reaction in which a single product is formed from two or
more reactants is known as a combination reaction.
COMBINATION REACTIONS

58. COMBINATION REACTIONS

59.  Calcium oxide reacts vigorously with water to produce
slaked lime (calcium hydroxide) releasing a large amount
of heat.
CaO + H2O > Ca(OH)2
Calcium hydroxide reacts slowly with the carbon dioxide in
air to form a thin layer of calcium carbonate on the walls.
Calcium carbonate is formed after two to three days of white
washing and gives a shiny finish to the walls.
It is interesting to note that the chemical formula for marble
is also CaCO3.
COMBINATION REACTIONS

61.  Reactions in which heat is released along with the formation of
products are called exothermic chemical reactions.
 Burning of natural gas and the decomposition of vegetable matter
into compost are also an example of an exothermic reaction.

65.  Endothermic reaction are those reactions in which heat is
absorbed.
 During digestion, food is broken down into
simpler substances.
For example, rice, potatoes and bread contain
carbohydrates. These carbohydrates are
broken down to form glucose. This glucose
combines with oxygen in the cells of our body
and provides energy. So Exothermic process S

67. When a product breaks up into its constituent
reactants the reaction is termed as decomposition
reaction.
Decomposition reactions
Heat
2FeSO4 (s) Fe2O3(s) + SO2 (g) + SO3 (g)

68.  In this reaction you can observe that a single reactant breaks
down to give simpler products. This is a decomposition reaction.
Decomposition reactions

69. 
 Ferrous sulphate crystals (FeSO4, 7H2O) lose water when heated
and the colour of the crystals changes.
 It then decomposes to ferric oxide (Fe2O3), sulphur dioxide (SO2)
and sulphur trioxide (SO3). Ferric oxide is a solid, while SO2 and
SO3 are gases.
Decomposition reactions

70. Decomposition reactions

71.  Decomposition of calcium carbonate to calcium oxide and carbon
dioxide on heating is an important decomposition reaction used
in various industries.

Decomposition reactions

72. Decomposition reactions

73.  Decomposition of calcium carbonate to calcium oxide and carbon
dioxide on heating is an important decomposition reaction used
in various industries.
 Calcium oxide is called lime or quick lime. It has many uses –
one is in the manufacture of cement.
 When a decomposition reaction is carried out by heating, it is
called thermal decomposition.
Decomposition reactions

74. Decomposition reactions

75. Decomposition reactions

76. Decomposition reactions

77.  Take a plastic mug. Drill two holes at its base and
fit rubber stoppers in these holes. Insert carbon
electrodes in these rubber stoppers. Connect
these electrodes to a 6 volt battery.
 Fill the mug with water such that the electrodes are
immersed.
 Add a few drops of dilute sulphuric acid to the
water.
 Take two test tubes filled with water and invert
them over the two carbon electrodes.
 Switch on the current and leave the apparatus
undisturbed for some time.
Electrolytic decomposition (electrolysis)

78. You will observe the formation of bubbles
at both the electrodes. These bubbles
displace water in the test tubes.
We will see hydrogen at the cathode and
oxygen at the anode and oxygen is
double of hydrogen in terms of volume.
Electrolytic decomposition (electrolysis)

79. white silver chloride turns grey in sunlight.
This is due to the decomposition of silver
chloride into silver and chlorine by light
2AgCl(s) >2Ag(s) + Cl2
(g)
Photolysis (thermal decompsition)

80. Halogen compounds decompose on exposure to sunlight
Photolytic reactions are used in black and white
photography
Photolysis (thermal decompsition)

81. l
The decomposition reactions require energy either in the
form of heat, light or electricity for breaking down the
reactants.
Reactions in which energy is absorbed are known as
endothermic reactions.
Decompsition Reactions

82. Decompsition Reactions

88. • In activity 1.7, gas collected in one of the test tubes is double of
the amount collected in the other because water gets
hydrolysed to release H2 and O2 gas.
• Here, after electrolysis two molecules of Hydrogen and one
molecule of oxygen gas are released; hence, the amount of
Hydrogen collected would be double than that of oxygen.

90. When a element displaces another element from its
respective compound of lower reactivity it is said to be
a displacement reaction
Displacement reactions

91. Displacement reactions

92. Displacement reactions

93. Displacement reactions

94. Displacement reactions

95. Displacement reactions

96. Displacement reactions

97. Examples of displacement reactions

98. Double Displacement reactions

99. Double Displacement reactions

100. Double Displacement reactions

101.  When a product breaks up into its constituent reactants the reaction is termed as
decomposition reaction.
Heat
2FeSO4 (s) >Fe2O3(s) + SO2 (g) + SO3 (g)
 In this reaction you can observe that a single reactant breaks down to give
simpler products. This is a decomposition reaction. Ferrous sulphate crystals
(FeSO4, 7H2O) lose water when heated and the colour of the crystals changes. It
then decomposes to ferric oxide (Fe2O3), sulphur dioxide (SO2) and sulphur
trioxide (SO3). Ferric oxide is a solid, while SO2 and SO3 are gases.
 Decomposition of calcium carbonate to calcium oxide and carbon dioxide on
heating is an important decomposition reaction used in various industries.
Calcium oxide is called lime or quick lime. It has many uses – one is in the
manufacture of cement. When a decomposition reaction is carried out by heating,
it is called thermal decomposition.
TYPES OF CHEMICAL REACTIONS

102. • Metals are described as chemical elements that readily lose
valence electrons to form positive ions (cations).
Examples: Aluminium, copper, iron, tin, gold.
• Around 90 of the total 118 elements are metals.
Metals

103. The property of metals by which they can be beaten
into thin sheets is called malleability.
Malleability

104. The property of metal by which it can be drawn into wires is
called ductility.
Ductility

105. Metals produce ringing sounds, they are said to be sonorous.
The materials other than metals are not sonorous
Sonorous

106. ● Hard and have a high tensile strength
● Solids at room temperature, except mercury, which is
liquid at room temperature.
● Sonorous
● Good conductors of heat and electricity
● Malleable, i.e., can be beaten into thin sheet
Physical Properties of Metals

107. ● Malleable, i.e., can be beaten into thin sheets
● Ductile, i.e., can be drawn into thin wires
● High melting and boiling points (except Cesium (Cs) and Gallium
(Ga))
● Dense, (except alkali metals). Osmium – highest density and
lithium – least density
● Lustrous
● Silver-grey in colour, (except gold and copper)
Physical Properties of Metals

108. • Non-metals are those elements, which do not exhibit the properties
of metals.
Examples: Carbon, Boron, etc.
Non-Metals

109. • Materials like coal and sulphur are soft and dull
in appearance.
• They break down into a powdery mass on tapping with a
hammer.
• They are not sonorous and are poor conductors of heat and
electricity.bThese materials are called non-metals.
• The examples of non-metals are sulphur, Carbon, oxygen,
phosphorus, etc.
Non-metals

112. • The name of the product formed in the reaction of sulphur
and oxygen is sulphur dioxide gas.
• When sulphur dioxide is dissolved in water sulphurous acid is
formed.
• The reaction can be given as follows:
Sulphur dioxide + Water →Sulphurous acid
SO2 + H2O  H2SO4
• The sulphurous acid turns blue litmus paper red.
Generally, oxides of non-metals are acidic in nature.
Reaction of Non-Metals with oxygen.

119. Reaction with Water
You observed that sodium reacts vigorously with water. Some
other metals do not do so.
For example, iron reacts with water slowly.

123. Reactions with Bases
• What does the ‘pop’ sound indicate?
• As before, the ‘pop’ sound indicates the presence of hydrogen
gas.
• Metals react with sodium hydroxide to produce hydrogen gas.

125. • Many non-metals react with bases to form salts.
• Bases are electron donors whereas non-metals are electron
acceptors. Thus, bases donate electrons to non-metals which
readily accepts them and form a salt.
• Reactions of non-metals with bases are complex.