The document discusses chemical reactions and equations, defining chemical reactions as transformations of substances that can be represented symbolically through equations. It explains various types of reactions, including combination, decomposition, displacement, double displacement, and redox reactions, while highlighting the need for balancing chemical equations to adhere to the conservation of mass. The document also touches on practical examples, the effects of reactions such as corrosion and rancidity, and methods of preventing deterioration of materials.

1. CHEMICAL
REACTIONS AND
EQUATIONS

2. CHEMICAL
REACTION
S
A chemical reaction is a process that
leads to the transformation of one set
of chemical substances to another.
Classically, chemical reactions
encompass changes that only involve
the positions of electrons in the forming
and breaking of chemical bonds
between atoms, with no change to the
nuclei, and can often be described by a
chemical equation.

3. CHEMICAL REACTIONS ARE
EVERYWHERE
COOKING RESPIRATION

4. CHEMICAL REACTIONS ARE EVERYWHERE
RUSTING
FORMATION OF
CURD

5. Indications of a Chemical Reaction
EVOLUTION OF
GAS
CHANGE IN
COLOUR
CHANGE
IN TEMPERATURE
FORMATION
OF PRECIPITATE

6. CHEMICAL EQUATION
A chemical equation is
the symbolic
representation of a
chemical reaction in the
form of symbols and
formulae.
 ex:-
Magnesium + Oxygen magnesium oxide
 The substances that undergo chemical change in
the reaction
(magnesium and oxygen) are the reactants.
 The new substances (magnesium oxide) formed
during the reactions is the product.

7. WORD EQUATION
A word – equation shows change of reactants to products through an
arrow placed between them.
 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 reactions.

8. SKELETAL EQUATION
Skeletal equations are those equations which show the
reactant and product formed without balancing them.
Example :-
Mg + O2 MgO
H2 + O2 H2O
 It is also known as unbalanced equation.

11. BALANCED EQUATION
 The chemical equation needs to be balanced so that it follows the law of
conservation of mass.
 “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.”
 A chemical equation is called a balanced chemical equation when the number of
the different atoms of elements in the reactants side is equal to that of the
products side.

12. Examine the number of atoms of different
elements on both sides of the arrow.
Zn + H2SO4 ZnSO4 + H2

13. How to balance a equation
Step I: To balance a chemical equation, first draw boxes around each
formula. Do not change anything inside the boxes while balancing the
equation.

14.  Step II: List the number of atoms of different elements present in
the unbalanced equation.

15. Step III: It is often convenient to start balancing with the compound that contains
the maximum number of atoms. It may be a reactant or a product.
To balance the oxygen atoms –
To equalise the number of atoms, it must be remembered that we cannot alter the
formulae of the compounds or elements involved in the reactions. For example, to
balance oxygen atoms we can put coefficient ‘4’ as 4 H2O and not H2O4 or (H2O)4.

16. Step IV: Fe and H atoms are still not balanced. Pick any of these
elements to proceed further.

17. Step V: Examine the above equation and pick up the third
element which is not balanced. You find that only one element
is left to be balanced, that is, iron.

18. Step VI: Finally, to check the correctness of the balanced equation, we
count atoms of each element on both sides of the equation.
The numbers of atoms of elements on both sides of the equation are equal.
This equation is now balanced. This method of balancing chemical equations is
called hit-and-trial method as we make trials to balance the equation by using
the smallest whole number coefficient.

19. WRITING SYMBOLS OF PHYSICAL STATES
 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.
 Sometimes the reaction conditions , such as temperature, pressure,
catalyst etc are indicated above or below the arrow in the equation.

21. Write the balanced equation for the following
chemical reactions.
(i) Hydrogen + Chlorine → Hydrogen chloride
(ii) Barium chloride + Aluminium sulphate → Barium sulphate +Aluminium chloride
(iii) Sodium + Water → Sodium hydroxide + Hydrogen
H2 + Cl2 HCl
BaCl2 + Al2(SO4)3 BaSO4 + AlCl3
Na + H2O NaOH + H2

22. Write the balanced equation for the following
chemical reactions.
(i) Hydrogen + Chlorine → Hydrogen chloride
(ii) Barium chloride + Aluminium sulphate → Barium sulphate +Aluminium chloride
(iii) Sodium + Water → Sodium hydroxide + Hydrogen
H2 + Cl2 2 HCl
3BaCl2 + Al2(SO4)3 3BaSO4 + 2AlCl3
2Na + 2 H2O 2NaOH + H2

23. Write a balanced chemical equation with state
symbols for the following reactions.
 (i) Solutions of barium chloride and sodium sulphate in water react to give
insoluble barium sulphate and the solution of sodium chloride.
 (ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in
water) to produce sodium chloride solution and water.

24. Write a balanced chemical equation with state
symbols for the following reactions.
 (i) Solutions of barium chloride and sodium sulphate in water react to give
insoluble barium sulphate and the solution of sodium chloride.
BaCl2 (aq) + Na2SO4 (aq) BaSO4 (s) + 2NaCl (aq)
 (ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in
water) to produce sodium chloride solution and water.
NaOH (aq) + HCl (aq) NaCl (aq) + H2O (l)

25. Reactions in which heat is released along with the formation of
products are called exothermic chemical reactions.
Examples of an exothermic reaction
 Burning of natural gas
CH4(g) + 2O2 (g) CO2 (g) + 2H2O (g)
Respiration
C6H12O6(aq) + 6O2(aq) 6CO2(aq) + 6H2O(l) + energy
Decomposition of vegetable matter into compost
Endothermic reactions are those reactions in which heat is
absorbed

26. TYPES OF CHEMICAL REACTIONS
THERE ARE 5 TYPES OF REACTIONS:-
Combination reactions
Decomposition reaction
Displacement reactions
Double displacement reactions
Redox reactions

27. COMBINATION REACTIONS
 When two or more reactants combine to form a single product it is called a
combination reaction.
 Calcium oxide(quick lime) reacts vigorously with water to produce slaked lime
(calcium hydroxide) releasing a large amount of heat.
CaO (s) + H2O (l) Ca (OH)2 (aq)
Calcium oxide (quick lime) calcium hydroxide (slaked lime)
 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.

28. More examples –
 Burning of Coal
C(s) + O2(g) CO2(g)
 Formation of water
2H2(g) + O2(g) 2H2O(g)

29. Decomposition reactions
 When a single reactant breaks up into two or more products the reaction is termed as
decomposition reaction.
 The decomposition reactions require energy either in the form of heat, light or electricity
for breaking down the reactants. There are three types of decomposition reactions based on
the type of energy used to break down the reactant.
1) Thermal decomposition reaction
2) Photochemical decomposition reaction (Photolysis)
3) Electric decomposition reaction (Electrolysis)

30. Thermal Decomposition reaction
 When a single reactant breaks up into two or more products due to heating, the reaction is
termed as thermal decomposition reaction.
 Example
2FeSO4 (s) Fe2O3(s) + SO2 (g) + SO3 (g)
Ferrous sulphate Ferric oxide sulphur dioxide sulphur trioxide
 Ferrous sulphate crystals (FeSO4 .7H2O) lose water when heated and the colour of the
crystals changes (light green changes to white and then changes to brown). 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. A smell of burning sulphur is
obtained due to formation of sulphur dioxide gas.

31. Thermal decomposition
 Example 3: Decomposition of lead nitrate
2Pb(NO3)2(s) 2PbO(s) + 4NO2(g) + O2(g)
Lead nitrate lead oxide nitrogen dioxide oxygen
In this reaction, colourless lead nitrate when heated forms yellow lead oxide and
brown fumes due to the formation of nitrogen dioxide gas is observed.

32. Photolysis (Photochemical decomposition reaction)
 When a single reactant breaks up into two or more products due to energy obtained from
light, the reaction is termed as photochemical decomposition reaction .Halogen
compounds decompose on exposure to sunlight . Photolytic reactions are used in black
and white photography.
silver chloride silver chlorine
(white) (greyish white) (yellowish green)
silver bromide silver bromine
(pale yellow) (greyish white) (reddish brown)

33. Electrolytic decomposition (electrolysis)
 When a single reactant breaks up into two or more products due to the
passage of electricity, the reaction is termed as electrolytic
decomposition reaction.
 Example
2H2O electricity 2 H2 + O2

34. Electrolytic decomposition (electrolysis)
 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. 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 hydrogen is double of oxygen in
terms of volume.

35. Displacement reactions
 When a more reactive element displaces a less reactive element from its
compound, it is known as a displacement reaction.
In the above reaction Zn is more reactive Cu, so Zn displaces Cu from
CuSO4.
In the above reaction Zn is more reactive than Ag , so Zn displaces Ag
from its solution.
Zinc copper sulphate zinc sulphate copper
Silver nitrate zinc silver zinc nitrate

36. Displacement reaction
 Example 3 – Iron displaces copper from copper sulphate to give iron sulphate and
copper. (Iron is more reactive than copper)​

37. Displacement reaction
 Example 4: Zinc displaces hydrogen from hydrogen sulphide to form zinc sulphide
and hydrogen gas

38. Displacement reaction
Example 5 : Lead is more reactive than copper. Lead displaces copper from copper
chloride to form lead chloride and copper.
Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
Lead copper chloride lead chloride copper

39. Double displacement reactions
 Reactions in which there is an exchange of ions between the reactants are
called double displacement reactions.
 These reactions are also called precipitation reactions as an insoluble
substance is formed which is known as a precipitate. Any reaction that
produces a precipitate can be called a precipitation reaction.
Barium chloride sodium sulphate barium sulphate sodium chloride
(white precipitate)
BaCl2(aq) + Na2SO4(aq) BaSO4(s) + 2NaCl(aq)

40. Double displacement reactions
 Example 2 – Silver nitrate reacts with sodium chloride to form sodium nitrate and a
white precipitate of silver chloride

41. Double displacement reactions
Example 3
Lead nitrate and potassium iodide react to form potassium nitrate and lead iodide.
Lead iodide is formed as a yellow precipitate (yellow solid) in this reaction
Pb(NO3 )2 (aq) + 2KI (aq) 2 KNO3 (aq) + PbI2 (s)
lead nitrate potassium iodide potassium nitrate lead iodide (yellow precipitate)

42. NEUTRALIZATION REACTION
When an acid and a base react to form a salt and water, the reaction is
known as a neutralization reaction.
A neutralization is a type of double replacement reaction.
 In the reaction, H+ and OH- combine to form H2O or water molecules.
For example
HCl + NaOH NaCl + H2O
Hydrochloric acid sodium hydroxide sodium chloride water
(acid ) (base) (salt)

43. Virtual lab video links
1) Combination reaction
https://www.youtube.com/watch?v=EHlVFgg-3Kc
2) Decomposition reaction
https://www.youtube.com/watch?v=2WqQ8sQ4pPE
3)Displacement reaction
4) Double displacement reaction
https://www.youtube.com/watch?v=pTsWD0lyRfI

44. Oxidation & reduction (redox reaction)
 If a substance gains oxygen during a reaction, it is said to be oxidised. If a
substance loses oxygen during a reaction, it is said to be reduced.
 If one reactant gets oxidised while the other reactant gets reduced during a
reaction, the reaction is called an oxidation-reduction reaction or redox
reaction.
 If a substance gains oxygen or loses hydrogen during a reaction, it is
oxidised. If a substance loses oxygen or gains hydrogen during a reaction, it
is reduced.

45. Oxidation is
1) Addition of oxygen
2) Removal of hydrogen
Any chemical substance following any these is said
to be oxidised.
Reduction is
1) Removal of oxygen
2) Addition of hydrogen
Any chemical substance following any of these rules is said to be reduced.
A reaction involving both oxidation and reduction process , occurring simultaneously is known
as redox reaction or oxidation and reduction reaction.

46.  In the above redox reaction, zinc oxide ZnO is getting reduced (loss of oxygen) to
zinc Zn . Carbon C is getting oxidised (gain of oxygen) to carbon monoxide CO
 In the above redox reaction, manganese (IV) oxide MnO2 is getting reduced (loss
of oxygen) to MnCl2 manganese (II) chloride
 In the same reaction, HCl hydrochloric acid is getting oxidised (loss of hydrogen) to
chlorine molecule Cl2

47. Corrosion
 Corrosion is a natural process, which converts a refined metal to a more stable form, such
as its oxide, hydroxide, or sulfide.
 When a metal is attacked by substances around it such as moisture, acids, etc., it is said to
corrode .The brown coating on iron,black coating on silver and the green coating on copper
are other examples of corrosion.
 Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made
of metals, specially those of iron. Corrosion of iron is a serious problem.
 Every year an enormous amount of money is spent to replace damaged iron.
 The rusting of iron can be prevented by painting , oiling , galvanizing, anodizing etc.
 Galvanization is a method of protecting steel and iron from rusting by coating them with a
thin layer of zinc.

48. Rancidity
 When fats and oils are oxidized, they become rancid and
their smell and taste change. Usually substances which
prevent oxidation (antioxidants) are added to foods
containing fats and oil.
Keeping food in air tight containers helps to slow down
oxidation. This prevents the food from becoming rancid
and hence chips are flushed with nitrogen gas to prevent it
from rancidity.

49. Thank you