Chemical Reactions and Equations

Chemical Reaction:
• The transformation of chemical substance into another chemical substance is known
as Chemical Reaction. For example: Rusting of iron, the setting of milk into curd,
digestion of food, respiration.
• Example: The burning of magnesium in the air to form magnesium oxide is an
example of a chemical reaction.
2Mg(s) + O2(g) △→ 2MgO(s)
• When copper sulphate reacts with iron, two new substances, ferrous sulphate and
copper are formed.
Fe(s) +CuSO4(aq) → FeSo4(aq) +Cu(s)

• Reactant: Substances which take part in a chemical reaction are called reactants.
Example: Mg and O2, Fe and CuSO4
• Product: New substance formed after a chemical reaction is called a product.
Example: MgO, FeSO4 and Cu

Characteristics of Chemical Reactions:

a)Evolution of a gas
b)Change in colour
c)Change in state
d)Change in temperature
e)Formation of a precipitate
a) Evolution of gas: In many chemical reactions, gas is evolved sometimes with an
effervescence. For, example when zinc metal reacts with hydrogen chloride,
hydrogen gas is evolved.
Zn + 2HCl → ZnCl2 + H2↑
mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g) ↑

b) Change of colour: Sometimes, a colour change occurs during a chemical reaction

due to the formation of a new product. For example, when silver chloride is exposed
to sunlight it produces black metallic coloured silver with the evolution of chlorine
gas. 2AgCl → 2Ag + Cl2
 c) Change of state: In some reactions, the change of state is observed in the
formation of products from reactants. For example, sodium chloride (aqueous) when
reacts with silver nitrate (aqueous) forms sodium nitrate (aqueous) and silver
chloride (solid white precipitate).
NaCl (aq) + AgNO3 (aq) → NaNO3 (aq) + AgCl (white ppt)
d) Change in temperature: The chemical reaction between quick lime water to form
slaked lime is characterized by a change in temperature (which is a rise in
temperature).
The chemical reaction between zinc granules and dilute sulphuric acid is also
characterised by a change in temperature (which is a rise in temperature).
e) Formation of precipitate: In certain chemical reactions, an insoluble solid
substance called precipitate is formed. For example, silver nitrate reacts with sodium
bromide to form a yellow precipitate of silver bromide.
AgNO3 + NaBr → NaNO3 + AgBr ↓
BaCl2(aq) + H2SO4(aq) → BaSO4(s) (ppt) + 2HCl(aq)

Chamical Reaction and Equation :

Chemical Equation:
• Representation of chemical reaction using symbols and formulae of the substances is
called Chemical Equation.
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.
• When hydrogen reacts with oxygen, it gives water. This reaction can be represented
by the following chemical equation:
Hydrogen + Oxygen → Water
H2 + O2 → H2O

Writing chemical equations

Representation of a chemical reaction in terms of symbols and chemical formulae of the
reactants and products is known as a chemical equation.The reactants is always written in
left sight and the product in written in right sight.
• For solids, the symbol is “(s)”.
• For liquids, it is “(l)”.
• For gases, it is “(g)”.
• For aqueous solutions, it is “(aq)”.
• For gas produced in the reaction, it is represented by “(↑)”.
• For precipitate formed in the reaction, it is represented by “(↓)”.
Types of chemical reactions
chemical reactions are grouped into multiple categories Few examples are:
● Combination
● Decomposition
● Displacement
● Double displacement
● Redox
● Endothermic
● Exothermic
● Precipitation
● Neutralisation

Combination Reaction:
• The Reactions in which two or more reactants combine to form one product are
called Combination Reaction. (A + B → AB)
• Examples: -When magnesium is burnt in the air (oxygen), magnesium oxide is
formed. In this reaction, magnesium is combined with oxygen.
Mg(s) + O2(g) → 2MgO(s)
Magnesium + Oxygen → Magnesium Oxide
• 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
• When calcium oxide (quick lime) is reacted with water, calcium hydroxide (slaked
lime) is formed.
Cao + H2O → Ca (OH)2
Decomposition Reaction:
• Reactions in which one compound decomposes in two or more compounds or
elements are known as Decomposition Reaction.
• A decomposition reaction is just the opposite of combination reaction. (AB → A + B)
• Examples: When calcium carbonate is heated, it decomposes into calcium oxide and
carbon dioxide.
CaCO3(s) →− CaO(s) + CO2(g)
Calcium carbonate → Calcium oxide + Carbon dioxide
• When ferric hydroxide is heated, it decomposes into ferric oxide and water
2Fe(OH)3(s) △→ Fe2O3(s) + 3H2O(l)
Thermal Decomposition: The decomposition of a substance on heating is known as
Thermal Decomposition.

CuSo4→ CuO + So3 ( copper sulphate→ cupric oxide +sulphur trioxide)

2FeSo4→ Fe2o3 + So3 + So2 (ferrous sulphate →Ferric oxide +Sulphur dioxide +Sulphur trioxide)

Electrolytic Decomposition: Reactions in which compounds decompose into simpler

compounds because of passing of 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 → 2H2 + O2

Photolysis or Photo Decomposition Reaction: Reactions in which a compound

decomposes because of sunlight are known as Photolysis or Photo Decomposition Reaction.
Example: When H202 put in sunlight, it decomposes into water.
H2O2 (Sunlight)→ H2o

Displacement Reaction:
• The chemical reactions in which a more reactive element displaces a less reactive
element from a compound is known as Displacement Reactions. (A + CD → C + AD)
• When zinc reacts with hydrochloric acid, it gives hydrogen gas and zinc chloride.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
 • When zinc reacts with copper sulphate, it forms zinc sulphate and copper metal.
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Double Displacement Reaction:

• Reactions in which ions are exchanged between two reactants forming new
compounds are called Double Displacement Reactions. (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) (Precipitate) + 2NaCl(aq)
• When sodium hydroxide (a base) reacts with hydrochloric acid, sodium chloride and
water are formed.
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

Precipitation Reaction:
• The reaction in which precipitate(substance in soluble in water) is formed by the
mixing of the aqueous solution of two salts is called Precipitation Reaction.
• Examples: When the solution of silver nitrate reacts with the solution of sodium
chloride, it gives silver chloride and sodium nitrate.
AgNo3(aq) + NaCl(aq) → AgCl ↓ (aq) + NaNo3(aq)
BaCl2 (aq) + Na2SO4 (aq) → BaSO4 ↓ (s) + NaCl (aq)

Neutralization Reaction:
• The reaction in which an acid reacts with a base to form salt and water by an
exchange of ions is called d Neutralization Reaction.
• Examples: When the solution of sodium hydroxide reacts with the solution of
hydrochloric acid, sodium chloride and water is formed.
NaOH(aq) + Hcl(aq) → Nacl (aq) + H2o(aq)
Oxidation:
• Addition of oxygen or removal of hydrogen from a compound is known as Oxidation.
Elements or compounds in which oxygen is added or hydrogen is removed are called
to be Oxidized.
• Example: Rusting of iron Reaction showing the formation of rust when reacted with
water in presence of oxygen (atmosphere).
4Fe + 3O₂+ XH₂O → 2Fe₂O₃.XH₂O
Oxidizing agent:
• The substance which gives oxygen for oxidation is called an Oxidizing agent.
• The substance which removes hydrogen is also called an Oxidizing agent.
Reduction:
• The reaction which involves addition of hydrogen or removal of oxygen from a
substance is called reduction reaction.
• Example: Photosynthesis is a reduction reaction.
Reducing agent:
• The substance which gives hydrogen for reduction is called a Reducing agent.
• The substance which removes oxygen is also called a Reducing agent.

Exothermic reaction:
• Chemical reactions in which heat energy released are known as exothermic reactions.
Burning of methane gas in air produces heat energy.
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)+ heat
• All combustion reactions are exothermic reactions.

Endothermic reaction:
• Reactions which involve absorption of energy are known as endothermic reactions.
• Example:Formation of nitric oxide from nitrogen and oxygen.
N2 (g)+ O2 (g) + Heat → 2NO

Combustion reaction:
• Burning of a substance in the presence of oxygen which resulting release of energy is
called combustion.
• Example: Combustion of ethylene gas.
C2H4 + 3O2 → 2CO2 + 2H2O

Reactivity Series
• The reactivity series is a list of metals arranged in the order of their decreasing
activities.
Reaction of Metals With Water:
Cold Sodium, potassium and calcium readily react with cold water, to produce
Water hydroxides.

Hot Magnesium does not react with cold water, but reacts with hot water to produce
Water hydroxide.

ZInc and aluminium react only with steam, to produce oxides. Iron's reaction with
Steam cold water is very slow and can take days. But when iron reacts with steam, then
iron oxide is formed.
When a metal reacts with water, it forms metal hydroxide. Hydrogen gas is also produced in
this reaction. Following is the general equation for this reaction.

Metal + Water ⇨ Metal hydroxide + Hydrogen

Examples:
(i) 2Na + 2H2O → 2NaOH + H2 + Heat
(ii) Ca + 2H2O → Ca(OH)2 + H2
(iii) Mg + 2H2O → Mg(OH)2 + H2
(iv) 2Al + 3H2O → Al2O3 + 3H2
(v) 3Fe + 4H2O → Fe3O4 + 4H2

Reaction of metals with acids (Dilute)

• Metal + Dilute acid → Salt + H2

• Cu, Ag, Hg do not react with dil. acids.
• Examples:
(i) Fe + 2HCl → FeCl2 + H2
(ii) Mg + 2HCl → MgCl2+ H2

(iii) Zn + 2HCl → ZnCl2 + H2

(iv) 2Al + 6HCl → 2AlCl3 + 3H2

Reaction of Metals with Solutions of other Metal Salts

• Metal A + Salt solution B → Salt solution A + Metal B

• Reactive metals can displace less reactive metals from their compounds in solution
form.
Fe + CuSO4→ FeSO4 + Cu

Reaction of Metals with Non-metals

• Reactivity of elements is the tendency to attain a completely filled valence shell.
 • Atoms of the metals lose electrons from their valence shell to form cation. Atom of
the non-metals gain electrons in the valence shell to form anion.

E.g.: Formation of NaCl

Na → Na+ + e-
2, 8, 1 2, 8
Sodium cation
Cl + e- → Cl-
2, 8, 7 2, 8, 8
Reaction of metals with air
• Metals combine with oxygen to form metal oxide.
• Metal + O2 → Metal oxide
Examples:
(i) 2Cu + O2 → 2CuO
Copper oxide (black)
(ii) 4Al + 3O2 → 2Al2O3
Aluminium oxide
(iii) 2Mg + O 2 → 2MgO
• The reactivity of different metals is different with O2.
• Na and K react so vigorously that they catch fire if kept in open so they are kept
immersed in kerosene.

Reversible Reactions
• It is said that the reaction can occur in both directions: the forward reaction (which
forms the products) and the reverse reaction(which forms the reactants)
• When writing chemical equations for reversible reactions, two arrows are used to
indicate the forward and reverse reactions
• Each one is drawn with just half an arrowhead – the top one points to the right, and
the bottom one points to the left: ⇌
• Example-The reaction for the Haber process which is the production of ammonia
from hydrogen and nitrogen:
N2 + 3H2 ⇌ 2NH3

Catalytic Reactions
• A catalyst is a substance that is used to speed up a chemical reaction but it is not
consumed by the reaction.
• Catalysts that increase the rate of a chemical reaction are positive catalysts.
• Catalysts that decrease the rate of reaction are negative catalysts.

Oxides
 • An oxide that combines with water to give an acid is termed as an acidic oxide.
• The oxide that gives a base in water is known as a basic oxide.
• An amphoteric solution is a substance that can chemically react as either acid or
base.
• However, it is also possible for an oxide to be neither acidic nor basic, but is a neutral
oxide.

Acidic Oxides
• Acidic oxides are the oxides of non-metals (Groups 14-17) and these acid
anhydrides form acids with water:
• Sulfurous Acid- SO2+H2O→H2SO3
• Sulfuric Acid - SO3+H2O→H2SO4

Basic Oxides
• Generally Group 1 and Group 2 elements form bases called base anhydrides or basic
oxides e.g., K2O(s)+H2O(l)→2KOH(aq) ,K2O(s)+H2O(l)→2KOH(aq)
• Basic oxides are the oxides of metals. If soluble in water, they react with water to
produce hydroxides (alkalis) e.g.,

CaO+H2O→Ca(OH)2

MgO+H2O→Mg(OH)2

Amphoteric Oxides
• Amphoteric oxides are metal oxide compounds that can behave as either an acid or a
base in a chemical reaction.
• These compounds are able to accept protons when acting as a base, or donate
protons when acting as an acid.
• Example- Lead oxide (PbO), Aluminum oxide (Al2O3), Vanadium (V) pentoxide (V2O5),
Manganese oxide (MnO2), Ferric oxide – Iron (III) oxide (Fe2O3), Zinc oxide (ZnO),
Aluminum oxide (Al2O3).
• Ferric oxide reacts with NaOH in the following reaction, forming a weak base
precipitate.
Fe2O3(s) + 6NaOH(aq) → 2Fe (OH)3(s) + 6Na+(aq)
• Na+ is a very weak acid that will form an aqueous solution.
• Ferric oxide reacts with HCl in the following reaction, forming an aqueous solution:
Fe2O3(s) + 6HCl(aq) → 2FeCl3(aq) + 3H2O(l)

Neutral Oxides-Neutral oxides show neither basic nor acidic properties and hence do
not form salts when reacted with acids or bases, e.g., carbon monoxide (CO); nitrous oxide
(N2O); nitric oxide (NO), etc., are neutral oxides.