Anything which occupies space or volume, has mass and can be perceived by our senses is called matter.

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Matter around Us
Matter
 Anything which occupies space or volume, has mass and can be perceived by our senses is called
matter.
 All materials are made up of matter, and matter consists of particles such as atoms, ions or molecules.
 The three states of matter—solid, liquid and gaseous—are based on the differences in physical
properties such as mass, volume, shape, rigidity, density and arrangement of particles.
Kinetic Theory of Matter
The kinetic theory of matter states that all matter is composed of particles which
 Have intermolecular spaces between them.
 Attract each other with a force.
 Are in a continuous random motion.
Properties of Matter
Solid State Liquid State Gaseous State
Mass Solids have a definite
mass.
Liquids have a
definite mass.
Gases have a
definite mass.
Volume,
Shape,
Rigidity
Solids have a definite
shape. They maintain
their shape even when
they are subjected to an
external force, i.e. they
are rigid.
Liquids do not have
a fixed shape but
have a fixed
volume. Liquids take
up the shape of the
container in which
they are poured.
They are less rigid.
Gases neither have
a definite shape nor
have a definite
volume. They fill up
the container
completely. They
are not rigid.
On the basis
of kinetic
theory –
Intermolecul
ar space
The space between the
particles is very less.
The space between
the particles is
slightly more as
compared to solids
but still very less.
The particles of
liquids can slip and
slide over each
other.
The particles are
much farther apart
from one another as
compared to solids
and liquids. They
have a very
disorderly
arrangement of
particles compared
to solids and liquids.

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On the basis
of kinetic
theory –
Force of
attraction
The force of attraction
between the particles is
strong. Thus, particles in
a solid are closely
packed.
The force of
attraction between
the particles is
strong enough to
hold the particles
together but not
strong enough to
hold the particles in
a fixed position.
Thus, particles in a
liquid are not as
closely packed as in
solids.
The force of
attraction between
the particles is
negligible; hence,
particles of a gas
move freely in all
directions.
Gases can thus mix
or diffuse into other
gases.
On the basis
of kinetic
theory –
Movement
Solid particles vibrate
only about their mean
position.
Liquid particles are
free to move around
in the liquid only.
They can slip and
slide over each
other.
Gaseous particles
move with high
speed in all
directions and can
exert pressure on
the walls of the
container.
Density Solids have high density.
This is because the
number of particles in a
solid is more and the
intermolecular space is
minimum.
Liquids have less
density as
compared to solids
because the number
of particles is less
and the
intermolecular
space is more.
Gases have least
density as the
number of particles
is least and the
intermolecular
space is maximum.
Free surface Solids have an infinite
free surface.
Liquids have one
upper free surface.
Gases do not have
any free surface.
Miscibility or
Diffusibility
Solids do not diffuse with
other solid particles.
Liquids may diffuse
with other liquid
particles.
Gaseous particles
rapidly diffuse with
other gaseous
particles.
Compressibi
lity
Solids cannot be
compressed.
Liquids cannot be
compressed much.
The compressibility
of liquids is almost
negligible.
Gases can be
compressed easily.
Examples: LPG
cylinders used at
home and CNG
cylinders used in
vehicles.
Some solids may change
their shape when an
external force is applied,
but when that force is
removed, they can regain
their original shape. This
shows that some solids
are elastic.
Liquids show a
property called
viscosity. More
viscous liquids flow
slowly, while less
viscous liquids flow
easily.
__

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Change of State (Interconversion of Matter)
The phenomenon of change from one state of matter to another and then back to the original state is
called the interconversion of states of matter.
It is affected by changes in conditions such as
1. Changing the temperature.
2. Increasing or decreasing the pressure.
3. Changing both temperature and pressure.
Interconversion of states of matter

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Term Definition Process of change (at
particular temperature)
Melting The process of change of state from solid
to liquid on absorbing heat at a particular
temperature and at one atmospheric
pressure is called melting or fusion.
Melting (Solid  Liquid)
Vaporisation The process of conversion of a substance
from the liquid state to the gaseous state
at any temperature below its boiling point
is called evaporation or vaporisation.
Evaporation is a surface phenomenon.
Evaporation (Liquid  Gas)
Liquefaction The process of change from the gaseous
state to the liquid state at a particular
temperature is called liquefaction.
Liquefaction (Gas  Liquid)
Solidification The process of change of matter from the
liquid state to the solid state at a particular
temperature is called freezing or
solidification.
Freezing (Liquid  Solid)
Sublimation A change of state of a substance directly
from solid to gas without changing into a
liquid state (or vice versa) is called
sublimation.
Sublimation (Solid  Gas)
Melting point The constant temperature at which a solid
becomes liquid upon absorbing heat under
normal pressure is called the melting
point of that solid.
Melting point (Solid → Liquid)
Boiling point The constant temperature at which a liquid
becomes gas upon absorbing heat under
normal pressure is called the boiling
point of that liquid
Boiling point (Liquid → Gas)
Liquefaction
point
The constant temperature at which a gas
becomes liquid under normal pressure is
called the liquefaction point of that liquid.
Liquefaction point (Gas →
Liquid)
Freezing point The constant temperature at which a liquid
changes into a solid by giving out heat
energy is called the freezing point of that
liquid.
Freezing point (Liquid →
Solid)

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Interconversion of Matter on the basis of Kinetic Theory
Melting Vaporisation
On heating solids, the temperature rises.
 Particles gain energy.
 Intermolecular force of attraction
decreases.
 Intermolecular space between the
molecules of a solid increases.
 At the melting point, particles of
solid become free from the fixed
position and get converted into a
liquid.
On heating liquids, the temperature rises.
 Particles gain energy.
 Intermolecular force of attraction
decreases.
 Intermolecular space between the
molecules of a liquid increases.
 At the boiling point, particles become
completely free and convert into a
gas/vapour.
Liquefaction Solidification
On cooling gases, the temperature falls.
 Particles lose energy.
 Intermolecular attraction
increases.
 Intermolecular space between the
particles/molecules of gases
decreases.
 At liquefaction point, particles
slow down, come close and get
converted into a liquid.
On cooling liquids, the temperature falls.
 Particles lose energy.
 Intermolecular attraction increases.
 Intermolecular space between the
particles/molecules of gases
decreases.
 At solidification/freezing point,
particles slow down, come close and
get converted into a solid.
Dalton’s Atomic Theory, 1808
In 1808, the English chemist John Dalton gave a systematic idea about the structure of an atom. His ideas
are grouped together and known as Dalton’s atomic theory.
Postulates of Dalton’s Atomic Theory
1. All matter is made up of very tiny, indivisible and indestructible particles called atoms.
2. Atoms can neither be created nor be destroyed.
3. Atoms of same elements are alike in all respects.
4. Atoms combine in the ratio of small whole numbers to form compounds or molecules.
Modern Atomic Theory
 An atom comprises three fundamental sub-atomic particles—electron, proton and neutron.
 Atoms of the same element may not be alike in all respects.
 Isotopes are atoms of the same element differing in properties.

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Atoms and Molecules
Atom
 An atom is the basic unit of matter.
 An atom may or may not have an independent existence.
 An atom is further divisible into protons, neutrons and electrons.
Molecules
 Atoms of the same element or different elements can join together to form molecules.
 A molecule is the smallest particle of a substance which can normally exist on its own.
 A molecule retains the physical and chemical properties of the substance.
Atoms – Structure
Basic Structure of an Atom
 An atom consists of three fundamental or sub-atomic particles—proton, neutron and electron.
 The central core or the nucleus of every atom consists of protons and neutrons. Electrons revolve
around the nucleus in different orbits.
Fundamental particles of an atom
Atomic Number and Atomic Mass Number
Atomic Number (Z)
The number of protons in an atom of an element is called the atomic number of the element. It is denoted
by the letter Z.
Z = p = e
(Atomic no.) (No. of protons) (No. of electrons)
Mass Number (A)
The sum of the number of protons and the number of neutrons present in the nucleus of an atom of an
element is called the mass number of that element. It is denoted by the letter A.
A = p + e
(Mass no.) (No. of protons) (No. of electrons)

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Electronic Configuration
Orbits
Electrons revolve around the nucleus in an imaginary path called an orbit or shell.
 The maximum number of electrons which can be present in any shell of an atom is given or
represented by the formula 2n2
.
 Here, n is the principal quantum number which is equal to the number of shells as counted from the
nucleus.
Shell
designation
Shell number
(n)
Formula
2n2
Maximum number of
electrons in each shell
K-shell 1 2 × (1)2
2
L-shell 2 2 × (2)2
8
M-shell 3 2 × (3)2
18
N-shell 4 2 × (4)2
32
 The outermost shell cannot hold more than 8 electrons.
 As soon as the maximum capacity of a shell is satisfied, a new shell is filled.
Electronic Configuration of Sodium
Atomic number of sodium = 11. Thus, we know that an atom of sodium contains 11 electrons.
The orbit-wise distribution of electrons in a sodium atom is as follows:
Orbit number of
electrons
Maximum number of electrons
K-shell 2n2
= 2 × 12
= 2 electrons
L-shell 2n2
= 2 × 22
= 8 electrons
M-shell Remaining = 1 electron
Thus, the electronic configuration of sodium is (2,8,1).
Relative Atomic Mass
The relative atomic mass of an element is defined as the ratio of the average mass of an atom of the
element to 1/12th
of the mass of one carbon-12 atom.
Atomic mass or relative atomic mass is simply the number of times one atom of an element is heavier
than either the mass of an atom of hydrogen or the 1/12th
mass of an atom of carbon-12.
Relative Molecular Mass
The relative molecular mass of an element or compound is the ratio of mass of one molecule of a
substance to the mass of 1/12th
the mass of one atom of carbon-12.
 Molecular mass or relative atomic mass is simply the number of times one molecule of an element is
heavier than either the mass of an atom of hydrogen or the 1/12th
mass of an atom of carbon-12.
Isotopes and their application
 Atoms of the same elements differing in the number of neutrons in their nuclei are known as isotopes.
Thus, isotopes of an element have the same atomic number but different atomic mass number.
 Isotopes of an element have similar chemical properties but different physical properties.

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Formation of Compounds from Atoms
1. Stable and unstable Electronic Configuration
Stable electronic configuration
Noble gases have stable electronic configuration.
Their valance or outermost shell is completely filled.
They have 2 electrons (He) or 8 electrons (Ne, Ar, Kr, Xe, Rn) in their outer shell.
They do not gain, lose or share electrons.
Unstable electronic configuration
Atoms of other elements have unstable electronic configuration.
Their valance or outermost shell is incompletely filled.
They tend to attain stable electronic configuration of the nearest inert gas by
 Gaining or losing electrons (electron transfer).
 Sharing electrons.
2. Atoms Combine to Form Compounds
Electron transfer – Electrovalency
Sodium (Na) atom
 Electronic configuration = 2,8,1
 Nearest inert gas = Neon [2,8]
 It loses one electron from the outermost
shell (valence shell) to attain stability.
Chlorine (Cl) atom
 Electronic configuration = 2,8,7
 Nearest inert gas = Argon [2,8,8]
 It gains one electron in the outermost
shell (valence shell) to attain stability.
Sharing of electrons – Covalency
Oxygen (O) atom
 Electronic configuration = 2,6
 Nearest inert gas = Neon [2,8]
 Needs two electrons to attain stability.
Oxygen (O) atom
 Electronic configuration = 2,6
 Nearest inert gas = Neon [2,8]
 Needs two electrons to attain stability.

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Classification of Matter (Elements, Compounds and Mixtures)
Matter may be broadly classified into elements, compounds and mixtures.
Element
 An element is a pure substance composed of only one kind of atom.
 An element cannot be broken down into two or more simple substances by any physical or chemical
means.
Characteristics of an Element
1. An element is made up of only one kind of atom.
2. An element is a pure and homogeneous substance.
3. An element has fixed melting and boiling points.
4. An atom is the smallest particle of an element which takes part in a chemical reaction.
5. An element may chemically react with another element or compound.
6. An element can occur in the solid, liquid or gaseous state.
Classification of Elements
 Have metallic
lustre.
 Are good
conductors of
heat and
electricity.
 Are malleable
and ductile.
 Are solids.
 Contain one kind
of atom
(monoatomic).
 Do not have
lustre.
 Are bad
conductors of
heat and
electricity.
 Are neither
malleable nor
ductile.
 Are solids, liquids
and gases.
 Contain two kinds
of atoms
(monoatomic or
diatomic).
 Properties are
mid-way between
metals and non-
metals.
 Contain one kind
of atom
(monoatomic).
 Are gaseous in
nature.
 Are chemically
inert.
 Occur in free
state in traces in
the atmosphere.
 Contain one kind
of atom
(monoatomic).

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Compound
 A compound is a pure substance composed of two or more elements combined chemically in a fixed
proportion by mass.
Characteristics of Compounds
1. Components in a compound are present in a definite proportion.
2. They have a homogeneous composition.
3. Particles in a compound are of one kind.
4. A compound is made up of one or more atoms of the same or different elements.
5. In a compound, the elements are present in a fixed ratio by mass.
6. A compound can be divided into simpler substances by a chemical process.
7. The physical and chemical properties of a compound are completely different from those of its
constituents.
A Comparative Study between Elements and Compounds
Element Compound
1. It is a pure substance which cannot
be converted into simpler substances
by any physical or chemical means.
1. It is a pure substance made up of two
or more elements combined
chemically in a fixed ratio.
2. It is made up of atoms of only one
kind.
2. It is made up of two or more different
kinds of atoms.
3. The molecules are made up of one or
more atoms.
3. The molecules are made up of two or
more atoms.
4. Elements cannot be broken down into
two or more simpler substances by
any physical or chemical means.
4. A compound can be divided into
simpler substances only by chemical
means.
5. Elements have their own set of
properties.
5. Properties of compounds are
different from their constituent
elements.

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Mixtures
 A mixture is defined as matter composed of two or more substances whose particles are in contact but
are not chemically combined and have not lost their individual properties.
 The properties of a mixture vary according to the proportions of the components present in it.
Types of Mixtures
Homogeneous mixtures: A mixture which has uniform composition and properties throughout its mass is
called a homogeneous mixture. Examples: Sugar solution, salt solution
Heterogeneous mixture: A mixture which has different composition and properties in different parts of its
mass is called a heterogeneous mixture. Examples: Sand mixed with salt, sugar in oil
Types of mixture Homogeneous Heterogeneous
Solid in solid Alloys (Bronze - Cu, Zn, Sn) Gun powder (charcoal,
sulphur, nitre)
Solid in liquid Iodine in alcohol, sugar in
carbon disulphide, sugar in
water, salt in water
Sugar in oil, sand in water
Liquid in solid Amalgam (Hg + Au) Water in sponge
Liquid in liquid Methanol in water, acetone
in water
Oil in water, kerosene in
water
Gas in liquid HCl in water Helium in water
Liquid in gas Moisture in air Mist, fog
Gas in gas Pure air Air in industries
A Comparative Study between Compounds and Mixtures
Compound Mixture
1. It is obtained by the chemical
combination of more than one
element.
1. It is obtained by the physical
combination of either elements,
compounds or both.
2. The composition of elements
present in a compound is fixed.
2. The composition of elements present
in a mixture is not fixed.
3. The properties of a compound are
different from those of its elements.
3. It shows the properties of all its
constituent elements.
4. Its constituents can be separated by
using only chemical and
electrochemical methods.
4. Its constituents can be separated
using physical methods.
5. A compound is always
homogeneous in nature.
5. The mixtures can be homogeneous
or heterogeneous.

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Methods of Separation
 Separation of Solid–Solid Mixtures
Sublimation
To separate the mixture of sublimable solid X + non-sublimable solid Y.
Magnetic separation
To separate the mixture of magnetic solid X + non-magnetic solid Y.
Solvent extraction method
To separate the mixture of insoluble solid X + soluble solid Y.
Chromatography – For Complex Mixtures
To separate the mixture of different solid constituents in a liquid constituent.
 Separation of Solid–Liquid Mixtures
Filtration
To separate the mixture of insoluble solid X from liquid component Y.
Sedimentation and Decantation
Insoluble solid X from liquid component Y.
Evaporation
To separate the mixture of soluble solid X from liquid component Y.
Distillation
To separate the mixture of soluble solid X from liquid component Y.
 Separation of Liquid–Liquid Mixtures
Separating funnel
To separate the mixture of immiscible heavier liquid X from immiscible lighter liquid Y.
Fractional distillation
To separate the mixture of miscible liquid X with lower boiling point and miscible liquid Y with higher
boiling point.

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Basic Chemistry – Terminology and Reactions
Symbol
A symbol is a short form of the name of an element.
The specific abbreviation used to denote the name of an element is called its symbol.
Representation of a Symbol
Dalton’s Symbol, 1808
 Dalton was the first scientist to use figurative symbols for atoms of some of the elements. His symbols
represented the ‘element’ as well as ‘one atom’ of that element.
Drawbacks of Dalton’s Symbols
 Dalton's symbols for elements were difficult to draw and inconvenient to use. Thus, they are not used
any more.
Berzelius Symbol, 1814
In 1814, the Swedish Chemist Jöns Jakob Berzelius devised a system using letters of the alphabet. He
put forward certain points for presentation.
1. In most cases, the first letter of the name of an element was taken as the symbol for that element and
written in capitals.
Name Symbol
Carbon C
2. In some cases, the initial letter of the name in capital and its second letter in small were used.
Name Symbol
Calcium Ca
3. The symbols for some elements were derived from their Latin names.
English name of
the element
Latin name of the
element
Symbol
Sodium Natrium Na
Potassium Kalium K
Silver Argentum Ag
 The method suggested by Berzelius forms the basis of the IUPAC (International Union of Pure and
Applied Chemistry) system of chemical symbols and formulae.

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Radicals and Valency
What are Ions or Radicals?
Two or more non-metals which collectively accept or donate one or more electrons and become
negatively or positively charged in the process are called radicals.
Types of Ions or Radicals
Radicals are either positively charged or negatively charged.
Positively charged radicals (ions) are called cations, and negatively charged radicals (ions) are called
anions.
Valency
It is the number of hydrogen atoms which can combine with or displace one atom of the element or
radical.
Number of atoms of
Hydrogen (H)
Combining
element
Molecule Valency
1 1 atom of Chlorine
(Cl)
Hydrogen chloride
(HCl)
1
2 1 atom of Oxygen
(O)
Water
(H2O)
2
Variable Valency
Sometimes, the same element may exhibit one valency in one compound and another valency in other
compound. This property is called variable valency.
Example
Element Symbol Valencies exhibited
(variable valencies)
Copper Cu 1, 2 Cu1+
, Cu2+
Valency Chart
List of common electrovalent
positive ions or radicals
List of common electrovalent negative ions or radicals
1. Monovalent electropositive
ions
Cuprous Cu+ Copper [I]
2. Bivalent electropositive ions
Cupric Cu
2+
Copper [II]
3. Trivalent electropositive ions
Aluminium Al
3+
4. Tetra-positive ions
Plumbic Pb
4+
1. Monovalent electronegative ions
Acetate CH3COO−
2. Bivalent electronegative ions
Sulphate SO4
2–
3. Trivalent electronegative ions
Phosphide
4. Tetravalent electronegative ions
Carbide C
4–

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Molecular Formula
A molecular formula, also known as a chemical formula, is a combination of elemental symbols and
subscript numbers which is used to show the composition of a compound.
Writing the Molecular Formula
Step 1: Write the symbol of a basic radical (element with positive valency) to the LHS and that of the acid
radical (element with negative valency) to the RHS.
Step 2: Write the valency number of each of the respective radicals at the right hand top of
its symbol.
Step 3: Interchange the valency number. Ignore the (+) and () signs.
Step 4: Write the interchanged number.
Step 5: Write the compound’s formula.
Chemical Equation
The representation of a chemical reaction with the help of chemical formulae of the reactants and
products is a chemical equation.
The reaction can be represented by either a word equation or a chemical equation using symbols and
formulae.
Word equation:
Chemical equation:
How are Chemical Equations Represented?
 In a chemical reaction, the reactants are written on the LHS and the products on the RHS of the
equation.
 An arrow (→) pointing towards the products is inserted between the reactants and the products. It also
represents the direction of the reaction.
 A single arrow (→) indicates the direction in which the reaction proceeds.
 A double arrow ( ) indicates a reversible reaction, i.e. products recombine to form reactants.
 A plus sign (+) is inserted between two or more reactants and products formed.
 A chemical reaction can be characterised by factors such as change of state, change in colour,
evolution of a gas, change in temperature, formation of a precipitate and evolution of heat, light or
sound.
Balancing an Equation
 In a balanced chemical equation, the total number of atoms of each element in the reactants on the
LHS of the equation is the same as the number of atoms in the products formed on the RHS of the
equation.
 The total mass of the reactants is equal to the total mass of the products, or the number of atoms of
each element before the reaction and after the reaction is equal.

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Steps involved in Balancing a Chemical Equation
Consider the chemical reaction between magnesium and oxygen to understand the steps involved in
balancing a chemical equation.
1. Let us first write the word equation for this reaction.
Magnesium + Oxygen → Magnesium oxide
2. Write the chemical equation for the reaction between magnesium and oxygen.
Mg + O2 → MgO
3. Count the number of times an element occurs on both LHS and RHS in this equation.
Mg + O2 → MgO
Component Reactant Product
Magnesium 1 1
Oxygen 2 1
This is an unbalanced equation.
4.
 Choose a reactant or product which has the maximum number of atoms in it. In this equation, we shall
select MgO, i.e. magnesium oxide and the element oxygen in it.
 To balance the oxygen atoms, let us multiply the magnesium oxide molecule by 2 on the RHS.
The equation can now be expressed as
Mg + O2 → 2MgO
Component Reactant Product
Magnesium 1 2
Oxygen 2 2
5. There are two oxygen atoms on either side of the equation, but one magnesium atom on the reactant's
side and two on the product's side. Therefore, multiply the magnesium atom by 2 on the LHS.
Component Reactant Product
Magnesium 2 2
Oxygen 2 2
The balanced equation is 2Mg + O2 → 2MgO
 The number of atoms of each element of reactants = The number of atoms of each element of
products

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Chemical Reaction
Chemical reactions involve the transfer of matter from one substance to another substance during a
chemical change.
Main Types of Chemical Reactions
1. Direct
combination or
synthesis
A chemical reaction in which two or more substances combine to form a
single product is called a combination reaction or synthesis.
2. Decomposition A chemical reaction in which a single compound splits into two or more
simple substances is called a decomposition reaction.
3. Thermal
decomposition
During thermal decomposition, a chemical compound breaks into
simpler compounds. The simpler compounds do not reunite to form the
original compound on cooling.
4. Reversible
reaction
A reaction in which the direction of a chemical change can be easily
reversed by changing the conditions under which the reaction occurs is
called a reversible reaction.
5. Thermal
dissociation
A reaction in which a substance dissociates into two or more simpler
substances on the application of heat is called a thermal dissociation
reaction. It is a reversible reaction.
6. Displacement Reactions in which the more reactive element displaces the less
reactive element from its compound are called displacement reactions.
7. Double
displacement
Reactions in which ions of the reactants exchange places to form two
new compounds are called double displacement reactions.
In double displacement reactions, the two reactants taking part are
generally water soluble, and one of the products is soluble and the other
being insoluble separates out as a solid.
8. Double
decomposition
A type of chemical change in which two compounds in a solution react
to form two new compounds by the mutual exchange of radicals.
Usually, a solid is formed as a result of the reaction.
These reactions are of two types—precipitation reactions and
neutralisation reactions.
Precipitation reactions
The insoluble solid formed during double displacement reactions is
called a precipitate. Reactions in which a precipitate is formed as one of
the products are also called precipitation reactions.
Neutralisation reactions
The reaction between an acid and a base to form a salt and water is
called a neutralisation reaction.
9. Catalytic
reaction
A chemical reaction which involves the use of a catalyst.
Catalyst: It is a compound which alters the rate of a reaction but does
not take part in a chemical reaction.
Examples: V2O5, Fe2O3, finely divided iron (Fe), platinum (Pt)

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10. Exothermic
and
Endothermic
reactions
Exothermic reaction
Chemical reactions which proceed with the evolution of heat energy
are called exothermic reactions.
A + B → C + Δ
Endothermic reactions
Chemical reactions which proceed with the absorption of heat energy
are called exothermic reactions.
11. Oxidation and
reduction
reactions
(redox)
Oxidation Reactions
Reactions which involve the addition of oxygen or the removal of
hydrogen are called oxidation reactions.
What is an oxidising agent?
The substance which loses oxygen or an electronegative radical is
called an oxidising agent. The substance which gains hydrogen or an
electropositive radical is also called an oxidising agent.
Examples: Oxygen, chlorine
Reduction reactions
The addition of hydrogen to a substance is called reduction. The
removal of oxygen from a substance is also called reduction. Reactions
which involve the addition of hydrogen or removal of oxygen are called
reduction reactions.
What is a reducing agent?
The substance which loses hydrogen or an electronegative radical is
called a reducing agent. The substance which gains oxygen or an
electronegative radical is also called a reducing agent.
Examples: Hydrogen, carbon monoxide, hydrogen sulphide
Redox reactions
Oxidation and reduction reactions occur together. When oxidation
occurs in one substance, reduction occurs in the other substance. Such
reactions are called redox reactions.
When an atom or a group of atoms loses electrons, oxidation occurs. In
this reaction, hydrogen sulphide is oxidised to sulphur.
Differences between Oxidation and Reduction
Oxidation Reduction
1. Addition of oxygen
2. Removal of hydrogen
3. Loss of electrons
1. Removal of oxygen
2. Addition of hydrogen
3. Gain of electrons
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