Matter in Our Surroundings Class 9 Science Revision Notes 2026-27

Matter in Our Surroundings Class 9 Science Revision Notes : Understanding the chapter Matter in Our Surroundings is very important for students of Central Board of Secondary Education because it builds the basic foundation of Science. These matter in our surroundings class 9 science notes are prepared in simple and easy language so students can understand all concepts without confusion. In this chapter, students learn about matter, particles of matter, states of matter, evaporation, diffusion, and changes between solid, liquid, and gas. These topics are important for school exams as well as future science studies.

The matter in our surroundings class 9 notes are based on the latest syllabus of the CBSE Board and help students revise quickly before exams. Many students also search for matter in our surroundings class 9 science notes pdf free download to study anytime on mobile or computer. These notes include important definitions, examples from daily life, short explanations, and key points for quick revision.

Along with notes, students can also practice matter in our surroundings class 9 questions answers to improve their understanding and score better marks in exams. These Class 9 Science notes are useful for homework, revision, and exam preparation. The chapter also explains scientific concepts like kinetic energy, intermolecular force, melting point, boiling point, and latent heat in a very easy way. Sometimes students may feel this chapter is little difficult, but with regular study and practice it becomes much more easier.

What Is Matter?

Everything you see, touch, or breathe is made of matter. The book on your desk, the water in your glass, the air filling your lungs all of these are matter. But what exactly defines matter in science?

Matter is anything that occupies space, has mass, and offers resistance to force. A balloon filled with air occupies space and has measurable weight confirming that even invisible substances like air qualify as matter.

This single concept forms the foundation of Class 9 Chemistry, and understanding it deeply will help you grasp everything from chemical reactions to real-world phenomena like sweating, cooking, and water purification.

Physical Nature of Matter

Matter Is Made of Tiny Particles

Modern science confirms that all matter is composed of extremely small units called atoms and molecules. These particles are:

• Incredibly small a few crystals of potassium permanganate (KMnO₄) can colour hundreds of litres of water, demonstrating that a single crystal contains millions of particles.
• Always moving particles in any state of matter are in constant motion. This motion increases with temperature.
• Separated by spaces when sugar dissolves in water, the sugar molecules occupy the spaces between water molecules, which is why the total volume does not increase.
• Attracted to each other forces of cohesion bind particles together. The stronger this force, the harder it is to break the substance (iron > ice > chalk).

Evidence: Diffusion and Brownian Motion

Diffusion is the intermixing of particles of two different substances due to their spontaneous movement from regions of higher concentration to lower concentration.

Example: When a perfume bottle is opened in one corner of a room, its scent spreads throughout the room within minutes, as the fragrant particles diffuse through the air.

Brownian Motion was first described by botanist Robert Brown in 1827. He observed that pollen grains suspended in water moved in a continuous, random, zig-zag pattern under a microscope. This movement occurs because the much smaller, invisible water molecules constantly collide with the pollen grains from all sides at different speeds providing direct evidence that liquid particles are in perpetual motion.

Note: Brownian motion increases with temperature, because higher temperature means faster particle movement, resulting in more energetic and frequent collisions.

Classification of Matter: The Three States

On the basis of physical properties, all matter is classified into three primary states:

1. Solids

A solid has a definite shape, definite mass, and definite volume.

Examples: Ice, wood, iron, chalk, coal.

Properties of Solids:

• Negligible compressibility particles are packed closely together.
• Very strong intermolecular forces of attraction.
• High density compared to liquids and gases.
• Do not flow; maintain their shape under normal conditions.
• Very slow diffusion between solid particles.

2. Liquids

A liquid has a definite mass and volume but no definite shape. It takes the shape of its container.

Examples: Water, alcohol, mercury, milk.

Properties of Liquids:

• Slightly compressible.
• Intermolecular forces weaker than in solids.
• Can flow classified as fluids.
• Diffuse into other liquids, though more slowly than gases.
• Expand more than solids on heating.

3. Gases

A gas has a definite mass, but neither definite shape nor definite volume. It fills the entire container it is placed in.

Examples: Oxygen (O₂), nitrogen (N₂), hydrogen (H₂).

Properties of Gases:

• Highly compressible used in LPG cylinders, oxygen tanks, and CNG vehicles.
• Negligible intermolecular forces.
• Very low density.
• Diffuse rapidly to form homogeneous mixtures.
• Gaseous particles move randomly and exert pressure on container walls.

Comparative Table: Three States of Matter

Liquids and gases are both classified as fluids because they have the ability to flow.

Interconversion of States of Matter

The state of any substance can be changed by altering temperature or pressure. This is called interconversion of states of matter.

By Changing Temperature

Solid → Liquid: Melting (Fusion)

When a solid absorbs heat energy, the kinetic energy of its particles increases. Eventually, the particles overcome the intermolecular forces holding them in fixed positions and begin to move freely the solid melts.

• Melting Point of ice = 0°C (273.16 K)
• Latent Heat of Fusion of ice = 3.34 × 10⁵ J/kg

Water particles at 0°C have more energy than ice particles at the same temperature, because water has already absorbed the latent heat of fusion.

Liquid → Solid: Freezing (Solidification)

When a liquid releases heat, its particles slow down and intermolecular forces draw them back into fixed positions. The liquid freezes.

• Freezing point = Melting point numerically (0°C for water).

Liquid → Gas: Boiling (Vaporisation)

When a liquid is heated to its boiling point, particles gain enough energy to completely overcome intermolecular attraction and escape into the gaseous state.

• Boiling Point of water = 100°C (373 K)
• Latent Heat of Vaporisation of water = 22.5 × 10⁵ J/kg

Steam particles at 100°C carry more energy than water particles at the same temperature because steam has absorbed extra energy as latent heat.

Gas → Liquid: Condensation (Liquefaction)

When a gas loses heat, its particles slow down and are drawn together into the liquid state. This is called condensation.

• Condensation point = Boiling point (100°C for water).

Solid → Gas (Directly): Sublimation

Some solids convert directly to vapour on heating, without passing through the liquid state. This is sublimation.

Sublimable substances: Ammonium chloride (NH₄Cl), iodine, camphor, naphthalene, dry ice (solid CO₂).

• The substance that sublimes is called the sublime.
• The solid formed by cooling its vapours is the sublimate.

By Changing Pressure

Increasing pressure forces gas particles closer together, reducing the intermolecular spaces until the gas liquefies.

Example: Carbon dioxide gas (CO₂) can be liquefied by compressing it to 70 times atmospheric pressure. Solid CO₂ is called Dry Ice used to deep-freeze food and keep ice cream cold.

When pressure is lowered, the boiling point of a liquid also decreases, enabling faster vaporisation.

Evaporation

Evaporation is the conversion of a liquid into vapour at any temperature below its boiling point. Unlike boiling, it is a surface phenomenon only the topmost particles (those with the highest kinetic energy) escape into the vapour phase.

Factors Affecting the Rate of Evaporation

Cooling Effect of Evaporation

When a liquid evaporates, it absorbs latent heat of vaporisation from its surroundings, lowering the temperature of the surface it touches.

Practical examples:

• Applying spirit (ether) on the palm causes a cooling sensation.
• Sweating is the body's natural cooling mechanism perspiration evaporates and removes heat from the skin surface.
• Cotton clothes in summer are preferred because cotton absorbs sweat efficiently, allowing evaporation and keeping the body cool.
• Water droplets on a cold glass: Water vapour in the surrounding air loses energy when it contacts the cold glass surface and condenses back into liquid droplets.

Evaporation vs. Boiling

Other States of Matter

Plasma

Plasma is a fourth state of matter consisting of super-energised, ionised gas particles. When electrical energy passes through a gas, the particles become ionised, forming plasma that glows with a characteristic colour depending on the type of gas.

Examples: Neon sign bulbs (neon gas), fluorescent tubes (helium gas).

The Sun and other stars exist in the plasma state, which is why they glow with such intense light.

Bose-Einstein Condensate (BEC)

BEC is formed by cooling a gas of extremely low density (about one-hundred-thousandth the density of normal air) to temperatures just above absolute zero. At these super-low temperatures, particles lose almost all kinetic energy and essentially merge into the same quantum state.

Pure Substances

A pure substance is a homogeneous material containing particles of only one kind, with a definite and fixed set of properties.

Examples: Iron, silver, oxygen, sulphur, carbon dioxide, water.

Characteristics of a Pure Substance:

• Homogeneous in nature.
• Has a fixed, unique set of physical and chemical properties.
• Composition cannot be changed by any physical method.

Elements

An element is a pure substance that cannot be broken down into simpler substances by chemical means. Elements are composed of atoms with the same atomic number.

Examples: Hydrogen, oxygen, copper, nitrogen, mercury.

Classification of Elements:

• By Physical State at Room Temperature:
  • 2 liquid elements: Mercury (metal) and Bromine (non-metal)
  • 11 gaseous elements: H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, Rn
  • Remaining 102 elements: Solids
• Metals vs. Non-Metals: There are 93 metals and 22 non-metals. Non-metals include solids (carbon, sulphur), one liquid (bromine), and gases (H₂, O₂, N₂, F₂, Cl₂) plus noble gases (He, Ne, Ar, Kr, Xe, Rn).
• Metalloids: Elements that exhibit properties of both metals and non-metals neither good conductors nor insulators, functioning as semiconductors. Examples: Boron (B), Silicon (Si), Germanium (Ge). Also called semi-metals.
• Normal vs. Radioactive Elements: Elements with atomic numbers 1–82 are generally stable (normal). Elements from atomic number 83 onwards are radioactive and emit harmful radiation.

Hydrogen is the lightest element in the periodic table.

Compounds

A compound is a pure substance composed of two or more elements chemically combined in a fixed ratio by weight, separable only by chemical means.

Example: Water (H₂O) hydrogen and oxygen in a 1:8 ratio by weight, decomposable only by electrolysis.

Compounds can be classified as:

• Acids: Sulphuric acid, hydrochloric acid, nitric acid, formic acid.
• Bases: Sodium hydroxide, potassium hydroxide, zinc hydroxide, calcium hydroxide.
• Salts: Ammonium chloride, zinc sulphate, lead nitrate, calcium carbonate. Salts are formed from the chemical reaction between acids and bases.

Mixtures

A mixture is formed when two or more substances are combined without undergoing any chemical change, each retaining its individual properties.

Types of Mixtures

1. Homogeneous Mixtures (Solutions) Constituents are uniformly distributed throughout. No visible boundaries between components. Examples: Salt in water, copper sulphate solution, alloys (brass, bronze).

2. Heterogeneous Mixtures Constituents are not uniformly distributed; distinct boundaries are visible. Examples: Sand and salt, iron powder and sulphur, muddy water.

Differences: Mixtures vs. Compounds

Why Air Is a Mixture (Not a Compound)

• Its composition varies with altitude and location (less oxygen at higher altitudes; more CO₂ in industrial areas).
• Its components (O₂, N₂, H₂O vapour) can be separated by physical methods such as fractional distillation.
• No heat or light is evolved when its components mix.
• Each component retains its properties oxygen supports combustion; CO₂ turns lime water milky.

Why Water Is a Compound (Not a Mixture)

• Its composition is always fixed: 1 part hydrogen : 8 parts oxygen by weight.
• It cannot be separated by physical means only by electrolysis.
• Heat is released when hydrogen and oxygen combine.
• Its properties are completely different from its components (hydrogen is combustible; oxygen supports combustion yet water extinguishes fire).

Why Alloys Are Mixtures (Not Compounds)

• The ratio of constituent metals can vary (e.g., brass can be 58% or 60% copper with the remainder zinc both are still brass).
• Individual metals retain their chemical properties within the alloy.

Solutions

A solution is a homogeneous mixture of two or more substances. It consists of:

• Solvent the larger component (e.g., water in salt water, turpentine in paint).
• Solute the smaller component (e.g., salt in salt water, carbon dioxide in soda water).

Types of Solutions by State

Types by Saturation

• Saturated Solution: Holds the maximum possible amount of solute at a given temperature.
• Unsaturated Solution: Contains less solute than the maximum capacity.
• Supersaturated Solution: Contains more solute than a saturated solution (achieved by careful cooling of a hot saturated solution).

Concentration of a Solution

The concentration of a solution is the amount of solute dissolved in a given quantity of solution.

Mass Percentage:

Concentration (%) = (Mass of Solute / Mass of Solution) × 100

Volume Percentage (for liquid solutes):

Concentration (%) = (Volume of Solute / Volume of Solution) × 100

Concentration is a dimensionless percentage it has no units.

True Solutions, Colloids, and Suspensions

Three types of mixtures can be distinguished based on the particle size of the dispersed substance:

True Solution

• Particle size: < 1 nm (10⁻⁹ m)
• Transparent and clear no light scattering.
• Particles do not settle; cannot be filtered.
• Examples: Salt in water, sugar in water.

Colloidal Solution

• Particle size: 1 nm – 100 nm (10⁻⁷ to 10⁻⁵ cm)
• Translucent; particles scatter light Tyndall Effect.
• Particles do not settle; pass through filter paper but not semipermeable membranes.
• Heterogeneous in nature.
• Examples: Milk, blood, ink, starch solution, toothpaste, jelly, fog, mist.

Suspension

• Particle size: > 100 nm (> 10⁻⁵ cm)
• Opaque; heavy scattering.
• Particles settle on standing; can be filtered.
• Heterogeneous.
• Examples: Muddy water, chalk in water, paint.

Tyndall Effect

When a strong beam of light passes through a colloidal solution, the colloidal particles scatter the light, making the beam's path visible as a glowing cone (called the Tyndall cone). True solutions do not show this effect.

Everyday examples of Tyndall Effect:

• Visibility of sunlight beams through dusty rooms or forest canopies.
• Headlight beams appearing visible in foggy weather.

Classification of Colloids

Colloidal solutions can be separated using the process of centrifugation.

Comparison Table

Separation Techniques for Mixtures

Mixtures are separated based on differences in physical properties such as density, boiling point, solubility, volatility, and the ability to sublime or diffuse.

1. Evaporation

Used to separate a non-volatile solid from a liquid solvent.

Application: Recovering salt from saltwater, separating copper sulphate from water.

Method: The mixture is heated in a china dish. The liquid evaporates, leaving the solid residue behind.

2. Centrifugation

Separates fine or colloidal particles suspended in a liquid by spinning the mixture at high speed.

Principle: Denser (heavier) particles are thrown outward (to the bottom), while lighter particles stay near the top.

Applications:

• Separating cream from milk in dairy plants.
• Separating blood components in blood banks.
• Testing urine samples in diagnostic labs.
• Removing water from wet clothes in spin dryers.

3. Separating Funnel

Used to separate two immiscible liquids (liquids that do not dissolve in each other) based on their density difference.

Examples of immiscible pairs:

Industrial Application: Separating slag (waste) from molten metal during iron extraction in a blast furnace.

4. Sublimation

Separates mixtures where one component converts directly from solid to vapour without melting.

Examples:

Method: The mixture is heated in a china dish under an inverted funnel. The sublimable substance forms vapours that condense on the cooler funnel walls; the non-sublimable residue remains in the dish.

5. Chromatography

Separates dissolved constituents of a mixture based on their different rates of movement over an absorbent material (like filter paper) in a solvent.

The name comes from the Greek word "Kroma" (colour) the technique was first used to separate plant pigments.

Factors governing separation:

• Relative solubility of each component in the solvent.
• Relative affinity of each component for the adsorbent medium.

Applications:

• Separating colours from dyes and inks.
• Separating amino acids in biochemistry.
• Detecting drugs in blood samples.
• Separating sugar from urine in medical diagnostics.

Advantages:

• Works with very small sample quantities.
• No wastage of the substances being separated.

6. Distillation

Heating a liquid to form vapour and then cooling the vapour to recover it as a liquid. The recovered liquid is called the distillate.

Used when: recovering both solvent and solute from a solution (e.g., getting both salt and pure water from salt solution).

7. Fractional Distillation

Used when two miscible liquids have similar but distinct boiling points (difference typically less than 10°C). A fractionating column is used, which preferentially condenses the vapour with the higher boiling point.

Examples:

Major Industrial Application: Separation of gases from air air is first purified, then liquefied, then fractionally distilled to obtain liquid nitrogen (–195°C) and liquid oxygen (–183°C).

Physical and Chemical Changes

Physical Changes

A physical change alters the physical properties (state, texture, colour, shape, magnetism) of matter without changing its chemical composition (molecular structure).

Note:

• No new substance is formed.
• The change is temporary and reversible.
• No net gain or loss of energy.
• Mass of the substance remains unchanged.

Examples: Melting of ice, evaporation of water, dissolving salt in water, breaking glass, bending a metal wire, making ice cream, sublimation of camphor.

Chemical Changes

A chemical change alters the molecular composition of matter, forming one or more entirely new substances with different properties.

Main Point:

• New substances are always formed.
• The change is permanent and irreversible.
• Energy is always absorbed or released (heat, light, or electricity).
• Apparent change in weight of individual substances (though total mass is conserved).

Examples:

Other common chemical changes: Digestion of food, curdling of milk, ripening of fruit, photosynthesis, baking of cake, burning of charcoal, formation of wine, formation of biogas (gobar gas), decomposition of water by electrolysis.

Physical vs. Chemical Changes Summary Table

City Water Supply Application of Separation Techniques

River water used for drinking is purified through a multi-stage process that applies several separation techniques:

1. Sedimentation Water is stored in large tanks; alum (loading agent) is added to make suspended impurities settle faster.
2. Filtration Water passes through beds of sand, charcoal, and gravel, which trap remaining suspended particles.
3. Sterilisation (Chlorination) Bleaching powder or chlorine gas is added to kill pathogenic microorganisms (bacteria causing typhoid, cholera, jaundice, dysentery).
4. The purified water is then pumped into overhead tanks for city distribution.

Matter in Our Surroundings Class 9 Science Questions and Answers 

Q1. Define matter. Give three examples each of matter and non-matter.

Answer: Matter is anything that occupies space, possesses mass, and offers resistance to applied force.

Examples of matter: Wood, water, iron, air, sugar, stone.

Things that are NOT matter: Light, heat, sound, gravity, emotions these do not occupy space and have no mass.

Scientific note: Although light has momentum and energy, it has no rest mass, so it does not qualify as matter by the classical definition used in Class 9.

Q2. What is Brownian motion? What does it prove about matter?

Answer: Brownian motion is the continuous, random, zig-zag movement of small particles suspended in a liquid or gas, caused by constant collisions from surrounding molecules.

Discovered by: Robert Brown (1827) observed pollen grains in water through a microscope.

What it proves:

• Particles of matter (liquid or gas) are in constant motion.
• These particles are extremely small and invisible.
• Temperature increase intensifies Brownian motion, confirming that kinetic energy increases with temperature.

Q3. Explain the latent heat of fusion and latent heat of vaporisation with their values for water.

Answer:

Latent Heat of Fusion is the amount of heat required to convert 1 kg of a solid into liquid at its melting point and atmospheric pressure, without any change in temperature.

• For ice: 3.34 × 10⁵ J/kg
• This explains why water at 0°C has more energy than ice at 0°C water has already absorbed this latent heat.

Latent Heat of Vaporisation is the amount of heat required to convert 1 kg of a liquid into vapour at its boiling point and atmospheric pressure, without any change in temperature.

• For water: 22.5 × 10⁵ J/kg
• This explains why steam at 100°C causes more severe burns than boiling water at 100°C steam carries this additional latent heat energy.

The word "latent" comes from Latin, meaning hidden, because this heat is stored invisibly in the substance without raising its temperature.

Q4. Why does evaporation cause cooling? Explain with two everyday examples.

Answer: When a liquid evaporates, the particles at its surface that have the highest kinetic energy escape into the vapour phase. These fast-moving particles take their energy with them, leaving behind slower (cooler) particles. The remaining liquid therefore drops in temperature causing a cooling effect.

Examples:

1. Sweating: When the body overheats, sweat glands release perspiration. As sweat evaporates from the skin surface, it absorbs latent heat of vaporisation from the body, cooling it down. This is the body's natural thermoregulation mechanism.
2. Spirit on the palm: If a few drops of spirit (ether or alcohol) are placed on the palm, the hand instantly feels cold. Spirit has a very low boiling point and high vapour pressure, so it evaporates rapidly, absorbing heat from the skin.

Q5. Distinguish between a true solution, a colloidal solution, and a suspension with one example each.

Answer:

Q6. What is the Tyndall Effect? Give two natural examples.

Answer: The Tyndall Effect is the scattering of a beam of light by colloidal particles, making the path of the beam visible as a glowing cone called the Tyndall cone. It occurs because colloidal particles (size 1–100 nm) are large enough to scatter light.

Natural examples:

1. When sunlight passes through a gap in curtains into a dusty room, the beam of light becomes visible due to scattering by dust particles (a colloid in air).
2. Headlights of vehicles appear as visible beams in foggy conditions fog is a liquid-in-gas colloid (aerosol) that scatters the headlight beam.

True solutions do not show the Tyndall Effect because their particles (< 1 nm) are too small to scatter visible light.

Q7. What is chromatography? List any four of its applications.

Answer: Chromatography is a technique for separating the dissolved components of a mixture by exploiting their different rates of movement over an absorbent medium (such as filter paper) when carried by a solvent. Components separate because they have different solubilities in the solvent and different affinities for the adsorbent.

The technique was named using the Greek word Kroma (colour) because it was first applied to separate plant pigments.

Applications:

1. Separating different dyes or pigments from a coloured mixture (e.g., different colours in ink).
2. Separating and identifying amino acids in biochemical research.
3. Detecting and isolating drugs from blood or urine samples.
4. Separating sugar from urine samples in medical diagnostics.

Q8. Why are alloys considered mixtures even though they are homogeneous?

Answer: Alloys are classified as mixtures (not compounds) for two key reasons:

1. Variable composition: The ratio of constituent metals in an alloy can vary within a range. For example, brass can contain anywhere from 58% to 65% copper with the rest being zinc yet it is still called brass. A compound, by definition, always has a fixed ratio of elements.
2. Retained individual properties: The component metals in an alloy retain their own chemical properties. For instance, in brass, the zinc component reacts with dilute sulphuric acid to form zinc sulphate and hydrogen gas, while the copper component does not react. This demonstrates that each metal retains its chemical identity.

Q9. Describe the principle, method, and applications of centrifugation.

Answer:

Principle: When a suspension or colloidal solution is whirled at very high speed, the centrifugal force pushes denser (heavier) particles outward to the bottom of the tube, while lighter particles stay near the top. This separation would otherwise require much longer settling times under gravity alone.

Method:

1. The mixture (e.g., full-cream milk) is placed in a test tube balanced in the centrifuge.
2. The centrifuge spins at high speed, swinging the tube to a horizontal position.
3. The heavy protein and carbohydrate particles are flung to the bottom, while lighter fat particles concentrate near the top.
4. The layers are then separated.

Applications:

• Separating cream from milk in dairy industries.
• Separating components of blood in blood banks (red cells, white cells, platelets, plasma).
• Analysing urine samples in pathological laboratories.
• Removing water from wet clothes in spin dryers.

Q10. Distinguish between physical and chemical changes with two examples of each.

Answer:

A physical change is reversible and involves no change in molecular composition. No new substance is formed.

Examples:

• Ice melting into water (reversible by cooling).
• Dissolving sugar in water (sugar can be recovered by evaporation).

A chemical change is irreversible and involves the formation of new substances with different properties. Energy is exchanged in the process.

Examples:

• Burning of magnesium in air → magnesium oxide (white ash); cannot revert to magnesium metal by simple cooling.
• Rusting of iron → iron oxide (rust); the iron cannot be restored without a chemical process.

Key test for a chemical change: Does the product have different properties from the original? Is it irreversible under ordinary conditions? If yes to both, it is a chemical change.