Is Matter Around Us Pure Complete Class 9 Science Revision Notes 2026-27

Understanding the chapter “Is Matter Around Us Pure” is very important for every student of Central Board of Secondary Education because it explains the basic concepts of pure substances, mixtures, solutions, suspension, and colloids in an easy way. These Is Matter Around Us Pure Complete Class 9 Science notes are prepared according to the latest syllabus of CBSE Board and help students learn concepts. In this chapter, students study different types of matter, physical and chemical changes, separation techniques, and properties of substances used in daily life.

These class 9 is matter around us pure notes are useful for quick revision, homework, exam preparation, and concept clarity. Students looking for is matter around us pure complete class 9 science notes pdf or is matter around us pure notes pdf download can use these notes to revise all important topics in short and easy points. The notes also include important definitions, examples, NCERT based explanations, and key questions that are frequently asked in school exams.

The chapter is matter around us pure complete class 9 science notes chapter 2 is one of the scoring chapters in CBSE class 9 Science. These Class 9 Science notes are written in a student-friendly style so every learner can understand the concepts properly without confusion. Some topics may look little difficult at first, but with regular practice and reading, students can easily improve there understanding and perform better in exams.

Is Everything Around Us Pure?

Take a moment to look around you. The glass of water on your desk, the air you breathe, the soil in your garden do you think any of these are truly pure? In everyday language, "pure" often just means clean or unadulterated. But in science, the word carries a much more precise and demanding meaning.

This chapter "Is Matter Around Us Pure?" is one of the most concept-rich chapters in Class 9 Chemistry. It builds directly on your knowledge of matter's physical nature and asks a deeper question: what is the actual composition of the substances around us? Are they made of one type of particle, or many? Can they be separated by physical or chemical means?

By the end of this guide, you will be able to:

• Clearly distinguish between pure substances (elements and compounds) and mixtures.
• Understand how solutions, colloids, and suspensions differ and why this matters.
• Explain the Tyndall Effect and relate it to everyday phenomena.
• Apply the correct separation technique for any given mixture.
• Distinguish confidently between physical and chemical changes.

Pure Substances

A pure substance is a homogeneous material that contains particles of only one kind and has a definite, fixed set of physical and chemical properties.

Examples: Iron (Fe), silver (Ag), oxygen (O₂), sulphur (S), carbon dioxide (CO₂).

Each of these is a pure substance because every particle within it is identical same type, same properties, no impurities mixed in.

Characteristics of a Pure Substance

1. Homogeneous in Nature A pure substance has a completely uniform composition throughout no part of it differs from another. There are no visible boundaries, layers, or variations.

2. Definite and Fixed Set of Properties Every pure substance has its own unique, unchanging set of physical properties (like melting point, boiling point, density, solubility) and chemical properties. These properties are always reproducible and reliable. For example, pure water always boils at exactly 100°C at atmospheric pressure.

3. Composition Cannot Be Changed by Physical Means No matter how you grind, heat (without chemical reaction), dissolve, or filter a pure substance, its fundamental chemical composition remains unchanged. Altering its composition requires a chemical process.

Classification of Matter A Bird's Eye View

MATTER

• Pure Substances (only one type of particle)
  • Elements (cannot be broken down chemically)
  • Compounds (two or more elements, fixed ratio, chemical combination)
• Mixtures (more than one substance present)
  • Homogeneous Mixtures → Solutions
  • Heterogeneous Mixtures → Suspensions & Colloids

Elements

An element is a pure substance that cannot be broken down into two or more simpler substances by any chemical means. All atoms in an element have the same atomic number.

Examples: Hydrogen (H), oxygen (O), nitrogen (N), copper (Cu), zinc (Zn), tin (Sn), lead (Pb), mercury (Hg).

Classification of Elements

A. By Physical State at Room Temperature

B. Metals vs. Non-Metals

There are 93 metals and 22 non-metals among the known elements.

Metals:

• Mostly solids at room temperature.
• Good conductors of heat and electricity.
• Malleable and ductile.
• Only mercury is a liquid metal all other metals are solids.

Non-Metals (22 total):

• 10 are solids: Boron (B), Carbon (C), Silicon (Si), Phosphorus (P), Sulphur (S), Selenium (Se), Arsenic (As), Tellurium (Te), Iodine (I), Astatine (At).
• 1 is liquid: Bromine (Br) the only non-metal liquid at room temperature.
• 5 are chemically active gases: Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine.
• 6 are chemically inactive (noble/rare) gases: Helium, Neon, Argon, Krypton, Xenon, Radon.

Hydrogen is the lightest element in the periodic table.

C. Metalloids (Semi-Metals)

Some elements display intermediate properties they look like metals but are brittle like non-metals. They are neither fully conducting (like metals) nor fully insulating (like non-metals), making them semiconductors. These are called metalloids or semi-metals.

Key examples: Boron (B), Silicon (Si), Germanium (Ge)

Silicon, in particular, is the backbone of modern electronics and computer chips a real-world testament to why understanding metalloids matters.

D. Normal vs. Radioactive Elements

• Normal elements (atomic numbers 1–82): Do not emit harmful radiation under ordinary conditions.
• Radioactive elements (atomic numbers 83–112, and 114, 116, 118): Spontaneously emit harmful radiation (alpha, beta, or gamma rays). Examples include uranium (92), radium (88), and plutonium (94).

Compounds

A compound is a pure substance composed of two or more elements chemically combined in a definite, fixed ratio by weight. Its components can only be separated by chemical means not physical ones.

The elements present in a compound are called its constituents or components.

Example: Water (H₂O) hydrogen and oxygen combined in a 1:8 ratio by weight. Water can only be broken back into hydrogen and oxygen by electrolysis (passing electric current through it). No amount of filtering, boiling, or distillation will separate water into hydrogen and oxygen.

Classification of Compounds

Compounds are further classified into three broad categories:

Acids compounds that produce hydrogen ions in solution and typically taste sour. Examples: Sulphuric acid (H₂SO₄), hydrochloric acid (HCl), nitric acid (HNO₃), formic acid (HCOOH).

Bases compounds that produce hydroxide ions in solution and typically feel slippery. Examples: Sodium hydroxide (NaOH), potassium hydroxide (KOH), zinc hydroxide, calcium hydroxide.

Salts ionic compounds formed by the chemical reaction between an acid and a base (neutralisation reaction). Examples: Ammonium chloride (NH₄Cl), zinc sulphate (ZnSO₄), lead nitrate, calcium carbonate (CaCO₃).

Mixtures

Definition

A mixture is formed when two or more substances (elements, compounds, or both) are combined in any proportion without undergoing any chemical change, such that each substance retains its individual physical and chemical properties.

In a mixture, no new substance is created. No fixed ratio is required. No energy is released or absorbed during formation.

Types of Mixtures

1. Homogeneous Mixtures

A mixture in which all components are uniformly distributed throughout no visible boundaries, and every portion of the mixture has the same composition.

Examples: Sugar in water, copper sulphate solution, common salt solution, alloys (brass, bronze, bell metal).

Note: All solutions are homogeneous mixtures, but not all homogeneous mixtures are liquids alloys are solid homogeneous mixtures.

2. Heterogeneous Mixtures

A mixture in which components are not uniformly distributed visible or detectable boundaries exist between them.

Examples: Sand and salt, iron powder and sulphur powder, soil, muddy water.

Differences: Mixtures vs. Compounds

Why Air Is a Mixture Not a Compound

Many students wonder: is air a compound (like water) or a mixture? The evidence firmly places air in the mixture category:

1. Variable composition: The percentage of oxygen decreases at higher altitudes. Industrial areas have higher concentrations of CO₂ and pollutants. If air were a compound, its composition would be fixed everywhere.

2. Separable by physical methods: The components of air nitrogen, oxygen, argon, CO₂ can be separated by liquefaction followed by fractional distillation. A compound cannot be separated physically.

3. No chemical reaction on mixing: When the gases that make up air (O₂, N₂, H₂O vapour, CO₂) are mixed together, no heat or light is produced. Compound formation always involves energy exchange.

4. Constituents retain individual properties: Oxygen in air still supports combustion; carbon dioxide still turns lime water milky. In a compound, the properties of constituents disappear entirely.

Why Water Is a Compound Not a Mixture

Conversely, water is a compound, not a mixture of hydrogen and oxygen:

1. Fixed composition: Pure water always contains hydrogen and oxygen in a 1:8 ratio by weight everywhere on Earth, always.

2. Cannot be separated physically: Water cannot be separated into hydrogen and oxygen by boiling, filtration, or any other physical process. Electrolysis (a chemical/electrochemical process) is required.

3. Energy is released during formation: When hydrogen and oxygen combine to form water, a large amount of heat and light energy is released a clear chemical reaction signature.

4. Properties are completely different: Hydrogen is a combustible gas (it burns). Oxygen supports combustion. Yet water their compound actively extinguishes fire. The product's properties are nothing like its components.

Why Alloys Are Considered Mixtures

Alloys are homogeneous mixtures of metals, yet they cannot always be easily separated by ordinary physical methods. Despite this, they are classified as mixtures because:

1. Variable composition: The ratio of constituent metals can vary within a range. Brass (copper + zinc) can contain 58%–65% copper yet it is still called brass. A compound must always have a fixed ratio.

2. Constituents retain chemical properties: In brass treated with dilute sulphuric acid, the zinc reacts to form zinc sulphate and hydrogen gas, but the copper remains unreacted proving each metal retains its own chemical identity within the alloy.

Common alloys and their compositions:

Solutions

A solution is a homogeneous mixture of two or more substances. While we commonly think of solutions as liquids, a solution can also be solid (alloys) or gaseous (air).

Components of a Solution

Every solution has two fundamental components:

Solvent the component present in larger proportion; the dissolving medium.

Solvent = LARGER component

Solute the component present in smaller proportion; the substance that gets dissolved.

Solute = SMALLER component

Examples:

• In copper sulphate solution: water is the solvent, copper sulphate is the solute.
• In paints: turpentine oil is the solvent, the pigments are the solutes.
• In tincture of iodine: ethyl alcohol is the solvent, iodine is the solute.
• In carbonated drinks: water is the solvent, carbon dioxide is the solute.

Types of Solutions by State

Types of Solutions by Saturation

• Saturated Solution: A solution that has dissolved the maximum possible amount of solute at a given temperature. Adding more solute will not dissolve it simply settles at the bottom.
• Unsaturated Solution: A solution containing less dissolved solute than its maximum capacity. More solute can still be dissolved.
• Supersaturated Solution: A solution that contains more dissolved solute than a normal saturated solution would allow at a given temperature. This is achieved by dissolving solute in a hot solvent and then cooling it very slowly and carefully. Supersaturated solutions are unstable any disturbance causes excess solute to crystallise out.

Practical example: Honey is a supersaturated solution of sugars in water. Over time, sugar crystals form at the bottom this is why honey sometimes "granulates."

Concentration of a Solution

Concentration is the amount of solute dissolved per unit quantity of solution. It is most commonly expressed as a percentage.

Mass Percentage (for solid solute in liquid solvent):

Concentration (%) = [Mass of Solute (g) / Mass of Solution (g)] × 100 where: Mass of Solution = Mass of Solute + Mass of Solvent

Volume Percentage (for liquid solute in liquid solvent):

Concentration (%) = [Volume of Solute (mL) / Volume of Solution (mL)] × 100 where: Volume of Solution = Volume of Solute + Volume of Solvent

Critical note: Concentration is always calculated per 100 g (or 100 mL) of solution NOT per 100 g of solvent alone. This is a very common source of errors in exams.

Concentration is a dimensionless percentage it has no units.

True Solutions

A true solution is a homogeneous mixture in which solute particles are broken down to such a fine (molecular) level that they cannot be seen even under the most powerful optical microscope.

Examples: Common salt solution, sugar solution, copper sulphate solution.

Characteristics of a True Solution

1. Transparent and Clear: Light passes through without scattering. You can read text behind a true solution.

2. Ultra-fine Particle Size: The diameter of dissolved particles is 1 nm (10⁻⁹ m) or less essentially molecular-sized.

3. Passes Through Filter Paper: Because the particle size is far smaller than the pores of filter paper, true solutions pass completely through without any residue.

4. Homogeneous: Composition is identical throughout.

5. Particles Do Not Settle: Provided temperature remains constant, solute particles remain permanently in solution they never settle to the bottom.

6. Solute Is Recoverable: The dissolved solute can be recovered from a true solution by evaporation or crystallisation.

Suspensions

A suspension is a heterogeneous mixture in which insoluble particles of solute are spread (dispersed) throughout a solvent. The particle size in a suspension is more than 10⁻⁵ cm in diameter.

Examples:

• Muddy water (sand and clay particles in water).
• Slaked lime suspension used for whitewashing.
• Paints (dye particles dispersed in turpentine oil).

Characteristics of Suspensions

1. Large Particle Size: Greater than 10⁻⁵ cm in diameter visible to the naked eye.

2. Settles on Standing: When left undisturbed, suspended particles settle to the bottom under gravity. This settling process is called sedimentation.

3. Filterable: Because the particle size is larger than the pores of filter paper, suspension particles are retained by the filter paper and separated from the solvent.

4. Heterogeneous: Composition is not uniform throughout.

5. Scatters Light: The relatively large particles scatter light significantly, making the suspension appear cloudy or opaque.

Colloidal Solutions

A colloidal solution is a heterogeneous mixture in which particle size is intermediate between 10⁻⁷ cm and 10⁻⁵ cm such that the particles neither dissolve completely nor settle down.

In a colloidal solution:

• The dispersed particles are called the dispersed phase.
• The medium in which they are dispersed is called the continuous phase or dispersing medium.

Examples of Colloidal Solutions

Blood, milk, writing ink, jelly, starch solution, gum solution, toothpaste, soap solution, liquid detergents, mist, fog.

Characteristics of Colloidal Solutions

1. Intermediate Particle Size: 10⁻⁷ cm to 10⁻⁵ cm.

2. Visible Under Microscope: Colloidal particles are visible under a powerful microscope but not with the naked eye.

3. Do Not Settle: Unlike suspensions, colloidal particles remain permanently dispersed and do not settle over time.

4. Pass Through Filter Paper: Colloidal particles are small enough to pass through ordinary filter paper (though they are retained by semipermeable membranes).

5. Scatter Light (Tyndall Effect): When a beam of light passes through a colloidal solution, the path becomes visible due to light scattering by colloidal particles.

6. Translucent (Not Transparent): Colloidal solutions allow some light through but scatter enough to appear cloudy they are not fully transparent like true solutions.

7. Electrically Charged Particles: Colloidal particles carry electric charges (all of the same sign), which keeps them mutually repelled and prevents them from clumping together and settling.

8. Heterogeneous in Nature: Despite appearing uniform, colloidal solutions are technically heterogeneous at the microscopic level.

Colloidal solutions can be separated by the process of centrifugation.

The Tyndall Effect

The Tyndall Effect is the phenomenon in which a beam of light passing through a colloidal solution is scattered by colloidal particles, making the path of the light beam visible as a glowing cone called the Tyndall cone.

The Experiment

A wooden box is fitted with a convex lens on one side and a microscope on the other. A beaker of soap solution (a colloid) is placed inside. When a powerful beam of light is focused through the lens into the soap solution and observed through the microscope, individual colloidal particles appear surrounded by a cone of bluish scattered light the Tyndall cone.

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

Everyday Examples of the Tyndall Effect

• A beam of sunlight entering a dusty room through a gap in curtains becomes visible as a glowing shaft dust particles in air scatter the light.
• Vehicle headlights appear as visible beams in foggy weather fog (water droplets in air) is a colloid that scatters light.
• Light from a projector becomes visible as a beam in a smoky or dusty hall.
• The blue colour of the sky itself is partly explained by light scattering by atmospheric particles (Rayleigh scattering a related phenomenon).

Classification of Colloids

Colloids are classified based on the physical state of the dispersed phase and the dispersing medium:

Comparison: True Solution vs. Colloidal Solution vs. Suspension

Separation of Mixtures

Different mixtures are separated by exploiting differences in the physical properties of their components including density, boiling point, volatility, solubility, ability to sublime, and ability to diffuse.

1. Evaporation

Used for: Separating a non-volatile solid from its liquid solvent.

Principle: The liquid (solvent) is heated and converts to vapour, leaving the dissolved solid behind.

Suitable mixtures:

Method: The solution is poured into a china dish and heated gently on a sand bath. Water evaporates slowly, leaving behind the dry solid residue.

Limitation: Evaporation is not suitable when the solid would decompose on heating, or when both the solid and liquid are needed as separate recoverable products.

2. Centrifugation

Used for: Separating finely suspended or colloidal particles from a liquid.

Principle: When a mixture is spun at very high speed, denser (heavier) particles are thrown to the bottom of the tube (outward), while lighter particles remain near the top much faster than gravity-driven sedimentation.

Method:

1. Place the mixture in a centrifuge tube.
2. Balance it with an identical tube on the opposite side.
3. Spin at high speed. Heavier particles move to the bottom; lighter particles collect at the top.
4. Carefully decant or pipette the separated layers.

Applications:

3. Separating Funnel

Used for: Separating two immiscible liquids (liquids that do not dissolve in each other) based on their density difference.

Apparatus: A pear-shaped glass funnel with a tap at the bottom.

Principle: The denser liquid sinks to the bottom layer; the lighter liquid floats on top. The tap allows each layer to be drained separately.

Common immiscible liquid pairs:

Method:

1. Close the tap; clamp the funnel vertically.
2. Pour the immiscible mixture in; let it stand for 30+ minutes.
3. Two distinct layers form heavier at the bottom.
4. Open the tap gently; drain the bottom layer into a beaker.
5. Once the interface passes the tap, close it and collect the top layer in a separate beaker.

Industrial Application: During the extraction of iron in a blast furnace, the denser molten iron sinks to the bottom while the lighter slag floats on top they are tapped out from different outlets using this density-difference principle.

4. Sublimation

Used for: Separating a mixture where one solid converts directly to vapour (sublimes) on heating, without passing through the liquid phase. The vapour is then cooled to re-solidify separately.

Sublimable substances: Dry ice (solid CO₂), naphthalene, anthracene, iodine, ammonium chloride (NH₄Cl), camphor.

Mixtures separable by sublimation:

Method:

1. Place the mixture in a china dish; heat over a low Bunsen flame.
2. Place an inverted glass funnel over the dish; plug the stem with cotton wool to slow the vapour.
3. The sublimable component converts to dense vapour, which rises and condenses as a fine powder on the cooler funnel walls.
4. When no more vapour is produced, lift the funnel and scrape off the pure sublimate.
5. The non-sublimable component remains as residue in the dish.

5. Chromatography

Used for: Separating dissolved components of a mixture based on their different affinities for an adsorbent material and their different rates of movement through that material when carried by a solvent.

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

Principle: Each component of a mixture adsorbs to the adsorbent (e.g., filter paper) to a different degree. Components with higher solubility in the solvent or lower affinity for the adsorbent travel farther up the paper; components with lower solubility or higher affinity for the adsorbent travel less.

Paper chromatography method (for separating ink components):

1. Take a 22 cm × 5 cm strip of filter paper; mark an ink spot 2 cm from the bottom.
2. Suspend the strip in a tall cylinder with a small amount of water at the bottom just enough to touch the lower edge, but not the ink spot.
3. Cover the cylinder to prevent evaporation. Leave undisturbed for about an hour.
4. Water travels up the paper by capillary action, dissolving and carrying the ink components at different rates.
5. Remove and dry the paper a chromatogram showing separated colour bands is visible.

Applications of Chromatography:

• Separating colours/pigments from dyes and inks.
• Identifying and separating amino acids in biochemistry.
• Separating sugar from urine in medical testing.
• Detecting and isolating drugs from blood samples.
• Quality control testing in the pharmaceutical and food industries.

Advantages:

• Can be performed with very small sample quantities.
• No sample is wasted during separation.

6. Distillation

Used for: Separating a dissolved solid from its solvent (recovering both), or separating two miscible liquids with sufficiently different boiling points.

Principle: The liquid is heated to form vapour; the vapour is then cooled and condensed back into liquid in a separate vessel. The condensed liquid is called the distillate.

Liquid → [Heat] → Vapour → [Cool via Liebig Condenser] → Distillate

Liebig Condenser: A long glass tube surrounded by a wider water jacket. Cold tap water circulating through the jacket absorbs heat from the vapour passing through the inner tube, causing condensation.

When to use simple distillation:

• Separating a non-volatile solid from a liquid (e.g., recovering both salt and pure water from saltwater).
• Separating two miscible liquids with a boiling point difference of more than 25°C.

7. Fractional Distillation

Used for: Separating two or more miscible liquids whose boiling points are close to each other (difference less than 10°C).

Why simple distillation fails: When boiling points are very close, both liquids tend to vapourise simultaneously in similar proportions making clean separation impossible with a simple still.

Solution: A fractionating column is inserted between the flask and the condenser. The column provides resistance to vapour flow and gives the effect of repeated distillation higher-boiling-point vapour condenses within the column and drips back, while lower-boiling-point vapour passes on to the condenser.

Common fractional distillation separations:

Major industrial application Separation of gases from air:

Air is first purified, liquefied, and then fractionally distilled to obtain its component gases:

Step 1 Purification:

• Air is washed through water to remove dust particles.
• Washed air is passed through dilute caustic soda solution to remove CO₂, SO₂, and H₂S.
• Dried by passing through pipes cooled below −20°C moisture freezes out.

Step 2 Liquefaction:

• Purified air is compressed to 200× atmospheric pressure (which raises its temperature).
• Compressed hot air is cooled through a water-cooled tank.
• Cooled compressed air is rapidly expanded through a fine jet sudden pressure drop causes temperature to plummet (Joule-Thomson effect).
• The cycle repeats until the air liquefies at approximately −200°C.

Step 3 Fractional Distillation:

• Liquid air is gradually warmed to −195°C nitrogen (boiling point −195°C) boils off first.
• Nitrogen gas is collected, compressed, and stored in steel cylinders.
• Remaining liquid oxygen (boiling point −183°C) is then collected, converted to gas, and stored.

City Water Supply Applied Separation Science

Municipal water treatment applies multiple separation techniques in sequence:

Stage 1 Sedimentation: River water is held in large tanks. Alum is added as a coagulant it causes fine suspended particles to clump together into larger aggregates that settle faster (a process called loading). Heavy impurities settle at the bottom.

Stage 2 Filtration: The partially clarified water is passed through layered beds of sand, charcoal, and gravel. Fine suspended particles are trapped in these layers. The water then passes through a sand filter for further purification.

Stage 3 Sterilisation (Chlorination): Filtered water still contains harmful microorganisms (bacteria causing typhoid, cholera, dysentery, jaundice). Bleaching powder or chlorine gas is added chlorine kills the pathogenic bacteria, making the water safe to drink.

The purified water is then pumped to elevated overhead tanks for gravity-fed distribution throughout the city.

Physical and Chemical Changes

Physical Changes

A physical change is one that alters a specific physical property of matter such as state (solid, liquid, gas), texture, colour, shape, magnetism, or electrical properties without any change in its molecular composition (chemical identity).

Main pointers:

1. No new substance is formed. The molecules remain chemically identical. Example: When ice melts into water, only the state changes every molecule is still H₂O.

2. Temporary and reversible. Removing the cause reverses the change. Example: Water formed from ice can be frozen back into ice by cooling it below 0°C.

3. No net energy change. Energy absorbed to cause the change equals energy released upon reversal. Example: 1 g of water at 100°C requires 2260 J to become steam; 1 g of steam at 100°C releases exactly 2260 J on condensing back to water.

4. Mass is conserved. No matter is added or removed only energy changes hands.

Examples of physical changes:

More common examples: Formation of dew, evaporation of water, crystallisation of sugar, ringing of an electric bell, breaking glass, making ice cream, melting of wax, bending a glass tube by heating, sublimation of camphor.

Chemical Changes

Definition: A chemical change is one that alters the molecular composition of a material, forming one or more entirely new substances with different properties.

Key characteristics:

1. New substances are always formed. The product has a different molecular composition from the starting material. Example: Heating sugar → carbon (black solid) + steam. Carbon and water are chemically different from sugar.

2. Permanent and irreversible. The change cannot be reversed simply by altering temperature or pressure. Example: Charcoal produced from burning sugar cannot be converted back into sugar.

3. Apparent change in mass. The mass of the visible product differs from the starting material (though total mass of all products including gases equals the starting mass law of conservation of mass).

4. Energy is always exchanged. Heat, light, or electrical energy is absorbed or released as chemical bonds break and form.

Key examples of chemical changes:

More common examples: Burning of wood or charcoal, digestion of food, curdling of milk, formation of biogas (gobar gas), burning of petrol or diesel, rusting of iron, ripening of fruit, clotting of blood, baking of cake, photosynthesis, formation of wine, butter turning rancid, fading of dyed cloth, decomposition of water by electrolysis, formation of water from hydrogen and oxygen.

Physical vs. Chemical Changes Complete Comparison

Class 9 Science Is Matter Around Us Pure Questions and Answers

Q1. What is a pure substance? Give three examples and explain why each qualifies.

Answer: A pure substance is a homogeneous material that contains particles of only one kind and has a definite, fixed set of properties that cannot be altered by any physical process.

Examples and reasoning:

1. Iron (Fe): Every particle is an iron atom. Its melting point (1538°C), density (7.87 g/cm³), and magnetic properties are always the same. No physical process can alter its composition.
2. Oxygen (O₂): All particles are oxygen molecules. Its properties colourless, odourless gas, supports combustion, boiling point −183°C are fixed and reproducible worldwide.
3. Water (H₂O): Always composed of hydrogen and oxygen in a 1:8 mass ratio. Its boiling point (100°C at 1 atm) and density (1 g/cm³ at 4°C) are always the same for pure samples, regardless of source.

Q2. Distinguish between elements and compounds with two examples each.

Answer:

Expanded example: Copper cannot be broken into anything simpler by any chemical reaction it is an element. Water, however, can be decomposed into hydrogen gas (H₂) and oxygen gas (O₂) by electrolysis confirming it is a compound.

Q3. What are metalloids? Name three metalloids and explain their significance.

Answer: Metalloids (also called semi-metals) are elements that display properties intermediate between metals and non-metals. They appear shiny like metals but are brittle like non-metals. Their electrical conductivity is intermediate they are neither full conductors (like metals) nor full insulators (like non-metals). This makes them semiconductors.

The three key metalloids:

1. Boron (B): Used in semiconductors, glass (borosilicate glass), and neutron absorbers in nuclear reactors.
2. Silicon (Si): The most important metalloid the basis of all modern computer chips, solar cells, and electronic components. "Silicon Valley" is named after it.
3. Germanium (Ge): Used in early transistors and still used in fibre-optic cables and infrared optics.

Their semiconductor properties arise because their electrical conductivity can be precisely controlled by adding tiny amounts of other elements (doping) essential for electronics manufacturing.

Q4. Explain with reasons why alloys are classified as mixtures even though they appear homogeneous and cannot always be separated by simple physical methods.

Answer: The classification of alloys as mixtures (rather than compounds) rests on two scientifically sound arguments:

Reason 1 Variable composition: In a compound, elements are always present in a fixed ratio. Alloys do not obey this rule. Brass, for example, is classified as brass whether it contains 58% or 65% copper the ratio can vary within a range. This variable composition is the hallmark of a mixture.

Reason 2 Constituent metals retain individual chemical properties: In a compound, the properties of individual constituents disappear entirely. In brass (copper + zinc), if treated with dilute sulphuric acid, the zinc component reacts to form zinc sulphate and hydrogen gas, while the copper component remains completely unreacted. Each metal behaves exactly as it would in isolation proving they have not chemically combined.

The fact that alloys are homogeneous and cannot easily be separated physically is irrelevant to the classification what matters is whether the constituents are chemically combined (compound) or not (mixture).

Q5. Differentiate between a true solution, a colloidal solution, and a suspension. Give one example of each.

Answer:

Q6. What is the Tyndall Effect? Explain with a laboratory experiment and two natural observations.

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 is observed only in colloidal solutions, not in true solutions, because colloidal particles (1–100 nm) are large enough to scatter visible light, whereas particles in true solutions (< 1 nm) are too small.

Laboratory Experiment: A wooden box fitted with a convex lens on one side and a microscope on the other is used. A beaker of soap solution (a colloid) is placed inside. A powerful light source is directed through the convex lens, forming a narrow parallel beam through the soap solution. Viewed through the microscope, individual colloidal particles appear surrounded by a cone of bluish scattered light the Tyndall cone.

Natural Observations:

1. When sunlight enters a dusty room through a gap in curtains, the beam becomes clearly visible as a shaft of light dust particles (colloid in air) scatter the sunlight.
2. Vehicle headlights appear as bright visible beams in fog fog (tiny water droplets dispersed in air, an aerosol colloid) scatters the light along the entire beam path.

Q7. Describe the process of fractional distillation of air to obtain nitrogen and oxygen.

Answer: Fractional distillation of air is carried out in three stages:

Stage 1 Purification: Air contains CO₂, H₂S, SO₂, and dust as impurities. It is first washed through water (removing dust), then passed through dilute caustic soda solution (removing acidic gases), and finally dried by passing through pipes kept below −20°C (freezing out moisture).

Stage 2 Liquefaction: Purified air is compressed to 200× atmospheric pressure, then cooled in a water-cooled tank. The cooled compressed air is forced through a fine jet sudden expansion causes a sharp temperature drop (Joule-Thomson effect). The cooled air rises and cools the next batch of incoming compressed air. After many repeated cycles, the temperature drops enough for the air to liquefy (at approximately −200°C).

Stage 3 Fractional Distillation: Liquid air (mainly nitrogen and oxygen) is gradually warmed. Since liquid nitrogen boils at −195°C and liquid oxygen boils at −183°C, nitrogen vapourises first (at the lower temperature) and is collected, compressed, and stored in steel cylinders. The remaining liquid oxygen is then also converted to gas and stored in cylinders.

Q8. Explain with an example how paper chromatography works. What determines how far each component moves?

Answer: Paper chromatography separates dissolved components of a mixture based on their differing affinities for the adsorbent (filter paper) and their differing solubilities in the solvent.

Example Separating ink components:

A spot of ink is placed 2 cm from the bottom of a filter paper strip. The strip is suspended in a cylinder with just enough water at the base to wet the lower edge (but not reach the ink spot). Water rises by capillary action, reaches the ink spot, dissolves its components, and carries them upward through the paper. Components with:

• Higher solubility in water and lower affinity for the paper → travel farther up.
• Lower solubility in water and higher affinity for the paper → travel less distance, stay closer to the original spot.

After an hour, distinct coloured bands appear at different heights this is the chromatogram. Each band represents a separated component of the ink.

Two factors governing distance travelled:

1. Relative solubility of each component in the solvent (mobile phase).
2. Relative affinity of each component for the adsorbent medium (stationary phase the filter paper).

Q9. Give four differences between physical and chemical changes with one example for each.

Answer:

Q10. A student dissolves 20 g of salt in 180 g of water. Calculate the concentration of the solution by mass percentage.

Answer:

Given:

• Mass of solute (salt) = 20 g
• Mass of solvent (water) = 180 g
• Mass of solution = 20 + 180 = 200 g

Formula:

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

Calculation:

Concentration = (20 / 200) × 100 = 10%

The concentration of the solution is 10% by mass.

This means that every 100 g of this salt solution contains 10 g of dissolved salt and 90 g of water.

Common mistake to avoid: Do NOT divide by the mass of solvent alone (180 g). The formula always uses the total mass of the solution (solute + solvent) in the denominator.