Detailed Notes: Chapter 2 – Matter and Its States

2.1 Matter: An Introduction

• Definition: Anything that has mass and occupies space (volume) is called matter.

• Characteristics of Particles of Matter:

  • Particles are continuously moving.

  • They have space between them (interparticle or intermolecular space).

  • They attract each other (intermolecular force of attraction).

2.2 States of Matter

Scientists now recognize five states of matter:

1. Solid State:

   • Definite shape and volume.

   • Negligible compressibility.

   • High rigidity; particles are tightly packed.

   • Particles vibrate about their fixed positions.

2. Liquid State:

   • Definite volume but no definite shape. They take the shape of their container.

   • Slight compressibility.

   • They can flow and are considered fluids.

   • Particles are less tightly packed than in solids and can move freely.

3. Gaseous State:

   • No definite shape or volume. They completely fill the container they are in.

   • Highly compressible.

   • Particles are far apart and move randomly at high speeds.

   • They are fluids.

4. Plasma:

   • The fourth state of matter, consisting of super energetic and super excited particles in the form of ionised gases.

   • Examples: The glow in fluorescent tubes (contains helium or other gases) and neon sign bulbs (contains neon gas). The Sun and stars glow due to plasma.

   • It occurs at very high temperatures.

5. Bose-Einstein Condensate (BEC):

   • The fifth state of matter.

   • Formed by cooling a gas of extremely low density to super low temperatures.

2.3 Interconversion of States of Matter

States of matter are interconvertible by changing temperature and pressure.

1. Effect of Temperature:

• Melting/Fusion: The process of a solid changing into a liquid at its melting point.

• Vaporization/Boiling: The process of a liquid changing into a vapour at its boiling point.

• Sublimation: The process of a solid changing directly into a gas without passing through the liquid state (e.g., Naphthalene, Iodine, Camphor).

• Latent Heat of Fusion: The heat energy required to change 1 kg of a solid into a liquid at its melting point.

• Latent Heat of Vaporization: The heat energy required to change 1 kg of a liquid into vapour at its boiling point.

• Key Points:

  • Melting point of ice = 273 K (0°C).

  • Boiling point of water = 373 K (100°C).

  • The presence of an impurity lowers the melting point and raises the boiling point.

  • The rate of evaporation increases with an increase in surface area, temperature, and wind speed, and decreases with an increase in humidity.

2. Effect of Pressure:

• Applying pressure and reducing temperature can liquefy gases.

• Increasing pressure raises the boiling point of a liquid.

• Dry Ice: Solid carbon dioxide (CO₂) sublimes directly to gas at 1 atm pressure without becoming liquid.

2.4 Classification of Matter

Based on chemical composition, matter is classified as follows:

                    Matter
                      |
        -------------------------------
        |                             |
    Pure Substance                 Mixture
    (Fixed composition)        (Variable composition)
        |                             |
    ------------               ----------------
    |          |               |              |
Element    Compound      Homogeneous    Heterogeneous
(1 type of atom) (2+ elements (Uniform)     (Non-uniform)
                 in fixed ratio)             e.g., Sand
  e.g., Fe, O₂      e.g., H₂O, CO₂          in water

A. Pure Substances

• Have a fixed chemical composition and characteristic properties.

• Elements: Cannot be broken down into simpler substances by chemical means. (e.g., Gold, Oxygen).

• Compounds: Formed when two or more elements combine chemically in a fixed proportion. Their properties are different from their constituent elements (e.g., Water H₂O).

B. Mixtures

• Contain two or more elements or compounds mixed in any proportion.

• Constituents can be separated by physical methods.

• Homogeneous Mixture: Has a uniform composition throughout.

  • Examples: Salt solution, Sugar solution, Air, Alloys.

• Heterogeneous Mixture: Does not have a uniform composition.

  • Examples: Mixture of salt and sand, Soil.

Types of Mixtures in Detail:

1. True Solutions:

   • Homogeneous mixtures.

   • Particle size < 1 nm.

   • Stable; particles do not settle down.

   • Do not scatter light (No Tyndall effect).

   • Example: Tincture of Iodine (Iodine in alcohol).

2. Suspensions:

   • Heterogeneous mixtures.

   • Particle size > 100 nm. Visible to naked eye.

   • Unstable; particles settle down on standing.

   • Show Tyndall effect.

   • Can be separated by filtration.

   • Example: Mixture of sand in water.

3. Colloids (Colloidal Solutions):

   • Heterogeneous mixtures.

   • Particle size between 1 nm and 100 nm.

   • Stable (particles show Brownian movement).

   • Show Tyndall effect.

   • Cannot be separated by filtration but can be by centrifugation.

   • Examples: Milk (emulsion), Fog (aerosol), Cheese (gel), Smoke (aerosol).

2.5 Separation Techniques

• For Heterogeneous Mixtures:

  • Handpicking: Removing stones from pulses.

  • Threshing: Separating grains from husks.

  • Winnowing: Separating lighter husk from heavier grains using wind.

  • Sieving: Separating particles of different sizes using a sieve.

  • Sedimentation & Decantation: Allowing insoluble solids to settle and then pouring out the liquid.

  • Filtration: Using a filter paper to separate an insoluble solid from a liquid.

• For Homogeneous Mixtures:

  • Evaporation: To separate a volatile component (solvent) from a non-volatile solute (e.g., salt from seawater).

  • Centrifugation: Using centrifugal force to separate components based on density (e.g., in dairies, diagnostic labs).

  • Sublimation: To separate a sublimate from a non-sublimate (e.g., Ammonium chloride from sand).

  • Crystallization: The process of obtaining pure solid in the form of crystals (e.g., purification of salt).

  • Distillation: To separate two miscible liquids based on differences in their boiling points.

  • Chromatography: Used to separate and purify components of a mixture (especially coloured ones) based on their distribution between a stationary and a mobile phase.

2.6 Metals and Non-Metals

A. Metals

• Tend to form positive ions (cations) by losing electrons.

• Physical Properties: Solid (except Mercury), lustrous, malleable, ductile, good conductors of heat and electricity.

• Chemical Properties:

  • With Oxygen: Form basic oxides (mostly). Some like Al₂O₃ and ZnO are amphoteric.

  • With Water:

    • Vigorous (Na, K): React with cold water.

    • Less Vigorous (Ca): Reacts with cold water.

    • With Hot Water/Steam (Mg, Al, Zn, Fe):

    • No Reaction (Pb, Cu, Ag, Au):

  • With Acids: Reactive metals (above H in reactivity series) displace H₂ gas.

  • With Bases: Some metals (like Al, Zn) react with bases to produce H₂ gas.

B. Non-Metals

• Tend to form negative ions (anions) by gaining electrons.

• Physical Properties: Can be solid, liquid (Bromine), or gas; generally not lustrous (except Iodine, Diamond), non-malleable, non-ductile, poor conductors (except Graphite).

• Chemical Properties:

  • Form acidic or neutral oxides (e.g., CO₂ is acidic, CO is neutral).

  • Generally, do not react with water.

  • Phosphorus is stored in water to prevent contact with air.

C. Reactivity Series
K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au
(Most Reactive to Least Reactive)

• A more reactive metal can displace a less reactive metal from its salt solution.

D. Metalloids

• Possess properties of both metals and non-metals.

• Examples: Germanium (Ge), Silicon (Si), Arsenic (As).

2.7 Chemical Changes & Reactions

• Physical Change: Change in physical properties. No new substance is formed. (e.g., Melting of wax, Boiling of water).

• Chemical Change: One or more new substances are formed. (e.g., Burning of candle, Rusting of iron).

Types of Chemical Reactions:

1. Combination Reaction: A + B → C (e.g., CaO + H₂O → Ca(OH)₂)

2. Decomposition Reaction: AB → A + B (e.g., 2HgO → 2Hg + O₂)

3. Displacement Reaction: A + BC → AC + B (e.g., Zn + CuSO₄ → ZnSO₄ + Cu)

4. Double Displacement Reaction: AB + CD → AD + CB (e.g., AgNO₃ + NaCl → AgCl + NaNO₃)

5. Oxidation and Reduction (Redox):

   • Oxidation: Addition of Oxygen / Removal of Hydrogen.

   • Reduction: Removal of Oxygen / Addition of Hydrogen.

   • Example: CuO + H₂ → Cu + H₂O (Here, H₂ is oxidised, CuO is reduced).

2.8 Corrosion & Prevention

• Corrosion: The slow process of deterioration of a metal due to reaction with air and moisture.

  • Examples: Rusting of iron (forms Fe₂O₃.xH₂O), tarnishing of silver (forms Ag₂S), green coating on copper (basic copper carbonate, CuCO₃.Cu(OH)₂).

• Prevention: Painting, greasing, galvanization (coating with Zn), electroplating, alloying.

2.9 Alloys

• A homogeneous mixture of two or more metals, or a metal and a non-metal.

• Purpose: To improve hardness, prevent corrosion, and lower melting points.

• Common Alloys:

  • Brass: Cu + Zn

  • Bronze: Cu + Sn

  • Steel: Fe + C

  • Stainless Steel: Fe + Cr + Ni + C

  • Solder: Pb + Sn (low melting point)

  • Amalgam: Hg + other metal(s) (e.g., for dental fillings)

2.10 Atomic Structure

• Atom: The smallest particle of an element that can take part in a chemical reaction.

• Subatomic Particles:

  • Electron (e⁻): Discovered by J.J. Thomson. Negatively charged, revolves around the nucleus.

  • Proton (p⁺): Discovered by Goldstein. Positively charged, resides in the nucleus.

  • Neutron (n⁰): Discovered by Chadwick. Neutral, resides in the nucleus.

• Atomic Number (Z): Number of protons in an atom. (Z = p = e in a neutral atom).

• Mass Number (A): Sum of protons and neutrons. (A = p + n).

• Isotopes: Atoms of the same element with the same atomic number but different mass numbers (e.g., ¹H¹, ¹H², ¹H³).

• Rules for Electron Filling:

  • Aufbau Principle: Orbitals are filled in order of increasing energy.

  • Pauli’s Exclusion Principle: An orbital can hold a maximum of two electrons with opposite spins.

  • Hund’s Rule: Every orbital in a subshell is singly occupied before any orbital is doubly occupied.

Of course. Here are the comprehensive and detailed notes for the topic “Matter and Its States,” compiled and structured from the uploaded document for CTET (Class 6-8) Science preparation.

Detailed Notes: Chapter 2 – Matter and Its States

2.1 Matter: Definition and Characteristics

• Definition: Anything that has mass and occupies space (volume) is called matter.

• Characteristics of Particles of Matter:

  1. They are constantly moving: Particles possess kinetic energy.

  2. They have space between them: This is called interparticle or intermolecular space. The amount of space determines the state of matter (greatest in gases, least in solids).

  3. They attract each other: These are intermolecular forces of attraction. The strength of these forces is highest in solids and lowest in gases.

2.2 The Five States of Matter

2.3 Interconversion of States of Matter

States of matter can be changed into one another by altering temperature and pressure.

A. Effect of Temperature

• Melting/Fusion: The process of a solid changing into a liquid at its melting point.

  • Melting Point of Ice = 273 K (0°C).

  • Latent Heat of Fusion: The amount of heat energy required to change 1 kg of a solid into a liquid at its melting point.

• Vaporization/Boiling: The process of a liquid changing into a vapour (gas) at its boiling point.

  • Boiling Point of Water = 373 K (100°C).

  • Latent Heat of Vaporization: The amount of heat energy required to change 1 kg of a liquid into vapour at its boiling point.

• Sublimation: The process where a solid changes directly into a gas without passing through the liquid state.

  • Sublimates: Substances that undergo sublimation (e.g., Ammonium chloride, Camphor, Naphthalene, Iodine).

• Effect of Impurity: The presence of an impurity lowers the melting point and raises the boiling point.

• Evaporation: The phenomenon of a liquid changing into vapour at any temperature below its boiling point.

  • Factors increasing evaporation: Increase in surface area, temperature, wind speed; decrease in humidity.

B. Effect of Pressure

• By applying pressure and reducing temperature, a gas can be liquefied.

• Increasing pressure raises the boiling point of a liquid.

• Dry Ice: Solid carbon dioxide (CO₂) gets converted directly into gaseous state at 1 atm pressure without becoming a liquid.

2.4 Classification of Matter

A. Pure Substances

• Have a fixed chemical composition and characteristic properties.

• Elements: Cannot be broken down into simpler substances by chemical or physical methods. They consist of only one type of atom.

• Compounds: Formed when two or more elements combine chemically in a fixed proportion. The properties of a compound are entirely different from those of its constituent elements.

B. Mixtures

• Contain two or more elements or compounds mixed in any proportion.

• Constituents can be separated by physical methods.

• They do not have a fixed melting or boiling point.

2.5 Types of Mixtures in Detail

Common Types of Colloids:

2.6 Separation Techniques of Mixtures

• For Heterogeneous Mixtures:

  • Handpicking: Removing stones from pulses.

  • Threshing: Separating grains from husks by beating.

  • Winnowing: Separating lighter husk from heavier grains using wind.

  • Sieving: Separating particles of different sizes using a sieve.

  • Sedimentation & Decantation: Allowing insoluble solids to settle (sedimentation) and then pouring out the liquid (decantation).

  • Filtration: Using a filter paper to separate an insoluble solid from a liquid.

• For Homogeneous Mixtures:

  • Evaporation: To separate a volatile solvent from a non-volatile solute (e.g., salt from seawater).

  • Centrifugation: Using centrifugal force to separate components based on density (e.g., separating butter from cream, blood testing).

  • Sublimation: To separate a sublimate from a non-sublimate (e.g., separating Ammonium chloride from sand).

  • Crystallization: The process of obtaining a pure solid in the form of crystals from its solution (e.g., purification of salt).

  • Distillation: To separate two miscible liquids based on differences in their boiling points.

  • Chromatography: Used to separate and purify components (especially coloured ones) based on their distribution between a stationary and a mobile phase (e.g., separating pigments from a plant leaf).

2.7 Metals and Non-Metals

A. Metals

• Tend to form positive ions (cations) by losing electrons.

• Physical Properties:

  • State: Solid at room temperature (except Mercury).

  • Lustre: Shiny appearance.

  • Malleability: Can be beaten into sheets (Gold and Silver are the most malleable).

  • Ductility: Can be drawn into wires (Gold and Silver are the most ductile).

  • Conductivity: Good conductors of heat and electricity (Silver > Copper > Gold > Aluminium).

  • Hardness: Generally hard (except Sodium, Potassium which are soft and can be cut with a knife).

  • Sonority: Produce sound when struck.

• Chemical Properties:

  1. Reaction with Oxygen: Form basic oxides (mostly). Some like Al₂O₃ and ZnO are amphoteric (act as both acid and base).

     • Sodium and Potassium are so reactive they catch fire in air, hence stored in kerosene.

     • Silver and Gold do not react with oxygen even at high temperatures.

  2. Reaction with Water:

     • Vigorous (Na, K): React with cold water, reaction is exothermic enough to ignite H₂.

     • Less Vigorous (Ca): Reacts with cold water.

     • With Hot Water/Steam (Mg, Al, Zn, Fe):

     • No Reaction (Pb, Cu, Ag, Au):

  3. Reaction with Acids: Metals more reactive than Hydrogen displace H₂ gas from dilute acids.

     • Mg + 2HCl → MgCl₂ + H₂↑

     • Copper does not react with dilute HCl but reacts with HNO₃ and H₂SO₄.

  4. Reaction with Bases: Some metals (e.g., Al, Zn) react with strong bases to produce Hydrogen gas.

B. Non-Metals

• Tend to form negative ions (anions) by gaining electrons.

• Physical Properties:

  • State: Can be solid (Sulphur), liquid (Bromine), or gas (Oxygen, Nitrogen).

  • Lustre: Generally not lustrous (except Iodine and Diamond).

  • Malleability/Ductility: Non-malleable and non-ductile (brittle if solid).

  • Conductivity: Poor conductors of heat and electricity (except Graphite, a form of carbon).

• Chemical Properties:

  • Form acidic oxides (e.g., SO₂, CO₂) or neutral oxides (e.g., CO, NO).

  • Generally, do not react with water.

  • Phosphorus is a very reactive non-metal and catches fire in air, so it is stored in water.

C. Reactivity Series
The series which arranges metals in the order of their decreasing reactivity.
K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au
(Most Reactive to Least Reactive)

• A more reactive metal can displace a less reactive metal from its salt solution.

  • Fe + CuSO₄ → FeSO₄ + Cu (Occurs, as Fe is more reactive than Cu)

  • Cu + AgNO₃ → No reaction (Does not occur, as Cu is less reactive than Ag)

D. Metalloids

• Elements that possess properties of both metals and non-metals.

• Examples: Germanium (Ge), Silicon (Si), Arsenic (As).

2.8 Chemical Changes & Reactions

• Physical Change: A change in which only the physical properties (shape, size, state) alter. No new substance is formed. It is generally reversible.

  • Examples: Melting of ice, Boiling of water, Breaking of glass.

• Chemical Change: A change in which one or more new substances with new chemical properties are formed. It is accompanied by a chemical reaction and is often irreversible.

  • Examples: Burning of wood, Rusting of iron, Cooking of food, Digestion.

Types of Chemical Reactions:

1. Combination Reaction: Two or more substances combine to form a single product.

   • A + B → C

   • Example: CaO (Quicklime) + H₂O → Ca(OH)₂ (Slaked lime) + Heat (Exothermic)

2. Decomposition Reaction: A single compound breaks down into two or more simpler substances.

   • AB → A + B

   • Example: 2HgO (s) → 2Hg (l) + O₂ (g)

3. Displacement Reaction: A more reactive element displaces a less reactive element from its compound.

   • A + BC → AC + B

   • Example: Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) + Cu (s)

4. Double Displacement Reaction: Ions of two compounds exchange places to form two new compounds.

   • AB + CD → AD + CB

   • Example: AgNO₃ (aq) + NaCl (aq) → AgCl (s) ↓ + NaNO₃ (aq) (Precipitation Reaction)

5. Oxidation and Reduction (Redox):

   • Oxidation: Addition of Oxygen or Removal of Hydrogen.

   • Reduction: Removal of Oxygen or Addition of Hydrogen.

   • Example: CuO + H₂ → Cu + H₂O

     • Here, H₂ is oxidised to H₂O (addition of O).

     • CuO is reduced to Cu (removal of O).

   • Rancidity: The oxidation of oils and fats in food, leading to a bad smell and taste.

   • Corrosion: The slow oxidation of metals when exposed to air and moisture (a type of redox reaction).

2.9 Corrosion and its Prevention

• Corrosion: The gradual destruction of a metal by chemical or electrochemical reaction with its environment.

  • Rusting of Iron: Formation of a reddish-brown layer of hydrated ferric oxide (Fe₂O₃.xH₂O).

  • Tarnishing of Silver: Formation of a black layer of silver sulphide (Ag₂S).

  • Green Coating on Copper: Formation of a basic copper carbonate (CuCO₃.Cu(OH)₂).

• Prevention Methods:

  • Painting/Greasing/Oiling: Creates a barrier between the metal and the atmosphere.

  • Galvanization: Coating iron with a layer of zinc.

  • Electroplating: Coating a metal with another metal using electrolysis (e.g., chromium plating).

  • Alloying: Mixing a metal with other metals or non-metals to change its properties.

2.10 Alloys

• A homogeneous mixture of two or more metals, or a metal and a non-metal.

• Purpose: To enhance properties like hardness, strength, resistance to corrosion, and lower melting points.

2.11 Atomic Structure

• Atom: The smallest particle of an element that can take part in a chemical reaction.

• Fundamental Particles:

  Particle	Symbol	Discoverer	Charge	Location
  Electron	e⁻	J.J. Thomson (1897)	-1	Revolves around nucleus
  Proton	p⁺	Goldstein (1919)	+1	Inside nucleus
  Neutron	n⁰	Chadwick (1932)	0	Inside nucleus

• Atomic Number (Z): The number of protons in the nucleus of an atom. For a neutral atom, Z = Number of protons = Number of electrons.

• Mass Number (A): The sum of the number of protons and neutrons in the nucleus. A = Number of protons (Z) + Number of neutrons (N).

• An atom is represented as: z X^A (e.g., 8 O^16).

Atomic Species:

• Isotopes: Atoms of the same element having the same atomic number (Z) but different mass numbers (A).

  • Example: 1H¹ (Protium), 1H² (Deuterium), 1H³ (Tritium). They have the same number of protons but different neutrons.

Rules for Electron Configuration:

1. Aufbau Principle: Electrons are filled in orbitals in the order of increasing energy levels.

   • Order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s...

2. Pauli’s Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite spins.

3. Hund’s Rule: Every orbital in a subshell is singly occupied before any orbital is doubly occupied.