CLASS 10 • SCIENCE

CHEMISTRY PASSING PACKAGE

Most Important Repeated Concepts

Complete Chemistry Revision Notes

Definitions • Reactions • Formulas • Important Questions

04Chapters

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ExamPreparation

01. Chemical Reactions and Equations

Definitions • Balanced Equations • Types of Reactions

Key Definitions

Chemical Reaction: Reactants are converted into products.

Chemical Equation: Symbolic representation of a chemical reaction.

Balanced Chemical Equation: Same number of atoms on LHS and RHS; follows the Law of Conservation of Mass.

Important Balanced Chemical Equations

H₂ + Cl₂ → 2HCl

2Na + 2H₂O → 2NaOH + H₂

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

CaO + H₂O → Ca(OH)₂ + Heat

CH₄ + 2O₂ → CO₂ + 2H₂O

CaCO₃ → CaO + CO₂

Mg + 2HCl → MgCl₂ + H₂

Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl

3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + 2AlCl₃

Pb + CuCl₂ → PbCl₂ + Cu

2FeSO₄ → Fe₂O₃ + SO₂ + SO₃

Zn + CuSO₄ → ZnSO₄ + Cu

ZnCO₃ → ZnO + CO₂

2AgCl → 2Ag + Cl₂

2H₂ + O₂ → 2H₂O

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

N₂ + 3H₂ → 2NH₃

NaCl + AgNO₃ → AgCl + NaNO₃

Types of Chemical Reactions

1. Combination Reaction

Two or more reactants combine to form a single product.

CaO + H₂O → Ca(OH)₂ + Heat

C + O₂ → CO₂

2H₂ + O₂ → 2H₂O

2. Decomposition Reaction

A single reactant breaks down into two or more products.

Thermal Decomposition

2FeSO₄ → Fe₂O₃ + SO₂ + SO₃

CaCO₃ → CaO + CO₂

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

Electrolytic Decomposition

2H₂O → 2H₂ + O₂

Cathode: H₂ | Anode: O₂

Photolytic Decomposition

2AgCl → 2Ag + Cl₂

2AgBr → 2Ag + Br₂

3. Oxidation, Reduction and Redox Reactions

Oxidation: Addition of oxygen.
Reduction: Removal of oxygen.
Redox Reaction: Oxidation and reduction occur simultaneously in the same reaction.

2Cu + O₂ → 2CuO

ZnO + C → Zn + CO

CuO + H₂ → Cu + H₂O

Chemical Reaction	Oxidised Substance	Reduced Substance
Na + O₂ → Na₂O	Na	O₂
CuO + H₂ → Cu + H₂O	H₂	CuO
ZnO + C → Zn + CO	C	ZnO

4. Displacement and Double Displacement Reactions

Displacement Reaction: More reactive element displaces a less reactive element.

Fe + CuSO₄ → FeSO₄ + Cu

Fe₂O₃ + 2Al → Al₂O₃ + 2Fe

Double Displacement Reaction: Exchange of ions between reactants.

Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl

Precipitation Reaction: Formation of an insoluble substance.

Precipitate: BaSO₄ (Barium sulphate). Ions involved: Ba²⁺ and SO₄²⁻.

Energy Changes, Corrosion and Rancidity

Exothermic Reaction

Heat is released.

CaO + H₂O → Ca(OH)₂ + Heat

Endothermic Reaction

Heat is absorbed.

CaCO₃ → CaO + CO₂

Corrosion: Metal reacts with atmospheric substances to form undesirable compounds. Example: Rusting of iron.

Rancidity: Change in smell or taste of oil/fat-containing foods due to oxidation.

Prevention of Rancidity: Use antioxidants, refrigerate food, flush packaging with nitrogen gas and store in airtight containers.

02. Acids, Bases and Salts

Properties • pH Scale • Chemical Reactions • Important Compounds

Properties of Acids and Bases

Property	Acids	Bases
Taste	Sour	Bitter
Ions Produced	H⁺ ions in aqueous solution	OH⁻ ions in aqueous solution
Litmus Test	Blue to Red	Red to Blue
pH Range	Below 7	Above 7
Examples	HCl, H₂SO₄, HNO₃	NaOH, KOH, Ca(OH)₂

Key Reactions of Acids and Bases

1. Reaction with Metals

Acid + Metal → Salt + Hydrogen gas

Example: Zn + H₂SO₄ → ZnSO₄ + H₂

2. Reaction with Metal Carbonates

Acid + Metal Carbonate → Salt + Water + CO₂

Example: HCl + NaHCO₃ → NaCl + H₂O + CO₂

3. Reaction with Metal Oxides

Acid + Metal Oxide → Salt + Water

Example: 2HCl + ZnO → ZnCl₂ + H₂O

4. Neutralization Reaction

Acid + Base → Salt + Water

Example: NaOH + HCl → NaCl + H₂O

Naturally Occurring Acids

Source	Natural Acid
Vinegar	Acetic acid
Orange / Lemon	Citric acid
Tamarind	Tartaric acid
Tomato	Oxalic acid
Sour milk (Curd)	Lactic acid
Nettle sting / Ant sting	Methanoic acid

Important pH Values and Definitions

Blood pH
7.2 to 7.8

Acid Rain
Below 5.6

Concentrated Acid: High concentration of H⁺ ions.
Dilute Acid: Low concentration of H⁺ ions.
Strong Acid: Completely dissociates in water.
Weak Acid: Incompletely dissociates in water.

Tooth Decay: Bacteria in the mouth convert sugar into acid, causing tooth decay. Basic toothpaste neutralizes acid to prevent decay.

pH Classification of Salts

Salt Type	Combination	pH	Example
Neutral Salt	Strong Acid + Strong Base	7	NaCl
Acidic Salt	Strong Acid + Weak Base	Below 7	NH₄Cl
Basic Salt	Weak Acid + Strong Base	Above 7	Na₂CO₃

Chlor-Alkali Process

2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂

Brine Solution: Aqueous solution of sodium chloride.

Product	Uses
NaOH	Soaps, detergents and paper making
Cl₂	Water treatment, swimming pools and PVC manufacture
H₂	Rocket fuel and margarine

Important Chemical Compounds

Common Name	Chemical Name	Formula	Uses
Washing Soda	Sodium carbonate	Na₂CO₃	Manufacture of borax, softening hard water
Baking Soda	Sodium hydrogen carbonate	NaHCO₃	Antacid, baking powder
Bleaching Powder	Calcium oxychloride	CaOCl₂	Bleaching clothes, disinfecting water
Plaster of Paris	Calcium sulphate hemihydrate	CaSO₄·½H₂O	Plastering fractured bones, toys and decorative items

Water of Crystallisation: Fixed number of water molecules present in one formula unit of a salt. Example: CuSO₄·5H₂O has 5 water molecules.

Marble and Limestone: Calcium carbonate (CaCO₃).

Soil Treatment: Farmers treat acidic soil with slaked lime [Ca(OH)₂] when soil acidity increases.

03. Metals and Non-Metals

Physical Properties • Chemical Properties • Extraction • Corrosion

Physical and Chemical Properties

Property	Metals	Non-Metals
Sonorous	Yes, produces ringing sound	No
Ductility	Ductile; can be drawn into wires	Non-ductile
Malleability	Malleable; can be beaten into sheets	Non-malleable
Lustre	Lustrous	Generally non-lustrous
Conductivity	Good conductors of heat and electricity	Poor conductors
Reaction with Dilute Acids	Liberate H₂ gas	Do not liberate H₂ gas
Nature of Oxides	Generally basic	Generally acidic
Electrochemical Nature	Electropositive	Electronegative

Special Concepts and Ionic Compounds

Amphoteric Oxides: Metal oxides that react with both acids and bases to produce salt and water. Examples: Al₂O₃ and ZnO.

Ionic Compounds: Compounds formed by the transfer of electrons from a metal to a non-metal.

Properties of Ionic Compounds

Hard and brittle.
High melting and boiling points due to strong inter-ionic electrostatic forces.
Soluble in water and insoluble in organic solvents.
Conduct electricity in molten or solution state, but not in solid state because ions are fixed.

Extraction of Zinc

1. Roasting (Sulphide Ore)

2ZnS + 3O₂ → 2ZnO + 2SO₂

Heating sulphide ore strongly in the presence of excess air.

2. Calcination (Carbonate Ore)

ZnCO₃ → ZnO + CO₂

Heating carbonate ore strongly in the presence of limited air.

3. Reduction

ZnO + C → Zn + CO

Reducing metal oxide using carbon or coke.

Corrosion and Prevention

Corrosion: Attack on metals by atmospheric substances to form undesirable compounds.

Metal	Corrosion Product
Iron	Rust
Copper	Green coating of copper carbonate
Silver	Black coating of silver sulphide

Methods of Prevention

Painting
Oiling and greasing
Galvanising
Chrome plating
Anodising
Alloy formation

Galvanising: Process of coating iron or steel with a thin layer of zinc to prevent rusting.

Alloys

Alloys: Homogeneous mixtures of two or more metals or a metal and a non-metal.

Alloy	Composition
Brass	Copper + Zinc
Bronze	Copper + Tin
Solder	Lead + Tin

Liquid Elements at Room Temperature:

Mercury – Liquid metal.
Bromine – Liquid non-metal.

Important Give-Reason Questions

Q1. Why are metals used in making cooking vessels?

Metals are good conductors of heat, have high melting points and possess high malleability.

Q2. Why is sodium metal stored in kerosene?

Sodium reacts vigorously with atmospheric oxygen and moisture, but does not react with kerosene.

Q3. Why does calcium not catch fire when it reacts with water?

The reaction of calcium with water is less violent; heat evolved is insufficient for hydrogen to catch fire.

Q4. Why do ionic compounds have high melting and boiling points?

They have strong electrostatic forces of attraction between oppositely charged ions, requiring high energy to break.

Q5. Why is H₂ gas not liberated when a metal reacts with concentrated HNO₃?

Concentrated HNO₃ is a strong oxidizing agent; it oxidizes the H₂ produced into water (H₂O).

Q6. Why are ZnO and Al₂O₃ called amphoteric oxides?

They react with both acids and bases to yield salt and water.

Q7. Why is gold used to make jewellery?

Gold is lustrous, highly ductile, malleable and extremely unreactive.

Q8. Why do solid ionic compounds not conduct electricity?

In solid state, ions are held in fixed positions and cannot move freely.

04. Carbon and Its Compounds

Covalent Bonding • Hydrocarbons • Functional Groups • Reactions

Bonding and Versatile Nature of Carbon

Covalent Bond: Sharing of electrons between two atoms.

Ionic Bond: Transfer of electrons from one atom to another.

Catenation: Carbon's unique ability to form bonds with other carbon atoms, creating long chains or rings.

Tetravalency: Carbon has 4 valence electrons and can bond with 4 other atoms.

Why C⁴⁺ cannot form: Removing 4 electrons requires a massive amount of energy.

Why C⁴⁻ cannot form: 6 protons in the nucleus cannot hold 10 electrons easily.

Classification of Hydrocarbons

Type	Bonding	General Formula	First Member	Flame
Alkanes	Single bond	CₙH₂ₙ₊₂	Methane (CH₄)	Clean blue
Alkenes	Double bond	CₙH₂ₙ	Ethene (C₂H₄)	Yellow smoky
Alkynes	Triple bond	CₙH₂ₙ₋₂	Ethyne (C₂H₂)	Yellow smoky

Functional Groups

Functional Group	Symbol
Alcohol	–OH
Aldehyde	–CHO
Ketone	C=O
Carboxylic Acid	–COOH

Structural Isomers: Compounds with the same molecular formula but different structural formulas. Example: n-butane and isobutane.

Chemical Reactions of Carbon Compounds

1. Oxidation Reaction

Alcohols are oxidised to carboxylic acids using alkaline KMnO₄ or acidified K₂Cr₂O₇.

2. Addition Reaction / Hydrogenation

Hydrogen is added to unsaturated hydrocarbons using Nickel (Ni) or Palladium catalyst to produce saturated fats.

Unsaturated oil + H₂ → Saturated fat

3. Substitution Reaction

Hydrogen atoms of a hydrocarbon are replaced by other atoms.

CH₄ + Cl₂ → CH₃Cl + HCl

4. Combustion Reaction

Burning carbon compounds in air produces CO₂, heat and light.

2CH₃OH + 3O₂ → 4H₂O + 2CO₂

Important Carbon Compounds and Processes

Ethanol: Used in alcoholic drinks, tincture iodine and medicines. Concentrated H₂SO₄ acts as a dehydrating agent, removing water from ethanol.

Ethanoic Acid (Glacial Acetic Acid): Pure ethanoic acid has a melting point of 290 K and freezes in winter, earning the name Glacial Acetic Acid.

Esterification: Reaction between carboxylic acid and alcohol to form pleasant-smelling esters used in perfumes and flavouring agents.

Saponification: Reaction of an alkaline base with long-chain carboxylic acid to form soap.

Soaps and Detergents

Soap	Detergent
Sodium salts of long-chain fatty acids.	Sodium/Potassium salts of sulphonic acids.
Forms micelles with oil and dirt.	Helps remove dirt and grease.
Does not lather in hard water.	Forms lather in both hard and soft water.

Important Molecular Formulae and Structures

Compound	Formula	Structural Description
Cyclohexane	C₆H₁₂	Cyclic alkane ring with single bonds
Benzene	C₆H₆	Aromatic ring with alternating double bonds
Propyne	C₃H₄	Three-carbon alkyne with one triple bond
Butane	C₄H₁₀	Four-carbon alkane
Propanoic Acid	C₃H₆O₂	Three-carbon carboxylic acid
Pentane	C₅H₁₂	Five-carbon alkane

Electron Dot Structure Requirements

Practise the electron dot structures of:

Methane (CH₄)
Ethene (C₂H₄)
Oxygen molecule (O₂)
Hydrogen molecule (H₂)
Nitrogen molecule (N₂)

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