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Chemical Reactions and Equations: Class 10 Chemistry Revision Notes

This lecture covers the foundational concepts of chemical reactions, distinguishing physical and chemical changes, and detailing reaction indicators. It demonstrates how to balance chemical equations using the hit-and-trial method while satisfying the law of conservation of mass. Finally, it explores four primary types of chemical reactions—combination, decomposition, displacement, and redox—along with real-life redox effects like rancidity and corrosion.

Science
English
· 12 topics
Checked against the video on 18 Oct 2026

Made by AI from the YouTube video Class 10 Science Chapter 1: Chemical Reactions and Equations | Full NCERT Covered | Gyan Shot (opens YouTube in a new tab) by Sunil Bhaiya Science, then checked against it. All credit for the teaching goes to the creator. Tap a topic’s time to confirm a fact in the video.

Chemical Reactions and Equations: Class 10 Chemistry Revision Notes

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This lecture covers the foundational concepts of chemical reactions, distinguishing physical and chemical changes, and detailing reaction indicators. It demonstrates how to balance chemical equations using the hit-and-trial method while satisfying the law of conservation of mass. Finally, it explores four primary types of chemical reactions—combination, decomposition, displacement, and redox—along with real-life redox effects like rancidity and corrosion.

Introduction, Physical vs Chemical Changes

A change is a process where the final state of a substance differs from its initial state.
In a physical change, only physical properties such as shape, size, or physical state change, while chemical composition remains identical.
- Tearing paper alters its shape and size, but its chemical identity remains paper.
- Melting ice converts solid water to liquid water without altering the H2O molecular composition.
In a chemical change, new substances with entirely different chemical compositions are formed.
- Burning paper or charcoal produces carbon dioxide gas and ash, which cannot be converted back to the original substance.
Must knowA chemical change is defined as a chemical reaction, where chemical bonds and atomic arrangements form new substances.
Rusting of iron is a chemical change because iron reacts with atmospheric oxygen and moisture to produce a new substance called rust.

Characteristics and Indicators of Chemical Reactions

Chemical reactions are identified via observable characteristics or indicators including color change, precipitate formation, gas evolution, state change, and temperature change.
Mixing colorless aqueous solutions of barium chloride and sodium sulphate immediately produces a white insoluble solid called a precipitate.
Must knowA precipitate is an insoluble solid substance formed as a result of a chemical reaction between solutions.
Burning carbon or charcoal in air consumes oxygen gas and releases carbon dioxide gas along with heat.
Chemical reactions involve energy changes classified into exothermic and endothermic reactions based on whether energy is released or absorbed.
Exothermic reactions release energy typically as heat or light, while endothermic reactions absorb energy from the surroundings.

Reaction Type	Energy Change	Key Examples
Exothermic	Energy released (heat/light)	Respiration, burning natural gas, composting, rusting
Endothermic	Energy absorbed (heat/light/electricity)	Photosynthesis, electrolysis of water, thermolysis of lead nitrate

Condition for Reactions & Writing Chemical Equations

A necessary condition for all chemical reactions is physical contact between the reacting substances so that chemical change can occur.
Word equations represent reactions using names of substances, with reactants on the left and products on the right separated by an arrow.
- Burning magnesium ribbon is represented as: Magnesium + Oxygen -> Magnesium oxide.
A skeletal equation is an unbalanced equation written with chemical formulae where the number of atoms of each element on both sides is not equal.
Must knowChemical equations must be balanced to obey the law of conservation of mass, which states that mass can neither be created nor destroyed in a chemical reaction.
Balancing ensures that the total number of atoms of each element remains identical before and after the reaction through atomic rearrangement.

Balancing Equations: Hit-and-Trial Method

In the hit-and-trial method, count each element's atoms in reactants and products, and balance the compound containing the maximum number of atoms first.
To balance Fe + H2O -> Fe3O4 + H2, start with oxygen in Fe3O4 by multiplying H2O by 4, then balance hydrogen and iron to yield 3Fe + 4H2O -> Fe3O4 + 4H2.
In the combustion of methane (CH4 + O2 -> CO2 + H2O), balancing hydrogen first by adding coefficient 2 to H2O and then balancing oxygen gives CH4 + 2O2 -> CO2 + 2H2O.
Stoichiometric coefficients are whole numbers placed before chemical formulae; subscripts within chemical formulae must never be altered during balancing.
Polyatomic ions that remain intact across the reaction can be balanced as single units to simplify calculations.
Must knowBalancing the reaction of zinc with phosphoric acid gives: 3Zn + 2H3PO4 -> Zn3(PO4)2 + 3H2.

Making Chemical Equations Informative

Balanced equations become more informative by denoting physical states: solid (s), liquid (l), gas (g), aqueous solution (aq), and precipitate (ppt. or downward arrow).
Aqueous (aq) denotes that a substance is dissolved in water as the solvent.
Reaction conditions such as temperature, pressure, catalysts, or light are indicated above or below the reaction arrow.
In the reaction 3Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g), the symbol (g) next to H2O specifies that water reacts in the form of steam.
Carbon monoxide gas reacts with hydrogen gas at an elevated pressure of 340 atm to produce methanol: CO(g) + 2H2(g) -> CH3OH(l).
Photosynthesis requires sunlight and chlorophyll over the arrow: 6CO2(aq) + 12H2O(l) -> C6H12O6(aq) + 6O2(aq) + 6H2O(l).

Combination Reactions

Must knowA combination reaction (or synthesis reaction) is a chemical reaction where two or more reactants combine to form a single product.
At room temperature (25 °C), magnesium reacts slowly with atmospheric oxygen to form a thin white protective layer of magnesium oxide (MgO) on its surface.
Magnesium ribbon must be cleaned with sandpaper before burning to remove this protective oxide layer and expose fresh metal.
When strongly heated, magnesium ribbon burns with a dazzling white flame emitting ultraviolet light, requiring eye protection like dark glasses.
The combustion of magnesium ribbon (2Mg + O2 -> 2MgO) is both a combination reaction and an exothermic reaction releasing intense heat and light.
Quicklime (calcium oxide, CaO) reacts vigorously with water to form slaked lime (calcium hydroxide, Ca(OH)2) in an exothermic reaction that boils water and produces hissing steam.
Calcium hydroxide solution applied to walls reacts with atmospheric carbon dioxide over 2 to 3 days to form a shiny white layer of calcium carbonate (CaCO3).

Decomposition Reactions: Thermal Decomposition

Must knowA decomposition reaction is a reaction where a single reactant breaks down into two or more simpler products when supplied with energy.
Thermal decomposition (thermolysis) uses heat to break down a compound.
Pale green hydrated ferrous sulphate crystals (FeSO4.7H2O) lose seven water molecules when heated, turning into white anhydrous ferrous sulphate (FeSO4).
Further strong heating of anhydrous FeSO4 produces reddish-brown ferric oxide (Fe2O3) solid along with sulphur dioxide (SO2) and sulphur trioxide (SO3) gases.
Sulphur dioxide gas is identified by its choking odor, which smells like burning sulphur or burnt matchsticks.
Safety precautions during boiling tube experiments include pointing the tube mouth away from oneself and others, using a test tube holder, and wafting fumes gently rather than sniffing directly.
Heating limestone (calcium carbonate, CaCO3) produces quicklime (CaO, used in cement manufacture) and carbon dioxide gas (CO2).
Heating white lead(II) nitrate powder [Pb(NO3)2] produces a yellow solid residue of lead(II) oxide (PbO), brown fumes of nitrogen dioxide (NO2), and oxygen gas (O2) accompanied by a crackling sound (decrepitation).

Electrolytic Decomposition (Electrolysis)

Electrolytic decomposition (electrolysis) uses electrical energy to decompose a compound into simpler substances.
Pure or distilled water is a poor conductor of electricity; a few drops of dilute sulphuric acid (H2SO4) or sodium chloride (NaCl) must be added to provide free ions for conductivity.
In the electrolysis of water, the electrode connected to the negative terminal is the cathode, while the one connected to the positive terminal is the anode.
Must knowWater decomposes during electrolysis into hydrogen gas at the cathode and oxygen gas at the anode in a 2:1 volume ratio (2H2O -> 2H2 + O2).
The volume of hydrogen collected is double that of oxygen because each water molecule contains two hydrogen atoms for every one oxygen atom.
Hydrogen is confirmed by a pop sound when a burning splinter is brought near it (combustible gas); oxygen rekindles or brightens a burning flame (supporter of combustion).

Photolytic Decomposition, Displacement Reactions

Photolytic decomposition (photolysis) is the breakdown of a substance caused by light energy (sunlight).
White silver chloride (AgCl) turns silvery-grey in sunlight as it decomposes into silver metal and greenish-yellow chlorine gas: 2AgCl -> 2Ag + Cl2.
Pale yellow silver bromide (AgBr) similarly decomposes into silver and reddish-brown bromine vapor: 2AgBr -> 2Ag + Br2.
Both AgCl and AgBr are photosensitive compounds historically used in black-and-white photography.
Must knowA displacement reaction occurs when a more reactive element displaces a less reactive element from its aqueous salt solution.
The reactivity series ranks metals in decreasing order of reactivity: K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au > Pt.
Iron displaces copper from blue copper sulphate solution, turning it pale green (FeSO4) while depositing reddish-brown copper on the iron nail (Fe + CuSO4 -> FeSO4 + Cu).
Zinc displaces copper from copper sulphate to yield colorless zinc sulphate solution (ZnSO4) and copper metal.
Copper placed in ferrous sulphate solution shows no reaction because copper is less reactive than iron.

Double Displacement Reactions

Must knowA double displacement reaction (metathesis) involves an mutual exchange of ions (cations and anions) between two aqueous reactant compounds.
Types of double displacement reactions include precipitation reactions, acid-base neutralization reactions, and gas-forming reactions.
Mixing aqueous sodium sulphate (Na2SO4) and barium chloride (BaCl2) produces a white precipitate of barium sulphate (BaSO4) and aqueous sodium chloride (NaCl).
Mixing potassium iodide (KI) and lead(II) nitrate [Pb(NO3)2] produces a bright yellow precipitate of lead(II) iodide (PbI2) and aqueous potassium nitrate (KNO3).
Mixing silver nitrate (AgNO3) and sodium chloride (NaCl) yields a white precipitate of silver chloride (AgCl) and aqueous sodium nitrate (NaNO3).
Passing hydrogen sulphide gas (H2S) through blue copper sulphate (CuSO4) solution yields a black precipitate of copper sulphide (CuS) and sulphuric acid (H2SO4).

Redox Reactions

A redox reaction is a chemical reaction where reduction and oxidation occur simultaneously.
Oxidation is defined as the addition of oxygen, the removal of hydrogen, or both.
Reduction is defined as the addition of hydrogen, the removal of oxygen, or both.
An oxidizing agent (oxidant) is the substance that gets reduced by supplying oxygen or removing hydrogen.
A reducing agent (reductant) is the substance that gets oxidized by removing oxygen or supplying hydrogen.
Heating reddish-brown copper in air oxidizes copper to black copper(II) oxide: 2Cu + O2 -> 2CuO (Cu is oxidized; O2 is reduced).
Passing hydrogen gas over heated black copper(II) oxide reverses the color back to reddish-brown: CuO + H2 -> Cu + H2O (CuO is reduced to Cu; H2 is oxidized to H2O).
All single displacement reactions and combustion reactions are redox reactions.

Effects of Oxidation in Daily Life: Rancidity and Corrosion

Rancidity is the aerial oxidation of fats and oils in food items, producing unpleasant smells and foul tastes.
Prevention of rancidity includes adding antioxidants (such as Vitamin C and Vitamin E), flushing chip packets with unreactive gases like nitrogen or helium, and refrigeration to slow oxidation rates.
Vacuum packing is avoided for fragile snacks like chips because the lack of cushion gas causes the chips to crush.
Must knowCorrosion is the slow deterioration of metal surfaces into more stable compounds like oxides, sulphides, or carbonates through the attack of atmospheric gases and moisture.
Rusting of iron forms a flaky reddish-brown hydrated iron oxide layer (rust).
Corrosion of copper produces a green protective surface layer called patina (basic copper carbonate).
Corrosion of silver forms a black tarnish layer of silver sulphide (Ag2S) upon reaction with airborne hydrogen sulphide.

Key terms

Precipitate: An insoluble solid that separates out from a liquid solution during a chemical reaction
Exothermic Reaction: A reaction that releases energy into the surroundings, primarily in the form of heat or light
Endothermic Reaction: A reaction that absorbs energy from the surroundings in the form of heat, light, or electricity
Skeletal Equation: An unbalanced chemical equation showing reactant and product formulae without equal numbers of atoms on both sides
Combination Reaction: A reaction in which two or more elements or compounds join together to form a single product
Thermal Decomposition: A breakdown reaction driven by heat energy (thermolysis)
Electrolysis: A decomposition reaction driven by an electric current passing through an electrolyte or acidified liquid
Photolysis: A decomposition reaction driven by the absorption of light energy
Displacement Reaction: A reaction where a more reactive element displaces a less reactive element from its compound
Double Displacement Reaction: A reaction in which two reacting ionic compounds exchange their respective cations and anions
Redox Reaction: A reaction involving simultaneous oxidation of one reactant and reduction of another
Oxidizing Agent: A reactant that causes another substance to be oxidized and is itself reduced
Reducing Agent: A reactant that causes another substance to be reduced and is itself oxidized
Rancidity: The spoilage of fats and oils in food via aerial oxidation, causing offensive odor and bad taste
Corrosion: The gradual destruction of metals caused by chemical reactions with atmospheric gases and moisture

Exam tips

Never alter subscripts in chemical formulae when balancing equations; only change the stoichiometric coefficients in front.
Identify the compound with the maximum number of atoms first and balance the element within it that has the highest atom count.
Single displacement reactions are always redox reactions, whereas double displacement precipitation reactions are generally non-redox ion exchanges.

Quick revision

Magnesium ribbon must be rubbed with sandpaper before ignition to scrape off the protective surface layer of magnesium oxide.
Magnesium burns in oxygen with an intense dazzling white flame emitting hazardous UV radiation.
Chemical equations must be balanced to conform with the law of conservation of mass.
Decomposition of pale green FeSO4.7H2O crystals produces white anhydrous FeSO4, followed by reddish-brown Fe2O3 solid and choking SO2 and SO3 gases.
In the electrolysis of water, hydrogen gas collects at the cathode and oxygen gas collects at the anode in a 2:1 volume ratio.
White silver chloride turns grey in sunlight due to photolytic decomposition into silver metal and chlorine gas.
Slaked lime solution [Ca(OH)2] applied to walls reacts with airborne CO2 to form a shiny white coating of calcium carbonate (CaCO3).
Iron nails immersed in blue CuSO4 solution become coated with reddish-brown copper while the liquid fades to pale green FeSO4.
Lead(II) nitrate decomposes on heating to yield yellow lead(II) oxide, pungent brown NO2 fumes, and O2 gas.
Antioxidants like vitamins C and E and inert flushing gases like nitrogen and helium prevent rancidity in packaged oily foods.

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