Class 10 Science Chapter 1 Notes: Chemical Reactions and Equations (2026-2027)

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H1 Title: Class 10 Science Chapter 1 Notes: Chemical Reactions and Equations (Ultimate 2026-2027 Board Revision Guide)

Navigating through the CBSE Class 10 Science curriculum requires an in-depth understanding of chemical transformations, symbolic equation balancing, reaction thermodynamics, and ionic displacement mechanisms. Chapter 1 of Class 10 Science, “Chemical Reactions and Equations”, forms the foundation of all senior secondary chemistry, industrial synthesis, metabolic biochemistry, and material science. It investigates the microscopic rearrangement of atoms during bond-breaking and bond-forming processes; explores the observable macroscopic indicators of chemical change; and details the stoichiometric balancing of mass under the Law of Conservation of Mass alongside real-world redox phenomena like corrosion and rancidity. To help students master every aspect of this high-weightage chapter, this comprehensive guide offers textbook-accurate, highly structured, and pedagogically sound responses strictly aligned with the latest CBSE evaluation standards.

Every concept presented in the official NCERT textbook—ranging from line-by-line theoretical elucidations and an exhaustive dissection of all 11 practical NCERT activities to zero-mark-loss balanced chemical equations and a concluding master revision cheat sheet—has been detailed with precision. Key scoring terms, state symbols, color shifts, and examiner cautions have been highlighted to ensure students secure a perfect score in their CBSE Board Examinations.

Chapter 1: Chemical Reactions and Equations

Master Concept Overview & Thematic Architecture

A chemical reaction is a process in which one or more substances (reactants) undergo chemical transformation through the breaking and making of bonds to form entirely new substances (products) possessing distinct physical and chemical properties.

Reaction CategoryDefining PhenomenonGeneral Symbolic FormatEnergy CharacteristicPrototypical Board Exam Example
CombinationTwo or more reactants coalesce into a single product$A + B \longrightarrow AB$Predominantly Exothermic$\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$
DecompositionA single compound splits into two or more simpler units$AB \overset{\Delta / h\nu / \text{elec}}{\longrightarrow} A + B$Exclusively Endothermic$2\text{FeSO}_4(s) \overset{\Delta}{\longrightarrow} \text{Fe}_2\text{O}_3(s) + \text{SO}_2(g) + \text{SO}_3(g)$
DisplacementA more reactive element displaces a less reactive element$A + BC \longrightarrow AC + B$Exothermic$\text{Fe}(s) + \text{CuSO}_4(aq) \longrightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$
Double DisplacementMutual exchange of ions between two ionic compounds$AB + CD \longrightarrow AD + CB$Iso-thermal / Precipitation$\text{Na}_2\text{SO}_4(aq) + \text{BaCl}_2(aq) \longrightarrow \text{BaSO}_4(s)\downarrow + 2\text{NaCl}(aq)$
RedoxSimultaneous transfer of electrons / oxygen / hydrogen$\text{Oxidant} + \text{Reductant} \longrightarrow \text{Products}$Variable$\text{CuO}(s) + \text{H}_2(g) \overset{\Delta}{\longrightarrow} \text{Cu}(s) + \text{H}_2\text{O}(l)$

🧠 Examiner’s Secret: A major source of mark loss in Board Examinations is writing bare chemical formulas without physical state symbols [$(s)$, $(l)$, $(g)$, $(aq)$] and reaction conditions [$\Delta$, catalysts, temperature, sunlight]. The CBSE Marking Scheme awards zero marks for unbalanced chemical equations, regardless of how accurate the accompanying descriptive text is.

Foundational Chemistry: Representation and Equation Balancing

Chemical Equation

A chemical equation is the symbolic representation of a chemical reaction using chemical formulas of reactants and products, accompanied by state symbols and stoichiometric coefficients.

In any chemical equation, the starting materials are written on the left-hand side (Reactants, LHS) with a plus sign ($+$) between them. The resulting substances are written on the right-hand side (Products, RHS). An arrow ($\longrightarrow$) points from reactants to products, indicating the direction of chemical change. Reaction conditions (such as heat indicated by $\Delta$, sunlight, catalysts, or pressure) are written above or below this arrow.

Indicators of a Chemical Reaction

A chemical reaction can be confirmed by observing one or more macroscopic changes:

  1. Change in State: Conversion of liquid reactants into solid precipitates or gases (e.g., reaction of ammonia gas with hydrogen chloride gas producing solid ammonium chloride fumes).
  2. Change in Colour: Shift in chromatic absorption (e.g., blue copper sulphate solution fading to pale green upon adding iron nails).
  3. Evolution of a Gas: Release of gaseous bubbles (e.g., effervescence of hydrogen gas when granulated zinc reacts with dilute sulphuric acid).
  4. Change in Temperature:
    • Exothermic: Temperature rises as thermal energy is liberated (e.g., adding water to quicklime).
    • Endothermic: Temperature drops as thermal energy is absorbed (e.g., dissolving barium hydroxide in ammonium chloride).
  5. Formation of a Precipitate: Separation of an insoluble solid phase from an aqueous solution (e.g., formation of bright yellow lead iodide precipitate).

The Law of Conservation of Mass & Balancing Technique

The Law of Conservation of Mass states that matter can neither be created nor destroyed in a chemical reaction; hence, the total mass of the products must equal the total mass of the reactants.

To satisfy this fundamental law, the total number of atoms of each individual element must remain identical on both sides of the skeletal equation. Balancing is achieved using the systematic Hit-and-Trial Method without ever altering the chemical formulas (subscripts) of compounds:

Step-by-Step Balancing Walkthrough: Reaction of Steam on Iron

  • Step 1: Write the Skeletal Word and Chemical Equation
    $$\text{Iron} + \text{Steam} \longrightarrow \text{Iron(II,III) Oxide} + \text{Hydrogen}$$
    $$\text{Fe} + \text{H}_2\text{O} \longrightarrow \text{Fe}_3\text{O}_4 + \text{H}_2$$
  • Step 2: Tabulate the Atom Inventory
ElementNumber of Atoms in Reactants (LHS)Number of Atoms in Products (RHS)
$\text{Fe}$13
$\text{H}$22
$\text{O}$14
  • Step 3: Balance the Element with Maximum Atoms ($\text{Oxygen}$)
    To balance Oxygen, place coefficient $4$ in front of $\text{H}_2\text{O}$ on LHS:
    $$\text{Fe} + 4\text{H}_2\text{O} \longrightarrow \text{Fe}_3\text{O}_4 + \text{H}_2$$
  • Step 4: Balance Hydrogen Atoms
    LHS now has $4 \times 2 = 8$ hydrogen atoms. Place coefficient $4$ in front of $\text{H}_2$ on RHS:
    $$\text{Fe} + 4\text{H}_2\text{O} \longrightarrow \text{Fe}_3\text{O}_4 + 4\text{H}_2$$
  • Step 5: Balance Iron Atoms
    RHS has $3$ iron atoms in $\text{Fe}_3\text{O}_4$. Place coefficient $3$ in front of $\text{Fe}$ on LHS:
    $$3\text{Fe} + 4\text{H}_2\text{O} \longrightarrow \text{Fe}_3\text{O}_4 + 4\text{H}_2$$
  • Step 6: Verify and Append State Symbols
    $$\text{LHS: } \text{Fe} = 3, \text{H} = 8, \text{O} = 4 \quad \Longleftrightarrow \quad \text{RHS: } \text{Fe} = 3, \text{H} = 8, \text{O} = 4$$
    Append physical states ($\text{H}_2\text{O}$ is steam, designated by $g$):
    $$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \longrightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)$$

Exhaustive NCERT Laboratory Activities Vault

Activity 1.1: Combustion of Magnesium Ribbon

  • Aim: To observe the combustion of magnesium ribbon in the presence of atmospheric oxygen.
  • Setup: A $3 – 4\text{ cm}$ piece of magnesium ribbon is held with a pair of tongs over a watch glass and ignited using a Bunsen burner.

Tongs Dazzling White Flame Watch Glass (White MgO Powder)

  • Mandatory Precaution: The magnesium ribbon must be rubbed thoroughly with sandpaper before ignition.
    • Scientific Reason: Magnesium is a highly reactive alkaline earth metal. Upon prolonged exposure to moist air, it forms a protective, unreactive layer of basic magnesium carbonate [a mixture of $\text{MgO}$ and $\text{MgCO}_3$] on its surface. Sanding removes this passivating coating, ensuring direct contact between metallic magnesium and atmospheric oxygen.
  • Observations:
    1. The magnesium ribbon burns with a blinding, dazzling white flame.
    2. It transforms into a powdery white residue collected in the watch glass.
  • Balanced Chemical Equation:
    $$2\text{Mg}(s) + \text{O}_2(g) \longrightarrow 2\text{MgO}(s)$$
  • Board Conclusion & Chemical Nature: The white powder is Magnesium Oxide ($\text{MgO}$). When dissolved in water, it forms magnesium hydroxide:
    $$\text{MgO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Mg(OH)}_2(aq)$$
    This solution turns red litmus paper blue, confirming that magnesium oxide is a basic metallic oxide.

Activity 1.2: Precipitation of Lead(II) Iodide

  • Aim: To observe the precipitation reaction between lead nitrate and potassium iodide solutions.
  • Setup: Take dilute lead nitrate solution in a test tube and add potassium iodide solution dropwise.
  • Observations:
    1. A brilliant, bright yellow precipitate forms instantaneously upon mixing the two clear, colourless solutions.
    2. The reaction mixture becomes opaque yellow and gradually settles at the bottom of the test tube.
  • Balanced Chemical Equation:
    $$\text{Pb(NO}_3)_2(aq) + 2\text{KI}(aq) \longrightarrow \text{PbI}_2(s)\downarrow + 2\text{KNO}_3(aq)$$
  • Board Conclusion: This is a classic example of both a Precipitation Reaction and a Double Displacement Reaction. The yellow insoluble solid is Lead(II) Iodide [$\text{PbI}_2$], while potassium nitrate [$\text{KNO}_3$] remains dissolved in the colourless aqueous solution.

Activity 1.3: Reaction of Zinc with Mineral Acids

  • Aim: To observe the evolution of hydrogen gas and associated thermal changes during the reaction of zinc with acids.
  • Setup: Place granulated zinc in a conical flask fitted with a single-bore cork and a glass delivery tube. Pour dilute sulphuric acid (or dilute hydrochloric acid) over the zinc.
  • Observations:
    1. Rapid effervescence occurs on the surface of the zinc granules with the evolution of a colourless, odourless gas.
    2. When the conical flask is touched from the bottom, it feels distinctly warm, indicating an increase in temperature.
    3. When a burning splinter is brought near the mouth of the delivery tube, the gas burns with a characteristic ‘pop’ sound, confirming it is hydrogen.
  • Balanced Chemical Equations:
    • With Dilute Sulphuric Acid:
      $$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)\uparrow + \text{Heat}$$
    • With Dilute Hydrochloric Acid:
      $$\text{Zn}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)\uparrow + \text{Heat}$$
  • Board Conclusion: The reaction is a Single Displacement Reaction (as Zinc is more reactive than Hydrogen) as well as an Exothermic Reaction ($\Delta H < 0$).

Activity 1.4: Slaking of Lime (Combination Reaction)

  • Aim: To observe the combination reaction between calcium oxide (quicklime) and water.
  • Setup: Take a small quantity of calcium oxide ($\text{CaO}$) in a beaker. Slowly add water to it.
  • Observations:
    1. The reaction is vigorous and produces a hissing sound.
    2. A large volume of thermal energy is liberated, causing the water to boil and the beaker to become extremely hot.
    3. The solid quicklime breaks apart, dissolving partially to produce a milky suspension that settles into a clear solution over time.
  • Balanced Chemical Equation:
    $$\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$
  • Board Terminology Note:
    • $\text{CaO}(s)$ is commonly called Quicklime (Chemical name: Calcium Oxide).
    • $\text{Ca(OH)}_2(aq)$ is commonly called Slaked Lime (Chemical name: Calcium Hydroxide).
    • The clear filtrate of slaked lime in water is called Limewater.

💡 Did You Know? (The Whitewashing Chemistry): When slaked lime [$\text{Ca(OH)}_2$] is applied to walls, it reacts slowly with atmospheric carbon dioxide over $2 – 3\text{ days}$ to form a thin, durable, lustrous layer of Calcium Carbonate [$\text{CaCO}_3$]:
$$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s) + \text{H}_2\text{O}(l)$$
Marble, limestone, chalk, eggshells, and seashells share this identical chemical formula: $\text{CaCO}_3$.

Activity 1.5: Thermal Decomposition of Ferrous Sulphate Crystals

  • Aim: To observe the thermal decomposition of ferrous sulphate heptahydrate crystals.
  • Setup: Place approximately $2\text{ g}$ of green ferrous sulphate crystals in a dry boiling tube. Heat the tube over the flame of a burner while keeping the mouth directed away from yourself and others.
  • Observations:
    1. Water droplets condense on the cooler upper inner walls of the boiling tube as water of crystallisation is driven off.
    2. The initial pale green colour of the crystals fades to anhydrous white, which then turns into a reddish-brown solid.
    3. A pungent, choking gas with the distinct smell of burning sulphur (resembling burning matches) is evolved.
  • Balanced Chemical Equations:
    • Step 1: Loss of Water of Crystallisation
      $$\text{FeSO}_4 \cdot 7\text{H}_2\text{O}(s) \xrightarrow{\Delta} \text{FeSO}_4(s) [\text{Dirty White}] + 7\text{H}_2\text{O}(g)$$
    • Step 2: Thermal Cleavage of Anhydrous Salt
      $$2\text{FeSO}_4(s) \xrightarrow{\Delta} \text{Fe}_2\text{O}_3(s) [\text{Reddish-Brown}] + \text{SO}_2(g) + \text{SO}_3(g)$$
  • Board Conclusion: This is a Thermal Decomposition Reaction. The reddish-brown residue is Ferric Oxide ($\text{Fe}_2\text{O}_3$). The gaseous byproducts are Sulphur Dioxide ($\text{SO}_2$) and Sulphur Trioxide ($\text{SO}_3$), both of which are acidic oxides that turn moist blue litmus paper red.

Activity 1.6: Thermal Decomposition of Lead Nitrate

  • Aim: To observe the thermal decomposition of lead(II) nitrate powder.
  • Setup: Heat about $2\text{ g}$ of white lead nitrate powder in a dry, heat-resistant boiling tube using a pair of test tube holders.
  • Observations:
    1. A sharp, crackling sound (decrepitation) is heard as the crystals shatter under heat.
    2. Copious, thick brown fumes are liberated from the boiling tube.
    3. A solid residue remains in the tube, which is reddish-brown when hot and turns bright yellow upon cooling.
  • Balanced Chemical Equation:
    $$2\text{Pb(NO}_3)_2(s) \xrightarrow{\Delta} 2\text{PbO}(s) [\text{Yellow}] + 4\text{NO}_2(g) [\text{Brown Fumes}] + \text{O}_2(g)$$
  • Gas Identification Protocols:
    • The brown gas is Nitrogen Dioxide ($\text{NO}_2$), which is toxic and acidic.
    • The accompanying colourless gas is Oxygen ($\text{O}_2$), confirmed by the rekindling of a glowing wooden splint held at the mouth of the tube.
    • The yellow residue sticking to the glass walls is Lead(II) Oxide ($\text{PbO}$).

Activity 1.7: Electrolysis of Water (Electrolytic Decomposition)

  • Aim: To demonstrate the electrolytic decomposition of water into its constituent elements and establish their volumetric ratio.
  • Setup: A plastic mug is fitted with two carbon (graphite) electrodes sealed through rubber stoppers at its base. The mug is filled with distilled water containing a few drops of dilute sulphuric acid ($\text{H}_2\text{SO}_4$). Two graduated test tubes filled with water are inverted over the electrodes. A $6\text{ V}$ DC power supply is connected.

2V: H₂ 1V: O₂ 6V DC Cathode (-) Anode (+)

  • Critical Operational Queries:
    • Why is dilute $\text{H}_2\text{SO}_4$ added? Pure distilled water is a covalent compound with an extremely low degree of self-ionisation, making it a poor electrical conductor. Adding a few drops of acid provides free, mobile hydronium [$\text{H}_3\text{O}^+$] and sulphate [$\text{SO}_4^{2-}$] ions, significantly increasing its electrical conductivity.
  • Observations:
    1. Continuous streams of gas bubbles form at both submerged electrodes, displacing water downward in the inverted test tubes.
    2. The rate of gas collection is unequal: the volume of gas collected at the Cathode (negative electrode) is exactly twice the volume of gas collected at the Anode (positive electrode).
  • Gas Confirmation Tests:
    • Cathode Gas: When a burning splint is brought near, it ignites with a sharp pop sound, confirming it is Hydrogen ($\text{H}_2$).
    • Anode Gas: When a glowing splint is brought near, the splint bursts into flame, confirming it is Oxygen ($\text{O}_2$).
  • Balanced Chemical Equation:
    $$2\text{H}_2\text{O}(l) \xrightarrow{\text{Electric Current}} 2\text{H}_2(g) [\text{Cathode}] + \text{O}_2(g) [\text{Anode}]$$
  • Theoretical Proof for 2:1 Ratio: A single molecule of water consists of two hydrogen atoms bonded to one oxygen atom ($2:1$ atomic ratio). During electrolysis, two moles of liquid water decompose into two moles of hydrogen gas and one mole of oxygen gas:
    $$V_{\text{H}2} : V{\text{O}_2} = 2 : 1$$

Activity 1.8: Photochemical Decomposition of Silver Salts

  • Aim: To observe the photolytic decomposition of silver chloride and silver bromide in sunlight.
  • Setup: Place about $2\text{ g}$ of pure white silver chloride ($\text{AgCl}$) on a clean china dish and expose it to natural sunlight for $1 – 2\text{ hours}$.
  • Observations:
    The pure white colour of silver chloride gradually turns into a distinct dull grey.
  • Balanced Chemical Equations:
    $$2\text{AgCl}(s) [\text{White}] \xrightarrow{\text{Sunlight}} 2\text{Ag}(s) [\text{Grey}] + \text{Cl}_2(g)$$
    $$2\text{AgBr}(s) [\text{Pale Yellow}] \xrightarrow{\text{Sunlight}} 2\text{Ag}(s) [\text{Grey}] + \text{Br}_2(g)$$
  • Commercial Board Application: Both silver chloride and silver bromide undergo photolysis and are widely used in traditional black-and-white photographic paper and films. They are always stored in dark-amber or opaque reagent bottles to prevent premature decomposition caused by ambient light.

Activity 1.9: Single Displacement Reaction (Iron and Copper Sulphate)

  • Aim: To observe the displacement of copper from copper sulphate solution by metallic iron.
  • Setup: Clean two iron nails with sandpaper. Suspend them in a test tube containing $10\text{ mL}$ of blue copper sulphate solution using a thread for $20\text{ minutes}$. Keep a third nail and an untouched test tube of copper sulphate solution aside as controls.
  • Observations:
    1. The deep blue colour of the copper sulphate solution gradually fades and transforms into a light pale-green colour.
    2. The submerged iron nails become coated with a thick, rough, reddish-brown deposit.
  • Balanced Chemical Equation:
    $$\text{Fe}(s) [\text{Grey}] + \text{CuSO}_4(aq) [\text{Blue}] \longrightarrow \text{FeSO}_4(aq) [\text{Pale Green}] + \text{Cu}(s) [\text{Reddish-Brown}]$$
  • Reactivity Series Justification: Iron lies higher than copper in the reactivity series of metals ($\text{Fe} > \text{Cu}$). Iron has a lower standard reduction potential, meaning it oxidises more readily by releasing two electrons to reduce copper ions:
    $$\text{Fe}(s) \longrightarrow \text{Fe}^{2+}(aq) + 2e^- \quad (\text{Oxidation})$$
    $$\text{Cu}^{2+}(aq) + 2e^- \longrightarrow \text{Cu}(s) \quad (\text{Reduction})$$

Activity 1.10: Double Displacement Reaction (Sodium Sulphate and Barium Chloride)

  • Aim: To observe the mutual exchange of ions between aqueous sodium sulphate and barium chloride.
  • Setup: Mix $3\text{ mL}$ of sodium sulphate solution with $3\text{ mL}$ of barium chloride solution in a clean test tube.
  • Observations:
    A dense, insoluble white precipitate forms instantaneously upon contact, remaining suspended before gradually settling to the bottom.
  • Balanced Molecular and Ionic Equations:
    • Molecular Equation:
      $$\text{Na}_2\text{SO}_4(aq) + \text{BaCl}_2(aq) \longrightarrow \text{BaSO}_4(s)\downarrow [\text{White ppt}] + 2\text{NaCl}(aq)$$
    • Net Ionic Equation:
      $$\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \longrightarrow \text{BaSO}_4(s)\downarrow$$
  • Board Classification: This is a Double Displacement Reaction accompanied by precipitation. The white precipitate is Barium Sulphate ($\text{BaSO}_4$), which is insoluble in water and dilute acids.

Activity 1.11: Redox Reactions (Oxidation and Reduction of Copper)

  • Aim: To observe the oxidation of copper powder to copper(II) oxide and its subsequent reduction back to metallic copper.
  • Setup: Heat about $1\text{ g}$ of reddish-brown copper powder in a china dish over a burner. Subsequently, pass dry hydrogen gas over the heated black substance.
  • Observations:
    • Phase 1 (Heating in air): The reddish-brown copper powder becomes coated with a uniform black surface layer.
    • Phase 2 (Passing $\text{H}_2$ gas): The black coating turns back into a shiny reddish-brown solid, while moisture droplets condense on the cooler parts of the setup.
  • Balanced Chemical Equations:
    • Phase 1 (Oxidation of Copper):
      $$2\text{Cu}(s) [\text{Reddish-Brown}] + \text{O}_2(g) \xrightarrow{\Delta} 2\text{CuO}(s) [\text{Black Copper(II) Oxide}]$$
    • Phase 2 (Reduction by Hydrogen):
      $$\text{CuO}(s) [\text{Black}] + \text{H}_2(g) \xrightarrow{\Delta} \text{Cu}(s) [\text{Reddish-Brown}] + \text{H}_2\text{O}(l)$$
  • Redox Species Identification (Phase 2):
    • Substance Oxidised: $\text{H}_2$ (Gains oxygen to form $\text{H}_2\text{O}$).
    • Substance Reduced: $\text{CuO}$ (Loses oxygen to form elemental $\text{Cu}$).
    • Oxidising Agent: $\text{CuO}$ (Supplies oxygen for oxidation).
    • Reducing Agent: $\text{H}_2$ (Removes oxygen from copper(II) oxide).

[👉 Also Read: Class 10 Science Chapter 2 Acids, Bases and Salts Notes]

In-Depth Taxonomic Breakdown of Chemical Reaction Types

Combination Reactions

A combination reaction is a reaction where two or more distinct reactants combine chemically to form a single product.

$$\text{General Scheme: } A + B \longrightarrow AB$$

Classification of Combination Reactions:

  1. Combination of Two Elements:
    • Combustion of Carbon: $\text{C}(s) + \text{O}_2(g) \longrightarrow \text{CO}_2(g) + \text{Heat}$
    • Synthesis of Water: $2\text{H}_2(g) + \text{O}_2(g) \longrightarrow 2\text{H}_2\text{O}(l)$
  2. Combination of an Element and a Compound:
    • Oxidation of Carbon Monoxide: $2\text{CO}(g) + \text{O}_2(g) \longrightarrow 2\text{CO}_2(g)$
    • Chlorination of Sulphur Dioxide: $\text{SO}_2(g) + \text{Cl}_2(g) \longrightarrow \text{SO}_2\text{Cl}_2(l)$
  3. Combination of Two Compounds:
    • Slaking of Quicklime: $\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$
    • Ammonium Chloride Formation: $\text{NH}_3(g) + \text{HCl}(g) \longrightarrow \text{NH}_4\text{Cl}(s)$

Exothermic vs. Endothermic Reactions

Exothermic reactions are chemical transformations that release thermal energy into their surroundings, causing the reaction mixture’s temperature to rise.

$$\text{Reactants} \longrightarrow \text{Products} + \text{Heat} \quad (\Delta H < 0)$$

Examples of Exothermic Reactions:

  1. Combustion of Natural Gas (Methane):
    $$\text{CH}_4(g) + 2\text{O}_2(g) \longrightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) + \text{Heat Energy}$$
  2. Cellular Respiration:
    Biological combustion in human cells:
    $$\text{C}6\text{H}{12}\text{O}_6(aq) + 6\text{O}_2(aq) \longrightarrow 6\text{CO}_2(aq) + 6\text{H}_2\text{O}(l) + \text{Energy (ATP)}$$
  3. Decomposition of Organic Vegetable Matter:
    The microbial breakdown of vegetable waste into nutrient-rich compost releases heat, making it an exothermic decomposition reaction.

Endothermic reactions are chemical reactions that absorb thermal, electrical, or radiant energy from their surroundings, causing the reaction mixture’s temperature to fall.

$$\text{Reactants} + \text{Energy} \longrightarrow \text{Products} \quad (\Delta H > 0)$$

Example: Barium hydroxide reacting with ammonium chloride:
$$\text{Ba(OH)}_2 \cdot 8\text{H}_2\text{O}(s) + 2\text{NH}_4\text{Cl}(s) \longrightarrow \text{BaCl}_2(aq) + 2\text{NH}_3(aq) + 10\text{H}_2\text{O}(l)$$
(This reaction absorbs so much thermal energy that the bottom of the flask can freeze to a wet wooden block).

Decomposition Reactions

A decomposition reaction is a chemical process where a single reactant breaks down into two or more simpler, smaller products upon the absorption of energy.

$$\text{General Scheme: } AB \xrightarrow{\text{Energy Input}} A + B$$

Decomposition reactions are classified by the form of energy supplied: Thermal (Δ) Heat Energy e.g., CaCO₃, FeSO₄ Electrolytic Electric Current e.g., H₂O, NaCl Photolytic (hν) Sunlight Energy e.g., AgCl, AgBr

Detailed Forms of Decomposition:

  1. Thermal Decomposition (Thermolysis): Driven by heat energy.
    • Limestone to Quicklime (Industrial manufacture of cement):
      $$\text{CaCO}_3(s) \xrightarrow{\Delta} \text{CaO}(s) + \text{CO}_2(g)$$
    • Ferrous Sulphate Breakdown:
      $$2\text{FeSO}_4(s) \xrightarrow{\Delta} \text{Fe}_2\text{O}_3(s) + \text{SO}_2(g) + \text{SO}_3(g)$$
  2. Electrolytic Decomposition (Electrolysis): Driven by electric current.
    • Electrolysis of Molten Alumina:
      $$2\text{Al}_2\text{O}_3(l) \xrightarrow{\text{Electric Current}} 4\text{Al}(l) + 3\text{O}_2(g)$$
  3. Photolytic Decomposition (Photolysis): Driven by photons (sunlight).
    • Photolysis of Hydrogen Peroxide:
      $$2\text{H}_2\text{O}_2(l) \xrightarrow{\text{Sunlight}} 2\text{H}_2\text{O}(l) + \text{O}_2(g)$$

Displacement Reactions

A displacement reaction is a chemical reaction in which a more chemically active element displaces a less reactive element from its aqueous salt solution.

$$\text{General Scheme: } A + BC \longrightarrow AC + B \quad (\text{Where } A \text{ is more reactive than } B)$$

Reactivity Order Governing Single Displacement:

$$\text{K} > \text{Na} > \text{Ca} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Pb} > [\text{H}] > \text{Cu} > \text{Hg} > \text{Ag} > \text{Au}$$

Prototypical Examples:

  1. Zinc Displacing Copper:
    Zinc is more electropositive than copper:
    $$\text{Zn}(s) + \text{CuSO}_4(aq) [\text{Blue}] \longrightarrow \text{ZnSO}_4(aq) [\text{Colourless}] + \text{Cu}(s) [\text{Reddish-Brown}]$$
  2. Lead Displacing Copper:
    Lead is more reactive than copper:
    $$\text{Pb}(s) + \text{CuCl}_2(aq) [\text{Greenish-Blue}] \longrightarrow \text{PbCl}_2(aq) [\text{Colourless}] + \text{Cu}(s) [\text{Reddish-Brown}]$$
  3. Non-Reactivity of Less Reactive Metals:
    $$\text{Cu}(s) + \text{FeSO}_4(aq) \longrightarrow \text{No Reaction}$$
    (Because Copper lies below Iron in the reactivity series, it cannot displace iron ions).

Double Displacement and Precipitation Reactions

A double displacement reaction is a chemical process involving the mutual exchange of ions between two aqueous compounds to form two completely new compounds.

$$\text{General Scheme: } A^+B^- + C^+D^- \longrightarrow A^+D^- + C^+B^-$$

Double displacement reactions generally fall into two categories:

  1. Precipitation Reactions: At least one of the resulting products is insoluble in water and separates as an opaque solid phase (a precipitate).
    • Formation of White Barium Sulphate:
      $$\text{BaCl}_2(aq) + \text{Na}_2\text{SO}_4(aq) \longrightarrow \text{BaSO}_4(s)\downarrow + 2\text{NaCl}(aq)$$
    • Formation of Bright Yellow Lead Iodide:
      $$\text{Pb(NO}_3)_2(aq) + 2\text{KI}(aq) \longrightarrow \text{PbI}_2(s)\downarrow + 2\text{KNO}_3(aq)$$
    • Formation of White Silver Chloride:
      $$\text{AgNO}_3(aq) + \text{NaCl}(aq) \longrightarrow \text{AgCl}(s)\downarrow + \text{NaNO}_3(aq)$$
  2. Neutralisation Reactions: An acid reacts with a base to form a soluble salt and water, with no precipitate formed:
    $$\text{HCl}(aq) + \text{NaOH}(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$

Oxidation, Reduction, and Redox Chemistry

A redox reaction is an electronic chemical transformation wherein one chemical species undergoes oxidation while another undergoes reduction simultaneously.

Analytical CriterionOxidation ProcessReduction Process
Oxygen TransferGain of Oxygen ($\text{O}$)Loss of Oxygen ($\text{O}$)
Hydrogen TransferLoss of Hydrogen ($\text{H}$)Gain of Hydrogen ($\text{H}$)
Electron TransferLoss of electrons ($\text{LEO}$)Gain of electrons ($\text{GER}$)
Oxidation StateIncrease in positive valenceDecrease in positive valence

Roles of Chemical Agents:

  • Oxidising Agent (Oxidant): The reactant that brings about oxidation by donating oxygen, accepting hydrogen, or gaining electrons. The oxidising agent is always the substance that gets reduced.
  • Reducing Agent (Reductant): The reactant that brings about reduction by removing oxygen, donating hydrogen, or losing electrons. The reducing agent is always the substance that gets oxidised.

High-Yield Board Redox Equations:

  1. Reaction of Copper(II) Oxide with Hydrogen Gas:
    $$\text{CuO}(s) + \text{H}_2(g) \xrightarrow{\Delta} \text{Cu}(s) + \text{H}_2\text{O}(l)$$
    • $\text{CuO} \longrightarrow \text{Cu}$ (Loss of oxygen $\implies$ Reduced, $\text{CuO}$ is the Oxidising Agent).
    • $\text{H}_2 \longrightarrow \text{H}_2\text{O}$ (Gain of oxygen $\implies$ Oxidised, $\text{H}_2$ is the Reducing Agent).
  2. Reaction of Zinc Oxide with Carbon:
    $$\text{ZnO}(s) + \text{C}(s) \xrightarrow{\Delta} \text{Zn}(s) + \text{CO}(g)$$
    • $\text{ZnO}$ loses oxygen $\implies$ Reduced ($\text{ZnO}$ is the Oxidising Agent).
    • $\text{C}$ gains oxygen $\implies$ Oxidised ($\text{C}$ is the Reducing Agent).
  3. Reaction of Manganese Dioxide with Hydrochloric Acid:
    $$\text{MnO}_2(s) + 4\text{HCl}(aq) \longrightarrow \text{MnCl}_2(aq) + 2\text{H}_2\text{O}(l) + \text{Cl}_2(g)$$
    • $\text{MnO}_2 \longrightarrow \text{MnCl}_2$ (Loss of oxygen $\implies$ Reduced, $\text{MnO}_2$ is the Oxidising Agent).
    • $\text{HCl} \longrightarrow \text{Cl}_2$ (Loss of hydrogen $\implies$ Oxidised, $\text{HCl}$ is the Reducing Agent).

Real-World Effects of Everyday Oxidation Reactions

Corrosion

Corrosion is the slow, progressive deterioration and eating away of refined metals by chemical or electrochemical attack from atmospheric gases, moisture, and acids.

The Chemistry of Rusting of Iron:

Rusting is the specific corrosion of iron. It requires the simultaneous presence of both oxygen and liquid water or water vapour.
$$\text{Unbalanced: } \text{Fe} + \text{O}_2 + \text{H}_2\text{O} \longrightarrow \text{Hydrated Iron(III) Oxide}$$
$$\text{Balanced: } 4\text{Fe}(s) + 3\text{O}_2(g) + 2x\text{H}_2\text{O}(l) \longrightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s) [\text{Rust}]$$
Rust is a soft, reddish-brown, porous, non-adherent flake that does not protect the underlying metal; instead, it falls away and exposes fresh iron to continued corrosion.

Corrosion in Other Metals:

  1. Tarnishing of Silver: Silver items turn black upon exposure to air due to trace hydrogen sulphide gas forming a surface layer of Silver Sulphide:
    $$2\text{Ag}(s) + \text{H}_2\text{S}(g) \longrightarrow \text{Ag}_2\text{S}(s) [\text{Black}] + \text{H}_2(g)$$
  2. Corrosion of Copper: Copper structures develop a protective green surface layer (patina) upon long-term exposure to moist air containing carbon dioxide, consisting of Basic Copper Carbonate:
    $$2\text{Cu}(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g) + \text{O}_2(g) \longrightarrow \text{CuCO}_3 \cdot \text{Cu(OH)}_2(s) [\text{Green Layer}]$$

Methods of Preventing Corrosion:

  • Barrier Protection: Painting, oiling, greasing, or applying enamel varnishes to block air and moisture.
  • Galvanisation: Coating iron or steel with a thin protective layer of molten Zinc ($\text{Zn}$). Because zinc is more reactive than iron, it oxidises preferentially to form a protective zinc oxide layer. Even if the coating is scratched, the remaining zinc continues to sacrifice itself to protect the exposed iron.
  • Alloying: Modifying the metal’s bulk properties (e.g., combining Iron with Chromium and Nickel produces Stainless Steel, which does not rust).
  • Electroplating: Depositing an unreactive, corrosion-resistant metal (like Tin, Nickel, or Chromium) onto the surface via electrolysis.

Rancidity

Rancidity is the slow oxidation of fats and oils in food items, producing volatile, foul-smelling, and bad-tasting carboxylic acids and aldehydes.

When foods containing unsaturated fatty acids or oils are exposed to atmospheric oxygen, airborne enzymes and light catalyse the cleavage of carbon-carbon double bonds, yielding malodorous volatile compounds that make the food unpalatable and toxic.

Methods to Inhibit and Prevent Rancidity:

  1. Flushing with Inert Gas: Packaging snack foods (such as potato chips) in bags flushed with unreactive Nitrogen gas ($\text{N}_2$) to displace oxygen and prevent oxidative decay.
  2. Addition of Synthetic Antioxidants: Adding edible antioxidants like BHA (Butylated Hydroxyanisole) and BHT (Butylated Hydroxytoluene), which oxidise preferentially and neutralise free radicals before the food fats can react.
  3. Vacuum Sealing: Removing all air from food packages to eliminate oxygen contact.
  4. Refrigeration: Storing food at low temperatures slows the reaction kinetics and enzymatic activity responsible for oxidation.
  5. Airtight Storage: Storing food in airtight containers away from direct sunlight blocks both atmospheric oxygen circulation and light-induced photolytic breakdown.

[👉 Also Read: Class 10 Science Chapter 3 Metals and Non-Metals Notes]

Master Revision Cheat Sheet

Reactant SystemReaction ClassificationObservable Changes & HallmarksBalanced Chemical Equation
$\text{Mg} + \text{O}_2$Combination / OxidationDazzling white flame; white powder forms$2\text{Mg}(s) + \text{O}_2(g) \longrightarrow 2\text{MgO}(s)$
$\text{CaO} + \text{H}_2\text{O}$Combination / ExothermicHissing noise; boiling temperature rise$\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$
$\text{FeSO}_4 \xrightarrow{\Delta}$Thermal DecompositionGreen to reddish-brown; burning sulphur odour$2\text{FeSO}_4(s) \xrightarrow{\Delta} \text{Fe}_2\text{O}_3(s) + \text{SO}_2(g) + \text{SO}_3(g)$
$\text{Pb(NO}_3)_2 \xrightarrow{\Delta}$Thermal DecompositionDecrepitation; yellow residue; brown fumes$2\text{Pb(NO}_3)_2(s) \xrightarrow{\Delta} 2\text{PbO}(s) + 4\text{NO}_2(g) + \text{O}_2(g)$
$\text{H}_2\text{O} \xrightarrow{\text{electrolysis}}$Electrolytic DecompositionVolume ratio $2:1$; cathode pop, anode flame$2\text{H}_2\text{O}(l) \xrightarrow{\text{elec}} 2\text{H}_2(g) + \text{O}_2(g)$
$\text{AgCl} \xrightarrow{h\nu}$Photolytic DecompositionPure white solid turns dull grey$2\text{AgCl}(s) \xrightarrow{\text{sunlight}} 2\text{Ag}(s) + \text{Cl}_2(g)$
$\text{Fe} + \text{CuSO}_4$Single DisplacementBlue solution turns pale green; red deposit$\text{Fe}(s) + \text{CuSO}_4(aq) \longrightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$
$\text{Na}_2\text{SO}_4 + \text{BaCl}_2$Double Displacement / PrecipitationDense, insoluble white precipitate forms$\text{Na}_2\text{SO}_4(aq) + \text{BaCl}_2(aq) \longrightarrow \text{BaSO}_4(s)\downarrow + 2\text{NaCl}(aq)$
$\text{Pb(NO}_3)_2 + \text{KI}$Double Displacement / PrecipitationInstant bright yellow precipitate forms$\text{Pb(NO}_3)_2(aq) + 2\text{KI}(aq) \longrightarrow \text{PbI}_2(s)\downarrow + 2\text{KNO}_3(aq)$
$\text{CuO} + \text{H}_2 \xrightarrow{\Delta}$Redox (Reduction of oxide)Black coating turns back to shiny reddish-brown$\text{CuO}(s) + \text{H}_2(g) \xrightarrow{\Delta} \text{Cu}(s) + \text{H}_2\text{O}(l)$
$\text{Zn} + \text{H}_2\text{SO}_4$Displacement / ExothermicGas bubbles; burning pop sound; warm flask$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)\uparrow$

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