NCERT Solutions Class 10 Science Chapter 1: Chemical Reactions and Equations

Part 1: Chapter Introduction & Core Concepts

Welcome to the definitive guide to Chemical Reactions and Equations. In our
daily lives, matter undergoes various changes—some temporary, others permanent.
Whether it is milk spoiling at room temperature, iron rusting in a humid
atmosphere, food getting digested in our body, or the simple act of breathing,
the fundamental nature and identity of the initial substances are transformed.

Whenever a chemical change occurs, we can say that a chemical reaction has taken
place.

This chapter lays the foundational bedrock for Chemistry in Class 10. You will
learn how to systematically identify chemical reactions, represent them
concisely using Chemical Equations, and balance them to satisfy the universal
Law of Conservation of Mass. Furthermore, you will explore a myriad of reaction
types including Combination, Decomposition, Displacement, Double Displacement,
and Redox Reactions, and see how oxidation processes like Corrosion and
Rancidity impact our everyday lives.

Part 2: Comprehensive Theory & Line-by-Line Notes

Identifying a Chemical Reaction How do we know a chemical reaction has

occurred? A chemical reaction is accompanied by certain observable changes. Any
of the following observations helps us determine whether a chemical reaction has
taken place:

• Change in state
• Change in colour
• Evolution of a gas
• Change in temperature

Chemical Equations Writing out chemical reactions in sentence form is

tedious. A more concise way is to use a Word-Equation.

• A word-equation shows the change of reactants to products through an arrow
  placed between them.
• Reactants: The substances that undergo chemical change in the reaction. They
  are written on the left-hand side (LHS) with a plus sign (+) between them.
• Products: The new substances formed during the reaction. They are written on
  the right-hand side (RHS) with a plus sign (+) between them.
• The arrowhead points towards the products, showing the direction of the
  reaction.

To make it even shorter, we use chemical formulae.

• Skeletal Chemical Equation: An equation that is unbalanced because the mass
  is not the same on both sides of the equation. Example:
  Mg + O_2 \rightarrow MgO.

Balanced Chemical Equations The Law of Conservation of Mass states that mass

can neither be created nor destroyed in a chemical reaction. Therefore, the
total mass of the elements present in the products of a chemical reaction has to
be equal to the total mass of the elements present in the reactants. The number
of atoms of each element must remain the same before and after a chemical
reaction.

To make an equation highly informative, we write the Physical States of the
components:

• Solid: (s)
• Liquid: (l)
• Gas: (g)
• Aqueous (solution in water): (aq) Reaction conditions like temperature,
  pressure, and catalysts are indicated above and/or below the arrow.

Types of Chemical Reactions

1. Combination Reaction

• A reaction in which a single product is formed from two or more reactants is
  known as a combination reaction.
• Example: Quick lime (Calcium oxide) reacts vigorously with water to form
  slaked lime. CaO_{(s)} + H_2O_{(l)} \rightarrow Ca(OH)_{2(aq)} + Heat
• Exothermic Chemical Reactions: Reactions in which heat is released along
  with the formation of products. Examples include the burning of natural gas
  (CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O), respiration, and the decomposition
  of vegetable matter into compost.

1. Decomposition Reaction

• A reaction in which a single reactant breaks down to give simpler products
  is called a decomposition reaction.
• Decomposition reactions require energy in the form of heat, light, or
  electricity.
• Thermal Decomposition: When a decomposition reaction is carried out by
  heating. (e.g., heating Limestone CaCO_3).
• Endothermic Reactions: Reactions in which energy is absorbed. All
  decomposition reactions are inherently endothermic.

1. Displacement Reaction

• A reaction in which a more reactive element displaces or removes another
  element from its compound.
• Example: Iron displacing Copper from Copper Sulphate solution.
  Fe_{(s)} + CuSO_{4(aq)} \rightarrow FeSO_{4(aq)} + Cu_{(s)}

1. Double Displacement Reaction

• Reactions in which there is an exchange of ions between the reactants are
  called double displacement reactions.
• Precipitate: An insoluble substance formed during a chemical reaction.
• Precipitation Reaction: Any reaction that produces a precipitate.
  Na_2SO_{4(aq)} + BaCl_{2(aq)} \rightarrow BaSO_{4(s)} + 2NaCl_{(aq)}

1. Oxidation and Reduction (Redox Reactions)

• Oxidation: If a substance gains oxygen or loses hydrogen during a reaction,
  it is said to be oxidised.
• Reduction: If a substance loses oxygen or gains hydrogen during a reaction,
  it is said to be reduced.
• Redox Reactions: Reactions in which one reactant gets oxidised while the
  other gets reduced.

Feature	Oxidation	Reduction
Oxygen	Gained	Lost
Hydrogen	Lost	Gained

Effects of Oxidation in Everyday Life

Corrosion

• When a metal is attacked by substances around it such as moisture, acids,
  etc., it is said to corrode and this process is called corrosion.
• Examples: Rusting of iron (reddish-brown coating), the black coating on
  silver, and the green coating on copper. It causes massive damage to car
  bodies, bridges, and ships.

Rancidity

• When fats and oils are oxidised, they become rancid and their smell and
  taste change.
• Prevention: Adding antioxidants to foods, storing food in airtight
  containers, or flushing bags of chips with non-reactive gases like nitrogen.

Part 3: NCERT Activities Master Breakdown

Activity 1.1: Burning of Magnesium Ribbon

• Aim: To observe the combustion of a magnesium ribbon in air.
• Material Required: Magnesium ribbon (3-4 cm), sandpaper, tongs, spirit
  lamp/burner, watch-glass.
• Theory: Magnesium reacts with atmospheric oxygen upon heating to form
  magnesium oxide.
• Procedure:
  1. Clean the magnesium ribbon by rubbing it with sandpaper (to remove the
     protective oxide layer).
  2. Hold it with tongs and burn it over a flame.
  3. Collect the ash in a watch-glass.
• Observation: The magnesium ribbon burns with a dazzling white flame.
• Results / Conclusion: Magnesium undergoes a combination reaction with oxygen
  to form a white powder, Magnesium Oxide (MgO).
• Changes Noted:
  • Physical Change: Solid metal turns into a white powder.
  • Chemical Change: Heat and light are evolved, and a new chemical
    substance (MgO) is formed.

Activity 1.2: Reaction of Lead Nitrate and Potassium Iodide

• Aim: To observe a precipitation (double displacement) reaction.
• Material Required: Lead nitrate solution, potassium iodide solution, test
  tube.
• Theory: Exchange of ions between lead nitrate and potassium iodide yields an
  insoluble precipitate.
• Procedure:
  1. Take lead nitrate solution in a test tube.
  2. Add potassium iodide solution to it.
• Observation: A bright yellow precipitate forms immediately.
• Results / Conclusion: The yellow precipitate is Lead Iodide (PbI_2). This is
  a double displacement reaction.
• Changes Noted:
  • Chemical Change: Formation of a yellow precipitate.

Activity 1.3: Action of Acid on Zinc

• Aim: To observe the evolution of gas and change in temperature during a
  chemical reaction.
• Material Required: Zinc granules, dilute Hydrochloric acid (HCl) or
  Sulphuric acid (H_2SO_4), conical flask/test tube.
• Theory: Active metals displace hydrogen from dilute acids.
• Procedure:
  1. Place a few zinc granules in a flask.
  2. Add dilute acid.
  3. Touch the flask.
• Observation: Bubbles form around the zinc granules. The flask becomes hot to
  the touch.
• Results / Conclusion: Hydrogen gas (H_2) is evolved. The reaction is highly
  exothermic.
• Changes Noted:
  • Chemical Change: Gas evolution (H_2), increase in temperature.

Activity 1.4: Reaction of Quick Lime with Water

• Aim: To observe an exothermic combination reaction.
• Material Required: Calcium oxide (quick lime), water, beaker.
• Theory: Quick lime reacts vigorously with water to form slaked lime.
• Procedure:
  1. Take a small amount of quick lime in a beaker.
  2. Slowly add water.
  3. Touch the beaker.
• Observation: The mixture produces a hissing sound, and the beaker becomes
  extremely hot.
• Results / Conclusion: Calcium oxide combines with water to form calcium
  hydroxide (Ca(OH)_2). It is an exothermic combination reaction.
• Changes Noted:
  • Chemical Change: Huge change in temperature (heat evolved).

Activity 1.5: Thermal Decomposition of Ferrous Sulphate

• Aim: To study the decomposition of ferrous sulphate on heating.
• Material Required: 2g Ferrous sulphate crystals, dry boiling tube, burner.
• Theory: Heating breaks a single compound into multiple products.
• Procedure:
  1. Note the initial green colour of the crystals.
  2. Heat the boiling tube over a flame.
  3. Observe the colour change and smell the gas carefully.
• Observation: The green colour of the crystals changes to reddish-brown. A
  characteristic odour of burning sulphur is detected.
• Results / Conclusion: Ferrous sulphate (FeSO_4) decomposes into solid ferric
  oxide (Fe_2O_3) and toxic gases sulphur dioxide (SO_2) and sulphur trioxide
  (SO_3).
• Changes Noted:
  • Physical Change: Colour change (green to reddish-brown).
  • Chemical Change: Evolution of suffocating gases.

Activity 1.6: Thermal Decomposition of Lead Nitrate

• Aim: To study the decomposition of lead nitrate.
• Material Required: 2g Lead nitrate powder, boiling tube, pair of tongs,
  burner.
• Theory: Lead nitrate decomposes upon heating.
• Procedure:
  1. Take lead nitrate in a boiling tube.
  2. Hold with tongs and heat over a flame.
• Observation: Emission of thick brown fumes.
• Results / Conclusion: The brown fumes are Nitrogen dioxide (NO_2). The salt
  breaks down into lead oxide, nitrogen dioxide, and oxygen.
• Changes Noted:
  • Chemical Change: Evolution of brown gas.

Activity 1.7: Electrolysis of Water

• Aim: To observe the electrolytic decomposition of water.
• Material Required: Plastic mug, rubber stoppers, carbon electrodes, 6-volt
  battery, water, dilute sulphuric acid, two test tubes.
• Theory: Passing electricity through acidified water breaks it down into
  hydrogen and oxygen gases.
• Procedure:
  1. Drill holes at the base of a mug, fit stoppers, and insert carbon
     electrodes.
  2. Connect to a 6V battery.
  3. Fill the mug with water and a few drops of dilute H_2SO_4.
  4. Invert water-filled test tubes over the electrodes.
  5. Switch on the current.
• Observation: Bubbles form at both electrodes, displacing water. The volume
  of gas collected in one tube is double that in the other.
• Results / Conclusion: Water (H_2O) decomposes into H_2 and O_2. The gas with
  double volume is Hydrogen. This is an electrolytic decomposition reaction.
• Changes Noted:
  • Chemical Change: Gas evolution via electricity.

Activity 1.8: Photolytic Decomposition of Silver Chloride

• Aim: To observe the effect of sunlight on silver chloride.
• Material Required: 2g silver chloride, china dish.
• Theory: Silver halides decompose in the presence of light (used in black and
  white photography).
• Procedure:
  1. Place white silver chloride in a china dish.
  2. Leave the dish in direct sunlight for some time.
• Observation: The white powder turns grey.
• Results / Conclusion: Silver chloride decomposes into grey silver metal and
  chlorine gas due to light. This is a photolytic decomposition reaction.
• Changes Noted:
  • Physical Change: Colour change (white to grey).

Activity (In-text): Endothermic Reaction Setup

• Aim: To observe an endothermic reaction.
• Material Required: 2g Barium hydroxide, 1g Ammonium chloride, test tube,
  glass rod.
• Procedure: Mix Barium hydroxide and Ammonium chloride in a test tube using a
  glass rod. Touch the bottom of the test tube.
• Observation: The bottom of the test tube feels extremely cold.
• Results / Conclusion: The reaction absorbs heat from the surroundings,
  proving it is a highly endothermic reaction.
• Changes Noted:
  • Chemical Change: Drastic drop in temperature.

Activity 1.9: Iron Nails in Copper Sulphate

• Aim: To observe a displacement reaction.
• Material Required: Three iron nails, sandpaper, two test tubes, copper
  sulphate solution, thread.
• Theory: A more reactive metal (Fe) displaces a less reactive metal (Cu) from
  its salt solution.
• Procedure:
  1. Clean iron nails with sandpaper.
  2. Pour 10 mL copper sulphate (blue) into test tubes A and B.
  3. Tie two iron nails and immerse them in test tube B for 20 minutes.
  4. Remove the nails and compare colours.
• Observation: The iron nails become coated with a brownish layer. The vibrant
  blue colour of the copper sulphate solution fades to a light green.
• Results / Conclusion: Iron displaces copper to form iron sulphate (light
  green). Copper metal (brown) is deposited on the nail.
• Changes Noted:
  • Chemical Change: Colour change in solution (blue to green) and on the
    solid (grey to brown).

Activity 1.10: Double Displacement / Precipitation

• Aim: To observe a double displacement reaction.
• Material Required: 3 mL sodium sulphate solution, 3 mL barium chloride
  solution, two test tubes.
• Theory: Mutual exchange of ions results in the formation of an insoluble
  precipitate.
• Procedure: Mix the sodium sulphate solution into the barium chloride
  solution.
• Observation: A white, insoluble substance forms instantly.
• Results / Conclusion: The white precipitate is Barium Sulphate (BaSO_4).
  Sodium chloride remains in the aqueous solution.
• Changes Noted:
  • Chemical Change: Formation of a white precipitate.

Activity 1.11: Oxidation of Copper

• Aim: To study the oxidation of copper powder.
• Material Required: 1g copper powder, china dish, burner.
• Theory: Copper reacts with atmospheric oxygen upon heating to form copper
  oxide.
• Procedure: Heat the china dish containing copper powder.
• Observation: The surface of the brown copper powder gets coated with a black
  substance.
• Results / Conclusion: Copper is oxidised to form black Copper(II) Oxide
  (CuO).
• Changes Noted:
  • Chemical Change: Colour change (brown to black) due to the addition of
    oxygen.

Group Activity: Exothermic vs Endothermic Profiling

• Aim: To classify reactions by measuring temperature changes.
• Procedure:
  1. Take 25 mL water in beakers A, B, C, and copper sulphate in D. Record
     initial temperatures.
  2. Add Potassium sulphate to A, Ammonium nitrate to B, Anhydrous copper
     sulphate to C, and fine iron filings to D.
  3. Stir and record final temperatures.
• Results / Conclusion:
  • Beaker A & B (Potassium sulphate & Ammonium nitrate dissolving):
    Temperature drops -> Endothermic.
  • Beaker C (Anhydrous CuSO_4 hydrating): Temperature rises -> Exothermic.
  • Beaker D (Fe + CuSO_4 displacement): Temperature rises -> Exothermic.

Part 4: Physics Derivations & Numerical Format (Balancing Equations)

Although there are no pure physics numericals in this chapter, chemistry
requires mathematical rigor to balance equations using the Hit-and-Trial Method.

Example Problem: Balance the skeletal equation for iron reacting with steam.
Given: Fe + H_2O \rightarrow Fe_3O_4 + H_2 To Find: The balanced chemical
equation. Formula/Method Used: Hit-and-Trial Method (Law of Conservation of
Mass).

Step-by-Step Solution:

1. Draw Boxes: Draw boxes around each formula. (Do not change anything inside
   the boxes).
2. List Atoms:
   • LHS: Fe = 1, H = 2, O = 1
   • RHS: Fe = 3, H = 2, O = 4
3. Balance Oxygen first (Compound with max atoms is Fe_3O_4):
   • LHS has 1 Oxygen, RHS has 4.
   • Multiply H_2O by 4.
   • Partly Balanced: Fe + 4H_2O \rightarrow Fe_3O_4 + H_2
4. Balance Hydrogen:
   • LHS now has 8 Hydrogen atoms (4 \times 2). RHS has 2.
   • Multiply H_2 on RHS by 4.
   • Partly Balanced: Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2
5. Balance Iron (Fe):
   • LHS has 1 Iron. RHS has 3.
   • Multiply Fe on LHS by 3.
   • Fully Balanced Equation: 3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2
6. Add Physical States:
   3Fe_{(s)} + 4H_2O_{(g)} \rightarrow Fe_3O_{4(s)} + 4H_{2(g)} (Note: (g) is
   used for water to indicate it is used in the form of steam).

Part 5: Biology & Core Diagrams Placeholder

To ensure full marks in the board exams, you must be able to draw and label the
experimental setups described in this chapter perfectly.

[🛑 DIAGRAM REQUIRED HERE: Insert detailed, labeled diagram of Burning of a
magnesium ribbon in air and collection of magnesium oxide in a watch-glass
(Fig 1.1) 🛑]

[🛑 DIAGRAM REQUIRED HERE: Insert detailed, labeled diagram of Formation of
hydrogen gas by the action of dilute sulphuric acid on zinc (Fig 1.2) 🛑]

[🛑 DIAGRAM REQUIRED HERE: Insert detailed, labeled diagram of Electrolysis of
water with 6V battery, graphite electrodes, and inverted test tubes (Fig 1.6) 🛑]

[🛑 DIAGRAM REQUIRED HERE: Insert detailed, labeled diagram of Iron nails dipped
in copper sulphate solution comparing colours before and after (Fig 1.8 a & b)
🛑]

Part 6: Topper’s Quick Revision & Formulas

Memorize these exact balanced chemical equations as they are repeatedly asked in
the CBSE Board Exams:

1. Combination Reactions:

• Burning of Magnesium: 2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)}
• Formation of Slaked Lime:
  CaO_{(s)} + H_2O_{(l)} \rightarrow Ca(OH)_{2(aq)} + Heat
• Whitewashing (Shiny Finish):
  Ca(OH){2(aq)} + CO{2(g)} \rightarrow CaCO_{3(s)} + H_2O_{(l)}
• Respiration (Exothermic):
  C_6H_{12}O_{6(aq)} + 6O_{2(aq)} \rightarrow 6CO_{2(aq)} + 6H_2O_{(l)} + Energy

1. Decomposition Reactions:

• Thermal (Ferrous Sulphate):
  2FeSO_{4(s)} \xrightarrow{Heat} Fe_2O_{3(s)} + SO_{2(g)} + SO_{3(g)}
• Thermal (Limestone): CaCO_{3(s)} \xrightarrow{Heat} CaO_{(s)} + CO_{2(g)}
• Thermal (Lead Nitrate):
  2Pb(NO_3){2(s)} \xrightarrow{Heat} 2PbO{(s)} + 4NO_{2(g)} + O_{2(g)}
• Photolytic (Black & White Photography):
  2AgCl_{(s)} \xrightarrow{Sunlight} 2Ag_{(s)} + Cl_{2(g)}
  2AgBr_{(s)} \xrightarrow{Sunlight} 2Ag_{(s)} + Br_{2(g)}

1. Displacement Reactions:

• Iron & Copper Sulphate:
  Fe_{(s)} + CuSO_{4(aq)} \rightarrow FeSO_{4(aq)} + Cu_{(s)}
• Zinc & Copper Sulphate:
  Zn_{(s)} + CuSO_{4(aq)} \rightarrow ZnSO_{4(aq)} + Cu_{(s)}

1. Double Displacement (Precipitation):

• Na_2SO_{4(aq)} + BaCl_{2(aq)} \rightarrow BaSO_{4(s)}\downarrow + 2NaCl_{(aq)}
  (Note: Downward arrow denotes a precipitate).

1. Redox Reactions:

• Oxidation of Copper: 2Cu + O_2 \xrightarrow{Heat} 2CuO
• Reduction of Copper Oxide: CuO + H_2 \xrightarrow{Heat} Cu + H_2O
• Manganese Dioxide & HCl: MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2
  (Here, HCl is oxidised to Cl_2, and MnO_2 is reduced to MnCl_2).

Pro Tip for Boards: Always write physical states (s, l, g, aq) in your final
answers. When defining a reaction type, always provide one balanced chemical
equation as an example, even if the question does not explicitly ask for it!