Chemical Reactions and Equations - Class 10 CBSE Notes

Introduction to Chemical Reactions and Equations

A chemical reaction is a process in which one or more substances (reactants) are transformed into new substances (products) with different chemical compositions and properties. During a chemical reaction, bonds between atoms break and new bonds form, resulting in the creation of entirely new materials.

What is a Chemical Equation?

A chemical equation is the symbolic representation of a chemical reaction using chemical formulas and symbols. It provides a concise way to describe what happens during a reaction, showing both the reactants and products.

Example:

Magnesium + Oxygen → Magnesium oxide2Mg(s) + O₂(g) → 2MgO(s)

Physical Change vs Chemical Change

Understanding the difference between physical and chemical changes is fundamental to chemistry.

Physical Changes

A physical change alters the physical properties of a substance without changing its chemical composition.

Characteristics:

• Identity of the substance remains the same
• Change can be reversed
• Heat change may or may not occur
• Only physical state or properties change

Examples:

• Freezing and melting of water
• Boiling and condensation
• Dissolving sugar in water
• Breaking glass

Chemical Changes

A chemical change transforms one or more substances into new substances with different chemical compositions.

Characteristics:

• Identity of original substance is lost
• Change cannot be easily reversed
• Generally accompanied by energy change
• Chemical composition changes
• Formation of new substances

Examples:

• Burning of wood
• Rusting of iron
• Cooking of food
• Digestion of food
• Combustion of fuel

Common Indicators and Signs of Chemical Reactions

Chemical reactions can be identified through observable changes:

1. Change in State

Physical state of substances changes during reaction.

Example: Formation of solid MgO from solid Mg and gaseous O₂

2. Change in Colour

Visual colour transformation indicates a chemical reaction.

Examples:

• Formation of brown rust on black iron nails
• Formation of yellow precipitate of lead iodide from colourless solutions

3. Evolution of Gas

Release of gas bubbles during reaction.

Examples:

• Evolution of H₂ gas when Zn reacts with dilute HCl
• Evolution of CO₂ gas during burning of carbon-containing fuels

4. Change in Temperature

Reactions are accompanied by temperature increase (exothermic) or decrease (endothermic).

Examples:

• Warming of flask when Zn reacts with H₂SO₄ (exothermic)
• Cooling when Ba(OH)₂ reacts with NH₄Cl (endothermic)

5. Formation of Precipitate

An insoluble solid forms in a solution.

Example: White precipitate of BaSO₄ when Na₂SO₄ reacts with BaCl₂

Balanced and Unbalanced Chemical Equations

Unbalanced Equation

An equation where the number of atoms of different elements on both sides are not equal.

Example:

Mg + O₂ → MgO (Unbalanced)

This is also called a skeletal equation.

Balanced Equation

An equation where the number of atoms of each element is equal on both sides.

Example:

2Mg + O₂ → 2MgO (Balanced)

Why Balance Chemical Equations?

Balancing is essential to fulfill the Law of Conservation of Mass, which states that matter can neither be created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

How to Balance Chemical Equations - Step by Step

Step 1: Write the word equation for the reaction

Step 2: Convert word equation to symbol equation using correct formulas

Step 3: Count the number of atoms of each element on both sides

Step 4: Start balancing with the most complex formula

Step 5: Use coefficients (numbers before formulas) to balance atoms

Step 6: Never change subscripts in formulas

Step 7: Verify that all atoms are balanced

Example: Balancing Mg + HCl → MgCl₂ + H₂

Solution: Place coefficient 2 before HCl

Mg + 2HCl → MgCl₂ + H₂

Now all atoms are balanced!

Related Helpful Resources

Types of Chemical Reactions with Examples

Chemical reactions are classified into several types based on how reactants transform into products:

1. Combination Reactions

Definition: Two or more substances combine to form a single new substance.

General Form: A + B → AB

Examples:

(a) Two elements combining:

2Mg(s) + O₂(g) → 2MgO(s)C(s) + O₂(g) → CO₂(g)2H₂(g) + O₂(g) → 2H₂O(l)

(b) Element + Compound:

2NO(g) + O₂(g) → 2NO₂(g)2CO(g) + O₂(g) → 2CO₂(g)

(c) Two compounds combining:

CaO(s) + H₂O(l) → Ca(OH)₂(s)Quick lime + Water → Slaked lime

Important Application: Formation of slaked lime is used in whitewashing. The Ca(OH)₂ reacts with atmospheric CO₂ to form CaCO₃ (marble), which gives walls a shiny finish.

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

2. Decomposition Reactions

Definition: A compound breaks down into two or more simpler substances.

General Form: AB → A + B

Decomposition reactions are the reverse of combination reactions and require energy in the form of heat, light, or electricity.

(a) Thermal Decomposition

Decomposition by heating.

Examples:

2FeSO₄(s) --Heat→ Fe₂O₃(s) + SO₂(g) + SO₃(g)

Ferrous sulphate (dirty white) → Ferric oxide (brown)

2Pb(NO₃)₂(s) --Heat→ 2PbO(s) + 4NO₂(g) + O₂(g)

Lead nitrate → Lead oxide + Nitrogen dioxide + Oxygen

CaCO₃(s) --Heat→ CaO(s) + CO₂(g)

Calcium carbonate → Calcium oxide (Quick lime) + Carbon dioxide

(b) Electrolytic Decomposition

Decomposition by passing electric current.

Example: Electrolysis of Water

2H₂O(l) --Electric current→ 2H₂(g) + O₂(g)

Volume of hydrogen collected is double that of oxygen, confirming the formula H₂O.

(c) Photolytic Decomposition

Decomposition in the presence of sunlight.

Examples:

2AgCl(s) --Sunlight→ 2Ag(s) + Cl₂(g)Silver chloride (white) → Silver (grey) + Chlorine2AgBr(s) --Sunlight→ 2Ag(s) + Br₂(g)Silver bromide (yellow) → Silver (grey) + Bromine

Application: This reaction forms the basis of black and white photography. Silver halides are kept in coloured bottles to prevent light exposure.

3. Displacement Reactions

Definition: A more reactive element displaces a less reactive element from its compound.

General Form: A + BC → AC + B

Activity Series of Metals

The reactivity series arranges metals in order of decreasing reactivity:

K (Potassium) ↑Na (Sodium) |Ba (Barium) |Ca (Calcium) | Most ReactiveMg (Magnesium) |Al (Aluminium) |Zn (Zinc) ↓Fe (Iron)Ni (Nickel)Sn (Tin)Pb (Lead)H (Hydrogen)Cu (Copper) ↑Hg (Mercury) |Ag (Silver) | Least ReactiveAu (Gold) ↓

Examples:

(a) Displacement of copper by iron:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)(Blue solution) → (Green solution) + (Reddish brown)

(b) Displacement of copper by zinc:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)(Blue solution) → (Colourless) + (Reddish brown)

(c) Displacement of less reactive non-metal:

2KI(aq) + Cl₂(g) → 2KCl(aq) + I₂(s)(Colourless) + Chlorine → Potassium chloride + (Violet)

Rule: A metal higher in the reactivity series can displace a metal lower in the series from its compound.

4. Double Displacement Reactions

Definition: Two compounds exchange their ions to form two new compounds.

General Form: AB + CD → AD + CB

This is also called metathesis reaction.

(a) Precipitation Reactions

Reactions where an insoluble solid (precipitate) is formed.

Examples:

Na₂SO₄(aq) + BaCl₂(aq) → 2NaCl(aq) + BaSO₄(s)↓Sodium sulphate + Barium chloride → Sodium chloride + Barium sulphate (white ppt)AgNO₃(aq) + NaCl(aq) → AgCl(s)↓ + NaNO₃(aq)Silver nitrate + Sodium chloride → Silver chloride (white ppt) + Sodium nitrate

(b) Neutralisation Reactions

Reactions between an acid and a base producing salt and water.

Examples:

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)Hydrochloric acid + Sodium hydroxide → Sodium chloride + WaterKOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l)Base + Acid → Salt + WaterBa(OH)₂(aq) + 2HCl(aq) → BaCl₂(aq) + 2H₂O(l)2NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2H₂O(l)

Application: Antacids contain bases that neutralise excess stomach acid, providing relief from acidity.

5. Oxidation and Reduction Reactions (Redox Reactions)

Classical Definitions

Oxidation:

• Addition of oxygen to a substance
• Removal of hydrogen from a substance

Reduction:

• Addition of hydrogen to a substance
• Removal of oxygen from a substance

Examples:

CuO(s) + H₂(g) → Cu(s) + H₂O(l)

Here:

• CuO is reduced (loses oxygen)
• H₂ is oxidised (gains oxygen)

ZnO(s) + C(s) → Zn(s) + CO(g)

Here:

• ZnO is reduced (loses oxygen)
• C is oxidised (gains oxygen)

Oxidising Agent: Substance that provides oxygen for oxidation or removes hydrogen. It gets reduced.

Reducing Agent: Substance that provides hydrogen for reduction or removes oxygen. It gets oxidised.

Modern Concept of Oxidation and Reduction (Electronic Concept)

Based on Electron Transfer

Oxidation: Loss of electrons (de-electronation)

Mg → Mg²⁺ + 2e⁻Na → Na⁺ + e⁻Cu → Cu²⁺ + 2e⁻

Reduction: Gain of electrons (electronation)

Cl + e⁻ → Cl⁻Zn²⁺ + 2e⁻ → ZnSn⁴⁺ + 2e⁻ → Sn²⁺

Oxidation Number (Oxidation State)

Definition: A positive or negative number representing the charge an atom appears to have when bonding electrons are counted according to specific rules.

Rules for Assigning Oxidation Numbers

1. Free elements: Oxidation number = 0
   • Examples: H₂, O₂, N₂, P₄, Ag
2. Monoatomic ions: Oxidation number = charge on ion
   • Ca²⁺ = +2, S²⁻ = -2
3. Oxygen: Usually -2 (except in peroxides where it's -1, and in OF₂ where it's +2)
4. Hydrogen: Usually +1 (except in metal hydrides where it's -1)
5. Group I metals: Always +1
6. Group II metals: Always +2
7. Halogens in binary compounds: Usually -1
8. Sum of oxidation numbers in a compound: = 0
9. Sum of oxidation numbers in a polyatomic ion: = charge on ion

Example Calculations

Find oxidation number of Mn in KMnO₄:

Let oxidation number of Mn = x

(+1) + x + 4(-2) = 0 1 + x - 8 = 0 x = +7

Find oxidation number of Cr in Cr₂O₇²⁻:

Let oxidation number of Cr = x

2x + 7(-2) = -2 2x - 14 = -2 2x = 12 x = +6

Find oxidation number of S in H₂SO₄:

2(+1) + x + 4(-2) = 0 2 + x - 8 = 0 x = +6

Redox Reactions in Terms of Oxidation Number

Oxidation: Increase in oxidation number 

Reduction: Decrease in oxidation number

Example:

2H₂ + O₂ → 2H₂O 0 0 +1 -2

• Hydrogen: 0 → +1 (oxidation)
• Oxygen: 0 → -2 (reduction)

 How to Balance Redox Equations Step by Step

Method 1: Oxidation Number Method

Steps:

1. Write the skeleton equation
2. Indicate oxidation numbers above all atoms
3. Identify elements undergoing oxidation number change
4. Calculate increase and decrease in oxidation number per atom
5. Equate total increase and decrease by multiplying with suitable coefficients
6. Balance all atoms except H and O
7. Balance O atoms (add H₂O if needed)
8. Balance H atoms (add H⁺ in acidic medium or OH⁻ in basic medium)

Example: Balance CuO + NH₃ → Cu + N₂ + H₂O

Step 1: Assign oxidation numbers

Cu⁺²O⁻² + N⁻³H₃⁺¹ → Cu⁰ + N₂⁰ + H₂⁺¹O⁻²

Step 2: Identify changes

• Cu: +2 → 0 (decrease by 2)
• N: -3 → 0 (increase by 3)

Step 3: Equate changes

3CuO + 2NH₃ → Cu + N₂ + H₂O

Step 4: Balance Cu and N

3CuO + 2NH₃ → 3Cu + N₂ + H₂O

Step 5: Balance H and O

3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O

Balanced equation!

Method 2: Ion-Electron Method (Half-Reaction Method)

Steps:

1. Write the ionic equation
2. Identify oxidation and reduction species
3. Split into two half-reactions
4. Balance atoms in each half-reaction
5. Balance charges by adding electrons
6. For acidic medium: Add H₂O to balance O, add H⁺ to balance H
7. For basic medium: Add OH⁻ and H₂O as needed
8. Multiply half-reactions to equalize electrons
9. Add half-reactions and cancel common terms

Example: Balance Cr₂O₇²⁻ + Fe²⁺ + H⁺ → Cr³⁺ + Fe³⁺ + H₂O

Step 1: Oxidation half-reaction

Fe²⁺ → Fe³⁺ + e⁻ ... (i)

Step 2: Reduction half-reaction

Cr₂O₇²⁻ → Cr³⁺

Balance Cr:

Cr₂O₇²⁻ → 2Cr³⁺

Balance O with H₂O:

Cr₂O₇²⁻ → 2Cr³⁺ + 7H₂O

Balance H with H⁺:

Cr₂O₇²⁻ + 14H⁺ → 2Cr³⁺ + 7H₂O

Balance charge with electrons:

Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O ... (ii)

Step 3: Multiply equation (i) by 6

6Fe²⁺ → 6Fe³⁺ + 6e⁻

Step 4: Add both equations

Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

Balanced equation!

Exothermic and Endothermic Reactions

Exothermic Reactions

Definition: Reactions that release heat energy.

Examples:

C(s) + O₂(g) → CO₂(g) + HeatCH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) + HeatCaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

Important: Respiration is an exothermic process:

C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + Energy

Endothermic Reactions

Definition: Reactions that absorb heat energy.

Examples:

N₂(g) + O₂(g) ⇌ 2NO(g) – HeatH₂(g) + I₂(g) ⇌ 2HI(g) – HeatC(s) + 2S(g) → CS₂(l) – Heat

Effects of Oxidation Reactions in Everyday Life

1. Respiration

Respiration is a vital biochemical oxidation process that releases energy in cells.

Process:

• Oxygen enters lungs → diffuses into blood → binds to haemoglobin
• Carried to cells throughout the body
• Glucose combustion occurs in cells

C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy

• Energy powers cellular functions, heart, muscles
• CO₂ and H₂O are expelled through breathing

2. Corrosion

Definition: Gradual deterioration of metals due to reaction with environmental substances (oxygen, water, acids).

Example: Rusting of Iron

4Fe(s) + 3O₂(g) + xH₂O → 2Fe₂O₃.xH₂O (rust)

Prevention Methods:

• Painting
• Oiling/greasing
• Galvanization (zinc coating)
• Electroplating
• Alloying (e.g., stainless steel)

Importance: Corrosion causes significant economic losses through damage to buildings, bridges, ships, and vehicles.

3. Rancidity

Definition: Oxidation of fats and oils in food, causing unpleasant taste and smell.

Prevention Methods:

• Store in airtight containers
• Refrigeration
• Adding antioxidants (e.g., vitamin E, BHA)
• Packaging in nitrogen atmosphere
• Keeping away from light

Example: Chips packets are flushed with nitrogen to prevent oxidation of oils.

4. Combustion

All fuels undergo combustion (oxidation) to release energy:

• Cooking gas (LPG - butane)
• Natural gas (methane)
• Petrol, diesel in vehicles
• Coal, wood for heating

Chemical Formulas - Quick Reference Table

Practice Problems for Class 10 Chemical Equations

Problem Set 1: Balancing Equations

Balance the following chemical equations:

1. Fe + O₂ → Fe₂O₃
2. CH₄ + O₂ → CO₂ + H₂O
3. Al + HCl → AlCl₃ + H₂
4. Ca + H₂O → Ca(OH)₂ + H₂
5. NH₃ + O₂ → NO + H₂O
6. C₆H₆ + O₂ → CO₂ + H₂O
7. SnO₂ + C → Sn + CO
8. KClO₃ → KCl + O₂
9. P₄ + O₂ → P₂O₅
10. Na + H₂O → NaOH + H₂

Problem Set 2: Identify Reaction Types

Identify the type of reaction:

1. 2KClO₃ → 2KCl + 3O₂
2. Zn + H₂SO₄ → ZnSO₄ + H₂
3. AgNO₃ + NaCl → AgCl + NaNO₃
4. 2H₂ + O₂ → 2H₂O
5. 2Mg + O₂ → 2MgO
6. CaCO₃ → CaO + CO₂
7. Fe + CuSO₄ → FeSO₄ + Cu
8. HCl + NaOH → NaCl + H₂O

Problem Set 3: Oxidation Numbers

Calculate oxidation numbers:

1. S in H₂SO₃
2. N in HNO₃
3. Cr in K₂Cr₂O₇
4. Mn in MnO₂
5. C in CO₃²⁻
6. P in H₃PO₄
7. S in Na₂S₂O₃

Problem Set 4: Redox Reactions

Identify oxidising and reducing agents:

1. CuO + H₂ → Cu + H₂O
2. MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O
3. 2Fe₂O₃ + 3C → 4Fe + 3CO₂
4. Zn + CuSO₄ → ZnSO₄ + Cu

Problem Set 5: Application-Based

1. Why is food cooked in aluminium vessels not stored in them overnight?
2. Why are silver chloride and silver bromide stored in coloured bottles?
3. What is the role of nitrogen in chips packets?
4. Why does the colour of copper sulphate solution change when an iron nail is dipped in it?
5. Explain why respiration is considered an exothermic reaction.

Answers to Practice Problems

Problem Set 1: Balancing

1. 4Fe + 3O₂ → 2Fe₂O₃
2. CH₄ + 2O₂ → CO₂ + 2H₂O
3. 2Al + 6HCl → 2AlCl₃ + 3H₂
4. Ca + 2H₂O → Ca(OH)₂ + H₂
5. 4NH₃ + 5O₂ → 4NO + 6H₂O
6. 2C₆H₆ + 15O₂ → 12CO₂ + 6H₂O
7. SnO₂ + 2C → Sn + 2CO
8. 2KClO₃ → 2KCl + 3O₂
9. P₄ + 5O₂ → 2P₂O₅
10. 2Na + 2H₂O → 2NaOH + H₂

Problem Set 2: Reaction Types

1. Decomposition
2. Displacement
3. Double displacement
4. Combination
5. Combination
6. Decomposition
7. Displacement
8. Double displacement (Neutralisation)

Problem Set 3: Oxidation Numbers

1. S = +4
2. N = +5
3. Cr = +6
4. Mn = +4
5. C = +4
6. P = +5
7. S = +2

Problem Set 4: Redox Agents

1. Oxidising agent: CuO; Reducing agent: H₂
2. Oxidising agent: MnO₂; Reducing agent: HCl
3. Oxidising agent: Fe₂O₃; Reducing agent: C
4. Oxidising agent: CuSO₄; Reducing agent: Zn

Summary

• Chemical reactions involve breaking and forming of bonds, creating new substances
• Chemical equations must be balanced to satisfy the law of conservation of mass
• Five main types of reactions: Combination, Decomposition, Displacement, Double Displacement, and Redox
• Oxidation = loss of electrons or gain of oxygen
• Reduction = gain of electrons or loss of oxygen
• Balancing redox equations requires oxidation number method or ion-electron method
• Indicators of chemical reactions: colour change, gas evolution, precipitate formation, temperature change
• Real-life applications: Respiration, combustion, corrosion, rancidity all involve chemical reactions

Conclusion

Chemical Reactions and Equations is fundamental to mastering chemistry. This chapter covers the classification of reactions, balancing techniques, redox processes, and real-world applications. Regular practice with balancing equations, identifying reaction types, and calculating oxidation numbers will strengthen your conceptual clarity and prepare you excellently for CBSE Class 10 examinations.