Grade 10 Chapter 3: Chemical Reactions and Equations

Introduction

Chemical reactions are processes in which substances transform into new substances with different properties. These transformations involve the breaking and formation of chemical bonds, resulting in the rearrangement of atoms. Chemical equations are symbolic representations of these reactions, showing the reactants, products, and their relative quantities. This chapter explores the fundamentals of chemical reactions, their types, and how to balance chemical equations.

Chemical Equations

Writing Chemical Equations

A chemical equation is a shorthand representation of a chemical reaction using chemical symbols and formulas. It shows:

1. Reactants: The starting substances, written on the left side of the equation
2. Products: The substances formed, written on the right side of the equation
3. Arrow (→): Indicates the direction of the reaction
4. Coefficients: Numbers placed before formulas to indicate the relative quantities
5. State symbols: (s) for solid, (l) for liquid, (g) for gas, (aq) for aqueous solution

Example: When magnesium burns in oxygen, it forms magnesium oxide. Word equation: Magnesium + Oxygen → Magnesium Oxide Chemical equation: 2Mg(s) + O₂(g) → 2MgO(s)

Balancing Chemical Equations

According to the Law of Conservation of Mass, matter cannot be created or destroyed in a chemical reaction. Therefore, the number of atoms of each element must be the same on both sides of a chemical equation.

Steps to balance a chemical equation:

1. Write the correct formulas for all reactants and products
2. Count the number of atoms of each element on both sides
3. Balance the equation by adding appropriate coefficients
4. Verify that the number of atoms of each element is the same on both sides

Example: Balancing the combustion of propane (C₃H₈)

Unbalanced: C₃H₈ + O₂ → CO₂ + H₂O

Step 1: Count the atoms - Reactants: 3 C, 8 H, 2 O - Products: 1 C, 2 H, 3 O

Step 2: Balance carbon atoms C₃H₈ + O₂ → 3CO₂ + H₂O - Reactants: 3 C, 8 H, 2 O - Products: 3 C, 2 H, 7 O

Step 3: Balance hydrogen atoms C₃H₈ + O₂ → 3CO₂ + 4H₂O - Reactants: 3 C, 8 H, 2 O - Products: 3 C, 8 H, 11 O

Step 4: Balance oxygen atoms C₃H₈ + 5O₂ → 3CO₂ + 4H₂O - Reactants: 3 C, 8 H, 10 O - Products: 3 C, 8 H, 10 O

The equation is now balanced.

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the nature of the change occurring:

1. Combination Reactions

In combination reactions, two or more substances combine to form a single product.

General form: A + B → AB

Examples: - 2Mg(s) + O₂(g) → 2MgO(s) - CaO(s) + H₂O(l) → Ca(OH)₂(aq) - C(s) + O₂(g) → CO₂(g)

2. Decomposition Reactions

In decomposition reactions, a single compound breaks down into two or more simpler substances.

General form: AB → A + B

Examples: - 2H₂O₂(l) → 2H₂O(l) + O₂(g) - CaCO₃(s) → CaO(s) + CO₂(g) - 2KClO₃(s) → 2KCl(s) + 3O₂(g)

Decomposition reactions often require energy in the form of heat, light, or electricity: - Thermal decomposition: Decomposition by heat - Photolysis: Decomposition by light - Electrolysis: Decomposition by electricity

3. Displacement Reactions

In displacement reactions, a more reactive element displaces a less reactive element from its compound.

Single Displacement Reactions

General form: A + BC → AC + B

Examples: - Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) - Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) - Cl₂(g) + 2KBr(aq) → 2KCl(aq) + Br₂(l)

The reactivity series of metals helps predict whether a displacement reaction will occur: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

A metal can displace any metal below it in the series from its salt solution.

Double Displacement Reactions

In double displacement reactions, ions of two compounds exchange places to form two new compounds.

General form: AB + CD → AD + CB

Examples: - AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq) - BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq) - HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Double displacement reactions often result in the formation of a precipitate, gas, or water.

4. Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between reactants.

• Oxidation: Loss of electrons, increase in oxidation number
• Reduction: Gain of electrons, decrease in oxidation number

Examples: - 2Mg(s) + O₂(g) → 2MgO(s) - Mg is oxidized (loses electrons): Mg → Mg²⁺ + 2e⁻ - O is reduced (gains electrons): O₂ + 4e⁻ → 2O²⁻

• CuO(s) + H₂(g) → Cu(s) + H₂O(l)
• Cu is reduced: Cu²⁺ + 2e⁻ → Cu
• H is oxidized: H₂ → 2H⁺ + 2e⁻

Oxidizing and Reducing Agents

• Oxidizing agent: Substance that causes oxidation (accepts electrons)
• Reducing agent: Substance that causes reduction (donates electrons)

In the reaction: 2Mg + O₂ → 2MgO - O₂ is the oxidizing agent (accepts electrons from Mg) - Mg is the reducing agent (donates electrons to O₂)

5. Neutralization Reactions

Neutralization reactions occur between acids and bases to form salt and water.

General form: Acid + Base → Salt + Water

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

6. Combustion Reactions

Combustion reactions involve the rapid oxidation of a substance, usually with oxygen, producing heat and light.

General form: Fuel + Oxygen → Carbon dioxide + Water + Energy

Example: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + Energy

Effects of Chemical Reactions

Physical Changes

Chemical reactions often produce observable physical changes:

1. Change in state: Solid to liquid, liquid to gas, etc.
2. Change in color: For example, copper sulfate solution turns blue when copper is present
3. Formation of precipitate: Insoluble solid formed in a solution
4. Evolution of gas: Bubbles or odor indicating gas formation
5. Change in temperature: Heat released (exothermic) or absorbed (endothermic)

Energy Changes

Chemical reactions involve energy changes:

1. Exothermic reactions: Release energy to the surroundings
2. Example: Combustion, neutralization, respiration

3. Energy diagram shows products at lower energy than reactants

4. Endothermic reactions: Absorb energy from the surroundings

5. Example: Photosynthesis, thermal decomposition
6. Energy diagram shows products at higher energy than reactants

Factors Affecting Chemical Reactions

Several factors influence the rate of chemical reactions:

1. Nature of reactants: Different substances have different reaction rates
2. Concentration: Higher concentration generally leads to faster reactions
3. Temperature: Higher temperature usually increases reaction rate
4. Surface area: Larger surface area leads to faster reactions
5. Catalysts: Substances that increase reaction rate without being consumed
6. Inhibitors: Substances that decrease reaction rate

Practical Applications of Chemical Reactions

Chemical reactions have numerous applications in our daily lives:

1. Industrial processes: Manufacturing of chemicals, fertilizers, plastics
2. Food preparation: Cooking, fermentation, baking
3. Medicine: Drug synthesis, diagnostic tests
4. Agriculture: Fertilizers, pesticides
5. Energy production: Combustion of fuels, batteries
6. Environmental processes: Waste treatment, pollution control

Chemical Reactions in Everyday Life

Corrosion

Corrosion is the gradual destruction of materials (usually metals) by chemical reactions with their environment. The most common example is the rusting of iron:

4Fe(s) + 3O₂(g) + 2H₂O(l) → 2Fe₂O₃·H₂O(s) (rust)

Methods to prevent corrosion: - Painting - Oiling or greasing - Galvanization (coating with zinc) - Alloying - Cathodic protection

Rancidity

Rancidity is the oxidation of fats and oils when exposed to air, light, moisture, or bacteria, resulting in unpleasant taste and odor.

Methods to prevent rancidity: - Antioxidants - Refrigeration - Vacuum packaging - Nitrogen flushing

Balancing Redox Reactions

Balancing redox reactions involves identifying the oxidation numbers and ensuring that electrons lost equal electrons gained.

Oxidation Number Method

1. Assign oxidation numbers to all atoms
2. Identify the atoms that change oxidation numbers
3. Calculate the change in oxidation number for each atom
4. Balance the electron transfer
5. Balance the rest of the equation

Example: Balancing Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ in acidic solution

Step 1: Assign oxidation numbers - Fe²⁺: +2 - Cr in Cr₂O₇²⁻: +6 - O in Cr₂O₇²⁻: -2 - Fe³⁺: +3 - Cr³⁺: +3

Step 2: Identify changes - Fe: +2 → +3 (oxidation, loss of 1 electron) - Cr: +6 → +3 (reduction, gain of 3 electrons)

Step 3: Balance electron transfer - 6Fe²⁺ → 6Fe³⁺ + 6e⁻ - Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

Step 4: Combine half-reactions 6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O

Summary

Chemical reactions involve the transformation of substances through the breaking and formation of chemical bonds. These reactions can be represented by balanced chemical equations that follow the Law of Conservation of Mass. Chemical reactions are classified into various types, including combination, decomposition, displacement, redox, neutralization, and combustion reactions. Understanding these reactions and their characteristics is essential for explaining natural phenomena and developing industrial processes.

Practice Questions

1. Balance the following chemical equations: a) Fe + O₂ → Fe₂O₃ b) C₃H₈ + O₂ → CO₂ + H₂O c) Al + HCl → AlCl₃ + H₂

2. Identify the type of reaction in each of the following: a) 2Na + Cl₂ → 2NaCl b) CaCO₃ → CaO + CO₂ c) Zn + CuSO₄ → ZnSO₄ + Cu d) HCl + NaOH → NaCl + H₂O

3. Explain the difference between exothermic and endothermic reactions with examples.

4. What is a redox reaction? Identify the oxidizing and reducing agents in the reaction: Zn + CuSO₄ → ZnSO₄ + Cu

5. Describe three methods to prevent the corrosion of iron.

6. What factors affect the rate of a chemical reaction? Explain how each factor influences the reaction rate.

7. Write balanced chemical equations for the following reactions: a) Combustion of methane (CH₄) b) Reaction of sodium hydroxide with hydrochloric acid c) Thermal decomposition of calcium carbonate

8. Explain the concept of a catalyst and give an example of a catalyzed reaction.

9. What is rancidity? How can it be prevented?

10. Balance the following redox reaction in acidic solution: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺

References: 1. Maharashtra State Board 10th Standard Science Syllabus 2025-26 2. NCERT Science Textbook for Class 10 3. Royal Society of Chemistry - Chemical Reactions and Equations 4. American Chemical Society - Types of Chemical Reactions