Study Notes: Chemical Reactions and Equations (Class 10)

Overview

Chemical reactions form the foundation of all chemistry topics in competitive exams. This topic bridges the gap between basic chemistry concepts and advanced problem-solving required in NSO. Students must master three core skills: identifying reaction types instantly, balancing equations using systematic methods, and recognizing oxidation-reduction processes.

NSO questions on this topic test both conceptual understanding and application skills. You'll encounter questions asking you to classify reactions, complete and balance equations, identify oxidizing and reducing agents, and predict products. The Achievers section often combines this topic with real-world applications like rusting, photosynthesis, and industrial processes.

Mastery requires memorizing common reaction patterns, practicing balancing techniques until they become automatic, and understanding electron transfer in redox reactions. This topic accounts for 8–12% of the Science section and frequently appears in multi-step problem-solving questions.

Key Concepts

• **Chemical Reaction Definition**: A process where reactants transform into products with different chemical properties, indicated by color change, temperature change, gas evolution, precipitate formation, or state change.

• **Chemical Equation**: A symbolic representation showing reactants on the left and products on the right, separated by an arrow. Must be balanced to follow the Law of Conservation of Mass (atoms are neither created nor destroyed).

• **Combination Reaction**: Two or more substances combine to form a single product (A + B → AB). Usually exothermic, releasing heat and light.

• **Decomposition Reaction**: A single compound breaks down into two or more simpler substances (AB → A + B). Requires energy input — thermal, electrical, or light decomposition.

• **Displacement Reaction**: A more reactive element displaces a less reactive element from its compound (A + BC → AC + B). Reactivity series determines which element displaces which.

• **Double Displacement Reaction**: Exchange of ions between two compounds (AB + CD → AD + CB). Often produces precipitate, gas, or water.

• **Oxidation-Reduction (Redox) Reaction**: Simultaneous oxidation (loss of electrons or gain of oxygen) and reduction (gain of electrons or loss of oxygen). The substance that gets oxidized is the reducing agent; the substance that gets reduced is the oxidizing agent.

• **Exothermic and Endothermic Reactions**: Exothermic reactions release energy (combustion, respiration); endothermic reactions absorb energy (photosynthesis, decomposition of calcium carbonate).

Formulas / Key Facts

**Balancing Equations Rule**: Total atoms of each element must be equal on both sides. Balance in order: metals → non-metals → hydrogen → oxygen.

**Law of Conservation of Mass**: Total mass of reactants = Total mass of products. No atoms are created or destroyed during a chemical reaction.

**Reactivity Series (Decreasing order)**: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au. A metal higher in the series displaces metals below it from their salt solutions.

**Common Combination Reactions**: 2Mg + O₂ → 2MgO; C + O₂ → CO₂; CaO + H₂O → Ca(OH)₂

**Common Decomposition Reactions**: 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃ (thermal); 2H₂O → 2H₂ + O₂ (electrolysis); 2AgBr → 2Ag + Br₂ (photolysis)

**Oxidation Definitions**: (1) Addition of oxygen, (2) Removal of hydrogen, (3) Loss of electrons. Example: 2Mg + O₂ → 2MgO (Mg oxidized)

**Reduction Definitions**: (1) Removal of oxygen, (2) Addition of hydrogen, (3) Gain of electrons. Example: CuO + H₂ → Cu + H₂O (CuO reduced)

**Rancidity**: Oxidation of fats and oils when exposed to air, causing unpleasant smell and taste. Prevented by antioxidants, refrigeration, or nitrogen packing.

Worked Examples

**Example 1: Balancing a Chemical Equation**

Balance: Fe + H₂O → Fe₃O₄ + H₂

Step 1: Count atoms — Fe (1 left, 3 right), H (2 left, 2 right), O (1 left, 4 right) Step 2: Balance Fe by putting 3 before Fe on left: 3Fe + H₂O → Fe₃O₄ + H₂ Step 3: Balance O by putting 4 before H₂O: 3Fe + 4H₂O → Fe₃O₄ + H₂ Step 4: Balance H by putting 4 before H₂: 3Fe + 4H₂O → Fe₃O₄ + 4H₂ Step 5: Verify — Fe (3=3), H (8=8), O (4=4) ✓

**Balanced equation**: 3Fe + 4H₂O → Fe₃O₄ + 4H₂

**Example 2: Identifying Reaction Type and Redox Process**

Classify and identify oxidation-reduction: CuO + H₂ → Cu + H₂O

Classification: This is a displacement reaction (hydrogen displaces copper from copper oxide).

Oxidation-Reduction Analysis:

• CuO loses oxygen → Reduction of copper (Cu²⁺ gains electrons to become Cu⁰)
• H₂ gains oxygen → Oxidation of hydrogen (H⁰ loses electrons to become H⁺)
• CuO is the oxidizing agent (causes oxidation of H₂)
• H₂ is the reducing agent (causes reduction of CuO)

**Example 3: Predicting Products**

Predict products and balance: Zn + CuSO₄ → ?

Step 1: Zn is more reactive than Cu (reactivity series), so displacement occurs Step 2: Zn will displace Cu: Zn + CuSO₄ → ZnSO₄ + Cu Step 3: Check balance — Already balanced (1 Zn, 1 Cu, 1 SO₄ on each side) Step 4: Reaction type — Displacement reaction

**Balanced equation**: Zn + CuSO₄ → ZnSO₄ + Cu

Common Mistakes

**Mistake**: Balancing by changing subscripts in chemical formulas (H₂O → H₂O₂). **Fix**: Only change coefficients in front of formulas, never alter subscripts within a formula. Changing subscripts changes the compound itself.

**Mistake**: Confusing oxidation and reduction — thinking oxidation means gain of oxygen only in all contexts. **Fix**: Remember all three definitions. In ionic equations, focus on electron transfer: oxidation = loss of electrons (OIL), reduction = gain of electrons (RIG).

**Mistake**: Forgetting to balance diatomic elements (H₂, O₂, N₂, Cl₂) when they appear as free elements. **Fix**: Always write these seven elements as diatomic molecules when they're not part of a compound.

**Mistake**: Identifying only the oxidized/reduced substance without naming the oxidizing/reducing agent. **Fix**: The substance oxidized IS the reducing agent; the substance reduced IS the oxidizing agent. They're not the same thing.

**Mistake**: Assuming all combination reactions are exothermic. **Fix**: Most are exothermic, but N₂ + O₂ → 2NO is endothermic. Read the question carefully for energy indicators.

Quick Reference

→ **Balancing priority**: Metals first, then non-metals, then H, finally O

→ **Redox memory**: LEO says GER (Loss of Electrons = Oxidation; Gain of Electrons = Reduction)

→ **Reaction type clues**: One product = combination; one reactant = decomposition; element + compound = displacement

→ **Precipitate symbol**: ↓ after formula; Gas symbol: ↑ after formula

→ **Reactivity check**: K–Na–Ca–Mg–Al–Zn–Fe–Pb–H–Cu–Hg–Ag–Au (decreasing)

→ **Oxidizing agent gets reduced; Reducing agent gets oxidized** — they're opposite roles in the same reaction