Important Questions on Atoms and Molecules for Class 10 Exam Preparation

Important Questions on Atoms and Molecules for Class 10 Exam Preparation

Preparing for your Class 10 Science exam? The chapter on Atoms and Molecules is a fundamental part of Chemistry that forms the basis for understanding matter, chemical reactions, and stoichiometry. To help you ace this topic, we’ve compiled a structured guide with key questions, detailed explanations, and exam-focused tips to ensure you grasp every concept thoroughly.

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This blog post is divided into five main sections, each covering a critical aspect of the chapter. Whether you’re revising laws of chemical combination, mole concepts, or numerical problems, this guide will equip you with the knowledge and confidence to tackle any question in your exam.

Understanding the Basics: Atoms and Molecules

Before diving into complex problems, it’s essential to build a strong foundation. This section covers the fundamental definitions, differences, and real-world examples of atoms and molecules.

What Are Atoms and Molecules?

• Atom: The smallest unit of an element that retains its chemical properties. For example, a single oxygen atom (O) or a carbon atom (C).
• Molecule: A group of two or more atoms chemically bonded together. For instance:

– O₂ (Oxygen gas) – A diatomic molecule of two oxygen atoms.
– H₂O (Water) – A triatomic molecule of two hydrogen atoms and one oxygen atom.
– Key Difference:
– Atoms are single particles of elements.
– Molecules are combinations of atoms (same or different elements).

Exam Tip: Memorize examples of monoatomic (He, Ne), diatomic (O₂, N₂), and polyatomic (H₂O, CO₂) molecules.

Laws of Chemical Combination

This chapter introduces two fundamental laws that govern chemical reactions. Understanding these is crucial for solving numerical problems.

1. Law of Conservation of Mass (Lavoisier, 1774)
– Statement: Mass is neither created nor destroyed in a chemical reaction.
– Example:
– Reaction: 2H₂ + O₂ → 2H₂O
– Mass Before: 4g (H₂) + 32g (O₂) = 36g
– Mass After: 36g (H₂O)
– Conclusion: Mass remains the same.

2. Law of Constant Proportions (Proust, 1799)
– Statement: A chemical compound always contains the same elements combined in a fixed ratio by mass.
– Example:
– Water (H₂O) always has H:O = 1:8 by mass, whether from a river or a lab.

Actionable Insight: Practice problems where you balance equations and verify these laws. For example:

• Question: 12g of carbon reacts with 32g of oxygen to form CO₂. What is the mass of CO₂ formed?
• Solution: Apply the Law of Conservation of Mass → 12g + 32g = 44g CO₂.

Dalton’s Atomic Theory

Dalton’s theory explains the behavior of atoms and molecules. Key postulates include:

1. Atoms are indivisible (later disproved by subatomic particles).
2. Atoms of the same element are identical in mass and properties.
3. Compounds are formed by the combination of atoms in fixed ratios.
4. Chemical reactions involve rearrangement of atoms (no creation/destruction).

Exam Focus: Questions may ask you to criticize Dalton’s theory (e.g., “Why is the first postulate incorrect?” → Answer: Discovery of electrons, protons, and neutrons).

Atomic Mass and Molecular Mass: Key Concepts

This section breaks down how to calculate atomic mass, molecular mass, and formula unit mass, along with their applications in real-world problems.

What Is Atomic Mass?

• Definition: The mass of one atom of an element compared to 1/12th the mass of a carbon-12 atom.
• Unit: Atomic Mass Unit (u).
• Examples:

– Hydrogen (H): 1 u
– Oxygen (O): 16 u
– Carbon (C): 12 u

Step-by-Step Calculation:

1. Check the periodic table for the atomic mass of each element.
2. For isotopes, use the average atomic mass (weighted by abundance).

– Example: Chlorine has two isotopes (³⁵Cl and ³⁷Cl) with abundances of 75% and 25%. Its atomic mass = (35 × 0.75) + (37 × 0.25) = 35.5 u.

Molecular Mass and Formula Unit Mass

– Molecular Mass: Sum of atomic masses of all atoms in a molecule.
– Example: CO₂ = 12 (C) + 2 × 16 (O) = 44 u.
– Formula Unit Mass: Used for ionic compounds (e.g., NaCl, CaCO₃).
– Example: NaCl = 23 (Na) + 35.5 (Cl) = 58.5 u.

Practice Problem:

• Question: Calculate the molecular mass of glucose (C₆H₁₂O₆).
• Solution:

– C₆ = 6 × 12 = 72 u
– H₁₂ = 12 × 1 = 12 u
– O₆ = 6 × 16 = 96 u
– Total = 72 + 12 + 96 = 180 u

Gram Atomic and Gram Molecular Mass

– Gram Atomic Mass (GAM): Atomic mass expressed in grams.
– Example: 12 g of carbon = 1 gram atomic mass of carbon.
– Gram Molecular Mass (GMM): Molecular mass expressed in grams.
– Example: 18 g of water (H₂O) = 1 gram molecular mass of water.

Exam Tip: Questions may ask:
– “How many grams are in 2 moles of CO₂?”
– Solution: 1 mole CO₂ = 44 g → 2 moles = 88 g.

The Mole Concept: Bridging Atoms and Grams

The mole concept is one of the most important (and often confusing) topics in this chapter. This section simplifies it with clear definitions, formulas, and problem-solving strategies.

What Is a Mole?

• Definition: A mole is the amount of substance that contains 6.022 × 10²³ particles (Avogadro’s number).
• Analogy: Just like 1 dozen = 12 items, 1 mole = 6.022 × 10²³ atoms/molecules.
• Example:

– 1 mole of oxygen (O₂) = 6.022 × 10²³ O₂ molecules = 32 g (since molecular mass of O₂ = 32 u).

Molar Mass and Its Calculations

– Molar Mass: Mass of 1 mole of a substance (numerically equal to atomic/molecular mass but in grams).
– Example:
– 1 mole of carbon (C) = 12 g.
– 1 mole of water (H₂O) = 18 g.

Key Formula:
[
text{Number of moles (n)} = frac{text{Given mass (g)}}{text{Molar mass (g/mol)}}
]

Problem-Solving Steps:

1. Find the molar mass of the substance.
2. Divide the given mass by the molar mass to get moles.
3. Multiply by Avogadro’s number if particles are asked.

Example Problem:

• Question: How many molecules are in 9 g of water (H₂O)?
• Solution:

1. Molar mass of H₂O = 18 g/mol.
2. Moles of H₂O = 9 g / 18 g/mol = 0.5 moles.
3. Molecules = 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules.

Mole Concept in Chemical Reactions

The mole concept helps balance equations and calculate reactant/product quantities.

Example:

• Reaction: N₂ + 3H₂ → 2NH₃
• Question: How many moles of NH₃ are produced from 2 moles of N₂?
• Solution:

– From the equation, 1 mole N₂ → 2 moles NH₃.
– So, 2 moles N₂ → 4 moles NH₃.

Actionable Insight: Always balance the equation first before applying the mole concept.

Numerical Problems: Step-by-Step Solutions

This section provides solved numerical problems covering all major concepts in the chapter. Practice these to master calculations.

Problem 1: Calculating Moles from Mass

Question: How many moles are in 54 g of water (H₂O)?
Solution:

1. Molar mass of H₂O = 18 g/mol.
2. Moles = Given mass / Molar mass = 54 g / 18 g/mol = 3 moles.

Problem 2: Finding Mass from Moles

Question: What is the mass of 0.5 moles of CO₂?
Solution:

1. Molar mass of CO₂ = 44 g/mol.
2. Mass = Moles × Molar mass = 0.5 × 44 = 22 g.

Problem 3: Using Avogadro’s Number

Question: How many atoms are in 11 g of CO₂?
Solution:

1. Molar mass of CO₂ = 44 g/mol.
2. Moles of CO₂ = 11 g / 44 g/mol = 0.25 moles.
3. Molecules of CO₂ = 0.25 × 6.022 × 10²³ = 1.5055 × 10²³ molecules.
4. Since CO₂ has 3 atoms per molecule, total atoms = 3 × 1.5055 × 10²³ = 4.5165 × 10²³ atoms.

Exam Tip: Break down multi-step problems into smaller, manageable parts.

Common Mistakes and Exam Strategies

Even the best students make mistakes. This section highlights frequent errors and proven exam strategies to maximize your score.

Mistake 1: Confusing Atomic Mass and Molar Mass

• Error: Using atomic mass (u) instead of molar mass (g/mol) in calculations.
• Fix:

– Atomic mass = Mass of 1 atom (in u).
– Molar mass = Mass of 1 mole (in grams).

Mistake 2: Ignoring Balanced Equations

• Error: Calculating moles without balancing the chemical equation.
• Example:

– Incorrect: H₂ + O₂ → H₂O (unbalanced).
– Correct: 2H₂ + O₂ → 2H₂O (balanced).
– Fix: Always balance the equation first before solving.

Exam Strategy: Time Management and Question Selection

1. Start with definitions (e.g., atom, molecule, mole) to build confidence.
2. Tackle numericals next (they carry more marks).
3. Leave long-answer questions (e.g., laws of chemical combination) for the end.
4. Double-check units (g, moles, u) to avoid silly mistakes.

Bonus Tip: Use formula sheets during revision to memorize key equations:

• Moles = Mass / Molar mass
• Particles = Moles × Avogadro’s number
• Mass = Moles × Molar mass