🧪 Atoms and Molecules
Complete Class 9 Science Notes | Chapter 3 | Deep Learning Guide
📚 Welcome to Your Complete Study Guide!
You've come to the right place to master Atoms and Molecules. This comprehensive guide covers everything from basic concepts to advanced numerical problems, designed specifically for Class 9 students following the NCERT curriculum.
🌍 1. Introduction – The Story of the Smallest Things
Look around you right now. The air you're breathing, the water you drink, your mobile phone, the book in your hands, and even your own body – everything you see, touch, or feel is made up of something incredibly tiny and invisible to the naked eye called atoms.
Imagine taking a piece of chocolate and cutting it into smaller and smaller pieces. You keep dividing it until you reach a point where you cannot divide it anymore without changing its fundamental nature. That final, indivisible particle is called an atom.
Word Origin: The word "atom" comes from the Greek word 'atomos', which means "indivisible" or "uncuttable." The Greeks believed that if you kept dividing matter, you would eventually reach a particle that could not be divided further.
Interestingly, thousands of years before modern science developed microscopes and sophisticated instruments, ancient Indian philosophers had already proposed similar ideas. Maharshi Kanad, an ancient Indian sage and philosopher, proposed that everything in the universe is made up of tiny, indivisible particles called Parmanu (paramāṇu). This shows that human curiosity about the fundamental nature of matter is as old as civilization itself!
Today, we know that atoms are not truly indivisible – they contain even smaller particles like protons, neutrons, and electrons. However, in chemical reactions, atoms behave as if they are indivisible units, which is why understanding atoms is fundamental to understanding chemistry.
⚛️ 2. The Early Ideas About Atoms
Before scientists could actually observe atoms through experiments, they had to imagine and reason about their existence. Both ancient Indian and Greek philosophers independently arrived at similar conclusions through logical thinking.
Ancient Philosophical Thinking
The Greek philosophers Democritus and Leucippus (around 400 BCE) proposed that if you kept dividing matter into smaller and smaller pieces, you would eventually reach a point where further division was impossible. They called these ultimate particles "atoms."
Similarly, in ancient India, Maharshi Kanad proposed the concept of Parmanu in his work Vaisheshika Sutra. He suggested that these particles were eternal, indestructible, and infinitesimally small. He even classified these particles based on the elements – earth, water, fire, air, and ether (akasha).
However, these were philosophical concepts based on logical reasoning rather than experimental evidence. It wasn't until the early 19th century that John Dalton, an English school teacher and scientist, gave the first scientific atomic theory based on experimental observations.
📘 3. Dalton's Atomic Theory (1808)
John Dalton's atomic theory, published in 1808, was revolutionary because it was the first theory to explain atoms in a scientific, systematic way. Dalton imagined atoms as solid, indivisible, perfectly spherical particles – similar to tiny billiard balls.
Dalton's Main Postulates (Explained in Detail)
| No. | Postulate | Explanation |
|---|---|---|
| 1 | All matter is made up of extremely small particles called atoms. | Everything in the universe, whether living or non-living, solid, liquid, or gas, is built from these tiny building blocks called atoms. |
| 2 | Atoms cannot be created, divided, or destroyed. | Atoms are permanent and eternal. In chemical reactions, they are neither created nor destroyed; they simply rearrange themselves to form new substances. |
| 3 | Atoms of a given element are identical in mass and properties. | Every hydrogen atom is identical to every other hydrogen atom. All gold atoms are the same. This uniformity is what gives elements their characteristic properties. |
| 4 | Atoms of different elements have different masses and properties. | An oxygen atom is different from a nitrogen atom in weight, size, and chemical behavior. This difference is what makes elements unique. |
| 5 | Atoms combine in fixed whole-number ratios to form compounds. | Water always has hydrogen and oxygen atoms in a 2:1 ratio (H₂O). This ratio never changes, which is why water always has the same properties. |
| 6 | Chemical reactions involve rearrangement of atoms. | In a chemical reaction, atoms simply exchange partners and rearrange. No atoms are created or destroyed in the process. |
Why Dalton's Theory Was Important
- Why chemical reactions follow fixed proportions – If atoms combine in definite ratios, then the substances they form will always have the same composition.
- Why mass is conserved in reactions – If atoms are neither created nor destroyed, then the total mass before and after a reaction must remain constant.
- Why elements have unique properties – Different atoms have different masses and behaviors, giving each element its characteristic properties.
- How compounds are formed – Atoms of different elements combine in simple whole-number ratios to form compounds.
⚛️ 8. What is a Molecule?
Most atoms in nature do not exist independently. They have a strong tendency to combine with other atoms (of the same or different elements) to form larger particles called molecules.
Definition: A molecule is the smallest particle of a substance (element or compound) that can exist independently and retains all the chemical properties of that substance.
Think of it this way: If atoms are like individual LEGO blocks, then molecules are the structures you build by connecting those blocks together. Just as different LEGO structures have different shapes and functions, different molecules have different properties.
Why Do Atoms Form Molecules?
Atoms combine to form molecules because they become more stable when bonded together. Except for noble gases (like helium, neon, argon), most atoms are unstable when alone. They achieve stability by sharing or transferring electrons with other atoms, forming chemical bonds.
Examples of Molecules
- Oxygen molecule (O₂): Two oxygen atoms combine to form one oxygen molecule that we breathe.
- Water molecule (H₂O): Two hydrogen atoms and one oxygen atom combine to form water.
- Carbon dioxide molecule (CO₂): One carbon atom and two oxygen atoms combine to form carbon dioxide.
- Glucose molecule (C₆H₁₂O₆): 6 carbon, 12 hydrogen, and 6 oxygen atoms combine to form glucose sugar.
Types of Molecules
Molecules can be classified into two main categories based on their composition:
(i) Molecules of Elements
These molecules are formed when atoms of the same element combine together. These molecules represent elements in their natural state.
| Element | Atomicity | Molecular Formula | Example |
|---|---|---|---|
| Hydrogen | Diatomic | H₂ | Two hydrogen atoms form one molecule |
| Oxygen | Diatomic | O₂ | Two oxygen atoms form one molecule |
| Nitrogen | Diatomic | N₂ | Two nitrogen atoms form one molecule |
| Chlorine | Diatomic | Cl₂ | Two chlorine atoms form one molecule |
| Ozone | Triatomic | O₃ | Three oxygen atoms form ozone |
| Phosphorus | Tetratomic | P₄ | Four phosphorus atoms form one molecule |
| Sulphur | Polyatomic | S₈ | Eight sulphur atoms form one molecule |
| Helium | Monoatomic | He | Exists as single atoms (noble gas) |
| Neon | Monoatomic | Ne | Exists as single atoms (noble gas) |
| Argon | Monoatomic | Ar | Exists as single atoms (noble gas) |
(ii) Molecules of Compounds
These molecules are formed when atoms of different elements combine together in a fixed ratio. These molecules represent chemical compounds.
| Compound | Molecular Formula | Constituent Atoms | Atomicity |
|---|---|---|---|
| Water | H₂O | 2 Hydrogen + 1 Oxygen | 3 |
| Carbon dioxide | CO₂ | 1 Carbon + 2 Oxygen | 3 |
| Ammonia | NH₃ | 1 Nitrogen + 3 Hydrogen | 4 |
| Methane | CH₄ | 1 Carbon + 4 Hydrogen | 5 |
| Sulphuric acid | H₂SO₄ | 2 Hydrogen + 1 Sulphur + 4 Oxygen | 7 |
| Glucose | C₆H₁₂O₆ | 6 Carbon + 12 Hydrogen + 6 Oxygen | 24 |
Atomicity: The number of atoms present in one molecule of an element or compound is called its atomicity. For example, O₂ has atomicity 2, O₃ has atomicity 3, and S₈ has atomicity 8.
🧮 9. Molecular Mass and Formula Unit Mass
Molecular Mass
The molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance.
Example 3: Calculate the molecular mass of water (H₂O)
Step 1: Identify the atoms present
Water has 2 hydrogen atoms and 1 oxygen atom
Step 2: Find atomic masses
- Atomic mass of H = 1 u
- Atomic mass of O = 16 u
Step 3: Calculate molecular mass
Molecular mass of H₂O = (2 × 1) + (1 × 16)
= 2 + 16 = 18 u
Example 4: Calculate the molecular mass of glucose (C₆H₁₂O₆)
Atoms present: 6 Carbon + 12 Hydrogen + 6 Oxygen
Atomic masses:
- C = 12 u
- H = 1 u
- O = 16 u
Calculation:
Molecular mass = (6 × 12) + (12 × 1) + (6 × 16)
= 72 + 12 + 96 = 180 u
Formula Unit Mass
For ionic compounds (like common salt - NaCl), we don't use the term "molecular mass" because they don't exist as discrete molecules. Instead, we use the term formula unit mass.
Formula Unit Mass: The sum of atomic masses of all atoms in the formula unit of an ionic compound.
Example 5: Calculate the formula unit mass of sodium chloride (NaCl)
Atoms present: 1 Sodium + 1 Chlorine
Atomic masses:
- Na = 23 u
- Cl = 35.5 u
Calculation:
Formula unit mass = 23 + 35.5 = 58.5 u
🔢 10. The Mole Concept – Counting Atoms
Imagine you're a shopkeeper who needs to count millions of rice grains or sugar crystals every day. Counting them one by one would be impossible! Instead, you would weigh them. Similarly, chemists need a convenient way to count the enormous number of atoms and molecules in chemical reactions.
This is where the concept of the mole comes in. A mole is simply a very large counting unit, similar to how a dozen means 12 or a gross means 144.
Definition: One mole is defined as the amount of substance that contains as many particles (atoms, molecules, or ions) as there are atoms in exactly 12 grams of carbon-12.
This number is called Avogadro's number (or Avogadro's constant), named after the Italian scientist Amedeo Avogadro.
(This is Avogadro's number)
How Large is Avogadro's Number?
To appreciate how incredibly large this number is, consider this:
If you counted 6.022 × 10²³ grains of rice at a rate of one grain per second, it would take you more than 19 trillion years to finish counting – that's longer than the age of the universe!
Understanding Molar Mass
The molar mass of a substance is the mass of one mole of that substance, expressed in grams.
This remarkable relationship means:
- 1 mole of hydrogen atoms (H) = 1 gram
- 1 mole of carbon atoms (C) = 12 grams
- 1 mole of oxygen atoms (O) = 16 grams
- 1 mole of water molecules (H₂O) = 18 grams
Important Formulas for Mole Calculations
Example 6: Calculate the number of moles in 46 grams of sodium (Na)
Given:
- Mass of sodium = 46 g
- Atomic mass of Na = 23 u
- Therefore, molar mass of Na = 23 g/mol
Solution:
Number of moles = Given mass / Molar mass
= 46 g / 23 g/mol
= 2 moles
Example 7: How many molecules are present in 18 grams of water?
Given:
- Mass of water = 18 g
- Molecular mass of H₂O = 18 u
- Molar mass of H₂O = 18 g/mol
Step 1: Calculate number of moles
Number of moles = 18 g / 18 g/mol = 1 mole
Step 2: Calculate number of molecules
Number of molecules = 1 mole × 6.022 × 10²³
= 6.022 × 10²³ molecules
Example 8: Calculate the mass of 0.5 moles of carbon dioxide (CO₂)
Given:
- Number of moles = 0.5
- Molecular mass of CO₂ = 12 + (2 × 16) = 44 u
- Molar mass of CO₂ = 44 g/mol
Solution:
Rearranging: Mass = Number of moles × Molar mass
= 0.5 × 44
= 22 grams
⚡ 11. Valency – The Combining Capacity
You might have noticed that hydrogen always forms H₂, water is always H₂O, and ammonia is always NH₃. Why do elements combine in these specific ratios? The answer lies in a property called valency.
Definition: Valency is the combining capacity of an element – the number of hydrogen atoms or twice the number of oxygen atoms that can combine with one atom of the element.
Think of valency as the number of "hands" an atom has available to hold onto other atoms. Some atoms have one hand (valency 1), some have two hands (valency 2), and so on.
Common Valencies of Elements
| Element | Symbol | Valency | Example Compound |
|---|---|---|---|
| Hydrogen | H | 1 | HCl, H₂O |
| Sodium | Na | 1 | NaCl, Na₂O |
| Potassium | K | 1 | KCl, K₂O |
| Magnesium | Mg | 2 | MgCl₂, MgO |
| Calcium | Ca | 2 | CaCl₂, CaO |
| Oxygen | O | 2 | H₂O, CO₂ |
| Aluminium | Al | 3 | AlCl₃, Al₂O₃ |
| Nitrogen | N | 3 | NH₃, N₂O₃ |
| Carbon | C | 4 | CH₄, CO₂ |
| Silicon | Si | 4 | SiO₂, SiCl₄ |
Valency of Radicals (Polyatomic Ions)
| Radical Name | Formula | Valency |
|---|---|---|
| Hydroxide | OH⁻ | 1 |
| Nitrate | NO₃⁻ | 1 |
| Carbonate | CO₃²⁻ | 2 |
| Sulphate | SO₄²⁻ | 2 |
| Phosphate | PO₄³⁻ | 3 |
| Ammonium | NH₄⁺ | 1 |
📝 12. Writing Chemical Formulas
A chemical formula represents the composition of a molecule using symbols and numbers. It tells us which elements are present and in what ratio.
Rules for Writing Chemical Formulas
- Write the symbols: Write the symbols of the elements or radicals present in the compound.
- Write valencies: Write the valency of each element/radical below its symbol.
- Exchange valencies: Exchange the valencies and write them as subscripts.
- Simplify if needed: If the subscripts have a common factor, divide them by that factor.
- For radicals: If the valency exchanged is greater than 1, enclose the radical in brackets with the subscript outside.
Example 9: Write the formula for magnesium chloride
Step 1: Write symbols and valencies
Mg (valency 2) and Cl (valency 1)
Step 2: Exchange valencies
Mg₁Cl₂ → MgCl₂
Note: We don't write the subscript "1", so it becomes MgCl₂
Example 10: Write the formula for aluminium oxide
Step 1: Write symbols and valencies
Al (valency 3) and O (valency 2)
Step 2: Exchange valencies
Al₂O₃ → Al₂O₃
Example 11: Write the formula for calcium hydroxide
Step 1: Write symbols and valencies
Ca (valency 2) and OH (valency 1)
Step 2: Exchange valencies
Since the hydroxide radical appears twice, enclose it in brackets
Ca₁(OH)₂ → Ca(OH)₂
Example 12: Write the formula for ammonium sulphate
Step 1: Write symbols and valencies
NH₄ (valency 1) and SO₄ (valency 2)
Step 2: Exchange valencies
(NH₄)₂(SO₄)₁ → (NH₄)₂SO₄
🎯 13. Practice Questions
Short Answer Questions
Question 1:
What is the difference between an atom and a molecule?
Question 2:
Define the law of conservation of mass with an example.
Question 3:
Calculate the molecular mass of sulphuric acid (H₂SO₄).
Question 4:
How many molecules are present in 32 grams of oxygen gas (O₂)?
Question 5:
Write the chemical formula for: (a) Sodium carbonate (b) Calcium hydroxide (c) Aluminium sulphate
Long Answer Questions
Question 6:
State and explain Dalton's atomic theory. Mention any two limitations of this theory.
Question 7:
Explain the law of constant proportions with suitable examples.
Question 8:
What is a mole? How is it related to Avogadro's number? Calculate the number of atoms in 0.5 moles of carbon.
Question 9:
Differentiate between atomicity and valency with examples.
Question 10:
Calculate the number of moles and molecules in 90 grams of water.
📌 14. Key Points to Remember
💡 15. Tips for Exam Success
For Theory Questions:
- Understand concepts, don't just memorize
- Practice writing definitions in your own words
- Always provide examples when explaining laws
- Draw diagrams wherever possible (especially for molecular structures)
For Numerical Problems:
- Write the given data clearly
- Mention the formula you're using
- Show all steps of calculation
- Always write units with your final answer
- Practice at least 10-15 numerical problems daily
Common Mistakes to Avoid:
- ❌ Confusing atomic mass with molar mass
- ❌ Forgetting to enclose radicals in brackets when writing formulas
- ❌ Writing incorrect subscripts in chemical formulas
- ❌ Not simplifying chemical formulas when possible
- ❌ Mixing up valency and atomicity
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🔬 16. Advanced Topics - Deep Dive into Atomic Structure
Subatomic Particles: Inside the Atom
While Dalton believed atoms were indivisible, later scientific discoveries revealed that atoms themselves are made up of even smaller particles called subatomic particles. The three main subatomic particles are:
| Particle | Symbol | Charge | Mass | Location | Discovered By |
|---|---|---|---|---|---|
| Proton | p⁺ | +1 | 1 u | Nucleus | Ernest Rutherford (1919) |
| Neutron | n⁰ | 0 (neutral) | 1 u | Nucleus | James Chadwick (1932) |
| Electron | e⁻ | -1 | 1/1836 u (negligible) | Orbits around nucleus | J.J. Thomson (1897) |
Structure of an Atom
An atom consists of two main regions:
1. Nucleus (Center of the Atom):
- Contains protons and neutrons
- Extremely small but very dense
- Contains almost all the mass of the atom
- Positively charged due to protons
2. Electron Cloud (Surrounding the Nucleus):
- Contains electrons moving in orbits or shells
- Takes up most of the space in an atom
- Negatively charged due to electrons
- Mass is negligible compared to nucleus
Important Relationships in Atomic Structure
Example 13: Carbon Atom Structure
Carbon has atomic number 6 and mass number 12
- Number of Protons: 6 (same as atomic number)
- Number of Electrons: 6 (same as number of protons)
- Number of Neutrons: 12 - 6 = 6
Therefore, a carbon atom has 6 protons, 6 neutrons, and 6 electrons.
Limitations of Dalton's Atomic Theory
While Dalton's atomic theory was groundbreaking, modern science has revealed several limitations:
| Dalton's Statement | Modern Understanding |
|---|---|
| Atoms are indivisible | Atoms can be divided into protons, neutrons, and electrons. Nuclear reactions can split atoms. |
| All atoms of an element are identical | Isotopes exist - atoms of the same element can have different masses (different numbers of neutrons). |
| Atoms of different elements have different masses | Some atoms of different elements can have the same mass (isobars). Example: Argon-40 and Calcium-40. |
| Atoms cannot be created or destroyed | Nuclear reactions (fission and fusion) can create or destroy atoms by converting mass to energy (E=mc²). |
| Atoms combine in simple whole number ratios | Complex organic compounds don't always follow simple ratios. Example: sugar (C₁₂H₂₂O₁₁) has a complex ratio. |
⚗️ 17. Comprehensive Guide to Chemical Combination Laws
Why Are These Laws Important?
Before Dalton's atomic theory, chemists had observed certain patterns in how substances combined. These observations led to the formulation of laws that still form the basis of modern chemistry. Understanding these laws helps us predict how substances will react and in what proportions.
Law of Conservation of Mass - Detailed Analysis
Statement: "In a chemical reaction, the total mass of products is always equal to the total mass of reactants. Mass is neither created nor destroyed."
Proposed by: Antoine Lavoisier (1789)
Also known as: Law of Indestructibility of Matter
Real-World Applications
Example 14: Rusting of Iron
Reaction: Iron + Oxygen → Iron Oxide (Rust)
If 112 grams of iron reacts with 48 grams of oxygen:
- Total mass of reactants = 112 + 48 = 160 grams
- Total mass of product (rust) = 160 grams
Conclusion: The mass before and after the reaction remains constant.
Example 15: Decomposition of Water
Reaction: 2H₂O → 2H₂ + O₂
If 36 grams of water is decomposed:
- Mass of water decomposed = 36 grams
- Mass of hydrogen produced = 4 grams
- Mass of oxygen produced = 32 grams
- Total mass of products = 4 + 32 = 36 grams
Verification: Mass of reactant = Mass of products ✓
Law of Constant Proportions - Detailed Analysis
Statement: "A pure chemical compound always contains the same elements combined in the same fixed proportion by mass, regardless of its source or method of preparation."
Proposed by: Joseph Proust (1799)
Also known as: Law of Definite Proportions or Law of Definite Composition
Detailed Examples
Example 16: Ammonia (NH₃)
Ammonia always contains nitrogen and hydrogen in the ratio 14:3 by mass.
| Sample | Mass of Nitrogen | Mass of Hydrogen | Ratio |
|---|---|---|---|
| Sample 1 | 28 g | 6 g | 28:6 = 14:3 |
| Sample 2 | 70 g | 15 g | 70:15 = 14:3 |
| Sample 3 | 140 g | 30 g | 140:30 = 14:3 |
No matter where the ammonia comes from, the ratio is always constant!
Example 17: Carbon Dioxide (CO₂)
Carbon dioxide always contains carbon and oxygen in the ratio 3:8 by mass.
Calculation:
- Atomic mass of C = 12 u
- Atomic mass of O = 16 u
- In CO₂: 1 carbon atom + 2 oxygen atoms
- Mass ratio = 12 : (2 × 16) = 12:32 = 3:8
This means in 44 grams of CO₂:
- Carbon = (3/11) × 44 = 12 grams
- Oxygen = (8/11) × 44 = 32 grams
🧪 18. Advanced Mole Concept Problems
Understanding the Mole - A Chemist's Dozen
Just as a baker uses "dozen" (12) to count eggs, chemists use "mole" to count atoms and molecules. The mole is simply a very large number used to count incredibly small particles.
Fun Fact: If you had a mole of rice grains, they would cover the entire surface of Earth to a depth of about 75 meters (246 feet)! That's how large Avogadro's number is.
Comprehensive Mole Problems
Example 18: Multi-Step Mole Calculation
Question: Calculate the number of atoms in 5 moles of oxygen molecules (O₂).
Solution:
Step 1: Find number of molecules
Number of molecules = 5 moles × 6.022 × 10²³
= 3.011 × 10²⁴ molecules
Step 2: Find number of atoms
Since each O₂ molecule contains 2 oxygen atoms:
Number of atoms = 3.011 × 10²⁴ × 2
= 6.022 × 10²⁴ atoms
Example 19: Converting Between Mass and Particles
Question: How many molecules are present in 88 grams of carbon dioxide (CO₂)?
Given:
- Mass of CO₂ = 88 g
- Molecular mass of CO₂ = 12 + (2 × 16) = 44 u
- Molar mass = 44 g/mol
Step 1: Calculate moles
Number of moles = 88 / 44 = 2 moles
Step 2: Calculate molecules
Number of molecules = 2 × 6.022 × 10²³
= 1.2044 × 10²⁴ molecules
Example 20: Reverse Calculation
Question: What mass of calcium contains 3.011 × 10²³ atoms?
Given:
- Number of atoms = 3.011 × 10²³
- Atomic mass of Ca = 40 u
- Molar mass = 40 g/mol
Step 1: Calculate moles
Number of moles = (3.011 × 10²³) / (6.022 × 10²³)
= 0.5 moles
Step 2: Calculate mass
Mass = 0.5 × 40 = 20 grams
Example 21: Complex Compound Calculation
Question: Calculate the number of oxygen atoms in 0.25 moles of sulphuric acid (H₂SO₄).
Given:
- Moles of H₂SO₄ = 0.25
- Each molecule has 4 oxygen atoms
Step 1: Find number of molecules
Number of molecules = 0.25 × 6.022 × 10²³
= 1.5055 × 10²³ molecules
Step 2: Find oxygen atoms
Number of O atoms = 1.5055 × 10²³ × 4
= 6.022 × 10²³ oxygen atoms
📐 19. Mastering Chemical Formula Writing
Advanced Formula Writing Techniques
Writing chemical formulas correctly is essential for understanding chemistry. Let's explore more complex examples and techniques.
Method: The Criss-Cross Method
Steps for Criss-Cross Method:
- Write symbols of elements/radicals side by side
- Write their valencies below the symbols
- Cross-over the valencies (first element's valency becomes second element's subscript)
- Simplify if possible by dividing by common factor
- For radicals with subscript > 1, use brackets
Example 22: Aluminium Sulphate
Elements/Radicals: Al (valency 3) and SO₄ (valency 2)
Step-by-step:
Al³⁺ and SO₄²⁻
Criss-cross: Al₂(SO₄)₃
Final Answer: Al₂(SO₄)₃
Note: We use brackets for SO₄ because its subscript is 3.
Example 23: Calcium Phosphate
Elements/Radicals: Ca (valency 2) and PO₄ (valency 3)
Criss-cross: Ca₃(PO₄)₂
Final Answer: Ca₃(PO₄)₂
Example 24: Magnesium Oxide (with simplification)
Elements: Mg (valency 2) and O (valency 2)
Before simplification: Mg₂O₂
After simplification: Divide both by 2
Final Answer: MgO
Practice Set: Write Formulas for These Compounds
| Compound Name | Elements/Radicals | Valencies | Formula |
|---|---|---|---|
| Sodium Sulphate | Na, SO₄ | 1, 2 | Na₂SO₄ |
| Potassium Nitrate | K, NO₃ | 1, 1 | KNO₃ |
| Magnesium Hydroxide | Mg, OH | 2, 1 | Mg(OH)₂ |
| Ferric Oxide | Fe, O | 3, 2 | Fe₂O₃ |
| Zinc Carbonate | Zn, CO₃ | 2, 2 | ZnCO₃ |
| Barium Chloride | Ba, Cl | 2, 1 | BaCl₂ |
🎓 20. Comprehensive Practice Questions
Numerical Problems (Advanced Level)
Question 11:
Calculate the molecular mass of: (a) Nitric acid (HNO₃) (b) Calcium carbonate (CaCO₃) (c) Ammonium sulphate [(NH₄)₂SO₄]
Question 12:
How many moles are present in: (a) 71 grams of chlorine gas (Cl₂)? (b) 100 grams of calcium carbonate (CaCO₃)?
Question 13:
Calculate the mass of: (a) 2.5 moles of oxygen atoms (b) 0.2 moles of water molecules (c) 5 moles of sodium atoms
Question 14:
How many atoms are present in: (a) 8 grams of oxygen gas (O₂)? (b) 52 grams of helium?
Question 15:
Calculate the number of molecules in: (a) 36 grams of water (b) 22 grams of carbon dioxide (c) 17 grams of ammonia
Theory Questions (Comprehensive)
Question 16:
Explain the law of conservation of mass with two examples from daily life.
Question 17:
What are isotopes? Why do some elements have non-whole number atomic masses? Explain with the example of chlorine.
Question 18:
Differentiate between: (a) Atom and molecule (b) Element and compound (c) Atomic mass and molecular mass (d) Mole and Avogadro's number
Question 19:
Describe the structure of an atom. What are the three subatomic particles? Where are they located and what are their charges?
Question 20:
Write the chemical formulas for the following compounds: (a) Magnesium hydroxide (b) Aluminium oxide (c) Sodium carbonate (d) Potassium permanganate (e) Copper sulphate
Application-Based Questions
Question 21:
A compound contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. Calculate the empirical formula. (Given: C=12, H=1, O=16)
Question 22:
In a reaction, 5.6 grams of iron reacts with oxygen to form 8 grams of iron oxide. Calculate the mass of oxygen that reacted. Verify the law of conservation of mass.
Question 23:
A sample of ammonia (NH₃) contains 14 grams of nitrogen. Calculate: (a) Mass of hydrogen in the sample (b) Total mass of ammonia (c) Number of moles of ammonia
✅ 21. Solutions to Practice Questions
Numerical Solutions
Solution to Question 11:
(a) HNO₃: 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 u
(b) CaCO₃: 40 + 12 + (3 × 16) = 40 + 12 + 48 = 100 u
(c) (NH₄)₂SO₄: (2 × 14) + (8 × 1) + 32 + (4 × 16) = 28 + 8 + 32 + 64 = 132 u
Solution to Question 12:
(a) Chlorine gas (Cl₂):
- Molecular mass of Cl₂ = 2 × 35.5 = 71 u
- Molar mass = 71 g/mol
- Moles = 71 / 71 = 1 mole
(b) CaCO₃:
- Molecular mass = 100 u (from Question 11)
- Molar mass = 100 g/mol
- Moles = 100 / 100 = 1 mole
Solution to Question 14:
(a) 8 grams of O₂:
- Molecular mass of O₂ = 32 u
- Moles = 8/32 = 0.25 moles
- Molecules = 0.25 × 6.022 × 10²³ = 1.5055 × 10²³
- Each molecule has 2 atoms
- Total atoms = 1.5055 × 10²³ × 2 = 3.011 × 10²³ atoms
🏆 22. Final Summary and Quick Revision
Most Important Formulas - Memorize These!
- Number of moles = Mass / Molar mass
- Number of particles = Moles × 6.022 × 10²³
- Mass Number = Protons + Neutrons
- Atomic Number = Protons = Electrons
- Molecular Mass = Sum of atomic masses of all atoms
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