Notes Class 9 Science Exploration Chapter 9 Atomic Foundation of Matter

🌊

Chapter Introduction: Matter & Its Properties

You already know that matter is made of atoms. But have you wondered — when hydrogen gas and oxygen gas combine to form water, why is water completely different from both gases?

🔥 Hydrogen (H₂)

Colourless gas. Highly combustible — catches fire easily. Lighter than air.

💨 Oxygen (O₂)

Colourless gas. Supports combustion — helps other things burn. Essential for breathing.

💡

Surprising Fact!

When H₂ and O₂ combine, they form water (H₂O) — a liquid that neither burns nor helps burning. It actually extinguishes fire! Yet the mass of water formed = sum of masses of H₂ and O₂ used. This leads us to the Law of Conservation of Mass.

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Key Questions for this Chapter

Does mass change during physical changes? What about chemical changes? How do atoms combine? What types of bonds form? How do we write chemical formulae?

⚖️

Law of Conservation of Mass (द्रव्यमान संरक्षण का नियम)

🔬 Experiments That Proved the Law

🧂 Activity 9.1 — Physical Change (Salt + Water)

When common salt is dissolved in water, the mass of the solution = mass of water + mass of salt. No mass is lost or gained during a physical change.

🧪 Activity 9.2 — Chemical Change (Vinegar + Baking Soda)

In Set-up 1 (open system): baking soda is added to vinegar — CO₂ gas escapes, so final mass < initial mass. Seems like mass is lost!
In Set-up 2 (closed system with balloon): the CO₂ gas inflates the balloon and is captured. Now final mass = initial mass ✅

💡

Lesson Learned!

If products are not allowed to escape (closed system), total mass before reaction = total mass after reaction. The apparent “loss” of mass in open systems is because gas products escape into the air.

Law of Conservation of Mass: Mass of Reactants = Mass of Products

Matter can neither be created nor destroyed in a chemical reaction.

🏆

Antoine Lavoisier — Father of Modern Chemistry (1789)

Lavoisier proposed this law. He conducted careful experiments where he weighed substances before and after reactions in sealed containers. He concluded: “In every operation, an equal quantity of matter exists both before and after the operation.”

📐 Solved Examples

Example 9.1 — Verify Conservation of Mass

Reactants: CaCO₃ = 4.0 g + HCl = 2.92 g

Total mass of reactants = 4.0 + 2.92 = 6.92 g

Products: CO₂ = 1.76 g + H₂O = 0.72 g + CaCl₂ = 4.44 g

Total mass of products = 1.76 + 0.72 + 4.44 = 6.92 g

Since mass of reactants = mass of products → Law of Conservation of Mass is obeyed ✅

Example 9.2 — Carbon + Oxygen → Carbon Dioxide

Given: 12 g C + 32 g O₂ → 44 g CO₂

Find: How much CO₂ is produced by 2.4 g of carbon?

1 g of C gives = 44/12 g of CO₂

2.4 g of C gives = (44/12) × 2.4 = 8.8 g CO₂

⚠️

Common Mistake!

If a student burns ethanol in an open beaker and sees no residue, it does NOT mean mass is violated. The carbon and hydrogen atoms from ethanol form CO₂ and H₂O gases that escape into the air. Total mass is still conserved — you need a closed system to verify!

📏

Law of Constant Proportions (निश्चित अनुपात का नियम)

After Lavoisier, French chemist Joseph Proust studied the composition of compounds and found something remarkable: a compound ALWAYS has the same elements in the SAME ratio by mass, no matter where it comes from!

💧

Water is Always Water — Wherever It Comes From!

Whether water comes from a river, a borewell, the ocean (after purification), or a lab — it always contains hydrogen and oxygen in the mass ratio of1:8. If you take 9 g of pure water, you always get 1 g H and 8 g O upon decomposition.

Law of Constant Proportions (Proust’s Law):

In a given compound, the elements are always present in a fixed ratio by mass — regardless of source or method of preparation.

🏺

Ancient Indian Connection — Cinnabar (Hingula)

Ancient civilisations, including India, used red pigment from rocks called hingula (cinnabar in Latin). Heating cinnabar always gave mercury and sulfur in the same ratio — 86.22% and 13.78% by mass. Grinding mercury and sulfur in this exact ratio also reformed cinnabar! This is a perfect example of the Law of Constant Proportions.

📐 Solved Examples

Example 9.3 — Sodium Chloride (NaCl)

NaCl contains Na and Cl in mass ratio 23 : 35.5

Given: 46 g of sodium reacts completely. How much chlorine is needed?

Mass of chlorine = (35.5 ÷ 23) × 46

= 71 g of Chlorine needed

🟢 Applies to Compounds

A compound always has fixed element ratios. NaCl always has Na and Cl in 23:35.5 ratio, no matter how it was made.

🔴 Does NOT apply to Mixtures

Air is a mixture — its composition varies (more or less O₂ in different places). Mixtures have no fixed ratio.

🔬

Joseph Louis Proust — Law of Definite Proportions

Proust was a French chemist famous for careful experimental work. He studied copper carbonate and showed it always contains copper, carbon, and oxygen in the same proportion by mass, regardless of how it was prepared or where it was found.

⚗️

Dalton’s Atomic Theory (1808)

The two laws of Conservation of Mass and Constant Proportions formed the foundation of John Dalton’s Atomic Theory. Dalton used these laws to explain atomic behavior scientifically.

📖

What is a Postulate?

A postulate is a fundamental assumption accepted as truth without formal proof, from which further ideas are developed.

📋 Dalton’s Postulates

All matter is made up of very tiny particles called atoms (परमाणु), which participate in chemical reactions.
Atoms are indivisible particles — they cannot be created or destroyed in a chemical reaction.
Atoms of a given element are identical in mass and chemical properties.
Atoms of different elements have different masses and chemical properties.
Atoms combine in the ratio of simple whole numbers to form compounds.
The relative number and kinds of atoms are constant in a given compound.

✅ How it explains Conservation of Mass

In a chemical reaction, atoms are not destroyed or created. They simply rearrange. So total number of atoms before = after → total mass is conserved.

✅ How it explains Constant Proportions

Atoms combine in fixed whole-number ratios. Since each atom has a definite mass, the same compound always has the same mass ratio of elements.

⚠️

Dalton Was Partially Wrong!

Later discoveries showed atoms ARE divisible (electrons, protons, neutrons). Atoms of the same element can have different masses (isotopes). But his theory was still revolutionary for its time and forms the foundation of chemistry!

🏆

John Dalton — Born in England

In 1793, Dalton moved to Manchester to teach mathematics, physics, and chemistry. He spent most of his life teaching and researching there. In 1808, he presented his atomic theory — a turning point in the study of matter.

🔬

How Atoms Combine — Molecules & Chemical Bonds

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Definition: Molecule (अणु)

A molecule is an electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance.

Examples: H₂ (two hydrogen atoms), Cl₂ (two chlorine atoms), H₂O (two hydrogen + one oxygen atom), HCl (one hydrogen + one chlorine atom).

💡

Helium is special!

Some elements like Helium (He) exist only as single atoms (monoatomic) because their outermost shell is already full (2 electrons). They don’t need to bond with any other atom!

🔗 Why Do Atoms Combine?

Atoms combine because they want to achieve a stable electronic configuration — 8 electrons in the outermost shell (2 for K-shell atoms like H and He). When atoms bond, the total energy of the system decreases, making it more stable.

🤝 Sharing of Electrons

Atoms share their valence electrons with another atom. Both atoms get to “count” the shared electrons. This forms aCovalent Bond.

↔️ Transfer of Electrons

One atom donates its valence electrons to another atom. Ions (charged particles) are formed. This forms anIonic Bond.

Chemical Bond (रासायनिक बंध) = The force that holds atoms together in a molecule or compound

🤝

Covalent Bond — Bonding by Sharing of Electrons

📌

Definition: Covalent Bond (सहसंयोजक बंध)

A covalent bond is formed by the sharing of one or more pairs of electrons between two atoms. Each atom contributes electrons to the shared pair, and the shared electrons are attracted to both nuclei.

🔵 Types of Covalent Bonds

Single Bond (—)

One pair of electrons shared between two atoms. Example: H—H, H—Cl, Cl—Cl

Double Bond (=)

Two pairs of electrons shared between two atoms. Example: O=O (oxygen molecule)

Triple Bond (≡)

Three pairs of electrons shared. Example: N≡N (nitrogen molecule)

⚗️ Formation of Molecules of Elements

🔵 Hydrogen Molecule (H₂)

H has 1 electron, needs 1 more to complete K-shell (max = 2). Two H atoms each share 1 electron → H—H (single bond). Formula: H₂

🟢 Chlorine Molecule (Cl₂)

Cl has 7 valence electrons, needs 1 more. Two Cl atoms each share 1 electron → Cl—Cl (single bond). Formula: Cl₂

H atom

→

H—H (H₂)

O atom

double bond

O=O

O₂ molecule

H₂ forms a single bond; O₂ forms a double bond

🧪 Formation of Compound Molecules

Compound	Atoms Involved	Electrons Needed Each	Bond Type	Formula
Hydrogen Chloride	H (1e⁻) + Cl (7e⁻)	H needs 1; Cl needs 1	Single (H—Cl)	HCl
Water	2H + O (6e⁻)	O needs 2; each H needs 1	2 single bonds	H₂O
Ammonia	N (5e⁻) + 3H	N needs 3; each H needs 1	3 single bonds	NH₃
Carbon dioxide	C (4e⁻) + 2O	C needs 4; each O needs 2	2 double bonds	CO₂

H—Cl H—O—H H—N—H (with H below N) O=C=O

Structural representations of HCl, H₂O, NH₃, CO₂

📛 Naming Covalent Compounds

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Prefix System for Covalent Compounds

mono(1), di(2), tri(3), tetra(4), penta(5), hexa(6), hepta(7), octa(8). First element keeps its name; second element ends in-ide. Mono- is usually omitted for the FIRST element.

Formula	IUPAC Name	Trick to Remember
CO	Carbon monoxide	mono- used for second element O
CO₂	Carbon dioxide	di = 2 oxygen atoms
CS₂	Carbon disulfide	di = 2 sulfur atoms
PCl₃	Phosphorus trichloride	tri = 3 chlorine atoms
SF₆	Sulfur hexafluoride	hexa = 6 fluorine atoms
N₂O₄	Dinitrogen tetroxide	both elements use prefix
H₂S	Hydrogen sulfide	No prefix before hydrogen!
H₂O	Water (common name)	Official: hydrogen monoxide
NH₃	Ammonia (common name)	Official: nitrogen trihydride

⚠️

Naming Rule: Drop the vowel if needed!

If a prefix ends in ‘a’ or ‘o’ AND the element name starts with a vowel → drop the last vowel. Example: penta + oxide = pentoxide (not pentaoxide), mono + oxide = monoxide (not monooxide).

⚡

Ionic Bond — Bonding by Transfer of Electrons

📌

Definition: Ionic Bond (आयनिक बंध)

An ionic bond is the electrostatic force of attraction between oppositely charged ions (cation and anion) that holds them together. It forms when one atom transfers electrons to another atom.

🧂 Classic Example: Formation of NaCl (Common Salt)

Na → Na⁺ (Sodium Cation)

Sodium (Z=11): Electronic config = 2, 8, 1. Has 1 valence electron — it donates this to get stable 2,8 configuration. After losing 1 electron: 11 protons, 10 electrons → Net charge = +1 → FormsNa⁺ cation

Cl → Cl⁻ (Chloride Anion)

Chlorine (Z=17): Electronic config = 2, 8, 7. Has 7 valence electrons — it accepts 1 to get stable 2,8,8 configuration. After gaining 1 electron: 17 protons, 18 electrons → Net charge = −1 → FormsCl⁻ anion

(2,8,1)

Sodium atom

→ e⁻ →

(2,8,7)

Chlorine atom

⟹

Na⁺

(2,8)

Cation (+)

Cl⁻

(2,8,8)

Anion (−)

→ NaCl

Formation of ionic compound NaCl — sodium donates, chlorine accepts

🔮

Cation vs Anion

Cation (धनायन): Positively charged ion (lost electrons). Metals form cations. Example: Na⁺, Ca²⁺, Fe³⁺

Anion (ऋणायन): Negatively charged ion (gained electrons). Non-metals form anions. Example: Cl⁻, O²⁻, S²⁻

Together, cations and anions are calledions (आयन).

🏗️ Crystal Structure of Ionic Compounds

Ionic compounds do not form simple molecules. Instead, they form three-dimensional (3-D) crystal structures. In NaCl, each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ is surrounded by 6 Na⁺, forming a regular repeating pattern called a crystal lattice.

📋 Naming Ionic Compounds

Name the cation first, then the anion.
Names of simple anions end with -ide (chloride, oxide, sulfide).
Polyatomic ions generally do NOT end with -ide (sulfate, carbonate, nitrate).
Examples: NaCl = Sodium chloride; CaO = Calcium oxide; Na₂S = Sodium sulfide

Type	Name of Ion	Formula	Valency/Charge
Common Cations (+)	Sodium	Na⁺	1
	Potassium	K⁺	1
	Calcium	Ca²⁺	2
	Magnesium	Mg²⁺	2
	Aluminium	Al³⁺	3
	Iron (Ferrous)	Fe²⁺	2
	Iron (Ferric)	Fe³⁺	3
Common Anions (−)	Chloride	Cl⁻	1
	Oxide	O²⁻	2
	Sulfide	S²⁻	2
	Hydroxide	OH⁻	1
	Sulfate	SO₄²⁻	2
Polyatomic Ions	Carbonate	CO₃²⁻	2
	Nitrate	NO₃⁻	1
	Hydrogencarbonate	HCO₃⁻	1
	Ammonium	NH₄⁺	1

🧮

Writing Chemical Formulae — The Criss-Cross Method

There’s a quick method to write chemical formulae using the criss-cross method: write the symbols and their valencies, then swap (cross) the valencies as subscripts.

📌

Important Rules

(1) If subscripts have a common factor, simplify (e.g., Mg₂O₂ → MgO). (2) Use brackets when more than one polyatomic ion is present: Mg(OH)₂, not MgOH₂. (3) Charges are NOT written in the final formula.

📝 Examples: Covalent Compounds

HClHvalency: 1 × Clvalency: 1→HCl

H₂SHval: 1 × Sval: 2→H₂S

CCl₄Cval: 4 × Clval: 1→CCl₄

📝 Examples: Ionic Compounds

Compound	Cation (charge)	Anion (charge)	Criss-cross	Final Formula
Calcium chloride	Ca (2+)	Cl (1−)	Ca¹Cl²	CaCl₂
Aluminium oxide	Al (3+)	O (2−)	Al²O³	Al₂O₃
Magnesium oxide	Mg (2+)	O (2−)	Mg²O² → simplify	MgO
Calcium carbonate	Ca (2+)	CO₃ (2−)	Ca¹(CO₃)¹ → simplify	CaCO₃
Magnesium hydroxide	Mg (2+)	OH (1−)	Mg¹(OH)²	Mg(OH)₂
Aluminium sulfate	Al (3+)	SO₄ (2−)	Al²(SO₄)³	Al₂(SO₄)₃
Aluminium hydroxide	Al (3+)	OH (1−)	Al¹(OH)³	Al(OH)₃ NOT AlOH₃

💡

Exam Trick: When to Use Brackets?

Use brackets ( ) ONLY when you have 2 or more polyatomic ions (ions with more than one atom like OH⁻, SO₄²⁻, CO₃²⁻). Single polyatomic ions don’t need brackets. Example: NaOH (no brackets needed — just one OH⁻), but Mg(OH)₂ (brackets needed — two OH⁻ ions).

🔭

Properties of Ionic vs Covalent Compounds

⚡ Ionic Compounds (NaCl, CuSO₄)

Generally soluble in water ✅
Generally insoluble in organic solvents (kerosene, petrol) ❌
Do NOT conduct electricity in solid state (ions are fixed) ❌
DO conduct electricity when dissolved in water (ions are free to move) ✅
Also conduct in molten (liquid) state ✅
High melting and boiling points (strong ionic bonds)
Form crystalline solids

🤝 Covalent Compounds (Camphor, Naphthalene, Sugar)

Most are insoluble in water (except some like sugar) ❌
Generally soluble in organic solvents (kerosene, petrol) ✅
Do NOT conduct electricity in any state (no ions formed) ❌
Even if dissolved in water (like sugar) — no ions in solution → no conductivity ❌
Low melting and boiling points (weak intermolecular forces)
Can be liquids or gases at room temperature

💡

Why does salt solution conduct electricity but sugar solution doesn’t?

NaCl (ionic) dissolves in water to give free Na⁺ and Cl⁻ ions. These ions can carry electric charge → conducts electricity. Sugar (covalent) dissolves but does NOT break into ions — just sugar molecules in water. No ions → no conductivity!

⚠️

Common Exam Mistake!

Ionic compounds DO NOT conduct electricity in the solid state even though they have ions. Why? Because in solid state, the ions are held in fixed positions in the crystal lattice and cannot move. Conductivity requires FREE-MOVING ions.

⚖️

Molecular Mass & Formula Unit Mass

🔵 Molecular Mass (आण्विक द्रव्यमान) — for Covalent Compounds

📌

Definition

Molecular mass = sum of atomic masses of ALL atoms present in ONE molecule of the compound. Unit:u (unified atomic mass unit)

Molecular Mass = Σ (Atomic mass of each atom × Number of that atom in molecule)

Water (H₂O)

H = 1 u; O = 16 u

MM = (1 × 2) + (16 × 1)

= 2 + 16

= 18 u

Carbon Dioxide (CO₂)

C = 12 u; O = 16 u

MM = (12 × 1) + (16 × 2)

= 12 + 32

= 44 u

🔶 Formula Unit Mass (सूत्र इकाई द्रव्यमान) — for Ionic Compounds

📌

Definition

Ionic compounds don’t form molecules — they form 3-D crystals. The simplest whole number ratio of ions is called aformula unit. The mass of one formula unit =formula unit mass.

Sodium Oxide (Na₂O)

Na = 23 u; O = 16 u

FUM = (23 × 2) + (16 × 1)

= 46 + 16

= 62 u

Calcium Nitrate Ca(NO₃)₂

Ca=40; N=14; O=16

FUM = 40 + [(14 + 48) × 2]

= 40 + 124

= 164 u

Compound	Type	Atoms/Ions	Calculation	Mass
H₂O (water)	Covalent → Molecular Mass	2H + 1O	(1×2)+(16×1)	18 u
CO₂	Covalent → Molecular Mass	1C + 2O	(12×1)+(16×2)	44 u
CH₄ (methane)	Covalent → Molecular Mass	1C + 4H	(12×1)+(1×4)	16 u
HNO₃ (nitric acid)	Covalent → Molecular Mass	1H+1N+3O	1+14+(16×3)	63 u
NaCl (common salt)	Ionic → Formula Unit Mass	1Na + 1Cl	23+35.5	58.5 u
Na₂O	Ionic → Formula Unit Mass	2Na + 1O	(23×2)+16	62 u
Mg(OH)₂	Ionic → Formula Unit Mass	1Mg+2O+2H	24+(16+1)×2	58 u

💡

Atomic Masses to Remember for Exams!

H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, K=39, Ca=40, Fe=56, Cu=64, Zn=65

📝 Quick Revision Summary

⚖️ Conservation of MassMass cannot be created or destroyed in a reaction. Total mass of reactants = total mass of products.

📏 Constant ProportionsElements in a compound always combine in a fixed mass ratio regardless of source. (Proust’s Law)

⚗️ Dalton’s TheoryAtoms are indivisible, combine in simple whole number ratios. Foundation of modern chemistry.

🤝 Covalent BondFormed by sharing of electron pairs. Results in molecules. E.g., H₂, O₂, H₂O, CO₂.

⚡ Ionic BondFormed by transfer of electrons. Forms cations (+) and anions (−). E.g., NaCl, CaCl₂.

🔗 Chemical BondForce holding atoms together. Atoms bond to achieve stable octet configuration.

🧮 Criss-Cross MethodSwap valencies as subscripts to write chemical formulae quickly. Simplify if common factor exists.

🔵 Ionic PropertiesSoluble in water, conduct electricity in solution/melt, high melting points, crystalline structure.

🤝 Covalent PropertiesSoluble in organic solvents, do NOT conduct electricity, low melting points.

⚗️ Molecular MassSum of atomic masses of all atoms in one molecule (for covalent compounds). Unit = u.

🔶 Formula Unit MassSum of masses in simplest ion ratio (for ionic compounds). Ionic compounds form crystals, not molecules.

🏷️ NamingCovalent: prefixes (mono, di, tri…) + -ide. Ionic: cation name first, anion ends in -ide.

📋 Important Exam Questions

Q1. State the Law of Conservation of Mass. Describe an experiment that demonstrates this law. (CBSE — 4 Marks)

Law: Matter can neither be created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

Experiment (Vinegar + Baking Soda):

Set-up: A conical flask with vinegar and a balloon filled with baking soda are weighed together. The balloon is fixed to the flask mouth. Initial mass is recorded. The baking soda is allowed to fall into vinegar — CO₂ is produced and captured in the balloon. The system is weighed again.

Observation: Initial mass = Final mass.

Conclusion: Mass is conserved in the chemical reaction. In an open system, if CO₂ escapes, apparent mass decreases — but total mass including the gas is still constant.

Q2. Explain the formation of NaCl with a diagram showing electron transfer. What type of bond is formed and why? (CBSE — 3 Marks)

Sodium (Na, Z=11): Electronic config = 2, 8, 1. Has 1 valence electron. To achieve stable octet (2,8), it loses this electron → Na becomes Na⁺ cation (11 protons, 10 electrons, net charge +1).

Chlorine (Cl, Z=17): Electronic config = 2, 8, 7. Has 7 valence electrons. To achieve stable octet (2,8,8), it gains 1 electron → Cl becomes Cl⁻ anion (17 protons, 18 electrons, net charge −1).

Bond Formation: The oppositely charged Na⁺ and Cl⁻ are held together by electrostatic force of attraction → this is an ionic bond (आयनिक बंध). It forms because Na (less than 4 valence electrons) donates its electron to Cl (more than 4 valence electrons), allowing both to achieve stability.

Q3. Write the chemical formulae for: (i) Aluminium nitrate (ii) Calcium carbonate (iii) Ferric oxide (iv) Magnesium hydroxide. (CBSE — 2 Marks)

(i) Aluminium nitrate: Al³⁺ and NO₃⁻ → Criss-cross: Al¹(NO₃)³ → Al(NO₃)₃

(ii) Calcium carbonate: Ca²⁺ and CO₃²⁻ → same valency: Ca¹(CO₃)¹ → simplify → CaCO₃

(iii) Ferric oxide: Fe³⁺ (ferric = Fe³⁺) and O²⁻ → Criss-cross: Fe²O³ → Fe₂O₃

(iv) Magnesium hydroxide: Mg²⁺ and OH⁻ → Criss-cross: Mg¹(OH)² → Mg(OH)₂

Q4. Distinguish between ionic and covalent compounds on the basis of: (i) Formation (ii) Solubility (iii) Electrical conductivity. (CBSE — 3 Marks)

(i) Formation: Ionic compounds form by transfer of electrons from one atom to another, forming ions (cations and anions). Covalent compounds form by sharing of electrons between atoms, forming molecules.

(ii) Solubility: Ionic compounds are generally soluble in water but insoluble in organic solvents like kerosene and petrol. Covalent compounds are generally insoluble in water but soluble in organic solvents.

(iii) Electrical conductivity: Ionic compounds do not conduct electricity in solid state (ions are fixed) but conduct in aqueous solution or molten state (ions are free to move). Covalent compounds generally do not conduct electricity in any state as they do not form ions.

Q5. Calculate the molecular mass of (i) Nitric acid HNO₃ (ii) Methane CH₄ and formula unit mass of (iii) KCl (iv) Ca(NO₃)₂. Given: H=1, C=12, N=14, O=16, K=39, Cl=35.5, Ca=40. (CBSE — 3 Marks)

(i) HNO₃: = 1 + 14 + (16×3) = 1 + 14 + 48 = 63 u

(ii) CH₄: = 12 + (1×4) = 12 + 4 = 16 u

(iii) KCl: = 39 + 35.5 = 74.5 u

(iv) Ca(NO₃)₂: = 40 + [(14 + 16×3) × 2] = 40 + [62 × 2] = 40 + 124 = 164 u

Notes Class 9 Science Exploration Chapter 9 Atomic Foundation of Matter

Amresh Academy

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