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chapter 8.Journey-Inside-the-Atom.pdf/or ONLINE Classes (ACADEMIC,IITJEE & NEET ) Contact 9491878325 // Subscribe samyans Eduhub youtube channel

Journey Inside the Atom Chapter 8 — Grade 9 Science Prepared by K Sandeep Swamy (M.Sc, B.Ed) For ONLINE Classes (ACADEMIC,IITJEE & NEET ) Contact 9491878325 Subscribe samyans Eduhub youtube channel Everything Around Us Is Made of Atoms All Matter Every living and non-living thing is ultimately made of atoms — the fundamental building blocks of matter. Incredibly Tiny Atoms are so small they cannot be seen with the naked eye — even the most powerful optical microscopes struggle to observe them. The Big Question Is the atom truly the smallest unit of matter, or can it be divided into even smaller particles? CHAPTER 8.1 Roots of Atomic Theory Where did the idea of the atom come from? Let's travel back thousands of years to find out. Ancient Ideas: Kanada, Leucippus & Democritus Acharya Kanada (India) Matter divided repeatedly reaches the smallest indivisible particle — the parmanu, recorded in the Vaisesika Sutras. Leucippus & Democritus (Greece) Called these particles atomos — Greek for "indivisible." Philosophical Origins Both ideas arose as philosophical concepts, not from experiments — science had not yet begun. Dalton's Atomic Theory (1808) First scientific description of atoms — based on experimental evidence, not philosophy Proposed: all matter is composed of indivisible particles called atoms Became the starting point for understanding atomic structure Raised new questions: What are atoms made of? What do they look like? Dalton's theory was revolutionary — it moved the atom from philosophy into science for the first time. CHAPTER 8.2 A Journey Through Atomic Models From Dalton's solid sphere to Bohr's energy levels — scientists gradually uncovered the true structure of the atom. Discovery of the Electron — J. J. Thomson (1897) Studied electric current through gases at very low pressure using a cathode ray tube Observed rays moving from cathode (−) to anode (+) — called cathode rays Concluded: streams of negatively charged particles (electrons) with mass much smaller than atoms Charge of electron: −1.602 × 10⁻¹⁹ C (taken as −1 by convention) Thomson's Plum Pudding Model Thomson proposed the first model of atomic structure to explain what was inside the atom. Atom = a sphere of positive charge with electrons distributed throughout Nicknamed the "plum pudding model" — like a watermelon: red pulp = positive charge, seeds = electrons First genuine attempt to explain how positive and negative charges stay balanced The Gold Foil Experiment (1911) The Setup Geiger & Marsden (under Rutherford) fired α-particles at a thin gold foil Expected Particles pass straight through — consistent with Thomson's spread-out model Observed! Most passed through, but some sharply deflected — a few even bounced back! Conclusion Thomson's model could NOT explain these astonishing results Rutherford's Planetary Model Positive charge is concentrated in a tiny, dense nucleus — not spread out Most of an atom is empty space — that's why most α-particles passed through undefl

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Journey Inside the Atom Chapter 8 — Grade 9 Science Prepared by K Sandeep Swamy (M.Sc, B.Ed) For ONLINE Classes (ACADEMIC,IITJEE & NEET ) Contact 9491878325 Subscribe samyans Eduhub youtube channel
Everything Around Us Is Made of Atoms All Matter Every living and non-living thing is ultimately made of atoms — the fundamental building blocks of matter. Incredibly Tiny Atoms are so small they cannot be seen with the naked eye — even the most powerful optical microscopes struggle to observe them. The Big Question Is the atom truly the smallest unit of matter, or can it be divided into even smaller particles?
CHAPTER 8.1 Roots of Atomic Theory Where did the idea of the atom come from? Let's travel back thousands of years to find out.
Ancient Ideas: Kanada, Leucippus & Democritus Acharya Kanada (India) Matter divided repeatedly reaches the smallest indivisible particle — the parmanu, recorded in the Vaisesika Sutras. Leucippus & Democritus (Greece) Called these particles atomos — Greek for "indivisible." Philosophical Origins Both ideas arose as philosophical concepts, not from experiments — science had not yet begun.
Dalton's Atomic Theory (1808) First scientific description of atoms — based on experimental evidence, not philosophy Proposed: all matter is composed of indivisible particles called atoms Became the starting point for understanding atomic structure Raised new questions: What are atoms made of? What do they look like? Dalton's theory was revolutionary — it moved the atom from philosophy into science for the first time.
CHAPTER 8.2 A Journey Through Atomic Models From Dalton's solid sphere to Bohr's energy levels — scientists gradually uncovered the true structure of the atom.
Discovery of the Electron — J. J. Thomson (1897) Studied electric current through gases at very low pressure using a cathode ray tube Observed rays moving from cathode (−) to anode (+) — called cathode rays Concluded: streams of negatively charged particles (electrons) with mass much smaller than atoms Charge of electron: −1.602 × 10⁻¹⁹ C (taken as −1 by convention)
Thomson's Plum Pudding Model Thomson proposed the first model of atomic structure to explain what was inside the atom. Atom = a sphere of positive charge with electrons distributed throughout Nicknamed the "plum pudding model" — like a watermelon: red pulp = positive charge, seeds = electrons First genuine attempt to explain how positive and negative charges stay balanced
The Gold Foil Experiment (1911) The Setup Geiger & Marsden (under Rutherford) fired α-particles at a thin gold foil Expected Particles pass straight through — consistent with Thomson's spread-out model Observed! Most passed through, but some sharply deflected — a few even bounced back! Conclusion Thomson's model could NOT explain these astonishing results
Rutherford's Planetary Model Positive charge is concentrated in a tiny, dense nucleus — not spread out Most of an atom is empty space — that's why most α-particles passed through undeflected Electrons revolve around the nucleus like planets around the Sun Nucleus is ~10⁵ times smaller than the atom Atom ≈ 10⁻¹⁰ m | Nucleus ≈ 10⁻¹⁵ m
Why Rutherford's Model Failed Energy Loss Problem An accelerating electron should continuously lose energy and spiral inward toward the nucleus. Atom Collapse? If electrons spiralled in, all atoms would collapse — but matter around us is clearly stable! A New Model Was Needed Rutherford's model could not explain atomic stability. Science needed a breakthrough…
Bohr's Model of the Atom (1913) Fixed Shells Electrons move in fixed circular paths called shells / orbits / energy levels (K, L, M, N… or n = 1, 2, 3, 4…) No Energy Loss While in a fixed shell, an electron does not radiate or lose energy — solving Rutherford's stability problem Energy Jumps Electrons absorb or release energy when jumping between shells — energy increases farther from the nucleus
Evolution of Atomic Models at a Glance 1 Dalton Indivisible solid sphere 2 Thomson Plum pudding model 3 Rutherford Nuclear planetary model 4 Bohr Fixed energy levels 5 Modern Quantum mechanical model Science advances step by step — each model improved on the last, revealing deeper truths about atomic structure.
CHAPTER 8.3 Subatomic Particles If atoms can be broken apart, what are they made of? Meet the three fundamental subatomic particles.
Discovery of the Proton & Neutron Proton (Rutherford) Positive charge in the nucleus; charge = +1; much heavier than the electron. Neutron (Chadwick, 1932) Neutral particle in the nucleus; mass ≈ proton; Chadwick won the Nobel Prize in 1935. Why Neutrons Matter Helium has 2 protons but 4× the mass of hydrogen — because it also has 2 neutrons! Particle Symbol Relative Charge Electron e⁻ −1 Proton p⁺ +1 Neutron n⁰ 0
CHAPTERS 8.4 – 8.6 Symbols, Atomic Number & Mass Number How do scientists label elements and count what's inside the nucleus?
Symbols of Elements 01 Pictorial Symbols (Dalton, 1803) Dalton introduced simple pictures to represent different elements — the first attempt at standardisation. 02 Alphabetic Symbols (Berzelius, 1813) Berzelius proposed using letters to represent elements — far more practical and universal. 03 IUPAC Standards (Today) Rules: first letter uppercase, second letter lowercase — e.g., Al, Co, Cl. Some symbols come from Latin/German: Fe (ferrum), Hg (hydrargyros), W (wolfram).
Atomic Number (Z) & Mass Number (A) Atomic Number (Z) = Number of protons in the nucleus. Uniquely identifies an element. Since atoms are neutral: protons = electrons. Mass Number (A) = Total number of nucleons (protons + neutrons) in the nucleus. Key Formula: A = p⁺ + n⁰ Therefore: Neutrons = A − Z
CHAPTER 8.7 Electron Distribution in Energy Levels How are electrons arranged inside an atom? The Bohr-Bury rules give us a clear set of guidelines.
Rules for Filling Electron Shells (Bohr-Bury Rules) 1 Maximum electrons per shell = 2n² K shell (n=1): max 2 electrons | L shell (n=2): max 8 | M shell (n=3): max 18 2 Outermost shell maximum = 8 No matter which shell is outermost, it can hold a maximum of 8 electrons (K- shell exception: max 2). 3 Shells fill in order: K → L → M → N Inner shells must be filled completely before electrons occupy the next shell.
Electronic Configuration: Examples Hydrogen (Z = 1) K = 1 Configuration: 1 Carbon (Z = 6) K = 2, L = 4 Configuration: 2, 4 Sodium (Z = 11) K = 2, L = 8, M = 1 Configuration: 2, 8, 1 Argon (Z = 18) K = 2, L = 8, M = 8 Configuration: 2, 8, 8 Notice how Argon (Z=18) has a complete outermost shell (8 electrons) — making it an extremely stable, unreactive noble gas!
CHAPTER 8.8 Valency — Combining Capacity of an Atom Why do some atoms bond with one partner while others bond with four? Valency holds the answer.
Valence Electrons & Valency The valence shell is the outermost shell, and its electrons are called valence electrons. Atoms strive to achieve a complete octet (8 electrons) or duplet (2 for He) for stability. Element Config. Valence e⁻ Valency Sodium (Na) 2, 8, 1 1 1 (loses 1e⁻) Oxygen (O) 2, 6 6 2 (gains 2e⁻) Carbon (C) 2, 4 4 4 (shares 4e⁻) Neon (Ne) 2, 8 8 0 (stable octet) Lose electrons Metals like Na lose valence electrons — valency = electrons lost Gain electrons Non-metals like O gain electrons — valency = electrons gained Share electrons Carbon shares electrons — valency = electrons shared
CHAPTER 8.9 A Deeper Look: Isotopes, Isobars & Average Atomic Mass Not all atoms of the same element are identical — and sometimes atoms of different elements share the same mass number.
Isotopes — Same Element, Different Mass Isotopes are atoms of the same element with the same atomic number (Z) but different mass numbers (A). The three isotopes of hydrogen: Isotopes have the same chemical properties (same valence electrons) but different physical properties. Protium ¹H ~99.98% — the most common form Deuterium ²H ~0.015% — used in heavy water Tritium ³H Trace amounts — radioactive
Applications of Isotopes ²³uU — Nuclear Energy Fuel in nuclear reactors to generate electricity safely and efficiently. v⁰Co — Cancer Treatment Radiation therapy for treating cancer tumours using cobalt-60. ¹³¹I — Thyroid Treatment Treatment of goitre and thyroid cancer — absorbed selectively by the thyroid gland. ¹tC — Carbon Dating Determines the age of ancient fossils and artefacts by measuring radioactive decay.
Average Atomic Mass — The Chlorine Example Why 35.5 u? Chlorine has two naturally occurring isotopes with different abundances in nature: Isotope Mass Abundance ³⁵Cl 35 u 75% ³⁷Cl 37 u 25% Simple average = (35 + 37) ÷ 2 = 36 u — this ignores abundance and is inaccurate! Weighted Average (Correct Method) (35 × 75/100) + (37 × 25/100) = 26.25 + 9.25 = 35.5 u The weighted average accurately reflects the mass of chlorine as it occurs in nature.
Isobars — Same Mass, Different Elements Isobars are atoms of different elements that share the same mass number (A) but have different atomic numbers (Z). Argon (Z = 18) Mass Number = 40 18 protons + 22 neutrons Potassium (Z = 19) Mass Number = 40 19 protons + 21 neutrons Calcium (Z = 20) Mass Number = 40 20 protons + 20 neutrons Unlike isotopes, isobars are completely different elements with different chemical properties — the same mass number is just a coincidence of nuclear composition.
Key Takeaways: Journey Inside the Atom 01 Three Subatomic Particles Atoms are made of electrons (−1), protons (+1), and neutrons (0) — each with a specific location and role. 02 Atomic Number & Mass Number Z = protons (identifies the element); A = protons + neutrons (total nucleons in nucleus). 03 Electron Shell Rules Shells fill in order K→L→M; max in outermost shell = 8; each shell holds 2n² electrons. 04 Valency = Electrons gained, lost, or shared to complete the octet and achieve stability. 05 Isotopes vs Isobars Isotopes: same Z, different A | Isobars: same A, different Z — two very different concepts! 06 Science Is Ongoing! The quantum mechanical model awaits you in higher grades — the atom still has secrets to reveal! 🔬