Structure Of Atoms Short Notes Class 9, Easy To Memorise!

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“Structure Of Atoms Short Notes Class 9” has all the important points and concepts from your NCERT chapter 4 Structure of atom.

This short notes will work best If you have read the chapter from your NCERT 5 to 10 times. This short notes are meant for memorisation and revision.

In no sense this notes are replacement of your NCERT textbook.

How Objects Get Charged

  • Rubbing two objects transfers charged particles between them.
  • This transfer makes objects electrically charged.
  • Charge comes from tiny particles inside atoms.

Are Atoms Divisible?

  • Yes, Atoms are divisible.
  • They contain smaller charged particles.

Key Discoveries

1886

E. Goldstein found canal rays in gas discharge.

  • These were positively charged rays.
  • Led to the discovery of the proton

1900

J.J Thomson identified the electron.

  • First known sub-atomic particle.
  • Negatively charged.

Properties of Electrons & Protons

ParticleSymbolChargeMass (approx.)Location in Atom
Electrone⁻–1Very small ≈, 1/2000 of a protonOuter region
Protonp⁺+11 unit =(2000× electron)Interior (nucleus)

Early Atomic Model

  • Atoms contain protons (+) and electrons (–).
  • Their charges balance → atom is neutral overall.
  • Electrons can be removed easily (e.g., by rubbing).
  • Protons stay fixed inside the atom.
Keyword Table
keyword Meaning
ChargeElectrical property causing attraction/repulsion
ElectronTiny negatively charged particle in atoms
ProtonPositively charged particle inside atom’s core
Canal raysPositive rays in gas tubes; led to proton discovery
Sub-atomicSmaller than an atom (e.g., electron, proton)
NeutralNo net charge (positive = negative charges)
AtomBasic unit of matter; contains protons & electrons

Atomic Models

Thomson’s Model of an Atom

Plum Pudding Model or watermelon model.

Main Idea

  • An atom = a sphere of positive charge.
  • Electrons are embedded inside like seeds or currants.

Two Key Rules

  • A positive charge is evenly distributed throughout the atom.
  • Total negative charge (electrons) = total positive charge
    • An atom is electrically neutral overall.

Explained atom neutrality.
Could not explain the later experiments.

Rutherford’s Model of an Atom

Structure Of Atoms Short Notes Class 9

The Gold Foil Experiment

  • Rutherford fired fast-moving alpha (α) particles at a thin gold foil.
  • Gold foil was only 1000 atoms thick – extremely thin.
  • Alpha particles = helium ions (charge +2, mass 4 u).

3 Key Observations

  • Most α-particles passed straight through → the atom is mostly an empty space.
  • Some α-particles are deflected at small angles → positive charge exists in a small region.
  • 1 in 12,000 particles bounced back (180°) → dense, heavy core present.

3 Major Conclusions

  • Most space inside an atom is empty.
  • Positive charge is concentrated in a tiny central region.
  • The nucleus holds nearly all the atom’s mass but is 100,000× smaller than the atom.

Nuclear Model Features

  • An atom has a positively charged nucleus at its center.
  • Electrons revolve around the nucleus in circular paths(orbits).
  • Nucleus size ≪ atom size (like a marble in a football stadium).

Drawback of the Model

  • Revolving electrons should lose energy and spiral into the nucleus.
  • This would make atoms unstable – but real atoms are stable.
  • → The model could not explain atomic stability.
Keyword Table
KeywordMeaning
Alpha (α) particleFast-moving helium ion (+2 charge, mass 4 u)
Gold foilExtremely thin metal sheet (≈1000 atoms thick) used in experiment
NucleusTiny, dense, positively charged center of the atom
DeflectionChange in path of α-particles due to repulsion from positive charge
ScatteringSpreading of α-particles in different directions after hitting foil
RevolveSpreading of α-particles in different directions after hitting the foil
Comparison: Thomson’s vs Rutherford’s Model
Revolve around the nucleusThomson’s Model (Plum Pudding)Rutherford’s Model (Nuclear)
Positive chargeSpread evenly throughout atomConcentrated in tiny nucleus
Electron positionEmbedded in positive sphereRevolve around nucleus
Atom structureSolid, uniform sphereMostly an empty space
Explains scattering?❌ No – cannot explain rebounding α-particles✅ Yes – explains all observations
Stability issueNot addressedMostly space

Also Read | Atoms And Molecules Short Notes

Bohr’s Model Of Atom

Why Bohr Proposed a New Model

  • Fixed Rutherford’s stability problem.
  • Explained why electrons don’t fall into the nucleus.
  • Introduced the idea of fixed energy levels.

Postulates of Bohr’s Atomic Model

  • Electrons move only in special fixed orbits (called discrete orbits).
  • While in these orbits, electrons do not lose energy.
  • → Atom stays stable.

Energy Levels Or Shells

  • Discrete orbits = energy levels or shells.
  • Named using letters: K, L, M, N…
  • Or numbers: n = 1, 2, 3, 4…
    • K-shell = n=1 (closest to nucleus, lowest energy)
    • L-shell = n=2
    • M-shell = n=3, and so on.

How It Solved Rutherford’s Problem

  • Electrons only radiate energy when jumping between orbits.
  • While staying in one orbit, no energy loss → no collapse into nucleus.
  • → Explains atomic stability.
Keyword Table
KeywordMeaning
Discrete orbitsFixed, special paths where electrons can move without losing energy
Energy levelsShells (K, L, M…) where electrons stay at specific energy values
Radiate energyRelease energy as light/heat (electrons do this only when changing orbits)
Stable atomAtom that does not collapse – electrons stay in fixed orbits
n = 1, 2, 3…Numbers representing energy levels (n=1 = innermost shell)
PostulateBasic rule or assumption in a scientific model
↔️ Comparison: Atomic Models Timeline
FeatureThomson (1897)Rutherford (1911)Bohr (1913)
Positive chargeSpread evenlyIn tiny nucleusIn tiny nucleus
Electron motionEmbedded (fixed)Revolve (unstable)Revolve in fixed orbits
Energy loss?Not discussedYes → should collapseNo loss in orbits → stable
Key ideaPlum puddingNuclear atomQuantized energy levels
Explains stability?❌❌In a tiny nucleus

How are Electrons Distributed in Different Orbits (Shells)?

Rule 1
Maximum Electrons per Shell

  • Formula: 2n²
    (n = shell number: 1, 2, 3…)
    • K-shell (n=1): 2 × 1² = 2 electrons
    • L-shell (n=2): 2 × 2² = 8 electrons
    • M-shell (n=3): 2 × 3² = 18 electrons
    • N-shell (n=4): 2 × 4² = 32 electrons

Rule 2
Outermost Shell Limit
(of electrons)

  • Max 8 electrons allowed in the outermost shell.
  • Called the octet rule.
  • Ensures atom stability.

Rule 3
Step-wise Filling

  • Inner shells must fill first.
  • Electrons go to the next shell only when the inner shell is full.
  • Example: L-shell fills only after the K-shell has 2 electrons.

Example: Oxygen (8 electrons)

  • K-shell: 2 electrons (full)
  • L-shell: 6 electrons (not full – needs 2 more for octet)

💡 Key Point: A shell can hold up to 2n² electrons, but if it is the outermost shell, it holds max 8 – even if 2n² allows more.

Comparison: Electron Capacity in Shells
Shell Namen-valueMax Electrons (2n²)Outermost Limit?
K-shelln = 12❌ No (inner shell)
L-shelln = 28✅ Yes (if outermost)
M-shelln = 318✅ Yes – but max 8 if outermost
N-shelln = 432✅ Yes – but max 8 if outermost

Valency

Valence Electrons

  • Electrons in the outermost shell of an atom.
  • Decide how an atom bonds with others.

Octet Rule (Stability Rule)

  • Atoms are stable when the outermost shell has 8 electrons.
  • Called a full octet.
  • Helium exception: stable with 2 electrons (only K-shell).

What is Valency?

  • The combining capacity of an atom.
  • The number of electrons it gains, loses, or shares to get a full outer shell.

How Atoms Achieve Octet

  • Lose electrons → if 1–3 in outer shell (e.g., Na loses 1).
  • Gain electrons → if 5–7 in outer shell (e.g., F gains 1).
  • Share electrons → in covalent bonds (e.g., H₂).
Calculating Valency
Outer ElectronsActionValency FormulaExample
1, 2, or 3Lose them= numberNa (1) → valency 1
5, 6, or 7Gain to 8= 8 – numberO (6) → valency 2
4Share= 4C → valency 4
8 (or 2 for He)None= 0Ne → valency 0
Examples
  • Hydrogen (H): 1 outer e⁻ → loses 1 → valency = 1
  • Magnesium (Mg): 2 outer e⁻ → loses 2 → valency = 2
  • Aluminium (Al): 3 outer e⁻ → loses 3 → valency = 3
  • Oxygen (O): 6 outer e⁻ → gains 2 → valency = 8 – 6 = 2
  • Fluorine (F): 7 outer e⁻ → gains 1 → valency = 8 – 7 = 1
  • Neon (Ne): 8 outer e⁻ → stable → valency = 0
Keyword Table
KeywordMeaning
Valence electronsElectrons in the outermost shell
ValencyHow many bonds does an atom form
Octet8 electrons in outermost shell (stable state)
InertUnreactive; full outer shell (valency = 0)
Combining capacityHow many bonds an atom can form
Shell fillingAtoms react to fill outer shell to 8 (or 2 for He)
Comparison: How Valency is Determined
SituationActionValency CalculationExample
Few outer electrons (1–3)Lose e⁻= number of e⁻Na → 1
Many outer electrons (5–7)Gain e⁻= 8 – number of e⁻Cl → 1
Half-full shell (4)Share e⁻= 4C → 4
Full shell (8 or 2 for He)No reaction= 0Ar → 0

Atomic Number and Mass Number

ATOMIC NUMBER

  • Atomic number = number of protons in the nucleus.
  • Denoted by the symbol ‘Z’.
  • Unique for each element – defines the element itself.

Key Facts

  • All atoms of the same element have the same atomic number (Z).
  • Change protons → change the element.
  • Hydrogen: Z = 1 (1 proton).
  • Carbon: Z = 6 (6 protons).
  • Neutral atom:
    Number of protons (Z) = Number of electrons.

Why It Matters

  • Determines element identity.
  • Never changes for a given element.

MASS NUMBER

  • Mass number = protons + neutrons in the nucleus.
  • Denoted by the symbol ‘A’.
  • Electrons have negligible mass → not counted.

Why Only Protons & Neutrons?

  • Both particles are heavy (~1 u each).
  • Together called nucleons (reside in the nucleus).
  • Electrons are ~2000× lighter → ignored for mass.


Examples

  • Carbon: 6 protons + 6 neutrons → A = 12
  • Aluminium: 13 protons + 14 neutrons → A = 27
  • Nitrogen: 7 protons + 7 neutrons → A = 14

Atomic Notation

  • Written as: ZAX{}^{A}_{Z}X
  • Example: 714N{}^{14}_{7}N → Nitrogen with A=14, Z=7

Key Rule

  • A = Z + number of neutrons
  • Neutrons = A – Z
Keyword Table
KeywordMeaning
Mass numberTotal protons + neutrons in nucleus (symbol A)
NucleonsCollective name for protons and neutrons
Atomic mass unit (u)Unit for atomic mass; proton/neutron ≈ 1 u
NeutronNeutral particle in nucleus; adds mass but no charge
NucleusCentral core holding protons and neutrons (99.9% of atom’s mass)

Isotopes

What Are Isotopes?

  • Atoms of the same element (same atomic number Z).
  • Have different mass numbers (A) due to different neutrons.
  • Same protons → same chemical behavior.
  • Different neutrons → different physical properties.

Key Properties

  • Chemical properties: Same (same electrons → same reactions).
  • Physical properties: Different (mass, density, boiling point).
  • Average atomic mass: Weighted average of all isotopes.
    • Chlorine example:
      (75% × 35 u) + (25% × 37 u) = 35.5 u
      → Not the mass of one atom, but the average of a large sample.

Hydrogen Isotopes (Best Example)

IsotopeSymbolProtonsNeutronsMass (A)Common Name
Protium¹H₁101Ordinary H
Deuterium²H₁ or D112Heavy H
Tritium³H₁ or T123Radioactive H
Other Examples
  • Carbon: ¹²C₆ (6n) and ¹⁴C₆ (8n) → used in carbon dating.
  • Chlorine: ³⁵Cl₁₇ (18n) and ³⁷Cl₁₇ (20n) → occur in 3:1 ratio.

Real-World Applications

  • Uranium-235: Nuclear reactor fuel.
  • Cobalt-60: Cancer radiation therapy.
  • Iodine-131: Treats thyroid disorders (goitre).
Keyword Table
KeywordMeaning
IsotopesAtoms of same element with same protons but different neutrons
Atomic number (Z)Number of protons (defines the element)
Mass number (A)Protons + neutrons in nucleus
NucleonsProtons + neutrons (particles in nucleus)
Average atomic massWeighted mean mass of all natural isotopes
ProtiumLightest hydrogen isotope (no neutrons)
DeuteriumHeavy hydrogen (1 neutron); used in heavy water
TritiumAtoms of the same element with the same protons but different neutrons

Isobars

What Are Isobars?

  • Atoms of different elements (different atomic numbers Z).
  • Have the same mass number (A).
  • Same total nucleons (protons + neutrons), but different proton/neutron split.

Key Properties

  • Chemical properties: Different (different elements → different reactions).
  • Physical properties: Different (mass distribution varies).
  • Nucleon count: Same total (A), but protons ≠ and neutrons in each.

Why They Exist

  • Different combinations of protons/neutrons can give the same total mass.
  • Example: 20p+20n = 40 and 18p+22n = 40 → both have A=40.

Example: Calcium & Argon

ElementSymbolProtons (Z)NeutronsMass Number (A)
Calcium⁴⁰Ca₂₀202040
Argon⁴⁰Ar₁₈182240

→ Same A = 40, but different elements (Z = 20 vs 18).

Keyword Table
KeywordMeaning
IsobarsAtoms of different elements with same mass number (A) but different atomic number (Z)
Mass number (A)Total protons + neutrons in nucleus
Atomic number (Z)Number of protons (defines the element)
NucleonsProtons + neutrons (particles in nucleus)
Different elementsAtoms with different Z → different chemical identity
↔️ Comparison: Isotopes vs Isobars
FeatureIsotopesIsobars
Atomic number (Z)Same (same element)Different (different elements)
Mass number (A)DifferentSame
ProtonsSameDifferent
NeutronsDifferentDifferent (adjusted to keep A same)
Chemical propertiesSame (same electrons)Different (different elements)
Example¹²C₆ and ¹⁴C₆ (both carbon)⁴⁰Ca₂₀ and ⁴⁰Ar₁₈ (calcium & argon)

💡 Quick Memory Trick:
Isotopes → same top (protons/Z)
Isobars → same weight (mass/A)

Conclusion : Structure Of Atoms Short Notes Class 9

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