Structure of the Atom — Class 9 (CBSE)

Dalton said atoms were the smallest possible particles. Within a hundred years, three physicists had broken open the atom and found smaller things inside. This chapter is the story of how — and how the inside of an atom turned out to be 99.99 % empty space.

1. The story — how we learned what's inside an atom

In Dalton's time (1808), the atom was considered the smallest indivisible particle. By 1932, three landmark experiments had revealed three sub-atomic particles inside it.

• 1897 — J.J. Thomson discovers the electron using cathode ray tubes.
• 1909 — Rutherford's gold-foil experiment reveals the nucleus, a tiny dense positive core.
• 1932 — James Chadwick discovers the neutron, the missing-mass particle.

Each experiment overturned the previous picture of the atom. The chapter is structured around these three discoveries and the models of the atom that emerged from them.

2. Sub-atomic particles — memorise the table

Key takeaways:

• Electrons are nearly massless compared with protons & neutrons.
• Protons & neutrons together (called nucleons) account for ~ 99.9 % of an atom's mass.
• Charges balance: an atom has equal protons and electrons → overall neutral.

3. Thomson's plum-pudding model (1904)

After discovering the electron, Thomson proposed:

• An atom is a sphere of uniform positive charge with electrons embedded in it like plums in a pudding (or seeds in a watermelon).
• The positive charge equals the total negative charge → atom is neutral.

It was the first attempt to picture the atom's interior. But it had a problem: nothing in the model predicted any specific behaviour or could be tested. That changed when Rutherford fired alpha particles at gold foil.

4. Rutherford's gold-foil experiment (1911)

The experiment

Rutherford and his students (Geiger and Marsden) shot fast-moving alpha particles (helium nuclei: 2 protons + 2 neutrons, charge +2) at a thin sheet of gold foil only a few atoms thick. Behind the foil, a fluorescent screen detected where each particle landed.

Observations

1. Most α-particles passed straight through with no deflection.
2. Some α-particles were deflected by small angles.
3. A very few α-particles (≈ 1 in 8000) bounced back at angles > 90°.

Rutherford famously said this last observation was "as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you."

Conclusions (the new atomic model)

1. Most of an atom is empty space — that's why most α-particles went straight through.
2. The atom has a tiny, dense, positively charged nucleus at its centre — that's what bounced α-particles back.
3. Electrons revolve around the nucleus like planets around the sun.

Size estimates from Rutherford's model

• Atomic radius: ≈ $10^{-10}$ m (1 Å).
• Nuclear radius: ≈ $10^{-15}$ m (1 fm).
• The atom is 100,000 times bigger than its nucleus.

If a nucleus were the size of a marble, the atom would be the size of a football stadium. The rest is electrons + empty space.

Drawbacks of Rutherford's model

• Classical physics predicted that orbiting electrons (being accelerated) should radiate energy and spiral into the nucleus within $10^{-8}$ s — but atoms are stable! Something was incomplete.
• Couldn't explain why hot gases emit specific colours (line spectra), not a continuous rainbow.

These problems opened the door for Bohr.

5. Bohr's model (1913)

Bohr fixed Rutherford's model with three postulates:

1. Electrons orbit the nucleus in fixed paths called shells or orbits (also called energy levels) — labelled K, L, M, N, … (or n = 1, 2, 3, 4, …).
2. Each shell has a specific energy. While an electron is in a shell, it does NOT radiate energy.
3. Energy is absorbed/emitted only when an electron jumps between shells — absorbed (going outward) or emitted (coming inward).

This explained why atoms are stable AND why spectral lines exist (each transition between shells corresponds to a specific energy = specific colour of light).

Maximum electrons per shell — Bohr-Bury rule

$\text{Max electrons in shell }n=2n^2$

Additional rules:

• The outermost shell can have a maximum of 8 electrons (the octet rule — only applies to outermost).
• An inner shell must be filled completely before electrons can fill the next outer shell.

Electronic configuration — how to write it

Example — Sodium (Z = 11):

• 11 electrons.
• K shell: 2 (filled).
• L shell: 8 (filled).
• M shell: 1 (left over).
• Configuration: 2, 8, 1.

Example — Chlorine (Z = 17):

• K: 2, L: 8, M: 7.
• Configuration: 2, 8, 7.

Example — Argon (Z = 18):

• K: 2, L: 8, M: 8.
• Configuration: 2, 8, 8 (filled — that's why argon is a noble gas, unreactive).

6. Atomic number and mass number

Two important integers describe any atom:

Atomic number ($Z$) = number of protons in the nucleus. (Also equals number of electrons in a neutral atom.) This identifies the element. Carbon is Z = 6. Always. Change Z and you have a different element.

Mass number ($A$) = number of protons + number of neutrons = total nucleons.

So: number of neutrons = $A-Z$.

Standard notation:

${}_Z^{A}X$

For example, ${}_6^{12}C$ means: carbon, atomic number 6, mass number 12. So 6 protons + 6 neutrons + 6 electrons.

Memorise the first 20 elements (in order)

This table is the single most useful thing to memorise in Class 9 chemistry. Most periodic-table questions, valency questions, ion-formation questions trace back here.

7. Valency from electron configuration

Valency = combining capacity of an atom = the number of electrons it needs to gain, lose, or share to reach a noble-gas configuration (full outermost shell, typically 8 electrons).

Three rules:

1. If outermost electrons ≤ 4 → valency = outermost electron count (atom LOSES them).
2. If outermost electrons > 4 → valency = 8 − outermost electron count (atom GAINS them).
3. If outermost shell is full (2 for He, 8 for others) → valency = 0 (inert / noble).

Examples

• Sodium (2, 8, 1): outermost = 1, so valency = 1. Tends to lose 1 e⁻ → Na⁺.
• Magnesium (2, 8, 2): outermost = 2, valency = 2. Loses 2 e⁻ → Mg²⁺.
• Aluminium (2, 8, 3): outermost = 3, valency = 3. Loses 3 e⁻ → Al³⁺.
• Carbon (2, 4): outermost = 4, valency = 4. Shares (in CH₄, CO₂).
• Nitrogen (2, 5): outermost = 5, valency = 8 − 5 = 3. Gains 3 (or shares) → N³⁻ / NH₃.
• Oxygen (2, 6): outermost = 6, valency = 8 − 6 = 2. Gains 2 → O²⁻.
• Chlorine (2, 8, 7): outermost = 7, valency = 8 − 7 = 1. Gains 1 → Cl⁻.
• Neon (2, 8): outermost = 8, valency = 0 (noble gas, inert).

8. Isotopes and isobars

Isotopes

Isotopes are atoms of the same element (same Z, same number of protons) but different mass numbers (different number of neutrons).

Examples:

• Hydrogen has 3 isotopes:

  • Protium: ${}_1^1\text{H}$ — 1 proton, 0 neutrons (most common).
  • Deuterium: ${}_1^2\text{H}$ — 1 proton, 1 neutron (heavy water D₂O).
  • Tritium: ${}_1^3\text{H}$ — 1 proton, 2 neutrons (radioactive, used in nuclear fusion).

• Carbon has 3 main isotopes:

  • ${}_6^{12}C$ — 6n (most common, ~99 %).
  • ${}_6^{13}C$ — 7n (~1 %).
  • ${}_6^{14}C$ — 8n (radioactive, used in carbon dating).

• Chlorine has 2 isotopes:

  • ${}_{17}^{35}Cl$ — 18n (~75 %).
  • ${}_{17}^{37}Cl$ — 20n (~25 %).
  • The "average atomic mass" 35.5 = 0.75 × 35 + 0.25 × 37.

Important: Isotopes have IDENTICAL chemical properties (because chemistry depends on electron config, which only depends on Z). But they have slightly different PHYSICAL properties (mass, density, BP — heavy water boils at 101.4 °C vs 100 °C for regular water).

Applications of isotopes

• Carbon-14: dating fossils and archaeological samples (half-life 5730 y).
• Iodine-131: treatment of hyperthyroidism (radioactive iodine concentrates in the thyroid).
• Cobalt-60: cancer radiation therapy.
• Uranium-235: nuclear fuel (fission reactor / atom bomb).
• Deuterium: nuclear fusion research, moderator in heavy-water reactors.

Isobars

Isobars are atoms of different elements (different Z) but the same mass number (same A).

Examples:

• ${}_{18}^{40}\text{Ar}$ and ${}_{20}^{40}\text{Ca}$. Both have A = 40 but different Z (18 vs 20).
• ${}_6^{14}\text{C}$ and ${}_7^{14}\text{N}$. Both have A = 14 but different Z.

Memorising:

• Isotopes: same protons (top number Z), different neutrons.
• Isobars: same A (mass), different elements.

9. Closing thought

The chapter is a journey from one wrong picture to another, each less wrong than the last:

1. Dalton (1808): solid indivisible billiard ball.
2. Thomson (1904): plum pudding (positive sphere with embedded electrons).
3. Rutherford (1911): tiny nucleus + electrons orbiting in empty space.
4. Bohr (1913): electrons in fixed energy shells.
5. Modern (1925+): electrons as 3D probability clouds (you'll meet quantum-mechanical orbitals in Class 11).

Each model explained more experimental data than the last while still being partly wrong. That's how science actually advances — by improving the picture, not by leaping to perfect truth. This is the most important meta-lesson of the chapter.

You can now look at any element on the periodic table and, from its atomic number alone, write its electronic configuration, predict its valency, guess its likely ion, and reason about its chemistry. That's an enormous superpower for someone who couldn't tell sodium from potassium two months ago.