Chapter 8: Journey Inside the Atom
Complete Study Notes Atomic Models, Subatomic Particles & Isotopes Exam-Ready Format
1. The Ancient Origins of Atomic Theory
Over 2,000 years ago, thinkers in both ancient India and ancient Greece independently asked: What is everything made up of?
Acharya Kanada (India)
Proposed that matter (dravya) divided repeatedly reaches a stage of smallest, indivisible particles called parmanus, recorded in the Vaisesika Sutras. Parmanus combine to form dyads (pairs) and triads (groups of three), building up the entire material universe.
Leucippus & Democritus (Greece)
Proposed similar indivisible particles called atomos (Greek for “indivisible”) — the origin of the modern word “atom”.
Following Dalton, scientists sought answers to: What are atoms made of? What would they look like? What makes atoms of one element different from another?
2. Evolution of Atomic Models
Discovery of radioactivity (certain elements emitting invisible energy/particles called radiation) in the late 19th century proved that atoms are not indivisible — they must contain smaller particles.
2.1 Thomson’s Model (1897) — “Plum Pudding”
J. J. Thomson studied conduction of electricity through gases at very low pressure using a cathode ray tube (two electrodes in a glass tube with high voltage). He observed rays moving from cathode (–) to anode (+), called cathode rays. Studying their behaviour in electric/magnetic fields, he concluded these are streams of negatively charged particles — much lighter than atoms — later named electrons.
Key finding: the nature of cathode rays was independent of the cathode material and the gas used — proving electrons are a fundamental component of all atoms. Charge of an electron = $-1.602 \times 10^{-19}\,\text{C}$, taken as $-1$ by convention.
Since atoms are neutral overall, Thomson had to explain where the positive charge sits. He proposed the atom as a sphere of positive charge with electrons embedded in it — like a pudding studded with plums, or seeds in a watermelon.
2.2 Testing Thomson’s Model — The Gold Foil Experiment (1911)
A beam of alpha (α) particles (positively charged; later shown to be helium nuclei — 2 protons + 2 neutrons) was fired at an extremely thin gold foil.
Expected (per Thomson’s model): particles pass straight through or deflect only slightly, since positive charge is spread evenly.
Observed: most particles passed straight through undeflected; some were sharply deflected; a few bounced straight back. This deflection is called scattering — hence also called the α-ray scattering experiment.
2.3 Rutherford’s Nuclear (Planetary) Model
Based on the gold foil results, Rutherford proposed:
- Most of an atom is empty space
- All positive charge and most mass is concentrated in a tiny, dense central region — the nucleus
- Electrons revolve around the nucleus like planets around the Sun — the planetary model
Discovery of the Proton
Rutherford showed the nucleus carries positive charge from particles called protons — much heavier than electrons, with charge equal and opposite to an electron’s. For a neutral atom: number of protons = number of electrons.
2.4 Bohr’s Model (1913) — Fixed Energy Levels
Niels Bohr proposed that electrons:
- Move only in fixed circular paths called stationary states, orbits, or shells — not randomly
- Each shell has a definite energy — hence called energy levels, labelled K, L, M, N… (or $n = 1, 2, 3, 4…$)
- Do not lose energy while moving in a fixed shell (this is the key postulate that fixes Rutherford’s stability problem)
- K-shell ($n=1$, closest to nucleus) has the least energy; energy increases with distance from the nucleus
- An electron changes shells only by absorbing or releasing energy exactly equal to the difference between the two shells’ energies
- Each shell holds only a limited number of electrons
Dalton (1808)
Solid indivisible sphere
Thomson (1897)
Plum pudding model
Rutherford (1911)
Nucleus + empty space
Bohr (1913)
Fixed energy shells
Modern Model
Quantum electron cloud
3. Discovery of the Neutron & the Three Subatomic Particles
A puzzle remained: a helium atom (2 protons) has about 4× the mass of a hydrogen atom (1 proton) — not 2×. This suggested another mass-contributing particle existed.
Discovered the neutron — mass nearly equal to a proton’s, but electrically neutral (no charge). Found in the nucleus of all atoms except ordinary hydrogen ($^1_1\text{H}$). Neutrons also help bind the nucleus together by countering the mutual repulsion between protons (the nuclear force) — this is why heavier atoms need proportionally more neutrons.
| Particle | Symbol | Relative Charge | Location | Discovered by |
|---|---|---|---|---|
| Electron | $e^-$ | −1 | Revolves around nucleus | J. J. Thomson (1897) |
| Proton | $p^+$ | +1 | Inside nucleus | Ernest Rutherford (1911) |
| Neutron | $n^0$ | 0 | Inside nucleus | James Chadwick (1932) |
4. Symbols of Elements
Dalton (1803) introduced pictorial symbols. Berzelius (1813) proposed deriving symbols from Latin names — giving us modern alphabetic symbols. Today, IUPAC approves all element names/symbols internationally.
Rules for Writing Symbols
- Often the first letter, or first two letters, of the element’s name
- First letter always capital; second letter (if any) always lowercase — e.g., Aluminium = Al (not AL), Cobalt = Co (not CO)
- Some symbols use a letter other than the second letter — e.g., Chlorine = Cl, Zinc = Zn
- Some come from Latin/Greek/German names — Iron = Fe (Latin ferrum), Mercury = Hg (Greek hydrargyros), Tungsten = W (German wolfram)
5. Atomic Number, Mass Number & Notation
Mass Number $(A)$ = Number of protons $+$ Number of neutrons (protons and neutrons together = nucleons)
Standard notation for an atom:
| Element | Symbol | Z | Protons | Neutrons | Electrons | K | L | M |
|---|---|---|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | 0 | 1 | 1 | – | – |
| Helium | He | 2 | 2 | 2 | 2 | 2 | – | – |
| Lithium | Li | 3 | 3 | 4 | 3 | 2 | 1 | – |
| Beryllium | Be | 4 | 4 | 5 | 4 | 2 | 2 | – |
| Boron | B | 5 | 5 | 6 | 5 | 2 | 3 | – |
| Carbon | C | 6 | 6 | 6 | 6 | 2 | 4 | – |
| Nitrogen | N | 7 | 7 | 7 | 7 | 2 | 5 | – |
| Oxygen | O | 8 | 8 | 8 | 8 | 2 | 6 | – |
| Fluorine | F | 9 | 9 | 10 | 9 | 2 | 7 | – |
| Neon | Ne | 10 | 10 | 10 | 10 | 2 | 8 | – |
| Sodium | Na | 11 | 11 | 12 | 11 | 2 | 8 | 1 |
| Magnesium | Mg | 12 | 12 | 12 | 12 | 2 | 8 | 2 |
| Aluminium | Al | 13 | 13 | 14 | 13 | 2 | 8 | 3 |
| Silicon | Si | 14 | 14 | 14 | 14 | 2 | 8 | 4 |
| Phosphorus | P | 15 | 15 | 16 | 15 | 2 | 8 | 5 |
| Sulfur | S | 16 | 16 | 16 | 16 | 2 | 8 | 6 |
| Chlorine | Cl | 17 | 17 | 18 | 17 | 2 | 8 | 7 |
| Argon | Ar | 18 | 18 | 22 | 18 | 2 | 8 | 8 |
5.1 Bohr–Bury Rules for Electron Distribution
- Maximum electrons in a shell $= 2n^2$ (where $n$ = shell number): K-shell ($n=1$) → 2; L-shell ($n=2$) → 8; M-shell ($n=3$) → 18
- Maximum electrons in the outermost shell = 8 (except when it’s the only shell, which holds max 2)
- Shells fill stepwise from innermost outward (K → L → M → N…); a shell fills completely before the next one starts
A: 12 electrons to distribute: K-shell takes 2 (full) → 10 remain; L-shell takes 8 (full) → 2 remain; M-shell takes the remaining 2.
Configuration: 2, 8, 2
6. Valency — Combining Capacity
Valency = number of electrons an atom gains, loses, or shares to complete its octet (achieve stability).
Fewer than 4 valence electrons
Tends to lose electrons. E.g., Sodium (2, 8, 1) loses 1 electron → valency 1.
More than 4 valence electrons
Tends to gain electrons. E.g., Oxygen (2, 6) gains 2 electrons → valency 2.
Exactly 4 valence electrons (e.g., Carbon: 2, 4) → tends to share 4 electrons → valency 4.
7. Isotopes and Isobars
7.1 Isotopes
Isotopes are atoms of the same element (same atomic number $Z$, same number of protons/electrons) but with different numbers of neutrons, hence different mass numbers $A$.
| Isotope | Notation | Protons | Neutrons | Electrons | Natural abundance |
|---|---|---|---|---|---|
| Protium | $^1_1\text{H}$ | 1 | 0 | 1 | ~99.98% |
| Deuterium | $^2_1\text{H}$ | 1 | 1 | 1 | ~0.015% |
| Tritium | $^3_1\text{H}$ | 1 | 2 | 1 | traces |
Applications of Isotopes
- $^{235}_{92}\text{U}$ — fuel in nuclear reactors (electricity generation)
- $^{60}_{27}\text{Co}$ — radiotherapy for cancer treatment
- $^{131}_{53}\text{I}$ — treatment of goitre and thyroid cancer
- $^{14}_{6}\text{C}$ — radiocarbon dating of fossils/artefacts in archaeology and geology
Average (Weighted) Atomic Mass
Since isotopes don’t occur in equal proportions, the atomic mass of an element is calculated as a weighted average, not a simple average.
7.2 Isobars
Isobars are atoms of different elements with the same mass number but different atomic numbers.
| Element | Atomic Number (Z) | Mass Number (A) |
|---|---|---|
| Argon | 18 | 40 |
| Potassium | 19 | 40 |
| Calcium | 20 | 40 |
8. Beyond Bohr — The Modern View
Bohr’s model, too, was eventually found incomplete. Electrons do not follow fixed, well-defined circular paths. The modern quantum mechanical model describes electrons as existing in “electron clouds” — regions around the nucleus where an electron is most likely to be found, rather than an exact, predictable path. (Detailed treatment is reserved for higher grades.)
9. Scientists to Remember
Discovered the electron via cathode ray studies; Nobel Prize in Physics, 1906; headed the Cavendish Laboratory, Cambridge, and mentored Ernest Rutherford.
Born in New Zealand; known as the “Father of Nuclear Physics”; discovered the atomic nucleus and the proton; Nobel Prize in Chemistry, 1908; his portrait appears on New Zealand’s $100 banknote.
Proposed fixed energy levels/shells to explain atomic stability; Nobel Prize in Physics, 1922.
Discovered the neutron in 1932 while working under Rutherford at the Cavendish Laboratory; Nobel Prize in Physics, 1935.
“Father of the Indian nuclear programme”; established the Tata Institute of Fundamental Research (TIFR) and the Bhabha Atomic Research Centre (BARC) for peaceful uses of atomic energy — electricity generation, agriculture, medical treatment.
10. Exam Question Bank
A. MCQ / Assertion–Reason (1 mark each)
- The gold foil experiment showed most alpha particles pass undeflected, proving that most of the atom is empty space (NOT evidence of neutrons — neutrons were discovered separately by Chadwick in 1932).[1]
- Assertion: Electrons do not fall into the nucleus in Bohr’s model. Reason: Electrons in a stationary state (fixed shell) do not lose energy. — Both true, R correctly explains A.[1]
- The maximum number of electrons in the M-shell is $2n^2 = 2(3)^2 = 18$.[1]
B. Short Answer (2–3 marks)
- State the postulates of Bohr’s model of the atom that explain the stability of atoms.[3]
- Differentiate between isotopes and isobars with one example each.[3]
- Why do atoms of the same element have the same chemical properties even when they are different isotopes?[2]
- Explain why the gold foil experiment’s results could not be explained by Thomson’s model.[3]
C. Numerical Problems
- An atom has atomic number 26 and 56 nucleons. Find the number of electrons, protons, and neutrons.[2]
- An atom has 18 neutrons and atomic number 17. Find its mass number.[2]
- Bromine has two isotopes: $^{79}_{35}\text{Br}$ (49.7%) and $^{81}_{35}\text{Br}$ (50.3%). Calculate its average atomic mass.[3]
- An atom $^{23}A$ has 11 electrons. Find the number of neutrons in it.[2]
D. Long Answer / HOTS (4–5 marks)
- Trace the historical development of atomic models from Dalton to the modern quantum mechanical model, highlighting the key experimental evidence that led to each revision.[5]
- Explain, with reasoning, why Rutherford’s model failed to explain atomic stability, and how Bohr’s model resolved this issue.[4]
- For an element with mass number 35 and 18 neutrons: (i) find protons and electrons, (ii) find atomic number, (iii) identify the element, (iv) write its electronic configuration, (v) state its valency.[5]
- Explain why isotopes of an element have the same valency but the weighted average atomic mass of the element is often a non-whole number.[4]
At a Glance — Chapter Summary
- Atomic theory origins: Acharya Kanada’s parmanu and Democritus’s atomos — imaginative, not experimental. Dalton (1808) gave the first scientific atomic theory.
- Model evolution: Dalton (indivisible sphere) → Thomson (plum pudding — electrons in positive sphere) → Rutherford (nucleus, mostly empty space, but unstable) → Bohr (fixed energy shells, stable) → Modern (quantum/electron cloud model).
- Gold foil experiment (Geiger–Marsden, under Rutherford) → discovered the nucleus via α-particle scattering.
- Three subatomic particles: electron ($e^-$, −1, Thomson), proton ($p^+$, +1, Rutherford), neutron ($n^0$, 0, Chadwick).
- Atomic number ($Z$) = number of protons; Mass number ($A$) = protons + neutrons (nucleons).
- Bohr–Bury rule: max electrons per shell = $2n^2$; outermost shell max = 8.
- Valency = electrons lost/gained/shared to complete the octet.
- Isotopes: same element, same $Z$, different $A$ (different neutrons) — same chemical, different physical properties.
- Isobars: different elements, same $A$, different $Z$.
- Average atomic mass = weighted average based on natural isotopic abundance.