What is everything made of? This question, asked by humans for thousands of years, finds its most fundamental modern answer in the structure of the atom. An atom is the smallest unit of matter that retains the chemical identity of an element — and yet atoms themselves are composed of even smaller particles: protons, neutrons, and electrons.
The story of how we discovered atomic structure is one of the greatest detective stories in science. Over just 50 years (1897–1932), scientists went from no knowledge of subatomic particles to a complete quantum mechanical picture of the atom. For CBSE Class 9 Chemistry Chapter 4 (Structure of Atom), we study three key atomic models (Thomson, Rutherford, Bohr), the discovery of subatomic particles, electron shell configurations, atomic number and mass number, and the remarkable phenomena of isotopes, isobars, and isotones.
Mass: 1.673×10⁻²⁷ kg
Relative mass: 1 u
Location: Nucleus
Discovered: Goldstein, 1886
Symbol: p (or p⁺)
Mass: 1.675×10⁻²⁷ kg
Relative mass: 1 u
Location: Nucleus
Discovered: Chadwick, 1932
Symbol: n (or n⁰)
Mass: 9.109×10⁻³¹ kg
Relative mass: 1/1837 u
Location: Shells (orbits)
Discovered: J.J. Thomson, 1897
Symbol: e (or e⁻)
- 1. Discovery of Subatomic Particles
- 2. Thomson’s Atomic Model (Plum Pudding Model)
- 3. Rutherford’s Nuclear Model — Gold Foil Experiment
- 4. Bohr’s Atomic Model — Postulates and Electron Shells
- 5. Electron Distribution in Shells — The 2n² Rule
- 6. Atomic Number, Mass Number, and Valency
- 7. Isotopes, Isobars, and Isotones
- 8. Solved Numerical Problems
- 9. Frequently Asked Questions (FAQ)
1. Discovery of Subatomic Particles
Key observations:
• The ray traveled in straight lines from cathode to anode regardless of the gas used.
• The ray was deflected toward the positive plate in an electric field → the ray consisted of negatively charged particles.
• The same particles were obtained regardless of which metal the cathode was made of → these particles are a universal constituent of all matter.
Conclusion: Thomson discovered negatively charged particles called electrons. Charge of electron: −1.6 × 10⁻¹⁹ C. Mass of electron: 9.1 × 10⁻³¹ kg (about 1/1837 times the mass of a hydrogen atom).
Discovery of Proton — Goldstein (1886): Using a perforated cathode in a discharge tube, Goldstein observed rays moving from anode to cathode (opposite to cathode rays). These were positively charged particles. The lightest positive ion was obtained from hydrogen gas — a particle with mass 1 u and charge +1, called the proton.
Discovery of Neutron — Chadwick (1932): James Chadwick bombarded beryllium atoms with alpha particles and detected electrically neutral particles with mass approximately equal to a proton. These were neutrons. This completed the set of three fundamental subatomic particles.
2. Thomson’s Atomic Model (1904) — The “Plum Pudding Model”
(1) An atom is a sphere of uniform positive charge with a radius of approximately 10⁻¹⁰ m.
(2) Electrons are embedded throughout the positive sphere, like plums (raisins) in a plum pudding (or watermelon seeds in a watermelon).
(3) The total positive charge equals the total negative charge from the embedded electrons, making the atom electrically neutral overall.
Success: Explained why atoms are electrically neutral. Explained the existence of electrons in atoms.
3. Rutherford’s Nuclear Model — Gold Foil Experiment (1911)
Observations and Conclusions:
• Most alpha particles (99.99%) passed straight through: → Atom is mostly empty space.
• Some deflected at small angles: → There is a positive charge concentrated in the atom that repels the positive alpha particles.
• Very few (about 1 in 20,000) bounced back at nearly 180°: → The positive charge and most of the mass are concentrated in a tiny, dense core called the nucleus.
Rutherford’s atomic model:
(1) Every atom has a nucleus at the centre — a tiny, dense, positively charged region containing all the protons (and neutrons, discovered later).
(2) The nucleus has a radius of about 10⁻¹⁵ m — only about 1/100,000th of the atom’s radius (10⁻¹⁰ m). The nucleus is to the atom as a marble is to a football stadium.
(3) Electrons orbit the nucleus at a relatively large distance, like planets around the sun.
(4) Most of the atom is empty space.
(1) Stability problem: According to classical physics, an electron moving in a circular orbit (accelerating) must continuously emit electromagnetic radiation, losing energy. It should spiral into the nucleus in about 10⁻&sup8; seconds — but atoms are stable. Rutherford’s model could not explain atomic stability.
(2) No electron arrangement: It did not describe how electrons are arranged around the nucleus.
(3) No explanation of atomic spectra: Could not explain the discrete line spectra observed for hydrogen and other elements.
4. Bohr’s Atomic Model (1913) — Postulates and Electron Shells
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Fixed Circular Orbits (Stationary States)Electrons revolve around the nucleus in fixed circular paths called orbits or shells. These are designated K, L, M, N (or n = 1, 2, 3, 4). An electron in a fixed orbit does not emit or absorb energy — its energy remains constant. These states are called stationary states.
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Quantised Energy LevelsEach shell corresponds to a definite fixed energy level. The energy of a shell increases with its distance from the nucleus: E₁ < E₂ < E₃ < E₄ (K < L < M < N). Electrons cannot exist between shells — energy levels are quantised (discrete).
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Energy Transitions (Emission and Absorption)An electron can jump from one orbit to another by absorbing or emitting a definite quantity of energy (a photon). If an electron jumps from a higher energy shell to a lower one, it emits a photon with energy equal to the difference between the two shell energies: E = E₂ − E₁ = hν (where h = Planck’s constant, ν = frequency of light). This explains the line spectra of atoms.
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Angular Momentum is QuantisedThe angular momentum of an electron in any permitted orbit is a whole-number multiple of h/2π: mvr = nh/2π, where m = electron mass, v = velocity, r = orbit radius, n = principal quantum number (1, 2, 3...). This is what restricts electrons to discrete orbits.
(1) Works only for hydrogen and hydrogen-like ions (He⁺, Li²⁺) — fails for multi-electron atoms.
(2) Cannot explain the fine structure of spectral lines (splitting in magnetic fields — Zeeman effect).
(3) Does not incorporate wave-particle duality of electrons (later described by quantum mechanics/de Broglie).
5. Electron Distribution in Shells — The 2n² Rule
The second-to-last shell can never hold more than 18 electrons. Electrons fill inner shells first.
| Element | Symbol | Atomic No. (Z) | K(n=1) | L(n=2) | M(n=3) | N(n=4) | Config. |
|---|---|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | — | — | — | 1 |
| Helium | He | 2 | 2 | — | — | — | 2 |
| Lithium | Li | 3 | 2 | 1 | — | — | 2,1 |
| Carbon | C | 6 | 2 | 4 | — | — | 2,4 |
| Nitrogen | N | 7 | 2 | 5 | — | — | 2,5 |
| Oxygen | O | 8 | 2 | 6 | — | — | 2,6 |
| Neon | Ne | 10 | 2 | 8 | — | — | 2,8 |
| Sodium | Na | 11 | 2 | 8 | 1 | — | 2,8,1 |
| Magnesium | Mg | 12 | 2 | 8 | 2 | — | 2,8,2 |
| Aluminium | Al | 13 | 2 | 8 | 3 | — | 2,8,3 |
| Silicon | Si | 14 | 2 | 8 | 4 | — | 2,8,4 |
| Chlorine | Cl | 17 | 2 | 8 | 7 | — | 2,8,7 |
| Argon | Ar | 18 | 2 | 8 | 8 | — | 2,8,8 |
| Potassium | K | 19 | 2 | 8 | 8 | 1 | 2,8,8,1 |
| Calcium | Ca | 20 | 2 | 8 | 8 | 2 | 2,8,8,2 |
6. Atomic Number, Mass Number, and Valency
• Definition: The number of protons in the nucleus of an atom.
• Z uniquely identifies an element: all atoms of carbon have Z=6; all atoms of oxygen have Z=8.
• In a neutral atom: number of protons = number of electrons = Z.
• Symbol: Z (from German “Zahl” = number)
Mass Number (A):
• Definition: The total number of protons and neutrons in the nucleus of an atom.
• A = Z + N, where N = number of neutrons.
• Also called nucleon number (protons and neutrons are collectively called nucleons).
Number of Neutrons (N):
• N = A − Z
• Neutrons have no charge and contribute to mass but not to chemical properties.
Standard notation:
An atom of element X with atomic number Z and mass number A is written as: ¹ₔX (mass number as superscript, atomic number as subscript before the symbol). Example: ¹⁶₈C = Carbon atom with A=12, Z=6 (6 protons, 6 neutrons, 6 electrons).
Valency from electron configuration:
Valency = number of electrons in the outermost (valence) shell, OR
If the outermost shell has >4 electrons: Valency = 8 − (outermost electrons) [tendency to gain electrons]
Examples: Na (2,8,1) → valency = 1. Cl (2,8,7) → valency = 8−7 = 1. O (2,6) → valency = 8−6 = 2. N (2,5) → valency = 3. Noble gases (full outer shell) → valency = 0.
7. Isotopes, Isobars, and Isotones
Examples:
• Hydrogen isotopes: ¹H (protium, 0 neutrons), ²H (deuterium, 1 neutron), ³H (tritium, 2 neutrons) — all have Z=1.
• Carbon: ¹²C (6n), ¹⁶C (8n) — both have Z=6.
• Chlorine: ²⁵Cl (Z=17, A=35, 18n) and ³⁷Cl (Z=17, A=37, 20n).
Same chemical properties (same Z, same electron config.). Different physical properties (different mass). Chlorine’s average atomic mass of 35.5 u is due to the natural mixture of ²⁵Cl (75%) and ³⁷Cl (25%).
Uses of isotopes: ¹⁶C (carbon-14) in radiocarbon dating. ¹³¹I (iodine-131) in thyroid treatment. U-235 in nuclear fuel. Co-60 in cancer radiotherapy.
Examples:
• ⁶₄Ca (Z=20, N=24) and ⁶€Ar (Z=18, N=22) — both have A=40.
• ²⁶Fe (Z=26) and ²⁶Ni (Z=28) — both have A=58? (No, let’s use a classic): ³⁸Ar (Z=18) and ³⁸K (Z=19) and ³⁸Ca (Z=20) — all have A=40.
• ¹⁶C (Z=6) and ¹⁶N (Z=7) — both have A=16? (No.) Classic: ²⁶Mg (Z=12, A=24) and ²⁶Na (A=24, Z=11).
Different chemical properties (different Z, different electron configuration). Same mass number only.
Examples:
• ¹⁶C (Z=6, A=12, N=6) and ¹⁶N (Z=7, A=13? No.) Let’s be clear:
• ¯⁶C (Z=6, A=13, N=7) and ¯⁶N (Z=7, A=14, N=7) — both have 7 neutrons.
• ¹⁶⁴C (Z=6, A=14, N=8) and ¹⁶⁵N (Z=7, A=15, N=8) — both have 8 neutrons.
• ¹⁶⁴Si (Z=14, N=14) and ¹⁶⁵P (Z=15, N=14).
Note: Isotones have the same number of neutrons but are atoms of different elements (different Z and different A). This is the least tested of the three in CBSE exams but appears in higher-order questions.
| Feature | Isotopes | Isobars | Isotones |
|---|---|---|---|
| Atomic Number (Z) | Same | Different | Different |
| Mass Number (A) | Different | Same | Different |
| Neutrons (N) | Different | Different | Same |
| Same element? | Yes | No | No |
| Chemical properties | Same | Different | Different |
| Classic example | ¹H, ²H, ³H | ⁶€Ar and ⁶€Ca | ¹⁶⁴C and ¹⁶⁵N (both 8n) |
8. Solved Numerical Problems
Electrons = Z = 15 (neutral atom)
Neutrons = A − Z = 31 − 15 = 16
Electron configuration: 15 electrons → K=2, L=8, M=5 (since 2+8=10; 15−10=5 for M)
Configuration: 2, 8, 5
Valency: outermost shell has 5 electrons → Valency = 5 (or 8−5=3, as phosphorus can show valency 3 or 5)
Mass number A = Z + N = 17 + 18 = 35
Valency: outermost shell has 7 electrons → Valency = 8 − 7 = 1
Element: Z = 17 → Chlorine (Cl)
(a) ¹²C (Z=6) and ¹²N (Z=7) (b) ¹H (Z=1, A=1) and ²H (Z=1, A=2) (c) ¹³C (Z=6, A=13) and ¹⁶N (Z=7, A=14)
(b) ¹H: Z=1, A=1, N=0; ²H: Z=1, A=2, N=1 → Same Z (=1), different A → Isotopes
(c) ¹³C: Z=6, A=13, N=7; ¹⁶N: Z=7, A=14, N=7 → Same N (=7), different Z → Isotones
Filling for Ca (20 electrons): K=2, L=8 (2+8=10), M=8 (but M max is 18; however CBSE rule: outermost shell max 8)
After K=2 and L=8, remaining = 20−10 = 10. Since outermost cannot exceed 8: M=8, N=2
Configuration: 2, 8, 8, 2
= (35 × 0.75) + (37 × 0.25)
= 26.25 + 9.25
Explore Related CBSE Class 9 Science Guides
9. Frequently Asked Questions (FAQ)
Bohr’s atomic model (proposed by Niels Bohr in 1913) describes electrons orbiting the nucleus in fixed circular paths called shells or energy levels (K, L, M, N corresponding to n = 1, 2, 3, 4).
Four main postulates:
(1) Fixed circular orbits: Electrons revolve around the nucleus in fixed circular paths called stationary orbits or shells. An electron in a fixed orbit does not emit energy.
(2) Quantised energy levels: Each shell has a definite fixed energy. Energy increases with shell number: K < L < M < N.
(3) Energy transitions: When an electron jumps from a higher shell to a lower shell, it emits energy (photon). When it absorbs energy, it jumps to a higher shell. Energy emitted/absorbed = E₂ − E₁ = hν. This explains atomic spectra (line spectra).
(4) Quantised angular momentum: The angular momentum of an electron in any allowed orbit = nh/2π (where n = 1, 2, 3... and h = Planck’s constant).
Success: Explained the stability of atoms and the discrete line spectra of hydrogen. Limitation: Only works for hydrogen and single-electron ions; fails for multi-electron atoms.
Rutherford’s alpha particle scattering experiment (1911):
A narrow beam of alpha particles (α²⁺ ions = helium nuclei) was directed at an extremely thin gold foil (about 1000 atoms thick). A zinc sulphide detection screen surrounded the foil.
Observations:
(1) Most alpha particles (99.99%+) passed straight through the gold foil without any deflection.
(2) Some alpha particles were deflected at small angles.
(3) Very rarely (about 1 in 20,000), an alpha particle bounced back at nearly 180° (backscattered).
Conclusions (Rutherford’s Nuclear Model):
(1) Since most particles passed through: The atom is mostly empty space.
(2) Since some were deflected: There is a concentrated positive charge that repels positive alpha particles.
(3) Since some bounced straight back: All the positive charge and most of the mass is concentrated in a tiny, dense region at the centre called the nucleus.
(4) The nucleus is extremely small (~10⁻¹⁵ m) compared to the atom (~10⁻¹⁰ m).
(5) Electrons orbit the nucleus at a large distance (like planets around the sun).
Atomic Number (Z):
• Definition: The number of protons in the nucleus of an atom.
• Z uniquely identifies the element. All atoms of carbon have Z=6; all atoms of nitrogen have Z=7.
• In a neutral atom: Z = number of protons = number of electrons.
• Example: Oxygen has Z=8 (8 protons, 8 electrons in neutral atom).
Mass Number (A):
• Definition: The total number of protons AND neutrons in the nucleus.
• A = Z + N (where N = number of neutrons). Also called nucleon number.
• Example: Oxygen-16 has A=16, Z=8, N=16−8=8 neutrons.
Number of Neutrons: N = A − Z
Key distinction: Two atoms of the same element always have the same Z (same number of protons) but can have different A (different numbers of neutrons) — these are called isotopes. Standard notation: ¹⁶₈X means element X with mass number 12 and atomic number 6 (carbon-12).
Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A). They have the same number of protons and electrons but different numbers of neutrons.
Examples:
(1) Hydrogen isotopes: Protium (¹H, 0 neutrons), Deuterium (²H or D, 1 neutron), Tritium (³H or T, 2 neutrons). All have Z=1.
(2) Carbon isotopes: C-12 (¹²C, 6 neutrons, most abundant) and C-14 (¹⁶C, 8 neutrons, radioactive).
(3) Chlorine isotopes: Cl-35 (75% natural abundance) and Cl-37 (25%). Average mass = 35.5 u.
(4) Uranium isotopes: U-235 (fissile, used in nuclear reactors) and U-238 (most abundant).
Chemical properties of isotopes: Identical (same Z = same electron configuration = same chemistry).
Physical properties: Different (different mass, different melting/boiling point for heavy isotopes).
Uses of isotopes:
• C-14 (Carbon-14): Radiocarbon dating of ancient artifacts, fossils, and archaeological finds (half-life = 5,730 years).
• I-131 (Iodine-131): Treatment of thyroid cancer and hyperthyroidism.
• Co-60 (Cobalt-60): Cancer radiotherapy (radiation oncology).
• U-235: Fuel in nuclear reactors and nuclear weapons (fission).
• D₂O (Heavy water, using deuterium): Moderator in nuclear reactors.
Electron configuration describes how electrons are arranged in the various shells (K, L, M, N) of an atom, starting from the innermost shell.
Rules:
(1) 2n² rule: Maximum electrons in shell n = 2n². So K(n=1) = 2, L(n=2) = 8, M(n=3) = 18, N(n=4) = 32.
(2) Outermost shell limit: The outermost (valence) shell can have maximum 8 electrons (for CBSE Class 9 — this is the octet rule).
(3) Fill inner shells first before proceeding to outer shells.
Step-by-step example for Sodium (Na, Z=11):
• Total electrons = 11
• K shell: fill up to 2 → K=2 (9 remaining)
• L shell: fill up to 8 → L=8 (1 remaining)
• M shell: remaining 1 electron → M=1
• Configuration: 2, 8, 1 • Valency: 1 (1 electron in outermost shell)
Valency from electron configuration:
• If outermost electrons ≤ 4: Valency = outermost electrons (tendency to lose)
• If outermost electrons > 4: Valency = 8 − outermost electrons (tendency to gain)
• Noble gases (full outer shell): Valency = 0
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