AFNS Chemistry Chapter 2 — Atomic Structure
TOPIC 1: Subatomic Particles and Discovery of the Atom
Discovery Timeline — Important Scientists
| Scientist / Year | Contribution | Key Point |
| John Dalton (1803) | Dalton’s Atomic Theory: All matter made of indivisible atoms. Atoms of same element identical. Atoms combine in simple ratios. | First scientific atomic theory. Cannot explain electrons. |
| J.J. Thomson (1897) | Discovered the ELECTRON using cathode ray tube. Proposed ‘Plum Pudding’ model: electrons embedded in positive sphere. | Electron charge-to-mass ratio measured. First subatomic particle discovered. |
| Ernest Rutherford (1911) | Gold foil experiment → discovered the NUCLEUS. Most atom is empty space. Tiny, dense, positive nucleus. | Planetary model — electrons orbit nucleus. |
| Niels Bohr (1913) | Bohr’s model: electrons move in fixed circular orbits (energy levels) at specific energies. Electrons absorb/emit specific energy to move between levels. | Explained hydrogen emission spectrum. ⭐ Most tested model. |
| James Chadwick (1932) | Discovered the NEUTRON. Bombardment of beryllium with alpha particles. | Completed the three-subatomic-particle model. |
| Erwin Schrödinger (1926) | Wave mechanical (quantum mechanical) model. Electron described by wave function. Electrons in orbitals (probability regions), not fixed orbits. | Modern model — most accurate. Orbitals have specific shapes. |
Subatomic Particles
| Particle | Symbol | Charge | Relative Mass | Location |
| Proton | p⁺ | +1 | 1 amu | Nucleus |
| Neutron | n⁰ | 0 | 1 amu | Nucleus |
| Electron | e⁻ | −1 | 1/1836 amu (≈ 0) | Shells/orbitals around nucleus |
| KEY FACT | Atomic Number (Z) = number of PROTONS in nucleus = number of electrons (in neutral atom) Mass Number (A) = number of protons + number of neutrons Number of neutrons = A − Z Notation: ᴬ_Z X → e.g. ¹²₆C (Mass number=12, Atomic number=6, so neutrons=6) In neutral atom: protons = electrons In positive ion (cation): protons > electrons (lost electrons) In negative ion (anion): protons < electrons (gained electrons) |
| 📌 WORKED EXAMPLE Q: For ²³₁₁Na: find protons, neutrons, electrons in the neutral atom and in Na⁺. A: Atomic number Z=11 → protons=11 Mass number A=23 → neutrons = 23−11 = 12 Neutral Na: electrons = 11 (= protons) Na⁺ (lost 1 electron): electrons = 11−1 = 10 |
TOPIC 2: Isotopes, Isobars, and Isotones
| Term | Definition | Same | Different | Example |
| Isotopes | Same element, different mass numbers | Atomic number (Z) Number of protons | Mass number (A) Number of neutrons | ¹H, ²H (deuterium), ³H (tritium) ¹²C and ¹⁴C (carbon-14) ³⁵Cl and ³⁷Cl |
| Isobars | Different elements, same mass number | Mass number (A) | Atomic number (Z) Chemical properties | ⁴⁰₁₈Ar and ⁴⁰₂₀Ca ¹⁴₆C and ¹⁴₇N |
| Isotones | Different elements, same number of neutrons | Number of neutrons | Atomic number (Z) Mass number (A) | ¹⁴₆C (n=8) and ¹⁵₇N (n=8) |
| KEY FACT | Isotopes KEY facts: • Same element (same Z, same proton number, same chemical properties) • Different neutron numbers → different physical properties (density, melting point) • Radioactive isotopes used in medicine: ¹³¹I (thyroid cancer treatment), ⁹⁹ᵐTc (nuclear imaging), ¹⁴C (carbon dating) • Relative atomic mass = weighted average of isotope masses Example: Cl has ³⁵Cl (75%) and ³⁷Cl (25%): Ar = 0.75×35 + 0.25×37 = 26.25+9.25 = 35.5 |
TOPIC 3: Bohr’s Atomic Model — Most Tested in AFNS
Bohr’s Postulates
Niels Bohr (1913) proposed his model specifically for the hydrogen atom. He assumed electrons move in FIXED, CIRCULAR orbits (shells/energy levels) around the nucleus — like planets around the sun.
- Postulate 1 — Electrons move in circular orbits around the nucleus. These orbits are called STATIONARY STATES or ENERGY LEVELS (shells). Each shell is designated by principal quantum number n = 1, 2, 3, 4… (or K, L, M, N shells).
- Postulate 2 — Not all orbits are allowed. Only orbits where the angular momentum (mvr) is an integer multiple of h/2π are permitted: mvr = nh/2π (where h = Planck’s constant, n = 1,2,3…)
- Postulate 3 — As long as an electron remains in its orbit, it does NOT radiate energy (it is in a stationary state). Contradicts classical physics which says accelerating charges must radiate.
- Postulate 4 — An electron CAN jump from one orbit to another by ABSORBING or EMITTING energy equal to the DIFFERENCE between the two energy levels: ΔE = E₂ − E₁ = hν (where h = Planck’s constant, ν = frequency of light)
| KEY FACT | Energy of Bohr orbit: En = −13.6/n² eV (for hydrogen) n=1: E₁ = −13.6 eV (ground state — lowest energy, most stable) n=2: E₂ = −3.4 eV n=3: E₃ = −1.51 eV n=∞: E = 0 (ionisation — electron completely removed) Negative sign means electron is BOUND to nucleus (energy needed to remove it) Ionisation energy of hydrogen = 13.6 eV (energy to remove electron from n=1) |
Hydrogen Spectrum — Spectral Series
When hydrogen atoms are excited (given energy), electrons jump to higher levels. When they fall back, they emit light at specific wavelengths — the hydrogen emission spectrum. Different series correspond to electrons falling to different final energy levels:
| Series Name | Final Level (n) | Region of Spectrum | Transition |
| Lyman series | n=1 | Ultraviolet (UV) ⭐ | n=2,3,4… → n=1 |
| Balmer series | n=2 | Visible light ⭐ (most important) | n=3,4,5… → n=2 |
| Paschen series | n=3 | Infrared (IR) | n=4,5,6… → n=3 |
| Brackett series | n=4 | Infrared (IR) | n=5,6,7… → n=4 |
| Pfund series | n=5 | Far infrared | n=6,7,8… → n=5 |
| QUICK TIP | Memory trick for spectral series: ‘Little Boys Playing Ball Players’ Lyman (UV) | Balmer (Visible) | Paschen (IR) | Brackett (IR) | Pfund (Far IR) Balmer series = VISIBLE light → most commonly asked. It has 4 lines in visible spectrum: Red (n=3→2), Blue-green (n=4→2), Blue (n=5→2), Violet (n=6→2) |
Limitations of Bohr’s Model
- Cannot explain spectra of multi-electron atoms (only works for hydrogen-like species: H, He⁺, Li²⁺)
- Cannot explain fine structure of spectral lines (splitting seen at higher resolution)
- Cannot explain the Stark effect (splitting of lines in electric field) or Zeeman effect (splitting in magnetic field)
- Does not follow Heisenberg’s Uncertainty Principle — assumes exact position AND momentum of electron (cannot know both simultaneously)
- Does not account for wave nature of electrons (de Broglie relation: λ = h/mv)
TOPIC 4: Quantum Numbers — Electronic Address of an Electron
In the modern (wave mechanical) model, every electron in an atom has a unique set of FOUR quantum numbers that completely describe its state — its ‘address’ in the atom.
| Quantum Number | Symbol | Values | What It Describes |
| Principal quantum number | n | 1, 2, 3, 4, 5… (positive integers) | Main energy level (shell). Determines energy and size of orbital. n=1 (K shell), n=2 (L), n=3 (M), n=4 (N) Max electrons in shell = 2n² |
| Azimuthal (Angular Momentum) quantum number | l | 0 to (n−1) For n=3: l=0,1,2 | Subshell (sublevel) and shape of orbital. l=0 → s (spherical) l=1 → p (dumbbell) ⭐ l=2 → d (cloverleaf) ⭐ l=3 → f (complex shape) |
| Magnetic quantum number | mₗ | −l to +l (includes 0) For l=1: mₗ=−1,0,+1 | Orientation of orbital in space. Number of mₗ values = number of orbitals in subshell s: 1 orbital | p: 3 orbitals ⭐ d: 5 orbitals | f: 7 orbitals |
| Spin quantum number | mₛ | Only +½ or −½ | Spin of electron (clockwise or anticlockwise). Each orbital can hold MAX 2 electrons (one +½, one −½) ⭐ Pauli Exclusion Principle: no two electrons same all 4 quantum numbers |
| KEY FACT | Orbital capacity and shapes: s subshell: l=0, 1 orbital, max 2 electrons — SPHERICAL shape ⭐ p subshell: l=1, 3 orbitals (px,py,pz), max 6 electrons — DUMBBELL shape ⭐ d subshell: l=2, 5 orbitals, max 10 electrons — CLOVERLEAF shape ⭐ f subshell: l=3, 7 orbitals, max 14 electrons — complex shape Max electrons per shell: n=1→2 | n=2→8 | n=3→18 | n=4→32 Formula: 2n² |
| 📌 WORKED EXAMPLE Q: For the 3rd shell (n=3), what values are possible for l and mₗ? A: n=3, so l = 0, 1, 2 • l=0 (3s): mₗ = 0 (1 orbital) • l=1 (3p): mₗ = −1, 0, +1 (3 orbitals) • l=2 (3d): mₗ = −2, −1, 0, +1, +2 (5 orbitals) Total orbitals in n=3 shell = 1+3+5 = 9 orbitals Max electrons = 9×2 = 18 (= 2n² = 2×3² = 18 ✓) |