Atomic Structure (RRB)

An atom is the smallest particle of an element that retains its chemical properties.

Sub-atomic particles:

• Proton (p⁺): positive charge, mass ≈ 1 atomic mass unit (u). In nucleus.
• Neutron (n⁰): no charge, mass ≈ 1 u. In nucleus.
• Electron (e⁻): negative charge, mass ≈ 1/1836 u (essentially zero). Orbits the nucleus in shells.

Atomic number (Z) = number of protons. Defines the element.

Mass number (A) = protons + neutrons.

Number of electrons = Z (in a neutral atom).

Notation: ⁻ᴬZₓ — e.g. ¹²₆C means carbon with 6 protons, 6 neutrons.

Electron shells (Bohr model):
Shells K (1st), L (2nd), M (3rd), N (4th) … hold max 2n² electrons:

• K: 2
• L: 8
• M: 18 (commonly fills 8 first)
• N: 32

Valence electrons = electrons in the outermost shell. Determine chemistry.

• Noble gases (group 18) have full outer shells → unreactive.
• Other atoms try to achieve a full shell by gaining, losing, or sharing electrons.

Ions:

• An atom that loses electrons → positive ion (cation). E.g. Na → Na⁺.
• An atom that gains electrons → negative ion (anion). E.g. Cl + e⁻ → Cl⁻.

Isotopes: atoms of the same element with different numbers of neutrons. Same Z, different A.

• ¹H, ²H (deuterium), ³H (tritium) — all hydrogen.
• ¹²C, ¹³C, ¹⁴C — carbon. ¹⁴C is radioactive, used in carbon-dating.
• ²³⁵U and ²³⁸U — uranium. Only ²³⁵U is fissile.
• ¹³¹I and ¹³³I — used in medical imaging and treatment.

Isobars: different elements with same mass number, e.g. ⁴⁰Ar and ⁴⁰Ca.

Isotones: different elements with same number of neutrons.

Radioactivity: spontaneous decay of unstable nuclei → emission of α (helium nucleus), β (electron), or γ (high-energy photon) radiation.

Half-life: time for half of a radioactive sample to decay. ¹⁴C has half-life 5730 years; ²³⁵U has 7 × 10⁸ years.

Evolution of the atomic model:

1. Dalton (1808): atoms are indivisible. (Wrong — they have sub-atomic parts.)

2. Thomson (1897): discovered the electron (cathode-ray experiment). Proposed the "plum pudding" model — positive cloud with embedded electrons.

3. Rutherford (1911): alpha-scattering experiment — most α particles passed through gold foil, but some bounced back. Concluded: atoms have a tiny dense positive nucleus, with electrons orbiting like planets. Most of the atom is empty space.

4. Bohr (1913): electrons move in fixed energy orbits around the nucleus. Energy is released/absorbed as electrons jump between orbits. Explains the discrete spectra of hydrogen.

5. Quantum mechanical model (1920s): electrons exist in orbitals — probability regions, not fixed paths. Position and momentum can't be known exactly together (Heisenberg uncertainty principle).

Electronic configuration examples:

• Hydrogen (Z=1): K shell has 1 electron → 1.
• Carbon (Z=6): 2, 4.
• Oxygen (Z=8): 2, 6.
• Sodium (Z=11): 2, 8, 1 — that 1 outermost electron makes it reactive.
• Chlorine (Z=17): 2, 8, 7 — 1 short of full shell, eager to gain 1 → Cl⁻.

Octet rule: atoms tend to gain/lose/share electrons to achieve 8 valence electrons (like a noble gas). Sodium loses 1 → Na⁺; Chlorine gains 1 → Cl⁻. Together: NaCl.

Why noble gases are stable: full outermost shell → no tendency to gain/lose electrons.

Modern uses of atomic concepts:

• Carbon-14 dating: ratio of ¹⁴C to ¹²C in fossils tells age (up to ~50,000 years).
• Radioisotope therapy: ⁶⁰Co for cancer (gamma source).
• Smoke detectors: ²⁴¹Am (alpha emitter) ionizes air; smoke disrupts current → alarm.
• Nuclear reactors: ²³⁵U fission produces heat → steam → electricity.
• PET scans: ¹⁸F-labelled glucose injected → cancer cells absorb more → image.

Common exam trap: confusing mass number (A) with atomic number (Z) and atomic mass (the weighted average across isotopes, often non-integer). E.g. atomic mass of chlorine is 35.5 u because natural Cl is a mix of ³⁵Cl and ³⁷Cl.