Atomic Structure Essentials

Overview

This lecture explains the historical development of atomic structure, key experiments, atomic models (Thomson, Rutherford, Bohr), subatomic particles, electron configuration (Bohr–Bury rules), valency, ions, atomic and mass numbers, isotopes and isobars, with examples and calculations.

Early Ideas of the Atom

• 500 BC (Kanada) and 400 BC (Democritus): Matter is made of indivisible particles called atoms.
• Dalton (1800s): Revived atom as indivisible; later experiments showed divisibility.

Discovery of Subatomic Particles

• Electron: J. J. Thomson (1897) via cathode rays; negatively charged, very light.
• Canal rays: E. Goldstein (1886) observed positive rays (anode rays); led to proton discovery.
• Proton: Attributed to Rutherford’s work following canal rays (credit typically later).
• Neutron: James Chadwick (1932); neutral, mass ~ proton; in nucleus.

Thomson’s Model (Plum Pudding Model)

• Atom is a positively charged sphere with electrons embedded within.
• Total positive charge equals total negative charge; atom overall neutral.
• Model failed to explain later scattering results (Rutherford).

Rutherford’s Alpha Particle Scattering (Gold Foil) Experiment

• Setup: Thin gold foil, alpha particles (+2 charge, ~4 amu), fluorescent screen.
• Observations:
  • ~99% particles passed straight through.
  • Some deflected by small angles.
  • About 1 in 12,000 rebounded (~180°).
• Conclusions:
  • Most atomic space is empty.
  • Positive charge concentrated in tiny central nucleus.
  • Nucleus is ~10^5 times smaller than atom; contains almost all mass.
  • Electrons revolve around nucleus; planetary model.

Limitations of Rutherford’s Model

• Maxwell’s electromagnetism: Accelerating charged particles radiate energy.
• Electrons in circular motion would lose energy and spiral into nucleus quickly.
• Could not explain atomic stability.

Bohr’s Model (1913)

• Electrons revolve only in fixed circular orbits (energy levels) around nucleus.
• Orbits labeled K, L, M, N (also 1, 2, 3, 4).
• Each orbit has fixed energy; while in a given orbit, an electron neither gains nor loses energy.
• Solved stability issue by quantized energy levels.

Bohr–Bury Rules (Electron Distribution)

• Maximum electrons in a shell = 2n^2 (n = shell number: K=1, L=2, M=3, N=4).
  • K: 2; L: 8; M: 18; N: 32.
• Outer (valence) shell cannot have more than 8 electrons, even if capacity allows more.
• Valence shell: Outermost occupied shell. Valence electrons: Electrons in valence shell.

Electron Configuration (Z = 1–20)

• Fill shells as per 2, 8, 18, 32 with outermost shell ≤ 8.
• Examples:
  • H: 2,1 → 1
  • He: 2
  • Li: 2,1
  • C: 2,4
  • O: 2,6
  • Ne: 2,8 (stable octet)
  • Na: 2,8,1
  • Mg: 2,8,2
  • Al: 2,8,3
  • Si: 2,8,4
  • P: 2,8,5
  • S: 2,8,6
  • Cl: 2,8,7
  • Ar: 2,8,8 (stable octet)
  • K: 2,8,8,1
  • Ca: 2,8,8,2

Stability and Octet Rule

• Atoms with 8 electrons in valence shell are very stable (noble gases: Ne, Ar, etc.).
• Most atoms tend to achieve octet by losing, gaining, or sharing electrons.
• Exceptions: H, He, Li, Be, B (often follow duplet for stability).

Valency

• Definition: Number of electrons lost, gained, or shared to attain a stable configuration (octet/duplet).
• Trends:
  • Na (2,8,1): loses 1 → valency 1.
  • Mg (2,8,2): loses 2 → valency 2.
  • Al (2,8,3): loses 3 → valency 3.
  • Si (2,8,4): typically shares 4 → valency 4.
  • P (2,8,5): gains 3 or shares → valency 3.
  • S (2,8,6): gains 2 → valency 2.
  • Cl (2,8,7): gains 1 → valency 1.
  • Noble gases (e.g., Ar): valency 0.
  • Duplet cases (H, He, Li, Be, B): target 2 in valence shell.

Ions Formation

• Losing electrons → cations (positive charge), e.g., Na → Na+ (2,8).
• Gaining electrons → anions (negative charge), e.g., Cl → Cl− (2,8,8).
• Charge magnitude equals number of electrons lost/gained.

Atomic Number and Mass Number

• Atomic number (Z): Number of protons; identifies the element.
• Mass number (A): Number of protons + number of neutrons (nucleons).
• Electrons do not contribute significantly to mass.
• Relation: A − Z = number of neutrons.
• Neutral atom: number of electrons = number of protons.

Isotopes

• Definition: Atoms of the same element (same Z) with different mass numbers (different neutrons).
• Hydrogen isotopes:
  • Protium: 1^1H, Z=1, A=1, neutrons=0.
  • Deuterium: 2^1H (D), Z=1, A=2, neutrons=1.
  • Tritium: 3^1H (T), Z=1, A=3, neutrons=2.
• Chemical properties: Similar (same valence electrons).
• Physical properties: Different (mass-dependent).
• Applications:
  • Uranium isotope as nuclear reactor fuel.
  • Cobalt isotope in cancer treatment.
  • Iodine isotope in goiter treatment.

Fractional Atomic Mass

• Occurs due to natural isotopic mixture with different abundances.
• Example (chlorine): 35 (75%) and 37 (25%) → average atomic mass = (35×75/100) + (37×25/100) = 35.5 u.

Isobars

• Definition: Atoms of different elements (different Z) with the same mass number (A).
• Example: 40^18Ar and 40^20Ca.
• Chemical properties: Different (different electronic configurations).

Key Terms & Definitions

• Electron: Negatively charged subatomic particle; very small mass; in shells.
• Proton: Positively charged particle in nucleus; mass ~1 amu.
• Neutron: Neutral particle in nucleus; mass ~1 amu.
• Nucleus: Small, dense, positively charged center containing protons and neutrons.
• Alpha particle: He2+ ion; +2 charge; ~4 amu; used in scattering experiments.
• Valence shell: Outermost occupied electron shell.
• Valence electrons: Electrons in the valence shell.
• Valency: Number of electrons lost/gained/shared to achieve stability.
• Atomic number (Z): Number of protons in nucleus.
• Mass number (A): Total number of nucleons (protons + neutrons).
• Isotopes: Same Z, different A.
• Isobars: Same A, different Z.
• Octet rule: Stability associated with 8 valence electrons.
• Duplet rule: Stability associated with 2 valence electrons (H, He, Li, Be, B).

Structured Summary

Sample Calculations and Configurations

• A − Z = neutrons; example Na: 23 − 11 = 12 neutrons.
• Chlorine average atomic mass: 35.5 u from 75% 35 and 25% 37.
• Mg2+: Start Mg (2,8,2) → lose 2 e− → (2,8); M shell electrons = 0.
• Cl−: Start Cl (2,8,7) → gain 1 e− → (2,8,8); valence electrons = 8.
• K (Z=19): 2,8,8,1; electrons in L shell = 8.

Action Items / Next Steps

• Memorize atomic numbers (1–20) and corresponding electron configurations.
• Practice drawing atomic structures (nucleus with protons/neutrons; shells with electrons).
• Solve exercises on:
  • Identifying Z, A, number of protons, neutrons, electrons from nuclide notation.
  • Forming ions and writing electron configurations.
  • Calculating average atomic mass from isotopic abundances.
• Review definitions: valency, isotopes, isobars, octet/duplet rules.