2.1 Discovery of Sub-atomic Particles
- Electron: Discovered by J.J. Thomson (1897) using cathode ray experiment; charge-to-mass ratio measured.
- Proton: Discovered by Rutherford (1917); positively charged particle in nucleus.
- Neutron: Discovered by James Chadwick (1932); neutral particle in nucleus.
2.2 Atomic Models
- Thomson’s Model: “Plum pudding” model; electrons embedded in positively charged sphere.
- Rutherford Model: Small dense nucleus, electrons revolve around nucleus; most of atom is empty space.
- Limitations: Could not explain atomic spectra and stability of atom.
2.3 Developments Leading to Bohr’s Model of Atom
- Emission spectra: Line spectra of hydrogen indicate quantized energy levels.
- Limitations of Rutherford model: Accelerating electrons should radiate energy → collapse of atom.
- Need for quantized orbits to explain stable electron motion.
2.4 Bohr’s Model for Hydrogen Atom
- Postulates:
- Electrons revolve in stable orbits without radiating energy.
- Angular momentum quantization:
mvr=n2πh,n=1,2,3,…
- Electron emits/absorbs radiation when jumping between orbits:
ΔE=hν
- Energy levels:
En=−n213.6 eV
- Explains hydrogen emission spectrum (Balmer series).
2.5 Towards Quantum Mechanical Model of the Atom
- Light and electrons exhibit wave-particle duality.
- Classical mechanics insufficient for microscopic particles.
- de Broglie hypothesis: Electrons have wavelength
λ=mvh
- Leads to quantum conditions for electron motion.
2.6 Quantum Mechanical Model of Atom
- Schrödinger equation: Determines electron distribution (ψ) in an atom.
- Quantum numbers: Describe electron’s energy, shape, and orientation of orbitals:
- Principal quantum number (n) – energy level
- Azimuthal quantum number (l) – orbital shape
- Magnetic quantum number (m) – orientation of orbital
- Spin quantum number (s) – electron spin
- Electrons exist in orbitals rather than fixed paths (Bohr’s orbits).
- Explains atomic spectra, chemical behavior, and stability.