Discovery of Subatomic Particles

Introduction

• Focus on revision of the chapter covering key points in 1 hour.
• Significant for JEEs, scoring one easy question every year from this chapter.
• Covers gas filters, shadows, Thompson's atomic model, and magnetic radiations among others.

Discovery of Subatomic Particles

Electron

• Discovered using a discharge tube setup with cathode and anode under low pressure and high voltage.
• Zinc sulfide coating on anode showed glow indicating presence of 'negatively charged particles' – electrons.
• Key properties: straight-line movement and e/m ratio independent of the nature of gas in the tube.

Proton

• Discovered using a perforated cathode setup under similar low pressure and high voltage conditions.
• Positively charged particles moving towards the negative electrode – protons (H+ for hydrogen gas).
• Anode rays/canal rays, with properties like e/m ratio dependent on the nature of gas.

Neutron

• Discovered by bombarding beryllium with alpha particles (helium nuclei).
• Ejection of neutral particles – neutrons, with specific values for mass and charge.

Atomic Models

Thompson's Atomic Model

• Atom is spherical with non-uniformly distributed negative charges (compared to watermelon seeds).
• Positive charge uniformly distributed.
• Also called the plum pudding model.

Rutherford's Atomic Model

• Conducted the famous alpha ray scattering (gold foil) experiment.
• Most alpha particles pass straight, few deviate slightly, very few bounce back: indicating most of the atom is empty space, and positive charge is concentrated in a small nucleus.
• Proposed the concept of a nucleus and the relative size/scope comparison (~10^-15 meters for nucleus vs 10^-10 meters for the atom).
• Limitations: Did not explain electron distribution or Maxwell's theory.

Dual Nature of Light vs Particle Nature

• Electromagnetic radiations (gamma rays, x-rays, UV, visible rays) exhibit both particle (photon) and wave nature.
• Key electromagnetic properties: travel in vacuum, perpendicular electric and magnetic fields.

Planck's Quantum Theory

• Energy emission/absorption occurs in discrete packets called quanta (photon for light).
• Energy (E) proportional to frequency (ν): E = hν or E = hc/λ.

Photoelectric Effect

• Light of certain threshold frequency can eject electrons from a metal surface instantly with no time lag.
• Key observations: emission immediate, number of emitted electrons dependent on light intensity, kinetic energy dependent on frequency.
• Explained by Einstein’s formula: hν = hν₀ + K.E.

Spectrum

• Light spectrum is continuous, but atomic spectra (like hydrogen's) are discontinuous (line spectrum).
• Different series in hydrogen spectrum: Lyman (UV), Balmer (visible), Paschen, Brackett, Pfund, Humphrey (IR).
• Wave number formula: 1/λ = RZ² (1/n₁² - 1/n₂²), where n₁ < n₂.

Bohr's Atomic Model

• Spectra explained by electrons transitioning between fixed orbits without energy loss, emitting/absorbing quanta when jumping between orbits.
• Quantization of angular momentum: mvr = nh/2π.
• Formulae for radius, velocity, and energy of electron in nth orbit given.
• Limitations: Doesn't explain multi-electron spectra, Z-man effect, Stark effect, or fine spectra in high-resolution instruments.

Quantum Mechanical Model

• Proposes wave-particle duality for particles like electrons (de Broglie wavelength: λ = h/mv).
• Heisenberg's Uncertainty Principle: ΔxΔp ≥ h/4π.
• Electrons in orbitals described by probability densities.
• Schrödinger equation: Key quantum mechanical model for electron distribution.
• Probability of finding electron (ψ²) defines orbital regions.

Quantum Numbers

• Principal (n): shell number, energy level, shell size.
• Azimuthal (l): subshell type (s, p, d, f), shape.
• Magnetic (m): orientation of orbitals (-l to +l).
• Spin (s): electron spin direction (+1/2, -1/2).

Nodes

• Regions with zero probability of finding an electron.
• Radial nodes and angular nodes:
  • Radial node = n - l - 1.
  • Angular node = l.

Aufbau Principle

• Electrons fill orbitals in order of increasing energy (n + l rule).
  • E.g., 1s fills before 2s, 2p fills before 3s.

Pauli’s Exclusion Principle

• No two electrons in an atom can have the same set of four quantum numbers.

Hund's Rule

• Degenerate orbitals are singly occupied before pairing up with opposite spins.
  • E.g., in p orbitals (3 orientations: px, py, pz).

Special Configurations

• Chromium ([Ar] 3d⁵ 4s¹) and copper ([Ar] 3d¹⁰ 4s¹): more stable half-filled and fully filled d-orbitals.
  • Exchange energy and symmetrical distribution of electrons.

Conclusion

• Chapter overview of atomic structures, subatomic particles, and key principles/models in atomic theory essential for exams like JEE.
• Each principle/theory explained with focus on its contributions and limitations where applicable.